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Hybridisation is defined as the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies, resulting in the formation of new set of orbitals of equivalent energies and shape.

Let us take the example of carbon to understand what that means. The atomic number of carbon is 6, and hence its electronic configuration is 1s22s22p2.

2 unpaired electrons in carbon atom

Ground state of carbon

Clearly, there are only two unpaired electrons in carbon; therefore, carbon should form two bonds only. But in reality, carbon forms four covalent bonds.

Methane (CH4)


Carbon tetrachloride (CCl4)

Carbon tetrachlorde

The reason why carbon forms four covalent bonds is that one paired electron from orbital 2s becomes unpaired and jumps to the orbital 2p, which has ample space, without consuming excessive energy.

Excited state of carbon

Excited state of carbon

Now, carbon has 4 unpaired electrons.

In the excited state, carbon has 4 unpaired electrons — 1 in 2s and 3 in 2p orbitals — thereby explaining the tetravalent nature of carbon. The four orbitals can overlap with other incompletely filled orbitals such as hydrogen, chlorine to form four covalent bonds.

Although both s and p orbitals take part in bond formation, all four bonds in carbon are same in energy. Moreover, the bonds formed by 2p orbitals should have been inclined at 90° to one another and the bond formed by 2s orbitals would have been in any direction. However, all four bonds are inclined at an angle of 109°28' to one another. This leads to the following conclusion:

Atomic orbitals of slightly different energies combine to form new set of orbitals of equivalent energies and shape. The new orbitals are called hybrid orbitals and the concept is termed as hybridisation.

Note : On this page, z-axis is taken as molecular axis (internuclear axis) unless otherwise stated.

Formation of sp3 hybrids

formation of sp3 hybrids

Note : The three p-orbitals are at right angle to one another; therefore, they are denoted as px, py and pz. The px, py and pz are aribtrary symbols.

Salient Features of Hybridisation

  1. Type of hybridisation indicates the geometry of molecules.
  2. The hybridised orbitals are always equivalent in energy and shape.
  3. The bond formed by hybrid orbitals is much more stable than the bond formed by the pure atomic orbitals.

Types of Hybridisation

The different types of hybridisation are as under :

sp hybridisation

sp hybridisation involves mixing of one s orbital and one p orbital resulting in the formation of two equivalent sp hybrid orbitals . An sp hybrid orbital has 50% s character and 50% p character.

Study the following example which explains the formation of BeCl2 molecule :

The atomic number of Be is 4 and its electronic configuration is 1s22s2.

Be in its Ground State

Ground state of Be

In its excited state, an electron from 2s orbital of Be atom jumps to 2p orbital which results in two unpaired electrons in Be.

Be in its Excited State

 Excited state of Be

One 2s and one 2p orbitals get hybridised to form two sp hybrid orbitals. The sp hybrid orbitals then overlap with the 2p-orbital of chlorine to form BeCl2.

sp hybridisation

 sp hybrids of Be

Geometry of sp hybridised molecules

sp hybrid orbitals are collinear with an angle of 180°. sp hybridisation is also known as diagonal hybridisation. Formation of sp hybrids and BeCl2 molecule is given below :

Formation of sp hybrids

Formation of sp hybrids
formation of BeCl2

Some other examples of sp hybridisation are BeF2, BeH2 and C2H2 etc.

sp2 hybridisation

sp2 hybridisation involves mixing of one s-orbital and two p-orbitals resulting in the formation of three equivalent sp2 hybridised orbitals. Each sp2 hybridised orbital has 33.33% s-character and 66.67% p-character.

The ground state and excited state of boron, and formation of BCl3 are given below:

Boron in its ground state

Ground state of Boron

Boron in its excited state

Excited state of boron

sp2 hybridisation in BCl3

sp2 hybridisation in BCl3

Geometry of sp2 hybridised molecules

In sp2 hybridisation, all the three hybrid orbitals remain in the same plane making an angle of 120° with one another. This arrangement is known as trigonal planar arrangement, and therefore sp2 hybridisation is also known as trigonal hybridisation.

Formation of sp2 hybrids

Formation of sp2 hybrids

Formation of BCl3 molecule is given below.

Formation of BCl3

formation of BCl3

BCl3 (simplified view)

BCl3 bond angle 120 degree

sp3 Hybridisation

sp3 hybridisation involves mixing of one s-orbital and three p-orbitals resulting in the formation of four sp3 hybrid orbitals. Each sp3 hybrid orbital has 25% s character and 75% p character.

The example of carbon at the beginning of this section is an example of sp3 hybridisation. sp3 hybridisation is seen in molecules such as CH4, CCl4.

sp3 hybridisation in CH4

sp3 hybridisation example : CH4

Geometry of sp3 hybridised molecules

The four sp3 hybrid orbitals are directed towards the four corners of the regular tetrahedron and make an angle of 109°28' with one another. The sp3 hybridisation is also called a tetrahedral hybridisation.

Formation of CH4

Formation of methane

CH4 (simplified view)

bond angle 109 degree 28 minutes

Hybridisation of elements involving d Orbitals

The concept of hybridisation of elements involving d orbitals is similar to what we have learned so far. In this case, d orbital in addition to s and p orbitals also takes part in hybridisation. It should, however, be noted that the energy of 3d orbitals is comparable to 3s and 3p orbitals as well as 4s and 4p orbitals. Hence, the hybridisation involves either 3s, 3p and 3d orbitals; or 3d, 4s and 4p orbitals. The difference in energies of 3p and 4s orbitals is significant. Hence, hybridisation involving 3p, 3d and 4s orbitals is not possible.

Hybridisation Involving d Orbitals
Shape of molecules/ionsHybridisation typeExample
Square planardsp2[Ni(CN)4]2-
Trigonal bipyramidalsp3dPCl5
Square pyramidalsp3d2BrF5
Octahedralsp3d2, d2sp3SF6

sp3d and sp3d2 hybridisations are discussed below.

sp3d hybridisation

Molecules such as PCl5 undergo sp3d hybridisation. In sp3d hybridisation, five orbitals (one s, three p and one d orbitals) undergo hybridisation forming five sp3d hybrid orbitals. The geometry of such molecules is trigonal bipyramidal. The formation of PCl5 and its geometry are given below:

Formation of PCl5

sp3d hybridisation : PCl5

Trigonal bipyramidal (PCl5)

PCl5 shape : trigonal bipyramidal

You should note that all the bond angles in trigonal bipyramidal geometry are not equivalent. In PCl5, three of the chlorine atoms lie in one plane making an angle of 120° with one another; these atoms are called equatorial atoms and the bonds formed by them are called equatorial bonds. The remaining two chlorine atoms lie at right angle to the plane of equatorial bonds and are known as axial atoms and the bonds formed by them are called axial bonds. Axial bond pairs suffer more repulsion from the equatorial bond pairs; therefore, axial bonds are slightly longer and hence weaker than the equatorial bonds.

sp3d2 hybridisation

Molecules such as SF6 undergo sp3d2 hybridisation. The geometry of such molecules is known as octahedral.

Formation of SF6

sp3d2 hybridisation example

Octahedral geometry (SF6)

octahedral shape of SF6

Geometry of molecules containing at least one lone pair

The force of repulsion between a lone pair and a bond pair is more than the force of repulsion between two bond pairs. The order of repulsion between electron pairs is : lone pair-lone pair > lone pair-bond pair > bond pair-bond pair
Due to this reason, the geometry of molecules containing lone pairs is somewhat distorted.

Geometry of Ammonia (NH3)

The atomic number of nitrogen is 7; hence, its electronic configuration is 1s22s22px12py12pz1.
The four orbitals (one 2s and three 2p) undergo sp3 hybridisation forming four sp3 hybridised orbitals out of which one contains a lone pair and three contain one electron each and form three N-H sigma bonds.

Because of the presence of a lone pair, the molecule gets distorted and the bond angle is reduced to 107° from 109°28'. The geometry of such a molecule is known as trigonal (or triangular) pyramidal.

Geometry of H2O

The atomic number of oxygen is 8; hence, its electronic configuration is 1s22s22px22py12pz1.
The four orbitals (one 2s and three 2p) undergo sp3 hybridisation forming four sp3 hybridised orbitals out of which two contain one electron each and take part in bond formation with hydrogen, and the other two contain a pair of electrons (lone pair).

Because of the presence of two lone pairs, the bond angle in this case is reduced to 104.5° from 109°28'. The shape of such a molecule is known as V-shaped or bent.

Ammonia (NH3)

Shape of ammonia (NH3)

Water (H2O)

Shape of water (H2O)

Geometry of SF4

The atomic number of sulfur, which is the central atom, is 16; hence, its ground state electronic configuration is 1s22s22p63s23px23py13pz1, in the excited state, it is 1s22s22p63s23px13py13pz13d1.

It undergoes sp3d hybridisation forming five hybrid orbitals. The arrangement is trigonal bi-pyramidal. Depending on the position of lone pair, two possible structures of SF4 are:

SF4 (Less stable)

SF4 shape 1 : less stable


SF4 (More stable)

SF4 shape 2 : more stable


In structure (I), there are three lp-bp repulsions at 90° whereas in structure (II), there are only two lp-bp repulsions at 90° and two lp-bp repulsions at 120°. Hence, structure (II) is more stable, and the correct structure. This shape is known as see saw or distorted tetrahedron or folded square.

Geometry of ClF3

The arrangement of electron pairs in ClF3 is trigonal bi-pyramidal, like in SF4. But ClF3 has two lone pairs and three bond pairs. The following are the three possible structures of ClF3:

ClF3 (Most stable)

Most stable shape of ClF3



Shape of ClF3 : less stable



Shape of ClF3 : less stable


The structure (I) is the most stable one because it has least repulsion; hence, (I) is the correct geometry of ClF3. Such a shape is known as T-shape.

Important Conditions for Hybridisation

Read the following points before you make any assumptions on your own :

  1. Only the orbitals present in the valence shell of the atom are hybridised.
  2. The orbitals undergoing hybridisation should have almost equal energy.
  3. It is not necessary that all the half-filled orbitals must participate in hybridisation. Similarly, in some cases, even completely filled orbitals participate in hybridisation.
  4. Jumping of electron from ground state to excited state is not necessary.

Hybridisation − Carbon

Since no organic compound exists without carbon, hybridisation of carbon is worth discussing.

sp3 Hybridisation − Carbon

When a carbon atom is linked to four atoms, the hybridisation is sp3.

All the carbon atoms in methane and ethane are sp3 hybridised

sp3 hybridisation example, methane
sp3 Hybridisation of carbon example, ethane

sp2 Hybridisation − Carbon

When a carbon atom is attached to 3 atoms (one double bond and two single bonds), the carbon is said to be sp2 hybridised carbon.

Both the carbon atoms in ethene are sp2 hybridised

sp2 hybridised carbon : ethene

sp Hybridisation − Carbon

When a carbon atom is linked to two atoms (one triple bond and one single bond; or two double bonds), the carbon is said to be sp hybridised carbon.

Both the carbon atoms in ethyne are sp hybridised

sp hybridised carbon : ethyne

What is the hybridisation of each carbon atom?

What is the hybridisation of carbon
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