The following topics are covered :
- What is hybridisation.
- Salient features of hybridsation
- Types of hybridisation.
- Geometry of molecules
- Important conditions for hybridisation.
- Hybridisation of carbon.
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.
Clearly, there are only two unpaired electrons in carbon; therefore, carbon should form two bonds only. But in reality, carbon forms four covalent bonds.
Carbon tetrachloride (CCl4)
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
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.
Formation of sp3 hybrids
Salient Features of Hybridisation
- Type of hybridisation indicates the geometry of molecules.
- The hybridised orbitals are always equivalent in energy and shape.
- 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 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
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
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.
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
Some other examples of sp hybridisation are BeF2, BeH2 and C2H2 etc.
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
Boron in its excited state
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 BCl3 molecule is given below.
Formation of BCl3
BCl3 (simplified view)
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
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
CH4 (simplified view)
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.
|Shape of molecules/ions||Hybridisation type||Example|
sp3d and sp3d2 hybridisations are discussed below.
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
Trigonal bipyramidal (PCl5)
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.
Molecules such as SF6 undergo sp3d2 hybridisation. The geometry of such molecules is known as octahedral.
Formation of SF6
Octahedral geometry (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
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
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.
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 (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)
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 :
- Only the orbitals present in the valence shell of the atom are hybridised.
- The orbitals undergoing hybridisation should have almost equal energy.
- 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.
- 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
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
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
What is the hybridisation of each carbon atom?