Hybridization

Hybridization is a concept used in organic chemistry to explain chemical bonding in cases where the valence bond theory does not provide satisfactory clarification. This theory is especially useful to explain the covalent bonds in organic molecules.

Basically, hybridization is intermixing of atomic orbitals of different shapes and nearly the same energy to give the same number of hybrid orbitals of the same shape, equal energy and orientation such that there is minimum repulsion between these hybridized orbitals.

Types of Hybridization

There are different types of hybridization-based on the mixing of the orbitals.

• sp³ hybridization: When one s orbital and three p orbital from the same shell of atom mix together to form a new equivalent orbital then this is called sp³ hybridization.

• sp² hybridization: It is observed when one s orbital and two p orbitals undergo mixing of energy for equivalent orbitals.

• sp hybridization: When one s and one p orbital goes in the process of mixing of energy to form a new orbital such kind of hybridization is called sp hybridization. The molecules possessing sp hybridization used to have a linear shape with an angle of 180°.

The above are three basic hybridizations along with them there are other hybridizations based on the mixing of orbitals such as sp³d hybridization, sp³d² hybridization and sp³d² hybridization.

Explanation of Hybridization Through Examples

Example 1:  Consider an example of the simplest hydrocarbon molecular Methane. CH₄. According to experimental observations, the Methane molecule has 4 identical C-H bonds with equal length and equal bond energy. All the four hydrogen atoms are arranged in a manner such that the four hydrogen atoms form corners of a regular tetrahedron. 

                                                 Image: Structural Formula of Methane

Based on the valence theory, a covalent bond is formed between two atoms in a molecule when there is an overlapping of half-filled atomic orbitals containing unpaired electrons. In the case of the methane molecule, we first write down the electronic configuration of each atom - C and H

      Image: Electronic configuration of carbon and hydrogen for hybridization

Each carbon atom has two unpaired electrons (in the 2pₓ and 2pᵧ orbitals). Based on the valence theory, only two hydrogen molecules could be paired to the two unpaired electrons of the carbon atom and there will be a formation of only 2 C-H bonds in the molecule. This will lead to an incomplete octet in the 2nd orbital of the carbon molecule (2pz orbital is unfilled) and so the molecule should be unstable. However, we see that actually the methane molecule is extremely stable in nature and has 4 C-H bonds and not two. Thus, the valence theory doesn’t explain the covalent bond of the methane molecule. 

The hybridization concept explains the formation of identical 4 C-H bonds and the tetrahedral shape of the molecule. 

According to this concept, when a carbon atom reacts with a hydrogen atom, the electrons in the carbon atom initially go into an excited state as shown here:

    Image: Electronic configuration of carbon in the ground state and in the excited state

Post excitation, hybridization can be imagined as the process where these 4 excited s and the p orbitals combine together to give a homogenous mixture and divide themselves into 4 identical orbitals having identical energy. These new orbitals have been termed hybridized orbitals. Since there one s orbital and 3 p orbitals have combined to form the hybrid orbital, the hybridized orbitals are called sp³ orbitals. The energy of these hybrid orbitals lie in between the energy levels of the s and the p orbitals as shown here:

                                       Image: Formation of the hybridized orbital sp³

Each sp³ hybrid orbitals has one unpaired electron. Since these 4sp³ orbitals are identical in terms of energy, there is a tendency amongst these electrons to repel each other. To minimize the repulsion between electrons, the sp³ hybridized orbitals arrange themselves around the carbon nucleus in a tetrahedral arrangement. The resulting carbon atom is termed as sp³ hybridized carbon atom. 

                          Image: Tetrahedral arrangement of sp³ hybridized orbital

Overlap of each of the 4sp³ orbitals of the hybridized carbon atom with the s orbital of the hydrogen atoms leads to the formation of a methane molecule. The methane molecule can be shown as:

                              Image: sp³-s overlapping to form C-H bonding

It can be seen from the above that there are 4 identical sp³-s overlaps forming 4 identical C-H bonds which are consistent with the observations. Moreover, since these sp³ orbitals are oriented in the form of tetrahedrons, the geometry of the methane molecule is tetrahedral.

Thus, adding the concept of hybridization to the valence theory helps to understand the bonding in the methane molecule.

Example 2: The above example of methane had sp³ hybridization formed because of hybridization of 1 s and 3 p orbitals of the carbon atom. There are other types of hybridization when there are hybrid orbitals between 2 p orbitals and 1 s orbital called sp² hybridization. In case, there are hybrid orbitals between 1 s and 1 p orbitals, it is called sp hybridization. 

Let us consider the case of sp² hybridization. The structure of ethylene can be explained using the concept of sp² hybridization. The structure of the ethylene molecule observed is as:

                         Image: Electronic configuration of sp² orbital

Experimentally, the four carbon-hydrogen bonds in the ethylene molecule are identical and the geometry at each carbon atom in the ethylene molecule is planar trigonal.

Since the carbon atom has only 2 unpaired electrons, the valence bond theory cannot explain the formation of 4 bonds by each of the carbon atoms. Hence, we have to consider the excited state of both the carbon atoms in order that each carbon atom forms 4 bonds. 

We first consider the two carbon atoms and the double bond between them. 

For each of the excited carbon atoms, the one 2s orbital and two 2p orbitals (of the three 2p orbitals) form hybridization resulting in 3 hybrid orbitals called sp2 sp² orbitals. (1 s and 2 p orbitals). These 3 sp² orbitals try to be as distant from each other as possible and hence form a planar trigonal structure. The third 2p orbital in each of the carbon atoms does not participate in hybridization and remains as 2p orbital. 

                        Image: Structural Formula of C₂H₄

Each of the three sp² hybrid orbitals and the non-hybrid 2p orbital has 1 unpaired electron. To minimize repulsion of this non-hybrid 2p orbital with the 3 sp² orbitals, the 2p orbital stands perpendicular to each of the sp² hybrid orbitals. Hence, post-hybridization, the sp2 hybridized carbon atom looks as:

                                      Image:  sp2 hybridized carbon atom 

Each carbon atom in the ethylene molecule is bonded to two hydrogen atoms. Thus, overlap two sp²-hybridized orbitals with the 1s orbitals of two hydrogen atoms Also, the covalent C-C bond forms by overlapping of sp² orbitals of the two carbon atoms as:

                        Image: C-C bond forms by the overlapping of sp² orbitals

The two 2p orbitals of the carbon atoms overlap laterally to form a weak bond called a pi bond. 

Thus, the ethylene molecule is said to have sp2- sp² s bonds (4 C-H bonds), one sp²-sp² bond (C-C bond) and one p-p pi bond (C-C bond).   

Thus, the sp² hybridization theory explains the double bond, the trigonal planar structure in ethylene molecules. 

Example 3: Similarly, for a triple bond formation, like that of an acetylene molecule, there is sp hybridization between 1 s and 1 p orbital of the carbon atom. 

 Image: Structural Formula of C₂H₂

Here, there are 2 C-H bonds and a triple C-C bond. 

In each of the excited carbon atoms, one 2s and one 2p orbital form hybrid molecules called sp hybrid orbitals and the non-hybrid two 2p orbitals do not participate in hybridization. Because there are electron molecules in each of the orbitals, they tend to repel each other and the 2sp orbitals form a linear arrangement. The non-hybrid 2p orbital position themselves as far away as possible from each sp-hybridized orbital when perpendicular to each sp-hybridized orbital. So the resulting sp hybrid carbon atom looks like this:

                                                  Image:  sp hybridization 

The s orbitals of the hydrogen atom overlap with one sp hybrid orbital of each of the carbon atoms forming the 2 C-H bonds. The C-C covalent bond is formed by overlapping the sp-sp orbitals of the two-hybrid carbon atoms. In order to complete octet, the two non-hybrid 2p orbitals of each of the carbon atoms overlap laterally forming 2 pi bonds as shown:

Thus, sp hybridization explains the triple bond in acetylene molecules and the linear structure as well.

Nature of  the Types of Hybridization 

Hence, from the above text, we understand that hybridization is mathematically a concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. Also, an entirely new orbital formed is different from its components and hence being called a hybrid orbital.