Concept of Hybridisation involving s, p and d Orbitals

Introduction

In Chemistry, hybridisation is explained as mixing two atomic orbitals to make new hybridised orbitals. The development of this kind of hybrid orbitals and unique energy structures is frequently the outcome of this intermixing. Hybridisation is mainly carried out with atomic orbitals of a similar energy level. Nonetheless, both half-filled and filled orbitals can help participate in the procedure if their strength or energies are equivalent. On the other hand, the notion of hybridisation can be considered an addition to the valence bond theory, as it aids in understanding the bond formation, energy bond, and bond lengths.

What is the definition of Hybridisation?

The two atomic orbitals come together to make a hybrid orbital in a molecule. The power of particular atoms’ orbitals is redolent to offer orbitals of energy equivalent. Hybridisation is the term for this procedure. The atomic orbitals with similar energies are combined throughout the hybridisation procedure, which usually involves mixing two orbitals or two ‘p’ orbitals or combining an ‘s’ orbital with a ‘p’ orbital and an ‘s’ orbital with a ‘d’ orbital. Hybrid orbitals are one of the novel orbitals that result from this process. Hybrid orbitals are particularly valuable in describing atomic bonding characteristics and molecule geometry.

Let’s take a look at a carbon atom as an example. The valence-shell s orbital mixes along three valence-shell p orbitals to generate four single bonds. This combination produces four sp3 hybridisation mixes that are comparable. These will have a tetrahedral structure from the carbon bound to four distinct atoms.

Hybridisation’s key characteristics

The following are the key characteristics of hybridisation:

• The number of hybrid orbitals is the same as that of hybridised atomic orbitals
• The energy and geometry of the hybridised orbitals are always the same
• Pure atomic orbitals are less successful than hybrid orbitals at forming stable bonds
• These hybrid orbitals are oriented in space in the desired direction to minimise electron pair repulsion and maintain a stable configuration
• As a result, the kind of hybridisation shows the molecule’s shape

Hybridisation requires Certain Circumstances 

(i) The orbitals in the atom’s valence shell have become hybridised.

(ii) The energy of the orbitals undergoing hybridisation should be about identical.

(iii) Before hybridisation, electron promotion was not a need.

(iv) Only half-filled orbitals do not have to participate in hybridisation. Hybridisation can involve even full valence shell orbitals in rare circumstances.

Types of Hybridization

Hybridisation involving s, p, and d orbitals can take several forms. The following are the many forms of hybridisation:

(I) sp Hybridization: The mixing of one s and one p orbital produces two equivalent sp hybrid orbitals in this kind of hybridisation. If the hybrid orbitals lie along the z-axis, the orbitals s and Pz are acceptable for sp hybridisation. Each sp hybrid orbitals has 50% s-character and 50% p-character. A molecule with linear geometry has an sp-hybridized core atom and is connected straight to two additional central atoms. The term “diagonal hybridisation” refers to this sort of hybridisation.

With projecting positive lobes and extremely tiny negative lobes, the two sp hybrids point in opposing directions along the z-axis, allowing for more efficient overlapping and the development of stronger connections.

(II) sp2 hybridization: Ones and two p-orbitals are involved in this hybridization, resulting in three equivalent sp2 hybridized orbitals. The electronic configuration of the central boron atom in the BCl3 molecule, for example, is 1s22s22p1. One of the 2s electrons is promoted to the unoccupied 2p orbital in the excited state, leaving the boron with three unpaired electrons.

These are the three orbitals (one 2s and two 2p)

Three sp2 hybrid orbitals result from the hybridisation. The resulting three hybrid orbitals overlap with a trigonal planar arrangement. Chlorine’s 2p orbitals combine to generate three BCl molecule bonds. With a ClBCl link, the geometry is trigonal planar at a 120° angle.

(III) sp3 hybridisation: This sort of hybridisation may be discussed using the CH4 molecule as an example, where one s-orbital and three p-orbitals of the valence shell are mixed to generate four sp3 hybrid orbitals with equivalent energies and shapes. Each sp3 hybrid orbital has a 25 per cent s-character and a 75 per cent p-character. The four sp3 hybrid orbitals that result are pointed towards the tetrahedron’s four corners. The angle between the sp3 hybrid orbital and the sp3 hybrid orbital is 109.5°.