4.6.1. Types of Hybridisation

There are various types of hybridisation involving s, p and d orbitals. The different types of hybridisation are as under:

Hybridisation is a concept introduced by Linus Pauling to explain the geometry of molecules. It describes how atomic orbitals mix to form new, equivalent orbitals that are used in bonding. The key point is that hybrid orbitals are formed from s, p, and, in some cases, d orbitals.

Think of hybridisation like mixing different paint colors to create a new shade. Just as mixing blue and yellow paint can create green, combining atomic orbitals creates new hybrid orbitals that have distinct properties.

(i) sp hybridisation: This type of hybridisation involves the mixing of one s and one p orbital resulting in the formation of two equivalent sp hybrid orbitals. The suitable orbitals for sp hybridisation are s and pz, if the hybrid orbitals are to lie along the z-axis. Each sp hybrid orbital has 50% s-character and 50% p-character. Such a molecule in which the central atom is sp-hybridised and linked directly to two other central atoms possesses linear geometry. This type of hybridisation is also known as diagonal hybridisation. The two sp hybrids point in the opposite direction along the z-axis with projecting positive lobes and very small negative lobes, which provides more effective overlapping resulting in the formation of stronger bonds.

In sp hybridisation, one s and one p orbital combine to give two sp hybrid orbitals. These orbitals align 180 degrees apart, resulting in a linear shape. Because each sp hybrid orbital has equal proportions of s and p characteristics, the resulting molecule has distinct directionality and strength. For example, in beryllium chloride (BeCl2), the beryllium atom hybridizes its orbitals to form two sp hybrid orbitals that create strong sigma bonds with chlorine atoms at an angle of 180°.

Imagine two people standing back-to-back with their arms stretched out sideways. This position can be likened to the two sp hybrid orbitals directing outward. Each arm represents a bond to a chlorine atom, illustrating how the beryllium atom forms bonds in a linear fashion.

(ii) sp2 hybridisation: In this hybridisation, there is involvement of one s and two p-orbitals in order to form three equivalent sp2 hybridised orbitals. For example, in BCl3 molecule, the ground state electronic configuration of the central boron atom is 1s²2s²2p¹. In the excited state, one of the 2s electrons is promoted to a vacant 2p orbital as a result boron has three unpaired electrons. These three orbitals (one 2s and two 2p) hybridise to form three sp2 hybrid orbitals.

When a central atom, like boron in BCl3, undergoes sp2 hybridisation, it mixes one s and two p orbitals to create three sp2 hybrid orbitals. These are arranged in a trigonal planar geometry with 120-degree angles between each other, allowing for effective overlap with the 2p orbitals of chlorine atoms to form three strong sigma bonds.

Consider a triangular table where each corner represents an sp2 hybrid orbital. If three friends each sit at a corner of the table, they represent the bond angles formed between the boron atom and the chlorine atoms— all sitting at equal spacing to maximize distance from each other.

(iii) sp3 hybridisation: This type of hybridisation can be explained by taking the example of CH4 molecule in which there is mixing of one s-orbital and three p-orbitals of the valence shell to form four sp3 hybrid orbital of equivalent energies and shape.

In sp3 hybridisation, one s and three p orbitals combine to create four equivalent sp3 hybrid orbitals. These orbitals are arranged in a tetrahedral geometry, with angles of 109.5° between them. This arrangement allows for maximum distance between the electron pairs, minimizing repulsion and stabilizing the molecule. For instance, in methane (CH4), the carbon atom forms four sigma bonds with hydrogen atoms, creating a tetrahedral shape.

Think of the four sp3 orbitals as the hands of a person holding a box in each hand while spreading their arms outwards. Each arm represents a bond with hydrogen atoms, and the angles formed are the tetrahedral angles in CH4, showing how the shape effectively reduces electron repulsion.

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