Molecules have different forms. From their physical features to their chemical reactivity, there is a wealth of scientific evidence to support this claim. Small molecules, defined as those with a single core atom, have predictable structures. Valence shell electron pair repulsion is a fundamental concept in molecular structures (VSEPR). According to VSEPR, electron pairs repel each other to move as far away from each other as possible because they are made up of negatively charged particles. VSEPR distinguishes between electron group geometry, which describes the arrangement of electron groups (bonds and nonbonding electron pairs), and molecular geometry, which describes the arrangement of atoms in a molecule.

Any sort of bond—single, double, or triple—and lone electron pairs are the two types of electron groups. The initial step in using VSEPR on simple compounds is to count the number of electron groups surrounding the core atom. It’s important to remember that a multiple bond only counts as one electron group.

Linear Molecule

The notable property of linear molecules is that all of the electron groups are mutually repelled, resulting in their atoms aligning on a straight line with a bond angle of 180 degree.

Non-Linear Molecule

The repulsion between the electron groups in non-linear molecules is not the same; it is discovered that the repulsion between a lone pair and a bonding pair is stronger than that between two bonding pairs. The bent structure of the H2O molecule and the triangular pyramidal or trigonal planar form of the NH3 molecule are due to this.

Tetrahedral Molecules

When tetravalent atoms (such as carbon and silicon) form perfect covalency with atoms of the same element, tetrahedral molecules form. The four covalent bonds are symmetrically dispersed in three dimensions, with a 109.5 degree angle between any two bonding pairs.

Examples can be seen in  methane, CH4; silane, SiH4; and tetrachloromethane, CCl4

Hybridization

The concept of hybridization is that atomic orbitals merge to generate fresh hybridised orbitals, which affects molecule shape and bonding properties. The valence bond hypothesis is additionally expanded by hybridization. To further investigate this concept, we’ll use three different types of hydrocarbon molecules to demonstrate sp3, sp2, and sp hybridization. 

When two atomic orbitals combine to form a hybrid orbital in a molecule, the energy of individual atoms’ orbitals is redistributed to give orbitals of equivalent energy. Hybridization is the term for this procedure. The atomic orbitals with comparable energies are mixed together during the hybridization process, which usually involves the merging of two ‘s’ orbitals or two ‘p’ orbitals, or the mixing of an’ orbital with a ‘p’ orbital, as well as an’ orbital with a ‘d’ orbital. 

Types of Hybridization

Hybridization can be classed as sp3, sp2, sp, sp3d, sp3d2, sp3d3 depending on the types of orbitals involved in mixing. Let’s look at the many types of hybridization and some examples of each.

sp Hybridization

When one s and one p orbital in the same main shell of an atom combine to generate two new equivalent orbitals, this is known as sp hybridization. The sp hybridised orbitals are new orbitals that are developed. It produces 180-degree linear molecules.

• One’s orbital and one ‘p’ orbital of equal energy are mixed to form a new hybrid orbital known as a sp hybridised orbital.
• Diagonal hybridization is also the name for sp hybridization.
• Each sp hybridised orbital has the same proportion of s and p characters – 50 percent s and 50 percent p.

sp Hybridization examples:

•  Compounds of beryllium like BeF2, BeH2, BeCl2
• All compounds of carbon-containing triple bond like C2H2.

sp2 Hybridization

When one s and two p orbitals of the same shell of an atom combine to generate three equivalent orbitals, this is known as sp2 hybridization. sp2 hybrid orbitals are the new orbitals that have been created.

• Trigonal hybridization is another name for sp2 hybridization.
• It entails combining one’s’ orbital with two ‘p’ orbitals of equal energy to form the sp2 hybrid orbital.
• A trigonal symmetry blend of s and p orbitals is maintained at 120 degrees.
• All the three hybrid orbitals are present in the same plane and they form a 120° angle with one another. The hybrid orbitals have a 33.33 percent’s’ character and a 66.66 percent ‘p’ character in each of them.
• A triangle planar form is found in molecules in which the central atom is connected to three other atoms and is sp2 hybridised.

Examples of sp2 Hybridization

• The compounds of Boron i.e. BF3, BH3
• The compounds of carbon-containing a carbon-carbon double bond, Ethylene (C2H4) is also the example of this hybridization.

sp3 Hybridization

A tetrahedral hybridization, or sp3, occurs when one’s orbital and three ‘p’ orbitals belonging to the same shell of an atom combine to generate four new equivalent orbitals. sp3 hybrid orbitals are the new orbitals that have been generated.

• These are aimed at the four corners of a conventional tetrahedron and form a 109°28′ angle with each other.
• The sp3 hybrid orbitals have a 109.28 degree angle between them.
• Each sp3 hybrid orbital contains 25% s character and 75% p character.

Example of this hybridization are : ethane (C2H6), methane.

sp3d Hybridization

The mixing of 1s, 3p, and 1d orbitals to generate 5 sp3d hybridised orbitals of equal energy is known as sp3d hybridization. Their geometry is trigonal bipyramidal.

• The trigonal bipyramidal symmetry is formed by combining the s, p, and d orbitals.
• The equatorial orbitals are three hybrid orbitals in the horizontal plane that are inclined at a 120° angle to each other.
• The remaining two orbitals, known as axial orbitals, are located in the vertical plane at 90 degrees to the equatorial orbitals.

Example: Hybridization in Phosphorus pentachloride (PCl5)

Conclusion

All bond lengths become equal throughout the hybridization process. The valence shell electron pair repulsion theory helps explain bond angles (VSEPR theory). According to this idea, electron pairs resist each other, therefore electron pairs in bonds or lone pairs in orbitals around an atom want to keep as far away from each other as possible.