Orbital Overlap and Its Role in Covalent Bonding

Ionic interactions and covalent bonding are viewed as 2 basic concepts for the chemical bond in today's knowledge of the phenomenon. Ionic bonds include the traditional electrostatic interactions between point charges. Electron-sharing or donor-acceptor bonds, which involve two orbitals that are profoundly populated and one that is vacant, are the two most common ways covalent bonding is described.

Even though Pauli's repulsion is greater than the electrostatic repulsive force when the orbital overlap concept increases, both types of bond formations ignore interactions among electrons with the same spin. The involved atom sizes, valence electrons, and degree of orbital overlap are the distinguishing factors. Furthermore, a larger overlap leads to a relatively strong bond formation among the two atoms. Thus, the orbital overlap theory explained the combination of atoms by interfering with one another's orbitals, resulting in relatively low energy levels where valence electrons come together to constitute covalent bonds.

What is the Concept of Orbital Overlap?

An orbital overlap is the combination of orbitals by the collision of neighbouring atoms in the space of the same area, which occurs in chemical bonds leading to a bond formation facilitated by orbital overlap. The two atoms that are close to one another, penetrate each other's orbitals during the orbital process, creating a new hybridised orbital through which the electrons of the bonding pair are located. Because it is less energetic than the atomic orbital, this hybridised orbital is firm and has a low energy state. Thus, the partial fusion of the orbital explains the orbital overlap concept.

To further clarify the system, it should be noted that it takes place over an atomic orbital. An atomic orbital is a location within the atom's interior where there is a high likelihood of finding electrons. The two provided nuclei within the atoms are also drawn to one another by the enhanced electron density in a small area, which reduces their repulsive forces.

For example, a covalent bond between H and Cl is the end outcome of the response.

Orbital Overlap Theory

Chemical bonding and the atom's shape or geometry are governed by how orbitals are arranged, which are explained by the below two orbital overlap theories. Molecular orbital theory (MOT) or Valence Bond Theory (VBT) can both be utilised to describe how these orbitals are arranged. The VBT explains the electron pair's orbital overlap. s, p, and d orbitals make up the majority of atomic orbitals.

The VBT states that a σ bond will be developed when two s or p orbitals overlap head-to-head. A π bond is created when two concurrent p-orbitals overlap. Since a double bond contains both a σ and a π bond; a single bond will comprise a σ bond. The MOT explains how overlapping atomic orbitals create molecular orbitals. This theory states that a molecular orbital can only support a maximum of 2 electrons. To reduce the attraction among them, these charged particles possess opposite spin.

Difference Between VBT and MOT

Overlap of Atomic Orbitals

The atomic orbitals of 2 atoms overlap once they are nearer to one another. The overlapping of atomic orbitals can have positive, negative, or zero overlaps, relying upon these characteristics. The figure beneath shows the different configurations of the s and p-orbitals that lead to positive, negative, and zero overlaps.

• Positive atomic orbital overlap: Whenever the two involved atomic orbitals phase is identical, positive overlap takes place. Bonds are created as a consequence of this overlap.

• Negative atomic orbital overlap: Negative overlap occurs whenever the phases of the involved atomic orbitals oppose one another. Bond formation doesn't take place in this instance.

• Zero overlaps of atomic orbital: Zero atomic orbital overlaps occur while two intriguing orbitals do not overlap with each other in an orbital.

The orbital overlap diagram is shown below.

Structure of Orbital Overlap

Orbital Overlap in Cumulene Compounds

The existence of 2 main carbon atoms carrying 2 double bonds accounts for the orbital overlap in cumulene compounds' rigidity. Due to the sp hybridisation of such carbon atoms, 2 π bonds—one to near each carbon atom—are formed. Cumulene molecules thus possess linear geometry. Hybridisation of cumulene contains 9 (σ) and 2 (π) bonds.

Key Features of Orbital Overlap

• The phrase "atomic orbital overlap" is another name for orbital overlapping.

• Linus Pauling highlighted the significance of orbital overlap while characterising the molecular bond angles found during experimentation.

• The idea of orbital hybridisation also represents an additional development of orbital overlapping.

• Orbital overlap refers to the methodology whereby a partial merger of orbitals creates a completely novel hybridised orbital. The overlapping regions of the orbitals are called pi (π) and sigma (σ).

• Bond-forming orbitals must have the same orientation and mode in space.

• The pair of atoms involved, their size, and valence electrons all play a role in determining the degree of overlap level. Higher levels of overlap result in the atoms forming firmer bonds with one another.

Conclusion

So, bond chemistry helps analyse the orbital overlap that happens when two atoms mingle with molecules. With the help of bond chemistry, orbital overlap theories VBT and MOT explain the atomic orbital overlap, while VBT specifically explains the orbital hybridisation concept where MOT falls short.