Orbital Overlap

1.0Introduction

Orbital overlap refers to the concept in which atomic orbitals of two atoms come close and partially merge into the same region of space. This occurs when atoms approach each other during the formation of a bond. As the atomic orbitals overlap, their electrons with opposite spins pair up to form a covalent bond. This overlapping results in the formation of a new orbital, often referred to as a hybridised or molecular orbital, which is more stable due to its lower energy.

2.0What is Atomic Orbital Overlap?

When two atoms approach each other to form a covalent bond, their atomic orbitals interact in such a way that the overall energy of the system is minimized. As the atoms approach very closely, their orbitals partially overlap in the same region of space.

This partial merging of atomic orbitals is referred to as orbital overlapping or the overlapping of atomic orbitals. The extent of orbital overlap depends on three main factors:

1. Type of participating atoms
2. Size of the atoms
3. Number of valence electrons

A greater degree of overlap leads to a stronger covalent bond between atoms. According to the orbital overlap concept, atoms form covalent bonds by merging their orbitals such that valence electrons with opposite spins pair up in the shared orbital, resulting in a more stable, low-energy state.

3.0Types of Orbital Overlapping

1. Sigma (σ) Bond:
   Formed when orbitals overlap directly along the internuclear axis (the line joining the centers of two nuclei). This type of overlap is head-on and usually stronger.
2. Pi (π) Bond:
   Formed when orbitals overlap sideways, i.e., above and below the internuclear axis. Pi bonds are generally weaker than sigma bonds.

In both cases, the overlapping results in a lower energy state, where the valence electrons with opposite spins pair up in the shared orbital, leading to the formation of a stable covalent bond.

4.0s-s Overlap 

s-s Overlap refers to the axial (head-on) overlap between two half-filled s-orbitals of two atoms. This overlap occurs when both atoms have unpaired electrons with opposite spins in their respective s-orbitals.

(Formation of H₂ Molecule)

• The electronic configuration of a hydrogen atom is 1s¹.
• Each hydrogen atom has one half-filled 1s orbital containing a single unpaired electron.
• During the formation of an H₂ molecule, the 1s orbital of one hydrogen atom overlaps with the 1s orbital of another hydrogen atom.
• The unpaired electrons with opposite spins pair up, resulting in a covalent bond.

Features of s–s Overlap:

• This type of overlap is called S-S Overlap.
• The resulting bond is a sigma (σ) bond, as it is formed by axial overlap along the internuclear axis.
• The bond formed is strong and stable due to the direct overlapping of orbitals.

5.0p–p Overlap 

P–P Overlap refers to the axial (head-on) overlap between two half-filled p-orbitals of two atoms, each containing unpaired electrons with opposite spins.

(Formation of F₂ Molecule)

• The atomic number of fluorine is 9, and its electronic configuration is:
  1s², 2s², 2pₓ², 2pᵧ², 2pz¹
• Each fluorine atom has one half-filled 2p_z orbital, which contains a single unpaired electron.
• During the formation of an F₂ molecule, the 2p_z orbital of one fluorine atom overlaps coaxially with the 2p_z orbital of another fluorine atom.
• The unpaired electrons with opposite spins pair up to form a covalent sigma (σ) bond.

Features of p–p Overlap:

• The bond formed is called a P–P sigma bond (σ).
• P orbitals have a dumbbell shape, so the overlap is directional, giving the bond directional character.
• When one lobe of a p orbital overlaps, the other lobe shrinks, as electron density is concentrated in the bonding region.

6.0s–p Overlap 

The overlap between a half-filled s orbital of one atom and a half-filled p orbital of another atom, each containing electrons with opposite spins, is known as s–p overlap. A typical example of this is the formation of the hydrogen fluoride (HF) molecule.

Formation of the HF Molecule

• Hydrogen (Z = 1) has the electronic configuration: 1s¹
• Fluorine (Z = 9) has the electronic configuration: 1s², 2s², 2p²ₓ, 2p²ᵧ, 2p¹𝓏

During the formation of HF:

• The half-filled 1s orbital of hydrogen overlaps coaxially with the half-filled 2p𝓏 orbital of fluorine, resulting in a sigma (σ) bond through s–p overlap.

Nature of the HF Bond

• The bond formed between H and F is a polar covalent bond.
• This is due to the large electronegativity difference:
• • Fluorine: 4.0
  • Hydrogen: 2.1
• The shared pair of electrons is pulled closer to the fluorine atom, making:
• • Fluorine partially negative (δ⁻)
  • Hydrogen partially positive (δ⁺)

Features of s–p Overlap Orbital Types:

• Involves a half-filled s orbital and a half-filled p orbital.
•  Axial (head-on) overlap forming a sigma (σ) bond.
•  Only σ bond is formed (no π bond). Strong bond if the overlap is significant.
•  Occurs in polar molecules (e.g., HF, HCl) due to the electronegativity difference.
•  p orbital gives direction to the bond.

7.0Types of Orbital Overlap Based on Phase

 Positive Overlap:

• Occurs when the phases of the overlapping atomic orbitals are the same.
• This constructive interference leads to bond formation (usually a sigma or pi bond).
• Example: 1s–1s, 1s–2p (same phase lobes).

 Negative Overlap:

• Happens when the phases of the overlapping orbitals are opposite.
• This destructive interference results in cancellation, and no bond is formed.
• Instead, it may lead to an antibonding interaction.

 Zero Overlap:

• Takes place when the two orbitals do not overlap at all due to improper orientation.
• As a result, there is no interaction or bond formation.