Valence Bond Theory (VBT)

Valence Bond Theory (VBT) is a fundamental theory in chemistry that explains the formation of covalent bonds between atoms by the overlap of atomic orbitals. It provides a qualitative picture of molecular structure and is instrumental in explaining the shapes and properties of molecules. Developed by Linus Pauling in the early 20th century, VBT is one of the two major theories used for this purpose, the other being Molecular Orbital Theory (MOT).

1.0Necessity of Valence Bond (VB) Theory in Chemistry

The Valence Shell Electron Pair Repulsion (VSEPR) theory provides valuable insights into the geometry of simple molecules by considering the repulsion between electron pairs around a central atom. However, VSEPR theory primarily offers geometric predictions without delving into the theoretical foundations of molecular structure, and its application is somewhat limited to simpler systems. To address these limitations and provide a more comprehensive understanding of molecular architecture and bonding interactions, two advanced theories based on quantum mechanical principles have been developed: Valence Bond (VB) Theory and the Molecular Orbital Theory (MOT).

2.0Concept of Valence Bond Theory (VBT)

Valence Bond Theory (VBT) primarily shows that covalent bonds are formed through the overlapping of half-filled atomic orbitals from two bonding atoms, which allows paired electrons with opposite spins to occupy the newly overlapped space. 

This theory distinguishes between two types of bonds based on the orientation of this overlap. Sigma (σ) bonds are the result of the end-to-end overlapping of orbitals like s-s, s-p, or p-p, and are characterized by their cylindrical symmetry around the bond axis, which often contributes to the formation of single bonds in molecules. In contrast, pi (π) bonds arise from the side-to-side overlap of p-orbitals, adding additional bonding interactions in molecules that feature double or triple bonds, thus providing extra layers of bonding beyond the initial sigma bond.

3.0Main points of Valence Bond Theory

(a) Covalent Bond Formation: 

Covalent bonds form through the overlap of half-filled valence shell orbitals of bonding atoms. This overlap allows electrons with opposite spins from each atom to pair up, lowering the system's energy and creating a stable bond.

(b) Directional Overlap: 

Orbitals overlap most effectively when they approach each other from directions that maximize their spatial overlap. This maximization leads to stronger bonding interactions and defines the molecular geometry.

(c) Directional Nature of Covalent Bonds: 

The covalent bond has a directional character because the shape of the atomic orbitals dictates the orientation of the overlap. This directional nature is crucial for determining the molecular geometry and the spatial arrangement of atoms within a molecule.

(d) Bond Strength and Overlap: 

The strength of a covalent bond is directly proportional to the extent of the orbital overlap. Greater overlap allows for more shared electron density between the atoms, which results in a stronger bond.

4.0Factors Affecting Overlap

• Orbital Size: Larger orbitals tend to overlap less effectively due to their diffused nature, while smaller orbitals can overlap more compactly and effectively.
• Orbital Shape: The shape of the orbitals involved in bonding affects how well they can overlap. For example, p orbitals can overlap side-to-side to form pi bonds in addition to the end-to-end sigma bond formation.
• Orbital Orientation: The orientation of orbitals with respect to each other is critical. Proper alignment achieves optimal overlap. Misalignment can significantly reduce overlap and, thus, bond strength.
• Electron Density Distribution: The distribution of electron density within the orbitals influences overlap. Orbitals with greater electron density in the bonding region can form stronger bonds.

5.0Nature of Orbitals and Their Impact on Bond Strength

1. Effect of Principal Quantum Number (n) on Bond Strength:

• As the principal quantum number (n) increases, the orbitals become larger and more diffused. This leads to a decrease in bond strength because the overlap between orbitals becomes less effective. The bond strength generally follows the sequence:

                     1s - 1s > 1s - 2s > 2s - 2s > 2s - 3s > 3s - 3s

• This sequence illustrates that bonds involving lower energy level orbitals (with smaller n values) are stronger due to more effective overlap.

2. Directionality of Orbitals:

• Orbitals like p, d, and f are directional, which allows for more effective overlapping compared to the spherical s-orbitals that are non-directional.
• A bond formed by the overlapping of directional orbitals (such as p-p) is typically stronger than one formed by non-directional s-orbitals, assuming the principal quantum number (n) remains the same. The order of bond strength based on orbital type is:

p-p > s-p > s-s

3. Comparison Within the Same Principal Quantum Number:

• When the principal quantum number is the same, the effectiveness of overlap follows the order:

2p - 2p > 2s - 2p > 2s - 2s

• This order underscores that the more directional p-orbitals, even when paired with less directional s-orbitals, facilitate better overlap than s-s interactions.

6.0Nature of Overlapping

1. Types of Overlapping:

• Co-axial Overlapping: This type involves orbitals overlapping along the same axis, which maximizes electron density between the nuclei and creates a strong sigma (σ) bond.
• Collateral Overlapping: In this type, p-orbitals overlap side-to-side. This overlapping is less extensive compared to co-axial overlapping and results in the formation of pi (π) bonds.

2. Order of Strength in Co-axial Overlapping:

• The effectiveness of co-axial overlap, which leads to the formation of sigma bonds, also follows the order:

p-p > s-p > s-s

• This ranking illustrates that the more directional the orbitals involved in the co-axial overlap, the stronger the resulting sigma bond.

3. Types of Bonds Formed by Overlapping:

Sigma (σ) Bonds:

• Formed by the head-on (co-axial) overlapping of orbitals.
• Characterized by cylindrical symmetry around the bond axis, these bonds are generally stronger and serve as the primary structural bond in molecules.

Pi (π) Bonds:

• Formed by the side-to-side (collateral) overlapping of primarily p-orbitals.
• These bonds are usually found in conjunction with sigma bonds in double and triple bond scenarios, contributing less to bond strength but essential for the overall bonding structure.

7.0Hybridization

To explain the shapes of molecules that couldn't be described by the simple overlap of atomic orbitals, the concept of hybridization was introduced. This involves the mixing of different types of orbitals (s, p, d) on the same atom to form new orbitals of equivalent energy.

Read in detail: Hybridization

8.0Limitations of Valence Bond Theory

• Less effective at predicting properties of molecules with delocalized electrons, such as those found in aromatic compounds.
• Does not account for the energy differences between orbitals as well as MOT.