Valence Bond Theory

Valence Bond Theory (VBT) describes how covalent bonds form when atomic valence orbitals overlap in-phase, creating a region of increased electron density directly between the bonded nuclei. This fundamental concept underpins our understanding of molecular structure and reactivity. The orbitals involved in this overlap can be s, p, or sp^x hybrid orbitals.

The Energy Landscape of Bond Formation

Consider the simplest case: the formation of a hydrogen molecule (H₂) from two individual hydrogen atoms (H₁ and H₂), each with a 1s orbital in the same phase.

• At Great Distances (High Energy): When the atoms are far apart (typically more than 4 Å), their orbitals exhibit minimal interaction. There's little attractive force, resulting in a high-energy state. This is represented on the right side of a potential energy diagram.
• Approaching Equilibrium (Lower Energy): As the atoms draw closer, the electrons from each atom begin to experience the attractive pull of the opposing nucleus. This mutual attraction stabilizes the system, causing the potential energy to decrease.
• Optimal Overlap (Lowest Energy): The equilibrium bond length (r_e) is the precise distance where the attractive forces between electrons and nuclei perfectly balance the repulsive forces between the positively charged nuclei. At this optimal distance, the system achieves its lowest possible energy configuration, and the H-H sigma (σ) bond forms from the effective overlap of the 1s orbitals.
• Too Close (High Energy): If the atoms are forced to come too close together (left side of the diagram), the strong repulsion between their nuclei dominates, causing the energy to rise sharply.

Sigma (σ) and Pi (π) Bonds: Distinguishing Overlap

When describing the structures of alkanes, alkenes, alkynes, and other functional groups, it's crucial to identify the types of valence orbitals that overlap to form specific bonds:

• Sigma (σ) Bonds: These are formed by the direct, head-on (axial) in-phase overlap of s-like orbitals (such as 1s, 2s, or sp^x hybrid orbitals). Every single covalent bond is a σ bond. It's easy to remember: s for sigma (σ).
• Pi (π) Bonds: These result from the side-by-side (lateral) in-phase overlap of unhybridized p-type orbitals. Pi bonds are found in multiple bonds (double and triple bonds), in addition to a sigma bond. It's easy to remember: p for pi (π).

Bond Characteristics: Length and Strength

Every covalent bond possesses a characteristic length and strength, which are crucial determinants of molecular properties and reactivity:

• Bond Length: This is the average distance between the nuclei of two bonded atoms.
  • Typical C-C single bond: ~150 pm
  • Typical C=C double bond: ~130 pm
  • Typical C=O double bond: ~120 pm
  • Typical C-H bond: 100-110 pm
• Bond Strength (Bond Dissociation Energy): This is the energy required to break a specific bond. Stronger bonds require more energy to break.
  • C-I bond (e.g., in thyroid hormone): ~234 kJ/mol
  • Typical C-H bond: ~410 kJ/mol
  • C≡C triple bond: >800 kJ/mol

Typical Bond Energies and Lengths