Dispense sugli orbitali molecolari corso di chimica inorganica 1

This document introduces qualitative molecular orbital theory, essential for understanding chemical bonding in inorganic chemistry:

• H₂⁺ Ion (LCAO): Explains the formation of bonding (φ₁) and antibonding (φ₂) molecular orbitals from the linear combination of 1s atomic orbitals (Xa, Xb). Bonding results from concordant sign combination (Xa+Xb), leading to lower energy and increased electron density between nuclei. Antibonding results from discordant sign combination (Xa-Xb), creating a nodal plane perpendicular to the internuclear axis and higher energy.
• Sigma (σ) Bonds:
  • Formed by head-on overlap of atomic orbitals (s-s, s-p, p-p along internuclear axis).
  • Characteristic: absence of nodal planes containing or intersecting the internuclear axis.
  • Bonding σ orbitals increase electron density between nuclei; antibonding σ* orbitals have a nodal plane between nuclei, destabilizing the bond.
• Pi (π) Bonds:
  • Formed by sideways overlap of parallel p orbitals.
  • Characteristic: presence of a nodal plane containing the internuclear axis.
  • Bonding π orbitals increase electron density above and below the internuclear axis; antibonding π* orbitals have an additional nodal plane perpendicular to the internuclear axis, further destabilizing the bond.
• Orbital Overlap and Stability: The energy difference between bonding and antibonding molecular orbitals is proportional to the extent of atomic orbital overlap (Sab integral). Greater overlap leads to stronger bonds.
• D-Orbital Involvement: Discusses σ and π bond formation involving d orbitals, particularly in octahedral complexes. For example, dx²-y² and dz² participate in σ bonds, while dxy, dxz, dyz can participate in π bonds.
• Molecular Orbital Diagrams (Diatomic Molecules):
  • O₂: 2s and 2px orbitals do not mix significantly due to large energy difference. Paramagnetic due to two unpaired electrons in degenerate π* orbitals (bond order = 2).
  • F₂ and Ne₂: Similar diagrams to O₂. F₂ has a single bond (bond order = 1), Ne₂ has no bond (bond order = 0).
  • N₂, B₂, C₂: Exhibit significant s-p mixing (hybridization of AOs) due to smaller energy differences between 2s and 2p orbitals. This alters the MO energy ordering, placing π orbitals below σ(2p) orbitals.
    • N₂: Bond order = 3 (triple bond), consistent with high dissociation energy.
    • B₂: Paramagnetic due to two unpaired electrons in degenerate 1π orbitals (bond order = 1, reinterpreted based on experimental data).
    • C₂: Diamagnetic (bond order = 2, reinterpreted based on experimental data, likely two π bonds).
  • CO and NO (Heteronuclear): AO energies differ between atoms. Explains CO's ability to form metal complexes via σ-donation from the HOMO (5σ*) and π-retrodonation to the LUMO (2π*). NO exhibits similar behavior.

This comprehensive overview helps understand bond types, molecular properties, and chemical reactivity based on molecular orbital theory.