Bonding and Molecular Structure: Orbital Hybridization and Molecular Orbitals

Molecular Orbital Theory

  • Valence electrons are delocalized and exist in molecular orbitals spread over the entire molecule.

Valence Bond Theory

  • Valence electrons are localized between atoms or exist as lone pairs.
  • Orbitals overlap to form bonds, holding two electrons of opposite spin.

Valence Bond Terminology

  • Overlap: Two orbitals occupying the same space.
  • lp: Lone pair of electrons (non-bonding).
  • bp: Bonding pair of electrons (result of orbital overlap).
  • Central atom: The atom of concern in a molecule
  • Hybridization: The linear combination of atomic orbitals.
  • Hybrid orbital: Bonding orbitals from the mixing of atomic orbitals.
  • σ bond (sigma bond): Overlap of orbitals along the bond axis.
  • π bond (pi bond): Overlap of orbitals above and below the bond axis.
  • Single bond: One σ bond.
  • Double bond: One σ bond and one π bond.
  • Triple bond: One σ bond and two π bonds.

Hybrid Orbitals

  • Explain electron-pair bonds and molecular geometries via valence-bond theory.
  • Common types: sp, sp2, sp3, sp3d, sp3d2.

sp^3 Hybrid Orbitals

  • Steric number: 4
  • Electron-pair geometry: Tetrahedron
  • Bond angle: Near 109.5°
  • Four sp^3 hybrid orbitals form a tetrahedron around a central atom to minimize repulsion.

sp^2 Hybrid Orbitals

  • Steric number: 3
  • Electron-pair geometry: trigonal planar
  • Bond angle: Near 120°
  • The 3 sp^2 hybrid orbitals lie in the same plane, 120º apart each other with trigonal planar orbital geometry.

sp Hybrid Orbitals

  • Steric number: 2
  • Electron-pair geometry: linear
  • Bond angle: 180°
  • sp hybrid orbitals have identical shape but the two large lobes point in opposite directions.

Hybridization involving d orbitals

  • sp^3d (steric number 5) hybrid orbitals form trigonal bipyramidal orbital geometry.
  • Atoms in the third period and beyond can use d orbitals to form hybrid orbitals.

Multiple Bonds in Ethylene (C2H4)

  • Steric number of carbon is 3; sp^2 hybridization.
  • Each C-H bond is formed by the overlap of a sp^2 hybrid orbitals with a 1s atomic orbital on H.
  • One C-C bond forms from overlap of two sp^2 hybrid orbitals.
  • One C-C bond forms from overlap of two unhybridized 2p orbitals on each carbon by side-by-side fashion.

Sigma (σ) Bond Formation

  • Electron density lies along the bond axis.
  • Formed by overlap of two s orbitals or two p orbitals.

Pi (π) Bonding

  • Bonding occurs above and below the bond axis.
  • Electrons occupy space above and below the nuclei.

Sigma (σ) and Pi (π) Bonds

  • σ bond: High electron density distributed symmetrically along the bond axis; all single bonds are σ bonds.
  • π bond: High electron density between nuclei, concentrated above and below the bond axis; a double bond contains one σ bond and one π bond.

Formation of a Triple Bond (Acetylene)

  • Steric number of carbon is 2; sp hybridization.
  • Each C-H bond is formed by the overlap of a sp hybrid orbitals with a 1s atomic orbital on H.
  • One C-C bond (σ bond) forms from overlap of two sp hybrid orbitals.
  • Two C-C bonds (π bonds) form from overlap of two unhybridized 2p orbitals on each carbon by side-by-side fashion. The two π bonds are perpendicular to each other.

General Trends

  • π bonding requires the presence of a second-row element due to valence orbital size.
  • Carbon, nitrogen, and oxygen readily form π bonds.
  • Silicon forms σ (single) bonds with tetrahedral geometry.

Determining Bonding Picture

  • Determine the Lewis structure.
  • Use the Lewis structure to determine steric numbers and hybridizations.
  • Construct a σ bond framework.
  • Add in the π bonds.
  • Predict bond angles.