• Valence Shell Electrons wish to minimize the Electron Pair Repulsions between them.
• The repulsions between “independent regions of elecron density” around an atom produce “electron pair geometries”.
• Electron pair geometries (number of regions) include: linear(2), trigonal planar(3), tetrahedral(4), trigonal bipyramid(5) and octahedral(6).
• The electron pair geometries determine the approximate bond angles for a structure.
• Electron pair geometries are further broken down into “molecular geometry” subclasses. The names for these subclasses is determined by the 3-D arrangement of atoms.
• A multiple bond and a lone pair each count as one independent region of electron density.
• Inner core and d-electrons are not part of VSEPR.
• Lone pairs are larger than bond pairs and have larger VSEPR effects. They adopt positions with the largest bond angles.
• Larger central atoms have reduced VSEPR effects and smaller than expected bond angles.

• Polar bonds may or may not create polar molecules depending on the molecular geometry.
• Polar bonds can be treated a vector forces. If the polar bonds are equal and opposite they cancel out each others effect.
• CO₂ had polar covalent bonds but is overall non-polar.
• Determine the molecular geometry and all the polar bond present. Check to see if all polar bonds are equal in magnitude and balanced geometrically. If they are, the molecule is NONPOLAR overall. If not the molecule is POLAR.

• The structure of methane, CH₄, has equal bond lengths and 109.5° H-C-H angles.
• Yet the carbon has an electron configuration of [He]2s²2p². How does this atom form 4 equivalent C-H bonds?
• Pauling created a new bonding model in which atomic orbitals are mixed to create “hybrids” to give the correct geometry.
• Hybrid orbitals
  • have different shapes than unhybridized atomic orbitals
  • create sigma bonds or lone pairs only
  • have the same number of units as the number of atomic orbitals used to create them: s+p+p+p = 4 x sp³
  • of each type are identical in size, shape and energy
  • have discrete bond angles between them (see table in text)

• Single or “sigma” bonds are created from atomic orbital overlaps or hybrid orbital overlaps that are “end on”.
• p-orbitals that overlap side-to-side are called pi(’) bonds and crreate multiple bonds.
• Each multiple bond contains one sigma bond; the rest are pi.
• Example: H₂C=CH₂
  • Each carbon has three electron-density regions.
  • This makes each carbon trigonal planar and 3 x sp².
  • The hybrid orbitals create the C-H and one C-C bond.
  • Carbon has 3 x p orbitals; one is not used in hybridization.
  • Unhybridized p orbitals on each carbon overlap to create a ’-bond.

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