Chemical Bonding and VSEPR Theory
Valence shell electron pair repulsion theory (VSEPR) enables us to identify the shapes of
individual molecules, based upon electron-pair repulsions. VSEPR theory is based on the idea
that the geometry of a molecule or polyatomic ion is determined primarily by repulsion among
the pairs of electrons associated with a central atom. The pairs of electrons may be bonding pairs
(BP) or nonbonding pairs (also called lone pairs, LP). Only valence electrons of the central atom
influence the molecular shape in a meaningful way. Three types of repulsion can occur between
pairs of electrons in a molecule:
1. LP-LP repulsions
2. LP-BP repulsions
3. BP-BP repulsions
A molecule must minimize repulsions to be stable. The theory states that repulsion becomes zero
at ~115-120°. When repulsion cannot be avoided, the weaker repulsion (i.e. the one that causes
the smallest deviation from the ideal shape) is preferred. The lone pair-lone pair (LP-LP)
repulsion is stronger than the lone pair-bonding pair (LP-BP) repulsion, which in turn is stronger
than the bonding pair-bonding pair (BP-BP) repulsion. Hence, the weaker BP-BP repulsion is
preferred over BP-LP repulsion, which in turn is preferred over LP-BP repulsion.
The figure above illustrates trigonal planar (3 balloons), tetrahedral (4 balloons), trigonal
bipyramid (5 balloons), and octahedral (6 balloons) geometries. The balloons represent
electron domains (LP or BP).
Example
trigonal planar
geometry
Example
tetrahedral
geometry
The ideal geometry associated with 4 electron domains is that of a tetrahedron. If all four
electron domains involve bonding pairs (BP), that is, if the central atom is bonded to four
other atoms, the molecular shape is tetrahedral. If the tetrahedron includes one LP (NH3),
the molecular shape is that of a trigonal pyramid. The LP compresses the ideal bond
angle down from 109.5˚ to 107˚. If the tetrahedron includes two LPs (H2O), the
molecular shape is angled or bent. The LP compresses the ideal bond angle down from
109.5˚ to 104.5˚.
Phosphorus pentachloride, PCl5, possesses a trigonal bipyramid shape. What molecular
geometry results in SF4 (4 BP, 1 LP), ClF3 (3 BP, 2 LP), or I3-1 (2 BP, 3 LP)?
Sulfur hexfluoride, SF6, possesses an octahedral geometry. What molecular geometry
results in SbCl5-2 (5 BP, 1 LP), or XeF4 (4 BP, 2 LP)?
Chemical Bonding and Valence Bond Theory
Valence bond theory features the overlapping of atomic orbitals when atoms form a
chemical bond. Atomic orbitals from different atoms may overlap in two distinct ways,
producing sigma bonds () and pi () bonds. Sigma bonds, the stronger of the two types
of bonding, occur when two half-filled orbitals overlap head-on along the internuclear
line of the bonded atoms. Pi bonds occur when two orbitals overlap sideways. For
example, a bond between two s-orbital electrons is a sigma bond, because two spheres are
always coaxial. In terms of bond order, single bonds have one bond, double bonds
consist of one bond and one bond, and triple bonds contain one bond and two 
bonds.
Head-on overlap - a sigma () bond
Hybridized Orbitals
The basic idea of valence bond theory is that a covalent bond is formed by the overlap of
atomic orbitals. The two electrons of paired spin, which are shared by two bonded atoms,
lie in an atomic orbital of each of the two atoms. The greater the degree of overlap of the
atomic orbitals, the greater will be the degree of sharing and the stronger will be the
covalent bond between them. Often, however, the geometry of these orbitals is such that
effective overlap cannot occur in the known geometry of the molecule. Under these
circumstances, the atomic orbitals on an atom can reconfigure themselves into a different
configuration, and the reconfigured orbitals are said to be hybridized. Hybridization of
atomic orbitals was first proposed by Linus Pauling. The hybridization (mixing) process
affects the spatial arrangement and energies of the hybrid orbitals. Hybrid orbitals are
simple to envision, and they can be used to explain the observed shapes and bond orders
for p-block molecules. They also make the distinction between sigma and pi bonding
easy to understand.
Methane, CH4, makes use of sp3 hybridization, producing a tetrahedral molecule.
Ethylene, C2H4, makes use of sp2 hybridization in both carbon atoms. The sp2 hybrid
orbitals on each carbon atom point towards the corners of an equilateral triangle. The
carbon atoms are bonded together by the sigma-overlap of two sp2 hybridized orbitals.
The remaining two sp2 hybrid orbital on each carbon overlap the half-filled 1s atomic
orbitals of two hydrogen atoms. The remaining unhybridized p-orbital on one carbon
atom overlaps sideways with the unhybridized p-orbital on the other carbon atom,
producing a  bond.
Acetylene, C2H2, employs sp-hybridization in its carbon atoms. The sp hybrid orbitals on
each carbon atom point in opposite directions. The carbon atoms are bonded together by
the sigma-overlap of two sp hybridized orbitals. The remaining sp hybrid orbital on each
carbon overlaps a half-filled 1s atomic orbital of a hydrogen atom. The pair of
unhybridized p-orbitals on one carbon atom overlap sideways with those on the other
carbon atom, producing a pair of  bonds.
Explaining Trigonal Bipyramidal and Octahedral Molecular Geometries
Trigonal bipyramidal geometry is explained by hybridizing an s-orbital with three porbitals and a d-orbital. This is known as sp3d hybridization.
Octahedral geometry is explained by hybridizing an s-orbital with three p-orbitals and
two d-orbitals. This is known as sp3d2 hybridization.
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