Valence bond theory
Electrons are not simply dots
And bonds are not sticks
Learning objectives
 Describe principles of valence bond theory
 Predict hybridization of orbitals based on
Lewis dot structures and electronic
geometry
 Describe difference between sigma and pi
bonding
Taking it to the next level:
acknowledging orbitals
 VSEPR is quite successful in predicting
molecular shapes based on the simplistic
Lewis dot approach
 But our understanding of the atom has the
electrons occupying atomic orbitals
 How do we reconcile the observed shapes
of molecules with the atomic orbital picture
of atoms
Valence bond theory
 Valence bond theory is the simplest
approach to an orbital picture of covalent
bonds
 Each covalent bond is formed by an overlap
of atomic orbitals from each atom
 The individual orbital identity is retained
 The bond strength is proportional to the
amount of orbital overlap
Overlap of two 1s orbitals in H2
 Overlap of two 2p orbitals directed along the bond
axis (sigma bond)
 Overlap of p and s orbitals
Problems with tetrahedral bonds
 In CH4 the bonds are all equivalent and at
angles of 109.5°
 The 2p orbitals in C are at 90° - far from
optimum for overlap
 The ground state configuration is 2s22p2
 Reconcile these facts with the known
structure
Hybridization
 The wave mechanics permits mixing of the
atomic orbital set to produce “hybrid” orbitals
 Hybridization alters the shape and energy of
the original
 In the case of C, the differences between
the 2s and 2p are smoothed out and a
homogeneous collection of four sp3 hybrid
orbitals is produced
sp3 hybridization
 Formally, one of the 2s
electrons is promoted
to the empty 2p orbital
(an energy cost, which
is repaid on bond
formation)
 The four basis orbitals
are then “hybridized” to
yield the set of four sp3
Tetrahedral directions and sp3
hybrids
Valence bond picture of CH4
 Each C sp3 hybrid contains one electron
 Each H 1s contains one electron
Lone pairs occupy sp3 hybrid orbitals
 Valence bond picture of the tetrahedral electronic
geometry provides same results for the molecules
with lone pairs
Notes on hybridization
 The total number of orbitals is unchanged
 Four atomic orbitals (s + 3 x p) give four hybrid
orbitals (4 x sp3)
 The electron capacity remains unchanged
 There is one hybridization scheme for each
of the five electronic geometries
 The same hybridization scheme is always
used for a given electronic geometry
sp hybridization for linear geometry
 One s and one p orbital
sp2 hybridization for trigonal planar
 One s and two p
orbitals
Sigma and pi bonding
 The hybridized orbitals describe the
electronic geometry: bonds along the
internuclear axes (sigma bonds)
 The “unused” p orbitals overlap in a parallel
arrangement above and below the
internuclear axis (pi bonds)
Comparison of pi and sigma bonding
Pi bonding accounts for bond
multiplicity
 Two unused p orbitals in sp hybrid (linear
geometry)
 Two pi bonds
 N≡N triple bond (one sigma, two pi)
 One unused p orbital in sp2 hybrid (trigonal
planar geometry
 One pi bond
 C=C double bond (one sigma, one pi)
Valence bond picture of ethylene
H2C=CH2
 Sigma bonds between C and H (blue/red) and C
and C (blue)
 Six electrons around C
 Pi bond between C and C (green)
 Two electrons around C
Valence bond picture of acetylene
HC≡CH
 Sigma bonds between C and H (red and blue) and
C and C (blue)
 4 electrons around C
 Two pi bonds between C and C (green)
 4 electrons around C
Beyond coordination number 4
 Invoke empty d orbitals
(impossible for second row
elements)
 One d orbital for trigonal
bipyramidal
 Two d orbitals for octahedral
 Number of orbitals in
hybrid always equals
number of charge clouds
Trigonal bipyramid – sp3d
Octahedral –sp3d2
Shortcomings of valence bond
 The orbitals still maintain atomic identity
 Bonds are limited to two atoms
 Cannot accommodate the concept of
delocalized electrons – bonds covering
more than two atoms
 Problems with magnetic and spectroscopic
properties