1
Chapter 9: Covalent Bonding: Orbitals
9.1
The valence orbitals are the orbitals associated with the highest principal quantum level that
contains electrons on a given atom. Hybridization is a modification of the localized electron model
to account for the observation that atoms often seem to use special atomic orbitals in forming
molecules. sp3 hybridization gives tetrahedral set of orbitals, that are identical shape, each one
having a large lobe and a small lobe. Whenever a set of equivalent tetrahedral atomic orbitals is
required by an atom, this model assumes that the atom adopts a set of sp3 orbitals; the atom
becomes sp3 hybridized. It is really not surprising that an atom in a molecule might adopt a
different set of atomic orbitals (called hybrid orbitals) from those it has in the free state.
The sp2 hybridization gives a trigonal planar arrangement of atomic orbitals. There are
different types of bonds that can be formed during this hybridization. Sigma (σ) bond is a covalent
bond in which the electron pair is shared in an area centered in a line running between the atoms.
Another type of bond that can be formed is the pi bond. Pi (π) bond is a covalent bond in which
parallel p orbitals share an electron pair occupying the space above and below the line joining the
atoms. A double bond always consists of one σ bond, where the electron pair is located directly
between the atoms, and one π bond, where the shared pair occupies the space above and below the
σ bond.
Whenever an atom is surrounded by three effective pairs, a set of sp2 hybrid orbitals is required.
Example:
The sp hybridization gives linear planar arrangement of atomic orbitals. Two effective pairs around
an atom will require sp hybridization of that atom.
Example:
2
The dsp3 hybridization gives octahedral arrangement of atomic orbitals. It is formed from
one d orbital, one s orbital and three p orbitals. Note that a set of five effective pairs around a given
atom always requires a trigonal bipyramidal arrangement, which in turn requires dsp3 hybridization
of that atom.
Example:
The d2sp3 hybridization in which two d orbitals, one s orbital, and three p orbitals are
combined. Note that six electron pairs around an atom are always arranged into an octahedral shape
and require d2sp3 hybridization of the atom.
Example:
The description of a molecule using the localized electron model involves three distinct steps:
• Draw the Lewis structures(s)
• Determine the arrangement of electron pairs using VSEPR model
• Specify the hybrid orbitals needed to accommodate the electron pairs
Examples:
3
9.2
Molecular Orbital Model (MO) model is a model that regards a molecule as a collection of
nuclei and electrons, where the electrons are assumed to occupy orbitals much as they do in atoms,
but having the orbitals extend over the entire molecule. In this model the electrons are assumed to
be delocalized rather than always located between a given pair of atoms. Bonding will result if the
molecule has lower energy than the separated atoms.
A bonding molecular orbital is lower energy than the atomic orbitals of which it is
composed. An antibonding molecular molecule is higher in energy than the atomic orbitals of
which it is composed.
In the H2 molecule the molecular orbitals are represented as follows:
MO1= 1σΑ + 1 σ Β
MO2= 1 σ Α − 1 σ Β
Bond order = (number of bonding electrons-number of antibonding electrons)/2
NOTE:
A larger bond order means greater bond strength.
9.3
Homonuclear diatomic molecules are those composed of two identical atoms. In order to
participate in molecular orbitals, atomic orbitals must overlap in space. This means that only the
valence orbitals of the atoms contribute significantly to the molecular orbitals of a particular
molecule.
Let's consider the molecular orbitals from the head-on overlap. Note that the electrons in the
bonding MO are, as expected, concentrated between the nuclei, and the electrons in the antibonding
MO are concentrated outside the area between the area between the two nuclei. Also, both of these
MOs are σ molecular orbitals. He p orbitals that overlap in a parallel fashion also produce bonding
and antibonding orbitals. Since the electron probability lies above and below the line between the
nuclei, both orbitals are π molecular orbitals.
4
Paramagnetism is a type of induced magnetism, associated with unpaired electrons, that causes a
substance to be attracted into the inducing magnetic field. Diamagnetism causes the substance to be
repelled from the inducing magnetic field. Studies have shown the paramagnetism is associated
with unpaired electrons and diamagnetism is associated with paired electrons.
Examples:
9.4
Bonding of Heteronuclear (different atoms) in diatomic molecules is a special case that
involves molecules containing atoms adjacent to each other on the periodic table. Since the atoms
involved in such a molecule are so similar, we can use the molecular orbital diagram for
homonuclear molecules. When the two atoms of a diatomic molecule are very different, he energylevel diagram for homonuclear molecules can no longer be used. A new diagram must be devised
for each molecule. You can determine the polarity by electron probability distribution in the
bonding molecular orbital of the (example) HF molecule. Note the greater electron density is close
to the fluorine atom.
Examples: