Chris Khan 2007
 Valence Bond Theory assumes that electrons in a molecule occupy atomic orbitals of the individual atoms. Molecular Orbital
Theory assumes the formation of molecular orbitals from the atomic orbitals.
 To explain bonding, VB uses hypothetical hybrid orbitals, atomic orbitals obtained when two or more nonequivalent orbitals of
the same atom combine in preparation for covalent bond formation. Hybridization is the term applied to the mixing of atomic
orbitals in an atom to generate a set of hybrid orbitals. When considering hybridization, we usually consider hybridization of the
central atom, or just one atom. In CH4, we can represent C as [He]2s22p2 showing four valence electrons. Therefore, the
hybridization of C is sp3.
 NH3 is also sp3 hybridized because the four valence electrons fit in the same orbitals (the extra one compared to methane pairs in
a p-orbital).
 BeCl2’s Beryllium atom is sp hybridized because there are two valence electrons, both of which should pair into the 2s orbital but
do not (because of hybridization) and split up: one in the 2s and one in the 2p orbital.
 Boron in BF3 is sp2 hybridized because the three valence electrons split up into empty orbitals, one in each. One goes in the 2s,
and two go into their own orbitals in the 2p shell.
 Hybridization is used only to describe covalent bonding. It is the mixing of at least two nonequivalent atomic orbitals. The
number of hybrid orbitals generated is equal to the number of pure atomic orbitals that participate in the hybridization process. It
requires an input of energy.
 Procedure for Hybridizing Atomic Orbitals:
o Draw the Lewis Structure
o Predict the arrangement of electron pairs (bonding and lone pairs) using VSEPR
o Deduce hybridization by matching arrangement of electron pairs with those of the hybrid orbitals
 d-orbitals also hybridize when s and p-orbitals are full. SF6 is an example: the S atom has six valence electrons, and therefore
hybridizes as sp3d2, leaving only one electron in each orbital and three 3d orbitals empty.
 There are sigma bonds (σ) and pi bonds (π). Sigma bonds are covalent bonds formed by orbitals overlapping end to end, with
the electron density concentrated between the nuclei of the bonding atoms. Pi bonds are covalent bonds formed by sideways
overlapping orbitals with electron density concentrated above and below the plane of the nuclei of the bonding atoms. All single
bonds are sigma bonds. In a double or triple bond, the first bond is a sigma bond and the 2nd and 3rd bonds are pi bonds.
 Examples:
o Determine the hybridization state of Be in BeH2, Al in AlI3, and P in PF3. First draw the Lewis Structures of each
molecule. BeH2 is H-Be-H, so since there are two areas of electron density, the Be is sp-hybridized. AlI3’s Lewis
Structure is Al with three I’s singlely bonded, each having three lone pairs. Since there are three areas of electron
density around the Al, singlely bonded, it is sp 2 hybridized. PF3’s structure is similar, BUT, there is ANOTHER lone
pair, this time on the P atom, making four total areas of electron density. This makes PF3 sp3 hybridized.
o What is the hybridization of P in PBr5. The Lewis Structure is a P, singlely bonded to five Bromine atoms, each with
three lone pairs. Since there are five areas of electron density around the P, we would think sp4, BUT, p has only three
orbitals, so we have a sp3d.
o Describe the bonding in formaldehyde (CH2O). The Lewis Structure is a Carbon atom, singlely bonded to two hydrogen
atoms and doublely bonded to an oxygen atom, with two lone pairs. The O atom is sp2 hybridized, as is the C (3 areas of
electron density). This molecule is trigonal planar because three ligands around the Carbon. There are three sigma
bonds and one pi bond in this molecule.
 Molecular Orbital Theory describes covalent bonds in terms of molecular orbitals, which result from interaction of the atomic
orbitals of the bonding atoms and are associated with the entire molecule.
 A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An
antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed. In the
bonding orbital, the electron density is greatest between the nuclei of the bonding atoms. In the antibonding orbital, the electron
density decreases to zero between the nuclei.
 Constructive Interference and Destructive Interference tell us that we can add wavelengths or subtract them. The formation of
antibonding molecular orbitals corresponds to destructive interference while the formation of bonding molecular orbitals
corresponds to constructive interference. Stars denote antibonding molecular orbitals.
 In a sigma molecular orbital, the electron density is concentrated symmetrically around a line between the two nuclei of the
bonding atoms. In a pi molecular orbital, the electron density is concentrated above and below a line joining the two nuclei of
the bonding atom.
 To fill molecular orbitals with electron, first know that the number of molecular orbitals formed is always equal to the number of
atomic orbitals combined. Then, the more stable the bonding orbital, the less stable the corresponding antibonding orbital. The
filling of molecular orbitals proceeds from low to high energies. In stable molecule, the number of electrons in bonding
molecular orbitals is always greater than that in antibonding molecular orbitals because we place electrons first in the lowerenergy bonding molecular orbitals. Like an atomic orbital, each molecular orbital can accommodate up to two electrons with
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opposite spins in accordance with the Pauli Exclusion Principle. When electrons are added to molecular orbitals of the same
energy, the most stable arrangement is predicted by Hund’s rule; that is, electrons enter these molecular orbitals with parallel
spins. The number of electrons in the molecular orbitals is equal to the sum of all the electrons on the bonding atoms.
To evaluate the stability of species, we determine their bond order, which is ½ (number of electrons in bonding MO’s – the
number of electrons in antibonding MO’s). Bond order of 0 or less means that the molecule cannot exist. H 2+ has only one
electron, which is in the σ1s orbital. Therefore, the bond order is ½, saying that this species may be stable. H 2 has a bond order of
1, making it stable with one full covalent bond. The electron configuration of H2 is (σ1s)2. He2+ has three electrons, with two in
the σ1s and one in the σ1s*, making the bond order ½, similar to that of H 2+. He has four electrons, with two in the σ1s and two in
the σ1s* orbital. H2 is more stable than H2+ and He2+, which are more stable than He2.
Homonuclear Diatomic Molecules are diatomic molecules containing atoms of the same elements. When bonding involves porbitals, 2px, 2py, and 2pz orbitals are used. Energies increase as follows: σ1s < σ1s* < σ2s < σ2s* < π2py = π2pz < σ2px < π2py* = π2pz* <
σ2px*.
o The electron configuration of Li2 is (σ1s)2(σ1s*)2(σ2s)2 because there are six electrons, which fill in the shells order of
increasing energy.
o The electron configuration of C2 is (σ1s)2(σ1s*)2(σ2s)2(σ2s*)2(π2py)2(π2pz)2 because there are 12 electrons, which also fill in
the shells in order of increasing energy.
o The electron configuration of O2 is (σ1s)2(σ1s*)2(σ2s)2(σ2s*)2(σ2px)2(π2py)2(π2pz)2(π2py*)1(π2pz*)1 because there are 16
electrons, which fill in the shells in order of increasing energy.
 The bond order of O2 is 2, because there are six bonding orbital electrons minus two antibonding orbital
electrons, not counting the σ1s and σ2s orbitals.
Sometimes, molecules cannot be explained by a single structure.
o Examples are Benzene and Carbonate molecules.
o This is done by determining the hybridization state of the atoms present (VB THEORY) and then by the formation of
appropriate molecular orbitals (MO THEORY).
Benzene (C6H6)
o Planar Hexagonal Molecule with Carbons at each of six corners.
o All C-C bonds are equal in length and strength, like the C-H bonds.
o The CCC and HCC angles are all 120 degrees.
 Each Carbon is sp2 hybridized.
 It forms three sigma bonds with two adjacent carbon atoms and a hydrogen atom.
o Unlike the pi bonding orbitals in ethylene, those in benzene form delocalized molecular orbitals, which are not
confined between two adjacent bonding atoms, but actually extend over three or more atoms.
o Each 2pz orbital overlaps two
others, one on each adjacent
carbon atom.
 The interaction of
six 2pz orbitals leads
to the formation of
six pi molecular
orbitals,
three
bonding and three
antibonding.
o A benzene molecule in the
ground state has six electrons
in the three pi bonding
molecular
orbitals,
two
electrons with paired spins in
each orbital.
Carbonate (CO32-)
o Trigonal planar molecule.
o We have to assume the C atom is sp2 hybridized.
o The C atom forms sigma bonds with three O atoms.
 The unhybridized 2pz orbital of the C atom can overlap the 2pz orbitals of all three O atoms.
 This results in a delocalized molecular orbital that extends over all four nuclei in such a way that the
electron densities in the C-O bonds are all the same.
o Molecules with delocalized molecular orbitals are generally more stable than those containing molecular orbitals
extending over only two atoms.