CHEMISTRY SEMESTER ONE
UNIT 9
UNIT 9: ORBITALS AND THEORIES OF CHEMICAL BONDING
INTRODUCTION TO THIS UNIT:
Up until this point, you have learned some very practical theories of bonding that allow you to predict
many properties of molecules. Recall that theories are modified or new theories are developed to
address observations not explained by earlier theories. With this in mind, we will look at two more
theories of chemical bonding: valence bond (VB) theory and molecular orbital (MO) theory. These are a
little more complicated than VSEPR theory, but we use them because we need them. For many types of
bonding addressed at the first-year level, VSEPR theory works very well, so there is no reason for you to
discard it.
By now, you are comfortable with the idea of orbitals on atoms. The main idea behind valence bond
theory is overlap of these atomic orbitals. As orbitals overlap, electrons are shared between the two
atoms, and the bonded atoms that result are lower in energy than the two independent atoms that
formed the diatomic molecule.
This concept is easy to understand, but hybridization is somewhat more complex. For example, if you
consider the p atomic orbitals on carbon, px, py and pz, they are orthogonal to each other. If these
simple atomic orbitals were used to explain the bonding of carbon, the bond angles at carbon would be
90º. This does not fit experimental data so the concept of hybridization is introduced. Here, the s and p
orbitals are mixed together. This is the hybridization. The resulting hybrid orbitals (sp, sp2 and sp3) are
neither pure s nor p, but have intermediate character (just as a mule is neither a donkey nor a horse, but
a hybrid of the two). After hybridization new sets of orbitals result in bond angles of 180º, 120º and
109.5º which support the experimental data.
When orbitals overlap in an end-on fashion, a sigma (σ) bond results where electron density lies along
the internuclear axis joining the two atoms. Atoms that are sp or sp2 hybridized also contain
unhybridized p orbitals. If the atoms are small enough, the half-filled p orbitals can overlap in a side-on
fashion resulting in a pi (π) bond. The π bond forms in addition to the σ bond. If you look at some
models you will see that the electron density of the π bond lies around, not along, the internuclear axis.
This has the effect of restricting rotation of the groups connected by the bonds. Note that the methyl
groups of ethane can freely rotate with respect to each other, but the methylene groups of ethene
cannot.
Molecular orbital theory allows us to more accurately describe the bonding in molecules. Using valence
bond theory we said that the electrons were localized between the two atoms that were connected by
the bond. According to MO theory, atomic orbitals combine to form molecular orbitals, and the
electrons are delocalized over the molecule. Also the total number of atomic orbitals that combine
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CHEMISTRY SEMESTER ONE
UNIT 9
must equal the total number of molecular orbitals that form. These new molecular orbitals are
categorized as bonding orbitals and anti-bonding orbitals. Electrons in bonding orbitals serve to stabilize
the molecule so bonding orbitals are lower in energy than the atomic orbitals from which they were
formed. Electrons in antibonding orbitals serve to destabilize the molecule, and these orbitals are
higher in energy than the atomic orbitals from which they were formed. If the energies of the
combining atomic orbitals are similar, the orbital overlap is better and a stronger bond is formed. As the
difference in energies of the atomic orbitals increases, the weaker the bond.
Notice that MO theory accurately predicts that He2 does not form, but He2+ is stable. Extending MO
theory to the second-row diatomic molecules and their ions shows how useful this theory is. If you
consider the electronic structure of O2 using the valence bond approach, you would say that there is a π
bond and all electrons are paired up. Since all electrons are paired, the molecule should be diamagnetic.
It turns out that O2 is paramagnetic, and this is demonstrated using an MO diagram. Create the MO
diagram of O2 and fill the molecular orbitals with electrons. Just as you followed the Pauli Exclusion
Principle and Hund's rule when deriving the electronic configurations of atoms, do the same when filling
molecular orbitals starting with the lowest energy orbitals first. As you add the last two electrons to the
MO diagram of O2 you will notice that each of these electrons goes into an anti-bonding pi (π*) orbital.
Since the electrons are not paired, MO theory correctly predicts the magnetism of O2. As predicted by
valence bond theory and supported by experimental evidence, the bond order of O2 is two. This is
consistent with the MO diagram too; subtract the number anti-bonding electrons from bonding
electrons and divide the result by two. This will give the correct bond order.
READINGS:
(in Petrucci, 9th ed.) : Sections 9.1 – 9.3
OVERVIEW & STRUCTURE FOR YOUR NOTES:

Valence bond theory and orbital overlap

Hybridization

Multiple bonding and delocalized systems

Molecular orbital theory

Bond order

Molecular orbital diagrams for simple diatomic molecules

Application of molecular orbital theory to unsaturated and aromatic systems
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CHEMISTRY SEMESTER ONE
UNIT 9
ASSIGNMENT:
(in Petrucci, 9th ed.): From Chapter 9: # 2, 4, 8, 18, 20, 22, 42
CONCLUDING REMARKS:
The material presented in this chapter is abstract, but it provides more sophisticated theories to explain
bonding. To appreciate the difference in VB and MO theories, attempt to explain how these theories
differ, what aspects of bonding each address and the strengths of each theory. After doing this exercise,
you will have a better idea of which theory to use when addressing a specific question about bonding.
Molecular orbital theory is commonly encountered in advanced courses in spectroscopy and organic and
inorganic chemistry.
SUMMARY LEARNING OUTCOMES FROM THIS UNIT:

Understand
o
the concept of valence bond theory and orbital overlap model
o
the concept of hybridization to accommodate varying bond angles
o
application of valence bond theory to multiple bonding
o
application of valence bond theory to benzene and other delocalized systems
o
the three principles of molecular orbital theory – note one should add a fourth namely the
extent of bonding depends crucially on the extent to which the orbitals can overlap.
o
The concept of bond order and be able to calculate bond order for chemical bonds in simple
molecules
o
the reason for the non-existence of certain diatomic molecules such as Be2 because of their
zero bond order

Be able to construct molecular orbital diagrams for simple diatomic molecules

Understand

o
application of molecular orbital theory to simple heteronuclear diatomic molecules
o
application of molecular orbital theory to resonance structures and delocalized orbitals
noting that from n atomic orbitals we generate n molecular orbitals
application of molecular orbital theory to benzene
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CHEMISTRY SEMESTER ONE
UNIT 9
NANSLO Chemistry Core Units and Laboratory Experiments
by the North American Network of Science Labs Online,
a collaboration between WICHE, CCCS, and BCcampus
is licensed under a Creative Commons Attribution 3.0 Unported License;
based on a work at rwsl.nic.bc.ca.
Funded by a grant from EDUCAUSE through the Next Generation Learning Challenges.
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