In chemistry, valence bond theory is one of two basic theories, along with
molecular orbital theory, that were developed to use the methods of quantum
mechanics to explain chemical bonding. It focuses on how the atomic orbitals of
the dissociated atoms combine to give individual chemical bonds when a
molecule is formed. In contrast, molecular orbital theory has orbitals
that cover the whole molecule. History
In 1916, G. N. Lewis proposed that a chemical bond forms by the interaction
of two shared bonding electrons, with the representation of molecules as Lewis
structures. In 1927 the Heitler–London theory was formulated which for the
first time enabled the calculation of bonding properties of the hydrogen
molecule H2 based on quantum mechanical considerations. Specifically, Walter
Heitler determined how to use Schrödinger's wave equation to show how
two hydrogen atom wavefunctions join together, with plus, minus, and exchange
terms, to form a covalent bond. He then called up his associate Fritz London and
they worked out the details of the theory over the course of the night.
Later, Linus Pauling used the pair bonding ideas of Lewis together with
Heitler–London theory to develop two other key concepts in VB theory:
resonance and orbital hybridization. According to Charles Coulson, author of
the noted 1952 book Valence, this period marks the start of "modern valence bond
theory", as contrasted with older valence bond theories, which are
essentially electronic theories of valence couched in pre-wave-mechanical
terms. Resonance theory was criticized as imperfect by Soviet chemists during
the 1950s. Theory
According to this theory a covalent bond is formed between the two atoms by the
overlap of half filled valence atomic orbitals of each atom containing one
unpaired electron. A valence bond structure is similar to a Lewis
structure, but where a single Lewis structure cannot be written, several
valence bond structures are used. Each of these VB structures represents a
specific Lewis structure. This combination of valence bond structures
is the main point of resonance theory. Valence bond theory considers that the
overlapping atomic orbitals of the participating atoms form a chemical
bond. Because of the overlapping, it is most probable that electrons should be
in the bond region. Valence bond theory views bonds as weakly coupled orbitals.
Valence bond theory is typically easier to employ in ground state molecules. The
inner-shell orbitals and electrons remain essentially unchanged during the
formation of bonds. The overlapping atomic orbitals can
differ. The two types of overlapping orbitals are sigma and pi. Sigma bonds
occur when the orbitals of two shared electrons overlap head-to-head. Pi bonds
occur when two orbitals overlap when they are parallel. 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 sigma bond, double
bonds consist of one sigma bond and one pi bond, and triple bonds contain one
sigma bond and two pi bonds. However, the atomic orbitals for bonding may be
hybrids. Often, the bonding atomic orbitals have a character of several
possible types of orbitals. The methods to get an atomic orbital with the proper
character for the bonding is called hybridization.
Valence bond theory today Modern valence bond theory now
complements molecular orbital theory, which does not adhere to the valence
bond idea that electron pairs are localized between two specific atoms in
a molecule but that they are distributed in sets of molecular orbitals which can
extend over the entire molecule. Molecular orbital theory can predict
magnetic and ionization properties in a straightforward manner, while valence
bond theory gives similar results but is more complicated. Modern valence bond
theory views aromatic properties of molecules as due to spin coupling of the
π orbitals. This is essentially still the old idea of resonance between Kekulé
and Dewar structures. In contrast, molecular orbital theory views
aromaticity as delocalization of the π-electrons. Valence bond treatments are
restricted to relatively small molecules, largely due to the lack of
orthogonality between valence bond orbitals and between valence bond
structures, while molecular orbitals are orthogonal. On the other hand, valence
bond theory provides a much more accurate picture of the reorganization
of electronic charge that takes place when bonds are broken and formed during
the course of a chemical reaction. In particular, valence bond theory
correctly predicts the dissociation of homonuclear diatomic molecules into
separate atoms, while simple molecular orbital theory predicts dissociation
into a mixture of atoms and ions. For example the molecular orbital function
for dihydrogen is an equal mixture of the covalent and ionic valence bond
structures and so predicts incorrectly that the molecule would dissociate into
an equal mixture of hydrogen atoms and hydrogen positive and negative ions.
Modern valence bond theory replaces the overlapping atomic orbitals by
overlapping valence bond orbitals that are expanded over a large number of
basis functions, either centered each on one atom to give a classical valence
bond picture, or centered on all atoms in the molecule. The resulting energies
are more competitive with energies from calculations where electron correlation
is introduced based on a Hartree–Fock reference wavefunction. The most recent
text is by Shaik and Hiberty. Applications of valence bond theory
An important aspect of the VB theory is the condition of maximum overlap, which
leads to the formation of the strongest possible bonds. This theory is used to
explain the covalent bond formation in many molecules.
For example, in the case of the F2 molecule, the F−F bond is formed by the
overlap of pz orbitals of the two F atoms, each containing an unpaired
electron. Since the nature of the overlapping orbitals are different in H2
and F2 molecules, the bond strength and bond lengths differ between H2 and F2
molecules. In an HF molecule the covalent bond is
formed by the overlap of the 1s orbital of H and the 2pz orbital of F, each
containing an unpaired electron. Mutual sharing of electrons between H and F
results in a covalent bond in HF. See also
Modern valence bond theory References