Valence Bond Theory
Another topic that you’ll need to be familiar with for
the SAT II Chemistry test is that of valence bond theory. By now,
you are aware that two atoms will form a bond when there is orbital
overlap between them, and a maximum of two electrons can be present
in the overlapping orbitals.
Since the pair of electrons is attracted to both atomic
nuclei, a bond is formed, and as the extent of overlap increases,
the strength of the bond increases. The electronic energy drops as
the atoms approach each other, but it begins to increase again when
they become too close. This means there is an optimum distance,
the observed bond distance, at which the total energy is at a minimum.
Let’s delve a little more deeply into sigma bonds now
and describe them in more detail. As you know, sigma (s) bonds are
single bonds. They result from the overlap of two s orbitals,
an s and a p orbital, or two head-to-head p orbitals.
The electron density of a sigma bond is greatest along the axis of
the bond. Maximum overlap forms the strongest-possible sigma bond,
and the two atoms will arrange themselves to give the greatest-possible
orbital overlap. This is tricky with p orbitals
since they are directional along the x, y,
and z axes.
Hybrid orbitals result from a blending of
atomic orbitals (in other words, s and p orbitals) to
create orbitals that have energy that’s in between the energy of
the lone orbitals. Look at the methane molecule, for example: all
four of the C—H bonds are 109.5º apart, while nonbonded p orbitals
are only 90º apart.
The orbitals shown at the left of the figure are for a
nonbonded carbon atom, but once the carbon atom begins to bond with
other atoms (in this case hydrogen), the atomic orbitals hybridize,
and this changes their shape considerably. Notice how the first
set of figures form the sp3 atomic
orbital, the hybrid, and this leads to further hybridization.
Ammonia also has sp3
even though it has a lone pair.
Now let’s look more closely at pi bonds. As we mentioned
earlier in this chapter, pi (p) bonds result from the sideways overlap
of p orbitals, and pi orbitals are defined by the region
above and below an imaginary line connecting the nuclei of the two
atoms. Keep in mind that pi bonds never occur unless a sigma bond
has formed first, and they may form only if unhybridized p orbitals
remain on the bonded atoms. Also, they occur when sp or sp2 hybridization
is present on central atom but not sp3 hybridization.
Below, we show the formation of a set of sp2 orbitals.
This molecule would contain a double bond, like ethene. Notice again
how the first set of figures form the sp2 atomic orbital,
the hybrid, and the last figure shows full hybridization:
The set of p orbitals that are unhybridized
are not shown in this depiction:
A different view, which doesn’t show the hydrogens and
centers on the C atoms, shows the unhybridized p orbitals
that create the sideways overlap that’s necessary to create the
double pi bond:
Here’s how it looks with all the pieces put together:
Here is a table summarizing hybridization and structure: