Another topic that you’ll need to be familiar with for the
SAT IIChemistry test is that of valence bond theory. By now, you are awarethat two atoms will form a bond when there is orbital overlap betweenthem, and a maximum of two electrons can be present in the overlappingorbitals.
Since the pair of electrons is attracted toboth atomic nuclei, a bond is formed, and as the extent of overlapincreases, the strength of the bond increases. The electronic energydrops as the atoms approach each other, but it begins to increase againwhen 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 sigmabonds 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 axisof the bond. Maximum overlap forms the strongest-possible sigma bond,and the two atoms will arrange themselves to give the greatest-possibleorbital 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
porbitals) to create orbitals that have energy that’s in between theenergy 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 figureare for a nonbonded carbon atom, but once the carbon atom begins tobond with other atoms (in this case hydrogen), the atomic orbitalshybridize, and this changes their shape considerably. Notice how thefirst set of figures form the
sp3 atomic orbital, the hybrid, and this leads to further hybridization.
Ammonia also has
sp3 hybridization, even though it has a lone pair.
Multiple Bonding
Now let’s look more closely at pi bonds. As wementioned earlier in this chapter, pi (p) bonds result from thesideways overlap of
p orbitals, and pi orbitals are defined bythe region above and below an imaginary line connecting the nuclei ofthe two atoms. Keep in mind that pi bonds never occur unless a sigmabond 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:
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p orbitals that are unhybridized are not shown in this depiction:
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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:
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Here’s how it looks with all the pieces put together:
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Here is a table summarizing hybridization and structure:
