00:01
Now if it was that straightforward,
then you'd have a relatively limited number
of organic molecules that could be formed.
However, things become a little bit more complicated
in the case of carbon and, indeed, some other
elements as well. And this is where you have
a change in energy, formally, of atomic orbitals.
Do you remember? We were looking at s orbitals
in a single shell being of a lower energy
to p orbitals in a slightly higher energy
subshell. But this isn't the end of the story.
If we look at methane, which has four covalent
bonds evenly spaced in three dimensions following
a tetrahedral structure with an angle of approximately
109 degrees between any of the two bonds,
this suggests that each bond is somehow the
same. This cannot be possible if you have
one s orbital and three p orbitals with different
geometries in different directions. And this
is where the concept of hybridization comes
in. So the way in which you can have sigma bonds
formed which are equal in terms of their electron
distribution, resulting in a uniform tetrahedral
structure, is if the s and the p orbitals
found in the element carbon are somehow hybridized—hybridization,
in this context, meaning conversion of those
2s and 2p into one uniform atomic orbital
type.
01:43
So if we look at the current electron configuration
for carbon, we have our 1s2, 2s2, and then
we have two unpaired electrons in the p orbitals.
If I bring it forward, it's possible to bring
together those atomic orbitals to form a single
orbital type. So we have the two p orbitals
from carbon, and we have two electrons in
the 2s orbital. If we look at the 2 p orbitals,
you'll see that a tetrahedral structure for
our methane is not possible, because you have
a 90-degree angle and a 90-degree angle in
the other direction for the different p orbitals:
px, py, and pz. The reality is you have a
mix-up of the s and the p orbitals, and this
is very important, because in order for them
to be uniform, they have to be uniform in
terms not necessarily in terms of direction,
but in terms of geometry and also in terms
of energy.
So let's see if we can explain this using
electron configuration energy diagram. Because
they are the same, we get a distribution,
or methane-based structure, based on four
identical (in terms of energy and shape) orbitals.
03:04
So let's actually have a look at what that
means. So on the left-hand side, we have the
atomic orbitals for carbon that we are familiar
with. Bear in mind we have the s orbital—the
2s, which has got two electrons in it from
carbon—and we have the px, py, and pz. And
there are two… one electron in, say, the
px and one electron in the py. In hybridization
for carbon, what happens is all four of those
orbitals effectively come together and form
four (as you can see, in terms of shape) identical
hybrid orbitals. In this case, the degree
of hybridization can be defined by the number
of original atomic orbitals which is involved
in that hybridization. So if you pay attention
to the hybrid orbital that is formed, you'll
see it says "4x sp3." This is telling us that
there are four sp3 hybrid orbitals where,
in this case, the s orbital and the three
p orbitals all come together to create one
orbital class, hence the name sp3 hybridization.
What's happening from an energetic perspective,
as you can see here, is that in our original
case, we have two electrons in the 2s and
two unpaired in the 2p. And in the process
of hybridization, which occurs naturally,
you end up with four unpaired electrons of
an energy level which is somewhere in between
the 2p and the 2s. Energy levels are the same.
Shape is the same. Geometry depends on how
they're interacting with other atoms. So if
we take, for example, hydrogen and carbon,
where car… hydrogen has its single s orbital,
and we have an sp3 hybridized orbital from
carbon, you can see that they form a sigma molecular
orbital not unlike that one which I showed
you for hydrogen and fluorine. Note you still
have the sigma bond with the electron density
distributed over that green area that you
can see there, and the small lobe which is
out of phase with the hydrogen atomic orbital.
For methane, of course, in sp3 hybridized
carbon, there are four of these hybrid atomic
orbitals, and therefore, it's possible for
four hydrogens to bind to a single carbon.
Equally, it's possible for those four carbons
to bind with a defined uniform tetrahedral
structure with 109-degree bond angles between
each of those sigma bonds formed. They're evenly
spaced, as I've just said.
05:59
But it's not just restricted to carbon. As
you can see, oxygen can also be hybridized.
06:05
If you recall, oxygen is in group 6; therefore,
it has six electrons in its outer shell. That
includes the 2s and also the 2p. Prior to
hybridization, it has two electrons in the
2s, one paired set in the 2p, and then two
unpaired electrons in the 2p. After hybridization,
which is also uses the same nomenclature as
we saw for carbon, all of those orbitals are
brought together to form four sets of similar
sp3 hybridized orbitals. And we see evidence
of this in the structure of water. Note if
you look at the structure of water, which
is shown there on the board, we have 109-degree
(or approximately 109-degree) bonding angle
between the two hydrogens in H2O. Bear in
mind the oxygen and the hydrogen are bound
together by sigma molecular orbitals. Two of
the sp3 orbitals will be filled with a pair
of electrons, and two have one electron and
are ready to form sigma bonds with the hydrogens.
07:18
Hence, water has a bent shape, as you can
see here. And knowledge of hybridization of
orbitals isn't just something to talk about
from an academic perspective; it's very important
in terms of geometry. And as we'll see as
we move along further in this course and we
start looking at biological interactions of
small molecules, we'll see that geometry is
absolutely essential for the correct formation
of a useful pharmacophore for medicinal chemistry.