Chemistry Reference
In-Depth Information
2.6.2 Hybrid orbitals in carbon
2
p
y
1
)
, which shows that carbon has two unpaired
electrons, each in a 2
p
orbital. From our study of
bonding so far, we might predict that carbon will be
able to bond to two other atoms, i.e. it should be
divalent, though this would not lead to an octet of
electrons. Carbon is usually tetravalent and bonds to
up to four other atoms. Therefore, we need to modify
the model to explain this behaviour. This modification
is
hybridization
.
We start this section with a word of caution. Students
frequently find
hybridization
a rather difficult con-
cept to understand and appreciate. However, there is
no particular reason why this should be so.
Chemistry is an experimental science, and to
rationalize our observations we gradually develop and
invoke a number of rules and principles. Theories
may have to change as scientific data increase, and
as old principles cease to explain the facts. All of
the foregoing description of atomic and molecular
orbitals is a hypothesis for atomic and molecular
structure supported by experimental data. So far, the
description meets most of our needs and provides a
good rationalization of chemical behaviour. However,
it falls short in certain ways, and we have to invoke
a further modification to explain the facts. Here are
three observations based upon sound experimental
evidence, which are not accommodated by the above
description of bonding:
sp
3
hybrid orbitals
Methane
is a chemical combination of one carbon
atom and four hydrogen atoms. Each hydrogen atom
contributes one electron to a bond; so, logically,
carbon needs to provide four unpaired electrons to
allow formation of four
bonds. The ability of
carbon to bond to four other atoms requires unpairing
of the 2
s
2
electrons. We might consider promoting
one electron from a 2
s
orbital to the third, as
yet unoccupied, 2
p
orbital (Figure 2.8). This would
produce an excited-state carbon; since the 2
p
orbital
is of higher energy than the 2
s
orbital, the process
would require the input of energy. We could assume
that the ability to form extra bonds would more
than compensate for this proposed change. We now
have four unpaired electrons in separate orbitals, and
the electronic configuration of carbon has become
1
s
2
2
s
2
p
x
1
2
p
y
1
2
p
z
1
. Each electron can now form a
bond by pairing with the electron of a hydrogen atom.
However, this does not explain why methane
is tetrahedral and has four equivalent bonds. The
bond that utilizes the 2
s
electron would surely be
different from those that involve 2
p
electrons, and
the geometry of the molecule should somehow reflect
σ
•
The hydrocarbon
methane
(CH
4
) is tetrahedral in
shape with bond angles of about 109
◦
, and the
four C-H bonds are all equivalent and identical
in reactivity.
•
Ethylene
(ethene,
C
2
H
4
)
is
planar,
with
bond
angles of about 120
◦
, and it contains one
π
bond.
•
Acetylene
(ethyne, C
2
H
2
) is linear, i.e. bond angles
180
◦
, and it contains two
π
bonds.
None of these observations follows immediately from
the electronic configuration of carbon
(
1
s
2
2
s
2
2
p
x
1
carbon
excited-state carbon atom
2
p
2
p
2
s
2
s
1
s
1
s
promotion of a 2
s
electron
into a 2
p
orbital
Figure 2.8
Electronic configuration: excited-state carbon atom
