HYDROGEN BONDING
49
• Structures that are not equivalent do not contribute
equally; the more stable a structure is, the more it
is likely to contribute.
• Highly unstable structures make little contribution
and may be ignored.
Acceptable resonance structures can then be imagined as contributing to the overall electronic distribution in the molecule. By considering the properties of
contributing structures, we can also predict some of
the properties of the molecule. We imagine that the
molecule is not fully represented by a single structure,
but is better represented as a hybrid of its contributing
resonance forms. It is likely that the energy associated with the molecule is actually lower than that of
any contributing resonance form; therefore, the delocalization of electrons that resonance represents is
a stabilizing feature. The larger the number of stable resonance structures we can draw, the greater
the extent of delocalization. The difference in energy
between the actual molecule and that suggested by
the best of the resonance structures is termed the resonance energy or resonance stabilization energy.
This can usually only be an estimated amount.
The resonance terminology and the double-headed
arrow may give the impression that the structures
are rapidly interconverting. This is not true. We
must appreciate right from the start that resonance
structures are entirely hypothetical. They are our
(sometimes clumsy) attempt to write down on paper
what the bonding in the molecule might be like,
and they may depict only the extreme possibilities.
The molecule is presumably happily going about its
business in a form that we cannot easily depict.
Nevertheless, resonance structures are extremely
useful and do help us to explain chemical behaviour.
Let us look again at the simple examples shown
above and the consequences of our hypothetical
resonance structures.
We shall see that most of the reactions of
simple carbonyl compounds, like formaldehyde, are
a consequence of the presence of an electron-deficient
carbon atom. This is accounted for in resonance
theory by a contribution from the resonance structure
with charge separation (see Section 7.1). The second
example shows the so-called conjugate acid of
acetone, formed to some extent by treating acetone
with acid (see Section 7.1). Protonation in this way
typically activates acetone towards reaction, and we
C O
H
H
C O
H
H
formaldehyde
O C
CH 3
CH 3
H
O C
CH 3
CH 3
H
C C
H
C H
H
H
H
C C
H
C H
H
H
H
allylic cation
conjugate acid
of acetone
note that this resonance is only
possible if the atoms are coplanar
find that electron-rich reagents (nucleophiles) attack
the carbon atom (see Section 7.1). This is reasonable,
since we can show this carbon as positively charged
in the right-hand resonance structure. The third
example is the allylic cation. This is a reasonably
stable carbocation, and we attribute this to resonance
stabilization; this is particularly favourable in this
case, since both contributing resonance forms are
identical. We can visualize the allylic cation as an
entity in which the positive charge is delocalized over
the whole structure (strictly, it is the electrons that are
delocalized, but we are one short of a full complement
and it is the positive charge that dominates the
representation).
2.11 Hydrogen bonding
Hydrogen bonds (H-bonds) describe the weak attraction of a hydrogen atom bonded to an electronegative atom, such as oxygen or nitrogen, to the lone
pair electrons of another electronegative atom. These
bonds are different in nature from the covalent bonds
we have described; they are considerably weaker than
covalent bonds, but turn out to be surprisingly important in chemistry and biochemistry.
Let us consider a molecule possessing an O–H σ
bond. This bond is polar because hydrogen is less
electronegative than oxygen (see Section 2.7), and
49
• Structures that are not equivalent do not contribute
equally; the more stable a structure is, the more it
is likely to contribute.
• Highly unstable structures make little contribution
and may be ignored.
Acceptable resonance structures can then be imagined as contributing to the overall electronic distribution in the molecule. By considering the properties of
contributing structures, we can also predict some of
the properties of the molecule. We imagine that the
molecule is not fully represented by a single structure,
but is better represented as a hybrid of its contributing
resonance forms. It is likely that the energy associated with the molecule is actually lower than that of
any contributing resonance form; therefore, the delocalization of electrons that resonance represents is
a stabilizing feature. The larger the number of stable resonance structures we can draw, the greater
the extent of delocalization. The difference in energy
between the actual molecule and that suggested by
the best of the resonance structures is termed the resonance energy or resonance stabilization energy.
This can usually only be an estimated amount.
The resonance terminology and the double-headed
arrow may give the impression that the structures
are rapidly interconverting. This is not true. We
must appreciate right from the start that resonance
structures are entirely hypothetical. They are our
(sometimes clumsy) attempt to write down on paper
what the bonding in the molecule might be like,
and they may depict only the extreme possibilities.
The molecule is presumably happily going about its
business in a form that we cannot easily depict.
Nevertheless, resonance structures are extremely
useful and do help us to explain chemical behaviour.
Let us look again at the simple examples shown
above and the consequences of our hypothetical
resonance structures.
We shall see that most of the reactions of
simple carbonyl compounds, like formaldehyde, are
a consequence of the presence of an electron-deficient
carbon atom. This is accounted for in resonance
theory by a contribution from the resonance structure
with charge separation (see Section 7.1). The second
example shows the so-called conjugate acid of
acetone, formed to some extent by treating acetone
with acid (see Section 7.1). Protonation in this way
typically activates acetone towards reaction, and we
C O
H
H
C O
H
H
formaldehyde
O C
CH 3
CH 3
H
O C
CH 3
CH 3
H
C C
H
C H
H
H
H
C C
H
C H
H
H
H
allylic cation
conjugate acid
of acetone
note that this resonance is only
possible if the atoms are coplanar
find that electron-rich reagents (nucleophiles) attack
the carbon atom (see Section 7.1). This is reasonable,
since we can show this carbon as positively charged
in the right-hand resonance structure. The third
example is the allylic cation. This is a reasonably
stable carbocation, and we attribute this to resonance
stabilization; this is particularly favourable in this
case, since both contributing resonance forms are
identical. We can visualize the allylic cation as an
entity in which the positive charge is delocalized over
the whole structure (strictly, it is the electrons that are
delocalized, but we are one short of a full complement
and it is the positive charge that dominates the
representation).
2.11 Hydrogen bonding
Hydrogen bonds (H-bonds) describe the weak attraction of a hydrogen atom bonded to an electronegative atom, such as oxygen or nitrogen, to the lone
pair electrons of another electronegative atom. These
bonds are different in nature from the covalent bonds
we have described; they are considerably weaker than
covalent bonds, but turn out to be surprisingly important in chemistry and biochemistry.
Let us consider a molecule possessing an O–H σ
bond. This bond is polar because hydrogen is less
electronegative than oxygen (see Section 2.7), and