LEWIS FORMULAS, STRUCTURAL ISOMERISM,
AND RESONANCE STRUCTURES
LEARNING OBJECTIVES: To understand the uses and limitations of Lewis formulas, to introduce structural
isomerIsm, and to learn the basic concept of resonance structures.
CHARACTERISTICS OF LEWIS FORMULAS: Lewis formulas are structures that show the connectivity, or bonding
sequence of the atoms, indicating single, double, or triple bonds. They should also show any formal charges
and unshared electrons that might be present in the molecule. Additional examples of Lewis formulas follow.
H
H
O
H C
C
C H
H
Cl
H
O
H
H C
C
C
H
C3H5ClO
H
C3H5ClO
Cl
H H
O
H C C
C
Cl
H H
C3H5ClO
These examples were deliberately chosen because all three molecules shown have the same molecular formula,
but different connectivities, or bonding sequences. Such substances are called structural isomers, or sometimes
constitutional isomers.
Notice that only the first structure shows the unshared electrons of chlorine. In Lewis formulas of organic compounds,
it is customary to omit the lone electron pairs on the halogens unless there is a reason to show them explicitly.
Lewis formulas are mostly used for covalent substances, but occasionally they also show ionic bonds that might
be present in certain compounds.
H
H
N
H
The bond between nitrogen and
chlorine is ionic. All others are covalent.
Cl
H
H
H
O
C
C
O
The bond between oxygen and
sodium is ionic. All others are covalent.
Na
H
COMMON BONDING PATTERNS FOR FIRST AND SECOND ROW ELEMENTS: Once we write enough Lewis
formulas containing the elements of interest in organic chemistry, which are mostly the second row elements, we
find that certain bonding patterns occur over and over. Learning these patterns is useful when trying to write Lewis
formulas because they provide a convenient starting point. For example, in several of the structures given in the
previous section, we find that the carbon bonded to three hydrogens is a unit that occurs quite frequently. It is called
the methyl group, represented by CH3. It is so common that it is valid to write it as such in Lewis formulas, even
though it is in fact an abbreviated form, because everybody knows what it stands for.
H
H
O
C
C
H
O
O
or
CH3
C
O
It is equally valid to represent the acetate
ion by either of these formulas.
Other common bonding patterns are shown below.
HYDROGEN: Usually forms only one bond.
H
H
F
H
O
H
H
C
H
Cl
CARBON: Forms four bonds when neutral, but it can also have only three bonds by bearing a positive or a negative
charge. When it bears a negative charge it should also carry a pair of unshared electrons.
H
H
H
C
H
H
C
H
Cl
H
C
H3C
H3C
H
Neutral carbon
A carbocation has a central carbon with
an incomplete octet and a formal +1 charge.
CH3
CH3
C
C
H
CH3
C
C
A carbanion has a central carbon with an
unshared electron pair and a formal -1 charge.
CH3
NITROGEN: Forms three bonds and carries a lone pair of electrons when neutral. It can also form four bonds by
bearing a positive charge, in which case it carries no unshared electrons. Finally, it can also form two bonds as it
carries two unshared electron pairs and a negative charge.
H
H
N
H3C
C
H
H3C
H
H
N
N
H
H
H3C
N
Neutral nitrogen
H
C
H
N
H
H3C
H
C
H
Positively charged nitrogen
N
H
Negatively charged nitrogen
OXYGEN: Forms two bonds and carries two lone pairs when neutral. It can form three bonds with a positive charge,
or one bond with a negative charge. In each case it must carry the appropriate number of unshared electron pairs
to complete the octet.
H
H
O
H
H
water
O
O
H
hydronium ion
H
hydroxide ion
HALOGENS: Form one bond and carry three electron pairs when neutral. Can carry a negative charge with no
bonds. They are rarely seen with positive charges.
H
F
Cl
THIRD ROW ELEMENTS: They behave like their second-row counterparts, except that they can expand their
valence shells if needed.
O
Electron pairs on oxygen
S
H O S O H
are not shown for clarity.
H
H
O
Cl
Cl
P
Cl
Cl
Cl
Br
P
Br
Br
ELECTRON DEFICIENCY IN SECOND ROW ELEMENTS: One thing worth noting is that, in the second row, only
boron and carbon can form relatively stable species in which they bond with an incomplete octet. Examples have
already been discussed. Boron has no choice but to be electron deficient. Carbon can bond with a complete octet
or with an incomplete octet. Obviously bonding with a complete octet provides higher stability.
CH3
F
F
B
F
Boron has no choice but
to have an incomplete octet
H3C
C
CH3
An electron deficient
carbon in a carbocation
It is however very rare to observe species where nitrogen or oxygen bond with incomplete octets. Their high
electronegativity renders such situation high energy, and therefore very unstable. For all intents and purposes, avoid
writing formulas where oxygen or nitrogen are shown with incomplete octets, even if they carry a positive charge.
CH3
CH3
O
CH3
O
CH3
To write this structure without the lone pair
of electrons on oxygen is unacceptable.
H
H
N
H
This structure is unacceptable
and indeed it looks quite awkward
This species might exitst in the high energy environment
of a mass spectrometer, but it is not frequently observed
in common organic reactions.
RESONANCE STRUCTURES AND SOME LIMITATIONS OF LEWIS FORMULAS : Lewis formulas are misleading
in the sense that atoms and electrons are shown as being static. By being essentially two-dimensional representations
they also fail to give an accurate idea of the three-dimensional features of the molecule, such as actual bond angles
and topography of the molecular frame.
Furthermore, a given compound can have several valid Lewis formulas. For example CH3CNO can be represented
by at least three different but valid Lewis structures called resonance forms, or resonance structures, shown
below.
H
H
H C C N O
H
H C C N O
H
H
H C C N O
H
I
III
II
However, a stable compound such as the above does not exist in multiple states represented by structures I, or II,
or III. The compound exists in a single state called a hybrid of all three structures. That is, it contains contributions
of all three resonance forms, much like a person might have physical features inherited from each parent to varying
degrees.
In the resonance forms shown above the atoms remain in one place. The basic bonding pattern, or connectivity,
is the same in all structures, but some electrons have changed locations. This means that there are certain rules
for electron mobility that enable us to “push” electrons around to arrive from one resonance structure to another.
These rules will be examined in detail in a later paper.
ALL RESONANCE STRUCTURES MUST BE VALID LEWIS FORMULAS: By convention, we use double-headed
arrows to indicate that several resonance structures contribute to the same hybrid. Continuing with the example
we’ve been using, the resonance structures for CH3CNO should be written in this way if we want to emphasize that
they represent the same hybrid.
CH3
C
N
I
O
CH3
C
N
II
O
CH3
C
N
O
III
Do not confuse double-headed arrows with double arrows. A double arrow indicates that two or more species are
in equilibrium with each other and therefore have a separate existence. Double-headed arrows indicate resonance
structures that do not exist by themselves. They simply represent features that the actual molecule, the hybrid,
possesses to one extent or another.
When writing resonance structures keep in mind that THEY ALL MUST BE VALID LEWIS FORMULAS. The factors
that make up valid Lewis formulas are as follows.
1. Observe the rules of covalent bonding, including common patterns as discussed previously. Make sure to show
all single, double, and triple bonds.
2. Account for the total number of valence electrons being shared (from all the elements), including bonding and
nonbonding electrons. Make sure to show these nonbonding electrons.
3. Account for the net charge of the molecule or species, showing formal charges where they belong.
4.Observe the octet rule as much as possible, but also understand that there are instances where some atoms may
not fulfill this rule.
5. Avoid having unpaired electrons (single electrons with no partners) unless the total number of valence electrons
for all elements is an odd number. This is not a very frequent occurrence, but the following example shows a species
that could exist as a reaction intermediate in some high energy environments.
H
H
C
H
The total number of valence electrons being shared for all atoms is 4 from carbon and 3 from the three hydrogens,
for a total of 7. Because it is an odd number, it is impossible to have all these electrons paired. Therefore the
presence of a single electron cannot be avoided. Notice that there is no formal charge on carbon, since it has no
surplus or deficit of valence electrons.
RELATIVE ENERGIES OF RESONANCE STRUCTURES. From the examples given so far it can be seen that
some resonance forms are structurally equivalent and others are not. The potential energy associated with
equivalent Lewis structures is the same. If the Lewis structures are not equivalent, then the potential energy
associated with them is most likely different. This means that equivalent resonance structures are also equivalent
in stability and nonequivalent structures have different stabilities. This in turn means that equivalent structures
contribute equally to the hybrid and nonequivalent structures do not contribute equally to the hybrid.
In the example below, the two structures are equivalent. Therefore they make equal contributions to the hybrid.
O
O
equivalent in energy
CH3
CH3
O
O
It is very difficult to accurately represent the hybrid with drawings because it is a composite of all the resonance
contributors. Some representations such as the first one shown below are sometimes given. In this case the broken
line represents the electrons and the negative charge which are spread over three atoms (O-C-O). Another commonly
used representation of the hybrid is given on the right, showing that each oxygen atom shares a -1/2 charge.
However, none of them accurately conveys the true picture. For convenience, the unshared electrons are sometimes
omitted from some representations, just like it’s done with hydrogen atoms. This must be kept in mind when examining
the different structures.
O
CH 3
O -1/2
CH 3
O
two commonly used representations
for the acetate ion hybrid
O -1/2
There is a third resonance form that can be drawn for the acetate ion hybrid. The structure shown below is structurally
different from the ones shown above. This means that it is of different energy and therefore does not contribute to
the hybrid to the same extent as the others. In this case, it happens to be less stable than the other two and therefore
does not make a significant contribution to the hybrid.
O
CH 3
O
What factors dictate the relative stabilities of different resonance structures? Charge separation is an important
one. In the example just given there are too many charges present in the structure (even though the net charge is
still the same). This results in a structure with higher potential energy than others with fewer or no formal charges.
This structure is less stable and its contribution to the hybrid is probably minor.
Another important factor that increases potential energy (lowers stability) is the presence of atoms with an incomplete
octet. In the example below structure I has a carbon atom with a positive charge and therefore an incomplete octet.
Based on this criterion, this structure is less stable than structure II and makes a less significant contribution to the
hybrid.
H3C
H3C
O
C
H3C
H
C
O
H
H3C
I
II
Finally, another factor that comes into play when determining the relative energies of resonance structures is the
relative electronegativities of atoms that bear charges. More electronegative atoms are more comfortable with
negative charges. Less electronegative atoms are more comfortable with positive charges. In the example below,
structure I is less stable than II. The negative charge is on carbon, which is less electronegative than oxygen.
Therefore structure II makes a larger contribution to the hybrid.
O
H3C
C
I
O
C
H
H3C
H
C
II
C
H
H
Some texts refer to the different contributors as “important” or “unimportant.” This can be misleading, because
“importance” is context-dependent. In the example above, structure I is less important in terms of its contribution
to the hybrid, but in terms of reactivity, it is very important. It will be seen later that this structure provides a better
indication of how this species reacts with electrophiles in certain types of reactions.
Based on the charge separation criterion, structure II below might be labeled “unimportant,” or “less important”
because it is less stable than I. But if we are trying to assess the polarity of this molecule, structure II becomes very
important because it reveals that the carbon atom has positive character and the oxygen has negative character.
The hybrid representation on the right then portrays the molecule as a polar molecule, which structure I alone does
not give much indication of.
δ−
H
O
O
O
C
C
C
I
H
H
H
H
δ+ H
Hybrid representation
as a polar structure.
II
The following additional examples further illustrate the relative importance of the factors that contribute to structural
energy and stability.
H
N
H
O
H
CH3
H
O
N
C
Major contributor.
Neutral structure.
H
C
N
CH3
H
Minor contributor.
Charge separation.
CH3 C
O
Minor contributor. Incomplete octet on carbon,
but more important in tems of reactivity.
CH3
H
CH3
H
N
C
Minor contributor.
Incomplete octet
on carbon.
CH3 C
CH3
C
CH3
Major contributor.
All atoms have octets, even though
there is a positive charge on nitrogen
O
Major contributor. All atoms have octets,
even though there is a positive charge on oxygen.