This Lewis structure shows two
different types of bonds, single and double. Because it takes more energy to
break a double bond than a single bond, we say that a double bond is stronger
than a single bond. Double bonds also have a shorter bond length, the distance
between the nuclei of the two atoms in the bond, than single bonds do. Thus, if
the above Lewis structure for nitrate were correct, the nitrate polyatomic ion
would have one bond that is shorter and stronger than the other two.
This is not the case.
Laboratory analyses show all three of the bonds in the nitrate ion to be the
same strength and the same length. Interestingly, the behavior of the bonds
suggests they are longer than double bonds and shorter than single bonds. They
are also stronger than single bonds but not as strong as double bonds. In order
to explain how this is possible for the nitrate ion and for molecules and
polyatomic ions like it, the valence-bond model had to be expanded.
It is important to stress that
the nitrate ion is not really changing from one resonance structure to another,
but chemists find it useful, in an
intermediate stage in the process of developing a better description of the
nitrate ion, to think of it as if it
were doing so. In actuality, the ion behaves as if it were a blend of the three
We can draw a Lewis-like
structure that provides a better description of the actual character of the
nitrate ion by blending the resonance structures into a single resonance
The actual geometry of the
polyatomic ion is trigonal planar with bond angles of 120.
Click here to see a molecular model
of the nitrate ion.
Resonance Structures and the Resonance Hybrid
is possible whenever a Lewis structure has a multiple bond and an adjacent atom
with at least one lone pair. The following is the general form for resonance in
a structure of this type. The arrows show how you can think of the electrons
shifting as one resonance structure changes to another.
You can follow these steps to
write resonance structures.
Shift one of the lone pairs
on an adjacent atom down to form another bond.
Shift one of the bonds in a
double or triple bond up to form a lone pair. (You might find it useful to
draw arrows indicating the hypothetical shift of electrons.)
Draw additional resonance
structures by repeating this process for each adjacent atom with a lone
Separate the resonance
structures with double-headed arrows.
For example, the two resonance
structures for the formate ion, HCO2- are
To generate the second
resonance structure from the first, we imagine one lone pair dropping down to
form another bond, and pushing an adjacent bond off to form a lone pair. The
arrows show this hypothetical shift of electrons. These resonance
structures lead to the resonance hybrid below.
Click here to see a molecular model of
the formate ion.
This general procedure for
drawing resonance structures will not always lead to a reasonable resonance
structure. For example, fluorine atoms do not participate in resonance.
According to the valence-bond model, for a fluorine atom to form two bonds and
two lone pairs, it would have to lose an electron, a highly unlikely act for the
most electronegative element on the periodic table. Thus, any resonance
structure that includes a double bond to fluorine is not considered a reasonable
resonance structure. Thus, although fluoroethene, CH2CHF, has a
double bond and an adjacent atom with a lone pair (components that suggest the
possibility resonance), only one of its two hypothetical resonance structures is
Therefore, fluoroethene does
not have resonance, and the first structure above is the best description of a
There is a similar situation
with oxygen atoms. Although it is possible for oxygen atoms to have three bonds
and one lone pair, its is not likely that the second most electronegative
element would lose the electron necessary to make this possible. Thus, we will
eliminate resonance structures that have three bonds and a lone pair for an
oxygen atom. For example, formic acid, HCO2H, has a double bond and
an adjacent atom with a lone pair, so we might think that it has resonance. The
two resonance structures would be
The first Lewis structure is
reasonable, but the second one, with three bonds and a lone pair on an oxygen
atom, is not considered a reasonable resonance structure. Therefore, there is no
significant resonance for formic acid, and the first Lewis structure above is
the best description of its structure.
We will consider resonance a
possibility for molecules and polyatomic ions that have the following as part of
their Lewis structure.
Z can have more than one
X and Y can have lone
The X-Y bond can be a
triple bond. The Y-Z bond can be a double bond.
Z cannot be F with one bond
and three lone pairs or O with two bonds and two lone pairs.
here to see Example 1.
here to see Exercise 1.
Expanded Lewis Structure
When resonance is considered,
we add another step to our Lewis structure drawing procedure.
#8: Once you have
a reasonable Lewis structure, consider the possibility of resonance. If
resonance is possible, draw the reasonable resonance structures and the
resonance hybrid for the structure.
here to see Example 2.
here to see Exercise 2.
the Benzene Molecule
It is possible to have
resonance without the participation of lone pairs. The most important examples
of this are benzene, C6H6, and compounds that contain the
benzene ring. Benzenes six carbon atoms are linked to each other in a six-membered
ring. Its Lewis structure is often represented with three double bonds as shown
below, but chemists often simplify it by leaving off the elements symbols and
the carbon-hydrogen bonds.
The Lewis structures above
depict the benzene molecule as if it contained two types of C-C bonds, double
and single. In actuality, all of benzenes C-C bonds appear to be the same,
and we can explain why in terms of resonance. It is as if the benzene ring were resonating between the two structures
The resonance hybrid is
Structure #3 below. Because it is a bit tedious to draw all the dots, the
structure of the benzene molecule is often written as shown in Structure #4,
with the dotted lines represented by a circle.
In summary, Structures #1, #2, #3, and #4 are all used to describe benzene.
Structure #3 Structure #4