You are no doubt an expert at drawing out the structures for simple molecules like those shown below. In these molecules there are covalent bonds between the individual atoms and as you would expect these covalent bonds involve the sharing of a pair of electrons. The electrons in these covalent bonds are held firmly in place between nuclei of the two atoms involved in the bond.
However there are other molecules such as benzene which cannot be accurately drawn using single
structures like that for methane, water or carbon dioxide shown above. Some of the bonding
electrons in benzene
are not held in place between two nuclei
as in the simple molecules above but they are
delocalised.
These delocalised
electrons are found in a
variety of molecules and ions; not just benzene.
Delocalised electrons are generally pi(π)
electrons that are
free to move
over more than two atoms, this leads to a bit
of a problem! How do we draw the structure of benzene or any other molecule
with delocalised electrons
if it is not possible to draw the
position of the bonds holding the molecule together?
In a molecule such as ethene which contain one carbon carbon double
covalent bond (C=C). The double
covalent
bond between the carbon atoms consists of a sigma
bond formed by the complete overlap of atomic
orbitals and a pi(π) bond formed by the
partial overlap of p-orbitals on the carbon atoms. The
two electrons in each of the sigma and the pi(π) bond are held firmly
between two carbon atoms.
This makes it easy to draw out a molecular structure like the one above; we can easily show where
the valency electrons are in the molecule
by simply looking at the positions of the covalent bonds.
However in a molecule such
as benzene the pi(π) electrons are spread over the whole
molecule, that is to say they are delocalised
in clouds of electron density above and below the flat planar ring of carbon atoms, this is shown in the image below.
To try and draw the structure of a benzene molecule
we often draw out two separate structures, as shown below. However
none of these structures represents the actual structure of a benzene
molecule. The best that we can
say is that the actual structure of a benzene molecule
is a combination of the two structures shown below.
The two structures shown below for benzene were proposed by Kekulé, he suggested, incorrectly that
the two
structures of benzene shifted back and forth between the two so that it was impossible to isolate any one
particular structure. However his idea
was wrong and we now know that the actual structure is a combination of both structures; we
would say that the two
structures are in resonance with each other.
It is important to be
aware that it is due to the limitations in the method we use to draw out the structures of
molecules that
stops us accurately drawing the real structure of benzene. So how would you describe the structure
of a benzene molecule?
Well you would have to say that its actual structure is a combination
or a composite
of the two different resonance forms. It is worth mentioning that
the only difference between these two resonance forms is in the position of the C=C. The only problem is that all
the carbon-carbon bond lengths in a benzene molecule are all the same length and are intermediate in length between
carbon carbon single bond and carbon carbon double bonds in length, so the C=C and C-C bonds
in the molecules shown below do not actually exist!
It is not only in benzene that we encounter the problem of accurately drawing the structure of a particular molecule or ion. If it is possible to describe the structure of a substance in such a way that the only difference between them is in the position of the electrons, usually the pi(π) electrons then the actual structure will be a composite or hybrid of the separate resonance structures. We would say that the actual structure is a resonance hybrid of all the separate structures. Resonance is the movement or delocalisation of electrons, usually in pi bonds, that is bonds which result from the partial overlap of p-orbitals, within molecules and ions. The nuclei of the atoms involved in resonance do not move.
The carbonate ion (CO32-) is another example of a molecule that cannot be drawn out using a single structure. Since carbon atoms make four bonds and oxygen atoms form two covalent bonds we might initially try to draw out the structure of the carbonate ion as a molecule containing two C-O and one C=O, with two of the oxygen forming and ion with a negative charge as shown below. However all the C-O bond lengths in the carbonate ion are the same length, so it cannot consist of single C-O bonds and a C=O bond. The actual structure of the carbonate ion will be a composite of the three resonance structure shown below:
To move from one resonance form to another it is simply a matter of moving an electron pair. In the diagram below the red arrows show the movement of a pair of electrons as one resonance form is changed into another in the forward direction. However you should bear in mind that these electron shifts DO NOT actually take place, remember that the electrons in the ACTUAL carbonate ion are delocalised. The drawing of these resonance structures is simply a limitation of the method used to draw the structures of these ions and molecules.
If you are unsure on how to start drawing resonance structures for molecules or ions a good place to begin is with drawing out Lewis structures. A Lewis structure should show the positions of the valency electrons within the molecule or ion. So for the carbonate ion we have:
So to try and work out the structure of the carbonate ion following the outine below:
As another example of resonance consider aromatic amines. Aliphatic amines and ammonia are good bases due to the presence of the lone pair of electrons on the nitrogen atom which is able to form dative covalent bonds with hydrogen ion (H+) in acids. As a simple example consider ammonia (NH3) which can use it lone pair of electrons to form a dative covalent bond to a hydrogen ion, as shown in the image below:
However aromatic amines; such as phenylamine or aniline (C6H5NH2) are poor bases. In an aromatic amine the nitrogen atom is bonded directly to an aromatic ring, this means that the lone pair of electrons on the nitrogen atom can be delocalised through the aromatic ring and so are not readily available to form dative covalent bonds in the same way aliphatic amines can. This is outlined in the diagram below:
In aniline the NH2 group is attached to an aromatic ring. The aromatic ring contains 6 delocalised electrons, one from each carbon atom in the aromatic ring. The p-orbitals in the ring can merge with the lone pair of electrons on the nitrogen atom, which are also in a p-orbital. This means that the lone pair of electrons on the nitrogen atom is delocalised through the aromatic ring system and not available to form dative covalent bonds, that is aromatic amines are poor bases because the nitrogen lone pair is delocalised through the aromatic ring.
The delocalisation of the pi(π) electrons will lead to the same problems we had above when trying to draw the structure of a benzene molecule. Delocalisation implies that the electrons are not in fixed positions so trying to draw molecular structures and bonds will not be possible. Instead we have to try and show the structure of aniline by drawing a series of resonance structures, shown below. Recall that the actual structure of aniline will be a mix or composite of all the individual resonance structures.
