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RESONANCE STRUCTURES, ELECTRON MOBILITY
AND DELOCALIZATION
I. RESONANCE STRUCTURES. Lewis formulas are misleading in the sense that atoms and electrons are shown as being static.
But we know that a given compound can have several valid Lewis formulas. For example CH3CNO can be represented by at least three different Lewis structures called resonance forms, or resonance structures :
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, but some electrons have changed locations. The basic bonding pattern that is unique to a specific compound is made up of sigma bonds. This is called the connectivity. The connectivity is the same in all the resonance structures.
II. ELECTRON MOBILITY. Electrons, on the other hand, can be moved around. That is to say, they possess a certain degree of
mobility. Further examination of the examples above shows that sigma bonds don't move, for that would amount to destroying the
connectivity and therefore the molecule. Some of the electrons forming part of π -bonds and unshared electron pairs, however,
have been moved.
For example, in structure I a carbon and a nitrogen are joined by a triple bond (that is, a sigma bond and two pi bonds), whereas in structure II the same atoms are joined by a double bond (that is, one sigma bond and one pi bond). Somehow one of the electron pairs that makes up the triple bond in I has been moved to a different location in II, and even in III.
H C C N O
H
H
H C C N O
H
H
H C C N O
H
H
I II III
III. ELECTRON “PUSHING” AND THE CURVED ARROW FORMALISM. The movement of mobile electrons in chemical structures
and in reaction mechanisms is indicated using the curved arrow formalism. Small, curved arrows indicate the movement of electron pairs , be it from a bond or an unshared pair. For example, the movement of electrons used to arrive at structure II from I is:
H C C N O
H
H
H C C N O
H
H
I II
Likewise, a lone pair from oxygen in I and II has been moved in III. The total number of electrons in the various resonance forms cannot change, for that would change the identity of the species. So there must be certain rules by which mobile electrons can be moved to and from certain locations. The strength of the bond determines how mobile the electrons that make it up are. The following criteria apply:
- Sigma bonds are the strongest type of bond and represent the “glue” that holds the atoms together. Therefore, such electrons are rarely moved. There are, however, some exceptions.
2. The π -bonds in double or triple bonds can move without destroying the connectivity , since the σ-bond remains untouched.
However, a certain energy investment must be made to break the π-bond before these electrons can be moved.
- Unshared electron pairs (or lone pairs) represent the most mobile type of electrons. They do not form part of any bonds and therefore no energy has to be spent breaking a bond before they can be moved.
Therefore, the order of elctron mobility and/or availability for reactions is:
Unshared electron pairs > π-bond electrons > σ-bond electrons.
Now, where can those electrons go? Can they move in any direction, or are there rules that govern their movement as well? There is a set of rules, but the most important thing to keep in mind is that electrons move towards areas of lower electron density, and/or away from negative charges.
This example illustrates how a lone pair of electrons (from carbon) can be moved to make a new π-bond (between carbons), and
how a π-bond (between carbon and oxygen) can be moved to make a new lone pair (on oxygen):
We have observed the two rules for moving electrons in resonance structures. We only moved electron pairs to adjacent, or neighboring positions. The Lewis structure that resulted from steps 1 and 2 is valid, and the net charge in both structures is -1.
Using the same example, but moving electrons in a different way, illustrates how such movement would result in invalid Lewis formulas, and therefore is unacceptable. The resulting structure violates several conventions. First, the central carbon has five bonds and therefore violates the octet rule. Second, the overall charge of the second structure is different from the first.
C
C CH 3
O
H
H
C
C CH 3
O
H
H
1
2
C
C CH 3
O
H H
C
C CH 3
O
H H
Likewise in the example below. First, electrons are moving towards an area of high electron density (a negative charge). Second, the octet rule is violated for carbon in the resulting structure.
C
C CH 3
O
H H
C
C CH 3
O
H H
Additional examples illustrate more specific rules that apply to electron movement when writing resonance structures:
(a) Unshared electron pairs (lone pairs) located on a given atom can only move to an adjacent position to make a new π-bond to
the next atom.
As the electrons from the nitrogen lone pair move towards the neighboring carbon to make a new π-bond, the π-electrons making
up the C=O bond must be displaced towards the oxygen to avoid ending up with five bonds to the central carbon.
H N H
CH 3
CH 3
H N H
CH 3
CH 3
(b) Unless there is a positive charge on the next atom (carbon above), other electrons will have to be displaced to preserve the
octet rule. In resonance structures these are almost always π-electrons.
H N H
O
CH 3
1 (^2) H
N H
O
CH 3
(c) As can be seen above, π-electrons can move towards one of the atoms they bond to form a new lone pair. In the example
above, the -electrons from the C=O bond moved towards the oxygen to form a new lone pair. Another example is:
H 3 C CH 3
O
H 3 C CH 3
O
Since conjugation brings up electron delocalization, it follows that the more extensive the conjugated system, the more stable the molecule (i.e. the lower its potential energy). If there are positive or negative charges, they also spread out as a result of resonance.The corollary is that the more resonance forms one can write for a given system, the more stable it is.
Examine the following examples and write as many resonance structures as you can for each to further explore these points:
more stable than
more stable than or^ or or
more stable than more stable than
In the last example, the positive charge and the π-bonds are delocalized around the entire ring in the first structure, and over three
carbon atoms in the second. The third structure has no delocalization of charge or electrons because no resonance forms are possible.
I (^) II III IV V
I
I II
VI. DELOCALIZATION SETUPS. There are specific structural features that bring up electron or charge delocalization. The presence of a conjugated system is one of them. Other common arrangements are:
(a) The presence of a positive charge next to a π-bond. The positive charge can be on one of the atoms that make up the π-bond,
or on an adjacent atom.
(b) The presence of a positive charge next to an atom bearing lone pairs of electrons.
H 3 C CH 3
OH
H 3 C CH 3
OH
H 3 C C O H 3 C C O
N N O N^ N^ O
(c) The presence of a π-bond next to an atom bearing lone pairs of electrons.
H (^) N
O
CH 3
CH 3
1
2
H (^) N
O
CH 3
CH 3
(^2) O 1 O
