Resonance Structures, Electron Mobility and Delocalization in Organic Chemistry, Lecture notes of Chemistry

Download Resonance Structures, Electron Mobility and Delocalization in Organic Chemistry and more Lecture notes Chemistry in PDF only on Docsity! 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: 1. 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. 3. 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. 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 CH3 CH3 H N H CH3 CH3 (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 CH3 1 2 H N H O CH3 (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: H3C CH3 O H3C CH3 O (d) π-electrons can also move to an adjacent position to make new π-bond. Once again, the octet rule must be observed: One of the most common examples of this feature is observed when writing resonance forms for benzene and similar rings. V. DELOCALIZATION, RESONANCE, AND CONJUGATION. The presence of alternating π-bonds and σ-bonds as in benzene is known as a conjugated system, or conjugated π-bonds. Conjugated systems can extend across the entire molecule, as in benzene, or they can comprise only part of a molecule. A conjugated system always starts and ends with a π-bond (i.e. an sp2 or an sp-hybridized atom). The atoms that form part of a conjugated system in the examples below are shown in blue, and the ones that do not are shown in red. Most of the times it is sp3 hybridized atoms that break a conjugated system. 1 2 3 benzene Practically every time there are π-bonds in a molecule, especially if they form part of a conjugated system, there is a possibility for having resonance structures, that is, several valid Lewis formulas for the same compound. What resonance forms show is that there is electron delocalization, and sometimes charge delocalization. All the examples we have seen so far show that electrons move around and are not static, that is, they are delocalized. Charge delocalization is a stabilizing force because it spreads energy over a larger area rather than keeping it confined to a small area. Since electrons are charges, the presence of delocalized electrons brings extra stability to a system compared to a similar system where electrons are localized. 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