After all these posts about resonance, I thought it would be good to have a post summarizing what’s been discussed so far.
One of the key skills in analyzing the reactivity of a molecule is to be able to figure out where the electrons are.
As I wrote here, if we’re dealing with single bonds, it’s a relatively straightforward matter of figuring out the differences in electronegativities.
However if multiple bonds (π bonds) are present, then we start to run into a little problem: there can be multiple ways to draw the same molecule – and thus, figuring out electron densities on a given molecule is not always as straightforward.
This situation – when we can draw two (or more) forms of the same molecule that differ only in the placement of their electrons – is called resonance, and the different structures we draw are called resonance forms.
How can we “find” resonance forms for a given molecule? It’s possible to do it through trial-and-error, but one surefire way is to do so is to apply the curved arrow formalism, which is a way of depicting the “movement” of electrons.
Here’s an important point about resonance forms. It is tempting (and very wrong!) to think that these resonance forms are in “equilibrium” between each other. Avoid this common mistake!
Instead, the “true” state of the molecule will be a “hybrid” of these resonance forms.
In this case both resonance forms are equal in energy, so the “hybrid” is a 1:1 mixture of the two. However this is only rarely the case, for instance in this ketone below, which has 3 different resonance forms.
Not all resonance forms will be of equal significance. How do we evaluate how important they are?
Resonance forms become less significant as the number of charges are increased. For example, in the ketone above, the resonance form with zero charges will be the most significant. (Note however, that each resonance form has a net charge of zero.
Resonance forms where all atoms have full octets will be more significant than resonance forms where atom(s) lack a full octet. Importantly, it’s a good general rule never to place less than a full octet on nitrogen or oxygen, as in the example above, right. Since these atoms are highly electronegative, these resonance forms are extremely unstable and will be insignificant.
Negative charges are most stable when placed on the least basic atom. There are four main trends to consider here:
- Electronegativity: across a row of the periodic table, negative charge becomes more stable as electronegativity is increased.
- Polarizability: down a column of the periodic table, negative charge becomes more stable as polarizability increases
- Electron withdrawing groups stabilize negative charge through inductive effects.
- Hybridization: negative charge becomes more stable as the s-character of the atom is increased. sp (most stable) > sp2 > sp3 (least stable
Note that stability is the opposite of basicity.
When dealing with positive charges, the resonance form where all octets are filled will be the best (see principle #2). Between resonance forms where there is a positive charge that has less than a full octet (that is, a carbocation), then follow these principles:
- Place positive charge on the most substituted carbon
- Avoid placing positive charge adjacent to electron withdrawing groups if possible
- Place positive charge preferentally on sp3 > sp2 > sp
When double bonds are connected to an atom with a lone pair of electrons, the molecule will have a significant resonance form where there is negative charge on the adjacent carbon.
When double bonds are connected to a polarized π bond, the molecule will have a significant resonance form where there is positive charge on the adjacent carbon.
For now, that does it for a summary of the important themes in resonance. Next stop (after a post about some common mistakes) will be to apply these principles to chemical reactivity.