Bonding, Structure, and Resonance
In Summary: Evaluating Resonance Structures
Last updated: November 26th, 2019 |
The Four Key Factors In Evaluating Resonance Structures
Not all resonance forms are of equal importance. So how do we evaluate how “important” each resonance structure is?
As we’ve seen in previous posts, four key factors that determine the importance of resonance structures in organic chemistry are:
- Rule #1: Minimize charges
- Rule #2: Full octets are favored
- Rule #3: How stable are the negative charges?
- Rule #4: How stable are the positive charges?
Today, let’s summarize everything we’ve learned about resonance structures in this unit.
Table of Contents
- Recall The Three “Legal” Electron-Pushing Arrow “Moves” Used For Interconverting Resonance Structures
- Remember That Resonance Structures Are Not In Equilibrium With Each Other – They Represent Contributions To An Overall Resonance “Hybrid”
- Not All Resonance Forms Are Of Equal Significance. So How Do We Evaluate How “Important” Each One Is?
- Rule #1: Neutral Resonance Structures Are More “Important” Than Charged Resonance Structures
- Rule #2: Full Octets Are Preferable To Empty Octets (And Never, Ever Have Empty Octets On Oxygen or Nitrogen!)
- Rule #3: Place Negative Charges On The Atom Best Able To Stabilize It (i.e. The Least Basic Atom)
- Rule #4: Place Any Empty Octets On the Atoms Best Able To Stabilize Them (i.e. Carbon And Not Oxygen Or Nitrogen)
- An Application Of Resonance: “Pi Donation”
- A Second Application Of Resonance: “Pi Acceptors”
- Quiz Yourself! (On Evaluating Resonance Structures)
1. Recall The Three “Legal” Electron-Pushing Arrow “Moves” Used For Interconverting Resonance Structures
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 distribute electrons on the same molecule (i.e. different resonance forms). Therefore, in order to understand electron density on a molecule where pi bonds are present, we must first understand the importance of its various 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.
There are three “legal” ways to move electrons using curved arrows: from pi bond to lone pair, from lone pair to pi bond, and from pi bond to pi bond:
2. Remember That Resonance Structures Are Not In Equilibrium With Each Other – They Represent Contributions To An Overall Resonance “Hybrid”
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.
For example in the acetate and allyl cation examples below, the “true” structure of the molecule is represented through a 50:50 combination of the two resonance forms.
3. Not All Resonance Forms Are Of Equal Significance. So How Do We Evaluate How “Important” Each One Is?
In the case of the acetate ion and the allyl cation, 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.
Take the ketone below (acetone, or “propanone”) for which we can draw 3 different resonance forms.
In cases like these, how do we evaluate the relative importance of each resonance form?
Resonance forms become less significant as the number of charges are increased (see earlier post).
For example, in the ketone above, the resonance form with zero formal charges will be the most significant.
How do we know? We can measure the physical properties of the molecule (e.g. boiling points, solvent properties, conductivity) and see if it’s more consistent with a charged species or a neutral compound.
All the physical properties of propanone (acetone) are consistent with it being a (mostly) neutral molecule. For example acetone has a boiling point of 56°C, significantly lower than water, and a freezing point of –95°C. It doesn’t dissolve charged species (like NaCl) nearly as well as water does. And it isn’t a particularly good conductor of electricity.
That isn’t to say that the “second-best” resonance form doesn’t play some role. Acetone is much higher-boiling than butane (–1°C) which has a similar molecular weight due to the dipole-dipole Van der Waals attractive forces, and as we’ll see later, the “second-best” resonance form can yield an important clue as to the reactivity of a molecule.
5. Rule #2: Full Octets Are Preferable To Empty Octets (And Never, Ever Have Empty Octets On Oxygen or Nitrogen!)
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 acetone example (above right). Since these atoms are highly electronegative, these resonance forms are extremely unstable and will be insignificant.
6. Rule #3: Place Negative Charges On The Atom Best Able To Stabilize It (i.e. The Least Basic Atom)
Given that neutral resonance structures are preferred overall, when a resonance structure absolutely must bear a negative charge somewhere, place it on the atom best able to stabilize that charge. Since, in essence, “basicity is the opposite of stability”, this is the same as saying, “put the negative charge on the least basic atom”.
The good news here is that if you understand the factors that affect acidity, you also understand the factors which stabilize negative charge.
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 again that stability is the opposite of basicity.
7. Rule #4: Place Any Empty Octets On the Atoms Best Able To Stabilize Them (i.e. Carbon And Not Oxygen Or Nitrogen)
As we said above, full octets are best. However, when dealing with a resonance structure where there absolutely must be an atom with less than a full octet, then follow these principles:
- Place the empty octet on carbon, never oxygen or nitrogen
- Place place the empty octet on the most substituted carbon (remember carbocation stability)
- Avoid placing positive charge adjacent to electron withdrawing groups if possible
- Place positive charge preferentially on alkyl carbocations as opposed to alkenyl or (especially) alkynyl carbons.
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 due to a phenomenon called, “pi donation“. This becomes particularly important once you start learning about reactions of pi bonds.
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. This phenomenon is known as “pi-accepting” behavior, and these groups are known as “pi acceptors“.
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.