# What is Resonance?

##### Add yours
Resonance Structures/Assigning Formal Charge

Tip: This isn't the place to ask a question because the teacher can't reply.

## Key Questions

• Resonance structures are imaginary Lewis structures which represent a compound that cannot be shown with one Lewis structure.

Consider ethanoate anion, for instance. It can be shown with the Lewis structure below. The negative charge is on the blue oxygen, and the red oxygen forms a double bond to the carbon. (NB: I used colors for oxygens to differentiate between them. In reality, they are absolutely identical)

But who forbids me to draw the structure with the negative charge on the red oxygen, and the blue oxygen forming a double bond to the carbon? Noone does!

But which one of the 2 Lewis structures is correct?! The answer is none of them. If we measure the C-O bond length and the charge density around oxygens, we will find that both C-O bonds are absolutely identical in length, and the negative charge is evenly spread around oxygens (see the electrostatic map below). Moreover, the length of the C-O bond is in between the length of a single and double C-O bonds.

Well, that means that none of the Lewis structures represent the actual structure of the molecule. And it's impossible to show a 1.5 bond with a Lewis structure. So, how to show that the 2 C-O bonds are indetical and their multiplicities are 1.5 ??? Well, we can draw both structures, and show that the actual structure is an average of the two. It will look like this:

This fancy arrow shows that the real structure inherits something from the first structure, and something from the second structure. It's a hybrid of the two . But it doesn't oscillates between the two structures! It's not an equilibrium! The two structures do not exist ! These structures are only used to show the molecule on paper.

• #### Answer:

The rules for drawing resonance structures are:

#### Explanation:

• You can never move atoms.
• You can move only π electrons or lone pairs that are in $p$ orbitals.
• All resonance structures must have the same number of valence electrons.

You can never move atoms

If atoms move, we have isomers, not resonance contributors.

The structures above are resonance contributors, because only electrons have moved. All atoms are in the same position in each structure.

You can move only π electrons or lone pairs that are in $p$ orbitals

The moving electrons must be on a "donor atom." The "acceptor atom" must be next to the donor atom.

The acceptor atom must have a positive charge or be able to accept an electron pair.

Electrons move towards a positive charge or to a more electronegative atom.

In the first example above, a lone pair of electrons on $\text{O}$ moves toward a positive charge. It forms a π bond between $\text{O}$ and the adjacent atom.

In the second example, a pair of π electrons on $\text{C}$ moves toward a positive charge. It forms a π bond on the other side of the $\text{C}$ atom.

In the third example, a pair of π electrons moves onto the more electronegative $\text{O}$ atom. This results in the formation of formal charges.

All resonance structures must have the same number of valence electrons.

Electrons are not created or destroyed. You must have as many electrons in the structures that you create as there were in the starting structure.

The rule is violated above because structure E has 12 valence electrons and structure F has 14 valence electrons. So E and F are not resonance structures (F also violates the octet rule).

• #### Answer:

Resonance structures should be used whenever the Lewis structure/molecule has resonance.

#### Explanation:

If a molecule has resonance and you write it as just a Lewis structure you're only representing a fraction of what that molecule looks like as a whole. Common ways of spotting resonance is a pi-bond next to a positive charge ie. any way of distributing charge to stabilize the molecule.

## Questions

• · 3 months ago
• · 4 months ago
• · 5 months ago
• · 7 months ago
• · 11 months ago
• · 1 year ago
• · 1 year ago
• · 1 year ago
• · 1 year ago
• · 1 year ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago
• · 2 years ago