Lecture 8 – Resonance Structures

Lecture Preview

• What is resonance?
• How can we visualize resonance?
• What is electron delocalization and how does it explain resonance?
• How does resonance affect molecular geometry?

Lecture Content (Part 1)

In the previous lecture, we dived very deep into the covalent bonding hole. Here is a list of concepts we introduced last time. Don’t worry if there are still a good amount of confusion – the next discussion will wrap everything up. It’s hard to introduce concepts one by one when each concept is so tightly connected with other concepts. So, hang in there, and hopefully next discussion will help.

• Lewis structures
  • Total count of electrons (TC)
  • Formal count of electrons (FC)
  • The octet rule (TC = 8)
  • Formal charges (q* = V – FC)
  • Drawing Lewis structures

• Molecular geometries (VSEPR)
  • Electron domains = covalent bond groups + lone electron pairs
  • VSEPR, hybridization, and Lewis structures

• Visualizing 3D structures
  • Using molview.com to visualize
  • Connecting Lewis structures, molecular geometry, and hybridization together

This lecture, we will introduce one singular concept, while the rest of the lecture will be practice. This concept is difficult to understand, but with enough practice and exposure, we will be very familiar with it. This concept is called resonance.

Resonance Structures

Let’s get one thing straight first.

As far as we are concerned, resonance only affects electrons in pi bonds or p orbitals.

Recall in the previous lecture where we explored how to draw Lewis structures, we used $N_2O$ as an example and ended up with 2 structures. We decided that the right structure is more “appropriate”, but mentioned that both structures are all possible structural representations of $N_2O$.

Essentially, the molecule $N_2O$ cannot be accurately expressed by one Lewis structure – it has 2 Lewis structure expressions. Those 2 structures both contribute into the actual structure of $N_2O$. Those contributing structures are called resonance structures.

Converting between Resonance Structures

So, before we get into visualizing resonance, let’s see how Lewis structures can help us convert between resonance structures. We need to use electron-pushing arrows.

Electron-pushing arrows are arrows that describe electron movements. It basically describes the following:

• Electrons move from one place…
• to another place next to it…
• forming or breaking bonds.

So, theoretically, electrons can move from being a lone pair to becoming a covalent bond. Alternatively, they can move from being a covalent bond to being a lone pair. Alternatively, they can move from being a covalent bond to being another covalent bond.

Remember: each arrow represents the movement of a pair of electrons (2 electrons), not 1 singular electron.

Let’s use the $N_2O$ example again. Let’s try to convert the left structure to the right structure.

Side note: for optimal learning, draw alongside!

• Let’s first move the electron pair from the left nitrogen and move it to form the triple bond, forming the intermediate structure on the right. Let’s look at the analysis table below to check if there are any problems with the intermediate structure.

• Notice that for the middle N, we have 10 total electrons – a direct violation of the octet rule. As mentioned before, despite being not applicable for a lot of elements, the octet rule does fully apply on elements like C, H, O, and N, and mostly apply on halogens and 3rd row elements. So, this intermediate won’t stay long – it will convert swiftly to the next structure.
• To achieve that, the NO double bond becomes a single bond, with the 2 electrons from the broken bond moving to the oxygen lone pair.

Understandably, the intermediate structure looks “cleaner” because the formal charges are 0 among all atoms. However, when drawing Lewis structures, the octet rule (for applicable atoms) takes priority, then we move to formal charges. Recall the 4 rules for formal charges from the previous lecture:

So, in the end, we can write our resonance structure conversion as follows:

Conversely, we can convert the other way:

Realistically, we will draw resonance structure conversions using the above 2 diagrams. No intermediate structure is needed. Below is a Lewis structure visualization of $N_2O$ resonance structure conversion.

Resonance Contributors and Molecular Structure

We have already established that the Lewis structure of $N_2O$ is not an “either-or” scenario. Rather, it’s a “both-and” scenario. To formalize this, we can conclude the following:

Resonance structures all contribute to the accurate representation of a molecule.

In the example of $N_2O$, it basically means:

• The NN bond order is not either 3 (triple bond) or 2 (double bond). Rather, it has a bond order of around 2.5 (~2.5 bond).
• The NO bond order is not either 2 (double bond) or 1 (single bond). Rather, it has a bond order of around 1.5 (~1.5 bond).

As a result, if we basically add up the structures of the resonance contributors and take the average, we end up with a more accurate representation of the molecule.

This is what we should see whenever we see a nice molecule with resonance structures. We call those resonance structures “contributors”. Depending on the molecule, there can be many contributors. And depending on the stability of the contributors, we can further divide them into major and minor contributors.

• Major contributors have more favorable Lewis structures (recall the 4 formal charge rules).
• Minor contributors have less favorable Lewis structures.

For the $N_2O$ structure above, we can assign major and minor contributors.

Major Contributor

Minor Contributor

Refer to the formal charges rules here, and see why this is the case.

• For the major contributor, the negative formal charge is positioned at oxygen. For the minor contributor, the negative formal charge is positioned at the left nitrogen.
• Per rule 3, elements with higher electronegativity handle negative formal charges better, thus, the structure with oxygen having the negative formal charge becomes the major contributor.
• Since for both structures have a positive formal charge for the center nitrogen, there is no comparison to be made.

As a result…

• The NN bond order is not either 3 (triple bond) or 2 (double bond). Rather, it has a bond order of around 2.6.
• The NO bond order is not either 2 (double bond) or 1 (single bond). Rather, it has a bond order of around 1.4.

Let’s look at another example.

Using Example 3, we introduced a new covalent bond property – the ability to rotate things around it. We didn’t mention it before because it’s not relevant and there’s too much information being introduced already. Now we’re here, let’s talk about it.

Intramolecular structures can rotate along covalent bonds under certain conditions. The conditions being:

• The bond is a single bond (a quick glance at the Lewis structure).
• The bond does not participate in resonance (recall that resonance only involves p orbitals and pi bonds, something a single bond does not have).

As bond orders increase (1.5, 2, 3, etc.), it becomes harder to rotate intramolecular structures along that particular covalent bond. For simplicity’s sake, we will assume that:

• If the conditions above are not met, the structure may not rotate freely.

So, if we look at the $N_2O$ example again, the NO bond cannot rotate because despite being described as a single bond on the Lewis structure, the bond participates in resonance, so it’s more of a ~1.5 bond – no free rotation.

Again, don’t worry about it too much if there is confusion. We have plenty of dedicated practice coming up.

Resonance Orbital Visualization and Hybridization

To get a better understanding of resonance structures, we need to visualize. Recall our previous visualizations.

Also, recall our discussions regarding hybridization.

Notice that now that we have more concepts introduced, we can tie them together, slowly piecing the puzzle pieces to find out the result.

Let’s, once again, look at the $N_2O$ example: what are the visualizations of each contributor like?

Looking at the animations for 30 seconds, we might see a problem.

• On the left side structure, oxygen has 4 electron domains. According to our previous lecture, 4 electron domains correspond to sp3/tetrahedral properties. However, the visualization has sp2/trigonal planar properties
• Similarly, on the right side structure, the left nitrogen has 3 electron domains. According to our previous lecture, 3 electron domains correspond to sp2/trigonal planar properties. However, the visualization has sp/linear properties.

What is happening? The answer lies in how those resonance structures convert between each other.

Observe the electron-pushing arrows moving the electrons from one place to another, forming and breaking bonds during the process.

Since we know that the more accurate structure is a blend of both contributors, let’s try blending the two visualizations together. The blend will be a bit messy, so we’ll clean it up as well (getting rid of extra electrons from the blending).

Let’s freeze frame this structure.

This is the structure we are looking for – the accurate representation of $N_2O$ structure visualized. If we want to talk about molecular geometry and hybridization, this is the structure we should have in mind. So, according to this visualization, we have:

This is the correct hybridization and electron geometry we look for, because the visualization is accurate.

Also, observe that in resonance structure conversions, all electron movements are confined within p orbitals – no sp/sp2/sp3 orbitals or s orbitals are involved. Again, it’s very important to keep in mind that:

Resonance activity involves only p orbitals and pi bonds.

• If the Lewis structure seems to indicate that an atom is sp3 hybridized but it participates in resonance that forms a double bond, it’s probably sp2 hybridized – the lone pair stays in the p orbital to participate in resonance.
• If the Lewis structure seems to indicate that an atom is sp2 hybridized but it participates in resonance that forms a triple bond, it’s probably sp hybridized – the 2 lone pairs stay in the p orbitals, with one participating in bonding and the other in resonance.

Electron Delocalization

In the end, Lewis structures don’t really convey resonance really well. We need to find them ourselves and go out of our way to think of their hybridizations and molecular geometries. However, Lewis structures have more pros than cons, so we’ll stick with them.

Another way to think about resonance is electron delocalization in pi bonds.

Electron delocalization basically means the electrons are free to go wherever they want.

We saw an example of delocalization in metallic bonds back in Lecture 6, where electrons are freely moving amongst the sea of metal nuclei.

In that lecture, we introduced the concept as:

Electrons in a molecule, ion, or solid metal are not associated with a single atom or a covalent bond.

Now, if we think of resonance as “delocalization among pi bonds”, this visualization makes more sense.

The electrons in the dashed area are delocalized, meaning that can flow freely within the connected areas. Because they are flowing in where pi bond usually form, we call this “delocalization among pi bonds”. And since electrons all freely flow in the pi bond range, it forces the atoms to adapt certain hybridizations and geometries.

We will see resonances a lot in organic chemistry, so remember this concept well.

Summary

Here, we will tidy up our resonance structures and conclude what we need to know, what structures we are expected to draw, and how we should think of resonance structures.

Regarding drawing Lewis structures

When asked to draw Lewis structures, we should draw the most favorable structure. This means that:

1. The octet rule for C, H, O, N, F should be 100% satisfied with no exceptions.
   • Hydrogen is more of a duet, but we’ll count them together regardless.
   • Other atoms will mostly follow the octet rule.
2. With the octet rule fulfilled, the formal charge distributions need to be optimal. The preferences are ordered as follows.

Regarding resonance structures

After drawing the Lewis structure, observe if resonance is possible. If it is possible (and not trivial), draw the resonance structure(s) as well, following the rules below.

• Indicate the structures as resonance structures by writing $\leftrightarrow$ between the structures.
• Add brackets [] at the start and end as well.

$ [ \text{resonance structure 1} \leftrightarrow \text{resonance structure 2} ] $

Here’s a confusing convention – despite being a more accurate description of the molecule structure, we do not usually draw the structure with dashed lines. The reason behind this is in organic chemistry, we will dive deep into electron-pushing arrows. Dashed lines mess up our electron-pushing arrows and also make us disoriented on whether the atom fulfills the octet rule or not.

Pros	Cons
Accurately describes molecular geometry	Does not conveniently show octet rule fulfillment
Fits in visualization and the concept of delocalization nicely	Does not show formal charge and bad for electron-pushing arrow notation

Because we, in the future, will value electron-pushing arrow notation more than accurate description of molecular geometry (organic chemistry stuff), we will stick with the traditional Lewis structures.

However, whenever we are drawing structures with resonance, we should always visualize them in our minds (or, more preferably, our scratch paper) with dashed lines and delocalized electrons.

We will have a lot of practice coming up.

At this point, hopefully, we should start to see some kind of connection between a lot of concepts from this chapter. We won’t be talking about this connection until the unit reviews, but try organizing all the different connections we have. This is a fundamental part for organic chemistry. Since organic chemistry is all about how structures can affect reactivity, Lewis structures (by extension, molecular geometry, hybridization, movement freedoms, etc.) play an important role in it.

Lecture Content (Part 2)

Once again, this part of the lecture (ideally having 30-45 minutes left) will be more exercises. This time, it will be about resonance structures. Resonance structures are harder to visualize, so we’ll stick with our trusty paper and pen (or tablet).

We will be presented with 10 molecules, like last lecture. This time, our tasks are as follows.

• Task 1: Draw all possible non-trivial resonance structures of each molecule. Some can have 2 resonance structures, while some can have more.
  • Electron-pushing arrows are highly recommended because we will be using them a lot in organic chemistry.
  • Make sure that the octet rule is fulfilled for each atom. We will not be dealing with octet rule exceptions in this exercise.

• Task 2: Label the molecular geometries and hybridizations of the highlighted atoms.
  • Remember, electrons participating in resonance always involve p orbitals or pi bonds.
  • Note: Not all the highlighted atoms have electrons participating in resonance.

• Task 3: Try to visualize the delocalized electrons by using dashed lines to represent delocalization.
  • Think about how it affects the bond order of each covalent bond group.

The exercises are also available on the worksheet.

Resonance Structure Exercises

As usual, for better understanding, feel free to use Molview to visualize the 3D structure. However, as of 2025, the 3D visualizations of structures on Molview does not represent resonance.

Let’s start with some easy ones. The number of resonance structures (including the one presented) is in parenthesis.

From here on out, the molecules will be a bit more complicated. However, those molecules are quite important. Some molecules can have more than 3 resonance structures. In that case, find 3 resonance structures (including the one presented).

Molecules 9 and 10 are from the previous lecture. In the previous lecture, we didn’t really consider how resonance can affect molecular geometry because we didn’t learn about resonance structures yet. Consequently, we only used VSEPR to predict molecular geometry. Now that we learned about resonance structures, some molecular geometries we predicted last lecture should be different.

Molecules 5-10 are important biochemical molecules used as medication or found in our bodies.

• Thymine is one of the four nucleotide bases in the nucleic acid of DNA. Disruption of the nucleotide bases can cause mutations in the DNA.

• Histidine is an essential amino acid that is used in the biosynthesis of proteins. Initially thought essential only for infants, it has now been shown in longer-term studies to be essential for adults also.

• Nicotiamide is a derivative of vitamin B3 found in food and used as a dietary supplement and medication. As a dietary supplement, it is used to treat vitamin B3 (nicotinic acid) deficiency. As a medication, it can be used to treat acne. In biochemistry, especially metabolism, it is a part of the molecule NAD+/NADH – an essential electron carrier in our cells.

• Serotonin is a monoamine neurotransmitter. Its biological function is complex, touching on diverse functions including mood, cognition, reward, learning, memory, and numerous physiological processes such as vomiting and vasoconstriction.

• Aspirin is an NSAID (non-steroidal anti-inflammatory drug) used for pain relief, anti-inflammatory effects, and blood thinning (antiplatelet properties). It works by inhibiting cyclooxygenase (COX) enzymes, reducing prostaglandin production.

• Dopamine plays a major role in the reward system, affecting motivation, addiction, and learning. Low dopamine levels are associated with Parkinson’s disease and depression. High dopamine activity is linked to schizophrenia and certain addictions.