Resonance

Notice that the more likely structure for the nitrite anion (NO2–) in the figure below may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O bonds:

In the left structure above the blue oxygen is making a double bond to nitrogen, while in the right structure the red oxygen is making the double bond to nitrogen.

If nitrite ions do indeed contain a single and a double bond, we would expect the two bond lengths to be different. A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms (review in the bond length and strength page). Experiments show, however, that both N–O bonds in NO2− have the same strength and length, and are identical in all other properties.

It is not possible to write a single Lewis structure for NO2− in which nitrogen has an octet and both bonds are equivalent. Instead, we use the concept of resonance: if two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures. The actual distribution of electrons in each of the nitrogen-oxygen bonds in NO2− is the average between a double bond and a single bond. We call the individual Lewis structures resonance forms. The actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the individual resonance forms. A double-headed arrow between Lewis structures indicates that they are resonance forms.

Let’s now practice how we can push electrons between the two structures, if you are yet to look at curly arrows and arrow pushing please check out the link here.

The resonance hybrid for NO2– is therefore a combination of both of these structures. Think about placing the two structures on top of each other, this is what our hybrid should look like:

Since the resonance hybrid is a combination of the two resonance structures we see that we have a -1/2 formal charge on each of the oxygens and the bonds between the N and Os is 1 and 1/2 or 1.5 (or 3/2).

We should remember that a molecule described as a resonance hybrid never possesses an electronic structure described by either resonance form. It does not fluctuate between resonance forms; rather, the actual electronic structure is always the average of that shown by all resonance forms. George Wheland, one of the pioneers of resonance theory, used a historical analogy to describe the relationship between resonance forms and resonance hybrids. A medieval traveler, having never before seen a rhinoceros, described it as a hybrid of a dragon and a unicorn because it had many properties in common with both. Just as a rhinoceros is neither a dragon sometimes nor a unicorn at other times, a resonance hybrid is neither of its resonance forms at any given time. Like a rhinoceros, it is a real entity that experimental evidence has shown to exist. It has some characteristics in common with its resonance forms, but the resonance forms themselves are convenient, imaginary images (like the unicorn and the dragon).

The carbonate anion, CO32−, provides a second example of resonance:

There are three equivalent resonance structures for CO32-. In the image above we have colour coordinated the oxygens so you can see how the bonding is changing. Looking at each oxygen (either the blue, red or green) we can see that in one resonance structure either the red/blue or green oxygen is making a double bond, while in the other two structures that specific oxygen is making a single bond and has a negative formal charge. To determine the resonance hybrid it is very important to look at the three resonance structures and think about them on top of each other when creating the resonance hybrid.

One oxygen atom must have a double bond to carbon to complete the octet on the central atom in each resonance structure. All oxygen atoms, however, are equivalent, and the double bond could form from any one of the three atoms. This gives rise to three resonance forms of the carbonate ion (CO32-). Because we can write three identical resonance structures, we know that the actual arrangement of electrons in the carbonate ion is the average of the three structures. Again, experiments show that all three C–O bonds are exactly the same.

Before we draw the hybrid let’s first practice arrow pushing again by showing how the three structures can interconvert:

In the above question, you looked at how the left structure went to the middle structure and how the middle structure became the right structure. Now which of the following options would show how the right structure becomes the middle structure?

Now that we have drawn all the three resonance structures, think about these three resonance structures on top of each other when deciding which is the resonance hybrid. It is important that you always draw out all the best resonance contributors before determining the hybrid.

The online Lewis Structure Maker includes many examples to practice drawing resonance structures as well.