A reasonable Lewis structure for the nitrate polyatomic ion, NO3−, is below.
This Lewis structure shows two different types of bonds, single and double. Because it takes more energy to break a double bond than a single bond, we say that a double bond is stronger than a single bond. Double bonds also have a shorter bond length, the distance between the nuclei of the two atoms in the bond, than single bonds do. Thus, if the above Lewis structure for nitrate were correct, the nitrate polyatomic ion would have one bond that is shorter and stronger than the other two.
This is not the case. Laboratory analyses show all three of the bonds in the nitrate ion to be the same strength and the same length. Interestingly, the behavior of the bonds suggests they are longer than double bonds and shorter than single bonds. They are also stronger than single bonds but not as strong as double bonds. In order to explain how this is possible for the nitrate ion and for molecules and polyatomic ions like it, the valence-bond model had to be expanded.
The model now allows us to view certain molecules and polyatomic ions as if they were able to resonate between two or more different structures. For example, the nitrate ion can be viewed as if it resonates between the three different structures below. Each of these structures is called a resonance structure. The hypothetical switching from one resonance structure to another is called resonance, and the convention is to separate the resonance structures with double headed arrows.
It is important to stress that the nitrate ion is not really changing from one resonance structure to another, but chemists find it useful, in an intermediate stage in the process of developing a better description of the nitrate ion, to think of it as if it were doing so. In actuality, the ion behaves as if it were a blend of the three resonance structures.
We can draw a Lewis-like structure that provides a better description of the actual character of the nitrate ion by blending the resonance structures into a single resonance hybrid:
Draw the skeletal structure, using solid lines for the bonds that are found in all of the resonance structures.
Where there is sometimes a bond and sometimes not, draw a dotted line.
Draw only those lone pairs that are found on every one of the resonance structures. (Leave off the lone pairs that are on one or more resonance structure but not on all of them.)
The resonance hybrid for the nitrate polyatomic ion is
The actual geometry of the polyatomic ion is trigonal planar with bond angles of 120°.
Resonance Structures and the Resonance Hybrid
Resonance is possible whenever a Lewis structure has a multiple bond and an adjacent atom with at least one lone pair. The following is the general form for resonance in a structure of this type. The arrows show how you can think of the electrons shifting as one resonance structure changes to another.
You can follow these steps to write resonance structures.
Shift one of the lone pairs on an adjacent atom down to form another bond.
Shift one of the bonds in a double or triple bond up to form a lone pair. (You might find it useful to draw arrows indicating the hypothetical shift of electrons.)
Draw additional resonance structures by repeating this process for each adjacent atom with a lone pair.
Separate the resonance structures with double-headed arrows.
For example, the two resonance structures for the formate ion, HCO2− are
To generate the second resonance structure from the first, we imagine one lone pair dropping down to form another bond, and pushing an adjacent bond off to form a lone pair. The arrows show this hypothetical shift of electrons. These resonance structures lead to the resonance hybrid below.
This general procedure for drawing resonance structures will not always lead to a reasonable resonance structure. For example, fluorine atoms do not participate in resonance. According to the valence-bond model, for a fluorine atom to form two bonds and two lone pairs, it would have to lose an electron, a highly unlikely act for the most electronegative element on the periodic table. Thus any resonance structure that includes a double bond to fluorine is not considered a reasonable resonance structure. Thus, although fluoroethene, CH2CHF, has a double bond and an adjacent atom with a lone pair (components that suggest the possibility resonance), only one of its two hypothetical resonance structures is reasonable:
The first structure is reasonable, but the second structure does not contribute to the resonance hybrid in a significant way. Therefore, fluoroethene does not have resonance, and the first structure above is the best description of a CH2CHF molecule.
There is a similar situation with oxygen atoms. Although it is possible for oxygen atoms to have three bonds and one lone pair, it is not likely that the second most electronegative element would lose the electron necessary to make this possible. Thus we will eliminate resonance structures that have three bonds and a lone pair for an oxygen atom. For example, formic acid, HCO2H, has a double bond and an adjacent atom with a lone pair, so we might think that it has resonance. The two resonance structures would be
The first Lewis structure is reasonable, but the second one, with three bonds and a lone pair on an oxygen atom, is not considered a reasonable resonance structure. Therefore, there is no significant resonance for formic acid, and the first Lewis structure above is the best description of its structure.
We will consider resonance a possibility for molecules and polyatomic ions that have the following as part of their Lewis structure.
Z can have more than one lone pair.
X and Y can have lone pairs.
The X-Y bond can be a triple bond. The Y-Z bond can be a double bond.
Z cannot be F with one bond and three lone pairs or O with two bonds and two lone pairs.
EXAMPLE 1 – Drawing Resonance Structures:
A reasonable Lewis structure for H2NCOCH3 is below. Predict whether it would have resonance. If it does, draw all of the reasonable resonance structures and the resonance hybrid.
The structure has a double bond and an adjacent atom with a lone pair, so it could have resonance. The possible resonance structures are below.
The first structure has the most common bonding pattern for all of its atoms, so it is a reasonable Lewis structure. Although the less common bonding patterns for the oxygen and nitrogen atoms in the second structure suggest that it is less stable than the first structure, we still consider it to be a reasonable resonance structure.
Using the two resonance structures above as a guide, we get the following resonance hybrid.
Expanded Lewis Structure Drawing Procedure
When resonance is considered, we add another step to our Lewis structure drawing procedure.
Once you have a reasonable Lewis structure, consider the possibility of resonance. If resonance is possible, draw the reasonable resonance structures and the resonance hybrid for the structure.
EXAMPLE 2 – Drawing Resonance Structures:
Draw a reasonable Lewis structure for the oxalate ion, C2O42−. The structure is best described in terms of resonance, so draw all of its reasonable resonance structures and the resonance hybrid that summarizes these structures. Ionic compounds containing the oxalate ion have many uses, including the bleaching and cleaning of textiles.
Step 1: C2O42− valence e− = 2(4) + 4(6) + 2 = 34
(Remember to add the two electrons for the −2 overall charge.)
Step 2: Oxygen atoms rarely bond to each other, but carbon atoms do, so we place the carbon atoms in the center of the structure and attach the oxygen atoms to them. If we put two oxygen atoms on each carbon atom, we are more likely to get a final structure that satisfies the requirement for four bonds to each carbon atom.
Step 3: e- remaining = 34 − 5(2) = 24
Step 4: Because oxygen atoms commonly have one bond and three lone pairs, we try the experiment of placing the remaining electrons as three lone pairs on each oxygen atom. This leaves the carbon atoms with only six electrons each, so we know that we will need to convert lone pairs into bonds in Step 5.
Step 5: Because we are short four electrons (or two pairs) to provide octets for the carbon atoms, we convert two lone pairs into bonds.
Step 6: The carbon atoms and two of the oxygen atoms (the ones with two bonds and two lone pairs) have their most common bonding pattern. The oxygen atoms with one bond and three lone pairs lead us to Step 7.
Step 7: There are no reasonable alternatives.
Step 8: With its double bonds and adjacent atoms with lone pairs, our structure meets one of the requirements for resonance. To compose the resonance structures, we imagine the electron pairs shifting as shown by the small arrows below. It is as if a lone pair drops down to form another bond, pushing a bond off to form a lone pair. Remember that we do not believe this is really happening. We just find it is useful to think of resonance structures in this way.
We follow the guidelines to draw the resonance hybrid that summarizes these structures and provides the best description of the bonds in the oxalate ion:
Resonance and the Benzene Molecule
It is possible to have resonance without the participation of lone pairs. The most important examples of this are benzene, C6H6, and compounds that contain the benzene ring. Benzene’s six carbon atoms are linked to each other in a six-membered ring. Its Lewis structure is often represented with three double bonds as shown below, but chemists often simplify it by leaving off the element’s symbols and the carbon-hydrogen bonds.
The Lewis structures above depict the benzene molecule as if it contained two types of C-C bonds, double and single. In actuality, all of benzene’s C-C bonds appear to be the same, and we can explain why in terms of resonance. It is as if the benzene ring were resonating between the two structures below.
The resonance hybrid is Structure 3 below. Because it is a bit tedious to draw all the dots, the structure of the benzene molecule is often written as shown in Structure 4, with the dotted lines represented by a circle.
In summary, Structures 1, 2, 3, and 4 are all used to describe benzene.