Discussion: Consider the Lewis structure of the carbonate ion, CO32-. The Lewis structure for this ion has a carbon-oxygen double bond, and two carbon-oxygen single bonds. Each of the singly bonded oxygen atoms bears a formal charge of 1-. (Review the formal charge tutorial if needed.) But which of the three oxygens forms the double bond? There are three possibilities:
These structures are similar in that the have the same types of bonds and electron positions, but they are not identical. The position of the carbon-oxygen double bond makes them different. In structure A the double bond is with the top oxygen atom, in B with the right hand oxygen atom, and C with the left hand oxygen atom. These oxygen atoms are at different places in space, so these are different structures. Consider this analogy: when the hands on a clock are at a 90o angle, the time could be 3 o'clock, or 6:15. The angle between the hands stays the same, but because they point to different places in space, they indicate a different time. The position of the carbon-oxygen bond is like the fixed angle of the clock hands, but pointing to different places on the clock face.
When more than one Lewis structure can be drawn, the molecule or ion is said to have resonance. The individual Lewis structures are termed contributing resonance structures. Resonance is a common feature of many molecules and ions of interest in organic chemistry.
Which one of these three structures is the correct one? How could we tell? If structure A was correct, laboratory measurements would show one shorter bond (the carbon-oxygen double bond) and two longer bonds (the carbon-oxygen single bonds). Measurement of structures B and C would give the same results as well. As it turns out, laboratory measurements show that all three bonds are equal and between single and double bond length. This suggests that none of the Lewis structures we have drawn are correct. It further suggests that the actual structure has three equal carbon-oxygen bonds that are intermediate between single and double bonds.
Perhaps the three Lewis structures for carbonate ion are in rapid equilibrium. The structures are changing so quickly that all we see can measure is an average blur (structure D), instead of being able to detect individual structures. By analogy, consider a camera with the shutter left open. The picture would be a blur that looks like A, B and C all at the same time. Structures A, B and C have the same bonds and electron distribution, the only difference is the position of the bond. Thus the three structures have equal stability, and the three structures would occur to the same extent at equilibrium. The blur we see would look like 1/3 A, 1/3 B and 1/3 C. The bond lengths would be a blur as well; we would perceive them as being something between single and double bonds. The charge would shift so rapidly that we would see it on all oxygen atoms at once. Since each oxygen atom has a 1- charge in two of the three equilibrium contributors, each oxygen atom would appear to have, on average, a charge of 2/3-.
All laboratory experiments have failed to detect structures A, B and C. No matter what experiments are performed, analysis has always concluded that D is the best description for the structure of the carbonate ion. This suggests that A, B and C do not exist, and are not adequate descriptions for carbonate ion at any time. The actual structure is D, and not equilibrium between A, B and C.
That D appears to be a combination of A, B and C still appears to be a useful way to determine the actual structure of carbonate ion. The only problem is that A, B, C and D are not in equilibrium. Can they be interacting in some other way? The answer is that the true structure of carbonate ion appears to be a simultaneous hybrid of the three resonance contributors A, B and C. Structure D has features derived from A, B, and C, but is never just A or just B or just C. For example, D has some double bond character, and each oxygen has some negative charge. Hybrid structure D is termed the resonance hybrid structure. It is a much better representation of reality for carbonate ion than any of the individual contributing resonance structure. We use a double headed arrow to show that individual structures are related by resonance.
It is important to understand that structure D is a hybrid of A, B and C. It has features common to A, B and C, but carbonate ion does not rapidly interconvert between A, B and C. By analogy, consider the case of a mule. A mule is the offspring of a horse and donkey. It has features similar to both, such as long tail and mane, but it does not rapidly interconvert between the two. A mule is not a horse one second, then a donkey the next. (An animal that changed would certainly be the highlight of any zoo!)
Drawing contributing resonance structures. Resonance is an important feature of many organic molecules. It can have a profound influence on their structure, chemical reactions, and physical properties. Key to understanding resonance is the ability to draw contributing resonance structures and the resonance hybrid structure.
Resonance structures are simply alternate Lewis structures for a given ion or molecule. Thus we can draw all resonance structures by drawing all of the possible Lewis structures. However, it is not always easy to see what all these Lewis structures might be. A set of Lewis structures for a give ion or molecule must have the same number of electrons as the Lewis structures are constructed from the same atoms. The only difference between the Lewis structures is the placement of the electrons. The position of the atoms in space is held constant. We can use this facts to assist us in drawing resonance structures. Because the number of electrons is conserved, electrons taken away from one atom must appear somewhere else in the structure. When drawing resonance structures, it is most convenient to shift these electrons between adjacent atoms.
Electrons that can be moved between adjacent atoms in resonance structures are lone pairs or p electrons. (Unpaired electrons present in radicals will be considered in section 15.5 of the text.) These electrons reside on a "resonance donor atom." The "resonance acceptor atom" must be adjacent to the donor atom. The acceptor atoms must also have an open octet, be able to accommodate an expanded octet (text section 1.1C; most commonly this will be chlorine, bromine, sulfur and phosphorus) or have another electron pair (lone pair or p) that can be displaced. An atoms with a formal positive charge can also be an resonance acceptor atom, as long as the atom does not accept more electrons that it can normally accommodate.
We use curved arrows as a bookkeeping tool to indicate the electron changes that differentiate resonance structures. Although these are the same curved arrows we use to indicate that bond changes are occurring in a reaction mechanism, they do not mean exactly the same thing. A reaction step requires a finite amount of time to occur, whereas a shift between resonance structures never actually occurs. Recall that individual resonance structures do not exist, they are simply alternate Lewis structures for the same molecule or ion. In the example below, the curved arrow indicates that the lone pair of electrons on the left-hand carbon is moving to a position where it can be shared between the two carbons, thus becoming the carbon-carbon p bond. This use of curved arrows is often termed "electron pushing."
Because electrons are shifting around, the formal charge
will vary between resonance structures. The formal charge of an
that gains a pair of electrons through resonance becomes one unit more
negative. Conversely, the formal charge of an atom that shares a
pair of electrons that it did not share previously becomes one unit
Rules for drawing contributing resonance structures. Some rules must be considered when drawing resonance structures.
Rule 1: All resonance structures must have the same number of valence electrons. Electrons are not created or destroyed, nor are they lost or gained to other molecules or ions during resonance.
The rule is violated because structure E has 12 valence
(four bonding pairs and two lone pairs), whereas structure F
14 valence electrons (five bonding pairs and two lone pairs).
these cannot be resonance structures of the same ion. (Structure
also violates rule 2.)
Rule 2: The octet rule must be obeyed. Hydrogen may never have more than two valence electrons. Lithium through fluorine may never have more than eight valence electrons. (A carbon with five attachments, often called a "pentavalent carbon" has ten valence electrons. This is forbidden because carbon does not have the space in its orbitals to accommodate ten electrons. You should take care to avoid this common mistake made by inexperienced organic chemistry students.) Certain elements commonly encountered in organic chemistry may have ten valance electrons. These elements are in periods three and higher in the periodic table, and include chlorine, bromine, iodine, phosphorus and silicon. Of all the atoms commonly encountered in organic chemistry, only sulfur can routinely expand its octet to include twelve valence electrons. These atoms expand their octets so as to improve the importance of the resonance structure.
Resonance structure G is acceptable. Structure H is not acceptable because the carbon has ten valance electrons.
Structures I and J are both acceptable resonance
for bisulfate ion, the conjugate base of sulfuric acid. The
atom of structure J has 12 valence electrons, an expanded
Sulfur is a third row element, so an expanded octet is allowed.
important resonance structure because it maximizes the number of
bonds and minimizes the number atoms with a nonzero formal charge.
Rule 3: Nuclei do not change positions in space between resonance structures. Resonance structures differ only in the arrangement of valence electrons.
Structures K and L are both acceptable Lewis
but they are not related by resonance because the circled hydrogen atom
has changed position in space.
Exercises: Draw all reasonable resonance forms for the structures shown below. Use curved arrows to indicate electron pair changes.