# Resonance (chemistry)

For other uses, see Resonance (disambiguation).
Two of the contributing structures of nitrogen dioxide

In chemistry, resonance or mesomerism[1] is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis structure. A molecule or ion with such delocalized electrons is represented by several contributing structures[2] (also called resonance structures or canonical structures).

Each contributing structure can be represented by a Lewis structure, with only an integer number of covalent bonds between each pair of atoms within the structure.[3] Several Lewis structures are used collectively to describe the actual molecular structure, which is an approximate intermediate between the canonical forms called a resonance hybrid.[4] Contributing structures differ only in the position of electrons, not in the position of nuclei.

Electron delocalization lowers the potential energy of the substance and thus makes it more stable than any of the contributing structures. The difference between the potential energy of the actual structure and that of the contributing structure with the lowest potential energy is called the resonance energy[5] or delocalization energy.

Resonance is distinguished from isomerism. An isomer is a molecule with the same chemical formula but with different arrangements of atoms in space. Resonance contributors of a molecule, on the contrary, can only differ by the arrangements of electrons. Therefore, the resonance hybrid cannot be represented by a combination of isomers.[6]

## History

The concept first appeared in 1899 in Johannes Thiele's "Partial Valence Hypothesis" to explain the unusual stability of benzene which would not be expected from August Kekulé's structure proposed in 1865 with alternating single and double bonds. Benzene undergoes substitution reactions, rather than addition reactions as typical for alkenes. He proposed that the carbon-carbon bond in benzene is intermediate of a single and double bond.

The mechanism of resonance was introduced into quantum mechanics by Werner Heisenberg in 1926 in a discussion of the quantum states of the helium atom. He compared the structure of the helium atom with the classical system of resonating coupled harmonic oscillators.[4][7] In the classical system, the coupling produces two modes, one of which is lower in frequency than either of the uncoupled vibrations; quantum mechanically, this lower frequency is interpreted as a lower energy. Linus Pauling used this mechanism to explain the partial valence of molecules in 1928, and developed it further in a series of papers in 1931-33.[8][9] The alternative term mesomerism popular in German and French publications with the same meaning was introduced by Christopher Ingold in 1938, but did not catch on in the English literature. The current concept of mesomeric effect has taken on a related but different meaning. The double headed arrow was introduced by the German chemist Fritz Arndt who preferred the German phrase zwischenstufe or intermediate stage.

In the Soviet Union, resonance theory – especially as developed by Pauling – was attacked in the early 1950s as being contrary to the Marxist principles of dialectical materialism, and in June 1951 the Soviet Academy of Sciences under the leadership of Alexander Nesmeyanov convened a conference on the chemical structure of organic compounds, attended by 400 physicists, chemists, and philosophers, where "the pseudo-scientific essence of the theory of resonance was exposed and unmasked".[10]

## General characteristics of resonance

Molecules and ions with resonance (also called mesomerism) have the following basic characteristics:

Contributing structures of the carbonate ion
• They can be represented by several correct Lewis formulas, called "contributing structures", "resonance structures" or "canonical forms". The real structure is an intermediate of these structures represented by a resonance hybrid.
• The contributing structures are not isomers. They differ only in the position of electrons, not in the position of nuclei.
• Each Lewis formula must have the same number of valence electrons (and thus the same total charge), and the same number of unpaired electrons, if any.[11]
• Bonds that have different bond orders in different contributing structures do not have typical bond lengths. Measurements reveal intermediate bond lengths.
• The real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be.

## Misconception

It is a common misconception that resonance structures are actual transient states of the molecule, with the molecule oscillating between them or existing as an equilibrium between them. However these individual contributors cannot be observed in the actual resonance-stabilized molecule. Any molecule or ion exists in only one form – the resonance hybrid. Due to confusion with the physical meaning of the word resonance, as no elements actually appear to be resonating, it has been suggested that the term resonance be abandoned in favor of delocalization.[12] Resonance energy would thus become delocalization energy and a resonance structure becomes a contributing structure. The double headed arrows would be replaced by commas to illustrate a set of structures rather than suggesting that there is a reaction that converts among them.

## Use of contributing structures

In Lewis formulas, electrons are paired between atoms to form covalent bonds. Each single bond is made by two valence electrons, localized between the two bonded atoms. Each double bond has two additional localized π electrons, while each triple bond has four additional π electrons (two pairs) between the bonded atoms.

In molecules or ions that have a combination of one or more single and multiple bonds, often the exact position of the respective bonds cannot be indicated by a single Lewis structure. The π electrons appear to be in an intermediate position. To solve this problem, the concept of resonance is used, and the molecule is represented by several contributing structures, each showing a possible distribution of single and multiple bonds. The actual structure has a lowered overall energy and an intermediate bond order.

### Resonance hybrids

The actual structure of a molecule in the normal quantum state has the lowest possible value of total energy. This structure is called the "resonance hybrid" of that molecule. The resonance hybrid is the approximate intermediate of the contributing structures, but the overall energy is lower than each of the contributors, due to the resonance energy.[4]

### Major and minor contributors

One contributing structure may resemble the actual molecule more than another (in the sense of energy and stability). Structures with a low value of potential energy are more stable than those with high values and resemble the actual structure more. The most stable contributing structures are called major contributors. Energetically unfavourable and therefore less probable structures are minor contributors. Major contributors are generally structures:

• that obey as much as possible the octet rule (8 valence electrons around each atom rather than having deficiencies or surplus)
• that have a maximum number of covalent bonds
• that carry a minimum of charged atoms. If unlike charges are present their separation must be least while for like charges the separation must be maximum.
• with negative charge, if any, on the most electronegative atoms and positive charge, if any, on the most electropositive.

The greater the number of contributing structures, the more stable the molecule. This is because the more states at lower energy are available to the electrons in a particular molecule, the more stable the electrons are. Also the more volume electrons can occupy at lower energy the more stable the molecule is. We can also understand this concept by borrowing a concept of physics. As we know that charge dispersed is directly proportional to stability. Here, electrons can be termed as charged bodies and the more volume they occupy, more the charge gets dispersed ultimately leading to stability.[citation needed]

Equivalent contributors contribute equally to the actual structure; those with low potential energy (the major contributors) contribute more to the resonance hybrid than the less stable minor contributors. Especially when there is more than one major contributor, the resonance stabilization is high. High values of resonance energy are found in aromatic molecules.

### Contributing structures in diagrams

 ${\displaystyle {\ce {[S=C=N^{\ominus }<->\ ^{\ominus }\!S-C{\equiv }N]}}}$ Contributing structures of the thiocyanate ion, enclosed in square brackets.
Hybrid of the nitrate ion
Hybrid of benzene.

In diagrams, contributing structures are typically separated by double-headed arrows (↔). The arrow should not be confused with the right and left pointing equilibrium arrow (⇌). All structures together may be enclosed in large square brackets, to indicate they picture one single molecule or ion, not different species in a chemical equilibrium.

Alternatively to the use of resonance structures in diagrams, a hybrid diagram can be used. In a hybrid diagram, pi bond that are involved in resonance are usually pictured as curves [13] or dashed lines, indicating that these are partial rather than normal complete pi bonds. In benzene and other aromatic rings, the delocalized pi-electrons are sometimes pictured as a solid circle.[14]

## Examples

### Ionic-covalent molecules

The ozone molecule is represented by two resonance structures. In reality the two terminal oxygen atoms are equivalent and the hybrid structure is drawn on the right with a charge of −12 on both oxygen atoms and partial double bonds with a full and dashed line and bond order 1 12.[15][16]

For hypervalent molecules such as xenon difluoride, the rationalization described above can be applied to generate resonance structures to explain the bonding in such molecules. This has been shown by quantum chemical calculations to be the correct description instead of the common expanded octet model.

${\displaystyle {\ce {[{\mathsf {F-XeF^{-}<->F^{-}Xe-F}}]}}}$

### Aromatic molecules

Main article: Aromaticity

In benzene the two cyclohexatriene Kekulé structures first proposed by Kekulé are taken together as contributing structures to represent the total structure. In the hybrid structure on the right the dashed hexagon replaces three double bonds, and represents six electrons in a set of three molecular orbitals of π symmetry, with a nodal plane in the plane of the molecule.

In furan a lone pair of the oxygen atom interacts with the π orbitals of the carbon atoms. The curved arrows depict the permutation of delocalized π electrons, which results in different contributors.

### Electron-deficient molecules

Main article: Electron deficiency

The diborane molecule is described by resonance structures, each with electron-deficiency on different atoms. This reduces the electron-deficiency on each atom and stabilizes the molecule. Below are the resonance structures of an individual 3c-2e bond in diborane.

The allyl cation has two contributing structures with a positive charge on the terminal carbon atoms. In the hybrid structure their charge is +12. The full positive charge can also be depicted as delocalized among three carbon atoms.

## Reactive intermediates

Main article: Reactive intermediate

Often, reactive intermediates such as carbocations and free radicals have more delocalized structure than their parent reactants, giving rise to unexpected products. The classical example is allylic rearrangement. When 1 mole of HCl adds to 1 mole of 1,3-butadiene, in addition to the ordinarily expected product 3-chloro-1-butene, we also find 1-chloro-2-butene. Isotope labelling experiments have shown that what happens here is that the additional double bond shifts from 1,2 position to 2,3 position in some of the product. This and other evidence (such as NMR in superacid solutions) shows that the intermediate carbocation must have a highly delocalized structure, different from its mostly classical (delocalization exists but is small) parent molecule. This cation (an allylic cation) can be represented using resonance, as shown above.

This observation of greater delocalization in less stable molecules is quite general. The excited states of conjugated dienes are stabilised more by conjugation than their ground states, causing them to become organic dyes.

A well-studied example of delocalization that does not involve π electrons (hyperconjugation) can be observed in the non-classical 2-Norbornyl cation. Other examples are diborane (see above) and methanium (CH+
5
). These can be viewed as containing three-center two-electron bonds and are represented either by contributing structures involving rearrangement of σ electrons or by a special notation, a Y that has the three nuclei at its three points.

Delocalized electrons are important for several reasons; a major one is that an expected chemical reaction may not occur because the electrons delocalize to a more stable configuration, resulting in a reaction that happens at a different location. An example is the Friedel–Crafts alkylation of benzene with 1-chloro-2-methylpropane; the carbocation rearranges to a tert-butyl group stabilized by hyperconjugation, a particular form of delocalization. Delocalization leads to lengthening of wavelength of electron therefore decreases the energy.

## Bond lengths

Resonance structures of benzene

Comparing the two contributing structures of benzene, all single and double bonds are interchanged. Bond lengths can be measured, for example using X-ray diffraction. The average length of a C–C single bond is 154 pm; that of a C=C double bond is 133 pm. In localized cyclohexatriene, the carbon–carbon bonds should be alternating 154 and 133 pm. Instead, all carbon–carbon bonds in benzene are found to be about 139 pm, a bond length intermediate between single and double bond. This mixed single and double bond (or triple bond) character is typical for all molecules in which bonds have a different bond order in different contributing structures. Bond lengths can be compared using bond orders. For example, in cyclohexane the bond order is 1 while that in benzene is 1 + (3 ÷ 6) = 1 12. Consequently, benzene has more double bond character and hence has a shorter bond length than cyclohexane.

## Resonance energy

Every structure is associated with a certain quantity of energy, which determines the stability of the molecule or ion (the lower energy, the greater stability). A resonance hybrid has a structure that is intermediate between the contributing structures; the total quantity of potential energy, however, is lower than the intermediate and the molecule is said to be "stabilized by resonance" or "resonance-stabilized". Hybrids are therefore always more stable than any of the contributing structures would be.[17] The difference between the potential energy of the actual structure (the resonance hybrid) and that of the contributing structure with the lowest potential energy is called the "resonance energy".[5]

### Resonance energy of benzene

Resonance (or delocalization) energy is the amount of energy needed to convert the true delocalized structure into that of the most stable contributing structure. The empirical resonance energy can be estimated by comparing the enthalpy change of hydrogenation of the real substance with that estimated for the contributing structure.

The complete hydrogenation of benzene to cyclohexane via 1,3-cyclohexadiene and cyclohexene is exothermic; 1 mole of benzene delivers 208.4 kJ (49.8 kcal).

Hydrogenation of one mole of double bonds delivers 119.7 kJ (28.6 kcal), as can be deduced from the last step, the hydrogenation of cyclohexene. In benzene, however, 23.4 kJ (5.6 kcal) are needed to hydrogenate one mole of double bonds. The difference, being 143.1 kJ (34.2 kcal), is the empirical resonance energy of benzene. Because 1,3-cyclohexadiene also has a small delocalization energy (7.6 kJ or 1.8 kcal/mol) the net resonance energy, relative to the localized cyclohexatriene, is a bit higher: 151 kJ or 36 kcal/mol. [18]

This measured resonance energy is also the difference between the hydrogenation energy of three 'non-resonance' double bonds and the measured hydrogenation energy:

(3 × 119.7) − 208.4 = 150.7 kJ/mol (36 kcal).[19]

## Resonance in valence bond (VB) theory

VB mixing diagram of benzene.[20] The A1g and B2u labels define the symmetries of the two states, as defined by the character table for the D6h symmetry group.

Resonance has a deeper significance in the mathematical formalism of valence bond theory (VB). Quantum mechanics requires that the wavefunction of a molecule obeys its observed symmetry. If a single contributing structure does not achieve this, resonance is invoked.

For example, in benzene, valence bond theory begins with the two Kekulé structures and constructs the actual wave function of the molecule as a linear superposition of the wave functions representing the two structures. As both Kekulé structures have equal energy, they are equal contributors to the overall structure – the superposition is an equally weighted average, or a 1:1 linear combination of the two – but this need not be the case. The symmetric combination gives the ground state while the antisymmetric combination gives the first excited state as shown.

In general, the superposition is written with undetermined coefficients, which are then variationally optimized to find the lowest possible energy for the given set of basis wave functions. When more contributing structures are included, the molecular wave function becomes more accurate and more excited states can be derived from different combinations of the contributing structures.

### Comparison with molecular orbital (MO) theory

In molecular orbital theory, the main alternative to valence bond theory, the equivalent of the symmetry-adapted linear combination role of resonance is the linear combination of atomic orbitals. In MO theory, the molecular orbitals (MOs) are approximated as sums of all the atomic orbitals (AOs) on all the atoms; there are as many MOs as AOs. Each AOi has a weighting coefficient ci that indicates the AO's contribution to a particular MO. For example, in benzene, the MO model gives us 6 π MOs which are combinations of the 2pz AOs on each of the 6 C atoms. Thus, each π MO is delocalized over the whole benzene molecule and any electron occupying an MO will be delocalized over the whole molecule. This MO interpretation has inspired the picture of the benzene ring as a hexagon with a circle inside. When describing benzene, the VB concept of localized sigma 'bonds' and the MO concept of 'delocalized' π electrons are frequently combined in elementary chemistry courses.

The resonance structures in the VB model are particularly useful in predicting the effect of substituents on π systems such as benzene. They lead to the models of resonance structures for an electron-withdrawing group and electron-releasing group on benzene. The utility of MO theory is that a quantitative indication of the charge from the π system on an atom can be obtained from the squares of the weighting coefficient ci on atom Ci. Charge qi ≈ c2
i
. The reason for squaring the coefficient is that if an electron is described by an AO, then the square of the AO gives the electron density. The AOs are adjusted (normalized) so that AO2 = 1, and qi ≈ (ciAOi)2 ≈ c2
i
. In benzene, qi = 1 on each C atom. With an electron-withdrawing group qi < 1 on the ortho and para C atoms and qi > 1 for an electron-releasing group.

### Coefficients

Weighting of the resonance structures in terms of their contribution to the overall structure can be calculated in multiple ways, using "Ab initio" methods derived from Valence Bond theory, or else from the Natural Bond Orbitals (NBO) approaches of Weinhold NBO5, or finally from empirical calculations based on the Hückel method. A Hückel method-based software for teaching resonance is available on the HuLiS Web site.

## Charge delocalization

In the case of ions it is common to speak about delocalized charge (charge delocalization). An example of delocalized charge in ions can be found in the carboxylate group, wherein the negative charge is centered equally on the two oxygen atoms. Charge delocalization in anions is an important factor determining their reactivity (generally: the higher the extent of delocalization the lower the reactivity) and, specifically, the acid strength of their conjugate acids. As a general rule, the better delocalized is the charge in an anion the stronger is its conjugate acid. For example, the negative charge in perchlorate anion (ClO
4
) is evenly distributed among the symmetrically oriented oxygen atoms (and a part of it is also kept by the central chlorine atom). This excellent charge delocalization combined with the high number of oxygen atoms (four) and high electronegativity of the central chlorine atom leads to perchloric acid being one of the strongest known acids with a pKa value of −10.[21] The extent of charge delocalization in an anion can be quantitatively expressed via the WAPS (weighted average positive sigma) parameter[22] parameter and an analogous WANS (weighted average negative sigma)[23][24] parameter is used for cations.

WAPS values of anions of common acids and WANS values of cations of common bases
Compound WAPS × 105 Compound WANS × 105
(C2F5SO2)2NH 2.0[25] Triphenylphosphine 2.1[23]
(CF3)3COH 3.6[25] Phenyl tetramethylguanidine 2.5[23]
Picric acid 4.3[22] Tripropylamine 2.6[23]
2,4-Dinitrophenol 4.9[22] MTBD (7-Methyl-triazabicyclodecene) 2.9[24]
Benzoic acid 7.1[22] DBU (1,8-Diazabicycloundec-7-ene) 3.0[24]
Phenol 8.8[25] TBD (Triazabicyclodecene) 3.5[24]
Acetic acid 16.1[22] N,N-Dimethylaniline 4.7[23]
HI 21.9[25] Pyridine 7.2[23]
HBr 29.1[25] Aniline 8.2[23]
HCl 35.9[22] Propylamine 8.9[23]

WAPS and WANS values are given in e/Å4. Larger values indicate more localized charge in the corresponding ion.