Aromaticity (chemistry)

From Wikipedia, the free encyclopedia

Two different resonance forms of benzene (top) combine to produce an average structure (bottom)

In organic chemistry, the term aromaticity is used to describe a cyclic (ring-shaped), planar (flat) molecule with a ring of resonance bonds that exhibits more stability than other geometric or connective arrangements with the same set of atoms. Aromatic molecules are very stable, and do not break apart easily to react with other substances. Organic compounds that are not aromatic are classified as aliphatic compounds—they might be cyclic, but only aromatic rings have special stability (low reactivity).

Since the most common aromatic compounds are derivatives of benzene (an aromatic hydrocarbon common in petroleum and its distillates), the word “aromatic” occasionally refers informally to benzene derivatives, and so it was first defined. Nevertheless, many non-benzene aromatic compounds exist. In living organisms, for example, the most common aromatic rings are the double-ringed bases in RNA and DNA. An aromatic functional group or other substituent is called an aryl group.

The earliest use of the term "aromatic" was in an article by August Wilhelm Hofmann in 1855.^[1] Hofmann used the term for a class of benzene compounds, many of which have odors (aromas), unlike pure saturated hydrocarbons. Aromaticity as a chemical property bears no general relationship with the olfactory properties of such compounds (how they smell), although in 1855, before the structure of benzene or organic compounds was understood, chemists like Hofmann were beginning to understand that odiferous molecules from plants, such as terpenes, had chemical properties that we recognize today are similar to unsaturated petroleum hydrocarbons like benzene.

In terms of the electronic nature of the molecule, aromaticity describes a conjugated system often made of alternating single and double bonds in a ring. This configuration allows for the electrons in the molecule's pi system to be delocalized around the ring, increasing the molecule's stability. The molecule cannot be represented by one structure, but rather a resonance hybrid of different structures, such as with the two resonance structures of benzene. These molecules cannot be found in either one of these representations, with the longer single bonds in one location and the shorter double bond in another (see Theory below). Rather, the molecule exhibits bond lengths in between those of single and double bonds. This commonly seen model of aromatic rings, namely the idea that benzene was formed from a six-membered carbon ring with alternating single and double bonds (cyclohexatriene), was developed by August Kekulé (see History below). The model for benzene consists of two resonance forms, which corresponds to the double and single bonds superimposing to produce six one-and-a-half bonds. Benzene is a more stable molecule than would be expected without accounting for charge delocalization.

Theory

As is standard for resonance diagrams, the use of a double-headed arrow indicates that two structures are not distinct entities but merely hypothetical possibilities. Neither is an accurate representation of the actual compound, which is best represented by a hybrid (average) of these structures. A C=C bond is shorter than a C−C bond. Benzene is a regular hexagon—it is planar and all six carbon–carbon bonds have the same length, which is intermediate between that of a single and that of a double bond.

In a cyclic molecule with three alternating double bonds, cyclohexatriene, the bond length of the single bond would be 1.54 Å and that of the double bond would be 1.34 Å. However, in a molecule of benzene, the length of each of the bonds is 1.40 Å, indicating it to be the average of single and double bond.^[2][3]

A better representation is that of the circular π-bond (Armstrong's inner cycle), in which the electron density is evenly distributed through a π-bond above and below the ring. This model more correctly represents the location of electron density within the aromatic ring.

The single bonds are formed from overlap of hybridized atomic sp²-orbitals in line between the carbon nuclei—these are called σ-bonds. Double bonds consist of a σ-bond and a π-bond. The π-bonds are formed from overlap of atomic p-orbitals above and below the plane of the ring. The following diagram shows the positions of these p-orbitals:

Since they are out of the plane of the atoms, these orbitals can interact with each other freely, and become delocalized. This means that, instead of being tied to one atom of carbon, each electron is shared by all six in the ring. Thus, there are not enough electrons to form double bonds on all the carbon atoms, but the "extra" electrons strengthen all of the bonds on the ring equally. The resulting molecular orbital is considered to have π symmetry.

History

The term "aromatic"

The first known use of the word "aromatic" as a chemical term—namely, to apply to compounds that contain the phenyl group—occurs in an article by August Wilhelm Hofmann in 1855.^[1][4] If this is indeed the earliest introduction of the term, it is curious that Hofmann says nothing about why he introduced an adjective indicating olfactory character to apply to a group of chemical substances only some of which have notable aromas. Also, many of the most odoriferous organic substances known are terpenes, which are not aromatic in the chemical sense. But terpenes and benzenoid substances do have a chemical characteristic in common, namely higher unsaturation than many aliphatic compounds, and Hofmann may not have been making a distinction between the two categories. Many of the earliest-known examples of aromatic compounds, such as benzene and toluene, have distinctive pleasant smells. This property led to the term "aromatic" for this class of compounds, and hence the term "aromaticity" for the eventually discovered electronic property.^[5]

The structure of the benzene ring

Historic benzene formulae as proposed by August Kekulé in 1865.^[6]

In the 19th century chemists found it puzzling that benzene could be so unreactive toward addition reactions, given its presumed high degree of unsaturation. The cyclohexatriene structure for benzene was first proposed by August Kekulé in 1865.^[7][8] Most chemists were quick to accept this structure, since it accounted for most of the known isomeric relationships of aromatic chemistry. The hexagonal structure explains why only one isomer of benzene exists and why disubstituted compounds have three isomers.^[4]

Between 1897 and 1906, J. J. Thomson, the discoverer of the electron, proposed three equivalent electrons between each pair of carbon atoms in benzene. An explanation for the exceptional stability of benzene is conventionally attributed to Sir Robert Robinson, who was apparently the first (in 1925)^[9] to coin the term aromatic sextet as a group of six electrons that resists disruption.

In fact, this concept can be traced further back, via Ernest Crocker in 1922,^[10] to Henry Edward Armstrong, who in 1890 wrote "the (six) centric affinities act within a cycle...benzene may be represented by a double ring (sic) ... and when an additive compound is formed, the inner cycle of affinity suffers disruption, the contiguous carbon-atoms to which nothing has been attached of necessity acquire the ethylenic condition".^{[11][verification needed]}

Here, Armstrong is describing at least four modern concepts.^{[verification needed]} First, his "affinity" is better known nowadays as the electron, which was to be discovered only seven years later by J. J. Thomson. Second, he is describing electrophilic aromatic substitution, proceeding (third) through a Wheland intermediate, in which (fourth) the conjugation of the ring is broken. He introduced the symbol C centered on the ring as a shorthand for the inner cycle, thus anticipating Erich Clar's notation. It is argued that he also anticipated the nature of wave mechanics, since he recognized that his affinities had direction, not merely being point particles, and collectively having a distribution that could be altered by introducing substituents onto the benzene ring (much as the distribution of the electric charge in a body is altered by bringing it near to another body).

The quantum mechanical origins of this stability, or aromaticity, were first modelled by Hückel in 1931. He was the first to separate the bonding electrons into sigma and pi electrons.

Aromaticity of an arbitrary aromatic compound can be measured quantitatively by the nucleus-independent chemical shift (NICS) computational method^[12] and aromaticity percentage^[13] methods.

Characteristics of aromatic (aryl) compounds

An aromatic (or aryl) compound contains a set of covalently bound atoms with specific characteristics:

1. A delocalized conjugated π system, most commonly an arrangement of alternating single and double bonds
2. Coplanar structure, with all the contributing atoms in the same plane
3. Contributing atoms arranged in one or more rings
4. A number of π delocalized electrons that is even, but not a multiple of 4. That is, 4n + 2 π-electrons, where n = 0, 1, 2, 3, and so on. This is known as Hückel's rule.

According to Hückel's rule, if a molecule has 4n + 2 π-electrons, it is aromatic, but if it has 4n π-electrons and has characteristics 1–3 above, the molecule is said to be antiaromatic. Whereas benzene is aromatic (6 electrons, from 3 double bonds), cyclobutadiene is antiaromatic, since the number of π delocalized electrons is 4, which of course is a multiple of 4. The cyclobutadienide(2−) ion, however, is aromatic (6 electrons). An atom in an aromatic system can have other electrons that are not part of the system, and are therefore ignored for the 4n + 2 rule. In furan, the oxygen atom is sp² hybridized. One lone pair is in the π system and the other in the plane of the ring (analogous to the C–H bond in the other positions). There are 6 π-electrons, so furan is aromatic.

Aromatic molecules typically display enhanced chemical stability, compared to similar non-aromatic molecules. A molecule that can be aromatic will tend to change toward aromaticity, and the added stability changes the chemistry of the molecule. Aromatic compounds undergo electrophilic aromatic substitution and nucleophilic aromatic substitution reactions, but not electrophilic addition reactions as happens with carbon–carbon double bonds.

In the presence of a magnetic field, the circulating π-electrons in an aromatic molecule produce an aromatic ring current that induces an additional magnetic field, an important effect in nuclear magnetic resonance.^[14] The NMR signal of protons in the plane of an aromatic ring are shifted substantially further down-field than those on non-aromatic sp² carbons. This is an important way of detecting aromaticity. By the same mechanism, the signals of protons located near the ring axis are shifted upfield.

Aromatic molecules are able to interact with each other in so-called π–π stacking: The π systems form two parallel rings overlap in a "face-to-face" orientation. Aromatic molecules are also able to interact with each other in an "edge-to-face" orientation: The slight positive charge of the substituents on the ring atoms of one molecule are attracted to the slight negative charge of the aromatic system on another molecule.

Planar monocyclic molecules containing 4n π-electrons are called antiaromatic and are, in general, unstable. Molecules that could be antiaromatic will tend to change from this electronic or conformation, thereby becoming non-aromatic. For example, cyclooctatetraene (COT) distorts out of planarity, breaking π overlap between adjacent double bonds. Recent studies have determined that cyclobutadiene adopts an asymmetric, rectangular configuration in which single and double bonds indeed alternate, with no resonance; the single bonds are markedly longer than the double bonds, reducing unfavorable p-orbital overlap. This reduction of symmetry lifts the degeneracy of the two formerly non-bonding molecular orbitals, which by Hund's rule forces the two unpaired electrons into a new, weakly bonding orbital (and also creates a weakly antibonding orbital). Hence, cyclobutadiene is non-aromatic; the strain of the asymmetric configuration outweighs the anti-aromatic destabilization that would afflict the symmetric, square configuration.

Importance of aromatic compounds

Aromatic compounds play key roles in the biochemistry of all living things. The four aromatic amino acids histidine, phenylalanine, tryptophan, and tyrosine each serve as one of the 20 basic building-blocks of proteins. Further, all 5 nucleotides (adenine, thymine, cytosine, guanine, and uracil) that make up the sequence of the genetic code in DNA and RNA are aromatic purines or pyrimidines. The molecule heme contains an aromatic system with 22 π-electrons. Chlorophyll also has a similar aromatic system.

Aromatic compounds are important in industry. Key aromatic hydrocarbons of commercial interest are benzene, toluene, ortho-xylene and para-xylene. About 35 million tonnes are produced worldwide every year. They are extracted from complex mixtures obtained by the refining of oil or by distillation of coal tar, and are used to produce a range of important chemicals and polymers, including styrene, phenol, aniline, polyester and nylon.

Types of aromatic compounds

The overwhelming majority of aromatic compounds are compounds of carbon, but they need not be hydrocarbons.

Neutral homocyclics

Benzene, as well as most other annulenes (cyclodecapentaene excepted) with the formula C_nH_n where n is an even number, such as cyclotetradecaheptaene.

Heterocyclics

In heterocyclic aromatics (heteroaromatics), one or more of the atoms in the aromatic ring is of an element other than carbon. This can lessen the ring's aromaticity, and thus (as in the case of furan) increase its reactivity. Other examples include pyridine, pyrazine, pyrrole, imidazole, pyrazole, oxazole, thiophene, and their benzannulated analogs (benzimidazole, for example). In all these examples, the number of π-electrons is 6, due to the π-electrons from the double bonds as well as the two electrons from any lone pair that is in the p-orbital that is in the plane of the aromatic π system. For example, in pyridine, the five sp²-hybridized carbons each have a p-orbital that is perpendicular to the plane of the ring, and each of these p-orbitals contains one π-electron. Additionally, the nitrogen atom is also sp²-hybridized and has one electron in a p-orbital, which adds up to 6 p-electrons, thus making pyridine aromatic. The lone pair on the nitrogen is not part of the aromatic π system. Pyrrole and imidazole are both five membered aromatic rings that contain heteroatoms. In pyrrole, each of the four sp²-hybridized carbons contributes one π-electron, and the nitrogen atom is also sp²-hybridized and contributes two π-electrons from its lone pair, which occupies a p-orbital. In imidazole, both nitrogens are sp²-hybridized; the one in the double bond contributes one electron and the one which is not in the double bond and is in a lone pair contributes two electrons to the π system.^[15]

Fused aromatics and polycyclics

Polycyclic aromatic hydrocarbons are molecules containing two or more simple aromatic rings fused together by sharing two neighboring carbon atoms (see also simple aromatic rings). Examples are naphthalene, anthracene, and phenanthrene. In fused aromatics, not all carbon–carbon bonds are necessarily equivalent, as the electrons are not delocalized over the entire molecule. The aromaticity of these molecules can be explained using their orbital picture. Like benzene and other monocyclic aromatic molecules, polycyclics have a cyclic conjugated pi system with p-orbital overlap above and below the plane of the ring.^[15]

Substituted aromatics

Many chemical compounds are aromatic rings with other functional groups attached. Examples include trinitrotoluene (TNT), acetylsalicylic acid (aspirin), paracetamol, and the nucleotides of DNA.

Aromatic ions

Aromatic molecules need not be neutral molecules. Ions that satisfy Huckel’s rule of 4n + 2 π-electrons in a planar, cyclic, conjugated molecule are considered to be aromatic ions. For example, the cyclopentadienyl anion and the cycloheptatrienylium cation are both considered to be aromatic ions, and the azulene molecule can be approximated as a combination of both.

In order to convert the atom from sp³ to sp², a carbocation, carbanion, or carbon radical must be formed. These leave sp²-hybridized carbons that can partake in the π system of an aromatic molecule. Like neutral aromatic compounds, these compounds are stable and form easily. The cyclopentadienyl anion is formed very easily and thus 1,3-cyclopentadiene is a very acidic hydrocarbon with a pK_a of 16.^[15] Other examples of aromatic ions include the cyclopropenium cation (2 π-electrons) and cyclooctatetraenyl dianion (10 pi electrons).

Atypical aromatic compounds

Aromaticity also occurs in compounds that are not carbocyclic or heterocyclic; inorganic six-membered-ring compounds analogous to benzene have been synthesized. For example, borazine is a six-membered ring composed of alternating boron and nitrogen atoms, each with one hydrogen attached. It has a delocalized π system and undergoes electrophilic substitution reactions appropriate to aromatic rings rather than reactions expected of non-aromatic molecules.^[16]

Quite recently, the aromaticity of planar Si⁶⁻
₅ rings occurring in the Zintl phase Li₁₂Si₇ was experimentally evinced by Li solid-state NMR.^{[17][non-primary source needed]} Metal aromaticity is believed to exist in certain clusters of aluminium, for example.^{[citation needed]}

Homoaromaticity is the term used to describe systems where conjugation is interrupted by a single sp³ hybridized carbon atom.^[18]

Möbius aromaticity occurs when a cyclic system of molecular orbitals, formed from p_π atomic orbitals and populated in a closed shell by 4n (n is an integer) electrons, is given a single half-twist to form a Möbius strip. A π system with 4n electrons in a flat (non-twisted) ring would be antiaromatic, and therefore highly unstable, due to the symmetry of the combinations of p atomic orbitals. By twisting the ring, the symmetry of the system changes and becomes allowed (see also Möbius–Hückel concept for details). Because the twist can be left-handed or right-handed, the resulting Möbius aromatics are dissymmetric or chiral. But as of 2012, no Möbius aromatic molecules had been synthesized.^[19][20] Aromatics with two half-twists corresponding to the paradromic topologies were first suggested by Johann Listing.^[21] In carbo-benzene the ring bonds are extended with alkyne and allene groups.

Y-aromaticity is a concept developed to explain the extraordinary stability and high basicity of the guanidinium cation. Guanidinium is not a ring molecule, and is cross-conjugated rather than a linear π system, but is reported to have its six π-electrons delocalized over the whole molecule. The concept is controversial and some authors emphasize different effects.^[22][23][24]

σ-aromaticity refers to stabilization arising from the delocalization of sigma bonds. It is often invoked in cluster chemistry and is closely related to Wade's Rule.

Resonance (chemistry)

From Wikipedia, the free encyclopedia

The experimental geometry of nitrite anion, NO₂^–, shown on the right, is best rationalized by describing its structure as a resonance hybrid consisting of two major and equally important contributing forms.

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 (or forms)^[2] (also variously known as resonance structures (or forms), canonical structures, or, in older works or translations, mesomers), which collectively constitute a resonance hybrid. (The concept of a resonance hybrid is unrelated to orbital hybridization.)

Under the framework of valence bond theory, resonance is an extension of the idea that the bonding in a chemical species can be described by a Lewis structure. For many chemical species, a single Lewis structure, consisting of atoms obeying the octet rule, possibly bearing formal charges, and connected by bonds of positive integer order, is sufficient for describing the chemical bonding and rationalizing experimentally determined molecular properties like bond lengths, angles, and dipole moment.^[3] However, in some cases, more than one Lewis structure could be drawn, and experimental properties are inconsistent with any one structure. In order to address this type of situation, several resonance structures are considered together as an average, and the molecule is said to be represented by a resonance hybrid in which several Lewis structures are used collectively to describe its true structure. For instance, in NO₂^–, nitrite anion, the two N–O bond lengths are equal, even though no single Lewis structure has two N–O bonds with the same formal bond order. However, its measured structure is consistent with a description as a resonance hybrid of the two major contributing structures shown above: it has two equal N–O bonds of 125 pm, intermediate in length between a typical N–O single bond (145 pm in hydroxylamine, H₂N–OH) and N–O double bond (115 pm in nitronium ion, [O=N=O]⁺). According to the contributing structures, each N–O bond is an average of a formal single and formal double bond, leading to a true bond order of 1.5. By virtue of this averaging, the Lewis description of the bonding in NO₂^– is reconciled with the experimental fact that the anion has equivalent N–O bonds.

As suggested by this example, individual Lewis structures of a resonance hybrid are hypothetical mental constructs, and the molecular geometry and properties implied by each structure standing alone do not correspond to any real chemical species. To correctly interpret the molecular structure described by a resonance hybrid, all significant contributors of the resonance hybrid must be considered together, since the hybrid represents the actual molecule as their "average," with bond lengths and other structural parameters taking on intermediate values compared to those expected for the individual Lewis structures of the contributors, were they to exist as "real" chemical entities.^[4] The contributing structures differ only in the formal apportionment of electrons to the atoms, and not in the actual physically and chemically significant electron or spin density. While contributing structures will differ in formal bond orders and, possibly, in formal charge assignments, the contributing structures of a resonance hybrid all represent a single chemical species. In particular, each contributing structure represents the same molecular geometry (position of atomic nuclei) and electron/spin distribution.^[5]

Because electron delocalization lowers the potential energy of a system, any species represented by a resonance hybrid is more stable than any of the (hypothetical) contributing structures. The difference in potential energy between the actual species and the (computed) energy of the contributing structure with the lowest potential energy is called the resonance energy^[6] or delocalization energy. The magnitude of the resonance energy depends on assumptions made about the hypothetical "non-stabilized" species and the computational methods used and does not represent a measurable physical quantity, although comparisons of resonance energies computed under similar assumptions and conditions may be chemically meaningful.

Resonance is to be distinguished from isomerism. Isomers are molecules with the same chemical formula but are distinct chemical species with different arrangements of atomic nuclei in space, or even different distributions of electrons or different spin states (see electromerism and spin isomerism). Resonance contributors of a molecule, on the other hand, can only differ in the way electrons are formally assigned to atoms in the Lewis structure depictions of the molecule. Specifically, when a molecular structure is said to be represented by a resonance hybrid, it does not mean that electrons or atomic nuclei of the molecule are "resonating" or shifting back and forth between several sets of positions, each one represented by a Lewis structure. Rather, it means that the set of contributing structures represents an intermediate structure (a weighed average of the contributors), with a single, well-defined geometry and distribution of electrons. It is incorrect to regard resonance hybrids as rapidly interconverting isomers, even though the term "resonance" might evoke such an image.^[7] As described below, the term "resonance" originated as a classical physics analogy for a quantum mechanical phenomenon, so it should not be construed too literally. Symbolically, the double headed arrow ${\displaystyle {\ce {A<->B}}}$ is used to indicate that A and B are resonance forms of a single chemical species (as opposed to an equilibrium arrow, e.g., ${\displaystyle {\ce {A <=> B}}}$.

In the language of molecular orbital theory, molecules described by resonance hybrids often correspond to ones containing several p orbitals spread across adjacent atoms that can participate in bonding through an extended π system.

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.^[8] 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 resonance proposal also helped explain the number of isomers of benzene derivatives. For example, Kekulé's structure would predict four dibromobenzene isomers, including two ortho isomers with the brominated carbons joined by either a single or a double bond. In reality there are only three dibromobenzene isomers and only one is ortho, in agreement with the idea that there is only type of carbon-carbon bond, intermediate between a single and a double bond.^[9]

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][10] 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.^[11][12] The alternative term mesomerism popular in German and French publications with the same meaning was introduced by C. K. 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".^[13]

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. Bonds that have different formal bond orders in different contributing structures have intermediate bond lengths, in between the bond lengths expected for the higher bond order and lower bond order. Less frequently, the same is true for bond angles. (See the structure of [OCNCO]⁺, for instance.^[14])
• The contributing structures are not isomers. They differ only in the bookkeeping of electrons, not where the nuclei or electrons actually reside.
• Each Lewis formula must have the same number of valence electrons, and thus the same total charge.
• Each Lewis formula must have the same number of unpaired electrons (if any), and thus the same total spin and spin multiplicity.
• The real structure has a lower total potential energy and is therefore more stable than each of the contributing structures (interpreted to represent hypothetical species with localized electron pairs).

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. A non-chemical analogy is illustrative: one can describe the characteristics of a real animal, the narwhal, in terms of the characteristics of two mythical creatures: the unicorn, a creature with a single horn on its head, and the leviathan, a large, whale-like creature. The narwhal is not a creature that goes back and forth between being a unicorn and being a leviathan, nor do the unicorn and leviathan have any physical existence outside the collective human imagination. Nevertheless, describing the narwhal in terms of these imaginary creatures provides a reasonably good description of its physical characteristics.

Due to confusion with the physical meaning of the word resonance, as no entities actually physically "resonate," it has been suggested that the term resonance be abandoned in favor of delocalization.^[15] 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, as arrows of any type may suggest to beginning students that a chemical change is taking place.

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 or ion 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 favorable 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 formally charged atoms, with the separation for unlike and like charges minimized and maximized, respectively;
• that place negative charge, if any, on the most electronegative atoms and positive charge, if any, on the most electropositive;
• that are not forced to deviate substantially from idealized bond lengths and angles (e.g., the relative unimportance of Dewar-type resonance contributors for benzene);
• that maintain aromatic substructures locally (while avoiding anti-aromatic ones; see Clar sextet and biphenylene).

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

${\displaystyle {\ce {[S=C=N^{\ominus }<->\ ^{\ominus }\!S-C{\equiv }N]}}}$

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 ^[16] 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.^[17]

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 −​¹⁄₂ on both oxygen atoms and partial double bonds with a full and dashed line and bond order ​1 ¹⁄₂.^[18][19]

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.^{[citation needed]}

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

Aromatic molecules

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

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 +​¹⁄₂. The full positive charge can also be depicted as delocalized among three carbon atoms.

Reactive intermediates

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 ¹⁄₂. 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.^[20] 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".^[6]

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. ^[21]

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).^[22]

Resonance in valence bond (VB) theory

VB mixing diagram of benzene.^[23] The A_1g and
B_2u labels define the symmetries of the two states,
as defined by the character table for the D_6h
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 AO_i has a weighting coefficient c_i 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 2p_z 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 c_i on atom C_i. Charge q_i ≈ 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 q_i ≈ (c_iAO_i)² ≈ c²
_i. In benzene, q_i = 1 on each C atom. With an electron-withdrawing group q_i < 1 on the ortho and para C atoms and q_i > 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 pK_a value of −10.^[24] The extent of charge delocalization in an anion can be quantitatively expressed via the WAPS (weighted average positive sigma) parameter^[25] parameter and an analogous WANS (weighted average negative sigma)^[26][27] 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[28]	Triphenylphosphine	2.1[26]
(CF3)3COH	3.6[28]	Phenyl tetramethylguanidine	2.5[26]
Picric acid	4.3[25]	Tripropylamine	2.6[26]
2,4-Dinitrophenol	4.9[25]	MTBD (7-Methyl-triazabicyclodecene)	2.9[27]
Benzoic acid	7.1[25]	DBU (1,8-Diazabicycloundec-7-ene)	3.0[27]
Phenol	8.8[28]	TBD (Triazabicyclodecene)	3.5[27]
Acetic acid	16.1[25]	N,N-Dimethylaniline	4.7[26]
HI	21.9[28]	Pyridine	7.2[26]
HBr	29.1[28]	Aniline	8.2[26]
HCl	35.9[25]	Propylamine	8.9[26]

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

Lewis acids and bases

From Wikipedia, the free encyclopedia

Diagram of some Lewis bases and acids

A Lewis acid is a chemical species that contains an empty orbital which is capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that has a filled orbital containing an electron pair which is not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH₃ is a Lewis base, because it can donate its lone pair of electrons. Trimethylborane (Me₃B) is a Lewis acid as it is capable of accepting a lone pair. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond.^[1] In the context of a specific chemical reaction between NH₃ and Me₃B, the lone pair from NH₃ will form a dative bond with the empty orbital of Me₃B to form an adduct NH₃•BMe₃. The terminology refers to the contributions of Gilbert N. Lewis.^[2]

Depicting adducts

In many cases, the interaction between the Lewis base and Lewis acid in a complex is indicated by an arrow indicating the Lewis base donating electrons toward the Lewis acid using the notation of a dative bond—for example, Me₃B←NH₃. Some sources indicate the Lewis base with a pair of dots (the explicit electrons being donated), which allows consistent representation of the transition from the base itself to the complex with the acid:

Me₃B +  :NH₃ → Me₃B:NH₃

A center dot may also be used to represent a Lewis adduct, such as Me₃B•NH₃. Another example is boron trifluoride diethyl etherate, BF₃•Et₂O. The center dot is also used to represent hydrate coordination in various crystals, as in MgSO₄·7H₂O for hydrated magnesium sulfate. In general, however, the donor–acceptor bond is viewed as simply somewhere along a continuum between idealized covalent bonding and ionic bonding.^[3]

Examples

Major structural changes accompany binding of the Lewis base to the coordinatively unsaturated, planar Lewis acid BF₃.

Classically, the term "Lewis acid" is restricted to trigonal planar species with an empty p orbital, such as BR₃ where R can be an organic substituent or a halide.^{[citation needed]} For the purposes of discussion, even complex compounds such as Et₃Al₂Cl₃ and AlCl₃ are treated as trigonal planar Lewis acids. Metal ions such as Na⁺, Mg²⁺, and Ce³⁺, which are invariably complexed with additional ligands, are often sources of coordinatively unsaturated derivatives that form Lewis adducts upon reaction with a Lewis base. Other reactions might simply be referred to as "acid-catalyzed" reactions. Some compounds, such as H₂O, are both Lewis acids and Lewis bases, because they can either accept a pair of electrons or donate a pair of electrons, depending upon the reaction.

Lewis acids are diverse. Simplest are those that react directly with the Lewis base. But more common are those that undergo a reaction prior to forming the adduct.

• Examples of Lewis acids based on the general definition of electron pair acceptor include:
  • the proton (H⁺) and acidic compounds onium ions, such as NH₄⁺ and H₃O⁺
  • high oxidation state transition metal cations, e.g., Fe³⁺;
  • other metal cations, such as Li⁺ and Mg²⁺, often as their aquo or ether complexes,
  • trigonal planar species, such as BF₃ and carbocations H₃C⁺
  • pentahalides of phosphorus, arsenic, and antimony
  • electron poor π-systems, such as enones and tetracyanoethylenes.

Again, the description of a Lewis acid is often used loosely. For example, in solution, bare protons do not exist.

Simple Lewis acids

Some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior:

BF₃ + F⁻ → BF₄⁻

In this adduct, all four fluoride centres (or more accurately, ligands) are equivalent.

BF₃ + OMe₂ → BF₃OMe₂

Both BF₄⁻ and BF₃OMe₂ are Lewis base adducts of boron trifluoride.

In many cases, the adducts violate the octet rule, such as the triiodide anion:

I₂ + I⁻ → I₃⁻

The variability of the colors of iodine solutions reflects the variable abilities of the solvent to form adducts with the Lewis acid I₂.

In some cases, the Lewis acids are capable of binding two Lewis bases, a famous example being the formation of hexafluorosilicate:

SiF₄ + 2 F⁻ → SiF₆²⁻

Complex Lewis acids

Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Well known cases are the aluminium trihalides, which are widely viewed as Lewis acids. Aluminium trihalides, unlike the boron trihalides, do not exist in the form AlX₃, but as aggregates and polymers that must be degraded by the Lewis base.^[4] A simpler case is the formation of adducts of borane. Monomeric BH₃ does not exist appreciably, so the adducts of borane are generated by degradation of diborane:

B₂H₆ + 2 H⁻ → 2 BH₄⁻

In this case, an intermediate B₂H₇⁻ can be isolated.

Many metal complexes serve as Lewis acids, but usually only after dissociating a more weakly bound Lewis base, often water.

[Mg(H₂O)₆]²⁺ + 6 NH₃ → [Mg(NH₃)₆]²⁺ + 6 H₂O

H⁺ as Lewis acid

The proton (H⁺) ^[5] is one of the strongest but is also one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is heavily solvated (bound to solvent). With this simplification in mind, acid-base reactions can be viewed as the formation of adducts:

• H⁺ + NH₃ → NH₄⁺
• H⁺ + OH⁻ → H₂O

Applications of Lewis acids

A typical example of a Lewis acid in action is in the Friedel–Crafts alkylation reaction.^[3] The key step is the acceptance by AlCl₃ of a chloride ion lone-pair, forming AlCl₄⁻ and creating the strongly acidic, that is, electrophilic, carbonium ion.

RCl +AlCl₃ → R⁺ + AlCl₄⁻

Lewis bases

A Lewis base is an atomic or molecular species where the highest occupied molecular orbital (HOMO) is highly localized. Typical Lewis bases are conventional amines such as ammonia and alkyl amines. Other common Lewis bases include pyridine and its derivatives. Some of the main classes of Lewis bases are

• amines of the formula NH_3−xR_x where R = alkyl or aryl. Related to these are pyridine and its derivatives.
• phosphines of the formula PR_3−xA_x, where R = alkyl, A = aryl.
• compounds of O, S, Se and Te in oxidation state 2, including water, ethers, ketones

The most common Lewis bases are anions. The strength of Lewis basicity correlates with the pK_a of the parent acid: acids with high pK_a's give good Lewis bases. As usual, a weaker acid has a stronger conjugate base.

• Examples of Lewis bases based on the general definition of electron pair donor include:
  • simple anions, such as H⁻ and F⁻.
  • other lone-pair-containing species, such as H₂O, NH₃, HO⁻, and CH₃⁻
  • complex anions, such as sulfate
  • electron rich π-system Lewis bases, such as ethyne, ethene, and benzene

The strength of Lewis bases have been evaluated for various Lewis acids, such as I₂, SbCl₅, and BF₃.<sup>[6]

Heats of binding of various bases to BF₃</sup>

Lewis base	Donor atom	Enthalpy of complexation (kJ/mol)
Et3N	N	135
quinuclidine	N	150
pyridine	N	128
Acetonitrile	N	60
Et2O	O	78.8
THF	O	90.4
acetone	O	76.0
EtOAc	O	75.5
DMA	O	112
DMSO	O	105
Tetrahydrothiophene	S	51.6
Trimethylphosphine	P	97.3

Applications of Lewis bases

Nearly all electron pair donors that form compounds by binding transition elements can be viewed as a collections of the Lewis bases – or ligands. Thus a large application of Lewis bases is to modify the activity and selectivity of metal catalysts. Chiral Lewis bases thus confer chirality on a catalyst, enabling asymmetric catalysis, which is useful for the production of pharmaceuticals.

Many Lewis bases are "multidentate," that is they can form several bonds to the Lewis acid. These multidentate Lewis bases are called chelating agents.

Hard and soft classification

Lewis acids and bases are commonly classified according to their hardness or softness. In this context hard implies small and nonpolarizable and soft indicates larger atoms that are more polarizable.

• typical hard acids: H⁺, alkali/alkaline earth metal cations, boranes, Zn²⁺
• typical soft acids: Ag⁺, Mo(0), Ni(0), Pt²⁺
• typical hard bases: ammonia and amines, water, carboxylates, fluoride and chloride
• typical soft bases: organophosphines, thioethers, carbon monoxide, iodide

For example, an amine will displace phosphine from the adduct with the acid BF₃. In the same way, bases could be classified. For example, bases donating a lone pair from an oxygen atom are harder than bases donating through a nitrogen atom. Although the classification was never quantified it proved to be very useful in predicting the strength of adduct formation, using the key concepts that hard acid — hard base and soft acid — soft base interactions are stronger than hard acid — soft base or soft acid — hard base interactions. Later investigation of the thermodynamics of the interaction suggested that hard—hard interactions are enthalpy favored, whereas soft—soft are entropy favored.

ECW model

The ECW model is quantitative model that describes and predicts the strength of Lewis acid base interactions, −ΔH . The model assigned E and C parameters to many Lewis acids and bases. Each acid is characterized by an E_A and a C_A. Each base is likewise characterized by its own E_B and C_B. The E and C parameters refer, respectively, to the electrostatic and covalent contributions to the strength of the bonds that the acid and base will form. The equation is

−ΔH = E_AE_B + C_AC_B + W

The W term represents a constant energy contribution for acid–base reaction such as the cleavage of a dimeric acid or base. The equation predicts reversal of acids and base strengths. The graphical presentations of the equation show that there is no single order of Lewis base strengths or Lewis acid strengths.^[7]

There is no single order of Lewis base strengths

Cramer–Bopp plots show graphically using the E and C parameters of the ECW model that there is no one single order of Lewis base strengths (or acid strengths).^[8] Single property or variable scales are limited to a small range of acids or bases.

History

MO diagram depicting the formation of a dative covalent bond between two atoms.

The concept originated with Gilbert N. Lewis who studied chemical bonding. In 1923, Lewis wrote An acid substance is one which can employ an electron lone pair from another molecule in completing the stable group of one of its own atoms.^[9][10] The Brønsted–Lowry acid–base theory was published in the same year. The two theories are distinct but complementary. A Lewis base is also a Brønsted–Lowry base, but a Lewis acid doesn't need to be a Brønsted–Lowry acid. The classification into hard and soft acids and bases (HSAB theory) followed in 1963. The strength of Lewis acid-base interactions, as measured by the standard enthalpy of formation of an adduct can be predicted by the Drago–Wayland two-parameter equation.

Reformulation of Lewis theory

Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons. When each atom contributed one electron to the bond it was called a covalent bond. When both electrons come from one of the atoms it was called a dative covalent bond or coordinate bond. The distinction is not very clear-cut. For example, in the formation of an ammonium ion from ammonia and hydrogen the ammonia molecule donates a pair of electrons to the proton;^[5] the identity of the electrons is lost in the ammonium ion that is formed. Nevertheless, Lewis suggested that an electron-pair donor be classified as a base and an electron-pair acceptor be classified as acid.

A more modern definition of a Lewis acid is an atomic or molecular species with a localized empty atomic or molecular orbital of low energy. This lowest energy molecular orbital (LUMO) can accommodate a pair of electrons.

Comparison with Brønsted–Lowry theory

A Lewis base is often a Brønsted–Lowry base as it can donate a pair of electrons to H⁺;^[5] the proton is a Lewis acid as it can accept a pair of electrons. The conjugate base of a Brønsted–Lowry acid is also a Lewis base as loss of H⁺ from the acid leaves those electrons which were used for the A—H bond as a lone pair on the conjugate base. However, a Lewis base can be very difficult to protonate, yet still react with a Lewis acid. For example, carbon monoxide is a very weak Brønsted–Lowry base but it forms a strong adduct with BF₃.

In another comparison of Lewis and Brønsted–Lowry acidity by Brown and Kanner,^[11] 2,6-di-t-butylpyridine reacts to form the hydrochloride salt with HCl but does not react with BF₃. This example demonstrates that steric factors, in addition to electron configuration factors, play a role in determining the strength of the interaction between the bulky di-t-butylpyridine and tiny proton.

A Brønsted–Lowry acid is a proton donor, not an electron-pair acceptor.