Conjugated system

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Cinnamaldehyde is a naturally-occurring compound that has a conjugated system Cinnamaldehyde acsv.svg
Cinnamaldehyde is a naturally-occurring compound that has a conjugated system
penta-1,3-diene is a molecule with a conjugated system 1,3-pentadien.svg
penta-1,3-diene is a molecule with a conjugated system
Diazomethane conjugated pi-system Diazomethane-pi-system.png
Diazomethane conjugated pi-system

In theoretical chemistry, a conjugated system is a system of connected p-orbitals with delocalized electrons in a molecule, which in general lowers the overall energy of the molecule and increases stability. It is conventionally represented as having alternating single and multiple bonds. Lone pairs, radicals or carbenium ions may be part of the system, which may be cyclic, acyclic, linear or mixed. The term "conjugated" was coined in 1899 by the German chemist Johannes Thiele. [1]

Contents

Conjugation is the overlap of one p-orbital with another across an adjacent σ bond (in transition metals, d-orbitals can be involved). [2] [lower-alpha 1]

A conjugated system has a region of overlapping p-orbitals, bridging the interjacent locations that simple diagrams illustrate as not having a π bond. They allow a delocalization of π electrons across all the adjacent aligned p-orbitals. [3] The π electrons do not belong to a single bond or atom, but rather to a group of atoms.

Molecules containing conjugated systems of orbitals and electrons are called conjugated molecules, which have overlapping p orbitals on three or more atoms. Some simple organic conjugated molecules are 1,3-butadiene, benzene, and allylic carbocations. [4] The largest conjugated systems are found in graphene, graphite, conductive polymers and carbon nanotubes.

Chemical bonding in conjugated systems

Some prototypical examples of species with delocalized bonding. Top row: pyridine, furan, tropylium cation. Second row: allyl radical, acetate ion, acrolein. Atoms involved are in bold red, while electrons involved in delocalized bonding are in blue. (Particular attention should be paid to the involvement or non-involvement of "non-bonding" electrons.) Delocalization-updated.png
Some prototypical examples of species with delocalized bonding. Top row: pyridine, furan, tropylium cation. Second row: allyl radical, acetate ion, acrolein. Atoms involved are in bold red, while electrons involved in delocalized bonding are in blue. (Particular attention should be paid to the involvement or non-involvement of "non-bonding" electrons.)

Conjugation is possible by means of alternating single and double bonds in which each atom supplies a p orbital perpendicular to the plane of the molecule. However, that is not the only way for conjugation to take place. As long as each contiguous atom in a chain has an available p orbital, the system can be considered conjugated. For example, furan is a five-membered ring with two alternating double bonds flanking an oxygen. The oxygen has two lone pairs, one of which occupies a p orbital perpendicular to the ring on that position, thereby maintaining the conjugation of that five-membered ring by overlap with the perpendicular p orbital on each of the adjacent carbon atoms. The other lone pair remains in plane and does not participate in conjugation.

In general, any sp2 or sp-hybridized carbon or heteroatom, including ones bearing an empty orbital or lone pair orbital, can participate in conjugated systems, though lone pairs do not always participate in a conjugated system. For example, in pyridine, the nitrogen atom already participates in the conjugated system through a formal double bond with an adjacent carbon, so the lone pair remains in the plane of the ring in an sp2 hybrid orbital and does not participate in the conjugation. A requirement for conjugation is orbital overlap; thus, the conjugated system must be planar (or nearly so). As a consequence, lone pairs which do participate in conjugated systems will occupy orbitals of pure p character instead of spn hybrid orbitals typical for nonconjugated lone pairs.

The p system of furan and lone pairs. Note that one of the oxygen lone pairs participates in conjugation in a p orbital, while the other lone pair is in an sp hybridized orbital in the plane of the molecule and not part of the p system. The participation of six electrons in the p system makes furan aromatic (see below). Furan-pi-system.png
The π system of furan and lone pairs. Note that one of the oxygen lone pairs participates in conjugation in a p orbital, while the other lone pair is in an sp hybridized orbital in the plane of the molecule and not part of the π system. The participation of six electrons in the π system makes furan aromatic (see below).

A common model for the treatment of conjugated molecules is a composite valence bond / Hückel molecular orbital theory (VB/HMOT) treatment, in which the σ framework of the molecule is separated from the π system (or systems) of the molecule (see the article on the sigma-pi and equivalent-orbital models for this model and an alternative treatment). Although σ bonding can be treated using a delocalized approach as well, it is generally the π bonding that is being considered when delocalized bonding is invoked in the context of simple organic molecules.

Sigma (σ) framework: The σ framework is described by a strictly localized bonding scheme and consists of σ bonds formed from the interactions between sp3-, sp2-, and sp-hybridized atomic orbitals on the main group elements (and 1s atomic orbitals on hydrogen), together with localized lone pairs derived from filled, nonbonding hybrid orbitals. The interaction that results in σ bonding takes the form of head-to-head overlap of the larger lobe of each hybrid orbital (or the single spherical lobe of a hydrogen 1s orbital). Each atomic orbital contributes one electron when the orbitals overlap pairwise to form two-electron σ bonds, or two electrons when the orbital constitutes a lone pair. These localized orbitals (bonding and non-bonding) are all located in the plane of the molecule, with σ bonds mainly localized between nuclei along the internuclear axis.

Pi (π) system or systems: Orthogonal to the σ framework described above, π bonding occurs above and below the plane of the molecule where σ bonding takes place. The π system(s) of the molecule are formed by the interaction of unhybridized p atomic orbitals on atoms employing sp2- and sp-hybridization. The interaction that results in π bonding takes place between p orbitals that are adjacent by virtue of a σ bond joining the atoms and takes the form of side-to-side overlap of the two equally large lobes that make up each p orbital. Atoms that are sp3-hybridized do not have an unhybridized p orbital available for participation in π bonding and their presence necessarily terminates a π system or separates two π systems. A basis p orbital that takes part in a π system can contribute one electron (which corresponds to half of a formal "double bond"), two electrons (which corresponds to a delocalized "lone pair"), or zero electrons (which corresponds to a formally "empty" orbital). Bonding for π systems formed from the overlap of more than two p orbitals is handled using the Hückel approach to obtain a zeroth order (qualitative) approximation of the π symmetry molecular orbitals that result from delocalized π bonding.

Using the s/p-separation scheme to describe bonding, the Lewis resonance structures of a molecule like diazomethane can be translated into a bonding picture consisting of p-systems and localized lone pairs superimposed on a localized framework of s-bonds. Diazomethane-bonding scheme.png
Using the σ/π-separation scheme to describe bonding, the Lewis resonance structures of a molecule like diazomethane can be translated into a bonding picture consisting of π-systems and localized lone pairs superimposed on a localized framework of σ-bonds.

This simple model for chemical bonding is successful for the description of most normal-valence molecules consisting of only s- and p-block elements, although systems that involve electron-deficient bonding, including nonclassical carbocations, lithium and boron clusters, and hypervalent centers require significant modifications in which σ bonds are also allowed to delocalize and are perhaps better treated with canonical molecular orbitals that are delocalized over the entire molecule. Likewise, d- and f-block organometallics are also inadequately described by this simple model. Bonds in strained small rings (such as cyclopropane or epoxide) are not well-described by strict σ/π separation, as bonding between atoms in the ring consists of "bent bonds" or "banana bonds" that are bowed outward and are intermediate in nature between σ and π bonds. Nevertheless, organic chemists frequently use the language of this model to rationalize the structure and reactivity of typical organic compounds.

Electrons in conjugated π systems are shared by all adjacent sp2- and sp-hybridized atoms that contribute overlapping, parallel p atomic orbitals. As such, the atoms and π-electrons involved behave as one large bonded system. These systems are often referred to 'n-center k-electron π-bonds,' compactly denoted by the symbol Πk
n
, to emphasize this behavior. For example, the delocalized π electrons in acetate anion and benzene are said to be involved in Π4
3
and Π6
6
systems, respectively (see the article on three-center four-electron bonding ). It is important to recognize that, generally speaking, these multi-center bonds correspond to the occupation of several molecular orbitals (MOs) with varying degrees of bonding or non-bonding character (filling of orbitals with antibonding character is uncommon). Each one is occupied by one or two electrons in accordance with the Aufbau principle and Hund's rule. Cartoons showing overlapping p orbitals, like the one for benzene below, show the basis p atomic orbitals before they are combined to form molecular orbitals. In compliance with the Pauli exclusion principle, overlapping p orbitals do not result in the formation of one large MO containing more than two electrons.

Hückel MO theory is commonly used approach to obtain a zeroth order picture of delocalized π molecular orbitals, including the mathematical sign of the wavefunction at various parts of the molecule and the locations of nodal planes. It is particularly easy to apply for conjugated hydrocarbons and provides a reasonable approximation as long as the molecule is assumed to be planar with good overlap of p orbitals.

Stabilization energy

The quantitative estimation of stabilization from conjugation is notoriously contentious and depends on the implicit assumptions that are made when comparing reference systems or reactions. The energy of stabilization is known as the resonance energy when formally defined as the difference in energy between the real chemical species and the hypothetical species featuring localized π bonding that corresponds to the most stable resonance form. [5] This energy cannot be measured, and a precise definition accepted by most chemists will probably remain elusive. Nevertheless, some broad statements can be made. In general, stabilization is more significant for cationic systems than neutral ones. For buta-1,3-diene, a crude measure of stabilization is the activation energy for rotation of the C2-C3 bond. This places the resonance stabilization at around 6 kcal/mol. [6] Comparison of heats of hydrogenation of 1,4-pentadiene and 1,3-pentadiene estimates a slightly more modest value of 3.5 kcal/mol. [7] For comparison, allyl cation has a gas-phase rotation barrier of around 38 kcal/mol, [8] a much greater penalty for loss of conjugation. Comparison of hydride ion affinities of propyl cation and allyl cation, corrected for inductive effects, results in a considerably lower estimate of the resonance energy at 20–22 kcal/mol. [9] Nevertheless, it is clear that conjugation stabilizes allyl cation to a much greater extent than buta-1,3-diene. In contrast to the usually minor effect of neutral conjugation, aromatic stabilization can be considerable. Estimates for the resonance energy of benzene range from around 36–73 kcal/mol. [10]

Homoconjugation weakens the double bond character of the C=O bond, resulting in a lower IR frequency. Homoconjugation.png
Homoconjugation weakens the double bond character of the C=O bond, resulting in a lower IR frequency.

There are also other types of interactions that generalize the idea of interacting p orbitals in a conjugated system. The concept of hyperconjugation holds that certain σ bonds can also delocalize into a low-lying unoccupied orbital of a π system or an unoccupied p orbital. Hyperconjugation is commonly invoked to explain the stability of alkyl substituted radicals and carbocations. Hyperconjugation is less important for species in which all atoms satisfy the octet rule, but a recent computational study supports hyperconjugation as the origin of the increased stability of alkenes with a higher degree of substitution (Zaitsev's rule). [11]

Homoconjugation [12] is an overlap of two π-systems separated by a non-conjugating group, such as CH2. Unambiguous examples are comparatively rare in neutral systems, due to a comparatively minor energetic benefit that is easily overridden by a variety of other factors; however, they are common in cationic systems in which a large energetic benefit can be derived from delocalization of positive charge (see the article on homoaromaticity for details.). [13] Neutral systems generally require constrained geometries favoring interaction to produce significant degrees of homoconjugation. [14] In the example below, the carbonyl stretching frequencies of the IR spectra of the respective compounds demonstrate homoconjugation, or lack thereof, in the neutral ground state molecules.

Due to the partial π character of formally σ bonds in a cyclopropane ring, evidence for transmission of "conjugation" through cyclopropanes has also been obtained. [15]

Two appropriately aligned π systems whose ends meet at right angles can engage in spiroconjugation [16] or in homoconjugation across the spiro atom.

Delocalization of negative charge in a generic carboxylate anion, derived from an organic carboxylic acid (cf. acetic acid), and the corresponding vinylogous carboxylate anion (the "vinylog/vinylogue" of the carboxylate anion), where a vinyl group now separates the charged oxygen from the carbonyl (
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C=O) group. The validity of the theoretical concept of vinylogy is supported by the pKa of such vinylogs, which approach that of the analogous carboxylic acid. Vinylogous.png
Delocalization of negative charge in a generic carboxylate anion, derived from an organic carboxylic acid (cf. acetic acid), and the corresponding vinylogous carboxylate anion (the "vinylog/vinylogue" of the carboxylate anion), where a vinyl group now separates the charged oxygen from the carbonyl (C=O) group. The validity of the theoretical concept of vinylogy is supported by the pKa of such vinylogs, which approach that of the analogous carboxylic acid.

Vinylogy is the extension of a functional group through a conjugated organic bonding system, which transmits electronic effects. [17]

Conjugated cyclic compounds

Basis p orbitals of benzene. Benzene Orbitals.svg
Basis p orbitals of benzene.
Benzene p molecular orbitals according to Huckel theory. Molecular orbitals are frequently described as linear combinations of atomic orbitals, whose coefficients are indicated here by the size and shading of the orbital lobes. Benzene MO diagram.png
Benzene π molecular orbitals according to Hückel theory. Molecular orbitals are frequently described as linear combinations of atomic orbitals, whose coefficients are indicated here by the size and shading of the orbital lobes.

Cyclic compounds can be partly or completely conjugated. Annulenes, completely conjugated monocyclic hydrocarbons, may be aromatic, nonaromatic or antiaromatic.

Aromatic compounds

Compounds that have a monocyclic, planar conjugated system containing (4n + 2) π-electrons for whole numbers n are aromatic and exhibit an unusual stability. The classic example benzene has a system of six π electrons, which, together with the planar ring of C–C σ bonds containing 12 electrons and radial C–H σ bonds containing six electrons, forms the thermodynamically and kinetically stable benzene ring , the common core of the benzenoid aromatic compounds. For benzene itself, there are two equivalent conjugated contributing Lewis structures (the so-called Kekulé structures) that predominate. [18] [19] The true electronic structure is therefore a quantum-mechanical combination (resonance hybrid) of these contributors, which results in the experimentally observed C–C bonds which are intermediate between single and double bonds and of equal strength and length. In the molecular orbital picture, the six p atomic orbitals of benzene combine to give six molecular orbitals. Three of these orbitals, which lie at lower energies than the isolated p orbital and are therefore net bonding in character (one molecular orbital is strongly bonding, while the other two are equal in energy but bonding to a lesser extent) are occupied by six electrons, while three destabilized orbitals of overall antibonding character remain unoccupied. The result is strong thermodynamic and kinetic aromatic stabilization. Both models describe rings of π electron density above and below the framework of C–C σ bonds.

Nonaromatic and antiaromatic compounds

Cyclooctatetraene. Adjacent double bonds are not coplanar. The double bonds are therefore not conjugated. All-Z-Cyclooctatetraene 3D skeletal formula.svg
Cyclooctatetraene. Adjacent double bonds are not coplanar. The double bonds are therefore not conjugated.

Not all compounds with alternating double and single bonds are aromatic. Cyclooctatetraene, for example, possesses alternating single and double bonds. The molecule typically adopts a "tub" conformation. Because the p orbitals of the molecule do not align themselves well in this non-planar molecule, the π bonds are essentially isolated and not conjugated. The lack of conjugation allows the 8 π electron molecule to avoid antiaromaticity, a destabilizing effect associated with cyclic, conjugated systems containing 4n π (n = 0, 1, 2, ...) electrons. This effect is due to the placement of two electrons into two degenerate nonbonding (or nearly nonbonding) orbitals of the molecule, which, in addition to drastically reducing the thermodynamic stabilization of delocalization, would either force the molecule to take on triplet diradical character, or cause it to undergo Jahn-Teller distortion to relieve the degeneracy. This has the effect of greatly increasing the kinetic reactivity of the molecule. Because of the lack of long-range interactions, cyclooctatetraene takes on a nonplanar conformation and is nonaromatic in character, behaving as a typical alkene. In contrast, derivatives of the cyclooctatetraene dication and dianion have been found to be planar experimentally, in accord with the prediction that they are stabilized aromatic systems with 6 and 10 π electrons, respectively. Because antiaromaticity is a property that molecules try to avoid whenever possible, only a few experimentally observed species are believed to be antiaromatic. Cyclobutadiene and cyclopentadienyl cation are commonly cited as examples of antiaromatic systems.

In pigments

In a conjugated pi-system, electrons are able to capture certain photons as the electrons resonate along a certain distance of p-orbitals - similar to how a radio antenna detects photons along its length. Typically, the more conjugated (longer) the pi-system is, the longer the wavelength of photon can be captured. Compounds whose molecules contain a sufficient number of conjugated bonds can absorb light in the visible region, and therefore appear colorful to the eye, usually appearing yellow or red. [20]

Many dyes make use of conjugated electron systems to absorb visible light, giving rise to strong colors. For example, the long conjugated hydrocarbon chain in beta-carotene leads to its strong orange color. When an electron in the system absorbs a photon of light of the right wavelength, it can be promoted to a higher energy level. A simple model of the energy levels is provided by the quantum-mechanical problem of a one-dimensional particle in a box of length L, representing the movement of a π electron along a long conjugated chain of carbon atoms. In this model the lowest possible absorption energy corresponds to the energy difference between the highest occupied molecular orbital (HOMO) and the lowest unoccupied molecular orbital (LUMO). For a chain of n C=C bonds or 2n carbon atoms in the molecular ground state, there are 2n π electrons occupying n molecular orbitals, so that the energy gap is [21]

Since the box length L increases approximately linearly with the number of C=C bonds n, this means that the energy ΔE of a photon absorbed in the HOMO–LUMO transition is approximately proportional to 1/n. The photon wavelength λ = hcE is then approximately proportional to n. Although this model is very approximate, λ does in general increase with n (or L) for similar molecules. For example, the HOMO–LUMO absorption wavelengths for conjugated butadiene, hexatriene and octatetraene are 217 nm, 252 nm and 304 nm respectively. [22] However, for good numerical agreement of the particle in a box model with experiment, the single-bond/double-bond bond length alternations of the polyenes must be taken into account. [23] Alternatively, one can use the Hückel method which is also designed to model the electronic structure of conjugated systems.

Many electronic transitions in conjugated π-systems are from a predominantly bonding molecular orbital (MO) to a predominantly antibonding MO (π to π*), but electrons from non-bonding lone pairs can also be promoted to a π-system MO (n to π*) as often happens in charge-transfer complexes. A HOMO to LUMO transition is made by an electron if it is allowed by the selection rules for electromagnetic transitions. Conjugated systems of fewer than eight conjugated double bonds absorb only in the ultraviolet region and are colorless to the human eye. With every double bond added, the system absorbs photons of longer wavelength (and lower energy), and the compound ranges from yellow to red in color. Compounds that are blue or green typically do not rely on conjugated double bonds alone.

This absorption of light in the ultraviolet to visible spectrum can be quantified using ultraviolet–visible spectroscopy, and forms the basis for the entire field of photochemistry.

Conjugated systems that are widely used for synthetic pigments and dyes are diazo and azo compounds and phthalocyanine compounds.

Phthalocyanine compounds

Conjugated systems not only have low energy excitations in the visible spectral region but they also accept or donate electrons easily. Phthalocyanines, which, like Phthalocyanine Blue BN and Phthalocyanine Green G, often contain a transition metal ion, exchange an electron with the complexed transition metal ion that easily changes its oxidation state. Pigments and dyes like these are charge-transfer complexes.

Copper phthalocyanine Copper phthalocyanine.svg
Copper phthalocyanine

Porphyrins and similar compounds

Porphyrins have conjugated molecular ring systems (macrocycles) that appear in many enzymes of biological systems. As a ligand, porphyrin forms numerous complexes with metallic ions like iron in hemoglobin that colors blood red. Hemoglobin transports oxygen to the cells of our bodies. Porphyrin–metal complexes often have strong colors. A similar molecular structural ring unit called chlorin is similarly complexed with magnesium instead of iron when forming part of the most common forms of chlorophyll molecules, giving them a green color. Another similar macrocycle unit is corrin, which complexes with cobalt when forming part of cobalamin molecules, constituting Vitamin B12, which is intensely red. The corrin unit has six conjugated double bonds but is not conjugated all the way around its macrocycle ring.

Heme b.svg C-3 position Chlorophyll a.svg Cobalamin.svg
Heme group of hemoglobin The chlorin section of the chlorophyll a molecule. The green box shows a group that varies between chlorophyll types. Cobalamin structure includes a corrin macrocycle.

Chromophores

Conjugated systems form the basis of chromophores, which are light-absorbing parts of a molecule that can cause a compound to be colored. Such chromophores are often present in various organic compounds and sometimes present in polymers that are colored or glow in the dark. Chromophores often consist of a series of conjugated bonds and/or ring systems, commonly aromatic, which can include C–C, C=C, C=O, or N=N bonds.

Chemical structure of beta-carotene. The eleven conjugated double bonds that form the chromophore of the molecule are highlighted in red. Beta-Carotene conjugation.svg
Chemical structure of beta-carotene. The eleven conjugated double bonds that form the chromophore of the molecule are highlighted in red.

Conjugated chromophores are found in many organic compounds including azo dyes (also artificial food additives), compounds in fruits and vegetables (lycopene and anthocyanidins), photoreceptors of the eye, and some pharmaceutical compounds such as the following:

This polyene antimycotic called Amphotericin B has a conjugated system with seven double bonds acting as a chromophore that absorbs strongly in the ultraviolet-visible spectrum, giving it a yellow color Amphotericin B new.svg
This polyene antimycotic called Amphotericin B has a conjugated system with seven double bonds acting as a chromophore that absorbs strongly in the ultraviolet–visible spectrum, giving it a yellow color

Conjugated polymer nanoparticles (PDots) are assembled from hydrophobic fluorescent conjugated polymers, along with amphiphilic polymers to provide water solubility. Pdots are important labels for single-molecule fluorescence microscopy, based on high brightness, lack of blinking or dark fraction, and slow photobleaching. [24] [25]

See also

Notes

  1. For the purposes of this article, we are primarily concerned with delocalized orbitals with π-symmetry. This is in line with the typical usage of 'conjugated system' to refer to π (and not σ) delocalization. Canonical molecular orbitals are fully delocalized, so in a sense, all electrons involved in bonding, including ones making up the σ bonds and lone pairs, are delocalized throughout the molecule. However, while treating π electrons as delocalized yields many useful insights into chemical reactivity, treatment of σ and nonbonding electrons in the same way is generally less profitable, except in cases of multicenter σ-bonding as found in cluster compounds of Li and B. Moreover, the added complexity tends to impede chemical intuition. Hence, for most organic molecules, chemists commonly use a localized orbital model to describe the σ-bonds and lone pairs, while superimposing delocalized molecular orbitals to describe the π-bonding. This view has the added advantage that there is a clear correspondence between the Lewis structure of a molecule and the orbitals used to describe its bonding.

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<span class="mw-page-title-main">Covalent bond</span> Chemical bond by sharing of electron pairs

A covalent bond is a chemical bond that involves the sharing of electrons to form electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs. The stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding. For many molecules, the sharing of electrons allows each atom to attain the equivalent of a full valence shell, corresponding to a stable electronic configuration. In organic chemistry, covalent bonding is much more common than ionic bonding.

In chemistry, the mesomeric effect is a property of substituents or functional groups in a chemical compound. It is defined as the polarity produced in the molecule by the interaction of two pi bonds or between a pi bond and lone pair of electrons present on an adjacent atom. This change in electron arrangement results in the formation of resonance structures that hybridize into the molecule's true structure. The pi electrons then move away from or toward a particular substituent group. The mesomeric effect is stronger in compounds with a lower ionization potential. This is because the electron transfer states will have lower energies.

<span class="mw-page-title-main">Aromaticity</span> Chemical property

In organic chemistry, aromaticity is a chemical property describing the way in which a conjugated ring of unsaturated bonds, lone pairs, or empty orbitals exhibits a stabilization stronger than would be expected by the stabilization of conjugation alone. The earliest use of the term was in an article by August Wilhelm Hofmann in 1855. There is no general relationship between aromaticity as a chemical property and the olfactory properties of such compounds.

In chemistry, resonance, also called mesomerism, is a way of describing bonding in certain molecules or polyatomic ions by the combination of several contributing structures into a resonance hybrid in valence bond theory. It has particular value for analyzing delocalized electrons where the bonding cannot be expressed by one single Lewis structure. The resonance hybrid is the accurate structure for a molecule or ion; it is an average of the theoretical contributing structures.

<span class="mw-page-title-main">Lone pair</span> Pair of valence electrons which are not shared with another atom in a covalent bond

In chemistry, a lone pair refers to a pair of valence electrons that are not shared with another atom in a covalent bond and is sometimes called an unshared pair or non-bonding pair. Lone pairs are found in the outermost electron shell of atoms. They can be identified by using a Lewis structure. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of electrons in lone pairs plus the number of electrons in bonds equals the number of valence electrons around an atom.

In chemistry, molecular orbital theory is a method for describing the electronic structure of molecules using quantum mechanics. It was proposed early in the 20th century.

<span class="mw-page-title-main">Delocalized electron</span> Electrons that are not associated with a single atom or covalent bond

In chemistry, delocalized electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or a covalent bond.

<span class="mw-page-title-main">Hückel's rule</span> Method of determining aromaticity in organic molecules

In organic chemistry, Hückel's rule predicts that a planar ring molecule will have aromatic properties if it has 4n + 2 π-electrons, where n is a non-negative integer. The quantum mechanical basis for its formulation was first worked out by physical chemist Erich Hückel in 1931. The succinct expression as the 4n + 2 rule has been attributed to W. v. E. Doering (1951), although several authors were using this form at around the same time.

Antiaromaticity is a chemical property of a cyclic molecule with a π electron system that has higher energy, i.e., it is less stable due to the presence of 4n delocalised electrons in it, as opposed to aromaticity. Unlike aromatic compounds, which follow Hückel's rule and are highly stable, antiaromatic compounds are highly unstable and highly reactive. To avoid the instability of antiaromaticity, molecules may change shape, becoming non-planar and therefore breaking some of the π interactions. In contrast to the diamagnetic ring current present in aromatic compounds, antiaromatic compounds have a paramagnetic ring current, which can be observed by NMR spectroscopy.

<span class="mw-page-title-main">Sigma bond</span> Covalent chemical bond

In chemistry, sigma bonds or sigma overlap are the strongest type of covalent chemical bond. They are formed by head-on overlapping between atomic orbitals along the internuclear axis. Sigma bonding is most simply defined for diatomic molecules using the language and tools of symmetry groups. In this formal approach, a σ-bond is symmetrical with respect to rotation about the bond axis. By this definition, common forms of sigma bonds are s+s, pz+pz, s+pz and dz2+dz2 . Quantum theory also indicates that molecular orbitals (MO) of identical symmetry actually mix or hybridize. As a practical consequence of this mixing of diatomic molecules, the wavefunctions s+s and pz+pz molecular orbitals become blended. The extent of this mixing depends on the relative energies of the MOs of like symmetry.

In chemistry, orbital hybridisation is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. For example, in a carbon atom which forms four single bonds, the valence-shell s orbital combines with three valence-shell p orbitals to form four equivalent sp3 mixtures in a tetrahedral arrangement around the carbon to bond to four different atoms. Hybrid orbitals are useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. Usually hybrid orbitals are formed by mixing atomic orbitals of comparable energies.

<span class="mw-page-title-main">Cyclooctadecanonaene</span> Chemical compound

Cyclooctadecanonaene or [18]annulene is an organic compound with chemical formula C
18
H
18
. It belongs to the class of highly conjugated compounds known as annulenes and is aromatic. The usual isomer that [18]annulene refers to is the most stable one, containing six interior hydrogens and twelve exterior ones, with the nine formal double bonds in the cis,trans,trans,cis,trans,trans,cis,trans,trans configuration. It is reported to be a red-brown crystalline solid.

<span class="mw-page-title-main">Hyperconjugation</span> Concept in organic chemistry

In organic chemistry, hyperconjugation refers to the delocalization of electrons with the participation of bonds of primarily σ-character. Usually, hyperconjugation involves the interaction of the electrons in a sigma (σ) orbital with an adjacent unpopulated non-bonding p or antibonding σ* or π* orbitals to give a pair of extended molecular orbitals. However, sometimes, low-lying antibonding σ* orbitals may also interact with filled orbitals of lone pair character (n) in what is termed negative hyperconjugation. Increased electron delocalization associated with hyperconjugation increases the stability of the system. In particular, the new orbital with bonding character is stabilized, resulting in an overall stabilization of the molecule. Only electrons in bonds that are in the β position can have this sort of direct stabilizing effect — donating from a sigma bond on an atom to an orbital in another atom directly attached to it. However, extended versions of hyperconjugation can be important as well. The Baker–Nathan effect, sometimes used synonymously for hyperconjugation, is a specific application of it to certain chemical reactions or types of structures.

The Hückel method or Hückel molecular orbital theory, proposed by Erich Hückel in 1930, is a simple method for calculating molecular orbitals as linear combinations of atomic orbitals. The theory predicts the molecular orbitals for π-electrons in π-delocalized molecules, such as ethylene, benzene, butadiene, and pyridine. It provides the theoretical basis for Hückel's rule that cyclic, planar molecules or ions with π-electrons are aromatic. It was later extended to conjugated molecules such as pyridine, pyrrole and furan that contain atoms other than carbon and hydrogen (heteroatoms). A more dramatic extension of the method to include σ-electrons, known as the extended Hückel method (EHM), was developed by Roald Hoffmann. The extended Hückel method gives some degree of quantitative accuracy for organic molecules in general and was used to provide computational justification for the Woodward–Hoffmann rules. To distinguish the original approach from Hoffmann's extension, the Hückel method is also known as the simple Hückel method (SHM).

<span class="mw-page-title-main">Antibonding molecular orbital</span> Molecular orbital which weakens chemical bonding

In theoretical chemistry, an antibonding orbital is a type of molecular orbital that weakens the chemical bond between two atoms and helps to raise the energy of the molecule relative to the separated atoms. Such an orbital has one or more nodes in the bonding region between the nuclei. The density of the electrons in the orbital is concentrated outside the bonding region and acts to pull one nucleus away from the other and tends to cause mutual repulsion between the two atoms. This is in contrast to a bonding molecular orbital, which has a lower energy than that of the separate atoms, and is responsible for chemical bonds.

<span class="mw-page-title-main">Homoaromaticity</span> Organic molecular structure

Homoaromaticity, in organic chemistry, refers to a special case of aromaticity in which conjugation is interrupted by a single sp3 hybridized carbon atom. Although this sp3 center disrupts the continuous overlap of p-orbitals, traditionally thought to be a requirement for aromaticity, considerable thermodynamic stability and many of the spectroscopic, magnetic, and chemical properties associated with aromatic compounds are still observed for such compounds. This formal discontinuity is apparently bridged by p-orbital overlap, maintaining a contiguous cycle of π electrons that is responsible for this preserved chemical stability.

Physical organic chemistry, a term coined by Louis Hammett in 1940, refers to a discipline of organic chemistry that focuses on the relationship between chemical structures and reactivity, in particular, applying experimental tools of physical chemistry to the study of organic molecules. Specific focal points of study include the rates of organic reactions, the relative chemical stabilities of the starting materials, reactive intermediates, transition states, and products of chemical reactions, and non-covalent aspects of solvation and molecular interactions that influence chemical reactivity. Such studies provide theoretical and practical frameworks to understand how changes in structure in solution or solid-state contexts impact reaction mechanism and rate for each organic reaction of interest.

A non-bonding orbital, also known as non-bonding molecular orbital (NBMO), is a molecular orbital whose occupation by electrons neither increases nor decreases the bond order between the involved atoms. Non-bonding orbitals are often designated by the letter n in molecular orbital diagrams and electron transition notations. Non-bonding orbitals are the equivalent in molecular orbital theory of the lone pairs in Lewis structures. The energy level of a non-bonding orbital is typically in between the lower energy of a valence shell bonding orbital and the higher energy of a corresponding antibonding orbital. As such, a non-bonding orbital with electrons would commonly be a HOMO.

<span class="mw-page-title-main">Radical (chemistry)</span> Atom, molecule, or ion that has an unpaired valence electron; typically highly reactive

In chemistry, a radical, also known as a free radical, is an atom, molecule, or ion that has at least one unpaired valence electron. With some exceptions, these unpaired electrons make radicals highly chemically reactive. Many radicals spontaneously dimerize. Most organic radicals have short lifetimes.

<span class="mw-page-title-main">Stereoelectronic effect</span> Affect on molecular properties due to spatial arrangement of electron orbitals

In chemistry, primarily organic and computational chemistry, a stereoelectronic effect is an effect on molecular geometry, reactivity, or physical properties due to spatial relationships in the molecules' electronic structure, in particular the interaction between atomic and/or molecular orbitals. Phrased differently, stereoelectronic effects can also be defined as the geometric constraints placed on the ground and/or transition states of molecules that arise from considerations of orbital overlap. Thus, a stereoelectronic effect explains a particular molecular property or reactivity by invoking stabilizing or destabilizing interactions that depend on the relative orientations of electrons in space.

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  13. Some orbital overlap is possible even between bonds separated by one (or more) CH2 because the bonding electrons occupy orbitals which are quantum-mechanical functions and extend indefinitely in space. Macroscopic drawings and models with sharp boundaries are misleading because they do not show this aspect.
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