# Conjugated system

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Cinnamaldehyde is a naturally-occurring compound that has a conjugated system
1,3-pentadiene is a molecule with a conjugated system

In 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 increase 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]

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

A conjugated system has a region of overlapping p orbitals, bridging the interjacent locations that simple diagrams illustate 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.

The largest conjugated systems are found in graphene, graphite, conductive polymers, and carbon nanotubes.

## Mechanism

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. 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 and an oxygen in position 1. Oxygen has two lone pairs, one of which occupies a p-orbital on that position, thereby maintaining the conjugation of that five-membered ring. The presence of a nitrogen in the ring or groups α to the ring like a carbonyl group (C=O), an imine group (C=N), a vinyl group (C=C), or an anion will also suffice as a source of p orbitals to maintain conjugation. A requirement for conjugation is orbital overlap; thus, the conjugated system must be planar (or nearly so).

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.[4]

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 isolated p orbital or unoccupied orbital of a π system. It is commonly invoked to explain the stability of alkyl substituted radicals and carbocations. Homoconjugation[5] 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.).[6] Neutral systems generally require constrained geometries favoring interaction to produce significant degrees of homoconjugation.[7] 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.

Two appropriately aligned π systems whose ends meet at right angles can engage in spiroconjugation.[8]

## Conjugated cyclic compounds

Basis p orbitals of benzene.
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, non-aromatic or anti-aromatic.

### 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, 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,[9][10] so that the true electronic structure is a quantum-mechanical combination (resonance hybrid) of these contributors, resulting 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, 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.

### Non-aromatic compounds

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 the effect is so unfavorable, 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.

## In pigments

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[11]

${\displaystyle E_{n+1}-E_{n}={\frac {(2n+1)\hbar ^{2}\pi ^{2}}{2mL^{2}}}}$

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.[12] 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.[13]

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 Phthalo Blue and Phthalo Green, 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

### 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 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.

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

## References and notes

1. ^ Thiele, Johannes (1899). "Zur Kenntnis der ungesättigten Verbindungen" [[Contribution] to our knowledge of unsaturated compounds]. Justus Liebig's Annalen der Chemie (in German). 306: 87–142. On p. 90, Thiele coined the term "conjugated": "Ein solches System benachbarter Doppelbindungen mit ausgeglichenen inneren Partialvalenzen sei als conjugirt bezeichnet." (Such a system of adjacent double bonds with equalized inner partial valences shall be termed "conjugated".)
2. ^
3. ^ March Jerry; (1985). Advanced Organic Chemistry reactions, mechanisms and structure (3rd ed.). New York: John Wiley & Sons, inc. ISBN 0-471-85472-7
4. ^ For the purposes of this article, we are primarily concerned with delocalized orbitals with π symmetry. 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, and the added complexity impedes chemical intuition. Hence, chemists often implicitly use a (semi-)localized orbital model to describe lone pairs and the σ-bond framework of organic molecules. Any delocalized π system is then superimposed upon the localized skeleton.
5. ^ IUPAC Gold Book – homoconjugation
6. ^ 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.
7. ^ Scott, L. T. (1986-01-01). "Cyclic homoconjugation in neutral organic molecules". Pure and Applied Chemistry. 58 (1): 105–110. doi:10.1351/pac198658010105. ISSN 1365-3075.
8. ^ Maslak, Przemyslaw (May 1994). "Spiroconjugation: An added dimensi in the design of organic molecular materials". Advanced Materials. 6 (5): 405–407. doi:10.1002/adma.19940060515. ISSN 0935-9648.
9. ^ Rashid, Zahid; van Lenthe, Joop H. (March 2011). "Generation of Kekulé valence structures and the corresponding valence bond wave function". Journal of Computational Chemistry. 32 (4): 696–708. doi:10.1002/jcc.21655. ISSN 1096-987X. PMID 20941739.
10. ^ While the two Kekulé resonance forms contribute to most (>90%) of the π bond energy, there are also a number of other minor contributors to the wavefunction in the valence bond treatment, including the three Dewar resonance forms, and even smaller contributions from various ionic and singlet diradical forms. See article by Rashid and van Lenthe for a recent computational treatment.
11. ^ P. Atkins and J. de Paula Physical Chemistry (8th ed., W.H.Freeman 2006), p.281 ISBN 0-7167-8759-8
12. ^ Atkins and de Paula p.398
13. ^ Autschbach, Jochen (November 2007). "Why the Particle-in-a-Box Model Works Well for Cyanine Dyes but Not for Conjugated Polyenes". Journal of Chemical Education. 84 (11): 1840. doi:10.1021/ed084p1840. ISSN 0021-9584.