Inductive Effect

In chemistry and physics, the inductive effect is an experimentally observable effect of the transmission of charge through a chain of atoms in a molecule by electrostatic induction.[1] The net polar effect exerted by a substituent is a combination of this inductive effect and the mesomeric effect. The electron cloud in a σ-bond between two unlike atoms is not uniform and is slightly displaced towards the more electronegative of the two atoms. This causes a permanent state of bond polarization, where the more electronegative atom has a slight negative charge (δ–) and the other atom has a slight positive charge (δ+). If the electronegative atom is then joined to a chain of atoms, usually carbon, the positive charge is relayed to the other atoms in the chain. This is the electron-withdrawing inductive effect, also known as the − I effect. Some groups, such as the alkyl group are less electron-withdrawing than hydrogen and are therefore considered as electron-releasing. This is electron releasing character and is indicated by the + I effect. As the induced change in polarity is less than the original polarity, the inductive effect rapidly dies out, and is significant only over a short distance. The inductive effect is permanent but feeble, as it involves the shift of strongly held σ-bond electrons, and other stronger factors may overshadow this effect. The inductive effect may be caused by some molecules also. Relative inductive effects have been experimentally measured with reference to hydrogen.

Inductive effects can be measured through the Hammett equation. Inductive Effect can also be used to determine whether a molecule is stable or unstable depending on the charge present on the atom under consideration and the type of groups bonded to it. For example, if an atom has a positive charge and is attached to a −I group its charge becomes 'amplified' and the molecule becomes more unstable than if I-effect was not taken into consideration. Similarly, if an atom has a negative charge and is attached to a +I group its charge becomes 'amplified' and the molecule becomes more unstable than if I-effect was not taken into consideration. But, contrary to the above two cases, if an atom has a negative charge and is attached to a −I group its charge becomes 'de-amplified' and the molecule becomes more stable than if I-effect was not taken into consideration. Similarly, if an atom has a positive charge and is attached to a +I group its charge becomes 'de-amplified' and the molecule becomes more stable than if I-effect was not taken into consideration. The explanation for the above is given by the

fact that more charge on an atom decreases stability and less charge on an atom increases stability.

Contents
[hide]
   

1 Definition 2 Applications 3 See also 4 References

[edit] Definition
Polarity induced in a covalent bond due to the difference in electronegativities of the bonded atoms is called the inductive effect.

[edit] Applications




 

Aliphatic carboxylic acids. The strength of a carboxylic acid depends on the extent of its ionization: the more ionized it is, the stronger it is. As an acid becomes stronger, the numerical value of its pKa drops. In aliphatic acids, the electron-releasing inductive effect of the methyl group increases the electron density on oxygen and thus hinders the breaking of the O-H bond, which consequently reduces the ionization. Greater ionization in formic acid when compared to acetic acid makes formic acid (pKa=3.75) stronger than acetic acid (pKa=4.76). Monochloroacetic acid (pKa=2.82), though, is stronger than formic acid, since the electron-withdrawing effect of chlorine promotes ionization. Aromatic carboxylic acids. In benzoic acid, the carbon atoms which are present in the ring are sp2 hybridised.As a result, benzoic acid(pKa=4.20) is a stronger acid than cyclohexane carboxylic acid(pKa=4.87). Also, electron-withdrawing groups substituted at the ortho and para positions, enhance the acid strength. Dioic acids. Since the carboxyl group is itself an electron-withdrawing group, the dioic acids are, in general, stronger than their monocarboxyl analogues. In the so-called Baker–Nathan effect the observed order in electron-releasing alkyl substituents is apparently reversed.

[edit] See also

[2]

Inductive

effect

Induction of charge due to less or more electronegative element is known as inductive effect. It occurs till four carbon atom and maximum at first carbon atom due to closeness impact.

Application

of

inductive

effect

(a) Acidic strength of Carboxylic acid

As the 'S' character increases electronegativity of atom also increases so that electro negativity order is sp . >sp2> sp3

(b) Basic strength of amines

II pd. elements are more basic than III pd. elements because in II pd elements vacant d- orbitals are absent. H2O > H2S and NH3 > PH3 due to vacant d orbitals

H2O

<

NH3

For same period elements consider electronegativity order. Less electro negative element is more basic. R3N is less basic than R2NH in H2

O because of the steric hinderance and solvation effect caused by the three bulky ‗R‘ groups.

(c) Reactivity of Carbonyl Compounds

HCHO > CH3CHO > CH3COCH3

carbonyl compounds give nucleophilic addition reactions where primary attack of nucleophile takes place.

As the size of alkyl group increases stearic hinderance comes into play thus reactivity decreases.

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The factors affecting the electron availability of a compound might reasonably be thought to have far-reaching consequences upon its reactivity with various compounds. For example, an area of high electron density is unlikely to be attacked by OH-, but an area of low electron density is likely to be far more susceptible to attack by the same reagent.

Inductive Effects

So far, when considering covalent single bonds, we have though of the electrons as being in between the two nuclei of the atoms involved in the bond. However, what was not stated is that the electron density is shifted towards the more electronegative of the pair, i.e. this can be thought of as the electrons spending more time nearer one of the atoms.

In the alkyl halide above, the C-F bond is polarized towards fluorine (meaning that there is more electron density towards that end of the bond). This imbalance of charge can be represented as above by the use of δ+ and δ-. This effect is due to the greater electronegativity of the fluorine over the carbon. It can also be visualised as a contribution to the overall structure from this resonance form:(Note that though it does involve charge separation, usually a high energy process, the negative charge is stabilised by being on the highly electronegative fluorine atom.)

The inductive effect diminishes through a greater number of bonds. i.e.:

In the alkyl halide above, the greatest inductive effect experienced is on carbon 1, followed by carbon 2, 3, and 4 in order. This can be thought of as the fact that carbon 1 will be left slightly electron deficient, so in order to rectify the loss, it pulls some electron density over from carbon 2. However, the effect is very slight beyond carbon 2. Inductive effects work through sigma bonds, and can push electrons in either direction with respect to carbon. i.e. metals (e.g. magnesium, lithium) inductively donate electrons (because

they are electropositive), and electronegative elements such as chlorine, fluorine, and oxygen, inductively withdraw electrons.
Field Effects

An effect which is similar in nature to the inductive effect operates through the space surrounding the molecule, or (if in solution) through the solvent molecules that surround it. This is known as a field effect.

Inductive Effect on Acid and Base Strengths When an atom or group of atoms is substituted for another atom or group of atoms in a molecule the distribution of electron density changes. The effect of the new atom or group of atoms on this electron density distribution is termed the "inductive effect" of the atom or group. If the new atom or group of atoms (the substituent) is more electronegative then the one that it replaced then electrons will be withdrawn from other parts of the molecule toward the new substituent. On the other hand if the new substituent is less electronegative than the group that it replaced then the electron density will increase in the rest of the molecule. Changes in substituents can have a profound effect on acidities and basicities of molecules and the relative strengths of structurally related acids and bases can be predicted on the basis of inductive effects of substituents. A more electronegative substituent X will increase the acidity of an oxy (or hydroxy) acid XO-H by greater withdrawal of electron density from the oxygen-hydrogen bond. This both weakens the O-H bond and increases the positive charge on hydrogen. A similar effect is present for the conjugate base XO1-; here the presence of a more electronegative substituent stabilizes the anion by withdrawing charge from the oxygen and effectively delocalizing the charge over more of the molecule. The pKa values for the X(O) (OH) acids on the following page illustrate this notion as do the pKa values n m of the larger series of the following simple acids: pKa Cl-OH 7.5 Br-OH 8.7 I-OH 10.7 HO-OH 11.8 H-OH 15.7 CH OH 16.6 3 Note the effect of the CH - group versus H-. 3 The following series show a similar trend even though the substituent is not bonded directly to the ionizing hydroxyl group. pKa F-CH C(O)OH 2.7 F-SO -OH > Cl-SO -OH > HO-SO -OH 2 2 2 2 Cl-CH C(O)OH 2.8 2 Br-CH C(O)OH 2.9 F P(O)(OH) > FP(O)(OH) > HOP(O)(OH) 2 2 2 2 I-CH C(O)OH 3.0 2

H-CH C(O)OH 4.7 2 CH -CH C(O)OH 4.9 3 2 Similar effects exist for bases. For example base strengths vary in the order NH > H NNH > ClNH > Cl NH > NF , etc. 2 2 2 2 3 pK values for acids of the type XO (OH) a n m n=0 HOCl 7.4 HOBr 8.7 HOI 10.7 n=1 HOClO 1.94 HONO 3.3 (HO) PO 2.16 7.21 12.3 3 (HO) AsO 2.25 7.77 11.6 3 n=2 HOClO -1 2 HOIO 0.7 3 HONO -1.4 2 n=3 HOClO -10 3

3

pKa values for hydrated metal ions, [M(OH2)n]m+
Mm+ pKa Mm+ pKa Th4+ 3.2 Al3+ 5.0 Sc3+ 4.3 Y3+ 7.7 Cr3+ 4.0 La3+ 8.5 Fe3+ 2.2 Mg2+ 11.4 Cr2+ 10.0 Ca2+ 12.8 Mn2+ 10.6 Sr2+ 13.3 Fe2+ 9.5 Ba2+ 13.5 Co2+ 9.6 Ni2+ 9.9 Zn2+ 9.0 Li+ 13.6 Ag+ 12.0 Na+ 14.2 Tl+ 13.2 K+ 14.5 pKa=-log K for [M(OH2)n]m+ + H2O = [M(OH2)n-1OH](m-1)+ + H3O+
How to recognize a Lewis Acid or Lewis Base For the average person there is probably most uncertainity about what molecules can act as Lewis acids. The following points should be helpful.

1) Molecules for which a simple Lewis structure indicates an atom does not have four pairs of electrons in its valence shell (an octet) will behave as Lewis acids. The electron deficient atom will have a vacant p orbital, although in reality there may be a pi interaction with one or more of the atoms bonded to it (BX for example). 3 2) Molecules whose Lewis structures indicate an atom to have an octet as a result of the formation of one or more multiple bonds will often function as Lewis acids. Examples are CO , SO , SO . 2 3 2 3) Molecules that have central atoms that are capable of expanding their valence shell to include more than four pairs of sigma-bonding electron pairs, i.e., use hybrid orbitals that involve d orbitals, are Lewis acids. Examples are PCl , SF , I (recall the 5 4 2 hybridization model for I ), ICl . Note, however, that alternative models can be used to 3 3
-

describe the interaction of these acids with bases. A notable example is I in which the 2 molecular orbital approach involves the interaction of the sigma antibonding orbital of the I with the donor orbital of the base. The acceptor orbital will normally be the 2 LUMO. Lewis bases are typically easy to recognize. In principle any molecule in which there is an atom with a nonbonding, or lone, pair of electrons can function as a Lewis base. In practice there are certain types of molecules where the energies of the orbitals containing these nonbonding electrons are too low in energy to interact appreciably with most Lewis acids. Examples of such molecules are fluorine containing molecules. In some instances sigma bonding electrons are basic; group III (13) metal hydrides such as BH , BH and AlH are examples. The resulting bonds are electron deficient, 3 4 3
-

3-center, 2-electron bonds. There are some molecules that have more than one basic site (there are termed ambidentate ligands when referring to coordination compounds). Examples are NCS-, NO -, P[N(CH ) ] . In molecular orbital terms, if the HOMO is
2323

nonbonding it will most often be the donor orbital but the nature of the acid can result in a lower energy orbital functioning as the donor orbital. Some molecules can function as either an acid or a base, for example SO , and in a 2 few instances as both acid and base. Examples of the latter are BH , molecules in 3 which there is intramolecular hydrogen bonding, and molecules such as (CH ) B(CH ) N(CH ) in which cyclization occurs by formation of a B-N bond. 3 2 2 3 3 2

Inductive Effects of Alkyl Groups
From $1 Table of contents 1. 2. 3. 4. 1. Introduction 2. The Effect of an Electron Withdrawing Group on a Benzene Ring 3. Halogens: A Special Case 4. Outside links

5. 5. References 6. 6. Problems 7. 7. Contributors

A substituent on a benzene ring can effect the placement of additional substituents on that ring during Electrophilic Aromatic Substitution. How do we know where an additonal substituent will most likely be placed? The answer to this is through inductive and resonance effects. Inductive effects are directly correlated with electronegativity. Substituents can either be meta directing or ortho-para directing.
1. 2. 3. 4. 5. 6. 7. 1. Introduction 2. The Effect of an Electron Withdrawing Group on a Benzene Ring 3. Halogens: A Special Case 4. Outside links 5. References 6. Problems 7. Contributors Introduction

The three general positions of a disubstituted benzene ring are ortho, meta and para.

The Effect of an Electron Donating Groups on a Benzene Ring The first scenario for adding an electrophile to a monosubstituted benzene ring is when the substituent is an electron donating group. Electron donating groups are alkyl groups, phenyl groups or substituents that have a lone pair of electrons on the atom directly bonded to the ring. Electron donating groups are donating by induction (Activating and Deactivating Benzene Rings) and resonance. Examples of electron donating groups: -CH3, -OCH3, -OH, -NH2

Electron donating groups cause the second subtituent to add on to the para or ortho position on the benzene ring. The reason for this can be explained by the different carbocation resonance structures of the ortho, meta and para positions.

Some electron donating groups have an extra resonance form in which there is a double bond between the atom and the carbon on the benzene. This is a very stable resonance form. This is due to directing effects of substituents in conjugation wih the benzene ring (see sec. 16.3). When the electrophile is added to the ortho position, three different resonance forms are possible. Carbocation forms 1 and 2 are secondary carbocations, but position 3 forms a tertiary carbocation and the positive charge is on the carbon directly attached to the electron donating group, which is the most stable. This carbocation is also stablized by the electrons from the electron donating group. More stable intermediates (the carbocation) have lower transition state energies and therefore a faster reaction rate, forming more of this product. This is the reason that the ortho position is one of the major products.

If the electrophile is added to the monosubstituted benzene ring in the para position one of the three resonance forms of the carbocations will be a tertiary carbocation. This carbocation intermediate is the same as the one formed from ortho substitution. For the meta substituted carbocation resonance structures, there are three possible resonance froms that are secondary carbocations. These forms are not as stable as the tertiary carbocation form in the ortho and para substituted carbocations.

Therefore, the two major products of the reaction of a monosubstituted benzene ring with an electron donating group and additional electrophile are the ortho and para positions. It's important to note that the para product is slightly more common than the ortho product due to steric hindrance. H-NMR spectroscopy can be used to determine whether or not a compound has a second substituent at the ortho or para position. At the ortho position there are four distinct signals, but for the para position there are only two signals because the molecule is symmetrical.

Electron donating groups on a benzene ring are said to be activating, because they increase the rate of the second substitution so that it is higher than that of standard benzene. To summarize, electron donating groups are said to be ortho/para directing and they are activators.
The Effect of an Electron Withdrawing Group on a Benzene Ring

The other circumstance is when you have add an additional electrophile to a monosubstituted benzene ring with an electron withdrawing group on it. Electron withdrawing groups have an atom with a slight positive or full positive charge directly attached to a benzene ring. Examples of electron withdrawing groups: -CF3, -COOH, -CN Electron withdrawing groups only have one major product, the second substituent adds in the meta position. Again, this can be explained by the resonance forms of the carbocation intermediates.

When the second electrophile is added on to the benzene ring in the ortho position, the same three resonance forms of the carbocation are produced. Again, one form is a tertiary carbocation with the positive charge on the carbon directly attached to the electron withdrawing group. Unlike in the case with an electron donating group, this resonance form is much less stable. This is due to the electron withdrawing group pulling away electrons from the carbon, creating an

even stronger positive charge. This situation holds true for the para substituted tertiary carbocation resonace form as well. For the meta position, all the carbocations formed are secondary. Although these are not entirely stable, they are more favored than the resonance forms of the ortho and para positions.

The major product of a monosubstituted benzene ring with an electron withdrawing group and an additional electrophile is a product with meta substitution. In contrast to electron donating groups, electron withdrawing groups are deactivating. This means that the rate of the second substitution is lower than that of standard benzene. Table of Substituents
Ortho-Para Directing Meta Directing

Strong Activating

Moderately Activating

Weakly Activating

Weakly Deactivating

Moderately Deactivating

Strongly Deactivating

-NH2
-NHR -OH -OCH3

-NHCOR
-OCOR

-CH3
-phenyl

-F
-Cl -Br -I

-COH
-COCH3 -COOCH3 -SO3H

-NO2
-CF3 -CCl3

Halogens: A Special Case

Halogens are very electronegative. This means that inductively they are electron withdrawing. However, because of their ability to donate a lone pair of electrons in resonance forms, they are activators and ortho/para directing. Resonance forms win out in directing. Because they are electron withdrawing, halogens are very weak activators. (See sec 16.3) To summarize, electron withdrawing groups are meta directors and they are deactivators.
Outside links
 

http://en.wikipedia.org/wiki/Inductive_effect http://chemistry2.csudh.edu/pendavis/elaromsust.html

References 1. Bohm, S., and O. Exner. "Interaction of two functional groups through the benzene ring: Theory and Experiment." Journal of Computational Chemistry (2008) (p. 1) http://www.lib.berkeley.edu 2. Brown, William H., Foote, Christopher S., Iverson, Brent L. Organic Chemisty. 4th ed. Belmont, CA: Thomson Learning Inc./ Brooks/Cole, 2005. (pp. 868-872) 3. Schore, Neil E., Vollhardt, Peter C. Organic Chemistry, Structure and Function. 5th ed. New York: W.H. Freeman & Company, 2007. (pp. 724-728)

In strict definition, an experimentally observable effect (on rates of reaction, etc.) of the transmission of charge through a chain of atoms by electrostatic induction. A theoretical distinction may be made between the field effect, and the inductive effect as models for the Coulomb interaction between a given site within a molecular entity and a remote unipole or dipole within the same entity. The experimental distinction between the two effects has proved difficult, except for molecules of peculiar geometry, which may exhibit ‗reversed field effects‘. Ordinarily the inductive effect and the field effect are influenced in the same direction by structural changes in the molecule and the distinction between them is not clear. This situation has led many authors to include the field effect in the term ‗inductive effect‘. Thus the separation of values into inductive and resonance components does not imply the exclusive operation of a through-bonds route for the transmission of the non-conjugative part of the substituent effect. To indicate the all-inclusive use of the term inductive, the phrase ‗so-called inductive effect‘ is sometimes used. Certain modern theoretical approaches suggest that the ‗so-called inductive effect‘ reflects a field effect rather than through-bonds transmission. See also mesomeric effect, polar effect. 1994, 66, 1124

IUPAC Compendium of Chemical Terminology 2nd Edition (1997

Problems

Contributors

Resonance Effect: 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 formula. A molecule or ion with such delocalized electrons is represented by several contributing structures [2] (also called resonance structures or canonical forms). Each contributing structure can be represented by a Lewis structure, with only an integer number of covalent bonds between each pair of atoms within the structure.[3] Several Lewis structures are used collectively to describe the actual molecular structure. However these individual contributors cannot be observed in the actual resonance-stabilized molecule; the molecule does not oscillate back and forth between the contributing structures, as might be assumed from the word "resonance". The actual structure is an approximate intermediate between the canonical forms, but its overall energy is lower than each of the contributors. This intermediate form between different contributing structures is called a resonance hybrid.[4] Contributing structures differ only in the position of electrons, not in the position of nuclei. Resonance is a key component of valence bond theory. Electron delocalization lowers the potential energy of the substance and thus makes it more stable than any of the contributing structures. The difference between the potential energy of the actual structure and that of the contributing structure with the lowest potential energy is called the resonance energy[5] or delocalization energy. Resonance is distinguished from tautomerism and conformational isomerism, which involve the formation of isomers, thus the rearrangement of the nuclear positions.

Contents
[hide]
 

  

 

1 General characteristics of resonance 2 Use of contributing structures o 2.1 Resonance hybrids o 2.2 Major and minor contributors o 2.3 Contributing structures in diagrams 3 Bond lengths 4 Resonance energy o 4.1 Resonance energy of benzene 5 Resonance in quantum mechanics o 5.1 Molecular orbital (MO) versus valence bond (VB) theory o 5.2 Coefficients 6 History 7 Representations o 7.1 Reactive intermediates

  

8 See also 9 External links 10 References

[edit] 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". However, the real structure is not a rapid interconversion of contributing structures. Several Lewis structures are used together, because none of them exactly represents the actual structure. To represent the intermediate, a resonance hybrid is used instead. The contributing structures are not isomers. They differ only in the position of electrons, not in the position of nuclei. Each Lewis formula must have the same number of valence electrons (and thus the same total charge), and the same number of unpaired electrons, if any.[6] Bonds that have different bond orders in different contributing structures do not have typical bond lengths. Measurements reveal intermediate bond lengths. The real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be.

[edit] Use of contributing structures
In Lewis formulas, covalent bonds are represented in accordance with the valence bond theory. 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 in the Lewis formula cannot be indicated. The π electrons appear to be delocalized and the multiple bonds could be in different positions. In those cases the molecule cannot be represented by one single Lewis formula. To solve this problem, in valence

bond theory 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 molecular orbital theory already includes the concept of delocalized electrons and therefore has no need of the concept of resonance. None of the contributing structures is considered to represent the actual structure, since bonds that have a different bond order in different contributing structures do not have, if measured, a bond length that is typical for a normal single or multiple bond. Moreover, the overall energy of the actual structure is lowered with the resonance energy.
[edit] Resonance hybrids

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

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

that obey as much as possible the octet rule (8 valence electrons around each atom rather than having deficiencies or surplus) that have a maximum number of covalent bonds that carry a minimum of charged atoms with negative charge, if any, on the most electronegative atoms and positive charge, if any, on the most electropositive.

The greater the number of contributing structures, the more stable the molecule. This is because the more states at lower energy are available to the electrons in a particular molecule, the more stable the electrons are. Also the more volume electrons can occupy at lower energy the more stable the molecule is. 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.

[edit] Contributing structures in diagrams

Contributing structures of the thiocyanate ion, enclosed in square brackets.

Hybrid of the nitrate ion

Hybrid of benzene.

In diagrams, contributing structures are typically separated by double-headed arrows ( ). The arrow should not be confused with the right and left pointing equilibrium arrow ( ). All structures together may be enclosed in large square brackets, to indicate they picture one single molecule or ion, not different species in a chemical equilibrium. Alternatively to the use of resonance structures in diagrams, a hybrid diagram can be used. In a hybrid diagram, pi bonds that are involved in resonance are usually pictured as curves [7] 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.[8]

[edit] 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.

[edit] 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. Hybrids are therefore always more stable than any of the contributing structures would be.[9] The molecule is sometimes said to be "stabilized by resonance" or "resonance-stabilized," but the stabilization derives from electron delocalization, of which "resonance" is only a description. Delocalization of the π-electrons lowers the orbital energies, imparting this stability. The difference between the potential energy of the actual structure (the resonance hybrid) and that of the contributing structure with the lowest potential energy is called the "resonance energy".[5]
[edit] 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 benzene delivers 208.4 kJ (49.8 kcal).

Hydrogenation of one double bond 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 double bond. 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. [10]

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 = 151 kJ (36 kcal).[11]

Note: The values used here are from the article of Wiberg, Nakaji, Morgan (1993). Values from other sources may differ.

[edit] Resonance in quantum mechanics
Resonance has a deeper significance in the mathematical formalism of valence bond theory (VB). When a molecule cannot be represented by the standard tools of valence bond theory (promotion, hybridisation, orbital overlap, sigma and π bond formation) because no single structure predicted by VB can account for all the properties of the molecule, one invokes the concept of resonance. Valence bond theory gives us a model for benzene where each carbon atom makes two sigma bonds with its neighbouring carbon atoms and one with a hydrogen atom. But since carbon is tetravalent, it has the ability to form one more bond. In VB it can form this extra bond with either of the neighbouring carbon atoms, giving rise to the familiar Kekulé ring structure. But this cannot account for all carbon-carbon bond lengths being equal in benzene. A solution is to write the actual wavefunction of the molecule as a linear superposition of the two possible Kekulé structures (or rather the wavefunctions representing these structures), creating a wavefunction that is neither of its components but rather a superposition of them. In benzene both Kekulé structures have equal energy and 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. 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 wavefunctions. This is taken to be the best approximation that can be made to the real structure, though a better one may be made with addition of more structures.
[edit] Molecular orbital (MO) versus valence bond (VB) theory

In molecular orbital theory, the main alternative to valence bond theory, resonance often (but not always) translates to a delocalization of electrons in π orbitals (which are a separate concept from π bonds in VB). For example, in benzene, the MO model gives us 6 π electrons completely delocalized over all 6 carbon atoms, thus contributing something like half-bonds. 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.
[edit] Coefficients

Weighting of the of 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.

[edit] History
The concept 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][12] Linus Pauling used this analogy to introduce his resonance theory in 1928. [13] 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. The alternative term mesomerism popular in German and French publications with the same meaning was introduced by Christopher Ingold in 1938, but did not catch on in the English literature. The current concept of mesomeric effect has taken on a related but different meaning. The double headed arrow was introduced by the German chemist Fritz Arndt who preferred the German phrase zwischenstufe or intermediate stage. In the Soviet Union, resonance theory — especially as developed by Linus 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".[14] Due to confusion with the physical meaning of the word resonance, as no elements actually appear to be resonating, it has been suggested that the term resonance be abandoned in favor of delocalization.[15] Resonance energy would become delocalization energy and a resonance structure becomes a contributing structure. The double headed arrows would be replaced by commas.

[edit] Representations
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 1.5.[16][17]

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.

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.

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

[edit] Reactive intermediates Main article: Reactive intermediate

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

This observation of greater delocalization in less stable molecules is quite general. The excited states of conjugated dienes are stabilised more by conjugation than their ground states, causing them to become organic dyes. A well-studied example of delocalization that does not involve π electrons (hyperconjugation) can be observed in the non-classical ion norbornyl cation. Other examples are diborane and methanium (CH5+). These can be viewed as containing 3-center-2-electron bonds and are represented either by contributing structures involving rearrangement of sigma electrons or by a special notation, a Y that has the three nuclei at its three points. A mesomeric effect is an electron redistribution that occurs via a pi orbital, quite often via conjugated systems. A good example of this effect is seen in the carbonyl group:

The properties of the carbonyl are not properly explained by the classical image (far left), nor by the dipole (centre), which represents the total relocation of the pi electrons; the "real" structure is a combination of these, shown on the far right. There is an inductive effect (indicated by the arrow) in the "real" structure, but it is smaller than the mesomeric effect. This is because the sigma electrons are less readily polarizable than pi electrons, hence they are not shifted as far. The mesomeric effect can be transmitted along conjugated systems if, for example, a carbonyl group is conjugated with a C=C bond:

Therefore, as above, there is an electron deficiency at the carbon next to oxygen, and also at the one indicated in the C3 position. Mesomeric effects are much better transmitted through bonds than the inductive effect;in the C3 position, the effect of the the carbonyl mesomeric influence is still noticeable. The stabilisation that results from delocalisation of charge through a mesomeric effect, can be an important influencing factor in the formation of the ion itself:

Hence we note that phenol is much more acidic than a straight-chain alcohol (i.e. ethanol pKa = 15.9). It is the stabilisation of the negative charge of the phenoxide ion that is a very important factor in the acidity of phenol. Mesomeric (and inductive) effects are permanent effects that are present in the ground state of the molecule. They are therefore present in the physical properties of the molecule.

Resonance effect or mesomeric effect
(1) The effect in which π electrons are transferred from a multiple bond to an atom, or from a multiple bond to a single covalent bond or lone pair (s) of electrons from an atom to the adjacent single covalent bond is called mesomeric effect or simply as M-effect. In case of the compound with conjugated system of double bonds, the mesomeric effect is transmitted through whole of the conjugated system and thus the effect may better be known as conjugative effect. (2) Groups which have the capacity to increase the electron density of the rest of the molecule are said to have +M effect. Such groups possess lone pairs of electrons. Groups

which decrease the electron density of the rest of the molecule by withdrawing electron pairs are said to have –M effect, e.g., (a) The groups which donate electrons to the double bond or to a conjugated system are said to have +M effect or +R effect. +M effect groups :

(b) The groups which withdraw electrons from the double bond or from a conjugated system towards itself due to resonance are said to have –M effect or –R effect. –M effect groups :

(3) The inductive and mesomeric effects, when present together, may act in the same direction or oppose each other. The mesomeric effect is more powerful than the former. For example, in vinyl chloride due to – I effect the chlorine atom should develop a negative charge but on account of mesomeric effect it has positive charge.

Application of resonance effect : It explains, (1) Low reactivity of aryl and vinyl halides, (2) The acidic nature of carboxylic acids, (3) Basic character comparison of ethylamine and aniline, (4) The stability of some free radicals, carbocations and carbanions. Difference between Resonance and Mesomerism : Although both resonance and mesomerism represent the same phenomenon, they differ in the following respect : Resonance involves all types of electron displacements while mesomerism is noticeable only in those cases where a multiple bond is in conjugation with a multiple bond or lone pair of electron. Example :

Both (i) and (ii) are the examples of mesomerism and resonance effect. Let us consider the following example . Such an electron displacement is the example of resonance only (not the mesomerism).

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Steric effects :arise from the fact that each atom within a molecule occupies a certain
amount of space. If atoms are brought too close together, there is an associated cost in energy due to overlapping electron clouds (Pauli or Born repulsion), and this may affect the molecule's preferred shape (conformation) and reactivity.

Contents
[hide]


1 Types of steric effects o 1.1 Steric hindrance o 1.2 Other types of steric effects

    

2 Steric effects vs. electronic effects 3 Significance 4 See also 5 References 6 External links

[edit] Types of steric effects
[edit] Steric hindrance

Steric hindrance or steric resistance occurs when the size of groups within a molecule prevents chemical reactions that are observed in related smaller molecules. Although steric hindrance is sometimes a problem (it prevents SN2 reactions with tertiary substrates from taking place), it can also be a very useful tool, and is often exploited by chemists to change the reactivity pattern of a molecule by stopping unwanted side-reactions (steric protection). Steric hindrance between adjacent groups can also restrict torsional bond angles. However, hyperconjugation has been suggested as an explanation for the preference of the staggered conformation of ethane because the steric hindrance of the small hydrogen atom is far too small. [1] [2] This is the effect responsible for the observed shape of rotaxanes.

Regioselective dimethoxytritylation of the primary 5'-hydroxyl group of thymidine in the presence of a free secondary 3'-hydroxy group as a result of steric hindrance due to the dimethoxytrityl group and the ribose ring (Py = pyridine).[3] [edit] Other types of steric effects

Steric shielding occurs when a charged group on a molecule is seemingly weakened or spatially shielded by less charged (or oppositely charged) atoms, including counterions in solution (Debye shielding). In some cases, for an atom to interact with sterically shielded atoms, it would have to approach from a vicinity where there is less shielding, thus controlling where and from what direction a molecular interaction can take place.

Steric attraction occurs when molecules have shapes or geometries that are optimized for interaction with one another. In these cases molecules will react with each other most often in specific arrangements. Chain crossing: A chain, ring, or a set of rings cannot change from one conformation to another if it would require a chain (or ring - a ring is a cyclic chain) to pass through itself or another chain. This is responsible for the shape of catenanes and molecular knots. Steric repulsions between different parts of molecular system were found of key importance to govern the direction of transition metal mediated transformations and catalysis. Steric effect can even induce a mechanism switch in the catalytic reaction.[4]

[edit] Steric effects vs. electronic effects
The structure, properties, and reactivity of a molecule is dependent on straight forward bonding interactions including covalent bonds, ionic bonds, hydrogen bonds and lesser forms of bonding. This bonding supplies a basic molecular skeleton that is modified by repulsive forces. These repulsive forces include the steric interactions described above. Basic bonding and steric are at times insufficient to explain many structures, properties, and reactivity. Thus steric effects are often contrasted and complemented by electronic effects implying the influence of effects such as induction, conjunction, orbital symmetry, electrostatic interactions, and spin state. There are more esoteric electronic effects but these are among the most important when considering structure and chemical reactivity. A special computational procedure was developed to separate electronic and steric effects of an arbitrary group in the molecule and to reveal their influence on structure and reactivity.[5]

[edit] Significance
Understanding steric effects is critical to chemistry, biochemistry and pharmacology. In chemistry, steric effects are nearly universal and affect the rates and energies of most chemical reactions to varying degrees. In biochemistry, steric effects are often exploited in naturally occurring molecules such as enzymes, where the catalytic site may be buried within a large protein structure. In pharmacology, steric effects determine how and at what rate a drug will interact with its target bio-molecules.

Hyperconjugation: In organic chemistry, hyperconjugation is the interaction of
the electrons in a sigma bond (usually C–H or C–C) with an adjacent empty (or partially filled) non-bonding p-orbital or antibonding π orbital or filled π orbital, to give an extended molecular orbital that increases the stability of the system.[1][2] Only electrons in bonds that are β to the positively charged carbon can stabilize a carbocation by hyperconjugation.

Contents
[hide]
 

  

1 History 2 Applications o 2.1 Effect on chemical properties o 2.2 Hyperconjugation in unsaturated compounds o 2.3 Stabilization of 1,3-butadiyne and 1,3-butadiene o 2.4 Trends in hyperconjugation o 2.5 Hyperconjugation: Gronert vs. Schleyer o 2.6 Rotational barrier of ethane 3 External links 4 See also 5 References

[edit] History
The term was introduced in 1939 by Robert S. Mulliken[3] in the course of his work on UV spectroscopy of conjugated molecules. Mulliken observed that on adding alkyl groups to alkenes the spectra shifted to longer wavelengths. This bathochromic shift is well known in regular conjugated compounds such as butadiene. He was also the first to attribute the lower heat of hydrogenation for these substituted compounds (compared to those without substitution) to hyperconjugation. An effect predating the 1939 hyperconjugation concept is the Baker–Nathan effect reported in 1935.

[edit] Applications
Hyperconjugation can be used for rationalizing a variety of other chemical phenomena, including the anomeric effect, the gauche effect, the rotational barrier of ethane, the beta-silicon effect, the vibrational frequency of exocyclic carbonyl groups, and the relative stability of substituted carbocations and substituted carbon centred radicals. Hyperconjugation is proposed by quantum mechanical modeling to be the correct explanation for the preference of the staggered conformation rather than the old textbook notion of steric hindrance.[4][5]
[edit] Effect on chemical properties

Hyperconjugation affects several properties.[6][7]
1. Bond length: Hyperconjugation is suggested as a key factor in shortening of sigma bonds (σ bonds). For example, the single C–C bonds in 1,3-butadiene and methylacetylene are approximately 1.46 angstrom in length, much less than the value of around 1.54 Å found in saturated hydrocarbons. This is due mainly to hyperconjugation that gives partial double-bond character of the bond.

2. Dipole moments: The large increase in dipole moment of 1,1,1-trichloroethane as compared with chloroform can be attributed to hyperconjugated structures. 3. The heat of formation of molecules with hyperconjugation are greater than sum of their bond energies and the heats of hydrogenation per double bond are less than the heat of hydrogenation of ethylene. 4. Stability of carbocations: (CH3)3C+ > (CH3)2CH+ > (CH3)CH2+ > CH3+ The C–C σ bond adjacent to the cation is free to rotate, and as it does so, the three C–H σ bonds of the methyl group in turn undergoes the stabilization interaction. The more adjacent C-H bonds are, the larger hyperconjugation stabilization is. [edit] Hyperconjugation in unsaturated compounds

Early studies in hyperconjugation were performed by Kistiakowsky et al. Their work, first published in 1937, was intended as a preliminary progress report of thermochemical studies of energy changes during addition reactions of various unsaturated and cyclic compounds. This pioneering work would lead many to investigate the group‘s puzzling findings. Kistiakowsky and fellow researchers collected heats of hydrogenation data during gas-phase reactions of various species containing one double bond. When comparing the monosubstituted alkene compounds propylene, 1-butene, 1-heptene, isopropylethylene, t-butyl ethylene, and neopentylethylene they found that the respective methyl, ethyl, n-amyl, isopropyl, t-butyl, and neopentyl groups are equally effective in stabilizing the adjacent alkene. The overall range of the ΔH values for these compounds was only 0.8 kcal/mol.[8]

A portion of Kistiakowsky‘s work involved a comparison of other unsaturated compounds in the form of CH2=CH(CH2)n-CH=CH2 (n=0,1,2). These experiments revealed an important result; when n=0, there is an effect of conjugation to the molecule where the ΔH value is lowered by 3.5 kcal. This is likened to the addition of two alkyl groups into ethylene. Kistiakowsky also investigated open chain systems, where the largest value of heat liberated was found to be during the addition to a molecule in the 1,4-position. Cyclic molecules proved to be the most problematic, as it was found that the strain of the molecule would have to be considered. The strain of five-membered rings increased with a decrease degree of unsaturation. This was a surprising result that was further investigated in later work with cyclic acid anhydrides and lactones. Cyclic molecules like benzene and its derivatives were also studied, as their behaviors were different from other unsaturated compounds.[8]

Despite the thoroughness of Kistiakowsky‘s work, it was not complete and needed further evidence to back up his findings. His work was a crucial first step to the beginnings of the ideas of hyperconjugation and conjugation effects.
[edit] Stabilization of 1,3-butadiyne and 1,3-butadiene

The conjugation of 1,3-butadiene was first evaluated by Kistiakowsky, a conjugative contribution of 3.5 kcal/mol was found based on the energetic comparison of hydrogenation between conjugated species and unconjugated analogues.[8] Rogers et al., who used the method first applied by Kistiakowsky, reported that the conjugation stabilization of 1,3-butadiyne was zero, as the difference of ΔhydH between first and second hydrogenation was zero. The heats of hydrogenation (ΔhydH) were obtained by computational MP2 quantum chemistry method.[9]

Another group led by Houk et al.,[10] suggested the methods employed by Rogers and Kistiakowsky was inappropriate, because that comparisons of heats of hydrogenation evaluate not only conjugation effects but also other structural and electronic differences. They obtained 70.6 kcal/mol and -70.4 kcal/mol for the first and second hydrogenation respectively by ab initio calculation, which confirmed Rogers‘ data. However, they interpreted the data differently by taking into account the hyperconjugation stabilization. To quantify hyperconjugation effect, they designed the following isodesmic reactions in 1-butyne and 1-butene.

Deleting the hyperconjugative interactions gives virtual states which have energies that are 4.9 and 2.4 kcal/mol higher than those of 1-butyne and 1-butene, respectively. Employment these virtual states results in a 9.6 kcal/mol conjugative stabilization for 1,3-butadiyne and 8.5 kcal/mol for 1,3-butadiene.

[edit] Trends in hyperconjugation

A relatively recent work (2006) by Fernández and Frenking (2006) summarized the trends in hyperconjugation among various groups of acyclic molecules, using energy decomposition analysis or EDA. Fernández and Frenking define this type of analysis as "...a method that uses only the pi orbitals of the interacting fragments in the geometry of the molecule for estimating pi interactions.[11]" For this type of analysis, the formation of bonds between various molecular moieties is a combination of three component terms. ΔEelstat represents what Fernández and Frenking call a molecule‘s ―quasiclassical electrostatic attractions.[11]‖ The second term, ΔEPauli, represents the molecule‘s Pauli repulsion. ΔEorb, the third term, represents stabilizing interactions between orbitals, and is defined as the sum of ΔEpi and ΔEsigma. The total energy of interaction, ΔEint, is the result of the sum of the 3 terms.[11] A group whose ΔEpi values were very thoroughly analyzed were a group of enones that varied in substituent.

Fernández and Frenking reported that the methyl, hydroxyl and amino substituents resulted in a decrease in ΔEpi from the parent 2-propenal. Conversely, halide substituents of increasing atomic mass resulted in increasing ΔEpi. Because both the enone study and Hammett analysis study substituent effects (although in different species), Fernández and Frenking felt that comparing the two to investigate possible trends might yield significant insight into their own results. They observed a linear relationship between the ΔEpi values for the substituted enones and the corresponding Hammett constants. The slope of the graph was found to be -51.67, with a correlation coefficient of -0.97 and a standard deviation of 0.54.[11] Fernández and Frenking conclude from this data that ..."the electronic effects of the substituents R on pi conjugation in homo- and heteroconjugated systems is similar and thus appears to be rather independent of the nature of the conjugating system.".[11][12]
[edit] Hyperconjugation: Gronert vs. Schleyer

Gronert (see Gronert model) [13][14] proposed a 1,3 repulsive interaction, otherwise known as a geminal repulsion in place of hyperconjugation. This model explains differences in bond strengths based on differential steric strain relief as a result of bond cleavage. The key point of Gronert's model is that 1,3 repulsions are the major factor in determining stability of C–C of C– H bonds in alkanes. This broad overarching supposition is based on several already existing assumptions:
1. The heats of formation of alkanes are only determined by 1,2 bonding interactions and 1,3 repulsive interactions.

2. All C–H bonding interactions provide the same stabilization. 3. All C–C bonding interactions provide the same stabilization. 4. The 1,3 repulsive interactions can be grouped into C–C–C, C–C–H, and H–C–H interactions.

Gronert's work is a logical step from work done 50 years previous by Dunitz, Schomaker, Bauld, Wiberg, Bickelhaupt, Ziegler and Schleyer. From the results of these groups, Gronert makes a leap of faith to assume that 1,3 repulsive interactions are not uniform and vary in magnitude based on what groups are involved. Gronert's Method for Evaluating Alkane, Cycloalkane, Alkene and Alkyl radical heats of formation:
∆Hf = nC–CEC–C + nC=CEC=C + nC–HEC–H + nC–C–CEC–C–C + nC–C–HEC–C–H + nH–C–HEH–C–H – f(C,H)

where
n = number of each type of interaction or atom E = stabilization/destabilization per interaction F(C,H) = (170.6 + EC)nC + 52.1nH EC = free parameter (correction term for electron pairing in atomic carbon).

The final term converts to heat of formation from values that are fundamentally atomization energies (170.6 kcal/mole for gaseous carbon and 52.1 kcal/mole for hydrogen atoms). There are several important justifications for Gronert's model:
1. Significant geminal repulsion is already expected because groups are separated by less than the combination of their van der Waals radii and there are no bonding interactions. Computational methods also agree that they are important and of the proper magnitude. 2. It ís already well-accepted that 1,3 repulsive interactions are important in determining structure. 3. Branching has a strong effect on stabilities of alkanes, not just the BDE. No current evidence that differences in bond strengths are only controlled by factors exclusive to the resulting radical. His method addresses the stability of the alkane and alkyl radical. 4. Model depends on interactions observed in many systems and affects both structure and reactivity. This is based on the theory that close-range, nonbonded interactions are repulsive, i.e. Steric strain.

Gronert claims that his model successfully reproduces accepted data without invoking hyperconjugation and can perhaps explain well-established trends. However, his analysis involves geminal repulsion absolutely replacing hyperconjugation as a reasonable alternative explanation.[clarification needed] Schleyer's model has several marked differences from Gronert's. He uses a new isodesmic additivity design that in his view faithfully reproduces heats of formation for many alkanes,

alkenes, alkynes and alkyl radicals. All 1,3 interactions are stabilizing so they support branching and hyperconjugation. All adjustable parameters originate from assumption that the magnitude of stabilizations effects at a specific carbon are eased when more than one substituent contributes:
∆Hf = base – 2.15n(CH2) – 1,3CCC branching attraction – hyperconjugation

Schleyer notes several advantages of his approach in comparison to Gronert's:
1. Gronert's derivation method arbitrarily set some parameters and adjusted the others as best-fit averages of experimental hydrocarbon heats of formation. 2. Gronert's derived C–C and C–H bond energy values are higher than those accepted in the literature. 3. Gronert uses 7 adjustable parameters whereas Schleyer uses only 4. Four is the minimum chemically plausible number of parameters, and the added flexibility of additional terms is not necessarily an improvement of general theory. 4. Schleyer's single attractive geminal term is sufficient to reproduce data satisfactorily. 5. Well-established theories of branching, hyperconjugation and attenuation.[clarification needed] 6. Schleyer's method depends only on energetic relationships between the simplest hydrocarbon molecules. [edit] Rotational barrier of ethane

An instance where hyperconjugation may be overlooked as a possible chemical explanation is in rationalizing the rotational barrier of ethane. It had been accepted as early as the 1930s that the staggered conformations of ethane were more stable than the eclipsed conformation. Wilson had proven that the energy barrier between any pair of eclipsed and staggered conformations was approximately 3 kcal/mol, and the generally accepted rationale for this was the unfavorable steric interactions between hydrogen atoms.

Staggered (left) and Eclipsed (right)

In their 2001 paper, however, Pophristic and Goodman[4] revealed that this explanation may be too simplistic.[15] Goodman focused on three principal physical factors: hyperconjugative interactions, exchange repulsion defined by the Pauli exclusion principle, as well as electrostatic interactions (Coulomb interactions). By comparing a traditional ethane molecule and a hypothetical ethane molecule with all exchange repulsions removed, potential curves were

prepared by plotting torsional angle versus energy for each molecule. The analysis of the curves determined that the staggered conformation had no connection to the amount of electrostatic repulsions within the molecule. These results demonstrate that Coulombic forces do not explain the favored staggered conformations, despite the fact that central bond stretching decreases electrostatic interactions.[4] Goodman also conducted studies to determine the contribution of vicinal (between two methyl groups) vs. geminal (between the atoms in a single methyl group) interactions to hyperconjugation. In separate experiments, the geminal and vicinal interactions were removed, and the most stable conformer for each interaction was deduced.[4]
Calculated torsional angle of ethane with deleted hyperconjugative effects Deleted interaction None All hyperconjugation Vicinal hyperconjugation Geminal hyperconjugation Torsional angle Corresponding conformer 60° 0° 0° 60° Staggered Eclipsed Eclipsed Staggered

From these experiments, it can be concluded that hyperconjugative effects delocalize charge and stabilize the molecule. Further, it is the vicinal hyperconjugative effects that keep the molecule in the staggered conformation.[4] Thanks to this work, the following model of the stabilization of the staggered conformation of ethane is now more accepted:

Hyperconjugation can also explain several other phenomena whose explanations may also not be as intuitive as that for the rotational barrier of ethane.[15] One such example is the explanations for certain Lewis structures. The Lewis structure for an ammonium ion indicates a positive charge on the nitrogen atom. In reality, however, the hydrogens are more electropositive than is

nitrogen, and thus are the actual carriers of the positive charge. We know this intuitively because bases remove the protons as opposed to the nitrogen atom.[15] It should be noted that the matter of the rotational barrier of ethane is not settled within the scientific community. An analysis within quantitative molecular orbital theory shows that 2orbital-4-electron (steric) repulsions are dominant over hyperconjugation.[16] A valence bond theory study also emphasizes the importance of steric effects.[17]

[edit] External links
 

Further reading Advanced hyperconjugation

[4] Hyper conjugation or no bond resonance or Baker-Nathan Effect

Stability order of alkenes

Prob. Write down the reactivity order towards electrophilic substitution reaction (good Puzzle)

Toluene is maximum reactive due to maximum hyperconjugation which develop maximum negative charge on the benzene ring and accelerates attack on benzene nucleus.

In elimination reactions like dehydration and dehydrohalogenation product are formed according to saytzeff's rule.

Effects of hyperconjugation (a) Bond Length : Like resonance, hyperconjugation also affects bond lengths because during the process the single bond in compound acquires some double bonded character and vice-versa. e.g. C — C bond length in propene is as compared to in Ehtylene.

(b) Dipole moment : Since hyperconjugation causes these charge developments, it also affects the dipole moment of the molecule. (c) Stability of carbonium ions is Tertiary > Secondary > Primary Above order of stability can be explained by hyperconjugation. In general greater the number of hydrogen atoms attached to , the more hyperconjugative forms are formed and thus greater is the stability of carbonium ions.

(d) Stability of Free radicals : Stability of Free radicals can also be explained as that of carbonium ion

(e) Orientation influence of methyl group : The o,p–directing influence of the methyl group in methyl benzene is attributed partly to inductive and partly of hyperconjugation effect.

The role of hyperconjugation in o,p–directing influence of methyl group is evidenced by the part that nitration of p-iso propyl toluene and p-tert-butyl toluene form the product in which —NO2 group is introduced in the ortho position with respect to methyl group and not to isopropyl or t- butyl group although the latter groups are more electron donating than methyl groups.

i.e., The substitution takes place contrary to inductive effect. Actually this constitutes an example where hyperconjugation over powers inductive effect.

[5] Electromeric effect It is the complete transfer of of a multiple bond towards one of the bonded atoms at the demand of an attacking reagent. The transfer of electrons takes place towards the more electronegative of the two bonded atoms. For example, when an addition reaction takes place at a carbonyl group (>C = O), the double bond are shifted at O-atom because it is more electronegative than carbon : of the

In Ethylene (ethene) molecule, shifting of p electrons may take place at any of the doubly bonded Catoms since both the atoms are identical :

However, in Propene the shifting of electrons takes place at carbon atom no.1 because of the +I effect of methyl group :

Electromeric effect (shifting of p electrons) is a temporary effect and takes place only at the demand of an attacking reagent during the course of a chemical reaction. In short it is termed as the E effect. If the I-effect and E - effect oppose each other then usually the E - effect predominates, i.e.,

Application of electromeric effect : The mechanisms of several organic reactions particularly the addition reactions, are explained by the help of electromeric effect. [6] Steric effects Whenever a chemical reaction between two compounds takes place, directly or indirectly it results in the formation of bond between the atoms of these two compounds. The bonding atoms of the two reacting compounds are, in fact, the active centres. These will react (form the bond) with one another only if they come with in the range of each other, i.e., within the attraction of each other. If we surround one of them with such mass of other atoms or groups that the other reacting atom is unable to force its way in, the reaction may not at all take place or may take place only slowly. Such a hindrance due to spatial crowding (crowding a space) is called steric hindrance. However, the spatial crowding may not always hinder a reaction, sometimes it may facilitate a reaction. Therefore, the term steric effect is better than the term steric hindrance.The phenomenon of steric effect was first identified by Hofmann in 1872. It may have sufficient influence on the physical and chemical properties of a molecule and may be defined as modification in molecular properties resulting from a spatial crowding of a reacting

atom in a molecule. A good example of steric effect is discussed here: The tertiary amine with the name Tri methylamine reacts with methyl iodide to form a quaternary salt tetra methyl ammonium iodide, but if the alkyl groups in tertiary amine are large, it does not react with methyliodide and so, does not form the quaternary ammonium salt :

Another good example is that of esterification between a carboxylic acid and an alcohol. Bulkier the alkyl group in acid or alcohol, slower is esterification. [7] Reactivity of various chemical reactions 1. Electrophilic addition reaction

Alkene is more reactive than alkyne towards addition reaction because of thick of alkyne. There is no addition reaction in case of benzene because resonance occurs in the molecule which stabilises the cloud

in case

2. Nucleophilic addition reaction As the - I effect or - M effect increases, + ve charge on > C = O group increases so that speed of primary

attack of nucleophile on carbon also increases.

3. Nucleophilic Substitution reaction

“Weaker the base better the leaving group” and conjugate base of strong acid always weak.

Resonance or Back bonding of electrons in vacant orbital always creates double bond character in between C – X due to this, chemical reactivity decreases. C6H5ClCH2 = CH–Cl,CH3–CH=CH–Cl and CH3–CCl=CH2 these all are less reactive halide due to resonance. While C6H5CH2Cl and Cl – CH2 – CH = CH2 are very reactive due to absence of resonance and high stability of intermediate carbocation. As the +ve charge on carbon attached with halide, increases speed of attacking nature of nucleophile also increases.

4. Electrophilic substitution reaction Chemical reactivity of aromatic compounds is decided by activating or deactivating series of group. Activating groups are more reactive towards electrophilic substitution reaction because they develop – ve charge on the benzene ring. These are o + p directing in nature. While deactivating groups are less reactive because they develop +ve charge on benzene nucleus and these are meta directing in nature.

“Deactivating groups are always less reactive than activating groups.”

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Keto-enol tautomerism. Left the keto form; right the enol.

In organic chemistry, keto-enol tautomerism refers to a chemical equilibrium between a keto form (a ketone or an aldehyde) and an enol (an alcohol). The enol and keto forms are said to be tautomers of each other. The interconversion of the two forms involves the movement of a proton and the shifting of bonding electrons; hence, the isomerism qualifies as tautomerism. A compound containing a carbonyl group (C=O) is normally in rapid equilibrium with an enol tautomer, which contains a pair of doubly bonded carbon atoms adjacent to a hydroxyl (−OH) group, C=C-OH. The keto form predominates at equilibrium for most ketones. Nonetheless, the enol form is important for some reactions. Furthermore, the deprotonated intermediate in the interconversion of the two forms, referred to as an enolate anion, is important in carbonyl chemistry, in large part because it is a strong nucleophile. Normally, the keto-enol tautomerization chemical equilibrium is highly thermodynamically driven, and at room temperature the equilibrium heavily favors the formation of the keto form. A classic example for favoring the keto form can be seen in the equilibrium between vinyl alcohol and acetaldehyde (K = [enol]/[keto] ≈ 3 x10−7). However, it is reported that in the case of vinyl alcohol, formation of a stabilized enol form can be accomplished by controlling the water concentration in the system and utilizing the kinetic favorability of the deuterium produced

kinetic isotope effect (kH+/kD+ = 4.75, kH2O/kD2O = 12). Deuterium stabilization can be accomplished through hydrolysis of a ketene precursor in the presence of a slight stoichiometric excess of heavy water (D2O). Studies show that the tautomerization process is significantly inhibited at ambient temperatures ( kt ≈ 10−6 M/s), and the half life of the enol form can easily be increased to t1/2 = 42 minutes for first order hydrolysis kinetics.[1]

Contents
[hide]
   

1 Mechanism 2 Erlenmeyer rule 3 Significance in biochemistry o 3.1 DNA 4 References

[edit] Mechanism
The conversion of an acid catalyzed enol to the keto form proceeds by a two step mechanism in an aqueous acidic solution. First, the exposed electrons of the C=C double bond of the enol are donated to a hydronium ion (H3O+). This addition follows Markovnikov's rule, thus the proton is added to the carbon with more hydrogens. This is a concerted step with the oxygen in the hydroxyl group donating electrons to produce the eventual carbonyl group. Second, the oxygen in a water molecule donates electrons to the hydrogen in the hydroxyl group, thus relieving the positive charge on the electronegative oxygen atom.

http://en.wikipedia.org/wiki/File:Keto-enol.svg

[edit] Erlenmeyer rule
One of the early investigators into keto-enol tautomerism was Richard August Carl Emil Erlenmeyer. His Erlenmeyer rule (developed in 1880) states that all alcohols in which the hydroxyl group is attached directly to a double-bonded carbon atom become aldehydes or ketones. This occurs because the keto form is, in general, more stable than its enol tautomer. As the lower energy form, the keto form is favored at equilibrium.

[edit] Significance in biochemistry
Keto-enol tautomerism is important in several areas of biochemistry. The high phosphatetransfer potential of phosphoenolpyruvate results from the fact that the phosphorylated compound is "trapped" in the less stable enol form, whereas after dephosphorylation it can assume the keto form. Rare enol tautomers of the bases guanine and thymine can lead to mutation because of their altered base-pairing properties[citation needed]. In certain aromatic compounds such as phenols, the enol is important due to the aromatic character of the enol but not the keto form. Melting the naphthalene derivative 1,4dihydroxynaphthalene 1 at 200 °C results in a 2:1 mixture with the keto form 2. Heating the keto form in benzene at 120°C for three days also affords a mixture (1:1 with first-order reaction kinetics). The keto product is kinetically stable and reverts back to the enol in presence of a base. The keto form can be obtained in a pure form by stirring the keto form in trifluoroacetic acid and toluene (1:9 ratio) followed recrystallisation from isopropyl ether.[2]

When the enol form is complexed with chromium tricarbonyl, complete conversion to the keto form accelerated and occurs even at room temperature in benzene.
[edit] DNA

In deoxyribonucleic acids(DNA), the nucleotide bases are in keto form. However, James Watson and Francis Crick first believed them to be in the enol tautomeric form, delaying the solution of the structure for several months.[3]

"Aromatic compound" redirects here. For meanings related to odor, see aroma compound.

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

In organic chemistry, the structures of some rings of atoms are unexpectedly stable. Aromaticity is a chemical property in which a conjugated ring of unsaturated bonds, lone pairs, or empty orbitals exhibit a stabilization stronger than would be expected by the stabilization of conjugation alone. It can also be considered a manifestation of cyclic delocalization and of resonance.[1][2][3] This is usually considered to be because electrons are free to cycle around circular arrangements of atoms which are alternately single- and double-bonded to one another. These bonds may be seen as a hybrid of a single bond and a double bond, each bond in the ring identical to every other. 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 Kekulé (see History section below). The model for benzene consists of two resonance forms, which corresponds to the double and single bonds superimposing to give rise to six one-and-a-half bonds. Benzene is a more stable molecule than would be expected without accounting for charge delocalization.

Contents
[hide]
    



1 Theory 2 History 3 Characteristics of aromatic (aryl) compounds 4 Importance of aromatic compounds 5 Types of aromatic compounds o 5.1 Heterocyclics o 5.2 Polycyclics o 5.3 Substituted aromatics o 5.4 Atypical aromatic compounds 6 See also



7 References

[edit] Theory

As is standard for resonance diagrams, a double-headed arrow is used to indicate that the 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, which can be seen at right. A C=C bond is shorter than a C−C bond, but benzene is perfectly hexagonal—all six carbon-carbon bonds have the same length, intermediate between that of a single and that of a double bond. 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 with electrons 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 has π symmetry.

[edit] History
The first known use of the word "aromatic" as a chemical term—namely, to apply to compounds that contain the phenyl radical—occurs in an article by August Wilhelm Hofmann in 1855.[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. It is the case, however, that 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 indices than many aliphatic compounds, and Hofmann may not have been making a distinction between the two categories. The cyclohexatriene structure for benzene was first proposed by August Kekulé in 1865. Over the next few decades, most chemists readily accepted this structure, since it accounted for most of the known isomeric relationships of aromatic chemistry. However, it was always puzzling that this purportedly highly unsaturated molecule was so unreactive toward addition reactions. The discoverer of the electron J. J. Thomson, between 1897 and 1906 placed three equivalent electrons between each carbon atom in benzene. An explanation for the exceptional stability of benzene is conventionally attributed to Sir Robert Robinson, who was apparently the first (in 1925)[5] 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,[6] to Henry Edward Armstrong, who in 1890, in an article entitled The structure of cycloid hydrocarbons, 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.[7] Here, Armstrong is describing at least four modern concepts. First, his "affinity" is better known nowadays as the electron, which was only to be discovered 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 Eric 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.

[edit] 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 number of π electrons, where n=0, 1, 2, 3, and so on. This is known as Hückel's Rule.

Whereas benzene is aromatic (6 electrons, from 3 double bonds), cyclobutadiene is not, 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 C-H bond on the other positions). There are 6 π electrons, so furan is aromatic. Aromatic molecules typically display enhanced chemical stability, compared to similar nonaromatic molecules. A molecule that can be aromatic will tend to alter its electronic or conformational structure to be in this situation. This extra 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 carboncarbon double bonds. 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. The circulating π electrons in an aromatic molecule produce ring currents that oppose the applied magnetic field in NMR.[8] 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 up-field.

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, destabilized. Molecules that could be antiaromatic will tend to alter their electronic or conformational structure to avoid this situation, thereby becoming non-aromatic. For example, cyclooctatetraene (COT) distorts itself out of planarity, breaking π overlap between adjacent double bonds. Relatively recently, cyclobutadiene was discovered to adopt an asymmetric, rectangular configuration in which single and double bonds indeed alternate; there is no resonance and the single bonds are markedly longer than the double bonds, reducing unfavorable p-orbital overlap. Hence, cyclobutadiene is non-aromatic; the strain of the asymmetric configuration outweighs the anti-aromatic destabilization that would afflict the symmetric, square configuration.

[edit] Importance of aromatic compounds
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. Other aromatic compounds play key roles in the biochemistry of all living things. 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. As well as that, the molecule heme contains an aromatic system with 22 π electrons. Chlorophyll also has a similar aromatic system.

[edit] Types of aromatic compounds
The overwhelming majority of aromatic compounds are compounds of carbon, but they need not be hydrocarbons.
[edit] Heterocyclics

In heterocyclic aromatics (heteroaromats), 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, imidazole, pyrazole, oxazole, thiophene, and their benzannulated analogs (benzimidazole, for example).

[edit] 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.
[edit] Substituted aromatics

Many chemical compounds are aromatic rings with other things attached. Examples include trinitrotoluene (TNT), acetylsalicylic acid (aspirin), paracetamol, and the nucleotides of DNA.
[edit] Atypical aromatic compounds

Aromaticity is found in ions as well: the cyclopropenyl cation (2e system), the cyclopentadienyl anion (6e system), the tropylium ion (6e) and the cyclooctatetraene dianion (10e). Aromatic properties have been attributed to non-benzenoid compounds such as tropone. Aromatic properties are tested to the limit in a class of compounds called cyclophanes. A special case of aromaticity is found in homoaromaticity where conjugation is interrupted by a single sp³ hybridized carbon atom. When carbon in benzene is replaced by other elements in borabenzene, silabenzene, germanabenzene, stannabenzene, phosphorine or pyrylium salts the aromaticity is still retained. Aromaticity also occurs in compounds that are not carbon-based at all. Inorganic 6 membered ring compounds analogous to benzene have been synthesized. Silicazine (Si6H6) and borazine (B3N3H6) are structurally analogous to benzene, with the carbon atoms replaced by another element or elements. In borazine, the boron and nitrogen atoms alternate around the ring. Metal aromaticity is believed to exist in certain metal clusters of aluminium. Möbius aromaticity occurs when a cyclic system of molecular orbitals, formed from pπ atomic orbitals and populated in a closed shell by 4n (n is an integer) electrons, is given a single half-twist to correspond to a Möbius strip. Because the twist can be left-handed or right-handed, the resulting Möbius aromatics are dissymmetric or chiral. Up to now there is no doubtless proof that a Möbius aromatic molecule was synthesized.[9][10] Aromatics with two half-twists corresponding to the paradromic topologies, first suggested by Johann Listing, have been proposed by Rzepa in 2005.[11] In carbo-benzene the ring bonds are extended with alkyne and allene groups.

[edit] See also
    

Aromatic hydrocarbons Aromatic amine PAH Simple aromatic ring Pi interaction

"Aromatic compound" redirects here. For meanings related to odor, see aroma compound.

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

In organic chemistry, the structures of some rings of atoms are unexpectedly stable. Aromaticity is a chemical property in which a conjugated ring of unsaturated bonds, lone pairs, or empty orbitals exhibit a stabilization stronger than would be expected by the stabilization of conjugation alone. It can also be considered a manifestation of cyclic delocalization and of resonance.[1][2][3] This is usually considered to be because electrons are free to cycle around circular arrangements of atoms which are alternately single- and double-bonded to one another. These bonds may be seen as a hybrid of a single bond and a double bond, each bond in the ring identical to every other. 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 Kekulé (see History section below). The model for benzene consists of two resonance forms, which corresponds to the double and single bonds superimposing to give rise to six one-and-a-half bonds. Benzene is a more stable molecule than would be expected without accounting for charge delocalization.

Contents
[hide]
    

1 Theory 2 History 3 Characteristics of aromatic (aryl) compounds 4 Importance of aromatic compounds 5 Types of aromatic compounds o 5.1 Heterocyclics o 5.2 Polycyclics o 5.3 Substituted aromatics o 5.4 Atypical aromatic compounds

 

6 See also 7 References

[edit] Theory

As is standard for resonance diagrams, a double-headed arrow is used to indicate that the 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, which can be seen at right. A C=C bond is shorter than a C−C bond, but benzene is perfectly hexagonal—all six carbon-carbon bonds have the same length, intermediate between that of a single and that of a double bond. 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 with electrons 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 has π symmetry.

[edit] History
The first known use of the word "aromatic" as a chemical term—namely, to apply to compounds that contain the phenyl radical—occurs in an article by August Wilhelm Hofmann in 1855.[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. It is the case, however, that 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 indices than many aliphatic compounds, and Hofmann may not have been making a distinction between the two categories. The cyclohexatriene structure for benzene was first proposed by August Kekulé in 1865. Over the next few decades, most chemists readily accepted this structure, since it accounted for most of the known isomeric relationships of aromatic chemistry. However, it was always puzzling that this purportedly highly unsaturated molecule was so unreactive toward addition reactions. The discoverer of the electron J. J. Thomson, between 1897 and 1906 placed three equivalent electrons between each carbon atom in benzene. An explanation for the exceptional stability of benzene is conventionally attributed to Sir Robert Robinson, who was apparently the first (in 1925)[5] 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,[6] to Henry Edward Armstrong, who in 1890, in an article entitled The structure of cycloid hydrocarbons, 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.[7] Here, Armstrong is describing at least four modern concepts. First, his "affinity" is better known nowadays as the electron, which was only to be discovered 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 Eric 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.

[edit] 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 number of π electrons, where n=0, 1, 2, 3, and so on. This is known as Hückel's Rule.

Whereas benzene is aromatic (6 electrons, from 3 double bonds), cyclobutadiene is not, 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 C-H bond on the other positions). There are 6 π electrons, so furan is aromatic. Aromatic molecules typically display enhanced chemical stability, compared to similar nonaromatic molecules. A molecule that can be aromatic will tend to alter its electronic or conformational structure to be in this situation. This extra 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 carboncarbon double bonds. 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. The circulating π electrons in an aromatic molecule produce ring currents that oppose the applied magnetic field in NMR.[8] 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 up-field.

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, destabilized. Molecules that could be antiaromatic will tend to alter their electronic or conformational structure to avoid this situation, thereby becoming non-aromatic. For example, cyclooctatetraene (COT) distorts itself out of planarity, breaking π overlap between adjacent double bonds. Relatively recently, cyclobutadiene was discovered to adopt an asymmetric, rectangular configuration in which single and double bonds indeed alternate; there is no resonance and the single bonds are markedly longer than the double bonds, reducing unfavorable p-orbital overlap. Hence, cyclobutadiene is non-aromatic; the strain of the asymmetric configuration outweighs the anti-aromatic destabilization that would afflict the symmetric, square configuration.

[edit] Importance of aromatic compounds
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. Other aromatic compounds play key roles in the biochemistry of all living things. 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. As well as that, the molecule heme contains an aromatic system with 22 π electrons. Chlorophyll also has a similar aromatic system.

[edit] Types of aromatic compounds
The overwhelming majority of aromatic compounds are compounds of carbon, but they need not be hydrocarbons.
[edit] Heterocyclics

In heterocyclic aromatics (heteroaromats), 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, imidazole, pyrazole, oxazole, thiophene, and their benzannulated analogs (benzimidazole, for example).

[edit] 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.
[edit] Substituted aromatics

Many chemical compounds are aromatic rings with other things attached. Examples include trinitrotoluene (TNT), acetylsalicylic acid (aspirin), paracetamol, and the nucleotides of DNA.
[edit] Atypical aromatic compounds

Aromaticity is found in ions as well: the cyclopropenyl cation (2e system), the cyclopentadienyl anion (6e system), the tropylium ion (6e) and the cyclooctatetraene dianion (10e). Aromatic properties have been attributed to non-benzenoid compounds such as tropone. Aromatic properties are tested to the limit in a class of compounds called cyclophanes. A special case of aromaticity is found in homoaromaticity where conjugation is interrupted by a single sp³ hybridized carbon atom. When carbon in benzene is replaced by other elements in borabenzene, silabenzene, germanabenzene, stannabenzene, phosphorine or pyrylium salts the aromaticity is still retained. Aromaticity also occurs in compounds that are not carbon-based at all. Inorganic 6 membered ring compounds analogous to benzene have been synthesized. Silicazine (Si6H6) and borazine (B3N3H6) are structurally analogous to benzene, with the carbon atoms replaced by another element or elements. In borazine, the boron and nitrogen atoms alternate around the ring. Metal aromaticity is believed to exist in certain metal clusters of aluminium. Möbius aromaticity occurs when a cyclic system of molecular orbitals, formed from pπ atomic orbitals and populated in a closed shell by 4n (n is an integer) electrons, is given a single half-twist to correspond to a Möbius strip. Because the twist can be left-handed or right-handed, the resulting Möbius aromatics are dissymmetric or chiral. Up to now there is no doubtless proof that a Möbius aromatic molecule was synthesized.[9][10] Aromatics with two half-twists corresponding to the paradromic topologies, first suggested by Johann Listing, have been proposed by Rzepa in 2005.[11] In carbo-benzene the ring bonds are extended with alkyne and allene groups.

[edit] See also
"Aromatic compound" redirects here. For meanings related to odor, see aroma compound.

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

In organic chemistry, the structures of some rings of atoms are unexpectedly stable. Aromaticity is a chemical property in which a conjugated ring of unsaturated bonds, lone pairs, or empty orbitals exhibit a stabilization stronger than would be expected by the stabilization of conjugation alone. It can also be considered a manifestation of cyclic delocalization and of resonance.[1][2][3] This is usually considered to be because electrons are free to cycle around circular arrangements of atoms which are alternately single- and double-bonded to one another. These bonds may be seen as a hybrid of a single bond and a double bond, each bond in the ring identical to every other. 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 Kekulé (see History section below). The model for benzene consists of two resonance forms, which corresponds to the double and single bonds superimposing to give rise to six one-and-a-half bonds. Benzene is a more stable molecule than would be expected without accounting for charge delocalization.

Contents
[hide]
    



1 Theory 2 History 3 Characteristics of aromatic (aryl) compounds 4 Importance of aromatic compounds 5 Types of aromatic compounds o 5.1 Heterocyclics o 5.2 Polycyclics o 5.3 Substituted aromatics o 5.4 Atypical aromatic compounds 6 See also



7 References

[edit] Theory

As is standard for resonance diagrams, a double-headed arrow is used to indicate that the 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, which can be seen at right. A C=C bond is shorter than a C−C bond, but benzene is perfectly hexagonal—all six carbon-carbon bonds have the same length, intermediate between that of a single and that of a double bond. 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 with electrons 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 has π symmetry.

[edit] History
The first known use of the word "aromatic" as a chemical term—namely, to apply to compounds that contain the phenyl radical—occurs in an article by August Wilhelm Hofmann in 1855.[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. It is the case, however, that 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 indices than many aliphatic compounds, and Hofmann may not have been making a distinction between the two categories. The cyclohexatriene structure for benzene was first proposed by August Kekulé in 1865. Over the next few decades, most chemists readily accepted this structure, since it accounted for most of the known isomeric relationships of aromatic chemistry. However, it was always puzzling that this purportedly highly unsaturated molecule was so unreactive toward addition reactions. The discoverer of the electron J. J. Thomson, between 1897 and 1906 placed three equivalent electrons between each carbon atom in benzene. An explanation for the exceptional stability of benzene is conventionally attributed to Sir Robert Robinson, who was apparently the first (in 1925)[5] 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,[6] to Henry Edward Armstrong, who in 1890, in an article entitled The structure of cycloid hydrocarbons, 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.[7] Here, Armstrong is describing at least four modern concepts. First, his "affinity" is better known nowadays as the electron, which was only to be discovered 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 Eric 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.

[edit] 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 number of π electrons, where n=0, 1, 2, 3, and so on. This is known as Hückel's Rule.

Whereas benzene is aromatic (6 electrons, from 3 double bonds), cyclobutadiene is not, 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 C-H bond on the other positions). There are 6 π electrons, so furan is aromatic. Aromatic molecules typically display enhanced chemical stability, compared to similar nonaromatic molecules. A molecule that can be aromatic will tend to alter its electronic or conformational structure to be in this situation. This extra 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 carboncarbon double bonds. 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. The circulating π electrons in an aromatic molecule produce ring currents that oppose the applied magnetic field in NMR.[8] 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 up-field.

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, destabilized. Molecules that could be antiaromatic will tend to alter their electronic or conformational structure to avoid this situation, thereby becoming non-aromatic. For example, cyclooctatetraene (COT) distorts itself out of planarity, breaking π overlap between adjacent double bonds. Relatively recently, cyclobutadiene was discovered to adopt an asymmetric, rectangular configuration in which single and double bonds indeed alternate; there is no resonance and the single bonds are markedly longer than the double bonds, reducing unfavorable p-orbital overlap. Hence, cyclobutadiene is non-aromatic; the strain of the asymmetric configuration outweighs the anti-aromatic destabilization that would afflict the symmetric, square configuration.

[edit] Importance of aromatic compounds
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. Other aromatic compounds play key roles in the biochemistry of all living things. 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. As well as that, the molecule heme contains an aromatic system with 22 π electrons. Chlorophyll also has a similar aromatic system.

[edit] Types of aromatic compounds
The overwhelming majority of aromatic compounds are compounds of carbon, but they need not be hydrocarbons.
[edit] Heterocyclics

In heterocyclic aromatics (heteroaromats), 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, imidazole, pyrazole, oxazole, thiophene, and their benzannulated analogs (benzimidazole, for example).

[edit] 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.
[edit] Substituted aromatics

Many chemical compounds are aromatic rings with other things attached. Examples include trinitrotoluene (TNT), acetylsalicylic acid (aspirin), paracetamol, and the nucleotides of DNA.
[edit] Atypical aromatic compounds

Aromaticity is found in ions as well: the cyclopropenyl cation (2e system), the cyclopentadienyl anion (6e system), the tropylium ion (6e) and the cyclooctatetraene dianion (10e). Aromatic properties have been attributed to non-benzenoid compounds such as tropone. Aromatic properties are tested to the limit in a class of compounds called cyclophanes. A special case of aromaticity is found in homoaromaticity where conjugation is interrupted by a single sp³ hybridized carbon atom. When carbon in benzene is replaced by other elements in borabenzene, silabenzene, germanabenzene, stannabenzene, phosphorine or pyrylium salts the aromaticity is still retained. Aromaticity also occurs in compounds that are not carbon-based at all. Inorganic 6 membered ring compounds analogous to benzene have been synthesized. Silicazine (Si6H6) and borazine (B3N3H6) are structurally analogous to benzene, with the carbon atoms replaced by another element or elements. In borazine, the boron and nitrogen atoms alternate around the ring. Metal aromaticity is believed to exist in certain metal clusters of aluminium. Möbius aromaticity occurs when a cyclic system of molecular orbitals, formed from pπ atomic orbitals and populated in a closed shell by 4n (n is an integer) electrons, is given a single half-twist to correspond to a Möbius strip. Because the twist can be left-handed or right-handed, the resulting Möbius aromatics are dissymmetric or chiral. Up to now there is no doubtless proof that a Möbius aromatic molecule was synthesized.[9][10] Aromatics with two half-twists corresponding to the paradromic topologies, first suggested by Johann Listing, have been proposed by Rzepa in 2005.[11] In carbo-benzene the ring bonds are extended with alkyne and allene groups.