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Acids, bases & aromaticity

  1. Acids, Bases and Aromaticity Dr. G. Krishnaswamy Faculty DOS & R in Organic Chemistry Tumkur University Tumakuru
  2. Acids and Bases Produce H+ (as H3O+) ions in water Produce a negative ion (-) too Taste sour Corrode metals React with bases to form salts and water Neutral Substance : Water! General properties of Acids
  3. General properties of Bases Produce OH- ions in water Taste bitter, chalky Are electrolytes Feel soapy, slippery React with acids to form salts and water Neutral Substance : Water!
  4. Definitions of acids and bases: • Acids – produce H+ • Bases - produce OH- • Acids – donate H+ • Bases – accept H+ • Acids – accept e- pair • Bases – donate e- pair Arrehenius Concept Only in water Bronsted-Lowry Any solvent Lewis used in organic chemistry, wider range of substances
  5. Examples of Brønsted–Lowry acids and bases
  6. 6 Reactions of Brønsted-Lowry Acids and Bases • A Brønsted-Lowry acid base reaction results in the transfer of a proton from an acid to a base. • In an acid-base reaction, one bond is broken, and another one is formed. • The electron pair of the base B: forms a new bond to the proton of the acid. • The acid H—A loses a proton, leaving the electron pair in the H—A bond on A.
  7. Each acid has a conjugate base and each base has a conjugate acid. These conjugate pairs only differ by a proton. O H H H Br+ H O H H + Br Base Acid Conjugate Acid C. Base
  8. There are two ways to predict when a proton transfer reaction will occur: (1)Quantitative approach (comparing pKa values) (2)Qualitative approach (analyzing the structures of the acids) Using pKa Values to Compare Acidity • Acid strength is the tendency of an acid to donate a proton. • Acidity is measured by an equilibrium constant. • When a Brønsted-Lowry acid H—A is dissolved in water, an acid- base reaction occurs, and an equilibrium constant can be written for the reaction.
  9. 9 Because the concentration of the solvent H2O is essentially constant, the equation can be rearranged and a new equilibrium constant, called the acidity constant, Ka, can be defined. It is generally more convenient when describing acid strength to use “pKa” values than Ka values.
  10. Selected pKa Values
  11. The principle is this: The stronger the acid, the weaker will be its conjugate base. We can, therefore, relate the strength of a base to the pKa of its conjugate acid. The larger the pKa of the conjugate acid, the stronger is the base. Predicting the Strength of Bases
  12. Using the chart of pKa values, we can also predict the position of equilibrium for any acid-base reaction. The equilibrium will always favor formation of the weaker acid (higher pKa value). For example, consider the following acid-base reaction: Using pKa Values to Predict the Position of Equilibrium The equilibrium for this reaction will lean to the right side, favoring formation of the weaker acid.
  13. Qualitative approach (analyzing the structures of the acids) comparisons by analyzing and comparing their structures and without the use of pKa values. In order to compare acids without the use of pKa values, we must look at the conjugate base of each acid: Conjugate Base Stability If A- is very stable (weak base), then HA must be a strong acid. If, on the other hand, A- is very unstable (strong base), then HA must be a weak acid.
  14. Consider the deprotonation of HCl: Chlorine is an electronegative atom, and it can therefore stabilize a negative charge. The chloride ion (Cl-) is in fact very stable, and therefore, HCl is a strong acid. HCl can serve as a proton donor because the conjugate base left behind is stabilized. The following discussion will develop a methodical approach for comparing negative charge stability. Specifically, there are four factors to consider: (1)Bond strength and which atom is the charge on (2)resonance (3)Induction and (4)orbitals Factors Affecting the Stability of Negative Charges/Acidity
  15. Consider Halogen Acid Series Acid: H—F H—Cl H—Br H—I pKa: 3.2 -7 -9 -10 Stronger Bonds Weaker Bonds H—X Bond Strength Weaker Bonds  Stronger Acids
  16. Which atom is the charge on? The first factor involves comparing the atoms bearing the negative charge in each conjugate base. 17 Across a row of the periodic table, the acidity of H—A increases as the electronegativity of A increases. Positive or negative charge is stabilized when it is spread over a larger volume.
  17. For example, C- and O-appear in the same row of the periodic table. When two atoms are in the same row, electronegativity is the dominant effect. The story is different when comparing two atoms in the same column of the periodic table. For example, the acidity of water and hydrogen sulfide: In this example, we are comparing O- and S-, which appear in the same column of the periodic table. In such a case, electronegativity is not the dominant effect. Instead, the dominant effect is size. Sulfur is larger than oxygen and can therefore better stabilize a negative charge by spreading the charge over a larger volume of space. The HS- is more stable than HO-, and therefore, H2S is a stronger acid than H2O. We can verify this prediction by looking at pKa values (the pKa of H2S is 7.0, while the pKa of H2O is 15.7).
  18. Resonance Effects • Resonance is a second factor that influences acidity. • In the example below, when we compare the acidities of ethanol and acetic acid, we note that the latter is more acidic than the former. • When the conjugate bases of the two species are compared, it is evident that the conjugate base of acetic acid enjoys resonance stabilization, whereas that of ethanol does not.
  19. • Resonance delocalization makes CH3COO¯ more stable than CH3CH2O¯, so CH3COOH is a stronger acid than CH3CH2OH. • The acidity of H—A increases when the conjugate base A:¯ is resonance stabilized.
  20. Inductive Effects • An inductive effect is the pull of electron density through  bonds caused by electronegativity differences between atoms. • Also Called Electron Withdrawing Effect • In the example below, when we compare the acidities of ethanol and 2,2,2-trifluoroethanol, we note that the latter is more acidic than the former. • The reason for the increased acidity of 2,2,2-trifluoroethanol is that the three electronegative fluorine atoms stabilize the negatively charged conjugate base.
  21. • When electron density is pulled away from the negative charge through  bonds by very electronegative atoms, it is referred to as an electron withdrawing inductive effect. • More electronegative atoms stabilize regions of high electron density by an electron withdrawing inductive effect. • The more electronegative the atom and the closer it is to the site of the negative charge, the greater the effect. • The acidity of H—A increases with the presence of electron withdrawing groups in A.
  22. Hybridization or orbital effect More ‘s’ character in the orbital  more stable anion • The higher the percent of s-character of the hybrid orbital, the closer the lone pair is held to the nucleus, and the more stable the conjugate base.
  23. Summary of Factors Affecting the Stability of Negative Charges/acidity
  24. A convenient way to look at basicity is based on electron pair availability.... the more available the electrons, the more readily they can be donated to form a new bond to the proton and, and therefore the stronger base. Key factors that affect electron pair availability in a base Electronegativity. When comparing atoms within the same row of the periodic table, the more electronegative the atom donating the electrons is, the less willing it is to share those electrons with a proton, so the weaker the base. Increasing basicity
  25. Size. When comparing atoms within the same group of the periodic table, the larger the atom the weaker the H-X bond and the lower the electron density making it a weaker base. Increasing basicity Resonance. In the carboxylate ion, RCO2 - the negative charge is delocalised across 2 electronegative atoms which makes it the electrons less available than when they localised on a specific atom as in the alkoxide, RO-. Strong base Weak base
  26. Lewis Acids and Bases The Lewis definition of acids and bases is more general than the BrØnsted-Lowry definition. •A Lewis acid is an electron pair acceptor. •All BrØnsted-Lowry acids are also Lewis acids, but the reverse is not necessarily true. •Any species that is electron deficient and capable of accepting an electron pair is also a Lewis acid. •Common examples of Lewis acids (which are not BrØnsted-Lowry acids) include BF3 and AlCl3.
  27. •Lewis bases are structurally the same as BrØnsted-Lowry bases. Both have an available electron pair—a lone pair or an electron pair in a  bond. •A BrØnsted-Lowry base always donates this electron pair to a proton, but a Lewis base donates this electron pair to anything that is electron deficient. •A Lewis base is an electron pair donor.
  28. •Lewis acid-base reactions illustrate a general pattern in organic chemistry. Electron-rich species react with electron-poor species. •In the simple Lewis acid-base reaction one bond is formed and no bonds are broken. •This is illustrated in the reaction of BF3 with H2O. H2O donates an electron pair to BF3 to form a new bond. • A Lewis acid is also called an electrophile. • When a Lewis base reacts with an electrophile other than a proton, the Lewis base is also called a nucleophile. electrophile nucleophile
  29. The word steric is derived from ‘stereos’ meaning space. So this effect is manifested when two or more groups or atoms come in close proximity to each other (precisely within each other’s van der Waals radii (definition of van der Waals radii can be found in any standard textbook)) and result in a mutual repulsion. This makes the molecule unstable. Steric effects
  30. The usual physical clash between groups, almost always is accompanied by an electronic component as well. This is called stereoelectronic effect, which is not the same as the electronic effects discussed above and does not carry have an effect on some other part of the molecule like inductive and resonance effects. When the two atoms get to close, into each other’s van der Waal’s radii, the electron cloud surrounding each atom repel each other leading to a lot of destabilization. Steric effect affects different properties of molecules, like acidity, basicity and general reactivity.
  31. • Resonance involves delocalization of π electrons, leaving the σ bond untouched. • However in some cases, a σ bond and an adjacent π bond may get involved in resonance. • Such a delocalization is called as Hyper conjugation. • The electrons of the sigma bond between C and H are involved in delocalization. • In structure to the right: No bond between C and H due to migration of the sigma bond. Hence Hyper conjugation is also called as ‘NO BOND RESONANCE’. Hyper conjugation
  32.  When carbon possessing atleast one Hydrogen is attached to another carbon bearing an unshared orbital or to an unsaturated carbon, there are more than one structure (canonical forms) possible for the molecule.  These result from the overlap of the bonding electrons of C – H σ bond with 2p or π orbital of the adjacent carbon atom.  As a result, the bond between carbon and hydrogen does not exist anymore.
  33. • The shared pair of electrons is now borne by carbon alone and hydrogen is in its close proximity as proton. • The negative charge developed on the carbon gets delocalized by overlap with adjacent p orbital.
  34. This does not indicate that hydrogen is completely detached from the structure, but some degree of ionic character in the C – H bond and some single bond character between carbon – carbon double bond. Such an interaction is also referred to as ‘Heterovalent’ or ‘Sacrificial hyperconjugation’. This is so named because the contributing structure contains one two-electron bond less than the normal Lewis formula of the compound.
  35. Effect of Hyperconjugation on the chemical properties Alkyl cations and their relative stability • Carbocations have an electron deficient (positively charged) carbon. • The empty p orbital of this sp2 carbon can overlap with σ orbital of C – H bond of adjacent alkyl group (α C – H bond). • This overlap permits individual electrons to help bind together three nucleus: 2 carbon and 1 hydrogen. • The positive charge thus gets dispersed over large volume of space and is stabilized.
  36. The more the number of alkyl groups on the carbocation, more is the number of α C – H σ bonds and hence more are the possibilities for hyperconjugation which makes the carbocation more stable. The order of stability of the Carbocations is: 9 Hyperconjugable H 6 Hyperconjugable H 3 Hyperconjugable H No Hyperconjugable H 30alkyl carbocation > 20alkyl carbocation > 10alkyl carbocation > methyl carbocation
  37. Alkyl radicals and their relative stability: • Alkyl radicals have p-orbital of carbon occupied by an odd electron which gets delocalized over three nuclei (2 carbon and 1 hydrogen) by the overlap with the σ orbital of C – H bond of adjacent alkyl groups. • The bond is formed between 2 carbons and odd electron is held by the hydrogen atom. • The order of stability depends on the extent of delocalization. The greater the delocalization, the more is the stability which similar to alkyl cations.
  38. • Overlap of σ orbital of C – H bond with π orbital of adjacent C – C double bond gives rise to canonical structures. • Delocalization of electrons occurs over three nuclei and thus stabilizes the alkene. Alkenes and their stability More the substituent, more is the opportunity for hyperconjugation and more stable is the alkene.
  39. Alkene Number of Hyperconjugable Hydrogen (α to unsaturated function) CH2=CH2 0 (CH3)CH=CH2 3 (CH3)2C=CH2 (CH3)CH=CH(CH3) 6 (CH3)2C=CH(CH3) 9 (CH3)2C=C(CH3)2 12 Stability of alkenes will increase with increase in number of Hydrogen α to unsaturated system. Hyper conjugation leads to shortening of sigma (σ) bond. Eg. C – C bond in 1,3-butadiene and methylacetylene is 1.46 A0 in length, much less as compared to 1.54 A0 found in saturated hydrocarbons. Bond length
  40. Bredt's Rule stated that bridged ring systems like camphane and pinane cannot have a double bond at the bridgehead position. This rule came from observations on dehydration of alcohols in these ring systems. Bredt's Rule
  41. For example, the bicyclooctyl 3º-chloride shown below appears to be similar to tert-butyl chloride, but it does not undergo elimination, even when treated with a strong base (e.g. KOH or KOC4H9). There are six equivalent beta hydrogens that might be attacked by base (two of these are colored blue as a reference), so an E2 reaction seems plausible. The problem with this elimination is that the resulting double bond would be constrained in a severely twisted (nonplanar) configuration by the bridged structure of the carbon skeleton. Because a pi-bond cannot be formed, the hypothetical alkene does not exist. Bredt's Rule should not be applied blindly to all bridged ring systems. If large rings are present their conformational flexibility may permit good overlap of the p-orbitals of a double bond at a bridgehead.
  42. Benzene, first isolated by Michael Faraday in 1825 is the simplest and the ideal molecule to illustrate electron delocalization, resonance and aromaticity. Important milestones during structure elucidation of benzene include: • Friedrich Kekule’s (1866) proposal of cyclic equilibrating structures I and II which partially explained the existence of three isomers (instead of four) for disubstituted benzene • Hydrogenation of benzene to cyclohexane by Paul Sabatier (1901) which confirmed its cyclic structure. Debate over the structure of benzene came to an end in 1930s when X-ray and electron diffraction studies confirmed that it is a planar, regular hexagon in which all the carbon-carbon bond lengths are 1.39 Å, which is shorter than C-C single bond (1.54 Å), but slightly longer than C-C double bond (1.33 Å). Such a structure is possible only if all the carbon atoms have the same electron density, with π electrons delocalized over the entire skeleton of ring carbons. AROMATICITY
  43. Delocalization means possibility of new orbital overlap and additional stabilization of the system. The extra stability (in terms of energy) gained through delocalization is called delocalization energy or resonance energy. The heat of hydrogenation of cyclohexene has been experimentally determined to be 28.6 kcal/mol. If we consider C6H6 as just a cyclohexatriene, the heat of hydrogenation should be 3 x 28.6 kcal/mol = 85.8 kcal/mol. However, when the heat of hydrogenation was experimentally determined for benzene, it was found to be 49.8 kcal/mol. Since hydrogenation of cyclohexatriene and benzene both lead to cyclohexane, reason for the difference in their heat of hydrogenation should be due to the difference in their stabilities. From this, it is clear that benzene is 36 kcal/mol (ie. 85.8-49.8 kcal/mol) more stable than ‘cyclohexatriene’. i.e. benzene with six delocalized π electrons is 36 kcal/mol more stable than ‘cyclohexatriene’ with six localized π electrons. Here, 36 kcal/mol is the resonance energy of benzene (Heat of hydrogenation is the quantity of heat released when one mole of an unsaturated compound is hydrogenated).
  44. In 1931 the German physicist Erich Hückel carrie dout a series of mathematical calculations. Hückel’s rule is concerned with compounds containing one planar ring in which each atom has a p orbital as in benzene. Based on the analysis of a number of compounds with unusual resonance stabilization energies, the following characteristics have been accepted as criteria for aromaticity.  The molecule must be cyclic, planar with uninterrupted cloud of π electrons above and below the plane of the ring.  It should have 4n+2 π electrons. Where n = 0, 1, 2, 3 and so on Hückel’s Rule: The 4n+2 π Electron Rule
  45. Non aromatic compounds, as the name implies, are not aromatic due to reasons such as lack of planarity or disruption of delocalization. They may contain 4n or 4n+2 π electrons. Antiaromatic compounds are planar, cyclic, conjugated systems with an even number of pairs of electrons. Such compounds satisfy the first three criteria for aromaticity. i.e. they are planar, cyclic with an uninterrupted ring of p orbital bearing atoms. But they have an even number of pairs of π electrons (4n, n = 1, 2, 3 etc). It should be noted that an aromatic compound is more stable compared to an analogous cyclic compound with localized electrons, where as an antiaromatic compound is less stable compared to an analogous cyclic compound with localized electrons (in 4n+2 systems delocalization increases the stability where as in 4n systems, delocalization decreases stability)
  46. The bond angles of 120°in benzene suggests that C atoms are sp2 hybridized. An alternative representation therefore starts with a planar framework and considers overlap of the p orbitals (π electrons). MOLECULAR ORBITAL REPRESENTATION OF BENZENE (MO THEORY)
  47. Each MO can accommodate 2 electrons, so for benzene we see all electrons are paired and occupy low energy MO’s (bonding MO’s). All bonding MO’s are filled. Benzene is therefore said to have a closed bonding shell of delocalized π electrons and this accounts in part for the stability of benzene. The relative energies of p molecular orbitals in planar cyclic conjugated systems can be determined by a simplified approach developed by A. A. Frost in 1953. This involves the following steps: 1) Draw a circle 2) Place the ring (polygon representing the compound of interest) in the circle with one of its vertices pointing down. Each point where the polygon touches the circle represents an energy level. 3) Place the correct number of electrons in the orbitals, starting with the lowest energy orbital first, in accordance with Hund’s rule.
  48. Frost diagrams
  49. Points to remember while making predictions on aromaticity using Frost’s circle Aromatic compounds will have all occupied molecular orbitals completely filled where as antiaromatic compounds would have incompletely filled orbitals. The word annulene is incorporated into the class name for monocyclic compounds that can be represented by structures having alternating single and double bonds. The ring size of an annulene is indicated by a number in brackets. Thus, benzene is [6]annulene and cyclooctatetraene is [8]annulene. Hückel’s rule predicts that annulenes will be aromatic if their molecules have 4n +2 π electrons and have a planar carbon skeleton. The Annulenes
  50. During the 1960s, and largely as a result of research by F. Sondheimer, a number of large-ring annulenes were synthesized, and the predictions of Hückel’s rule were verified. Consider the [14], [16], [18], [20], [22], and [24]annulenes as examples. Of these, as Hückel’s rule predicts, the [14], [18], and [22]annulenes (4n+2, when n 3, 4, 5, respectively) have been found to be aromatic. The [16]annulene and the [24]annulene are not aromatic; they are antiaromatic. They are 4n compounds, not 4n +2 compounds:
  51. In [10]-annulene, there is considerable steric interaction between hydrogens at 1 and 6 positions. Further, a planar form (regular decagon) requires an angle of 144o between carbon atoms which is too large to accommodate in a sp2 framework. The system prefers a nonplanar conformation and is not aromatic (the fact that angle strain need NOT always be a problem in achieving planarity is evident from examples such as cyclooctatetraenyl dianion, which is stable and aromatic). Bridging C1 and C6 in [10]- annulene leads to the compound VII (Figure 9) which is reasonably planar with all the bond distances in the range of 1.37-1.42Ao and show aromaticity
  52. If a stabilized cyclic conjugated system (4n+2 e s) can be formed by bypassing one saturated atom, that lead to homoaromaticity. Compared to true aromatic systems, the net stabilization here may be low due to poorer overlap of orbitals. Cyclooctatrienyl cation (homotropylium ion) formed when cyclooctatetraene is dissolved in concentrated sulfuric acid is the best example to demonstrate homoaromaticity. Here, six electrons are spread over seven carbon atoms as in Tropylium cation. HOMOAROMATICITY