1 Benzene 1 NT ompound 87 has the formula 6 6, is known as benzene, and it is a hydrocarbon derived from petroleum distillates. Benzene is the parent compound for a class of compounds known as aromatic hydrocarbons, where the term aromatic refers to the special stability imparted by the six-π-electrons in benzene when they are confined to a ring. The bond length of a double bond (=) is shown to be shorter than that of a single bond (-) earlier in this chapter. In cyclohexadiene (89), there are two distinct bond lengths, a longer one for the - bonds and shorter ones for the = bonds. The π-bond is stronger, so the internuclear distance between the carbons should be smaller. The bond length of a standard = unit is about 133 pm and the bond length of a standard - unit is about 148 pm. This alternating pattern is used to draw 90, which has exaggerated bond lengths to emphasize the difference in length between = and - units. In reality, the structure of benzene is 87, not 90, and measurements show that all - bond lengths in benzene are identical, about 139 pm. The only explanation is that the bonding in benzene is different than in other cyclic alkenes. For this reason, it is not appropriate to call it cyclohexatriene.
2 Benzene: Lack of Reactivity 2 l l l The structure of benzene is 87, so it contains six π-electrons and is certainly electron rich. The π- bond of a simple alkene such as 91 has only two electrons, benzene is more electron rich than an alkene. In chemical reactions that depend upon electron donation (as in a Lewis base donating electrons to a positive center), the greater the electron density of a molecule, the easier it is to donate electrons. The π-electrons of an alkene (91) react as a base with l, donating the electrons to form a new - bond and forming carbocation 92. This is an acid-base reaction in which the alkene is the base. This electron donating reaction can be used to examine the electron donating ability of benzene relative to cyclohexene and cyclohexadiene. yclohexene reacts rapidly with l, as does cyclohexadiene. When benzene is mixed with l, however, there is no reaction. Although benzene has six p-electrons, benzene cannot donate electrons as well as a simple alkene with a single π-bond. Benzene is less reactive than a simple alkene. If benzene does not react with l, benzene must be a weaker base than the alkene. Benzene is weaker base (unable to donate electrons as efficiently) because the six π-electrons of benzene are delocalized on six carbons whereas the two π-electrons of an alkene are only distributed between two carbons. Benzene is resonance stabilized, making it less reactive.
3 Benzene: Aromaticity 3 87A 87B 87 Examination of 87 shows that all six carbons of benzene are sp 2 hybridized, which means each carbon has trigonal planar geometry (see 87A) and a p-orbital that is perpendicular to the plane of the carbon atoms as represented by 87B. Each of the six π-orbitals contain a π-electron. Six contiguous carbons with π-orbitals (every carbon with a π-orbital is adjacent to another carbon with a π- orbital) in a ring share electron density as indicated in 87B. The electrons in the π-orbitals are referred to as π-electrons. Benzene is a molecule composed of six carbon atoms, each with one hydrogen, forming a planar ring. The six π-orbitals that are perpendicular to the plane of the carbons contain a total of six π-electrons and these π- orbitals share the six electrons. The six π-electrons are delocalized on the 6 π-orbitals. This delocalization in a cyclic π-system is resonance, and when confined to a ring in this manner, this type of resonance is known as aromaticity. Benzene is less reactive and more stable due to aromaticity. The intense red color in the center of the ring in the electron density map (87), above and below, indicates the concentration of electron density associated with the aromatic π-cloud.
4 Benzene: Resonance Structures 4 87A 87D 94 If benzene is drawn with the π-bonds "localized" as in 87A, there is no indication of electron delocalization and this single structure does not adequately represent the structure of benzene. If the double bonds are "moved" from their position in 87A to give 87D, 87D is also inadequate since all the "double bonds" in 87A are single bonds in 87D, and vice-versa. The actual structure of benzene is not 87A or 87D, but the actual structure can be represented by both structures, called resonance contributors. Imagine the electrons shifting back and forth between 87A and 87D (the electrons are "moving" within the framework of π-orbitals) to represent the electron delocalization in benzene. Showing both structures, a double-headed arrow relates the resonance structures, and the actual structure of benzene is represented by both 87A and 87D. Benzene is sometimes represented as a six-membered ring with a circle in the middle (see 94) to indicate the resonance. Structure 94 is inadequate for most of the chemical reactions of benzene because chemical reactions are usually viewed in terms of electron transfer. The movement of pairs of electrons cannot be represented with 94, so the use of this representation is discouraged.
5 Phenyl Substituent 5 Ph Ph 3 95 l 96 A benzene ring can be attached to a carbon chain, however, and in such cases the benzene unit is a substituent. In 95, the benzene ring is attached to 5 of 2-octanone. This benzene substituent has the formula 6 5 and it is attached to the carbon chain by a carbon-carbon bond. The 6 5 unit is called phenyl, and the name of 95 is 5-phenyl-2-octanone. Drawing a benzene ring occupies a lot of space so a shorthand representation is used in many structures. In 96, the shorthand symbol "Ph" is used to represent a phenyl substituent. The Ph representation will be used often in this book, ompound 96 is named 5-chloro-2,6-diphenyloct-2-ene.
6 6 hapter 6. Acids, Bases, Nucleophiles and Electrophiles It is important to modify views of acids to include very weak acids. The functional groups introduced in chapter 5 must be considered as weak acids, including alcohols, ketones, terminal alkynes, and even primary and secondary amines. If these functional groups are weak acids, then an acid-base reaction requires the use of a powerful base to generate the corresponding conjugate base.
7 To begin, you should know: 7 ovalent σ-bonds. (chapter 3, section 3.3) Polarized covalent bonding between carbon and heteroatoms and hydrogen and a heteroatom. (chapter 3, section 3.7) π-bonds (chapter 5, sections 5.1, 5.2, and 5.9) Factors that influence bond strength. (chapter 3, sectin3.6 and chapter 5, section 5.4) What is K a and pk a? (chapter 2, section 2.4) What factors influence the equilibrium constant K a in an acid-base reaction? (chapter 2, section 2.4) What constitutes a strong acid or a strong base? What constitutes a weak acid or a weak base? (chapter 2, section 2.4) What is electron release and electron donation? (chapter 3, section 3.7) What is an inductive effect? (chapter 3, section 3.7) The definition of a Lewis acid and a Lewis base. (chapter 2, section 2.5) Structures and names of functional groups. (chapter 4, and chapter 5, sections 5.6 and 5.9) Acid-base properties of functional groups. (chapter 5, section 5.7) What is resonance? (chapter 5, section 5.9.)
8 When completed, you should know: 8 ommon organic acids are carboxylic acids, sulfonic acids, and alcohols. arboxylic acids are more acidic than alcohols, but sulfonic acids are more acidic than carboxylic acids. The equilibrium constant for an acid-base reaction that involves organic acids is the same as any other acid-base reaction, K a, where K a = products/reactants, and pk a = log K a. The relative strength of an acid depends on the strength of the base it reacts with. The free energy of reaction is indicative of the spontaneity of that reaction (endothermic or exothermic), and the activation energy is a measure of the transition state energy. ΔG = Δ TΔS. Stability of a conjugate base is linked to its reactivity, and greater stability can result from charge dispersal due to increased size of that species or resonance stability. In both cases, K a is larger. Inductive effects are very important for determining the variations in pk a with structural changes in carboxylic acids. As electron-withdrawing substituents are positioned further away from the carboxyl group, the effect diminishes. An electron-withdrawing group attached to oxygen in an unit makes the hydrogen more acidic, and an electron-releasing group attached to oxygen in an unit makes the hydrogen less acidic.
9 When completed, you should know: 9 The solvent can play a significant role in determining K a for an acid, particularly when comparing polar ionizing solvents with nonpolar solvents. ompounds that behave as an acid in the presence of a strong base and a base in the presence of a strong acid are called amphoteric compounds. Alcohols, ethers, aldehydes and ketones, and alkenes are Brønsted-Lowry and Lewis bases. ommon organic bases includes amines, phosphines and alkoxide anions. The basicity of an amine is measured using pk B. The basicity of amines is influenced by both the electron releasing effects of alkyl groups attached to nitrogen as well as by the steric hindrance imposed by those alkyl groups to anything approaching the nitrogen. An electronwithdrawing group on the nitrogen of an amine will diminish its basicity. Lewis acids are electron pair acceptors and Lewis bases are electron pair donors. Electrophiles are carbon molecules that react with electron donating compounds. Nucleophiles are molecules that donate electrons to carbon.
10 Acid-Base Equilibrium 10 K! a!! A- + :B B- + :A acid base conjugate acid! conjugate base A classical acid-base reaction reacts an acid (A:) with a base (B:) to give a conjugate acid (B:) and a conjugate base (A:) as the products. This is an equilibrium reaction and the equilibrium constant K is given the special designator K a. If K a is used for this reaction, then pk a = log K a In this reaction, K a = and K a = [B][A:] / [A][B:] and pk a = log K a Note that pk a and K a are inversely proportional and K a = 10 pka
11 Acid-Base Equilibrium 11 K! a!! A- + :B B- + :A acid base conjugate acid! conjugate base Assume that the larger the value of K a (the smaller the pk a ) will correlate with a stronger acid. If K a is small (large pk a ), the equilibrium lies to the left, which essentially means that A: does not react with the base. The term weaker acid is a relative term, but a smaller K a (a larger pk a ) will correlate with a weaker acid.
12 Acid-Base Equilibrium: Reverse Reaction 12 K! a!! A- + :B B- + :A acid base conjugate acid! conjugate base The reverse reaction B: + A: to give A: + B: is also an acid-base reaction where B: is the acid (labeled conjugate acid) and A: is the base (labeled conjugate base). If the pk a of A: is 2.5 and the pk a of B: is 12.1, then the equilibrium constant is relatively large for the "forward" reaction and the reaction proceeds to the right. If the pk a of A: is 4.6 and the pk a of B: is 4.7, then K a is close to unity and there will be a mixture of both acids and both bases.
13 What is a Brønsted-Lowry Acid? 13 Normally, a generic acid is written as +. This is simply a proton, of course, which is a hydrogen atom with no electrons and a formal charge of +1. If + is taken to be an acid, then any hydrogen atom that is polarized δ+ should be acidic to some degree. This positive polarization is most often observed when hydrogen is attached to a heteroatom, particularly or S, and to a lesser extent N. Alcohols and carboxylic acids are defined as functional groups that have an - unit. These are - acids. Amines may have a N- unit, which is a N- acid.
14 rganic acids Alcohols versus arboxylic Acids
15 - Acids: Alcohols The - unit in an alcohol is polarized as shown in 1, and that proton is acidic. Most alcohols, which have pk a values of 16-18, but there is a notable exception. Methanol (1) has a pk a of 15.2 whereas water has a pk a of The conjugate base of this reaction is an alkoxide, R, 2 The base that reacts with the alcohol should generate a conjugate acid with a pk a that is greater than acid conjugate base Na + N 2 + N 2 Na 1 base conjugate acid 2 15
16 - Acids: Alcohols The hydronium ion is a quite potent acid, much stronger than methanol, and the equilibrium is shifted to the left, which means that K a is much smaller. If water is a weak base relative to sodium amide, then K a is large for the reaction with sodium amide and small for the reaction with water. The strength of the acid depends on the strength of the base. Methanol is a very weak acid when water is the base but it is a stronger acid when sodium amide is the base. acid 1 conjugate base + + base 2 16 conjugate acid 3
17 - Acids: arboxylic Acids & Sulfonic Acids arboxylic acids (4) and sulfonic acids (5) have pk a values generally in the range of 1 5, depending on the nature of the "R" group. In 5 the polarized S= unit leads to a larger δ oxygen in the unit, which makes the attached hydrogen more positive, which makes 5 slightly more acidic. Sufonic acids are 3-5 units more acidic - mostly due to increased stability of the conjugate base. R R S
18 The onjugate Bases 18 R R R R S R S R S R S The sulfonate anion has more resonance contributors: more charge dispersal and less reactive, so Ka for the sulfonic acid is larger - i.e., more acidic
19 - Acids: Formic Acid versus Methanol Formic acid (,) is a specific example of a carboxylic acid, with a pk a of Methanol has a pk a value of about When formic acid reacts with sodium amide (the base), removal of the proton leads to the formate anion (50) as the conjugate base. whereas similar reaction of methanol ( 3 ) gives methoxide ( 3, 60). The resonance stability of 50 makes that K a larger and the lack of charge dispersal for methoxide, 60, makes K a smaller. 19 Na + + Na + N N Na + Na + N N 2 methoxide formate anion formate anion
20 Factors That Influence the Strength of a Brønsted-Lowry Acid 20 Stability of the onjugate Base The size of the conjugate base and whether or not it is stabilized by resonance are two important factors that influence charge dispersal. If the charge of a conjugate base is dispersed over a greater area, it is more stable (lower reactivity). The resonance stability of an ion is determined by the extent of delocalization. The size of the anion is important, as when two different size conjugate bases are compared: iodide from I versus fluoride from F. In general, anything that leads to a more stable product will shift the equilibrium to products and a larger K a.
21 Factors That Influence the Strength of a Brønsted-Lowry Acid Solvent Effects: In Water Typical pk a values are based on their reaction in water, which means that water is the base in those reactions. If a different solvent is used, the pk a is different. If another base is added to the water solution, the pk a may be different. Remember that acid strength depends on the strength of the base, which influences the position of the equilibrium and K a. In the reaction of acetic acid (ethanoic acid, 18) and water, the conjugate acid is 3 + and the conjugate base is the acetate anion (24): water reacts as the base, but water is also the solvent and it has a profound effect on the course of the reaction
22 Factors That Influence the Strength of a Brønsted-Lowry Acid Water is a very polar molecule, and it is capable of solvating and separating ions. nce the reaction begins, water will solvate the acetate ion (24) and the hydronium ion ( 3 + ), which serves to separate the developing ions. Solvation means that solvent surrounds each ion. Separation of the ions pushes the reaction to the right since it facilitates formation of the products (higher concentration of products and larger K a ). The net result is that acetic acid is a stronger acid in water than in a solvent that cannot generate a polarized transition state to assist the ionization
23 Factors That Influence the Strength of a Brønsted-Lowry Acid 23 Solvent Effects: In ether If one molar equivalent of water is used as a base in a reaction with 18, but diethyl ether is used as a solvent, acetic acid is a weaker acid. As a solvent, ether does not separate ions, so ionization of 23 to form 24 and the hydronium ion will be slower than in water. Because the solvent does not facilitate charge separation nor stabilize the ionic products, acetic acid is a weaker acid (smaller K a, larger pk a ) in ether than in the water solvent. Water facilitates ionization and charge separation more efficiently than other solvents and acetic acid is a stronger acid in water than in diethyl ether Et Et Et Et + Et Et
24 rganic acids Structural variations in carboxylic acids Sulfonic acids
25 Factors That Influence the Strength of a Brønsted-Lowry Acid 25 Structural Variations. Inductive Effects In a carboxylic acid, that is the sp 2 carbon of a carbonyl group, but that is the sp 3 carbon of an alkyl group in alcohols. The nature of these groups has an influence on the pk a of proton in the - unit of the acid. An inductive effect occurs when electron density is donated towards the acidic proton or pulled away from the acidic proton. In general, "pushing" electron density towards the proton (electron releasing) makes the - bond stronger and that proton is less acidic. onversely, "pulling" electron density away from the proton (electron withdrawing) makes the - bond weaker and the proton is more acidic.
26 Inductive Effects: Formic versus Acetic δ δ+ 18A δ Not a dipole moment - electron redistribution δ+ Formic acid has a pk a of 3.75 and acetic acid has a pk a of Relative to hydrogen, the carbon group is electron releasing, because carbon is slightly more electronegative than hydrogen. This difference in electronegativity makes the carbon slightly electron releasing (see 18A) in the direction of the arrow. This effect occurs through the bonds by pushing and pulling electron density and it is called a through-bond inductive effect, graphically illustrated in 18A. In 6A and 18B,, the red area indicates higher electron density and the blue area lower electron density. 6A 18B
27 Inductive Effects: Acetic versus hloroacetic 3 18 l B 19 ompare acetic acid (18, pk a 4.76) and chloroacetic acid (19, pk a 2.87), with a chlorine on the sp 3 α-carbon. The acidic hydrogen and the chlorine are close together in space so there is an intramolecular hydrogen bond. This effect in 19A is described as a through-bond effect. This -bond makes the - bond more polar ( is more δ+) and more acidic. The effect in 19B is a through-space effect due to the attraction of the chlorine for the acidic hydrogen. This through-space effect is much more important than the through-bond effect described above. l 19A l 19B