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1、高等有机化学Advanced Organic Chemistry,CONTENTS,Chapter 1 Chemical Bonding Chapter 2 Carbocations Carbanions Free Radicals Carbenes and Nitrenes Chapter 3 Mechanisms and Methodsof Determining Them Chapter 4 Acids and Bases Chapter 5 Effects of Structure and Medium on Reactivity,Chapter 6 Aliphatic Substit
2、ution: Nucleophilic and Organometallic Chapter 7 Aromatic Substitution, Electrophilic, Nucleophilic and Organometallic Chapter 8 Aliphatic, Alkenyl, and Alkynyl Substitution, Electrophilic and Organometallic Chapter 9 Substitution Reactions: Free Radicals,Chapter 10 Addition to CarbonCarbon Multiple
3、 Bonds Chapter 11 Addition to CarbonHetero Multiple Bonds Chapter 12 Eliminations Chapter 13 Rearrangements Chapter 14 Oxidations and Reductions,References,Michael B. Smith, Marchs Advanced Organic Chemistry, 6th edition, Wiley-interscience, A John Wiley the loss of two electrons by any individual m
4、olecule almost never occurs). A photoelectron spectrum therefore consists of a series of bands, each corresponding to an orbital of a different energy. The spectrum gives a direct experimental picture of all the orbitals present, in order of their energies, provided that radiation of sufficiently hi
5、gh energy is used.,Broad bands usually correspond to strongly bonding electrons and narrow bands to weakly bonding or nonbonding electrons.,A typical spectrum is that of N2, shown in Fig. 1.8.The N2 molecule has the electronic structure shown in Fig. 1.9.,The two 2s orbitals of the nitrogen atoms co
6、mbine to give the two orbitals marked 1 (bonding) and 2 (antibonding), while the six 2p orbitals combine to give six orbitals, three of which (marked 3, 4, and 5) are bonding. The three antibonding orbitals (not indicated in Fig. 1.9) are unoccupied. Electrons ejected from orbital 1 are not found in
7、 Fig. 1.8 because the ionization potential of these electrons is greater than the energy of the light used (they can be seen when higher energy light is used). The broad band in Fig. 1.8 (the individual peaks within this band are caused by different vibrational levels) corresponds to the four electr
8、ons in the degenerate orbitals 3 and 4. The triple bond of N2 is therefore composed of these two orbitals and orbital 1.,The bands corresponding to orbitals 2 and 5 are narrow; hence these orbitals contribute little to the bonding and may be regarded as the two unshared pairs of N=N. Note that this
9、result is contrary to that expected from a naive consideration of orbital roverlaps, where it would be expected that the two unshared pairs would be those of orbitals 1 and 2, resulting from the verlap of the filled 2s orbitals, and that the triple bond would be composed of orbitals 3, 4, and 5, res
10、ulting from overlap of the p orbitals.,1.1.4 Electronic Structures Of Molecules,For each molecule, ion, or free radical that has only localized electrons, it is possible to draw an electronic formula, called a Lewis structure, that shows the location of these electrons. Only the valence electrons ar
11、e shown. Valence electrons may be found in covalent bonds connecting two atoms or they may be unshared. The student must be able to draw these structures correctly, since the position of electrons changes in the course of a reaction, and it is necessary to know where the electrons are initially befo
12、re one can follow where they are going.,1.1.5 Electronegativity,The electron cloud that bonds two atoms is not symmetrical (with respect to the plane that is the perpendicular bisector of the bond) except when the two atoms are the same and have the same substituents. The cloud is necessarily distor
13、ted toward one side of the bond or the other, depending on which atom (nucleus plus electrons) maintains the greater attraction for the cloud. This attraction is called electronegativity,Electronegativity information can be obtained from NMR spectra. In the absence of a magnetically anisotropic grou
14、p the chemical shift of a 1H or a 13C nucleus is approximately proportional to the electron density around it and hence to the electronegativity of the atom or group to which it is attached. The greater the electronegativity of the atom or group, the lower the electron density around the proton, and
15、 the further downfield the chemical shift. An example of the use of this correlation is found in the variation of chemical shift of the ring protons in the series toluene, ethylbenzene, isopropylbenzene, tert-butylbenzene,The dipole moment is a property of the molecule that results from charge separ
16、ations like those discussed above. However, it is not possible to measure the dipole moment of an individual bond within a molecule; we can measure only the total moment of the molecule, which is the vectorial sum of the individual bond moments.,1.1.6 Dipole Moment,1.1.7 Inductive And Field Effects,
17、The CC bond in ethane has no polarity because it connects two equivalent atoms. However, the CC bond in chloroethane is polarized by the presence of the electronegative chlorine atom. This polarization is actually the sum of two effects. In the first of these, the C-1 atom, having been deprived of s
18、ome of its electron density by the greater electronegativity of Cl, is partially compensated by drawing the CC electrons closer to itself, resulting in a polarization of this bond and a slightly positive charge on the C-2 atom. This polarization of one bond caused by the polarization of an adjacent
19、bond is called the inductive effect.,The effect is greatest for adjacent bonds but may also be felt farther away; thus the polarization of the CC bond causes a (slight) polarization of the three methyl CH bonds. The other effect operates not through bonds, but directly through space or solvent molec
20、ules, and is called the field effect. It is often very difficult to separate the two kinds of effect, but it has been done in a number of cases, generally by taking advantage of the fact that the field effect depends on the geometry of the molecule but the inductive effect depends only on the nature
21、 of the bonds. For example,Functional groups can be classified as electron-withdrawing (-I) or electrondonating (+I) groups relative to hydrogen. This means, for example, that NO2, a I group, will draw electrons to itself more than a hydrogen atom would if itoccupied the same position in the molecul
22、e.,1.1.8 Bond Distances,1.1.9 Bond Angles,1.1.10 Bond Energies,1.2 Delocalized Chemical Bonding,These compounds contain one or more bonding orbitals that are not restricted to two atoms, but that are spread out over three or more. Such bonding is said to be delocalized.,For planar unsaturated and ar
23、omatic molecules, many molecular-orbital calculations (MO calculations) have been made by treating the s and p electrons separately. It is assumed that the s orbitals can be treated as localized bonds and the,One type of MO calculation that includes all electrons is called ab initio. Despite the nam
24、e (which means from first principles) this type does involve assumptions, though not very many. It requires a large amount of computer time, especially for molecules that contain more than about five or six atoms other than hydrogen. Treatments that use certain simplifying assumptions (but still inc
25、lude all electrons) are called semiempirical methods. One of the first of these was called CNDO (Complete Neglect of Differential Overlap), but as computers have become more powerful, this has been superseded by more modern methods, including MINDO/3 (Modified Intermediate Neglect of Differential Ov
26、erlap), MNDO(Modified Neglect of Diatomic Overlap), and AM1 (Austin Model 1), all of which were introduced by M.J. Dewar and co-workers. Semiempirical calculations are generally regarded as less accurate than ab initio methods, but are much faster and cheaper. Indeed, calculations for some very larg
27、e molecules are possible only with the semiempirical methods.,Molecular-orbital calculations, whether by ab initio or semiempirical methods, can be used to obtain structures (bond distances and angles), energies (e.g., heats of formation), dipole moments, ionization energies, and other properties of
28、 molecules, ions, and radicals: not only of stable ones, but also of those so unstable that these properties cannot be obtained from experimental measurements.,Many of these calculations have been performed on transition states; this is the only way to get this information, since transition states a
29、re not, in general, directly observable. Of course, it is not possible to check data obtained for unstable molecules and transition states against any experimental values, so that the reliability of the various MO methods for these cases is always a question. However, our confidence in them does inc
30、rease when (1) different MO methods give similar results, and (2) a particular MO method works well for cases that can be checked against experimental methods.,Kinds of Molecules That Have Delocalized Bonds,There are four main types of structure that exhibit delocalization: 1. Double (or Triple) Bon
31、ds in Conjugation.,2. Double (or Triple) Bonds in Conjugation with a p Orbital on an Adjacent Atom. Where a p orbital is on an atom adjacent to a double bond, there are three parallel p orbitals that overlap. As previously noted, it is a general rule that the overlap of n atomic orbitals creates n m
32、olecular orbitals, so overlap of a p orbital with an adjacent double bond gives rise to three new orbitals, as shown in Fig. 2.4.,The middle orbital is a nonbonding orbital of zero bonding energy. The central carbon atom does not participate in the nonbonding orbital.,There are three cases: the orig
33、inal p orbital may have contained two, one, or no electrons. Since the original double bond contributes two electrons, the total number of electrons accommodated by the new orbitals is four, three, or two. A typical example of the first situation is vinyl chloride CH2CHCl. Although the p orbital of
34、the chlorine atom is filled, it still overlaps with the double bond(see 10). The four electrons occupy the two molecular orbitals of lowest energies. This is our first example of resonance involving overlap between unfilled orbitals and a filled orbital. Canonical forms for vinyl chloride are shown
35、in 11.,Kinds of Molecules That Have Delocalized Bonds,There are four main types of structure that exhibit delocalization: 1.Double (or Triple) Bonds in Conjugation,2. Double (or Triple) Bonds in Conjugation with a p Orbital on an Adjacent Atom.,There are three cases: the original p orbital may have
36、contained two, one, or no electrons. Since the original double bond contributes two electrons, the total number of electrons accommodated by the new orbitals is four, three, or two.,3. -Allyl and Other -Complexes. In the presence of transition metals, delocalized cations are stabilized by donating e
37、lectrons to the metal.,Ligands can therefore be categorized as -ligands according to their electron donation to the metal. A hydrogen atom (as in 14) or a halogen ligand (as in 13) are 1 ligands and an amine (NR3), a phosphine (PR3, as in 13, 14, and 18), CO (as in 16 or 17), an ether (OR2) or a thi
38、oether (SR2) are 2 ligands.,Hydrocarbon ligands include alkyl (as the methyl in 15) or aryl with a Cmetal bond (1), alkenes or carbenes (2), p-allyl (3), conjugated dienes such as 1,3-butadiene (4), cyclopentadienyl (5, as in 15), and arenes or benzene (6). Note that in the formation of 14 from 13,
39、the two electron donor alkene displaces a two-electron donor phosphine. Other typical complexes include chromium hexacarbonyl Cr(CO)6 (16), with six 2-CO ligands; 6 -C6H6Cr(CO)3 (18), and tetrakistriphenylphosphinopalladium(0), 17, with four 2 -phosphine ligands.,4. Hyperconjugation. The type of del
40、ocalization called hyperconjugation, isdiscussed on p. 95.We will find examples of delocalization that cannot be strictly classified asbelonging to any of these types.,Cross Conjugation,In a cross-conjugated compound, three groups are present, two of which are not conjugated with each other, althoug
41、h each is conjugated with the third. Some examples are benzophenone (21), triene 22 and divinyl ether 23. Using the molecular-orbital method, we find that the overlap of six p orbitals in 22 gives six molecular orbitals, of which the three bonding orbitals are shown in Fig. 2.5, along with their ene
42、rgies.,The Rules of Resonance,1. All the canonical forms must be bona fide Lewis structures (see p. 14). For example, none of them may have a carbon with five bonds. 2. The positions of the nuclei must be the same in all the structures. This means that when we draw the various canonical forms, all w
43、e are doing is putting in the electrons in different ways. For this reason, shorthand ways of representing resonance are easy to devise:,3. All atoms taking part in the resonance, that is, covered by delocalized electrons, must lie in a plane or nearly so. This, of course, does not apply to atoms th
44、at have the same bonding in all the canonical forms. The reason for planarity is maximum overlap of the p orbitals. 4. All canonical forms must have the same number of unpaired electrons. Thus .CH2CHCHCH2. is not a valid canonical form for butadiene. 5. The energy of the actual molecule is lower tha
45、n that of any form, obviously. Therefore, delocalization is a stabilizing phenomenon.,6. All canonical forms do not contribute equally to the true molecule. Each form contributes in proportion to its stability, the most stable form contributing most. Thus, for ethylene, the form tCH2CH2 has such a h
46、igh energy compared to CH2CH2 that it essentially does not contribute at all. We have seen the argument that such structures do not contribute even in such cases as butadiene.36 Equivalent canonical forms, such as 1 and 2, contribute equally. The greater the number of significant structures that can
47、 be written and the more nearly equal they are, the greater the resonance energy, other things being equal.,The Resonance Effect,Steric Inhibition of Resonance and the Influences of Strain,Rule 3 states that all the atoms covered by delocalized electrons must lie in a plane or nearly so. Many exampl
48、es are known where resonance is reduced or prevented because the atoms are sterically forced out of planarity.,1.3 Bonding Weaker than Covalent,In the first two sections, we discussed the structure of molecules each of which is an aggregate of atoms in a distinct three-dimensional (3D) arrangement h
49、eld together by bonds with energies on the order of 50100 kcal mol1 (200400 kJ mol1). There are also very weak attractive forces between molecules, on the order of a few tenths of a kilocalorie per mole.,These forces, called van der Waals forces, are caused by electrostatic attractions, such as those between dipole and dipole, induced dipole, and induced dipole, and are responsible for liquefaction of gases at sufficiently low temperatures. The bonding discussed in this chapter has energies of the order of 210 kcal mol-1 (940 kJ mol-1), i