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New Findings in the Electrophilic Addition of Halogens to Olefins.

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New Findings in the Electrophilic Addition of Halogens to Olefins
Rainer Herges*
According to the textbooks, both the mechanism and the
theory of electrophilic addition of halogens to olefins are well
understood.“] It is all the more surprising that there is still some
groundwork to be done in this field, as was demonstrated by
recent synthetic, spectroscopic, and theoretical publications.
As early as 1937, Roberts and KimballLZ1explained the anti
selectivity of the bromination of olefins by invoking a bridged
bromonium ion as intermediate. Thirty years later, such intermediates were detected under “stable ion conditions” by N M R
spectroscopy.[31The halonium ions of chlorine,[41 bromine,”]
and iodine“] with adamantylideneadamantane are even isolable
as stable salts, because the subsequent nucleophilic attack of the
counterion is prevented by steric hindrance in the substrate. It
is generally accepted that a n complex precedes formation of
such halonium
and it is supposed that cation-stabilizing substituents (e.g., phenyl) at the olefin lead to distorted or
even nonbridged cations (Scheme 1 ) .
n complex
halonium ion
Scheme 1
In the case of adamantylideneadamantane and some other
sterically hindered olefins, the free olefin and bromine are in
rapid equilibrium with the n complex and the bromonium
ion.[’. l o ,
The equilibrium constant for the formation of the n
complex from olefin and bromine strongly depends on the substituents. Moreover, it was demonstrated by N M R that there is
a rapid. degenerate transfer of Br’ from the bromonium ion to
the free olefin in the adamantylideneadamantane system
(Scheme 2) . I h . ‘*]According to a b initio calculations the transfer
proceeds via a “spiro” transition state.
Whereas crystal structures are available for bromonium and
iodonium salts,[’3-61little experimental data exists for the J[
complexes. According to calculations by Cossi et al.[’41 and
earlier theoretical
the n complex of ethene and
[*] P I - I \ . - L ~ Dr.
~ . R. Herges
Inatitul fur Organische Chemie der Universitdt Erlangen-Nurnberg
Henkestraw 42, D-91054 Erlangen (Germany)
Telefiix. l n t . code (9131)X5-9132
e-mail . herpestlr
Scheme 2
bromine is about 3-4 kcalmol-’ more stable than the starting
materials. The axis joining the nuclei in the halogen molecule is
perpendicular to the molecular plane of ethene and intersects
the midpoint of the double bond. No charge is transferred during complex formation.[’ 6] The frequently used term “chargetransfer complex” is therefore not correct. Since dispersion energy is the main driving force for binding, “van der Waals
complex” would be a better description. Compared to the n
complex, the bromonium ion is mucher higher in energy. In
vacuo (and with bromide as the counterion) the free energy
difference is 40 kcalmol-’. Under these conditions the bromonium ion is not likely to be formed by simple thermal activation.
However, polar media like methanol lower the free energy of
reaction to about + 10 kcal molbecause the very polar
bromonium ion is much better solvated than the nonpolar TI
A nonpolar environment and low temperatures should therefore allow the n complex to be observed as a long-lived species.
Indeed, as far back as the seventies, the n complex of chlorine
and ethene was detected by IR spectroscopy in a nitrogen matrix
at 20 K.[”.
A totally different and probably revolutionary
method for the structural investigation of such labile molecular
complexes in the gas phase was reported recently by Legon et al.
in a series of papers: Both components (halogen and Lewis base)
are expanded separately through a fast-mixing nozzle into the
evacuated Fabry-Perot cavity of a Fourier transform microwave spectrometer.[’91 The gases mix, and the supersonic
and collision-free expansion results in the molecular complexes
being frozen with very low internal energies. Thus, subsequent
reaction to form the usual products observed under normal
conditions is prohibited. Even complexes such as H,P . . CI,
can be investigated-the normally explosive course of reaction
is avoided.
The rotational spectra d o not only give structural information; the intermolecular (quadratic) force constants are related
to the centrifugal distortion constants, and the charge distribution can be derived from the nuclear quadrupole coupling constants. Table 1 shows an overview of halogen/Lewis base complexes investigated recently.[201
Table 1. Bond lengths r( B . ' . X) and force constants X , for the B . . X bond in the
halogeniLewis base complexes B . . . XY. characterized by rotational spectroscopy.
B . . XY
r(B. . . X)
k. [Nm-'1
4.2124.332 [a]
6.186.34 [a]
Hopf et al. recently reported a fascinating, sterically hindered
system that, in solution at room temperature, prevents further
reaction of the TI complex and bromonium ion to form the usual
dihalogenated products.[321 With 1,2,3-tri-tert-butyl substituted
butadiene, the cationic intermediate is stabilized by loss of an a
proton. Thus, the overall course of reaction corresponds to that
observed for aromatic electrophilic substitutions, which also
proceed via a TI complex and cationic intermediate (Scheme 3).
In the bromination of 2,2,5,5-tetramethyl-3-hexene,the intermediate bromonium ion is stabilized by loss of a proton.
Subsequent HBr elimination yields 2,3-di-tevt-butyl-l,3-butadiene.[331
Scheme 3
The examples show that long-known reactions are also worthy of closer inspection. Modern theoretical, analytical, and
synthetic methods not only open novel fields of application,
they are also able to contribute to the basic knowledge of mechanistic chemistry.
[a] Dependent on the isotopic composition
German version: A n g w . Chem. 1995, 107, 57
From a mechanistic point of view, the complexes of ethene
and ethyne are most interesting. The axis joining the nuclei in
the halogen molecule is perpendicular to and intersects the midpoint of the double o r triple bond (Fig. 1). The structural
parameters of both the halogen and ethene or ethyne are not
significantly changed by complexation. The small force constants k, (Table 1) also indicate a very weak interaction between
TI bond and
The nuclear quadrupole coupling constants of the chlorine atoms differ only slightly and are almost
equal to the values observed for free chlorine. This means that
there is only a small perturbation in the charge distribution on
moving from the separate components to the complex. No significant polarization within the bound halogen is observed.
Hence, the microwave investigations provide an impressive confirmation of the properties predicted earlier by theoretical calculations and matrix isolation.
: 3.128 (3.037)
3.165 (3.124)
Fig. 1. Structures of the x complexes of chlorine with ethene [21] (left) and ethyne
A (ab initio values (311 in parentheses).
[22] (right). Bond lengths In
'i: VCH Verlugsgi~.~ellscIi~~~
m h H , 0-69451 Weinheim.1995
Keywords: alkene- halogen complexes alkyne- halogen complexes - electrophilic addition reaction mechanisms
[l] For a very detailed review of the mechanism of the bromination of olefins, see
M.-F. Ruasse. Adv. Phys. Org. Chem. 1993. 28. 207.
[2] I. Roberts, G. E. Kimball, J. A m . Chem. SOC. 1937,59,947
[3] a) G. A. Olah, J. M. Bollinger, J. Brinich, J. Am. Chem. Soc. 1968, 90, 2587; b)
G. A. Olah. P. Schilling. P. W. Westerman, H. C. Lin, ibid. 1974, 96, 3581.
[4] J. H. Wieringa. J. Strating. H . Wynberg, Tetrahedron Lett. 1970, 4579.
[5] J. Strating. J. H. Wieringa. H. Wynberg. J. Chem. SOC.Chem. Cummun 1969,
[6] R. S. Brown, R. W. Nagorski, A. J. Bennet, R. E. D. McClung. G. H. M.
Aarts. M Klobukowski, R. McDonald, B. D. Santasiero, J. Am. Chem. Sor.
1994, 116, 2448
(71 For the first unambiguous spectroscopic evidence for the formation of a x
complex as intermediate in the electrophilic addition of chlorine and bromine
to olefins. see a) J. E. Dubois. F. Garnier. Spectrorhim. Acto Purl A 1967, 23,
2288; b) R. S. Brown, H. Slebocka-Tilk. A. J. Bennet. G. Bellucci, R. Bianchini. R. Ambrosetti. J. Am. Chem. Sue. 1990, 112. 6310; c) G. Bellucci. R.
Bianchini R. J. Ambrosetti. ihid. 1985. 107.2464; d) S. Fukuzumi, J. K. Kochi,
J. Org. Chem. 1981, 46. 4116.
[8] From the very strong dependance of the reaction rate, Dubois and Garnier
conclude that the nucleophilic attack of the counterion or the solvent takes
place at the stage of the bromonium ion, rather than at that of the x complex.
The unusually strong solvent isotope effect also suggests that the transition
state of the rate-determining step is strongly polarized and most probably
corresponds to the formation of the bomonium ion from t h e n complex: a) J. E.
Dubois. F. Garnier, J. Chem. Sue. Chem. Commun. 1968. 241 ; b) F. Gamier,
R. H. Donnay. J. E. Dubois. ibid. 1971, 829.
[9] There are (mostly arene-substituted) olefins that are probably not chlorinated
via cationic intermediates, since the cis products are formed stereospifically.
Not much is known about the mechanism of this "direct" addition. which
probably involves the cis collapse of the x complex: a) S. J. Cristol, E R.
Stermitz, P. S. Ramey, J. Am. Chrm. Soc. 1958, 78, 4939; b) P. B. D. de la
Mare. N. V. Klassen. R. Koenigsberger, J. Chrm. SOC.1961. 5285; c) P. B. D.
de la Mare, R. Koenigsberger. ibid. 1964. 5327; d) R. C. Fahey. C. Schubert,
J. Am. Cheni. SOC.1965, 87, 5172; e) M. D. Johnson, E. N. Trachtenberg, J.
Chtmr. Sot.. 5 1968. 1018.
0570-0833/95/01(11-0052 $ 10.00 + .25/0
Anges.. Chem. Inr. Ed. Eng/. 1995, 34, N o . I
G Bellucci. K . Bianchini. C. Chiappe. V. R. Gadgil. A. P. Marchand. J. Urg
"hcrli 1993. 58. 1575.
G Bellucci. R Bianchini. C. Chiappe. F. Marioni. R. Ambrosetti. R. S . Brown.
H Slehock;i-Tilk. J. .41m Cliiwi. Soc. 1989. 111. 2640.
Brumoiiiuiii IOIIS of cyclohexene and cyclopentene generated by solvolysis can
d s o trmslki- B r i to other olefins: R. S. Brown. K. Gedye. H. Slebocka-Tilk.
J. M . Biischek. K. R . Kopecky. J A m . Clwm. SO<.1984. 106, 4515.
H . Slrhock;i-Tilk. K. G. Ball. K. S. Brown. J. Am. Cheni. Soc. 1985. 107. 4504.
M Cow. M. Persico. J. Tomasi. J .4m. Chmr. Soc. 1994, 116. 5373.
a ) E Kochanski. Quunfiun T/ic~)r),
Cliriiz. Riwi.1. 1980- 1982 1981. 2. 177-91,
h) I. Prisette. (i.
Seger. E. Kochanski. J Am. Chori. S i x . 1978. 100. 6942: c)
K . A. Pvirier. P. G . Mezey. K. Yates. I. G. Csizmadia. J. Mol. Srrucr. 1981. 88,
153. d ) A . c'. Hopkinson. M. H. Lien. K. Yates, I. G . Csizmadia. T l i ~ o rC/iiiii.
4c i(i 1977. 34. 785. ci S. Kimabe. T. Minato, S. Inaedki, J. Chon. Sn<.Chrni.
c',nlnil//ll. 1988. 532
According i o calculations by E. Kochanski. less than 0.04 electrons are translerrcd. The m i i n contribution to the stabilization of 3.7 kcalmol-' is due to the
di\per\ion energy: see ref. [15a]
L Fredin. B Nelander. J, Mol. Srrucr. 1973. 16, 205.
The intermedi.icy of the complex at room temperature in solution was also
denionstreteci by UV spectroscopy. ref. [ 7 ] ;see also a ) R. S. Brown, H. Slebock;i-Tilk. A J. Bcnnet. G. Bellucci. R. Bianchini. R. Ambrosetti. J A m .
(%miS o < . 1990. I/'. 6310. b) S. Fukuzumi. J. K. Kochi. J. Org. Cli(vn. 1981.
46.4I lO
A. c'. Lsgoii.
A. Kego. J Choii. So(c Firrorlii)
1990. 86. 1915.
[20] Some HX/Lewis base complexes were investigalcd hq i l i c wme method: a )
A. C. Legon, D. G. Lister. H. E. Warner J. .4m. CIrcni. S o , . 1992. //4.8177. b)
A. C. Legon. C A. Kego. A. L. Wallwork, J , C ' l i r m Plri 5 . 1992. Y7. 3051).
1211 H. 1. Bloemink. K . Hinds, A. C. Legon. J. C Thorn. .I C % P ~ Soi.. <'IIIw.
Coinmun 1994. 1321
[22] H. I. Bloemink. K. Hinds. A. C. Legon. J. C. Thorn. ~ ' l i ~ wPhi
i . 5 . Ll,rr. 1994.
223, 162.
[23] a)W. Jiger, Y. Xu, M. C. L. Gerry. J Pli),,\. C l r ~ w1993. V 7 . 3685: b ) S BImco.
A. C Legon. J. C. Thorn, J. Cliwii. S O I .Fiiriu/(i! I t u i i s 111 press.
1241 A. C. Legon. D G. Lister. .I.C. Thorn. J. C/rcm. SO( ~ ' l r ~ Coninrun.
1251 A . C. Legon. H. E. Warner, J. Chein. Phi.\. 1993. YH. 3x27
1261 A. C. Legon, J. C. Thorn. J. C % i w i . Soc. Furiirlui. Trim. 1993, SY. 4157
[27] H. Bloemink. K. Hinds, A. C'. Legon. J. C. Thorn. AI?,WII.Clicvr. 1994. 106.
1517; A n p i i , . Clieiii. I n [ . Ed. Dig/. 1994. 33.1512.
[ZS] H. I . Bloemink. K . Hinds. A . C. Legon. I. C. Thorn. .I Clieirr. S O ( . Cho?~.
Conimun. 1994. 1229.
[29] S. Blanco. A. C. Legon. J. C. Thorn, J Chenr. Soc. F ~ K / ( Trun.\.
1994. 90.
[30] Cf.: the force constant k , of a C-C bond is about 450 N m I .
[31] R. Herges. unpublished, MP2!6-31G* level.
[32] H. Hopf. R . Hiinel, P. G. Jones. P. Bubenitschck. .Aii!yii. C l u w 1994. 106,
1444; Atigrir. Cliivm lnr. E d E q / . 1994. 33. 1269.
[33] G. Bellucci. R. Bianchini. C. Chiappe. D. Lenoir. /?//I ( WI/CWIJW 011 P/i~%l-\k~/
Oryonre Chenii.\trj, IIUP,4CI Aug. 39. Padow. I t a l q . 1994.
Mononuclear Tris(arene)-Metal Complexes
Ulrich Zenneck"
The first x-arene complex, bis(q6-benzene)chromium, was
synthesized and characterized by Fischer and Hafner almost
forty years ago.[" Are any remarkable discoveries in this area of
chemistry still possible today? Since that pioneering work numerous attempts have been made to vary the bonding mode and
the number of arene ligands per metal atom. Although the former was achieved after some time, a mononuclear complex with
three arene ligands was synthesized for the first time only recentI Y . [ ~ ] Considering the intensity of research in organometallic
chemistry, this is quite remarkable and calls for closer examination.
If the electronic demand of a metal center is low, the cyclic
delocalization of the x electrons in the arene Iigands may be
interrupted such that only a portion of these electrons interact
with the metal. Arenes can coordinate, for example, as q4 ligands with loss of planarity, if this assures that the 18 valence
electron (VE) configuration of the transition metal complex is
not exceeded. This occurs relatively easily for condensed arenes,
but stable complexes with monocyclic arene ligands having
analogous t14 coordination are also known and have been structurally ~haracterized.'~]
The preferred q4 coordination of polycyclic ligands can be explained by the difference in the resonance
energies of the x electrons in the q4-coordinated arenes, which
are especially unfavorable for benzene and its derivati~es.[~l
[*] Prof. Dr. U . Zenneck
Institiit f i r Anorganische Chemie der Universitit Erlangen-Nurnberg
Eperlandatrassc 1. D-91058 Erlangen (Germany)
Telefax: lrit code + (9131)857367
q4-naphthalene complexes the bound 4x unit is electronically
almost completely separated from the 6x unit, which is reflected
in the spectroscopic and structural properties. Thus these compounds can be considered benzannelated 1.3-cyclohexadiene
complexes; the NMR and UV spectra of 1,3-cyclohexadiene
complexes closely resemble those of the related subunits in q4naphthalene complexes. The dihedral angles formed between
the bound and not bound portions of the ligands are also in the
same range for both types of compounds.[51
Generally the stability of q4-arene complexes increase within
a transition metal triad with the atomic number of the metal
center. The decomposition points of the (q4-benzene)(q6-benzene)metal(o) complexes (metal = Fe, Ru, 0 s ) serve as an example. The iron complex decomposes at roughly - 50 C, the
ruthenium complex at O T , and the dynamic behavior of the
osmium complex can be observed by NMR spectroscopy at up
to 100"C.[61In addition 3 d metals sometimes tend to avoid
forming q4-arene complexes; in contrast to the heavier homologues they give rise to paramagnetic 20 VE complexes with a
ligand bound if fashion. Thus bis(hexamethy1benzene)iron is a
20 VE sandwich complex with two unpaired electrons, whereas
the diamagnetic bis(hexamethylben2ene)ruthenium is a typical
x-Arene complexes can be synthesized by three principle
routes: ligand exchange maintaining the oxidation state of the
metal center (suitable for cationic complexes), ligand exchange
under reducing conditions (for cationic. anionic. and neutral
complexes), and direct synthesis by cocondensation of metal
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