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DALTON PERSPECTIVE J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 2939 Homonuclear multiple bonding in heavier main group elements Philip P. Power Department of Chemistry, University of California, Davis, California 95616, USA Recent developments in multiple bonding between heavier main group atoms are the focus of this review. Emphasis is placed on compounds with homonuclear bonds. It is clear that the Group 15 derivatives REER (E 5 P, As, Sb or Bi; R 5 alkyl or aryl ligand) display double bonding throughout the group.For the Group 14 species R2EER2 (E 5 Si, Ge, Sn or Pb, R 5 organo or related group), it is argued that, at present, only the silicon and certain germanium derivatives merit designation as ‘dimetallenes’. Data for multiply bonded heavier Group 13 compounds are currently very scarce. Nonetheless, the available structures of compounds such as (MR)n (M 5 Al, Ga, In or Tl; R 5 alkyl or aryl group; n 5 1–6) indicate weakness of the M]M interaction especially for the gallium, indium and thallium compounds where monomeric species are obtained readily.The M]M bond order in the dimers RMMR is apparently less than 1 but can be increased by reduction to give [RMMR]22 but it is probable that the overall M]M bond order remains less than 2. 1 Introduction The stabilization of heavier main group element compounds having multiple bonding has been a central research theme in organometallic chemistry for almost 30 years.Much of this work has focused on compounds of Groups 141 and 15,2 although heavier Group 13 element derivatives have attracted increasing attention within the past decade.3 It was recognized from the early work, particularly in Group 14, that the multiple bonding of the heavier elements diVers from that seen for the lightest group members. This diVerence was first illustrated experimentally by Lappert and co-workers through the synthesis 4 and structure 5 of the landmark compound [Sn{CH- (SiMe3)2}2]2. This was the first isolable species in which there was a possibility of multiple bonding between two heavier main group elements.In the solid state (Fig. 1) it has a trans-bent dimeric structure, with a pyramidal metal geometry and an out-of-plane angle (d, see below) of 418. The Sn]Sn distance, 2.768(1) Å 5b is ca. 0.03 Å shorter than the 2.80 Å in elemental tin 6 and is very similar to the 2.764(2) Å in (SnPh3)2.7 Nonetheless, the compound is dissociated 4,5 in solution [DH = 12.8 kcal mol21, DS = 33 cal K21 mol21 (cal = 4.184 J)] 8 into stannanediyl monomers which exist in the singlet form and its chemistry is Philip Power received a B.A.from Trinity College Dublin in 1974 and a D.Phil., under the supervision of M. F. Lappert, from the University of Sussex in 1977. After post-doctoral studies with R. H. Holm at Stanford he joined the faculty at the University of California, Davis where he is currently Professor of Chemistry.His main research interests involve the structural chemistry of organoalkali metal and organocopper compounds, low-coordinate transition-metal chemistry, multiple bonding in main group chemistry and the development of new ligands for the stabilization of low-co-ordination numbers, unusual oxidation states and multiple bonding in both transition-metal and heavier main group compounds. He is a recipient of fellowships from the A. P. Sloan and Alexander von Humboldt foundations.consistent with a monomeric structure.9 The germanium analog [Ge{CH(SiMe3)2}2]2 5b,10 also is monomeric in solution and dimeric in the solid with a Ge]Ge bond length of 2.347(2) Å which is ca. 0.09 Å less than a Ge]Ge single bond.6 The out-ofplane angle at germanium is 328 (cf. 418 for tin). Evidently, the Ge]Ge interaction is stronger than the corresponding one in its tin congener. In addition to these early results there were also extremely important achievements in heavier Main Group 15 chemistry that provided a foretaste of developments to come.The first stable phosphabenzene was reported as early as 1966.11 In the 1970s the isolation and characterization of stable phosphaimines, 12 phosphaalkenes 13 and phosphaalkynes 14 were published. (It should be borne in mind, however, that related compounds involving multiple bonding between sulfur and carbon or nitrogen or oxygen had been already well established, in some cases, for many decades.) These discoveries heralded the explosive growth in the 1980s when a large variety of stable molecular compounds with multiple bonding between two heavier main group elements or between a heavier and lighter main group element were reported.Examples include Si]] C,15 Si]] Si,16 Ge]] C,17 Ge]] Ge,10,18 Si]] N,19 Ge]] N,20 Sn]] C,21 Si]] P,22 Ge]] P,23 Sn]] P,24 P]] P,25 P]] As,26 As]] As,27 P]] B,28,29 As]] B30 and As]] ] C bonds.31 Recent work in the 1990s has aVorded more examples of dimeric tin species 32–35 related to Lappert’s original [Sn{CH- (SiMe3)2}2]2 compound. These and related germanium18,35–40 and lead 33 analogues are listed in Table 1.5b,10,18,35–40 Structural details of the anion [{Sn(C6H3Trip2-2,6)}2]241 (Trip = C6H2Pri 3- 2,4,6), a singly reduced valence isomer of a ‘distannyne’, are also given in Table 1.It has a formal bond order of 1.5 and a Sn]Sn bond length of ca. 2.81 Å. The trigermanium ring compounds [{Ge(SiBut 3)2}2Ge(SiBut 3)2] 42a [{Ge(SiBut 3)}3]- [BPh4] 42b and [Ge(C6H3Mes2-2,6)]3 ? (Mes = C6H2Me3-2,4,6) 42c are also included.In addition, the first structural details for a (PbR2)2 compound, [Pb{Si(SiMe3)3}{C6H2(CF3)3-2,4,6}]2,33 with a Pb]Pb interaction, are listed. Besides these compounds, heteronuclear heavier Group 14 derivatives with Si]] S,43 Ge]] O,44a Ge]] S,44b,45,46 Ge]] Se,44c,46 Fig. 1 A drawing of the structure of [Sn{CH(SiMe3)2}2]2 illustrating its trans-bent configuration (ref. 5b)2940 J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 Table 1 Selected structural data for diorgano-germanium(II), -tin(II) and -lead(II) dimers and related species Compound [Ge(C6H3Et2-2,6)2]2 [Ge(C6HMe3-2,3,4-But-6)2]2 [Ge(SiMePri 2)2]2 [Ge(SiPri 3)2]2 [Ge(Mes)(C6H3Pri 2-2,6)]2 [Ge{CH(SiMe3)2}2]2 [GeCl(C6H3Mes2-2,6)]2 {GeN(But)(CH2)3N(But)SiN(But)(CH2)2N(But)}2 [{Ge(SiBut 3)2}2Ge(SiBut 3)2] [{Ge(SiBut 3)}3][BPh4] [Ge(C6H3Mes2-2,6)]3 ? K[{Ge(C6H3Mes2-2,6)}3] [Sn{CH(SiMe3)2}2]2 [Sn{Si(SiMe3)3}2]2 [Sn{C6H2(CF3)3-2,4,6}{Si(SiMe3)3}]2 [Sn(C6HMe3-2,3,4-But-6)2]2 [Sn{C6H2(CF3)3-2,4,6}2]2 [K(THF)6][{SnC6H3Trip2-2,6}2] [Pb{C6H2(CF3)3-2,4,6}{Si(SiMe3)3}]2 M]M/Å 2.213(2) 2.2521(8) 2.267(1) 2.298(1) 2.301(1) 2.347(2) 2.443(2) 2.451(2) 2.239(4) 2.226(4) 2.35(7) 2.422(2) 2.768(1) 2.8247(6) 2.833(1) 2.910(1) 3.639(1) 2.8123(9) 3.537(1) d*/8 12 000 36 32 39 41.3 ———— 41 28.6 41.5 21.4, 64.4 46 95.20 40.8 g*/8 10 20.4 6.5 16.4 700 42.3 ————0 63.2 0 —000 Ref. 18 36 37 37 38 5b, 10 39 40 42a 42b 42c 42c 5b 32 33 34 35 41 33 * The angles d and g are represented below: M M d M g Ge]] Te,44d,46 Sn]] S,47,48 Sn]] Se 47–49 and Sn]] Te 49 multiple bonds have been characterized. The compound (h5-C6H5)(CO)2MoGe- C6H3Mes2-2,6, in which there is a Mo]Ge triple bond, has also been reported.50 The characterization of the first stable Sb]] Sb51 and Bi]] Bi 52 bonds has been described and structural information for these and their arsenic and mixed Group 15 congeners 26,27,51–54 is in Table 2.Triply bonded phosphorus– transition-metal compounds have also been characterized.55 Another development of the 1990s has been the expansion of the range of multiple bonds to aluminum and gallium.3a Some structural parameters for homonuclear Group 13 multiple bonded compounds are included in Table 3.56–74 In addition, stable complexes featuring heteronuclear Ga]] Se,75 Ga]] Te 75 and In]] Se 76 double bonds have been synthesized.Table 3 includes structural data for the radical species [R2MMR2]2 [M = Al or Ga; R = CH(SiMe3)2 or Trip] which have a formal M]M bond order of 1.5 and the ring compound [{Ga(C6H3Mes2-2,6)}3]22 which has a 2 p-electron, three-membered ring with a formal Ga]Ga bond order of 1.33. It is related to the above-mentioned, isoelectronic cyclotrigermanium cation [{Ge(SiBut 3)3}3]1.42b The radical [(But 3Si)GaGa(SiBut 3)2]? 63 with a formal Ga]Ga bond order of 1.5 and the dimer Na2[{Ga(C6H3Trip2-2,6)}2],74 which was stated to have a Ga]Ga triple bond, are also listed.More recently, it was reported that the related species (OC)4FeGa- (C6H3Trip2-2,6) 77 also had a Ga]Fe triple bond. The bonding description of the last two compounds, in particular, has generated considerable discussion. Some say that Table 2 Selected structural parameters for organo-arsenic, -antimony and -bismuth double bonded compounds Compound [Mes*PAs{CH(SiMe3)2}] [MesPAs(C6H3Trip2-2,6)] [Mes*AsAs{CH(SiMe3)2}] [AsC(SiMe3)3]2 [As(C6H3Trip2-2,6)]2 [MesPSb(C6H3Trip2-2,6)] [Sb{C6H2[CH(SiMe3)2]3-2,4,6}]2 [Sb(C6H2Trip2-2,6)]2 [Bi{C6H2[CH(SiMe3)2]3-2,4,6}]2 E]E/Å 2.124(2) 2.134(2) 2.224(2) 2.244(1) 2.285(3) 2.335(2) 2.642(1) 2.664(2) 2.8206(8) E]E]C*/8 101.2(2), 96.4(2) 96.7(2), 101.5(2) 93.6(3), 99.9(3) 106.3(2) 96.40(22) 95.7(3), 100.9(2) 101.4(1) 98.58(32) 100.5(2) Ref. 26 53 27 54 53 53 51 53 52 * Where two angles are given they are listed in the order in which the atoms appear in the formula.the Ga]Ga bond is a double one on the basis of its bond length or its trans-bent geometry, DFT78 and ab initio calculations.79a The possibility of single bonding has also been suggested.79b In contrast, others have supported triple bonding on theoretical grounds.80,81 The claim for Fe]Ga triple bonding has also been questioned 82 on the basis of DFT calculations and a comparison of IR data with related aluminum83 and indium 84 compounds which imply essentially negligible p-bonding between the Group 13 element and the transition metal.The bonding of the Group 14 compounds1 in Table 1 has also been the subject of considerable discussion (see below). This is based on the fact that none of the tin or lead compounds has a distance much shorter than a single bond and all are dissociated to monomers in solution. They deviate from the planar geometries expected in ethylene analogs and have large out-of-plane angles at the metals.Three of the eight germanium dimers have similar characteristics. As a result of this behavior (which is at such variance to that normally seen in alkenes) several interesting questions arise. For example, is it justified to refer to the tin compounds in Table 1 as distannenes, thereby suggesting a behavior normally associated with corresponding alkenes? Is it expedient to refer to Na2[{Ga(C6H3- Trip2-2,6)}2] as a (triply-bonded) digallyne when it has a Ga]Ga distance similar to some Ga]Ga single bonds 61,62 as well as a non-linear structure? The object of the rest of this perspective is to examine some current views of these questions.The discussion is focused primarily on homonuclear bonds to avoid the complicating eVect of ionic factors associated with heteronuclear species. Furthermore, emphasis is given to recent developments in compounds of the third, fourth and fifth rows where deviations in the structural and chemical properties from compounds of the first row are the most marked. 2 Phosphorus, Arsenic, Antimony and Bismuth Multiple Bonding It is, perhaps, easiest to begin with the least controversial of the heavier main group multiply bonded compounds, the Group 15 derivatives of formula REER (E = P, As, Sb or Bi; R = alkyl or aryl group).2 Currently, there are structural data for about 20 diphosphenes (which generally have trans-planar structures) as well as several h1-diphosphene transition-metal complexes inJ.Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 2941 Table 3 Metal–metal bond lengths and torsion angles in structurally characterized Group 13 tetraorganodimetallanes and related species and their reduced analogs Compound [Al{CH(SiMe3)2}2]2 (AlTrip2)2 [Ga{CH(SiMe3)2}2]2 (GaTrip2)2 [GaCl{Si(SiMe3)3}2]2 [Ga{CH(SiMe3)2}DPPD]2 c [Ga{CH(SiMe3)2}N3Ph]2 [Ga(TMP)2]2 d [GaB4H4(CSiMe3)2]2 [GaN(But)CHCHN(But)]2 [GaI(C6H3Trip2-2,6)]2 [(But 3Si)GaGa(SiBut)3)2] [In{CH(SiMe3)2}2]2 (InTrip2)2 [In{C6H2(CF3)3-2,4,6}2]2 [In{N(But)SiMe2}2N(But)]2 [In(SiBut 3)2]2 [Tl{Si(SiMe3)3}2]2 [Tl(SiBut 3)2]2 [{Al{CH(SiMe3)2}2}2]2 [{Ga{CH(SiMe3)2}2}2]2 [{AlTrip2}2]2 [{GaTrip2}2]2 Na2[{Ga(C6H3Mes2-2,6)}3] Na2[{Ga(C6H3Trip2-2,6)}2] M]M/Å 2.660(1) 2.647(3) 2.541(1) 2.515(3) 2.505(4) 2.44(1) 2.457(9) 2.525(1) 2.340(2) 2.333(1) 2.511(3) 2.420(1) 2.828(1) 2.775(2) 2.744(2) 2.768(1) 2.922(1) 2.914(5) 2.966(2) 2.53(1) 2.401(1) 2.470(2) 2.343(2) 2.441(1) 2.319(3) M]C/Å 1.982(3) 1.996(3) 1.995(5) 2.008(7) 2.395(5) b 1.995(5) 1.977(2) 1.901(4) e — 1.836(4) e 1.994(9) — 2.194(5) 2.184(7) 2.21(1) 2.12(1)–2.29(8) e 2.778(4) b 2.675(2) b 2.789(12) b 2.040(5) 2.059(4) 2.021(1) 2.038(10) 2.037(3) 2.044(20) d a/8 ª0 44.8 ª0 43.8 ——— 31 — 90 0— ª0 47.8 85.9(5) — 90 78.1 90 00 1.4 15.5 —— Ref. 56 57 58a 59 60a 58b 58c 60 61 62 64 63 65 66 67 68 69 70 69 72 71 57 59 73 74 a Angle between the perpendiculars to the M]C2 co-ordination planes. b M]Si distances. c DPPD = 1,3-Diphenylpropane-1,3-dionate. d TMP = 2,2,6,6- Tetramethylpiperidine.e M]N distances. which the P]P double bond is conserved.2d,e There are also many metallodiphosphenes in which the phosphorus organic substituent is replaced by a transition-metal fragment.2d The range of P]P distances in organodiphosphenes lies between 1.985(2) 85 and 2.034(1) Å.25 This may be compared to the approximate single P]P bond distance of 2.22 Å.86a Thus, the percentage diVerence between singly and doubly bonded P]P moieties lies between ca. 7.5 and 9.8%. This impressive margin probably reflects not only P]P p-bond formation, but also a change in s hybridization at phosphorus which is thought to account for up to half the observed contraction.87 The shortening is ca. 50–60% of the 15.2% diVerence observed between hydrazine N2H4 (N]N 1.45 Å) 86b and the diimine N2H2 (N]N 1.23 Å).86c It seems probable that this margin is enhanced in view of the long single N]N s bond which is thought to be caused by interelectronic repulsion. Nonetheless, there seems to be little doubt that the P]P p bond is much weaker than an N]N p bond and this view is supported by spectroscopic and thermochemical data in Table 4 88 which features a comparison of the s- and p-bond strengths of some homonuclear maingroup element–element bonds.88,89a It can be seen that there is more than 50% decrease in p-bond strength between the first Table 4 Relative energies (kcal mol21) of s and p bonds in homonuclear main group diatomic species a B]B Al]Al Ga]Ga In]In Tl]Tl 70 b 36 b 32 b 23 b 2 b C]C Si]Si Ge]Ge Sn]Sn Pb]Pb 81/62 a 47/28 a 39/26 a 35/11 c 23(33)/— c,d N]N P]P As]As Sb]Sb Bi]Bi 38/94 a 48/34 a 35/28 a 31/20 c 21/10 c,e a These values are generated by using methods in ref. 88 which are in part abstracted from spectroscopic data in ref. 89(a). b These values were obtained from ref. 89(a) and they represent single bonds. c These data were abstracted from ref. 89(a) and from single bond values in ref. 89(b). d The Pb]Pb single bond value (in parentheses) is from ref. 89(b); the first value is from 89(a) and corresponds to a ‘double’ bond. e These values were estimated from the 41 kcal mol21 value for the diatomic Bi2 [see ref. 89(a)] and by assuming an approximate 2 : 1 ratio for the strengths of the s and p bonds. and second row. It is also notable that the p-bond strengths are greater in Group 15 than in Group 14, possibly as a result of their smaller size.By using bond dissociation energies for the diatomic triply bonded Group 15 molecules 89a and single bond strengths, it is possible to distinguish the s and p components of the bonding within this group. For phosphorus, although the P]P p bond is much weaker than the N]N p bond, the s:p bond strength ratio (48 : 34) in a P]P double bond is comparable to the ratio (81 : 62) in carbon. From this perspective, the P]P double bond is correctly regarded a full-fledged double bond.The P]P p-bond strength, which may be represented by the rotational barrier for the E–Z isomerization, was calculated 90 to be ca. 33.5 kcal mol21. This value is in good agreement with the 34 kcal mol21 estimated in Table 4. Laser irradiation studies of the E–Z isomerization in (PMes*)2 (Mes* = C6H2But 3-2,4,6) aVorded a lower energy barrier, ca. 20 kcal mol21.91 The s:p bond strength in arsenic is 35 : 28. As expected, both bonds are weaker than their phosphorus counterparts, nevertheless, the relative p-bond strength is slightly higher for arsenic.The bond lengths in diarsenes are in the range 2.224(2) 27–2.285(3) Å.53 With 2.44 Å as the As]As single bond distance,86d this represents a shortening of 6.3–8.0%. This margin is slightly less than that observed in the diphosphenes. Nonetheless, it is fully consistent with the presence of As]] As double bonds. The currently known range of structurally characterized distibenes and dibismuthenes is limited to two of the former51,53 and just one of the latter.52 The Sb]Sb distances in the two distibenes are 2.642(1) 51 and 2.664(2) 53 Å which represent shortenings of 6.94 and 6.16% respectively.86e It may be noted that both s and p bonds of the Sb]Sb double bond are weaker than in the arsenic case but the weakening is greater in the case of the p bond.A similar trend is observed for the dibismuthene. 52 In this case the Bi]Bi p-bond strength is estimated to be only about half that of the Sb]Sb p bond.In the dibismuthene the amount of shortening is 5.7%.86f Clearly, the percentage shortening diminishes in these heavier metals. This is consistent with the lower p-bond energies in Table 4.2942 J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 3 Silicon, Germanium, Tin and Lead Multiple Bonding Moving now to the Group 14 elements, it can be seen that the estimated p-bond strengths are less than those in Group 15. In contrast, the s-bond strengths for the Group 14 elements are generally higher (the P]P and Si]Si bonds have about equal strength) than in the Group 15 elements.The stronger p bonds in the pnictides may be explained primarily on the basis of the smaller size of these elements, whereas the weakness of the Group 15 s bonds has often been thought to be due to interelectronic repulsions between the lone pairs which is most pronounced in the case of the very weak N]N single bond. In Group 14 homonuclear double bonding primarily concerns the species R2MMR2 (M = Si–Pb; R = alkyl or aryl group).For silicon about a dozen compounds have been structurally characterized.1i The Si]Si distances vary from 2.251(1) 92a to 2.138(1) Å 92b which represents a shortening in the range 7.3–8.6%. As with the Lappert ditin compound some disilenes adopt trans-bent structures but the outof- plane bending is very much less (maximum published value 188) 93 and many have essentially planar geometry like most alkenes. 94 The activation enthalpy for cis–trans isomerization, which is considered to be a measure of Si]Si p-bond strength, ranges from 25.4 to 30.3 kcal mol21.95,1 j This is ca. 50% of the p-bond strength in alkenes and in good agreement with the value in Table 4. In one case, the compound [Si(Mes)C6H2{CH(SiMe3)2}3-2,4,6]2,96a,b there is dissociation to monomers under relatively mild conditions. The DHdiss is ca. 26 kcal mol21. The dimer has a lengthened Si]Si doubly bonded distance of 2.228(1) Å with out-of-plane angles in the case of the Z isomer of 9.4(3) and 14.6(3)8.The structures of eight (GeR2)2 compounds are known (Table 1). The Ge]Ge distances and out-of-plane angles range from 2.213(2) 18 to 2.451(2) Å40 and from 036,37 to 42.38.40 Thus, the Ge]Ge bond shortenings range between 9.3 and 0%. Two of the three compounds having the longest Ge]Ge bonds are dissociated to monomers in solution, which underlines the weakness of their association. Another feature of the germanium compounds is that the majority (five out of eight) show substantial out-of-plane angles and all display either a substantial out-of-plane or twist angle.The geometrical distortion therefore is greater than in the silicon compounds. The data in Table 4 indicate a p-bond strength of ca. 26 kcal mol21 which is close to that predicted theoretically 97 but higher than the enthalpy of activation (ca. 22 kcal mol21) of the interconversion of the E–Z isomers of [Ge(Mes)(C6H3Pri 2-2,6)]2.38 The lower value of this compound is in harmony with the elongation [Ge]Ge = 2.301(1) Å] and high out-of-plane angle of 368 which are probably caused by steric crowding.Furthermore, the 22 kcal mol21 barrier is less than the ca. 26 kcal mol21 in its silicon analog [Si(Mes)(C6H3Pri 2-2,6)]2.38 A recent variable temperature UV/ VIS spectroscopic study 96c of [Ge(MeS)(C6H2{CH(SiMe3)2}3)]2 (crystal structure currently unknown) yielded an enthalpy of dissociation of 14.7 kcal mol21, about half that of the silicon analog.96a,b Considerable Ge]Ge multiple bond character is also seen in the cyclotrigermanium cation [{Ge(SiBut)3}3]1 which has an average Ge]Ge distance of 2.226(4) Å.42b Short Ge]Ge distances, average 2.35(7) Å, are also observed in the cyclotrigermanium radical [{Ge(C6H3Mes2-2,6)}3]?.42c A sharp decrease in the p-bond energy is predicted in Table 4 upon descending the Group 14 elements from germanium to tin.The Sn]Sn p-bond strength is estimated to be just 11 kcal mol21 (cf. Ge]Ge p-bond strength = 26 kcal mol21) whereas the strength of an Sn]Sn s bond is marginally less than that of its Ge]Ge counterpart (35 vs. 39 kcal mol21). The decrease in p-bond strength is more abrupt than that in the corresponding Group 15 elements between arsenic and antimony. It is diYcult to explain this diVerence on the basis of sizes since antimony and tin have similar radii as do arsenic and germanium.The weakness of the Sn]Sn bonding in the five (SnR2)2 structures in Table 1 is supported by the fact that all compounds are dissociated in solution and the shortest tin–tin distances are similar to that of a single bond.5b,32,33 The longest Sn]Sn interaction is in [Sn{C6H2(CF3)3-2,4,6}2]2,35 Sn]Sn = 3.639(1) Å, which is ca. 0.8 Å longer than a single bond. The trans-bent structure is observed in all compounds but there is no correlation between the out-of-plane angle and Sn]Sn bond length in the compounds seen here.The structure of one dimeric (PbR2)2 species with a Pb]Pb interaction is currently available.33 In that compound, [Pb{C6- H2(CF3)3-2,4,6}{Si(SiMe3)3}]2, the Pb]Pb distance is 3.537(1) Å and the out-of-plane angle is 40.88. This distance is much longer than the single Pb]Pb bond in (PbMe3)2 [2.88(3) Å] 98 which indicates that the Pb]Pb interaction is very weak. This is in agreement with the data in Table 4, which predict a Pb]Pb bond energy of 23 kcal mol21 89a in the dimer Pb2 (putatively a Pb]Pb double bond) much less than the corresponding value (35 kcal mol21) for its tin analog.A problem arises, however, when this value (23 kcal mol21) is compared to that for a single Pb]Pb bond in (PbMe3)2 (33 kcal mol21).89b,98 This implies that a single Pb]Pb bond is stronger than a multiple one. This apparently absurd result is consistent with the observation that the ‘multiple’ Pb]Pb bond in [Pb{C6H2(CF3)3-2,4,6}{Si- (SiMe3)3}]2 33 is much longer than a single bond whereas the analogous tin species [Sn{C6H2(CF3)3-2,4,6}{Si(SiMe3)3}]2 has a bond length similar to a single Sn]Sn bond.33 One other class of homonuclear multiply bonded heavier Main Group 14 species needs to be considered.These are compounds of formula RMMR (M = Si, Ge, Sn or Pb; R = organo group), which are heavier analogues of the alkynes. Unfortunately, experimental details of only one related compound currently exist.The radical ion [{Sn(C6H3Trip2-2,6)}2]2, which is a singly reduced analogue of a neutral tin species of the formula (SnR)2, has a Sn]Sn distance near 2.80 Å and an Sn]Sn]C angle of ca. 958 (Fig. 2).41 This structure suggests that there is a lone pair at each tin occupying orbitals high in s character with p orbitals being used for Sn]Sn and Sn]C s bonding. The remaining p orbitals, one at each tin, may overlap to form a p orbital occupied by a single electron.Formally, the bond order is 1.5 but the Sn]Sn distance is the same as a single bond. Removal of the p electron should result in a slightly longer Sn]Sn bond but the trans-bent structure with lone pairs at each tin should be preserved. Fig. 2 Schematic drawing of [K(THF)6][{Sn(C6H3Trip2-2,6)}2]2 (ref. 41)J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 2943 4 Boron, Aluminum, Gallium, Indium and Thallium Multiple Bonding Homonuclear multiple bonding between the Group 13 elements is a relatively recent development.3a Historically, such compounds were more noted for their electron deficiency and consequent absence of suYcient numbers of electrons to form multiple bonds.This was particularly true for derivatives of boron. Nonetheless, there were a number of reports that indicated that the problem of electron deficiency could be remedied. For example, multiply bonded B]B moieties were generated in various metal complexes of reduced six-membered quasi-aromatic rings.99 In addition, the structures of a number of three-membered delocalized ring compounds containing B]B units have been synthesized.100 In these, the stabilized ring p orbitals provide the impetus for the delocalization incorporating the B]B unit.Boron–boron bonds as short as 1.58 Å (cf. single B]B bond ca. 1.71 Å) have been reported.101 Short B]B distances (ca. 1.6 Å) have also been observed in the structures of several transition-metal borides which contain onedimensional polyacene-type boron chains as part of a threedimensional metallic lattice.102 The common feature of these compounds is that B]B units with adjacent empty p orbitals as part of rings or chains are reduced to form multiple bonds between the boron atoms.The reduction of acyclic, molecular B]B bonded species had to await the synthesis of suitable compounds for reduction. The simplest are the species (BR2)2 (R = alkyl or aryl groups) which are only stable if the substituents are large.They were first reported in 1980 103 and it was shown that they could undergo a 1-electron reduction to the species [(BR2)2]2.104 Solution EPR data showed that the unpaired electron occupied a p orbital formed by overlap of two adjacent boron p orbitals for a formal B]B bond order of 1.5. However, further reduction of tetraalkyl diboron species was not achieved. Tetraaryl diboron compounds permitted double reduction as seen in the dianion [Mes2BB(Ph)Mes]22.105a It has a shortened B]B bond of 1.636(11) Å and an almost planar B2{Cipso}4 array consistent with the presence of a formal B]B double bond.The dianion [{B(NMe2)Ph}2]22 has a similar B]B distance of 1.631(9) Å as well as a planar core.105b The B]B bonds are ca. 0.07–0.08 Å shorter than in their neutral precursors. The shortening is not as great as that in corresponding ethylene species possibly as a result of coulombic repulsion. A comparison with [{B[C- (SiMe3)2]Mes}2]22, which has a much longer single bond [B]B 1.859(6) Å] between two negatively charged borons, seems to bear out this view.106 Theoretical data for diborane(4) dianions are scant but one paper suggests that the B]B p bond may be quite strong.107a Extension of these methods to aluminum and gallium has resulted in the synthesis of several monoreduced anions (Table 3) of tetraorganodimetallanes similar to the corresponding boron species.57,59,71,72 Stable multiply bonded dianion products from further reduction have not yet been obtained.The monoreduced species have a formal M]M (M = Al or Ga) bond order of 1.5. The M]M bond shortening is in the range 0.13– 0.18 Å (ca. 7–8%) and the torsion angle between the metal coordination planes is decreased. The multiple bonding is not complicated by the presence of associated alkali-metal counter cations which are separated from the anion by solvent coordination. The solution EPR spectra of the monoanions show that the unpaired electron is equally coupled to two metal nuclei and that the magnitude of the couplings is consistent with the location of the electron in a p orbital. For the aluminum compounds there is good agreement with theoretical data 107b which predict an Al]Al distance of 2.478 Å and a torsion angle of 6.38 in the hypothetical compound [(AlPh2)2]2; cf.Al]Al 2.470(2) Å and torsion angle 1.48 in [(AlTrip2)2]2. Reduction of the substituted m-terphenyl gallium dihalides GaCl2(C6H3Mes2-2,6) or GaCl2(C6H3Trip2-2,6) gave Na2- [{GaC6H3Mes2-2,6}3] 73 or Na2[{GaC6H3Trip2-2,6}2],74 which were the first examples of two new compound classes.The former species has a core composed of a Ga3 ring with equal Ga]Ga distances near 2.44 Å. It may be regarded as a delocalized 2 p-electron system that conforms to the Hückel rule. It has a formal Ga]Ga bond order of 1.33. This conclusion is supported by calculations,73b although the Ga]Ga distances (ca. 2.5 Å) in hypothetical species such as Na2[(GaH)3] or K2[(GaH)3] are significantly longer than those measured experimentally.73b This may indicate that Na1–aromatic ring interactions could play a role in shortening the Ga]Ga bonds. The digallium compound Na2[{GaC6H3Trip2-2,6}2] also has no precedent in heavier main group element chemistry.74 It crystallizes as an ion triple with interactions between the Na1 ions and the ortho-aryl substituents (Fig. 3). The dianion is isoelectronic to the corresponding, unknown, neutral germanium compound and it was stated to have a Ga]Ga triple bond.However, its geometry is not linear and Ga]Ga]C angles of ca. 1318 are observed. The Ga]Ga bond, 2.319(3) Å, is the shortest reported to date, but it can be seen from Table 3 that this distance is similar to the 2.343(2) Å 59 in the anion [(GaTrip2)2]2 (Ga]Ga bond order 1.5) or the 2.340(2) Å in the less crowded species [GaB4H4(CSiMe3)2]2 61 (Ga]Ga bond order 1) or the 2.333(1) Å in [GaN(But)CHCHN(But)]2 62 (Ga]Ga bond order 1).Unfortunately, no stable dimeric, neutral compounds of the type RMMR (M = Al or Ga; R = alkyl or aryl) solvent or separated ion pairs of the type [M9Ln]2[RMMR] (M9 = alkali metal; L = Lewis base donor) are known at present. Structural data for these would be of interest since the Ga]Ga bond in the dianion [{GaC6H3Trip2-2,6}2]22 would be illuminated by the synthesis of the corresponding unreduced species, the putative dimer [Ga(C6H3Trip2-2,6)]2 and the influence (if any) of the metal counter cations on the Ga]Ga bond distance could also be determined.The M]M bonds in compounds of the general formula (MR)n (M = Al or Ga; R = alkyl or aryl; n = 4–6) are usually weaker than single bonds. Evidence for this comes from structural studies 108–117 of species such as [Al(h5-C5Me5)]4,108 [Ga{C(SiMe3)3}]4,109a,b [Ga(h5-C5Me5)]6 110 and monomeric GaTpBut 2 [ TpBut 2 = tris(3,5-di-tert-butylpyrazolyl)hydroborate] 111 (Table 5).The tetramers, which have electron deficient M4 frameworks, display somewhat lengthened Al]Al and Ga]Ga bonds of ca. 2.77 and 2.69 Å and [Al(h5-C5Me5)]4,108a [Ga- {C(SiMe3)3}]4 and [Ga{C(SiMe2Et)}]4 109c dissociate to monomers in dilute benzene solution 109 and in the vapor phase. However, the recently synthesized tetramer [Al(SiBut 3)]4 108b has a shorter average Al]Al distance of 2.604(4) Å which demon- Fig. 3 Schematic drawing of Na2[{Ga(C6H3Trip2-2,6)}2] (ref. 74)2944 J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 Table 5 Selected structural data for some metal–metal bonded and related compounds of the formula (MR)n (M = Al, Ga, In or Tl; R = alkyl or aryl group; n = 1–6) Compound [Al(h5-C5Me5)]4 (solid) [Al(h5-C5Me5)] (vapor) [Al(SiBut 3)]4 (solid) [Ga(h5-C5Me5)]6 (solid) [Ga(h5-C5Me5)] (vapor) [Ga{C(SiMe3)3}]4 (solid) [Ga{C(SiMe3)3}] (vapor) GaTpBut 2 (solid) [In(h5-C5Me5)]6 (solid) [In(h5-C5Me5)] (vapor) [In{C(SiMe3)3}]4 (solid) [In{h5-C5(CH2Ph)5}]2 (solid) [In(C6H3Trip2-2,6)] (solid) [Tl{C(SiMe3)3}]4 (solid) [Tl{h5-C5(CH5Ph)5}]2 (solid) [Tl(h5-C5Me5)] (vapor) [Tl(C6H3Trip2-2,6)] (solid) [Tl{N(SiMe3)C6H3Pri 2-2,6}] (solid) Structure type Tetrahedral Al4 Monomer Tetrahedral Al4 Distorted octahedral Ga6 Monomer Tetrahedral Ga4 Monomer Monomer Octahedral In6 Monomer Tetrahedral In4 In]In trans-bent Monomer Tetrahedral Tl4 Tl]Tl trans-bent Monomer Monomer Tetramer, planar Tl4 M]M/Å 2.769(4) — 2.604(4) 4.073(2), 4.173(3) — 2.688(6) —— 3.942(1)–3.963(1) — 3.002(1) 3.631(2) — 3.322(1)–3.638(1) 3.632(1) —— 4.06 Ref. 108a 117a 108b 110 117b 109a 109b 111 112 117c 113b 114a 116a 113c 114b 117d 116b 117e strates the very important role of the electronic and steric properties of the ligand in determining bond strengths. The hexamer [Ga(h5-C5Me5)]6 has much longer Ga]Ga distances of 4.073(2) and 4.173(3) Å.110 An unusual gallium species, [(But 3Si)- GaGa(SiBut 3)2] (Fig. 4), which features a relatively short Ga]Ga bond of 2.420(1) Å, has also been reported.63 The formal Ga]Ga bond order is 1.5. The Ga]Ga bond is composed of a 2-electron s bond and the unpaired electron occupies a p orbital with coupling to two diVerent galliums. A number of indium(I) and thallium(I) compounds with metal–metal interactions are known. The metal–metal bonding appears to be very weak in all currently known cases (Table 5). This is exemplified by the indium structures of the hexameric [In(h5-C5Me5)]6 112 (In]In ca. 3.95 Å), tetrameric [In{C- (SiMe3)3}]4 113 [In]In 3.002(1) Å], and the dimers [M{h5- C5(CH2Ph)5}]2 114 (M = In or Tl; In]In and Tl]Tl 3.63 Å) which have In]In interactions that are much longer than the distances in the In]In singly bonded compounds (InR2)2 [2.768(1)–2.828(1) Å]. Indeed, it has been possible to crystallize monomeric species such as MTpBut 2 (M = In 115a or Tl 115b) and one-co-ordinate metal species of formula [M(C6H3Trip2-2,6)] (M = In 116a or Tl 116b).It is also notable that all the [M(h5- C5Me5)] (M = Al, Ga, In or Tl) derivatives are monomeric in the vapor phase.112,117 Unfortunately, no estimates of Group 13 p-bond strengths are available. However, it is probable that they are less than values given for the p bonds in Group 14 or 15 compounds owing to the larger Group 13 element sizes. If it is assumed that there is optimized orbital overlap in the Group 13 species, Fig. 4 Schematic drawing of [(But 3Si)GaGa(SiBut 3)2] (ref. 63) values for p bonds that are ca. half to two thirds those of s bonds seem warranted by extrapolation of the Group 14 and 15 data. 5 Bonding and Bond Order Bonding It is already clear that the bonding in the heavier main group element compounds diVers fundamentally from that in the lighter members of the group. Perhaps only in the heavier Group 15 derivatives are the simple bonding models used for the lighter nitrogen congeners usefully applicable.In the Group 13 and 14 compounds, however, there are major distinctions in the molecular architecture between the heaviest and lightest element derivatives. These geometrical changes, which mirror a changed electronic structure, have necessitated a diVerent bonding description of the compounds. It is these bonding models, or, more accurately, the interpretation of them, that give rise to controversy. Prior to discussing these it is perhaps worth quoting a definition for a chemical bond provided by Pauling who said that:118 ‘there is a chemical bond between two atoms or groups of atoms in case that the forces acting between them are such as to lead to an aggregate of suYcient stability to make it convenient for the chemist to consider it as an independent molecular species’.The advantage of this definition is that it does not assume any theory, not even the existence of electrons or orbitals. In essence, chemical bonding is assumed to exist when it is justified by the physical and chemical behavior of the species in question. If the Pauling criterion is applied to the Group 15 molecules in Table 2, it is clear, beyond doubt, that they are stable molecular entities of the formula REER (E = P, As, Sb or Bi).There is no dissociation to monomeric units of formula ER either in solution or in the vapor phase. In addition, the E]E bonds are ca. 6–10% shorter than corresponding single bonds, and there are substantial barriers to cis–trans isomerization.2 There is good agreement between calculations 2,78,119 and experimental findings and electron density measurements on a diphosphene have confirmed the existence of strong s and p bonds.120 Apart from these data there is an extensive chemistry of compounds related to those in Table 2 which show that E]E bonding is retained in many diVerent reactions.2 When the simple bonding definition is applied to the heavier diorgano–Group 14 element species (Table 1) a more complex picture emerges.121 Clearly, the disilenes are stable chemical species and chemical 1 and NMR data 122 support their double bonding.Even where one of them has been shown to dissoci-J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 2945 ate,96 the enthalpy of dissociation is suYciently high that the molecule remains mostly dimeric in solution at room temperature. With germanium, however, two39,40 of the eight (GeR2)2 species have Ge]Ge bond lengths similar to a Ge]Ge single bond and two5b,10,40 are monomers in solution.Five compounds18,36 –38 do not dissociate readily, have short Ge]Ge bonds with significant rotation barriers, undergo reactions that retain the Ge]Ge moiety,1h,k,121 and are, essentially, germanium analogs of alkenes. The structures of the five (SnR2)2 compounds show that three compounds5b,32,33 have an Sn]Sn distance comparable to the Sn]Sn single bond (2.80 Å) in elemental tin. In the other two, the Sn]Sn distances are significantly longer.34,35 More importantly, all the compounds are dissociated to monomers in solution at room temperature.Since the enthalpy of association of [Sn{CH(SiMe3)2}2]2 is 12.8 kcal mol21, and this compound 8 has the shortest Sn]Sn bond in the (SnR2)2 series, it may be assumed that the other four compounds have enthalpies of dissociation of the order of ca. 10 kcal mol21 which is somewhat greater in strength than an average hydrogen bond but is considerably less than the 35 kcal mol21 predicted for a single bond in Table 4.123 It is also noteworthy that the chemistry of the tin compounds is usually consistent with a monomeric stannanediyl formulation seen in compounds that are monomeric in the solid, e.g.[Sn{C(SiMe3)2(CH2)2C(SiMe3)2}] 124a or [Sn(C6H2But 3-2,4,6)2].124b Products with Sn]Sn bonds are rarely obtained from such reactions. Exceptions are the reaction of [Sn{CH(SiMe3)2}2]2 with HCCPh,125a the reaction of [Sn{C6H2(CF3)3-2,4,6}2] with MesN3,125b and the reaction of (SnTrip2)2 with tellurium.125c The latter diaryl is unusual in that it is the only diorganotin(II) species that exists as a dimer in solution.126 It is stable only at low temperature (ca. 270 8C) however, and it readily converts to the cyclic trimer (SnTrip2)3 above 0 8C. The structure of the dimer would be of great interest since it may display a short Sn]Sn bond and a less pyramidal metal geometry. The sole dilead compound, [Pb{C6H2(CF3)3-2,4,6}{Si(SiMe3)3}]2, has a Pb]Pb distance that is ca. 0.65 Å longer than that predicted for a single bond.98 Clearly, the Pb]Pb interaction is very weak even in the solid state. The compound, not surprisingly, is monomeric in solution. It is clear from their physical properties, structures and chemical behavior that the weakly associated germanium, tin or lead species bear little resemblance to alkenes. As a result the terms digermenes, distannenes, or diplumbenes hardly seem justified.These names often appear to be used for convenience 127 rather than accuracy. Since they behave as simple monomeric dialkyl or diaryl metal compounds in solution, the IUPAC terms germane-, stannane- or plumbane-diyls are obviously more descriptive. If they are weakly associated in the solid state, the terms bis-(germane-, stannane- or plumbane-diyls), which have been used in some of the recent literature, seem apt.33,35 The terms digermylene, distannylene and diplumbylene are, apparently, inappropriate.Such nomenclature is used 128 to describe the putatively singly bonded valence isomer of the corresponding dimetallyne (Fig. 5). The term distannylene may be an appropriate description of the (as yet unisolated) neutral analog of the anion [Sn(C6H3Trip2-2,6)2]241 that has a lone pair at each of the tin atoms which appear to be connected by a single bond. The Pauling definition may also be applied to the lower valent Group 13 compounds.It has already been seen that the Fig. 5 Disilylyne and disilylene forms of a compound of the general formula (SiR)2 R Si Si R R Si Si R •• •• disilylyne disilylene currently known (Table 5) neutral, lower valent organic derivatives of the Group 13 metals are usually weakly associated, often becoming monomers in the vapor phase or in solution. By analogy with the lower valent Group 14 species the terms digallene, diindene or dithallene for dissociating RMMR species are not indicative of their structures.For the monomeric species, the IUPAC name gallanediyl, indanediyl, etc. seems suYciently descriptive, and for the dimeric species bis(metallanediyl) is more representative of their structures than the term dimetallene. Bonding models The various bonding models for the heavier Group 14 and Group 15 compounds have been described in a recent review.121 Accordingly, these models are only briefly discussed here. As expected, the bonding in the Group 15 species is the most straightforward.In simplistic terms, the p bond is formed from parallel overlap of p orbitals and the s bond results in the nitrogen case from overlap of two orbitals with approximate sp2 hybridization. As the group is descended the s character in the s bond decreases and that in the lone pair increases so that in the heaviest antimony and bismuth compounds, the lone pair is predominantly s in character.78 In the Group 14 compounds a molecular orbital (MO) view128a,b of the M]M bonding in R2MMR2 dimers (M = Si, Ge, Sn or Pb) is that upon descending the group there is increased mixing (Fig. 6) of an antibonding M]M s* orbital and a p orbital due to the lowering of the p–s* energy gap. This is closely related to the critical orbital interactions in the pyramidalization 129 (which is a second-order Jahn–Teller eVect 130) of AH3 systems such as NH3, CH3 2 or PH3 and the bending of AH2 molecules such as H2O or H2S.The mixing is particularly marked in the case of tin and lead dimers and to a lesser extent in the silicon and germanium analogues. The increased mixing of the s* orbital results in stabilization of the original p orbital but weakens that p bond by increasing its lone pair character. In eVect, it comes to resemble an n lone pair orbital more than a p bond (n = non-bonding, lone pair orbital). The s bond is also weakened since the orbital orientations are not as favorable in the pyramidalized species as in the planar one.In eVect, bonding electron density is lowered between the two Group 14 atoms. In the tin and lead compounds in Table 1, the bonding interaction is weakened to such an extent that the compounds are dissociated to monomers in solution at room temperature. An appealing aspect of the MO bonding model is that it readily explains why the use of more electronegative (EN) substituents usually results in dissociation to monomers. Increasing the EN of the substituent (e.g.by using amide 131 instead of alkyl or aryl groups) or decreasing the EN of the central (Group 14) atom reduces the p–s* energy gap and gives s-bonding orbitals that are more located on the substituent atom while the s* orbitals are biased toward the central atom. This gives a greater interaction between the p and s* orbitals which increases the stabilization of the lone pair relative to the p orbital and increases the likelihood of a monomeric structure. On the other hand, more electropositive substituents, e.g.silyl groups,32,33,37 decrease the Fig. 6 Schematic drawing of the mixing of the p and s* molecular orbitals in a heavier group 14 species p s n *2946 J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 stability of the lone pair form and favor a p-bonded dimeric structure. The trans-bending in the heavier element compounds has also been considered 132 in light of the splittings of the singlet– triplet states (DEST) based on calculations on the hydrogen derivatives (MH2)2 (M = Si, Ge or Sn).According to this criterion the double bonds will be unstable with respect to the trans-bent distortion if the sum of the DEST values of the fragments is more than one half the total bend energy. It can be seen from Table 4 that the total bond energy drops rapidly descending the group. Moreover, the singlet–triplet (ST) splitting increases 132 so that the trans-bent structure is normally observed.The type of bonding occurring in such compounds has also been described as two weak semipolar dative bonds between two singlet metallanediyls, the so-called ‘pawpaw’ bond.133 This approach emphasizes the important role of ST separation in the bending; the larger the ST splitting the more trans-bent the structure becomes. This view of the M]M interaction is similar to the donor– acceptor model (Fig. 7) proposed originally by Lappert and coworkers. 134 In eVect, the metals are held together by weak (cf. 6.4 kcal mol21) 8 polarized dative bonds in which electron density is distributed asymmetrically. Ab initio MO calculations 5b,135 on the (MH2)2 (M = Ge or Sn) model compounds have aVorded an M]M dissociation energy of ca. 31.0 kcal mol21 for the Ge species and 21.5 kcal mol21 for the tin analog. These values are much less than the experimental M]M single bond dissociation energies of compounds such as (MH3)2 and (MMe3)2 (M = Ge or Sn) which are in the range of ca. 50 to 70 kcal mol21.136 However, the calculated value for (SnH2)2 137 is almost double the enthalpy of dissociation for the sterically hindered compound [Sn{CH(SiMe3)2}2]2. Apparently, steric eVects are Fig. 7 Donor–acceptor bonding model for [Sn{CH(SiMe3)2}2]2 Sn Sn R R R R Fig. 8 Schematic drawing of the unshared electron-pair resonance in (SnR2)2 (ref. 138) Sn+ Sn R R R R R Sn- Sn+ R R R - Fig. 9 Calculated MO energy and correlation diagram for (GeMe)2 (ref. 79) very important in determining the strength of the M]M bond. Density functional theoretical (DFT) studies 137 have indicated a significant influence of intra- and inter-atomic Pauli repulsion on trans-bending and p-bond strength in the heavier (MH2)2 species (M = Si, Ge, Sn or Pb). Moreover, quite high intrinsic p-bond strengths were calculated, 38 kcal mol21 (Ge), and 32 kcal mol21 (Sn). This study also showed that in the heavier elements the dissociation energy is less for double than for single bonds but that this was due to the fact that preparation energy for the singlet fragments and the interatomic Pauli repulsion was higher for the heavier elements.Pauling 138 also proposed that the Sn]Sn single bond distance in [Sn{CH- (SiMe3)2}2]2 can be explained by assuming that an unshared electron pair resonates between two tin atoms that are connected by a single bond (Fig. 8). A theoretical study of fractional bond orders in (MH2)2 (M = Si, Ge or Sn) molecules 139 revealed bond orders that were less than two but greater than one, and were in reasonable agreement with the bond order formula of Pauling.140 The MO explanation of the bonding in heavier Main Group 14 analogues of alkynes is similar to that for alkenes.The bending from the linear geometry and the lengthening of the element–element bonds occurs for the same reason it occurs in the alkene analogues. As the energies of the p and s* levels become closer going down the group the two levels interact and there is increased mixing of a s* antibonding orbital into one of the p orbitals which causes it to become more stable and assume non-bonding n character.Calculations 79 on the hypothetical species (GeMe)2 show that the trans-bent form is more stable than the linear form by ca. 15 kcal mol21. In the higher energy linear form the two p orbitals are degenerate (eu symmetry) and the s bond (ag symmetry) is lower in energy (Fig. 9). As long as the geometry remains linear the s* orbital cannot mix with the p levels as they are mutually orthogonal.With bending, this restriction is removed, and one of the original p orbitals (15bu) becomes progressively more stable as the amount of mixing increases. The mainly non-bonding, lonepair character of this orbital is indicated by the contour (Fig. 10) which shows regions of electron density opposite the methyl groups. In contrast, the energies of the remaining p orbital (6au) and the s orbital (15ag) are increased, indicating that these bonds are also weakened by the bending process.If this MO view of the bonding is accepted then it is clear that in the linear configuration the bond order is 3 since there are three bonding MOs each occupied by an electron pair. However, distortion of the geometry toward trans-bent and mixing of antibonding character from the s* level reduces the bond order below 3 by converting one of the three bonding orbitals originally associated with the Ge]Ge bond to an essentially non-bonding one.In eVect, the formal bond order is reduced from 3 to 2, or less than 2 if the weakening of the remaining s and p bonds is taken Fig. 10 Contour diagram of the 15bu lone pair MO in trans-(GeMe)2 obtained with a 6–31G* basis set [ref. 79(a)]J. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 2947 into account. Accordingly, the Ge]Ge bond (2.167 Å) in the lowest energy trans-bent configuration is considerably longer than that calculated (2.014 Å) for the triply bonded linear form.Interestingly, if this distance is compared to 2.417 Å calculated (with the same basis set) for the single Ge]Ge bond in (GeH3)2 and a line is drawn between the two points on a semilog plot based on the Pauling relationship between bond length and bond order,140 then the bond length for the trans-structure, 2.167 Å, aVords a bond order of 1.98. Calculations on (SnMe)2 141 show a similar pattern.In the linear distannyne the Sn]Sn distance is 2.432 Å whereas the lower (by 34 kcal mol21) energy configuration has a trans-bent structure (Sn]Sn]C 1258) and an Sn]Sn distance of 2.673 Å. This gives a Sn]Sn bond order of 1.46. However, it is clear that steric eVects can play a large role in determining the Sn]Sn bond length and longer Sn]Sn distances are seen in practice. The hypothetical compound (SiMe)2 142 (Si]Si ca. 2.07 Å) also has a bent geometry and is 14.4 kcal mol21 more stable than the linear form.The calculated bond order is 2.17. Several theoretical studies have been made on aluminum and gallium Main Group 13 species that have possible multiple M]M interactions.121 Some data are given in Table 6.78,79,81,143,144 In addition, the influence of the ligand geometry on InI]InI and TlI]TlI bonding has been investigated 145,146 via extended Hückel theory. The latter studies 146 showed that the weak M]M interaction occurs between empty p levels and the occupied s orbitals. Variation of the angle in the trans-bent structures such as (TlH)2 or [Tl(h5-C5H5)]2 indicated that interaction is almost non-bonding (or repulsive) in the linear arrangement but that upon trans-bending to a Tl]Tl]ligand angle of ca. 1208 there is a significant increase in the overlap population.145 Attention has been drawn to the similarity between the Group 14 and 15 metallanediyl fragments, e.g. :SnR2 and :InR, which only diVer in the number of acceptor orbitals.146 Thus, the trans-bending in dimers (MR)2 (M = In or Tl) corresponds to mixing of the p-acceptor orbitals into the in-phase and out-of-plane lone pair orbitals.In MO terms there is mixing of the antibonding s* level into a p level very similar to that already described for the Group 14 compounds. The valence bond description involves a polarized donor–acceptor bonding model like that proposed for the tin(II) dialkyl. The In ? ? ? In 146 or Tl ? ? ? Tl 146,147 interactions are quite weak and are weaker than the Sn ? ? ? Sn interaction in [Sn{CH(SiMe3)2}2]2.8 This weakness is exemplified by the M]M distances (ca. 3.63 Å) in the dimers [M{h5-C5- (CH2Ph)5}]2 (M = In or Tl) which are much longer than normal In]In (ca. 2.8 Å) or Tl]Tl (ca. 3.0 Å) single bonds (cf. Table 3). In the lighter species (AlH)2 and (GaH)2 the trans-bent form is also calculated to be lower in energy than the linear forms.143,144 The Ga]Ga distance, 2.656 Å is ca. 0.14 Å longer than the normal Ga]Ga single bonds given in Table 3, whereas the calculated Al]Al bond length is very similar to that of a single bond. A number of calculations have been carried out on the doubly reduced gallium analogues of these compounds since Table 6 Selected calculated structural parameters for some low-valent derivatives of aluminum and gallium Compound (AlH)2 (GaH)2 [(GaH)2]22 [(GaMe)2]22 Na2[(GaMe)2] Li2[(GaMe)2] Na2[(GaPh)2] Na2[{Ga(C6H3Ph2-2,6)}2] Structure Linear trans-bent Linear trans-bent Linear trans-bent trans-bent trans-bent Linear trans-bent trans-bent trans-bent M]M/Å 2.298 2.613 2.2512 2.656 2.214 2.4568 2.5221 2.5082 2.161 2.388 2.461 2.362 M]M]C/8 180 120 180 120.4 180 125.2 123.8 126.0 180 132.5 126.1 124.3 Ref. 143 143 144 144 81 81 81 81 79 79 78 78 the structure of Na2[{Ga(C6H3Trip2-2,6)}2] was published.74 Calculations using density functional quantum mechanical methods on Na2[(GaMe)2] or [(GaMe)2]22 aVord a trans-bent structure with long Ga]Ga distances of 2.508 and 2.522 Å.It was concluded 81 that there was no Ga]Ga bond length–bond order relationship and that the bonding in the experimental molecule 74 was between triple and double in character despite the relatively long bond length. The bond order was adduced on the basis of localized rather than canonical MOs. It may be noted, however, that extended Hückel MO calculations (see below)79a show that the three canonical MOs associated with the Ga]Ga bonds (of which one is essentially non-bonding) are already mainly localized 79 on the digallium moiety.The calculated distances are considerably longer than those experimentally observed in Na2[{Ga(C6H3Trip2-2,6)}2]. A possible explanation for this discrepancy comes from DFT calculations on the hypothetical compound Na2[{Ga(C6H3Ph2-2,6)}2] which suggest that the alkali metals may play a significant role in shortening the Ga]Ga bond since the potential energy curve as a function Ga]Ga distance is relatively flat in this region.78 Thus, the Ga]Ga distance in Na2[(GaPh)2] is 2.46 Å whereas the Ga]Ga distance in Na2[{Ga(C6H3Ph2-2,6)}2], which has non-covalent Na1 ? ? ?o-Ph interactions, is only 2.36 Å, i.e. 0.1 Å shorter. The relatively shallow potential energy curve for the Ga]Ga bond is also suggestive of overall bond weakness. The DFT calculations also led to the conclusion that the Ga]Ga bond is a double one.78 Extended Hückel MO calculations 79 on Li2[(GaMe)2], in which the [(GaMe)2]22 ion is isoelectronic to the (GeMe)2 compound discussed above, reveal a similar pattern and Ga]Ga bond order of 2 or less.In this case [cf. (GeMe)2, Fig. 10] the HOMO (15bu) is the lone pair orbital (i.e. the n2 combination) and the HOMO-1 (7au) is the p orbital which is believed to have been stabilized by the Li1 ions. The Ga]Ga bond length in the lowest energy trans-bent form was calculated to be 2.388 Å and the Ga]Ga]C angle is 132.58.The linear, triply bonded valence isomer lies ca. 7.5 kcal mol21 higher in energy and has a much shorter Ga]Ga bond of 2.161 Å. The calculated angle is similar to that observed experimentally (ca. 1318) although the Ga]Ga distance is about 0.06 Å longer than the experimental one. The calculations also indicated that the bending of the (GaMe)2 array in Li2[(GaMe)2] results in a weakening of the remaining p and s bonds. A recent paper, however, reported 80 that a natural bond order analysis of the anion [(GaH)2]22 led to the conclusion that, although one of the p bonds is ‘slipped’, ‘three bonds are .. . obtained’. The contour diagram for the ‘slipped’ p bond, however, can also be interpreted mainly in terms of an n lone pair orbital where the maxima of electron density are not located in the region between the gallium nuclei.79a 6 Summary and Conclusions DiVering interpretations of the structural, spectroscopic and theoretical data have given rise to the current debate.One interpretation says that the weak interactions frequently observed or calculated for heavier Main Group 13 or 14 analogs of compounds such as alkenes or alkynes fully merit designation as multiple bonds, with a nomenclature that emphasizes their relationship to their lighter analogs. Another holds that the bond lengths observed for these compounds can, in many cases, be greater than the lengths of normal single bonds and the interactions are frequently so weak that the ‘multiple’ bonds eVectively do not exist in solution. In short, the relationship between bond order and bond strength and length becomes increasingly meaningless if any interaction, no matter how weak, can be qualified as a ‘bond’.Not surprisingly, each side can cite theoretical data that support their position. However, the theoretical data have proved their usefulness in the sense that they demonstrate that pyramidalization or bending of the geometries at the heavier elements usually leads to a significantly weakened2948 J.Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 element–element bond in comparison to the idealized planar or linear configuration. Thus the theoretical data provide bond lengths for the idealized linear or planar species against which the significantly longer bonds in the real molecules can be compared. It is, therefore, to be hoped that further experimental investigation will provide deeper insight on the factors that influence the strength of these bonds.In the long run, it is probable that the actual physical behavior and chemical properties of these compounds will determine the bonding model that is most practical for chemists to use. More specific observations and comments as well as suggestions for further experiments are given below. Group 14 and 15 compounds At present, it seems clear that in the Group 15 element compounds there can be relatively strong double bonding for all elements of this group.In such compounds the physical and chemical characteristics of double bonding are maintained in the solution and solid states. In Group 14, however, no currently known compounds of formula (SnR2)2 or (PbR2)2 remain dimerized in solution at room temperature. Some germanium species also have this characteristic. The M]M bonding in these weakly associated compounds is much weaker than single bonding and their chemistry is more consistent with their formulation as monomers.The terms digermenes, distannenes or diplumbenes provide an inadequate description of the metal– metal bond. Similarly, the term dimetallyne seems inappropriate for neutral molecules of the type RMMR where the MM distances are close to that of a single bond and there is a strongly bent M]M]C angle {cf. the highly bent structure of the anion [{Sn(C6H3Trip2-2,6)}2]2}.41 It has been correctly said that amongst Group 14 elements the carbon species are in reality the ‘exotic’ compounds 121 in the type of bonds they form and that it is the trans-bent structures of the dimers (MR2)2 or (MR)2 (M = Si, Ge, Sn or Pb) that are normal.The multiple bonding model (i.e. distinct s and p bonds) used for alkenes and alkynes is so ingrained in the chemical consciousness that it is diYcult to think of their heavier analogues without citing it. However, the continuing emphasis of bonding models and terminologies appropriate for elements of the first row has tended to obscure the uniqueness of the bonding in the heavier element compounds.In short, the bonding in these elements is inherently interesting and distinct, and does not require an often misleading analogy with lighter congeners to emphasize its importance. Group 13 compounds In the Group 13 elements a similar pattern to that seen in the Group 14 elements (albeit with weaker bonding) is rapidly emerging, although data are scant. Stable (BR)2 and [(BR)2]22 compounds are currently unknown but calculations 148 indicate that only (BH)2 has a double bond with a triplet ground state. In contrast, the association of species of formula (MR)2 (M = Al, Ga, In or Tl) appears to be relatively weak and, except in the hypothetical aluminum compound (AlH)2, which has a singly bonded Al]Al distance, the M]M bonds are longer than normal single bonds.The long bonds and trans-bent geometry observed 114 for the organo-indium and -thallium dimers imply ever weaker bonding than that seen in related Group 14 species.In no sense do they resemble doubly bonded ‘dimetallenes’ either in their physical properties or their chemistry. Recent calculations 79a on the dimeric organogallium species (GaMe)2 also indicate a trans-bent structure and a long Ga]Ga distance of 2.676 Å which is ca. 0.16 Å longer than a single bond and incompatible with a bond order of 2. Irrespective of how the bonding in the RMMR (M = Al, Ga, In or Tl) dimer is seen, i.e. whether it involves the bending of the geometry and mixing of a s* orbital into a p level or that it is composed of weak donor–acceptor bonds that may be similar in strength to hydrogen bonds, there is little doubt that the bond order is much less than 2 (less than 1 in the case of the indium and thallium derivatives) and the generation of a further bond by double reduction to give the [RMMR]22 ion may be insuYcient to generate a bond order of 3.In this sense, the terms dimetallene and ‘dimetallyne’ are unrepresentative of their bonding. s Electron participation The most striking aspect of the currently available experimental and theoretical data for the heavier Group 13, 14 and 15 compounds is how profoundly the lower tendency of the s electrons to participate in bonding aVects the molecular configurations in homonuclear species.147b In the Group 15 elements it seems possible to form double bonds with p orbitals in compounds of formula REER (E = P, As, Sb or Bi) with little s electron participation (they remain primarily lone pair in character) 78 so that double bonding is observed even for the heavier elements antimony and bismuth.In Group 14 elements, however, full- fledged doubly bonded molecules that are formally analogous to ethylene would require substantial s orbital participation in the double bonding which becomes less energetically favored upon descending the group.In eVect, the s electrons display an increasing preference to remain in essentially non-bonded, lone pair orbitals which is, of course, just another manifestation of the so-called inert pair eVect. In Group 13 this behavior is seen in the increasing stability of the bent form of the [RMMR]22 ions, or weak association in the neutral molecules RMMR (M = Al, Ga, In or Tl). 7 Future Work Multiple bonding that is stable to dissociation in solution at room temperature has yet to be observed for In]In, Tl]Tl, Sn]Sn or Pb]Pb compounds.The synthesis of homonuclear derivatives of these elements with stable multiple bonds remains, therefore, an exciting synthetic challenge. As has also been seen, no stable dimeric compounds of formula RMMR (M = Al or Ga) are known currently. The synthesis of these would be of particular interest since they would help resolve the nature of the bond in their reduced M92[RMMR] (M9 = alkali metal) analogues. Furthermore, the synthesis of solvent separated ion pairs of the type [MLn]2[RMMR] will also illuminate the role played by ligand–alkali metal p interactions in determining M]M bond lengths which has been suggested theoretically, 78 but which remains undefined experimentally.Apart from one tentative estimate,3a there are no experimental data for p-bond strengths for homonuclear multiple bonding in the heavier Group 13 elements. Other families of compounds also beckon. There is a scarcity of stable compounds with heteronuclear multiple bonding, for instance heavier Group 13–15 compounds with valence multiple bonds. Stable, neutral homonuclear heavier Group 14 alkyne analogues and their doubly reduced derivatives (isoelectronic to neutral Group 15 species) also remain unknown.Data on these may throw light on the tendency to isolate the singly reduced anion [{Sn(C6H3Trip2- 2,6)}2]2 in preference to the neutral analogue.41 This tendency implies the presence of a low-lying unoccupied orbital in the neutral species and, possibly, a weak Sn]Sn bond order of 1 or less.The increasing use of electropositive substituents such as silyl groups32,42,63,108b to enhance M]M bond strengths may also be expected in future experiments. Bond lengths Inevitably, future investigations will rely on various spectroscopies and X-ray diVraction data. The latter technique, in particular, has proved of key importance but it can lead to erroneous conclusions if used in isolation. The observation of short bond lengths and deviations from idealized geometries often provides reliable information on the nature and strength of the bonds involved.Nonetheless, in some cases the unusual geometries and unsuitable orbital orientations or energies canJ. Chem. Soc., Dalton Trans., 1998, Pages 2939–2951 2949 allow a fairly close approach of atoms without aVording a bond of the expected strength, as in [Sn{CH(SiMe3)2}2]2.5 Also, some bond distances in relatively weak bonds may be easily shortened by other interactions, owing to a relatively flat potential curve in certain ranges, cf.calculations on Na2[{Ga(C6H3Ph2- 2,6)}2] 78 which suggest that the interactions of the Na1 with aryl rings may have shortened the Ga]Ga distance by as much as 0.1 Å. In essence, such eVects tend to shorten the intermetallic distance and increase the apparent bond order rather than the opposite. 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Engl., 1986, 25, 642; G. E. Herberich, B. Hessner and M. Hostalek, J. Organomet.Chem., 1988, 355, 473; J. H. Davis, E. Sinn and R. M. Grimes, J. Am. Chem. Soc., 1989, 111, 4784; G. E. Herberich, C. Ganter, L. Weseman and R. Boese, Angew. Chem., Int. Ed. Engl., 1990, 29, 912. 100 J. J. Eisch, Adv. Organomet. Chem., 1996, 39, 355. 101 H. Meyer, G. Schmidt-Lukasch, G. Baum, W. Massa and A. Berndt, Z. Naturforsch., Teil B, 1988, 43, 801. 102 Ref. 6, p. 1052; R. M. Minyaev and R. HoVmann, Chem. Mater., 1991, 3, 547. 103 W. BiVar, H. Nöth and H. Pommerening, Angew.Chem., Int. Ed. Engl., 1980, 19, 56; K. Schlüter and A. Berndt, Angew. Chem., Int. Ed. Engl., 1980, 19, 57. 104 H. Klusik and A. Berndt, Angew. Chem., Int. Ed. Engl., 1981, 20, 870; H. Klusik and A. Berndt, J. Organomet. Chem., 1981, 222, C25. 105 (a) A. Moezzi, M. M. Olmstead and P. P. Power, J. Am. Chem. Soc., 1992, 114, 2715; (b) A. Moezzi, R. A. Bartlett and P. P. Power, Angew. Chem., Int. Ed. Engl., 1992, 31, 1082. 106 M. 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Kobayashi and S. Nagase, Organometallics, 1997, 16, 2489. 143 Z. Palágyi, R. S. Grev and H. F. Schaefer, III, J. Am. Chem. Soc., 1993, 115, 1936. 144 Z. Palágyi and H. F. Schaefer, III, Chem. Phys. Lett., 1993, 203, 195; G. Treboux and J.-C. Barthelat, J. Am. Chem. Soc., 1993, 115, 4870. 145 C. Janiak and R. HoVmann, Angew. Chem., Int. Ed. Engl., 1989, 28, 1688; C. Janiak and R. HoVmann, J. Am. Chem. Soc., 1990, 112, 5924. 146 P. H. M. Budzelaar and J. Boersma, Recl. Trav. Chim. Pays-Bas, 1990, 109, 187. 147 (a) P. Schwerdtfeger, Inorg. Chem., 1991, 30, 1660; (b) P. Schwerdtfeger, G. A. Heath, M. Dolg and M. A. Bennett, J. Am. Chem. Soc., 1992, 114, 7518. 148 C. Jouany, J. C. Barthelat and J. P. Dandey, Chem. Phys. Lett., 1987, 36, 52; D. R. Armstrong, Theor. Chim. Acta, 1981, 60, 159; J. D. Dill, P. v. R. Schleyer and J. Pople, J. Am. Chem. Soc., 1975, 97, 3402. Received 24th March 1998; Paper 8/02281J

ISSN:1477-9226    

DOI:10.1039/a802281j

出版商:RSC    

年代:1998

数据来源: RSC

摘要:

DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1998, Pages 2953–2954 2953 1,3-Bis(silyl)cyclodisilazane: synthesis and crystal structure Bettina Jaschke, Regine Herbst-Irmer, Uwe Klingebiel,* Peter Neugebauer and Thomas Pape Institut für Anorganische Chemie der Universität, Tammannstr. 4, D-37077 Göttingen, Germany. E-mail: uklinge@gwdg.de The reaction between (Me3C)2MeSiNHLi and SiCl4 aVorded R–NH–SiCl3 1, and (R–NH)2SiCl2 2; 1 reacted with BuLi with formation of the tetrachlorocyclodisilazane (RN–SiCl2)2 3, R = SiMe(CMe3)2, hydrogenation of which with LiAlH4 resulted in the first cyclodisilazane (RN–SiH2)2 4. Compared to cyclosilazanes bearing organic substituents, only very few Si–N ring systems in which some or all of the substituents are inorganic are known.1–5 We have reported the first examples of cyclodisilazanes bearing silyl groups at the nitrogen and fluorine substituents at the silicon atoms.3,4 These compounds display the opposite structural features to their organic-substituted counterparts, i.e.in cyclodisilazanes the Si– N–Si angles are smaller while the N–Si–N angles are larger than 908, and the observed endocyclic Si–N bonds are shorter than the exocyclic bonds. Theoretical calculations for model compounds demonstrate that fluorination of the parent compound, i.e. (H2Si–NH)2 æÆ (F2Si–NH)2, leads to shortening of the endocyclic Si–N bonds from 173 to 170 pm and, thus, to stabilisation of the ring system.4 Furthermore silylation at the nitrogen atoms leads to a decrease in the endocyclic Si–N–Si angles.The combined eVect of the fluorine atoms at the silicon and of two silyl groups at the nitrogen atoms leads to the formation of the “smallest” (Si–N)2 four-membered ring systems known so far and to an overall shortening of the transannular Si ? ? ? Si distance. The shortest Si ? ? ? Si contact has been predicted for (F2Si–NSiH3)2 (242 pm) and found for [F2Si–NSi(CMe3)2Ph]2 (237.6 pm).These contacts are equal to or only slightly longer than Si–Si bonds in disilanes.5 1,3-Bis(silyl)cyclodisilazanes (R3Si–N–SiH2) are still unknown. We were interested in the structural features of these compounds and report the synthesis and molecular structure of the first cyclodisilazane that bears hydrogen atoms at the silicon and silyl groups at the nitrogen atoms. Knowing the stabilisation eVect of bulky silyl groups from the isolation of iminosilanes, 6 monomeric aminoalanes,6 or silyldiazenes,7 we used the di-tert-butylmethylsilyl group as substituent. Two routes lead to the formation of the cyclosilazanes (Scheme 1).†‡ Lithium di-tert-butylmethylsilylamide 8 reacts with SiCl4 in a molar ratio 1 : 1 or 1 : 0.5 to give the silylaminotrichlorosilane 1 or the bis(silyl)dichlorosilane 2.In the reaction of 1 with BuLi a lithium derivative is obtained, which is thermally unstable. The cyclodisilazane 3 is formed by LiCl elimination. Ring closure with formation of 3 also occurs in the reaction of the dilithium derivative of 2 and SiCl4.Compound 3 was found to react with LiAlH4 to give the 1,3- bis(silyl)cyclodisilazane 4. Fig. 1 shows the crystal structure of 4.§ Compound 4 crystallises in the space group P21/c with half of a molecule in the asymmetric unit. The sum of the angles around the nitrogen atoms is 359.58. The ellipsoids of the ring Si atoms showed a long displacement perpendicular to the ring plane.Therefore the SiH2 group is refined over two positions (Fig. 1 shows the Si- and H-atom position with higher occupancies). This leads to a great unreliability of the bond lengths and angles involved and they should not be discussed. Standard molecular orbital calculations4,11 were carried out for the parent 1,3-cyclodisilazane and the 1,3-bis(silyl)cyclodisilazane, which serve as close models for the synthesised † Preparative details. Compound 1: Di-tert-butylmethylsilylamine (0.5 mol, 86.5 g) in n-hexane (350 ml) was metallated with 1 equivalent of nbutyllithium (15% in n-hexane).The suspension was cooled to 0 8C and added to a solution of tetrachlorosilane (0.5 mol, 84.9 g) in n-hexane (100 ml). The mixture was warmed to room temperature and heated to reflux for 1 h. After removal of lithium chloride by filtration (1022 Torr) 1 was purified by distillation in vacuo (1022 Torr). Yield 91%, bp 50 8C (0.01 Torr); FI-MS: m/z (%) = 305 (95) [M]1, EI-MS m/z (%) = 248 (20) [M 2 C(CH3)3]1. Compound 2: 2 was synthesised like 1.The molar ratio was 2:1 in this case (0.5 mol, 86.5 g di-tert-butylmethylsilylamine, 0.25 mol, 42.5 g tetrachlorosilane), 2 was also purified by distillation. Yield 73%, bp 135 8C (0.01 Torr); EI-MS: m/z (%) = 385 (40) [M 2 C(CH3)3]1. Compound 3: 1 (80.1 mol, 27.9 g) was dissolved in n-hexane (50 ml) and cooled to 210 8C, 1 equivalent of BuLi (15% in n-hexane) was added slowly. After stirring for 1 h at 25 8C the reaction mixture was warmed to room temperature.Lithium chloride was filtered oV and after removal of the solvent 3 was crystallised in n-hexane. Yield 40%, mp 127 8C; FI-MS: m/z (%) = 540 (100) [M]1, EI-MS: m/z (%) = 483 (35) [M 2 C(CH3)3]1. Compound 4: a solution of 3 (0.02 mol, 10.8 g) in diethyl ether (50 ml) was slowly added to a suspension of LiAlH4 (0.02 mol, 0.76 g) in diethyl ether (25 ml). The mixture was heated to reflux for 16 h. After separation from the solid components by filtration 4 was obtained by distillation and crystallisation in n-hexane.Yield 42%, mp 96 8C; EI-MS: m/z (%) = 402 (5) [M]1, 345 (58) [M 2 C(CH3)3]1; IR: n& = 2132.7 cm–1 (SiH). The isolated compounds are analytically pure, air stable, but moisture sensitive. ‡ NMR data: 1H NMR (CDCl3, 250 MHz, SiMe4), 1: d 0.27 (s, SiMe, 3 H), 1.00 (s, SiCMe3, 18 H), 1.64 (s, NH, 1 H); 13C NMR (CDCl3, 250 MHz, SiMe4), d 27.83 (s, SiCH3), 20.16 (s, SiCC3), 27.70 (s, SiCC3); 29Si NMR (CDCl3, 250 MHz, SiMe4), d 224.69 (s, SiCl3), 12.40 (s, SiC). 2: 1H NMR (CDCl3, 250 MHz, SiMe4), d 0.25 (s, SiMe, 6 H), 0.99 (s, SiCMe3, 36 H), 1.25 (s, NH, 2 H); 13C NMR (CDCl3, 250 MHz, SiMe4), d 27.74 (s, SiCH3), 20.25 (s, SiCC3), 27.91 (s, SiCC3); 29Si NMR (CDCl3, 250 MHz, SiMe4), d 228.53 (s, SiCl2), 10.25 (s, SiC). 3: 1H NMR (CDCl3, 250 MHz, SiMe4), d 0.31 (s, SiCMe, 6 H), 1.08 (s, SiCMe3, 36 H); 13C NMR (CDCl3, 250 MHz, SiMe4), d 25.78 (s, SiCH3), 20.73 (s, SiCC3), 28.49 (s, SiCC3); 29Si NMR (CDCl3, 250 MHz, SiMe4), d 237.67 (s, SiCl2), 10.90 (s, SiC). 4: 1H NMR (CDCl3, 250 MHz, SiMe4), d 0.11 (s, SiMe, 6 H), 0.98 (s, SiCMe3, 36 H), 5.45 (s, SiH2, 4 H); 13C NMR (CDCl3, 250 MHz, SiMe4), d 27.67 (s, SiCH3), 21.14 (s, SiCC3), 28.13 (s, SiCC3); 29Si NMR (CDCl3, 250 MHz, SiMe4), d 220.67 (SiH2, 1JSiH = 230.2 Hz), 6.28 (s, SiC). § X-ray structure determination of 4: data were collected at 2140 8C on a Stoe-Siemens-Huber diVractometer with CCD area detector and monochromated Mo-Ka radiation (l = 71.073 pm).The structure was solved by direct methods.9 All non-hydrogen atoms were refined anisotropically. 10 For the hydrogen atoms the riding model was used. The disordered SiH2 group is refined with distance restraints and restraints for the anisotropic displacement parameters [occupancies: 0.61(2) : 0.39(2)]. Crystal data for 4: C18H46N2Si4, Mr = 402.93, monoclinic, space group P21/c, a = 871.8(2), b = 1236.8(2), c = 1245.9(2) pm, b = 107.25(1)8, U = 1.2830(4) nm3, Z = 2, r(calc.) = 1.043 Mg m23, m = 0.236 mm21, 16319 reflections collected of which 2619 were unique (Rint = 0.0354); 2619 data and 23 restraints used for the refinement of 139 parameters, R1 = 0.0396 for I > 2s(I), wR2 = 0.1124 for all data.CCDC reference number 186/1118. See http://www.rsc.org/suppdata/ dt/1998/2953/, for crystallographic files in .cif format.2954 J. Chem. Soc., Dalton Trans., 1998, Pages 2953–2954 Table 1 Calculated and measured geometrical parameters of cyclodisilazane NH Si Si HN H2 H2 SiH3 H3Si N Si Si N H2 H2 N Si Si N SiMe(CMe3)2 (Me3C)2MeSi H2 H2 Si ? ? ? Si Si–N (endo) N–Si (exo) Si–N–Si (ring) Calc. 254.2 173.5 — 94.2 Calc. 249.7 174.6 172.4 91.3 Measured 243–245 173.4 173.3 89.0 compound 4; Si–H bonds of 147.4 to 147.8 pm were obtained. It was found that substitution of the hydrogen atoms on the nitrogen atoms by silyl groups decreases the Si–N–Si ring angles Scheme 1 and therefore shortens significantly the Si ? ? ? Si distance. The bulkiness of the di-tert-butylmethyl groups in 4 explains the discrepancy of 5 pm in the endocyclic Si ? ? ? Si distance found experimentally in 4. Important calculated and measured geometrical parameters of cyclodisilazanes (bond lengths in pm; bond angles in 8) are in Table 1.Acknowledgements This work was supported by the Fonds der Chemischen Industrie. References 1 J. Haiduc and D. B. Sowerby, The Chemistry of Inorganic Homo- and Heterocycles, Academic Press, London, 1987. 2 H. Bürger, M. Schulze and U. Wannagat, Inorg. Nucl. Chem. Lett., 1967, 3, 43. 3 B. Tecklenburg, U. Klingebiel, M. Noltemeyer and D. Schmidt- Bäse, Z. Naturforsch., Teil B, 1992, 47, 855. 4 T. Müller, Y. Apeloig, I. Hemme, U. Klingebiel and M. Noltemeyer, J. Organomet. Chem., 1995, 494, 133. 5 Ch. Brönneke, R. Herbst-Irmer, U. Klingebiel, P. Neugebauer, M. Schäfer and H. Oberhammer, Chem. Ber./Recl., 1997, 130, 835. 6 J. Niesmann, U. Klingebiel, M. Noltemeyer and R. Boese, Chem. Commun., 1997, 365. 7 H. Witte-Abel, U. Klingebiel and M. Noltemeyer, Chem. Commun., 1997, 771. 8 L. Ruwisch, U. Klingebiel, S. Rudolph, R. Herbst-Irmer and M. Noltemeyer, Chem. Ber., 1996, 129, 823. 9 G. M. Sheldrick, SHELXS 90, Acta Crystallogr., Sect. A, 1990, 46, 467. 10 G. M. Sheldrick, SHELXL 93, University of Göttingen, 1993. 11 V. Metail, S. Joanteguy, A. Chrostowska-Senio, G. Pfister- Guillouzo, A. Systermans and J. L. Ripoll, Inorg. Chem., 1997, 36, 1482. Received 7th August 1998; Communication 8/06267F Fig. 1 Molecular structure of compound 4, selected bond lengths (pm) and angles (8): Si(2)–N(1) 172.8(4), Si(2)–N(1A) 174.0(4), Si(1)– N(1) 173.3(1); Si(1)–N(1)–Si(2A) 129.3(2), Si(2)–N(1)–Si(1) 140.5(2), Si(2)–N(1)–Si(2A) 89.7(2), N(1)–Si(1)–C(1) 106.4(1), N(1)–Si(1)–C(2) 109.0(1), N(1)–Si(2)–N(1A) 90.3(2).

ISSN:1477-9226    

DOI:10.1039/a806267f

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年代:1998

数据来源: RSC

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DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1998, Pages 2955–2956 2955 Preparation, structure and preliminary magnetic studies of tri- and tetra-nuclear cobalt–lanthanide carboxylate complexes † Yong Cui, Jiu-Tong Chen, De-Liang Long, Fa-Kun Zheng, Wen-Dan Cheng and Jin-Shun Huang * State Key Laboratory of Structural Chemistry, Fujian Institute of Research on the Structure of Matter, Chinese Academy of Sciences, Fuzhou, Fujian, 350002, P. R. China Carboxylate-bridged cobalt–lanthanide complexes [Co2Nd- (O2CCMe3)6(C9H7N)2(NO3)] and [Co2Er2(O2CCMe3)8- (C9H7N)2(NO3)2] have been prepared and structurally characterised by X-ray diVraction; both of them are found to be antiferromagnetically coupled.With the aim of clarifying the role of the exchange interactions between 4f and 3d metal ions modifying the properties of magnetic materials containing rare-earth metals, in the last few years, an increasing interest has been given to the magnetic properties of molecular complexes comprising simultaneously lanthanide and transition metal ions.1–8 This interest has been essentially focused on the CuII–GdIII couple 1–6 which has been found to be directly ferromagnetic. Recent studies have revealed that the magnitude of the ferromagnetic exchange interaction between the diVerent metal centres is exponentially dependent on the Cu]Gd distance.5,7 By comparison, the magnetic interactions between other 3d–4f mixtures have been poorly explored,8,9 and in fact only very few structurally characterised complexes are known which contain such combinations of metals.8–11 Herein we described the preparation, crystal structure and preliminary magnetic studies of [Co2Nd(O2CCMe3)6- (C9H7N)2(NO3)] 1 and [Co2Er2(O2CCMe3)8(C9H7N)2(NO3)2] 2.To the best of our knowledge, the present cases are the first two examples of carboxylate-bridged discrete heterometallic cobalt–lanthanide complexes 9,10 and are new model complexes for magnetic investigation. A mixture of Co(NO3)2?6H2O (0.58 g, 2 mmol), Nd(NO3)3? 6H2O (0.44 g, 1 mmol) and pivalic acid (0.61 g, 6 mmol) was dissolved in EtOH (80 mmol), followed by addition of quinoline (0.24 mL).The resultant mixture was refluxed for 16 h, filtered while hot, and then concentrated to 25 mL. The filtrate was left at room temperature and red-brown prismatic crystals of 1 were deposited in 35% yield after two weeks.‡ X-Ray crystallography § has established that complex 1 consists of a discrete trinuclear [Co2Nd(O2CCMe3)6(C9H7N)2(NO3)] molecule as shown in Fig. 1. Two cobalt atoms are each co-ordinated by three carboxylate oxygen atoms and a quinoline molecule to form distorted tetrahedrons where the most distorted angles are 96.6(2) and 96.0(2)8 for O(12)]Co(1)]N(1) and O(42)]Co(2)] N(2), respectively. Each terminal cobalt atom is connected to the central neodymium atom by three bridging pivalate ligands with Co ? ? ? Nd separations of 4.0668(7) and 4.0700(7) Å.The arrangement of the three metals is quasi-linear with a Co(1)] Nd]Co(2) angle of 134.10(2)8. Besides the six carboxylate oxygen atoms, the eight-co-ordination sphere of the neodymium † Supplementary data available: cell parameters for [Co2Ce(O2CCMe3)6(C9H7N)2(NO3)] and [Co2Y2(O2CCMe3)8(C9H7N)2(NO3)2], Nd and Er co-ordination polyhedra in complexes 1 and 2, and plots of meff vs. T for complexes 1 and 2. For direct electronic access see http://www.rsc.org/suppdata/dt/1998/2955/, otherwise available from BLDSC (No.SUP 57420, 5 pp.) or the RSC Library. See Instructions for Authors, 1998, Issue 1 (http://www.rsc.org/dalton). ‡ Satisfactory elemental analyses were obtained for both compounds. (1: Found C, 48.3; H, 5.9; N, 3.4. Calc. for C48H68Co2N3NdO15: C, 48.5; H, 5.8; N, 3.5%. 2: Found: C, 42.3; H, 5.3; N, 3.2. Calc. for C58H86Co2Er2N4O22: C, 42.4; H, 5.3; N, 3.4%.) Non-SI unit employed: mB ª9.274 × 10224 J T21. atom is completed by a chelating nitrate anion. The coordination polyhedron around the neodymium atom is irregular, with the chief distortion being caused by the presence of a small bidentate nitrate anion, but it still can be described as a dodecahedron with the two trapezia defined by O(11), O(1), O(2), O(51) and O(31), O(21), O(61), O(41) intersecting at an angle of 86.38.The analogous reaction of hydrated cobalt nitrate with hydrated erbium nitrate leads to a quite diVerent, crystallographically characterised product 2.Here a centrosymmetric tetranuclear Co2Er2 complex is formed, the inversion center residing on the midpoint of the Er–Er vector, as shown in Fig. 2. Again two four-co-ordinated cobalt atoms, each having a terminal quinoline ligand, are each triply connected to the adjacent Er atom via three bridging pivalate groups. The structure of the dinuclear Co–Er subunits is very similar to the structure of Co–Nd moieties in 1, but the Co ? ? ? Er distance of 3.9127(9) Å is slightly shorter than the Co ? ? ? Nd distance, as a consequence of the smaller size of the erbium atom.A pair of Er atoms in two centrosymmetrically related binuclear Co–Er subunits are linked together by two pivalate groups in a syn–syn bridging fashion with an Er ? ? ? Er distance of 5.2017(5) Å. As a result, the sequence of atoms Co]Er]Er]Co makes a broken line with a Co]Er]Er angle of 137.05(2)8. The Er atom is sevenco- ordinate, with five of the sites filled by five pivalate oxygen atoms and two sites by a chelating nitrate anion.The geometry of the Er atom is based on a pentagonal bipyramid, with the equatorial plane defined by O(1), O(2), O(21), O(11) and O(42A) and the axial positions occupied by O(31) and O(41) [O(31)]Er]O(41) = 176.6(2)8]. The change in structure from 1 to 2 probably results from the reduction in the ionic radius in going from Nd to Er, which is reflected in the lower co-ordination number of the lanthanide in 2.Similar eVects have been found previously.3 Elemental analysis and cell determination suggest that the Ce analogue has structure 1, whilst the yttrium one has structure 2 (see SUP 57420). It should be noted that quinoline plays an important role in the formation of such 3d–4f mixed metal complexes, especially for those 3d metal ions having tetrahedral coordination sites. Detailed discussion will be given in subsequent reports. The magnetic properties of 1 and 2 in the solid state have been investigated at 1.0 T, over the temperature range 5–280 K (see SUP 57420).For complex 1, meff per molecule decreases § Crystal data for 1 C48H68Co2N3NdO15: M = 1189.15, triclinic, space group P1� , a = 11.9406(2), b = 12.6970(2), c = 22.1401(1) Å, a = 98.209(1), b = 98.083(1), g = 117.24(2)8, U = 2871.74(7) Å3, Z = 2, Dc = 1.375 g cm23, m = 1.522 mm21, F(000) = 1222. Crystal data for 2 C58H86Co2Er2N4O22: M = 1643.69, triclinic, space group P1� , a = 10.9792(3), b = 12.4595(5), c = 13.7959(2) Å, a = 85.5964(8), b = 84.5640(7), g = 73.050(7)8, U = 1794.72(5) Å3, Z = 1, Dc = 1.521 g cm23, m = 2.838 mm21, F(000) = 828.Data collection and processing: T = 293 K, R = 0.038 (wR = 0.081) for 6505 observed reflections [2q < 45.58, F2 > 2s(F2)] and S = 1.05 for 1, and R = 0.050 (wR = 0.120) for 5140 observed reflections [2q < 50.08, F2 > 2s(F2)] and S = 1.05 for 2. For complex 1 the tertbutyl groups of the pivalate ligands involving C(12) and C(52) were treated as having rotational disorder around the C(11)]C(12) or C(51)]C(52) bond with three methyl groups each occupying two half-weighted sites, while those tert-butyl groups involving C(32) and C(62) were wholly split into two parts each having site occupancy factor of 0.50.In the latter case the terminal C]C bonds in each pivalate ligand were restrained to be similar with a standard deviation of 0.03 Å. CCDC reference number 186/1115.See http://www.rsc.org/suppdata/dt/1998/2955/for crystallographic files in .cif format.2956 J. Chem. Soc., Dalton Trans., 1998, Pages 2955–2956 Fig. 1 Molecular structure of complex 1. Only one part of the disordered tert-butyl groups is presented. Hydrogen atoms are omitted for claritylected interatomic distances (Å): Nd]O(1) 2.565(5), Nd]O(2) 2.538(3), Nd]O(11) 2.358(3), Nd]O(21) 2.385(5), Nd]O(31) 2.385(4), Nd]O(41) 2.408(4), Nd]O(51) 2.385(3), Nd]O(61) 2.428(4), Co(1)]O(12) 1.943(4), Co(1)]O(22) 1.964(3), Co(1)]O(32) 1.940(3), Co(1)]N(1) 2.081(4), Co(2)]O(42) 1.948(4), Co(2)]O(52) 1.939(4), Co(2)]O(62) 1.948(4), Co(2)]N(2) 2.092(4) Fig. 2 Molecular structure of complex 2. Hydrogen atoms are omitted for clarity. Selected interatomic distances (Å): Er]O(1) 2.412(5), Er]O(2) 2.404(6), Er]O(11) 2.229(6), Er]O(21) 2.287(5), Er]O(31) 2.231(5), Er]O(41) 2.174(5), Er]O(42Ai) 2.248(5), Co]O(12) 1.946(6), Co]O(22) 1.940(5), Co]O(32) 1.967(5), Co]N(2) 2.077(6).Symmetry code: i 2 2 x, 2y, 3 2 z gradually from 6.82 at 280 K to 5.63 mB at 5 K. This behavior is clearly characteristic of intramolecular antiferromagnetic coupling. Complex 2 is also antiferromagnetically coupled: meff per molecule is 14.19 mB at 274 K, which declines to 10.95 mB at 5 K. At room temperature, meff per Co21 ion is expected to be 3.83 mB;12 for Er31 and Nd31 ions, meff is ª9.6 and 3.5 mB,12 respectively. For CoII 2ErIII 2 and CoII 2NdIII complexes with noninteracting metal ions, meff is expected to be ª14.6 and 6.6 mB, respectively.The experimental values of meff are thus close to the corresponding non-interacting values, indicating that all interactions are quite weak. Acknowledgements The research was supported by grants from the State Key Laboratory of Coordination Chemistry at Nanjing University and the National Science Foundation of China. References 1 C. Benelli, A. Caneschi, D. Gatteschi, O. Guillou and L.Pardi, Inorg. Chem., 1990, 29, 1751 and refs. therein. 2 N. Matsumoto, M. Sakamoto, H. Tamaki, H. Okawa and S. Kida, Chem. Lett., 1989, 853; S. Wang, S. J. Trepanier and M. J. Wagner, Inorg. Chem., 1993, 32, 833; M. Andruk, I. Ramade, E. Codjovi, O. Guillou, O. Kahn and J. C. Trombe, J. Am. Chem. Soc., 1993, 115, 1822; I. Ramade, O. Kahn, Y. Jeannin and F. Robert, Inorg. Chem., 1997, 36, 930 and refs. therein. 3 A. J. Blake, R. O. Gould, C. M. Grant, P. E. Y. Milne, S.Parsons and R. E. P. Winpenny, J. Chem. Soc., Dalton Trans., 1997, 485 and refs. therein. 4 J.-P. Costes, F. Dahan, A. Dupuis and J.-P. Laurent, Inorg. Chem., 1997, 36, 3429 and refs. therein. 5 C. Benelli, A. J. Blake, P. E. Y. Milne, J. M. Rawson and R. E. P. Winpenny, Chem. Eur. J., 1995, 1, 614. 6 X.-M. Chen, M.-L. Tong, Y.-L. Wu and Y.-J. Luo, J. Chem. Soc., Dalton Trans., 1996, 2181; X-M. Chen, Y.-L. Wu, Y. X. Tong and X.-Y. Huang, J. Chem. Soc., Dalton Trans., 1996, 2443. 7 J. L. Sanz, R. Ruiz, A. Gleizes, F. Lloret, J. Faus, M. Julve, J. J. Borrás-Almenar and Y. Journaux, Inorg. Chem., 1996, 35, 7384. 8 E. K. Brechin, S. G. Harris, S. Parsons and R. E. P. Winpenny, J. Chem. Soc., Dalton Trans., 1997, 1665; J. P. Costes, F. Dahan, A. Dupuis and J.-P. Laurent, Inorg. Chem., 1997, 36, 4284. 9 S. Decurtins, M. Gross, H. W. Schmalle and S. Ferlay, Inorg. Chem., 1998, 37, 2443. 10 N. Sakagami and K. Okamoto, Chem. Lett., 1998, 201; N. Sakagami, M. Tsunekawa, T.Konno and K. Okamoto, Chem. Lett., 1997, 575; A. Gonzalez, A. Beltran and A. L. Bail, Acta Crystallogr., Sect. C, 1991, 47, 1624; D. M. L. Goodgame, T. E. Muller and D. J. Williams, Polyhedron, 1992, 11, 1513. 11 G. B. Deacon, C. M. Forsthy, W. C. Patalinghug, A. H. White, A. Dietrich and H. Schumann, Aust. J. Chem., 1992, 45, 567; L. F. Lindoy, H. C. Lip, H. W. Louie, M. G. B. Drew and M. J. Hudson, J. Chem. Soc., Chem. Commun., 1977, 778; J. M. Boncella and R. A. Andersen, J. Chem. Soc., Chem. Commun., 1984, 809; W. J. Evans, L. Bloom, J. W. Grate, L. A. Hughes, W. E. Hunter and J. L. Atwood, Inorg. Chem., 1985, 24, 4620; J. P. White, H. Deng, E. P. Boyd, J. Galluci and S. G. Shore, Inorg. Chem., 1994, 33, 1685; D. M. L. Goodgame, S. Menzer, A. T. Ross and D. J. Williams, J. Chem. Soc., Chem. Commun., 1994, 2605; D. Deng, X. Zheng, C. Qian, J. Sun, A. Dormond, D. Baudry and M. Visseaux; J. Chem. Soc., Dalton Trans., 1994, 1665; Y. Yukawa, S. Igarashi, A. Yamano and S. Sato, Chem. Commun., 1997, 711. 12 A. T. Casey, S. Mitra, in Theory and Applications of Molecular Paramagnetism, eds. E. A. Boudreaux and L. N. Mulay, John Wiley & Sons, New York, 1976, pp. 271–316. Received 2nd June 1998; Communication 8/04132F

ISSN:1477-9226    

DOI:10.1039/a804132f

出版商:RSC    

年代:1998

数据来源: RSC

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DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1998, Pages 2957–2959 2957 Kinetics of phosphodiester hydrolysis by lanthanide ions in weakly basic solutions Paola Gómez-Tagle and Anatoly K. Yatsimirsky *,† Facultad de Química, Universidad Nacional Autónoma de México, 04510 México D.F., México Lanthanide(III) cations in aqueous Bis-tris propane buVer remained stable at elevated pH values and at pH 9.0 catalyzed the hydrolysis of several phosphodiester substrates 103 times more eYciently than in neutral solutions.Much attention is currently focused on lanthanide (Ln) catalysis in phosphodiester hydrolysis.1 Although in the majority of known systems the active forms of Ln catalysts were not identified, most probably these are hydroxo complexes of Ln(III) cations, as one can conclude from available pHdependencies of rates of metal-promoted phosphate ester hydrolysis reactions [RNA hydrolysis by LaCl3,1j phosphate triester hydrolysis by a La(III) macrocyle,2 phosphodiester hydrolysis by complexes of Co(III),3 Zn(II),4 Ni(II)5 and Cu(II),5 phosphonate ester hydrolysis by La(III) 6].There is, however, a certain diYculty in generation and stabilization of such complexes in solution. From one side, Ln(III) aqua-cations possess high pKa values of about 9 7 and appreciable amounts of their hydroxo complexes are formed only in basic solutions. On the other hand, Ln(III) hydroxides have small solubility products of about 10220–10222 mol4 dm212 and readily co-precipitate with simple inorganic ions like nitrate or chloride that further reduce their solubilities.7a As a result, solutions of Ln(III) salts become unstable at pH slightly above 7 and the hydroxo complexes can be produced only at very low concentrations.A possible solution to this problem would be the use of appropriate ligands, which would protect a Ln(III) cation from precipitation, but would not decrease too much its electrophilicity, necessary to generate coordinated OH anions and to activate the substrate.Hay and Govan2 reported a 100-fold increase in activity of a La(III) macrocyclic complex in the hydrolysis of 2,4-dinitrophenyl diethyl phosphate on going from pH 7 to 9.5; no hydroxide precipitation was observed in basic solutions, but the substrate employed was highly activated. Schneider and co-workers1e,g tested eVects of various ligands on the hydrolytic reactivity of Eu31 towards bis(p-nitrophenyl) phosphate (BNPP) (a typical phosphodiester substrate with very low intrinsic reactivity) at pH 7.0, and found that amines, aminoethers and polyols produced minor eVects, but carboxylic acid ligands strongly inhibited the hydrolysis.It seems from these results that neutral ligands, e.g. aminoalcohols, could function as non-deactivating stabilizers of Ln(III) cations at elevated pH values. We tested several ligands for their ability to prevent precipitation of 2 mM Ln(III) cations (Ln = La, Nd, Pr or Eu) in the range pH 7–9 including ethylendiamine, 1,3-diaminopropane, polyamines, Bis-tris [2,2-bis(hydroxymethyl)-2,29,20-nitrilotriethanol], TRIS [tris(hydroxymethyl)aminomethane] and Bistris propane {BTP, 1,3-bis[tris(hydroxymethyl)methylamino]- propane}; and obtained satisfactory results with the latter.A minimum concentration of BTP suYcient to prevent Ln(III) precipitation over the whole pH range was 10 mM. The pro- † E-mail: anatoli@servidor.unam.mx cedure of solution preparation was as follows: to the required amount of BTP solution of pH slightly higher than the final one, an aliquot of a Ln(III) stock solution was added and the pH of the solution, which was still ca. 0.2 higher than desired, was adjusted with perchloric acid. We found it important to always follow this method of solution preparation in order to get reproducible results. Perchlorate salts of Ln(III) cations (Strem) were employed. The stock solutions of lanthanides were standardized by volumetric titration with H4EDTA.8 Kinetics of hydrolysis was followed spectrophotometrically by the appearance of p-nitrophenolate at 400 nm in the cases of BNPP and p-nitrophenylphosphate (NPP) as substrates and by the appearance of phenol at 280, 285 and 290 nm in the case of diphenylphosphate (DPP).The kinetics of hydrolysis of NPP and DPP was strictly first-order and the respective rate constants (kobs) were calculated by the integral method or (for very slow reactions) from initial rates.During the BNPP hydrolysis the liberation of 2 equivalents of p-nitrophenol per 1 substrate equivalent was observed, but the reaction kinetics deviated noticeably from the simple first-order rate law due to insuf- ficiently fast hydrolysis of the intermediate NPP. The kinetic curves were fitted to a respective two-exponential equation derived from the scheme of two successive pseudo-first-order steps of hydrolysis of BNPP to NPP and further to an inorganic phosphate with liberation of p-nitrophenolate at each step.‡ The correctness of the scheme was confirmed by comparison of calculated rate constants with those found from the initial rates (kobs of the first step) and from the kinetics of NPP hydrolysis (kobs of the second step).The sole phosphate product of BNPP hydrolysis was the inorganic phosphate quantitatively determined in the reaction mixture by a standard molybdate/ hydrazine colorimetric method9 [no color was developed with unhydrolyzed BNPP; analysis of a mixture containing 0.04 mM BNPP, 2 mM La(ClO4)3 and 20 mM BTP after 24 h at 25 8C showed the presence of 0.046 ± 0.005 mM H3PO4 and 0.083 ± 0.002 mM p-nitrophenol]. Fig. 1 shows the plots of kobs of the first step of BNPP hydrolysis at room temperature vs. BTP concentration in the presence of 2 mM La(III) at two diVerent pH values. The shape of the plots most probably reflects the formation of higher unreactive BTP/La(III) complexes on going to increased ligand concentrations (see below).At a fixed ligand concentration kobs was directly proportional to the La(III) concentration in the range 1–14 mM and increased on going to higher pH values. Fig. 2 shows the pH-dependence of kobs of the first step of BNPP hydrolysis at fixed ligand and metal concentrations. A rapid increase in the reaction rate in the range pH 8–9 is indicative of formation of the reactive hydroxo complexes of the type [La(BTP)(OH)n]3 2 n.The plot in Fig. 2 fits eqn. (1) kobs = kobs 0/(1 1 [H1]2/Ka) (1) ‡ The equation has the form At = eNP[BNPP]0{21(kobs,1e2kobs,2t 1 (kobs,1 2 2kobs,2)e2kobs,1t/(kobs,2 2 kobs,1)} where At is the absorbance at time t, eNP and [BNPP]0 are the molar absorptivity of p-nitrophenolate and initial substrate concentration respectively, kobs,1 and kobs,2 are the rate constants of the cleavage of BNPP and NPP respectively.2958 J. Chem. Soc., Dalton Trans., 1998, Pages 2957–2959 Table 1 Kinetics of phosphate ester hydrolysis by Ln(III)/BTP at 25 8C and pH 9.0 Substrate Bis(p-nitrophenyl)phosphate (BNPP) p-Nitrophenyl phosphate (NPP) Diphenyl phosphate (DPP) Metal None La Nd Pr Eu None La None La k2 a/dm3 mol21 s21 0.22 0.061 0.087 0.030 1.4 × 1023 Ea/kJ mol21 55 ± 6 50 ± 6 97 ± 2 kobs b/s21 6.9 × 10211 c 2.2 × 1023 6.1 × 1024 8.7 × 1024 3.0 × 1024 1.7 × 1029 d 4.1 × 1023 5 × 10215 e 1.4 × 1025 a Second-order rate constants based on total metal concentration at fixed 20 mM BTP concentration; relative error ±5%.b At 10 mM Ln(III) and 20 mM BPT. c The sum of pseudo-first-order rate constants of spontaneous hydrolysis (1.1 × 10211 s21) 1i and of alkaline hydrolysis (kOH = 5.8 × 1026 dm3 mol21 s21 at 25 8C)10 of BNPP at pH 9.0 and 25 8C. d First-order rate constant of the hydrolysis of NPP dianion at 25 8C. 11 e Calculated for 25 8C as described in ref. 1e. which corresponds to a reaction scheme where the reactive form is a dihydroxo complex (n = 2) § with kobs 0 = (5.0 ± 0.1) × 1024 s21 [the pseudo-first-order rate constant of hydrolysis of BTPP by a dihydroxo complex of La(III)] and pKa = 17.66 ± 0.03 [Ka corresponds to dissociation of two aqua ligands: M(H2O)2 M(OH)2 1 2H1].A negligible contribution of an aqua La(III)/BTP complex under given conditions is quite expected. Indeed, reported1i for 2 mM La(III) at 25 8C and pH 7.0 a value of kobs = 1.4 × 1027 s21 is 4 × 103 smaller than kobs 0.Mononuclear dihydroxo complexes of La(III) were not reported,7a but the strict first-order in La(III) indicates that most probably the reactive species is a mononuclear BTP/La(III) complex rather than one of known polyhydroxo polynuclear complexes.1j,7a Fig. 1 Pseudo-first-order rate constants of the hydrolysis of BNPP at 25 8C in the presence of 2 mM La(ClO4)3 as a function of BTP concentration. Solid squares, pH 8.5; open squares, pH 9.0. Fig. 2 Pseudo-first-order rate constants of the hydrolysis of BNPP at 25 8C in the presence of 2 mM La(ClO4)3 and 10 mM BTP as a function of pH.Solid line is the theoretical fit to eqn. (1). § Fitting of the results in Fig. 2 to a complete equation which corresponds to a reaction scheme where the aqua and monohydroxo complexes also contribute to the observed reactivity gives statistically insignificant values for the rate constants of these forms. Another possibility is that the reactive form is a coordinated alcoholate of the ligand, but this is not consistent with the formation of inorganic phosphate as the reaction product.The determination of stoichiometry and stability constants of Ln(III)/BTP complexes is in progress. Some conclusions can be inferred, however, from the kinetic results. As mentioned above, kinetic first-order in La(III) and pH-dependence represented by eqn. (1) indicate the reactive species to be a mononuclear dihydroxo complex. The inhibitory eVect of an excess of BTP, Fig. 1, should reflect the coordination of at least a second ligand molecule since 2 mM lanthanide(III) perchlorate alone precipitated at pH > 8. Inspection of molecular models shows that BTP can function as a hexadentate ligand through two N and four O atoms. Therefore, it probably forms only 1 : 1 and 1: 2 metal–ligand complexes, the first of which is the reactive species. In this case plots in Fig. 1 should be bellshaped, but we can observe only their decreasing hyperbolic parts due to the necessity of adding enough ligand to prevent precipitation of La(OH)3.Fitting of these plots to the respective theoretical equation shows that the first stepwise stability constant must be >103 dm3 mol21 and the second stability constants equal 220 and 85 dm3 mol21 at pH 8.5 and 9.0 respectively. A decrease in stability constant on going to higher pH probably reflects the competition with hydroxo ligands. Similar results were obtained for other lanthanides studied (Nd, Pr or Eu).Table 1 collects the second-order rate constants (k2) for all metals found from the slopes of kobs vs. total metal concentration profiles (ranging from 1 to 10 mM) at pH 9.0 (the highest pH value at which all metals did not precipitate in the presence of BTP) and 25 8C. Also the pseudo-first-order constants at 10 mM Ln concentration are given for comparison with related published results. In the case of La(III) the activation energy was determined from the Arrenius plot of kobs in the range 25–50 8C.The kinetics of the hydrolysis of NPP was studied where the only purpose was to confirm the reaction scheme used for analysis of the integral kinetics of BNPP (see above). One can see from the last column in Table 1 that NPP hydrolysis proceeds two times faster than BNNP hydrolysis and has practically the same activation energy, that is, the relative reactivity of these two substrates is independent of temperature.With the most active metal La(III) the kinetics of hydrolysis of a less reactive substrate DPP was studied. Similar kinetic behavior was observed, but the rate constant decreased ca. 200 times and the activation energy was considerably higher with this substrate, Table 1. The kinetics of phosphate diester hydrolysis by the Ln(III)/ BTP system is substantially diVerent from that by Ln(III) aqua ions. First, no “saturation” at metal concentrations at about 10 mM as reported for all Ln(III) aqua ions 1e–h is observed.Most probably this is due to the lower charge of the metal species (11 instead of 13) since phosphate binding to the metal is predominantly electrostatic.1h In agreement with this we observed aJ. Chem. Soc., Dalton Trans., 1998, Pages 2957–2959 2959 very weak dependence of kobs on the ionic strength in contrast to a strong dependence reported for Ln(III) aqua ions.1h Secondly, the order of activities of lanthanides (La > Nd ª Pr > Eu) in the Ln(III)/BTP system is opposite to that (La < Pr < Eu) for Ln(III) aqua ions.1h In the case of the Ln(III)/ BTP system the activity trend parallels the trend in pKa values of Ln(III) aqua ions, which increase in the order Eu < Nd < Pr < La.7b The same trend is typical for esterolytic activity of hydroxo cations towards p-nitrophenyl acetate as the substrate. 12 Evidently, in both cases the dominant factor is the nucleophilic reactivity of the metal-coordinated OH anion, while for aqua cations of Ln(III) the dominant factor is electrophilic assistance.1h Finally, the catalytic activity of Ln(III)/BTP complexes in a weakly basic media is much higher than that of Ln(III) ions in neutral solutions.Comparisons of kobs values given in the last column of Table 1 with those published in the literature show that although the activity of the Ln(III)/BTP system for Ln = La does not reach that of Th(IV) (2.82 × 1022 s21 at 37 8C),13 it is essentially similar to that of La(III)/H2O2 (4.8 × 1023 s21 at 25 8C)1i which is the most active at the moment among all Ln(III)-based systems.The value of kobs for La(III) extrapolated to 50 8C is calculated to be 1.25 × 1022 s21 by using the activation energy given in Table 1; the activation energy is 103 times higher than that in neutral solutions “saturated” with La(III) and 15 times higher than with most active Er(III) species at the same temperature.1h Comparison with spontaneous hydrolysis shows acceleration factors of ca. 108 for BNPP and DPP and ca. 106 for NPP, Table 1, which make the system practically useful for applications at room temperature. It is also worth noting that the fairly simple kinetic behavior of the Ln(III)/BTP system makes it suitable for further mechanistic study. Acknowledgements The work was supported by DGAPA-UNAM, Project IN 106495. References 1 (a) R. Breslow and D. L. Huang, Proc. Natl. Acad. Sci.USA, 1991, 88, 4080; (b) J. R. Morrow, L. A. Buttrey, V. M. Shelton and K. A. Berback, J. Am. Chem. Soc., 1992, 114, 1903; (c) M. Komiyama, K. Matsumura and Y. Matsumoto, J. Chem. Soc., Chem. Commun., 1992, 640; (d ) N. Takeda, M. Irisawa and M. Komiyama, J. Chem. Soc., Chem. Commun., 1994, 2773; (e) H.-J. Schneider, J. Rammo and R. Hettich, Angew. Chem., Int. Ed. Engl., 1993, 32, 1716; ( f ) J. Rammo, R. Hettich, A. Roigk and H.-J. Schneider, Chem. Commun., 1996, 105; (g) K.G. Ragunathan and H.-J. Schneider, Angew. Chem., Int. Ed. Engl., 1996, 35, 1219; (h) A. Roigk, R. Hettich and H.-J. Schneider, Inorg. Chem., 1998, 37, 751; (i) B. K. Takasaki and J. Chin, J. Am. Chem. Soc., 1993, 115, 9337; ( j) P. Hurst, B. K. Takasaki and J. Chin, J. Am. Chem. Soc., 1996, 118, 9982. 2 R. W. Hay and N. Govan, J. Chem. Soc., Chem. Commun., 1990, 714. 3 J. Chin, M. Banaszczyk, V. Jubian and X. Zou, J. Am. Chem. Soc., 1989, 111, 186. 4 T. Koike and E. Kimura, J. Am. Chem. Soc., 1991, 113, 8935; C. Bazzicalupi, A. Bencini, A. Bianchi, V. Fusi, C. Giorgi, P. Paoletti, B. Valtacoli and D. Zanchi, Inorg. Chem., 1997, 36, 2784. 5 M. A. De Rosch and W. C. Trogler, Inorg. Chem., 1990, 29, 2409. 6 A. Tsubouchi and T. C. Bruice, J. Am. Chem. Soc., 1995, 117, 7399. 7 (a) C. F. Baes and R. E. Mesmer, The Hydrolysis of Cations, Wiley, New York, 1976; (b) R. M. Smith and A. E. Martell. Critical Stability Constants, Plenum Press, New York, 1976, vol. 4. 8 S. J. Lyle and Md. M. Rahman, Talanta, 1963, 10, 1177. 9 G. Charlot, Les Méthodes de la Chimie Analytique, Masson et Cie, Paris, 4th edn., 1961. 10 J. A. A. Ketelaar and H. R. Gersmann, Recl. Trav. Chim. Pays-Bas, 1958, 77, 973. 11 A. J. Kirby and W. P. Jencks, J. Am. Chem. Soc., 1965, 87, 3209. 12 R. W. Hay and R. Bembi, Inorg. Chim. Acta, 1982, 64, L179. 13 R. A. Moss, J. Zhang and K. Bracken, Chem. Commun., 1997, 1639. Received 5th August 1998; Communication 8/06198J

ISSN:1477-9226    

DOI:10.1039/a806198j

出版商:RSC    

年代:1998

数据来源: RSC

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DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1998, Pages 2961–2963 2961 Dinuclear copper(II) complexes that promote hydrolysis of GpppG, a model for the 59-cap of mRNA† Kevin P. McCue, David A. Voss, Jr., Christian Marks and Janet R. Morrow* Chemistry Department, State University of New York at Buffalo, Amherst, New York 14260-3000, USA The hydrolysis of the monoribonucleotide GpppG, a model compound for the 59-cap structure of mRNA, by dinuclear Cu(II) complexes of triazacyclononane was 100-fold more rapid than in the presense of the analogous mononuclear complex; a first-order or second-order dependence on the catalyst was observed for two diVerent dinuclear complexes.There is much interest in the development of metal ion complexes that catalyze the cleavage of RNA.1–3 Such complexes when attached to a recognition agent such as an antisense oligonucleotide catalyze the sequence-specific cleavage of RNA,4–9 and it is proposed that these catalysts may be useful in the selective inactivation of mRNA.The hydrolytic cleavage or transesterification of the RNA phosphate diester backbone 1–9 by metal ion complexes has been studied extensively. An alternate approach which may be useful for the inactivation of mRNA entails the hydrolysis of the 59-cap structure of mRNA.10 Baker has demonstrated that several Cu(II) complexes hydrolyze the 59-cap structure both free in solution 11 and when attached to oligonucleotides.12 Inert lanthanide(III) complexes are also eVective in promoting cleavage of the 59-cap structure.13 Recently, we found that lanthanide(III) complexes promote cleavage more eYciently in the presence of an equivalent of a second metal ion.14 This result has prompted us to study dinuclear metal ion complexes as catalysts for cleavage of GpppG, a model for the 59-cap structure of mRNA (Fig. 1). Our studies here suggest that the mechanism of hydrolysis of * E-mail: jmorrow@acsu.buValo.edu † Supplementary data available: potentiometric titrations, plots of the log of pseudo-first-order rate constants versus log[(Cu2L)41] and the energy minimized structure of [Cu2L3]41 with a bridging pyrophosphate.For direct electronic access see http://www.rsc.org/suppdata/dt/ 1998/2961/, otherwise available from BLDSC (No. SUP 57417, 8 pp.) or the RSC Library. See Instructions for Authors, 1998, Issue 1 (http:// www.rsc.org/dalton). GpppG is distinctly diVerent from that of other phosphorus(V) substrates.Depending on the dinuclear complex employed as catalyst, a first-order or a second-order dependence on catalyst is observed in the hydrolysis of GpppG. Cu(II) complexes of linked triazacyclononane ligands were chosen for our initial studies of dinuclear catalysts (see below). The triazacyclononane ligand L1 binds strongly to transition metal ions 15 and recent studies have shown that similar ligands readily bind two Cu(II) ions.16–18 In addition, Zn(II) and Cu(II) complexes of triazacyclononane catalyze the hydrolytic cleavage of RNA and phosphate diester hydrolysis.19–21 The two dinucleating ligands were prepared ‡ by treatment of the linkers a,a9-dibromo-m-xylene or a,a9-dibromo-p-xylene with 2 equivalents of the N,N9-bis(p-tolylsulfonyl)-1,4,7-triazacyclononane22 and standard deprotection conditions were used.23 Potentiometric titrations of solutions containing a 2 : 1 ratio of Cu(NO3)2 to L2 or L3 (1.0 mM Cu21, 0.5 mM ligand, 0.1 M NaCl) showed well defined inflections at 6 equivalents and 8 equivalents of base (supplementary Figs. 1 and 2, SUP 57417), similar to data for other dinuclear Cu(II) complexes with two triazacyclononane ligands.16,17 This data suggests that L2 and L3 both bind two Cu(II) ions and the predominant species at a 2 : 1 ratio of Cu(II) to ligand at pH 5 is a dinuclear complex [Cu2L]41 while at neutral pH the predominant species is a bis-hydroxide dinuclear Cu(II) complex [Cu2L(OH)2]21.16,17 Hydrolysis of the capped monoribonucleotide, GpppG, by Cu(II) complexes of L1–L3 was examined. Disappearance of GpppG was monitored by use of an HPLC assay 13 and the sole products detected were GMP (guanosine 59-monophosphate) ‡ The ligands L2 and L3 were prepared by treating a,a9-dibromom- xylene or -p-xylene with 2 equivalents of the N,N9-bis(p-tolylsulfonyl)- 1,4,7-triazacyclononane in acetonitrile with a 2–3 fold excess of triethylamine.The mixture was refluxed under nitrogen for 24 h and the ligands were purified by use of silica gel chromatography (2% methanol in chloroform).The ligands were deprotected (ref. 23) and the HBr salts of L1, L2 and L3 were analyzed by use of FAB-MS and 1H NMR. Fig. 12962 J. Chem. Soc., Dalton Trans., 1998, Pages 2961–2963 Scheme 1 and GDP (guanosine 59-diphosphate). A pseudo-first-order rate constant of 2.3 × 1027 s21 was determined (half-life of 840 h) for the hydrolysis of GpppG by 0.25 mM [CuL1]21 at pH 7.3 and 37 8C, with 20 mM hepes buVer.Interestingly, this rate constant is approximately 60-fold lower than other mononuclear Cu(II) complexes that promote hydrolysis of GpppG under similar conditions.11 Both dinuclear Cu(II) complexes hydrolyzed GpppG more rapidly than did [CuL1]21. Pseudo- first-order rate constants of 3.6 × 1025 and 2.2 × 1025 s21 (halflives of 5.3 and 8.8 h) were obtained for the hydrolysis of GpppG in solutions containing 0.25 mM Cu(II) and 0.125 mM L2 or L3, respectively at pH 7.3 and 37 8C.Thus hydrolysis of GpppG is some 100-fold more rapid per Cu(II) ion in the dinuclear complexes than it is for the monomeric [CuL1]21 complex under similar conditions. Under similar conditions but in the absence of catalyst, only 3% of the GpppG was hydrolyzed over a period of 5 d. Further kinetic studies were conducted to study the mechanism of hydrolysis of GpppG by the dinuclear complexes. Hydrolysis of GpppG by both dinuclear Cu(II) complexes was first order in GpppG.§ Hydrolysis of GpppG by the dinuclear Cu(II) complex of L3 was first-order in complex in the § With a ten-fold excess of the dinuclear Cu(II) complexes of L2 or L3, plots of log of the concentration of GpppG versus time were linear for greater than four half-lives.concentration range 0.10 to 0.50 mM with a second-order rate constant of 0.10 M21 s21.¶ In contrast, the hydrolysis of GpppG by the dinuclear Cu(II) complex of L2 was second-order in complex for concentrations ranging from 0.030 to 0.210 mM with an apparent third-order rate constant of 730 M22 s21.|| How might the two Cu(II) centers in the dinuclear complexes of L2 and L3 cooperatively promote hydrolysis of GpppG? Dinuclear metal ion complexes hydrolyze phosphate esters and RNA by double Lewis acid activation 24–35 with one metal ion binding to the incoming nucleophile, the second metal ion binding to the leaving group and both metal ions binding to the phosphate diester. Hydrolysis of GpppG by two metal ions more likely proceeds through interaction of the two metal ions (Scheme 1) at two diVerent phosphate groups similar to the mechanism proposed for nucleoside triphosphate hydrolysis by two metal ions.36,37 In this scheme, one metal ion delivers the nucleophile and the second metal ion binds to the GDP leaving group through one or both phosphates. Modeling studies suggest that the two dinuclear Cu(II) complexes bind diVerently to GpppG.Molecular mechanics calculations** were carried out using five-coordinate Cu(II) complexes with three nitrogen donors and two water molecules.38 Bridging ligands were incorporated by replacement of water molecules. These structures were minimized and energies of the dinuclear complexes with and without bridging ligands were compared.39 Strain ¶ Supplementary data (Fig. 3) contains a plot of log k versus log of the concentration of the dinuclear Cu(II) complex of L3 giving a slope of 1.1.|| Supplementary data (Fig. 4) contains a plot of log k versus log of the concentration of the dinuclear Cu(II) complex of L2 giving a slope of 2.1. ** All calculations were done with Hyperchem 3.0 (Autodesk Inc, Sausalito, CA). The starting structure for the five-coordinate Cu(II) triazacyclononane complexes was obtained by replacing the two bromide ligands in the CuL1Br2 complex (ref. 38) with water ligands. Ideal Cu–N bond distances were obtained from the crystal structure and the Cu–O distance was set to 2.00 Å.Stretching parameters were set to 5.0 mdyne Å21 in order to maintain the macrocycle dimensions. M–L bending parameters were added to maintain the square pyramidal geometry. Torsional parameters about the metal–ligand bond were set to zero so ligand interactions would dictate the overall structure except for torsional interactions where the metal was the terminal atom. For these cases a C4 type atom replaced the metal.After energy minimization, final Cu–ligand bond lengths, bond angles and torsional angles were all within 0.03 Å, 1.3 and 2.08, respectively of those in the crystal structure.J. Chem. Soc., Dalton Trans., 1998, Pages 2961–2963 2963 energy increased only slightly (less than 7%) when phosphate or pyrophosphate are incorporated as bridging ligands into the dinuclear Cu(II) complex of L2. In contrast, the dinuclear Cu(II) complex of L3 could not bind a bridging phosphate ligand without large increases in strain energy due to distortions induced in the aromatic linker. However, a pyrophosphate ligand bridged the two Cu(II) ions without substantial strain being introduced (14%) into the complex.†† This result and the first-order dependence on the dinuclear complex are consistent with the dinuclear Cu(II) complex of L3 promoting GpppG hydrolysis by the mechanism shown in Scheme 1.In contrast, the dinuclear Cu(II) complex of L2 may promote hydrolysis through either binding to a single or to two diVerent phosphate groups of GpppG. One possible mechanism which is consistent with the rate law has the dinuclear Cu(II) complex of L2 binding through a single phosphate of GpppG.Thus 2 equivalents of dinuclear complex are required, one to activate the phosphate group undergoing nucleophilic attack and one to bind to the leaving group. This mechanism is consistent with the manner in which [Cu2L2]41 binds to small molecules. For example [Cu2L2(OH)2]21 contains two briding hydroxide ligands and the complex has a short Cu–Cu distance.18 Other mechanisms are possible and studies are underway to further characterize binding of the dinuclear complexes to GpppG.In summary, we have shown for the first time that dinuclear metal ion complexes eYciently hydrolyze GpppG, a model substrate for the 59-cap structure of RNA. Catalytic properties of the complexes vary dramatically with diVerent linkers for the triazacyclononane ligands.Future studies will focus on further delineating the mechanism of hydrolysis of phosphoric anhydrides with dinuclear metal ion complexes and the design of new linkers to more precisely position the two metal ion centers for the hydrolysis of phosphoric anhydrides. Acknowledgements J. R. M. thanks the Alfred P. Sloan foundation for a fellowship and the National Science Foundation for support. References 1 For a review see, S. Kuusela and H.Lonnberg, Metal Ions in Biological Systems, eds. H. Sigel and A. Sigel, Dekker, NY, 1996, vol. 32, ch. 7. 2 J. R. Morrow, ibid., vol. 33, ch. 19. 3 R. Haner and J. Hall, Antisense Nucleic Acid Drug Dev., 1997, 7, 423. 4 C. A. Stein and J. S. Cohen, Cancer Res., 1988, 48, 2659. 5 D. Magda, R. A. Miller, J. L. Sessler and B. L. Iverson, J. Am. Chem. Soc., 1994, 116, 7439. 6 J. Bashkin, E. I. Frolova and U. Sampath, ibid., 1994, 116, 5981. 7 J. Hall, D. Husken, U. Pieles and H.E. Moser, Chem. Biol., 1994, 1, 185. †† See supplementary data for an example of a minimized structure of a dinuclear complex with a bridging pyrophosphate. 8 D. Magda, M. Wright, S. Crofts, A. Lin and J. L. Sessler, J. Am. Chem. Soc., 1997, 119, 6947. 9 J. Hall, D. Husken and R. Haner, Nucleic Acids Res., 1996, 24, 3522. 10 B. F. Baker, Antisense Research and Applications, eds. S. T. Crooke and B. LeBleu, CRC Press, Boca Raton, FL, 1993, pp. 37–53. 11 B. F. Baker, J.Am. Chem. Soc., 1993, 115, 3378. 12 B. F. Baker, K. Ramasamy and J. Kiely, Bioorg. Med. Chem. Lett., 1996, 6, 1647. 13 B. F. Baker, H. Khalili, N. Wei and J. R. Morrow, J. Am. Chem. Soc., 1997, 38, 8749. 14 D. Epstein, H. Khalili, B. F. Baker and J. R. Morrow, manuscript in preparation. 15 A. E. Martell and R. M. Smith, Critical Stability Constants, Plenum, New York, 1982, vol. 5. 16 X. Zhang, W.-Y. Hsieh, T. N. Margulis and L. J. Zompa, Inorg. Chem., 1995, 34, 2883. 17 R.Haidar, M. Ipek, B. DasGupta, M. Yousaf and L. J. Zompa, ibid., 1997, 36, 3125. 18 L. J. Farrugia, P. A. Lovatt and R. D. Peacock, J. Chem. Soc., Dalton Trans., 1997, 911. 19 V. M. Shelton and J. R. Morrow, Inorg. Chem., 1991, 30, 4295. 20 K. A. Deal and J. N. Burstyn, Inorg. Chem., 1996, 35, 2792. 21 M. J. Young and J. Chin, J. Am. Chem. Soc., 1995, 117, 10 577. 22 K. Wieghart, I. Tolksdorf and W. Herrmann, Inorg. Chem., 1985, 24, 1230. 23 J. L. Sessler, J. W. Sibert, A. K.Burrell, V. Lynch, J. T. Markert and C. L. Wooten, Inorg. Chem., 1993, 32, 4277. 24 D. R. Jones, L. F. Lindoy and A. M. Sargeson, J. Am. Chem. Soc., 1984, 106, 7807. 25 D. H. Vance and A. W. Czarnick, J. Am. Chem. Soc., 1993, 115, 12 165. 26 E. A. Kesicki, M. A. DeRosch, L. H. Freeman, C. L. Walton, D. F. Harvey and W. C. Trogler, Inorg. Chem., 1993, 32, 5851. 27 D. Wahnon, A.-M. Lebuis and J. Chin, Angew. Chem., Int., Ed. Engl., 1995, 34, 2412. 28 N. H. Williams and J. Chin, Chem. Commun., 1996, 131. 29 W. H. Chapman, Jr. and R. Breslow, J. Am. Chem. Soc., 1995, 117, 5462. 30 T. Koike, M. Inoue, E. Kimura and M. Shiro, J. Am. Chem. Soc., 1996, 118, 3091. 31 A. Tsubouchi and T. C. Bruice, J. Am. Chem. Soc., 1995, 117, 7399. 32 M. Wall, R. C. Hynes and J. Chin, Angew. Chem., Int. Ed. Engl., 1993, 32, 1633. 33 M. Irisawa, N. Takeda and M. Komiyama, J. Chem. Soc., Chem. Commun., 1995, 1221. 34 T. A. Steitz and J. A. Steitz, Proc. Natl. Acad. Sci. USA, 1993, 90, 6498. 35 N. Strater, W. N. Lipscomb, T. Klabunde and B. Krebs, Angew. Chem., Int. Ed. Engl., 1996, 35, 2024. 36 H. Sigel, F. Hofstetter, R. B. Martin, R. M. Milburn, V. Scheller- Krattiger and K. H. Scheller, J. Am. Chem. Soc., 1984, 106, 7935. 37 R. M. Milburn, M. Gautam-Basak, R. Tribolet and H. Sigel, J. Am. Chem. Soc., 1985, 107, 3315. 38 R. D. Bereman, M. R. Churchill, P. M. Schaber and M. Winkler, Inorg. Chem., 1979, 18, 3122. 39 P. E. Jurek, A. E. Martell, R. J. Motekaitis and R. D. Hancock, Inorg. Chem., 1995, 34, 1823. Received 28th July 1998; Communication 8/05916K

ISSN:1477-9226    

DOI:10.1039/a805916k

出版商:RSC    

年代:1998

数据来源: RSC

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DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1998, 2965–2967 2965 Macropolyhedral boron-containing cluster chemistry. The [S2B18H19]2 anion, and the reversible dismantling and regeneration of an apical boron cluster site with cluster connectivity six Tomás¡ Jelínek,a,b Ivana Cisar¡ová,c Bohumil S¡ tíbr,b John D. Kennedy a and Mark Thornton-Pett a a The School of Chemistry of the University of Leeds, Leeds, UK LS2 9JT b The Institute of Inorganic Chemistry of the Academy of Sciences of the Czech Republic, 25068 R¡ ez¡-by-Prague, The Czech Republic c The Faculty of Natural Sciences of Charles’ University, Hlavova 2030, 12842 Prague 2, The Czech Republic Received 24th July 1998, Accepted 24th July 1998 The double-cluster [S2B18H19]2 anion, from the [syn- B18H21]2 anion and elemental sulfur, has an eleven-vertex arachno-type subcluster with a {B7} hexagonal planar-based pyramidal feature that reversibly disassembles and reassembles upon protonation followed by deprotonation.There is contemporary continuing interest in the generation of new geometries in boron-containing cluster chemistry.1 In the development of this context we have described the nineteenvertex macropolyhedral thiaborane cluster compound [S2B17- H17(SMe2)] (compound 1, schematic IA), obtained from the heat-induced autofusion of [SB8H10(SMe2)].2 Compound 1 has an unusual cluster structure in that one of its subclusters, formally arachno ten-vertex (IB), has an apical boron atom with the high cluster connectivity of six (vertex BH in IB).This open arachno-type subcluster geometry IB is quite diVerent from the conventional 3 arachno ten-vertex geometry as exhibited, for example, by the arachno [6-SB9H12]2 anion (IC).4 From this isolated result it has not been clear whether the six-connectivity {BH} component of this new cluster shape IB is (a) inherently stable, and has not yet been seen in an isolated single-cluster compound because the necessary synthetic route is not yet discovered, or (b) it is a kinetic artefact trapped by the constraints of the reaction coordinate that generates the intercluster linkage.2 Of relevance here, we have found now a second example of this type of high cluster connectivity.It occurs in an eleven-vertex arachno-type subcluster in the twenty-vertex [S2B18H19]2 macropolyhedral thiaborane monoanion (compound 2, [N(PPh3)2]1 salt). The route to anion 2 involves a completely diVerent synthetic strategy to that for compound 1, and, furthermore, it appears that the observed six-connectivity feature can reversibly disassemble and reassemble via a simple protonation–deprotonation sequence.S S SMe2 S B H H IA IB H H neutral 1 S IC syn-B18H22 (360 mg, 1.64 mmol) in thf (25 ml) was deprotonated with excess NaH, elemental sulfur (400 mg) was added, the mixture stirred for 12 h, and then heated at reflux for 3 h. Water was added, the thf was evaporated (water pump), and the solution was filtered and then precipitated with a slight excess of [N(PPh3)2]Cl.The pale yellow precipitate, the [N(PPh3)2]1 salt of the yellow anion [S2B18H19]2 (compound 2), was purified by column chromatography (silica gel, CHCl3) (RF 0.25 by analytical TLC on silufol with CH2Cl2 as liquid phase), yield 48%. The other principal product was the well recognised [nido-7- SB10H11]2 anion. Compound 2 is characterised by single-crystal X-ray diVraction analysis (Fig. 1 and schematic IIA) * and NMR spectroscopy.† It is seen to consist of a conventionally shaped3 nido eleven-vertex {B10S} subcluster that is fused, with two boron atoms in common, with a {B10S} unit of a previously unrecognised eleven-vertex arachno-type. The shape of this latter unit (IIB), formally based on an (as yet hypothetical) [SB10H13]2 anion, resembles a conventional ten-vertex arachno cluster (e.g. IC) except that an additional boron atom is incorporated in the open face, thus generating an eleven-vertex cluster and a cluster connectivity of six for the adjacent apical boron atom (vertex BH in IIB).This geometry of type IIB is Fig. 1 Crystallographically determined molecular structure of the [S2B18H19]2 anion 2. Selected interatomic distances (pm) are: from B(1) to B(2) 185.2(11), to B(3) 1.851(11), to B(4) 1.952(12), to B(5) 1.869(11), to B(6) 1.883(11), and to B(7) 1.913(11); B(2)–B(3) is 1.770(12), B(3)– B(4) 1.781(11), B(4)–B(5) 1.735(10), B(5)–B(6) 1.709(11), B(6)–B(7) 1.734(11) and B(7)–B(2) is 1.787(12); from S(11) to B(8) 1.964(10), to B(9) 1.945(10), and to B(10) 1.940(9).There was 50 : 50 {S(99)}: {BH(119)H(109,119)} crystallographic disorder within the conventionally structured (primed numbering) subcluster.2966 J. Chem. Soc., Dalton Trans., 1998, 2965–2967 previously unobserved. It is quite diVerent from the more conventional 5 eleven-vertex arachno geometries of structures IIC and IID.The essential planar-based hexagonal pyramidal {B7} unit is an unprecedented feature of both 1 and 2, and does not figure in classical 3 borane building-block philosophies. The 11B and 1H NMR chemical shifts † of species 2 are tentatively but reasonably assigned and are consistent with the molecular formulation. The two common boron atoms B(5) and B(6) do not have BH(exo) bondings, and, in accord with this, two singlets are discernible in the 11B spectrum, at d(11B) = 23.9 and 214.8.An interesting feature of the 11B spectrum consists of four unusually high-field resonances between d(11B) 247 and 250 which arise from the hexagonal basal B(2), B(3), B(4) and B(7) positions of the novel eleven-vertex subcluster of geometrical type IIB. Upon reprotonation with concentrated H2SO4, monitoring by 11B NMR spectroscopy shows that (a) these four very high-field resonances are lost, (b) a completely new, more compact, spectrum is obtained,‡ and (c) there is now only one singlet in the 11B NMR spectrum, at d(11B) = 213.1. We propose that these features result from the loss of the hexagonal pyramidal feature and the generation of a neutral conjugate acid [S2B18H20] (compound 3) which consists of two more conventional (though mutually diVerent) eleven-vertex arachno subclusters IIC and IID joined as in schematic III.There is a precedent for both cluster types in thiaborane cluster chemistry,5 and the NMR assignments‡ are not inconsistent with this formulation.In this scheme, the common vertex, designated B in IIA and BH in III (see also Scheme 1), is protonated on conversion of anion 2 to neutral 3. Concomitant with this, there is a diamond–square–diamond (DSD) intracluster flexing (bold lines in Scheme 1), and the twenty-vertex macropolyhedral unit thereby acquires a somewhat reduced inter-subcluster intimacy. Interestingly, deprotonation of neutral 3 with tmnda (N,N,N9,N9-tetramethylnaphthalene-1,8-diamine) does not immediately regenerate the starting anion 2, but immediately and quantitatively generates a species 4 which exhibits a 1:2:2:2:2:2:1:2:2:2 relative intensity pattern in its 11B NMR spectrum.§ This may result from the conversion of neutral 3 to a fluxional anion 4.Such a fluxionality could for example arise from an arachno(IIC)/arachno(IID)�arachno(IID)/ arachno(IIC) interconversion as in Scheme 1: clusters of types S B S IIB IIC H S B S IIA anion 2 H H H S IID IIC and IID are very closely related, and deprotonation of the right-hand subcluster of structure III and adjustment of the two open-face connectivities to the vertex designated BH would readily achieve interconversion of the two subcluster types.Over 3 h at 294–297 K in CDCl3 solution, this fluxional species 4 reverts quantitatively to the original anion 2 with its unusual hexagonal pyramidal feature. This reversion entails an unusual movement of hydrogen (presumably as a proton) from an open-face bridging position to an apical site, and a reversal of the original DSD process (Scheme 1).In sum, the overall protonation–deprotonation sequence appears to result in the deconstruction and the reconstruction of the very unusual hexagonal pyramidal feature. This second incidence of a {BH} vertex wi a cluster connectivity to six other boron atoms in the context of an open cluster, but now one which can apparently be readily disassembled and reassembled, reinforces ideas 2 that the hexagonal pyramidal {B7} unit may be of fundamental significance and can therefore be used as a building block in future borane-based architecture.We are currently attempting to elucidate further the two intermediates 3 and 4 associated with this fundamental cluster disassembly–reassembly sequence, as well as attempting to devise entries into other systems that may exhibit this type of feature. Acknowledgements Contribution no. 72 from the R¡ ez¡-Leeds Anglo–Czech Polyhedral Collaboration (ACPC).We thank the EPSRC (Grant nos. F78323, J56929 and K05818) and the Grant Agency of the Academy of Sciences of the Czech Republic (Grant no. A 403 2701) for support, the Royal Society (London), the Czech Academy of Sciences, and Borax Research (now Borax Europe Ltd.) for assistance with reciprocal visits, and Simon A. Barrett for kind assistance with NMR spectroscopy. Scheme 1 B S H S H B S H S H B S H S H H H protonation DSD regeneration DSD deprotonation anion 2 fluxionality neutral 3 H anion 4 H S BH S H H H III neutral 3J.Chem. Soc., Dalton Trans., 1998, 2965–2967 2967 Notes and references * Crystals of [N(PPh3)2][S2B18H19], C36H49B18NP2S2, M = 816.40, from CHCl3–OEt2, triclinic, space group P1� , a = 1092.28(8), b = 1421.90(12), c = 1532.04(10) pm, a = 70.977(6), b = 88.202(6), g = 80.321(6)8, U = 2216.8(3) Å3, Z = 2, T = 200(2) K, 5851 independent reflections collected on a Stoe STADI4 diVractometer in the range 3.05 < q < 608 were used in calculations after Lorentz-polarisation and absorption corrections (m = 1.97 mm21, based on azimuthal y-scans).The borane anion possesses a pseudo-mirror plane [passing through atoms B(1), B(9), S(11), B(19), B(39) and B(109)] which causes a 50 : 50 disorder {S(99)} : {B(119)H(109,119)} within the eleven-vertex subcluster. Thus the 9- and 10-positions of this subcluster were refined as 50 :50 B:S atoms. Final wR2 = 0.1232 for all unique data, conventional R = 0.0446 for F values of 4941 reflections with Fo 2 > 2s(Fo 2).CCDC reference number 186/1101. See http://www.rsc.org/suppdata/dt/1998/2965/ for crystallographic files in .cif format. † 11B and 1H NMR data for anion 2 [formally the nido-99-thiaundecaborano-( 79,89 : 5,6)-iso-(116kc·VII Ò)-arachno-11-thiaundecaboranate- (12) anion]; CD3CN, 294–297 K (Note: [PPh4]1 salt, not [N(PPh3)2]1 salt) {ordered as: tentative assignment d(11B) relative to X 32.083971 MHz [d(1H) of directly attached hydrogen]}: BH(29) 116.5 [13.81], BH(49) 13.0 [13.43], BH(109) 11.9 [12.83], BH(1) 10.9 [13.91], B(5) ca. 23.9 [conjuncto position], BH(59) ca. 23.9 [12.74], BH(9) 28.6 [13.13], B(6) 214.8 [conjuncto position], BH(8) 215.3 [12.33], BH(10) 215.8 [12.28], BH(119) 217.6 [11.13], BH(69) 221.1 [11.51], BH(39) 223.2 [20.38], BH(19) 223.6 [11.805], BH(7) ca. 247.4 [10.52], BH(4) 247.4 [10.11], BH(2) 248.7 [10.37], BH(3) 249.7 [10.15], with m-H(7,8), (4,10) and (109,119) at d(1H) 21.45, 21.70 and 21.79 respectively; assignments by homo- and hetero-nuclear 11B and 1H NMR experiments. ‡ 11B and 1H NMR data for neutral 3 [formally m-(79,8)-arachno-99- thiadecaborano-(69 : 7)-nido-10-thiaundecaborane], CDCl3, 294–297 K {ordered, assigned and referenced as above}: BH(3) 115.9 [13.69], BH(29) 115.3 [14.55], BH(5) 111.1 [13.57], BH(49) 1 9.1 [13.65], BH(11) 17.3 [13.39], BH(8) 15.6 [13.13], BH(6) 21.9 [13.09], BH(59) 25.1 [12.58], BH(8) 25.95 [12.60], BH(109) 28.5 [12.41], BH(1) and BH(69) both ca. 211.6 [12.98 and 12.79], B(69) 213.1 [conjuncto position], BH(2) and BH(39) both ca. 219.5 [12.47 and 11.61], BH(7) 222.2 [11.08], BH(4) 224.2 [11.54], BH(19) 234.9 [11.52], and m-H(79,89), (8,9) and (59,109) at d(1H) 11.00, 20.88 and 21.14 respectively. § d(11B) NMR values (and relative intensities) for the fluxional anion 4, which exhibits time-average two-fold symmetry (see Scheme 1): CDCl3, 294–297 K: ca. 14.8 (2BH), ca. 14.8 (1B), 12.3 (2BH), 25.3 (2BH), 26.8 (2BH), 27.9 (2BH), 210.0 (1BH), 211.1 (2BH), 238.0 (2BH) and 241.6 (2BH); all had doublet structures arising from couplings 1J(11B–1H) in the range ca. 135 to ca. 165 Hz, except for the resonance of relative intensity 1B at d(11B) ca. 14.8; 1H-{11B} NMR work was precluded because of the relatively rapid reversion of 4 to regenerate 2. 1 See, for example, J. Bould, J. D. Kennedy and M. Thornton-Pett, J. Chem. Soc., Dalton Trans., 1992, 563; J.D. Kennedy and B. S. tíbr, in Current Topics in the Chemistry of Boron, ed. G. W. Kabalka, Royal Society of Chemistry, Cambridge, 1994, pp. 285–292; B. S. tíbr, J. D. Kennedy, E. Drdáková and M. Thornton-Pett, J. Chem. Soc., Dalton Trans., 1994, 229; J. D. Kennedy, in The Borane–Carborane–Carbocation Continuum, ed. J. Casanova, Wiley, New York, 1998, ch. 3, pp. 85–116. 2 P. Kaur, J. Holub, N. P. Rath, J. Bould, L. Barton, B. S. tíbr and J. D. Kennedy, Chem. Commun., 1996, 273. 3 R. E. Williams, Adv. Inorg. Chem. Radiochem., 1976, 18, 64; K. Wade, Adv. Inorg. Chem. Radiochem., 1976, 18, 1. 4 K. Nestor, X. L. R. Fontaine, N. N. Greenwood, J. D. Kennedy and M. Thornton-Pett, J. Chem. Soc., Dalton Trans., 1991, 2657; see also B. S. tíbr, J. Holub, T. Jelínek, X. L. R. Fontaine, J. Fusek, J. D. Kennedy and M. Thornton-Pett, J. Chem. Soc., Dalton Trans., 1996, 1741. 5 J. Holub, A. E. Wille, B. S. tíbr, P. J. Carroll and L. G. Sneddon, Inorg. Chem., 1994, 33, 4920; T. Jelínek, J. D. Kennedy, B. S.tíbr and M. Thornton-Pett, Angew. Chem., Int. Ed. Engl., 1994, 33, 1599; J. Chem. Soc., Chem. Commun., 1995, 1665; P. Kaur, J. D. Kennedy, M. Thornton-Pett, T. Jelínek and B. S. tíbr, J. Chem. Soc., Dalton Trans., 1996, 1775. Communication 8/0579

ISSN:1477-9226    

DOI:10.1039/a805791e

出版商:RSC    

年代:1998

数据来源: RSC

摘要:

DALTON FULL PAPER J. Chem. Soc., Dalton Trans., 1998, 2969–2972 2969 Syntheses and structures of new copper(I) selenide clusters [Cu32Se16(PPh3)12], [Cu52Se26(PPh3)16] and [Cu72Se36(PPh3)20] † Andreas Eichhöfer and Dieter Fenske * Institut für Anorganische Chemie, Universität Karlsruhe, Engesserstraße Geb. Nr. 30.45, D-76128 Karlsruhe, Germany Received 15th May 1998, Accepted 20th July 1998 The reaction of Se(SiMe3)2 with copper(I) acetate in the presence of PPh3 yielded the new clusters [Cu32Se16(PPh3)12] 2, [Cu52Se26(PPh3)16] 3 and [Cu72Se36(PPh3)20] 4 depending on the conditions.The molecular structures of 2, 3 and 4 have been characterised by X-ray crystallography and compared with the known cluster [Cu146Se73(PPh3)30] 1 obtained from the same reaction under diVerent conditions. Introduction The synthesis of large metal-containing clusters is of considerable interest because these compounds are good model systems for investigating the chemical and physical properties of solid state materials. This is particularly important for nanoparticulate binary metal chacogenides which show quantum eVects such as size-dependent band gaps.1 We have been successful in preparing a large number of copper selenide clusters using the reaction of CuX (X = Cl or OAc) with Se(SiMe3)2 and PR3 (R = organic group) outlined in eqn.(1). Some of the compounds we have structurally charac- 2CuX1yPR31Se(SiMe3)2 THF or Et2O 22SiMe3X [Cu2Se(PR3)y]n (1) terised are as follows: [Cu12Se6(PnPr3)8], [Cu26Se13(PEt2Ph)14], [Cu30Se15(PiPr3)12], [Cu36Se18(PtBu3)12], [Cu44Se22(PEt2Ph)18] and [Cu70Se35(PEt3)22].A noteworthy example is the structure of [Cu146Se73(PPh3)30] 1 one of the largest known structurally characterised clusters.2b From the range of products listed above it can clearly be seen that the phosphine has a profound role in determining the size of the copper selenide cluster. This may be attributed both to steric eVects due to the characteristic Tolman cone angle and to kinetic factors arising from stabilisation of copper/phosphine/ acetate clusters.We now wish to report that other factors such as ligand concentration, temperature and solvent also influence cluster formation. For this study we have conducted a detailed investigation of the CuOAc/Se(SiMe3)2/PPh3 system from which we have previously isolated 1. Experimental Standard Schlenk techniques were employed throughout the syntheses using a double-manifold vacuum line with high purity dry nitrogen.The solvents diethyl ether, tetrahydrofuran and diglyme were dried over sodium–benzophenone and distilled under nitrogen. CuOAc3 and Se(SiMe3)2 4 were prepared according to standard literature procedures. Syntheses [Cu32Se16(PPh3)12] 2. PPh3 (1.73 g, 6.6 mmol) was added to a suspension of copper(I) acetate (0.27 g, 2.2 mmol) in THF (40 † Dedicated to Professor Brian F. G. Johnson on the occasion of his 60th birthday.mL). A clear colourless solution was formed. Upon the addition of Se(SiMe3)2 (0.25 mL, 1.1 mmol) the solution gradually turned yellow, and subsequently dark brown. After 1 d, 20 mL of diethyl ether were added to a solution cooled to 4 8C. After 3 d at this temperature, dark yellow crystalline plates of 2 were formed. Yield: 54% (Found: C, 39.02; H, 2.90. C216H180Cu32- P12Se16 requires C, 40.26; H, 2.80%). [Cu52Se26(PPh3)16] 3. PPh3 (1.73 g, 6.6 mmol) was added to a suspension of copper(I) acetate (0.27 g, 2.2 mmol) in THF (40 mL).A clear colourless solution was formed. Upon addition of Se(SiMe3)2 (0.25 mL, 1.1 mmol), the solution gradually turned yellow, and subsequently dark brown. After 1 d, 20 mL of diethyl ether were added and the solution was kept at room temperature. After 2 d, dark yellow crystalline plates of 2 were formed in low yield, and after 2 additional days, black crystals of 3 began to grow as the major product of the reaction.Yield: 50% (Found: C, 37.1; H, 2.80. C288H240Cu52P16Se26 requires C, 36.2; H, 2.5%). [Cu72Se32(PPh3)20] 4. PPh3 (1.28 g, 4.9 mmol) was added to a suspension of copper(I) acetate (0.20 g, 1.63 mmol) in diglyme (30 mL). On addition of Se(SiMe3)2 (0.25 mL, 1.1 mmol) the solution gradually turned clear, changing colour from yellow to dark brown. After 3 d, black crystals of 4 were formed. Yield: 51% (Found: C, 35.49; H, 3.33. C360H300Cu72P20Se36 requires C, 34.15; H, 2.37%).Two diglyme molecules were localised and refined in the X-ray structure analysis, but there is evidence for the existence of other highly disordered solvent molecules which may be responsible for the high carbon and hydrogen content found in the elemental analysis. X-Ray structural analyses Single crystal X-ray structural analyses of compounds 2–4 were performed using a Stoe-IPDS diVractometer (Mo-Ka radiation) equipped with an image plate area detector and a rotating anode.Structure solution and refinement were carried out using SHELXS 865 and SHELXL 936 software using direct methods techniques. The weighting scheme applied was of the form w = 1/[s2(F2) 1 (aP)2 1 bP] [a,b = refined variables, P = 1 3 – (Fo 2,0) 1 2 3 – Fc 2]. All calculations were performed on a Silicon Graphics INDY computer. Molecular diagrams were prepared using the SCHAKAL 92 program.7 Compound 2. C216H180Cu32P12Se16?4C4H8O, M = 6730.6, triclinic, space group P1� , a = 19.319(4), b = 21.531(4), c = 33.931(7) Å, a = 79.58(3), b = 80.11(3), g = 74.37(3)8, U =2970 J.Chem. Soc., Dalton Trans., 1998, 2969–2972 13254(5) Å3, T = 200 K, Z = 2, Dc = 1.675 g cm23, F(000) = 6496, m(Mo-Ka) = 4.817 mm21, 2qmax = 558, 31788 independent reflections measured (Rint = 0.0859) and 22364 I > 2s(I). The structure was solved by direct methods and refined on F2. All Se, Cu and P atoms were refined anisotropically, all C atoms isotropically and positions for H atoms were calculated except those of THF to yield R = 0.0823, wR = 0.2275, S = 0.991.Compound 3. C288H240Cu52P16Se26?2.5C4H8O, M = 9721.53, triclinic, space group P1� , a = 22.336(5), b = 23.203(5), c = 37.894(8) Å, a = 91.00(3), b = 106.95(3), g = 99.06(3)8, U = 18510(6) Å3, T = 200 K, Z = 2, Dc = 1.741 g cm23, F(000) = 9350, m(Mo-Ka) = 5.580 mm21, 2qmax = 51.88, 57286 independent reflections measured (Rint = 0.0892) and 39806 I > 2s(I). The structure was solved by direct methods and refined on F2.All Se, Cu and P atoms were refined anisotropically, all C atoms isotropically and positions for H atoms were calculated except those of THF to yield R = 0.0823, wR = 0.2373, S = 1.023. Compound 4. C360H300Cu72P20Se36?2C6H16O3, M = 12931.18, triclinic, space group P1� , a = 25.632(5), b = 33.806(7), c = 34.904(7) Å, a = 72.69(2), b = 87.16(2), g = 72.64(2)8, U = 27534(1) Å3, T = 190 K, Z = 2, Dc = 1.560 g cm23, F(000) = 12440, m(Mo-Ka) = 5.185 mm21, 2qmax = 458, 66463 independent reflections measured (Rint = 0.073) and 45921 I > 2s(I).The structure was solved by direct methods and refined on F2. All Se, Cu and P atoms were refined anisotropically, all C atoms isotropically to yield R = 0.0829, wR = 0.2271, S = 0.939. CCDC reference number 186/1098. See http://www.rsc.org/suppdata/dt/1998/2969/ for crystallographic files in .cif format. Results and discussion We have previously shown that copper(I) acetate reacts with 5 equivalents of PPh3 and 0.5 equivalent of Se(SiMe3)2 in THF at room temperature to yield [Cu146Se73(PPh3)30] 1,2b after a number of days.We now find that similar reactions with only 3 equivalents of PPh3 produce dark coloured solutions. Addition of diethyl ether and cooling to 4 8C leads to the crystallisation of [Cu32Se16(PPh3)12] 2. Keeping these mixed solvent solutions at room temperature leads to the formation of [Cu52Se26- (PPh3)16] 3. [Cu72Se36(PPh3)20] 4 crystallises from the same reactions if diglyme is used instead of THF as a solvent.Compound 2 crystallises as dark yellow plates in the triclinic Fig. 1 Molecular structure of [Cu32Se16(PPh3)12] 2. Cu11: blue; Se22: red; P: green. Cu–Cu contacts and carbon atoms are omitted for clarity. Cu, Se and P atoms are labelled with blue, red and green numbers respectively. space group P1� . Fig. 1 shows thure with the exception of the carbon atoms. The selenium atoms Se1–Se16 form a flattened polyhedron compromising of non-bonded Se3 triangles (Se–Se 3.984 to 4.645 Å). There are four distinct copper environments.Twenty-four of the copper atoms cap the selenium triangles, of these sixteen (Cu1, Cu2, Cu3, Cu4, Cu6, Cu11, Cu15, Cu16, Cu19, Cu21, Cu22, Cu24, Cu27, Cu28, Cu29, Cu31) are bonded only to selenium producing a distorted trigonal coordination geometry. These copper atoms lie below the plane formed by the selenium atoms to which they are coordinated such that they are within the selenium polygon.A further eight (Cu5, Cu8, Cu12, Cu14, Cu17, Cu18, Cu20, Cu25) are also bonded to a phosphine ligand (P2, P3, P4, P6, P7, P8, P9, P11) giving a tetrahedral environment with the copper atoms situated on the exterior of the selenium faces. Four copper atoms (Cu7, Cu10, Cu13, Cu23) are coordinated in a quasilinear fashion to two selenium atoms (Se3–Cu23–Se15 156.13, Se13–Cu7–Se10 159.28, Se1–Cu10–Se2 170.2, Se9–Cu13–Se14 170.528).The four remaining copper atoms Cu9, Cu26, Cu30 and Cu32 are bonded in a distorted trigonal environment to two selenium atoms and one phosphorus atom from PPh3. The Cu– Se distances range from 2.295 to 2.930 Å. For the selenium atoms, the coordination numbers five (Se3, Se4, Se10, Se11, Se12, Se13) and six (Se1, Se2, Se5, Se6, Se7, Se8, Se9, Se14) are found. Compound 3 forms large black or smaller brown crystals which crystallise in the triclinic space group P1� (Fig. 2).By analogy with 2, the framework in 3 can also be seen to be made up from edge-sharing Se3 triangles. The selenium atoms form an approximately spherical polyhedron with three additional atoms (Se1, Se8, Se18) located within the cavity. The Se–Se distances range from 3.803 to 4.707 Å and are essentially nonbonding. Thirty-six of the fifty-two copper atoms are bonded only to the selenium atoms (Cu–Se 2.19–2.91 Å) and are located within the lattice, exhibiting either a linear (Cu12, Cu28, Cu36, Cu41, Cu43), distorted trigonal (Cu1, Cu3, Cu4, Cu7, Cu10, Cu11, Cu15, Cu18, Cu21, Cu24, Cu25, Cu26, Cu29, Cu30, Cu33, Cu35, Cu38, Cu39, Cu40, Cu42, Cu44, Cu46, Cu47, Cu48, Cu49, Cu50, Cu52) or a distorted tetrahedral structure (Cu8, Cu13, Cu20, Cu51).Coordination numbers for the selenium atoms vary from four (Se26) to five (Se3, Se4, Se7, Se9, Fig. 2 Molecular structure of [Cu52Se26(PPh3)16] 3. Cu11: blue; Se22: red; P: green. Cu–Cu contacts and carbon atoms are omitted for clarity.Cu, Se and P atoms are labelled with blue, red and green numbers respectively.J. Chem. Soc., Dalton Trans., 1998, 2969–2972 2971 Se11, Se12, Se14, Se15, Se16, Se17, Se21, Se22, Se23, Se24, Se25), six (Se5, Se8, Se10, Se13, Se19), seven (Se1, Se2, Se6, Se18) and eight (Se20). Those copper atoms which are bonded to the sixteen phosphine ligands are either coordinated in a nearly trigonal planar fashion, for those that are bonded to one phosphorus atom and two selenium atoms (Cu2, Cu5, Cu9, Cu14, Cu22, Cu23, Cu34, Cu37, Cu45), or in a distorted tetrahedral way for those that are coordinated by three selenium atoms and one phosphorus atom (Cu6, Cu16, Cu17, Cu19, Cu27, Cu31, Cu32).As in structure 2 copper atoms which are coordinated by three selenium atoms only are shifted to the interior of the cluster cage. Copper atoms additionally bonded to one phosphorus atom are located on the cluster surface. In contrast to the wide range of Cu–Se distances, the Cu–P bonds all lie within a narrow range (2.215–2.284 Å).Compound 4 forms black crystals in the triclinic space group P1� . The thirty-six selenium atoms build up a nearly trigonal prismatic polyhedron with trigonal selenium faces (Fig. 3). In contrast to 2 and 3, 4 possesses a symmetrical layer-type structure with three layers each containing ten, sixteen and ten selenium atoms. The packing of the layers is nearly hexagonal, of the type ABA.Only Se30 and Se32 do not follow this trend. The non-bonding Se–Se distances range from 3.81 to 4.71 Å, and the distance between the layers is approximately 3.6 Å. The holes in the selenium lattice are filled with seventy-two copper atoms (Fig. 4). As in 2 and 3, there are three coordination environments for the copper atoms bonded only to selenium. Cu4 is linearly coordinated, forty copper atoms (Cu1, Cu2, Cu3, Cu7, Cu8, Cu9, Cu11, Cu12, Cu13, Cu14, Cu15, Cu17, Cu19, Cu20, Cu21, Cu22, Cu23, Cu24, Cu25, Cu27, Cu29, Cu30, Cu34, Cu38, Cu41, Cu42, Cu43, Cu46, Cu47, Cu49, Fig. 3 View perpendicular (above) and parallel (below) to the hexagonal ABA layering of the selenium atoms in 4. Non-bonding Se–Se contacts are shown only in the layers. Distance between the layers ca. 3.61 Å. Cu50, Cu54, Cu55, Cu58, Cu59, Cu61, Cu63, Cu65, Cu71, Cu72) exhibit a distorted trigonal environment and eleven copper atoms (Cu6, Cu16, Cu33, Cu36, Cu37, Cu39, Cu44, Cu51, Cu52, Cu57, Cu70) are coordinated in a distorted tetrahedral fashion (Cu–Se 2.354–2.921 Å).For those copper atoms bonded to selenium and phosphorus atoms, six copper atoms (Cu5, Cu18, Cu31, Cu32, Cu35, Cu69) exhibit the coordination number three (2 Se, 1 P atom), and fourteen copper atoms (Cu10, Cu26, Cu28, Cu40, Cu45, Cu48, Cu53, Cu56, Cu60, Cu62, Cu64, Cu66, Cu67, Cu68) are coordinated in a distorted tetrahedral geometry by three selenium atoms and one phosphorus atom of the PPh3 ligands (Cu–P 2.211–2.242 Å).The copper atoms coordinated to phosphine ligands are shifted to the surface of the cluster in contrast to pure trigonal coordinated ones, as was also observed for 2 and 3. If we assume a maximum Cu–Se distance for a bonding interaction of 2.92 Å we find a wide spread of selenium coordination numbers ranging from four to ten: four (Se11, Se4), five (Se1, Se5, Se7, Se10, Se22, Se26, Se27, Se31, Se36), six (Se2, Se3, Se4, Se6, Se8, Se9, Se24, Se28, Se29, Se30, Se32, Se33, Se34, Se35), seven (Se13, Se14, Se15, Se17, Se18, Se21, Se23, Se25), eight (Se16), nine (Se19) and ten (Se19).We have also previously observed a large range of selenium coordination numbers in other copper selenide clusters. The larger the cluster size the higher the average coordination numbers for selenium atoms leading to more dense copper selenium packing. The [Cu70Se35(PEt3)21] 2a structure is related to 4 simply by removal of a Cu2Se unit (Cu18, Cu35, Se30) which caps one of the corners of the CuSe polyhedron.Furthermore the two selenium atoms (Se30 and Se32) in this fragment are not part of the close-packed selenium array. This structural motif, in which a rectangular selenium frame is grafted onto one edge of the triangle is also present in 1 therefore 4 is in fact a fragment of the larger cluster. Conclusion In summary the Se cages for 2 and 3 diVer significantly from the symmetrical ones found in 4 and 1. This suggests that 2 and 3 cannot be intermediate structures in the formation of the layered clusters 4 and 1.Whilst there is a clear similarity between the structures of 4 and 1 no relationship between 2 and 3 is evident. Further, one can find analogies between all of the Fig. 4 Molecular structure of [Cu72Se36(PPh3)20] 4. Cu11: blue; Se22: red; P: green. Cu–Cu contacts and carbon atoms are omitted for clarity. Cu, Se and P atoms are labelled with blue, red and green numbers respectively.2972 J.Chem. Soc., Dalton Trans., 1998, 2969–2972 clusters in the coordination modes of the copper atoms. In common with other copper selenide clusters, the copper atoms prefer a distorted trigonal planar coordination of selenium atoms. At the cluster surfaces the copper atoms coordinate the selenium atoms in a slightly distorted triangular fashion, with some shifts towards the interior of the selenium polyhedron. Copper atoms that are linked to a phosphine ligand adopt the predicted tetrahedral geometry and are located well outside the Se lattice.The preference of trigonal coordination can also be seen from the fact that copper atoms in b-Cu2Se leave their ideal tetdral sites in the transition from a cubic high temperature phase to the lower symmetry a-Cu2Se phase to take up a more trigonal coordination.8 Although the final structure determination of the low temperature a-Cu2Se phase is still a point of discussion the best refinement of the data is achieved by assuming a cubic lattice of selenium atoms with the copper atoms occupying the tetrahedral and trigonal holes in a ratio 0.37 to 0.63.9 In contrast, 4 and 1 possess a hexagonal lattice of selenium atoms, a structure type known for the low temperature phase of Cu2S.10 Therefore it is possible that these cluster molecules display a new structure type of an as yet unknown Cu2Se polymorph. Hence cluster formation of 2 and 3 can be seen to be more like an unordered spontaneous aggregation process whereas the crystallisation of 4 and 1 is comparable to the formation of an ordered solid state structure.The results presented herein show the complex dependence of these reactions on a range of parameters. This knowledge will assist us in the design of rational syntheses for even larger clusters. Acknowledgements We are grateful to the Deutsche Forschungsgemeinschaft (SFB 195) and to the Bundesministerium für Bildung und Forschung (BMBF) for financial support of this work. References 1 H. Weller, Angew. Chem., 1996, 108, 1159; A. P. Alivisatos, Science, 1996, 271, 933; E. Kim, Y. Xia and G. M. Whitesides, Nature (London), 1995, 376, 581; M. Trau, D. A. Saville and I. A. Aksay, Science, 1996, 272, 706. 2 (a) D. Fenske and H. Krautscheid, Angew. Chem., 1990, 102, 1513; (b) D. Fenske, H. Krautscheid, G. Baum and M. Semmelmann, Angew. Chem., 1993, 105, 1364; (c) S. Dehnen, A. Schäfer, D. Fenske and R. Ahlrichs, Angew. Chem., 1994, 106, 786; (d ) S. Dehnen and D. Fenske, Chem. Eur. J., 1996, 2, 1407; (e) A. Deveson, S. Dehnen and D. Fenske, J. Chem. Soc., Dalton Trans., 1997, 4491. 3 D. A. Edwards and R. Richards, J. Chem. Soc., Dalton Trans., 1973, 2463. 4 H. Krautscheid, Ph.D. Thesis, Universität Karlsruhe, 1991. 5 G. M. Sheldrick, SHELXS 86, Program for the Solution of Crystal Structures, University of Göttingen, 1986. 6 G. M. Sheldrick, SHELXL 93, Program for Crystal Structure Determination, University of Göttingen, 1993. 7 E. Keller, SCHAKAL 92, A Computer Program for the Graphic Representation of Molecular and Crystallographic Models, University of Freiburg, 1992. 8 S. Kashida and J. Akai, J. Phys. C: Solid State Phys., 1988, 21, 5329. 9 K. Yamamoto and S. Kashida, J. Solid State Chem., 1991, 93, 202. 10 H. T. Evans, jun., Z. Kristallogr., 1979, 150, 299. Paper 8/03675F

ISSN:1477-9226    

DOI:10.1039/a803675f

出版商:RSC    

年代:1998

数据来源: RSC

摘要:

DALTON FULL PAPER J. Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 2973 Theoretical and experimental studies of the protonated terpyridine cation. Ab initio quantum mechanics calculations, and crystal structures of two diVerent ion pairs formed between protonated terpyridine cations and nitratolanthanate(III) anions † Michael G. B. Drew,*a Michael J. Hudson,a Peter B. Iveson,a Mark L. Russell,a Jan-Olov Liljenzin,b Mats Skålberg,b Lena Spjuth b and Charles Madic c a Department of Chemistry, University of Reading, PO Box 224, Whiteknights, Reading, UK RG6 6AD.E-Mail: m.g.b.drew@reading.ac.uk b Department of Nuclear Chemistry, Chalmers University of Technology, S-412 96 Goteborg, Sweden c Commissariat à l’Energie Atomique, Direction du Cycle du Combustible, B.P. 171, 30207 Bagnols-sur-Cèze, Cedex, France Ab initio quantum mechanics calculations have been carried out on all possible conformations of the terpyridine ligand and its mono- and di-protonated forms.Results show that the lowest energy form of the ligand is when the N]C]C]N torsion angles are trans but that in the protonated forms, the cis arrangement is prevalent being stabilised by intramolecular N]H? ? ? N hydrogen bonds. Results are consistent with the experimental crystal structure data found in the literature and also with two crystal structures reported here in which two diVerent ion pairs formed between protonated terpyridine cations and lanthanate(III) nitrate anions have been prepared and analysed structurally.Compound 1 consists of discrete diprotonated terpyridine cations and hexanitratolanthanate anions, namely 3[H2terpy]212[La(NO3)6]32?3H2O. Water molecules are present as hydrogen bond acceptors in the diprotonated terpyridyl cavities. Each lanthanum atom is 12-co-ordinate and the La]O bond lengths vary between 2.609(11) and 2.765(10) Å. Compound 2 consists of a diprotonated [H2terpy]21 cation together with a [Sm(terpy)(NO3)4]2 anion, and a NO3 2 anion which is present as a hydrogen bond acceptor in the diprotonated terpyridyl cavity.The samarium atom is 11-co-ordinate, the Sm]O bond lengths vary between 2.494(5) and 2.742(5) Å while the Sm]N bond lengths vary between 2.626(4) and 2.650(5) Å. Introduction There is much current interest in the separation of lanthanide( III) and actinide(III) complexes by solvent extraction routes.1 Various oligoamines have been shown to co-extract lanthanides and actinides from nitric acid solutions into an organic phase.Among these are the oligoamine ligands 2,4,6-tris(4-tert-butyl- 2-pyridyl)-1,3,5-triazine 2 ligand L1 and 2,29:69,20-terpyridine L2. It is thought that the species extracted at low levels of acidity are simple metal co-ordination complexes, for which there is plenty of structural evidence.3 It is clear that the lanthanides (and by implication the similarly sized actinides) can fit into the tridentate cavity of the terpyridyl moiety.Thus we have previously reported the structure of Ce(NO3)4L1 (ref. 2) and in addition complexes of L2 with lanthanides are well known [e.g. LnCl3(L2)?xH2O (Ln = La, Ce, Nd, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu or Y)].4 For Ln = La–Nd, x = 5, for Ln = Tb–Lu, x = 4 and for Ln = Sm or Gd there are between four and five water molecules in the co-ordination sphere. Also isolated was a dimeric species in which the Sm is co-ordinated to three N atoms from L2, two bridging and two terminal chloride anions and one water molecule.Although direct co-ordination of the metal ions to the polyamine ligands is thought to be necessary for the selective extraction of actinides over the lanthanides, the need still remains to establish the nature of the species which may be formed at higher acid concentrations, especially those in which the lanthanide or actinide ion exists as an anion such as [Ln(NO3)6]32 or the [LnL(NO3)4]2 ion-pair. It has been demonstrated that terpyridine-type ligands in synergistic combination † Non-SI unit employed: cal = 4.184 J.with 2-bromodecanoic acid show some selectivity for An(III) over Ln(III) ions.5 The species involved in these extractions are not known with any certainty. We have adopted a theoretical and experimental approach to the identification of these species. We are using quantum mechanical methods to investigate the conformational preferences of the ligand and also the likely structures of the protonated forms.In addition we are trying to prepare solid complexes in order to provide evidence for the kind of species which is involved in this type of extraction. In this paper we report the results from our quantum mechanical calculations and also the crystal structures of two diVerent ion pairs of diprotonated terpyridine and anionic lanthanide nitrates (1 and 2) both of which contain free diprotonated terpyridine moieties unco-ordinated to a metal ion.The asymmetric unit of 1 contains three diprotonated terpyridyl cations, two hexanitratolanthanate(III) anions and three water molecules, 3[H2L2]212[La(NO3)6]32?3H2O while the asymmetric unit of 2 contains an [(H2L2)(NO3)]1 cation together with a [SmL2- (NO3)4]2 anion. Experimental The compounds Ln(NO3)3?6H2O (99.99%), Sm(NO3)3?6H2O (99.99%) and 2,29:69,20-terpyridine (L2) (98%) were used as received from Aldrich. 2-Bromodecanoic acid (98%) was purchased from Fluka and used without further purification.Preparation of complex 1 This compound was prepared from [H2L2][NO3]2 which was initially prepared as a solid: L2 (0.25 g, 0.001 mol) was dissolved2974 J. Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 Table 1 Crystal data and structure refinement for compounds 1 and 2 Formula Empirical formula MT /K l/Å Crystal system Space group a/Å b/Å c/Å b/8 U/Å3 Z, Dc/g cm23 m/mm21 F(000) Size/mm q Range/8 hkl Range Reflections observed, Rint Unique reflections Weighting scheme a, b * Reflections, parameters Final indices [I > 2s(I)] R1, wR2 All data R1, wR2 Largest diVerence peak, hole/e Å23 1 3[H2L2]212[La(NO3)6]32?3H2O C45H45La2N21O39 1781.84 293(2) 0.710 73 Monoclinic P21 10.826(9) 30.00(2) 11.140(9) 104.76(1) 3498(5) 2, 1.692 1.316 1776 0.2 × 0.2 × 0.5 2.82 to 25.07 0 < h < 11, 235 < k < 35, 213 < l < 12 8750, 0.0508 8570 0.144, 35.48 8570, 965 0.0582, 0.1779 0.0657, 0.1901 0.928, 21.615 2 [(H2L2)(NO3)]1[Sm(L2)(NO3)4]2 C30H24SmN11O15 928.95 293(2) 0.710 73 Monoclinic P21/n 17.29(2) 10.230(9) 20.31(2) 106.53(1) 3445(6) 4, 1.791 1.796 1852 0.25 × 0.25 × 0.20 1.83 to 25.94 0 < h < 21, 212 < k < 12, 224 < l < 23 9872, 0.0289 5838 0.181, 25.94 5838, 515 0.0352, 0.0971 0.0514, 0.1080 1.116, 20.984 * Weighting scheme w = 1/[2(Fo 2) 1 (aP)2 1 bP], where P = (Fo 2 1 2Fc 2)/3.in methanol (1 ml). The addition of 69% AnalaR nitric acid (0.14 ml, 0.0022 mol) to a stirred solution of the ligand resulted in the formation of a precipitate (yield 82%).6 After the addition of methanol (2 ml), the solid was filtered oV and dried under vacuum over calcium chloride. Then [H2L2][NO3]2 (0.05 g) was stirred for 10 min in acetonitrile (20 ml), heated to reflux and another solution containing La(NO3)3? 6H2O (0.0498 g, 0.0001 mol) dissolved in acetonitrile (15 ml) at ca. 40 8C was added dropwise to the stirred ligand solution. After the addition, the solution was stirred for 10 min and then allowed to stand at ambient temperature.Crystals suitable for X-ray analysis were deposited overnight at room temperature, 3[H2L2]212[La(NO3)6]32?3H2O (Found: C, 30.1; H, 2.7; N, 16.2. Calc. for C45H45La2N21O39: C, 30.3; H, 2.5; N, 16.5%). Preparation of complex 2 A solution of Sm(NO3)3?6H2O (0.077 g, 0.0002 mol) in acetonitrile (5 ml) was heated to around 40 8C and then added dropwise to a stirred solution of L2 (0.051 g, 0.0002 mol) and 2-bromodecanoic acid (0.087 g, 0.000 35 mol) in acetonitrile (5 ml), also at around 40 8C.A small amount of precipitate appeared on mixing and a further four quantities of acetonitrile (5 ml) were added to dissolve the solid. After a few minutes of heating the solution was allowed to cool down slowly in an oil bath and crystals were formed on standing overnight, [(H2L2)- (NO3)]1[Sm(L2)(NO3)4]2 (Found: C, 38.5; H, 2.55; N, 16.3. C30H24N11O15Sm requires C, 38.8; H, 2.6; N, 16.6%).Crystallography Crystal data for complexes 1 and 2 are given in Table 1, together with refinement details. Data for both crystals were collected with Mo-Ka radiation using the MARresearch Image Plate System. The crystals were positioned at 75 mm for 1 and 70 mm for 2 from the Image Plate. 95 Frames were measured at 28 intervals with a counting time of 2 min. Data analysis was carried out with the XDS program.7 The structures were solved using direct methods with the SHELXS program.8 In both structures the non-hydrogen atoms were refined with anisotropic thermal parameters.The locations of the nitrogen atoms in the rings were selected via thermal parameters and confirmed from the structure refinement by comparison with other assignments which gave higher R values and unreasonable thermal parameters. The hydrogen atoms bonded to carbon were included in geometric positions and given thermal parameters equivalent to 1.2 times those of the atom to which they were attached.An empirical absorption correction was made for both structures using the DIFABS program.9 In 2, the two extra protons in the cation were readily observed in a Fourierdi Verence map bonded to N(11) and N(31) and successfully included in the refinement with no constraints. By contrast in 1 these hydrogen atoms bonded to nitrogen (and also those bonded to the water oxygen atoms) could not be located definitively although positive areas of electron density were located in appropriate positions. In order to establish whether the data were of suYcient quality so that the hydrogen atoms bonded to nitrogen should be locatable (and that therefore we should draw the conclusion from their absence that they were not positioned on the nitrogen atom) we looked at the Fourierdi Verence map for the hydrogen atoms bonded to carbon in these cations.In the first 200 peaks, only 12 of the 33 hydrogen atoms could be located, i.e.peaks were within 0.5 Å of calculated positions. We conclude that the data are not of suYcient quality so that we would necessarily expect to find definitive positions for the hydrogen atoms bonded to nitrogen and we therefore conclude that their likely positions are equivalent to those in 2 [bonded to atoms N(11) and N(31)], which is consistent with the formation of hydrogen bonds to the water molecule. Therefore the hydrogens in the cations A, B and C were included in calculated positions bonded to these outer nitrogen atoms.The assignment of absolute structure in 1 was carried out by comparisons of two structures with opposite signs for y coordinates. The structure with the lowest R value was chosen. Both structures were refined on F2 till convergence using SHELXL.10 All calculations were carried out on a Silicon Graphics R4000 Workstation at the University of Reading. Relevant bond lengths in each structure are shown in Table 2. The hydrogen bonds are shown in Table 3.CCDC reference number 186/1074. See http://www.rsc.org/suppdata/dt/1998/2973/ for crystallographic files in .cif format.J. Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 2975 Results and Discussion Structure of 3[H2L2]212[La(NO3)6]32?3H2O 1 The asymmetric unit of 1 contains three diprotonated terpyridyl cations [H2L2]21, two [La(NO3)6]32 anions and three water molecules. Thus the protonated terpyridine ligand is not coordinated to the lanthanate(III) nitrate anion.Instead the two terminal pyridine nitrogen atoms are protonated. Each protonated terpyridyl cation encapsulates a water molecule with which it forms two strong hydrogen bonds. Two parts of the structure are shown in Figs. 1 and 2. Fig. 1 shows the environment of one protonated terpyridyl cation and Fig. 2 of one lanthanate(III) anion. It is noteworthy that the environments of the other cations and anions are almost identical such that there are three cations in the asymmetric unit together with just three water molecules each of which is similarly encapsulated.There are no additional water molecules in the unit cell, indicating that the role of the water molecules in the structure is to stabilise the cations by forming hydrogen bonds and also to stabilise the packing by additional hydrogen bonds to the anions. While it is not impossible to imagine a structure in which the lanthanate nitrate anions are hydrogen bonded directly to the diprotonated cation, clearly the present arrangement in which the interaction between cation and anion is mediated by the water molecules is more favoured.Table 2 Bond lengths (Å) for compounds 1 and 2 Compound 1 La(1)]O(13) La(1)]O(32) La(1)]O(22) La(1)]O(14) La(1)]O(37) La(1)]O(33) La(1)]O(26) La(1)]O(27) La(1)]O(18) La(1)]O(24) La(1)]O(36) La(1)]O(16) 2.609(11) 2.649(11) 2.663(11) 2.665(13) 2.672(12) 2.674(10) 2.683(11) 2.702(13) 2.726(11) 2.734(12) 2.741(12) 2.758(14) La(2)]O(45) La(2)]O(56) La(2)]O(52) La(2)]O(58) La(2)]O(66) La(2)]O(41) La(2)]O(47) La(2)]O(53) La(2)]O(43) La(2)]O(63) La(2)]O(68) La(2)]O(64) 2.635(11) 2.653(12) 2.666(10) 2.669(11) 2.692(11) 2.702(10) 2.706(11) 2.702(9) 2.733(12) 2.725(11) 2.762(12) 2.765(10) Compound 2 Sm(1)]O(72) Sm(1)]O(62) Sm(1)]O(51) Sm(1)]O(53) Sm(1)]O(83) Sm(1)]N(21B) 2.494(5) 2.537(5) 2.535(4) 2.554(4) 2.559(4) 2.626(4) Sm(1)]O(82) Sm(1)]N(11B) Sm(1)]N(31B) Sm(1)]O(74) Sm(1)]O(63) 2.627(5) 2.635(5) 2.650(5) 2.690(5) 2.742(5) Table 3 Hydrogen bond distances (Å) in compounds 1 and 2 Compound 1 O(1A) ? ? ? O(63) O(1A) ? ? ? N(31AIII) O(1A) ? ? ? N(11AIII) O(1B) ? ? ? O(66II) O(1B) ? ? ? N(11B) O(1B) ? ? ? N(31B) O(1B) ? ? ? O(16II) O(1C) ? ? ? O(38II) O(1C) ? ? ? N(31CIV) O(1C) ? ? ? N(11CIV) O(1C) ? ? ? O(58) O(1C) ? ? ? O(47) 2.812(18) 2.814(19) 2.823(20) 2.800(18) 2.861(16) 2.800(16) 3.018(19) 2.799(18) 2.827(18) 2.860(21) 3.035(19) 3.142(18) O(1A) ? ? ? O(16I) N(11A) ? ? ? N(21A) N(21A) ? ? ? N(31A) O(1B) ? ? ? O(17I) O(1B) ? ? ? O(67) N(11B) ? ? ? N(21B) N(21B) ? ? ? N(31B) O(1C) ? ? ? O(36II) O(1C) ? ? ? O(41) N(11C) ? ? ? N(21C) N(21C) ? ? ? N(31C) 3.005(19) 2.691(17) 2.691(19) 2.965(22) 3.213(20) 2.702(18) 2.712(18) 3.210(18) 3.255(19) 2.660(21) 2.651(20) Compound 2 N(11A) ? ? ? O(303) N(11A) ? ? ? O(302) N(21A) ? ? ? O(303) 2.750(14) 3.095(16) 3.101(14) N(31A) ? ? ? O(303) N(21A) ? ? ? N(11A) N(21A) ? ? ? N(31A) 2.767(14) 2.677(13) 2.648(13) Symmetry elements: I 1 1 x, y, z; II 1 2 x, 0.5 1 y, 2z; III x, y, 1 1 z; IV x 2 1, y, z 2 1.While there are literally hundreds of examples of L2 complexed to metals recorded in the Cambridge Crystallographic Database (CCDS), there is only one previously reported example of the [H2L2]21 cation, namely in the salt 2[H2L2]21- [Tb(OH2)8]?7Cl2.11 The nitrogen atoms in the outer pyridine rings are protonated and a chloride anion was located in the diprotonated cavity.This structure containing [H2L2]21 together with those established in the four structures presented in our previous work2 for [HnL1]n1 suggest that the diprotonated terpyridyl species always attracts hydrogen bonding species into the tridentate cavity. In 1 in addition to the above three [H2L2]21?H2O moieties the asymmetric unit contains two [Ln(NO3)6]32 anions. The Ln]O distances in this structure vary between 2.609(11) and 2.765(10) Å. Although there are numerous structures in the CCDS with Ln]O (nitrate) bonds, there are only three previously reported structures with [Ln(NO3)6]32 anions.For these anions and the two in the present work, the average Ln]O distance is 2.67 Å [n = 60, s(n 2 1) = 0.044 Å]. It has been suggested previously that an analysis of the metal co-ordination sphere for this type Fig. 1 The structure of compound 1 showing the environment of one [H2L2]21 cation with a water encapsulated within the cavity and forming N]H? ? ? O hydrogen bonds.In addition, the water molecule forms hydrogen bonds with oxygen atoms from the nitrates in the [La(NO3)6]32 anions. All hydrogen bonds are shown as dotted lines. The environments of the other two [H2L2]21 cations in the asymmetric unit are similar. Lanthanum yellow, oxygen red, carbon green, hydrogen yellow, nitrogen purple. Fig. 2 The structure of compound 1 showing the environment of one [La(NO3)6]32 anion forming hydrogen bonds to the three water molecules which are each encapsulated within [H2L2]21 cations.All hydrogen bonds are shown as dotted lines. The environment of the second [La(NO3)6]32 anion in the unit cell is similar. Lanthanum yellow, oxygen red, carbon green, hydrogen yellow, nitrogen purple.2976 J. Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 of structure can be simplified by considering a bidentate ligand of small ‘bite’ such as nitrate to occupy only one site of a coordination polyhedron rather than two.The new metal coordination sphere would have a lower co-ordination number which is easier to analyse.12 Thus, for [Ln(NO3)6]32 the arrangement of nitrogen atoms around Ln should be close to an octahedral geometry. This is indeed the case for the environment around Ln(2) in which the N]Ln]N angles only vary between 82(1) and 97(1)8 and indeed there are six O]Ln(2)]O angles within 88 of 1808. There are larger deviations, however, around Ln(1), the N]Ln]N angles varying between 71(1) and 117(1)8 and while there are three O]Ln]O angles within 58 of 1808, the other potentially trans angles are ca. 1608. It is not clear why there should be such big geometrical diVerences between the two anions though it seems likely that the geometry could well be aVected by the significant numbers of hydrogen bonds found in the unit cell. Non-bonded nitrate oxygen atoms and bonded nitrate oxygen atoms in both anions are involved to approximately the same extent in hydrogen bonding with the water molecules in the terpyridyl cavity.This has the eVect of weakening the corresponding Ln]O bonds, for example, there are 10 bonds longer than 2.7 Å in the anions; however it is not clear why the geometry of one anion should be so diVerent from that of the other. The hydrogen bond dimensions are shown in Table 3. The distances between each water atom and the two protonated nitrogen atoms with which hydrogen bonds are formed are all between 2.80 and 2.87 Å.In addition, each water molecule forms several hydrogen bonds to the nitrate oxygen atoms, one strong at a distance of ca. 2.80 Å and several weaker (at distances of 3.0–3.3 Å). In Fig. 1 we show the arrangement of one water molecule which forms hydrogen bonds to two diVerent anions. The cations are approximately planar with N]C]C]N torsion angles of 20.1(2), 3.4(2)8 in ligand A, 1.7(1), 20.9(1)8 in ligand B and 20.5(2), 26.4(2)8 in ligand C, respectively. It is diYcult to establish the exact hydrogen bond pattern around each oxygen as there are so many short contacts [four for O(1A), six for O(1B), seven for O(1C)] less than 3.26 Å.It seems likely that the positions of the hydrogen atoms are fluxional and that weak interactions can occur to any of the close nitrate oxygen atoms. In Fig. 2 the environment around one anion which forms hydrogen bonds to all three [H2L2]21?H2O moieties is shown. It is interesting that one of the nitrate oxygen atoms forms hydrogen bonds to two diVerent water molecules.The distances between the protonated N(11) and N(31) atoms and the central unprotonated N(21) atom vary between 2.65(2) and 2.71(2) Å while the corresponding N]H? ? ? N angles are all between 106 and 1088. These angles are similar to those found in the structure of 2,29:69,20-terpyridinium trifluoromethanesulfonate (Fig. 7) which contains a monoprotonated terpyridyl cation with an N]N distance of 2.65 Å and an N]H? ? ? N angle of 1038, dimensions which are considered to be indicative of weak intramolecular hydrogen bonds.13 The corresponding N]N distances in the salt 2[H2L2]21[Tb(OH2)8]31?7Cl2 are also very similar, varying between 2.61 and 2.67 Å.11 Structure of [H2L2]21[SmL2(NO3)4]2[NO3]2 2 The structure of 2 is shown in Fig. 3.The asymmetric unit contains a [Sm(L2)(NO3)4]2 anion and an [H2L2]21 cation with an encapsulated NO3 2 anion. The stoichiometry and indeed the structure of the [Sm(L2)(NO3)4]2 anion is similar to that recently observed in the structure of [La(L2)2(NO3)2]1[La(L2)- (NO3)4].14 In both structures the metal atoms in the anion are 11-co-ordinate.There are nine other structures in the Cambridge Database containing Sm]O (nitrate) bonds. However, the Sm(III) ion is co-ordinated only to oxygen atoms in the surveyed structures but to both oxygen and nitrogen atoms in our structure. Six of these previously reported structures contain a 10-co-ordinate Sm(III) ion.Only one of the reported structures contains Sm(III) in both the anion and the cation and in both ions the metal is 11-co-ordinate.15 The Sm]O bond lengths in our structure vary between 2.494(5) and 2.742(5) Å. The average Sm]O (nitrate) distance from all the surveyed structures is 2.56 Å [n = 56, s(n 2 1) = 0.093 Å]. It is diYcult to explain why there is such a broad range in Sm]O distances in our structure. It is interesting to note that the oxygen atom involved in the longest Sm]O bond is trans to the shortest Sm]N bond which involves the nitrogen atom in the central pyridine ring.The three Sm]N bond lengths are all longer than those observed in the two previously published structures containing L2 with Sm.4 In the first of these structures the Sm(III) is either eight- or nine-co-ordinate depending on the cation. Thus, the formula can be described as Sm(L2)(Cl)(H2O)4,5. In the second structure the Sm(III) is eight-co-ordinate, the metal being co-ordinated to three L2 N atoms, four Cl atoms and one water molecule.The Sm]N bond lengths vary between 2.56 and 2.59 Å in both of these structures, while in 2 the Sm(III) ion is 11-co-ordinate and the bond lengths vary between 2.62 and 2.65 Å. The increased Sm]N distances in 2 are clearly a consequence of the higher co-ordination number compared to the other structures. The terpyridine co-ordinated to the Sm(III) ion is almost planar, the N]C]C]N dihedral angles are 5.6(1) and 3.7(1)8.The terpyridyl cation shows more distortion from planarity with corresponding dihedral angles of 23.6(1) and 11.0(1)8. The nitrate ion in the terpyridyl cavity is hydrogen bonded to the N]H protons through one oxygen only. As with the diprotonated terpyridine in 1, it is also possible to describe the N]H? ? ? N interactions as weak intramolecular hydrogen bonds. Theoretical structural analysis of L2, [HL2]1 and [H2L2]21 The terpyridyl ligand L2. There are three possible conformations for L2 which can be characterised by the N]C]C]N torsion angles as tt (trans,trans), ct (cis,trans) and cc (cis,cis). Fig. 3 The structure of compound 2 showing both the [H2L2]21 cation hydrogen bonded to a nitrate anion together with the [Sm(NO3)4L2] anion. Samarium brown, oxygen red, carbon green, hydrogen yellow, nitrogen purple.J. Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 2977 Table 4 Results from quantum mechanics calculations on terpy, [Hterpy]1 and [H2terpy]21.Geometry optimisation was carried out using the 6-31G** basis set. Energies in au (= 627.509 kcal mol21) Protonated Conformation Compound L2 [HL2]1 [H2L2]21 nitrogen N(11) N(21) N(31) N(11), N(21) N(11), N(31) N(21), N(31) cis,cis 2737.798 (cc) 2738.202 (H-cc1) 2738.213 (H-cc2) a 2738.457 (2H-cc1) 2738.486 (2H-cc2) c cis,trans 2737.807 (ct) 2738.203 (H-ct1) 2738.206 (H-ct2) 2738.189 (H-ct3) 2738.443 (2H-ct1) b 2738.484 (2H-ct2) 2738.443 (2H-ct3) b trans,trans 2737.818 (tt) 2738.195 (H-tt1) 2738.195 (H-tt2) a 2738.443 (2H-tt1) 2738.479 (2H-tt2) c a Structures with N(31) protonated are equivalent to structures with N(11) protonated.b Structures unstable to geometry optimisation. The 2H-tt1 structure was obtained. c Structures with N(21) and N(31) protonated are equivalent to structures with N(21) and N(11) protonated. We have analysed these conformations for L2, [HL2]1 and [H2L2]21 using the GAUSSIAN 94 program.16 Starting models were built using the CERIUS2 software 17 and the three rings were made approximately coplanar but no symmetry was imposed.Structures were then optimised using the 6-31G* basis set.18 Results are summarised in Table 4. For the neutral ligand L2, it was found that the order of energies was tt < ct < cc and this is consistent with the fact that the tt conformation is observed in the crystal structure of L2 19 and also the 49-phenyl,20 4-aniline 21 and 49-NMe2 22 derivatives. A structure with a boron cage in the 49 position viz 49-(closo-ocarboranyl) terpyridine also has this tt conformation.23 This trans conformation of pyridine rings is also found in quaterpyridine 24 and a para substituted sexipyridine.25 The reasons for this order of conformational preference is clear from the geometry of the optimised structures (Fig. 4). In the tt form, the N]C]C]N torsion angles are both 180.08. However in the ct form, while the trans torsion angle at 2176.18 shows that the rings are close to being coplanar, the cis torsion angle at 243.08 shows that the rings are very much twisted away from planarity. This twist is caused by the repulsion between the two ortho hydrogen atoms and possibly also from electron– electron repulsion between the lone pairs on the nitrogen atoms.This pattern of conformational change is also observed in the cc form where the torsion angles are 247.9, 47.98. Clearly, the cis arrangement reduces conjugation between the two rings and is particularly destabilised by repulsions between the ortho hydrogen atoms.These calculations were carried out on the free ligands but it is possible that the conformational preferences may well change in the presence of hydrogen bond donors or acceptors. It is interesting that the tt form is the lowest energy conformation of neutral terpyridyl which is unsuitable for tridentate complexation with a metal atom. However it is possible that this conformation could be stabilised in polar solvents by the formation of intermolecular hydrogen bonds.Thus a water molecule could enter the cavity with the same arrangement as found for the [H2L2]21?H2O cation found in 1 but with donor O]H? ? ? N instead of acceptor O ? ? ?H]N hydrogen bonds [see Fig. 5(a)]. A precedence for this proposed structure is found for N(21) N(31) N(11) N(21) N(31) N(11) N(21) N(31) N(11) cis, cis cis, trans trans, trans a crystal structure of terpyridyl co-crystallised with [SnPh3- (NCS)(H2O)].26 This is the only crystal structure where an oligopyridine contains adjacent pyridine ligands in the cis conformation, but the conformation is stabilised by the formation of two hydrogen bonds to the water molecule which is situated in the terpyridine cavity as well as being bonded to the tin. In order to calculate the eVect of the formation of this L2?H2O complex, a model structure was built using CERIUS2 [Fig. 5(a)] and subjected to optimisation in GAUSSIAN 94.Results show that the energy of the complex was 11.29 kcal mol21, lower than that of the L2 and H2O separated at infinity. Fig. 4 The three conformations of L2, cis,cis (cc), cis,trans (ct) and trans,trans (tt). Energies after geometry optimisation (au) and torsion angles (8) were for cc 2737.798, 247.9, 47.9; for ct 2737.807, 243.0, 2176.1 and for tt 2737.818, 180, 180.2978 J. Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 Water adducts built with the ct and tt structures proved not to be stable to geometry optimisation, no doubt because of the repulsions from adjacent C]H groups.In the cc L?H2O complex, the structure maintains Cs symmetry, though this was not imposed. The N]C]C]N torsion angles at 42.0, 242.08 are only slightly reduced (by 5.98) from values in the unhydrated form. Fig. 5 The stabilisation of the cis,cis conformation by intramolecular hydrogen bonds with water molecules. Geometry optimisation via GAUSSIAN 94.(a) The cc form accepting hydrogen bonds from a water molecule; N]C]C]N torsion angles are 242.0, 41.98. (b) The 2Hcc2 form donating hydrogen bonds to a water molecule; N]C]C]N torsion angles are 218.7, 18.78. Hydrogen bonds shown as dotted lines. Carbon green, hydrogen yellow, nitrogen purple, oxygen red. The N ? ? ? H and O ? ? ? H distances are 2.2, 3.15 Å respectively with an O]H? ? ? N angle of 177.38. This contrasts with torsion angles of 29.0, 218.08 and N ? ? ? O distances of 2.80, 2.76 Å in ref. 26. Monoprotonated [HL2]1. When L2 is monoprotonated, then there are seven possible structures (Fig. 6) depending on the conformation and which nitrogen atom is protonated. There are two cis,cis structures; H-cc1, H-cc2 depending on whether N(11) or N(21) is protonated [N(11) is equivalent to N(31)]; three cis,trans structures H-ct1, H-ct2, H-ct3 dependent upon whether N(11), N(21) or N(31) is protonated; and two trans, trans structures H-tt1, H-tt2 dependent upon whether N(11) or N(21) is protonated.These seven structures were built with CERIUS2 and geometry optimised with GAUSSIAN 94 and the resulting structures are shown in Fig. 6. As with terpyridyl L2, the energies of [HL2]1 can be correlated with two dominant structural features, one favourable and one unfavourable. The lowest energy structures contain at least one pair of mutually cis nitrogen atoms of which one is protonated so that an energetically favourable intramolecular hydrogen bond interaction can be formed.Thus H-cc2 has the lowest energy with two such interactions, next lowest are H-cc1, H-ct1 and H-ct2 with one such interaction and the three highest energies are H-ct3, H-tt1 and H-tt2 with no such interactions. The unfavourable interaction occurs when the N]H is cis to one or two C]H bonds leading to H ? ? ? H repulsion. Two of these interactions are found in H-tt2, and one in H-tt1, H-ct3 and H-ct2.It is interesting that this latter H-ct2 structure also has a favourable hydrogen bond interaction which more than compensates for this repulsion as the energy is one of the lowest found. Unfavourable ortho C]H interactions are also found in several of the structures. The geometry in the structures (Fig. 6) follow a regular logical pattern. For the cis N]C]C]N torsions, the angle is close to zero when a hydrogen bond is formed (e.g. 5.1 in H-cc1, 20.1 in H-cc2, 6.8 in H-ct1, 20.38 in H-ct2) but otherwise is twisted significantly to relieve steric strain (e.g. 228.6 in H-cc1, 237.28 in H-ct3). For the trans N]C]C]N torsions, the angle is only close to 1808 in H-tt1 where there are no ortho repulsions but in all the other structures the absolute value of the angle ranges from 148.8 to 155.88. There are other geometric changes concomitant with the formation of the intramolecular N]H? ? ? N hydrogen bonds. Thus in H-ct2 the Fig. 6 The seven structures of [HL2]1 which are characterised by the conformation and the nitrogen that is protonated.Energies after geometry optimisation (au) and torsion angles (8) were for H-cc1 2738.202, 228.6, 5.1; for H-cc2 2738.215, 20.1, 20.1; for H-ct1 2738.203, 6.8, 2163.8; for H-ct2 2738.206, 20.3, 155.8; for H-ct3 2738.189, 237.2, 2154.8; for H-tt1 2738.195, 2174.0, 2148.8 and for H-tt2 2738.195, 2150.0, 2150.0. Hydrogen bonds shown as dotted lines. Carbon green, hydrogen yellow, nitrogen purple.J.Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 2979 two C]C]N angles to the central pyridine atom are 115.18 to ring 1 to facilitate the hydrogen bond and 119.18 to ring 3 where there is no such hydrogen bond. The H ? ? ? N distances are 2.07 Å in H-cc1, but 2.16 Å in H-ct2 and 2.12 Å in H-ct1, 2.10 Å in H-ct2. The N]H? ? ? N angles range from 105 to 1088, rather small compared to the usual angle for a hydrogen bond, but the calculations presented here clearly indicate that the hydrogen bonds are significant. There is a significant change in the N]C bond lengths when the nitrogen is protonated and increases of ca. 0.01 Å are observed. Thus in H-cc2, the C]N distances are 1.338 Å compared to 1.326 and 1.315 Å where the N is unprotonated. There is very little experimental evidence on [HL2]1 that can be correlated with these calculations. There is only one crystal structure containing [HL2]1, viz the salt with [CF3SO3]213 and this has the H-ct1 structure.However the structure of diproton- Fig. 7 The structure of [HL2]1[CF3SO3]2 showing the intermolecular hydrogen bond between N]H and an oxygen atom in the anion.13 Sulfur brown, oxygen red, carbon green, fluorine light green, nitrogen purple, hydrogen yellow. ated quinquepyridinium which has the cis,trans,trans,cis conformations shows the outer two nitrogens to be protonated so that the structure contains two moieties equivalent to the H-ct1 structure.27 As is apparent from the energy values in Table 4, this H-ct1 structure is not the lowest energy structure, but unlike the more favourable H-cc2 and H-ct2 structures, it retains the possibility of being able to form intermolecular hydrogen bonds.Indeed in the two crystal structures, this is precisely what is found, that the protonated nitrogen atom forms an intramolecular hydrogen bond but also an intermolecular hydrogen bond to the anion in the crystal. The anion is situated well away from the cavity (Fig. 7) and calculations show that this anion (or indeed any other) could not form stable hydrogen bonds to the protonated central nitrogen atom as occurs in H-cc2 and H-ct2. We conclude, therefore, that in the presence of solvents and/or anions which can accept hydrogen bonds that the H-ct1 structure is more likely to be found in preference to H-cc2 and H-ct2 despite these having lower energies in the gas phase. Diprotonated [H2L2]21. For [H2L2]21 there are also seven different structures.There are two with the cis,cis conformation called 2H-cc1 [with N(11) and N(21) protonated], and 2H-cc2 [with N(11) and N(31) protonated]; three with the cis,trans conformation; 2H-ct1 [N(11), N(21) protonated], 2H-ct2 [N(11), N(31) protonated] and 2H-ct3 [N(21), N(31) protonated] and two with the trans,trans conformation; 2H-tt1 [N(11), N(21) protonated] and 2H-tt2 [N(11), N(31) protonated]. Models were built with CERIUS2 and then geometry optimised with GAUSSIAN 94 and the results are given in Table 4 and illustrated in Fig. 8. The same two structural features are crucial to determining the relative energy values, thus lower energy structures contain intramolecular hydrogen bonds and higher energy structures N]H? ? ?C]H repulsions. There are more of these latter types of repulsion in the diprotonated structures than in the monoprotonated structures and there is also the possibility of N]H? ? ?N]H repulsions. The C]H? ? ? C]H repulsions are present in all structures but seem to be less important, possibly because the electrostatic repulsions between hydrogen atoms is less.Because of these ortho–ortho repulsions only three structures contain intramolecular hydrogen bonds 2H-cc2, 2H-ct2 and 2H-cc1 and these together with 2H-tt2 have the lowest energies. The lowest energy conformation (2H-cc2) is also observed in the two crystal structures reported above. In this structure the Fig. 8 The seven structures of [H2L2]21 which are characterised by the conformation and the nitrogen that is protonated.Energies after geometry optimisation (au) and torsion angles (8) were for 2H-cc1 2738.456, 0.6, 58.7; for 2H-cc2 2738.486, 228.3, 228.3; for 2H-ct1 unstable; 2H-ct2 2738.484, 12.8, 2139.1; for 2H-ct3 unstable; for 2H-tt1 2738.444, 2130.5, 2155.1 and for H-tt2 2738.479, 2145.0, 2145.0. Hydrogen bonds shown as dotted lines. Carbon green, hydrogen yellow, nitrogen purple.2980 J.Chem. Soc., Dalton Trans., 1998, Pages 2973–2980 N]C]C]N torsion angles are 228.38 and the N]H? ? ? N distance is 2.43 Å. This compares with torsion angles of 0.4, 6.0; 23.3, 0.5; 21.9, 0.18 in 1 and 10.8, 23.78 in 2 and distances of 2.28–2.34 Å in the two structures (though it must be borne in mind that these are calculated hydrogen positions from the crystal structure determinations) and indeed to the 2.16 Å observed in the 1H-cc1 structure. The distance in 2H-ct2 is also large by comparison at 2.37 Å.It seems likely that the ligand in the crystal structures is more planar because of the intermolecular hydrogen bonds formed to the nitrate. In order to test this proposition, we carried out a GAUSSIAN 94 geometry optimisation of a water molecule contained within the cavity of the 2H-cc2 structure [Fig. 5(b)]. On optimisation the two torsion angles had decreased to 18.7 and 218.78 and the two H ? ? ?N distances reduced to 2.13, 2.13 Å. Results show that the energy of the complex was 18.72 kcal mol21, lower than that of [H2L2]21 and H2O separated at infinity.We conclude that in the case of the diprotonated [H2L2]21 the gas phase preference for the 2H-cc2 structure is enhanced in solution and that this is likely to be the only diprotonated species present. The conformation found in the 2H-cc1 structure is particularly interesting as it contains a N]H? ? ?N]H repulsion which is alleviated with a torsion angle of 58.78 while the second cis interaction contains a N]H? ? ? N attraction and the torsion angle remains close to zero at 0.68.Of the other structures 2Hct1 and 2H-ct3 proved not to be stable to geometry optimisation and reverted to the 2H-tt1 structure. The 2H-tt2 structure contains no hydrogen bonds but has a relatively low energy because there are no significant ortho repulsions. The 2H-tt1 structure has the highest energy of all because of significant repulsions between N]H and C]H in the cavity.Conclusion Theoretical calculations have shown that the lowest energy conformation of the terpyridyl ligand is the trans,trans form. However cis forms are stabilised on protonation by the formation of weak intramolecular N]H? ? ? N hydrogen bonds. It is likely that the lowest energy forms in solution for the protonated and diprotonated forms are the H-ct1 and 2H-cc2. These results are consistent with evidence from crystal structure determinations presented here and from the literature.It is suggested that any theory of extraction based on the ion-pair mechanism must be consistent with the presence of these cations both of which will be stabilised by the formation of intermolecular hydrogen bonds to solvent and/or accompanying anions. Acknowledgements We are grateful for the financial support by the EC Research Programme Contract NEWPART (FI41-CT-96-0010) and the Swedish Nuclear and Waste Management Co., SKB. We would also like to thank the EPSRC and the University of Reading for funding of the image-plate system.The use of the Origin 2000 at the High Performance Computer Centre at the University of Reading is gratefully acknowledged. References 1 K. L. Nash, Solvent Extraction and Ion Exchange, 1993, 11, 729. 2 G. Y. S. Chan, M. G. B. Drew, M. J. Hudson, N. S. Isaacs, P. Byers and C. Madic, Polyhedron, 1996, 15, 3385. 3 E. C. Constable, Adv. Inorg. Chem., 1986, 30, 69. 4 C. J. Kepert, L. Weimin, B.W. Skelton and A. H. White, Aust. J. Chem., 1994, 47, 365. 5 I. Hagström, L. Spjuth, Å. Enarsson, J. O. Liljenzin, M. Skålberg, M. J. Hudson, P. B. Iveson, P. Y. Cordier, C. Hill and C. Madic, Solvent Extraction and Ion Exchange, in the press. 6 A referee has commented that adding concentrated nitric acid to a methanol solution is highly dangerous. For this reason, as stated, we used extremely small quantities. 7 W. Kabsch, J. Appl. Crystallogr., 1988, 21, 916. 8 G. M. Sheldrick, SHELXS, program for structure determination, University of Göttingen, 1997. 9 N. Walker and D. Stuart, DIFABS, Acta Crystallogr., Sect. A, 1983, 39, 158. 10 G. M. Sheldrick, SHELXL, program for crystal structure refinement, University of Göttingen, 1997. 11 C. J. Kepert, B. W. Skelton and A. H. White, Aust. J. Chem., 1994, 47, 391. 12 J. G. Bergman and F. A. Cotton, Inorg. Chem., 1966, 5, 1208. 13 A. Hergold-Brundic, Z. Popovic and D. Matkovic-Calogovic, Acta Crystallogr., Sect. C, 1996, 52, 3154. 14 M. Frechette and C. Bensimon, Inorg. Chem., 1995, 34, 3520. 15 J. H. Burns, Inorg. Chem., 1979, 18, 3044. 16 M. J. Frisch, G. W. Trucks, H. B. Schlegel, P. M. W. Gill, B. G. Johnson, M. A. Robb, J. R. Cheeseman, T. A. Keith, G. A. Petersson, J. A. Montgomery, K. Raghavachari, M. A. Al-Laham, V. G. Zakrzewski, J. V. Ortiz, J. B. Foresman, J. Cioslowski, B. B. Stefanov, A. Nanayakkara, M. Challalcombe, C. Y. Peng, P. Y. Ayala, W. Chen, M. W. Wong, J. L. Andrews, E. S. Replogle, R. Gomperts, R. L. Martin, D. L. Fox, J. S. Binkley, D. J. Defrees, J. Baker, J. P. Stewart, M. Head-Gordon, C. Gonzalez and J. A. Pople, GAUSSIAN 94, Revision A1, Gaussian, Inc., Pittsburgh, PA, 1995. 17 CERIUS2 software, Molecular Simulations Inc., San Diego, CA, 1997. 18 For all the structures reported here, geometry optimisation was carried out with no constraints and no imposed symmetry. It is possible particularly in cases where intramolecular ‘bonds’ are present that basis-set superposition errors may not be negligible but a detailed investigation is beyond the scope of this paper. 19 C. A. Bessel, R. F. See, D. L. Jameson, M. R. Churchill and K. J. Takeuchi, J. Chem. Soc., Dalton Trans., 1992, 3223. 20 E. C. Constable, J. Lewis, M. C. Liptrot and P. R. Raithby, Inorg. Chim. Acta, 1990, 178, 47. 21 G. D. Storrier, S. B. Colbran and D. C. Craig, J. Chem. Soc., Dalton Trans., 1997, 3011. 22 E. C. Constable, A. M. W. C. Thompson, D. A. Tocher and M. A. M. Daniels, New. J. Chem., 1992, 16, 855. 23 D. Armspach, E. C. Constable, C. D. Housecroft, M. Neuburger and M. Zehnder, New. J. Chem., 1996, 20, 331. 24 E. C. Constable, S. M. Elder, J. Healy and D. A. Tocher, J. Chem. Soc., Dalton Trans., 1990, 1669. 25 K. T. Potts, K. A. G. Raiford and M. Keshavarz, J. Am. Chem. Soc., 1993, 115, 2793. 26 L. Prasad and F. E. Smith, Acta Crystallogr., Sect. B, 1982, 38, 1815. 27 E. C. Constable, S. M. Elder, J. V. Walker, P. D. Wood and D. A. Tocher, J. Chem. Soc., Chem. Commun., 1992, 229. Received 31st March 1998; Paper 8/02458H

ISSN:1477-9226    

DOI:10.1039/a802458h

出版商:RSC    

年代:1998

数据来源: RSC

摘要:

DALTON FULL PAPER J. Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 2981 Electronic eVects in the nickel-catalysed hydrocyanation of styrene applying chelating phosphorus ligands with large bite angles ‡ Wolfgang Goertz,a Wilhelm Keim,a Dieter Vogt,*,†,a Ulli Englert,b Maarten D. K. Boele,c Lars A. van der Veen,c Paul C. J. Kamerc and Piet W. N. M. van Leeuwen c a Institut für Technische Chemie und Petrolchemie, RWTH Aachen, Templergraben 55, 52056 Aachen, Germany b Institut für Anorganische Chemie, RWTH Aachen, Templergraben 55, 52056 Aachen, Germany c J.H. van’t Hoff Research Institute, Department of Inorganic Chemistry, University of Amsterdam, Nieuwe Achtergracht 166, 1018 WV Amsterdam, The Netherlands Chelating phosphorus ligands with a rigid backbone and a large natural bite angle were applied in the nickelcatalysed hydrocyanation of styrene. The para substituents in the diphenylphosphanyl moiety of the 4,6-bis- (diphenylphosphanyl)-2,8-dimethylphenoxathiine (Thixantphos) ligands were varied and their electronic eVects on the activity and selectivity of the catalytic experiments were investigated.The activity of the nickel complexes decreased when electron-donating substituents lead to a more basic phosphorus while electron-withdrawing substituents led to a higher activity. The results of variable temperature 31P-{1H} NMR experiments on the in situ catalysts are discussed in relationship to the catalytic performance. 4,6-Bis(diphenylphosphanyl)-2,8- dimethylphenoxathiine (Thixantphos) L1d and the complexes [NiCl2L1d] 1 and [Ni(CN)2L1a] 2 (p-Me2N on phenyl) have been characterised by single-crystal X-ray diVraction. Complex 2 represents the first crystal structure of a monomeric dicyanonickel(II) complex with a P]P chelating ligand. The geometries of ligand L1d and complex 1 were predicted by molecular mechanics calculations. The hydrocyanation of butadiene in the presence of nickel complexes, known as the DuPont ADN Process, is one of the most prominent examples of homogeneous catalysis on an industrial scale.1 Many investigations have been undertaken to understand its mechanism and to improve the activity and selectivity.2 For some decades the process has been carried out with monodentate phosphites which are bound only very weakly to the zerovalent metal centre.3 For this reason, one major drawback in this process arises from the formation of inactive nickel cyanides when the monodentate ligands dissociate and the catalyst metal is exposed to an excess of HCN.Only recently bidentate ligands have drawn some attention.4 They are able to stabilise the active species to a much higher extent due to the chelate eVect. It was reported by Pringle and co-workers 5 that diphosphites form complexes with nickel and palladium, that are active in hydrocyanation catalysis and highly resistant towards oxidation as well. Striking results have been obtained in the asymmetric hydrocyanation of vinylarenes by using chiral diphosphinites derived from sugar backbones.6 Enantiomeric excesses up to 95% combined with high activities were achieved.7 It was suggested that the enantioselectivity of the reaction is determined in the reductive elimination step of the catalytic cycle which is strongly influenced by the electronic properties of the ligand.In preliminary communications, we have reported on the successful application of diphosphines in the hydrocyanation of styrene 8 and long-chain, non-activated olefins.9 These newly developed Xantphos ligands 10 have large natural bite angles and rigid backbones.The compounds are based on heterocyclic xanthene-like aromatics. They were designed to stabilise geometries with P]Ni]P angles larger than 1008, shown to have † E-Mail: dieter.vogt@post.rwth-aachen.de ‡ Supplementary data available: computational and other details. For direct electronic access see http://www.rsc.org/suppdata/dt/1998/2981/, otherwise available from BLDSC (No.SUP 57410, 5 pp.) or the RSC Library. See Instructions for Authors, 1998, Issue 1 (http://www.rsc.org/ dalton). a strong influence on catalyst selectivity in manifold reactions. 10,11 By modifying the ligand backbone the bite angle can be precisely adjusted to various geometries. For the catalytic hydrocyanation we proposed that ligands with fixed large bite angles would (a) disfavour inactive square planar nickel(II) dicyano species, and (b) stabilise active tetrahedral nickel(0) complexes.8 Next to the geometry of the ligands, their electronic properties play a crucial role for the complex stability and have a great influence on elementary steps of the catalytic cycle.12 Ligands with electron-withdrawing substituents at the phosphorus atoms bear a low basicity and can easily compensate the high electron density of the d10 configurated zerovalent nickel by back donation into non-occupied orbitals.A nickel which is electron-poor due to ligand properties will readily co-ordinate an olefinic substrate. Moreover, kinetic studies by RajanBabu6 and co-workers indicate that the rate of reductive elimination is increased by electron-withdrawing substituents at the ligand. For these reasons, phosphites with three electron-withdrawing oxygen substituents are much more favourable for hydrocyanation reactions than phosphines. However, we could demonstrate that Xantphos type diphosphines are quite comparably active to monophosphites.8,9 In this paper we describe the application of electronically tuned Thixantphos type ligands L1a–L1g and L2b (Scheme 1) in the hydrocyanation of styrene. Substituents were introduced in the para positions of the phenyl rings only in order to keep steric eVects as small as possible.In this way it should be possible to study the electronic eVects independently from the bite angle. Natural bite angles for nickel complexes and other ligand geometries were calculated by molecular modeling, using an augmented TRIPOS force field with SYBYL software.13 The natural bite angle is defined as the preferred chelation angle determined only by ligand-backbone constraints and not by metal valence angles.14 We have studied the crystal structures of one free Thixantphos L1d (R = H) and two nickel complexes [NiCl2L1d] 1 and [Ni(CN)2L1a] 2 to support the suggested relationship between the geometry and the reactivity of2982 J.Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 catalytically active intermediates. We have also carried out 31P- {1H} NMR investigations with the in situ catalysts which provide useful information on the performance and deactivation pathways for the various nickel complexes. Results and Discussion Mechanistic aspects The catalytic cycle for the hydrocyanation of olefins applying monodentate phosphorus ligands P as shown in Scheme 2 has been proposed by Tolman et al.15 Intermediates and catalytically active species have been isolated or observed and character- Scheme 1 Thixantphos ligands and their calculated natural bite angles O P P S R R R R O P P S O O O O L2b, bn = 120.4° Me2N MeO Me H F Cl CF3 L1a L1b L1c L1d L1e L1f L1g R 118.6 113.8 112.9 112.8 112.8 113.4 114.5 P P bn /° Scheme 2 Proposed catalytic cycle for hydrocyanation VII III VI 90/180° square-planar VII IV III, V II trigonal trigonal-bipyramidal Angle P - Ni - P Geometry Species 120° 120° 109° tetrahedral I, IV,VI I V II P Ni P P P R R NC CN Ni P P P H CN Ni H P P CN Ni P P H R CN Ni P P R P Ni P P CN Ni P NC P H2 HCN P P P HCN ised by various spectroscopic methods.16 Ligand dissociation from a tetrahedrally co-ordinated zerovalent nickel I generates a trigonal, highly reactive NiL3 species II, which undergoes rapid oxidative addition of HCN. The resulting complex III with a hydride and a cyano ligand co-ordinating in trans positions can easily lose a second monodentate ligand P to yield the tetrahedral species IV.Depending on the steric bulk of the ligands P the hydridocyano intermediate IV can also adopt a square planar structure. The cycle is continued by olefin coordination V and insertion of the substrate into the nickel– hydrogen bond giving another tetrahedral intermediate VI. Similar to IV it should be possible that the cyanoalkyl species VI can have a square planar geometry as well.In the last step, reductive elimination takes place to form the product nitrile and the cycle is completed by generating the active complex II again. Deactivation of the catalyst is often noticed by the formation of dicyanonickel(II) species VII, preferably with the cyano ligands occupying trans positions at the metal. All other intermediates of the catalytic cycle favour P]Ni]P angles of either 1098 together with tetrahedral geometry or 1208 with trigonal structures.When monodentate ligands are used the complexes are very flexible to structural changes, but also highly susceptible to dissociation and catalyst deactivation. Consequently, our conception was focussed on ligands with large natural bite angles and rigid backbones. These Xantphos bidentates should suppress the formation of dicyano complexes VII, stabilise the substantially tetrahedral intermediates IV and VI and thus enhance the reductive elimination and the overall catalysis.17 Since our aim was the detailed investigation of electronic properties, we selected the phenoxathiine backbone and the Thixantphos ligand L1d respectively for further modification.Synthesis and molecular modeling Next to a whole series of Thixantphos diphosphines L1a–L1g that were modified only at the para positions of the diphenylphosphanyl moiety, a new diphosphonite ligand L2b was synthesized (Scheme 1). This latter ligand is much less basic at the phosphorus atoms because of the biphenyl oxygen substituents. The diphosphonite ligand L2b is readily accessible by the reaction of the amido phosphonito phenoxathiine L2a with 2 equivalents of 2,29-dihydroxy-1,19-biphenyl (Scheme 3).The geometries of the ligands were simulated by molecular modeling based on force field calculations.13 Since the diphosphines L1a–L1g are only modified at the para positions of the phenyl substituents, the calculated natural bite angles for nickel complexes are within a very narrow range of 112.8 to 114.58, except for L1a.A steric influence on the hydrocyanation results can be excluded. Only ligand L1a with dimethylamino groups requires a somewhat larger angle of 118.68. For the diphosphonite ligand L2b we calculated a bite angle of 120.48. We suppose that this value is enforced by the limited flexible ring structures of the biphenoxy groups and their steric demand. Crystal structures The molecular structure of ligand L1d is drawn in Fig. 1, together with the atom numbering scheme, while selected bond lengths and angles are given in Table 1. The structure clearly shows that only very little adjustment of the ligand is necessary Scheme 3 (i) toluene, 90 8C, 15 h O P P Et2N NEt2 NEt2 Et2N S OH OH ( i ) - 4 HNEt2 L2b L2a + 2J. Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 2983 to form a chelate complex since the orientation of the diphenylphosphanyl moieties is nearly ideal for metal co-ordination. The non-bonding P ? ? ? P distance is 4.02 Å.While molecular modeling calculations suggest a bending in the phenoxathiine moiety, the two aromatic rings in the backbone of Thixantphos L1d are coplanar to each other. Probably, the planarity in the solid state arises from crystal structure packing eVects. Simulated structures are just idealised gas phase molecules with no interaction to other compounds. In contrast, the dihedral angle between the two phenyl planes in the backbone of the related Xantphos ligand (CMe2 instead of S, H instead of Me) is 1588, determined by X-ray analysis. 10,18 p-Stacking interactions between the phenyl substituents at the phosphorus atoms are thought to be responsible for this bending.19,20 Despite the numerous bidentate phosphorus ligands known in the literature, only a few dichloro nickel complexes have been characterised by X-ray analysis. Most common geometries to be found for [NiCl2(P]P)] are tetrahedral 21 and square planar.22 The molecular structure and numbering scheme for complex [NiCl2L1d] 1 are given in Fig. 2. Selected bond lengths, nonbonding distances and angles are presented in Table 2. The compound is paramagnetic which made recording of NMR spectra impossible. Complex 1 was simulated by molecular mechanics and predicted to have a distorted tetrahedral geometry around the metal. The calculated P]Ni]P angle is 110.68 which is in very good agreement with the X-ray determined angle of 109.52(6)8.The backbone bending also could be predicted with a C]S]C angle of 97.48 and a C]O]C angle of 114.68 [X-ray: 98.2(3) and 116.2(5)8 respectively]. The dihedral angle between the two aromatic rings in the phenoxathiine moiety is 147.7(3)8. The bond distances Ni]P [2.307(2) and 2.304(2) Å] appear quite normal compared with the values of 2.28 Å found for Fig. 1 Molecular structure of ligand 1d Table 1 Selected bond lengths (Å) and angles (8) for ligand L1d with estimated standard deviations in parentheses S]C(6) P(1)]C(2) P(1)]C(30) P(2)]C(40) O]C(1) C(4)]C(13) C(6)]S]C(7) C(2)]P(1)]C(20) C(11)]P(2)]C(40) C(1)]O]C(12) O]C(1)]C(6) O]C(12)]C(11) S]C(6)]C(5) P(2)]C(50)]C(55) 1.763(3) 1.846(3) 1.823(3) 1.824(3) 1.371(3) 1.510(4) 101.4(1) 100.5(1) 101.6(1) 124.3(2) 123.4(2) 114.4(2) 117.3(2) 126.5(2) S]C(7) P(1)]C(20) P(2)]C(11) P(2)]C(50) O]C(12) C(9)]C(14) P(2)]C(40)]C(45) C(20)]P(1)]C(30) P(2)]C(11)]C(10) O]C(1)]C(2) O]C(12)]C(7) S]C(6)]C(1) P(2)]C(50)]C(51) P(2)]C(40)]C(41) 1.760(3) 1.841(3) 1.836(3) 1.824(3) 1.375(3) 1.518(5) 117.3(3) 103.2(1) 125.1(2) 115.9(2) 124.4(2) 123.4(2) 115.6(2) 124.9(2) [NiCl2(PPh3)2] 23 and 2.333 Å (average) for [NiBr2(PPh3)2].24 Although the bite angle P]Ni]P of 109.52(6)8 is ideal for a tetrahedral geometry around the nickel, the structure is signifi- cantly distorted with a Cl]Ni]Cl angle of 132.37(7)8 due to the lone-pair repulsion between the two chlorine atoms.A similar geometry was found in the complexes [NiCl2(pop)] 25 [pop = 2,29-bis(diphenylphosphino)diethyl ether] and [NiCl2(diop)] 26 [diop = 4,5-bis(diphenylphosphinomethyl)-2,2-dimethyl-1,3-dioxolane] with Cl]Ni]Cl angles of 127.1(2) and 1308 respectively.A discussion of distortion of nickel(II) tetrahedra in terms of ligand field theory has been given by Venanzi.27 The complex [Ni(CN)2L1a] 2 represents the first known crystal structure of a monomeric dicyanonickel(II) complex with a P]P chelating ligand (Fig. 3 and Table 3). Even complexes of the type [Ni(CN)2L2] (L = monodentate phosphorus ligand) have been determined by X-ray analysis very rarely.28 Related dimeric species with bidentate ligands were described by Holah et al.,29 [Ni2(CN)2(dppm)2], and by Manojlovic-Muir et al.,30 [Ni2(CN)4(Me2PCH2PMe2)2], only the latter being characterised by X-ray diVraction. While the calculated natural bite angle for the ligand L1a is 118.68, the P]Ni]P angle found in the complex 2 is widened to 151.51(5)8.This value is out of the flexibility range31 for Xantphos type ligands, which is typically about 358 for an excess strain energy of 15 kJ mol21.10 The two cyano groups occupy trans positions at the nickel with a C]Ni]C angle of 161.0(2)8. Considering the small Ni]O distance of 2.72 Å and the four angles P]Ni]C in the range of 89.8(1) to 95.9(1)8, the geometry around the metal can be described as distorted square pyramidal with the oxygen on Fig. 2 Molecular structure of complex 1 Table 2 Selected bond lengths (Å) and angles (8) for complex 1 with estimated standard deviations in parentheses Ni]Cl(1) Ni]P(1) S]C(6) P(1)]C(2) P(1)]C(30) O]C(12) P(1) ? ? ? P(2) Cl(1)]Ni]Cl(2) Cl(1)]Ni]P(2) Cl(2)]Ni]P(2) C(6)]S]C(7) Ni]P(1)]C(20) C(2)]P(1)]C(20) C(20)]P(1)]C(30) 2.191(2) 2.307(2) 1.767(6) 1.829(6) 1.823(6) 1.397(6) 3.77 132.37(7) 102.88(7) 102.08(7) 98.2(3) 108.3(2) 103.0(3) 104.7(3) Ni]Cl(2) Ni]P(2) S]C(7) P(1)]C(20) O]C(1) C(4)]C(13) Ni ? ? ?O Cl(1)]Ni]P(1) Cl(2)]Ni]P(1) P(1)]Ni]P(2) Ni]P(1)]C(2) Ni]P(1)]C(30) C(2)]P(1)]C(30) C(1)]O]C(12) 2.206(2) 2.304(2) 1.758(6) 1.826(6) 1.386(6) 1.497(8) 3.37 105.82(7) 103.03(7) 109.52(6) 116.8(2) 116.8(2) 105.7(3) 116.2(5)2984 J.Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 top. However, an interaction between the nickel and the oxygen cannot be predicted from the crystal structure only. For nickel complexes with ether ligands found in the Cambridge Crystallographic Database the average Ni]O distance is 2.15 Å.The metal–phosphorus and –carbon bond lengths are in good agreement with those observed in other dicyano d8-nickel complexes. 32 Both of the Ni]C]N linkages are slightly distorted from ideal linearity with angles of 173.9(4) and 172.1(4)8. The backbone of the ligand is bent by the dihedral angle of 33.6(3)8. Ligand L1a is the most basic out of the series of diphosphines and the only one that formed a dicyano complex with nickel.We assign that to the electron-donating capability towards the Ni(CN)2 system. Catalytic hydrocyanation of styrene Reactions of the Thixantphos ligands L1a–L1g and L2b with 1 equivalent [Ni(cod)2] yield active catalyst precursors for the hydrocyanation of styrene (Table 4). The compounds L1a–L1c are more basic at the phosphorus atoms than the unsubstituted ligand L1d, caused by a positive mesomeric eVect of the amino and methoxy groups and a positive inductive eVect of the methyl group.Less basic than Thixantphos L1d are the halogen derivatives L1e–L1g (with R = F, Cl or CF3). Nevertheless, the negative inductive eVect of Cl is reduced by a positive mesomeric eVect. The electron-rich diphosphines L1a–L1c (entries 1–3, Table 4) lead to lower activities in catalytic experiments than the unsubstituted Thixantphos L1d. Best results are obtained with the most electron-poor diphosphine L1g (entry 7, R = CF3). The Fig. 3 Molecular structure of complex 2 Table 3 Selected bond lengths (Å) and angles (8) for complex 2 with estimated standard deviations in parentheses Ni]P(1) Ni]C(1) S]C(15) P(1)]C(11) P(1)]C(40) P(2)]C(50) O]C(10) N(1)]C(1) N(3)]C(33) N(5)]C(53) P(1) ? ? ? P(2) P(1)]Ni]P(2) P(1)]Ni]C(1) P(2)]Ni]C(1) C(15)]S]C(16) Ni]C(1)]N(1) Ni]P(1)]C(11) Ni]P(1)]C(40) Ni]P(2)]C(50) 2.208(1) 1.863(5) 1.764(5) 1.840(4) 1.808(5) 1.806(4) 1.410(5) 1.154(6) 1.379(6) 1.363(6) 4.27 151.51(5) 90.5(1) 89.8(1) 98.5(2) 173.9(4) 105.2(1) 115.9(1) 117.1(1) Ni]P(2) Ni]C(2) S]C(16) P(1)]C(30) P(2)]C(20) P(2)]C(60) O]C(21) N(2)]C(2) N(3)]C(36) N(5)]C(56) Ni]O C(1)]Ni]C(2) P(1)]Ni]C(2) P(2)]Ni]C(2) C(10)]O]C(21) Ni]C(2)]N(2) Ni]P(1)]C(30) Ni]P(2)]C(20) Ni]P(2)]C(60) 2.197(1) 1.861(5) 1.760(5) 1.813(4) 1.829(4) 1.795(4) 1.385(5) 1.145(6) 1.475(7) 1.442(6) 2.72 161.0(2) 95.9(1) 93.0(1) 117.3(3) 172.1(4) 122.4(2) 105.9(2) 118.9(2) other two halogen derivatives L1e and L1f perform less successfully than expected (entries 5 and 6).A plausible explanation may be that ligand decomposition has occurred by reaction of the aryl halogenides with styrene mediated by nickel species. Similar Heck type reactivity is best known with palladium.33 We have observed the same deactivation mechanism with nickel complexes in the presence of chlorinated or brominated aromatic substrates. The fluoride in ligand L1g is not bound directly to an aromatic ring and therefore inert towards reaction with the metal of the catalyst.Moderate yield combined with a very high selectivity is obtained in the presence of the diphosphonite ligand L2b (entry 8). This unprecedented behaviour will be explained later. With styrene as a substrate, the branched nitrile is the strongly favoured product over the linear one. This regioselectivity is attributed to stabilisation of the branched alkyl intermediate by a h3-benzyl interaction of the nickel.2 Still, the selectivity induced by the Thixantphos ligands L1a–L1g and L2b (>99%) is significantly higher than those obtained with common diphosphines (up to 95%) 8 and the commercial o-tolyl phosphite system (91%),34 and comparable to that of diphosphonites based on sugar backbones.35 NMR characterisation of nickel complexes Various in situ catalyst solutions were investigated by means of 31P-{1H} NMR spectroscopy.Surprisingly, the addition of either 1 or 2 equivalents of Thixantphos diphosphine ligands L1a–L1g to toluene solutions of [Ni(cod)2] resulted in similar spectra.The complexity of the peak pattern observed can only be explained by multiple couplings between various phosphorus atoms. The species which are responsible for this are the bis(chelate) complexes [Ni(P]P)2] 3a–3g (Scheme 4). The 31P-{1H} NMR spectrum of complex 3d (P]P = Thixantphos L1d) was simulated with g-NMR software 36 (Fig. 4). The peak pattern could most accurately be simulated by an AA9XX9 system with four large coupling constants of 50 Hz and two smaller ones of 12.5 Hz (Table 5).Therefore the structure of complex 3d is supposed to be distorted tetrahedral. The complex peak pattern at 273 K is caused by a highly asymmetric geometry of the bis(chelate) complex 3d. We conclude that due to the non-planarity of the ligand backbones there are at least two diVerent co-ordination modes (Scheme 5), which make the phosphorus nuclei Pa and Pb magnetically inequivalent. With a fixed roof-like co-ordination of the first Scheme 4 Formation of bis-chelate nickel complexes P Ni P P P 3a–3g L1a–L1g [Ni(cod)2] Table 4 Catalytic hydrocyanation of styrene a Entry 12345678 Ligand L1a L1b L1c L1d L1e L1f L1g L2b Conversion b (%) 16 33 55 77 46 75 98 49 Yield b (%) 12 22 48 70 38 52 90 47 Selectivity b (%) 76 66 87 92 83 70 92 97 Regioselectivity b (%) >99 99.5 99.8 99.4 99.4 99.3 99.0 99.4 a Reaction conditions: toluene (2 cm3), nickel : ligand : styrene :HCN = 1 : 1.05 : 20 : 25; 60 8C, 16 h, preformation time = 30 min.b Conversion (based on the substrate); yield, selectivity and regioselectivity determined by temperature-controlled GC analysis. Regioselectivity is defined as the percentage of branched nitrile.J. Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 2985 equivalent of ligand the backbone of the second diphosphine can be folded either to the left or to the right side. These two species are enantiomeric complexes. However, since there are two diastereotopic phosphorus atoms, the 31P-{1H} NMR spectrum shows two diVerent chemical shifts for Pa and Pb at d 11.7 and 15.0 respectively.Subsequently, the bis(chelate) complex [NiL1d] 3d was investigated in detail by variable temperature 31P-{1H} NMR as shown in Fig. 5. On heating the sample from room temperature to 371 K, line broadening occurs caused by a fluxional process and the phosphorus atoms become magnetically equivalent, with a coalescence temperature of 353 K.For this exchange process a butterfly flipping of the ligand backbones as well as ligand dissociation can be proposed. At higher temperatures this process is fast in comparison to the n(Pa) 2 n(Pb) frequency separation. Especially since the bis(chelate) complex is as catalytically active as the monochelate complex [Ni(cod)L1d], ligand dissociation cannot be excluded as a reason for the fluxionality. From the recorded and simulated 31P-{1H} NMR spectra, the rate constants (k) for the fluxional process have been determined for complex 3d.In the Eyring plot in Fig. 6 the rate constants (k) for phosphorus exchange are given as a function of the NMR sample temperature for 3d. From the Eyring equation and the Eyring plot, the thermodynamic values for the activation of the dynamic process were calculated. The assumption used is that DH‡ and DS‡ are constant over the temperature range (293–371 K) employed. The enthalpy of activation DH‡ is 61.4 kJ mol21 and the entropy of activation Fig. 4 Recorded (left, 273 K, toluene-d8) and simulated (right) 31P- {1H} NMR spectra of [NiL1d 2] 3d Scheme 5 Possible coordination modes and exchange processes of a bis-chelate complex (phenyl groups are omitted for clarity) Pb¢ Pa¢ Pa¢ Pb¢ Pb Pa Ni Pa¢ Pb¢ Pb Pa Ni ligand exchange backbone flipping Table 5 Chemical shifts and J(PP) values of the NMR simulation for [NiL1d 2] 3d J/Hz Nucleus P1 P2 P3 P4 d 15.00 11.70 15.00 11.70 1 50 50 12.5 2 12.5 50 3 50 DS‡ is 216.7 K21 mol21.The latter value is relatively small, indicative of an intramolecular rearrangement process, thus a butterfly flipping of the ligand backbones. For a temperature of 293 K, the free energy value DG‡ = 66.3 kJ mol21 has been determined from the Gibbs equation DG‡ = DH‡ 2 TDS‡. When our investigations were extended to the Thixantphos diphosphonite L2b a similar behaviour was observed. The 31P- {1H} NMR spectrum ([2H8]toluene, 298 K) of an in situ catalyst solution containing equimolar amounts of [Ni(cod)2] and ligand L2b shows a singlet at d 196.6 5 min after mixing, resulting from the monochelate [Ni(cod)L2b].After 30 min, about 80% of the ligand has already formed the bis(chelate) complex [NiL2b 2] 4 showing two pseudo-triplet signals at d 187.5 and 177.6 (Fig. 7). Surprisingly, NMR simulation for complex 4 requires equal Fig. 5 Variable temperature 31P-{1H} NMR spectrum (toluene-d8) of [NiL1d 2] 3d, recorded (left) and simulated (right)2986 J.Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 2J(PaPa9), 2J(PaPb), 2J(PaPb9), 2J(PbPb9) couplings of 27 Hz suggesting a structure of high symmetry. The bis(chelate) complex 4 is very stable. Heating to 373 K does not change the signals in the 31P-{1H} NMR spectrum. Furthermore, 4 is catalytically inactive up to 423 K. Together with the results obtained in the hydrocyanation of styrene, the NMR spectra of the catalyst solutions allow a deeper insight into the pathways of catalyst activity and deactivation for the various ligands (Scheme 6).The chelating Thixantphos ligands L1a–L1g and L2b preferably form bis(chelates) via route C, even if they are added substoichiometrically. In the case of the diphosphines L1a–L1g the bis(chelate) nickel complexes 3a–3g are still catalytically active since dissociation of the ligand occurs even at room temperature (route D). In contrast, the bis(chelate) complex [NiL2b 2] 4 is uncommonly stable and therefore not active in catalysis.High conversion of the substrate is only possible when the preformation time is short and the active catalyst precursor [Ni(cod)L2b] is still Fig. 6 Eyring plot: data obtained from variable-temperature 31P-{1H} NMR of [NiL1d 2] 3d Fig. 7 31P-{1H} NMR spectrum (toluene-d8, 298 K) of in situ catalyst preformation with [Ni(cod)2] and 1 equivalent of ligand L2b, recorded 30 minutes after mixing Scheme 6 Pathways of in situ catalyst preformation P Ni P P P P Ni(cod) P P P 2 P P D C B A catalysis [Ni(cod)2] present in large amounts (routes A and B).Regarding a preformation time of 30 min and a deactivation of about 80% of the nickel, the moderate values for conversion and yield in the presence of ligand L2b given in Table 4 must be reconsidered. Taking all this into account, L2b shows high catalytic activity due to electronic properties. Conclusion A series of electronically tuned Thixantphos ligands was applied to the nickel catalysed hydrocyanation of styrene. Electron-withdrawing substituents at the phosphorus atoms make the ligands more acidic and lead to best activities.The concept of large bite angles combined with rigid backbones is demonstrated by their importance for intermediates of the catalytic cycle and is supported by both crystal structures and 31P-{1H} NMR spectroscopy. A diphosphonite Thixantphos ligand with comparable geometry forms highly active nickel complexes.Bis(chelate) complexes of this ligand are very stable and catalytically inactive. Experimental Computational details Molecular mechanics calculations were performed on a Silicon Graphics Indigo2 workstation using SYBYL software and a modified TRIPOS force field.13 Natural bite angles were calculated using a method similar to that described by Casey and Whiteker,14 based on a Ni]P bond length of 2.177 Å and a P]Ni]P bending force constant of 0 kcal mol21 degree22 (cal = 4.184 J).Syntheses All preparations were carried out under an atmosphere of puri- fied argon using standard Schlenk techniques. Solvents were dried and freshly distilled prior to use. 2,8-Dimethylphenoxathiine and the diphosphines L1a–L1g were prepared as described in recent publications.10 The compound [Ni(cod)2] was synthesized according to literature methods.37 The NMR spectra were recorded on a Bruker DPX300 spectrometer (1H, 300; 13C, 75; 31P, 121 MHz).Hydrogen cyanide. CAUTION: HCN is a highly toxic, volatile liquid (b.p. 27 8C) that is susceptible to exothermic and uncontrolled polymerisation in the presence of basic catalysts. It should be handled only in a well ventilated fume hood and by teams of at least two technically qualified persons who have received appropriate medical training for treating HCN poisoning. Sensible precautions include also the use of HCN monitoring equipment. It can be generated by the addition of H2SO4 to sodium cyanide. Uninhibited HCN should be stored at a temperature lower than its melting point (213 8C). 4,6-Bis(diethylaminophosphino)-2,8-dimethylphenoxathiine L2a. 2,8-Dimethylphenoxathiine (3.26 g, 14.3 mmol) and N,N,N9,N9-tetramethylethane-1,2-diamine (4.14 g) were dissolved in diethyl ether (50 cm3) and cooled to 230 K. n-Butyllithium (14.3 cm3 of a 2.5 M solution in hexanes, 35.7 mmol) was added slowly, giving a bright yellow solution. The mixture was stirred at room temperature for 16 h and then added to a solution of chlorobis(diethylamino)phosphine (6.32 g, 35.7 mmol) in pentane (30 cm3) at 230 K.After stirring for 16 h at room temperature and evaporating the solvents a crude yellow product was obtained. The compound L2a was recrystallised from pentane, yield 3.90 g (48%) (Found: C, 61.6; H, 8.60; N, 9.3. C30H50N4OP2S requires C, 62.5; H, 8.75; N, 9.7%). NMR (C6D6): 1H, d 7.28 (s, 2 H), 6.74 (s, 2 H), 3.21 (m, 16 H, CH2), 2.09 (s, 6 H, CH3) and 1.13 (t, 12 H, 3J 7.0 Hz); 13C-{1H},J.Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 2987 d 151.0, 132.8, 131.0, 127.7, 127.1, 118.9, 43.5 (CH2), 20.7 (CH3) and 14.8 (CH2CH3); 31P-{1H}, d 90.4. 4,6-Bis(dibenzo[d,f ][1,3,2]dioxaphosphepino-1-yl)-2,8- dimethylphenoxathiine L2b. Compound L2a (577 mg, 1.0 mmol) and dihydroxybiphenyl (372 mg, 2.0 mmol) were dissolved in toluene (5 cm3) and the solution was stirred at 383 K for 15 h. The mixture was allowed to evaporate to dryness to give a bright yellow crystalline product, yield 650 mg (99%), m.p.>260 8C (Found: C, 69.1; H, 4.25. C38H26O5P2S requires C, 69.5; H. 4.00%). NMR (CDCl3): 1H, d 7.33 (d, 4J 2.1, 2 H), 7.30 (d, 4J 2.4, 2 H), 7.18–7.09 (m, 8 H), 6.96 (d, 4J 1.8 Hz, 2 H), 6.93– 6.90 (m, 6 H) and 2.01 (s, 6 H, CH3); 13C-{1H}, d 152.2, 151.4, 134.0, 132.0, 130.0, 129.8, 129.4, 129.0, 128.8, 122.1, 119.4 and 20.5 (CH3); 31P-{1H}, d 178.3. [NiCl2L1d] 1.Ligand L1d (853 mg, 1.43 mmol) and nickel(II) dichloride hexahydrate (340 mg, 1.43 mmol) were suspended in benzene (10 cm3) and stirred at 323 K for 3 h. Evaporating the solvent gave a reddish brown solid in quantitative yield. Red crystals suitable for X-ray analysis were obtained from boiling benzene, m.p. >260 8C (Found: C, 64.0; H, 4.37. C38H30Cl2Ni- OP2S requires C, 62.8; H, 4.16%). [Ni(CN)2L1a] 2. One large blue crystal of complex 2 was isolated from a toluene solution (2 cm3) of a catalytic experiment after standing for several days at room temperature.Before catalysis, the solution contained [Ni(cod)2] (17.9 mg, 0.065 mmol), ligand L1a (53.6 mg, 0.070 mmol), styrene (135.4 mg, 1.3 mmol) and hydrogen cyanide (42 mg, 1.625 mmol). Catalytic hydrocyanation of styrene In a typical experiment, a bright yellow 0.065 mM solution of [Ni(cod)2] in toluene (2 cm3) was added to a Schlenk tube containing a stirring bar and 1.05 equivalents of ligand.The mixture was stirred for 30 min to ensure complete formation of the catalyst precursor. Then styrene (1.3 mmol) was added. The solution was cooled to 220 K, liquid HCN (63 ml, 1.625 mmol) was added at once and the tube was placed in a heating bath. After 16 h at 333 K the excess of HCN was removed by a gentle stream of argon, solid particles were removed by centrifugation and the remaining solution was analysed by temperaturecontrolled gas chromatography. Crystallography Colourless crystals of ligand L1d were obtained from boiling acetone, red crystals of complex 1 from boiling benzene.Blue crystals of complex 2 were isolated from a catalysis experiment applying ligand L1a as described earlier. Data were collected on a CAD-4 diVractometer (graphite monochromator) and w–2q scans. Structure solution was by direct methods. Full-matrix least-squares refinement on F was carried out with anisotropic displacement parameters for all non-hydrogen atoms.Hydrogen atoms were placed in idealised positions with isotropic displacement parameters of U(H) = 1.3B(C) and allowed to ride on their C atoms. Calculations were performed using the SDP system of programs.38 L1d. C38H30OP2S, Mr 596.67, monoclinic, space group P21/c, a = 11.904(2), b = 21.377(2), c = 13.285(2) Å, b = 112.18(1)8, U = 3130.4(7) Å3, Dc = 1.266 g cm23, Z = 4, F(000) 1248, l(Cu]Ka) = 1.541 84 Å. Crystal size 0.40 × 0.25 × 0.20 mm. A total of 9557 reflections were collected at 293 K in the range 5.0 < q < 75.08, corresponding to 4989 unique data with I > 1.0s(I), which were used for further computations.An empirical absorption correction based on y scans was applied. Refinement converged with 499 parameters using a statistical weighting scheme at values of R = 0.062 and R9 = 0.063 with a goodness of fit of 1.280. Complex 1. C38H30Cl2NiOP2S?3C6H6, Mr 960.63 (including solvent of crystallisation), monoclinic, space group P21/c, a = 19.755(8), b = 19.314(8), c = 13.072(3) Å, b = 103.14(3)8, U = 4857(3) Å3, Dc = 1.314 g cm23, Z = 4, F(000) 2000, l(Mo- Ka) = 0.710 73 Å. Crystal size 0.40 × 0.15 × 0.20 mm.A total of 12 207 reflections were collected at 203 K in the range 2.0 < q < 25.08, corresponding to 5601 unique data with I > 1.0s(I), which were used for further computations. An empirical absorption correction based on y scans was applied. Refinement converged with 568 parameters using a statistical weighting scheme at values of R = 0.085 and R9 = 0.065 with a goodness of fit of 1.173.Complex 2. C48H50N6NiOP2S?2C7H8, Mr 1063.98 (including solvent of crystallisation), triclinic, space group P1, a = 11.436(5), b = 12.496(3), c = 21.699(8) Å, a = 77.20(2), b = 86.18(3), g = 65.86(2)8, U = 2758(2) Å3, Dc = 1.282 g cm23, Z = 2, F(000) 1124, l(Mo-Ka) = 0.710 73 Å. Crystal size 0.52 × 0.48 × 0.28 mm. A total of 10 824 reflections were collected at 203 K in the range 2.0 < q < 25.08, corresponding to 7164 unique data with I > 1.0s(I), which were used for further computations. Refinement converged with 658 parameters using a statistical weighting scheme at values of R = 0.081 and R9 = 0.072 with a goodness of fit of 1.414.CCDC reference number 186/1065. See http://www.rsc.org/suppdata/dt/1998/2981/ for crystallographic files in .cif format. Acknowledgements Financial support from the E.U. Human Capital and Mobility Program (MMCOS Network) and the Dutch Foundation for Chemical Research and Foundation for Technological Sciences (SON/STW) is gratefully acknowledged.References 1 K. Huthmacher and S. Krill, in Applied Homogeneous Catalysis with Organometallic Compounds, eds. B. Cornils and W. A. Hermann, VCH, Weinheim, 1996, p. 465. 2 R. J. McKinney, in Homogeneous Catalysis, ed. G. W. Parshall, Wiley, New York, 1992, p. 42. 3 W. C. Drinkard (DuPont), US Pat., 3 766 237, 1973; M. Rapoport (DuPont), US Pat., 4 371 474, 1983; W.Tam (DuPont), US Pat., 5 543 536, 1996. 4 K. A. Kreutzer and W. Tam (DuPont), US Pat., 5 512 696, 1996; A. I. Breikss (DuPont), US Pat., 5 523 453, 1996; K. A. Kreutzer and W. Tam (DuPont), US Pat., 5 663 369, 1997; W. Tam, K. A. Kreutzer and R. J. McKinney (DuPont), US Pat., 5 688 986, 1997. 5 M. J. Baker, K. N. Harrison, A. G. Orpen, P. G. Pringle and G. Shaw, J. Chem. Soc., Chem. Commun., 1991, 803; M. J. Baker and P. G. Pringle, J. Chem. Soc., Chem. Commun., 1991, 1292. 6 A. L. Casalnuovo, T. V. RajanBabu, T. A. Ayers and T. H. Warren, J. Am. Chem. Soc., 1994, 116, 9869. 7 T. V. RajanBabu and A. L. Casalnuovo, J. Am. Chem. Soc., 1996, 118, 6325. 8 M. Kranenburg, P. C. J. Kamer, P. W. N. M. van Leeuwen, D. Vogt, and W. Keim, J. Chem. Soc., Chem. Commun., 1995, 2177. 9 W. Goertz, P. C. J. Kamer, P. W. N. M. van Leeuwen and D. Vogt. Chem. Commun., 1997, 1521. 10 M. Kranenburg, Y. E. M. van der Burgt, P. C. J. Kamer, P. W.N. M. van Leeuwen, K. Goubitz and J. Fraanje, Organometallics, 1995, 14, 3081; M. Kranenburg, P. C. J. Kamer, P. W. N. M. van Leeuwen and B. Chaudret, Chem Commun., 1997, 373; M. Kranenburg, J. G. P. Delis, P. C. J. Kamer, P. W. N. M. van Leeuwen, K. Vrieze, A. L. Spek, K. Goubitz and J. Fraanje, J. Chem. Soc., Dalton Trans., 1997, 1839. 11 M. Kranenburg, P. C. J. Kamer and P. W. N. M. van Leeuwen, Eur. J. Inorg. Chem., 1998, 1, 25; 2, 155. 12 C. A. Tolman, J. Chem. Educ., 1986, 63, 199. 13 SYBYL, version 6.3, TRIPOS Associates, St. Louis, MO, 1996. 14 C. P. Casey and G. T. Whiteker, Isr. J. Chem., 1990, 30, 299. 15 C. A. Tolman, R. J. McKinney, W. C. Seidel, J. D. Druliner and W. R. Stevens, Adv. Catal., 1985, 33, 1.2988 J. Chem. Soc., Dalton Trans., 1998, Pages 2981–2988 16 C. A. Tolman, J. Am. Chem. Soc., 1970, 92, 2956; Inorg. Chem., 1971, 10, 1540; J. Am. Chem. Soc., 1972, 94, 2994; J. D. Druliner, A. D. English, J. P. Jesson, P. Meakin and C.A. Tolman, J. Am. Chem. Soc., 1976, 98, 2156; W. R. Jackson and C. G. Lovel, Aust. J. Chem., 1982, 35, 2053; C. A. Tolman, W. C. Seidel and L. W. Gosser, Organometallics, 1983, 2, 1391; J. E. Bäckvall and O. S. Andell, Organometallics, 1986, 5, 2350; R. J. McKinney and D. C. Roe, J. Am. Chem. Soc., 1986, 108, 5167. 17 J. M. Brown and P. J. Guiry, Inorg. Chim. Acta, 1994, 220, 249. 18 S. Hillebrand, J. Bruckmann, C. Krüger and M. W. Haenel, Tetrahedron Lett., 1995, 36, 75. 19 M. Kranenburg, Ph.D. Thesis, University of Amsterdam, 1996. 20 C. A. Hunter and J. K. M. Sanders, J. Am. Chem. Soc., 1990, 112, 5525. 21 M. D. Fryzuk, P. A. McNeil, S. J. Rettig, A. S. Secco and J. Trotter, Organometallics, 1982, 1, 918. 22 S.-T. Liu, G.-J. Liu, C.-H. Yieh, M.-C. Cheng and S.-M. Peng, J. Organomet. Chem., 1990, 387, 83; R. Busby, M. B. Hursthouse, P. S. Jarrett, C. W. Lehmann, K. M. A. Malik and C. Philips, J. Chem. Soc., Dalton Trans., 1993, 3767; A. L. Spek, B. P. van Eijck, R. J. F. Jans and G. van Koten, Acta Crystallogr., Sect. C, 1987, 43, 1878; F. Bachechi and L. Zambonelli, Acta Crystallogr., Sect. C, 1992, 48, 788. 23 G. Garton, D. E. Henn, N. M. Powell and L. M. Venanzi, J. Chem. Soc., 1963, 3625. 24 J. A. J. Jarvis, R. H. B. Mais and P. G. Owston, J. Chem. Soc. A, 1968, 1473. 25 P. T. Greene and L. Sacconi, J. Chem. Soc. A, 1970, 866. 26 V. Gramlich and G. Consiglio, Helv. Chim. Acta, 1979, 62, 1016. 27 L. M. Venanzi, J. Chem. Soc., 1958, 719. 28 B. Corain, M. Basato, G. Favero, P. Rosano and G. Valle, Inorg. Chim. Acta, 1984, 85, L27; H. Hope, M. M. Olmstead, P. P. Power and M. Viggiano, Inorg. Chem., 1984, 23, 326. 29 D. G. Holah, A. N. Hughes and N. I. Khan, Can. J. Chem., 1984, 62, 1016. 30 L. Manojlovic-Muir, K. W. Muir and M.-A. Rennie, Acta Crystallogr., Sect. C, 1995, 51, 1533. 31 C. P. Casey, G. T. Whiteker, M. G. Melville, L. M. Petrovich, J. A. Gavney and D. R. Powell, J. Am. Chem. Soc., 1992, 114, 5535. 32 H. M. Powell, D. J. Watkin and J. B. Wilford, J. Chem. Soc. A, 1971, 1803; J. K. Stalick and J. A. Ibers, Inorg. Chem., 1969, 8, 1090. 33 A. de Meijere and F. E. Meyer, Angew. Chem., Int. Ed. Engl., 1994, 33, 2379. 34 C. A. Tolman, W. C. Seidel, J. D. Druliner and P. J. Domaille, Organometallics, 1984, 3, 33. 35 T. V. RajanBabu and A. L Casalnuovo, J. Am. Chem. Soc., 1992, 114, 6265. 36 g-NMR, version 3.6, Ivory Soft, Cherwell Scientific, Oxford, 1996. 37 B. Bogdanovic, M. Kröner and G. Wilke, Liebigs Ann. Chem., 1966, 699, 1; R. A. Schunn, Inorg. Synth., 1974, 15, 5. 38 SDP program (Structure Determination Package), V5.0, B. A. Frenz & Associates, Inc., College Station, TX, 1989. Received 23rd March 1998; Paper 8/02269K

ISSN:1477-9226    

DOI:10.1039/a802269k

出版商:RSC    

年代:1998

数据来源: RSC

摘要:

DALTON FULL PAPER J. Chem. Soc., Dalton Trans., 1998, Pages 2989–2993 2989 Synthesis, crystal structure and electrochemical properties of [NBu4][Ni(mdt)2]: a potential precursor for new materials (mdt 5 1,3-dithiole-4,5-dithiolate) Yvonne S. J. Veldhuizen,a (the late) Nora Veldman,b Anthony L. Spek,†,b Patrick Cassoux,c Roger Carlier,d Martijn J. J. Mulder,a Jaap G. Haasnoot *,a and Jan Reedijk a a Leiden Institute of Chemistry, Gorlaeus Laboratories, Leiden University, PO Box 9502, 2300 RA Leiden, The Netherlands b Bijvoet Center for Biomolecular Research, Utrecht University, Padualaan 8, 3584 CH Utrecht, The Netherlands c Laboratoire de Chimie de Coordination du CNRS, 205 route de Narbonne, 31077 Toulouse Cedex, France d Laboratoire d’Electrochimie, UA 439, Université de Rennes 1, Beaulieu, 35042 Rennes Cedex, France Reaction of methylenedithio-1,3-dithiol-2-one with sodium methanolate resulted in the anionic ligand mdt22 (mdt = 1,3-dithiole-4,5-dithiolate) which has been co-ordinated to nickel. The salt [NBu4][Ni(mdt)2] was successfully crystallised and its crystal structure showed that the Ni(mdt)2 unit is significantly distorted from planarity.Electrochemical studies showed that [NBu4][Ni(mdt)2] can be further oxidised in two steps via the neutral [Ni(mdt)2] complex to a partially oxidised positively charged complex. The neutral complex has also been successfully synthesized by oxidation with TCNQ. The redox potential of the oxidation from the neutral to the partially oxidised positive Ni(mdt)2 complex is relatively low, 10.60 V, making this compound a promising precursor for the preparation of conducting materials, based on derived cation radical salts. 1 Introduction The crystallisation of the first partially oxidised Ni(dmit)2 compound (dmit = 4,5-disulfanyl-1,3-dithiole-2-thionate), [NBu4]2- [Ni(dmit)2]7?2MeCN, by Valade et al. in 1983,1 initiated a real impetus to the investigation of transition-metal complexes of sulfur donor ligands as possible conducting materials.Variants were synthesized by substitution of the cation, the ligand and the metal. Substitution of Bu4N1 by other cations resulted in a rich variety of [Cation]x[M(dmit)2] combinations (0 < x < 1; M = Ni, Pd or Pt), which were extensively reviewed in 1991 and 1992.2,3 The Bu4N cation has been substituted by mostly spherical tetraalkylammonium-type cations and alkali metals. A few compounds using planar cations such as guanidinium, acridinium and phenazinium have also been synthesized.4 Among the M(dmit)2 compounds are seven superconductors of which a-[EDT-TTF][Ni(dmit)2] [ EDT-TTF = ethylenedithiotetrathiafulvalene] is the only compound which becomes superconducting at ambient pressure at a Tc of 1.3 K.5 A variation on the dmit ligand can be made by substituting one or more of the sulfur atoms by the larger selenium atoms.The greater size and polarisability of selenium might diminish the Coulomb repulsion in a system.The selenium-containing ligands dmise and dsit are among the most studied (dmise = 4,5-disulfanyl-1,3-dithiole-2-selenonate; dsit = 4,5-diselanyl- 1,3-dithiole-2-thionate). Comparison of the conducting properties of [NMe4][Ni(dmise)2]2 6 and [NMe4][Ni(dsit)2]2 7 with the properties of [NMe4][Ni(dmit)2]2 8 shows, however, that in these compounds the substitution does not improve the conducting properties of the compounds: both selenium compounds show a semiconducting behaviour, while the dmit compound is a metallic conductor down to 100 K.On the contrary, the con- † To whom correspondence pertaining to the crystallographic studies should be addressed. ducting properties of [NHMe3][Ni(dmise)2]2 and [Me3NH]- [Ni(dmit)2]2 are very similar.9,10 Tight-binding band-structure calculations show that the selenium atoms in the dmise compound give this compound a three-dimensional electronic structure, whereas the dmit compound has only a onedimensional electronic structure.So, in this example, the larger selenium atoms improve the dimensionality of the conducting properties of the compound. Apparently, the conducting properties of Ni(dmit)2 can be improved as well as deteriorated by substituting sulfur atoms by selenium atoms. Another way for the preparation of new complexes as possible precursors for new inorganic molecular conductors is the synthesis of the inorganic analogues of successful donor molecules. As a d8 metal ion (M21) is isolobal to a (C2 41) unit,11 substitution of the C]] C bond of a donor molecule by a d8 metal ion results in the isolobal inorganic analogue of the donor molecule.A successful example of this strategy is the complex2990 J. Chem. Soc., Dalton Trans., 1998, Pages 2989–2993 M(dddt)2, first synthesized in 1985.12 This complex is the inorganic analogue of the successful donor BEDT-TTF [bis(ethylenedithio)tetrathiafulvalene], which is the precursor donor for a large number of organic superconductors.13 Several metallic conductors with M(dddt)2 have been reported, among which [Ni(dddt)2]3[HSO4]2 is the best synthesized with this complex so far [s(RT) = 60–300 S cm21, metallic behaviour down to 25 K].14 Here, the synthesis and characterisation of [NBu4][Ni(mdt)2] (mdt = 1,3-dithiole-4,5-dithiolate), a new variation of a bis- (dithiolene) nickel complex, is reported; Ni(mdt)2 is isolobal to the organic donor BMDT-TTF [bis(methylenedithio)tetrathiafulvalene], first synthesized in 1984.15 One of the most interesting conductors synthesized with this donor is [BMDTTTF] 2[Au(CN)2] [s(RT) = 300 S cm21, metallic behaviour down to 80 K].16 The complex [Ni(mdt)2] can also be considered as a variation on [Ni(dddt)2] and their properties may therefore be expected to be comparable. Parts of these results have been published as a preliminary communication.17 2 Experimental 2.1 Synthesis of [NBu4][Ni(mdt)2] The route to the synthesis of [NBu4][Ni(mdt)2] is shown in Scheme 1.All reactions were carried out under a dinitrogen atmosphere. Ligand precursors 1 and 2 were synthesized by a combination of the methods described by Papavassiliou et al.18 and Nigrey et al.19 The final product was synthesized according to the method described by Faulmann et al.20 for M(dddt)2 compounds. Methylenedithio-1,3-dithiole-2-thione 1. The complex [NBu4]2- [Zn(dmit)2] (4.72 g, 5 mmol) was dissolved in acetone p.a.(75 cm3). Dibromomethane (ca. 7 cm3 100 mmol) was added and the resulting solution brought to reflux and heated for 4 h. A change from red to brown occurred during the reaction and a yellow-orange solid precipitated. The mixture was concentrated under vacuum until almost dry after cooling. The residue was washed with hot dichloromethane to extract the product until the washings were colourless (approximately 300 cm3 of dichloromethane were necessary).The yellow-brown solution was concentrated under vacuum to a small volume. The product was filtered oV and washed with methanol. Purity was checked by infrared spectroscopy for the presence of Bu4N1 (characteristic pattern between 3000 and 2800 cm21). If necessary the product was further purified by suspending and stirring it in a small amount of methanol, filtration and drying. Yield ª65%. IR (KBr): 3002w, 2930w, 1741w, 1667s, 1603m, 1495m, 1402w, 1388w, 1190w, 1072m, 962m, 876m, 852m, 743m, 710w, 688w, 552w, 536w, 506w, 472m, 400m, 376w and 338w cm21. Scheme 1 Methylenedithio-1,3-dithiol-2-one 2.Compound 1 (1.05 g, 5 mmol) was dissolved in a mixture of chloroform (125 cm3), glacial acetic acid (50 cm3) and demineralised water (10 cm3). After the solution was brought to reflux, Hg(O2CMe)2 (1.60 g, 5 mmol) was carefully added to the boiling mixture. The resulting mixture was stirred under reflux for 20 h, while black HgS precipitated. After cooling the black precipitate was filtered oV and the filtrate concentrated to dryness under vacuum.The yellow product was recrystallised from dichloromethane and methanol. Yield ª72%. IR (KBr): 1466m, 1054m, 1036s, 1011m, 952m, 890w, 683w, 508m, 479w, 453w, 392w and 309w cm21. [NBu4][Ni(mdt)2]. Compound 2 (0.194 g, 1 mmol) was suspended in methanol p.a. (4 cm3) after which a solution (6 cm3) of 1 M NaOMe was added slowly. After 1 h of stirring all 2 had reacted and a clear brown solution was formed.The compound NiCl2?6H2O (0.119 g, 0.5 mmol) in methanol (10 cm3), was added using a transfer tube, resulting in a change to orange-red. After another hour of stirring a solution of NBu4Br (0.17 g, 0.5 mmol) in methanol (5 cm3) was added using a transfer tube. The product crystallised by standing overnight at 218 8C. The orange-brown crystals were filtered oV, washed with PriOH and dried under vacuum. Yield >85%. Recrystallisation from acetone–PriOH (1: 1) resulted in single crystals suitable for Xray determination.The oxidation state of 12 is in accord with the C]] C stretching vibration at 1330 cm21 in the infrared spectrum and was confirmed by elemental analysis (C, H, N, S) performed by the University College Dublin. IR (KBr): 2956m, 2920w, 2867m, 1472m, 1443w, 1435w, 1410w, 1385w, 1330s, 1030w, 927w, 739m, 455m and 352m cm21 [Calc. (Found) for C22H40NNiS8: C, 41.69 (41.72); H, 6.36 (6.34); N, 2.21 (2.02); S, 40.47 (39.44%)]. 2.2 Synthesis of [Ni(mdt)2] The complex [NBu4][Ni(mdt)2] (0.158 g, 0.25 mmol) was dissolved in benzonitrile (10 cm3) at 80 8C and TCNQ (0.051 g, 0.25 mmol) was dissolved in boiling acetonitrile (10 cm3). The hot TCNQ solution was added to the hot [NBu4][Ni(mdt)2] solution using a transfer tube, resulting in a change from orange-brown from [NBu4][Ni(mdt)2] to green from TCNQ2. After cooling to room temperature the black precipitate of [Ni(mdt)2] was filtered oV, washed with hot acetonitrile and acetone and dried under vacuum.Yield >95%. The neutral oxidation state was shown by the strong C]] C stretching vibration at 1195 cm21 in the infrared spectrum and confirmed by elemental analysis (C, H, S) performed by the Laboratoire de Chimie de Coordination for CNRS in Toulouse. IR (KBr): 2923w, 2522w, 2081w, 1338m, 1195s, 1054m, 964m, 684m, 553w, 466s and 364m cm21 [Calc. (Found) for C6H4NiS8: C, 18.42 (19.13); H, 1.03 (0.65); S, 65.55 (66.30%)]. Attempts to synthesize single crystals of [Ni(mdt)2] by slow interdiVusion of solutions of [NBu4][Ni(mdt)2] and TCNQ resulted in hair-thin needles, not suitable for X-ray determination. 2.3 X-Ray data collection and structure determination of [NBu4][Ni(mdt)2] X-Ray data were collected on an Enraf-Nonius CAD-4T/ rotating anode diVractometer for a dark, cut to size blockshaped crystal (0.25 × 0.25 × 0.25 mm). Numerical data are collected in Table 1. Unit-cell dimensions were derived from the SET4 setting angles 21 (10 < q < 148) and checked for higher lattice symmetry with LEPAGE.22 The structure was solved by automated Patterson techniques (DIRDIF 92 23) and refined on F2 by full-matrix least-squares techniques (SHELXL 93 24).All non-hydrogen atoms were refined with anisotropic thermal parameters. Hydrogen atoms were introduced at calculated positions and refined riding on their carrier atoms with fixed isotropic displacement parameters related to the value of theJ.Chem. Soc., Dalton Trans., 1998, Pages 2989–2993 2991 equivalent isotropic displacement parameter of the atom they are attached to, by a factor 1.5 for CH3 and 1.2 for the other hydrogen atoms, respectively. The Flack absolute structure parameter refined to 0.02 (0.06). CCDC reference number 186/1075. 2.4 Electrochemical experiments The redox properties of [NBu4][Ni(mdt)2] and the possibility to synthesize partially oxidised Ni(mdt)2 compounds were studied by electrochemical experiments. Cyclic voltammetry (CV) was carried out in an air-tight three-electrode cell, using laboratory-made, microcomputer-controlled instrumentation with “interrupt method” ohmic resistance compensation.25 A platinum wire auxiliary electrode was used in conjunction with a platinum disc working electrode (Tacussel EDI rotating electrode, 1 mm diameter).The measurement was performed on a 1023 M solution of [NBu4][Ni(mdt)2] in nitrobenzene with NBu4BF4 (0.1 M) as supporting electrolyte, carried out in the region 20.70 to 0.85 V, starting at 20.20 V, with a potential scan rate of 0.1 V s21.All potentials were referred to a saturated calomel electrode (SCE) separated from the solution by a bridge compartment filled with the same solvent and supporting electrolyte solution as in the cell. A linear voltammetry experiment was carried out with a 2 mm platinum disc working electrode, at a rotation speed of 1000 min21, and a scan rate of 5 mV s21.Coulometric measurements were performed on a platinum electrode with a large surface area at a potential slightly higher than the considered oxidation potential. A thin-layer CV measurement was performed to study the possibility of the synthesis of a partially oxidised Ni(mdt)2 compound. With this special technique the CV experiment is carried out on a thin layer of solution, with an accurately known volume and concentration. The principle of the measurement is that during the diVerent oxidation and reduction processes all the electroactive material present in the thin layer undergoes a charge transfer. Integration of the peak surfaces results in the exact amount of electrons transferred during the diVerent reactions.26 The thin-layer CV experiment was performed on a 2.5 × 1024 M solution of [NBu4][Ni(mdt)2] in dichloromethane with NBu4PF6 (1 M) as supporting electrolyte.A platinum electrode with a surface of 1 mm2 and a Fig. 1 An ORTEP 50% probability plot (PLATON)27 of the crystallographically independent unit of [NBu4][Ni(mdt)2] with the labelling scheme of the heavy atoms. Hydrogen atoms have been omitted for clarity. Fig. 2 Side view showing the distortion from planarity of the Ni(mdt)2 anion. potential scan rate of 2 mV s21 were applied and a EG&E potentiostat, Princeton Applied Research model 362, was used. 3 Results and discussion 3.1 Crystal structure The asymmetric unit of the crystal structure contains one Ni(mdt)2 unit and one Bu4N cation.Fig. 1 shows the labelling scheme used for the heavy atoms in the anion and the cation. It can be seen that the torsion angle of one of the butyl groups deviates from the expected antiperiplanar orientation. The torsion angle C(19)]C(20)]C(21)]C(22) is 260.2(8)8, i.e a synclinal orientation, which is probably due to packing eVects. Table 2 lists the bond distances and bond angles of the Ni(mdt)2 unit.The observed Ni]S bond lengths vary from 2.154(2) to 2.170(2) Å, slightly larger than those found in [NBu4][Ni(dddt)2], 2.130(6)–2.145(9) Å,28 but are similar to those in [NBu4][Ni(dmit)2], 2.151(3)–2.160(3) Å.29 The large deviation from 1808 for the angles S(3)]Ni]S(6) and S(4)]Ni]S(5), 173.37(10) and 166.53(10)8 respectively, indicates a large distortion from a square-planar configuration for the NiS4 co-ordination geometry. This distortion is clearly illustrated by the side view of Ni(mdt)2 presented in Fig. 2.The twist between the planes S(3)]Ni]S(4) and S(5)]Ni]S(6) is 14.77(13)8, showing the distorting towards a tetrahedral geometry. Furthermore, the nickel centre is forming a plane together with the fragment S(5)]S(6)]C(4)]C(5)]S(7)]S(8) [maximum Fig. 3 Unit-cell contents of [NBu4][Ni(mdt)2] down the a and the c axis. Hydrogen atoms have been omitted for clarity. Table 1 Crystallographic data for [NBu4][Ni(mdt)2] Formula M Space group Crystal system a/Å b/Å c/Å U/Å3 Z (formula units) Dc/g cm23 m/cm21 F(000) T/K qmin,qmax Total data Total unique data (n) Observed data No.parameters (p) R1b wR2c Sd C22H40NNiS8 633.79 Pna21 (no. 33) Orthorhombic 17.2633(6) 18.4274(9) 9.0116(8) 2866.8(3) 4 1.4685(2) 12.7 a 1340 150 1.6, 26.5 5766 2949 2071 [I > 2s(I)] 294 0.050 0.0825 1.00 a Graphite monochromated Mo-Ka radiation; l = 0.710 73 Å. b R1 = S||Fo| 2 |Fc ||/S|Fo|. c wR2 = [Sw(Fo 2 2 Fc 2)2/Sw(Fo 2)2]� �� ; w21 = s2(Fo 2) 1 (0.0233P)2, where P = (Fo 2 1 2Fc 2)/3.d S = [Sw(Fo 2 2 Fc 2)2/(n 2 p)]� �� .2992 J. Chem. Soc., Dalton Trans., 1998, Pages 2989–2993 atomic deviation 0.038(2) Å for S(6)], whereas a second plane is formed in the molecule by the fragment S(1)]S(2)]C(2)] C(3)]S(3)]S(4) [maximuiation 0.063(2) Å for S(2)]. The plane with S(1) is tilted 24.84(9)8 from that with Ni(1), showing the large deviation from planarity of the anion. An inclination of 6.18 between the two ligands in Ni(dmit)2 has been reported for [NBu4][Ni(dmit)2], but a large distortion as here has not been found in any comparable nickel dithiolene complex and donor molecule.However, the distortion may not be large enough to prohibit large orbital overlap and therefore good conducting properties in derived compounds. Moreover, this is the first crystal structure containing Ni(mdt)2 and it is not clear at present whether the deviation is due to packing eVects, or to internal tensions in the complex.Fig. 3 shows the unit cell viewed down the a and the c axis. For clarity the four molecules on the right-hand side in Fig. 3(a) have been translated one step in the c direction. No stacking or dimerisation of Ni(mdt)2 units is present in the structure. The shortest Ni ? ? ? Ni distance found is 8.8806(19) Å, from the crystallographically independent Ni(mdt)2 unit towards the Ni atom of the unit with a symmetry operation 2x, 2y, 0.5 1 z.The shortest intermolecular S ? ? ? S distance present is between the atoms S(1) and S(8) [with a symmetry operation of 0.5 2x, 0.5 1 y, 0.5 1 z on S(8)] and indicated with a dotted line in Fig. 3(b). This distance is 3.779(3) Å, which is slightly larger than the sum of the van der Waals radii,30 i.e. 3.70 Å. Also the Fig. 4 Cyclic voltammogram of [NBu4][Ni(mdt)2] in nitrobenzene; * = start of the scan. Table 2 Bond distances (Å) and angles (8) of the Ni(mdt)2 unit in [NBu4][Ni(mdt)2] (estimated standard deviations in parentheses) Ni]S(3) Ni]S(4) Ni]S(5) Ni]S(6) S(1)]C(1) S(1)]C(2) S(2)]C(1) S(2)]C(3) S(3)]C(2) S(3)]Ni]S(4) S(3)]Ni]S(5) S(3)]Ni]S(6) S(4)]Ni]S(5) S(4)]Ni]S(6) S(5)]Ni]S(6) C(1)]S(1)]C(2) C(1)]S(2)]C(3) Ni]S(3)]C(2) Ni]S(4)]C(3) Ni]S(5)]C(4) Ni]S(6)]C(5) C(4)]S(7)]C(6) C(5)]S(8)]C(6) 2.167(2) 2.154(2) 2.155(2) 2.170(2) 1.807(8) 1.744(8) 1.833(8) 1.732(8) 1.729(8) 92.52(8) 88.08(8) 173.37(10) 166.53(10) 88.27(9) 92.69(8) 95.0(3) 94.6(4) 102.2(3) 102.8(3) 102.8(3) 102.3(3) 93.5(4) 93.9(3) S(4)]C(3) S(5)]C(4) S(6)]C(5) S(7)]C(4) S(7)]C(6) S(8)]C(5) S(8)]C(6) C(2)]C(3) C(4)]C(5) S(1)]C(1)]S(2) S(1)]C(2)]S(3) S(1)]C(2)]C(3) S(3)]C(2)]C(3) S(2)]C(3)]S(4) S(2)]C(3)]C(2) S(4)]C(3)]C(2) S(5)]C(4)]S(7) S(5)]C(4)]C(5) S(7)]C(4)]C(5) S(6)]C(5)]S(8) S(6)]C(5)]C(4) S(8)]C(5)]C(4) S(7)]C(6)]S(8) 1.723(8) 1.727(8) 1.722(7) 1.756(7) 1.821(8) 1.759(8) 1.800(7) 1.372(11) 1.349(10) 108.1(4) 122.0(4) 117.1(6) 120.7(6) 122.9(5) 116.8(6) 120.1(6) 122.2(4) 120.7(6) 117.0(6) 121.5(4) 121.4(6) 117.0(5) 108.9(4) structures of [NBu4][Ni(dmit)2] and [NBu4][Ni(dddt)2] do not show any S ? ? ? S interactions shorter than 3.70 Å.28,29 3.2 Electrochemical properties Fig. 4 shows the cyclic voltammogram of [NBu4][Ni(mdt)2] in nitrobenzene recorded between 20.70 and 0.85 V at a potential scan rate of 0.1 V s21. The voltammogram shows the presence of three redox couples in the recorded region. The first reversible wave (A1, C1) at E2� 1 = 20.55 V vs.SCE corresponds to the redox couple [Ni(mdt)2]22–[Ni(mdt)2]2, while the second reversible wave (A2, C2) at E2� 1 = 0.08 V vs. SCE corresponds to the redox couple [Ni(mdt)2]2–[Ni(mdt)2]. Linear voltammetry showed that the number of electrons transferred in the two reactions is the same and coulometry showed that one electron is transferred in each. The third wave (A3, C3) in the cyclic voltammogram at E2� 1 = 0.60 V vs. SCE shows that the neutral complex [Ni(mdt)2] can be further oxidised to a cationic species. However, the amount of electrons transferred in this oxidation could not be determined by either linear voltammetry, because of deposition of the product on the electrode, or by coulometry, because of the poor solubility of [Ni(mdt)2] in nitrobenzene.The amount of electrons transferred was therefore determined by a thin-layer CV experiment, carried out in dichloromethane with NBu4PF6 as a supporting electrolyte. This experiment con- firmed the transfer of one electron in the first two reactions and showed that 0.5 electron is transferred during the third reaction.A comparison of the electrochemical properties of [NBu4]- [Ni(mdt)2], [Ni(dddt)2], [Ni(pddt)2] and [Ni(mtdt)2] is presented in Table 3 [pddt = propane-1,3-diyldithioethylene-1,2-dithiolate; mtdt = 1,2-bis(methylsulfanyl)-1,2-dithiolate].31,32 So far [Ni(mdt)2] and [Ni(dddt)2] are the only complexes for which oxidation to a partially oxidised positively charged complex has been reported; [Ni(mdt)2] even shows a lower redox potential for this oxidation step, indicating that it should be easier to synthesize the positively charged Ni(mdt)2 than the Ni(dddt)2 complex. 4 Conclusion A new nickel bis(dithiolene) complex, [Ni(mdt)2], has been successfully synthesized. The crystal structure of [NBu4][Ni(mdt)2] shows that the Ni(mdt)2 unit is largely distorted from planarity. However, the distortion may not be large enough to prohibit considerable orbital overlap and therewith good conducting properties in derived oxidised compounds.An electrochemical study showed that [NBu4][Ni(mdt)2] can be further oxidised to a partially oxidised positively charged complex. Comparison of the electrochemical properties of [NBu4][Ni(mdt)2] and of [NBu4][Ni(dddt)2] shows that the redox potential for the final oxidation step is lower for Ni(mdt)2 than for Ni(dddt)2. This indicates that it should be easier to synthesize the positively charged Ni(mdt)2 than the Ni(dddt)2 complex, making [Ni(mdt)2] a promising precursor for future conducting materials.Electrooxidation experiments to synthesize single crystals of partially oxidised Ni(mdt)2 complexes are in progress. Table 3 Comparison of electrochemical properties of several nickel bis(dithiolene) complexes vs. SCE (0 < x < 1) E2� 1 /V Complex [Ni(mdt)2] [Ni(dddt)2] [Ni(pddt)2] [Ni(mtdt)2] Solvent, electrolyte Nitrobenzene, NBu4BF4 Nitrobenzene, NEt4BF4 dmf, NEt4ClO4 dmf, NBu4ClO4 22/12 20.55 20.79 20.66 20.73 12/0 10.08 20.11 10.24 10.04 0/x1 10.60 10.82 —— Ref. This work 20 31 32J.Chem. Soc., Dalton Trans., 1998, Pages 2989–2993 2993 5 Acknowledgements Technical assistance by D. de Montauzon in carrying out cyclic voltammetry studies is gratefully acknowledged. This work was supported in part (A. L. S., N. V.) by the Netherlands Foundation of Chemical Research (SON) with financial aid from the Netherlands Organisation for Scientific Research (Y.S. J. V.) and in part (Y. S. J. V., M. J. J. M., J. G. H., J. R.) by WFMO (Werkgroep Fundamenteel Materialen Onderzoek of Leiden University). Financial support by the European Community, allowing exchange of preliminary results with several European colleagues, under Contract ERBCHRXCT920080 is thankfully acknowledged. Also support and sponsorship concerted by COST Action D4/0001/95 (Chemistry of Molecular Materials) is kindly acknowledged.References 1 L. Valade, M. Bousseau, A. Gleizes and P. Cassoux, J. Chem. Soc., Chem. Commun., 1983, 110. 2 P. Cassoux, L. Valade, H. Kobayashi, A. Kobayashi, R. A. Clark and A. E. Underhill, Coord. Chem. Rev., 1991, 110, 115. 3 R.-M. Olk, B. Olk, W. Dietzsch, R. Kirmse and E. Hoyer, Coord. Chem. Rev., 1992, 117, 99. 4 Y. S. J. Veldhuizen, N. Veldman, A. L. Spek, C. Faulmann, J. G. Haasnoot and J. Reedijk, Inorg. Chem., 1995, 34, 140; Y. S. J. Veldhuizen, Ph.D.Thesis, Leiden University, 1997. 5 H. Tajima, M. Inokuchi, A. Kobayashi, T. Ohta, R. Kato, H. Kobayashi and H. Kuroda, Chem. Lett., 1993, 1235. 6 J. P. Cornelissen, D. Reefman, J. G. Haasnoot, A. L. Spek and J. Reedijk, Recl. Trav. 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ISSN:1477-9226    

DOI:10.1039/a804034f

出版商:RSC    

年代:1998

数据来源: RSC