摘要:

DALTON PERSPECTIVE J. Chem. Soc., Dalton Trans., 1999, 1519–1529 1519 The bite angle makes the diVerence: a practical ligand parameter for diphosphine ligands Peter Dierkes * and Piet W. N. M. van Leeuwen Institute for Molecular Chemistry, Nieuwe Achtergracht 166, NL-1018 WV Amsterdam, The Netherlands Received 7th October 1998, Accepted 5th February 1999 Over the past twenty years, a correlation between the P–M– P bite angle in diphosphine complexes and selectivity has been observed in various catalytic reactions such as hydroformylation, hydrocyanation and cross coupling.The large number of examples indicates that this correlation is not fortuitous. In order better to understand the underlying principles of the bite angle effect, we have first analysed crystal structures available in the Cambridge Crystallographic Database. Systematic searches indicate that for many bidentate diphosphine ligands the P–M–P angles concentrate in surprisingly small ranges, even if complexes of different metals in various oxidation states are considered.Several examples in the literature show that continuous electronic changes associated with changing bite angles cannot only be verified by different spectroscopic techniques, but also explained on a theoretical level (Walsh diagrams). The ligand bite angle is a useful parameter for the explanation of observed rates and selectivities and likewise for the design of ligands for new catalytic reactions. 1 Introduction Homogeneous catalysis has reached a state of maturity that allows its application in organic synthesis and industry. The optimisation of known procedures and the eYcient development of new catalytic systems call for a more systematic approach to ligand design. Several empirical ligand parameters have been suggested to predict catalyst performance. Many of them can help to understand qualitatively why a ligand leads to the observed rate and/or selectivity. A ligand parameter that can easily be evaluated from low-level computer modelling and/ or systematic crystal structure analyses could be a valuable tool to determine which ligand to use for a reaction or how to modify a given ligand to gain more control over a catalytic reaction.More than 1000 citations have been made in the last ten years to Tolman’s review article 1,2 introducing the cone angle concept for phosphine ligands. A large number of studies have been reported since on more advanced parameters for (mono)phosphine ligands.Models using two or three parameters can be used to predict properties ranging from pKa values to various spectroscopic properties of phosphines/phosphites and the corresponding metal complexes.3–5 While Tolman’s cone angle concept is widely accepted for monodentate ligands, the extension to bidentate ligands appears to be less straightforward. Diphosphine ligands, however, oVer more control over regio- and stereo-selectivity in many catalytic reactions.The major diVerence between monoand bi-dentate ligands is the ligand backbone, a scaVolding which keeps two phosphorus donor atoms at a specific distance (Scheme 1). The distance is ligand specific and, together with the flexibility of the backbone, an important characteristic of a ligand. A standardised bite angle, with defined M–P bond lengths and a “metal” atom that does not prefer any specific P–M–P angle would appear to be the most convenient way of comparing bidentate ligands systematically. Scheme 1 The P ? ? ? P distance is determined by backbone constraints.It can be measured most easily as a standardised ligand bite angle using a dummy atom to ensure the right orientation of the two donor atoms. Peter Dierkes studied chemistry at Münster University and Imperial College in London. After taking a Ph.D. in Professor Dehnicke’s group in Marburg, Germany, 1994, he spent two years as a postdoc in Professor Osborn’s laboratories in Strasbourg, and is currently working at the University of Amsterdam towards his Habilitation.Research interests include organometallic chemistry, the preparation of chiral ligands, mechanistic studies of catalytic reactions and molecular modelling on different levels of theory. Piet W. N. M. van Leeuwen is professor of homogeneous catalysis at the University of Amsterdam. He did his Ph.D. in co-ordination chemistry at the University of Leyden. He spent much of his career at Shell Research in Amsterdam.His research aims at the development of novel transition-metal homogeneous catalysts, using the full range of available tools and techniques. Peter Dierkes Piet W. N. M. van Leeuwen1520 J. Chem. Soc., Dalton Trans., 1999, 1519–1529 Of the large number of conformations “free” ligands can adopt, few contain the phosphorus atoms with the correct orientation to allow bidentate co-ordination to a single metal centre. The lone pairs of electrons have to point in the direction of the metal centre.The easiest way to find these conformations is the introduction of a dummy metal atom. If the same dummy–phosphorus bond length is used for all ligands, the calculated bite angles are a function of the non-bonded P ? ? ?P distance. We will describe the “bite angles” obtained in this way as ligand bite angle, as opposed to P–M–P “bite” angles measured in crystal structures. The evaluation of parameters aVecting catalyst eYciencies is not always straightforward.Catalytic reactions are sequences of elementary reactions. Variation of a single parameter may promote one step but slow down another. Equilibria involving catalytically active species can be shifted by small changes in reaction conditions.6,7 The overall eVect may not be very characteristic. 8 The potential formation of diVerent catalytic species has to be taken into account when a series of similar ligands are compared.With increasing ligand bite angles, the formation of trans complexes or dimeric species becomes more likely.9 Increasing flexibility of a ligand backbone raises the chance of an arm-oV h1 co-ordination. The latter may explain the sometimes drastically diVerent eYciency of Ph2P(CH2)4PPh2, dppb, compared to Ph2P(CH2)3PPh2, dppp, observed in various reactions. Despite these limitations, many examples show that the ligand bite angle is related to catalytic eYciency in a number of reactions.Early examples are the platinum–diphosphine– tin catalysed hydroformylation 10,11 or palladium catalysed cross coupling reactions of Grignard reagents with organic halides.12,13 In recent years, a correlation between ligand bite angles and catalyst selectivities has been observed in rhodium catalysed hydroformylation,14–16 nickel catalysed hydrocyanation 17,18 and even Diels–Alder reactions.19 Some catalytic reactions, for example nickel–diphosphine catalysed hydrocyanation, only work if ligands with very large bite angles (>1008) are employed.In other reactions, high enantioselectivities with specific substrates require ligands with large bite angles.20–22 The P–M–P angle found in transition metal complexes is a compromise between the ligand’s preferred bite angle and the one preferred by the metal centre. The former is mainly determined by constraints imposed by the ligand backbone and by steric repulsion between substituents on the phosphorus atoms and/or the backbone.Electronic eVects seem to have a more indirect influence by changing the preferred metal–phosphorus bond length. The metal preferred bite angle, on the other hand, is mainly determined by electronic requirements, i.e. the nature and number of d orbitals involved in forming the molecular orbitals. Other ligands attached to the metal centre can influence the bite angle if they are very bulky or if they have a strong influence on the metal orbitals (p-bonding ligands for example).Ligands enforcing unusually small or large bite angles have long been employed to synthesize transition metal complexes that are not stable when ligands with classic ligand bite angles are used. p Complexes with dioxygen are an early example,23 and Hofmann et al.24 synthesized [Ni{h2-(C,O)-Ph2C]] C]] O}- (P–P)] (P–P = bidentate diphosphine) complexes after correctly predicting an increasing stability of h2-(C,O) compared to h2- (C,C) co-ordination with decreasing bite angles.For the calculation of ligand bite angles, either molecular modelling or P ? ? ? P distances determined from crystal structures can be used. Molecular modelling has been used to calculate “natural” bite angles, ligand bite angles calculated using a “rhodium” dummy atom and fixed Rh–P distances of 2.315 Å.25 Crystallographic data can be retrieved from the Cambridge Crystallographic Database. With the large number of crystal structures available, average values can be calculated for many ligands.This decreases errors due to packing eVects or other “errors” in the crystal structures. Bite angles are a function of the M–P bond length. In order to obtain a meaningful comparison between diVerent ligands, the measured (or calculated) angles have to be standardised to one, defined M–P bond length. Standardised bite angles can be calculated from the P ? ? ? P distance (Table 1). In the following, a discussion of ligand and metal preferences for certain bite angles will be followed by examples illustrating the use of similar ligands with diVerent bite angles in various catalytic reactions. 2 The ligand preferred bite angle Statistical analyses of crystal structures The examination of crystal structures is a good starting point to gain more insight into ligand and metal preferences contributing to the actually measured bite angle. The P–M–P bite angles of transition metal complexes containing (P–P)M fragments documented in the Cambridge Crystallographic Database CSD have been examined for a series of bidentate diphosphine ligands P–P.No restrictions were imposed on the nature of the transition metal M, its oxidation state or other ligands co-ordinated to the same metal centre. The results are summarised in Table 2. Despite the rather crude filtering,26,27 the angles concentrate in a narrow range for most ligands; the standard deviations, a measure for the broadness of a distribution around the mean value, lie between 1.5 and 38 (see Fig. 1). The standard deviations are similar to those Müller and Mingos28 found for Tolman cone angles of monophosphine ligands. This is remarkable, as one would expect ligands with Fig. 1 Number of (P–P)M fragments with P–P = dppm (r), dppe (j), dppp (m) and dppf (×) found in a CSD search. For details of structure searches see Table 2. Table 1 The bite angle of a ligand with a given P ? ? ? P distance strongly depends on the M–P bond length: the bite angle for a defined bond length can be calculated from the P ? ? ? P distance [P–M–P = 2arcsin(rP ? ? ?P/2rM–Pnew)] or a measured bite angle and the M–P-distances rP ? ? ?P is the P ? ? ? P distance and M–Pnew the new bond length)]J.Chem. Soc., Dalton Trans., 1999, 1519–1529 1521 Table 2 Diphosphine ligand bite angles (8) calculated with force field methods compared to average P–M–P angles calculated from crystal structures retrieved from the CSD X-ray, average value Ligand dppm dpp-benzene dppe dppe-Bud dppp dppp-Med dppp-Bud dppb dppf BINAP DIOP DUPHOS-Me BISBI NORPHOS TRANSPHOS T-BDCP DPEphos Xantphos Thixantphos Sixantphos DBFphos P–M–P2.315 a/ 8 71.71 (1.60) 83.04 (2.73) 85.03 (3.11) 87.23 (1.26) 91.08 (4.00) 89.87 (3.86) 99.33 (0.74) 97.70 (5.15) 95.60 (4.34) 92.43 (2.6) 97.63 (4.72) 82.61 122.18 (14.35) 102.51 107.12 105.36 P–M–P/8 71.53 (2.44) 81.95 (3.25) 82.55 (3.65) 89.74 (2.49) 91.56 (3.70) 90.51 (2.27) 100.35 (1.53) 97.07 (2.84) 98.74 (3.42) 92.77 (1.95) 100.0 (4.3) 84.7 68 119.64 (43.48) 104, 131.9, 175.7 g 101.46 i 104.64 i 104.28 i rM–P,av (CSD)/Å 2.32 (0.09) 2.34 (0.12) 2.38 (0.13) 2.27 (0.06) 2.29 (0.09) 2.30 (0.06) 2.30 (0.02) 2.38 (0.11) 2.26 (0.07) 2.31 (0.04) 2.28 (0.02) 2.27 2.34 (0.12) 2.33 i 2.36 i 2.34 i Molecular modelling, bn/ 8 78.1,b 84.4 (70–95) c 86.2 b 98.6 b 102.2 (90–120) f 122.6 (101–148) f 112.6 (92–155) c 123 (110–145) c 111.2 h 107.6 (93–131) c 102.2 (86–120) f 111.7 (97–135) f 109.4 (94–130) f 108.7 (93–132) f 131.1 (117–147) f The bite angles given are based on crystal structure data retrieved from the October 1997 version of the CSD.The data have been filtered: the structure contains a d- or f-block metal, “R < 10”, the “error free” and the “no disorder” options were enabled, and the coordinates had to be available in the database. If a structure contained a complex with more than one diphosphine ligand bound to a metal centre, or if more than one molecule was present in an elementary cell, or if a structure was determined more than once, all entries were processed.Ligands carrying substituents other than H on the phenyl rings or the bridge and ligands with a mono- or tri-dentate co-ordination have not been used. The structures of the ligands are shown in Tables 3 and 4. For crystal structure data, the values given in parentheses are the standard deviations; for modelled bite angles they indicate the range of bite angles a ligand can accommodate with no more than 3 kcal mol21 strain energy (flexibility range).a Standardised ligand bite angle with M–P distances of 2.315 Å calculated from the P ? ? ? P distance as found in the CSD {=2arcsin1– 2(rP ? ? ?P/2.315)}. b Ref. 63. c Ref. 14. d dppe- Bu = tBu2P(CH2)2PtBu2, dppp-Me = Me2P(CH2)3PMe2 and dppp-Bu = tBu2P(CH2)3PtBu2. e Averages calculated excluding structures of tetrahedral nickel(II) complexes [(ca. 1048, rM–P ª 2.30 Å) and complexes of CuI, AgI and AuI with d10 metal centres 111.428 (2.47) (rM–P 2.32 Å), see Fig. 2]. f Ref. 15. g Cited in ref. 25. h Calculated from 112.38 25 with a Rh–P distance of 2.30 Å. i Bidentate co-ordination with no M–O interaction.32 flexible backbones, such as dppb, to adopt large ranges of bite angles. The narrow distribution of bite angles observed in mononuclear complexes are an indication that the P–M–P angle in monomeric complexes containing (P–P)M fragments with small P–M–P angles is predominantly determined by the P ? ? ? P distance defined by the ligand backbone.If a (P–P)M complex is formed the measured bite angle reflects the ligand preferences rather than the metal requirements. If metal and ligand requirements do not match di- or poly-nuclear complexes are formed preferentially. Ligands with bite angles above 1008 appear to be more flexible. Unfortunately, the number of crystal structures with these ligands is too small to allow a meaningful statistical treatment.The good agreement, however, between the ligand bite angles obtained from molecular modelling and the normalised P–M–P angles found in crystal structures is encouraging. Caution should be applied when ligands with potential donor atoms in the backbone are analysed. Weak interactions between a third ligand atom and the metal centre with M–X distances around 3 Å can lead to considerably larger P–M–P angles in crystal structures (even above 1508) of dppf 29 or Xantphos-type 30 ligands.An example may help to illustrate the contributions of the metal and the ligand preferences to the observed bite angle. A more detailed analysis of the angles found for 1,19-bis(diphenylphosphino) ferrocene (dppf) ligands yielded interesting results. The ferrocene backbone is generally assumed to be very flexible: the bite angle can be increased by either opening the angle between the two Cp planes or by increasing the torsion angle along the axis described by the two centroids of the Cp rings.31 Fig. 2 illustrates that, in most complexes, the bite angle is determined by the ligand preference, roughly 968. This corresponds to a nearly coplanar orientation of the cyclopentadienyl rings with a staggered conformation (“ideal” staggered torsion angle 368). If, however, the metal has a strong preference to form other angles the ligand will adapt to them. The angles of 102 and 1058 found for tetrahedral nickel complexes are roughly the average values between ligand (968) and metal (tetrahedral angle 1098) preferences.Complexes of CuI, AgI and AuI and other d10 metal ions have a strong tendency to form linear (L–M–L 1808) or trigonal planar (L–M–L 1208) geometries. The Walsh diagram (see below) indicates a sharp increase of energy with decreasing bite angles. Consequently, the bite angles are larger. Interestingly, the average M–P distance (2.34 Å, b2.315 = 111.658) is not significantly diVerent to that in other dppf complexes (2.32 Å, b2.315 = 99.338).This can be interpreted in terms of flexibility of the ferrocene backbone. A study of palladium–TCNE complexes with diVerent ligands based on the more rigid xanthene based backbone shows that elongation of the Pd–P bond is another way of compromising between the ligand and the metal preferred bite angle.32 Bite angles from computer modelling Computer modelled geometries can be used to estimate ligand bite angles.The obvious advantage is that no crystal structure is required. The calculations can even be performed before ligands are synthesized. If computer modelling is employed to design new ligands it is more important to calculate a correct trend rather than perfect geometries. Bite angles of all ligands in a series should thus be modelled with the same program and the same parameter set. Various molecular modelling packages using diVerent force fields and force field parameters have been used to calculate ligand bite angles.For the Xantphos ligand (Table 3), the structure calculated with force field methods was found to be closer to the crystal structure than that obtained1522 J. Chem. Soc., Dalton Trans., 1999, 1519–1529 with AM1 or PM3 semiempirical calculations.32 The force field parameters available for the ligand part of transition metal diphosphine complexes are suYciently good for our purposes. The parameterisation of a metal atom in a transition metal complex with several diVerent ligands bound to the central core is more diYcult.A closer examination of the force field parameters that have been used to calculate bite angles of diphosphine Fig. 2 Distribution of ligand bite angles vs. Cipso–centroid–centroid9– Cipso9 torsion for transition metal–1,19-bis(diphenylphosphino)ferrocene complexes found in the CSD. The torsion angles indicated are absolute values retrieved from the CSD.The distance between the centroids of the two cyclopentadienyl rings has been constrained to 3 < r < 3.5 Å to exclude values for h1-bound dppf ligands (average value in h2-bound dppf ligands 3.287, minimum 3.206, maximum 3.327, standard deviation 0.02 Å). In most complexes the dppf ligand assumes a staggered conformation with a bite angle around 968 and a torsion angle between 30 and 408; the torsion angle for the staggered conformation of ferrocene is 368. Interestingly, the angle between the two Cp rings decreases with increasing bite angle.ligands in rhodium complexes 15,25 shows that the rhodium atom is eVectively reduced to a “dummy metal atom”. The force constant for the P–M–P bond angle is reduced to 0, and a high energy constraint fixes the “Rh”–P bond length to 2.315 Å, a typical distance found in crystal structures. Despite the strongly simplified parameters, the agreement of the calculated ligand bite angles with average values from crystal structures is very good (Table 2).The metal atom cannot be eliminated completely. A “dummy metal atom” is necessary to pull the lone pairs of electrons and the substituents on the phosphorus atoms in the right direction and to simulate the constraints imposed on the ligand backbone by the metal atom in a complex, more precisely the eVect of two M–P bonds with a given length and orientation (Scheme 1). It may be useful to keep some points in mind when estimating ligand geometries using force field methods: many programs tend to overestimate attractive coplanar p-stacking interactions of aromatic groups and repulsive interactions of alkyl groups on the ligand.Ligand bite angles may thus be calculated as being too small in structures with coplanar phenyl groups on the two phosphorus atoms (e.g. dppe) and too large for ligands carrying bulky alkyl substituents. Owing to leverage eVects, small errors in the backbone geometry can lead to larger errors in the calculated ligand bite angle.A second parameter can be used to describe the rigidity of the ligand backbone: Casey defined a “flexibility range”,25 the range of bite angles a ligand can adopt if conformations with energies slightly above that of the minimised structure are considered. It can be estimated from a computed potential energy Table 3 Ligand bite angles calculated by molecular modelling, within parentheses the range accessible with DE calculated to be <3 kcal mol21 (flexibility range) Values calculated or cited in: a ref. 15; b ref. 14; c see Table 2. The original authors have used rhodium as a “dummy metal atom” (rRh–P = 2.315 Å) and the Sybyl (Tripos force field) a or Macromodel (Amber force field) b,c program packages to calculate the angles.J. Chem. Soc., Dalton Trans., 1999, 1519–1529 1523 diagram. The flexibility range has been defined as the range of bite angles accessible within 3 kcal mol21 of the minimum energy.Fig. 3 shows the energies of three ligands as a function of the bite angle. The values are calculated by minimising the energies of a series of structures with the P–M–P angle constrained to diVerent values. A steeper curve indicates a less flexible ligand. Interestingly, the distribution of bite angles in crystal structures is indeed broader for ligands with a larger calculated flexibility range (Table 1). 3 The metal preferred bite angle It is important to know the “preferred” bite angles of crucial intermediates in the catalytic cycle to decide which ligands to use for a catalytic reaction.A first approach can be based on complex geometries. In a rough approximation, the metal preferred bite angle for cis-co-ordinated bidentate ligands is 908 in a square planar or octahedral complex, 1098 in a tetrahedral complex and 1208 for the bis-equatorial co-ordination in a trigonal-bipyramidal complex. These geometries are frequently found for intermediates in reactions catalysed by diphosphine– transition metal complexes.The metal preferred bite angle changes during a catalytic reaction. For a more accurate indication of metal preferred bite angles, structures calculated by high-level computer modelling are useful. Most complexes modelled by ab initio or density functional methods have been extremely simplified in order to reduce the computation time necessary to optimise the structures. This is an advantage for this discussion.In “real” complexes the ligand positions are determined by a mixture of metal orbital requirements and inter-/intra-ligand steric interactions. Bulky ligands require space. In modelled structures, with protons representing the bulky groups, steric interactions are limited. The position of the PH3 ligands is consequently a very good indication of metal orbital preferences. An early (1978) and very good example to illustrate the value of computer modelled structures is the calculation of energies for Pt-group metal–diphosphine complexes as a function of the P–M–P angle.Thorn and HoVmann33 analysed the addition of hydrogen to ethylene catalysed by platinum diphosphine complexes (Scheme 2). Extended Hückel calculations showed a bite angle increasing from 958 in the olefin–hydride complex to 1108 Fig. 3 Calculated (Sybyl 6.4) flexibility of the DUPHOS (j), Xantphos (r) and BISBI (m) ligands. The flexibility range is determined by calculating the energy of structures minimised with constrained P–M–P angles.Scheme 2 The platinum-catalysed hydrogenation of ethylene calculated by Thorn and HoVmann.33 in the transition state during the insertion of the hydride ligand into the Pt–C bond. The implied conclusion is that ligands with bite angles around 1108 should stabilise intermediates between the h2-olefin complex and the h1-ethyl complex. Recent calculations on ab initio/density function levels indicate a bite angle around 1018 in the transition structure.34 Thirteen years after HoVmann’s original publication ethene– hydride and agostic ethyl complexes were indeed isolated: “The size of the chelating diphosphine ligand is shown to control the extent of transfer of a hydrogen atom from the b-carbon of the coordinated ethyl group to platinum with the smaller diphosphines favouring transfer to the metal.” A P–Pt–P angle of 104.88 was determined in the crystal structure of [Pt(h3-Hethylene)( dppp-Bu)]1 (Scheme 3).This complex is indeed an intermediate between an olefin–hydride complex and the alkyl complex, the product of a hydride insertion reaction.35 The relationship between the P–M–P bite angle and electronic properties of the metal centre has been investigated since by several authors. In transition metal complexes the symmetry of the molecular orbitals and the extent to which atomic orbitals contribute to them depends on the angle between the ligands.Walsh diagrams are a convenient way to visualise the energy of the molecular orbitals as a function of the P–M–P angle. A qualitative diagram calculated by Otsuka 36 (Fig. 4) for example helps to understand the preferred geometries in L–Pt0– L and L–PtII–L complexes with predominantly s-bonding ligands L. In d10 complexes with monodentate ligands five MOs are occupied. The sum of their energies is lowest if the L–Pt–L Fig. 4 Walsh diagram for PtL2 complexes with predominantly s-bond ligands L.Reprinted from ref. 36 with permission from Elsevier Science. Scheme 31524 J. Chem. Soc., Dalton Trans., 1999, 1519–1529 angle is close to 1808. If two electrons are removed, the dg*/2b1 orbital is empty. The d8 complexes thus prefer a geometry with 3a1 being the HOMO. If the L–Pt–L angle in a d10 complex is constrained to close to 908 the energy of the electrons in the non-bonding 2b1 orbital is very high and the reduction potential of the complex is raised considerably.The energy diagram explains why linear geometries are preferred by d10 and square-planar ones by d8 metal centres. The relation between the energy of the two electrons in the nonbonding orbital and the bite angle is nicely demonstrated when [PtCl2(P–P)] complexes are reduced with sodium amalgam in thf. If two (t-Bu)2P units linked by a (CH2)3 chain (P–M–P2.315 998) are employed a dimeric complex with a weak Pt–Pt bond is obtained. If the ligand bite angle is reduced by employing a (CH2)2 spacer (P–M–P2.315 878) “the enhanced reactivity of the [Pt(P–P)] species precludes formation of the dinuclear compound.” The dihydride complex isolated instead is probably the product of the reaction with thf.36 Peter Hofmann39 has used Walsh diagrams to explain the diVerent reactivity of platinum bis(phosphine) complexes {generated in situ from [Pt(PEt3)3] and [Pt(PMe3)2(PhCH]] CHPh)], respectively} towards hexakis(trifluoromethyl)benzene.The unit [Pt(PEt3)2] forms an h2 adduct,37 whereas [Pt(PMe3)2] inserts into the aromatic ring to form a seven-membered metallacyclic ring (Scheme 4).38 Hofmann explains the diVerent reactivities with a Walsh diagram very similar to the one shown in Fig. 4. The energy of the b2 orbital increases with decreasing P–M–P angle while that of the 3a1 orbital decreases at the same time. The smaller cone angle of PMe3 corresponds to a smaller P–Pt–P angle in the complex. According to Fig. 4, the energy of the p-symmetric b2 orbital should be higher, and the energy of the s-symmetric 3a1 orbital lower than for the PEt3 complex.More electrons will consequently move from the ligand C–C s bond to the empty 4a1 orbital, and the 2 electrons occupying the 2b1 orbital in the neutral platinum(0) complex move towards the ligand p* orbitals when the complex is formed. The overall eVect is a weaker ligand C]] C bond due to electrons “moving” from the s to the p* orbital.In the PMe3 complex this electron migration is strong enough to break the bond. The C6(CF3)6 ring is opened to form the observed metallacycle.39 Even though tridentate ligands are used in the following example, the conclusions are very similar to those discussed for diphosphine complexes. Dubois and co-workers 40 examined the acidity of the hydride ligand and the catalyst activity in [Pd(H)(P–P–P)] complexes [P–P–P = Ph2P(CH2)nPPh(CH2)m- PPh2, n,m = 2 or 3]. Extended Hückel calculations indicate that, in diphosphine complexes with small ligand bite angles, electron density is shifted onto the hydride ligand. With increasing ligand bite angles the hydride ligand becomes more acidic. The Scheme 4 results can, qualitatively, be understood with the same Walsh diagram (Fig. 4). Finally, the higher energy of the non-bonding electron pair in the resulting d10 complexes with smaller ligand bite angles explains the relationship between bite angles and the ease of oxidative addition 41 and reductive elimination 42 reactions of palladium diphosphine complexes (see below). 4 Spectroscopic and electrochemical changes related to bite angles The changes of energy levels associated with changes in ligand bite angles can also be measured directly.Changing energy levels should lead to changes in the spectroscopic properties of the metal complexes. Even though the qualitative discussion above does not allow quantitative predictions, consistent trends can be expected if a series of similar ligands is investigated.The absorption maxima in UV spectra of the p–p* transition of (dppm, dppe, dppp) platinum 7,8-benzoquinoline complexes are blue-shifted with increasing ligand bite angles. The emission lifetime measured at 77 K increases in the same order.43 In a study of nickel and palladium [M(P–P)2][BF4]2 complexes the UV/vis absorption maximum lmax and the half-wave potential E2� 1 (NiII/I, PdII,0) were found to increase with the ligand bite angle.44 These examples show that the eVect of the ligand bite angle on a metal centre can indeed be measured directly by spectroscopy.The most impressive examples are recent correlations of transition metal NMR spectra with ligand bite or with Tolman angles. Metal NMR shifts can be a sensitive probe to electronic changes on a metal centre.45 The 103Rh NMR shift of a series of [Rh(hfacac)(P–P)] (hfacac = hexafluoroacetylacetonate) complexes is related to the P–Rh–P bite angles in a direct linear correlation.46 The 103Rh signal is shifted downfield with increasing ligand bite angle.This is a direct proof for the electronic eVect of ligand bite angles. A correlation between metal NMR shifts and another ligand parameter has previously been observed for Tolman cone angles of mono-phosphine/ -phosphite ligands and 187Os/57Fe NMR shifts in h6-areneosmium or cyclopentadienyliron complexes respectively.47,48 5 Ligand bite angles and reaction energy profiles A ligand with a restrained ligand bite angle can be expected to change the energy profile of a reaction.During oxidative addition, insertion or reductive elimination reactions the coordination number of the metal centre and consequently the L–M–L angle changes. The angle of the transition state is likely between that of the reactant and that of the product complex. Enzymes often provide “pockets” which force the substrate into a transition state-like geometry.Even though most ligand backbones do not really change the substrate geometry considerably, they can force the catalyst into a transition state-like shape. If a ligand bite angle is constrained to a value close to that expected for the transition state of a rction the energy of the reactant complex is higher than that of an unrestrained complex. The energy of the transition state should be relatively unaVected and the activation energy for this step should be lower.The reaction step is accelerated and, if it is rate determining, the catalytic cycle runs more smoothly with a higher frequency. The reverse is true if the ligand bite angle is close to that of the reactant or product complex. The energy of the transition state will be higher, and the energy of the starting material is less aVected. Owing to the increased activation energy 49 the reaction is slower (Fig. 5). A good example is the nickel catalysed hydrocyanation of olefins (Scheme 5).With monophosphine ligands square-planar nickel dicyanides are formed (P–Ni–P ca. 908) and the catalyst deactivates. If diphosphine ligands with large ligand bite angles are used the catalytic pathway is favoured.17,18J. Chem. Soc., Dalton Trans., 1999, 1519–1529 1525 Fig. 5 The activation energy (dotted line) as a function of the L–M–L bite angle (qualitative picture). 6 Bite angle eVects in catalysis Hydroformylation The hydroformylation reaction is one of the most important applications of homogeneous catalysis in industry.From the first rhodium–phosphine catalysts found by Wilkinson and coworkers in the late sixties 50,51 it was only a short step to their application in industry (Scheme 6). Higher activities and selectivities obtained using rhodium catalysts compared to cobalt catalysts more than compensate the higher price of rhodium; cobalt catalysts are still widely used for the hydroformylation of > C4 olefins. The development of catalysts with even higher selectivities is one of the goals of current research in this area.While the thermodynamically favoured products are branched aldehydes, linear aldehydes are commercially more interesting. In 1987 Kodak Eastman patented a BISBI-based rhodium catalyst with a high selectivity towards linear aldehydes.52 To explain the selectivity, Casey and Whiteker 25 looked at bite angles of several ligands that form catalysts with a preference for linear aldehydes.They found a very good correlation between the ligand bite angle and the catalyst selectivity (Fig. 6).14 Ligand (natural) bite angles were the basis of the first series of ligands developed by molecular modelling in our group. The members of the Xantphos family have a rigid backbone in common, which keeps the two phosphorus donor atoms at a distance of typically around 4 Å, corresponding to ligand bite angles between 100 and 1308. The ligands have very similar Scheme 5 Simplified catalytic cycle of the nickel catalysed hydrocyanation reaction.electronic properties. Table 4 summarises the results of hydroformylation reactions with 1-octene as a substrate. The correlation between the ligand bite angles and the selectivity is good. The selectivity towards the linear aldehyde is even higher than that of the system using BISBI.15 The reason for the correlation between the ligand bite angle and catalytic selectivity is the subject of an ongoing discussion.The first hypothesis was that a bis-equatorial co-ordination of the diphosphine ligand in a trigonal-bipyramidal intermediate increases the selectivity towards the linear product. The ideal angle for such a co-ordination is 1208 (in a complex with the third ligand in the plane being similar to the other two). Mechanistic studies have so far concentrated on the hydride insertion step (see Scheme 6). If a 1-alkyl intermediate is formed the linear aldehyde will result, whereas the 2-alkyl intermediate leads to the branched product (Fig. 7). An explanation based purely on steric interactions was ruled Scheme 6 Simplified catalytic cycle of the rhodium catalysed hydroformylation reaction.1526 J. Chem. Soc., Dalton Trans., 1999, 1519–1529 Table 4 Hydroformylation of 1-octene at 40 8C: bn is the calculated ligand bite angle (flexibility range in parentheses), t.o.f. the turnover frequency (mol alkene mol Rh21 h21), hydrogenation products were not observed.Conditions: CO :H2 = 1 : 1, substrate :Rh = 674 : 1, [Rh] = 1.78 mM. Data from ref. 15. The large bite angle of DBFphos probably prevents the formation of a monomeric (P–P)Rh complex Linear : Product (%) Ligand DPEphos Sixantphos Thixantphos Xantphos DBFphos BISBI bn/8 102.2 (86–120) 108.7 (93–132) 109.4 (94–130) 111.7 (97–135) 131.1 (117–147) 122.6 (101–148) branched 10.5 : 1 35.0 : 1 47.6 : 1 57.1 : 1 3.4:1 58.2 : 1 Linear 91.3 96.3 97.0 98.3 76.1 95.5 Isomerisation 0 <1 10 1.6 2.9 t.o.f. 5 4.4 13.2 10 1.9 30 out because molecular mechanics energy calculations on “guessed” transition states for the hydride insertion step predict the wrong selectivities.53 Similarly, a purely electronic explanation is unlikely. Hydroformylation experiments with electronically modified dppe, T-BDCP and BISBI ligands gave no conclusive picture.54 Experiments with electronically modified Thixantphos ligands show that the ligand’s preference for a bisequatorial co-ordination increases with decreasing phosphine basicity, while the regioselectivity is almost unaVected.56 There are two preliminary conclusions.“The regioselectivity of hydroformylation is governed by a complex web of electronic and steric eVects that have so far defied unraveling”,54 and it may be possible that the regioselectivity is determined earlier in the cycle, namely the olefin addition to a four-co-ordinate [Rh(CO)H(P–P)] intermediate.55 Palladium catalysed cross coupling reactions Palladium catalysed cross coupling reactions are probably the most widespread examples for the application of homogeneous catalysis on a laboratory scale, and the industrial importance is growing rapidly.The broad range of coupling partners allows one to find conditions for almost any two reactants to be Fig. 6 Hydroformylation of 1-hexene with rhodium–diphosphine complexes: plot of % n-aldehyde vs. calculated ligand bite angle of diphosphine (BISBI, T-BDCP, DIOP, dppe from top to bottom).Horizontal bars indicate the range of bite angles accessible with <3 kcal additional calculated strain energy. Reprinted with permission from ref. 14. Copyright 1992 American Chemical Society. Fig. 7 The current discussion on the selectivity determining step in hydroformylation reactions concentrates on the alkene addition step, the [Rh(CO)(alkyl)(P–P)] intermediate and the transition state of the hydride insertion reaction leading to its formation (P–P = Xantphos, backbone and hydrogen atoms omitted).The 1-alkyl chain (right) can avoid steric interactions more easily than the 2-alkyl group. coupled. While most reactions are performed with palladium complexes bearing monodentate phosphine ligands e.g. [Pd- (PPh3)4], the use of bidentate ligands can enhance rate and/or selectivity. As an example, the eVect of diphosphine ligand bite angles (Table 5) on palladium catalysed cross coupling reaction of alkylmagnesium reagents with aryl halides has been studied by several authors. The reaction is assumed to begin with an oxidative addition of the aryl halide component to a zerovalent palladium–diphosphine species, followed by a transmetallation step that yields a [Pd(alkyl)(aryl)(P–P)] species.The cycle is closed by the reductive elimination of the alkylated aryl product (Scheme 7). Hayashi et al.12,13 found an increased reaction rate with increasing ligand bite angles, dppe (P–M–P 858) forming the slowest and least selective catalyst and dppf (P–M–P 968) the best.The selectivity decreases again if ligands with bite angles above 1028 are employed56 (Table 6). An aspect that has been neglected somewhat in the original papers is the strong dependence of the reaction rate on the ligand bite angle. In a first approach it is tempting to explain reaction rates with steric interactions between substrate and ligand. The larger the ligand bite angle the stronger is the substrate–ligand interaction (Scheme 8) and the slower the rate.However, the experiments show that the reverse is true. A closer examination of the reaction energy profile gives a likely explanation. Extended Hückel calculations 49 indicate a higher activation energy for the reductive elimination step if a diphosphine ligand bite angle is constrained to 908 compared to the Scheme 7 The catalytic cycle of a palladium catalysed Grignard cross coupling reaction.J.Chem. Soc., Dalton Trans., 1999, 1519–1529 1527 Table 5 Average ligand bite angles compiled from crystal structures (standard deviation in parentheses, see Table 1 for details) Table 6 Cross-coupling of 2-butylmagnesium chloride with bromobenzene in diethyl ether. Conditions: 0.04 mmol catalyst, 8 mmol Grignard reagent and 4 mmol bromobenzene in 20 ml ether, T = 20 8C Conversion b Product (%) Ligand dppec dpppc dppbc dppf DPEphos Sixantphos Thixantphos Xantphos Bite angle/8 85 (P–M–P2.315) 91 (P–M–P2.315) 98 (P–M–P2.315) 96 (P–M–P2.315) 102.2 (bn) 108.7 (bn) 109.4 (bn) 111.7 (bn) t.o.f.a n.d.n.d. n.d. 79 181 36 24 24 t/h 48 24 822 16 16 16 (%) 4 67 98 100 100 58.8 36.5 23.6 Linear 0 69 51 95 98 67 51 41 Branched 0 31 25 21 17 17 19 Biphenyl n.d. n.d. n.d. 31 16 32 40 n.d. = Not determined; ligand bite angles from Table 2. a Initial turnover frequency [mol (mol Pd)21 h21], determined after 5 min of reaction time. b Conversions based on bromobenzene. c Results from ref. 13. activation energy in a system with an unconstrained bite angle (Fig. 5). Portnoy and Milstein 41 observed that the oxidative addition of aryl chlorides to (P–P)Pd0 species is faster with decreasing ligand bite angles, and Brown and Guiry42 concluded that the rate of reductive elimination in palladium catalysed Grignard cross coupling reactions increases with the bite angle. Earlier, Yamamoto and co-workers 57 had found that ethane was Scheme 8 The interaction between the substrate (represented as dark rectangle) co-ordinated to a metal centre and the (chiral) ligand’s substituents (circles) increases with the L–M–L bite angle.eliminated from [NiMe2(dppp)] 46 times faster than from [NiMe2(dppe)] (ligand bite angles of 91 and 858, respectively. Using DIOP (988), the elimination of RCN from [Pd(R)(CN)- (P–P)] (R = CH2SiMe3) is 10000 times faster than that using the dppe ligand (858).58 These results can be understood with the Walsh diagram (Fig. 4) discussed above. The oxidative addition is supported by ligands having smaller bite angles, because they increase the electron density on the (P–P)Pd0 metal centre. Larger bite angles which go along with decreased electron density on the metal make the reductive elimination easier. A correlation between increasing diphosphine ligand bite angles and rate or selectivity has also been observed in cobaltor nickel-catalysed cross coupling reactions, for example in the coupling of arylboronic acids with aryl halides 59 or in the reaction of catechol–borane with various dienes.60 An important point to keep in mind when discussing rates and selectivities in Pd-catalysed reactions is the catalyst stability.Palladium(0) species tend to form metallic palladium (palladium black), especially if reactions are performed at higher temperatures. This decreases the activity of the catalyst if metallic palladium does not catalyse the reaction and the1528 J.Chem. Soc., Dalton Trans., 1999, 1519–1529 selectivity if it does. The catalyst concentration might also be diminished by equilibria involving the starting material used,6 and solvent eVects, albeit poorly understood, have a strong influence in many catalytic reactions. The overall eVect of a single parameter such as the ligand bite angle may become ambiguous if the catalyst concentration and its activity are both a function of this parameter.The bite angle eVect in palladium catalysed reactions of aryl halides with amines is a good example.8 The ambiguous overall picture is probably due to a combination of elementary steps aVected diVerently by changes of the ligand bite angle. Enantioselective reactions From a mechanistic point of view, reactions involving chiral centres are very similar to those not involving chirality. Factors determining rate, selectivity and the energy profile should be the same for both.A reaction is enantioselective if interactions between ligand and substrate during the catalytic reaction favour the formation of one of the product enantiomers. Interactions require proximity. Parts of the ligand have to be close to the active site of the catalyst. In most bidentate ligands the orientation of the interacting parts of the ligand with regard to the substrate co-ordinated to the metal atom is related to the ligand bite angle (Scheme 8). Trost et al.61 have designed a modular ligand system for allylic alkylation reactions: two 2-(diphenylphosphino)benzoic acid groups are attached to a chiral backbone.Palladium complexes of these ligands are very selective catalysts for the enantioselective substitution of small allylic substrates. The chiral pocket formed upon complexation to a metal centre can be fine tuned by varying the chiral diol/diamine used to form the backbone (Scheme 9). It was postulated that the opening of the bite angle is necessary for high chiral recognition 61 in allylic alkylation reactions.For one of the ligands the P–Pd–P angle in a [Pd(C3H5)(P–P)] complex was determined to be 110.58.62 From studying Corey–Pauling–Koltun (CPK) models, Trost and Murphy 20 had earlier concluded that “a larger ring of a chelating bidentate ligand leads to greater embracing of the allyl fragment by the metal template and consequently higher asymmetric induction.” Similarly, the regioselectivity in the (non-asymmetric) alkylation of 2-hexenyl acetate increases with an increasing ligand bite angle.63,64 Hayashi et al.65 compared the selectivity of chiral ferrocene and ruthenocene based ligands BPPFA (P–M–P angle 98.798) and BPPRA (P–M–P angle 100.478) in asymmetric silylation and cyclisation reactions.Despite the small diVerence, the enantioselectivities obtained with BPPRA were up to 40% higher. In a series of spirobis(oxazoline) ligands the enantioselectivity in copper-catalysed Diels–Alder reactions was shown to increase with the (calculated) ligand bite angle.19 There are many asymmetric reactions where the bite angle eVect might be important.Asymmetric hydrogenation is an example for a reaction of which the ligand bite angle has not been examined explicitly. Burk et al.66–69 have prepared a series Scheme 9 A modular system to build ligands for asymmetric allylic alkylation reactions. of ligands with two phospholane units carrying chiral groups attached to diVerent backbones. Even though the substrates tested are not exactly the same, the selectivities observed clearly decrease with increasing ligand bite angle (Table 7, Scheme 10).It would certainly be wrong to generalise that ligands with bite angles of 858 form better hydrogenation catalysts than those with larger angles. BINAP and DIOP for example, with ligand bite angles of 92 and 988, can be as selective as DUPHOS. To develop new phospholane based ligands for the hydrogenation of olefins ligands with bite angles between 85 and 958 would probably be the best choice. 7 Conclusion The ligand bite angle is a property that can be calculated from a series of crystal structures or from molecular modelling. Systematic searches in the Cambridge Crystallographic Database show that the P–M–P angles concentrate in narrow ranges for most transition metal complexes containing (P–P)M fragments. The average values calculated from crystal structures are in good agreement with ligand bite angles calculated with molecular mechanics, even if simplified parameters are employed for a “dummy metal atom” connecting the two phosphorus atoms.The ligand bite angle correlates with various spectroscopic properties of metal–diphosphine complexes and with the regio- or stereo-selectivity in a variety of catalytic reactions. In rhodium catalysed hydroformylation reactions, ligands with bite angles >1008 favour the formation of linear aldehydes.Scheme 10 Bisphospholane ligands with diVerent bite angles. Table 7 The eVect of the ligand bite angle on the enantioselectivity in rhodium catalysed homogenous hydrogenation reactions e.e. Ligand DUPHOS BPE BPP BPFc DuThixantphos P–M–Pa/8 83 85 98 96 104 R = Ph 99 91.4 60 30 Me 98 85 H 64 Ref. 68 68 66 69 72 a For consistency reasons, all bite angles are those of non-chiral analogs (PPh2, see Scheme 9). Measured bite angles for DUPHOS and BPE in [Rh(COD)(P–P)]SbF6 complexes are 84.72 68 and 83.38,67 respectively [r(Rh–Pav) 2.26 Å for both].J.Chem. Soc., Dalton Trans., 1999, 1519–1529 1529 In many palladium catalysed reactions ligands such as dppf or dppp with bite angles around 1008 give the best results in terms of activity and selectivity. In the case of palladium this can be explained by transition states between (four-co-ordinate) palladium(II) species with P–Pd–P angles around 908 and twoor three-co-ordinate palladium(0) species with P–Pd–P angles of probably 120–1808.Although the hypothesis will have to be substantiated, the examples discussed in this article indicate that there is an optimum bite angle for many catalytic systems. A large number of (chiral) ligand backbones and of (chiral) PR2 groups has been developed over the last decades. They can be combined to tailor an almost indefinite number of ligands. The ligand bite angle concept can help to determine the right one. Sharpless has been quoted “If you give me high rates of turnover, I can probably give you the selectivity later .. .”.70 A better understanding of the bite angle eVect might be the key to both. 8 References 1 C. A. Tolman, Chem. Rev., 1977, 77, 313. 2 Science Citation Index Web of Science, www.isinet.com. 3 A. Fernandez, C. Reyes, M. R. Wilson, D. C. Woska, A. Prock and M. P. Giering, Organometallics, 1997, 16, 342. 4 S. Joerg, R. S. Drago and J. Sales, Organometallics, 1998, 17, 589. 5 E. C.Alyea and S. Q. Song, Comments Inorg. Chem., 1996, 18, 189. 6 C. Amatore, G. Broeker, A. Jutand and F. Khalil, J. Am. Chem. Soc., 1997, 119, 5176. 7 C. Amatore, A. Jutand and G. Meyer, Inorg. Chim. Acta, 1998, 273, 76. 8 B. C. Hamann and J. F. Hartwig, J. Am. Chem. Soc., 1998, 120, 3694. 9 E. M. Vogl, J. Bruckmann, C. Krüger and M. W. Haenel, J. Organomet. Chem., 1996, 520, 249. 10 T. Hayashi, Y. Kawabata, T. Isoyama and I. Ogata, Bull. Chem. Soc. Jpn., 1981, 54, 3438. 11 Y.Kawabata, T. Hayashi and I. Ogata, J. Chem. Soc., Chem. Commun., 1979, 462. 12 T. Hayashi, M. Konishi and M. Kumada, Tetrahedron Lett., 1979, 21, 1871. 13 T. Hayashi, M. Konishi, Y. Kobori, M. Kumada, T. Higuchi and K. Hirotsu, J. Am. Chem. Soc., 1984, 106, 158. 14 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 and 10680. 15 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. 16 K. Yamamoto, S. Momose, M. Funahashi, S. Ebata, H. Ohmura, H. Komatsu and M. Miyazawa, Chem. Lett., 1994, 189. 17 W. Goertz, P. C. J. Kamer, P. W. N. M. van Leeuwen and D. Vogt, Chem. Commun., 1997, 1521. 18 M. Kranenburg, P. C. J. Kamer, P. W. N. M. van Leeuwen, D. Vogt and W. Keim, J. Chem. Soc., Chem. Commun., 1995, 2177. 19 I. W. Davies, L. Gerena, L. Castonguay, C. H. Senanayake, R. D. Larsen, T. R.Verhoeven and P. J. Reider, Chem. Commun., 1996, 1753. 20 B. M. Trost and D. J. Murphy, Organometallics, 1985, 4, 1143. 21 P. Dierkes, S. Ramdeehul, L. Barloy, A. De Cian, J. Fischer, P. C. J. Kamer, P. W. N. M. van Leeuwen and J. A. Osborn, Angew. Chem., 1998, 110, 3299; Angew. Chem., Int. Ed., 1998, 37, 3116. 22 S. Ramdeehul, P. Dierkes, R. Aguado, P. C. J. Kamer, P. W. N. M. van Leeuwen and J. A. Osborn, Angew. Chem., 1998, 110, 3302; Angew. Chem., Int. Ed., 1998, 37, 3118. 23 T. Yoshida, K. Tatsumi and S. Otsuka, Pure Appl. Chem., 1980, 52, 713. 24 P. Hofmann, L. A. Perezmoya, O. Steigelmann and J. Riede, Organometallics, 1992, 11, 1167. 25 C. P. Casey and G. T. Whiteker, Isr. J. Chem., 1990, 30, 299. 26 F. H. Allen, O. Kennard, D. G. Watson, L. Brammer, A. G. Orpen and R. Taylor, J. Chem. Soc., Perkin Trans. 2, 1987, S1. 27 A. G. Orpen, L. Brammer, F. H. Allen, O. Kennard, D. G. Watson and R. Taylor, J. Chem. Soc., Dalton Trans., 1989, S1. 28 T. E. Müller and D.M. P. Mingos, Transition Met. Chem., 1995, 20, 533. 29 M. Sato, H. Shigeta, M. Sekino and S. Akasori, J. Organomet. Chem., 1993, 458, 199. 30 M. D. K. Boele and P. W. N. M van Leeuwen, unpublished work. 31 A. G. Avent, R. B. Bedford, P. A. Chaloner, S. Z. Dewa and P. B. Hitchcock, J. Chem. Soc., Dalton Trans., 1996, 4633. 32 M. Kranenburg, J. G. P. Delis, P. C. J. Kamer, P. W. N. M. van Leeuwen, K. Vrieze, N. Veldman, A. L. Spek, K. Goubitz and J. Fraanje, J. Chem.Soc., Dalton Trans., 1997, 1839. 33 D. L. Thorn and R. HoVmann, J. Am. Chem. Soc., 1978, 100, 2079. 34 B. B. Coussens, F. Buda, H. Oevering and R. J. Meier, Organometallics, 1998, 17, 795. 35 L. Mole, J. L. Spencer, N. Carr and A. G. Orpen, Organometallics, 1991, 10, 49. 36 S. Otsuka, J. Organomet. Chem., 1980, 200, 191. 37 J. Browning, M. Green, A. Laguna, L. E. Smart, J. L. Spencer and F. G. A. Stone, J. Chem. Soc., Chem. Commun., 1975, 723. 38 J. Browning, M. Green, B.R. Penfold, J. L. Spencer and F. G. A. Stone, J. Chem. Soc., Chem. Commun., 1973, 31. 39 P. Hofmann, H. Heiß and G. Müller, Z. Naturforsch., B: Chem. Sci., 1987, 42, 395. 40 S. A. Wander, A. Miedaner, B. C. Noll, R. M. Barkley and D. L. Dubois, Organometallics, 1996, 15, 3360. 41 M. Portnoy and D. Milstein, Organometallics, 1993, 12, 1665. 42 J. M. Brown and P. J. Guiry, Inorg. Chim. Acta, 1994, 220, 249. 43 J. De Priest, G. Y. Zheng, C. Woods, D. P. Rillema, N. A. Mikirova and M. E. Zandler, Inorg. Chim. Acta, 1997, 264, 287. 44 A. Miedaner, R. C. Haltiwanger and D. L. Dubois, Inorg. Chem., 1991, 30, 417. 45 S. Q. Song and E. C. Alyea, Comments Inorg. Chem., 1996, 18, 145. 46 K. Angermund, W. Baumann, E. Dinjus, R. Fornika, H. Görls, M. Kessler, C. Krüger, W. Leitner and F. Lutz, Chem. Eur. J., 1997, 3, 755. 47 A. G. Bell, W. Kozminski, A. Linden and W. von Philipsborn, Organometallics, 1996, 15, 3124. 48 V. Tedesco and W. von Philipsborn, Magn. Reson. Chem., 1996, 34, 373. 49 M. J. Calhorda, J. M. Brown and N. A. Cooley, Organometallics, 1991, 10, 1431. 50 C. K. Brown and G. Wilkinson, J. Chem. Soc. A, 1970, 2753. 51 D. Evans, J. A. Osborn and G. Wilkinson, J. Chem. Soc. A, 1968, 3133. 52 T. J. Devon, G. W. Phillips, T. A. Puckette, J. L. Stavinoha and J. J. Vanderbilt, World Patent, W087/07600, 1987. 53 C. P. Casey and L. M. Petrovich, J. Am. Chem. Soc., 1995, 117, 6007. 54 C. P. Casey, E. L. Paulsen, E. W. Beuttenmueller, B. R. Proft, L. M. Petrovich, B. A. Matter and D. R. Powell, J. Am. Chem. Soc., 1997, 119, 11817. 55 L. A. van der Veen, M. D. K. Boele, F. Bregman, P. C. J. Kamer, P. W. N. M. van Leeuwen, K. Goubitz, J. Fraanje, H. Schenk and C. Bo, J. Am. Chem. Soc., 1998, 120, 11616. 56 M. Kranenburg, P. C. J. Kamer and P. W. N. M. van Leeuwen, Eur. J. Inorg. Chem., 1998, 155. 57 T. Kohara, T. Yamamoto and A. Yamamoto, J. Organomet. Chem., 1980, 192, 265. 58 J. E. Marcone and K. G. Moloy, J. Am. Chem. Soc., 1998, 120, 8527. 59 S. Saito, S. Ohtani and N. Miyaura, J. Org. Chem., 1997, 62, 8024. 60 M. Zaidlewicz and J. Meller, Tetrahedron Lett., 1997, 38, 7279. 61 B. M. Trost, D. L. van Vranken and C. Bingel, J. Am. Chem. Soc., 1992, 114, 9327. 62 B. M. Trost, B. Breit, S. Peukert, J. Zambrano and J. W. Ziller, Angew. Chem., Int. Ed. Engl., 1995, 34, 2386. 63 M. Kranenburg, P. C. J. Kamer and P. W. N. M. van Leeuwen, Eur. J. Inorg. Chem., 1998, 25. 64 R. J. van Haaren, H. Oevering, P. W. N. M. van Leeuwen, P. C. J. Kamer, J. N. H. Reek, G. P. F. van Strijdonck and B. B. Coussens, submitted for publication. 65 T. Hayashi, A. Ohno, S. J. Lu, Y. Matsumoto, E. Fukuyo and K. Yanagi, J. Am. Chem. Soc., 1994, 116, 4221. 66 M. J. Burk, J. E. Feaster and R. L. Harlow, Tetrahedron: Asymmetry, 1991, 2, 569. 67 M. J. Burk, J. E. Feaster and R. L. Harlow, Organometallics, 1990, 9, 2653. 68 M. J. Burk, J. E. Feaster, W. A. Nugent and R. L. Harlow, J. Am. Chem. Soc., 1993, 115, 10125. 69 M. J. Burk and M. F. Gross, Tetrahedron Lett., 1994, 35, 9363. 70 S. Borman, Chem. Eng., 1996, 4th November, 37. 71 T. Hayashi, M. Kumada, T. Higuchi and K. Hirotsu, J. Organomet. Chem., 1987, 334, 195. 72 P. Dierkes, U. Nettekoven and P. W. N. M. van Leeuwen, unpublished work. Paper 8/07799A

ISSN:1477-9226    

DOI:10.1039/a807799a

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON PERSPECTIVE J. Chem. Soc., Dalton Trans., 1998, Pages 1525–1531 1525 Exploring the interactions of d-block elements with boron. A case for electronically unsaturated metallaborane clusters Thomas P. Fehlner Department of Chemistry and Biochemistry, University of Notre Dame, Notre Dame, IN 46556, USA The ability to synthesize metallaboranes of fixed cluster shape with varying numbers of metal atoms and metal identities highlights the unique ways in which the metals diVer from their maingroup counterparts in a cluster environment.Thus we have found that electronic unsaturation introduced by the use of early transition metals is expressed in an intriguing and novel manner in a metallaborane cluster. 1 Introduction Contemporary chemistry consists of a set of sub-disciplines each of which attracts scientists who identify with the particular section of nature subsumed. But it is one of the fascinating aspects of chemistry, as well as most scientific endeavors, that nature will not be constrained by a medieval approach to scientific farming that springs naturally from our limited intellectual capacities.Some of the most exciting developments arise from the recognition that an idea of one area fits very nicely with an idea or fact of another. This being the case, it is unfortunate that the most successful explanations of an aspect of chemistry often create barriers that hinder further After graduating from Siena College in 1959, Thomas P.Fehlner obtained his Ph.D. in physical chemistry in 1963 at the Johns Hopkins University and was a research associate the following year at the same institution. He joined the Department of Chemistry at the University of Notre Dame in 1964 and presently holds the Grace-Rupley chair of chemistry. His research has ranged from the characterization of unstable species by mass and photoelectron spectroscopies to the syntheses of new cluster compounds and thin films.He has held DAAD, Guggenheim and Japan Society for the Promotion of Science Fellowships and received an award for Distinguished Achievements in Boron Science in 1990. development. These are sometimes so imposing that when they are broken by paradigm defying chemists the expression of surprise (or disbelief) is extraordinary, e.g., Werner’s six-coordinate octahedral metal complexes, Lipscomb’s three-center two-electron bonds applied to ‘electron deficient’ boranes, Miller’s and Paulson’s syntheses of ferrocene and the analysis of its structure and bonding by Fischer, Wilkinson and Woodward, and Bartlett’s synthesis of a compound of xenon. Forays into non-traditional areas force chemists to forge intellectual bridges and stimulate conceptual development. 2 The Chemistry of Clusters Cluster chemistry is one such area. It is an area that overlays a substantial fraction of contemporary chemistry: organometallic, main-group inorganic, co-ordination chemistry, physical chemistry and solid-state chemistry.Cluster chemistry ranges from p-block to d-block element chemistry (boranes, polynuclear metal carbonyls), from early transition metals to beyond the late transition metals (zirconium clusters to gold clusters), from solution to solid-state chemistry (polynuclear Group 14 anions, Zintl phases), and from organometallic synthetic techniques to those of chemical physics or solid-state chemistry (chiral tetrahedral MM9M0C clusters by fragment substitution on M3C clusters, laser evaporated clusters, clusters excised from extended solid-state structures).It is a sign of the maturity of the area that many monographs 1–12 and review articles, too numerous to list here, address aspects of the topic. Another sign of maturity is the demonstrated success of the cluster electron counting rules initially presented in a usable form 25 years ago by Wade and Mingos.13–21 Subsequent development and modification allow a simple electron counting connection between stoichiometry and structure.This provides the same practical guideline to the working cluster chemist that the 8, 18 electron rules continue to provide the main group and organometallic chemist. The overlapping ideas of three (and four)-connect clusters, spherical and non-spherical clusters with intrinsic delocalized bonding, the isolobal principle, and the cluster fusion principle rationalize and interconnect the vast majority of main-group and transition metal clusters known and, at the same time, suggest the existence of an even larger number of possible clusters.9 That is, these rules imply the existence of (a) sets of compounds of a given structure in which isolobal main-group and metal fragments are systematically varied, (b) a set of positional isomers for a given compound stoichiometry, (c) isomeric forms which are described by the capping principle as well as the fusion ideas of Mingos.In addition (d) the rules allow one to speculate on the possible formation of cluster types for which no exemplars exist at present.Points (a) and (d) concern us here. Given the connection between metal clusters and main-group clusters, another class of compounds, those containing varied mixes of both maingroup and transition-metal fragments, should exist.22 The electron counting rule for delocalized clusters (an n atom closed1526 J. Chem. Soc., Dalton Trans., 1998, Pages 1525–1531 Fig. 1 A comparison of the orbital properties of the BH and CpCo fragments. The yellow box contains the valence orbitals not used in atom– ancillary ligand bonding. The blue and red orbitals within each box represent the s and p symmetry frontier orbitals, respectively BH antibonding CpCo antibonding CpCo bonding BH bonding cluster will possess 2n 1 2 cluster bonding electrons) suggests that main group–transition-metal clusters will be most abundant for fragments that possess three frontier orbitals containing two electrons.Indeed, the BH fragment is a three orbital– two electron fragment (Fig. 1) and it was the structure determinations of polyborane clusters that gave us the three-center bond,23 the perceptive analysis of geometry that gave us families of cluster shapes 24,25 and, ultimately, the electron counting rules mentioned above. The incorporation of metal fragments into a polyborane fragment generates metallaboranes and these species constitute a substantial class of compounds.7,26–31 Although cluster charge and shape, number of bridging hydrogens, and metal ancillary ligands allow considerable flexibility in the metal identity, the number and types of metallaboranes containing Group 8 metals [M(CO)3 fragments] and Group 9 metals [CpM fragments, Cp = h5-C5H5] far exceed those for all other fragments.Although one cannot exclude bias from limited synthetic trials, the fact is that both of these fragments are three orbital–two electron fragments isolobal with BH (Fig. 1). Although the frontier orbital properties of a main-group fragment are well defined, the five additional metal-based orbitals of a transition-metal fragment give it a flexible set of frontier orbitals with properties that vary depending on metal nuclear charge as well as ancillary ligand numbers, types and positions. Consequently, the bonding behavior exhibited by a given metal fragment type does not always correspond to that expected from a simple series of fragments.For example, it is well known that the fragment CpFe(CO)2 mimics CH3 (one orbital–one electron). From this one correctly concludes that CpMn(CO)2 mimics BH3 (one orbital–zero electron). However, in many instances CpMn(CO)2 is better represented as CH2 (two orbital–two electron).32 The isoelectronic connection is there (BH3 vs. CH2) but the consequences for bonding are different. The set of three filled orbitals that are low-lying and non-bonding for the iron fragment (as with CpCo in Fig. 1) increasingly participate in bonding interactions in going to Mn. It is this added variability relative to main-group fragments that makes metal clusters so fascinating albeit frustrating at times. For example, as shown in Fig. 2 for six-atom clusters, one can find examples of diVerent cluster shapes for the same electron count as well as the same shape for diVerent electron counts.8 Understanding the possibilities for a given shape or electron count is one thing.Controlling the outcome of a synthesis is quite a diVerent thing. But this is what the chemist requires in order to design syntheses of species with desired structure and properties. 3 Attractive Features of Metallaboranes One motivation for our research originates in the idea that the synthesis of metallaboranes provides a significant simplification of the metal cluster problem. The step-wise incorporation of borane fragments in place of metal fragments in a metal cluster has several beneficial eVects.First of all, the reaction chemistry is changed. As EBB > EMM, the potential energy surface associated with a metallaborane cluster core will have lower barriers than that of a polyborane cluster but higher barriers than that of a metal cluster {C2B4H6 undergoes skeletal rearrangement only at high temperatures whereas [Rh2Fe4(CO)16B]2 rearranges at room temperature 33}. This permits kinetic control to be achieved at convenient temperatures. It also suggests that longer sequences of reactions will be accessed so that intermediate products can be isolated.A set of compounds empirically connected by simple reactions is requisite for a basic understanding of the reaction chemistry. Second, we reason that the incorporation of a number of boron fragments into a metal cluster bonding network limits the behavior of the metal fragment. In a sense, the less flexible bonding capabilities of the borane fragment relative to a metal fragment can act as a cluster enforcer.That is, the greater the number of boron atoms, the greater the tendency for ‘normal’ cluster behavior. Then, by decreasing boron content for a given cluster shape, the more flexible metal bonding capabilities are Fig. 2 A comparison of six-atom clusters having (top row across) the same number of cluster valence electrons (cve) and diVerent core shapes, and (first and third columns) the same core shape and diVerent numbers of cluster valence electrons [Os6(CO)18]2-, 86 cve H2Os6(CO)18, 86 cve [Pt6(CO)12]2-, 86 cve Ni6Cp6, 90 cve Re6(CO)18(PMe)3, 90 cveJ. Chem.Soc., Dalton Trans., 1998, Pages 1525–1531 1527 given scope and the unique characteristics imparted by a (ML)n fragment, n = 1, 2, 3, . . ., if any, will be expressed. Third, an additional interesting situation is created by varying the identity of the metal for a specified cluster size, structure and metal/boron ratio.The borane cluster orbitals for a given fragment are of fixed energy whereas the metal orbital energies vary depending on metal and ligands. As we scan the metal orbital energies through the boron frontier orbital energies different sets of metal orbitals will match up and it is not at all clear a priori how the block of available metal orbitals will be partitioned. In essence, the borane fragment fixes the cluster order and we can then examine the perturbation of structure and reactivity as the number and types of metal fragments are separately varied.The last thought kindled an interest in the metals lying to the left of iron. Could we frustrate the cluster counting rules by blurring the orbital separation on the metal fragment implicit in the isolobal analogy and demanded by the borane fragment? In particular we wondered if a metal like Cr, which often tends to form paramagnetic ‘electron deficient’ organometallic compounds, 34 would likewise cause a breakdown in the cluster paradigm.Of course, other possible cluster responses to an insuYciency of electrons are possible. Clusters can reduce required electron counts by forming capped or fused clusters;20 a structural response that can be restricted by working with small clusters. Metal species can respond to a reduced number of valence electrons by forming multiple bonds;35 a structural response that would be evident in a two-metal cluster system if present.The objective of this contribution is to compare a pair of related metal fragments in a single cluster environment. Selected metallaboranes from Groups 5 and 6, when compared to closely related metallaboranes from Group 9, reveal unexpected structural and bonding eVects. We ascribe this unusual behavior to the presence of electronic unsaturation. 4 Synthesis The ideas described above cannot be tested unless one has the ability of making specific compounds which are related in the desired fashion.Many metallaboranes exist and nicely illustrate the first two points discussed in the preceding section. However, as already noted, nearly all the metallaboranes with more than one metal fragment contain metals from either Group 8 or 9. Most of these compounds are less reactive than their isolobal pure borane counterparts and separation of complex mixtures is readily accomplished. This fact undoubtedly enhanced the number of compounds characterized.With some notable exceptions,36 these compounds follow and nicely illustrate the consequences of the cluster electron counting rules combined with the isolobal principle. After shifting research emphasis from more physical pursuits to synthesis in the mid-1970’s, my group continually sought synthetic approaches that would be selective. We found that in instances where both borane and metal fragment source were comparably reactive, good yields of metallaboranes resulted. For this reason, monoboranes have been the boron reagents of choice in that the problems associated with activation of an intrinsically stable polyborane are avoided.Thus, we assemble our polyborane fragment on a metal fragment reminiscent of the manner that macrocyclic ligands are constructed. The activation of the metal fragment has been the sticking point. Our slow progress in the synthesis of metallaboranes from monoborane and various organometallic species can be illustrated by three examples.Reactions of iron formyl anions gave highly complex mixtures of hydrocarbyl clusters and ferraboranes all in very low yields.37 Reactions of cobalt phosphine complexes gave only cobaltaboranes but again mixtures and modest to low yields were encountered.38 The use of ‘lightly coordinated’ iron carbonyl led to high yields of ferraboranes with selectivity largely controlled by initial stoichiometry.39 All these examples reflect the diYculties we experienced in finding keys to unlock the metal fragment in the presence of monoboranes.As metallaboranes formed from the early transition metals probably would not be separable via, e.g. chromatographic techniques, selectivity became the most important issue in the chemistry which is the focus of this essay. The report by Ting and Messerle of the preparation of [{Cp*Ta}B2H6]2 from [{Cp*Ta}Cl2]2 and [BH4]2, (Cp* = h5- C5Me5) 40 reinforced by the results of Leach and co-workers on the apparently complex reaction of Mo and W monocyclopentadienyl hydrides and halides with tetrahydroborate,41 led us to initiate a more general investigation of the reactions of monocyclopentadienylmetal halide oligomers with monoboranes.Thus far, this approach has yielded new metallaboranes of Co,42 Rh,43 Cr,44 Mo,45 W46 and Ta47 and the generally high yields and clean chemistry have provided opportunities to study the reaction chemistry of these compounds. Both borane and tetrahydroborate yield new compounds and the similarities and diVerences between the two reagents reflect their diVering reduction and co-ordination properties as well as the properties of the Cp*M fragments.A key feature of the chemistry associated with neutral borane, e.g., BH3?THF, is its facile reaction with the metal halide to produce metallaboranes accompanied by the release, in all except one case of BH2Cl [equation (1)].42 Reaction condi- [Cp*CoCl]2 1 5 BH3?THF æÆ 2,4-{Cp*Co}2B3H7 1 2 BH2Cl (1) tions are mild and in most cases a single metallaborane product is formed in very good yield.The structure of 2,4-{Cp*Co}2- B3H7 is shown in Fig. 3. Note that borane serves a dual role: it removes the Cl, thereby activating the metal fragment, and provides the borane units for the B3H7 fragment. Occasionally it serves a third role in the prereduction of a monocyclopentadienylmetal halide to the lower oxidation state that yields the metallaborane. The last is not a necessary feature of the chemistry but is a synthetic convenience when the desired monocyclopentadienylmetal halide is only obtainable from a precursor in a higher oxidation state.Tetrahydroborate first acts as a pseudo-halide displacing the halide from the metal. The first formed metal tetrahydroborate (isolated in the case of Cr44) then converts into metallaboranes by loss of H2 and the formation of B]B bonds,40 e.g., equation (2),48 or by more complex processes.Fig. 3 The molecular structures of nido-B5H9, nido-2-{CpCo}B4H8, and nido-2,4-{Cp*Co}2B3H7 B5H9 2-{CpCo}B4H8 2,4-{Cp*Co}2B3H71528 J. Chem. Soc., Dalton Trans., 1998, Pages 1525–1531 [Cp*MoCl2]2 1 2[BH4]2 æÆ {Cp*MoCl}2(B2H6) 1 H2 1 2 Cl2 (2) Most of the metallaboranes isolated to date by this approach obey the cluster electron counting rules. These compounds constitute interesting new examples of metallaboranes but the focus of this essay is on a small set of compounds that stretch the cluster electron counting rules. Thus, the following sections feature selected compounds formally derived from five- or sixatom borane frameworks by subrogation of one or two BH fragments by metal fragments. 5 Unsaturated Clusters 2-{CpCo}B4H8 vs. 2-{Cp*TaCl2}B4H8 The cobaltaborane, 2-{CpCo}B4H8 (Fig. 3), one of many precedent-setting metallaboranes from Grimes’ laboratory,49 has a structure that conforms to expectations based on the isolobal analogy between BH and CpCo fragments.50 As illustrated in Fig. 1, of the six valence energy metal orbitals containing eight electrons which remain after forming the metal–ligand bonds, three containing two electrons serve in cluster bonding and three containing six electrons are cluster non-bonding. The similarity of the geometry of the B4H8 fragment relative to the same fragment within B5H9 itself (Fig. 3), as well as the positioning of the bridging hydrogens, provides structural corroboration of the isolobal analogy.In a geometric sense the tantalaborane, 2-{Cp*TaCl2}B4H8 (Fig. 4) is the partner of the cobaltaborane.47 That is, if one simply replaces the 14-electron CpCo fragment with the 12- electron Cp*TaCl2 fragment the structure of the observed compound is generated. The borane fragments are very similar. This creates a problem. That is, the valence energy metal-based orbitals of CpMnL2 (four orbitals containing six electrons; see above) give a Cp*TaCl2 fragment a set of three empty and one filled valence orbitals.If one presumes that the three highest energy empty orbitals will be used in cluster bonding 2-{Cp*TaCl2}B4H8 ends up two cluster bonding electrons short to accommodate the geometric structure displayed. The question is why? As this question involves the partitioning of the metal-based valence electrons, geometry, with or without formal electron counting, cannot provide an answer. One needs to dissect the cluster bonding in a systematic fashion.Approximate molecular orbital (MO) methods 51 provide the knife for this dissection using cobaltaborane, a molecule we consider well understood in terms of cluster bonding, as a control. As schematically illus- Fig. 4 The molecular structures of nido-B5H9, nido-2-{Cp*TaCl2}- B4H8, and closo-2,3-{Cp*MoCl}2B3H7. The methyl groups have been removed to improve the view of the core geometries B5H9 2-{Cp*TaCl2}B4H8 2,3-{Cp*MoCl}2B3H7 trated in Fig. 5, this MO analysis suggests the tantalum fragment eVectively provides two electrons to the B4H8 fragment as does CpCo in 2-{CpCo}B4H8.Contrary to initial expectations the fourth and lowest energy metal orbital of the CpML2 fragment, which is the only one filled for Cp*TaCl2, is involved in cluster bonding. But how does the Cp*TaCl2 act as a surrogate BH fragment? Based on overlap populations, all four of the metal-based valence orbitals of the Cp*TaCl2 fragment are substantially involved in binding the B4H8 fragment.However, two orbitals of Cp*TaCl2 are partitioned between a bonding interaction with the B4H8 valence orbitals and a metal-localized non-bonding orbital which is empty. Upon tracking down the corresponding MO of the latter in the cobaltaborane, we find it to be the HOMO which is of d symmetry relative to the principal symmetry axis of the CpCo fragment. Thus, the MO which is the LUMO in the case of the tantalaborane is the HOMO of the cobaltaborane.The two ‘missing electrons’ of {Cp*TaCl2}B4H8 eVectively come from the ‘lone pairs’ of the metal atom rather than from the cluster bonding network. This MO of the tantalaborane lies at higher energy and is unfilled because of the lower eVective nuclear charge of the earlier transition-metal atom and the perturbation of the metal orbitals by the Cl ligands. An alternative view of these MB4 compounds adds understanding. Almost all monometal metallaboranes can be viewed equally well as metal complexes with a ligand set that includes a borane.Thus, h4-C4H4 is viewed as a four-electron donor to a CpCo fragment as is the isoelectronic h4-B4H8 ligand [Fig. 6(a) and 6(b)]. As the latter compound is simply 1-{CpCo}B4H8 in cluster notation, the B4H8 fragment in 2-{CpCo}B4H8 [Fig. 6(c)] must be a four-electron donor to the CpCo fragment even though no organometallic analog exists for comparison. In eVect, the tantalaborane contains a 16-electron metal center and can be considered an analog of a compound like Cp2TiCl2.It behaves as one might expect if one had viewed it as a mimic of an early-transition-metal organometallic complex rather than an isolobal analog of a borane cluster. But, will the same be true of a dimetal cluster of the same shape? 2,4-{Cp*Co}2B3H7 vs. 1,5-{CpMoCl}2B3H7 The synthetic chemistry provides us with a cobalt complex, nido-2,4-{Cp*Co}2B3H7 (Fig. 3),52 which can be compared with a molybdenum complex, closo-2,3-{Cp*MoCl}2B3H7 (Fig. 4).45 Fig. 5 A MO correlation diagram for nido-2-{CpTaCl2}B4H8 showing the interactions of the valence orbitals of the CpTaCl2 fragment with those of the B4H8 fragment CpTaCl2 {CpTaCl2}B4H8 B4H8 37 36 33 32 21 19 11 10 9 7 22 23 24 25 EJ. Chem. Soc., Dalton Trans., 1998, Pages 1525–1531 1529 Again the geometry of the cobalt complex corresponds to that of the isolobal pentaborane(9) cluster, e.g., the B3H7 fragments are very similar and the CpCo fragment is isolobal with BH.Consider now the molybdenum cluster which also looks very similar to the cobalt cluster. A CpML fragment has five valence energy metal-based orbitals. Let us be more flexible now. We can see that the CpMoCl fragment can serve as a three-orbital– zero-, two- or four-electron fragment. Ah, it is simple you say, CpMoCl acts as a two-electron fragment like CpTaCl2 thereby meeting the requirements of a nido five-atom cluster. However, life becomes more complicated if we pay attention to the distinctive geometric diVerences between the two metallaboranes.A comparison of the structural parameters of the Co and Mo clusters defines the problem. The Co]Co distance in 2,3- {Cp*Co}2B3H7 (3.36 Å) confirms the lack of a direct M]M bonding interaction as required by its formulation as a nido cluster. The Mo]Mo distance in {Cp*MoCl}2B3H7 (3.096 Å) is long but still suggests a bonding interaction. Further, the B3 fragment in the molybdenum compound is significantly more open than that of the cobaltaborane (118 vs. 1018). Both observations are consistent with the geometrical expectations of a trigonal bipyramid, i.e., {CpMoCl}2B3H7 is a closo cluster. So the dilemma is the following. If each Mo fragment contributes three orbitals and two electrons to cluster bonding then the geometry should be analogous to that of the cobalt compound. However, it is not. If each Mo fragment contributes three orbitals and zero electrons then the cluster lacks two electrons relative to the number expected for the observed trigonalbipyramidal structure.Both Cp*MoCl fragments are in equivalent environments and it is not reasonable to consider one a two-electron donor and the other a zero-electron donor. Thus, a simple selection of the three orbitals to be used by the metal fragment from the five available provides no solution to the problem. Again, a molecular orbital analysis is helpful. In comparing the MO behaviors of the Co and Mo compounds (Fig. 7) we see two things. First, the structural distortion that brings the two Mo atoms closer together and opens the boron fragment causes a large splitting of a pair of orbitals which are Mo]Mo bonding and Mo]Mo antibonding, respectively. In the Mo complex, the latter ends up at high energy and empty. The pair of orbitals represent a net Mo]Mo bonding interaction. Both of the corresponding orbitals are at lower energy and filled in the case of the cobaltaborane.Second, just as in the case of the tantalaborane, one of the filled, metal-based cluster non-bonding orbitals of the Co compound is found at higher energy and empty in the Mo compound. Thus, the higher metal d orbital energies of the Mo compound vs. the Co compound play the same role they did in the tantalaborane but, in addition, a structural distortion creating a M]M bond leads to the destabilization of a second orbital. As one is hard put to assign a specific isolobal character to the Cp*MoCl fragment, we consider the molybdaborane an unsaturated metallaborane cluster with some similarities to a Fig. 6 Schematic drawings of the structures of CpCo(h4-C4H4) (a), CpCo(h4-B4H8) ]] ] nido-1-{CpCo}B4H8 (b), and nido-2-{CpCo}B4H8 C H H H C C H C Co B H H H B H B H H B Co H H Co H H B H B H H B B H H H (c) (b) (a) 16-electron transition-metal complex but with a closer connection to an unsaturated compound such as H2Os2(CO)10.8 However, in the molybdaborane there is no Mo]Mo double bond and the unsaturation is thought to be substantially spread over the cluster network.Clearly it is spread over the two metal centers but there is evidence that it is spread over the entire cluster framework as well. That is, saturated molybdaboranes, e.g. {Cp*Mo}2B5H9 41 (see below) have shorter M]M distances and longer M]B distances relative to {Cp*MoCl}2B3H7. This is consistent with the M]M d bonding and M]B antibonding nature of the MO which is filled in the former and empty in the latter.In our tight focus on the nido-MnB52n system, the existence of closely related ‘normal’ metallaboranes as controls is very important. A low formal electron count for a single compound is not unambiguous evidence of electronic unsaturation as even in main-group clusters geometric distortions can lead to nonstandard counts.20 These have been discussed before in various contexts, e.g., predicted distortions leading to a stabilization of Sin vs.[Sin]22 clusters.53 However, in the comparison of closely related MnB52n clusters one can distinguish the unusual unsaturated systems from the normal ones. 6 Reaction Chemistry The real proof, and usefulness, of delocalized cluster unsaturation must be found in the reaction chemistry of these molecules. Thus far, only the reaction chemistry of the first example of an unsaturated metallaborane, {Cp*Cr}2B4H8, has been investigated.44,54–56 The analysis of the bonding in this species is analogous to that of the Mo compound;57 it lacks two of the prescribed electrons required by its geometry.However, in this case we have no ‘normal’ transition metal analog with which to compare it and the interpretation of the molecular orbital calculations is not as firmly based. The chromaborane reacts selectively with a variety of substrates all of which are formal Lewis bases. The reactions established to date are shown in the reaction wheel in Fig. 8. Except for one, the products have at least the prescribed number of electrons for the cluster geometry displayed, i.e., the cluster bonding is saturated. Clearly this is one driving force for the reactions although there must be others because {Cp*Cr}2B4H8 has a rather selective reactivity. For example, the reaction of {Cp*Cr}2B4H8 with CS2 results in clean hydroboration of both C]S bonds and binding of the resulting H2CS2 fragment to the cluster.54 On the other hand, alkynes and alkenes fail to react.An interesting reaction is the one in which an electron is added to form the radical anion. Rapid electron exchange is observed Fig. 7 A MO correlation diagram for nido-2,4-{CpCo}2B3H7 and closo-2,3-{CpMoCl}2B3H7, showing the origins of the four-electron diVerence between the two cluster structures E {CpCo}2B3H7 {CpMoCl}2B3H7 45 44 43 42 41 40 39 38 44 45 46 47 48 491530 J. Chem. Soc., Dalton Trans., 1998, Pages 1525–1531 Fig. 8 Schematic diagram of the preparation of {Cp*Cr}2B4H8 and its reactions with main-group and transition-metal moieties Cr H H B H H H H H B B H B Cr Cr H B H H H H B H B Cr S C H H H B H H H H B B H B C O C O Cr H B H H H H B B H B Cr H H Fe(CO)3 Cr H B H H H B B H B Cr H H Co(CO)3 H B H H H H B H B H B B H Cr Cr B S Cr Cr H 44 cve 6 sep 42 cve 5 sep 46 cve 7 sep 44 cve 6 sep 44 cve 6 sep 44 cve 6 sep [Cp*CrCl]2 + BH3THF at r.t. –[Cp*CrCl2]2 + BH3•THF at 55 °C +CS2 at 50 °C + CO at 50 °C + Na–Hg at r.t.[{Cp*Cr}2B4H8]– + BH3•THF at 55 °C + Co2(CO)8 at r.t. + Fe2(CO)9 at 50 °C between the anion and the neutral precursor establishing the basic similarity of the two cluster structures.58 A particularly revealing set of reactions is the addition of isolobal BH and Fe(CO)3 fragments to unsaturated {Cp*Cr}2- B4H8 when compared to the addition of the same fragments to saturated {Cp*Mo}2B5H9.59 The BH fragment inserts into {Cp*Cr}2B4H8 to form {Cp*Cr}2B5H9 with a structure analogous to that of {Cp*Mo}2B5H9.The isolobal Fe(CO)3 fragment, on the other hand, simply co-ordinates to {Cp*Cr}2B4H8 with only minor changes in the cluster geometry of the latter.55 In co-ordinating the chromaborane, the iron fragment contributes two electrons to the {Cp*Cr}2B4H8 cluster bonding system making it saturated and, at the same time, receives four electrons thereby satisfying the 18-electron rule. The Fe(CO)3 fragment is only loosely connected to the bridging hydrogens as it swings back and forth between the two equivalent sites on the NMR time-scale.Addition of BH3 removes the iron fragment and produces {Cp*Cr}2B5H9. The analog {Cp*Mo}2B5H9 also adds Fe(CO)3 but the two electrons the iron fragment contributes to the saturated {Cp*Mo}2B5H9 cluster results in formation of a bicapped octahedron with the iron fragment fully incorporated into the octahedral core. This is perfectly in accord with expectations based on the electron counting rules.Addition of a twoelectron metal fragment to the unsaturated compound produces no major cluster structure modification whereas similar addition to the saturated cluster results in the expected core structural change. 7 Final Remarks It is a curious fact that neither {Cp*MoCl}2B3H7 nor {Cp*Cr}2- B4H8 adopt localized M]M multiple bonding to make up for the low electron counts. Metal–metal multiple bonding is a common response for dinuclear organometallic species and is found in one organometallic relative of {Cp*Cr}2B4H8 which has a Cr]Cr triple bond (Fig. 9).60 However, although metallaborane chemistry is clearly related to organometallic chemistry, we do not expect it to be identical. In fact, it is the diVerences that one finds meaningful. It is also true that the large majority of the known metallaboranes as well as those of the earlier transition metals obey the electron counting rules. It is only in the comparison with the geometric and electronic behavior of these ‘normal’ metallaboranes that the unusual properties of electronically unsaturated metallaboranes are revealed.This is work in progress and much remains to be accomplished. We must, of course, go on to systems containing three metal atoms but first many more metals remain to be explored in the MBn and M2Bn cluster networks. We fully expect more surprises on the way to gaining a better understanding of transition metal–main group cluster systems.After all, the full scope of organometallic chemistry was only revealed when compounds containing metal–carbon bonds for all metals became accessible. Fig. 9 A comparison of schematic drawings of the structures of {Cp*Cr}2B4H8 (42 cve) and {CpCr}2(CO)(C4Ph4) (44 cve). The latter can also be compared with {Cp*Cr(CO)}2B4H6 (44 cve) which is shown in Fig. 8 Cr H H B H H H H H B B H B Cr Cr C Ph Ph Ph C C Ph C Cr C OJ. Chem. Soc., Dalton Trans., 1998, Pages 1525–1531 1531 8 Acknowledgements The experimental eVorts of my co-workers, who are named in the references, created the chemistry described and it was carried out with the generous support of the National Science Foundation. 9 References 1 R. N. Grimes, Carboranes, Academic Press, New York, 1970. 2 R. N. Grimes, in Metal Interactions with Boron Clusters, ed. R. N. Grimes, Plenum, New York, 1982, p. 269. 3 E. L. Muetterties (Editor), Boron Hydride Chemistry, Academic Press, New York, 1975. 4 B. Gates, L. Guczi and H. Knözinger (Editors), Metal Clusters in Catalysis, Elsevier, New York, 1986, vol. 29. 5 J. D. Woollins, Non-Metal Rings, Cages and Clusters, Wiley, New York, 1988. 6 H. W. Roesky (Editor), Rings, Clusters and Polymers of Main Group and Transition Elements, Elsevier, Amsterdam, 1989. 7 C. E. Housecroft, Boranes and Metalloboranes, Ellis Horwood, Chichester, 1990. 8 D. F. Shriver, H. D. Kaesz and R. D. Adams (Editors), The Chemistry of Metal Cluster Complexes, VCH, New York, 1990. 9 D. M. P. Mingos and D. J. Wales, Introduction to Cluster Chemistry, Prentice Hall, New York, 1990. 10 C. E. Housecroft, Cluster Molecules of the p-Block Elements, Oxford University Press, Oxford, 1994. 11 M. H. Chisholm (Editor), Early Transition Metal clusters with p- Donor Ligands, VCH, New York, 1995. 12 D. M. P. Mingos (Editor), Structural and Electronic Paradigms in Cluster Chemistry, Springer, Berlin, 1997, vol. 87. 13 K. Wade, Electron Deficient Compounds, Nelson, London, 1971. 14 K. Wade, Inorg. Nucl. Chem. Lett., 1972, 8, 559. 15 K. Wade, New Scientist, 1974, 62, 615. 16 K. Wade, Adv. Inorg. Chem. Radiochem., 1976, 18, 1. 17 D. M. P. Mingos, Nature (London), 1972, 236, 99. 18 D. M. P. Mingos, Acc. Chem. Res., 1984, 17, 311. 19 D. M. P. Mingos, J. Chem. Soc., Chem. Commun., 1985, 1352. 20 D. M. P. Mingos and R. L. Johnston, Struct. Bonding (Berlin), 1987, 68, 29. 21 D. M P. Mingos and A. S. May, in The Chemistry of Metal Cluster Complexes, eds.D. F. Shriver, H. D. Kaesz and R. D. Adams, VCH, New York, 1990. 22 T. P. Fehlner (Editor), Inorganometallic Chemistry, Plenum, New York, 1992. 23 W. N. Lipscomb, Boron Hydrides, Benjamin, New York, 1963. 24 R. E. Williams, Inorg. Chem., 1971, 10, 210. 25 R. E. Williams, Adv. Inorg. Chem. Radiochem., 1976, 18, 67. 26 J. D. Kennedy, Prog. Inorg. Chem., 1984, 32, 519. 27 J. D. Kennedy, Prog. Inorg. Chem., 1986, 34, 211. 28 C. E. Housecroft and T. P. Fehlner, Adv.Organomet. Chem., 1982, 21, 57. 29 C. E. Housecroft, Coord. Chem. Rev., 1995, 143, 297. 30 C. E. Housecroft, Chem. Soc. Rev., 1995, 215. 31 R. N. Grimes, Pure Appl. Chem., 1982, 54, 43. 32 R. HoVmann, Science, 1981, 211, 995. 33 A. K. Bandyopadhyay, R. Khattar and T. P. Fehlner, Inorg. Chem., 1989, 28, 4434. 34 K. H. Theopold, Acc. Chem. Res., 1990, 23, 263. 35 F. A. Cotton and R. A. Walton, Multiple Bonds Between Metal Atoms, Wiley, New York, 1982. 36 D. N. Cox, D.M. P. Mingos and R. HoVmann, J. Chem. Soc., Dalton Trans., 1981, 1788. 37 J. C. Vites, C. E. Housecroft, C. Eigenbrot, M. L. Buhl, G. J. Long and T. P. Fehlner, J. Am. Chem. Soc., 1986, 108, 3304. 38 J. Feilong, T. P. Fehlner and A. L. Rheingold, J. Am. Chem. Soc., 1987, 109, 1860. 39 X. Meng, A. K. Bandyopadhyay, T. P. Fehlner and F.-W. Grevels, J. Organomet. Chem., 1990, 394, 15. 40 C. Ting and L. Messerle, J. Am. Chem. Soc., 1989, 111, 3449. 41 H. J. Bullick, P. D. Brebenik, M. L. H. Green, A. K. Hughes, J. B. Leach and P. C. McGowan, J. Chem. Soc., Dalton Trans., 1995, 67. 42 Y. Nishihara, K. J. Deck, M. Shang, T. P. Fehlner, B. S. Haggerty and A. L. Rheingold, Organometallics, 1994, 13, 4510. 43 X. Lei, M. Shang and T. P. Fehlner, J. Am. Chem. Soc., 1989, 120, 2686. 44 J. Ho, K. J. Deck, Y. Nishihara, M. Shang and T. P. Fehlner, J. Am. Chem. Soc., 1995, 117, 10 292. 45 S. Aldridge, M. Shang and T. P. Fehlner, J. Am. Chem. Soc., 1997, 119, 11 120. 46 A. S. Weller, M. Shang and T. P. Fehlner, 1998, unpublished work. 47 S. Aldridge, H. Hashimoto, M. Shang and T. P. Fehlner, Chem. Commun., 1998, 207. 48 S. Aldridge, M. Shang and T. P. Fehlner, J. Am. Chem. Soc., 1998, 120, 2586. 49 V. R. Miller, R. Weiss and R. N. Grimes, J. Am. Chem. Soc., 1977, 99, 5646. 50 L. G. Sneddon and D. Voet, J. Chem. Soc., Chem. Commun., 1976, 118. 51 T. A. Albright, J. K. Burdett and H.-H. Whangbo, Orbital Interactions in Chemistry, Wiley, New York, 1985. 52 Y. Nishihara, K. J. Deck, M. Shang and T. P. Fehlner, J. Am. Chem. Soc., 1993, 115, 12 224. 53 T. Slee, L. Zhenyang and D. M. P. Mingos, Inorg. Chem., 1989, 28, 2256. 54 H. Hashimoto, M. Shang and T. P. Fehlner, Organometallics, 1996, 15, 1963. 55 H. Hashimoto, M. Shang and T. P. Fehlner, J. Am. Chem. Soc., 1996, 118, 8164. 56 S. Aldridge, T. P. Fehlner and M. Shang, J. Am. Chem. Soc., 1997, 119, 2339. 57 T. P. Fehlner, Struct. Bonding (Berlin), 1997, 87, 112. 58 K. Kawamura and T. P. Fehlner, Organometallics, 1998, in the press. 59 S. Aldridge, H. Hashimoto, K. Kawamura, M. Shang and T. P. Fehlner, Inorg. Chem., 1998, 37, 928. 60 S. A. R. Knox, R. F. D. Stansfield, F. G. A. Stone, M. J. Winter and P. Woodward, J. Chem. Soc., Dalton Trans., 1982, 173. Received 3rd February 1998; Paper 8/00924D

ISSN:1477-9226    

DOI:10.1039/a800924d

出版商:RSC    

年代:1998

数据来源: RSC

摘要:

DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1999, 1531–1532 1531 Synthesis and characterization of K3Cu11Te16 from supercritical ethylenediamine Mehtap Emirdag, George L. Schimek and Joseph W. Kolis Department of Chemistry, H. L. Hunter Chemistry Laboratory, Clemson University, Clemson, SC 29634-1905, USA Received 8th December 1998, Accepted 23rd March 1999 Reaction of copper with potassium polytelluride, K2Te5, in supercritical ethylenediamine (305 8C, 2370 psi) led to a new compound with building blocks formed from Cu8- Te12 pentagonal dodecahedral cages made with Cu2Te3 pentagons; the dodecahedrons are connected by copper atoms and tritelluride fragments to generate tunnels filled with potassium.Supercritical fluids have been used successfully to synthesize novel compounds at relatively low temperature. We have been exploring the preparation of a number of ternary alkali metal transition metal sulfides 1 and selenides 2,3 by using supercritical amines as the reaction media.To further study the chemistry of transition metal chalcogenides in supercritical amines, we have begun to explore the chemistry of metal tellurides. Synthesis of ternary tellurides has accelerated lately, but there are still only a limited number of ternary alkali metal copper tellurides known, KCuTe,4 NaCuTe,5 KCu3Te2,6 NaCu3Te2,7 K2Cu5Te5,8 and K4Cu8Te11 9,10 all of which have been prepared via melt reactions. In this paper we report the synthesis and structural characterization of the new compound, K3Cu11Te16.The new compound, K3Cu11Te16, was prepared from the reaction of Cu powder, K2Te5 and Te powder in supercritical ethylenediamine (en).11 The compound is extremely air and thermally sensitive. This sensitivity inhibits much physical characterization. The structure of K3Cu11Te16 was studied by single crystal X-ray diVraction.12 The building block of the structure is a Cu8(Te2)6 pentagonal dodecahedral cage shown in Fig. 1. The pentagonal dodecahedron is made of planar Cu2Te3 pentagons [average deviation from planarity 0.04(1) Å] each containing a ditelluride edge. The pentagons are connected through the ditelluride with copper to form dodecahedrons. Each adjacent dodecahedral cage shares ditellurides to form columns parallel to the c axis. In addition to ditellurides, the structure has tri- Fig. 1 Structure of KCu8Te12 dodecahedral cage. Selected distances (Å): Te(3)–Te(5) 2.801(2), Te(3)–Cu(2) 2.589(2), Te(3)–Cu(2c) 2.589(2), Te(4)–Te(4c) 2.773(2); Te(4)–Cu(2a) 2.602(2), Te(4)–Cu(3a) 2.601(2), Te(5)–Cu(3b) 2.612(2), Te(5)–Cu(3c) 2.612(2), Te(6)–Te(6c) 2.834(2), Te(6)–Cu(3a) 2.670(2), Te(6)–Cu(3b) 2.670(2).Average angles within the cage: Te–Cu–Te 109(3) and Te–Te–Cu 108(4)8. tellurides as bridging groups between the dodecahedral cages along the a axis. The central atom of the tritelluride chain also coordinates to a bridging tetrahedral Cu atom which links the dodecahedra together.These tetrahedral copper atoms connect the cages along the b axis thereby generating tunnels that are filled with K1 cations. The average distance across the tunnel is around 6.5 Å. The unit cell view of K3Cu11Te16 is shown in Fig. 2. There are four unique copper atoms tetrahedrally coordinated by tellurium atoms. Bond distances between Cu(3) and Te2 22 ligand and Cu(2) and Te3 22 ligand range from 2.601(2) to 2.670(2) Å and 2.585(2) to 2.738(2) Å, respectively.Both Cu(3) and Cu(2) are part of the dodecahedron and they reside on general positions as do Te(2) and Te(4). Cu(1) is located on an mm2 site between two cages and connects clusters along the b axis. Cu(1) has two bonds to tritellurides at 2.634(2) Å and two bonds to ditellurides with an average bond distance of 2.693(4) Å. Cu(4) is also located between two cages and connects them together along the a axis. Cu(4) has two bonds to tritellurides and two bonds to ditellurides with bond distances of 2.597(1) and 2.587(1) Å, respectively.Cu(4) and Te(6) reside on 2-fold sites. The average Te–Te distances of the ditellurides in the cluster are 2.80(3) Å, while in the tritelluride they are 2.817(1) Å. There are two unique K1 cations in the title structure. K(1) is located on an mm2 site in the center of the dodecahedral cage and has twelve interactions with the Te atoms of the dodecahedron. K(1)–Te distances range from 3.647(1) to 3.951(1) Å.K(2) resides in the tunnel and has eight interactions with Te atoms. The average K(2)–Te distance is 3.75(3) Å. The formula of K3Cu11Te16 can be written more informatively as K3Cu11[(Te2)5(Te3)2]. If the formal charges of the Fig. 2 Packing diagram of K3Cu11Te16. Cu atoms are shown as red spheres, tellurium atoms as orange spheres, and potassium atoms as green spheres.1532 J. Chem. Soc., Dalton Trans., 1999, 1531–1532 ditelluride and tritelluride units are considered to be 22, and that of Cu as 11, the compound is electron precise, and should be a semiconductor. We were not able to do conductivity measurements due to the air-sensitivity of the compound, but the black shiny appearance of the crystals suggests that it is a narrow band gap semiconductor.The copper is presumably oxidized to Cu(I) via the oxidative cleavage of the polytelluride chain. 2Cu 1 Te5 22 æÆ 2Cu1 1 Te2 22 1 Te3 22 The title compound, K3Cu11Te16, is somewhat similar to K4Cu8Te11 as synthesized from molten polytelluride flux.9 Both compounds have the same Cu8Te12 building block.The primary diVerences between the two are that K4Cu8Te11 has large tunnels filled with six cations, and has only ditellurides in the structure. In the title compound the dodecahedra are only fused to each other through a shared Te2 22 in the z direction, creating a chain. Along the y direction these chains are linked via shared tetrahedral copper atoms while in the x direction another tetrahedral copper atom and a Te3 22 form the bridging units.In contrast, the dodecahedra in K4Cu8Te11 are more densely packed. Not only are there fused chains similar to K3Cu11Te16, but two such chains are also fused through another shared Te2 22 creating a column. This column connects to four other symmetry related columns through adjacent Cu–Te interactions and another Te2 22. This anionic arrangement creates large columnar channels that are filled with the counter cation potassium.We have never observed formation of the previously reported flux synthesized K4Cu8Te11 phase in numerous reactions in supercritical amines. Kanatzidis et al. have synthesized a number of alkali metal copper tellurides by using the flux method.10 For these reactions, they have observed a general trend that lower temperature and more acidic (longer-chained) Tex 22 (x = 4, 5) favor polychalcogenide compounds, while higher temperature and more basic (shorter-chained) Tex 22 (x = 2, 3) tend to form monochalcogenides.Based on our reaction conditions of relatively low temperature and long chain polytelluride (305 8C and K2Te5), it appears that our compound follows this trend. Acknowledgements We are indebted to the National Science Foundation for support of this work. Notes and references 1 M. Emirdag, G. L. Schimek and J. W. Kolis, Acta Crystallogr., Sect. A, in the press. 2 M. Emirdag, G. L. Schimek and J. W. Kolis, J. Solid State Chem., in the press. 3 M.Emirdag, G. L. Schimek and J. W. Kolis, J. Chem. Cryst., in the press. 4 Y. Park, Ph.D. Dissertation, Michigan State University, 1992. 5 G. Savelsberg and H. Schäfer, Z. Naturforsch., Teil B, 1978, 33, 370. 6 K. O. Klepp, J. Less-Common Met., 1987, 128, 79. 7 G. Savelsberg and H. Schäfer, Mater. Res. Bull., 1981, 16, 1291. 8 Y. Park, D. C. Degroot, J. Schindler, C. R. Kannewurf and M. G. Kanatzidis, Angew. Chem., Int. Ed. Engl., 1991, 30, 1325. 9 Y. Park and M. G. Kanatzidis, J. Chem. Mater., 1991, 3, 781. 10 X. Zhang, Y. Park, T. Hogan, J. L. Schindler, C. R. Kannewurf, S. Seong, T. Albright and M. G. Kanatzidis, J. Am. Chem. Soc., 1995, 117, 10300. 11 The synthesis of K3Cu11Te16 was achieved by combining 30 mg Cu powder (0.47 mmol), 52 mg K2Te5 (0.072 mmol), and 44 mg Te powder (0.34 mmol) in 0.7 ml of ethylenediamine (en) in a 1– 4 inch quartz tube. The tube was sealed under vacuum and placed in a high-pressure autoclave. The autoclave was counter-pressured to 2370 psi, placed in a furnace, and heated at 305 8C for one day.Tubes were opened under argon and washed with ethylenediamine and tetrahydrofuran using standard Schlenk techniques. After the crystals were dried under vacuum, they were placed in mineral oil. Shiny black column crystals were recovered with over 90% yield. They are extremely air sensitive. 12 Crystal data for K3Cu11Te16. The single crystal X-ray diVraction data for the title compound were collected at room temperature on a four circle Rigaku AFC7R diVractometer (Mo-Ka = 0.71073 Å, graphite monochromator). The unit cell was based on the indexing of 25 reflections with a 2q range of 36.89–44.328. Three standard reflections were measured every 100 reflections with decay = 21.6%. 2qmax = 508, M = 2857.90, gray column, 0.064 × 0.064 × 0.034 mm. Orthorhombic Imma, Z = 4, a = 23.642(3), b = 19.577(2), c = 6.958(2) Å, V = 3220.5(9) Å3, Dc = 5.89 g cm23, m(Mo-Ka) = 21.701 mm21. Refinement on F, 1255 observed reflections [I > 3s(I)], 80 parameters (all atoms anisotropic), R = 0.031, Rw = 0.037, S = 2.59, maximum and minimum residual electron density: 1.28 and 21.01 e Å23. The structure was solved with direct methods and refined by full matrix least-square techniques explained as in previous work.1 CCDC reference number 186/1408. Communication 9/02530H

ISSN:1477-9226    

DOI:10.1039/a902530h

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1999, 1533–1534 1533 A new thallium–thiolate coordination environment as found in the polymer [{Tl7(SthV)6}n][PF6]n (HSthV 5 tetrahydrofurfurylthiol) J. Elaine Barclay,a David J. Evans,*a Sian C. Davies,a David L. Hughes a and Piotr Sobota b a Nitrogen Fixation Laboratory, John Innes Centre, Norwich Research Park, Colney, Norwich, UK NR4 7UH. E-mail: dave.evans@bbsrc.ac.uk b Faculty of Chemistry, University of Wroclaw, 14F Joliot-Curie, 50-383 Wroclaw, Poland Received 19th March 1999, Accepted 9th April 1999 The crystal structure of the polymeric thallium–thiolate salt [{Tl7(Sthff)6}n][PF6]n (HSthff 5 tetrahydrofurfurylthiol) shows an unusual octahedrally coordinated thallium(I) atom linking novel Tl6S6 “prismane” units. Homoleptic thiolate complexes of thallium(I) crystallise as monomers, oligomers and polymers and exhibit a large structural variety in the solid state.Structurally characterised complexes include: monomeric [Tl(1,1-S2PEt2)] 1 and dimeric [Tl(1,1-S2CNR2)]2 (R = Me, Et , nPr, iPr, nBu or iBu); 2,3 discrete molecules of [Tl8(StBu)8]; 4,5 polymeric TlSPh which consists of cage-like structural units [Tl5(SPh)6]2 and [Tl7(SPh)6]1; 4,5 linked 1,3-Tl2S2 ring two-dimensional polymers5–8 for TlSCH2Ph, TlSC7H7, TlSC6H11 and Tl{2,4,6-SC6H2(CF3)3}?0.5diox (diox = 1,4-dioxane); dimeric [Tl{SSi(OtBu)3}]2, which also contains a 1,3-Tl2S2 ring; 9 and [{Tl[1,2-(m-S)2C6H4]}2]22 which contains Tl2S4 cages.10 Here we report a new structural motif for thallium(I) thiolates.Although first reported in the 1950’s 11–13 the use of tetrahydrofurfurylthiol † (HSthV), Fig. 1, as a proligand in coordination chemistry has not been explored. As part of a survey of the utility of the anion of HSthV as a ligand, we have studied its reaction with thallium hexafluorophosphate. Under anaerobic conditions, to a stirred solution of NaSthV‡ (0.48 g, 3.4 mmol) in methanol (35 ml) was added thallium(I) hexafluorophosphate (1.19 g, 3.4 mmol) to form immediately a yellow solution.The mixture was stirred for 10 min after which time a yellow precipitate had begun to form. The precipitate redissolved on gentle warming and on cooling to room temperature gave the product as yellow needle-like crystals (0.52 g, 50% based on Tl).§ Crystallography¶ showed a structure of polymeric chain cations of [Tl7(SthV)6]1, Fig. 2; the novel Tl6S6 “prismane”- like units are linked through thallium(I) atoms on opposite faces.The resultant Tl8S6 unit consists of a cube of thallium atoms and each face of the cube is capped by the sulfur atom of a thiolate ligand; the unit shows pseudo-m3m symmetry. There is a crystallographic centre of symmetry at the centre of each cube and a pseudo-centre at each linking thallium. Two independent cube units alternate along the cation chain; there are only very small diVerences in orientation between the two units.There are thus two types of thallium environment, one trigonal pyramidal (bonded to three S atoms) at the six corners of the “prismane”, and one octahedral (bonded to six S atoms) at the bridging positions. This is the first example of a thallium thiolate containing octahedrally coordinated thallium atoms.|| The Tl–S bonds involving the trigonal pyramidal Tl atoms, i.e. within the prismane units, are in the range 2.880(11)–3.086(11) Å, mean 2.97(2) Å; about the linking, octahedral Tl(7) atoms, the Tl–S distances are rather longer, 3.138(11)–3.302(12) Å, mean 3.24(2) Å.Each sulfur atom is coordinated to four thallium atoms and the carbon atom of a tetrahydrofurfuryl group in a square-pyramidal pattern. There is extensive disorder in the thiolate ligands, which do not conform to the pseudo-m3m symmetry. There is a chiral centre in each ligand and we were not able to resolve clearly the two possible configurations in any of the ligands.We do note, however, that each tetrahydrofurfuryl group is oriented with the ring tilted towards a thallium atom so that there are interactions between the thallium atom and the atom (which we designated an oxygen atom) at the 2-position in the ring with distances of 3.03(9) to 3.36(7) Å. There is also disorder in the positions of the fluorine atoms of the discrete PF6 2 anions. The cation chains lie parallel to the crystallographic c axis and the anions lie between pairs of chains.Each cation chain is linked to four others via the anions with close Tl ? ? ? F interactions, the shortest five of which are in the range 3.05(8) to 3.39(5) Å. The basic structural Tl–S motif in our complex, the capped cube, has similarities only, to our knowledge, in the two units that are bonded together in the complex polymer of Tl(SPh).4,5 In our polymeric chain, two almost identical cubes, each with little distortion from regular cubes, are linked alternately through opposite corners of the cubes.The polymer of Tl(SPh) is a three-dimensional lattice of units described as [Tl7(SPh)6]1 Fig. 1 Tetrahydrofurfurylthiol showing atom numbering. O SH 5 4 3 1 6 2 Fig. 2 Fragment of a cation chain showing the linking of Tl8S6 cube units through Tl(7) atoms at opposite corners of the cubes. Only the Tl, S and a-carbon atoms are shown; the sulfur atoms are hatched, the larger and smaller open circles represent the Tl and C atoms respectively, and the atom numbering scheme is indicated.The primed numbers indicate symmetry operations about inversion centres: 9 at 1 2 x, 2y, 2z and 0 at 1 2 x, 2y, 1 2 z.1534 J. Chem. Soc., Dalton Trans., 1999, 1533–1534 and [Tl5(SPh)6]2; the crystal system is cubic and both units lie on three-fold symmetry axes. The cationic unit is a capped cube but with one Tl corner void. The anion is less regular, it is basically a completely capped cube but with three of the thallium atoms displaced considerably from the regular cube; these three atoms are in fact the atoms that link the cations with the anions and are not included in the formula of the anion units.If one starts with the array of sulfur atoms, as suggested by Krebs and Brömmelhaus,4 the six sulfur atoms of the cation unit form an almost regular octahedron and seven of the eight faces of the octahedron are capped by thallium atoms. The octahedron of sulfur atoms in the anion is less regular, five of its faces are Tl-capped and the remaining three faces have rather oVset capping thallium atoms.In our Tl(SthV) polymer, the two S6 octahedra are close to regular and each has all eight faces capped by thallium atoms. Acknowledgements We thank the Biotechnology and Biological Sciences Research Council for funding. Notes and references † IUPAC name: (tetrahydrofuran-2-yl)methanethiol. ‡ HSthV was prepared by a method similar to that reported previously. 13 Found: C, 51.0; H, 8.4; S, 27.2.C5H10OS requires: C, 50.8; H, 8.5; S, 27.1%; nmax/cm21 (SH) 2555 (neat, KBr disc); dH(400 MHz; CDCl3) 3.85 (m, 3H, CH2{3} and CH{1}), 2.62 (m, 2H, CH2{6}), 1.8 (mm, 4H, CH2{4 1 5}), 1.47 (t, 1H, SH); dC(100 MHz, CDCl3) 79.9 {C1}, 68.4 {C3}, 30.2 {C5}, 29.5 {C6}, 25.9 {C4}. NaSthV was prepared from HSthV and sodium metal in tetrahydrofuran. § Found: C, 15.8; H, 2.3; S, 8.2. C30H54F6O6PS6Tl7 requires: C, 15.8; H, 2.4; S, 8.4%; lmax/nm (CH3OH) 224 (e/dm3 mol21 cm21 47600), 256 (sh) (13600) and 292 (8700); dH(400 MHz, CD3OD) 4.18 (m, 1H, CH{1}), 3.80 (mm, 2H, CH2{3}), 3.29 (m, 2H, CH2{6}), 1.60–2.14 (mm, 4H, CH2{4 1 5}); dC(100 MHz, CDCl3): 80.0 {C1}, 68.0 {C3}, 30.5 {C5}, 29.7 {C6}, 26.0 {C4}.¶ Crystal data: C30H54O6S6Tl7PF6, M = 2278.7, monoclinic, space group P21/n (equivalent to no. 14), a = 14.401(2), b = 23.661(2), c = 15.2958(12) Å, b = 97.404(9)8, V = 5168.4(10) Å3, Z = 4, Dc = 2.93 g cm23, F(000) = 4056, T = 293(1) K, m(Mo-Ka) = 220.7 cm21, l(Mo-Ka) = 0.71069 Å, 5054 reflections measured, 4819 unique, Rint = 0.053, wR2 = 0.133, R1 = 0.161 for all the data;14 R1 = 0.052 for the 1683 “observed” data. In the refinement, the Tl, S and P atoms were refined anisotropically, while the C, O and F atoms (most of which were disordered with partial site occupancy) were refined with isotropic thermal parameters.CCDC reference number 186/1419. See http:// www.rsc.org/suppdata/dt/1999/1533/ for crystallographic files in .cif format.|| It has been suggested,6 though not fully reported, that a similar coordination environment is seen in polymeric [Tl5(SC3H7)5]. 1 S. Esperas and S. Husebye, Acta Chem. Scand., Ser. A, 1974, 28, 1015. 2 L. Nilson and R. Hesse, Acta Chem. Scand., 1969, 23, 1951; P. Jennische, A. Olin and R. Hesse, Acta Chem. Scand., 1972, 26, 2799; P. Jennische and R. Hesse, Acta Chem. Scand., 1973, 27, 3531; H. Anacker-EickhoV, P. Jennische and R. Hesse, Acta Chem. Scand., Ser.A, 1975, 29, 51; H. Pritzkow and P. Jennische, Acta Chem. Scand., Ser. A, 1975, 29, 60; E. Efwing, H. Anacker-EickhoV, P. Jennische and R. Hesse, Acta Chem. Scand., Ser. A, 1976, 30, 335. 3 D. Coucouvanis, Prog. Inorg. Chem., 1979, 26, 301. 4 B. Krebs and A. Brömmelhaus, Angew. Chem., Int. Ed. Engl., 1989, 28, 1682. 5 B. Krebs and A. Brömmelhaus, Z. Anorg. Allg. Chem., 1991, 595, 167. 6 A. Brömmelhaus, A. Pinkerton and B. Krebs, Annual Meeting of the American Crystallographic Association, Toledo, OH, 1991, abstract p. 123. 7 B. Krebs, A. Brömmelhaus, B. Kersting and M. Nienhaus, Eur. J. Solid State Inorg. Chem., 1992, t29, 167. 8 D. Labahn, E. Pohl, R. Herbst-Irmer, D. Stalke, H. W. Roesky and G. M. Sheldrick, Chem. Ber., 1991, 124, 1127. 9 W. Wojnowski, K. Peters, E.-M. Peters and H. G. v. Schnering, Z. Anorg. Allg. Chem., 1985, 531, 147. 10 B. E. Bosch, M. Eisenhawer, B. Kersting, K. Kirschbaum, B. Krebs and D. M. Giolando, Inorg. Chem., 1996, 35, 6599. 11 J. H. Chapman and L. N. Owen, J. Chem. Soc., 1950, 579. 12 Y. K. Yuryev and E. G. Vendelshtein, Zh. Obshch. Khim., 1952, 22, 687; J. Gen. Chem. USSR (Engl. Transl.), 1952, 22, 751. 13 V. C. Barry and J. E. McCormick, Proc. R. Irish Acad., Sect. B, 1958, 59, 345. 14 G. M. Sheldrick, SHELXL, Program for Crystal Structure Refinement, University of Göttingen, 1993. Communication 9/02193K

ISSN:1477-9226    

DOI:10.1039/a902193k

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1999, 1535–1536 1535 A novel polymeric copper(I) complex with an unusual azide bridge. Synthesis and crystal structure of [Cu(pyza)(Ï-1,1,3-N3)]• (pyza 5 pyrazinecarboxamide) Mohamed A. S. Goher*a and Franz A. Mautner b a Department of Chemistry, Faculty of Science, Kuwait University, PO Box 5969 Safat, 13060 Kuwait b Institut fur Physikalische und Theoretische Chemie, Technische Universitat Graz, A-8010 Graz, Austria Received 22nd February 1999, Accepted 22nd March 1999 The first polymeric copper(I) complex with a Ï-1,1,3-N3 bridge [Cu(pyza)(Ï-1,1,3-N3)]• (pyza 5 pyrazinecarboxamide) has been synthesized and characterized.The azide ion, N3 2 is known to coordinate to metals in both terminal and bridging modes. As a bridging ligand it can link a pair of metal centers in either an end-on (m-1,1) or an end-toend (m-1,3) bonded fashion. The azide ligand may link a third metal atom giving rise to a m-1,1,3 mode (see below).Bi-, and poly-nuclear copper azide systems are of considerable interest due to the broad range of their structural and magnetic properties.1 However, the vast majority of studies have focused on copper(II) azide complexes.2 On the other hand, reports on copper(I) complexes are scarce and only two structures, namely those of [(Ph2P)2Cu(m-1,3-N3)Cu(PPh3)2] 3 and [Cu(m-pyz-N,N)(m-1,3-N3)]n (pyz = pyrazine),4 which exhibit m-1,3-N3 bridges, and that of [Cu2(m-Ph2Ppypz)2(m-1,1-N3)]- [ClO4]?Et2O [Ph2Ppypz = 2-(diphenylphosphino)-6-(pyrazol- 1-yl)pyridine] which possesses a m-1,1-N3 bridge,5 have been established.To our knowledge, there is as yet no known example of a polynuclear copper(I) complex with a m-1,1,3-N3 bridge. We report here the first structure of this kind in the polymeric copper(I) complex [Cu(pyza)(m-1,1,3-N3)]• (pyza = pyrazinecarboxamide). The reaction between copper(I) azide and pyrazinecarboxamide resulted in the formation of a deep red-brown complex in high yield.† Copper(I) azide has been rarely used in coordination chemistry due to its insolubility in common solvents.Signifi- cantly, concentrated solutions of sodium or potassium azide can be used to dissolve polymeric CuI(N3). Importantly, using this procedure allowed us to avoid the isolation of any CuII impurities, and single crystals of [Cu(N3)L]•‡ were grown by mixing with pyrazinecarboxamide in EtOH. This synthesis was found to be reproducible. The IR spectrum of [Cu(pyza)(m-1,1,3-N3)]• shows the characteristic asymmetric N3 stretching vibrations at 2070 cm21, which is substantially higher than that of the Ph2Ppypz complex with m-1,1-N3 bridges (2037 cm21) as well as those of the binuclear copper(I) complex of Ph3P (2055 cm21) or the polymeric complex of pyrazine (2041 cm21),4 both with m-1,3- N3 bridges.This value, however, along with the symmetric N3 stretch band at 1322 cm21 is indicative of an asymmetric azido ligand.The carbonyl and amide stretching vibrations appear almost at the same positions as in the spectrum of free pyrazinecarboxamide, suggesting that the CONH2 group is not involved in bond formation. The polymer [CuL(N3)]• exists as a 3-D network consisting of sheets of [CuI(N3)]• linked by pyrazinecarboxamide (Fig. 1). Within the [Cu(N3)]• sheets each copper(I) center is coordinated by three symmetry-related azide ligands, thus the azido groups behave as m-1,1,3 bridges forming a 2-D layer structure oriented parallel to the ab plane of the monoclinic unit cell.This arrangement leads to the formation of ten-membered Cu–NNN–Cu–N–Cu–NNN rings (Fig. 2). Owing to the rigid rod-like nature of (N3)2 each ten-membered ring adopts a pseudo-chair conformation, similar to that found in the structure of [Cu2(NCS)2(pyz)]•.6 The remaining fourth coordination site of each CuI center is occupied by a nitrogen atom [N(2)] of a pyrazinecarboxamide ligand.The 2-D layers are connected via N–H ? ? ? O hydrogen bonds formed by adjacent pyrazinecarboxamide molecules. The structural motif observed in the complex [Cu(N3)L]• represents both a new arrangement of polymeric CuI(N3) and also a new 3-D network. The preparation of this compound Fig. 1 A view of [Cu(N3)(L)]• down the b axis showing the arrangement of the CuI–azido-sublattice as 2-D layers oriented normal to the c axis. The 2-D layers are connected via N–H ? ? ? O hydrogen bonds formed by adjacent pyrazineamide ligands (L).1536 J.Chem. Soc., Dalton Trans., 1999, 1535–1536 illustrates a new and potentially versatile approach to the construction of uncharged inorganic coordination networks and we are currently pursuing this methodology towards the synthesis of such new materials. Acknowledgements Financial support by the Kuwait University Research Administration Project (SC097) and the Department of Chemistry Fig. 2 A view of the [CuI(N3)]• sheet showing the ten-membered pseudo-cyclohexane rings.Selected bond lengths (Å) and angles (8): Cu(1) ? ? ? Cu(1A)[Cu(1C)] 3.484(1), Cu(1) ? ? ? Cu(1B) 5.533(2), Cu(1)– N(4) 2.098(3); Cu(1)–N(4C) 2.030(3), Cu(1)–N(6E) 2.007(4), not shown: Cu(1)–N(2) 2.067(3) and N(3) ? ? ? O(1D) 2.890(5); N(2)–Cu(1)– N(4) 101.60(14), N(4)–Cu(1)–N(4C) 116.07(7), N(4C)–Cu(1)–N(2) 107.33(13), N(2)–Cu(1)–N(6E) 114.2(2), N(4C)–Cu(1)–N(6E) 114.5(2), N(4)–Cu(1)–N(6E) 102.63(14) [symmetry codes: A ��� 2x, y 2 ��� , ��� 2 z; B ��� 2 x, y 2 ��� , ��� 2 z; C ��� 2 x, y 1 ��� , ��� 2 z; D 3 2 x, 1 2 y, 2z; E ��� 2 x, ��� 1 y, ��� 2 z].General Facility Projects (Analab) are gratefully acknowledged. The authors thank Professor Kratky and Dr Belaj (Graz University) for the use of the STOE diVractometer. Notes and references † Preparation of [Cu(pyza)(m-1,1,3-N3)]•. To an aqueous suspension of CuN3 (2 mmol) a saturated solution of NaN3 was added until a clear solution was obtained.Pyrazinecarboxamide (3 mmol) dissolved in ethanol (10 ml) was then added and the final mixture allowed to stand for several days to deposit deep red-brown needle-like crystals of the complex. Yield, ca. 80% [Found (Calc.): C, 27.1 (26.26); H, 2.3 (2.21); N, 36.5 (36.73); Cu, 27.2 (27.78%)]. IR (KBr disc): n(N3) 2070, 1322 cm21. Electronic spectrum (solid Nujol mull): 409 (br), 628 (br) nm (CuIÆL CT). ‡ Crystal data. [Cu(pyza)(m-1,1,3-N3)]•, C5H5N6OCu, M = 228.69, monoclinic, space group P21/n, a = 5.528(2), b = 5.442(2), c = 25.999(9) Å, b = 90.59(3)8, V = 782.1(4) Å3, Z = 4, m(Mo-Ka) = 2.757 mm21, T = 298(2) K. 1543 unique reflections (Rint = 0.0242) were collected. At final convergence R1 [1296 data with I > 2s(I)] = 0.0427, wR2 (all 1528 data) = 0.1132 for 125 parameters. CCDC reference number 186/1396. See http://www.rsc.org/suppdata/dt/1999/1535/ for crystallographic files in .cif format. 1 E. I. Solomon, M. J. Baldwin and M. D. Lowery, Chem. Rev., 1992, 92, 521. 2 See for example, (a) L. K. Thompson, S. S. Tandon and M. E. Manuel, Inorg. Chem., 1995, 34, 2356; (b) J. H. Satcher, Jun., M. W. Droege, T. J. R. Weakly and R. T. Taylor, Inorg. Chem., 1995, 34, 3317; (c) A. Escuer, R. Vicente, F. A. Mautner and M. A. S. Goher, Inorg. Chem., 1997, 36, 1233. 3 R. F. Ziolo, A. P. Gaughan, Z. Dori, C. G. Pierpont and R. Eisenberg, Inorg. Chem., 1971, 10, 1289. 4 M. A. S. Goher and F. A. Mautner, Polyhedron, in the press. 5 S.-M. Kuang, Z.-Z. Zhang, Q.-G. Wang and T. C. W. Mak, J. Chem. Soc., Dalton Trans., 1997, 4477. 6 A. J. Blake, N. R. Champness, M. Crew, L. R. Hannon, S. Parsons and M. Sc

ISSN:1477-9226    

DOI:10.1039/a901425j

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1999, 1537–1538 1537 Syntheses and single-crystal structures of novel soluble phosphonato- and phosphinato-bridged titanium oxo alkoxides Gilles Guerrero,a Michael Mehring,a P. Hubert Mutin,*a Françoise Dahan b and André Vioux a a UMR CNRS 5637, Chimie Moléculaire et Organisation du Solide, Université de Montpellier II, cc 007, 34095 Montpellier cedex 5, France. E-mail: mutin@crit.univ-montp2.fr b UPR CNRS 8241, Laboratoire de Chimie de Coordination du CNRS, 205 Route de Narbonne, 31077 Toulouse cedex 4, France Received 25th March 1999, Accepted 9th April 1999 The reactions of PhP(O)(OH)2 or Ph2P(O)OH with Ti(OPri)4 in DMSO give the soluble tetranuclear complexes [Ti4(Ï3-O)(OPri)5(Ï-OPri)3(PhPO3)3]?DMSO 1 or [Ti(Ï3-O)- (OPri)(Ph2PO2)]4?0.5DMSO 2, the first examples of phosphonato- and phosphinato-bridged titanium oxo alkoxides, which have been characterised by single-crystal X-ray diffraction.Metal oxo alkoxides are not only models of the molecular species present at the initial stages of the sol–gel processing of metal alkoxides, but they also serve as molecular building blocks for the design of oxide materials by sol–gel processing and metal–organic chemical vapour deposition (MOCVD).1–3 The reactivity towards hydrolysis and condensation of metal oxo alkoxides is lower than in the parent alkoxides, and may be further decreased via replacement of some of the alkoxide ligands by carboxylate, b-diketonate or sulfonate ligands.Furthermore, the use of such ligands allows the introduction of organic functionalities for organic–inorganic hybrid materials applications.4,5 Recently, we proposed the use of organophosphorus acids as coupling molecules to prepare organic–inorganic hybrids by a sol–gel route, as the P–C bond is stable towards hydrolysis and P–OH groups readily condense with M–OR groups. A twostep synthesis was used, involving first the condensation between the organophosphorus acid and a metal alkoxide, followed by hydrolysis–condensation of the remaining alkoxy groups.6 When titanium isopropoxide and phenylphosphonic acid were used as precursors 31P NMR investigations pointed to the formation of a soluble intermediate.In the present work, we report the structural characterization of this intermediate and of another compound obtained starting from titanium isopropoxide and diphenylphosphinic acid (Scheme 1).To our knowledge, compounds 1 and 2 are the first examples of titanium oxo alkoxides modified by tridentate phosphonate (RPO3 22) or bidentate phosphinate (R2PO2 2) ligands, whereas modification by bidentate carboxylate groups has been extensively studied.7,8 Layered titanium phosphates and phosphonates are well known,9 and polymeric titanium alkoxo-phosphinates have been reported;10 however, very few molecular structures of titanium complexes with phosphato, 11,12 phosphonato 13,14 or phosphinato 15 ligands have been reported to date, whereas these ligands have been widely used to synthesise hybrid polynuclear oxo anions such as vanadates 16 and molybdates.17 The Ti–O–P bonds in 1 and 2 result from the condensation of P–OH and Ti–OPri groups.11 Most likely, the Ti–O–Ti bonds are formed by partial hydrolysis of Ti–OPri groups as nonhydrolytic condensation with elimination of an ether would require temperatures above 150 8C.In the case of 1, the sources of water could be DMSO and/or phenylphosphonic acid even after careful drying,15 although condensation of P–OH groups in the presence of Ti(OPri)4 cannot be completely ruled out.Nevertheless, in order to obtain 2 in a good yield, water had to be added to the reaction mixture.† The molecular structures of 1 and 2 and selected bond distances and angles are given in Figs. 1 and 2.‡ Compound 1 is made up of discrete clusters that consist of four titanium atoms, three tridentate phosphonato groups, three m-isopropxy groups and one m3-oxygen atom.The Ti–O–P core is best described as being based on a six-membered Ti3(m-OPri)3 ring adopting a chair conformation with alternating isopropoxy groups and six-coordinated Ti atoms; these Ti atoms are linked via a m3-O atom [Ti–m3-O–Ti 104.49(8)–106.08(8)8] leading to a Ti3O4 fragment. The markedly distorted octahedral geometry at Ti(1), Ti(2), and Ti(3) is built up by two bridging m-OPri groups [Ti–O 2.014(2)–2.044(2) Å], one terminal OPri group [Ti–O 1.770(2)–1.781(2) Å], two phosphonato-oxygen atoms [Ti–O 1.947(2)–1.984(2) Å] and the m3-oxygen atom which is trans to the terminal OPri groups [Ti–O 1.949(2)–1.976(2) Å].The trans O–Ti–O angles are in the range 162.27(8)–175.49(8)8. Each phosphonato group bridges two Ti atoms of the six membered ring and Ti(4) [Ti(4)–O 1.955(2)–2.095(2) Å]. In addition to these three phosphonato-oxygen atoms, the distorted octahedral geometry at Ti(4) [O–Ti–O trans angles 171.24(8)– 172.56(8)8] involves two terminal OPri groups [Ti(4)–O(18)/ O(19) 1.833(2)/1.786(2) Å] and a coordinated DMSO molecule [Ti(4)–O(4) 2.059(2) Å].The molecular structure of 2 is best described as a distorted cube consisting of a Ti4O4 core with the titanium and oxygen atoms occupying alternating corners. The Ti atoms show a distorted octahedral geometry [trans O–Ti–O angles 161.61(7)– 177.29(7)8] built up by one terminal OPri group [Ti–O 1.759(2)– 1.790(2) Å], two phosphinato oxygen atoms [Ti–O 2.002(2)– 2.040(2) Å] and three unsymmetrically bonded m3-oxygen atoms.The Ti–m3-O bonds trans to the isopropxy groups are Scheme 11538 J. Chem. Soc., Dalton Trans., 1999, 1537–1538 slighlty longer [Ti–m3-O 2.107(2)–2.163(2) Å] than those trans to the phosphinato oxygen atoms [Ti–m3-O 1.888(2)–1.938(2) Å]. Four sides of the Ti4O4 cube are capped by four phosphinato groups, each bridging two Ti atoms. This type of structure has been proposed for oxotitanium compounds modified by phosphonato 12 or carboxylato ligands.18 The basic structural arrangement of 1 and 2 is retained in solution, as demonstrated by the 31P NMR data in solution and Fig. 1 Molecular structure of 1. Thermal ellipsoids are shown at 60% probability. For clarity, hydrogens atoms are omitted and carbon atoms are represented by open circles. Selected interatomic distances (Å) and bond angles (8) are: Ti(1)–O(1) 1.953(2), Ti(1)–O(5) 1.976(2), Ti(1)– O(6) 2.014(2), Ti(1)–O(8) 1.781(2), Ti(4)–O(3) 1.955(2), Ti(4)–O(16) 2.095(2), Ti(4)–O(17) 2.040(2); O(1)–Ti(1)–O(7) 162.33(8), O(2)–Ti(1)– O(6) 165.22(8), O(5)–Ti(1)–O(8) 175.37(8), O(3)–Ti(4)–O(4) 172.56(8), O(16)–Ti(4)–O(18) 171.24(8), O(17)–Ti(4)–O(19) 171.83(8). Fig. 2 Molecular structure of 2. Thermal ellipsoids are shown at 60% probability. For clarity, hydrogens atoms are omitted and carbon atoms are represented by open circles. Selected interatomic distances (Å) and bond angles (8) are: Ti(1)–O(1) 1.903(2), Ti(1)–O(2) 2.160(2), Ti(1)– O(3) 1.938(2), Ti(1)–O(5) 2.004(2), Ti(1)–O(7) 2.014(2), Ti(1)–O(13) 1.765(2); O(1)–Ti(1)–O(7) 164.32(7), O(2)–Ti(1)–O(13) 177.29(7), O(3)–Ti(1)–O(5) 162.03(7).in the solid-state.† In the field of organic–inorganic hybrids, these clusters present two main interests: first, as intermediates in the sol–gel route to hybrids that we are developing, secondly, as novel building blocks for the preparation of nanostructured hybrid materials.Notes and references † Syntheses: [Ti4(m3-O)(OPri)5(m-OPri)3(PhPO3)3]?DMSO 1. Ti(OPri)4 (3.56 g, 12.52 mmol) was added to a solution of PhP(O)(OH)2 (1.00 g, 6.33 mmol) in 5 mL of dried DMSO, resulting in a cloudy mixture. After several hours a clear solution was obtained. Colourless crystals of 1 were obtained from this solution after several days. These crystals were filtered oV, washed with two 5 mL portions of Et2O and dried in vacuo giving 1.02 g [40% yield based on PhP(O)(OH)2].Colourless needles suitable for single-crystal X-Ray diVraction were recrystallised from a concentrated DMSO solution and mounted in inert oil. Calc. for C44H77O19P3STi4: C, 43.08; H, 6.78; P, 7.57; S, 2.61; Ti, 15.61. Found: C, 48.85; H, 6.52; P, 8.20; S, 3.22; Ti, 15.80. 31P NMR: d 8.45, 8.65 (CH2Cl2, 200 MHz), 8.0, 7.0 (solid-state, 400 MHz). [Ti(m3-O)(OPri)- (Ph2PO2)]4?0.5DMSO 2. To a solution of Ti(OPri)4 (1.00 g, 3.52 mmol) in 15 mL of dried DMSO was added dropwise a solution of Ph2P(O)OH (767 mg, 3.52 mmol) and H2O (34 mL, 1.88 mmol) in 15 mL of DMSO.After stirring at room temperature for 12 h, the resulting cloudy mixture was heated to 80 8C until a clear solution was obtained, which was then allowed to cool slowly to ambient temperature. The crystallised colourless needles were filtered oV, washed with 3 mL dried DMSO and dried in vacuo giving 720 mg (58% yield) of 2. X-Ray quality crystals were obtained from a dilute DMSO solution and mounted in inert oil.Calc. for C61H71O16.5P4S0.5Ti4: C, 52.35; H, 5.48; P, 8.85; S, 1.15; Ti, 13.68. Found: C, 51.98; H, 6.30; P, 10.10; S, 1.37; Ti, 15.10. 31P NMR: d 33.0 (CH2Cl2, 200 MHz), 33.3 (solid-state, 400 MHz). ‡ Data for both structures were collected on a STOE-IPDS with Mo-Ka radiation (l = 0.71073 Å). Crystal data for 1 : C44H77O19P3STi4 at 160 K, M = 1226.63, monoclinic, P21/n (no. 14), a = 11.6668(13), b = 13.9198(12), c = 35.885(3) Å, b = 90.533(15)8, V = 5827.5(10) Å3, Dc = 1.398 g cm23, Z = 4, m = 0.712 mm21.Of the 27893 reflections, 5980 unique reflections were used in the final least-squares refinement on Fo 2 for 640 variable parameters to yield wR(all) = 0.0549 and R(obs) = 0.0249 for 4478 observed reflections >4s(Fo). For 2 C61H71O16.5P4S0.5Ti4 at 180 K, M = 1399.69, monoclinic, P21/c (no. 14), a = 12.4983(10), b = 16.1317(14), c = 33.706(3) Å, b = 97.141(10)8, V = 6743.1(10) Å3, Dc = 1.379 g cm23, Z = 4, m = 0.630 mm21.Of the 52425 reflections, 10165 unique reflections were used in the final leastsquares refinement on Fo 2 for 793 variable parameters to yield wR(all) = 0.0607 and R(obs) = 0.0240 for 6923 observed reflections >4s(Fo). CCDC reference number 186/1420. See http://www.rsc.org/ suppdata/dt/1999/1537/ for crystallographic files in .cif format. 1 L. G. Hubert-Pfalzgraf, Polyhedron, 1994, 13, 1181. 2 R. C. Mehrotra and A. Singh, Chem.Soc. Rev., 1996, 25, 1. 3 V. W. Day, T. A. Eberspacher, W. G. Klemperer and C. W. Park, J. Am. Chem. Soc., 1993, 115, 8469. 4 P. Judeinstein and C. Sanchez, J. Mater. Chem., 1996, 6, 511. 5 U. Schubert, N. Hüsing and A. Lorenz, Chem. Mater., 1995, 7, 2010. 6 P. H. Mutin, C. Delenne, D. Medoukali, R. Corriu and A. Vioux, in Mater. Res. Soc. Symp. Proc., 1998, 519, 345. 7 R. Papiernik, L. G. Hubert-Pfalzgraf, J. Vaissermann and M. Goncalves, J. Chem. Soc., Dalton Trans., 1998, 2285. 8 T. J. Boyle, M. A. Todd, C. J. Tafoya and B. L. Scott, Inorg. Chem., 1998, 37, 5588. 9 A. Clearfield, Prog. Inorg. Chem., 1998, 47, 371. 10 G. H. Dahl and B. P. Block, Inorg. Chem., 1967, 6, 1439. 11 C. G. Lugmair and T. D. Tilley, Inorg. Chem., 1998, 37, 1821. 12 D. L. Thorn and R. L. Harlow, Inorg. Chem., 1992, 31, 3917. 13 J. R. Errington, J. Ridland, K. J. Willett, W. Clegg, R. A. Coxall and S. L. Heath, J. Organomet. Chem., 1998, 550, 473. 14 M. G. Walawalkar, S. Horchler, S. Dietrich, D. Chakraborty, H. W. Roesky, M. Schafer, H. G. Schmidt, G. M. Sheldrick and R. Murugavel, Organometallics, 1998, 17, 2865. 15 S. A. A. Shah, H. Dorn, J. Gindl, M. Noltemeyer, H.-G. Schmidt and H. W. Roesky, J. Organomet. Chem., 1998, 550, 1. 16 J. Zubieta, Comments Inorg. Chem., 1994, 16, 153. 17 Y. D. Chang and J. Zubieta, Inorg. Chim. Acta, 1996, 245, 177. 18 X. Lei, M. Shang and T. P. Fehlner, Organometallics, 1997, 16, 5289. Communication 9/02407G

ISSN:1477-9226    

DOI:10.1039/a902407g

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1999, 1539–1540 1539 A new route to (N)n-donor functionalised phosphines; novel homo- and hetero-nuclear complexes of a phosphino-substituted triazacyclononane ligand Scott E. Watkins, Xinhao Yang, Donald C. Craig and Stephen B. Colbran * School of Chemistry, University of New South Wales, Sydney, NSW 2052, Australia. E-mail: S.Colbran@unsw.edu.au Received 10th March 1999, Accepted 6th April 1999 A phosphino-substituted triazacyclononane ligand, oPtacn, has been prepared by a new, potentially versatile route, the complexes [CuI(oPtacn)][PF6], [PtIICl2{H- (oPtacn)}2][PF6]2 and [PtIICl2{(oPtacn)[CuII(OAc)]}2][PF6]2 made, and the crystal structure of [CuI(oPtacn)][PF6] determined.Heterobimetallic systems oVer prospects for advantageous synergistic eVects; for example, catalysis of olefin hydroformylation is markedly enhanced when a separate cobalt complex is added to a bis(phosphine)palladium catalyst.1 Discrete systems comprised of two diVerent metal ions stabilised by binucleating ligands have the potential to further enhance these cooperative eVects.Binucleating ligands possessing both (N)n and P-donor domains are among the more interesting because they can exhibit site selectivity for “hard” and “soft” metal ions.2,3 However, whereas the synthesis and coordination chemistry of functionalised derivatives of (N)n-donor macrocyclic ligands such as 1,4,7-triazacyclononane has attracted much recent attention,4 only a handful of nitrogen macrocycles with phosphine pendants have been reported.2,3,5,6 Moreover, most have flexible N-alkyl “arms” linking the (N)n and P-donor domains3,5,6 and, as a result, only mononuclear complexes with the metal ion bound by both the phosphine and the macrocycle have been isolated.Indeed the only crystallographically characterised metal complex of this type of ligand shows 1-(diphenylphosphinopropyl)-1,4,7-triazacyclononane coordinated to a zinc ion through the phosphine and all three amine groups.5 The formation of dimeric, heteronuclear complexes of a triphenylphosphine-pendant C-tethered cyclam has been reported.2 However, the synthesis of the ligand is lengthy, low in yield and not readily adaptable to other macrocycles.We present here an example of a new, potentially general route to (N)n-donor functionalised phosphine ligands via the reductive amination of secondary amine groups with (diphenylphosphino) benzaldehydes.This method should allow functionalisation of all or some of the secondary nitrogen atoms in a wide range of macrocycles and indeed ought to be applicable to any secondary amine. The new phosphino-macrocycle, oPtacn, was synthesised by the reductive amination of 1,4- diisopropyl-1,4,7-triazacyclononane 7a with 2-(diphenylphosphino) benzaldehyde 7b using sodium triacetoxyborohydride in 1,2-dichloroethane, Scheme 1. The product, oPtacn, was contaminated by some unreacted aldehyde but was easily purified by treatment with ammonium hexafluorophosphate in methanol.Recrystallisation from acetonitrile–diethyl ether gave the pure hemihydrate of the monoprotonated ligand hexafluorophosphate, [H(oPtacn)][PF6]?0.5H2O 1, in 72% yield.† Coordination complexes can be prepared directly from 1, Scheme 1. For example, reaction of 1 with one equivalent of [Cu(MeCN)4][PF6] gave [Cu(oPtacn)][PF6] 2 as a white powder which was recrystallised from acetonitrile–diethyl ether.Elemental analysis and the ES-mass spectrum confirm the formulation of 2.‡ The crystal structure, Fig. 1, reveals that oPtacn acts as a mononucleating N3P-donor ligand to the copper(I) ion.§ The coordination geometry for the copper centre is distorted tetrahedral (e.g. the sum of the six bond angles about the copper ion is 6318 compared to 6578 for a perfect tetrahedron) with three similar Cu–N bond lengths [2.150(4), 2.132(4), 2.106(3) Å] and a Cu–P bond length of 2.141(1) Å close to that observed in [CuL(PPh3)][BF4], (L = 1,3,5-triisopropyl-1,3,5- triazacyclohexane).8 Protonation of oPtacn protects the macrocyclic domain and enables selective binding of the phosphine group to second or third row transition metals.For example, reaction of [PtCl2(PhCN)2] with 1 (2 equivalents) gave mononuclear [PtCl2{H(oPtacn)}2][PF6]2, 3, in 87% yield.¶ Noteworthy data for the complex include: a molecular ion (M21) peak at 621 m/z in the ES mass spectrum; a significant downfield shift for the benzylic protons in the 1H NMR spectrum compared to the ligand salt as well as changes in the aromatic region, all consistent with coordination of the phosphine to the platinum ion; also in the 1H NMR spectrum, methylene resonances virtually unshifted from those for 1 reveal no change for the protonated macrocyclic domain; and a sharp singlet at d 15.44 flanked by 31P–195Pt satellites with a trans coupling constant (1JPPt 2597 Hz) Scheme 1 (i) Na[BH(OAc)3], 1,2-dichloroethane; (ii) [NH4][PF6], MeOH; (iii) [Cu(MeCN)4][PF6], CH2Cl2; (iv) [PtCl2(PhCN)2], CH2Cl2; (v) NEt3 1 [Cu2(OAc)4(H2O)2], MeCN.HC Ph2P O N N P N N N Cu N H 3 2+ 2+ P N N N Pt Cl H P N N N Cl H P N N N H P Pt Cl P Cl N N N Cu O O N N N Cu O O + + + (i) (ii) 1 2 4 (iii) (v) (iv)1540 J. Chem. Soc., Dalton Trans., 1999, 1539–1540 in the 31P NMR spectrum. In sum this evidence leads to the structure for 3 in Scheme 1.Addition of base frees the macrocyclic centres in 3 and allows selective formation of trimers. For example, reaction of 3 with triethylamine and [Cu2(OAc)4(H2O)2] gave [PtCl2{oPtacn- [Cu(OAc)]}2][PF6]2, 4, in 56% yield. The base is essential in the preparation of this heterometallic trimer; whilst the coordination of “soft” Pt(II) and Cu(I) ions by the phosphine group in oPtacn is relatively facile, coordination of the “hard” Cu(II) ion requires deprotonation of the macrocyclic domain and, therefore, is pH dependent.Partial analytical data for C, H and N and data for the Cu:P:Pt ratio agree with the formulation of 4, as does the ES mass spectrum which shows a prominent peak at 743 m/z for the molecular ion (M21).|| The UV/VIS spectrum reveals a band at 661 nm with a distinct low energy tail and an axial EPR spectrum is observed. These spectroscopic data closely match those of a similar crystallographically-characterised N3O2-coordinated copper(II) complex9 and are indicative for isolated mononuclear copper(II) centres with distorted square-pyramidal coordination.The structure proposed for 4 in Scheme 1 is consistent with these results. We have shown that reductive amination of nitrogen macrocycles with (diphenylphosphino)benzaldehydes provides a convenient synthesis of novel phosphino-substituted macrocycles. The ligating properties of multinucleating ligands of this type can be tuned by varying the relative orientation of the phosphine and macrocycle groups by changing from ortho to meta or para aryl substitution.Studies of other heterometallic oligomers, including those bridged by the obligatory binucleating meta-analogue of oPtacn [prepared from 3-(diphenylphosphino) benzaldehyde 7c] and, as well, with new ligands derived from other secondary amines, are underway. The challenge to demonstrate novel reactivities for these heterometallic oligomers remains. Acknowledgements We are grateful for support from the Australian Research Fig. 1 Drawing of the complex cation 2 showing the 20% thermal ellipsoids. Council and for an Australian Postgraduate Award (to S. E. W.). Notes and references † 1, [H(oPtacn)][PF6]?0.5H2O (Found: C, 57.91; H, 6.47; N, 6.60. C31H43N3P2F6?0.5H2O requires C, 57.94; H, 6.90; N, 6.54%); dP(CDCl3) 214.22 (s), 2143.78 [sept, J(PF) 707 Hz]; dH(CDCl3) 7.37–7.32 (8 H, m), 7.24–7.17 (5 H, m), 6.90–6.85 (1 H, m), 3.97 (2 H, d), 3.11 (2 H, sept), 3.00–2.93 (4 H, m), 2.79 (8 H, br s), 1.21 (6 H, d), 1.14 (6 H, d); m/z (ES-MS) 488 {[H(oPtacn)]1}. ‡ 2, [Cu(oPtacn)][PF6] (Found: C, 52.62; H, 6.41; N, 6.10.C31H42- N3P2F6Cu?0.5H2O requires C, 52.80; H, 6.15; N, 5.96%); dP[(CD3)2CO] 20.3 (br s), 2142.00 [sept J(PF) 707 Hz]; dH[(CD3)2CO] 7.65–7.50 (8 H, m), 7.50–7.45 (5 H, m), 6.94 (1 H, t), 3.88 (2 H, s), 3.39 (2 H, sept), 3.15–2.55 (12 H, m), 1.47 (6 H, d), 1.13 (6 H, d); m/z (ES-MS) 550 {[Cu(oPtacn)]1}. § Crystal data for 2: C31H42CuF6N3P2, M = 696.2, monoclinic, space group P21/c, a = 15.091(6), b = 14.532(3), c = 17.990(9) Å, b = 123.87(2)8, U = 3276(2) Å3, Z = 4, Dc = 1.41 g cm21, Enraf-Nonius CAD-4 diVractometer, m(Mo-Ka: l = 0.71073 Å) = 8.21 cm–1, F(000) = 1448.0, T = 294 K, final R = 0.037, Rw = 0.051 for 2821 observed data [I > 3s(I), 2q < 468].CCDC reference number 186/1414. See http:/www.rsc.org/suppdata/dt/1999/1539/ for crystallographic files in .cif format. ¶ 3, [PtCl2{H(oPtacn)}2][PF6]2 (Found: C, 48.07, H, 5.33, N, 5.13.C62H86N6P4Cl2F12Pt?H2O requires C, 48.06, H, 5.72, N, 5.42%); dP[(CD3)2CO] 15.44 [s, J(PPt) 2597], 2142.05 [sept, J(PF) 707 Hz]; dH[(CD3)2CO] 7.95–7.75 (10 H, m), 7.65–7.45 (14 H, m), 7.29 (2 H, t), 7.09 (2 H, q), 4.50 (4 H, s), 3.35 (4 H, sept), 3.27 (4 H, m), 3.10–2.60 (20 H, m), 1.30 (12 H, d), 1.18 (12 H, d); m/z (ES-MS) 1387 ([PtCl2{H- (oPtacn)}2 1 PF6]1), 621 ([PtCl2{H(oPtacn)}2]21). || 4, [PtCl2{(oPtacn)[Cu(OAc)]}2][PF6]2 (Found: C, 44.08, H, 4.99, N, 4.43.C66H90N6O4P4Cl2F12Cu2Pt?H2O requires C, 44.17, H, 5.17, N, 4.68%); inductively-coupled-plasma analysis (±10%): ratio Cu:P:Pt = 2.0 : 3.6 :0.9; lmax/nm (MeCN) 661 (e/dm3 mol21 cm21 797); EPR (MeCN, 77 K): g|| 2.25 (A|| 164 G), g^ 2.07; dP[(CD3)2CO] 2141.97 [sept, J(PF) 707 Hz]; m/z (ES-MS) 743 ([PtCl2- {(oPtacn)[Cu(OAc)]}2]21). 1 Y. Ishii, K. Miyashita, K. Kamita and M. Hidai, J. Am. Chem. Soc., 1997, 119, 6448. 2 E. Kimura, Y. Kodama, M.Shionoya and T. Koike, Inorg. Chim. Acta, 1996, 246, 151. 3 A. Carroy, C. R. Langick, J.-M. Lehn, K. E. Matthes and D. Parker, Helv. Chim. Acta, 1986, 69, 580. 4 K. P. Wainwright, Coord. Chem. Rev., 1997, 166, 35; L. Spiccia, B. Graham, M. T. W. Hearn, G. Lazarev, B. Moubaraki, K. S. Murray and E. R. T. Tiekink, J. Chem. Soc., Dalton Trans., 1997, 4089; S. Mahapatra, S. Kaderli, A. Llobet, Y.-M. Neuhold, T. Palanché, J. A. Halfern, V. G. Young, Jr., T. A. Kaden, L. Que, Jr., A. D. Zuberbühler and W. B. Tolman, Inorg. Chem., 1997, 36, 6343; L. J. Farrugia, P. A. Lovatt and R. D. Peacock, J. Chem. Soc., Dalton Trans., 1997, 911; A. Sokolowski, J. Müller, T. Weyhermüller, R. Schnepf, P. Hildebrandt, K. Hildenbrand, E. Bothe and K. Wieghardt, J. Am. Chem. Soc., 1997, 119, 8889; M. Di Vaira, F. Mani and P. Stoppioni, Inorg Chim. Acta, 1998, 273, 151. 5 D. Ellis, L. J. Farrugia, D. T. Hickman, P. A. Lovatt and R. D. Peacock, Chem. Commun., 1996, 1817. 6 H. Hope, M. Viggiano, B. Moezzi and P. P. Power, Inorg. Chem., 1984, 23, 2550. 7 (a) J. A. Halfern and W. B. Tolman, Inorg. Synth, 1998, 32, 75; (b) T. B. Rauchfuss and D. A. Wrobleski, Inorg Synth., 1982, 21, 175; (c) G. P. Schiemenz and H. Kaack, Liebigs Ann. Chem., 1973, 1480. 8 R. D. Köhn, G. Seifert and G. Kociok-Köhn, Chem. Ber., 1996, 129, 1327. 9 P. Bradford, R. C. Hynes, N. C. Payne and C. J. Willis, J. Am. Chem. Soc., 1990, 112, 2647. Communication 9/01885I

ISSN:1477-9226    

DOI:10.1039/a901885i

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON COMMUNICATION J. Chem. Soc., Dalton Trans., 1999, 1541–1542 1541 Compounds containing Î3,Û2-Sb] C bonds: synthesis and structural characterisation of the first stiba-enol, Mes*C(O)Sb] C(OH)Mes* (Mes* 5 C6H2But 3-2,4,6) and a 2,3-distibabutadiene, {Mes(Me3SiO)C] Sb}2 (Mes 5 C6H2Me3-2,4,6) † Cameron Jones,*a Jonathan W. Steed b and Ryan C. Thomasa a Department of Chemistry, University of Wales, CardiV, PO Box 912, Park Place, CardiV, UK CF1 3TB. E-mail: jonesca6@cardiV.ac.uk b Department of Chemistry, King’s College London, Strand, London, UK WC2R 2LS Received 8th April 1999, Accepted 20th April 1999 The reactions of [Li{Sb(SiMe3)2}] with RCOCl, R 5 C6H2But 3-2,4,6 (Mes*) or C6H2Me3-2,4,6 (Mes), afford mixtures of the 2,3-distibabutadienes, {R(Me3SiO)C]] Sb}2, and the 2-stiba-1,3-dionatolithium complexes, [Li{OC(R)- SbC(R)O}(DME)n], n 5 1 or 0.5, the latter of which (R 5 Mes*) can be protonated to give the first stiba-enol, Mes*C(O)Sb]] C(OH)Mes*, which has been structurally characterised.Since the preparation of the first thermally stable phosphaalkyne, P]] ] CBut, in 1981 the field of low coordination phosphorus chemistry has become well established.1 Not surprisingly, the chemistry of compounds containing As–C multiple bonds was slower to develop but is now relatively well explored.2 By contrast, there is a paucity of knowledge of analogous low coordination antimony compounds which probably results from their inherent thermal instability.In fact, to date there is only one structurally characterised example of a compound, {R(Me3SiO)C]] Sb}2 1 [R = C6H2But 3-2,4,6 (Mes*)], that contains largely localised Sb–C double bonds,3 though related compounds have recently been implicated as reactive intermediates in the formation of stibacycles.4 This remarkably stable compound was prepared in low yield from the reaction of Mes*COCl with [Li{Sb(SiMe3)2}], a surprising result considering that the analogous reaction of ButCOCl with [Li{Sb(SiMe3)2}] aVords a high yield of the delocalised 2- stiba-1,3-dionatolithium complex, [{[Li{OC(But)SbC(But)O}- (DME)0.5]2}•],5 the coordination chemistry of which we are currently investigating.6 Herein we report that a stibadionatolithium complex is, indeed, the major reaction product in the preparation of 1 and that a similar product mixture is obtained in the reaction of the less hindered acyl chloride MesCOCl (Mes = C6H2Me3-2,4,6) with [Li{Sb(SiMe3)2}]. In addition, this work has led to the synthesis and structural characterisation of the first stiba-enol which, in the solid state, contains a rare example of a localised Sb–C double bond.The product mixtures obtained from the treatment of [Li{Sb- (SiMe3)2}] with 1 equivalent of either Mes*COCl or MesCOCl were extracted with hexane to give the distibabutadienes, 1 and 2, in low yield (18% and 5% respectively) after concentration of the extracts (Scheme 1).The hexane insoluble fractions of the reaction mixtures were further extracted with diethyl ether aVording the 2-stiba-1,3-dionatolithium complexes, 3 and 4, in moderate yields (38% and 45%) after recrystallisation. Treatment of a diethyl ether solution of 3 with 1 equivalent of anhydrous HCl, followed by recrystallisation from diethyl ether gave red crystals of the light sensitive stiba-enol, 5, in high yield (97%). † Supplementary data available: synthetic and spectroscopic details for compounds 2–5.For direct electronic access see http://www.rsc.org/ suppdata/dt/1999/1541/, otherwise available from BLDSC (No. SUP 57542, 4 pp.) or the RSC Library. See Instructions for Authors, 1999, Issue 1 (http://www.rsc.org/dalton). Compound 5 is stable in toluene solutions for only 15 minutes at room temperature. As a result, spectroscopic data (see SUP 57542) for the compound were collected at 0 8C and are consistent with it existing predominantly in the enol form in this solvent (cf.its As analogue 5). Evidence for this suggestion comes from its 1H NMR spectrum which displays a low field resonance at d 18.48 in the region normally associated with strongly hydrogen bonded alcoholic protons. The symmetry of this spectrum also suggests that 5 possesses a fully delocalised structure in solution in which the alcoholic proton is undergoing a rapid exchange between the two oxygen centres of the molecule. In the solid state 5 is more thermally stable (decomp. 103 8C) and its crystal structure ‡ (Fig. 1) confirms that it exists in the enol form but with localised Sb(1)–C(20) and C(1)–O(1) double bonds, the former of which compares well with those in 1 [2.056(10) Å] 3 and 2 [2.066(5) Å] (see below) but is considerably shorter than normal Sb–C single bonds {e.g. 2.225 Fig. 1 Molecular structure of Mes*C(O)Sb]] C(OH)Mes* 5. Selected bond lengths (Å) and angles (8): Sb(1)–C(20) 2.078(3), Sb(1)–C(1) 2.192(3), O(1)–C(1) 1.248(4), O(2)–C(20) 1.319(4), O(2)–H(2) 0.96(5), O(1) ? ? ? H(2) 1.61(5); C(20)–Sb(1)–C(1) 91.31(12), O(2)–H(2)–O(1) 172(5), O(1)–C(1)–C(2) 120.9(3), O(1)–C(1)–Sb(1) 121.2(2), C(2)–C(1)– Sb(1) 118.0(2), O(2)–C(20)–C(21) 117.9(3), O(2)–C(20)–Sb(1) 124.7(2), C(21)–C(20)–Sb(1) 117.2(2).Scheme 1 Reagents and conditions: i, RCOCl, DME, 18 h; ii, R = Mes*, HCl, Et2O, 0 8C, 2 h. O Sb O Li R R C Sb Sb C OSiMe3 Me3SiO R R O Sb O H Mes* Mes* (DME)n [Li{Sb(SiMe3)2}] + 1 R = Mes* 2 R = Mes 3 R = Mes*, n = 1 4 R = Mes, n = 0.5 5 i ii1542 J.Chem. Soc., Dalton Trans., 1999, 1541–1542 (average) in [But 3Sb?Fe(CO)4] 7}. The acute nature of the C–Sb–C angle in 5 [91.31(12)8] probably results from a signifi- cant degree of s-character for the hetero-atom lone pair. This is a common feature of other low coordinate Group 15 systems (e.g. RE]] ER, E = N, P, As, Sb, Bi) and has been found to be augmented with increasing molecular weight of the Group 15 element.8 The alcoholic proton H(2) was located from diVerence maps and refined isotropically.It is bonded to O(2) and appears to have a strong H-bonded interaction with O(1), the angle O(1)–H(2)–O(2) being 172(5)8. As has been suggested for 1 3 the unusual stability of 5 can probably be attributed to a combination of the steric protection aVorded by its bulky aryl substituents and the conjugated nature of the system. The distibabutadiene 2 (decomp. 105 8C) is not as thermally stable as its more sterically protected counterpart 1 (decomp. 213 8C) but is nevertheless stable in air at ambient temperature for days. Its molecular structure ‡ (Fig. 2) is similar to that of 1 and shows it to exist in the trans- form with the atoms C(10), Sb(1), Sb(1)9 and C(10)9 being necessarily co-planar. The Sb–C bond length is close to those in 1 and 5 (see above) and as with the C–Sb–C angle in 5 the sharp Sb–Sb–C angles in 2 [92.99(13), cf. 94.7(3)8 in 1 3] can be explained by a high degree of s-character for the Sb lone pairs.The 2-stibadionato lithium complexes, 3 and 4, are considerably more stable (3 decomp. 170, 4 decomp. 103 8C) than the only other example of such a compound, [{[Li{OC(But)SbC- (But)O}(DME)0.5]2}•] 6 (decomp. 65 8C).5 No crystallographic data were obtained for 4 but in the solid state it probably consists of oxygen and lithium bridged dimeric units linked by nonchelating, bridging DME molecules, as has been found for 6 and a number of related 2-arsa- and 2-phospha-dionatolithium complexes.5 Compound 3 on the other hand is probably monomeric in the solid state and has its Li centre chelated by a Fig. 2 Molecular structure of {Mes(Me3SiO)C]] Sb}2 2.Selected bond lengths (Å) and angles (8): Sb(1)–C(10) 2.066(5), Sb(1)–Sb(1)9 2.8018(8), O(1)–C(10) 1.377(5), Si(1)–O(1) 1.697(3); C(10)– Sb(1)–Sb(1)9 92.99(13), O(1)–C(10)–C(1) 110.6(4), O(1)–C(10)–Sb(1) 125.2(3), C(1)–C(10)–Sb(1) 124.3(3). DME molecule, as is the case for its As counterpart.9 These diVerences in the degree of association between 3 and 4 would be expected considering the bulk of the aryl substituent in 3.As is the case for 6, the symmetry of the solution state 1H and 13C NMR spectra of 3 and 4 suggest that the ligand backbones of these complexes are delocalised. We are currently exploring the use of 2 and 5 as ligands in inorganic synthesis and the utility of 3 and 4 as reagents for the transfer of the 2-stiba-1,3-dionate fragments onto other metal centres.We are also investigating the mechanisms of formation of 1–4. The results of these investigations will form the basis of forthcoming publications. Acknowledgements We gratefully acknowledge financial support from the EPSRC (studentship for R. C. T.). Notes and references ‡ Crystal data for 5: C38H59O2Sb, M = 669.60, orthorhombic, space group Pcab, a = 11.4906(2), b = 20.0826(4), c = 31.2742(5) Å, V = 7216.9(2) Å3, Z = 8, Dc = 1.233 g cm23, F(000) = 2832, m = 7.94 cm21, crystal 0.20 × 0.20 × 0.10 mm, radiation Mo-Ka (l = 0.71070 Å), T = 100(2) K, 50378 reflections collected.For 2: C26H40O2Sb2Si2, M = 684.26, monoclinic, space group P21/c, a = 10.704(2), b = 14.043(3), c = 10.889(2) Å, b = 109.57(3)8, V = 1542.2(5) Å3, Z = 2, Dc = 1.473 g cm23, F(000) = 1368, m = 18.48 cm21, crystal 0.20 × 0.20 × 0.10 mm, radiation Mo-Ka (l = 0.71070 Å), T = 100(2) K, 13007 reflections collected. All crystallographic measurements were made using an Enraf-Nonius Kappa-CCD diVractometer.Both structures were solved by direct methods and refined on F 2 by full matrix least squares (SHELX97)10 using all unique data. All non-hydrogen atoms are anisotropic with H-atoms [except H(2) in 5] included in calculated positions (riding model). Absorption corrections were carried out using Scalepack. 11 Final R (on F) were 0.0426 (5) and 0.0331 (2) and wR (on F 2) were 0.0838 (5) and 0.0940 (2) for I > 2s(I).CCDC reference number 186/1431. See http://www.rsc.org/suppdata/dt/1999/1541/ for crystallographic files in .cif format. 1 K. B. Dillon, F. Mathey and J. F. Nixon, in Phosphorus: The Carbon Copy, Wiley, Chichester, 1998, and refs. therein. 2 L. Weber, Chem. Ber., 1996, 129, 367 and refs. therein. 3 P. B. Hitchcock, C. Jones and J. F. Nixon, Angew. Chem., Int. Ed. Engl., 1995, 34, 492. 4 P. C. Andrews, C. L. Raston, B. W. Skelton, V.-A. Tolhurst and A. H. White, Chem. Commun., 1998, 575. 5 J. Durkin, D. E. Hibbs, P. B. Hitchcock, M. B. Hursthouse, C. Jones, J. Jones, K. M. A. Malik, J. F. Nixon and G. Parry, J. Chem. Soc., Dalton Trans., 1996, 3277 and refs. therein. 6 S. J. Black, D. E. Hibbs, M. B. Hursthouse, C. Jones and J. W. Steed, Chem. Commun., 1998, 2199. 7 A. L. Rheingold and M. E. Fountain, Acta Crystallogr., Sect. B, 1985, 41, 1162. 8 N. Tokitoh, Y. Arai, T. Sasamori, R. Okazaki, S. Nagase, H. Uekusa and Y. Ohashi, J. Am. Chem. Soc., 1998, 120, 433; N. Tokitoh, Y. Arai, R. Okazaki and S. Nagase, Science, 1997, 277, 78; N. C. Norman, Polyhedron, 1993, 12, 2431. 9 C. Jones and R. C. Thomas, unpublished work. 10 G. M. Sheldrick, SHELX97, University of Göttingen, 1997. 11 Z. Otwinowski and W. Minor, in Methods in Enzymology, ed. C. W. Carter and R. M. Sweet, Academic Press, New York, 1996. Communication 9/02791B

ISSN:1477-9226    

DOI:10.1039/a902791b

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON FULL PAPER J. Chem. Soc., Dalton Trans., 1999, 1543–1554 1543 Synthesis and electrochemistry of platinum complexes of hydroquinon-2-ylmethyl- and p-benzoquinon-2-ylmethyldiphenylphosphine Seri Bima Sembiring,a Stephen B. Colbran *b and Donald C. Craig b a Department of Chemistry, Faculty of Mathematics and Natural Science, University of North Sumatra, Medan, Indonesia b School of Chemistry, University of New South Wales, Sydney, NSW 2052, Australia. E-mail: s.colbran@unsw.edu.au Received 22nd January 1999, Accepted 22nd March 1999 Platinum(II) complexes of new (hydroquinon-2-ylmethyl)diphenylphosphine (PPh2thqH2) and diphenyl(quinon-2- ylmethyl)phosphine (PPh2tq) ligands have been studied. Reaction of (2,5-dimethoxybenzyl)diphenylphosphine (PPh2dmb) with [PtCl2(PhCN)2] aVorded [PtCl2(PPh2dmb)2] 1a.Metathesis reactions gave the bromo (1b) and iodo (1c) congeners. Deprotection of the hydroquinone groups in 1a using boron tribromide followed by treatment with base produced the O,P-chelated hydroquinonate phosphine complex, cis-[Pt(O,P-PPh2thqH)2] 2.Reaction of 2 with hydrobromic acid aVorded cis-[PtBr2(PPh2thqH)2] 3 which can be oxidised to give the quinone phosphine complex cis-[PtBr2(PPh2tq)2] 4. Unlike previously reported quinone phosphine complexes, 4 is robust and stable to hydrolysis; its reduction with excess of zinc and dilute hydrobromic acid produced cis-[Pt(O,P-PPh2thqH)2(ZnBr2)] 5. Crystal structure analyses of 2?2dmf, 4?0.5dcm and 5?2dmf were performed.The electrochemistries of 1b, 3, and 4 have been characterised by cyclic voltammetry and controlled potential electrolyses. Cyclic voltammograms of 4 in the presence of dilute hydrobromic acid exhibit a four-electron cathodic process, attributed to reduction to 3; those of 3 show an anodic process attributed to oxidation to 4. The electrochemistry of 4 under aprotic conditions is extraordinary. Although there are two well separated, pendant quinone substituents only a single one-electron reduction process is observed.The reduction aVords a radical species (6~2) which has been characterised by cyclic voltammetry, and by EPR and UV/Vis/NIR spectroscopy. It is argued from the available data that 6~2 is a novel platinum(IV) complex with bound hydroquinonate and semiquinonate (sq) groups, namely [PtBr2(O,P-PPh2thq)(O,P-PPh2tsq?)]2, and a possible mechanism for its remarkable formation is discussed. This paper describes the preparation and electrochemistry of some transition metal (specifically platinum) complexes of novel phosphine ligands with p-hydroquinone or p-quinone substituents.p-Quinones along with their redox products, p-semiquinones and p-hydroquinones, comprise perhaps the quintessential organic electron and hydrogen transfer systems.1 For example, electron transfer reactions between transition metal centres and p-quinone cofactors are vital for all life, occurring in key biological processes as diverse as the oxidative maintenance of biological amine levels,2,3 tissue (collagen and elastin) formation,3,4 photosynthesis 5,6 and aerobic (mitochondrial) respiration.6,7 Nevertheless, in comparison to the extensive co-ordination chemistry for chelate-stabilised, o-(hydro/semi)quinone s-donor ligands,8,9 there are relatively few transition metal s-complexes of p-(hydro/semi)quinone ligands,10 thus giving some incentive for this study.However, the main impetus was provided by the observations that small amounts of a p-quinone promoter can prolong the eYcacy and lifetime of certain homogeneous transition metal catalysts.Examples include Shell’s olefin-oligomerisation process 11 and the palladium-catalysed alternating copolymerisation of alkenes and carbon monoxide to aVord polyketones with unusual, useful properties of potential commercial value.12 In both processes, phosphine ligands are used to stabilise the catalytically active transition metal centres. The exact role(s) of the p-quinone promoter has not been established, but possibilities for its action include as a hydride acceptor, as an oxidant of “dead-end” lower oxidation-state species to regenerate active catalyst, and as oxidant promoting more of the catalytically active species.Recent results revealing that p-quinones can be used as hydrogen acceptors in cycles for transition metalcatalysed oxidations of organic substrates provide further incentive for this work.13 Moreover, p-quinones are commonly used in organic synthesis as organic dehydrogenation/oxidation reagents.1,14 Taken together the above results suggest potential uses for transition metal complexes of quinone phosphine ligands, including, for example, in homogeneous oxidations of organic substrates. In response, we instigated studies of the palladium and platinum co-ordination chemistry of the first phosphine ligands with p-hydroquinone or p-quinone substituents. Some early results have been communicated.15–19 Under rigorously aprotic conditions, oxidation of the hydroquinone phosphine complexes [MX2(PPh2hqH2)2] (M = Pd or Pt; PPh2hqH2 = 2-diphenylphosphino-1,4-hydroquinone; X = Cl, Br or I) produced the quinone phosphine complexes [MX2(PPh2q)2] (PPh2q = 2-diphenylphosphino-1,4-benzoquinone). The complexes are electrochemically active; for example, four-electron, two-proton reduction of the quinone phosphine complexes [MX2(PPh2q)2] aVorded the chelated hydroquinonate phosphine complexes cis-[M(O,P-PPh2hqH)2].15,20 However, our previously reported quinone phosphine complexes are susceptible to hydrolytic loss of the quinone group, probably because of the direct attachment of the quinone groups to the phosphorus atoms.16,19,20 This is a major limitation to the further development and use of these systems.For example, chemically reversible interconversion between the hydroquinone/quinone redox states is expected in buVered acidic aqueous solutions but can not be observed because the quinone phosphine ligands decompose.20 A more (hydrolytically) stable quinone1544 J.Chem. Soc., Dalton Trans., 1999, 1543–1554 phosphine ligand system was therefore sought. We thought that an intervening methylene group would isolate the phosphorus and redox-active hydroquinone/quinone centres leading to increased stability for the quinone phosphine complexes.To this end, we now report a study of the platinum co-ordination chemistry of new p-hydroquinone-/p-quinone phosphine ligands. Results and discussion Synthetic studies Ligands. Complexes of the ligands (hydroquinone-2-ylmethyl) diphenylphosphine (PPh2thqH2) and (p-benzoquinone- 2-ylmethyl)diphenylphosphine (PPh2tq, Chart 1) were targeted for this study. The entry point to these was provided by the protected precursor (2,5-dimethoxybenzyl)diphenylphosphine (PPh2dmb, Chart 1) which was prepared as follows. 2,5-Dimethoxybenzylbromide, obtained by methylation of 2,5-dihydroxytoluene with iodomethane–potassium carbonate followed by bromination of the intermediary 2,5-dimethoxytoluene with Nbromosuccinamide, was treated with magnesium filings that had been precrushed dry under nitrogen [precrushing is essential to avoid the Wurtz coupled product 21] to give the 2,5-dimethoxybenzyl Grignard reagent which was treated with chlorodiphenylphosphine at 280 8C to aVord PPh2dmb in excellent yield (88%).No attempt was made to obtain the free hydroquinone and quinone phosphines PPh2thqH2 and PPh2tq respectively, because we anticipated that complexes of hydroquinone/ quinone phosphines would be best prepared by first forming complexes of phosphines with a protected hydroquinone/ quinone substituent (PPh2dmb in the present case) and only then deprotecting the hydroquinone/quinone substituents(s). Platinum(II) complexes. Platinum(II) complexes of PPh2dmb were easily prepared.Reactions of two equivalents of PPh2dmb with [PtCl2(PhCN)2] in dichloromethane (dcm) gave good yields of the phosphine complex [PtCl2(PPh2dmb)2] 1a an oV- white solid, 95%, eqn. (1). Metathesis reactions of 1a in acetone [PtCl2(PhCN)2] 1 2 PPh2dmb æÆ [PtCl2(PPh2dmb)2] 1 2 PhCN (1) with a large excess of potassium bromide or potassium iodide aVorded the bromo (1b) and iodo (1c) analogues in near quantitative yields, eqn. (2). [PtCl2(PPh2dmb)2] 1 2 KX æÆ [PtX2(PPh2dmb)2] 1 2 KCl (2) The elemental analyses of complexes 1a–1c were all consistent with their respective formulations.The 31P-{1H} and 1H NMR spectra reveal that they were isolated as mixtures of trans and cis stereoisomers. The peaks for the cis and trans isomers in the 31P-{1H} NMR spectra were directly assigned from the 31P–195Pt coupling constants {for [PtX2(PR3)2] complexes, 1J(PPt) typically is less than 2800 Hz for trans isomers and greater than 3000 Hz for cis isomers}.22–24 Notably, in 1H NMR spectra, the a-methylene resonances of the trans isomers [where J(PP) is large (>500 Hz)] appear as “virtual” 1:2:1 triplets whereas “filled in” 1 : 1 doublets were observed for the cis isomers [where J(PP) is typically small (<80 Hz)].25 Assign- Chart 1 OH OH Ph2P O O Ph2P OMe OMe Ph2P PPh2thqH2 6 6 3 PPh2tq 4 4 3 3 4 PPh2dmb 6 ments of the other peaks were made on the basis of their relative intensities compared to those of the methylene resonances.The relative amounts of cis and trans isomers deduced from 31P-{1H} and 1H NMR spectra were consistent in all cases. Treatment of complex 1a with BBr3 followed by methanolic sodium carbonate gave a pale yellow powder which analyses as cis-[Pt(O,P-PPh2thqH)2]?1.5H2O (2?1.5H2O), Scheme 1. The complex is poorly soluble and dimethylformamide (dmf) was the only solvent found in which it readily dissolved. Although it initially dissolved in dimethyl sulfoxide, within a few seconds a yellow fibrous powder began to precipitate.Elemental analysis of this is in accord with the formulation 2?2H2O?2(CH3)2SO suggesting that 2 forms a dimethyl sulfoxide solvate hydrate that, unusually, is insoluble in dimethyl sulfoxide. Formation of the hydrate provides a convenient method for purification of 2. The electrospray mass spectrum of 2 shows a prominent peak at m/z 810 for the [M 1 H]1 ion. A broad hydroxy band at 3283 cm21 is seen in the IR spectrum.The NMR spectra are unexceptional and confirm that 2 has cis stereochemistry: for example, the 1H NMR spectrum reveals a “filled in” cis methylene doublet at d 4.11 and the 31P-{1H} NMR spectrum exhibits a singlet at d 28.08 flanked by a satellite doublet with a cis 31P–195Pt coupling constant of 3822 Hz. Although not tested, it is likely that the initial product of the reaction of 1a with BBr3 followed by methanol quenching of the borate intermediates is trans-[PtBr2(PPh2thqH2)2], and that reaction of this with sodium carbonate leads to closure of the O,P-chelate rings to aVord 2?nH2O.A single crystal analysis of complex 2?2dmf was undertaken on pale yellow crystals grown from diethyl ether saturated–dmf solution. The molecular structure and labelling scheme is shown in Fig. 1. The platinum(II) ion is planar and cis-coordinated by two O,P-chelate hydroquinonate phosphine ligands. In the crystal structure, a molecule of dmf hydrogen bonds to the terminal hydroquinonate OH group of each ligand [O (dmf) ? ? ? O (hydroquinonate) 2.62, 2.70 Å].The significant distortion from normal square planar geometry around the platinum ion [P–Pt–P 101.0(1) and O–Pt–O = 79.7(1)8] arises because of the constrained ligand “bite” [(P–Pt–O)average 89.8(1)8]. Other key bond lengths and angles (Table 1) are unexceptional and similar to those found in other platinum(II) complexes with a cis-O2P2 donor set.24 Reaction of complex 2 with hydrobromic acid gave cis- [PtBr2(PPh2thqH2)2] 3 in good isolated yield (>80%), Scheme 1.Analytical and spectroscopic data for 3 accord with its formulation: for example, IR spectra show a broad hydroxy band at 3272 cm21, the 31P{-1H} NMR spectrum exhibits a singlet at d 12.18 flanked by a satellite doublet with a cis 31P–195Pt coupling constant of 3720 Hz, and the 1H NMR spectrum reveals a cis methylene doublet at d 4.30 with flanking satellites due to 195Pt coupling.Treating a suspension of 3 in methanol with excess of sodium carbonate gave back 2 in near quantitative yield, showing that 2 and 3 can be cleanly interconverted, Scheme 1. No evidence was found for the trans analogues of 2 and 3 in these reactions. The results reveal that binding of the Table 1 Selected bond lengths (Å) and angles (8) for cis-[Pt(O,P-PPh2- thqH)2]?2dmf (2?2dmf) with estimated standard deviations (e.s.d.s.) in parentheses Pt–PA Pt–O1A Pt–PB Pt–O1B PA–C7A PA–C8A PA–C14A PA–Pt–O1A PA–Pt–PB PA–Pt–O1B 2.229(1) 2.044(3) 2.232(1) 2.030(3) 1.830(4) 1.815(3) 1.817(3) 89.5(1) 101.0(1) 168.7(1) O1A–C1A C1B–C2B PB–C7B PB–C8B PB–C14B O1B–C1B O2B–C4B O1A–Pt–PB O1A–Pt–O1B PB–Pt–O1B 1.340(5) 1.399(6) 1.830(4) 1.815(3) 1.818(3) 1.334(5) 1.379(5) 168.6(1) 79.7(1) 90.0(1)J.Chem. Soc., Dalton Trans., 1999, 1543–1554 1545 hydroquinone groups and O,P-chelate formation can be pHcontrolled, and that the cis stereochemistry of the platinum ion is unchanged during these reversible transformations.This behaviour is much the same as found for transition metal complexes of other potentially O,P-chelate, alcohol- or phenolsubstituted phosphine ligands.23,24 Phosphines undergo Michael addition reactions with quinones, making these groups incompatible in an uncomplexed, “free” ligand.27 Complex formation prevents these reactions by “tying up” the phosphorus lone pair, and there should be no impediment to preparation of a complex with quinone phosphine ligands.In accord with these expectations, the p-quinone phosphine complex [PtBr2(PPh2tq)2] 4 was obtained by oxidation of 3 with two equivalents of strongly oxidising 2,3- dichloro-5,6-dicyano-1,4-benzoquinone (DDQ) in 97% yield, Scheme 1. Partial elemental analytical data for 4 are consistent with its formulation. Notable spectroscopic data include the Fig. 1 An ORTEP26 plot of cis-[Pt(O,P-PPh2thqH)2 2. The thermal ellipsoids are drawn at the 10% probability level.strong peak at m/z 887 for the [M 2 Br]1 ion in the electrospray mass spectrum, a strong quinone C]] O band at 1657 cm21 in the IR spectrum, a singlet at d 10.23 flanked by a satellite doublet with a cis 31P–195Pt coupling constant of 3685 Hz in the 31P- {1H} NMR spectrum and a methylene doublet at d 4.07 (flanked by satellite peaks due to 195Pt–H coupling) in the 1H NMR spectrum, again indicative of cis stereochemistry. In order to test the stability of 4 to hydrolysis, a d6-acetone solution of the complex was treated with a few drops of 0.1 M hydrobromic acid and the fate of 4 monitored by 1H and 31P- {1H} NMR spectroscopy.The solution was stable. No decomposition or other reaction of 4 was observed. An X-ray analysis of light brown crystals of complex 4?0.5CH2Cl2 obtained from dcm solution saturated with vapour from a 1: 1 mixture of diethyl ether–pentane confirms that the platinum(II) ion is co-ordinated by two cisoid toluquinone phosphine and two bromo ligands, Fig. 2. The distorted square planar geometry about the platinum ion [Br–Pt–Br 86.6(1); P–Pt–P 98.8(1); Br–Pt–P 85.9, 89.0(1)8] is unremarkable for cisbis( phosphine)platinum(II) complexes, as are the metal–ligand bond lengths (see Table 2).23,24 The quinone substituents lie on opposite sides of the platinum co-ordination plane and display the expected quinonoid C–C, C]] C and C]] O distances 1 (Table 2); they are unequivocally quinones.Quinones can be reduced by zinc under mildly acidic conditions to the corresponding hydroquinones.1 The oxidation of complex 3 (by DDQ) to 4 has already been described. In order to establish the chemical reversibility of the 3/4 couple, we treated a solution of 4 in acetone with excess of zinc powder and 0.1 M hydrobromic acid. After 2 h the excess of zinc was removed by filtration and the product precipitated with water. Recrystallisation from dmf–diethyl ether aVorded bimetallic cis-[Pt(O,P-PPh2thqH)2(ZnBr2)]?2dmf (5?2dmf); the dimer formally is the adduct of the Lewis acid ZnBr2 and 2 (Scheme 1).Only a small sample of 5 was prepared and a partial elemental analysis for C, H was not obtained. However, ICP analysis reveals a 2:1:1 ratio for P :Pt : Zn, and the electrospray mass Scheme 1 O O Pt Br P Ph2 P Ph2 Br O O Ph2 P O HO Ph2 P Pt O OH Zn Br Br OH OH Pt Br P Ph2 P Ph2 Br OH OH Ph2 P O HO Ph2 P Pt O OH OMe OMe Pt Cl P Ph2 Cl Ph2P OMe OMe Zn, HBr(aq), acetone HBr(aq), acetone Na2CO3, MeOH DDQ, acetone 2 4 5 3 HBr(aq), acetone 1a (i) BBr3, CH2Cl2 (ii) Na2CO3, MeOH1546 J.Chem. Soc., Dalton Trans., 1999, 1543–1554 spectrum shows prominent peaks at m/z 1031, 516, 952 and 809 corresponding to the [M 2 H]1, [M]21, [M–Br]1 and [M– ZnBr2]1 ions. The 31P-{1H} NMR spectrum exhibits a singlet at d 29.66 flanked by a cis satellite doublet [1J (31P–195Pt) 3851 Hz] and the 1H NMR spectrum shows the requisite peaks for two equivalent cisoid hydroquinonate phosphine ligands and for the dmf solvate molecules. The structure of the metal dimer 5 was established by X-ray crystallography, Fig. 3, and reveals that the ortho-oxygen atoms of the hydroquinonate groups bridge the platinum and zinc centres. Analogously to the structure of 2 (see above), the platinum(II) ion in 5 is O,P-co-ordinated by two cisoid hydroquinonate phosphine (PPh2thqH) ligands [the sum of the bond angles about the platinum ion (Table 3) is 6958 compared to 7208 for a perfect square planar centre; the bond angles P–Pt–P 100.08 and O-Pt-O 76.28 are little changed from those in 2].The zinc ion is co-ordinated by two bromo ligands and by the two bridging ortho-oxygen atoms of the hydroquinonate phosphine ligands in distorted-tetrahedral fashion [the sum of the bond angles about the zinc ion of 652.38 compares with 6578 for a perfect tetrahedron, but the individual bond angles (Table 3) are considerably diVerent from 109.58, e.g.Br–Zn–Br 116.0(1)8 Fig. 2 An ORTEP plot of cis-[PtBr2(PPh2tq)2] 4. Details as in Fig. 1. Table 2 Selected bond lengths (Å) and angles (8) for cis-[PtBr2- (PPh2tq)2]?0.5dcm (4?0.5dcm) with e.s.d.s. in parentheses Pt–Br1 Pt–Br2 PA–C1A PA–C8A PA–C14A O1A–C3A O2A–C6A C1A–C2A C2A–C3A C2A–C7A C3A–C4A C4A–C5A C5A–C6A C6A–C7A Br1–Pt–Br2 Br1–Pt–PA Br1–Pt–PB 2.477(1) 2.469(1) 1.855(7) 1.833(5) 1.835(4) 1.204(8) 1.211(8) 1.497(9) 1.497(10) 1.338(8) 1.477(10) 1.335(10) 1.448(11) 1.468(10) 86.6(1) 85.9(1) 174.3(1) Pt–PA Pt–PB PB–C1B PB–C8B PB–C14B O1B–C3B O2B–C6B C1B–C2B C2B–C3B C2B–C7B C3B–C4B C4B–C5B C5B–C6B C6B–C7B Br2–Pt–PA Br2–Pt–PB PA–Pt–PB 2.255(2) 2.254(2) 1.840(7) 1.824(5) 1.815(4) 1.219(8) 1.213(8) 1.504(9) 1.490(9) 1.317(9) 1.451(11) 1.335(10) 1.475(10) 1.481(10) 169.7(1) 89.0(1) 98.9(1) and O-Zn-O 78.8(2)8].The bond lengths (Table 3) are not remarkable,24 and are similar to the corresponding distances for 2. In the crystal structure each of the terminal hydroquinonate OH groups is hydrogen bonded to a molecule of dmf [there are two dmf per molecule of 5; O (dmf) ? ? ? O (hydroquinonate) 2.56, 2.66 Å]. The cobalt(II) iodide adduct of a bis(phosphinoenolate)nickel(II) complex, cis-[Ni{O,P-[Ph2- PCH]] ? ? ?C(]] ? ? ?O)(4-MeC6H4)]}2(CoI2)],28 exhibits square planar nickel and pseudotetrahedral cobalt centres giving overall a structure comparable to that of 2.The preparation of complex 5 from 4 confirms that the latter complex undergoes four-electron reduction to produce a hydroquinone phosphine complex.With the benefit of hindsight, it is not surprising that 5 was isolated rather than 3 which was targeted. A large excess of zinc powder was used, which would consume the hydrobromic acid and form zinc bromide. Under the final conditions pertaining in the reaction, no acid and an excess of zinc and zinc bromide, loss of the bromo ligands from 3 is expected and 5 is a logical product.The reverse reaction, 5 to 3, goes cleanly. Addition of a drop of 0.1 M hydrobromic acid to 5 in d6-acetone aVorded 3 in quantitative yield (as judged by NMR spectroscopy). Electrochemical studies The electrochemistries of complexes 1b, 3 and 4 have been characterised by cyclic voltammetric and bulk controlledpotential electrolysis experiments. Cyclic voltammograms were recorded in anhydrous dcm or dmf with 0.1 M tetrabutylammonium hexafluorophosphate at a platinum disc working Fig. 3 An ORTEP plot of cis-[Pt(O,P-PPh2thqH)2(ZnBr2)] 5. Details as in Fig. 1. Table 3 Selected bond lengths (Å) and angles (8) for cis-[Pt(O,PPPH2thqH) 2(ZnBr2)]?2dmf (5?2dmf) with e.s.d.s. in parentheses Pt–PA Pt–O1A Pt–PB Pt–O1B Zn–Br1 Zn–Br2 Zn–O1A Zn–O1B PA–Pt–O1A PA–Pt–PB PA–Pt–O1B O1A–Pt–PB O1A–Pt–O1B PB–Pt–O1B 2.217(3) 2.080(5) 2.219(2) 2.060(6) 2.349(2) 2.341(2) 2.014(6) 2.011(6) 91.7(2) 100.0(1) 167.5(2) 168.1(2) 76.2(2) 92.2(2) PA–C7A PA–C8A PA–C14A O1A–C1A PB–C7B PB–C8B PB–C14B O1B–C1B Br1–Zn–Br2 Br1–Zn–O1A Br1–Zn–O1B Br2–Zn–O1A Br2–Zn–O1B O1A–Zn–O1B 1.829(9) 1.807(6) 1.808(6) 1.362(10) 1.800(9) 1.803(6) 1.819(6) 1.382(10) 116.0(1) 105.7(2) 120.5(2) 123.2(2) 108.2(2) 78.8(2)J.Chem. Soc., Dalton Trans., 1999, 1543–1554 1547 electrode. All potentials are quoted relative to the ferrocenium– ferrocene couple. Except where otherwise stated, all description and potentials refer to measurements at 295 K run at a scan rate (n) of 100 mV s21.The following information, from voltammograms recorded under these conditions, is given for comparison with the electrochemistry of the platinum complexes. p-Hydroquinone29 (H2hq) in dmf exhibits an irreversible anodic peak at 10.46 V for oxidation to p-benzoquinone (q)-H1 and a daughter cathodic peak at 20.76 V for reduction of q-H1 to give back q. p-Benzoquinone 29 features a reversible one-electron quinone–semiquinone radical anion (q–sq~2) couple at 20.91 V in dcm and at 20.87 V in dmf followed at more negative potential by a quasireversible one-electron semiquinone anion–hydroquinone dianion (sq~2–hq22) couple at ca. 21.45 V [DEp = 165 mV cf. DEp(Fc1–Fc) = 86 mV] in dcm and at 21.69 V [DEp = 220 mV cf. DEp(Fc1–Fc) = 60 mV] in dmf. Bromide ion 30 (as 2.0 × 1023 M Et4NBr) exhibits an anodic peak at 10.40 V in dcm and at 10.32 V in dmf for the irreversible oxidation of Br2 to Br3 2 followed by the quasireversible Br3 2–Br2 couple at 10.62 V [DEp = 210 mV] in dcm and at 10.61 V [DEp = 190 mV] in dmf.Cyclic voltammograms of complex 1b in dcm reveal an irreversible reduction at 22.1 V and two broad irreversible oxidation processes at 10.82 and 11.05 V followed by a quasireversible couple at ca. 11.3 V. By comparison with the electrochemistry of [Pt(PR3)2X2] complexes where PR3 is an electrochemically inactive phosphine and X is a halide ligand, the irreversible reduction for 1b is attributed to the PtII–Pt0 couple, and the two irreversible anodic peaks to metal-centred oxidations of the cis and trans isomers, respectively.31 The quasireversible couple at more positive potential is assigned to oxidation of the dimethoxyphenyl substituents.32 A cyclic voltammogram of complex 3 in dmf solution revealed a broad anodic peak at 10.65 V attributed to the four-electron oxidation of the hydroquinone substituents 29 to aVord 4-H1, eqns.(3) and (4), and a daughter cathodic peak at [PtBr2(PPh2thqH2)2] æÆ [PtBr2(PPh2tq)2] 1 4 e2 1 4 H1 (3) [PtBr2(PPh2tq)2] 1 2 H1 æÆ [PtBr2(PPh2tq-H)2]21 (4) 20.74 V for reduction of this product.The synthesis of 4 by oxidation of 2 with 2 equivalents of DDQ provides support for this assignment. Figs. 4 and 5 show cyclic voltammograms of complex 4 recorded in dcm and dmf, respectively, at a freshly polished platinum working electrode. The initial scans are clearly similar in both solvents. However, subsequent scans in dcm are complicated by stripping and adsorption behaviour resulting from deposition of reduction product(s) onto the electrode, behaviour not observed in dmf.In both solvents, scans to negative potentials display an irreversible reduction process, PI (at 20.86 V in dcm and at 20.76 V in dmf), which upon scan inversion gives rise to two irreversible oxidation processes, PIII and PIV (at 10.05 and 10.23 V, respectively, in dcm and at 20.26 and 10.02 V, respectively, in dmf). In dcm, PIV has a cathodic counterpart, PV [DEp = 230 mV; i(PV)/i(PIV) ª 0.4 with n = 100 mV s21], e.g.Fig. 4. In dmf, PV is barely discernible, e.g. see Fig. 5. Peaks PII–PV are not present unless PI is traversed first. In some cases in dmf, shoulders for the peaks of a weak quasireversible couple can be discerned in the tail of PI. Cyclic voltammograms were also recorded at scan rates between 50 mV s21 and 10 V s21 in both solvents (e.g. Fig. 6) and in dcm at 278 8C. As the scan rate increases PI moves to more negative potential, and plots of peak current versus n1/2 are linear for PI– PIV.At ambient temperature the following general changes are noted as the scan rate (n) is increased. (1) The reduction process shows some chemical reversibility with PII, the anodic counterpart of PI, growing in relative magnitude as the scan rate is increased. (2) The relative magnitudes of PIII and PIV change with PIII growing at the expense of PIV. At 278 8C the voltammograms display PI–PIII only, with the PI–PII couple becoming more chemically reversible and PIII relatively smaller at higher scan rates.Accurate measurements in both dcm or dmf reveal that the first oxidation exhibited by Br2 ion occurs to positive potential of PIV, and cyclic voltammograms of equimolar complex 4 and Br2 ion in dcm exhibit PI–PV well clear of the peaks for the Br2– Br3 2 oxidation and the Br3 2–Br2 couple. The observations are noteworthy because they indicate that neither PIII nor PIV originates from oxidation of Br2 ion.Moreover in dcm solutions peaks for oxidation of Br2 ion were never observed (e.g. see Fig. 6); Br2 ion is not a product of PI–PIV. However, in dmf solution a broad anodic peak was observed at ª10.35 V in the tail of the anodic solvent discharge (marked by the asterisk in Fig. 7), which is close to the potential measured for the Br2–Br3 2 oxidation. We ascribe the broad 10.35 V peak to oxidation of a small amount of bromide ion produced by preceding solvation Fig. 4 Cyclic voltammogram (four cycles) of complex 4 in dcm–0.1 M [NBun 4][PF6] at a freshly polished platinum-disc working electrode; scan rate = 100 mV s21 and temperature = 295 K. Fig. 5 Cyclic voltammogram (four cycles) of complex 4 in dmf–0.1 M [NBun 4][PF6]. Other details as in Fig. 4.1548 J. Chem. Soc., Dalton Trans., 1999, 1543–1554 of 4 by dmf, the chemical step–electrochemical step (CE) mechanism33 described by eqns. (5) and (6).Consistent with C: [PtBr2(PPh2tq)2] 1 dmf [PtBr2-n(dmf)n(PPh2tq)2]n1 (n = 1 or 2) 1 nBr2 (5) E: 3 Br2 æÆ Br3 2 1 2 e2 (6) this interpretation, the peak is present in initial scans to positive potentials, prior to scanning through PI–PV, and the current from the oxidation remains about the same whether PI–PV are traversed or not. Conductivity measurements indicate that 4 is a non-electrolyte in dmf solution and NMR spectra of 4 in d7-dmf show peaks for a single species only.Taken together these results imply that 4 is unsolvated [i.e. the equilibrium (5) lies firmly to the left] with solvation only occurring, within the CV timescale, when an electrode potential suYcient to oxidise Br2 ion is applied, eqn. (6). Fig. 6 Cyclic voltammograms of complex 4 in dmf–0.1 M [NBun 4]- [PF6] at 295 K and diVerent scan rates. The platinum-disc working electrode was polished between the recording of each voltammogram. The arrows mark the potentials of the Br2 æÆ Br3 2 and Br3 2 æÆ Br2 oxidations under identical conditions.Fig. 7 Cyclic voltammograms of complex 4 in dmf–0.4 M [NBun 4][PF6] at 295 K prior to (a) and after (b) controlled potential electrolysis at 21.2 V. Controlled-potential coulometry reveals that PI is an oneelectron process. Two bulk electrolyses of complex 4 (1.0 mM) in dmf at a Pt-gauze electrode held at 21.2 V were performed and caused the solution to change from pale yellow to dark redgrey and consumed 0.98 and 1.03 Faraday mol21.A third bulk electrolysis at 21.0 V and half the concentration of 4 (0.5 mM) consumed 0.89 Faraday mol21. In cyclic voltammograms recorded after exhaustive electrolysis initial cathodic scans reveal no peak out to 22.3 V, i.e. PI disappears, but PIII and PIV appear in the reverse anodic scans, Fig. 7. Bulk electrolysis of fully reduced 4 in dmf at 10.27 V, i.e. positive of PIV, consumed 1.12 Faraday mol21, i.e. PIII and PIV together consume one electron per mol.The voltammogram after this reoxidation showed peaks PI, PIII and PIV, albeit with PI broadened compared to that in cyclic voltammograms of 4 at a freshly polished platinum working electrode. The reduction product of complex 4 was further characterised by UV/Visible/NIR, EPR and NMR spectroscopic studies. Fig. 8 displays Visible/NIR spectra of 4 prior to and after exhaustive reduction at PI. Both 4 and its reduction product display a strong UV band [at 308 (e ª 2790) and at 315 nm (e ª 6070 M21 cm21), respectively]. Whereas 4 only exhibits a weak band in the visible region at 785 nm (e ª 15 M21 cm21), the reduction product shows a broad band at 850 nm (e ª 300 M21 cm21) with shoulders at ca. 790 and 960 nm and a peak at 490 nm (e ª 395 M21 cm21) on the tail of the intense UV band. The radical nature of the reduction product of complex 4 was confirmed by EPR spectroscopy. First a solution of 4 in d2-dcm was treated with excess of cobaltocene [4 equivalents; E1/2 (CoCp2–CoCp2 1) = 21.35 V in dmf].A brown solid immediately precipitated, a result consistent with the deposition of the reduction product observed in cyclic voltammograms of 4 in dcm. The suspension of the reduction product in d2-dcm exhibited a strong isotropic EPR signal at g = 2.002 which slowly decayed over several days whilst an anisotropic signal at g = 5.2 slowly appeared. Next 4 in d4–dmf was titrated with cobaltocene. Addition of 1 equivalent gave a black solution with a reddish hue.The 1H NMR spectrum of the solution exhibited only peaks for cobaltocenium ion and the protio impurities in the solvent and the EPR spectrum of the frozen solution at 77 K revealed a broad isotropic peak at g = 2.003, Fig. 9. Hyper- fine coupling was not resolved. The intensity of the EPR signal did not change upon addition of further cobaltocene (up to 5 equivalents). Solutions produced by exhaustive reductive bulk electrolysis of 4 exhibited the same EPR signal.The insolubility of the radical product in dcm and other solvents with a suitable Fig. 8 The Vis/NIR spectra of complex 4 in dmf–0.4 M [NBun 4][PF6] at 295 K prior to (– – –) and after (——) controlled potential electrolysis at 21.2 V.J. Chem. Soc., Dalton Trans., 1999, 1543–1554 1549 IR window prevented measurement of its IR spectrum over the region for quinone C]] O bands. Attempts to isolate solid product for further analysis were thwarted by decomposition, consistent with the radical product being extremely oxygen and/or moisture sensitive.The above data reveal the following. Peak PI can confidently be assigned to a quinone-centred reduction of complex 4 by comparison with the cyclic voltammograms of 1b (devoid of electrochemistry in this region) and p-benzoquinone (q–sq~2 couple at about the same potential). The dependence of PII on scan rate indicates that a fast chemical step depletes the semiquinone product of PI (4~2) and generates a product radical species (6~2; we assign this as an anion because bromide ion is not produced).Both peaks PIII and PIV remain in cyclic voltammograms of 6~2 (whether produced by exhaustive reductive electrolysis or by cobaltocene reduction of 4). Clearly these anodic peaks are due to 6~2. Cooling a sample slows down chemical steps and recording a cyclic voltammogram at low temperature can be thought of as an alternative to scanning at very fast scan rates.Only PI–PIII are observed in voltammograms of complex 4 at low temperature and PIII is therefore assigned to irreversible oxidation of 6~2. Peak PIV can not arise from further oxidation of the product(s) produced at PIII as this would consume more than one electron per mol. The results are consistent, however, with pre-equilibration between 6~2, which is oxidised at PIII, and a second species, which is oxidised at PIV, with the preequilibrium favouring 6~2 and being rapid at ambient temperature and slow compared to the CV timescale at low temperature (278 8C).Possibilities for the chemical step include solvation or isomerisation reactions. That the reduction of complex 4 is firmly an one-electron process and produces a radical product (6~2) that can not be further reduced (no cathodic processes before the tail of the cathodic discharge at 22.3 V and no further reduction with excess of cobaltocene, see above) is (at first) puzzling! There are two quinone groups in 4 and a q–sq~2 couple is anticipated for each one.29,34 If the two quinone groups were non-interacting, the q–sq~2 couples would be coincident and appear as a single two-electron process.For closer quinone groups electrostatic eVects and electronic delocalisation might lead to consecutive, individual q–sq~2 couples, but even in the limit of full electronic delocalisation between the redox centres the couples should be separated by @500 mV.34 For example, the helical bis(quinone) I displays consecutive reversible q–sq~2 couples separated by 470 mV.34a,b This is the largest separation of the q–sq~2 couples Fig. 9 The EPR spectrum of the product (6~2) from reduction of complex 4 with cobaltocene (1 equivalent). Conditions: dmf glass, T = 77 K, n = 9.497 GHz. reported for a bis(quinone) and arises because the unpaired electron in the monoanion I~2 is fully delocalised over both quinone centres. Inspection of crude molecular models of 4,35 starting with the parameters from its crystal structure and leaving the cis-PtBr2P2 core in fixed position, indicates that rotations about the P–Pt bonds and the bonds to the benzylic carbon atoms lead to minimum attainable distances between the quinone rings of 5.3 Å with the rings parallel and 4.2 Å with the rings perpendicular.The quinone groups remain too far apart for significant p interaction although a small splitting of the two q–sq~2 couples is expected due to electrostatic eVects.The absence of any reduction processes for 6~2 (to over 1.5 V more negative potential than the q–sq~2 couple for 4) therefore points to both quinone groups in 4 being involved in the single chemically irreversible, one-electron reduction process. How can the addition of a single electron aVect the two widely separated quinone groups? Scheme 2 presents a possible mechanism for reduction of complex 4, a solution to this enigma. The first step is oneelectron reduction of a quinone substituent, i.e.the expected q–sq~2 couple.† The nucleophilicity of semiquinones has been measured to increase by up to six orders of magnitude compared to the parent quinones.36 The second step is addition of the nucleophilic semiquinone substituent to the electrophilic platinum(II) centre 8 to aVord {4*}~2. This is the anticipated first step in a ligand exchange reaction at a square planar platinum(II) centre, and such a reaction would normally proceed by loss of a leaving ligand from {4*}~2 (e.g.Br2 ion). However, the platinum centre is now electron rich (bound by the semiquinonate anion) and we posit that the conveniently placed, second quinone substituent oxidatively adds to the platinum centre to aVord 6~2, which is suggested to be an octahedral platinum(IV) (hydroquinonate)(semiquinonate) species, namely [PtBr2(O,P-PPh2thq)(O,P-PPh2tsq?)]2. The platinum(III) tautomer, [PtBr2(O,P-PPh2tsq?)2]2, is discounted for 6~2 on the basis of the isotropic EPR signal without 195Pt hyperfine coupling [we also note that genuine platinum(III) complexes are rare 37].The proposed structure for 6~2 neatly accounts for the EPR spectrum, a ligand-centred radical, and the intense, broad band in the Vis/NIR spectrum (which is atypical for a simple platinum complex 37 or for a simple semiquinone 38); it could arise from either an intervalence charge transfer transition (nIT in Scheme 2) or a ligand-to-metal charge transfer (hq22ÆPtIV) transition.9 Moreover, the proposed structure is entirely analogous to those of recently described [CoIII(N–N)(catecholate)( o-semiquinonate)] (N–N = dinitrogen chelate ligand) complexes which display both intervalance charge transfer and ligand-to-metal charge transfer {catecholateÆCoIII} transitions as well as characteristic temperature-dependent, spin transitions to [CoII(N–N)(o-semiquinonate)2] species.9 The major diVerence between the structures of these cobalt species O OO O I † If the quinone substituents in complex 4 are suYciently isolated for simultaneous reduction of both to give 422, a diradical with two semiquinone substituents, then addition of the semiquinone substituents to the metal centre and the cross-reduction of parent 4 (within the CV timescale) would lead to 6~2 and the observed overall one-electron reduction stoichiometry.Scheme 2 shows only one of several isomers possible for [PtBr2(O,P-PPh2thq)(O,P-PPh2tsq?)]2 6~2.1550 J.Chem. Soc., Dalton Trans., 1999, 1543.1554 Scheme 2 PPh2 PtII Br O O Ph2P Br PPh2 PtII Br O O Br PPh2 PtII Br Br O O Ph2P O O PtIV Br Br O O Ph2P O O PPh2 O O PtIV Br Br O O Ph2P PPh2 O O Ph2P O O {4*}. . ¡¾ e. 4 hn IT e. E1/2 (q/sq. .) 4. . 6. . and that proposed for 6~2 is that the 5 oxygen atoms (see Chart 1) in 6~2 are not bound to the metal centre and remain susceptible to protonation and subsequent reaction(s), perhaps accounting for the observed instability of this reduction product.Fig. 10 shows the cyclic voltammogram of complex 4 in dmf in the presence of dilute hydrobromic acid. On adding the acid a new reduction process (PI*) with a peak current �£7.5 times larger than that for PI (i.e. prior to addition of acid) appears at 21.05 V. Also, PII.PV disappear and a new anodic peak (PII*) is found at 20.62 V (compare Figs. 5 and 10). The Randles. Sevc¢§ik eqn. (7) 33 describes the peak current in a cyclic voltamip = 2(2.69 ¡¿ 105)n3/2Co .D1/2n1/2 (7) mogram, where Co .is the bulk concentration of the species undergoing oxidation/reduction, D its diVusion coeYcient and n the number of electrons transferred. In the present experiment only the latter parameter changes on adding the acid. The increase in the peak current indicates, therefore, a transition from an one- to a four-electron reduction process (43/2 = 8). Accordingly, PI* is attributed to reduction of 4 to aVord 3, i.e.the reverse of eqn. (3), and PII* to the reverse process. This is consistent with the voltammetry displayed by benzoquinones under acidic conditions and with the preparation of 5 from 4 (see above). Semiquinones disproportionate under protic conditions, often at near to diVusion controlled rates,39 explaining the switch from one- to four-electron reduction behaviour when acid is added to a solution of 4.29 Conclusions Synthetic strategies to the hydroquinone diphenylphosphine (PPh2thqH2) complexes of platinum(II) have been developed.The co-ordination of the hydroquinone groups to the platinum centre in these complexes can be reversibly controlled by pH adjustment. There are few surprises here; other studies of transition metal complexes with potentially O,P-chelate alcohol or phenol-substituted phosphine ligands have demonstrated similar pH control of oxygen co-ordination to the metal centre.23,24 The quinone phosphine complex 4 is easily prepared by oxidation of the hydroquinone phosphine precursor 3 and is not susceptible to hydrolytic loss othe two quinone groups, unlike Fig. 10 Cyclic voltammogram of complex 4 and ferrocene in dmf.0.1 M [NBun 4][PF6] at a freshly polished platinum-disc working electrode after addition of 5% v/v 0.1 M hydrobromic acid; scan rate = 100 mV s21 and temperature = 295 K.J. Chem.Soc., Dalton Trans., 1999, 1543–1554 1551 related complexes where the quinone substituents are directly bonded to the phosphorus atoms.16,19,20 The stability of 4 suggests other transition metal complexes of quinone phosphines should be readily available and we anticipate rich chemistry for these.Most remarkable is the electrochemistry of complex 4. Under protic conditions, chemical and electrochemical reduction of 4 is a four-electron process (two electrons per quinone group) and produces hydroquinone phosphine complexes (e.g. 2 and 5). This behaviour is as expected.In stark contrast, 4 is cleanly reduced under aprotic conditions by one electron to a radical product (6~2) that can not be further reduced. The dichotomy between one-electron reduction and the absence of quinone-centred electrochemistry in the product leads us to argue that reduction of 4 is accompanied by co-ordination of both quinone substituents to the platinum ion. It is proposed that a concomitant, extraordinary redistribution of electrons between the ligand and metal redox centres leads to 6~2 with a platinum(IV) ion bound by one hydroquinonate and one semiquinonate ligand (i.e., the addition of one electron to 4 causes two-electron oxidation of platinum and net three-electron reduction of the quinone groups!).The mechanism, Scheme 2, follows logically from the increased nucleophilicity of the semiquinone substituent produced by quinone-centred reduction of 4, and the suggested structure of 6~2 neatly accounts for its reactivity, and electrochemical and spectroscopic properties.We believe that the formation of 6~2 demonstrates the proclivity of quinone pendants, upon reduction to their semiquinone or hydroquinone anion counterparts, to bind and transfer electrons with a transition metal centre.8–10 In closing, we reemphasise that such processes are of pivotal biological importance,2–7 although the present work in no way attempts to model biological systems. Experimental Reactions were routinely carried out under an atmosphere of dry dinitrogen using standard Schlenk and cannula techniques. Solvents were distilled from the appropriate drying agent under dinitrogen immediately prior to use: dcm and acetonitrile from P2O5 and then from CaH2; acetone from KMnO4 and then from anhydrous B2O3; hexanes from sodium wire; diethyl ether and tetrahydrofuran from sodium benzophenone ketyl; dmf was dried over calcium hydride and then twice distilled under reduced pressure; methanol and ethanol were distilled from magnesium turnings.Chemicals were obtained from commercial sources (usually Aldrich) and used as obtained. Microanalyses for C, H and N were performed by the University of New South Wales microanalytical service. Inductively coupled plasma (ICP) analyses for other elements (P, Pt and Zn) employed a GBC Integra ICP-AES multi-channel instrument. Prior to analysis, samples were dried at 35 8C for 48 h under vacuum (0.2 mmHg) over phosphorus pentaoxide. Quoted melting points are uncorrected.The EI and ES mass spectra were recorded using a VG Quattro mass spectrometer; the carrier stream for ES was 1% acetic acid in 1 : 1 acetonitrile– water. The 1H and 13C NMR spectra were obtained in the designated solvents on a Bruker AC300F (300 MHz) instrument, 31P NMR spectra on a Bruker ACP300 spectrometer operating at 121.46 MHz and were referenced relative to external 85% phosphoric acid. The EPR spectra were recorded on a Bruker EMX 10 spectrometer, IR spectra as paraYn mulls on a Perkin-Elmer 580B spectrometer and electronic spectra using a CARY 5 spectrophotometer in the dual beam mode.Electrochemical measurements were recorded using a Pine Instrument Co. AFCBP1 Bipotentiostat interfaced to and controlled by a Pentium computer. Data were transferred to a Power Macintosh computer for processing using the IGORPRO 2.0TM software.40 For CV measurements, a standard three electrode configuration was used with a quasi-reference electrode comprised of a commercial Ag–AgCl mini-reference electrode (Cyprus Systems, Inc. EE008) but filled with the electrolyte solution to be used in the experiment [rather than AgCl saturated 3 M KCl(aq) solution], a freshly polished platinum disc (1 mm diameter) working electrode and a platinum wire as the auxiliary electrode. Freshly polished platinum working electrodes were prepared from commercial mini-electrodes (Cyprus Systems, Inc.EE041) by grinding with SiC emery paper (600 mesh), then successively polishing with 6 and 1 mm diamond slurries, and finally with 0.2 mm alumina slurry.Between each grinding and polishing step, and after final polishing, the electrode was sonicated in doubly distilled water for 5 min. The electrodes were then rinsed with the solvent to be used and thoroughly dried. The solvents used for electrochemical measurements, dcm and dmf, were highest quality anhydrous grade sealed under argon (Aldrich) and were used as obtained.The support electrolyte was 0.1 M [NBun 4][PF6]. Solutions were deoxygenated by flushing with high purity nitrogen (presaturated with solvent) and then blanketed with a cover of nitrogen for the duration of the experiment. An electrochemical scan of the solvent electrolyte system was always recorded before the addition of the compound to ensure that there were no spurious signals. All potentials are quoted relative to the ferrocenium–ferrocene (Fc1–Fc) couple which was measured in situ as an internal reference.Controlled potential coulometry was carried out in a conventional three-compartment “H”-cell adapted so that it could be loaded and sealed under an inert atmosphere (high purity dinitrogen). An Ag–AgCl quasi-reference electrode (the same as used in CV experiments) was placed in the working compartment along with the platinum gauze (5 × 2 cm2) working electrode and a platinum disc mini-electrode for running CV experiments.The counter electrode was a platinum gauze (4 × 2 cm2) separated from the working compartment by two fineporosity glass frits. During the electrolyses the solutions in the working compartment were stirred magnetically with a Tefloncoated stirring bar. The concentration of the substrate was 1.0 mM and the support electrolyte was 0.4 M [NBun 4][PF6]. Two test electrolyses were performed using these conditions: oxidation of ferrocene to ferrocenium ion and reduction of cobaltocenium hexafluorophosphate to cobaltocene consumed 1.04 and 0.99 Faraday mol21 respectively.Preparations (2,5-Dimethoxbenzyl)diphenylphosphine (PPh2dmb). Step 1: 2,5-dimethoxytoluene. Iodomethane (33 mL, 0.53 mol) was added by syringe to a mechanically stirred mixture of 2,5- dihydroxytoluene (30 g, 0.24 mol) and potassium carbonate (73.5 g, 0.24 mol) in degassed dmf under a dinitrogen atmosphere. The solution mixture was stirred for 16 h. Cooled water (0 8C) was added to the resulting pink solution until all the potassium carbonate had dissolved.The red solution was extracted with diethyl ether several times and the combined extracts washed successively with 2.5 M sodium hydroxide and water. The diethyl ether phase was dried over anhydrous magnesium sulfate and the solvent removed giving a red liquid. Distillation under 20 mmHg pressure at 120 8C gave a colourless liquid, 2,5-dimethoxytoluene (17.50 g, 50%). The compound was identified from its 1NMR spectrum:41 dH (CDCl3) 6.83–6.73 (3 H, m, C6H3), 3.84 (3 H, s, OCH3), 3.81 (3 H, s, OCH3) and 2.31 (3 H, s, CH3).Step 2: 2,5-dimethoxybenzyl bromide. CAUTION: This product irritates the skin and eyes and must be handled with due care. A solution of 2,5-dimethoxytoluene (17.50 g, 0.12 mol) in carbon tetrachloride (25 mL) was added to a suspension of N-bromosuccinamide (22.5 g, 0.12 mmol) and benzoyl peroxide (10 mg) in carbon tetrachloride (200 mL). The solution mixture was mechanically stirred and irradiated with a 200 watt lamp for 5 h.The hot solution was filtered and the volume of the1552 J. Chem. Soc., Dalton Trans., 1999, 1543–1554 solution reduced to ca. 100 mL. The resulting solution was cooled (ice–methanol bath) and gave the product as an oV white solid (14 g, 50% ), mp 75 8C (lit.42: 75–76 8C) (Found: C, 46.90; H, 4.75. Calc. for C9H11BrO2: C, 46.75; H, 4.76%). dh (CDCl3) 6.91 (2 H, m, C6H3), 6.82 (1 H, s, C6H3), 4.54 (2 H, s, CH2), 3.85 (3 H, s, OCH3) and 3.77 (3 H, s, OCH3).Step 3: PPh2dmb. Magnesium turnings (8 g, 0.33 mol) were stirred dry in a 250 mL Schlenk flask under a dinitrogen atmosphere for 48 h. Grey crushed magnesium powder was produced. 21 The flask was then connected to a pressure equalising dropping funnel. Freshly distilled diethyl ether (50 mL) was added to the magnesium powder, the mixture was cooled to 0 8C and a solution of 2,5-dimethoxybenzyl bromide (3 g, 13 mmol) in diethyl ether (100 mL) added dropwise to the centre of the vortex created by rapid stirring.Addition of the 2,5-dimethoxybenzyl bromide solution took 2 h. The mixture was stirred for 2 h and then the solution was filtered via a cannula into a solution of chlorodiphenylphosphine (2.33 mL, 13 mmol) in diethyl ether at 0 8C. A white precipitate formed immediately. After stirring for 16 h the reaction mixture was quenched with aqueous ammonium chloride (3 g in 50 mL of water).The diethyl ether phase was collected and dried with magnesium sulfate. Removing the solvent gave an oily residue. Recrystallisation from methanol yielded the product as a white solid (3.85 g, 88%), mp 50 8C (Found: C, 75.20; H, 6.50. C21H21O2P requires C, 75.00; H, 6.25%). dH (CDCl3) 7.45–7.31 (10 H, m, Ph), 6.74 [1 H, d, J(HH) 9, C6H3], 6.68 [1 H, dd, J(HH) 9 and 2 Hz, C6H3], 6.39 (1 H, m, C6H3), 3.69 (3 H, s, OCH3), 3.57 (3 H, s, OCH3) and 3.44 (2 H, s, CH2). dP (CDCl3) 211.65 (s).m/z (EI-MS) 336 (M1). [PtX2(PPh2dmb)2] complexes. X = Cl (1a). A solution of [PtCl2(PhCN)2] (0.17 g, 0.36 mmol) in dcm (3 mL) was added to a solution of PPh2dmb (0.25 g, 0.74 mmol) in dcm (10 mL) and the resulting clear solution stirred for ca. 15 min. The solvent was then reduced to ca. 3 mL. Dropwise addition of diethyl ether (40 mL) precipitated an oV white solid (0.32 g, 95%), mp 248 8C (Found: C, 53.77; H, 4.62. C42H42Cl2O4P2Pt requires C, 53.73; H, 4.48%); NMR spectra show two isomers.dH (CDCl3), cis-[PtCl2(PPh2dmb)2] (95%) 7.81 (2 H, m, C6H3), 7.18 (8 H, m, Ph), 7.09 (8 H, m, Ph), 6.94 (4 H, m, Ph), 6.81 [2 H, d, J(HH) 9, C6H3], 6.41 [2 H, d, J(HH) 9, C6H3], 4.16 [4 H, d, J(PH) 12 Hz, CH2], 3.90 (6 H, s, OCH3) and 2.98 (6 H, s, OCH3); and trans- [PtCl2(PPh2dmb)2] (5%) 7.60 (8 H, m, Ph), 7.38 (8 H, m, Ph), 7.30 (6 H, m, Ph and C6H3), 6.65 (2 H, m, C6H3), 6.52 [2 H, d, J(HH) 9, C6H3], 4.23 [4 H, t, J(PH) 8 Hz, CH2], 3.51 (6 H, s, OCH3) and 3.03 (6 H, s, OCH3).dP (CDCl3) 11.53 [s, J(31P-195Pt) 3780, cis isomer] and 15.99 [s, J(31P–195Pt) 2565 Hz, trans isomer]. X = Br (1b). A mixture of [PtCl2(PPh2dmb)2] (0.10 g, 0.1 mmol) and sodium bromide (0.10 g, 1.00 mmol) was stirred in dcm (25 mL) for 2 h. The undissolved solid was removed by filtration and the solvent removed from the filtrate to produce a pale yellow solid (0.09 g, 80%), mp 246 8C (Found: C, 49.14; H, 4.48. C42H42Br2O4P2Pt requires C, 49.07; H, 4.09%); NMR spectra show two isomers.dH (CDCl3), cis-[PtBr2- (PPh2dmb)2] (95%) 7.80 (2 H, m, C6H3), 7.18 (8 H, m, Ph), 7.09 (8 H, m, Ph), 6.94 (4 H, m, Ph), 6.83 [2 H, dd, J(PH) 12 and 3, C6H3], 6.46 [2 H, d, J(HH) 9, C6H3, 2H], 4.32 [4 H, d, J(PH) 12 Hz, CH2], 3.91 (6 H, s, OCH3) and 2.98 (6 H, s, OCH3); and trans-[PtBr2(PPh2dmb)2] (5%) 7.56 (8 H, m, Ph), 7.37 (8 H, m, Ph), 7.29 (6 H, m, Ph and C6H3), 6.68 (2 H, m, C6H3), 6.53 [2 H, d, J(HH) 9, C6H3], 4.24 [4 H, t, J(PH) 8 Hz, CH2], 3.53 (6 H, s, OCH3) and 3.08 (6 H, s, OCH3).dP (CDCl3) 10.97 [s, J(31P– 195Pt) 3725, cis isomer] and 15.30 [s, J(31P–195Pt) 2485 Hz, trans isomer]. X = I (1c). Reaction of [PtCl2(PPh2dmb)2] (0.10 g, 0.10 mmol) with potassium iodide (0.20 g, 1.2 mmol) using the method outlined for [PtBr2(PPh2dmb)2] gave the product as a red-brown solid (0.09 g, 80%) mp 230 8C (Found: C, 44.65; H, 3.39. C42H42I2O4P2Pt requires C, 44.96; H, 3.75%); NMR spectra show two isomers.dH (CDCl3), cis-[PtI2(PPh2dmb)2] (65%) 7.75 (2 H, m, C6H3), 7.15 (8 H, m, Ph), 7.09 (8 H, m, Ph), 6.92 (4 H, m, Ph), 6.82 (2 H, m, C6H3), 6.46 [2 H, d, J(HH) 9, C6H3], 4.52 [4 H, d, J(PH) 12 Hz, CH2], 3.90 (6 H, s, OCH3) and 2.98 (6 H, s, OCH3); and trans-[PtI2(PPh2dmb)2] (35%) 7.56 (8 H, m, Ph), 7.40 (8 H m, Ph), 7.29 (6 H, m, Ph and C6H3), 6.68 (2 H, m, C6H3), 6.52 [2 H, d, J(HH) 9, C6H3], 4.28 [4 H, t, J(PH) 9 Hz, CH2], 3.52 (6 H, s, OCH3) and 3.10 (6 H, s, OCH3). dP (CDCl3) 6.66 [s, J(31P–195Pt) 3545, cis isomer] and 1.55 [s, J(31P–195Pt) 2395 Hz, trans isomer].[Pt(O,P-PPh2thqH)2] 2. Boron tribromide (1.5 mL, 15 mmol) was added to a stirred solution of [PtCl2(PPh2dmb)2] (0.50 g, 0.11 mmol), in dcm (25 mL) at 0 8C under a dinitrogen atmosphere. The solution changed from light yellow to red brown on addition of the boron tribromide. After 16 h the solvent was removed. The resulting foamy yellow solid was treated with distilled methanol (15 mL) and then sodium carbonate (1.5 g) was added.After 3 h much light yellow precipitate had formed. This was collected by filtration, washed with water to remove the excess of sodium carbonate, rinsed with diethyl ether and air dried to yield the product, as a light yellow solid (420 mg, 96%), mp 180 8C (decomp.) (Found: C, 54.43; H, 4.06. C38H32O4P2Pt?1.5H2O requires C, 54.55; H, 4.19%). This compound dissolved in (CH3)2SO to produce a clear light yellow solution, but almost immediately a yellow solvate hydrate precipitated.[Found: C, 50.74; H, 4.69. C38H32O4P2Pt?2(CH3)2SO? 2H2O requires C, 50.35; H, 4.79%]. The solvate hydrate and the original product both dissolved in dmf. m/z (ES-MS, dissolved in dmf) 810 (M 1 H). dH [(CD3)2NCDO] 8.07 (1 H, s, OH), 8.01 (1 H, s, OH), 7.35–7.19 (20 H, m, Ph), 6.62 [2 H, d, J(HH) 6, C6H3], 6.43 [2 H, d, J(HH) 6 Hz, C6H3], 6.16 (2 H, s, C6H3) and 4.11 [4 H, d, J(PH) 13.5 Hz CH2]. dP [(CD3)2NCDO] 28.08 [s, J(31P–195Pt) 3822 Hz].[PtBr2(PPh2thqH2)2] 3. Concentrated hydrobromic acid (0.5 mL) was added to a suspension of [Pt(O,P-PPh2thqH)2] (0.21 g, 0.26 mmol) in acetone (20 mL). The solution was stirred for 15 min, then the volume of the solvent was reduced to 5 mL and water added to give the product, [PtBr2(PPh2thqH2)2], as an oV white solid (0.21 g, 81%), mp 264 8C (Found: C, 47.90; H, 3.83. C38H34Br2O4P2Pt?C3H6O requires C, 47.81; H, 3.89%). dH [(CD3)2CO] 8.30 (2 H, s, OH), 7.75 (2 H, s, C6H3), 7.44 (2 H, s, OH), 7.21 (12 H, m, Ph), 6.93 (8 H, m, Ph), 6.60 [2 H, d, J(HH) 9, C6H3], 6.47 [2 H, d, J(HH) 9, C6H3] and 4.30 [4 H, d, J(PH) 12 Hz, CH2].dP [(CD3)2CO)] 12.18 [s, J(31P–195Pt) 3720 Hz]. n& max/cm21 (OH) 3272s and 3261s (paraYn mull). [PtBr2(PPh2tq)2] 4. A solution of DDQ (45 mg, 0.20 mmol) in acetone (2 mL) was added to a solution of [PtBr2(PPh2thqH2) 2] (0.10 g, 0.10 mmol) in acetone (15 mL), and the resulting solution stirred for 15 min. The solvent was then removed and the residue extracted with dcm (20 mL).The solvent was removed from the extract giving the product as a redbrown solid (0.10 g, 97%), mp 140 8C (Found: C, 47.60; H, 3.21. C38H30Br2O4P2Pt requires C, 47.16; 3.10%). m/z (ES-MS) 887 [M 2 Br]1. dH (CDCl3) 7.45 (8 H, m, Ph), 7.32 (4 H, m, Ph), 7.14 (8 H, m, Ph), 7.09 (2 H, m, C6H3), 6.65 [2 H, dd, J(HH) 10 and 2.5, C6H3], 6.53 [d, J(HH) 10, C6H3] and 4.07 [d, J(PH) 13 Hz, CH2]. dP (CDCl3) 10.23 [s, J(31P–195Pt) 3685 Hz].Conductivity: LM (1.0 mM in dmf) = 3.7 S cm2 mol21. n& max/cm21 (CO) 1657s (paraYn mull). Reduction of [PtBr2(PPh2tq)2] 4 with cobaltocene. Several experiments were performed, all in sealable EPR/NMR tubes which could be fitted to a vacuum manifold. In a typical experiment, a sealable NMR tube was charged with [PtBr2- (PPh2tq)2] (20 mg, 0.02 mmol) and cobaltocene (3.9 mg, 0.02 mmol) and attached to the vacuum manifold. Freeze-thawJ. Chem. Soc., Dalton Trans., 1999, 1543¡V1554 1553degassed CD2Cl2 transferred to the tube by cooling it withliquid nitrogen whilst under vacuum.The NMR tube wassealed and the CD2Cl2 thawed. A brown precipitate in a clearyellow solution formed. The 1H NMR spectrum showed onlypeaks for protio-solvent impurities (d 5.32) and cobaltoceniumion (d 5.88). The EPR spectrum showed a strong isotropicsignal at g = 2.002. In experiments using dmf as the solvent,deoxygenated dmf was introduced to the evacuated EPR tubevia cannula.[Pt(O,P-PPh2thqH)2(ZnBr2)] 5.Hydrobromic acid (3 mL, 0.1M) was added dropwise to a stirred suspension of zinc powder(200 mg) in a solution of [PtBr2(PPh2tq)2] (30 mg, 0.03 mmol)in acetone (25 mL). After 1 h the mixture was filtered to removethe excess of zinc and the solvent concentrated to ca. 10 mL.Water was added to precipitate a faun powder which was collectedby filtration, washed with water and diethyl ether, anddried in vacuo (30 mg, 90%). ICP analysis: P :Pt :Zn ratio 2:1:1.m/z (ES-MS) 516 (M21), 809 (M 2 ZnBr2]1), 952([M 2 Br]1) and 1031 ([M 2 H]1). dH [(CD3)2CO] 7.55¡V7.27(20 H, m, Ph), 6.50 (2 H, m, C6H3), 6.18 (2 H, br, C6H3), 5.61[2 H, d, J(PH) 10, C6H3] and 3.86 [4 H, d, J(PH) 13 Hz, CH2].dP [(CD3)2CO] 29.66 [s, J(31P¡V195Pt) 3851 Hz]. Crystals of thissample were obtained on recrystallisation from dmf under anatmosphere of diethyl ether, and the formulation of 5 rests onthe above data, on a crystal structure analysis of 5?2dmf and onthe reaction of 5 with hydrobromic acid (see next).Reaction of [Pt(O,P-PPh2thqH)2(ZnBr2)] 5 with hydrobromicacid.A sample of [Pt(O,P-PPh2thqH)(ZnBr2] (25 mg) was dissolvedin d6-acetone (0.5 mL) in an NMR tube. After 1H and31P-{1H} NMR spectra had been acquired, 0.1 M hydrobromicacid (ca. 0.05 mL) was added, the solution shaken and theNMR spectra re-run. The 1H and 31P-{1H} NMR spectrarevealed that 5 cleanly reacted to give 3.CrystallographyCrystal data for [Pt(O,P-PPh2thqH)2]?2dmf (2?2dmf).Pt(C19-H16O2P)2?(2C3H7NO), M 955.9, triclinic, space group P1, a9.958(8), b 14.336(14), c 15.438(9) , a 101.40(4), b 93.85(4), g108.79(3)8, V 2025(3) 3, Dc 1.57 g cm23, Z 2, m(Mo-Ka) 36.25cm21, T = 294 K. Crystal size 0.10 ¡Ñ 0.12 ¡Ñ 0.19 mm, 2qmax 488,minimum and maximum transmission factors 0.50 and 0.70.The number of reflections was 5424 considered observed[I > 3s(I)] out of 6350 unique data, with Rmerge 0.013 for equivalentreflections.Final residuals R, R9 were 0.023, 0.032 for theobserved data.Crystal data for [PtBr2(PPh2tq)2]?0.5CH2Cl2 (4?0.5CH2Cl2).Pt(C19H15O2P)2Br2?(0.5CH2Cl2), M 1010.0, monoclinic, spacegroup P21/c, a 10.908(5), b 17.329(5), c 20.523(10) , b105.06(2)8, V 3746(3) 3, Dc 1.79 g cm23, Z 4, m(Mo-Ka) 60.90cm21, T = 294 K. Crystal size 0.09 ¡Ñ 0.13 ¡Ñ 0.19 mm, 2qmax 448,minimum and maximum transmission factors 0.53 and 0.76.The number of reflections was 3303 considered observed[I > 3s(I)] out of 4878 unique data, with Rmerge 0.018 for equivalentreflections.Final residuals R, R9 were 0.024, 0.031 for theobserved data.Crystal data for [Pt(O,P-PPh2thqH)2(ZnBr2)]?2dmf (5?2dmf).Pt(C19H16O2P)2ZnBr2?(2C3H7NO), M 1181.1, monoclinic,space group P21/c, a 13.030(9), b 30.701(13), c 12.460(9) , b113.81(3)8, V 4560(5) 3, Dc 1.72 g cm23, Z 4, m(Mo-Ka) 54.85cm21, T = 294 K. Crystal size 0.09 ¡Ñ 0.19 ¡Ñ 0.20 mm, 2qmax 448,minimum and maximum transmission factors 0.38 and 0.51.The number of reflections was 3569 considered observed[I > 3s(I)] out of 5572 unique data, with Rmerge 0.017 forequivalent reflections.Final residuals R, R9 were 0.033, 0.042for the observed data.CCDC reference number 186/1399.See http://www.rsc.org/suppdata/dt/1999/1543/ for crystallographicfiles in .cif format.AcknowledgementsS. B. S. thanks the University of North Sumatra for a period ofleave.References1 Z. Rappoport and S. Patai (Editors), The Chemistry of theQuinonoid Compounds, Wiley, New York, 1988, vol.II, Parts 1 and2; S. Patai (Editor), The Chemistry of the Quinoid Compounds,Wiley, New York, 1974, vol. I, Parts 1 and 2.2 D. M. Dooley, M. A. McGuirl, D. E. Brown, P. N. Turoski,W. S. McIntire and P. F. Knowles, Nature (London), 1991, 349,262; D. M. Dooley, R. A. Scott, P. F. Knowles, C. M. Colangelo,M. A. McGuirl and D. E. Brown, J. Am. Chem. Soc., 1998, 120,2599.3 J. P. Kliman, Chem. Rev., 1996, 96, 2541; W. S. McIntire, Annu.Rev.Nutrition, 1998, 18, 145.4 S. X. Wang, M. Mure, K. F. Medzihradsky, A. L. Burlingame,D. E. Brown, D. M. Dooley, A. J. Smith, H. M. Kagan and J. P.Klinman, Science, 1996, 273, 1076; L. I. Smithmungo and H. M.Kagan, Matrix Biology, 1998, 16, 387.5 J. Deisenhofer, O. Epp, K. Micki, R. Huber and H. Michel, Nature(London), 1985, 318, 618; H. Kurreck and M. Huber, Angew. Chem.,Int. Ed. Engl., 1995, 34, 849.6 D. G. Nichols and S. J. Ferguson, Bioenergetics 2, Academic Press,New York, 1992.7 S.Iwata, J. W. Lee, K. Okada, J. K. Lee, M. Iwata, B. Rasmussen,T. A. Link, S. Ramaswamy and B. K. Jap, Science, 1998, 281, 64;U. Brandt, Biochim. Biophys. Acta., Bioenerg., 1998, 1364, 85.8 C. G. Pierpont and R. M. Buchanan, Coord. Chem. Rev., 1981, 38,45; C. G. Pierpont and C. Lange, Prog. Inorg. Chem., 1994, 41, 331.9 D. M. Adams, A. Dei, A. L. Rheingold and D. N. Hendrickson,J. Am. Chem. Soc., 1993, 115, 8221; O. S. Jung, D. H. Jo, Y. A. Lee,Y. S. Sohn and C.G. Pierpont, Angew. Chem., Int. Ed. Engl., 1996,35, 1694; D. M. Adams and D. N. Hendrickson, J. Am. Chem. Soc.,1996, 118, 11515; O.-S. Jung, D. H. Jo, Y.-A. Lee, B. J. Conklin andC. G. Pierpont, Inorg. Chem., 1997, 36, 19.10 R. A. Klein, P. Witte, R. Vanbelzen, J. Fraanje, K. Goubitz,M. Numan, H. Schenk, J. M. Ernsting and C. I. Elsevier, Eur. J.Inorg. Chem., 1998, 319; T. E. Keyes, R. J. Forster, P. M. Jayaweera,C. G. Coates, J. J. Mcgarvey and J. G. Vos, Inorg. Chem., 1998, 37,5922; M.Handa, T. Nakao, M. Mikuriya, T. Kotera, R. Nukadaand K. Kasuga, Inorg. Chem., 1998, 37, 149; T. P. Vaid, E. B.Lobkovsky and P. T. Wolczanski, J. Am. Chem. Soc., 1997, 119,8742; R. A. Klein, C. J. Elsevier and F. Hartl, Organometallics, 1997,16, 1284; A. Kunzel, M. Sokolow, F.-Q. Liu, H. W. Roesky,M. Noltemeyer, H.-G. Schmidt and I. Usn, J. Chem. Soc., DaltonTrans., 1996, 913; S. L. Scott, A. Bakac and J. H. Espenson, J. Am.Chem. Soc., 1992, 114, 4605; T. van der Graaf, D.J. Stufkens,J. Vichov and A. Vlc£¾ek, Jr., 1991, 401, 305; S. Ernst, P. Hael,J. Jordanov, W. Kaim, V. Kasack and E. Roth, J. Am. Chem. Soc.,1989, 111, 1733; M. Hanaya and M. Iwaizumi, Organometallics,1989, 8, 672; A. Vlc£¾ek, Jr., J. Klma and A. A. Vlc£¾ek, Inorg. Chim.Acta, 1983, 69, 191; J.-T. M. Tuchagues and D. N. Hendrickson,Inorg. Chem., 1983, 22, 2545.11 W. Keim, Chem.-Ing.-Tech., 1984, 56, 850; personal communication.12 E. Drent and H. M. Budzelaar, Chem.Rev., 1996, 96, 663.13 K. Itami, A. Palmgren, A. Thorarensen and J. E. Backvall, J. Org.Chem., 1998, 63, 6466; H. Grennberg and J. E. Backvall, Chem. Eur.J., 1998, 4, 1083; A. Palmgren, A. Thorarensen and J. E. Backvall,J. Org. Chem. 1998, 63, 3764; C. Jonasson and J. E. Backvall,Tetrahedron Lett., 1998, 39, 3601; K. Bergstad, H. Grennberg andJ. E. Backvall, Organometallics, 1998, 17, 45; J. Vicente, A. Arcas,D. Bauutista and G. B. Shul¡¦pin, J. Chem. Soc., Dalton Trans., 1997,1505; R.Z. Chen and J. J. Pinatello, Environ. Sci. Technol., 1997, 31,2399; T. Yokota, S. Fujibayashi, Y. Nishiyama, S. Sakaguchi andY. Ishii, J. Mol. Catal. A., 1996, 114, 113; C. Jia, P. Muller andM. Mimoun, J. Mol. Catal. A., 1995, 101, 127; G. B. Shul¡¦pin,M. M. Bochkova and G. V. Nizova, J. Chem. Soc., Perkin Trans. 2,1995, 1465.14 M. Hudlicky, ACS Monogr., 1990, 186.15 S. B. Sembiring, S. B. Colbran and D. C. Craig, Inorg. Chem., 1995,34, 761.16 S. B. Sembiring, S. B.Colbran, D. C. Craig and M. L. Scudder,J. Chem. Soc., Dalton Trans., 1995, 3731.1554 J. Chem. Soc., Dalton Trans., 1999, 1543–1554 17 S. B. Sembiring, S. B. Colbran, R. Bishop and A. D. Rae, Inorg. Chim. Acta, 1995, 228, 109. 18 S. B. Sembiring, S. B. Colbran and L. R. Hanton, Inorg. Chim. Acta, 1992, 202, 67. 19 S. B. Sembiring, S. B. Colbran and D. C. Craig, Inorg. Chim. Acta, 1990, 176, 225. 20 S. B. Sembiring, Zainal, S. B. Colbran and D. C. Craig, manuscript in preparation. 21 K. V. Baker, J. M. Brown, N. Hughes, A. J. Skarnulis and A. Sexton, J. Org. Chem., 1991, 56, 698. 22 J. A. Rahn, D. O. O’Donnell, A. R. Palmer and J. H. Nelson, Inorg. Chem., 1989, 28, 2631. 23 H. D. Empsall, B. L. Shaw and B. L. Turtle, J. Chem. Soc., Dalton Trans., 1976, 1500; T. B. Rauchfuss, Inorg. Chem., 1977, 16, 2967; H. D. Empsall, P. N. Heys and B. L. Shaw, J. Chem. Soc., Dalton Trans., 1978, 257; P. Braunstein, D. Matt, D. Nobel, F. Balegroune, S.-E. Bouaoud, D. Grandjean and J. Fischer, J. Chem. Soc., Dalton Trans., 1988, 353; A. W. G. Platt and P. Pringle, J. Chem. Soc., Dalton Trans., 1989, 1193; P. Bergamini, S. Sostero, O. Traveso, T. J. Kemp and P. Pringle, J. Chem. Soc., Dalton Trans., 1989, 2017; L. Kollar and G. S. Zalontai, J. Organomet. Chem., 1991, 421, 341; E. M. Georgiev, H. tom Dieck, G. Hahn, G. Petrov and M. Kirilov, J. Chem. Soc., Dalton Trans., 1992, 1311. 24 C. D. Montgomery, N. C. Payne and C. J. Willis, Inorg. Chem., 1987, 26, 519; N. W. Alcock, A. W. G. Platt and P. Pringle, J. Chem. Soc., Dalton Trans., 1987, 2273; Inorg. Chim. Acta, 1987, 128, 215; J. Chem. Soc., Dalton Trans., 1989, 139; A. Varshney, M. L. Webster and G. M. Gray, Inorg. Chem., 1992, 31, 2580. 25 F. H. Allen and S. N. Sze, J. Chem. Soc. A, 1971, 2054. 26 C. K. Johnson, ORTEP II, report ORNL-5138, Oak Ridge National Laboratory, Oak Ridge, TN, 1976. 27 F. Rameriz, Pure Appl. Chem., 1964, 9, 337. 28 J. Andrieu, P. Braunstein, M. Drillon, Y. Dusausoy, F. Ingold, P. Rabu, A. Tiripicchio and F. Ugozzoli, Inorg. Chem., 1996, 35, 5986. 29 B. R. Eggins and J. Q. Chambers, J. Electrochem. Soc., 1970, 117, 186; N. Gupta and H. Linshitz, J. Am. Chem. Soc., 1997, 119, 6384; J. Q. Chambers in ref. 1, vol. I, ch. 14 and vol. II, ch. 12. 30 P. Castellonese and G. Launay, Bull. Soc. Chim. Fr., 1978, 7/8, I317; B. E. Conway, Y. Phillips and S. Y. Qian, J. Chem. Soc., Faraday Trans., 1995, 283. 31 G. Mazzocchin, G. Bontempelli, M. Nicolini and B. Crociani, Inorg. Chim. Acta, 1976, 18, 159; J. A. Davies and V. Uma, Inorg. Chim. Acta, 1983, 76, L305; J. Electroanal. Chem., 1983, 158, 13; G.-A. Mazzocchin and G. Bontempelli, J. Electroanal. Chem., 1984, 179, 269; J. A. Davies and V. Uma, J. Electroanal. Chem., 1984, 179, 273. 32 A. Zweig, W. G. Hodgson and W. H. Jura, J. Am. Chem. Soc., 1964, 86, 4124; V. Le Berre, L. Angely, J. Simonet, G. Mousset and M. Bellec, J. Electroanal. Chem., 1987, 218, 173. 33 A. J. Bard and L. R. Faulkner, Electrochemical Methods: Fundamentals and Applications, Wiley, New York, 1980. 34 (a) B. Yang, L. Liu, T. J. Katz, C. A. Liberko and L. R. Miller, J. Am. Chem. Soc., 1991, 113, 8993; (b) A. L. Sargent, J. Almlof and C. A. Liberko, J. Phys. Chem., 1994, 98, 6114; (c) S. Hunig, K. Sinzger, M. Kemmer, U. Langohr, H. Rieder, S. Soderholm, J. U. Vonschutz and H. C. Wolf, Eur. J. Org. Chem., 1998, 9, 1977. 35 CS Chem3DProTM V3.2, Cambridgesoft Corp., Cambridge, MA, 1996. 36 P. S. Rao and E. Hayon, J. Phys. Chem., 1973, 77, 2274. 37 D. M. Roundhill, Comprehensive Coordination Chemistry, Pergamon, Oxford, 1987, vol. 5, ch. 52, pp. 351–532. 38 T. Itoh, Chem. Rev., 1995, 95, 2351. 39 E. T. Denisov and I. V. Khudyakov, Chem. Rev., 1987, 87, 1313. 40 IGORPRO 2.0TM, Wavemetrics Inc., Lake Oswego, OR, 1995. 41 K. Yoshida, M. Shigi and T Fueno, J. Org. Chem., 1975, 40, 63. 42 A. T. Shulgin and E. M. Gal, J. Chem. Soc., 1953, 1316. Paper 9/00610I

ISSN:1477-9226    

DOI:10.1039/a900610i

出版商:RSC    

年代:1999

数据来源: RSC

摘要:

DALTON FULL PAPER J. Chem. Soc., Dalton Trans., 1998, Pages 1551–1556 1551 Magnetostructural behaviour of the complex [MnL(H2O)2]Cl2?4H2O at variable temperature studied by electron spin resonance (L 5 2,13- dimethyl-3,6,9,12,18-pentaazabicyclo[12.3.1]octadeca-1(18),2,12,14,16- pentaene) Omar Jiménez-Sandoval,a Daniel Ramírez-Rosales,b María del Jesús Rosales-Hoz,c Martha Elena Sosa-Torres *,†,a and Rafael Zamorano-Ulloa b a División de Estudios de Posgrado, Facultad de Química, Universidad Nacional Autónoma de México, Ciudad Universitaria, México, D.F. 04510, México b Depto. de Física, Escuela Superior de Física y Matemáticas del IPN, Edif. 9, U. P. Zacatenco, Col. San Pedro Zacatenco, México, D. F. 07738, México c Depto. de Química, Centro de Investigación y Estudios Avanzados del IPN, Apartado Postal 14-740, México, D. F. 07000, México The molecular and crystal structure of the complex [MnL(H2O)2]Cl2?4H2O 1 (L = 2,13-dimethyl-3,6,9,12,18- pentaazabicyclo[12.3.1]octadeca-1(18),2,12,14,16-pentaene) have been determined, the solid state and solution electronic spectra recorded and the thermogravimetric analysis as well as an ESR analysis at diVerent temperatures performed.The structure shows that the cation displays a distorted pentagonal-bipyrimidal co-ordinated geometry, with the macrocycle in the pentagonal plane and two water molecules in the axial positions. From the UV/VIS spectra it can be seen that the electronic structure of 1 is very sensitive to surroundings.The 300 K ESR spectrum of 1 consists of five fine-structure lines centred at g = 2.111, showing anisotropy. A sequence of spectra neatly shows that the compound has a clear magnetic dependence on temperature. Spectral analysis and theoretical calculations give the best 300 K zero-field splitting parameters as D = 0.07 cm21, E = 0.008 cm21, l = E/D = 0.1142. The 77 K zero-field splitting parameters increase to D = 0.074 cm21 and E = 0.012 cm21, thus indicating an increasing rhombic distortion as the temperature decreases.The Q-band spectra at 300 and at 77 K are isotropic, and the zero-field eVects are very small. The theoretical Q-band spectra were calculated on the basis of the X-band parameters. The temperature variation of the crystal-field parameters is interpreted as a smooth magnetostructural temperature dependence of the compound. The booming development that the chemistry of macrocyclic ligand complexes has experienced in the last few decades is well founded on the interesting structural, thermodynamic, kinetic and magnetic properties that such compounds display.Another key factor in their development and study is their possible usage as models for important biological systems that involve metal ions co-ordinated to macrocycles.1 Moreover, the co-ordination chemistry of manganese has been important for the better understanding of the O2 evolution mechanism in the photosynthetic process.2 Additional interest bears on totally or partly unsaturated macrocyclic compounds, the lack of flexibility of which results in a restriction of their possible co-ordination modes to metal ions, thus forcing the metal in some cases to accommodate uncommon geometries3 and promoting an enhancement of the stabilising macrocyclic eVect.4 The 15-membered, pentadentate macrocycle 2,13-dimethyl- 3,6,9,12,18-pentaazabicyclo[12.3.1]octadeca-1(18),2,12,14,16- pentaene, L, has been found to adopt a planar conformation, imposing a pentagonal-based geometry on metal ions of diVerent electronic configurations,5 leading to the stabilisation of seven-co-ordinate species.6 Here we report the synthesis, molecular and crystal structure, solid-state and solution electronic spectra, thermogravimetric analysis, and variable-temperature ESR study at X and Q band of the compound [MnL(H2O)2]Cl2?4H2O 1.The structure of the [MnL(H2O)2]21 cation, as its PF6 salt, has been reported,7 however no details were given nor was the precision of the X-ray determination evaluated. A more recent † E-mail: mest@servidor.unam.mx report,8 on the [MnL(H2O)2]Cl[ClO4] complex, is mainly concerned with other features and provides little discussion on the crystal structure of the manganese cation.On the other hand, the ESR spectrum of [MnL]Cl2?6H2O was the subject of an earlier report,9 however only room-temperature results were provided.Experimental Synthesis The compound was obtained by reaction of stoichiometric amounts of 2,6-diacetylpyridine (0.75 g, 4.6 mmol), 3,6-diazaoctane- 1,8-diamine (0.675 g, 4.6 mmol), and MnCl2?4H2O (0.9 g, 4.5 mmol) by using a slight modification of the synthesis reported by Humanes.10 The amine was added dropwise to a hot solution of the metal salt and diacetylpyridine in water (12.5 cm3). After a 3 h reflux the hot reaction mixture was filtered and the dark brown tarry residue discarded.The deep orange solution was then allowed to cool, rendering deep orange crystals of the compound, suitable for X-ray diVractometry [71% yield, m.p. = 267 8C (decomp.)] (Found: C, 35.46; H, 7.00; N, 13.62. C15H35Cl2MnN5O6 requires C, 35.51; H, N N H3C N CH3 N N H H L1552 J. Chem. Soc., Dalton Trans., 1998, Pages 1551–1556 6.95; N, 13.80%). IR (KBr) 11: 3358s [n(OH)], 3270s [n(NH)], 2906m [nasym(CH2)], 2852s [nsym(CH3)], 1648s [n(C]] N)], 1584m [n(C]] C)], 1458m [dasym(CH3)], 1376m [dsym(CH3)] and 546mw (OH, co-ordinated water).Crystallography Unit cell dimensions with estimated standard deviations were obtained from least-squares refinements of the setting angles of 25 carefully centred reflections. Two standard reflections monitored periodically showed no change during the data collection. A summary of important crystallographic data is presented in Table 1. Corrections were made for Lorentz-polarisation eVects. The atomic scattering factors were taken from refs. 12 and 13. The structure was solved by direct methods by using the CRYSTALS package 14 and refined by full-matrix least-squares cycles. Anisotropic thermal parameters were introduced for all non-hydrogen atoms. Hydrogen atoms were found on diVerence electron-density maps and refined isotropically. CCDC reference number 186/928. ESR measurements Since the single crystals used for the X-ray studies were not of suitable size for ESR measurements, these studies were carried out on polycrystalline samples at X and Q band on a JEOL JES-RES3X spectrometer, operating at 100 KHz and equipped with a laboratory-made X-band low-temperature accessory for the 77 K experiments. The manganese(II) complex shows a broad unresolved singlet centred at around g = 2.0.In order to obtain resolved spectra, the compound was magnetically diluted into a diamagnetic rhodium matrix {cis[RhCl2- (cyclam)]Cl} to a final concentration of 9.52% (w/w).The ESR X-band spectra were recorded at variable temperature, ranging from 300 down to 77 K. The g value was calculated from the accurate measurements of magnetic field and frequency parameters. The Q-band spectra were recorded using a JEOL ES-SQ4 microwave cavity and a ES-UTQ3 Q-band variabletemperature system. Additionally, dilutions of 1 in dimethyl sulfoxide were prepared and studied at 77 K, at X-band frequency. Theoretical spectra were calculated by means of a set of programs, run in MATLAB, especially developed for this purpose by our group.15 The method includes matrix diagonalisation and numerical exact eigenvalues, eigenfunction solutions, transition probabilities and transition fields, along the lines of the calculations made by Dowsing and Gibson,16 Griscom and Griscom,17 and more recently Pilbrow18 and Mabbs and Collison.19 Other measurements The elemental analyses (C, H and N) were carried out on a Perkin-Elmer 240B microanalyser at University College London.Infrared spectra of KBr pellets of the complex were recorded on a Perkin-Elmer 599-B spectrophotometer, in the range 4000–200 cm21, electronic spectra of the solid sample on a Cary 5E UV/VIS-NIR spectrophotometer, aqueous solutions on a Hewlett-Packard 8452A diode-array spectrophotometer. Thermogravimetric analysis was performed on a 951 Du Pont Thermogravimetric analyser. Magnetic susceptibility was measured on a Faraday balance at room temperature. Diamagnetic corrections were made by using Pascal’s constants.The set-up was calibrated with Hg[Co(SCN)4] as standard. Results and Discussion The compound was formulated as [MnL]Cl2?6H2O when first synthesized by Alexander et al.,20 considering the two chlorine atoms co-ordinated to the manganese(II) ion in CH3NO2 solutions, but dissociated in water. A second formulation, containing the [MnL(H2O)2]21 cation, was proposed by Drew et al.21 on the basis of spectroscopic and conductance studies.This cation was later found in the crystal structure of the related compound [MnL(H2O)2][PF6]2, however, as previously pointed out, only very scarce structural data were provided.7 The thermogravimetric analysis for the previously reported [MnL(H2O)2][PF6]2 compound7 does not show the same pattern as that of our complex, Fig. 1. The latter is very clear and reveals a single weight loss (21.4%, 40–117 8C) before decomposition, corresponding to all the six H2O molecules (21.3%), without diVerentiation between co-ordination and crystallisation water (even though the sample was heated at 5 8C min21).In contrast, the two co-ordinated H2O molecules in the PF6 salt 7 are lost at lower temperatures (80–100 8C). Compound 1 has a magnetic moment of 6.0 mB (mB ª 9.27 × 10224 J T21) and hence is a high-spin system. It is light yellow in aqueous solution, but gives strong absorptions at 294, 256 and 230 nm in the UV region. These bands are probably associated with charge-transfer transitions of the pyridineimine site, as pointed out previously.21 There is also a weak band (e = 140 dm3 mol21 cm21) in the visible region, at 400 nm.The ground term for a d5 configuration is the orbital singlet 6S, which cannot be split by a crystal field of any symmetry. The absence of any other spin sextet terms requires that all transitions in high-spin d5 complexes are spin- as well as Laporteforbidden. However, if the orbital angular momentum is not completely quenched, the spin–orbit interaction mixes the ground state with the first excited states and then otherwise forbidden transitions have small probabilities. Therefore, if present, they will generally be very weak,22 in complete agreement with our experimental observations.Hence, 1 is proposed to have as the ground term the orbital singlet 6S mixed with excited states. The ESR spectroscopic g value is then expected to show some anisotropy and be centred around 2.0, since it is a spin S = 5 2 – weak crystal-field system.The visible absorption obtained in the solid state at room temperature appears at 445 nm, shifted from that in aqueous solution. This shift of 45 nm indicates that the form of 1 in aqueous solution is substantially diVerent from that in the solid state, showing that the electronic structure is very sensitive to the surroundings. Crystal structure The system comprises the [MnL(H2O)2]21 cation (Fig. 2), two chloride counter ions and four crystallisation water molecules. Important lengths and angles are collected in Table 2. The cationic unit shows a slightly distorted pentagonal-bipyrimidal co-ordination geometry with the macrocycle in the pentagonal plane and two water molecules in the axial positions. The angle between the planes formed by the five nitrogen and the manganese atom and the plane including the metal and the two oxygen atoms is 89.88. The pentagonal plane is slightly distorted with N(1) being Fig. 1 Thermogram of [MnL(H2O)2]Cl2?4H2O, run under N2 and at 5 8C min21J. Chem. Soc., Dalton Trans., 1998, Pages 1551–1556 1553 0.19 Å above it, a larger deviation from planarity than that observed in a similar manganese complex derived from a 15- membered N3O2 macrocycle.23 By contrast, another manganese complex derived from a larger 17-membered macrocycle 21 shows larger deviations from planarity. The angles within the co-ordination sphere are close to expected values.All O (co-ordinated water)]Mn]N (macrocycle) angles are within 68 of 908 and the O(1)]Mn]O(2) angle Table 1 Summary of crystallographic data of [MnL(H2O)2]Cl2?4H2O Formula M Crystal symmetry Space group Crystal size/mm a/Å b/Å c/Å b/8 U/Å3 ZF (000) DiVractometer Radiation (l/Å) m/cm21 Dc/g cm23 Scan type Scan range/8 q Limits/8 T Octants collected No. data collected No. unique data No. unique data used Rint Absorption correction R = S(||Fo| 2 |Fc||)/S|Fo| R9 = [Sw(|Fo| 2 |Fc|)2/SwFo 2]� �� Goodness of fit s No.variables Drmin, Drmax/e Å23 C15H35Cl2MnN5O6 507.31 Monoclinic P21/n 0.5 × 0.5 × 0.3 11.130(1) 8.613(1) 24.883(5) 95.22(5) 2373.3(5) 4 1068 Enraf-Nonius CAD4 Mo-Ka (0.710 69) 7.97 1.42 w–2q 0.8 1 0.345 tan q 1–25 Room temperature 0–13, 210 to 0, 228 to 28 4229 3891 3500 [(Fo)2 > 3s(Fo)2] 1.78 DIFABS12 (minimum = 0.84, maximum = 1.13) 0.028 0.031 0.97 411 20.97, 0.97 Table 2 Selected bond lengths (Å), angles (8) and torsion angles (8) for [MnL(H2O)2]Cl2?4H2O Mn]O(1) Mn]N(1) Mn]N(7) Mn]N(13) O(1)]Mn]O(2) O(2)]Mn]N(1) O(2)]Mn]N(4) O(1)]Mn]N(7) N(1)]Mn]N(7) O(1)]Mn]N(10) N(1)]Mn]N(10) N(7)]Mn]N(10) O(2)]Mn]N(13) N(4)]Mn]N(13) N(10)]Mn]N(13) Mn]N(1)]C(15) Mn]N(4)]C(3) C(3)]N(4)]C(5) Mn]N(7)]C(8) Mn]N(10)]C(9) C(9)]N(10)]C(11) Mn]N(13)]C(14) Mn]N(1)]C(2)]C(3) Mn]N(1)]C(15)]C(14) Mn]N(4)]C(3)]C(2) Mn]N(4)]C(5)]C(6) Mn]N(7)]C(6)]C(5) Mn]N(7)]C(8)]C(9) Mn]N(10)]C(9)]C(8) Mn]N(10)]C(11)]C(12) 2.259(2) 2.298(2) 2.254(2) 2.316(2) 174.99(9) 87.32(8) 92.33(8) 89.01(8) 143.21(8) 96.34(8) 147.14(8) 69.59(7) 96.48(9) 147.62(8) 72.20(8) 109.7(2) 116.0(2) 123.1(2) 119.7(2) 120.0(2) 121.6(2) 108.7(2) 252.72 242.89 22.60 2.11 3.63 23.28 211.15 219.70 Mn]O(2) Mn]N(4) Mn]N(10) O(1)]Mn]N(1) O(1)]Mn]N(4) N(1)]Mn]N(4) O(2)]Mn]N(7) N(4)]Mn]N(7) O(2)]Mn]N(10) N(4)]Mn]N(10) O(1)]Mn]N(13) N(1)]Mn]N(13) N(7)]Mn]N(13) Mn]N(1)]C(2) C(2)]N(1)]C(15) Mn]N(4)]C(5) Mn]N(7)]C(6) C(6)]N(7)]C(8) Mn]N(10)]C(11) Mn]N(13)]C(12) C(12)]N(13)]C(14) Mn]N(13)]C(12)]C(11) Mn]N(13)]C(14)]C(15) N(1)]C(2)]C(3)]N(4) N(4)]C(5)]C(6)]N(7) N(7)]C(8)]C(9)]N(10) N(10)]C(11)]C(12)]N(13) N(13)]C(14)]C(15)]N(1) 2.249(2) 2.286(2) 2.289(2) 88.00(8) 84.63(8) 73.44(8) 93.68(8) 69.77(7) 88.55(8) 139.32(8) 84.09(9) 75.91(9) 140.09(8) 107.8(2) 114.3(2) 120.9(2) 119.2(2) 121.0(2) 117.1(2) 108.4(2) 115.0(3) 219.70 243.12 249.91 23.70 9.35 247.94 59.99 is 174.99(9)8.The pyridine ring forms an angle of 1.08 with the MnN5 plane, smaller than in the N3O2 complex, where the value of 5.78 was believed to be due to a greater diYculty in accommodating the manganese ion in the macrocycle cavity.23 The conformation of the macrocycle can be appreciated from the torsion angles shown in Table 2. The two five-membered rings, which include the two imine bonds [Mn]N(4)]C(5)]C(6)] N(7) and Mn]N(7)]C(8)]N(10)], are nearly planar with maximum deviations of 0.06 Å.The remaining five-membered rings, as expected, are not planar and show atoms at about 0.3 Å above and below the square planes.There are some significant diVerences in the Mn]N distances with values ranging from 2.254(2) to 2.316(2) Å, the shortest bond corresponding to Mn]N(7), the only pyridine nitrogen atom. This bond was also the shortest in the MnN3O2 compound referred to above 23 and in the iron complex derived from the same quinquedentate ligand,24 presumably due to the diVerent hybridisation of the nitrogen atom.These Mn]N distances in 1 seem to be, even considering experimental error, longer than in the mentioned manganese and iron complexes of similar macrocycle size. The manganese–oxygen distance [2.249(2) and 2.259(2) Å] have similar values to those observed for other Mn]N (macrocycle) systems, and are shorter than those observed for Mn]O (macrocycle) equatorial bonds in the MnN3O2 complex.21 X-Band ESR results at room temperature The solid line in Fig. 3(a) shows the ESR X-band powder spectrum at 300 K of complex 1 when diluted magnetically to 9.52% (w/w) in a diamagnetiium matrix. The spectrum shows five well defined lines from zero to 7500 G (G = 1024 T), but not hyperfine structure. The g value of this quintuplet is 2.111. The linewidth increases with increasing magnetic field, from 350 G for the leftmost line to 620 G for the rightmost line. Powder ESR spectra of this type have been reported previously9,25–27 for manganese(II) seven-co-ordinated complexes.Such spectra have been interpreted using the spin Hamiltonian treatment of Dowsing and Gibson,16 Griscom and Griscom 17 and Bleaney and Ingram28 and more recently Mabbs and Collison 19 and it is written in equation (1) where the first term is the Zeeman H = gBS?H 1 D[Sz 2 2 (1/3)S(S 1 1)] 1 E(Sx 2 2 Sy 2) 1 A S?I (1) interaction and the second and third terms correspond to the crystal-field splitting and the nuclear hyperfine interaction, respectively.The last term will not be considered any further Fig. 2 Molecular structure of [MnL(H2O)2]Cl2?4H2O1554 J. Chem. Soc., Dalton Trans., 1998, Pages 1551–1556 since the complex did not show hyperfine splitting; D and E are the usual axial and rhombic crystal field parameters, respectively.16–19,29 It is customary to define the parameter l = E/D, which can take values between zero and 1/3, where the extreme value l = 0 represents the totally axial case of the tensor D and the other extreme value l = 1/3 represents the maximum rhombic distortion, E = D/3, of the crystal-field tensor.Values of l outside this range reproduce the cases already included in the 0 < l < 1/3 range.16–19,29 Electron spin systems with S > ��� for which l and D are so small that they can be taken as zero give rise to ESR spectra centred around g = 2. Systems with large axial crystal fields with D > 0.25 cm21 and l = 0 show an intense feature at g = 6 and a small feature at g = 2.24,27,28 On the other hand, systems with large D and rhombicities with l in the vicinity of 1/3 give rise to ESR spectra with sharp resonances at g = 4.28.26–28 Complex 1 shows five fine-structure lines centred at g = 2.111, which are compatible only with total electron spin S = 5 2 – and parameters D and l that are small, but not zero.The high-spin value deduced from ESR is consistent with that obtained from the roomtemperature susceptibility measurements mentioned above. In consequence, the crystal electric field suggested is weak.In addition, the absence of resonances at g = 6.0 and 4.28 indicates also a weak crystal field with l far from the extreme values zero and 1/3. The measured g value of 2.111 is considerably far from 2.0, expected for S = 5 2 – systems experiencing very small crystal fields. Griscom and Griscom 17 have demonstrated that when D starts to be comparable to the Zeeman term, i.e. (1/4)gBH or larger, the central fine transition 1��� �2��� deviates from g = 2.0 towards higher g values and splits into two transitions, until it reaches the value of g = 4.3 for rhombic crystal fields, or even the value g = 6.0 for large axial crystal fields. The X-band spectrum of complex 1 in Fig. 3(a) clearly shows partial splitting and shifting towards g values higher than 2.0. Hence, for this compound, a first estimate of D > (1/4)gBH, which implies that the crystal-field eVects are not only not negligible, but are of the order of the Zeeman term, and a complete solution of equation (1) is required.On the other hand, the field positions of the five absorptions of the 300 K spectrum are consistent with the positions in the Dowsing and Gibson (DG) plots16 for 0.055 < D < 0.076 cm21 and 0.1 < l < 0.15. From these values E is estimated to be in the range 0.0055–0.0076 cm21. These first estimates of D and E were taken into account to solve the spin Hamiltonian (1) exactly, by matrix diagonalisation, since the perturbation theory cannot be applied.A computer program, using the MATLAB package, was written for this purpose.15 Fig. 3 X-Band powder ESR spectra of the [MnL(H2O)2]Cl2?4H2O complex obtained at (a) 300 K, microwave frequency 9.450 GHz, measured g value 2.111 and (b) 77 K, microwave frequency 9.0945 GHz. Sweep field 8000 G, modulation amplitude 0.5 G, gain 1000, time constant 0.03 s Exact numeric eigenfunctions, eigenenergies, transition probabilities and transition magnetic fields were calculated for the magnetic field H, parallel and perpendicular to each of the principal X, Y and Z axes.The line spectrum so calculated agreed completely with the experimental one, following standard procedures.17,19 The deduced spin Hamiltonian parameters are D = 0.07 cm21, E = 0.008 cm21 and l = 0.1142. Hence, the crystal electric field is axial with a rhombic contribution, consistent with the distorted pentagonal-bipyrimidal geometry determined from the X-ray analysis.The absence of the expected hyperfine structure from the manganese nuclear spin I = 5 2 – in the spectrum indicates that the magnetic dilution of 9.52% (w/w) in the rhodium matrix still represents a high concentration and dipole–dipole interactions prevent resolution of the hyperfine structure.30 In addition, Birdy and Goodgame31 and Wagnon and Jackels 27 have interpreted the lack of hyperfine splitting as a consequence of a relatively large amount of zero-field splitting, due to the crystal field, as is indeed the case for compound 1.It is quite interesting that this seven-co-ordinated manganese( II) complex, being an odd-electron Kramer system, shows all its five fine structure transitions well resolved, and the outer transitions have not spread out substantially. It is well known that for Kramer ions, such as MnII, a substantial broadening of the outer fine-structure transitions results from distribution in D and E and other spin Hamiltonian parameters, due to internal microscopic strains which are intrinsic to the samples.17,18,30 Such eVects do not influence the 1��� �2��� transition of Kramer systems.In this case it seems that the molecular arrangement of 1 is so regular that the distribution of D and E, in contrast to many other systems, is very sharp. In this sense the regularity allowed the definition of the ESR fine-structure lines and the X-ray determination with high precision.X-Band ESR results at low temperature The X-band ESR spectrum of complex 1 at 77 K is shown in Fig. 3(b). Several changes are quite clear: a large, negative and broad (500 G) spectral feature a9 appears at practically zero magnetic field. Also a very intense peak b9 has grown with its maximum at the apparent position g = 3.66. These changes prompted us to determine how the low-temperature features developed. Therefore, spectra were taken at several intermediate temperatures, and a set of these is presented in Fig. 4. The gradual growth of a9 and b9 [Fig. 3(b)] is clearly seen to start at a temperature of only 278 K. In addition, the rightmost lines opened up gradually with the temperature decrease to 100 K, and the position of the rightmost line shifted from 6005 to 6350 G, with concomitant increase in the axial parameter D of ª58 G, its value being ª0.0056 cm21. A temperature dependence of the axial crystal-field parameter D has been observed for several other manganese(II) systems.32 Since the low-temperature a9 and b9 features overlap greatly Fig. 4 Sequence of spectra taken every 50 K, from 300 to 100 K, of [MnL(H2O)2]Cl2?4H2OJ. Chem. Soc., Dalton Trans., 1998, Pages 1551–1556 1555 with three out of the five 300 K lines their line shapes are masked; however, the g value of 3.66, i.e. the maximum of peak b9, can be taken as representative. The b9 absorption, along with the close to zero-field absorption a9, clearly seen at low temperatures, and the DG plots,16 allow us to estimate D as 0.074 cm21 and E close to (1/3)D.These parameters were used to calculate a low-temperature line spectrum, which is consistent with all the experimental features. All these gradual, neat changes in temperature (Fig. 4) indicate a transformation, diVerent from a phase transition, since for the latter to occur abrupt changes are expected.33 Instead we propose a slight magnetostructural dependence on te the behaviour of D and E with temperature, in turn, indicates that the crystal field becomes larger and with a higher rhombicity as the temperature decreases.The increase of the intensities of the a9 and b9 features with decreasing temperature is well accounted for by the Boltzmann population factor for paramagnetic systems.18,30 In order to examine whether the diamagnetic rhodium matrix has any eVect on the magnetic behaviour described above, ESR spectra of frozen solutions (77 K) of complex 1 in dimethyl sulfoxide were obtained at X-band frequency.These spectra were practically identical to those obtained for the diamagnetic rhodium matrix dilutions of 1, i.e. the magnetic changes with temperature already described are indeed intrinsic to the manganese compound. Q-Band ESR results at room and low temperature In order to confirm the evaluation of the spin Hamiltonian parameters, the Q-band spectra of complex 1, at both 300 and 77 K, were recorded (Fig. 5). Three important transitions appear at 10 800, 12 450 and 13 850 G. it is presumed that the other two fine-structure lines could appear at ª14 650 and ª15 340 G, beyond the reach of our spectrometer. The band at 12 450 G is assigned to the central 1��� �2��� transition, with g = 2.0. Theoretical Q-band line ESR spectra were calculated by using the X-band calculated parameters, and all the lines were very well reproduced, hence confirming our evaluations. In this case, the microwave quantum hnQ value increases to 1.22 cm21, while D and E remain constant.Under this condition the crystal-field eVects are about four times smaller than in X-band, approximating more to a perturbation of the Zeeman term. Neither a zero-field transition nor a strong transition around g = 4.0 can be expected; second-order eVects of the crystal field on the g value are not expected either, and the five fine-structure transition lines should be centred very close to g = 2.0.All these features are present in the 300 and 77 K Q-band ESR spectra in Fig. 5. At Q-band there is no variation of the spectrum with temperature, confirming that the X-band temperature changes are due to crystal-field variations. Fig. 5 Q-Band powder ESR spectra of the [MnL(H2O)2]Cl2?4H2O complex obtained at 300 K, microwave frequency 35.01 GHz, measured g value 2.00, and 77 K, microwave frequency 34.99 GHz.Sweep field 14 000 G, modulation amplitude 1.2 G, gain 3200, time constant 0.03 s Conclusion We conclude that the ESR spectra indicate that at room temperature the crystal field is mainly rhombic, with l = 0.1142, a result that is consistent with the distorted pentagonalbipyramidal co-ordination geometry, determined from our X-ray experiments. As the temperature decreases D increases slightly and a larger rhombicity appears, with l = 0.162 at 77 K. The Q-band results are fully consistent with the X-band determination.It is considered that the smooth crystal-field dependence on temperature is most likely related to small structural changes with temperature. Further, that if this is the case then a small torsion of the bonding angles of the axially coordinated water molecules in compound 1 is more likely to produce a slightly distorted crystal field than a distortion involving the pentagonal-base nitrogen bonds, for these seem to be more rigid and it would require more energy to distort them.The shifting of the 400 nm UV/VIS peak in solution to 445 nm in the solid state suggests that the electronic energy levels originating the transition have become closer, which is consistent with a geometrical change around the manganese(II) ion when passing from solution to the solid state. Therefore, the electronic structure of 1 seems to be very sensitive to the surroundings. Acknowledgements The authors are grateful to E.Basurto for his help with the software used in this work. M. E. S. T. and O. J. S. are grateful for the economical support to the DGAPA-UNAM research project IN213794. We would also like to thank Mrs. C. Vázquez for the thermogravimetric analysis. D. R. R. and R. Z. U. are grateful to CONACYT for the acquisition of the ESR spectrometer. M. J. R. is grateful to CONACYT for the acquisition of an X-ray diVractometer. References 1 D. R. Williams, The Metals of Life, Van Nostrand-Reinhold, London, 1971; I.Bertini, H. B. Gray, S. J. Lippard and J. S. Valentine, Bioinorganic Chemistry, University Science Books, Lill Valey, CA, 1994. 2 G. W. Brudvig and R. H. Crabtree, Proc. Natl. Acad. Sci. USA, 1986, 83, 4586. 3 D. L. Kepert, Prog. Inorg. Chem., 1979, 25, 41; M. G. B. Drew, A. H. bin Othman, W. E. Hill, P. McIlroy and S. M. Nelson, Inorg. Chim. Acta, 1975, 12, L25; M. G. B. Drew, J. De O. Cabral, M. F. Cabral, F. S. Esho and S. M. Nelson, J. Chem.Soc., Chem. Commun., 1979, 1033; M. M. Bishop, J. Lewis, T. D. O’Donoghue, P. R. Raithby and J. N. Ramsden, J. Chem. Soc., Dalton Trans., 1980, 1390. 4 L. F. Lindoy, The Chemistry of Macrocyclic Ligand Complexes, Cambridge University Press, Cambridge, 1989, p. 9. 5 S. M. Nelson, Pure Appl. Chem., 1980, 52, 2461. 6 S. M. Nelson and D. H. Busch, Inorg. Chem., 1969, 8, 1859; E. Fleischer and S. Hawkinson, J. Am. Chem. Soc., 1967, 89, 720; M. G. B. Drew and S. M. Nelson, Acta Crystallogr., Sect.A, 1975, 31, S140; D. H. Cook, D. E. Fenton, M. G. B. Drew, A. Rodgers, M. McCann and S. M. Nelson, J. Chem. Soc., Dalton Trans., 1979, 414; M. G. B. Drew, A. H. bin Othman, S. G. McFall and S. M. Nelson, J. Chem. Soc., Chem. Commun., 1975, 818; M. G. B. Drew, J. Grimshaw, P. D. A. McIlroy and S. M. Nelson, J. Chem. Soc., Dalton Trans., 1976, 1388; M. G. B. Drew, A. H. bin Othman, P. D. A. McIlroy and S. M. Nelson, Acta Crystallogr., Sect. B, 1976, 32, 1029; M.G. B. Drew, A. H. bin Othman and S. M. Nelson, J. Chem. Soc., Dalton Trans., 1976, 1394. 7 F. Begun, M. S. Khan, S. Z. Haider, K. M. A. Malik, F. K. Khan and M. B. Hursthouse, J. Bangladesh Acad. Sci., 1991, 15, 185. 8 Y. Nishida, N. Tanaka, A. Yamazaki, T. Tokii, N. Hashimoto, K. Ide and K. Iwasawa, Inorg. Chem., 1995, 34, 3616. 9 A. Van Heuvelen, M. D. Lundeen, H. G. Hamilton, jun. and M. D. Alexander, J. Chem. Phys., 1969, 50, 489. 10 M. M. Humanes, Ph.D. Dissertation, University of Lisbon, 1984. 11 K. Nakanishi, Infrared Absorption Spectroscopy, Holden-Day, San Francisco, 2nd edn., 1977, pp. 42–51; K. Nakamoto, Infrared and Raman Spectra of Inorganic and Coordination Compounds, Wiley, New York, 4th edn., 1986, pp. 227–229. 12 N. Walker and D. Stuart, Acta Crystallogr., Sect. A, 1983, 39, 158.1556 J. Chem. Soc., Dalton Trans., 1998, Pages 1551–1556 13 International Tables for X-Ray Crystallography, Kynoch Press, Birmingham, 1974, vol. 4. 14 D.J. Watkin, C. K. Prout, R. J. Carruthers, R. J. Betteridge and P. Betteridge, CRYSTALS, Issue 10, Chemical Crystallography Laboratory, Oxford, 1996. 15 E. Basurto-Uribe, Simulación de Espectros EPR de Compuestos Lantánidos, Master Degree Thesis, ESFM-IPN, México, 1996. 16 R. D. Dowsing and J. F. Gibson, J. Chem. Phys., 1969, 50, 294. 17 D. L. Griscom and R. E. Griscom, J. Chem. Phys., 1967, 47, 2711. 18 J. R. Pilbrow, Transition Ion Electron Paramagnetic Resonance, Clarendon Press, Oxford, 1990. 19 F. E. Mabbs and D. Collison, Electron Paramagnetic Resonance of d Transition Metal Compounds, Elsevier, Amsterdam, 1992. 20 M. D. Alexander, A. Van Heuvelen and H. G. Hamilton jun., Inorg. Nucl. Chem. Lett., 1970, 6, 445. 21 M. G. B. Drew, A. H. bin Othman, S. G. McFall, P. D. A. McIlroy and S. M. Nelson, J. Chem. Soc., Dalton Trans., 1977, 438. 22 A. B. P. Lever, Inorganic Electronic Spectroscopy, Elsevier, Amsterdam, 2nd end., 1984, p. 446. 23 M. G. B. Drew, A. H. bin Othman, S. G. McFall, P. D. A. McIlroy and S. M. Nelson, J. Chem. Soc., Dalton Trans., 1977, 1173. 24 M. G. B. Drew, A. H. bin Othamn and S. M. Nelson, J. Chem. Soc., Dalton Trans., 1994, 1394. 25 R. D. Dowsing, J. F. Gibson, D. M. L. Goodgame, D. Goodgame and P. J. Hayward, J. Chem. Soc. A, 1969, 1242. 26 J. E. Newton and S. C. Jackels, J. Coord. Chem., 1988, 19, 265. 27 B. K. Wagnon and S. C. Jackels, Inorg. Chem., 1989, 28, 1923. 28 B. Bleaney and D. J. E. Ingram, Proc. R. Sor. A, 1951, 205, 336. 29 R. D. Dowsing, J. F. Gibson, M. Goodgame and P. J. Haygood, J. Chem. Soc. A, 1970, 1133. 30 A. Abragam and B. Bleaney, Electron Paramagnetic Resonance of Transition Ions, Clarendon, Oxford, 1970. 31 R. B. Birdy and M. Goodgame, Inorg. Chem., 1979, 18, 172. 32 K. L. Wan, S. L. Hutton and J. E. Drumheller, J. Chem. Phys., 1987, 86, 3801; J. E. Drumheller and R. S. Rubins, J. Chem. Phys., 1986, 85, 1699; C. Benelli, D. Gatechi, C. Zanchini, R. J. Doedens, M. H. Dickman and L. C. Porter, Inorg. Chem., 1986, 25, 3453. 33 A. B. Pippard, Classical Thermodynamics, Cambridge University Press, Cambridge, 1966. Received 16th February 1998; Paper 8/01292J

ISSN:1477-9226    

DOI:10.1039/a801292j

出版商:RSC    

年代:1998

数据来源: RSC