Metal-Only Lewis Pairs with Transition Metal Lewis Bases - American

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Metal-Only Lewis Pairs with Transition Metal Lewis Bases Jürgen Bauer, Holger Braunschweig,* and Rian D. Dewhurst Institut für Anorganische Chemie, Julius-Maximilians-Universität Würzburg, Am Hubland, 97074 Würzburg, Germany but it would not be due to either of these reasons. Aside from its reversible dioxygen coordination, the remarkable reactivity of Vaska’s complex toward acids, halides, and dihydrogen was quickly established, yielding discrete, stable IrIII addition products.2 This meant that, providing one had access to an iridium chloride salt, synthetic laboratories worldwide were merely one step and less than half a day away from having an air-stable CONTENTS compound that required little special handling but still 1. Introduction A underwent oxidative addition reactions. Furthermore, the 1.1. The History of the Metal−Metal Dative reactions could be conveniently monitored via the carbonyl Bonding Concept A C−O stretching band in the IR spectrum and later by 31P NMR 1.2. What Is a Dative Bond and How Do We spectroscopy. No element−element bond was safe, although Recognize One? B the stability of Vaska’s complex, which had made it so 1.3. Scope of this Review B convenient to use, also limited its reactivity. Vaska’s complex 2. Complexes with d-Block Metal → s/p-Block Metal became a well-defined molecular platform for understanding Dative Bonding C the oxidative addition reactions that make both homogeneous 2.1. Complexes with s-Block Acceptors C and heterogeneous catalysis possible. Consequently, the term 2.2. Complexes with Group 13 Acceptors D “oxidative addition” is now ubiquitous in inorganic chemistry 2.3. Complexes with Group 14 Acceptors E and catalysis. However the underlying concept of this elemental 2.4. Complexes with Borderline or Supported dstep, transition metal Lewis basicity, appears very rarely in the Block → p-Block Bonding E literature. 3. Complexes with d-Block → d/f-Block Metal There are a number of ways a metal can break an element− Dative Bonding F element bond, ranging from the classically invoked concerted, 3.1. Introduction: Bridged and Other Borderline side-on cleavage to backside attack reminiscent of an organic Cases F SN2 reaction. For instance, the mechanism of oxidative addition 3.2. Complexes with Group 6 Metal Bases G of sp2-carbon−halide bonds can be completely different from 3.3. Complexes with Group 7 Metal Bases H that of sp3-carbon−halide bonds.3 For substrates with 3.4. Complexes with Group 8 Metal Bases H appropriate empty orbitals (e.g., boranes), prior M → E 3.5. Complexes with Group 9 Metal Bases K donation may precede a concerted oxidative addition, 3.6. Complexes with Group 10 Metal Bases L complicating matters.4 Many oxidative addition reactions are 3.7. Analysis of Structural Data N simply accepted as such: the manner of the bond breakage by 4. Conclusions O the metal is not analyzed. However, regardless of the Author Information P mechanism, the metal must act as a nucleophile or Lewis Notes P base. In light of this, it appears pertinent to consider how a Biographies P metal acts as a Lewis base, and how its Lewis basic properties Acknowledgments P can be tuned. We conceived this review in part to highlight the References Q Lewis basicity of transition metals in order to better understand what makes a good transition metal Lewis base. The ability of transition metal carbonyl complexes to act as 1. INTRODUCTION proton acceptors (Brønsted bases) was noted as early as 1928 by Hieber,5 while his later work uncovered a number of anionic 1.1. The History of the Metal−Metal Dative Bonding metal carbonyl complexes with strong nucleophilicity.6 Early Concept organometallic chemists slowly began to recognize that In 1961, a report in the Journal of the American Chemical Society transition metal complexes need not be charged in order to detailed the synthesis of a lemon-yellow compound asssigned as 1 be nucleophilic and that oxidative addition reactions were the organometallic complex trans-[IrCl(CO)(PPh3)2]. The linked to high electron density, and thus Lewis basicity, of the results contained in this communication were novel in a metal center. Consequently, a number of groups began to target number of ways. The authors, Vaska and DiLuzio, had reduced Lewis acid−base adducts of metals with small molecules.7−12 iridium(III) to iridium(I) using just triphenylphosphine and an alcohol. They had prepared an iridium carbonyl complex without the addition of carbon monoxide. The name Vaska was Received: January 3, 2012 to become one of the most well-known in inorganic chemistry, © XXXX American Chemical Society

A

dx.doi.org/10.1021/cr3000048 | Chem. Rev. XXXX, XXX, XXX−XXX

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however, in terms of transition metals, fragments which may act as Lewis bases in a complex may not be recognizable as such when free, or may be unstable without the presence of the Lewis acid. The gas-phase minimum-energy rupture concept21 is useful; however this requires that the complexes be treated computationally, very few of which have been. Theoretical calculations in general would assist greatly in the study of these molecules; however most of the complexes included herein have not been the subject of such a study. Mössbauer spectroscopy can also aid understanding;23 however this is limited to only a handful of metals and has not been performed on the complexes we include herein. Experimental electron density methods have been used to a similar end, but again, this technique is an uncommon one and has not yet been performed on the complexes in this review.24 X-ray absorption near edge structure (XANES) spectroscopy has recently been used by the group of Gabbaï to determine the oxidation states in Au−Sb complexes.25 This, although it requires synchrotron radiation, is a promising method for determining oxidation states of metals and thus the presence or absence of dative bonding. Overall, none of these experimental or theoretical methods have been applied extensively enough to suspected metal−metal dative-bonding complexes and are thus only of utility in discussing the individual complexes to which they have been applied. Consequently, our inclusion or exclusion of complexes in this review is based on formalism, despite its admittedly simplistic nature. Our rationale is discussed in the following section.

The result, reports of transition metal adducts of Lewis acids such as BF3, BH3, O2, SO2, and tetracyanoethylene, kick-started the concept of transition metal Lewis basicity (although syntheses of unsupported metal−borane adducts were later discounted,13−16 and the existence of such a complex is yet to be convincingly demonstrated). To the best of our knowledge, the first mention of metal-tometal “dative” bonding occurs in a 1964 report from Coffey, Lewis, and Nyholm, exemplified by the Ni−Ni interactions in bis(dimethylglyoximato)nickel(II).17 However the first definitive evidence for unsupported TM → TM bonding came in 1967, with the X-ray molecular structure of [(η5-C5H5)(OC)2Co→HgCl2] reported by Nowell and Russell.18 In the authors’ own words: (the) compound is best described as a Lewis acid−base complex between (η5-C5H5)Co(CO)2 and mercuric chloride. Metal−metal bonded complexes have long been the subject of intense interest.19 However, no efforts have been made to compile complexes with metal−metal dative bonds in which the base is a transition metal, although there are many scattered reports of such complexes. In many cases, these complexes were not described in terms of “Lewis base−Lewis acid” pairs in the original reports, and indeed, in some cases this description may reflect only a formalism and not the actual bonding interaction. 1.2. What Is a Dative Bond and How Do We Recognize One?

The IUPAC Goldbook20 defines a dative bond in the following way: “The coordination bond formed upon interaction between molecular species, one of which serves as a donor and the other as an acceptor of the electron pair to be shared in the complex formed, e.g., the N→B bond in H3N→BH3. In spite of the analogy of dative bonds with covalent bonds, in that both types imply sharing a common electron pair between two vicinal atoms, the former are distinguished by their significant polarity, lesser strength, and greater length. The distinctive feature of dative bonds is that their minimum-energy rupture in the gas phase or in inert solvent follows the heterolytic bond cleavage path.” This minimum-energy rupture concept, developed by Haaland, involves analysis of the energy of the products of breaking the bond in question by either homolytic or heterolytic means.21 Put simply, if homolytic bond rupture is more energetically favorable, the bond can be considered covalent; if heterolytic rupture is lower in energy, the bond can be considered dative. Distinguishing between covalent and dative bonding is sometimes thought of as determining where the electrons came from before the bond was formed; in contrast, Haaland’s concept looks instead at where the electrons go after the bond is broken. The former can in some cases be determined qualitatively from what reagents are added to a reaction vessel (e.g., if a known Lewis base is aded to a known Lewis acid), but the latter is a more generally applicable concept. When we address transition metals, the picture becomes more uncertain, because electronegativity differences, ionic/electrostatic, covalent, dative, and metallophilic/closed-shell22 components may all contribute to the overall bonding. This brings us to methods for identifying dative bonds. In the IUPAC definition alone, there are two possibilities for recognizing a dative bond: if one atom is a donor and one an acceptor of electrons, and if their minimum-energy rupture pathway is heterolytic. Recognizing fragments that are known to be donors or acceptors is one way to define a dative bond,

1.3. Scope of this Review

There exists a number of related families of complex that have been independently reviewed and will not be covered herein, such as transition metal Lewis base complexes of main group nonmetals (particularly boron),26−29 main group metal Lewis base complexes with other main group metals,30 and early−late metal complexes, which involve metal−metal multiple bonding, some of which is presumed to be dative.31 Complexes with ligands bridging the metal−metal bond will only be briefly addressed in this review, because we focus on complexes in which the dative bond is the sole element holding the metal centers together. Because it is difficult to definitively exclude the presence of bridging or supporting ligands by spectroscopy alone, we have also chosen to cover only crystallographically characterized complexes. However, this ignores the possibility that bridging ligand interactions may be present in the solution state. As such, we have used the Cambridge Crystallographic Database as the primary method for compiling the complexes herein, because this is the most efficient way to establish the presence of a metal−metal bond and the absence of bridging elements. Since a great variety of metal−metal bonded complexes exist, the definition of such a dative bond becomes somewhat blurry, particularly in the case of charged systems. Many of the examples in this review are cationic species containing group 11 and 12 elements as “acceptor” groups. In these cases, we feel it is justified to formally assign the positive charge to the group 11/12 element. For example, the positive charge in the relatively simple complex [{(η5-C5H5)(OC)2Rh}2Ag]+ is most logically assigned to the Ag center (Figure 2). If we depict dative bonds from the Rh centers to Ag, the oxidation states of the metals are formally RhI→AgI←RhI. A mixed dative− covalent description leads to a RhI→Ag0−RhII picture, while with two covalent bonds the complex must be formally RhII− B


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qualifying lithium compound is the Li-bridged dinuclear compound [{(Et3P)Pt(CCPh)2}2(μ-Li)2] from the group of Wrackmeyer.35

Figure 1. Portion of the periodic table showing the metallic elements and their roles (Lewis acid/Lewis base) in the compounds considered herein.

Ag−I−RhII. We favor the former description involving dative bonds for two reasons: (a) the complex and others like it are synthesized using the independently stable monovalent group 9 complexes [M(CO)2(η5-C5R5)] (M = Co, Rh, Ir; R = H, Me), known to be Lewis-basic, and (b) the RhII, Ag0, and Ag−I oxidation states are uncommon in organometallic complexes compared with the much more common RhI and AgI states. Note, throughout this chapter the metal−metal distances, dMM, are compared with the sum of the experimentally-derived covalent radii of the atoms involved, ∑covrad, as sourced from a survey of the Cambridge Crystallographic Database.32 This ratio is denoted by drel throughout.

Figure 3. Complexes featuring alkali metal fragments as Lewis acids.

When it comes to complexes of group 2 displaying unsupported linkages to transition metals, the number of compounds increases. However, most of these compounds again show a high proportion of electrostatic interaction. For example, the complex [CpCo(η3-C3H5)→MgBr(thf)2] (3), synthesized by Jonas (Figure 4),36 has a Lewis-basic fragment that is also negatively charged [CpCo(η3-C3H5)]−. Therefore, determination of the bonding situation is not clear-cut, being somewhere between Lewis donation and ionic bonding. Another example is the iridium complex [{Cp*IrH(PMe3)}2(μ-MgPh)2] from Bergman and Anderson,37 in which the two magnesium fragments adopt a bridging position between two Lewis-basic metal fragments. Nevertheless, to the best of our knowledge, there are only two related complexes that have unambiguously dative M−M bonds without generating charges on the fragments. These examples are the platinum complex [Cy3P)2Pt→BeCl2] (4) and its secondary product [(Cy 3 P) 2 Pt→BeClMe] (5), both reported by Braunschweig (see Figure 4).38 Density functional theory (DFT) calculations performed on these two complexes by Parameswaran and Frenking indicated the M→Be (M = Ni, Pd, Pt) bond strength to be similar to dative bonds between Be and ammonia. Additionally, the contribution of the orbital and electrostatic components to the overall stability of the complexes were found to be relatively similar, and both σ

2. COMPLEXES WITH d-BLOCK METAL → s/p-BLOCK METAL DATIVE BONDING 2.1. Complexes with s-Block Acceptors

s-Block metals reveal the least-pronounced tendency to form adducts with transition metal bases. However, there are a few crystallographically determined examples of group 1 elements with an unsupported bond to a late transition metal. The limitation for these compounds is that they all have as a common feature a negatively charged transition metal center and a group 1 cation such as Li+ or Na+. Therefore we propose that the character of these M−M bonds could be best described as a predominantly electrostatic interaction. This is in accordance with the description of the bonding situation by the authors who first reported these complexes. The two crystallographically characterized examples of sodium compounds are [Cp*(tmbp)Ru→Na] (1) by Mathey and Le Floch,33 and [(tBu3Si)2Cu→Na(thf)2] (2) by Lerner and coworkers (see Figure 3).34 The only known example of a

Figure 2. Example complex used to illustrate the oxidation state rationale used frequently herein. Here, and throughout this review, CO ligands are depicted as bonds terminated with dots “●”. C


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Figure 4. Complexes featuring alkaline earth metal fragments as Lewis acids.

(strong) and π (weak) M→Be dative bonds were present in the molecules.39 2.2. Complexes with Group 13 Acceptors

In contrast to the few known s-block Lewis pairs, there are many adducts with p-block acids. Unsurprisingly, row 13 metals are among the best-investigated acids in terms of adduct formation with transition metal complexes. Most of these adducts are known for the lightest metallic representative, aluminum. Among these adducts are platinum complexes such as [(Cy3P)2Pt→AlCl3] (6), [(ItBu)(Cy3P)Pt→AlCl3] (9), [(SIMes)(Cy3P)Pt→AlCl3] (10), or [(SIMes)2Pt→AlCl3] (7),40,41 as well as an example with palladium, [(ItBu)(Cy3P)Pd→AlCl3] (8), synthesized by Braunschweig (Figure 5).42 In addition to these fully characterized complexes is an early report of the (structurally unconfirmed) adduct [Cp(Me3P)2Rh→AlMe3] by Mayer.43 An isolobal anionic species, [Et4N][Cp(OC)2Fe→AlPh3], was synthesized by Burlitch and Hughes.44 More difficult to describe is the compound [Cp*(Me3P)(H)2Ir→AlPh3] by Bergman and Anderson.37 The bridging nature of the hydrogen atoms remains unclear; however it is less pronounced than in the complex [Cp2W(H)2(AlMe3)], where no direct transition metal-main group metal bond is apparent, as determined by the X-ray analyses of Caulton.45 However, DFT calculations performed by Lelj on this complex indicated that there is significant charge donation from the transition metal to the aluminum atom.46 In the case of gallium, most complexes featuring gallium− metal interactions bear the {GaCp*} moiety, which itself can be described as a Lewis base. The two neutral examples in which the gallium center acts as a Lewis acid are the complexes [Cp*(Cp*Ga)2Rh→GaCl3] (11) synthesized by Fischer47 and [(Cy3P)2Pt→GaCl3] (12) by Braunschweig (Figure 5).48 In addition to the neutral complexes, there is one example of an anionic complex, K[(OC)4Co→GaCl3] (K[13]), reported by Fischer and theoretically investigated by Frenking,49 whereas the complex Na2[(OC)4Fe→(GaCl3)2] (Na2[14]) from the group of Scheer serves as example of a dianionic bis(gallane) complex.50 Even fewer examples are known for indium, where the favored bonding situation is the indyl ligand, stabilized by main group bases. An example of a transition metal base-stabilized

Figure 5. Complexes featuring group 13 metal (Al, Ga) fragments as Lewis acids.

indyl ligand is the bimetallic cobalt complex [{(OC)4Co→}2(InCl2)]−, reported by Norman.51 However, examples of crystallographically characterized adducts are the dianionic complexes [Ph4P]2[(OC)5M→InCl3] (M = Cr, [Ph4P]2[15]; Mo, [Ph4P]2[16]; W, [Ph4P]2[17]; Figure 6) reported by Huttner.52 Another well-investigated dianionic complex is [(Me2N)3C]2[(CO)4Fe→InCl3] ([(Me2N)3C]2[18]), synthesized by Neumüller and calculated by Frenking.53 Examples of monoanionic indane complexes are the iron compound [Cp*Fe(C7H8)][Cp*(OC)2Fe→InCl3] ([Cp*Fe(C7H8)][19]) from the group of Aldridge54 and the D


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Fernández,58 again the question arises of how much of the bond character is due to electrostatic interaction between the anionic transition metal fragment and the thallium cation (see Figure 6). Frenking et al. presented in 2009 a theoretical approach for the understanding of the bonding situation in complexes of group 10 bases with group 13 acids.59 They introduced the socalled “reversed Dewar−Chatt−Duncanson” model, describing a σ donation from the transition metal center and a π backdonation from the group 13 element. This is the exact opposite of the “normal” Dewar−Chatt−Duncanson model, which is used to describe donor−acceptor complexes of transition metals with main group donors like alkenes or phosphines.60,61 The theoretical investigations include bis(phosphine) complexes of nickel, palladium, and platinum with acids of the entire group 13. 2.3. Complexes with Group 14 Acceptors

To conclude the p-block elements, the only relevant M−M Lewis adducts outside group 13 contain the heavier group 14 metals. There are, for instance, complexes with interactions of an unclear nature, such as the bridging SnCl2 moieties in [{(Me3P)3ClRh→}2(μ-SnCl2)] (24),62 reported by Chan and Marder, or [{(Cy3P)(OC)(I)Pt}2(μ-SnCl2)] (25), synthesized by Ros and Roulet (Figure 7).63 However, the structural

Figure 7. Complexes featuring group 14 metal fragments as Lewis acids. Figure 6. Complexes featuring group 13 metal (In, Tl) fragments as Lewis acids. Ar = 2,6-(2,4,6-Me3C6H2)2C6H3.

parameters of 25, namely, the tetrahedral geometry at tin, suggest the presence of covalent Pt−Sn bonds, whereas the analysis of the NMR spectroscopic parameters of the transition metal fragments indicate a dative interaction. Therefore, the description as a dative bond is again just one possibility. An unambiguous, albeit not crystallographically determined, example of an interaction of a Lewis-basic transition metal with the acid SnCl4 is the adduct [Cp(Me3P)2Co→SnCl4], reported by Werner.64

nickel complex [HNC7H13][Cp(OC)Ni→InBr3] ([HNC7H13][20]), which was a byproduct in the synthesis of transition metal-substituted indanes by Fischer.55 Further evidence for complexes featuring an In(III) ligand stabilized by a transition metal Lewis base is given in the aforementioned work of Burlitch and Hughes, who reported, in addition to the iron− aluminum adduct, a series of anionic metal complexes, for example, [(Ph3P)2N][(OC)4Co→InPh3], which were not characterized by single-crystal X-ray diffraction analyses.44 When it comes to the heaviest group 13 metal, thallium, again more examples are known. In contrast to the lighter homologues, the favored oxidation state of the thallium is I, whereas in all other group 13 adducts the acidic centers are found in the oxidation state III. However, the TlI cation is coordinated in cationic complexes [(PPh2py)3Pt→Tl][NO3] ([21][NO3], py = pyridine, Figure 6) reported by Catalano56 and [(ArNC)3Ni→Tl][OTf] ([22][OTf] Ar = 2,6-(2,4,6Me3C6H2)2C6H3) by Figueroa.57 In the case of the neutral complex [(Cl5C6)2Au→Tl(η6-toluene)] (23), synthesized by

2.4. Complexes with Borderline or Supported d-Block → p-Block Bonding

A related, more intensively studied field is the family of supported group 14 adducts displaying dative bonds from late transition metals to silicon or tin. Known examples of silatranes are the palladium and the platinum compounds by Wagler65,66 (see Figure 8). Related stannatranes have also been reported, together with additional complexes with gold as an acidic center, by Maron, Gabbai,̈ and Bourissou.67 The last example of a main group acid contains interactions of the bridging {PbCl2} moiety with two Cr fragments, of which the nature is unclear, in E


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Figure 8. Complexes featuring group 14 metal fragments as Lewis acids.

while My is used to define metals in the y oxidation state). Seyferth proposed that this interaction was a M8→M10 dative bond. At the prompting of referees, the authors also considered a second possibility: the cyclic ligands as cyclopentadien(thi-/ selen-)ones and the bridging ligands as neutral donors to both M and M′. This corresponds to a four-electron intramolecular reduction of the metals, now both formally zero-valent, with the ligands providing two electrons each. The addition of phosphine to the Pd center cleaves the Pd−Fe bond, which can be thought of as the Lewis base phosphine replacing the Lewis base iron, implying the Lewis acidity of the Pd center. Thus the former situation is favored. This problem of assigning meaningful and accurate oxidation states in metal complexes with conjugated ketone, thione, or imine ligands has been highlighted by the extensive work of Wieghardt and many others, and these complexes present a similar conundrum.74 However, since the metals are still formally in the same oxidation state, the direction of the electron donation, if any, is unclear. Further examples include neutral donor L-groups such as -PtBu2 and -NPPh3, providing a situation with zero-valent M10 and divalent M8 atoms. In this case the M0 center is clearly the more electron-rich and could be considered to be a Lewis donor to M8. These metallocene complexes thus provide ample warning against overinterpretation of the nature of the M10−M8 bond, particularly in bridged systems. A dinuclear Nd/Rh complex featuring amidopyridine ligands was reported by Kempe to possess possible M−M dative bonding (Figure 9, above right).75 The geometry can be described as a d8 square planar arrangement of the rhodium atom with an additional apical interaction with the Nd center, suggesting donation of electrons from the Rh dz2 orbital to Nd. A similar complex featuring a purported Pd→Nd interaction is also reported. 2-Phosphinopyridines, having only one atom between the two donor atoms P and N, often act as ditopic ligands to two different metals. Zhang, Che, and Mak reported a dinuclear Ir/ Cd complex in which the phosphine ligands were bound to Ir and the Cd was bound by both pyridine nitrogens (Figure 9, bottom left).76 Despite the ligands bridging the two metals, the Ir−Cd distance (ratio of M−M distance to sum of covalent radii: 0.976) is not noticeably shorter than in unsupported examples (vide infra). In the following sections, reports of monotopic behavior of these ligands are described, leading to complexes that appear in the solid state to be unsupported M− M Lewis adducts. One of the more convincing examples of bridged TM−TM dative bonding is a Zr2Pd trinuclear complex, in which two phosphinotrozircene “metalloligands” are bound to a single Pd0 atom (Figure 9, middle left).77 The Pd environment is reminiscent of bis(phosphine) complexes of Pd, known to be

the dianionic compound [Ph4P]2[{(OC)5Cr}2(μ-PbCl2)] (Figure 8), reported by Huttner.68

3. COMPLEXES WITH d-BLOCK → d/f-BLOCK METAL DATIVE BONDING 3.1. Introduction: Bridged and Other Borderline Cases

Metallocene complexes containing a second metal in the bridging segment, first published by Seyferth (Figure 9, above left), are classical examples of ambiguous metal−metal bonding.69−73 When the ligands, L, are chalcogen atoms, they can be considered either as alkoxide/thiolate/selenolate ligands to M10, in which case M10 is necessarily divalent (throughout the text the notation Mx is used to define metals of group x,

Figure 9. Selected ligand-bridged complexes with possible dative M→ M bonds. F


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excellent TM Lewis bases, except that the Pd geometry is nonlinear and unsually situated close to one Zr center. DFT calculations on an intermolecular model featuring one trozircene and one [Pd(PMe3)2] fragment (Figure 9, middle right) showed a definite interaction: the Zr−Pd distance in this case was calculated to be ca. 5% shorter than that of the bridged derivative. Although the authors indicate that the interaction is predominantly noncovalent, they state that dissociation of this bond costs 10.5 kcal/mol and that such an adduct may be isolable. Apart from the very recent use of XANES techniques by Gabbaı̈ and co-workers,25 there are few experimental diagnostic tools for determining whether a metal−metal bond is covalent, dative, or a mixture of the two. Thus we and the groups referenced herein have mostly relied on valence electron counting, oxidation states, and the known reactivity (i.e., Lewis basic/acidic) of the constituent fragments to identify M→M dative bonds. A consequence is that only the more clear-cut cases of M→M dative bonding will be covered systematically in this review, although other systems with “borderline” interactions exist. One notable case is Power’s recent syntheses of complexes with unsupported M−Fe bonds (M = Fe, Mn, Cr) protected by the bulky aryl group {C6H-2,6-(C6H2-2,4,6iPr3)2-3,5-iPr2} (Figure 9).78 While by our definition, these complexes would not contain formal M→M dative bonds, DFT calculations showed the HOMO of these species to contain a significant Fe→M dative bonding interaction. 3.2. Complexes with Group 6 Metal Bases

To the best of our knowledge, no qualifying TM−TM Lewis adducts are known featuring M3−5 metals, generally considered to be less electron-rich TM centers, as the base fragment. The rarity of genuine neutral Lewis-basic complexes of group 6 and 7 metals is presumably responsible for the dearth of Lewis adducts based thereon. A large disparity in the electron density of two bonded TM atoms leads us to assume some extent of dative bonding between the two tungsten atoms of complex 26, in which one W fragment is formally zero-valent with 16 valence electrons and the other is tetravalent with 10 valence electrons (Figure 10).79 The strongly σ-basic and weakly π-acidic ligand sphere around the W0 atom (i.e., four trialkyl phosphines, only one CO) is further evidence for the basicity of this fragment. However, the extremely short W−W distance (drel = 0.779) implies multiple bonding between the two metals, which the authors suggest could be either double or triple. This bond distance makes 26 a true outlier in the systems studied herein and might be better thought of as a nondative case. Two crystallographically characterized examples of M6→M11 adducts are shown in Figure 10. The assignment of a Mo→Ag dative bond from the protonated trinuclear Mo2Ag complex 27 is based more on formalism than evidence. The “Lewis donor” Mo fragment can be thought of as a zero-valent Mo center bound by a neutral boranaphthalene ring in which the boron atom provides one π-electron and is additionally stabilized by an exocyclic amine acting as a Lewis base to the boron atom.80 Balch and co-workers attempted to exploit the bridging ability of ditopic ligands connected by one-atom units to prepare a dinuclear MoAu complex.81 By adding [(Ph3P)AuCl] to a Mo0 complex with two bis(diphenylphosphino)methane ligands in air, the authors obtained instead the nonbridged cationic dinuclear complex [28][PF6] (Figure 10), in which one phosphorus atom of each ligand had been oxidized to a

Figure 10. Complexes featuring group 6 and 7 Lewis bases.

phosphine oxide. This oxidation extended the ligands by one atom, making them now capable of forming stable fivemembered chelates with one metal and disfavoring bridging coordination modes. The two CO ligands of the molecule appear slightly bent away from the Au atom, perhaps indicating a small amount of bridging character. In 1968, following on from the synthesis of metal complexes with (HgX) “ligands”, Lewis and co-workers reported the synthesis of Lewis acid−base adducts of Cr, Mo, W, and Fe with mercuric halides (HgX2).82 A large range of these complexes were prepared and, while not structurally characterized, were some of the earliest TM−TM complexes to be explicitly described as Lewis adducts. Similarly, Snow reported the simple addition of mercuric chloride to [MoG


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Figure 11. Complexes featuring group 8 Lewis bases, part one.

3.4. Complexes with Group 8 Metal Bases

(CO)3(mesitylene)] to give a Mo→Hg adduct 29 (Figure 10), which crystallizes as a dimer bridged by two extra HgCl2 units, held together by Hg···Cl interactions.83

The addition of acetonitrile to the complex [(H3N)xYbFe(CO)4] provided the YbFe adduct 31 (Figure 11).85 Two carbonyl oxygen atoms from each Fe center donate to the oxophilic Yb, resulting in polymeric ladders of octahedral Yb and trigonal bipyramidal Fe centers. The linearity of the carbonyl ligands prompted the authors to assume the absence of bridging elements and a more pure Fe→Yb dative bond. The anionic 18-valence-electron fragments [Fe(CO)4]2− and [FeH(CO)3L]− (L = CO, P(OMe)3) were used by Darensbourg and co-workers as Lewis bases for donating to the Lewis-acidic 16-valence electron fragments [M(CO)5] (L = Cr, W; Figure 11), producing the adducts [NEt4]2[32], [PPN][33], [PPN][34], and [PPN][35] (PPN = bis-

3.3. Complexes with Group 7 Metal Bases

Similarly to the Mo complex described above by Snow, cymantrene and its monophosphine derivative [(η5-C5H5)Mn(CO)2L] (L = CO, PPh3) form adducts with the divalent mercurial [Hg(O2CCF3)2].84 When L = PPh3, the adduct is stable, leading to the crystallographically characterized bis(trifluoroacetato)-bridged dimer 30 (Figure 10). When L = CO, the adduct is unstable and goes on to form a cymantrene derivative with a permercurated cyclopentadienyl ligand. H


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(triphenylphosphine)iminium).86 Although the metal−metal bond of [NEt4]2[32] is unbridged in the solid state, its 13C NMR spectrum shows one single resonance in the carbonyl region at 23 °C, which splits into three in a 1:4:4 ratio at −80 °C. This suggests the presence of extensive fluxionality of the carbonyls in solution, perhaps including bridging bonding modes. Group 8 metal ions are arguably not the most electron-rich transition metals; however TM−TM Lewis adducts featuring these elements as the basic fragment are the most wellrepresented such class. This can be attributed to the prolific work of Pomeroy and co-workers, who have used M8 carbonyl fragments to great effect in the synthesis of metal-only Lewis pairs. By liberating one carbonyl ligand from M6 hexacarbonyl complexes, it is possible to prepare the Lewis-acidic complexes [M(CO)5(thf)] (M = Cr, W). Pomeroy and co-workers found that a range of 18-valence-electron pentacoordinate group 8 complexes, when added to the group six complexes, form complexes with M8→M6 dative bonds, 36−45 (Figure 11).87−90 Each one of these complexes has a longer M−M distance than the sum of the experimental covalent radii of the elements (i.e., a drel value greater than unity), with M→Cr examples being notably longer than M→W examples. From these data, we can see that M8→M6 as well as M8→M8 (vide infra) adducts are the subsets with the longest M−M dative bonds of all complexes considered herein. The Lewis adduct complex [(Me3P)(OC)4Os→ReBr(CO)4] (46, Figure 12) and covalently bound complex [(Me3P)(OC)3BrOs−Re(CO)5] are a pair of constitutional isomers that presented a unique opportunity to contrast dative and covalent metal−metal bonds.91 The former complex has a slightly longer Os−Re distance (almost 1%) than the latter. While the difference is small, it is statistically significant, causing Pomeroy and co-workers to tentatively attribute this to the differing nature of the M−M bond. The observation suggests that dative metal−metal bonds are weaker than covalent ones; however, analysis of the drel values of complexes in this review shows the average M−M distance to be shorter than comparable covalent bonds (vide infra). A similar approach was used by Pomeroy to prepare complexes with M8→M8 dative bonds (47−51, 53, Figure 12).92,93 Zero-valent, pentacoordinate group 8 complexes were added to divalent bis(trichlorosilyl) or (chloro)(trichlorogermyl) group 8 complexes, providing the Lewis adducts. In all cases, one anionic group 14 ligand is situated trans to the M−M dative bond. A similar complex, 52, featuring a bis(trifluoroacetate) RuII Lewis acidic fragment was prepared by Funaioli and Fachinetti by adding carbon monoxide to the doubly bridged bis(trifluoroacetate) diruthenium complex [Ru2(μ:η2-O2CCF3)2(CO)6].94 This complex shows a distinctly shorter M−M bond, perhaps due to the strongly electronwithdrawing trifluoroacetate ligands on the Lewis-acidic metal center or the presence of a strongly π-acidic carbonyl ligand trans to the M−M bond. Curiously, despite the aptitude of group 11 metals to act as Lewis acids, there exists only two M8→M11 Lewis adduct complexes. Both involve, as above, 18-valence-electron (VE) M8 fragments. Complex 54 (Figure 12) was prepared by Trogler by the simple addition of silver trifluoroacetate to [Os(CO)3(PPh3)2].95 The X-ray structure of 54 revealed a distinct orientation of two CO ligands toward the silver atom; however, the CO ligands themselves appear linear and are thus presumably not involved electronically with the silver center.

Figure 12. Complexes featuring group 8 Lewis bases, part two.

Adduct 55, while also featuring an 18VE osmium Lewis base, is much more unusual given the structure of the acidic fragment, the cluster [AuOs3(μ2-Cl)(CO)10].96 The adduct is the result of treatment of the [Os4(μ-Cl)(CO)13]− anion with [Au(PPh3)]+. The net increase in the number of carbonyl ligands in the product suggests that the mechanism of the reaction is significantly complicated. The M8→M12←M8 dianions 56−59 (Figure 13) have a long history and have been prepared by a number of groups.97−101 Their dianionic character leads to considerable confusion regarding their electronic structure and the presence of dative bonding between the metals. While the synthetic route to these complexes leads to dipotassium salts K2[56/58/59], crystal data were obtained for [PPN]2[56], [Yb(py)5(NCMe)2][56], [Yb(DME)2(NCMe)2][56], [Na(THF)2][57], [PPN]2[58], and [Ph3PMe]2[59] (py = pyridine; DME = 1,2-dimethoxyethane). More clear-cut from a formalism viewpoint are the cationic M8→Hg←M8 adducts [59−62][BF4]2 (Figure 13), consisting of two neutral, divalent M8 metallocenophane bases I


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Figure 13. Complexes featuring group 8 Lewis bases, part three.

both forming dative bonds to a HgII center.102 The complexes are prepared by adding the neutral metallocenophane to [Hg(CN)2] with addition of tetrafluoroboric acid. In the earlier section on supported metal−metal dative bonds, the short-spacer ditopic 2-phosphinopyridine ligands were shown to be excellent at holding two metal centers together. However in a number of cases, these ligands, and an analogous morpholino-substituted phosphine ligand, act as monotopic or nearly monotopic ligands in the solid state, supporting Fe→Hg and Ru→Hg dative bonds in the complexes 63, 64, and [65][HgCl3] (Figure 13).103−105 In these three

complexes, either one or two weak N···Hg interactions can be found (2.730−2.832 Å). Complexes 66,106 67,107,108 68−70,109,110 and 71111 (Figure 13) all feature unambiguous M8→Hg dative bonds by virtue of their mercuric dihalide fragments, the X−Hg−X angle being in all cases significantly bent. However, complicating matters is the presence of weak intermolecular mercury−halide interactions in 68−70. The drel values for the complexes 65−71 fall into an astoundingly narrow range (0.964−0.973) considering that the metal and structure of the donor vary considerably. J


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Figure 14. Double Lewis adducts featuring M→M→M donations.

This appears unaffected by the presence or absence of intermolecular mercury−halide interactions. During the course of their studies on metal−metal dative bonding, Pomeroy and co-workers uncovered two highly unusual trinuclear complexes, which they described as having two M−M dative bonds in tandem (72, 73; Figure 14).112,113 The classification of the Os−Os bond as dative in both 72 and 73 was based, in part, on the observation of high-field 13C NMR signals corresponding to the carbonyl trans to the Os− Os bond (72, δ 165.6; 73, δ 162.9). These high-field signals had been found by Pomeroy in earlier dinuclear M−M dative bonded species. 3.5. Complexes with Group 9 Metal Bases

Derivatives of the [(η5-C5R5)M9(CO)L] (M = Co, Rh, Ir) fragment are strong transition metal Lewis bases; thus it is no surprise that this fragment features often in complexes with Lewis acids of groups 6 and 8−12. The M−M distances in these complexes are relatively long; only two have a drel value below 1.0. This appears not to depend on the M9 center present, the nature of the R group, nor the size or Lewis basicity of the L ligand. Iridium complexes 74−77114,115 (Figure 15) were prepared by the group of Pomeroy and use Lewis acidic fragments that feature in other M−M Lewis adducts from the same research group, [W(CO)5] and [OsX2(CO)3]. Considerable carbonyl fluxionality was noted for these complexes, suggesting that bridging CO ligands may be present in solution at room temperature. The unusual M9→M9 adduct 78 (Figure 15) was prepared by Del Paggio and co-workers by the reaction of [(η5-C5Me5)Ir(CO)2] with [Rh(μ-Cl){P(OiPr)3}]2, resulting in the exchange of one CO with one phosphine ligand between the two metals.116 Although both metals are monovalent group nine centers, the dative bond in this complex was assigned on the basis of the demonstrated basicity of [(η5-C5R5)M(CO)L] fragments, as well as by comparison with square planar, Vaska’stype complexes. Thus the Ir fragment can be thought of as replacing the Lewis-basic phosphine ligand in complexes of the form trans-[M9Cl(CO)(PR3)2]. In contrast, the M9−M10 bond in 79 (Figure 15) comfortably fits a dative description given the formal oxidation state of the Pt center (PtII). Trinuclear M9→Ag←M9 complexes [80][PF6], [81][PF6], and [82][BF4] were prepared by the addition of the basic fragment to Ag[PF6] or Ag[BF4] (Figure 16).117−119 While [80][PF6] and [82][BF4] both display drel values akin to those seen in Figure 15, the value for [81][PF6] is significantly lower (drel = 0.850). This could be attributed to the overall stronger Lewis basicity of PPh3 relative to CO and P(OiPr)3. Besides the [(η5-C5R5)M(CO)L] fragments, there are a number of other monovalent M9 Lewis bases that form


complexes with M 11 cations containing dative bonds. Complexes [83][PF6]120 and [84][PF6]121 (Figure 16) feature terminal hydrido ligands on the M9 center and either chelating or monodentate phosphine ligands. When a tetraphosphine ligand of the form “PP3” was used instead of “NP3” in [84][PF6], the hydrido ligand was found to be bridging. The carborane complexes 85122 and 86123 were prepared by the group of Reed in the course of their development of weakly coordinating anions. In each case an Ir→Ag bond is present, with the Ag+ ion being stabilized by weak interactions with B− Br or B−H bonds of the carborane counterion. The cationic complex [Ir(dppe)2][BF4] was used in the synthesis of the K


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dative interaction from the cobalt. Like many earlier complexes featuring mercuric dihalides as Lewis acidic fragments, long intermolecular Hg···Cl interactions were observed. Additionally, Balch and Catalano prepared the unusual tetranuclear complex 89 by adding excess mercuric chloride to [AuIr(CO)Cl(μ-dppm)2][PF6].126 3.6. Complexes with Group 10 Metal Bases

The Pt→Zr complex 90 (Figure 18) has a very low drel value (0.812), yet the complex contains no bridging elements.127


dicationic IrAg complex [87][BF4]2.124 If we consider the complex to be a simple adduct of Ir+ and Au+ fragments, we can assume an Ir→Au dative interaction is present. As mentioned earlier, the Co→Hg complex 88 (Figure 17) was perhaps the first definitive example of an unsupported M→


This makes the complex distinct from the only other unsupported adduct containing an early transition metal Lewis acid, the Fe→Yb complex 31 (drel = 0.944, Figure 11). Both fragments are stable independently; the demonstrated Lewis basicity of the zero-valent [Pt(PCy3)2] base and the Lewis acidity of the tetravalent complex [ZrCl4] allowed us to assign an unambiguous dative interaction in this case. The pentanuclear complexes 91−95 (Figure 18) were prepared by Ishihara and Matsumoto in 2007.128 Although it is difficult to assign oxidation states to the metal centers in these complexes, the synthesis of the complexes, that is, addition of half an equivalent of [MX′4]2− to a cationic diplatinum complex, led the authors to describe these complexes as Lewis adducts. The complexes 96,129 [97][ClO4],130 [98][ClO4],130 [99][BF4]3,131 [100][BF4]3,132 [NBu4][101],133 and [NBu4][102]133 (Figure 19) are all variations of complexes featuring square-planar PdII/PtII Lewis bases and M11 centers. Within this family there is considerable M→M distance variation: the drel values of the complexes range from 0.895 in the neutral 96 to 1.096 in the tricationic complex [99][BF4]3, suggesting that the overall charge of the complex may have some bearing on the M→M distance (this is explored further in the analysis section). Furthermore, [97][ClO4] and [98][ClO4] display one-dimensional polymeric structures with long Ag···Pt interactions.


M dative bond.18,125 The early structural evidence for this complex was provided in a communication in 1967, and a more accurate structure was provided in 1972. The Co−Hg distance is practically identical to the sum of the experimental covalent radii; however, the Cl−Hg−Cl angle (114.0(2)°) suggests a L


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A rare example of a complex with a nickel Lewis base is 104 (Figure 20), with a relatively large drel value of 1.023.135 The complex exists as a dimer with long Hg···I ineractions. The dicationic complexes [105][ClO4]2 and [106][ClO4]2, featuring Pt→Cd dative bonds, were prepared by Yamaguchi and Ito,136 and a related neutral complex 107 was reported 5 years

The Pt→Au complex [103][BF4] (Figure 19) was obtained from the reaction of [Pt(PPh3)4] and in situ-generated [Au(PPh3)][BF4].134 According to the drel value of this complex, the zero-valent nature of the Pt fragment appears not to intensify the Pt→Au bonding, compared with the other M10→M11 complexes described above. M


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bases to construct adducts with [HgX]+ fragments, resulting in the complexes [111][OTf],139 [112][OTf],139 and [113][BPh4] (Figure 20).140 The trinuclear complex [112][OTf] is prepared by salt elimination from a derivative of [111][OTf] with Na[Mn(CO)5], showing that the original M→M dative bond is strong enough not to be perturbed by strong nucleophiles. The Pt→Hg complexes 114141 and 115142 (Figure 20) resemble the aforementioned complexes with mercuric halides as the Lewis acidic components, in that both possess long Hg···Cl interactions. The drel values of these two complexes are unusually large, despite unimposing steric bulk of the Pt fragments. 3.7. Analysis of Structural Data

Given the enormous structural diversity of complexes treated in this review and the timespan over which they were published, the only representative piece of experimental data common to all complexes is the metal−metal distance. By dividing this number by the sum of the covalent radii (as derived from published X-ray structures in the Cambridge Crystallographic Database32), we obtain a value that compares the M−M distance with the distances expected for a single covalent bond between the same two atoms (drel). We realize that these covalent radii values come with limitations. The values were derived from determination of the radii of carbon, nitrogen, and oxygen, and these values were used for determining the radii of transition metals by surveying complexes with predominantly dative interactions between these three elements and the metal. The estimated standard deviation (esd) values of the radii for the transition metals are 0.04−0.09 Å, which must be added together to gain the esd of the ∑covrad value. Thus we urge caution in the interpretation of this analysis and stress that the drel value is not a measure of the dMM related to the “true” covalent radii of an element (a somewhat meaningless construct), but a relation of the dMM to the average metal− ligand distances in complexes with C-, N- or O-ligands. Nevertheless, these crystallographically derived covalent radii are the most appropriate values for us to use, and we feel important information can be gained from the analysis. As a first step, we sought to determine how the experimental M−M distance (dMM) correlates with the literature covalent radii (∑covrad). This is depicted in Figure 21 in three graphs: one colored according to the group number of the donor metal, one colored according to the group number of the acceptor metal, and one colored according to the charge on the complex. The correlation in this case is relatively weak (correlation coefficient = 0.659, see Table 1). The majority of complexes treated herein have a drel value below unity, and the average drel value is 0.980, indicating that in general, the M−M distance of the metal-only Lewis pairs considered in this review is around 2% shorter than that in complexes with single covalent metal− metal bonds. From Figure 21a,b, we can observe trends for the different metal triads in the periodic table for both donors and acceptors. For instance, group 10 metals create relatively short M−M distances irrespective of their role in the complex (i.e., donor or acceptor). Group 9 metals lead to very short M−M distances as donors but appear rarely as acceptors. These values are offset by a number of complexes with group 8 donor metals with relatively long M−M distances. In the Lewis acid role, group 6 and 8 metals lead to long M−M distances (drel almost exclusively above unity), while complexes with acceptor metals of groups 10−12 in general produce relatively short M−M


later by Forniés and Lalinde.137 All three complexes feature a Lewis-acidic [Cd(cyclen)]2+ (cyclen =1,4,7,10-tetraazacyclododecane) fragment and square-planar Pt centers. The dicationic complexes have lower drel values than 107. This contraction with increasing net charge is similar to that seen in the complexes in Figure 19 above. The Pt→Hg←Pt complexes [108−110][PF6]2 (Figure 20) are unusual in that the Lewis acidic Hg center does not sit at an apex of square-planar Pt fragments, the position where a filled dz2 orbital often can donate to a Lewis acid.138 Instead, the Hg is effectively the fourth ligand in two square-planar Pt complexes. Thus, we could envisage two alternative bonding situations with covalent bonds: one in which one positive charge is assigned to each Pt center instead, leading to a PtII− Hg−II−PtII description, and one where both charges are on the mercury giving a PtI−Hg0−PtI complex. Group 10 metal centers ligated by tetradentate NP3 or PP3 ligands have been shown above to be excellent Lewis bases. Reports from 1989 and 2001 use these transition metal Lewis N


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Table 1. Average drel Values for Subsets of the Complexes and Correlations between M−M Distance or drel Value and Chemical Parametersa entry 1 2 3 4 entry 5 6 7 8 9 10 11 12 13 14 15 16 17 18

avg drel

subset all complexes neutral complexes anionic complexes cationic complexes correlation

0.980 0.987 0.986 0.947 correlation coefficient

Overall Correlations dMM/∑covrad Correlations with Group Numbers drel/Δ(group number of TM) drel/|Δ(group number of TM)| drel/donor metal group number drel/acceptor metal group number Correlations with VE Counts drel/Δ(fragment VE count) drel/donor fragment VE count drel/acceptor fragment VE count dMM/donor fragment VE count dMM/acceptor fragment VE count Correlations with Coord. Numbers drel/donor coord. number drel/acceptor coord. number dMM/donor coord. number dMM/acceptor coord. number

0.659 −0.273 −0.265 0.002 −0.266 −0.590 0.322 0.705 0.244 0.726 0.129 0.522 0.186 0.772

a

The strongest correlations (correlation coefficient >0.7) are shown in bold.

bonds. Figure 21c groups the complexes by their net charge, showing that the data points of neutral (average drel = 0.987) and anionic (average drel = 0.986) complexes appear relatively evenly scattered above and below the equivalence line, while the points for cationic complexes lie exlusively below the line. This is reflected in the average drel value for cationic complexes (0.947) being ca. 4% lower than neutral and anionic complexes. Next we looked for correlations between the drel value of a complex and its chemical makeup. Thus, we sorted the complexes according to a number of different parameters and graphed each one separately against drel or dMM, quantifying the correlation (Table 1). Parameters were tested based on the group numbers of the metals (entries 6−9), the valence electron counts of the fragments (entries 10−14), and the coordination numbers of the fragments (entries 15−18; facecapping η5 or η6 ligands were counted as occupying three coordination sites). The best correlations were found between the drel value and the acceptor VE count (entry 12, coefficient = 0.705), the dMM value and the acceptor VE count (entry 14, coefficient = 0.726), and the dMM value and the acceptor coordination number (entry 18, coefficient = 0.772).

4. CONCLUSIONS In some ways, using the descriptors dative or covalent to describe a bond is like using a person’s birth certificate to determine where they currently live, both are concerned only with origin and not current reality. There is little real distinction between dative and covalent bonds, other than the origin of the two bonding electrons. Neither descriptor says anything about the current position of these electrons or the net partial charge on the two atoms when bound: dative bonds have degrees of polarity, as do covalent bonds. We use the concept here in the

Figure 21. Graphs of the sum of the literature covalent radii of the relevant atoms vs measured M−M distance: (a) grouped by donor metal group number; (b) grouped by acceptor metal group number; (c) grouped by net complex charge. Black diagonal lines indicate where ∑covrad = dMM (i.e., where drel = 1).

O


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degree with Professor Holger Braunschweig in the fields of late transition metal complexes as Lewis bases and the synthesis of unsupported organometallic Lewis adducts.

search for complexes with formal, unsupported dative bonds between metal atoms and thereby identify the best donor (Lewis base) and acceptor (Lewis acid) fragments. Obviously, those that reappear often in this review can be thought of as “good” donors (e.g., the group 9 fragment [(η5-C5R5)M(CO)L], [Pt(PCy3)2]) or acceptors (e.g., [BeX2], [AuL]+, [HgX2], [AlX3]). However, there may be many more that are underrepresented not because of poor donor or acceptor ability but for practical reasons: a fragment may be too sterically bulky to interact with its Lewis counterpart or may strongly promote bridging behavior in ligands and thus be excluded from this review. Alternatively, a fragment may simply be less convenient to prepare synthetically. In reality, very little is known about the M−M bonding in most of these complexes. Most of these complexes were prepared when quantum chemical calculations and computing ability were unable to deal with such large and complicated molecules. However, from measuring the M−M distances of the unsupported complexes in this review, one can see that these dative bonds are only very slightly shorter (ca. 2%) than those considered to be “covalent” and are thus practically indistinct. The key to gaining meaningful understanding of the dative metal−metal bond is theory and computation; however, it will require significant effort to determine whether a real difference exists between the bonding in metal-only Lewis pairs and conventional covalent metal−metal bonded complexes. The recent application of XANES to determine the oxidation ̈ is state of metals in metal−metal-bonded systems by Gabbai25 one promising experimental technique that may shed some light on the true bonding situation. We are following developments in this area with keen interest. Overall, we do not intend this review to be the “final word” on the subject of metal−metal dative bonding but merely an attempt to compile complexes that may have M−M dative bonding, bonding that in many cases was not recognized as such by the original authors. We hope that the review will spur lively debate on what constitutes a dative metal−metal bond and what does not.

Holger Braunschweig (born 1961 in Aachen) obtained his Ph.D. (1991) and Habilitation (1998) from the RWTH Aachen with P. Paetzold and stayed for a postdoctoral appointment with M. F. Lappert, FRS, at the University of Sussex, Brighton. After two years at Imperial College as Senior Lecturer and Reader, he moved to a chair for inorganic chemistry at the Julius-Maximilians-University Würzburg in 2002. In 2009, he was awarded the Gottfried Wilhelm Leibniz prize of the DFG, was elected as a member to the Bavarian Academy of Sciences, and became member of the National Academy of Sciences (Leopoldina) in 2011. His research interests lie in the area of boron chemistry, organometallic synthesis, and catalysis and are currently focused on borametallocenophanes, boron heterocycles, boron−boron multiple bonds, and transition metal complexes of boron.

AUTHOR INFORMATION Notes

The authors declare no competing financial interest. Biographies

Rian Dewhurst obtained his B.Sc. (Hons) (University of Canterbury, New Zealand) degree in 2002 and his Ph.D. (Australian National University) in 2006 after completion of his research with Professor Anthony F. Hill, for which he was awarded the J. G. Crawford Medal of the ANU. After a postdoctoral stay in the research group of Professor Guy Bertrand (2006−2007, University of California, Riverside), he took up an Alexander von Humboldt Postdoctoral Fellowship in the group of Professor Holger Braunschweig (2007− 2009, Julius-Maximilians-University of Würzburg). Since this time, Dr. Dewhurst has been an independent researcher at the JuliusMaximilians-University of Würzburg, funded by a grant from the Deutsche Forschungsgemeinschaft (DFG).

ACKNOWLEDGMENTS

Jürgen Bauer graduated with a diploma degree in chemistry from the Julius-Maximilians-University of Würzburg in 2009, for which he received the Faculty Prize. Supported by the Fonds der Chemischen Industrie with a fellowship, he is currently working toward his doctoral

Financial support from the Deutsche Forschungsgemeinschaft and the Fonds der Chemischen Industrie is gratefully P


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(43) Mayer, J. M.; Calabrese, J. C. Organometallics 1984, 3, 1292. (44) Burlitch, J. M.; Leonowicz, M. E.; Petersen, R. B.; Hughes, R. E. Inorg. Chem. 1979, 18, 1097. (45) Bruno, J. W.; Huffman, J. C.; Caulton, K. G. J. Am. Chem. Soc. 1984, 106, 444. (46) Amati, M.; Lelj, F. Can. J. Chem. 2009, 87, 1406. (47) Steinke, T.; Fischer, R. A.; Winter, M; Cokoja, M.; Gemel, C. Dalton Trans. 2005, 55. (48) Braunschweig, H.; Gruß, K.; Radacki, K. Inorg. Chem. 2008, 47, 8597. (49) Fischer, R. A.; Miehr, A.; Hoffmann, H.; Rogge, W.; Boehme, C.; Frenking, G.; Herdtweck, E. Z. Anorg. Allg. Chem. 1999, 625, 1466. (50) Leiner, E.; Hampe, O.; Scheer, M. Eur. J. Inorg. Chem. 2002, 584. (51) Clarkson, L. M.; McCrudden, K.; Norman, N. C.; Farrugia, L. J. Polyhedron 1990, 9, 2533. (52) Rutsch, P.; Renner, G.; Huttner, G.; Sandhö fer, S. Z. Naturforsch. 2002, 57b, 757. (53) Esser, M.; Neumüller, B.; Petz, W.; Uddin, J.; Frenking, G. Z. Anorg. Allg. Chem. 2000, 626, 915. (54) Aldridge, S.; Kays, D. L.; Bunn, N. R.; Coombs, N. D. Main Group Met. Chem. 2005, 28, 201. (55) Weiß, J.; Frank, A.; Herdtweck, E.; Nlate, S.; Mattner, M.; Fischer, R. A. Chem. Ber. 1996, 129, 297. (56) Catalano, V. J.; Bennett, B. L.; Noll, B. C.; Muratidis, S. J. Am. Chem. Soc. 2001, 123, 173. (57) Fox, B. J.; Millard, M. D.; DiPasquale, A. G.; Rheingold, A. R.; Figueroa, J. S. Angew. Chem., Int. Ed. 2009, 48, 3473. (58) Fernández, E. J.; Laguna, A.; López-de-Luzuriaga, J. M.; Monge, M.; Montiel, M.; Olmos, E. Inorg. Chem. 2007, 46, 2953. (59) Goedecke, C.; Hillebrecht, P.; Uhlemann, T.; Haunschild, R.; Frenking, G. Can. J. Chem. 2009, 87, 1470. (60) Dewar, M. Bull. Soc. Chim. Fr. 1951, 18, C79. (61) Chatt, J.; Duncanson, L. A. J. Chem. Soc. 1953, 2939. (62) Chan, D. M. T.; Marder, T. B. Angew. Chem., Int. Ed. 1988, 27, 442. (63) Béni, Z.; Ros, R.; Tassan, A.; Scopelliti, R.; Roulet, R. Dalton Trans. 2005, 315. (64) Dey, K.; Werner, H. J. Organomet. Chem. 1977, 137, C28. (65) Wagler, J.; Hill, A. F.; Heine, T. Eur. J. Inorg. Chem. 2008, 4225. (66) Wagler, J.; Brendler, E. Angew. Chem., Int. Ed. 2010, 49, 624. (67) Gualco, P.; Lin, T.-P.; Sircoglou, M.; Mercy, M.; Ladeira, S.; Bouhadir, G.; Prez, L. M.; Amgoune, A.; Maron, L.; Gabbaï, F. P.; Bourissou, D. Angew. Chem., Int. Ed. 2009, 48, 9892. (68) Rutsch, P.; Huttner, G. Z. Naturforsch. 2002, 57b, 25. (69) Seyferth, D.; Hames, B. W.; Rucker, T. G.; Cowie, M.; Dickson, R. S. Organometallics 1983, 2, 472. (70) Akabori, S.; Kumagai, T.; Shirahige, T.; Sato, S.; Kawazoe, K.; Tamura, C.; Sato, M. Organometallics 1987, 6, 526. (71) van Leeuwen, P. W. N. M.; Zuideveld, M. A.; Swennenhuis, B. H. G.; Freixa, Z.; Kamer, P. C. J.; Goubitz, K.; Fraanje, J.; Lutz, M.; Spek, A. L. J. Am. Chem. Soc. 2003, 125, 5523. (72) Mann, G.; Shelby, Q.; Roy, A. H.; Hartwig, J. F. Organometallics 2003, 22, 2775. (73) Metallinos, C.; Tremblay, D.; Barrett, F. B.; Taylor, N. J. J. Organomet. Chem. 2006, 691, 2044. (74) Chaudhuri, P.; Verani, C. N.; Bill, E.; Bothe, E.; Weyhermüller, T.; Wieghardt, K. J. Am. Chem. Soc. 2001, 123, 2213. (75) Kempe, R.; Noss, H.; Irrgang, T. J. Organomet. Chem. 2002, 647, 12. (76) Kuang, S.-M.; Xue, F.; Zhang, Z.-Z.; Xue, W.-M.; Che, C.-M.; Mak, T. C. W. J. Chem. Soc., Dalton Trans. 1997, 3409. (77) Büschel, S.; Jungton, A.-K.; Bannenberg, T.; Randoll, S.; Hrib, C. G.; Jones, P. G.; Tamm, M. Chem.Eur. J. 2009, 15, 2176. (78) Lei, H.; Guo, J.-D.; Fettinger, J. C.; Nagase, S.; Power, P. P. J. Am. Chem. Soc. 2010, 132, 17399. (79) Chisholm, M. H.; Folting, K.; Kramer, K. S.; Streib, W. E. Inorg. Chem. 1998, 37, 1549.

acknowledged. J.B. thanks the Fonds der Chemischen Industrie for a doctoral scholarship.

REFERENCES (1) Vaska, L.; DiLuzio, J. W. J. Am. Chem. Soc. 1961, 83, 2784. (2) Vaska, L.; DiLuzio, J. W. J. Am. Chem. Soc. 1962, 84, 679. (3) Mollar, C.; Besora, M.; Maseras, F.; Asensio, G.; Medio-Simón, M. Chem.Eur. J. 2010, 16, 13390. (4) Zeng, G.; Sakaki, S. Inorg. Chem. 2011, 50, 5290. (5) Hieber, W.; Sonnekalb, F. Ber. Deutsch. Chem. Ges. 1928, 61, 558. (6) Hieber, W. Adv. Organomet. Chem. 1970, 8, 1. (7) Vaska, L. Acc. Chem. Res. 1968, 1, 335. (8) Collman, J. P.; Roper, W. R. Adv. Organomet. Chem. 1968, 7, 53. (9) Shriver, D. F. Acc. Chem. Res. 1970, 3, 231. (10) Vaska, L. Inorg. Chim. Acta 1971, 5, 295. (11) Werner, H. Pure Appl. Chem. 1982, 54, 177. (12) Werner, H. Angew. Chem., Int. Ed. 1983, 22, 927. (13) Shriver, D. F. J. Am. Chem. Soc. 1963, 85, 3509. (14) Braunschweig, H.; Wagner, T. Chem. Ber. 1994, 127, 1613. (15) Braunschweig, H.; Wagner, T. Z. Naturforsch., B 1996, 51, 1618. (16) Braunschweig, H.; Kollann, C. Z. Naturforsch., B 1999, 54, 839. (17) Coffey, C. E.; Lewis, J.; Nyholm, R. S. J. Chem. Soc. 1964, 1741. (18) Nowell, I. N.; Russell, D. R. J. Chem. Soc., Chem. Commun. 1967, 817. (19) Structure and Bonding: Metal-Metal Bonding; Parkin, G., Ed.; Springer-Verlag: Berlin Heidelberg, 2010. (20) Compendium of Chemical Terminology (Goldbook); International Union of Pure and Applied Chemistry: 2011, Vol 2.3. (21) Haaland, A. Angew. Chem., Int. Ed. 1989, 28, 992. (22) Pyykkö, P. Chem. Rev. 1997, 97, 597. (23) Long, G. J. Mössbauer Spectroscopy Applied to Inorganic Chemistry; Fackler, J. P., Jr., Ed.; Plenum Press: New York, 1984. (24) Mebs, S.; Kalinowski, R.; Grabowsky, S.; Förster, D.; Kickbusch, R.; Justus, E.; Morgenroth, W.; Paulmann, C.; Luger, P.; Gabel, D.; Lentz, D. J. Phys. Chem. A 2011, 115, 1385. (25) Wade, C. R.; Lin, T.-P.; Nelson, R. C.; Mader, E. A.; Miller, J. T.; Gabbaï, F. P. J. Am. Chem. Soc. 2011, 133, 8948. (26) Hill, A. F.; Owen, G. R.; White, A. J. P.; Williams, D. J. Angew. Chem., Int. Ed. 1999, 38, 2759. (27) Braunschweig, H.; Dewhurst, R. D.; Schneider, A. Chem. Rev. 2010, 110, 3924. (28) Braunschweig, H.; Dewhurst, R. D. Dalton Trans. 2011, 549. (29) Amgoune, A.; Bourissou, D. Chem. Commun. 2011, 859. (30) Schulz, S.; Kuczkowski, A.; Schuchmann, D.; Flörke, U.; Nieger, M. Organometallics 2006, 25, 5487. (31) Gade, L. H. Angew. Chem., Int. Ed. 2000, 39, 2658. (32) Cordero, B.; Gómez, V.; Platero-Prats, A. E.; Revés, M.; Echeverría, J.; Cremades, E.; Barragán, F.; Alvarez, A. Dalton Trans. 2008, 2832. (33) Rosa, P.; Ricard, L.; Mathey, F.; Le Floch, P. Organometallics 2000, 19, 5247. (34) Lerner, H.-W.; Scholz, S.; Bolte, M. Organometallics 2001, 20, 575. (35) Sebald, A.; Wrackmeyer, B.; Theocharis, C. R.; Jones, W. J. Chem. Soc., Dalton Trans. 1984, 747. (36) Jonas, K.; Koepe, G.; Krüger, C. Angew. Chem., Int. Ed. 1986, 25, 923. (37) Golden, J. T.; Peterson, T. H.; Holland, P. H.; Bergman, R. G.; Andersen, A. R. J. Am. Chem. Soc. 1998, 120, 223. (38) Braunschweig, H.; Gruß, K.; Radacki, K. Angew. Chem., Int. Ed. 2011, 48, 4239. (39) Parameswaran, P.; Frenking, G. J. Phys. Chem. A 2010, 114, 8529. (40) Braunschweig, H.; Gruß, K.; Radacki, K. Angew. Chem., Int. Ed. 2007, 46, 7782. (41) Bauer, J.; Braunschweig, H.; Brenner, P.; Kraft, K.; Radacki, K.; Schwab, K. Chem.Eur. J. 2010, 16, 11985. (42) Bauer, J.; Braunschweig, H.; Damme, A.; Gruß, K.; Radacki, K. Chem. Commun. 2011, 47, 12783. Q


Chemical Reviews

Review

(80) Braunstein, P.; Cura, E.; Herberich, G. E. J. Chem. Soc., Dalton Trans. 2001, 1754. (81) Balch, A. L.; Noll, B. C.; Olmstead, M. M.; Toronto, D. V. Inorg. Chem. 1992, 31, 5226. (82) Edgar, K.; Johnson, B. F. J.; Lewis, J.; Wild, S. B. J. Chem. Soc. (A) 1968, 2851. (83) Ciplys, A. M.; Geue, R. J.; Snow, M. R. J. Chem. Soc., Dalton Trans. 1976, 35. (84) Kuzmina, L. G.; Ginzburg, A. G.; Struchkov, Y. T.; Kursanov, D. N. J. Organomet. Chem. 1983, 253, 329. (85) Deng, H.; Shore, S. G. J. Am. Chem. Soc. 1991, 113, 8538. (86) Arndt, L. W.; Darnesbourg, M. Y.; Delord, T.; Bancroft, B. T. J. Am. Chem. Soc. 1986, 108, 2617. (87) Einstein, F. W. B.; Jones, T.; Pomeroy, R. K.; Rushman, P. J. Am. Chem. Soc. 1984, 106, 2707. (88) Davis, H. B.; Einstein, F. W. B.; Glavina, P. G.; Jones, T.; Pomeroy, R. K.; Rushman, P. Organometallics 1989, 8, 1030. (89) Shipley, J. A.; Batchelor, R. J.; Einstein, F. W. B.; Pomeroy, R. K. Organometallics 1991, 10, 3620. (90) Jiang, F.; Jenkins, H. A.; Biradha, K.; Davis, H. B.; Pomeroy, R. K.; Zaworotko, M. J. Organometallics 2000, 19, 5049. (91) Einstein, F. W. B.; Jennings, M. C.; Krentz, R.; Pomeroy, R. K.; Rushman, P.; Willis, A. C. Inorg. Chem. 1987, 26, 1341. (92) Jiang, F.; Male, J. L.; Biradha, K.; Leong, W. K.; Pomeroy, R. K.; Zaworotko, M. J. Organometallics 1998, 17, 5810. (93) Einstein, F. W. B.; Pomeroy, R. K.; Rushman, P.; Willis, A. C. J. Chem. Soc., Chem. Commun. 1983, 854. (94) Funaioli, T.; Marchetti, F.; Fachinetti, G. Chem. Commun. 1999, 2043. (95) Song, L.; Trogler, W. C. Angew. Chem., Int. Ed. 1992, 31, 770. (96) Cathey, C.; Lewis, J.; Raithby, P. R.; Ramírez de Arellano, M. C. J. Chem. Soc., Dalton Trans. 1994, 3331. (97) Pierpont, C. G.; Sosinsky, B. A.; Shong, R. G. Inorg. Chem. 1982, 21, 3247. (98) Sosinsky, B. A.; Shong, R. G.; Fitzgerald, B. J.; Norem, N.; O’Rourke, C. Inorg. Chem. 1983, 22, 3124. (99) Alvarez, S.; Ferrer, M.; Reina, R.; Rossell, O.; Seco, M.; Solans, X. J. Organomet. Chem. 1989, 377, 291. (100) White, J. P., III; Deng, H.; Boyd, E. P.; Gallucci, J.; Shore, S. G. Inorg. Chem. 1994, 33, 1685. (101) Chun, S.-H.; Meyers, E. A.; Liu, F.-C.; Lim, S.; Shore, S. G. J. Organomet. Chem. 1998, 563, 23. (102) Watanabe, M.; Nagasawa, A.; Sato, M.; Motoyama, I.; Takayama, T. Bull. Chem. Soc. Jpn. 1998, 71, 1071. (103) Song, H.-B.; Wang, Q.-M.; Zhang, Z.-Z.; Mak, T. C. W. J. Organomet. Chem. 2000, 605, 15. (104) Song, H.-B.; Zhang, Z.-Z.; Mak, T. C. W. New J. Chem. 2002, 26, 113. (105) Chan, W.-H.; Zhang, Z.-Z.; Mak, T. C. W.; Che, C.-M. J. Chem. Soc., Dalton Trans. 1998, 803. (106) Coco, S.; Espinet, P.; Mayor, F.; Solans, X. J. Chem. Soc., Dalton Trans. 1991, 2503. (107) Khasnis, D. V.; Le Bozec, H.; Dixneuf, P. H.; Adams, R. D. Organometallics 1986, 5, 1772. (108) Le Bozec, H.; Dixneuf, P. H.; Adams, R. D. Organometallics 1984, 3, 1919. (109) Sato, S.; Sato, M.; Akabori, S. Bull. Chem. Soc. Jpn. 1989, 62, 532. (110) Akabori, S.; Sato, S.; Kawazoe, K.; Tamura, C.; Sato, M.; Habata, Y. Chem. Lett. 1987, 1783. (111) Sato, S.; Habata, Y.; Sato, M.; Akabori, S. Bull. Chem. Soc. Jpn. 1989, 62, 3963. (112) Batchelor, R. J.; Davis, H. B.; Einstein, F. W. B.; Pomeroy, R. K. J. Am. Chem. Soc. 1990, 112, 2036. (113) Batchelor, R. J.; Einstein, F. W. B.; Pomeroy, R. K.; Shipley, J. A. Inorg. Chem. 1992, 31, 3155. (114) Einstein, F. W. B.; Pomeroy, R. K.; Rushman, P.; Willis, A. C. Organometallics 1985, 4, 250.

(115) Jiang, F.; Biradha, K.; Leong, W. K.; Pomeroy, R. K.; Zaworotko, M. J. Can. J. Chem. 1999, 77, 1327. (116) Del Paggio, A. A.; Muetterties, E. L.; Heinekey, D. M.; Day, V. W.; Day, C. S. Organometallics 1986, 5, 575. (117) Connelly, N. G.; Lucy, A. R.; Payne, J. D.; Galas, A. M. R.; Geiger, W. E. J. Chem. Soc., Dalton Trans. 1983, 1879. (118) Connelly, N. G.; Lucy, A. R.; Galas, A. M. R. J. Chem. Soc., Chem. Commun. 1981, 43. (119) Einstein, F. W. B.; Jones, R. H.; Zhang, X.; Sutton, D. Can. J. Chem. 1989, 67, 1832. (120) Luke, M. A.; Mingos, D. M. P.; Sherman, D. J.; Wardle, R. W. M. Transition Met. Chem. 1987, 12, 37. (121) Bianchini, C.; Elsevier, C. J.; Ernsting, J. M.; Peruzzini, M.; Zanobini, F. Inorg. Chem. 1995, 34, 84. (122) Xie, Z.; Jelínek, T.; Bau, R.; Reed, C. A. J. Am. Chem. Soc. 1994, 116, 1907. (123) Liston, D. J.; Reed, C. A.; Eigenbrot, C. W.; Scheidt, W. R. Inorg. Chem. 1987, 26, 2739. (124) Casalnuovo, A. L.; Laska, T.; Nilsson, P. V.; Olofson, J.; Pignolet, L. H. Inorg. Chem. 1985, 24, 233. (125) Nowell, I. W.; Russell, D. R. J. Chem. Soc., Dalton Trans. 1972, 2393. (126) Balch, A. L.; Catalano, V. J. Inorg. Chem. 1992, 31, 2730. (127) Braunschweig, H.; Radacki, K.; Schwab, K. Chem. Commun. 2010, 46, 913. (128) Arai, S.; Ochiai, M.; Ishihara, K.; Matsumoto, K. Eur. J. Inorg. Chem. 2007, 2031. (129) Moret, M.-E.; Chen, P. J. Am. Chem. Soc. 2009, 131, 5675. (130) Yamaguchi, T.; Yamazaki, F.; Ito, T. J. Am. Chem. Soc. 2001, 123, 743. (131) Heckenroth, M.; Neels, A.; Garnier, M. G.; Aebi, P.; Ehlers, A. W.; Albrecht, M. Chem.Eur. J. 2009, 15, 9375. (132) Heckenroth, M.; Kluser, E.; Neels, A.; Albrecht, M. Angew. Chem., Int. Ed. 2007, 46, 6293. (133) Usón, R.; Forniés, J.; Tomás, M.; Ara, I.; Casas, J. M.; Martín, A. J. Chem. Soc., Dalton Trans. 1991, 2253. (134) Shan, H.; James, A.; Sharp, P. R. Inorg. Chem. 1998, 37, 5727. (135) Yamamoto, Y.; Ehara, K.; Takahashi, K.; Yamazaki, H. Bull. Chem. Soc. Jpn. 1991, 64, 3376. (136) Yamaguchi, T.; Yamazaki, F.; Ito, T. J. Am. Chem. Soc. 1999, 121, 7405. (137) Forniés, J.; Ibáñez, S.; Martín, A.; Gil, B.; Lalinde, E.; Moreno, M. T. Organometallics 2004, 23, 3963. (138) Tanase, T.; Yamamoto, Y.; Puddephatt, R. J. Organometallics 1996, 15, 1502. (139) Schuh, W.; Kopacka, H.; Wurst, K.; Peringer, P. Eur. J. Inorg. Chem. 2001, 2399. (140) Ghilardi, C. A.; Midollini, S.; Moneti, S.; Orlandini, A.; Scapacci, G.; Dakternieks, D. J. Chem. Soc., Chem. Commun. 1989, 1686. (141) Falvello, L. R.; Forniés, J.; Martín, A.; Navarro, R.; Sicilia, V.; Villarroya, P. Inorg. Chem. 1997, 36, 6166. (142) Ara, I.; Falvello, L. R.; Forniés, J.; Sicilia, V.; Villarroya, P. Organometallics 2000, 19, 3091.

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