Bond TheoryDOI: 10.1002/anie.200504322

To Boldly Pass the Metal–Metal Quadruple BondUdo Radius* and Frank Breher*

Dedicated to Professor HansgeorgSchn�ckel on the occasion of his 65thbirthday

Keywords:arene ligands · bond theory · metal–metal interactions ·multiple bonds

The theory of chemical bonding isintriguing and of fundamental impor-tance. Bonding concepts were devel-oped and rules were established overthe years to describe and subdividemolecular compounds consisting of ei-ther single, double, or triple bonds. Thelatter was believed to be the highestaccessible bond order for a long time.Although compounds that containquadruple bonds between transitionmetals were already prepared in the19th century, for example [Cr2(m-O2CMe)4(H2O)2]),

[1] it was not until1964, when Cotton et al. reported onthe crystal structure of K2[Re2Cl8]·2H2Ofeaturing a surprisingly short Re�Redistance of 2.24 /,[2] that such a quad-ruple bond between two transition-met-al atoms was confirmed unequivocally.The [Re2Cl8]

2� ion has become theprototype for this type of complexes,and a new era of inorganic chemistryand a rich chemistry has evolved aroundthis class of transition-metal com-plexes.[3]

The well-known qualitative descrip-tion of a s2p4d2 quadruple bond iselegant and deceptively simple (andfamiliar), but it is based on a one-electron model and the inherent as-sumption that four bonding orbitals aredoubly occupied. Today we know thatthis is not entirely true for weak inter-metallic bonds. The small overlap be-tween d orbitals results in a relativelyweak interaction, thus making a simple

molecular orbital (MO) description ofthe quadruple bond inappropriate.[4] Amore accurate, but less transparentdescription of the quadruple bond re-quires an approach that goes beyondsingle configuration methods inherentto simple MO formalisms. To properlydescribe systems already in the groundstate, treatments must include correla-tion effects. A bond order analysis of[Re2Cl8]

2� relying on CASSCF (com-plete active-space self-consistent field)calculations,[5] for example, has shown,that the Re�Re bond has an effective(calculated) bond order of 3.2 and thenet bond order contribution of the d

bond is about 0.5. The reason is mainly apartial occupation of the antibonding d*orbitals. Whether it should be called atriple or a quadruple bond or whetherthe d bond should be called a weak bondor half a bond is more or less a matter ofdefinition of the term “bond” or “bondorder”. The alternatives here are todescribe the bonding as a “weak” quad-ruple bond (four orbitals with relevantoverlap) or as a bond involving fourelectron pairs with an effective bondorder of about three.

Various orbitals contribute to a dif-ferent extent to metal–metal bonding, asdemonstrated by intriguing compoundssuch as the silicon analogue of analkyne, RSi�SiR, reported by Sekiguchiet al. or Wiberg et al.[6] and the corre-sponding germanium, tin, and lead com-pounds REER (E=Ge–Pb) comprisingbulky aryl ligands (Ar’ and Ar*)[7]

described not long before by Poweret al.[8] For the REER compounds, theterm alkyne analogue does not implythat each of the three valences availablefor the Group 14 element contributeequally to chemical bonding to retain atriple bond featuring an integer bondorder of three. As a result of the

gradually increasing, nonlinear, trans-bent geometries on descending thegroup, a considerable weakening ofone component of the degenerate p

bonding was suspected. Although thestructural data, as well as quantum-chemical calculations point towardsbond orders approaching three for Si,approximately two for Ge and Sn, andone for Pb, the correlation of bondlength and bond order is questionableand should be treated with care, espe-cially when counterions are involved.This has led to controversial discussionsin the past, for example, whether [Ar*-GaGaAr*]2� (formally isoelectronic toneutral RGeGeR) is the first experi-mentally proven gallium–gallium triplebond or not.[9] Although not capable ofdistinguishing between covalent andelectrostatic contributions, the best ex-perimental tool for a classification ofsuch bonds might be the force constantof the metal–metal bond.[10]

A milestone for multiple-bondchemistry involving transition metalswas recently reported by Power andco-workers, who impressively succeededin isolating “a stable compound withfivefold bonding between two chromi-um(i) centers”.[11] To achieve the highestbond order possible in an isolable com-pound, the number of ligands has to beminimized, since their binding reducesthe number of valence orbitals availableto form metal–metal bonds. Further-more, the steric requirements for theligand system is of crucial importancesince large ligands prevent oligomeriza-tion and undesirable bridging motifs areusually disfavored for steric reasons.Relying on their background in maingroup chemistry,[8] Power et al. havesynthesized and characterized a dinu-clear metal–metal-bonded complex withone ligand per metal atom in the ligand

[*] Priv.-Doz. Dr. U. Radius, Prof. Dr. F. BreherInstitut f.r Anorganische ChemieUniversit1t Karlsruhe (TH)Engesserstrasse 1576131 Karlsruhe (Germany)Fax: (+49)721-608-8440E-mail: radius@aoc1.uni-karlsruhe.de

breher@aoc1.uni-karlsruhe.de

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sphere. The reduction of [{Cr(m-Cl)Ar’}2]

[7] with a slight excess of potas-sium graphite afforded the thermallyrobust complex [Ar’CrCrAr’] (1) in41% yield.

Following the simplified picture de-veloped for [Re2Cl8]

2�, the bonding in a(hypothetical) linear complex[RCrCrR] (R=monoanionic ligand)can be described as a quintuple (ten-electron–two-center) Cr�Cr bondformed by a fivefold overlap betweenmetal d orbitals (Figure 1). Five electron

pairs play a dominant role in holding themetal atoms together, but it does notnecessarily imply that the bond order isfive or that the bonding is very strong.As pointed out for [Re2Cl8]

2�, theground state of the molecule possiblyinvolves mixing of other higher energyconfigurations with less bonding char-acter.

The molecular structure of [Ar’Cr-CrAr’] (1) (Figure 2) reveals a planar

core geometry with a Cr�Cr bond lengthof 1.8351(4) /. The chromium–chromi-um distance is slightly longer than thebond length reported for a quadruplybonded, but ligand-bridged CrII dimer[Cr2(m

2-OMe-5-MeC6H3)4] (1.828(2) /),which has the shortest metal–metalbond distance observed so far.[12] Theatoms C1, Cr1, Cr1A, and C1A of thecentral Cr2C2 unit in 1 are aligned in aplane, but deviate significantly fromlinearity (angle Cr1A-Cr1-C1102.78(3)8), adopting a trans-bent struc-ture. Each chromium atom is bonded tothe ipso-carbon atom (Cr1�C12.131(1) /) of the Ar’ ligand andthrough a somewhat weaker interaction(Cr1�C7A 2.2943(9) /) to an ipso-car-bon atom of a flanking dipp ring (dipp=

C6H3-2,6-iPr2) of the ligand. The inter-action of the dipp moiety with thechromium atom can also be consideredas that with a distorted h6-coordinatedarene ligand; the chromium–carbon dis-tances range from 2.29 to 2.97 /.

The simple bonding model givenabove, therefore, has to be adjusted tothe geometry observed. Additionally,mixing of orbitals can occur due to thelower symmetry of the complex. DFTsingle-point calculations on the experi-mentally verified structure presented bythe authors support the view that thereare five orbital interactions (one s, twop, and two d) between the CrI ions(Figure 3). HOMO and HOMO�1,which differ in energy by 0.41 eV, cor-respond to d bonds, and the LUMO,which is 2.01 eV higher in energy thanthe HOMO, corresponds to a d* orbital.The chromium–chromium s-binding or-bital (HOMO�2) that emerges from thesingle-point DFT calculations on 1 iscomparatively high in energy, above thetwo p-binding orbitals (HOMO�3 andHOMO�4). This might be due to sig-nificant orbital mixing and/or due toreduced overlap along the chromium–chromium axis. Although the authorsprovided first results on CASSCF cal-culations in the Supplementary Materialof the original article, the meaning ofthis feature for metal–metal bondingremains to be evaluated in a morethorough theoretical treatment of 1 ormodels of this compound.

Before entering into a further dis-cussion of the bonding in complex 1, it ismost instructive to briefly outline the

bonding situation in the dichromiumdimer, Cr2, which has attracted consid-erable interest in recent years.[13] Bycontinuing with the concept of minimiz-ing the number of metal ligands tomaximize the number of free metalvalence orbitals available to form met-al–metal bonds to an extreme, evenlarger bond orders than five should befeasible. The bare dimers of the open d-shell transition metals provide an op-portunity to examine multiple bondsbetween metal atoms in the absence ofligand effects. Among the dimers of thefirst transition series, Cr2 potentiallyprovides one of the most intriguingexamples of multiple metal–metal bond-ing. Since the chromium atom has ahigh-spin 7S ground state (a (3d)5(4s)1

valence electron configuration with oneelectron in each of six valence orbitals),closely spaced energy levels result. Thespin pairing of two ground-state atomsresults in a 1Sg

+ Cr2 molecule with avalence electron configuration (4ss)2-(3ds)2(3dp)4(3dd)4 comprising two s,two p, and two d bonds giving formalchromium–chromium bond order of six.The complexity of the bonding in Cr2,however, was for a long time a challengefor ab initio quantum chemistry becauseof this “hextuple” bond and the unusualshape of the Cr2 potential energy curve(Figure 4).

Figure 1. Schematic MO picture for linear [R-CrCr-R].

Figure 2. Molecular structure and numberingscheme of [Ar’CrCrAr’] (1).

Figure 3. Electron density surfaces and energiesfor the Cr–Cr frontier orbitals in [Ar’CrCrAr’] (1)[11a]

(reprinted with permission from Science 2005,310, 844. Copyright 2005 AAAS).

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At a short internuclear Cr�Cr dis-tance (re= 1.6788 /), the molecule ismultiply bonded and exhibits a relative-ly deep potential. As the atoms separate,the rapid loss of bonding between thecompact 3d orbitals is partially compen-sated by increased interatomic exchangestabilization, as well as by 4s–4s bond-ing. As a result, the potential curve risesmore slowly with increased internuclearseparation than would be expected,which renders the curve highly anhar-monic with a shelf or a shallow mini-mum at intermediate chromium–chro-mium separations. The nature of the Cr�Cr interaction changes qualitativelywith increased internuclear separation,from multiple 3d–3d bonding at shortdistances to single 4s–4s bonding (withthe 3d electron cores of the two atomsantiferromagnetically coupled to give asinglet state) at long distances. The Cr2

potential provides a beautiful illustra-tion of this phenomenon. Using highlycorrelated CASSCF and CASPT2 cal-culations, Anderson and Roos et al.[13d–h]

gave a qualitatively correct descriptionof the Cr2 potential energy surface andcalculated a total chromium–chromiumbond order of 4.4 at equilibrium dis-tance. Accordingly, the chromium–chro-mium bond consists of one s bond ofprimarily 4s and 3d contributions, two p

bonds totally 3d in character, two d

bonds totally 3d in character, and oneantiferromagnetically coupled electronpair. This assignment avoids the rathercounterintuitive description of two chro-mium–chromium s bonds. The distinc-

tion between antiferromagnetic cou-pling and bonding, however, is in thiscase not clearly defined.

As with Cr2, the chromium–chromi-um bond in the complex 1 presented byPower et al. might alternatively be de-scribed as a quadruple bond with twoantiferromagnetically coupled electronsresiding in chromium-localized orbitals.Magnetic measurements on [Ar’Cr-CrAr’] revealed a temperature-inde-pendent weak paramagnetism of0.000112(5) emu per mol Cr, which isconsistent with strongly coupled pairedelectrons (S= 0) and an first excitedstate (S= 1) relatively high in energy,without a significant population of S> 0states at room temperature. The mostcommon metrics used to gauge thequality of calculations of the electronicstructure of quadruple bonds are thegeometry-optimized metal–metal bondlength and the d–d* excitation energy.The electronic absorption spectrum of[Ar’CrCrAr’] displays a broad absorp-tion at 488 nm, which lies in the rangeobserved for d–d* transitions of com-pounds with metal–metal quadruplebonds.

Without doubt, the analysis of thebonding situation in [Ar’CrCrAr’] aswell as those for the iron and cobaltanalogues, briefly mentioned by Powerand co-workers, will be an interestingtopic among theoreticians. It is veryinteresting to note that the structurallyrelated [Ar’FeFeAr’] (+ 4 electrons) and[Ar’CoCoAr’] (+ 6 electrons) dimershave much longer metal–metal contacts(2.53 / (Fe�Fe) and 2.80 / (Co�Co)).Assuming the classical bonding picture,this would imply an increase of the bondlength of approximately 0.7 / due tooccupation of two antibonding d* orbi-tals. A bonding distance as long as2.53 / would result for a formal iron–iron triple bond!

Going one step further to f blockelements, the uranium atom holds, sim-ilarly to a chromium atom, six electronsin its valence shell. However, whereaschromium has exactly six valence orbi-tals, there are 16 such orbitals close inenergy available for uranium; that is theseven 5f, five 6d, one 7s, and three 7porbitals. Gagliardi and Roos recentlypresented calculations on hypotheticaldiuranium U2 using CASSCF calcula-tions.[14a] As the authors stated, the

maximum flexibility for describing elec-tronic structures and the capability ofhandling arbitrary spin as offered by theCASSCF method is important, because“we cannot assume anything concerningthe final number of paired electrons inU2”. Indeed, when considering heavyactinide elements, metal–metal bondsproved to be particularly complicated.Although the ground-state electronicconfiguration of an uranium atom is(5f)3(6d)1(7s)2 (quintet ground state),the energy cost of unpairing the 7selectrons is low, revealing in principlesix electrons available for each U atomto form chemical bonds. The bondingsituation for uranium, however, is con-siderably more complex than the situa-tion found for Cr2—in fact unique, asreported by Gagliardi and Roos.[14a]

Their calculations revealed thatthree (two-electron–two-center) elec-tron pair bonds are formed by hybridorbitals with predominantly 7s and 6dcharacter, that is one s-type and adegenerate set of p-type orbitals (Fig-ure 5, first row). Furthermore, two sin-gly occupied orbitals of s-type (6dsg)and of d-type (6ddg), which show mainlyd-orbital character, give rise to two one-electron–two-center bonds between thetwo uranium atoms. Two singly occupiedorbitals of d- and p-type (5fdg/5fpu and5fdu/5fpg, respectively), with predomi-nantly f character, form two additionalone-electron bonds. Finally, two elec-trons occupy a localized 5f orbital (5ffu

and 5ffg) equally distributed over bothuranium atoms with the electron spinsaligned parallel. To summarize, thebonding in U2 arises from three two-electron–two-center electron pair bonds(s + 2 p), four one-electron–two-centerbonds (s + p + 2 d), and two localizedelectrons. For U2, all single spins arepredicted to be parallel (“ferromagneti-cally” coupled) and the S= 3 septet stateis the ground state of the molecule.[15]

The calculated (CASPT2-SO, includingspin-orbit coupling, SO) equilibriumbond length of (2.43� 0.05) / and aharmonic vibrational frequency of265 cm�1 suggests comparability of thestrength of the U2 bond to that of othermultiple bonds between transition met-als. Overall, the unprecedented groundstate of U2 is expressed ass2p4s1d1d1p1f1f1 (with f1= localized 5forbital, S= 3, L= 11, ground state

Figure 4. Experimental and calculated poten-tial energy curves for Cr2 using CASSCF andCASPT2 according to Roos[13d] (reproducedfrom reference [13d] with permission of theInstitute of Organic Chemistry and Biochem-istry, Academy of Sciences of the CzechRepublic).

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71114), and is more complex than anyother known diatomic bond.

If two of the twelve electrons of U2

were removed, some simplifications ofthe electronic structure are predic-ted.[14b] Quantum-chemical calculations,based on multiconfigurational wavefunctions and including relativistic ef-fects, show that the U2

2+ system has alarge number of low-lying electronicstates with S= 0–2 and L ranging fromzero to ten. A bond length of approx-imately 2.30 / (cf. 2.43 / for neutralU2) is common for these states. Thelowest electronic state corresponds to anelectron configuration (sg)

2(pu)4 and

suggests a triple bond. The s orbital isa hybrid comprising 7s, 6ds, and 5fsatomic orbitals, and the pu orbitals aremainly 6d in character. The next fourelectrons, two localized on each of theuranium atoms, occupy 5fd and 5fforbitals, which are essentially nonbond-ing.

Recently, a related theoretical paperpresented model calculations on linearsinglet [HTh�ThH], which should be alikely candidate for the so far unknownmultiple, in this particular case triple,bond between f elements.[14c] The orbitalpicture and the bonding analyses sug-gest substantial f character in the Th�Th

bond in linear 1Sg-[HTh�ThH] bondingorbitals. The molecular orbitals of thiscompound correspond to s(Th�H)bonds, one s(Th�Th) bond(HOMO�1), and a double p(Th�Th)bond (HOMO) featuring up to 22% fcharacter in the bond. According to thecalculations the f-orbital participationstabilizes the linear geometry of themolecule.

The model of the two-electron–two-center bond, as introduced by G. N.Lewis in 1916, which features a singlebond to be formed by one pair ofelectrons, is one of the most importantconcepts in chemistry. This also coversmultiple bonds since they are regardedas composed of two, three or four two-electron components. The last few yearshave witnessed an in depth discussion onmultiple bonding between main groupelements, which has revealed that sim-ple concepts mainly emerging from thepeculiarities of carbon (or the secondperiod of elements) chemistry are notvalid for heavier elements. Apparently,the models, which are consistent andclear for the lighter main group ele-ments, become considerably more com-plicated than anticipated when appliedto their heavier counterparts. The sameholds true for the term “bond order”,

which is in fact more or less a matter ofdefinition—not trivial if reasonable atall. The decrease of orbital overlap andincrease of nonbonding lone-pair char-acter for molecules of multiply bondedmain group elements, and hence thedeviation from planarity (for R2EER2)or linearity (REER) has shown that theconcepts developed for elements of thesecond period are certainly not appro-priate to describe the nature of anelement–element “multiple” bond. Theinvolvement of d orbitals certainly com-plicates the situation due to moderateoverlap for d-type d orbitals. The workof Power et al., in unraveling the quin-tuple bond in trans-bent [Ar’CrCrAr’](1), has added a further dimension to theconcept of multiple bonding in transi-tion-metal chemistry. It is likely, how-ever, that the bonding situation in[Ar’CrCrAr’] is not fully understoodand that the trans-bending of this com-plex will add an additional level ofcomplexity to the analysis of bonding.The current state of the bond descrip-tion for 1 is that five pairs of overlappingorbitals are more or less involved inmetal–metal bonding; a situation thatchemists usually would describe as aquintuple bond.[16] A detailed theoret-ical analysis of the nature of the chem-ical bond in 1 will certainly be under-taken in the future and it is likely thatthere will be a renewed debate aboutmetal–metal multiple bonding. To fullyunderstand the bonding in this complexis of crucial significance and manyfundamental questions remain to beanswered. Whatever the conclusionturns out to be regarding the bondmultiplicity in the newly synthesizedcompound [Ar’CrCrAr’], this work un-doubtedly inspires scientists from atheoretical and experimental point ofview.

The greatest achievement here is thepreparative work, which has opened anew area previously considered non-existent. This work will certainly en-courage others to investigate transition-metal complexes bearing “ultralarge”ligands in more detail. The combinationof both the isolation of further com-pounds of the type [RMMR] on the onehand, and their precise characterizationon the other, may facilitate an accurateexperimental assessment of the bondingsituation and the bond strength in these

Figure 5. The molecular orbitals forming the chemical bond between two uranium atoms in U2.Orbital labels are given below each orbital, together with the number of electrons occupying thisorbital or pair of orbitals in the case of degeneracy[14a] (reprinted by permission from MacmillanPublishers Ltd: Nature 2005, 433, 848, copyright 2005).

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types of complexes. The main objectivefor experimental as well as theoreticalchemists is undoubtedly to work togeth-er to develop an understanding of thisunusual bonding situation between d-block metals—and similar compoundscomprising f-block metals are not out ofreach!

[1] E. Peligot, C. R. Hebd. Seances Acad.Sci. 1844, 19, 609.

[2] a) F. A. Cotton, N. F. Curtis, C. B. Har-ris, B. F. G. Johnson, S. J. Lippard, J. T.Mague, W. R. Robinson, J. S. Wood,Science 1964, 145, 1305; b) F. A. Cotton,N. F. Curtis, B. F. G. Johnson, W. R.Robinson, Inorg. Chem. 1965, 4, 326;c) F. A. Cotton, C. B. Harris, Inorg.Chem. 1965, 4, 330; d) F. A. Cotton,Inorg. Chem. 1965, 4, 334.

[3] F. A. Cotton, L. A. Murillo, R. A. Wal-ton, Multiple Bonds Between MetalAtoms, 3rd ed., Springer, Berlin, 2005.

[4] F. A. Cotton, D. G. Nocera, Acc. Chem.Res. 2000, 33, 483.

[5] a) L. Gagliardi, B. O. Roos, Inorg.Chem. 2003, 42, 1599; b) F. Ferrante, L.Gagliardi, B. E. Bursten, A. P. Sattlberg-er, Inorg. Chem. 2005, 44, 8476.

[6] a) A. Sekiguchi, R. Kinjo, M. Ichinohe,Science 2004, 305, 1755; b) N. Wiberg,W. Niedermayer, G. Fischer, H. NPth,M. Suter, Eur. J. Inorg. Chem. 2002,1066; c) N. Wiberg, S. K. Vasisht, G.Fischer, P. Mayer, Z. Anorg. Allg. Chem.2004, 630, 1823; d) M. Weidenbruch,Angew. Chem. 2005, 117, 518; Angew.Chem. Int. Ed. 2005, 44, 514 (Highlight);e) M. Weidenbruch in The Chemistry ofOrganic Silicon Compounds, Vol. 3(Ed.: Z. Rappoport, Y. Apeloig), Wiley,Chichester, 2001; f) N. Takagi, S. Na-gase, Eur. J. Inorg. Chem. 2002, 2775;Recently, Passmore and co-workers im-pressively succeeded in isolating theS2I4

2+ cation featuring a sulfur sulfur

bond with a high bond order, which iscomparable to that of RSiSiR: g) S.Brownridge, T. S. Cameron, H. Du, C.Knapp, R. KPppe, J. Passmore, J. M.Rautiainen, H. SchnPckel, Inorg. Chem.2005, 44, 1660; h) S. K. Ritter, Chem.Eng. News 2005, 83, 49.

[7] Ar’: -C6H3-2,6-dipp2 (dipp=C6H3-2,6-iPr2); Ar*: -C6H3-2,6-trip2 (trip=C6H2-2,4,6-iPr3).

[8] a) M. Stender, A. D. Phillips, R. J.Wright, P. P. Power, Angew. Chem.2002, 114, 1863; Angew. Chem. Int. Ed.2002, 41, 1785; b) A. D. Phillips, R. J.Wright, M. M. Olmstead, P. P. Power, J.Am. Chem. Soc. 2002, 124, 5930; c) L.Pu, B. Twamley, P. P. Power, J. Am.Chem. Soc. 2000, 122, 3524; d) P. P.Power, Appl. Organomet. Chem. 2005,19, 488; e) C. Cui, M. M. Olmstead, J. C.Fettinger, G. H. Spikes, P. P. Power, J.Am. Chem. Soc. 2005, 127, 17530; forselected reviews see: f) P. P. Power,Chem. Rev. 1999, 99, 3463; g) M. Wei-denbruch, Organometallics 2003, 22,4348; h) P. P. Power, Chem. Commun.2003, 2091.

[9] a) Y. Xie, R. S. Grev, J. Gu, H. F. Schae-fer, P. v. R. Schleyer, J. Su, X.-W. Li,G. H. Robinson, J. Am. Chem. Soc. 1998,120, 3773; b) G. H. Robinson, Chem.Commun. 2002, 2175; c) F. A. Cotton,A. H. Cowley, X. Feng, J. Am. Chem.Soc. 1998, 120, 1795; d) M. M. Olm-stead, R. S. Simons, P. P. Power, J. Am.Chem. Soc. 1997, 119, 11705; e) T. L.Allen, W. H. Fink, P. P. Power, J. Chem.Soc. Dalton Trans. 2000, 407; f) N. J.Hardman, R. J. Wright, A. D. Phillips,P. P. Power, J. Am. Chem. Soc. 2003, 125,2667; g) R. Ponec, G. Yuzhakov, X.GironUs, G. Frenking, Organometallics2004, 23, 1790; h) J. Grunenberg, N.Goldberg, J. Am. Chem. Soc. 2000, 122,6045; i) N. Takagi, M. W. Schmidt, S.Nagase, Organometallics 2001, 20, 1646.

[10] See for example: a) R. KPppe, H.SchnPckel, Z. Anorg. Allg. Chem. 2000,626, 1095; b) see also ref. [9h].

[11] a) T. Nguyen, A. D. Sutton, M. Brynda,J. C. Fettinger, G. J. Long, P. P. Power,Science 2005, 310, 844; b) G. Frenking,Science 2005, 310, 796.

[12] F. A. Cotton, S. A. Koch, M. Millar,Inorg. Chem. 1978, 17, 2084.

[13] For the potential energy curve of Cr2

see: a) S. M. Casey, D. G. Leopold, J.Phys. Chem. 1993, 97, 816; for theoret-ical calculations on Cr2 see for example:b) N. E. Schultz, Y. Zhao, D. G. Truhlar,J. Phys. Chem. A 2005, 109, 4388;c) E. A. Baudreaux, E. Baxter, Int. J.Quantum Chem. 2004, 100, 1170;d) B. O. Roos, Collect. Czech. Chem.Commun. 2003, 99, 265; e) G. L. Gutsev,C. W. Bauschlicher, Jr., J. Phys. Chem. A2003, 107, 4755; f) B. O. Roos, K. An-derson,Chem. Phys. Lett. 1995, 245, 215;g) K. Anderson, Chem. Phys. Lett. 1995,237, 212; h) K. Anderson, B. O. Roos, P.-/. Malmqvist, P.-O. Widmark, Chem.Phys. Lett. 1994, 230, 391; i) M. M.Goodgame, W. A. Goddard III, Phys.Rev. Lett. 1985, 54, 661.

[14] a) L. Gagliardi, B. Roos, Nature 2005,433, 848; b) L. Gagliardi, P. PyykkP,B. O. Roos, Phys. Chem. Chem. Phys.2005, 7, 2415; c) M. Stratka, P. PyykkP, J.Am. Chem. Soc. 2005, 127, 13090.

[15] Note the difference to Cr2 featuring“antiferromagnetic” coupling, whichcauses spin-pairing of the electrons andusually provides some additional smallbonding contributions and some extrastabilization. In the particular case ofU2, this ferromagnetic coupling can beattributed to favorable exchange stabi-lization, that is interaction between non-bonding 5f electrons and one-electronbonds.

[16] Bonding pictures might also be totallyunexpected, as experienced in the caseof U2; see also: M.-M. Rohmer, M.BUnard, Chem. Soc. Rev. 2001, 30, 340

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