摘要:

Some Properties of Metal Complexes Containing One Metal-Carbon BondBY H. A. 0. HILL, J. M. PRATT AND R. J. P. WILLIAMSInorganic Chemistry Laboratory, South Parks Road, OxfordReceived 16th January, 1969The physical and chemical properties, including electronic absorption, infra-red, and 'H n.m.r.spectra of cobalt(1II) complexes of the type CoA4XY, where .A4 can be corrin, a Schiffs base, ordimethylglyoxime and X a carbon ligand, are discussed in terms of the change in electron densityin the complex. It is shown that alkyl ligands act as strong donors, often giving rise to five-co-ordinate complexes. The relationship of the properties of these five-co-ordinate cobalt(III) complexesto those of analogous cobalt(lI) nickel(I1) and palladium(II) complexes is discussed and the differencebetween the cobalt-carbon bond in the five- and six-co-ordinate complexes considered.We are interested in the properties and influence of the metal-carbon bond incomplexes of the type MA4XY, where A4 represents a ligand or ligands co-ordinatedto the central metal through atoms other than carbon and arranged in a plane asshown in fig.1. The axial ligands X and Y can be varied in each complex; X isco-ordinated through carbon, e.g., CN, CH3, C6H5, whereas Y can be co-ordinatedthrough N, 0, P or C .FIG. 1.-Structure of complexes of thetype M&XY.The metal-carbon bond can be affected by the trans-ligand Y or the cis-ligand(s)A, and conversely the properties of Y and A4 can be altered by a change in the carbonligand X. We shall use various physical properties of A,, X and Y to uncover thenature of the co-operative interaction between the groups.We have shown that in monocyano derivatives of cobdamins and cobinamideswhere M is formally Co(III), A4 is the corrin ljgand and X = CN, the CN stretchingfrequency depend on the ligand Y as shown in table 1.We note that when Y = C2H5,the stretching frequency is close to that of free cyanide. From these results wemight conclude that as Y becomes a " better donor " to the cobalt, the trans-metal-carbon bond becomes weaker until perhaps the six-co-ordinate complex does notform at all. Similarly in other cobalt(1I.I) complexes, the physical properties of onetrans-ligand approach those of the " free "-figand as the donor character of the other16166 COMPLEXES WITH ONE METAL-CARBON BONDtrans-ligand increases. For example, in a series of cobalt(III) dimethylglyoximates,2Co(DH), pyridine X, the ‘H n.m.r.of the co-ordinated pyridine is close to that offree pyridine when X=CH3. In a series of c~balamins,~ the ‘H n.m.r. of the co-ordinated 5,6-dimethylbenziminazole (Bz) is at higher field when X is a strong donor.Does the bond length to the group Y also depend on X? In the only two examplesavailable, this is indeed the case ; ~yanocobalamin,~ Co-Bz 2.07 A, Co-corrin(average) 1 -905 ; 5’-deo~yadenosylcobalamin,~ Co-Bz 2-23 A, Co-corrin(average) 1 -94 A.TABLE DE DEPENDENCE OF THE CN STRETCHING FREQUENCY IN CYANOCOBALAMINS ANDCYANOCOBINAMIDES ON THE AXIAL LIGAND TRANS TO CN,igand axial 5,6-dimethyl- benziminazo,e OH- CN- HCrC- CHz=CH- ~~~~~~~i CHZ CH3CH?’_ free CN-ICNstretchingfrequency(cm-l) 2132 2130 2119 2110 2093 2091 2088 2082 2079Taking this process further it has been possible to prepare complexes with thefollowing A, ligands which have a variety of carbon ligands X but no sixth ligand;bis(acety1acetone)ethylenediimjne (BAE), bis(salicy1aldehyde)ethylenediimine(SALEN), bis(trifluoromethylacety1acetone)ethylenediimine * (BTFAE), anddimethylgly~xime.~ We have also concluded O from the temperature-dependenceof the absorption spectra and lH n.m.r.of cobinamides that there are three classesof cobinamides: (a) Those with relatively “weak” ligands X and Y, e.g., H20,NH3, whose ‘H n.m.r. and electronic absorption spectra are insensitive to temperature.(b) Those with one very strong ligand X, e.g., i-C3H7 whose lH n.m.r.and electronicabsorption spectra are virtually temperature independent. (c) Those with ligandsof intermediate strength, e.g., -CH=CH2. With such ligands there is a markeddependence of the lH n.m.r. and electronic absorption spectra with temperaturesuggesting the presence of two complexes in equilibrium.For various reasons which have been given in detail elsewhere,lO we considerthe complexes in class (c) to be an equilibrium mixture of six-co-ordinate cobaltcomplex as in class (a) and a complex in which the co-ordination of the cobalt ap-proaches more closely five-co-ordination as in class (b). We observe that as thedonor power of X is increased there is a discontinuity in the properties of the cobaltcomplexes as shown schematically in fig.2.The lH n.m.r. of the carbon ligand X, e.g., the Co-CH3 is also sensitive to bothY and A4 as shown in table 2. These effects are more difficult to interpret in termsof the M-C bond due to the influence of the metal and the A, ligands directly,i.e., through-space rather than along-the-bonds, on the ‘H-chemical shift. However,we see that in the limited series available with BAE as the A,-ligand, the chemicalshift of the Co--CH, increases with increasing basicity of the nitrogen ligand Y .All these data on the influence of the carbon ligand X on the properties of the A4and Y ligands and the effect of A, and Y in X-ligands can be interpreted in terms of achange in electron density on the cobalt.In a general sense the cobalt atom acts asa relay of electron density from one part of the molecules, X, Y or A4, to any other.If change in electron density is the principal factor determining the properties of theligands, we might expect to observe close relationships between the properties ofA4, Y and the carbon ligands. Several examples have been observed. Thus, thH . A . 0. HILL, J . M. PRATT AND R . J . P . WILLIAMS 167CN stretching frequency correlates 1 * lo with the IH chemical shift of the C-10hydrogen in the corrin ligand (fig. 3), which in turn correlates lo with the energy of the0-1 vibrational component of the first electronic transition (fig. 4). The latter cannotbe a general relationship but is expected here as the first excited state l4 of corrin hasa node at C-10 and therefore the energy of the transition reflects changes in electrondensity at C-10 in the ground state.a correlation betweenthe CN stretching frequency and the equilibrium constant for replacement of oneaxial ligand by another, e.g., H20 by CN or benziminazole by CN in the vitaminB12 series, fig. 5. When A4 is BAE, and Y is a para-substituted phenyl derivativethere is a correlation between the lH n.m.r. of the methine hydrogens in the planarSimilarly, there is-.IBRIUM. -/ 6-COORD IN ATEdonor strength of ligandFJG. 2.Cchematic representation of the changes in physical properties, e.g., chemical shift, energyof electronic transitions, with the change in the donor strength of the ligand.At low donor strengththe complexes are six-coordinate ; at high donor strength five-coordinate. In the intermediateregions the two types of complex are in equilibrium and can give rise to distinct physical properties,absorption spectra, or to a single timeaveraged property, n.ni.r. In the figure the continuous linesshow the region of donor strength in which one or other of the two species predominates.TABLE 2.-INFLUENCE OF AXIAL AND PLANAR LIGANDS ON THE CHEMICAL SHIFT OF THE COBALT-METHYL HYDROGENSplanar ligandcorrinwrrindimethyfglyoximeBAE l1SALENBTFAEtetrasulphonatopht halocyanineaetioporphyrin laxial ligand Ybenziminazolewatertriphenylphosphinep yridinepiperidine4methylpyridinepyridine4-cyanop yridinedimethylsulphoxidepiperidine4-methylpiperidinepyridine4-cyanopyridinewater?--C 4 H 3 t10.1410.248.829.187.757.667-467.437.417.887-257-166.976-966.9416.115.1168550COMPLEXES WITH ONE METAL-CARBON BONDDicyanocobinumide ysc / y Ethylcobalamin"in $ CObdam;n/x/w Methylcobahmin-2 1302 120n e k 2110W2; 9210020902080 0 3.95 4.00 4-a 4.10 4.15chemical shift C-10 H (7)FIG.3.--Chemical shift 7 of the C-10 hydrogen in cobdamins against the stretching frequency (cm-I)of the cyanide ion in cyanocobalamins and cobinamides.ligand and the equilibrium constant for formation of the six-co-ordinate complexwhen Y is pyridine (table 3). Here both properties are also related to the electron-donor character of the para-substituent.A similar correlation has been observedin the cobalt(I1I) dimethylglyoximates in which both alkyl and cyanide ligandsaffect the chemical shifts of the in-plane hydrogens in proportion to their Hammettcr para-substituent constants as do other non-carbon ligandsH . A . 0. HILL, J . M. PRATT A N D R . J . P . WILLIAMS 169From a consideration of all these properties in several series of complexes, weconclude that alkyl ligands have and cause properties which are the natural extensionof those of weaker donor ligands such as NH3. The alkyl ligand is just a very strongdonor. However, in certain cases the donation may be so marked as to prevent theformation of a six co-ordinate species and simultaneously this donation causes abreak in the gradation of the properties of the planar ligand (fig.2). We then askI1 1 I 1 I2080 2 roo 2120 2140cm-lFIG. 5-Gxrelation between the stretching frequency of co-ordinated CN- (in cyanocobalamins)and the formation constants for the substitution by CN- of H20 in cobinamides and of Bz in cobal-amins (reproduced by permission of the Chemical Society).TABLE 3.-cORRELATION BETWEEN THE CHEMICAL SHIFT OF THE METHENE HYDROGENS IN~-SUBSTTTUTED COBALT(ZII) BAE AND THE FORMATION CONSTANT FOR FORMATION OF THESIX-CO-ORDINATED PYRIDINE COMPLEX.para-substituent r-methine Kpy (1. mole-1)OCHJ 4.59 6.1 AO-8CH3 4.59 5.0 f0.6H 4-58 5.3 f0-7I 4.58 19 f2.5Br 4-57 20 f2.5CN 4.52 54 f7NO2 4.49 85f13the critical question : does the nature of the M-C bond also differ markedly whenthe change from six- to five-co-ordination occurs? We can see a possible answerto this question by a comparison of the cobalt(III) Complexes with those of loweroxidation states.Cobalt(II1) complexes are normally considered to adopt fairlyregular geometries and we would expect that, e.g., when X and Y = H20 or Cl-,the cobalt would lie in the plane of the A4 ligand. However, as the ligand X donatesmore charge to the cobalt so the formally Co(1II) metal ion has an electron densit170 COMPLEXES WITH ONE METAL-CARBON BONDcloser to the real charge of a lower oxidation state. Thus, we might expect that thecobalt ion should have properties more like those of a low-spin Co(U)d7, or evenlow-spin Co(I)d8.[This ambiguity when X is a covalent ligand is reflected in therelative weight of the resonance forms : CH; Co(III)++CH,. Co (II)-CH; Co(1).The relative weights could be very different in the five- and six-co-ordinate forms.]The absorption spectra of the five-co-ordinate forms of the cobalt(II1) corrins,cobalt(II), nickel(I1) and palladium(I1) corrins, have two main absorptionbands -450-470 mp and 308-320 mp, whereas in the six-co-ordinate forms thefirst two intense bands are 500-600 and 350-370mp. This similarity is alsoobserved l6 in the circular dichroism of the cobalt(II1) and cobalt(I1) complexes.The similar values of the equilibrium constants for substitution of the benziminazoleby H 2 0 on acidification of the cobalt(I1) complexes (pK,-2-5) and of the methyl-cobalt complex (pKa = 2.5) suggest that the electron density on the metal ion isapproximately the same in both cases.In other words, the effect of the carbon ligand bound to the cobalt(II1) has beentwo-fold: it has induced a lowering of the cobalt co-ordination number to thattypical of a lower oxidation state and it has also caused the ligand A4 to have thesame electronic characteristics and perhaps even structure, as judged by the absorp-tion spectra and the circular dichroism, as found in the lower oxidation state com-plexes.Clearly, this could markedly alter the polarity of the cobalt-carbon bond.We are studying the reactivity of these complexes to see if there is such aneffect.We now turn to the exact character of the change in geometry on going from thecobalt(II1) to the cobalt(I1) state.The marked reluctance of some cobalt(I1) com-plexes to co-ordinate two axial ligands, e.g., pyridine suggests that the cobalt(I1)ion in the monopyridine derivative does not lie in the plane of the A4 ligand(s).5-COORDINATEP I G n U r2d I s t o r ted-16-COORDI NATE-----distorted4twZWE!planar3FIG. 6.--Schematic representation of the energy of the four possible forms of five- and six-co-ordinatecomplexes.By inference, the cobalt atom in the five-co-ordinate alkyl cobalt(II1) complexes maybe similarly displaced. It is interesting that in an attempt l7 to calculate the effectof the ring current in phenylcobalt(II1) BAE on the chemical shifts of hydrogensforming an A2B2 system in the plane of the BAE ligand, results consistent with thoseobserved could be obtained only if the molecule was considerably distorted fromplanarity.The physical and chemical properties of these different series of cobalH . A . 0. HILL, J . M. PRATT AND R . 3 . P . WILLIAMS 171complexes of a given co-ordination number suggest that we should now considertwo possible structures for both five- and six-co-ordination, planar and distorted.The degree of distortion in A4 will depend in either case on the A4 and X ligands.Presumably, in the corrin and BAE complexes discussed above, the equilibriumbetween five- and six-co-ordinate species is that represented by 1 +3 in fig.6. Thoseligands which have a stronger tendency to remain planar, e.g., phthalocyanine anddimethylglyoxime, may be more reluctant to form five-co-ordinate species presumablybecause the distorted form is of higher energy, and so in such a case the equilibriummay be better represented by 2+3, in which case we would not expect the discontinuityillustrated in fig. 2.The above discussion is relevant to the reactivity at the cobalt atom. The transi-tion state for the replacement of the sixth ligand in CoA4XY must resemble the five-co-ordinate species which lies close to the ground state in the alkyl cobalt complexes.It is not surprising therefore that all their replacement reactions are fast.The unusual condition of the metal in these cobalt complexes should also beobserved in nickel(1V) l8 and iron(I1I) l3 alkyl complexes.The change in stereo-chemistry with change in oxidation state is related to that which occurs on uptakeof hydrogen or oxygen by d8 metal complexes which are known l9 to have interestingcatalytic properties.We thank the Medical Research Council and N.A.T.O. for financial support.' R. A. Firth, H. A. 0. Hill, J. M. Pratt, R. G. Thorp and R. J. P. Williams, J. Chem. SOC. A ,' H. A. 0. Hill and K. G. Morallee, J. Chem. SOC. A , 1969.1968, 2428.H. A. 0. Hill, B. E. Mann, J. M. Pratt and R. J. P. Williams, J. Chem. SUC. A, 1968, 564.D. C. Hodgkin, J. Linsey, R. A. Sparks, K. N. Trueblood and J. G. White, Proc. Roy. SOC. A ,1962, 266,494.G. Costa, G. Mestroni, G. Tauzher and L. Stefani, J. OrgammetaI. Chem., 1966, 6, 181.' P. G. Lenhert, Proc. Roy. SOC. A , 1968,303,45.' G. Costa, G. Mestroni and L. Stefani, J. Organometal. Chem., 1967, 7,493.* H. A. 0. Hill and I. D. McPherson, to be published.G. N. Schrauzer and R. J. Windgassen, J. Amer. Chem. SOC., 1966, 88,3738.J . Chem.Soc. A, 1968, 2419.1968, 11, 167.lo R. A. Firth, H. A. 0. Hill, B. E. Mann, J. M. Pratt, R. G. Thorp and R. J. P. Williams,'' H. A. 0. Hill, K. G. Morallee, G. Pellizer, G. Mestroni and G. Costa, J. Organometal. Chem.,l 2 P. Day, H. A. 0. Hill and M. G. Price, J. Chem. SOC. A, 1968, 90.l3 D. A. Clarke, R. Grigg and A. W. Johnson, Chem. Comm., 1966,208.l4 D. Day, Theor. Chim. Acta, 1967, 7,328.A. Eschenmoser, R. Scheffold, E. Bertele, M. Pesaro and H. Gschwend, Proc. Roy. SOC. A,1965,288,306.l6 R. A. Firth, H. A. 0. Hill, J. M. Pratt, R. J. P. Williams and W. R. Jackson, Biochem. 1967,6,2178. '' H. A. 0. Hill, K. G. Morallee and G. Pellizer, J. Chem. SOC., 1969, 2096.l8 R. Grigg, A. W. Johnson and G. Shelton, Chem. Comm., 1968, 1151.l9 L. Vaska, Accounts Chem. Res., 1968,1,335

ISSN:0366-9033    

DOI:10.1039/DF9694700165

出版商:RSC    

年代:1969

数据来源: RSC

摘要:

Formation of Cobalt-Carbon 0-BondsBY MICHAEL GREEN, J. SMITH AND P. A. TASKERDept. of Chemistry, University of York, YorkReceived 4th March, 1969The properties of N,N'-bis-(o-aminobenzylidene)-l,2-diaminoethanato-cobalt(II) and nickel(II)and various derivatives are described. The formation of stable cobalt-carbon o-bonds is discussedin terms of the energies of dzz and dxy orbitals.A stable a-cobalt-carbon occurs in vitamin B12 coenzyme (the corrin ring ofwhich is given as (I)). Although the majority of cobalt compounds are incapableof forming metal-alkyl bonds, there are a few macrocyclic systems from which stableCo-C links can be produced, viz., (II) to (VII) 2-7 :H3c7==N r? co+ N TcH.3mPnH3c?= NnN c o * J" H3PI.vm 17MICHAEL GREEN, J . SMITH A N D P .A . TASKER 173As all of these compounds contain conjugated systems and several an ‘‘ N4 ” system,it seemed likely that (VIII), N,N’-bis-(o-amin0benzylidene)- 1,Z-diaminoethanato-cobalt (II), “ cobalt amben ”, would also form metal-alkyl links, particularly in viewof its closeness to (VII), cobalt salen. (Nomenclature is explained in the appendix).The properties of cobalt amben (which has been reported briefly by Bailes and Calvin8)is described together with those of several of their derivatives.Many of the compounds listed undergo one electron reductions, so becomingisoelectronic with their nickel analogues. The properties of nickel amben andseveral derivatives are discussed. (Some of these compounds have already beenprepared by Pfeiffer’s g r o ~ p .~ )RESULTS AND DISCUSSIONThe following cobalt complexes were prepared: amben, ambtn, amaen andampen. They all appear to be monomeric, tetracoordjnate and low spin in thesolid form, in non-polar solvents and in pyridine, from comparison of electronicspectra and magnetic-subsceptibility measurements. Their magnetic moments liebetween 2.04 and 2.52 B.M. The following nickel complexes were also prepared :amben, ambtn, ambbuten, ambphen, ampen, ampbuten and Meamben. Similarity-of electronic spectra and measurement of magnetic susceptibilities lead to the con-clusion that they are all monomeric, tetracoordinate and diamagnetic as solids and insolution in non-polar solvents and in pyridine.Cobalt amben (which is in a formal oxidation state of (11)) can be reduced to itscobalt(1) form, e.g., using sodium sand; however, no evidence was found for theformation of any cobalt-alkyl link on treatment of the cobalt(1) systems with alkylhalides.Why were no cobalt-carbon bonds formed?Formation of a cobalt-carbon a-bond involves pairing of an electron of the 3d22orbital of the cobalt with one of an sp3 hybrid of a carbon atom. (The z-axis istaken to be the one perpendicular to the macrocyclic plane of the molecule. TheN4 or N202 systems are assumed to be square, so that the coordinating atoms lieon the x and y-axes). The cobalt complexes investigated here are low-spin, but thearrangement of orbitals, particularly the d22 and dxy, compounds is not known directly.However, Bosnich lo has shown that in nickel salpn the sequence runs dxz = dyz>dz2 > dxy > d , ~ - y ~ (in decreasing order of orbital stability).Theoretical treatmentsA B C DFIG. 1of d8 systems either agree with Bosnich‘s or run : dz2>dx2 = dyz> dxy>dx2-y~. Asboth sequences agree on the order : dz~>dxy>dx2-y2, and as the position of thed,,, dyz system is not crucial to the arguments put forward here, it will be assumed thatall the nickel complexes have the same arrangement as nickel salpn. The orbita174 FORMATION OF COBALT-CARBON 0-BONDSsequence will alter as in fig. 1 on going from a square planar d8 to an octahedral d6configuration (the change from left to right also being helped by the addition of axialligands). A d7 configuration could have therefore an orbital arrangement corre-sponding to A, B, C or D.Though formation of a cobalt-carbon link in cases Cor D merely involves simple pairing, in cases A and B an electron has also to bepromoted from the dz2 to the dxy orbital. Fig. 2 shows bonding schemes, which axeA B CI I I I I I' I1-11 dxy 1-1 1 1I - 1-1I \ i lI I II I II 1I - II I I I I II I dZ2 1 dxY i' \I \ / pP3] i [SP 31 :bp31 d,2 \, -'+ -1 ' , I t - 4dx, - dZz\w I I wFIG. 2much simplified. In cases B and C, a stable bond can be formed. However, Arepresents a case in which no bond is formed because the energy gained by bondformation would not be sufficient to offset that lost by having to promote an electronfrom the dzz to the dxy orbital.Thus, whether a bond can form or not depends(amongst other factors) on the energy separation between the dZ2 and dxy orbitals.It is suggested that this quantity is too great for cobalt-carbon bond formation incobalt amben which therefore corresponds to case A, in contrast to cobalt salenwhich must be of B or C type.Further evidence for this postulate comes from the failure of cobalt amben andits derivatives to react with oxygen. Unlike cobalt salen, for example, cobalt arnbendoes not take up oxygen either as a solid or in solution in dimethylformamide, whilecobalt ambtn, amaen and ampen do not change when left exposed to the air. Manyof the compounds listed in the Introduction (like most cobalt(I1) systems) reactreadily with oxygen giving either peroxy compounds or cobalt(II1) complexes, e.g.,Co l2 (VII), and C~(dmg)~, (11), respectively.In the fig. 2, [sp3] need notbe an orbital of a carbon atom from a formal point of view. It can equally wellrepresent a lone pair type orbital of an O2 system (if displaced downwards slightlybecause of a greater ionization potential). It is probably not a coincidence thereforethat cobalt amben does not react with oxygen.It is not sufficient merely to form a Co-C link. It is also important that it isnot too ionic in the sense, Co-C, otherwise heterolysis will occur. Ideally, forformation of a good covalent bond, the energies of the dz2 and sp3 orbitals should beequal. If 4 2 is too high in energy, as in case D which might represent a typicalammine system, a Co-C bond would be too ionic to exist.In contrast, a peroxygroup can tolerate more negative charge on it than an alkyl group can, so that it ispossible to form peroxy complexes, e.g., [(H3N)5C000Co(NH3)5]4+, in cases wherethe dz2 orbital is too high in energy for a cobalt-carbon bond to be stable. (Incidentally,the greater stabilization of the dz2 orbital in cases, A, B and C compared with Daccounts for the fact that systems in which a cobalt-carbon bond can be formed, canusually also be reduced to cobalt(1) d8 states.)To summarize therefore : if a stable Co-C link is to be formed the ligand systemmust stabilize the dz2 orbital, but not too much relative to the dxy orbital. The ligandslisted in the Introduction contain conjugated systems, and conjugated ligands arenormally high-field.This point is confirmed by the fact that several of the nickel(1I)complexes of these ligands are low spin, a property expected of d8 systems in whicha+ aMICHAEL GREEN, J . SMITH AND P . A . TASKER 175there is a large tetragonal field.13 Therefore in systems (I)-(VII), which all formCo-C links, the d orbitals would be expected to be well separated because of theligand field effect and thus the d22 orbital stabilized relative to the dX2-+ However,what evidence is there that in the amben compounds, the separation of the 4 2 anddxy orbitals is greater than in the ligand systems in which stable Co-C bonds occur?As these ligands vary considerably, the comparison will be limited to the one closestin structure to amben, i.e., salen.Data on weak transitions in the visible region are given in table 1 for nickelsalen and ampen compounds.The absorption in Ni salpn at 555 mp is attributed loto d,,,'~d,2-~2 excitations, and it seems reasonable to make a similar assignment forall the other ligands in the table. Though the dX,,*d,2-,,2 separation is greater innickel amben than in nickel salen, table 1 shows that this is not necessarily true forthe families as a whole.TABLE 1salen 540 mp l4 amben <481 mpsalpn 555 mp l osaltn 600 mp l4 ambtn 600 mpambuten 620 mpsalphen 540 mp l4 ambphen < 620 mpdt+d transitions in nickel salen, nickel amben and derivatives. (Unfortunately in nickel ambenand ambphen the d w d bands are obscured by stronger absorptions),The energy gap between the dxy and d , ~ - ~ 2 orbitals can also be assessed fromthe ability of nickel N202 systems to change from a low-spin planar complex to ahigh-spin octahedral one.Nickel salen, salpn, saltn and salphen are diamagneticand tetracoordinate as solids or in non-polar solvents. While nickel salen andsalpn remain so in pyridine, nickel saltn and nickel salphen, where the bridging unit(i.e., R, see appendix) is either longer or conjugated, become paramagnetic andhexac~ordinate.~~ These observations suggest a greater separation of the dx2 - y 2and dx,, orbitals in the amben family than in that of salen, as none of the nickelcomplexes investigated here, particularly those of ambtn, ambbuten and ambphenshow any appreciable changes in pyridine.This evidence could mean that in the amben compounds, relative to those ofsalen, the d,,, orbital is more low-lying (which might be quoted as evidence for asmaller separation of the dz2 and d,,, orbitals, rather than a larger), or that ambenexerts a larger ligand field, thus increasing d-d energy gaps in general.Thoughit is difficult to predict relative field-splitting strengths, the greater proton affinityof N- groups compared with 0-, and the positions of NH3 and OH2 in the spectro-chemical series are both compatible with the N4 system producing a larger field thanthe N202 one. So too is the fact that C-N bonds are longer than C-0, so thatrelative to their oxygen counterparts in salen, the two aliphatic imine nitrogen atomsin amben are forced closely to the coordinating metal. Moreover Yamada's Group l4has accounted for the difference in properties of nickel salen and nickel saltn mentionedearlier in terms of ligand field strengths.This evidence in favour of amben exertinga higher field than salen is useful here, as it is compatible with a correspondinglygreater separation of the d,,, and d2z orbitals.Direct evidence about the relative energies of the 4 2 orbitals in cobalt amben andcobalt salen can be obtained from redox potentials. Both compounds can be reducedelectrolytically to their cobalt(I) states, but it is easier to reduce cobalt(1I) salen byabout 0.1 V. The ionization potential d8+d7 (i.e., probably d-&+d,,,) is thus abou176 FORMATION OF COBALT-CARBON Cr-BONDS0-1 V greater for cobalt salen than cobalt amben, which at first sight implies that thedxy orbital is slightly more low-lying in the former.This is clearly too simple apicture however, as redox potentials indicate that the ionization potential, d7 +d6,is at least 0.6 V smaller for cobalt salen than cobalt amben. Cobalt salen can beoxidized to cobalt(I1I) salen compounds chemically and also electrolytically,when a half-wave potential of approximately 1.6 V is observed. In contrast, cobaltamben is remarkably resistant to oxidation chemically and electrolytically ; nocobalt(1II) compound is produced under conditions in which a cobalt(II1) salenderivative would be formed, nor is there any indication sign of electrochemicaloxidation at potentials up to 2.2 V.Unfortunately, ionization to a d6 state probablyinvolves loss of a dz2 electron and rearrangement to a (dxy)2(dyz)2(dxy)2 configuration.Moreover, one cannot be sure that the d7 configurations of cobalt amben and cobaltsalen are (dz~)2(dxy) as has been implied here. However, there is a reasonablecertainty that a Co(1) to Co(II1) redox process would involve the loss of two dxyelectrons. The process is at least 0.5 V more difficult in the cobalt amben systemthan in that of cobalt salen. Though this does not show conclusively that cobaltamben is case A, it does imply that the d,Z orbital is more low-lying in this compoundthan in cobalt salen.APPENDIXR2=H R1=H amb-R2=H R,=4 amp-R2=H R1=CH3 amaR2=CH3 R1=H Me-amb-Rphen/ \P. G. Lenhart and D. C. Hodgkin, Nature, 1961,192,937.G. N. Schrauzer and J. Kohnle, Chem. Ber., 1964,97,3056.G. Costa and G. Mestroni, Tetrahedron Letters, 1967, 4005.1968, 881.E. I. Ochiai and D. H. Busch, Chem. Comm., 1968,905.G. Costa, G. Mestroni, G. Tauzher and L. Stefani, J. Orgmometal. Chem., 1966,6,181.and C. Floriani, Chem. Comm., 1967, 139.4D. A. Clarke, D. Dolphin, R. Grigg, A. W. Johnson and H. A. Pimock, J. Chem SOC. C ,' G. Costa, G. Mestroni and L. Stefani, J. Organometal. Chenz., 1967, 7, 493 ; F. Calderazzo* R. H. Bailes and M. Calvin, J. Amer. Chem. SOC., 1947, 69,1886MICHAEL GREEN, J . SMITH AND P . A . TASKER 177P. Pfeiffer, T. Hesse, H. Pfitzner, W. Scholl and H. Thielert, J. prukt. Chem., 1937, 149,217.lo B. Bosnich, J. Amer. Chem. Suc., 1968,90,627.11 H. Basch and H. B. Gray, Inurg. Chem., 1967, 6, 365 ; F. A. Cotton and C. B. Harris, Inurg.Chem., 1967, 6,369 ; and references in both.F. Calderazzo, C. Floriani and J. J. Salzmann, J . Inurg. Nucl. Chem., Letters, 1966,2, 379.l3 G. Maki, J. Chem. Phys., 1958, 28,651.l4 S. Yamada, E. Ohno, Y. Kuge, A. Takeuchi, K. Tamanouchi and K. Iwasaki, Cuurd. Chem.l5 S . Yamada, H. Nishikawa and E. Yoshida, Proc. Jupun Acud., 1964, 40,211.Rev., 1968, 3, 247.

ISSN:0366-9033    

DOI:10.1039/DF9694700172

出版商:RSC    

年代:1969

数据来源: RSC

摘要:

Evidence for a -Donor Effect in Transition Metal-CarbonBonds from H-D Coupling ConstantsBY J. D. DUNCAN, J. C . GREEN, M. L. H. GREEN AND IS. A. MCLAUCHLANInorganic Chemistry Laboratory, South Parks Road, OxfordReceived 14th January, 1969The coupling constants JH-D of the CHzD group in the deuteriomethyl derivatives,Z - C ~ H ~ F ~ ( C ~ ) ~ C H ~ D , x-CSHs(Mo or W) (CO)3CH2D, Mn(CO),CH2D, x-C5HSNi(Ph3P)R(where R = CH2D, m- or p-C6H4CH2D), and XPt(Et,P),CH,D (where X=Br, NO3, or SCN),have been determined. The sign of JH-D in the iron compound has been shown to be negative.The magnitudes Of JH-H are discussed in terms of a x-donor effect from the metal to the neighbouringcarbon.Previous work on metal complexes containing simple organic groups as ligandshas suggested that the metal can considerably modify the chemical reactivity andproperties normally associated with the organic gr0up.l In an attempt to gainfurther information about the nature of these metal-ligand interactions, we haveprepared some transition metal-deuteriomethyl derivatives and have determinedthe values of JH+ of the CH2D group.EXPERIMENTALMATERIALSAll preparations and reactions were carried out under nitrogen or in vacuum.Allsolvents were thoroughly purified, dried and degassed before use.p-BROMO-M-MONODEUTERIOMETHYL TOLUENE.-The Grignard reagent p-BrC 6HaCH2 MgBr,prepared under strictly anhydrous conditions from magnesium and freshly distilledp-BrC6H4CH2Br, was treated with deuterium oxide (99.8 %). The mixture was stirredovernight and the deuteriomethyl derivative, p-BrCsH4CH2D was isolated in the normalmanner.m-Bromo-a-monodeuteriomethyl toluene was prepared in the same manner.PREPARATION OF THE MONODEUTERIOMETHYL DERIVATIVES.-Deuteriomethyl bromide(ca. 97 "/o) was carefully prepared starting from ketene and deuterium oxide. The acidCDH2COOD was isolated and the silver salt AgOCOCH2D was prepared from it. Thiswas treated with pure, dry bromine and the monodeuteriomethyl bromide was isolated inthe normal manner. The deuteriomethyl-transition metal derivatives were prepared eitherby reaction of the corresponding metal anion M- with CH2DBr (method A), or the Grignardreagent CH2DMgBr was prepared and reacted with a metal halide M-Cl (method B), orthe platinum complexes trans-X(Et3P),PtCH2D, where X = NO3 or SCN, were preparedby metathetical replacement, starting from the bromide trans-Br(Et3P),PtCH2D (method C).The methyl analogues of all the deuteriomethyl complexes have been previously reportedand the manner of preparation of the deuteriomethyl compounds, and the appropriatereference for the methyl complex are given in table 1.17DUNCAN, GREEN, GREEN AND MCLAUCHLAN 179TABLE 1method ofpreparation,compound reference MeCN- n-C5HSNi(PPh3)CH2D B, 3n-CS Hs Mo(C0)3 CH 2D A, 4 7 8.9Mn(C0)&H2D A, 5x ~ - C ~ H ~ W ( C ~ ) ~ C H ~ D A, 4 7 10.0BrHgCH2D B, 6tran~-BrPt(PEt~)~cH~D C B, 7 =F 10.0tran~-sCNPt(PEt~)~cH~D c, 7tran~-No~Pt(PEt~)~cH~D c, 7m-[n-C5HsNi(PPh3)]C6H4 .CH2D B, 3 -p-[n-CsH,Ni(PPh3)]C6H4 .CH2D B, 3 -m-BrC,H4. CH2D -p-BrC6H4. CH2D -X - C ~ H ~ F ~ ( C O ) ~ C H ~ D A, 4 - 8.3----PROTON MAGNETIC RESONANCE SPECTRAThe spectra were determined using a Japan Electron Optics Laboratory instrument at100 Hz. The double resonance experiments were carried out on a Varion H.A. 100 MHzmodified as previously described.2 The values of JH-D were determined at probe tempera-ture. An average of ten readings was made and the error of the determination is estimatedto be within 10'05 Hz. An internal check of the variation of field homogeneity was providedby the appearance of the CH2D tripIet.._ -20 Hz T- 5 HzFIG. 1 .---The 'H n.ni.r. spectrum of the CHzD group in, (a) trans-BrPt(Et,P),CHzD in MeCN and,(b) x-C5H5Fe(CO)&H2D in MeCN180 TRANSITION OF METAL-CARBON BONDSRESULTSThe deuteriomethyl complexes which have been prepared are shown in table 1,together with some 'H n.m.r.data. Examples of the spectra obtained are shownin fig. 1. Except for the platinum compounds, the full details of the lH n.m.r.spectra of the complexes are not given here since the spectra of the methyl derivativeshave been described elsewhere.The determination of JH-D of the complex n-C5H5Fe(CO),CH,D in twoconcentrations (approximately 20 and 50 weight %) showed that, within the experi-mental error, the values of JH-D were unchanged. However, as shown in table 1,JH-D is solvent sensitive and changes in JH+ as large as 0.7 Hz are observed.The relative sign of JH-D in the CH2D group of the complex n-C5H5Fe(CO),CH2Dhas been determined by 13C double resonance.Irradiation at the frequency corres-ponding to the low field band of the 13C central triplet caused a perturbation of thehigh-field band of the low-field triplet arising from coupling of the CH,D-hydrogenswith 13C (1.11 %) and the deuterium. Similarly, irradiation at the frequencycorresponding to the high field band of the central 13C triplet caused a perturbationof the low field band of the low field triplet of the CHzD hydrogens.DISCUSSIONThe 13C double resonance experiments show that the sign of JH--D is oppositeto that of Jc-D, and hence to JC-H. Since JC-H has been shown in 13CH3CN to havea positive sign the sign of JH-D in the iron complex n-CSH,Fe(C0)2CH2D isis likely to be negative.The sign of JH-H in a number of other methyl and methylenecompounds has been shown also to be negati~e.~ It is therefore a reasonable assump-tion that in all the deuteriomethyl-metal derivatives the sign of J H - H is negative,otherwise the range of value of JH-H of the metal derivatives would be much largerthan the range observed for other methyl and methylene derivatives.The large changes in values of JH.+ of some of the deuteriomethyl-metal complexeswith change in solvent is unexpected. They are not usually found for the JH-H ofsaturated hydrocarbons although changes of 0.25 Hz are observed for Jab inMeCHBr. CH,HbBr.l0 It may be assumed that in all the metal complexes theCHzD group rotates about the M-C axis sufficiently fast that only the averagevalue of JH-H is observed.It seems unlikely that different solvents could so restrictrotation that this would no longer be true. Most probably the solvent effects arisefrom differing degrees of interaction of the solvent molecules with the metal atoms.For example, it is known that MeCN, C6H6 and CS, may act as ligands to transitionmetals forming stable complexes.Pople and Bothner-By have developed a molecular Grbital theory of nuclear spincoupling between geminal hydrogen atoms.9 The theory provides a basis forexplanation of observed trends in the magnitudes of JH.+. The theory shows thatthe a- and n-effects of a substituent on a methyl group are independent and that ifthe substituent, relative to hydrogen, acts as either a a-withdrawer ( - I ) or a n-donorthe geminal coupling constant will increase (that is will become less negative).Conversely, a a-donor or n-acceptor substituent will cause JH-H to become morenegative.The values of JH--H for CH4,CH30H and CH3F are 3 12.4, -10.8 andT9.6 Hz re~pectively,~ here the increase in JH-H is attributed to the change ina-acceptor strength in the order, F>OH>H. The series, CH3F, CH3Cl, CH3Br,CH31 has values of ,IHvH = F9-6, 7 10.8, T 10.2 and 39.2 Hz respectively. Theorder here may be explained by change in a-acceptor power in the order, F>ClDUNCAN, GREEN, GREEN AND MCLAUCHLAN 181Br > I, together with the opposing trend of increase in n-donor power, I > Br > C1> F.The low values of JH-H in the compounds CH3CN and CH2(CN)2, JH-H = 'F 16.9and 18.7 Hz respectively are attributed largely to the n-acceptor properties ofthe CN group.gIt therefore follows that on the basis of this treatment we must explain the observedvdues of JH-H for the transition metal methyl complexes either by postulating thatthe metals are about as electro-negative as F,Cl or that they have a n-donor effect,which may be substantial, or both. There is little reliable evidence upon whichto base estimates of the a-electronegativity of transition metals in low oxidationstate compounds.Electronegativity values for the metals Mn, Fe, Ni, Mo, W, Hg,and Pt are, according to the Aldred and Rochow method of estimation, 1.60, 1.64,1-75, 1.30, 1.40, 1.44, and 1.44 respectively.There are also some lH n.m.r. datawhich are suggestive. Thus, the chemical shift of the CH3 hydrogens in a widevariety of transition metal methyl complexes occurs close to 10 z ; this suggests thatthe metals have similar electronegativity to the Me3Si group. Estimates of electro-negativities, from the relationship E = 0.64 6+ 1.78, where S is the chemical shiftseparation between the resonances of CH2 and CH3 hydrogens of transitionmetal ethyl complexes, give the values for the metals in n-C,H,Fe(CO),-,n-C5H,Mo(C0)3-, Mn(CO)5- of E = 1.98, 1-95 and 1.5 respectively. However,it is thought that, at the best, these values are very qualitative."It may be concluded that the a-electronegativity values of the metal in thecomplexes in table 1 probably lie near or below 2-0 and that the metals are certainlymuch less electronegative than atoms such as F, Cl or 0.If, therefore, the transitionmetals in the compounds have electronegativities lower than that of hydrogen, theincreased values (less negative) of JH-H of the metal compounds compared to thatof methane (JH-H = T 12.4 Hz) must be attributed to a n-donor effect from themetals to the methyl group.It is possible that there may be direct interaction between the filled metal d-orbitalsand the 1s-orbitals of the methyl hydrogens.The Pople and Bothner-By treatment i s only concerned with the changes in thecontributions of the hydrogen orbitals to the bonding and anti-bonding molecularorbitals. It does not distinguish whether these changes are brought about as aresult of interaction of the metal electrons with the cc-carbon or by direct interactionof the metal electrons with the hydrogens themselves.No distinctive trends in the values of JH-H for the different deuteriomethyl-metalcomplexes may be distinguished.However, it may be tentatively argued that theincrease in JHWH for the complexes, Mn(CO),CH,D < n-C,H,Fe(C0)2CH,D <n-C5H,Ni(Ph3P)CH2D, reflects the increasing n-donor power of the complexedmetal atoms. This is a function of the number and n-acceptor power of the otherligands. And it is generally thought that there is a change in n-acceptor powerin the order, (CO), > n-C5H5 : CO> Ph3P.Independent support for the proposal that there is a strong n-donor effect alongthe Ni-C bond in the complex n-C5H5Ni(Ph3P)CH2D comes for the values ofJH-H in the complexes m- and p-n-C5H5Ni(Ph3P)C6H4CH2D.The values forthe corresponding m- and p-CH2DC6H,Br are also given. The values of inthe nickel derivative are greater than those of the bromo-analogues (assuming anegative sign and only moderate solvent effects) indicating a metal n-donor effect.Further, the value of JH--H for the para-compound is probably greater (or not less)than that for the meta-compound. This is reasonable since n-donation effects areexpected to be greater to the para- than to the meta-position, while withdrawal isexpected to be less (cf. the values for the bromo-analogues)182 TRANSITION OF METAL-CARBON BONDSThe above proposal that in low oxidation state complexes transition metalshave available filled orbitals which can take part in x-donor, or other donor inter-actions is consistent with a wide range of other independent observations.Thus,preliminary studies on the photo-electron spectra of the complexes (CO),MnR,where R = CH3, H, COCF,, CF, or Mn(CO)S also suggest there to be a n-donor effectfrom the Mn to the CH3 group.12 The low stretching frequency in the complexesn-C,H,Fe(CO)LCOMe of the acyl C=O group and the protonation of these complexesforming cations which contain the system Fe=C(OH)Me shows the tendency of theiron to n-bond to the a-carbon.13 Further, there are the unusually low pK, of theacids M-CH,COOH, where M is, for example, Mn(CO),, n-C5HSFe(C0)2.14 Finally,there is the ease of loss of a hydride ion from the methyl group of M-ethylderivative^.^ These last two observations suggest that not only may metal electronsform n-symmetry interactions with alpha-carbon substituents but that the electronsmay also interact directly with beta-substituents causing considerable chemicalmodification.We thank Turner and Newall for financial support (to J.C. G.) and the ScienceResearch Council for a research grant (to J. D. D.).G. E. Coates, M. L. H. Green and K. Wade, OrganometalEic Compounh, (Methuen, London,K. A. McLauchlan, D. H. Whiffen and L. W. Reeves, Mol. Phys., 1966,10,131 ; and A. Charlesand W. McFarlane, Mol. Php., 1968, 14,299.H. Yamazaki, T. Nishido, Y. Matbumoto, S . Swmida and N. Hagihara, J. Oryanometal. Chem.,1966, 6,86.T . S. Piper and G. Wilkinson, J. Inorg. NucZ. Chem., 1956, 3, 104.W. Hieber and G. Wagner, 2. Naturforsch., 1957, 12b, 478.K. H. Slotta and K. R. Jacob, J.prakt. Chem., 1928, 120,272.1968), VOI. 2, pp. 209-217.’ J. Chatt and B. L. Shaw, J. Chem. SOC., 1959, 705. * G. Englert and S . Saupe, MoZ. CrystaZs, 1966, 1,503.J. A. Pople and A. A. Bothner-By, J. Chem. Phys., 1965,42,1339.lo H. Finegold, Proc. Chem. SOC., 1962, 213.l1 A. Davison, J. A. McCleverty and G. Wilkinson, J . Chem. Soc., 1963, 1133.l2 S. Evans, J. C. Green, M. L. H. Green, A. F. Orchard and D. W. Turner, Disc. Faraduy SOC.l 3 M. L. H. Green and C. R. Hurley, J. Organometal. Chem., 1967, 10, 188.l4 J. K. P. Ariyaratne, A. M. Bierrum, M. L. H. Green, M. Ishaq, C. K. Prout and M. G. Swan-i 5 M. L. H. Green and P. L. I. Nagy, J. Organometal. Chem., 1963,1,58.1969, in press.wick, J. Chem. SOC. A, in press ; idem, Chem. Comm., 1967,430

ISSN:0366-9033    

DOI:10.1039/DF9694700178

出版商:RSC    

年代:1969

数据来源: RSC

摘要:

Bonding in Hexame thyl-dialuminium and Related CompoundsBY K. A. LEVISON AND P. G. PERKINSDept. of Pure and Applied Chemistry, University of Strathclyde, Glasgow, C.l.Received 4th March, 1969A study of the electronic structures of A1Me3, AI2Me6, AIMeHt and some isomers of AlzMe2H4has been made by the S.C.M.O. method. Basis sets of 3s, 3p and diffuse and contracted 3d orbitalsplus all valence electrons were employed for the aluminium atoms. The 3d orbitals stabilize thedimeric species by direct, cross-ring G, x and 6 Al-A1 interactions : however, those of x type appearto be the most effective. Participation of the 3d orbitals in bridge bonding is more limited.The conformation of AlzMe,H, which seems to be energetically the most favoured is that inwhich the hydrogen atoms form the bridges between the monomer units whilst the methyl groupsare terminal and lie mutually trans.When the methyl group acts as a bridge, as in A12Me6, it isexpected that the two groups of three hydrogen atoms will lie in the eclipsed configuration.INTRODUCTIONConsiderable interest attaches to the electronic structures of trimethylaluminium,its dimer A1,Me6, and the related compounds in which methyl groups are replacedby hydrogen. This is because of (a) their relationship to the technically importantZiegler and Ziegler-Natta catalysts and (b) the problems posed by the bondingin such compounds.In polymeric organo-aluminium compounds the monomer units may be heldtogether by methyl groups, as in A12Me6,1 or by hydrogen atoms, as is thought to beso in the trimer (A1Me2H)3,1 or by bridging phenyl groups., The possibility of thevinyl group functioning as a bridge also exists.On the other hand, the correspondingboron compounds are usually monomeric.1In order to clarify the reasons for these phenomena it is desirable to study (a)the relative efficacy of Me, H, or C2H3, as bridging groups between aluminium atoms,(b) the importance of cross-ring bonding in the dimers, (c) the detailed orbital make-upof the M-X-M bridge and the M-M bonds, (d) the extent to which the emptyd-orbitals of aluminium participate in the bonding of the molecule and hence whetherthe dominant interactions involving d-orbitals are of Q, n, or &type. To theseends, we have carried out all- valency-electron self-consistent molecular orbitalcalculations on selected model compounds in the series.These comprised AlMe3, A12Me,, AlMeH, and three forms of the dimerAl,Me2H,, one in which both methyl groups function as bridges and the otherswhere they are terminal and cis or trans with respect to each other,METHODS OF CALCULATIONSThe structural parameters for all the models are given in table 1.In general theywere taken from Interatomic Distances but where the experimental lengths andangles were not available, reasonable values were assumed by comparison of themodel wit1 an analogous compound. In the dimers the terminal and bridginggroups are situated in two perpendicularly intersecting planes.I8184 BONDING I N HEXAMETHYL-DIALUMINIUMThe approach was based largely on the CNDO method of Pople, Santry, andSegal.4* This assumes that molecular orbitals for both 0 and 7c electrons may bewritten as linear combinations of orbitals sited on different atoms.In the presentwork the orbital co-efficients were calculated by diagonalizing a Fock interactionmatrix with elements,andF p v = ~ p v - 3 p ~ v Y I L Y ’ (2.2)The quantities Hpp, H,,, Ppp, Ppv and ypp, ypv are the diagonal and off-diagonalelements of the core Hamiltonian matrix, the density matrix and the electron repulsionTABLE 1 .-STRUCTURAL PARAMETERSmonomer dimerBOND LENGTHS (nm)0-200 0.200- 0.2240.1 60 0-1 60- 0.1 940.109 0.109BOND ANGLES (deg.)Ctem.-Al-Cterm. 120 1 24Ctem.,Al-Hterm. 1 20 124Hterm.-A1-Hterm.120 1 24Cbridge-A&Cbridge - 11097 H bridge-Al-HbSdge -H-C-H 109 109matrix respectively. Basis sets including or neglecting the aluminium d-orbitalswere chosen and in this context we considered it more satisfactory to use differentCoulomb repulsion integrals for s, p , and d orbitals rather than those based solelyon s-orbitals as previously e~nployed.~. Use of distinct integrals is importantbecause there is a significant numerical difference between the integrals ys,, y p p andYdd for a second-row atom. Both Hpp and Hpy are related to the valence-stateionization potentials, the latter through the Mulliken-Wolfsberg-Helmholtz equation.6Overlap integrals necessary for the calculations incorporated the form of orbitalsuggested by Burns since these have a non-zero exponent for the negative ion wherethe “extra ” electron is in a 3d orbital.The mutual disorientation of bases ondifferent atoms in the molecules requires that a set of matrix transformations becarriedout onall orbital pairs so as to resolve the 0, n,and 6 contributions to the bonds betweenthe two atoms. Because of this and the large number of possible overlap conibina-tions, it proved most convenient to perform the overall calculation in two parts.The first of these produced the starting Fock matrix, whilst the second carried outmatrix diagonalization and iteration to self-consistency. Most of the calculationswere carried out on the Newcastle University KDF9 and the Strathclyde UniversityICT1905 computers. The basis set required for Al,Me6, however, proved to be toolarge for the above machines and this molecule was therefore processed on the S.R.C.Chilton Atlas computerK.A . LEVISON AND P . G . PERKINS 185The atomic input parameters were valence-state ionization potentials and one-centre electron repulsion integrals. The former were calculated from atomic spectro-scopic data whilst the latter were obtained purely theoretically by integration afterexpanding the operator 1 /r12 in spherical harmonics.A conceptual difficulty arises with 3d orbitals because these are formally emptyin the neutral aluminium atom. An electron placed in such an orbital would, ineffect, be " outside " the neutral core of the atom and so the effective nuclear chargeacting on it, and hence its energy, should, as a first approximation, be that appropriateto a negative ion.However, in a compound, the surrounding ligands generally bringabout unshielding of the A1 nucleus with concomitant increase in Z* and contractionof the 3d radial function. A second, probably more realistic, value for Z* is thenthat for a d electron in the neutral atom (e.g., for the configuration ls22s22p63s23d).In the present calculations we considered both the above situations and table 2 liststhe input data for all atoms in the molecules.TABLE 2.-INPTJT PARAMETERS Ip AND Y p p (kT)atomH 1s c 2s2PA1 3s3P3d (Al")3d (Al-)Ifi1.2601.8761.0301.0920.57610-18910.0724YfiP1.9681.5051.3870.89250.63680.36310.2420RESULTS AND DISCUSSIONIt is convenient first to discuss points of interest which arise from the study of themodel system MeAlH,-Al,Me,H, and to consider features of this model whichbear on the interpretation of the electronic structure of AlMe, and Al,Me,.In thisinvestigation comparisons were made between the monomer and a number of dimericspecies based essentially on alternative methyl or hydrogen bridging.First it was necessary to try and establish the most energetically favourablepositions and orientations of the two methyl groups in the two dimers. When boththe methyl groups are in terminal positions the calculations showed that, althoughthe electronic energy factors (bonding plus repulsion) favour the cis dimethyl(I),the overall stability of the molecule is dominated by the internuclear repulsion factorandthis isappreciablysmallerwhen the methyl groups are trans as in (TI). Compound(111) was not studied but it would be expected to be less stable than (II) because theinternuclear repulsions would be greater than in the latter coupled with little or noconcomitant gain in electronic energy.In the methyl bridged model (IV), it is possible for the six hydrogens in the twomethyl groups to lie either eclipsed or staggered with respect to each other.Cal-culations on the two models showed that their nuclear and electronic energies aremutually compensating so that approximately the same overall molecular energyis obtained for each. This does depend on the actual extent to which the d orbitalsparticipate in the bonding but using contracted d-functions the eclipsed form isfavoured by -178 kJ.This result is cogent in fixing the orientation of the methylbridging groups in A12Me6.The relative efficacy of methyl or hydrogen bridging in the dimer was next com-pared in the models (11) and (IV). The results are interesting because they show,first, that the electronic energy o f eclipsed (IV) is lower than that of (11), whateve186 BONDING IN HEXAMETHYL-DIALUMINIUMthe choice of d orbital parameters (for the contracted d-function the energy differenceamounts to 12519 kJ). This means that methyl bridging is electronically more stable.However, in (II), the nuclear energy is lower than in eclipsed (IV) by 13689 kJ, withthe net result that hydrogen bridging should be favoured in this type of compound.MeMe \ / H \ /H /HMe’ \”’ \ ” At Atff’ \HA \MeA lII IHH19FIG.1.-Some isomers of Al,Me2H,.mIn the above calculations, d orbitals were included in every case. It is pertinentto enquire to what extent the d orbitals are involved in the bonding of such systems.Table 3 shows how the total energy of the molecules studied diminishes on inclusionof both diffuse and contracted d orbitals. The zero of energy is that for the casewith the d orbitals omitted. Participation of these orbitals in the dimer A12Me2H,is somewhat greater than in the monomer and this must be due, in some measure,to the Al-A1 cross-ring interactions which can be of 6, n, or S type. It is interestingthat, although the electronic energy of the methyl bridged model is greater than thatof the hydrogen bridged species, nevertheless the energy lowering by the d orbitalsis less in the former than in the latter.This is consistent with an increase in the cross-ring bonding interaction brought about by closer approach of the two Al atoms whena small atom like hydrogen forms the bridge.TABLE LO LOWERING OF GROUND STATE ENERGY BY d ORBITAL MIXING (kJ)d orbitals d orbitalsincluded* includedAlMeH2 33-94 130.81A12Me2H4 (H-bridged) 160.05 755.48AlMe3 62.21 225 -00* Z* appropriate to Al- used in the calculations.A12Me6 368.44 IThe magnitude of the Al-Al bonding components in the hydrogen bridgedmolecule is given in the orbital density matrix. These figures were taken from thecalculation employing a contracted d-function.All bond orders below 0.01 havebeen neglectedK . A. LEVISON A N D P . G . PERKINS 187The most important contributors to the cross ring bonding are the pu--pu andp,-p, components. However, the d,-p, (bond order 0.251) and d,-d, (bondorder 0.123) parts are also quite prominent and account for the increased d-stabiliza-tion energy of the dimer over the monomer. A 6 contribution also appears but isnot marked in this particular molecule. This is because the dX2-,2 orbital is mainlyinvolved in bonding to the bridging hydrogen atom (see table 4). Moreover dxy,TABLE 4.-PRINCIPAL COMPONENTS OF THE Al-A1 AND A1-Hbridge BONDS IN Al,Me2H4Al-A1 3dxz 3dyz 3dxy 3 4 2 3dx2-y2 3s 3Px 3PY 3Pz0.426- 0.1 51- 0.018 0.367- 0.251 0.039 - 0.1230.151 - - 0.054- - - - 0.014 -0.013 0.021 - - - - -- 0.017 0.141 - 0.030 - - 0.0543Pz 3dxz 3dyz 3dxy 3dz2 3dxZ-yZ 3Px 3PY0.460 - 0.375 0.162 0.031 0.015 0.036 0.148the second 6 bonding orbital (in the diatomic case), does not interact strongly withany other orbital.It is interesting also that du--du bonding appears to be insignificantin the Al-Al bond. The Al-H bridge bond is dominated by its 3p-1s and 3s-1scomponents, although the d orbitals do contribute to some extent.The overall charge distribution (table 5) is stable to variation in the d orbitalexponent and reveals that the aluminum atom is highly negatively charged. This isunexpected in view of the nature of the groups to which it is attached.TABLE 5.-ELECTRONIC CHARGE DENSITIES FOR THE METHYLALUMINIUM HYDRIDE SYSTEMatomAIMeH2 AILMe2H4 bridged by hydrogen A12MeiK bridged by methyl groups3d orbitals 3d orbitals 3d orbitals 3d orbitals 3d orbitals 3d orbitals 3d orbitals 3d orbitalsincluded* included omitted included* included omitted included* includedA1 3s 1.376A1 3p (total) 1.739A1 3d(total) 0.027total charge on A1 3.142@ridge 2s -Cbridge 2p -total charge on C -Hbridge Is -1.370 1.471 1.481 1.497 1.459 1-471 1.4851.695 1.850 1-803 1.726 1.803 1.761 1.4980.119 - 0.055 0.291 - 0.041 0.2783.184 3-321 3.339 3.514 3.262 3.273 3.261- - - 1.574 1.582 1.607- - - - 2.172 2.131 2.123- - - 3.746 3.713 3.730- 1.155 1.151 1.047 - - ---* Z* appropriate to Al- used in the calculations.HEXAMETHY L-DIALUMINIUMOnly the calculation employing the diffuse 3d function has been carried out todate.It is found (table 3) as before, that the dimeric species is stablized by the dorbitals more than twice as much as is the monomer AlMe,. This suggests thatdespite the greater distance apart of the aluminium atoms, there is a degree of Al-A1bonding when methyl rather than hydrogen acts as the bridging group.There is a strong drift of electrons to the metal atoms from the carbon atoms ofthe system (table 6) and this feature is even more marked in Al,Me6 than it is i188 BONDING IN HEXAMETHYL-DIALUMINIUMAI,Me2H,. The result is that the metal atoms carry an overall negative chargewhich is near unity.The A1-A1 bonding interactions (table 7) are not clearlyseparable into 0, n, and 6 contributions because of the destruction of diatomic sym-metry by the attached methyl groups. However, bond orders involving the 3py andatomTABLE 6.-ELECTRONIC CHARGE DENSITIES FOR AIMe3 AND &Me6AlMe3 &Me63d orbitals 3d orbitals 3d orbitals 3d orbitals 3d orbitalsomitted included* included omitted included*A1 3sA1 3p (total)A1 3d (total)total charge on A1Cbridge 2p (total)total charge onH 1s (on Cbridge)Cbridge 2s1.417 1.415 1 -407 1.5292-174 2.170 2.116 2.398- 0.048 0.193 -3.591 3.633 3.716 3.927I 1 -557- 2.069- 3.626I 1.104- -- -- -- -1.5402.3210-1 373.9981 -5652.0163-5811.1 10* Z* appropriate to Al- used in the calculations.3dzz orbitals are mainly of Q type, those involving 3p,, 3pz, 3dyz, and 3dxy are mainlyof n type, whilst the 3dx, orbitals bring about limited &bonding.The bond orderbetween the two 3dX2-,,2 orbitals resolves into both d, and C& components neitherof which is strongly dominant.TABLE 7.-PRINCIPAL COMPONENTS OF THE A-AI BOND IN &Me6A1 3s3s 0.1 503Px3py 0.19839,3dxz3dyz3dxy---I -3dzz 0.0403dx2 -,.2 0.0300-5420.464- - 0.065- - 0.199 - 0.033- 0.070 - -- 0.090 .- - - - 0.014 0.017I - - -I - 0.021 0.106 - -- - 0.023Although the overall bond is dominated by the A13px3px(~) and Al3py3p,(0)contributions, nevertheless the inclusion of the 3d orbitals appears to add consider-ably to its strength as witnessed by the 3p,3dY, and 3px3dXy bond orders (both 3p,3dR)which are 0.199 and 0.106 respectively.By contrast, the d--d interactions do notadd greatly to the stability of the system.TABLE 8.-PRINCIPAL COMPONENTS OF THE A1--Cbridge BOND IN AI,Me63s 3Px 3PY 3Pz 3dxz 3dyz 3dxy 3dZ2 3dx2--y2Cb2~ 0.130 0.246 0.195 - - - 0.080 0.049 0.029Cb2px 0.211 0.270 0.289 - - - 0.045 0.036 0.061Cb2py 0.082 0.038 0.022 - - - 0-021 - 0-01 7Cb2PZ - - - 0.070 0.029 - - 0.014 -The orbital-orbital bond orders between aluminium and a bridging carbon atomare listed in table 8. Here there are only Q and n components which are not separablebecause the basis functions are mutually disorientated with respect to the Al-K. A . LEVISON AND P. G . PERKINS 189internuclear axis. However, participation of the aluminium 3d orbitals in the bond,though not negligible, is much weaker than that of the aluminium 3s and 3p basissets. It is expected that a second calculation, in which the contracted 3d orbital isintroduced, will show the same relative trends. Hence, it seems that the 3d orbitalsof aluminium influence the dimerization of some compounds of the element bytheir contribution to a metal-metal bond rather than by bringing about any uniquemodification of the electronic structure of the bonds to the bridging units.One of us (K. A. L.) thanks the S R.C. for a maintenance grant.G .E. Coates and K. Wade, Organometallic Compounds, vol. 1 (Methuen, London, 1967), chap. 3.J. F. Malone and W. S. McDonald, Chem. Comm., 1967,4,444.Tables of Interatomic Distances and Configuration in Molecules and Ions, Chem. SOC. Spec.,Publ., 1958, no. 11.J. A. Pople, D. P. Santry and G. A. Segal, J. Chem. Phys., 1965, 43, S.129.D. P. Santry and G. A. Segal, J. Chem. Phys., 1967,47,158.R. S . Mulliken, J. Phys. Chem., 1952,56,295 ; M. Wolfsberg and L. Helmholtz, J. Chem. Phys.,1952, 20,837. ' G. Burns, J. Chem. Phys., 1964,41,1521. * C. E. Moore, Nat. Bur. Stand. circ. no. 467, 1949

ISSN:0366-9033    

DOI:10.1039/DF9694700183

出版商:RSC    

年代:1969

数据来源: RSC

摘要:

Complexes of Carbon Suboxides with Silver:New Chemistry of SilverBY E. T. BLUES AND D. BRYCE-SMITHDept. of Chemistry, The University, Whiteknights Park, Reading, BerkshireReceived 8th April, 1969Reaction of silver1 acetate with acetic anhydride at reffux under argon gives an explosive silvercomplex of empirical formula Ag3C302 (I). Pyrolysis of complex (I) gives carbon suboxide as themajor gaseous product. Decomposition of (I) by aqueous alkali, followed by acidification, gives somemalonic acid. The complex (I) isolated under strictly anhydrous conditions is diamagnetic, buttreatment with dilute acetic acid gives a complex (11) of the same empirical formula which exhibitsan unusual type of strong field-dependent antiferromagnet ism. Complexes (I) and (11) are oxidizedby oxygen at 20" in daylight, but not in the dark, to compIexes of empirical formula Ag3C303.Oxygen is taken up (reversibly) in the dark at -78".Molecular nitrogen is taken up photochemicallyat 20°, even in presence of oxygen. There is evidence that these complexes contain silver atom clust-ers. Certain of the silver complexes undergo ligand exchange with iodide and bromide ions to formsilver halides of empirical formula AgHal which exhibit field-dependent antiferromagnetism both assolids and in solution, These new forms of silver halides are thought to contain silver atcm clustersinherited from their precursors. Complexes akin to complex (I), but which differ in certain properties,are obtained at room temperature from many silver1 salts and acetic anhydride in the presence oftertiary amines such as pyridine or triethylamine.These may be complexes with CzO (ketenides),and complexes (I) and (11) may contain elements of a ketenide structure.~ ~ ~ ~~~The ligand properties of carbon suboxide C,02, are virtually unknown, althoughthere is said to be a reaction with haemog1obin.l Further, no simple carbonyls ofsilver have been described. We describe now the preparation and some propertiesof complexes of silver which appear to contain the elements of C302 and C20. Certainof these complexes are remarkable in that they appear to contain silver in an oxidationstate less than + I, are antiferromagnetic (but see text), undergo photo-oxidation byoxygen, and photochemically fix molecular nitrogen. Investigation of these complexesis far from complete but the preliminary results appear sufficiently unusual to justifydiscussion here.Although silver metal appears to react superficially with carbon suboxide atroom temperature, a more satisfactory indirect general procedure has been discoveredin the reaction of silver' salts (e.g., the acetate, nitrate," or phenylacetylide) withacetic anhydride above 100".Reactions to give similar, but not identical, complexesoccur readily at 20" under the catalytic influence of tertiary bases such as triethylamine.The complexes are yellow, and are usually formed in yields of 90 % and above. Theyare insoluble in water, concentrated aqueous ammonia, and all common organicsolvents, including pyridine, dimethylformaniide, dimethyl sulphoxide, and triethylphosphit e.EXPERIMENTALMATERIALSOxygen-free argon was dried by passage through a gIass spiral at - 78".Acetic anhydridewas fractionally distilled under argon, and an acid-free fraction was used. Silver salts were* The use of the nitrate is not recommended because of the tendency to form dangerously expZusivered mixed complexes containing nitrate, Silver acetate is satisfactory for general studies, althoughthe complexes from it are mildly explosive in an uncompressed state and must be handled with care.19E. T . BLUES AND D. BRYCE-SMITH 191of analytical reagent grades, recrystallized from water where appropriate. All were con-firmed to be diamagnetic.ANALYSIS AND SPECTRASilver estimation by direct ignition of an explosive complex was impracticable, but priortreatment of a complex with a mixture of acetic acid and pyridine (5 : 1) at 110" for 1 h,followed by evaporation to dryness, gives a residue which could be safely ignited to givesilver metal.Silver analyses by this procedure were consistent to within &0-1 %. Carbon,hydrogen and nitrogen were estimated by Mr. D. Abbott of this Department using an F. andM. Scientific 185C, H, and N analyzer (reproducibility f0.15 %.)Infra-red spectra of silver complexes and their derivatives as mulls in Nujol or hexachloro-butadiene were recorded in the region 400-4000 cm-l, with a Perkin-Elmer grating spectro-photometer (model 237 or 457). Infra-red spectra in the 40-400 cm-1 region were recordedby the Physicochemical Measurements Unit, Harwell.Magnetic susceptibilities weremeasured by the Gouy method using a variable-temperature apparatus (Newport Instru-ments).APPARATUSThe apparatus consisted of a 500 ml round-botomed three-necked flask equipped with areflux condenser, thermometer, and a combined gas-inlet and ground-glass stirrer guidesupporting a driven glass shaft fitted with a Teflon blade. The gas exit from the top of the con-denser was connected to a Nujol trap through a 10-cm pathlength cell with sodium chloridewindows. Provision was made for the composition of the effluent gas to be monitored byinfra-red analysis. Particular care was taken at all stages to avoid contamination by metallicimpurities which might have led to spurious magnetic results.PREPARATION OF SILVER COMPLEXESWITHOUT A CATALYSTSilver acetate (0.1 mole) suspended in acetic anhydride (3.0 mole) was stirred at reflux(ca.138") under argon for 16 h. (Neither air nor nitrogen may be regarded as inert:complexes prepared under nitrogen are brown, and have contained up to 1 % combinednitrogen.) A large excess of acetic anhydride is desirable because acetic acid is the majorby-product, and tends to inhibit the reaction. Small amounts of C3O2, COZY CO, and CH4were evolved during the first 30 min of the reaction, but subsequently only traces of C02were found. The canary-yellow insoluble product was separated by filtration under argonat 20°C. The clear yellow filtrate contained acetic acid (0.019 mole), methyl acetate (0.0007mole), and ethyl acetate (0.00006 mole) (g.1.c.).Evaporation of the filtrate at ca. 60"/10 mmyielded ca. 0-02 g of a brown residue, largely silver acetate (i.-r. spectrum).The yellow complex in the filter was divided into two portions. The first portion wasfreed from acetic anhydride in a stream of dry argon. This material [complex (I)] wasdiamagnetic (x = - 0-26 x c.g.s.u.) and had the empirical formula Ag3C302 (found:C, 9-3 ; H, -0.1 ; Ag, 82.0. Ag3C302 requires C, 9-3 ; Ag, 82.7 %). The infra-red spectrumis given in fig. 1.The second portion of the complex was washed consecutively with acetic acid and water,and dried under argon. This material [complex (II)] had the same empirical formula ascomplex (I), within analytical error (found: C, 9-4 ; H, 0.2 ; Ag, 82.0 %) but exhibited field-dependent antiferromagnetism.[At field-strengths of 7.3, 8-2 and 9.6 kilogauss, =1.70, 1.42 and 1.22 B.M. respectively : the plot of reciprocal subsceptibility against tempera-ture over the range 80-320°K for the iodide (see below) derived from complex (11) by treat-ment with 7 N potassium iodide, was characteristic of an antiferromagnetic exchange inter-action (fig. 5).] The infra-red spectrum was virtually identical with that of complex (I)(fig* 1)1 92 COMPLEXES OF CARBON SUBOXIDES WITH SILVER1.04000 3 0 0 0 2 0 0 0 1500 1000 5 0 0 250L . , , I , , . ,: , , I , ! N . , ,cm-'FIG. 1.-Infra-red absorption spectra of silver complexes from reaction of silver acetate with aceticanhydride at 138"C, as mulls in Nujol : A, complex (I) (diamagnetic) ; B, complex (11) (antiferro-magnetic).(N denotes Nujol absorption).0cm-'FIG. 2.-Infra-red absorption spectra of silver complexes from reactions of silver1 salts with aceticanhydride at 138"C, as mulls in Nujol: (-), complex from silver nitrate; (---), complex from silvertrifluoroacetate.cm-IFIG. 3.T-fra-red absorption spectrum of silver complex (IV), prepared by reaction of acetic anhydridewith silver1 oxide in pyridine at 20"C, as a mull in Nujol.Complexes have been obtained by similar procedures with acetic anhydride and silver1formate, -benzoate, -propionate, -sulphate, -carbonate, -nitrate, -trifluoroacetate, and-tetrafluoroborate : the last two salts were first dissolved in toluene.These complexes havenot yet been investigated in any detail but all have infra-red absorption similar (but notidentical with) to those of complexes (I) and (11) in the region 400-4000 cm-' (fig. 1-3), givecarbon suboxide when pyrolyzed under argon (see below) and are insoluble in all commonsolvents including 0.880 ammonia and pyridine.Silver' halides, silver1 cyanide and silver oxides do not react with acetic anhydride toform complexes under these conditions.FORMATION OF SILVER-CARBON SUBOXIDE COMPLEXES FROM SILVERI SALTS ANDTypically, a solution of silver' salt (0.1 mole) in pyridine or triethylamine (0.5 mole)was added to acetic anhydride (1.0 mole) under argon at 20°C. A bright yellow gel formedACETIC ANHYDRIDE, CATALYZED BY TERTIARY AMINEE.T. BLUES AND D. BRYCE-SMITH 193within oiie minute. (Silver nitrate, -trifiuoroacetate, and -tetrafluoroborate in pyridinedo not appear to react with acetic anhydride to form insoluble complexes, but they do SO intriethylamine.) The yellow gel was broken up by stirring, and after ca. 30 min the complexwas filtered off and washed with pyridine. Excess of pyridine was removed at room tempera-ture in a rotary evaporator. The complex (111) obtained from a suspension/solution ofsilver' oxide (0.05 mole) in pyridine (0.5 mole) and acetic anhydride (1.0 mole) under theseconditions was diamagnetic (x2O0 = - 0 . 2 4 ~ c.g.s.u.) and had the empirical formulaAg9C906 . Py, (found : C, 23.5 ; H, 1.4 ; Ag, 65-0 ; N, 3.8. Ag,CP06Py4 requires C, 23.4 ;H, 1.4; Ag, 65-1 ; N, 3.8 %).Treatment of complex (111) with 5 % aqueous acetic acidgave a pyridine-free diamagnetic complex (TV) having the empirical formula Ag7CsO4H3(found: C, 10.6; H, 0.4; N, 0-2; Ag. 82.3. Ag,Cs04& requires C, 10.6, H, 0.4; Ag,83.1 %). The main component appears to be Ag2C203 since reaction with CH3T gave theknown dimer of dimethylkeytone. Pyridine was also readily removed at 50" under reducedpressure. The infra-red spectrum of complex (TV) in the region 400-4000 cm-I is shown infig. 3.PHOTO-OXIDATION OF SILVER COMPLEXESA finely divided silver complex was agitated either in the dry state or in aqueous suspensionunder oxygen at 20°C and exposed to sunlight or light from a 250-W tungsten-filament lampfor a total period of ca.24 h. When a tungsten lamp was used, the reaction vessel wasimmersed in a thermostat at 20°C. Photo-oxidation of the diamagnetic complex (I)(Ag3C302) under these conditions gave a complex (V) with the empirical formula Ag,C3O3(found : C, 9.1 ; H, -0.1 ; Ag, 79.2 ; Ag3C303 requires C, 8.8 ; Ag, 79.4) which exhibitsfield-dependent antiferromagnetism (at field-strengths of 7.3,8-2 and 9.6 kilogauss, xfS x lo6 =6-12, 4.61 and 3.20 c.g.s.u. respectively.) The infra-red spectrum of complex (V) in theregion 400-4000 cm-I (fig. 4) closely reembles that of the precursor complex (I) (fig. 1) :it would appear that the carbon suboxide ligand is not directly oxidized in the above process.0FIG. .I.--Infra-red absorption spectrum of silver complex 0, prepared by photo-oxidation ofcompIex (I) (fig.l), as a mull in Nujol.REVERSIBLE THERMAL OXIDATION OF SILVER COMPLEX (I)The apparatus consisted of a 5-ml glass bulb wrapped in aluminium foil, to exclude light,connected by a capillary to a three-way tap leading to a vacuum pump, oxygen source,and a mercury-filled manometer. The bulb was evacuated, filled with oxygen at atmosphericpressure and cooled from 293 to 195°K by immersion in a slurry of solid carbon dioxide inacetone. The decrease in volume of the oxygen in the system at equiIibrium was noted.Complex (I) (4.22 g) in a finely divided state was added to the bulb and the above procedurewas repeated. The results of several experiments showed that the complex (I) (4.22 g)absorbed 10.2 mlf0.5 ml (con-.) of oxygen in the dark at 195"K, and evolved all this onwarming to 293°K.The absorption cycle was repeated three times without any significantloss. The amount absorbed suggests that reaction was largely confined to the surface of thesolid complex.PYROLYSIS OF COMPLEXESA small amount (ca. 10 mg) of a silver complex was heated in a thick-walled J?yrex tubeunder argon at 0.5 mm pressure until (at ca. 200") explosive decomposition occurred. Th194 COMPLEXES OF CARBON SUBOXIDES WITH SILVERgaseous decomposition products were passed into a gas cell for infra-red analysis or into theionization chamber of a mass spectrometer (A.E.I. model MS 10). The major gaseousproduct was carbon suboxide. Tertiary amines strongly catalyze the polymerization ofC301, and unless a complex is completely freed from tertiary amine prior to pyrolysis, allthe suboxide produced polymerizes within 1 min.For preparative purposes, amounts ofseveral grams of the complex can be safely decomposed to carbon suboxide when mixedwith an excess of clean sand.RESULTS AND DISCUSSIONEVIDENCE FOR CARBON SUBOXIDE AS A LIGANDThermal decomposition of the silver complexes (I) and (11) of empirical formulaAg3C302 under argon at ca. 200"/0.5 mm gives carbon suboxide as the major gaseousproduct, together with traces of CO and CO, and a black diamagnetic solid (containingsilver) of variable composition (typically: C-4, Ag, 94 %). Synthesis of carbonsuboxide from other primary fragments of decomposition is improbable as the onlyother gaseous products are CO and COz, which almost certainly are derived fromthe suboxide.(Carbon suboxide in polymerizing above 250°C splits off COY andbelow 250°C, COz : the latter process is catalyzed by the solid products from thermaldecomposition of the complexes and can be followed by infra-red analysis within thegas-cell.) If follows that the suboxide was either present as a discrete ligand or wasformed within the decomposing complex from other carbon oxide ligands, e.g., C,Oand CO. Further evidence comes from the production of some malonic acid followingdecomposition of the complexes with 2 N sodium hydroxide. The presence of carbonsuboxide as a discrete ligand is suggested by the infra-red spectra of the complexes (fig.1-4), notably the absorption at ca.1980 cm-l. An intense absorption band at ca. 2000cm-l is characteristic of metal carbonyls (cf. 2143 cm-1 for free carbon monoxide) andit would not be surprising that metal-carbon suboxide complexes (hitherto unknown)should also absorb in this region. Carbon suboxide itself shows intense absorptionat 2258 ~ m - l , ~ and bonding of the suboxide involving back-donation of electronsfrom the metal should shift this band to a lower frequency, as in carbonyls. [Theco-ordination geometry must clearly be different from that in carbonyls.] The smallnumber of absorption bands and the absence of absorption in the region 650-1900cm-' suggests, structurally-simple ligand species containing neither C-C nor C-0single bonds.The formation under mild conditions of carbon suboxide, or indeed anyligand containing a C3 fragment from acetic anhydride and silver acetate, neither ofwhich contains a C3 chain, is remarkable. Catalysis of the formation of these com-plexes by bases and inhibition by acids suggests that a Perkin-type self-condensationof acetic anhydride is involved. Although there is evidence for base-catalyzed self-condensation of acetic anhydride, such condensation normally occurs only underharsh condition^.^ Neither silver acetate nor the silver complexes show any excep-tional activity as catalysts for the Perkin condensation of benzaldehyde with aceticanhydride, the catalytic activity of silver acetate being similar to that of sodium acetate.Although the formation of small amounts of free carbon suboxide in the initial stagesof uncatalyzed preparations of the silver complexes suggests that formation of thefree suboxide may be involved, this cannot be the case in tertiary-amine-catalyzedpreparations, although pyrolysis of the complexes obtained gives carbon suboxideas the major gaseous product : tertiary amines vigorously catalyze the polymerizationof free carbon ~uboxide.~ If follows that ligand carbon suboxide must be formed bya series of reactions which occur within the coordinaion sphere of the silver: thE. T.BLUES AND D. BRYCE-SMITH 195driving force for the production of the suboxide may be preferential affinity of silverfor the product from each successive stage.A primary product obtained in the uncatalyzed formation of silver complexesfrom silver acetate and acetic anhydride is a canary-yellow material which analysesclosely for silver acetate and which has an infra-red spectrum closely similar to thatof silver acetate in the region 400-4000 cm-l.This new material differs strikinglyfrom normal silver acetate in being insoluble in all common solvents, including0.880 ammonia and pyridine. This observation suggests that the first step in theformation of the silver carbon suboxide complex is rearrangement of silver acetateto a form in which the silver is partly occluded from solvating species probably, weinfer, through association into a cluster structure, but we do not exclude a coatingphenomenon. The succeeding stages in the formation of the ligand are as yet obscure.From evidence (discussed later) that part at least of the silver in the carbon suboxidecomplexes has an oxidation state less than + 1, it appears that oxidation as well ascondensation (inferred from the production of 2 mol.of acetic acid per mol. of silveracetate) may be involved.X-ray powder photographs have shown that the complexes (I) and (11) containsimilar poorly-defined crystalline phases ; but the great insolubility (polymericstructure?) has made it impossible to obtain single crystals.In further work with Mr. M. J. Simons, Mr. A. Marriage and Dr. H. Herz, acomplex formed in a pyridine-catalyzed reaction has been obtained in a well-definedcrystalline form. Preliminary X-ray studies have indicated a tetragonal unit cell, anda unit molecular weight of 254+2, consistent with a " silver-ketenide " structureAg2C20 (M = 256).This result appears inconsistent with the empirical formulaA&C302 indicated by elemental analysis, and is under investigation. It does,however, suggest that the seemingly minor infra-red spectroscopic differences foundbetween complexes prepared with and without pyridine catalysis correspond to majorstructural differences. Further, complexes prepared bypyridine catalysis are normallydiamagnetic, and do not react with bromide ion to give an ammonia-insoluble form ofsilver bromide.CONSTITUTION. EVIDENCE FOR SILVER ATOM CLUSTERSEvidence that the silver complexes (I) and (11) having the common empiricalformula Ag,C,02, contain carbon suboxide as a discrete ligand has been discussed.If carbon suboxide is the only ligand species, the empirical formula implies thateither the suboxide is present as a triply-charged radical-anion (which seems highlyunlikely) or that part at least of the silver is present as Ago or is contained in a clusterin which the metal atoms have a fractional oxidation state less than + 1.In complex (I)the presence of radical-anions or unassociated Ago is ruled out by the observeddiamagnetism. However, metal-metal bonding is common in complexes containingatoms with an S , configuration. Although Ago complexes have never been described,their potential existence is foreshadowed by the considerable dissociation energy(163 kJ/mole) of Ag,,6 and by the stability of Cuo complexes and of gold clustercomplexes in which the gold atoms have fractional oxidation states of 1 /6, 1 /5, or 1 /3.*The antiferromagnetism of complex (II), obtained from the diamagnetic complex(I) by treatment with dilute acetic acid, could in principle arise by an exchange inter-action between C302 radical anions or weakly associated Ago, Ag", or Ag"'.Thepresence of Ag" or Ag"' appears excluded by the absence of any oxidizing propertiesand strong colour. Complexes (I) and (11) react with aqueous potassium iodidewithout liberation of iodine. Location of the magnetic properties in the silver rathe196 COMPLEXES OF CARBON SUBOXIDES WITH SILVERthan the C302 ligand is indicated by ligand exchange and displacement reactionswhich preserve the magnetic properties.Thus, reaction of complex (11) with 4-Naqueous sodium chloride for 4 h at 98" gave a mixture (analyzing approximately asAg,C12) from which silver chloride was removed by aqueous ammonia to give a black,strongly-antiferromagnetic, electrically-conducting material (VII), essentially silvermetal (found : C, 1.2 ; Ag, 98.3 %). The material (VI), on heating in air at ca. 350",abruptly lost carbon to give diamagnetic silver metal (found : Ag, 99-8 %), but themagnetic properties were largely unchanged after brief heating to 500" in argon.The magnetic susceptibility of the silver (VI) indicates that it contains the whole ofthe structural elements responsible for the field-dependent antiferromagnetism of itsprecursor. The results further indicate that only one-third of the silver in complex(11) is responsible for its antiferromagnetism and that two-thirds is present as, ortransformed to, silver'.It is as yet uncertain whether the small amount of carbonin material (VI) is constitutional or not. Further evidence for location of the magneticproperties in the metal comes from conversion of complexes (I) and (11) into newforms of silver bromide and iodide.4 0I-l , , , , I > , , I I I I I 3 l . n 0 I00 2 00 300temperature (OK)FIG. 5.-Temperature- and field-dependence of reciprocal magnetic susceptibility for cluster silveriodide (VLI). (A), H = 1.75 ; (B), H t= 3-75 ; (C), H = 5.15 ; (D), H = 5.80 ; (E), H = 6.70 ; (F),H f= 7.10 kilogauss.Complex (II) reacts with 7 N aqueoas potassium iodide at 10" to give a blacksuspension which soon dissolves to give a clear light-yellow (not iodine) solution.Dilution of the filtered solution gives a yellow precipitate (VII) substantially AgI(found : Ag, 44.8 ; I, 53.9 ; calc.for AgI ; Ag, 45.9 ; I, 54-1 %). The iodide (VII)is antiferromagnetic : the magnetic susceptibility both for the solid and solutions inaqueous potassium iodide is field-dependent, showing the presence of a polynucleaE. T. BLUES AND D. BRYCE-SMITH 197species having more than one paramagnetic centre. Although for the solid, peffdecreases with T (fig. 6) and the plot of reciprocal susceptibility against T from SO-320°K (fig. 5) is indicative of antiferromagnetic exchange, X increases slightly andalmost linearly as the temperature is lowered and there is no well-defined Curie point.This type of anomalous magnetic behaviour, usually found in crystals of low symmetryand in amorphous systems, has previously been explained in terms of an antiferro-magnetic exchange within each of two lattices and a ferromagnetic exchange betweenthem : the phenomenon has been termed metamagneti~m.~IWI.0 1 - 1 0 5100 2 00 3 0 0temperature (OK)FIG.6.-Temperature- and field-dependence of p+ff for silver in cluster silver iodide (W). (A), H =1.75 ; (B), H = 3.75 ; (C), H = 5.15 ; (D), H = 5.80; (E), H = 6.70 ; (F), H = 7-10 kilogauss.6 0The X-ray diffraction pattern of (VII) is apparently identical with that of normalsilver iodide containing a preponderance of the hexagonal form with a small amountof the cubic modification.The thermogram of (VII) from 0-450°C shows a singleendotherm at 147" corresponding to the hexagonal+cubic phase of silver iodide.However, the endotherm is ca. 30 % Iess than that shown by ordinary silver iodide,even on repeated recycling through the phase-change, and the accompanying colourchange is much less marked. Although freshly-prepared iodide (VII) is lighter incolour than ordinary silver iodide, it darkens in light much more rapidly. Thisdarkening is accompanied by a decrease in the magnetic susceptibility. Thesefindings suggest that iodide (VII) contains ca. 30 % of an amorphous, antiferro-magnetic form of approximate composition AgT.The new form is thermally stableup to at least 500°C.Treatment of complex (11) with 2.5 N potassium bromide at 98" for 4 h gives abrown precipitate analyzing approximately for Ag, oBr9. Treatment of the precipitatewith an excess of 0.880 ammonia dissolves out " ordinary " silver bromide leavingan antiferromagnetic material, analyzing forAg,Br,, which appears to be a mixture ofnovel antiferromagnetic forms of silver bromide and silver metal. The antiferro-magnetic silver bromide is unstable in contact with 0.880 ammonia and withinseveral hours gives ordinary diamagnetic silver bromide and a weakly-antiferro-magnetic black residue, substantially silver (found : Ag, 99.2 %).How then does the antiferromagnetic complex (11) differ from its diamagneticprecursor, complex (I)? The close similarity in their empirical formulae and theirinfra-red spectra in the region 400-4000 cm-l (fig.1) indicates that they contain thesame ligand and that the bonding of the ligand is essentially the same in both com-plexes. Significant differences are found in the far infra-red spectra within th198 COMPLEXES OF CARBON SUBOXIDES WITH SILVERregion 40-400 cm-1 : these might reflect differences in metal-metal bonding (fig. 7).X-ray powder photographs show that they contain similar but poorly-defined crystal-line phases. A tentative hypothesis in accord with the observations is that bothcomplexes (and their derivatives) contain silver clusters, but that the diamagneticclusters are more closely packed and/or ordered.The antiferromagnetic propertiesmay result from irregularities within individual clusters leading to imperfect com-pensation of electron spins. The cluster sizes are as yet unknown, but the magneticproperties suggest elements of lattice structure within clusters, consistent with largedegrees of association. Complexes (I) and (11) are graphite-like in their softness, sothe clustering could conceivably be of a type involving sheets of silver atoms.0'0 -1? I 140 100 2 0 0 3 0 0 4001.0 ' " " " ' I ' "Cm-lFIG. 7.-Far infra-red absorption spectra of wax discs of silver complexes (I), (-), and (11), (---).Chemical reactions of complexes (I) and (11) have tended to give about two-thirdsof the silver as ordinary Ag+ derivatives. Hence the complexes could be provisionallyformulated as Agz:+(AgC30,)2:-, where n is the degree of association of silver withinthe cluster structure ; but the weak affinity for amine and phosphine donors, and thepoor development of crystallinity, might suggest some greater interaction betweencationic and anionic silver than this simplified formulation would imply.We thank Mr. R. W. Matthews of this Department for the magnetic measurementsquoted in fig. 5 and 6, Dr. S. S. Adcock and Mr. S. Warren for some preliminarystudies, and Ethyl Corporation for financial support.H. Krepelka and V. Vebersek, Chem. Listy, 1949,43,25.A. R. Blake and A. F. Hyde, Trans. F a r h y SOC., 1964,691,1775.D. A. Long, F. S. Murfin and R. L. Williams, Proc. Roy. Soc. A, 1954,223,251.G. A. Ellestad, R. H. Evans and M. Kuntsmann, Chem. Comm., 1967,1069.C. Hagar, Beitrag zur Chemie des Kohlensuboxides (Zurich, 1961).M. Ackerman, F. E. Stafford and T. Drowart, J. Chem. Phys., 1960,33, 1784. ' (a) A. J. Layton, R. S. Nyholm, G. A. Pneumaticakis and M. L. Tobe, Nature, 1967,214,1111 ;(b) R. Nast, P. Lirst. G. Beck and J. Gremm, Chem. Ber., 1963,96,3302; (c) G. W. Watt andJ. W. Dawes, J. Inorg. Nuclear Chem., 1960,14, 32. * F. Cariati, L. Malatesta, L. Naldini and G. Simonetta, Chem. Comm., 1965, 212 ; 1966, 647 ;Inorg. chim. Acta, 1967, 24, 315 ; M. McPartlin, R. Mason and L. Malatesta, Chem. Comm.,1969,339.J. Becquerel and J. van den Handel, J. Phys. Radium, 1939,10, 10

ISSN:0366-9033    

DOI:10.1039/DF9694700190

出版商:RSC    

年代:1969

数据来源: RSC

摘要:

GENERAL DISCUSSIONProf. N, N. Greenwood (Newcastle upon Tyne) said: I would ask Perkins to whatextent is d orbital participation considered to be essential in stabilizing dimericAl,Me, on this model, and can the absence of such participation in BMe, be takenas an important factor in determining its inability to form a stable dimer?Dr. I. H. Hillier (University of Manchester) said : Could Perkins comment on thereliability of the CNDO method applied to the molecules considered in his paper, andon the 3d orbital exponent which should be used in such calculations ?Prof. P. G, Perkins (University of Strathclyde) said: In reply to Greenwood, wehave not yet analyzed the corresponding results for organoboron compounds butbelieve that the lack of participation of the 3d orbitals may well be the factor whichmilitates against formation of compounds such as B,Me,.In reply to Mrs.Green, one should indeed include further bases such as the 4s and4p atomic orbitals. However, the whole calculation is subject to a variational pro-cedure and it is the extent to which a set of new functions is likely to contribute whichdetermines their choice.In reply to Hillier and Urch, the extent to which the d orbitals are included isdetermined by the orbital exponent Z* and this should be determined during thecalculation by minimizing the energy with respect to it. However, such a procedurewill be very difficult in the CNDO framework and the present work therefore, as afirst approximation, takes two fixed cases. Each compound should be treated onits merits.In reply to Hill, the orbital exponent Z* which emerges from a variational calcula-tion need not have any precise physical meaning because it may relate to non-integralnumbers of electrons.This is true for simple systems (e.g., He atom) and so Z* isbest regarded as simply a variable parameter defining the 3d radial function.In reply to Hillier, in order to try and assess the reliability of the CNDO methodwe have made some comparison between it and an extended basis set ab initio calcula-tion on BH2 NH2. We find that the charge distribution is well reproduced by theCNDO method, and the eigenvalues are in satisfactory agreement. We believe atpresent that the latter method should be reasonably reliable for differential calculationswhere gross structural changes are involved.Prof.Theodore L. Brown (University of Illinois, Urbana) said: L. M. Ludwick inour laboratories has recently observed that the presumed 5-coordinate methylcobaltbis-dimethylglyoxime is in fact a dimer in methylene chloride. Dimerization probablyoccurs through the interaction of an oxygen from one molecule with the axial positionof the cobalt atom in a second species. Association in this manner is not uncommon.We have also been examining the kinetics of exchange reactions involvingreplacement of the axial ligand in the methylcobalt bisdimethylglyoxime system.Our evidence is not as yet extensive, but there is a suggestion in our results that theability of the axial ligand to act as a pi acceptor towards cobalt may be an importantfactor in coordination at this position.For example, triethylphosphite, whenbound in this axial position, does not exchange with excess ligand in methylenechloride, as evidence in the n.m.r. spectrum, at temperatures as high as +60°C.19200 GENERAL DISCUSSIONBy contrast, nitrogen and oxygen donors typically exchange much more rapidly,with coalescence temperatures of - 20 to - 50°C.Dr. H. A. 0. Hill (Oxford University) said: The suggestion made in the paper byDuncan et al. that the field metal d-orbitals interact directly with the 1s-orbitals ofthe methyl hydrogens may be most important. What effect would an increasingelectron density on a neighbouring metal atom acting directly on the methyl hydrogenshave on the JH--H coupling constants?Prof.Theodore I;. Brown (University of Illinois, Urbana) said: We have recentlyobserved the 35Cl quadrupole resonance spectrum of Zeise's salt. Absorptionswere observed, at 25"C, at 15.9, 20.10 and 20.37 MHz. The latter two absorptionsare ascribed to the cis chlorines, which are crystallographically distinct. The lowerfrequency absorption is assigned to the trans chlorine.The difference in quadrupole resonance frequencies of the two types of chlorineis substantial, and suggests that ethylene does indeed exert a trans influence insofaras this particular spectroscopic property is concerned. Chlorine quadrupoleresonance frequencies can generally be discussed satisfactorily in terms of p orbitalpopulations.The quadrupole resonance frequency is a measure of the differencein populations of the p,, orbital and the p x orbitals. In the present case, assumingthat the asymmetry parameter is small, the lower frequency for the trans chlorinecan be explained in terms of a more ionic Pt-Cl bond to the trans chlorine, as aresult of a strong c polarization by ethylene. There may, however, be an additionalfactor which would decrease the population difference in the p,, and pTc orbitals,namely, a n-donor action on the part of the trans chlorine toward platinum, resultingfrom the x-acceptor behaviour of ethylene. I suspect that this is a small effect, andthat most of the frequency difference must be ascribed to polarization in the bondsystem. The results are consistent with the observation of Denning regarding thePt-C1 stretching frequencies in Zeise's salt.Dr.M. G. Clark (University of Cambridge) said: An approach of the type usedby Green, Smith, and Tasker may be employed to study the stability of a 4-coordinatedsquare-planar system against the addition of two-electron donating ligands in eitheror both of the axial positions. The increase in coordination number is assumed tooccur without change in spin-state or equatorial bond length.Consider the effect of adding one or two axial ligands for each realistic orderingof the valence orbitals in a square-planar complex. We may work with the pointgroup C4u appropriate to the addition of one ligand, since the group D4h for theaddition of two is related to C4" by the addition of a centre of symmetry. Theaddition of either one or two ligands may be treated simultaneously, since in thelatter case only the centro-symmetric axial-ligand symmetry-orbital has correctsymmetry to interact with the orbitals of interest in the complex.The generalfeatures of the appropriate energy level diagrams are sketched in fig. 1, using CaUnotation. The diagrams may be considered as fragments of more complicatedmolecular-orbital diagrams, but since we require only qualitative features of theenergy levels they are adequate for our purpose, provided that it is borne in mind thatthey do not show all the interactions between levels explicitly. The only assumptionmade is that in the enlarged complex the antibonding a, orbital is more antibondingthan the bonding a, orbital is bonding.For a complex of a metal ion having configuration d" there are n+2 electronsto be assigned to the levels shown in the appropriate diagram. Table 1 shows foGENERAL DISCUSSION 201b*ale\\?\\ / - \ -FIG.1.TABLE 1ion n i c e e<ullow-spin d6 *high-spin d6 Ulow-spin d7 * *high-spin d7 Ulow-spin d8 S Shigh-spin d8 * .Ed9 s ss = stable against addition of ligands ; u = unstable against addition of ligands ; * = may or mayak**not be unstable depending on precise energetics.6<n<9 the results of considering whether or not the initially square-planar systemprofits energetically by taking on extra ligands under our assumptions. Not sur-prisingly, the results are indecisive in several cases ; but two definite points do emerge :(1) d9 and low-spin d8 ions are the obvious candidates for forming stable 4-coordinatedsquare-planar complexes, in agreement with observation ; (2) any truly 4coordinatedsquare-planar complexes of high-spin d6 and d7 ions should have al<e.Therarity of suitable examples makes the second prediction difficult to test; certainlythe mineral gillespite (BaFeSi,O, o), in which high-spin Fe2+ is square-planar co-ordinated by four oxygens, has a, <e.lThe results obtained above should be treated with considerable caution, since,apart from the simplicity of the arguments used, it is not obvious that other factorsR. G. Burns, M. G. Clark and A. J. Stone, Inorg.Chem., 1966,5,1268 ; M. G. Clark and R. G.Burns, J. Chem. SOC. A , 1967, 1034; M. G. Clark, G. M. Bancroft and A. J. Stone, J. Chem.Phys., 1967, 47, 4250202 GENERAL DISCUSSIONare negligible. For example, the iron in gillespite may have its coordination geometryto some extent forced upon it by the layer structure of the silicate framework.Dr. H. A. 0. Hill (Oxford University) said: The apparent reluctance of cobaltamben derivatives to give rise to alkyl cobalt(II1) derivatives may be caused by reasonsother than those mentioned in the paper by Greene et al. Except for ligand (V), onlyin the amben complexes will the nature of the coordinated nitrogen interfere with theplanarity of ligand when complexed. The steric requirements of the NH hydrogensmay cause a distortion of the complex such that in the cobalt(I1) derivative the struc-ture is intermediate between square-planar and tetrahedral.Even if formed, analkyl cobalt(l1l) derivative may be kinetically unstable. Though n-propylcobalaminis stable with respect to decomposition in solution, iso-propylcobalamin decomposesby homolytic fission of the cobalt-alkyl bond. This instability is presumably causedby steric interactions of the isopropyl ligand with the ring. Therefore it is possiblethat the alkyl cobalt amben complexes are unstable, or difficult to isolate, because ofsteric interaction of the alyl ligand with the non-planar amben ligand.Dr. Michael Green (University of York) said : In reply to Hill's remark about theplanarity of cobalt amben, conjugated '' N,-type " ligands are remarkable for theirtendency to form planar complexes even when geometric factors dictate otherwise.Weber has found that complexes of IX are planar not only when R is -(CH&--and -(CH2)3- but also when it is -(CH2),- in which case a tetrahedral configura-tion would be preferred stericalIy.We have found no significant variation either inthe visible/ultra-violet spectra or the magnetic properties of the cobalt ambenfamily (VIII), when R is -4CH2)2-, --(CH2)3 or --(CH&--. If a modification ingeometry occurred, the first complex would be expected to be planar and the lasttetrahedral. (The corresponding compounds from the cobalt salen family (VII) showdifferences in spectroscopic and magnetic properties compatible with such a modi-fication).Moreover, there is no notable difference between cobalt amben (VIII,R = -CH,CH,-), cobalt ambphen (VIII, R = 1,2-phenylene) and the cobaltomp pound,^ (X). It is difficult to see how these last two compounds could be dis-torted tetrahedrally.In view of this strong tendency for " N4 " systems to be planar and in view of theirgreater field strength compared with " Nz02 " ligands, any distortion in an alkylatedamben compound would be small and be most likely to involve the cobalt atom movingslightly out of the plane of the ligand which would remain flat as for CH,CoBAE whichHill mentions.Dr. H. A. 0. Hill (Oxford University) said: We have measured the e.p.r. spectraof the mono- andlor dipyridine complexes of the cobalt(I1) derivatives of ligandsJ.H. Weber, Inorg. Chem., 1967, 6,258.M. Hariharan and F. L. Urbach, Imrg. Chem., 1969,8, 556 and references therein.M. Green and P. A. Tasker, Chem. Cum., 1968,518GENERAL DISCUSSION 203I,l 11,2 131,l IV,3 VI,2 and VIII,2 in the paper by Greene et al. and in every case thespectra are interpretable in terms of a low-spin d7 complex in which the unpairedelectron is well-described by the d,Z orbital corresponding to position C and D in fig. 1.It would be of interest to examine the e.p.r. spectra of cobalt(I1) amben derivatives.Dr. Michael Green (University of York) (communicated) : Hill's data for (VII) areinteresting as they imply a reversal in the order of the dz2 and dx,, orbitals in going fromnickel salen to the cobalt salen/pyridine system.The presence of an axial ligandwill tend to cause changes A+B+C-+D, as we mention in our paper. As cobaltamben is reluctant to add pyridine as a ligand it will have a greater chance of being ofA-type than cobalt salen in pyridine. I agree the e.p.r. spectrum will be interesting.Dr. H. A. 0. Hill (Oxford University) said: A recent determination of the structureof CH3CoBAE, which has just been communicated to us prior to publication byProf. Rondaccio, University of Trieste, Italy,4 shows that the cobalt is indeed j u e -coordinate, with the BAE ligand planar and the cobalt lying 0.16 A above the plane ofthe BAE ligand. Thus, the suggestion made in our paper that in the five-coordintealkyl Co(II1) complexes, the cobalt atom lies above the plane of the ligands is confirmedin this case at least.Such a distribution may have important consequences forthe enzymatic role of the five-coordinate complexes. The importance may lie, notin the vacant coordination site trans- to the alkyl ligand but rather, that in the five-coordinate complex, the cobalt and its coordinated alkyl ligand lie out of the planeof the corrin ligand, making both the cobalt and the attached carbon atom moreaccessible for reaction with the substrate.Dr. J. F. Ogilvie and Dr. M. J. Newlands (Memorial University of Newfoundland,Canada) said : One of the most important applications of the theory of bonding inmetallo-organic compounds must be to elucidate the binding to metals of reactants,molecular intermediates and products in reactions which are catalyzed by materialscontaining metals, especially transition metals.A process of some current interestis the fixation of nitrogen in nature, in which molybdenum, perhaps cobalt, and parti-cularly iron are believed to be invol~ed.~ Inorganic nitrogen complexes of iron havealready been prepared.6* A commonly postulated intermediate in the ammonia-producing process is di-imide.5 However, di-imide, HN=NH, has been shown toinhibit fixation of nitrogen by, for instance, Clostridium pasteurianum.In connection with our consideration of multiple bonding in silicon and germaniumimine derivatives * we have been led to suggest that an imine intermediate, M=NHHN-M, rather than M . .. HN-NH . . . M, containing nitrogen atoms stronglybound to the appropriate metal atom M held rigidly in some location by some co-ordinating molecule such as a porphin type of chelating agent. An arylimino complexof rhenium has been prepared which has a relatively short lo strong bond betweennitrogen and rhenium atoms, indicating that such a strongly covalent linkage isindeed possible for transition metals.Cockle, Hill, Pratt and Williams, Biochern. Biophys. Acta, 1969, 177, 686.Hill, Morallee and Pellizer, to be published.Hill, MacFarlane and Williams to be published.S. Brucker, M. Calligaris, G. Nardin and L. Randaccio, to be published.R. Murray and D. C. Smith, Courd. Chem. Rev., 1968,3,429.A. Sacco and M. Aresta, Chem. Comm., 1968, 1223.' G. M. Bancroft, M. J. Mays and B. E. Prater, this discussion. * J. F. Ogilvie and M. J. Newlands, Trans. Farday SOC., in press.J. Chatt, J. D., Garforth, N. P. Johnson and G. A. Rowe, J. Chem. SOC., 1964, 1012.lo D. Bright and J. A. Ibers, Inorg. Chern. 1968,1,1099204 GENERAL DISCUSSIONProf. N. N. Greenwood (Newcastle upon Tyne) said: With regard to Bryce Smith’spaper, what is the origin of the antiferromagnetism in the dl0 compounds formulatedas Ag*Br and Ag*I? If there are, indeed, small domains of metal-metal clusterswould it not be expected that the compounds would be superparamagnetic ratherthan antiferromagnetic since the volume exchange interaction energy is known toapproach thermal energies for particles smaller than - 100 A and so magnetic orderingcannot occur ?Prof.D. Bryce-Smith (University of Reading) said. In reply to Greenwood, Ishould state fist of all that we have used the term “ antiferrornagnetism ” for con-venience because it presently seems to provide the closest description of the observedmagnetic behaviour; but we accept that the behaviour is not classically antiferro-magnetic. Since Ag‘ is isoelectronic with PdO, we agree that Agl clusters might atfirst sight exhibit paramagnetic properties similar to those of palladium metal. Thedifferent behaviour observed might be explained in several ways, one of which is thatwe could be dealing with sublattices of Ago and Ag”.Small domains of metal clusters would lead to superparamagnetism only if thesewere individually ferro- or ferrimagnetic.Antiferromagnetic domains would notresult in superparamagnetism. Tt is not certain whether metamagnetic domains,i.e., domains in which both ferro- and antiferromagnetic interactions occur, wouldresult in a form of superparamagnetism or what the bulk magnetic properties of asubstance containing such domains would be. Certainly, the forms of the curvesshown in fig. 5 and 6 are inconsistent wih normal superparamagnetic behaviour.Dr. M. G. Clark (University of Cambridge) said: I would ask Bryce-Smith if it ispossible that the new complexes may have polymeric or macromolecular structureswith the carbon suboxide, bromide, or iodide acting as bridging ligands. This wouldhelp to explain their solid-state magnetic properties and lack of crystalline form,but difficulties might be met with their solution properties and the new form of silver“Ag* ”-Prof. D. Bryce-Smith (University of Reading) said. In reply to Clark, the greatinsolubility of the silver complexes is certainly consistent with polymeric structures.The magnetic properties of the suboxide complexes, in so far as a comparison can bemade, appear to parallel those of the derived iodides, a fact which suggests that theseproperties arise from a common structural feature involving the silver atoms ratherthan the ligands. For this reason, we believe that super exchange does not contributein a major way to the magnetic properties, although we do not exclude the possibilitythat C302 units can chemically link cluster centres

ISSN:0366-9033    

DOI:10.1039/DF9694700199

出版商:RSC    

年代:1969

数据来源: RSC

摘要:

AUTHOR INDEX *Anderson, W. P., 37.Bancroft, G. M., 136, 146, 148Blues, D. T., 190.Bor, G., 64, 65, 68.Braterman, P. S., 121.Brown, T. L., 37, 199,200.Bryce-Smith, D., 190, 204Buttery, H. J., 48.Campbell, C. H., 107.Canadine, R. M., 27, 62.Chambers, D. B., 157.Coates, G. E., 157.Cotton, F. A., 79, 108.Clark, M. G., 144, 145, 147, 200, 204.Dahl, L. F., 93.Dakkouri, M., 149.Duncan, J. D., 178.Evans, S., 112.Gamlen, G. A., 7.Glocking, F., 157.Greatrex, R., 126.Green, J. C., 112, 178.Green, M., 110, 172, 203.Green, M. L. H., 112, 178,202.Greenwood, N. N., 68, 126, 146, 147, 148, 199,Haase, J., 149.Hamnett, A., 107.Hill, H. A. O., 110, 148, 165, 200, 202, 203.Hillier, I. H., 27, 62, 199.Ibers, J. A., 84, 110.Keeling, G., 48.Kettle, S. F. A., 48.204.Levison, K. A., 183.McLauchlan, K. A., 178.McWeeny, R., 20.Manojlovid-Muir, L., 84.Mason, R., 20, 59.Mays, M. J., 136, 145.Merk, W., 71.Miller, J. R., 63, 65.Muir, K. W., 84.Murrell, J. N., 59, 62.Newlands, M. J., 203.Ogilvie, J. F., 203.Orchard, A. F., 60, 112, 144.Paul, I., 48.Perkins, P. G., 183, 199.Pettit, R., 71, 107.Prater, B. E., 136.Pratt, J. M., 165.Reddy, M. L. N., 53.Rest, A. J., 69, 146.Smith, J., 172.Stamper, P. J., 48.Strouse, C. E., 93.Sugahara, H., 71.Tasker, P. A., 172.Towl, A. D. C., 20.Turner, D. W., 112.Urch, D. S., 53,63,64,66,68,69,107,110.Wagstaff, K. F., 108.Walker, A. P., 121.Williams, R. J. P., 165.Wristers, J., 71.Zeil, W., 149.* The references in heavy type indicate papers submitted for discussion

ISSN:0366-9033    

DOI:10.1039/DF9694700205

出版商:RSC    

年代:1969

数据来源: RSC