Yves Jean - Molecular Orbitals of Transition Metal Complexes (793957), страница 4
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Finally, the metal can bind justto the radical centre (X-type ligand), giving an η1 -C5 H5 coordination(1-9). In this latter case, one can no longer describe it as a π complex,since the metal centre interacts with only one of the ring carbon atoms,with which it forms a σ bond.This diversity of coordination behaviour is also found for otherconjugated polyenes. Thus, butadiene can act as an L2 ligand if theelectrons of the two π bonds are involved (η4 -butadiene, 1-10) or as anL ligand involving a single π bond (η2 -butadiene, 1-11).In the same way, benzene can bind in the η6 (L3 ligand, 1-12), η4 (L2ligand, 1-13), or η2 modes (L ligand, 1-14) (see Exercise 1.5).
In the η4 andη2 coordination modes, the six carbon atoms become non-equivalentM2)1-11 (η2 , L)MMM1-12 (η6 , L3 )1-13 (η4 , L2 )1-14 (η2 , L)Electron count in a complex: the covalent model(four or two, respectively, are bound to the metal), perturbing theπ-electron conjugation.
As a result, the ring becomes non-planar.A slightly different case arises for ligands which can bind to a metalcentre through several different sites without any conjugation of theelectrons involved. These ligands are said to be polydentate (bidentate,tridentate, . . . ), in contrast to monodentate ligands such as PR 3 , CR 3 ,etc. For example, 1,2-bis(dimethylephosphino)ethane is a bidentate ligand, since it can bind through its phosphino sites (1-15). As each ofthese has a lone pair, it behaves as an L2 ligand towards the metal.
1,2dioxyethane (O-CH2 -CH2 -O) is also a bidentate ligand (1-16), but eachoxygen atom supplies only one electron to the metal (an X2 ligand).H2 CCH2Me2 PPM e 2M1-15 (L2 )4This property is readily explained bymolecular orbital theory: two electrons must∗ orbitals. Thebe placed in two degenerate πoomost favourable arrangement contains oneelectron in each orbital, with their spinsparallel (a triplet state).CCR2R11-17a (L)R2R11-17b (X2 )H2 CCH2OOM1-16 (X2 )We end this section by discussing several ligands whose classificationas L- or X-type can create difficulties.The usual Lewis structure for the dioxygen molecule, O2 , showsa double bond and two lone pairs on each oxygen atom. One mighttherefore conclude that it is an L-type ligand, and that the coordinationwould be either η1 (through a lone pair) or η2 (involving the π bond).However, this Lewis structure, in which all the elctrons are paired, isnot satisfactory since the magnetic moment measured experimentallyshows that there are two unpaired electrons with parallel spin (the groundstate is a triplet).4 This is why O2 behaves as an X2 ligand rather than anL ligand.Carbene ligands, CR 1 R2 , provide another example.
These speciescontain two electrons on the carbon atom that do not participate in theformation of the C-R1 and C-R2 bonds. Depending on the nature of theR1 and R2 atoms or groups, the ground state is either diamagnetic, inwhich case the two nonbonding electrons are paired, forming a lonepair on the carbon atom (1-17a), or paramagnetic, in which case the twoelectrons are unpaired, giving a triplet state 1-17b).
In the first case, it islogical to consider the carbene as an L-type ligand, whereas it is an X2ligand in the second case.These two ways of describing a CR 1 R2 ligand are indeed used, andthis distinction is at the origin of the two families of carbene complexesin organometallic chemistry: the Fischer-type (L) and the Schrock-type(X2 ) carbenes. We shall return to this difference and offer an orbitalinterpretation (Chapter 4, § 4.3).Setting the scene1.1.1.4.
Bridging ligandsClClMMX ligandL ligand1-18OX ligandOX ligandMM1-19In bimetallic complexes, some ligands can be ‘bridging’, that is, boundsimultaneously to the two metal centres. These cases are indicated by thenomenclature µ. If one considers a bridging chlorine atom (M2 (µ-Cl),1-18), it behaves as an X ligand towards the first metal centre, thanks to itsunpaired electron, but as an L ligand towards the second, thanks to oneof its lone pairs (the roles of the two metal centres can, of course, beinterchanged). Overall, the chlorine atom is therefore an LX ligand; itsupplies three electrons to the pair of metal centres. Other ligands inwhich an atom has an unpaired electron and at least one lone pair, suchas OR, SR, NR 2 , PR 2 , etc. are analogous.
A bridging oxygen atom is anX-type ligand towards each of the two metal centres, since it has twounpaired electrons (1-19), so it therefore acts as an X2 ligand overall.1.1.2. Electron count and the 18-electron ruleOnce the nature of the ligands has been established, the second stageof our analysis of the electronic structure of transition metal complexeswill require us to count the number of electrons around the metal andthen to assign them, in a formal way, either to the metal or to the ligands.In what follows, we shall consider complexes written as [MLℓ Xx ]q , inwhich the metal M is bound to ℓ ligands of L type and to x ligands ofX type, the overall charge being q.1.1.2.1.
Total number of electrons, the 18-electron rule5There are two other transition series thatcorrespond to the filling of the 4f(lanthanides) and 5f (actinides) sub-shells.Each ligand L supplies two electrons to the metal’s environment, whileeach ligand X supplies a single electron. The total number of electronssupplied by the ligands is therefore equal to 2ℓ + x.
Only the valenceelectrons are considered for the transition metal, as, following the spiritof Lewis theory, we assume that core electrons play a negligible role.In what follows, we shall limit our analysis to transition elements corresponding to the progressive filling of the 3d, 4d, and 5d sub-shells(the transition metals of the d block, see Table 1.1). The valenceelectron configuration of these elements is of the type nda (n + 1)sb ,where n equals 3, 4, or 5, for the first, second, or third transition series,respectively.5 The metal therefore supplies (a + b) electrons.
We notethat some authors do not consider zinc to be a transition element, as itsd sub-shell is full (valence-electron configuration 3d10 4s2 ). This remarkalso applies to cadmium (Cd, 4d10 5s2 ) and to mercury (Hg, 5d10 6s2 ).When we take account of the overall charge q of the complex, thetotal number of valence electrons, Nt , is:Nt = m + 2ℓ + x − q(1.1)Electron count in a complex: the covalent modelTable 1.1. Electron configuration and number of valence electrons, m, for the d-block transition metals1st seriesSc3d1 4s2Ti3d2 4s2V3d3 4s2Cr3d5 4s1Mn3d5 4s2Fe3d6 4s2Co3d7 4s2Ni3d8 4s2Cu3d10 4s1Zn3d10 4s22nd seriesY4d1 5s2Zr4d2 5s2Nb4d4 5s1Mo4d5 5s1Tc4d5 5s2Ru4d7 5s1Rh4d8 5s1Pd4d10 5s0Ag4d10 5s1Cd4d10 5s23rd seriesLu5d1 6s2Hf5d2 6s2Ta5d3 6s2W5d4 6s2Re5d5 6s2Os5d6 6s2Ir5d7 6s2Pt5d9 6s1Au5d10 6s1Hg5d10 6s2m3456789101112Several examples of the application of this rule are given below:Complexm2ℓxqNt[Fe(CO)5 ][Ir(CO)(Cl)(PPh3 )2 ][Mn(CO)6 ]+[Ni(CN)5 ]3−[Zn(Cl)4 ]2−[V(Cl)4 ][Cr(CO)3 (η6 -C6 H6 )][Fe(η5 -C5 H5 )2 ][Cu(η5 -C5 H5 )(PMe3 )][Zr(η5 -C5 H5 )2 (CH3 )]+[Ti(PR 3 )2 (Cl)3 (CH3 )][W(PR 3 )2 (CO)3 (η2 -H2 )][Ir(PR 3 )2 (Cl)(H)2 ][Ni(H2 O)6 ]2+89710125681144691010612000128684124120105440213403000+1−3−20000+1000+2181618181891818181412181620By analogy with the octet rule, it has been proposed that a transitionmetal tends to be surrounded by the number of valence electrons equalto that of the following rare gas (electron configuration nd10 (n + 1)s2(n + 1)p6 ).
One thereby obtains the 18-electron rule, for which we shallprovide a first theoretical justification in this chapter (§ 1.6.3). However,in light of the examples given above, one must note that there are manyexceptions to this rule; we shall analyse them in greater detail in thefollowing chapters.1.1.2.2.
Oxidation stateIn order to determine the oxidation state of the metal in the complex,one performs a fictitious dissociation of all the ligands, supposing thatSetting the sceneTable 1.2. The Allred–Rochow electronegativity scale: (a) for thetransition metals and (b) for the light elements(a)Sc1.20Y1.11Lu1.14Ti1.32Zr1.22Hf1.23V1.45Nb1.23Ta1.33Cr1.56Mo1.30W1.40Mn1.60Tc1.36Re1.46Fe1.64Ru1.42Os1.52Co1.70Rh1.45Ir1.55(b)H2.2Li1.0Na1.0Be1.5Mg1.2B2.0Al1.5C2.5Si1.7N3.1P2.1O3.5S2.4F4.1Cl2.8Ni1.75Pd1.35Pt1.44Cu1.75Ag1.42Au1.41Zn1.66Cd1.46Hg1.44each of them, either L or X, takes with it the electron pair that createdthe metal–ligand bond. The remaining charge on the metal after thisdecomposition is the oxidation state of the metal in the complex.
Thisdistribution of the electrons, which ‘assigns’ the bond pair to the ligand,can be partially justified when one notes that this latter is usually amore electronegative entity than is the transition metal (see Table 1.2,the Allred–Rochow electronegativity scale). The metal–ligand bondsare therefore polarized, and the electron pair is more strongly localizedon the ligand than on the metal. To assign the two electrons of the bondjust to the ligand is, however, a formal distribution, which exaggeratesthe tendency linked to the difference in electronegativity.In the fictitious dissociation that we are considering, a ligand L leaveswith the two electrons that it had supplied, so the number of electronson the metal is not changed in any way. However, an X-type ligand,which had supplied only a single electron to make the bond, leaves in itsanionic form X− , carrying the two electrons from the bond with it.
Ittherefore ‘removes’ an electron from the metal, that is, it oxidizes it byone unit. The result of this dissociation is therefore written:[MLℓ Xx ]q → ℓL + xX− + M(x+q)(1.2)The oxidation state (no) of the metal in the complex is therefore equalto the algebraic sum of the number of X-type ligands and the charge onthe complex:no = x + q(1.3)Electron count in a complex: the covalent modelIn a widely used notation to specify the oxidation state of the metal ina complex, the chemical symbol of the metal is followed by the oxidationstate written in Roman letters (Mn(I), Fe(II), Cr(III), etc.).ExamplesComplexxqnoOxidation state[Fe(CO)5 ][Ir(CO)(Cl)(PPh3 )2 ][Mn(CO)6 ]+[Ni(CN)5 ]3−[Zn(Cl)4 ]2−[V(Cl)4 ][Cr(CO)3 (η6 -C6 H6 )][Fe(η5 -C5 H5 )2 ][Cu(η5 -C5 H5 )(PMe3 )][Zr(η5 -C5 H5 )2 (CH3 )]+[Ti(PR3 )2 (Cl)3 (CH3 )][W(PR3 )2 (CO)3 (η2 -H2 )][Ir(PR3 )2 (Cl)(H)2 ][Ni(H2 O)6 ]2+0105440213403000+1−3−20000+1000+201122402144032Fe(0)Ir(I)Mn(I)Ni(II)Zn(II)V(IV)Cr(0)Fe(II)Cu(I)Zr(IV)Ti(IV)W(0)Ir(III)Ni(II)In bimetallic complexes, the oxidation state is calculated bysupposing that the metal–metal bond(s), if any, is/are broken homolytically.
This procedure is justified by the fact that the electronegativities ofthe two metal centres are equal if they are identical, or similar in heteronuclear complexes (see Table 1.2(a)). The presence of one or more bondsbetween the metals therefore has no effect on their oxidation state. Forexample, the complex [Mo(Cl)2 (PR3 )2 ]2 can initially be considered to bedecomposed into two monometallic neutral fragments [Mo(Cl)2 (PR3 )2 ]in which the oxidation state of molybdenum is +2.In conclusion, we note that the oxidation state must not be equatedto the real charge on the metal in the complex, as it is obtained from aformal distribution of the electrons between the metal and the ligands.1.1.2.3.
dn Configuration of a metalThe oxidation state of the metal, which supplies m valence electrons, isequal to no after complex formation. The formal number of electronsremaining on the metal, n, is therefore given by the relationship:n = m − no(1.4)Setting the sceneWe are considering here n electrons which are not involved inthe formation of metal–ligand bonds, in other words ‘nonbonding’electrons. The electron configuration of the metal in the complex isrepresented as dn .ExamplesComplexnomConfiguration[Fe(CO)5 ][Ir(CO)(Cl)(PPh3 )2 ][Mn(CO)6 ]+[Ni(CN)5 ]3−[Zn(Cl)4 ]2−[V(Cl)4 ][Cr(CO)3 (η6 -C6 H6 )][Fe(η5 -C5 H5 )2 ][Cu(η5 -C5 H5 )(PMe3 )][Zr(η5 -C5 H5 )2 (CH3 )]+[Ti(PR3 )2 (Cl)3 (CH3 )][W(PR3 )2 (CO)3 (η2 -H2 )][Ir(PR3 )2 (Cl)(H)2 ][Ni(H2 O)6 ]2+0112240214403+2897101256811446910d8d8d6d8d10d1d6d6d10d0d0d6d6d8This notation might seem surprising at first sight, as it implies thatall the nonbonding electrons on the metal occupy d-type atomic orbitals(AO).
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