Yves Jean - Molecular Orbitals of Transition Metal Complexes (793957), страница 9
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In general, the π orbital is higherin energy than the MO that describe the σ bonds, while the π ∗ orbital islower than the σ ∗ MO (§1.3.4). Although none of these π or π ∗ orbitalsis nonbonding, in contrast to the p orbitals of AH2 , AH, and A ligands,their energy level is neither sufficiently low (π) nor sufficiently high(π ∗ ) for one to be able, a priori, to neglect their role in the metal–ligandinteraction.Carbon monoxide, CO, also called the carbonyl ligand, is an examplewhich illustrates these points nicely. The essential features of its electronic structure are shown in Figure 1.7.
The highest occupied orbitalis a nonbonding σ orbital, mainly concentrated on the carbon atomand polarized in the direction away from the oxygen atom. This orbital,which describes the lone pair on the carbon atom, is the one whichallows a σM−−CO bond to be formed (L-type ligand). The two πCO bond≡O are lower in energy.ing orbitals associated with the π bonds in C≡They are mainly concentrated on oxygen, as that atom is more electronegative than carbon. The lowest empty orbitals are the antibonding∗ orbitals, which have a larger contribution from carbon than oxygen.πCOThese four orbitals can lead to π-type interactions with orbitals of suitable symmetry on the metal, similar to those we have already seen withthe nonbonding p orbitals of the AH2 and AH molecules.We shall therefore have to study a set of five orbitals (one σ , two π,and two π ∗ ) when we wish to analyse the metal–carbonyl bond.121.5.2.5.
π complexesIn the examples studied so far, the ligand is bound to the metal centreby only one of its atoms. The situation is different when several atomsof the ligand are bound in an equivalent manner to the metal centre (ηxcoordination). This is the case for π complexes, in which the π systemof the ligand is oriented towards the metal. All the π orbitals of theligand, both occupied and empty, must now be considered to describethe metal–ligand bonds. As an example, we shall treat an η2 -ethylenecomplex in detail in Chapter 3 (§ 3.4).Setting the scene1.6.
Initial orbital approach to MLℓ complexesThe shape and the energy of the molecular orbitals of a complexdepend on the number of ligands and their geometrical arrangementaround the metal. It is possible to obtain some important informationon the MO without defining the particular complex studied. The purpose of this paragraph is thus to derive the general characteristics of theorbital structure which do not depend (or depend only slightly) on thecomplex.1.6.1. Simplified interaction diagramWe shall consider, for simplicity, a complex in which the metal issurrounded by ℓ identical ligands, each with just a single orbital thatcan take part in the metal–ligand interaction (a σ interaction, § 1.5.1).A simplified diagram for the interaction between the ℓ ligand orbitalsand the nine atomic orbitals on the metal (five d orbitals, one s orbital,and three p orbitals, without distinction) is given in Figure 1.8.
In thisdiagram, the metal orbitals are placed higher in energy than those on theligands, since the latter are more electronegative. The ℓ ligand orbitalsinteract with ℓ metal orbitals, to form ℓ bonding MO and the associatedℓ antibonding MO. There are therefore (9 − ℓ) nonbonding orbitals thatremain on the metal.antibondingMO∆E –9metalAO(9 – )nonbondingMOligandorbitalsFigure 1.8. Simplified diagram for theinteraction of the atomic orbitals on a metalcentre and the ℓ ligands which surround it (σinteractions only).bondingMOInitial orbital approach to MLℓ complexesGiven the relative energies for the participating orbitals (§ 1.3.2), wecan make the following points:M–H *M–H1-391-401.
The bonding MO, which describe the σM−Lig bonds, are mainlyconcentrated on the ligand orbitals. An example is given for anM-H bond involving the metal z2 orbital (1-39).2. The corresponding antibonding MO are mainly concentrated onthe metal orbitals (1-40).3. The nonbonding MO are orbitals that are localized on the metalcentre. A more detailed analysis of the orbital structure of complexes (Chapter 2) will show that usually, but not always, thesenonbonding MO are pure d orbitals, or orbitals in which theprincipal component is of d type.1.6.2. Strong-field and weak-field complexesThe separation between the energy levels (E − , Figure 1.8) of the∗) is directly linked to thenonbonding and antibonding MO (σM−Ligstrength of the interaction between the ligand orbitals and those onthe metal. The stronger this interaction, the more the antibondingorbitals are destabilized, so the larger the energy gap E − .
When themetal–ligand interaction is strong, E − is large and one refers to strongfield complexes; in contrast, when E − is small, one refers to weak-fieldcomplexes.1.6.3. Electronic configuration and the 18-electron ruleIn so far as the electronic occupation of the molecular orbitals isconcerned, the stability of an MLℓ complex is generally maximized whenthe bonding MO, of which there are ℓ, together with the nonbondingMO, of which there are (9 − ℓ), are doubly occupied, but when theℓ antibonding MO remain empty (Figure 1.8).
The bonding MO describethe M-Lig bonds and the nonbonding MO represent lone pairs on themetal. In this situation, the total number of electrons is:Nt = (2 × ℓ) + 2 × (9 − ℓ) = 18(1.13)In this way, we rationalize the 18-electron rule that was previouslydiscussed with reference to the valence electronic structure of the nearestnoble gas (§ 1.1.2.1).The electrons that occupy the nonbonding MO are not used to formmetal–ligand bonds. They therefore correspond to the n electrons that‘remain’ on the metal in the classical counting scheme (§ 1.1.2.3).
The dnnotation for the electronic configuration of the metal assumes, however,Setting the scenethat all the occupied nonbonding orbitals on the metal are of d type (seeChapter 2).Even if the 18-electron rule is often obeyed, we must not forgetthat there are many exceptions. There are some complexes that havefewer than 18 electrons. For example, the [M(Lig)4 ] complexes thatadopt a ‘square-planar’ geometry (four ligands at the vertices of asquare whose centre is occupied by the metal) with a d8 electronicconfiguration (e.g.: [Ir(CO)(Cl)(PPh3 )2 ], § 1.1.2.1) are 16-electron complexes. Since the bonding MO that describe the bonds are doublyoccupied, this analysis shows that one of the nonbonding orbitalsin Figure 1.8 is empty, and a more detailed study of the electronicstructure is necessary to understand this result (Chapter 2, §2.2).These 16-electron complexes are stable, but often reactive towardsother molecules since they have a tendency to form 18-electron complexes, by binding other ligands.
For example, Wilkinson’s catalyst[Rh(PPh3 )3 Cl] is used industrially for the catalytic hydrogenation ofolefins (see Exercise 1.3).There are also some complexes with more than 18 electrons, such as[Ni(H2 O)6 ]2+ which possesses 20 electrons (§ 1.1.2.1). Some antibonding MO must therefore be occupied, which can happen only if theyare sufficiently low in energy, as is the case in weak-field complexes.Organometallic complexes, characterized by the presence of one or several metal–carbon bonds, are strong-field complexes. It is therefore rarefor them to possess more than 18 electrons.1.6.4. Analogy with the octet ruleIn the same way, one can construct a simplified interaction diagram forAHn (or ARn ) molecules in which A is an element from the second orthird row of the periodic table (C, Si, N, P, O, S, etc.).
There are thusfour valence orbitals on the central atom: one s AO, and three p AO. Theσ interactions with the orbitals on the atoms bound to A lead to theformation of n bonding MO, n antibonding MO and (4 − n) nonbondingMO. If the bonding and nonbonding MO are doubly occupied, thenumber of electrons Nt is equal toNt = (2 × n) + 2 × (4 − n) = 8(1.14)We have therefore derived the octet rule. As examples, we canquote:CH4NH3OH2FHn=4n=3n=2n=14 bonding MO, 4 antibonding MO, 0 nonbonding MO3 bonding MO, 3 antibonding MO, 1 nonbonding MO2 bonding MO, 2 antibonding MO, 2 nonbonding MO1 bonding MO, 1 antibonding MO, 3 nonbonding MOExercisesIt is straightforward to verify that the number of bonding,nonbonding, and antibonding MO predicted by this simple model agreeconsistently with the detailed orbital structures of the NH3 , OH2 , andHF molecules (§ 1.1.4–1.1.6).Exercises1.1What are the two coordination modes of the allyl ligand,H2 C-CH-CH2 , for which a Lewis structure is given below? In eachcase, categorize the ligand as Lℓ Xx .HCHHCCHH1.2Write each of the following complexes as [MLℓ Xx ]q , give theoxidation state of the metal no, the electronic configurationdn , and the total number of electrons Nt .
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