Limitations of the 18-Electron Rule

You will see that this reduces to Eq. 2.2 and so the two methods of electron counting are equivalent. 

There are many cases in which the electron count for a stable complex is not 18 examples are MeTiCIa, 8e Me2NbCl3, lOe WMe, 12e Pt(PCy3)2, 14e [M(H20)6] (M = V, 15e Cr, 16e Mn, 17e Fe, 18e), CoCp2, 19e and NiCp2, 20e. For the 18e rule to be useful, we need to be able to predict when it will be obeyed and when it will not. 

As shown in 2.13 and 2.14, we often write M—X to signify the covalent bond, but L- -M for the coordinate bond, as an indication that both electrons are regarded as coming from the ligand L. 

For the 18e rule to be useful, we need to be able to predict when it will be obeyed and when it will not. 

TABLE 23 The Metals that can Adopt a 16e Square Planar Configuration 

An important class of late metal complexes prefers 16e to 18e, because one of the nine orbitals is very high lying and usually empty. This can happen for the d metals of groups 8-11 (Table 2.4). Group 8 shows the least and group 11 the highest tendency to become 16e. When these metals are 16e, they normally become square planar, as in RhClI, IrCl(CO)L2, PdCl2L2, [PtCU], and [AuMeJ- (L = PR3). 

No rule is useful for main group elements for example, SiMe4 is 8e PFs, lOe SFe, 12e HgMe2,14e MeHg(bipy)+, 16e [I(py)2], 20e [SbFe] , 22e and IF7,24e. Although early metal (f complexes can have electron counts well below 18e (e.g., 8e TiMe4 and 12e WMee), an ambiguity 

The / block metals have seven / orbitals to fill before they even start on the d orbitals, and so they are essentially never able to bind a sufficient number of ligands to raise the electron count to the full count of 32e some examples are U(cot)2,22e, and Cp2LuMe, 28e. TTie stoichiometry of an / block complex instead tends to be decided by steric saturation of the space around the metal. Although coordination numbers of 8 and 9 are most common, a CN as high as 15 has been reported for a thorium aminodiboranate, [Th(H3BNMe2BH3)4].  

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We will define the rules that allow us to understand and construct bimetallic complexes, then we will provide the most salient features of metallic clusters including electron counting (limits of the 18-electron rule and Wade s rules) and the isolobal analogy. 

It should also be noted that for all the transition metals the effect of a positive charge on the metal is to lower the energy of all the orbitals. More important, however, is the increase in the separation between the 3d orbitals and the 4, 4p orbitals. From these observations we may, in a very qualitative manner, understand the occurrence and limitations of the 18-electron rule, as follows. 

This mode of calculation has been called the EAN rule (effective atomic number rule). It is valid for arbitrary metal clusters (closo and others) if the number of electrons is sufficient to assign one electron pair for every M-M connecting line between adjacent atoms, and if the octet rule or the 18-electron rule is fulfilled for main group elements or for transition group elements, respectively. The number of bonds b calculated in this way is a limiting value the number of polyhedron edges in the cluster can be greater than or equal to b, but never smaller. If it is equal, the cluster is electron precise. 

Although very reliable for Os3 clusters, the 18-electron rule is of limited use for larger ones. In the case of Os3 clusters with added transition metal atoms (Section VIII) one cannot simply apply the rule assuming that each short M—M contact involves a two-electron bond, but despite that the rule works well in most cases. 

The 18-electron rule can be a useful guide to stable organometallic compounds, especially when p-acceptor ligands are present, although it has the limitations referred to in Topic H9. Compounds 3, 5 and 6 obey this rule, but 1 without p bonding ligands has an electron count of only 12. Metallocenes [M(h5 C5H5)2 ] are known for the 3d series elements V-Ni, with 15-20 valence... 

Both of these reactions involve use of a strong reducing agent. In recent years, reduction of carbonyl complexes has been pushed to its limit with the synthesis of highly reduced anions such as [Mn(CO)4p, [Cr(CO)4] , [V(CO)5] , and [Ti(CO)5]. 2 t Since [MnlCOl ]", Cr(CO), and [V(CO)J were known to be relatively stable, there was some expectation, based on the 18-electron rule, that it would be possible to synthesize [Ti(CO)5] , even though Ti(CO)y is unknown. Often the expectation that a product should exist is a long way from synthetic success. The involved synthesis of [Ti(CO)(J illustrates the point.25 The overall reaction is... 

It has come to the point where some chemists, faced with a compound that does not exhibit a normal ll NMR spectrum, discard it and move on to a more promising project. In effect, our addiction to NMR spectroscopy has allowed the 18-electron rule to become a set of blinders limiting our view. 

Mechanistic studies have been used to attempt to explain the rapid initiation rate of the 18-electron pyridine solvates. Using the reported initiation rate of 4a [6], the upper limit of free energy of activation was determined to be 15.45kcalmol" at 5°C using the Eyring equation [39], Similar to the Grubbs-Hoveyda complexes, these precatalysts have multiple potential pathways by which they can initiate (e.g., interchange or dissociative). An associative mechanism can be ruled out due to the coordinatively saturated, six-coordinate nature of these complexes. 

The ability of an atom in a molecular entity to expand its valence shell beyond the limits of the Lewis octet rule. Hyper-valent compounds are common for the second and subsequent row elements in groups 15 -18 of the periodic table. Hyperva-lent bonding implies a transfer of the electrons from the central (hypervalent) atom to the nonbonding molecular orbitals which it forms with (usually more electronegative) ligands. A typical example of a hypervalent bond is a linear three-center, four-electron bond, e.g., that of the Fap-P-Fap fragment of PF5. 

It can be seen also that the 18-electron rule will tend to break down when the first-row transition metals are in a high oxidation state. There may also be steric limitations to the formation of 18-electron complexes in high oxidation states - consider for example the hypothetical cation [Cr(CO)9] . Steric limitations may also explain why vanadium forms the monomeric, paramagnetic and 17-electron hexacarbonyl, V(CO)6, rather than the dimeric complex [V(CO)6]2 which would be diamagnetic and would obey the 18-electron rule. 

With metal clusters it is even harder than in other fields of inorganic chemistry to substantiate theoretical results by energy measurements. Only two such measurements have come to the attention of the author — the photoelectron spectrum of [CpFe(C0)]4 370) andbond energy determinations in 03(00)9CX-compounds 187). However, a considerable number of papers deal with metal-metal bonding in, and the symmetry properties of, clusters as related to their stoichiometry and their electron count. These studies have confirmed the wide apphcability of the simple 18-electron rule in predicting metal-metal bonds and structures, but they have also led to an understanding of the limits of this rule for clusters with more than four metal atoms. 

Limitations of the 8- and 18-electron rule localized electron-deficient compounds... 

Limitations of the eight- and 18-electron rule delocalized bonding... 

In the chapters that follow you will find numerous exercises in counting electrons for clusters - elaborations of the eight- and 18-electron rules for these complex structures. The same factors that cause the eight- and 18-electron rules to fail will similarly limit cluster-counting rules based upon them. Like these fundamental rules, even when satisfied, the cluster-counting rules yield no detailed information on electronic structure. Hence, the bolder student occasionally asks, Why count by which he or she means Of what real value are these counting exercises if little is learned about where the electrons really are ... 

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