18 valence electron metal-like compounds

Boron is a unique and exciting element. Over the years it has proved a constant challenge and stimulus not only to preparative chemists and theoreticians, but also to industrial chemists and technologists. It is the only non-metal in Group 13 of the periodic table and shows many similarities to its neighbour, carbon, and its diagonal relative, silicon. Thus, like C and Si, it shows a marked propensity to form covalent, molecular compounds, but it differs sharply from them in having one less valence electron than the number of valence orbitals, a situation sometimes referred to as electron deficiency . This has a dominant effect on its chemistry. 

The alkali metals also release their valence electrons when they dissolve in liquid ammonia, but the outcome is different. Instead of reducing the ammonia, the electrons occupy cavities formed by groups of NH3 molecules and give ink-blue metal-ammonia solutions (Fig. 14.14). These solutions of solvated electrons (and cations of the metal) are often used to reduce organic compounds. As the metal concentration is increased, the blue gives way to a metallic bronze, and the solutions begin to conduct electricity like liquid metals. 

In this way we come to class III complexes, i.e. complexes in which the two sites are indistinguishable and the element has a non-integral oxidation state (delocalized valence). Usually one divides this class in two subclasses. In class IIIA the delocalization of the valence electrons takes place within a cluster of equivalent metal ions only. An example is the [NbgCli2] ion in which there are six equivalent metal ions with oxidation state + 2.33. In class IIIB the delocalization is over the whole lattice. Examples are the linear chain compound K2Pt(CN)4.Bro.3o. 3H2O with a final oxidation state for platinum of 2.30, and three-dimensional bronzes like Na WOg. 

Since these structures are formed by filling the open spaces in the diamond and wurtzite structures, they have high atomic densities. This implies high valence electron densities and therefore considerable stability which is manifested by high melting points and elastic stiffnesses. They behave more like metal-metalloid compounds than like pure metals. That is, like covalent compounds embedded in metals. 

Armed with the band masses and relative positions of the band centres one may easily sketch the likely valence electron energies in actinide metals and compounds. This is useful because it is easy to understand, involves no computation and gives a rough idea of what to expect from large machine computations of electronic structure. 

An s-block element has a low ionization energy, which means its outermost electrons can be lost easily. A Group 1 element is likely to form +1 ions, such as Li+, Na+, and K+. Group 2 elements similarly form +2 ions, such as Mg2+, Ca2+, and Ba2+. An s-block element is likely to be a reactive metal with all the features that the name metal implies (Fig. 1.48, Table 1.4). Because ionization energies are lowest at the bottom of each group, and the elements there lose their valence electrons most easily, the heavy elements cesium and barium react most vigorously of all s-block elements. They have to be kept stored out of contact with air and water. The alkali metals have few direct uses as materials but are enormously important as compounds. 

Both A1 and B are in Group 3 of the periodic table and have three valence electrons in their outer shell. These elements can form three bonds. However, there is still room for a fourth bond. For example in BF3, boron is surrounded by six electrons (three bonds containing two electrons each). However, boron s valence shell can accommodate eight electrons and so a fourth bond is possible if the fourth group can provide both electrons for the new bond. Since both boron and aluminium are in Group 3 of the periodic table, they are electropositive and will react with electron-rich molecules so as to obtain this fourth bond. Many transition metal compounds can also act like Lewis acids (e.g. TiCl4 and SnCl4). 

All elements of the group form Zinti compounds with electropositive metals. Continuous networks of covenantally bonded atoms are generally found, rather than the clusters common with group 14. For example, NaAl and NaTl have tetrahedral diamond-like networks of A1 or Tl, which can be understood on the basis that A1 and Tl- have the same valence electron count as carbon. 

The correlation between the valence electron counts and the stabilities of intermetallic phases and stmctures were also espoused by others, like the physical chemists Neds N. Engel (b. 1904) and Leo Brewer (1919-2005), although Hume-Rothery found their result somewhat controversial. The Engel-Brewer theory asserts that the crystal stmctures of transition metals and their intermetallic compounds are determined solely by the number of valence s and p electrons. For example, Engel suggested in 1949 that the BCC stmcture correlated with (where n is the total number of valence... 

For binary metal carbonyl compounds, the 18-electron mle is a very useful concept. Stable metal complexes will be formed when the metal has 18 electrons in its valence shell (metal valence electrons -H 2 electrons from each CO ligand). Since Tc(0) has 7 valence electrons, the neutral monomeric species Tc(CO) cannot be stable, but ions like [Tc(CO)6]" or [Tc(CO)5] attain a total of 18 electrons. In the neutral molecule, it will dimerize to Tc2(CO)io in order to obey the 18-electron rule. The formation of a Tc-Tc bond adds an electron on each Tc atom. This 18-electron mle is quite useful to predict the stmctures of the metal binary carbonyl compounds. 

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