Boron electron-deficient molecules

Localized Bonds. Because boron hydrides have more valence orbitals than valence electrons, they have often been called electron-deficient molecules. This electron deficiency is partiy responsible for the great interest surrounding borane chemistry and molecular stmcture. The stmcture of even the simplest boron hydride, diborane(6) [19287-45-7] 2 6 sufficientiy challenging that it was debated for years before finally being resolved (57) in favor of the hydrogen bridged stmcture shown. 

As has already been mentioned, boron halides are electron-deficient molecules. As a result, they tend to act as strong Lewis acids by accepting electron pairs from many types of Lewis bases to form stable acid-base adducts. Electron donors such as ammonia, pyridine, amines, ethers, and many other types of compounds form stable adducts. In behaving as strong Lewis acids, the boron halides act as acid catalysts for several important types of organic reactions (see Chapter 9). 

The metal halide thus functions in similar manner to the proton and may be considered to be an acidic catalyst (cf. Luder and Zuffanti, 19). The catalyst-olefin complex differs in one significant respect from the product formed by the addition of the proton (or the corresponding acid) to the olefin the halide catalyst is a neutral but electronically deficient molecule and combines with the pi electrons of the double bond to form a coordinate bond between the carbon atom and the aluminum or boron. On the other hand, the addition of the positive proton to the double bond results in the formation of a true (covalent) link between carbon and hydrogen. In other words, the complex, while it contains an electron-deficient (hence, positive) carbon atom, is in itself electronically neutral the product of the addition of a proton to the alkene contains a similar carbon atom but is itself electrically positive. It has been suggested (Whitmore and Meunier, 20) that this difference is related to the fact that metal halide catalysts tend to yield much higher polymers than do the acid (proton) catalysts. 

Evidence for the reverse process, donation of electron density from the nucleophilic dimer atom to an electron-deficient molecule, also exists. Konecny and Doren theoretically found that borane (BH3) will dissociatively adsorb on Si(100)-2x1 [293]. While much of the reaction is barrierless, they note an interaction between the boron atom and the nucleophilic atom of the Si dimer during the dissociation process. Cao and Hamers have demonstrated experimentally that the electron density of the nucleophilic dimer atom can be donated to the empty orbital of boron trifluoride (BF3) [278]. XPS on a clean Si(100)-2 x 1 surface at 190 indicates that BF3 dissociates into BF2(a) and F(a) species. However, when BF3 is exposed on a Si(100)-2 x 1 surface previously covered with a saturation dose of trimethylamine, little B-F dissociation occurs, as evidenced by the photoelectron spectrum. They conclude that BF3 molecularly adsorbs to the nucleophilic dimer atom and DFT calculations indicate that the most energetically favorable product is a surface-mediated donor-acceptor complex (trimethylamine-Si-Si-BF3) as shown in Figure 5.19. 

The group 3A elements—B, Al, Ga, In, and T1—are metals except for boron, which is a semimetal. Boron is a semiconductor and forms molecular compounds. Boranes, such as diborane (B2H6), are electron-deficient molecules that contain three-center, two-electron bonds (B-H-B). 

Because of the number and complexity of the boranes and their derivatives we shall not attempt to describe all their structures in detail. Apart from B2H6 most of the boranes have boron skeletons which are usually described as icosahedral fragments. (They could equally well be related to the more recently discovered carboranes B C2H +2 since the boron-carbon skeletons in these compounds are highly symmetrical triangulated polyhedra (Table 24.8) which include the icosahedron.) The great theoretical interest of the boranes stems from the fact that they are electron-deficient molecules, that is, there are not sufficient valence electrons to bond together all the atoms by normal electron-pair bonds. In these molecules the number of atomic orbitals (1 for each H and 4 for each B) is greater than the total number of valence electrons ... 

Electron-Deficient Molecules Gaseous molecules containing either beryllium or boron as the central atom are often electron deficient that is, they have/ewer... 

When the three electron groups are bonding groups, the molecular shape is trigonal planar (AX3). Boron trifluoride (BF3), another electron-deficient molecule, is an example. It has six electrons around the central B atom in three single bonds to F atoms. The nuclei lie in a plane, and each F—B—F angle is 120° ... 

To illustrate electron-deficient molecules, consider in Scheme 2.9 the molecules that can be constructed from beryllium (Be) and boron (B) with as many hydrogen... 

In the beryllium chloride molecule (BeCl2(g)), the beryllium atom has only four electrons in its valence shell (Figure 4-31). The molecule is described as electron deficient. The boron trichloride molecule is also electron deficient the central boron atom has only six electrons in its valence shell (Figure 4-32). A related example of an electron-deficient molecule is aluminium trichloride, AICI3. The aluminium atom has only six electrons in its valence shell. All these molecules have incomplete octets. 

Typical examples of this class of Lewis acids are electron-deficient molecules such as the halides of boron, beryllium and aluminium, for example, BCI3, BeCl2 and AICI3. 

Boron trifluoride is an electron-deficient molecule. It has only six electrons in its outer shell. The three bonding pairs of electrons repel each other equally, so the F—B—F bond angles are 120° (Figure 4.18). We describe the shape of the molecule as trigonal planar. Trigonal means having three angles . 

Usually, when applied to molecular closed-shell ground states, the various localization methods lead to orbitals that are concentrated either around individual nuclei (for example inner-shell orbitals not very different from those in the free atoms) or in the valence regions (for example lone-pair orbitals, mainly on one centre, and bond-pair orbitals, confined mainly to two adjacent centres). Naturally, there are exceptions, in which such a high degree of localization cannot be attained (notably in electron-deficient molecules like the boron hydrides, and in conjugated systems), but in such cases the remaining delocalization is associated with very particular molecular properties. 

Boranes are typical species with electron-deficient bonds, where a chemical bond has more centers than electrons. The smallest molecule showing this property is diborane. Each of the two B-H-B bonds (shown in Figure 2-60a) contains only two electrons, while the molecular orbital extends over three atoms. A correct representation has to represent the delocalization of the two electrons over three atom centers as shown in Figure 2-60b. Figure 2-60c shows another type of electron-deficient bond. In boron cage compounds, boron-boron bonds share their electron pair with the unoccupied atom orbital of a third boron atom [86]. These types of bonds cannot be accommodated in a single VB model of two-electron/ two-centered bonds. 

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