Boron electron-deficient compounds

This chapter summarizes recent developments in the expanding field of electron-deficient compounds having from three up to 13 skeletal boron and carbon atoms. In particular, the focus will be on the transition of classical organoboranes into non-classical compounds. Therefore, we first want to briefly review electron counting rules and bonding characteristics of these classes. For a more thorough discussion see Chapter 1 by King and Schleyer. 

Being an electron deficient compound, boron trifluoride forms complexes with Lewis bases and compounds that have unshared pair(s) of electrons. With ammonia, it forms boron trifluoride ammonia. Similar coordination compounds are formed with monoethylamine, BF3-NH2C2H5 diethyl ether, CH3CH20(BF3)CH2CH3 and methanol, BF3—OHCH3. It forms a sohd complex HNO3-2BF3 with concentrated nitric acid. 

This term is used to describe a three-centre bond - usually of the type E-H-E and non-linear, in contrast to the other variety - which we imagine to be formed by three orbitals (one on each atom) and two electrons (one furnished by the central H atom). Such bonding is invoked in electron-deficient compounds where there are insufficient electrons to form the requisite number of two-centre bonds in a Lewis/VB treatment. Electron-deficient compounds abound in boron chemistry the classic case of diborane B2H6 will be discussed in detail. 

Boranes, carboranes, and metallocarboranes are electron-deficient compounds in which B-H-B three-center two-electron (3c-2e) bridged bonds are found. The B-H-B bond results from the overlap of two boron sp3 hybrid orbitals and the Is orbital of the hydrogen atom, which will be discussed in Chapter 13. 

Borane has the same structure as a carbocation. The boron is sjr hybridized, with trigonal planar geometry, and has an empty p orbital. Although neutral, it is electron deficient because there are only six electrons around the boron. It is a strong Lewis acid. An electron-deficient compound often employs unusual bonding to alleviate somewhat its instability. In the case of borane, two molecules combine to form one molecule of diborane ... 

Borane is an electron-deficient compound. It has only six valence electrons, so the boron atom in BH3 cannot have an octet. Acquiring an octet is the driving force for the unusual bonding structures ( banana bonds, for example) found in boron compounds. As an electron-deficient compound, BH3 is a strong electrophile, capable of adding to a double bond. This hydroboration of the double bond is thought to occur in one step, with the boron atom adding to the less substituted end of the double bond, as shown in Mechanism 8-6. 

Lipscomb, W. N. (1963). Boron Hydrides. New York Benjamin. Wade, K. (1971). Electron Deficient Compounds. London Nelson. 

Metallacarbaborane (also called metallacarborane) an electron-deficient compound, usually a polyhedral cluster comprising an array of boron-hydride (BH), carbon-hydride (CH), and metal (ML, where L = ligand) fragments the inclusion of a substituted carbon fragment CR (R = alkyl, aryl, or trimethylsilyl) in place of a CH unit is common... 

Hydrogen bridges between the beryllium atoms produce a polymeric structure for BeH2, as shown in Fig. 18.6. The localized electron model describes this bonding by assuming that only one electron pair is available to bind each Be—H—Be cluster. This is called a three-center bond, since one electron pair is shared among three atoms. Three-center bonds have also been postulated to explain the bonding in other electron-deficient compounds (compounds where there are fewer electron pairs than bonds), such as the boron hydrides (see Section 18.5). 

Formally subvalent compounds of boron containing a boron-boron single bond are intermediate in structural complexity between simple monoboron derivatives and the polyhedral electron-deficient compounds of the element. The properties of such compounds, particularly the simple derivatives of the B2X4 type, have attracted the attention of several groups of workers since Stock s initial discovery of BjCL some 45 years ago (96). These materials provide the simplest examples of catenation in boron chemistry and offer suitable systems in which to study the properties of the covalent B—B bond and the characteristic chemistry of compounds containing this linkage. 

The systematic principles of boron hydride structures ahd chemistry are the principal subjects of the present review. There are several reasons why these principles became clear such a long time after the discovery of these compounds (a) most of the compounds must be handled in grease-free vacuum line systems (b) some of the boron hydrides are unstable at ordinary temperatures, explosive on contact with air, and toxic (c) the structures are based on principles, still incompletely developed, of electron-deficient compounds and (d) location of the hydrogen atoms is a crucial part of the structure determinations, unlike the situation in hydrocarbons, and had to be done for the most part in X-ray diffraction studies of single crystals grown at low temperatures. 

The recent report by W. H. Bauer and S. E. Wiberley [Abstr. ISSrd Meeting Am. Chem. Soc., San Francisco, p. 13L, (1958)] and by J. A. Hammond (private communication) of a possibly electron deficient compound B4H12O has suggested the investigation of valence rules for this and similar compounds. Assume that Bp Hp +gOn satisfies the same bonding rules as the boron hydrides. Then the equations of orbital... 

Compounds of boron and hydrogen, such as B2H6 and B loH h —> Electron-deficient compounds with three-center, two-electron B-H-B bonds... 

The boron hydrides are called electron-deficient compounds because they are easily reduced by hydrogen. Incorrect boron hydrides are electron-deficient because they lack the electrons required to fill the bonding and non-bonding molecular orbitals. 

Moffitt and Ballhausen,41 and the application of the molecular orbital theory to electron-deficient compounds, particularly the hydrides of boron, by Eberhardt, Crawford, and Lipscomb has been fully described by these authors in their original papers.14 The present review will, therefore, be devoted entirely to those developments in the molecular orbital theory which have been associated with its application to the electronic spectra of unsaturated hydrocarbons. These developments form in themselves a relatively coherent story, the main lines of which are now clear, and it seems a suitable moment at which to put the history of the subject into perspective. Before doing this, however, it will be convenient to outline the general premises of the molecular orbital theory. 

A classic account of the boron hydrides and related electron-deficient compounds. 

The most general view of acids and bases was advanced by G. N. Lewis. In this model, acids are substances which have an affinity for lone electron pairs, and bases are substances which possess lone electron pairs. Water and ammonia are the most common substances which possess lone electron pairs, and therefore behave as bases in the Lewis scheme. The reaction of silver ion, Ag with cyanide ion, CN , and boron trifluoride, BF3 (an electron-deficient compound), with ammonia, NH3, are two examples of Lewis acid-base reactions. The Lewis acid-base concept is most useful in chemical reactions in nonaqueous solvents. We will not find it useful in our study of ionic equilibria in water. 

Boron has a small atomic radius and a relatively high ionization energy. In consequence its chemistry is largely covalent and it is generaUy classed as a metalloid. It forms a large number of volatile hydrides, some of which have the uncommon bonding characteristic of electron-deficient compounds. It also forms a weakly acidic oxide. In some ways, boron resembles siUcon (see diagonal relationship). 

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