Covalent bonding Hybrid orbitals Lewis

Covalent bond formation in ethylene, (a) Lewis structure, (b) overlap of sf hybrid orbitals forms a sigma [Pg.25]

The B atom inBFs is 5p -hybridized [M Section 9.4]. The vacant, unhybridized Ip orbital accepts the pair of electrons from NH3. Thus, BF3 functions as an acid according to the Lewis definition, even though it does not contain an ionizable proton. A coordinate covalent bond [M Section 8.8] is formed between the B and N atoms. In fact, every Lewis acid-base reaction results in the formation of a coordinate covalent bond. 

Soon after the quantum revolution of the mid 1920s, Linus Pauling and John C. Slater expanded Lewis s localized electronic-structural concepts with the introduction of directed covalency in which bond directionality was achieved by the hybridization of atomic orbitals.1 For normal and hypovalent molecules, Pauling and Slater proposed that sp" hybrid orbitals are involved in forming shared-electron-pair bonds. Time has proven this proposal to be remarkably robust, as has been demonstrated by many examples in Chapter 3. 

Valence bond theory pictures bonding in complex ions as arising from coordinate covalent bonding between Lewis bases (ligands) and Lewis acids (metal ions). Ligand lone pairs occupy hybridized metal-ion orbitals to form complex ions with characteristic shapes. 

The idea of using organotin compounds as ionophores was based on the fact that since, like carbon, tin forms covalent bonds via sp hybridization, and with the presence of empty d orbitals, it can coordinate with up to three extra electron-donating substituents, such as Lewis-basic anions. It was Selwyn, in 1970,9.10 ujal (ook advantage of this property and showed clearly the direct role of the trimethyltin, tri-n-propyltin, tri-n-butyltin, and triphenyltin chlorides on the active chloride transport mediated in mitochondrial membranes, as shown in Figure 3.4.5. It was also shown in this study that the mediation is based on chloride-hydroxide antiporter transport. This fact was verified many years later, as Simon showed, based on NMR and other studies, that indeed these compounds act as neutral carriers in liquid polymeric membranes.  

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

Boron is the first member of Group 13 elements. It is a non-metal and forms only covalent compounds. It exhibits an oxidation state +3 in all its compounds. The electron configuration of boron is ns np and boron is said to form three covalent bonds using sp hybrid orbitals. The compounds of boron are electron deficient and accept a pair of electrons (Lewis acids). The bonding in certain boron compounds is of considerable theoretical interest. 

D) BF3 is a trigonal-planar molecule because electrons can be found in only three places in the valence shell of the boron atom. As a result, the boron atom is sp hybridized, which leaves an empty 2p orbital on the boron atom. BF3 can therefore act as an electron-pair acceptor, or Fewis acid. It can use the empty 2p orbital to pick up a pair of nonbonding electrons from a Fewis base to form a covalent bond. BF3 therefore reacts with Lewis bases such as NH3 to form acid-base complexes in which all of the atoms have a filled shell of valence electrons. 

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