Covalent bonding Hybrid orbitals Lewis

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. 

Trimethylboron is an example of one type of Lewis acid. This molecule has trigonal planar geometry in which the boron atom is s hybridized with a vacant 2 p orbital perpendicular to the plane of the molecule (Figure 21-11. Recall from Chapter 9 that atoms tend to use all their valence s and p orbitals to form covalent bonds. Second-row elements such as boron and nitrogen are most stable when surrounded by eight valence electrons divided among covalent bonds and lone pairs. The boron atom in B (CH ) can use its vacant 2 p orbital to form a fourth covalent bond to a new partner, provided that the new partner supplies both electrons. Trimethyl boron is a Lewis acid because it forms an additional bond by accepting a pair of electrons from some other chemical species. 

The simplest type of Lewis acid-base reaction is the combination of a Lewis acid and a Lewis base to form a compound called an adduct. The reaction of ammonia and trimethyl boron is an example. A new bond forms between boron and nitrogen, with both electrons supplied by the lone pair of ammonia (see Figure 21-21. Forming an adduct with ammonia allows boron to use all of its valence orbitals to form covalent bonds. As this occurs, the geometry about the boron atom changes from trigonal planar to tetrahedral, and the hybrid description of the boron valence orbitals changes from s p lo s p ... 

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. 

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. 

The quantum content of current theories of chemical cohesion is, in reality, close to nil. The conceptual model of covalent bonding still amounts to one or more pairs of electrons, situated between two atomic nuclei, with paired spins, and confined to the region in which hybrid orbitals of the two atoms overlap. The bond strength depends on the degree of overlap. This model is simply a paraphrase of the 19th century concept of atomic valencies, with the incorporation of the electron-pair conjectures of Lewis and Langmuir. Hybrid orbitals came to be introduced to substitute for spatially oriented elliptic orbits, but in fact, these one-electron orbits are spin-free. The orbitals are next interpreted as if they were atomic wave functions with non-radial nodes at the nuclear position. Both assumptions are misleading. 

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. ... 

All resonance structures for the same molecule must have the same sigma framework (w sigma bonds form from the head on overlap of hybridized orbitals). Furthermore, they must be correct w Lewis structures with the same number of electrons (and consequent charge) as well as the same number of unpaired electrons. Resonance structures with arbitrary separation of charge are unimportant, as are those with fewer covalent bonds. These unimportant resonance structures only contribute minimally (or not at all) to the overall bonding description however, they are important in some cases such as for a w carbonyl group. 

In Section 10.4 we saw that the B atom in BF3 is xp -hybridized. The vacant, unhybridized 2p orbital accepts the pair of electrons from NH3. So BF3 functions as an acid according to the Lewis definition, even though it does not contain an ionizable proton. Note that a coordinate covalent bond is formed between the B and N atoms, as is the case in all Lewis acid-base reactions. 

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. 

Although the title has an almost magical sound to it, the nature of the chemical bond was truly the domain Pauling began to explore. He formulated the concept of hybridization to explain how localized atomic orbitals best overlap to form two-electron bonds. The Kossel-Lewis-Langmuir picture explained ionic and covalent bonding in terms of the octet rule. An interesting question was... 

A FIGURE 23.27 MetaHigand bond formation. The ligand acts as a Lewis base by donating its nonbonding electron pair to a hybrid orbital on the metal ion. The bond that results is strongly polar with some covalent character. 

It is useflil to show the valence bond representations of the complexes [CoFe] and [Co(NH3)6], which can then be compared with representations from the crystal field and molecular orbital theories to be discussed later. First, we must know from experiment that [CoF ] contains four unpaired electrons, whereas [Co(NH3)g] has all of its electrons paired. Each of the ligands, as Lewis bases, contributes a pair of electrons to form a coordinate covalent bond. The valence bond theory designations of the electronic structures are shown in Figure 2.7. The bonding is described as being covalent. Appropriate combinations of metal atomic orbitals are blended together to give a new set of orbitals, called hybrid orbitals. 

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

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. 

Covalent bond formation in ethylene (CHjCHj). (a) Lewis structure, (b) Overlap oisp hybrid orbitals on adjacent carbons forms a C—C cr bond (see Figure 1.18), and overlap of carbon sp hybrid orbitals on carbons with 1 orbitals on hydrogens gives C—H [Pg.73]

As noted above (equations 9 and 10), each pair of valence NHOs /ia, /ib leads to a complementary pair of valence bond (/)ab) and antibond ( ab) orbitals. Although the latter orbitals play no role in the elementary Lewis picture, their importance was emphasized by Lennard-Jones and Mulliken in the treatment of homonuclear diatomic molecules. Since valence antibonds represent the residual atomic valence-shell capacity that is not saturated by covalent bond formation, they are generally found to play the leading role in noncovalent interactions and delocalization effects beyond the Lewis structure picture. Indeed, it may be said that the NBO treatment of bond-antibond interactions constitutes its most unique and characteristic contribution toward extending the Lewis structure concepts of valence theory. Although the NBO hybrids and polarization coefficients are chosen to minimize the role of antibonds, the final non-zero weighting of non-Lewis orbitals reflects their essential contribution to wavefunction delocalization. 

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