Covalent molecules orbital hybridization

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

The VSEPR theory is only one way in which the molecular geometry of molecules may be determined. Another way involves the valence bond theory. The valence bond theory describes covalent bonding as the mixing of atomic orbitals to form a new kind of orbital, a hybrid orbital. Hybrid orbitals are atomic orbitals formed as a result of mixing the atomic orbitals of the atoms involved in the covalent bond. The number of hybrid orbitals formed is the same as the number of atomic orbitals mixed, and the type of hybrid orbital formed depends on the types of atomic orbital mixed. Figure 11.7 shows the hybrid orbitals resulting from the mixing of s, p, and d orbitals. 

The water molecule is composed of two hydrogen atoms covalently bonded to an oxygen atom with tetrahedral (sp3) electron orbital hybridization. As a result, two lobes of the oxygen sp3 orbital contain pairs of unshared electrons, giving rise to a dipole in the molecule as a whole. The presence of an electric dipole in the water molecule allows it to solvate charged ions because the water dipoles can orient to form energetically favorable electrostatic interactions with charged ions. 

Write the Lewis formulas and predict the hybrid orbitals and the shapes of these polyatomic ions and covalent molecules (a) HgCl2 (b) BF3 (c) Bp4 (d) SbCl5 (e) SbFg". (a) What is the hybridization of each C in these molecules (i) H2C=0 (ii) HC=N (iii) CH3CH2CH3 (iv) ketene, H2C=C=0. (b) Describe the shape of each molecule. The following fluorides of xenon have been well charac-... 

There is no quantum-mechanical evidence for spatially directed bonds between the atoms in a molecule. Directed valency is an assumption, made in analogy with the classical definition of molecular frameworks, stabilized by rigid links between atoms. Attempts to rationalize the occurrence of these presumed covalent bonds resulted in the notion of orbital hybridization, probably the single most misleading concept of theoretical chemistry. As chemistry is traditionally introduced at the elementary level by medium of atomic orbitals, chemists are conditioned to equate molecular shape with orbital hybridization, and reluctant to consider alternative models. Here is another attempt to reconsider the issue in balanced perspective. 

Write the Lewis formulas and predict the hybrid orbitals and the shapes of these polyatomic ions and covalent molecules (a) HgCl2 (b) BFj (c) BF4- (d) SbClj (e) SbF -. 

Both experimental and theoretical studies of the seven-membered zirconacyclo-cumulenes have been reported [43, 45, 46]. The stability of the seven-membered zirconacyclocumulene has been ascribed to the interaction between one of the Zr t/ orbital with one terminal a orbital and the in-plane ti orbital of the cumulene, forming a type covalent bonding interaction [46]. The molecule orbital analysis shown in Fig. 8 is consistent with this conclusion (Fig. 8, HOMO-2). In HOMO-4, the Zr orbital overlaps with the sp hybridized orbital of Cl, forming the Zr-Cla covalent bond. 

Like Charges Repel It is the repulsion of the electrons in covalent bonds of the valence shell of a molecule that is central to the valence shell electron pair repulsion model for explaining molecular geometry. And, although it is not so obvious, this same factor underlies the explanations of molecular geometry that come from orbital hybridization because these repulsions are taken into account in calculating the orientations of the hybrid orbitals. 

The concepts behind the hybrid orbital model are very flexible with regard to the number of orbitals that can be combined, as shown in Table 7.2. When only one p orbital is included in the scheme, we form sp hybridized atoms, and when two p orbitals are included, the hybrid orbitals are sp. For atoms with valence electrons in the = 3 shell or higher, hybridization can also include d orbitals. Thus sp d hybridization yields five orbitals, and when six orbitals are included we have sp d hybrids. All of these schemes can be used to help explain bonding in covalent molecules. 

Thus the hydrogen molecule contains an unpaired electron in 1 s, and the oxygen atom contains unpaired electrons in 2py and 2px. In forming the water molecule, two hybrid 3p (linear) orbitals are formed, yielding two covalent H—O bonds. The length of the O—H bond is 0.99 A. (The radii of the atomic orbitals are 0.53 A for the 1 s orbital of hydrogen, 0.45 A for the 2p orbital of oxygen.)... 

In the corresponding model calculations, performed at the LSD approximation level, two Al atoms were allowed to interact with a tetradecaheptaene molecule. The Al atoms form covalent bonds with the polyene, and the Al(3p) and Al(3.v) orbitals hybridize with the molecular n orbitals [59]. The Al(3p) orbitals are found to overlap with the n orbitals of two adjacent carbon atoms in the polyene without changing the planarity of the system. The double bond character in the vicinity of the Al atom, however, is lost. 

Molecular orbital (MO) theory describes covalent bond formation as arising from a mathematical combination of atomic orbitals (wave functions) on different atoms to form molecular orbitals, so called because they belong to the entire molecule rather than to an individual atom. Just as an atomic orbital, whether un hybridized or hybridized, describes a region of space around an atom where an electron is likely to be found, so a molecular orbital describes a region of space in a molecule where electrons are most likely to be found. 

We saw in the last chapter how covalent bonds between atoms are described, and we looked at the valence bond model, which uses hybrid orbitals to account for the observed shapes of organic molecules. Before going on to a systematic study of organic chemistry, however, we still need to review a few fundamental topics. In particular, we need to look more closely at how electrons are distributed in covalent bonds and at some of the consequences that arise when the electrons in a bond are not shared equally between atoms. 

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. 

Remember from Chapter 6 that energy is released when a bond forms. Consequently, atoms that form covalent bonds tend to use all their valence s and p orbitals to make as many bonds as possible. We might expect the S p -hybridized aluminum atom to form a fourth bond with its unused 3 p orbital. A fourth bond does not form in A1 (C2 115)3 because the carbon atoms bonded to aluminum have neither orbitals nor electrons available for additional bond formation. The potential to form a fourth bond makes triethylaluminum a very reactive molecule. 

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. 

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. 

---