Covalent bonds resonance descriptions

The difference in electronegativity of chlorine and phosphorus is 0.9, which corresponds to 18% of partial ionic character. Accordingly, an alternative description of the PCI5 molecule is that the phosphorus atom forms four covalent bonds, using only the four orbitals of the outer shell, and one ionic bond to Cl , and that the four covalent bonds resonate among the five positions, so that each chlorine atom is held by a bond with 80% covalent and 20% ionic character. 

It is difficult to give a localized orbital description of the bonding in a period 3 hypervalent molecule that is based only on the central atom 3s and 3p orbitals and the ligand orbitals, that is, a description that is consistent with the octet rule. One attempt to do this postulated a new type of bond called a three-center, four-electron (3c,4e) bond. We discuss this type of bond in Box 9.2, where we show that it is not a particularly useful concept. Pauling introduced another way to describe the bonding in these molecules, namely, in terms of resonance structures such as 3 and 4 in which there are only four covalent bonds. The implication of this description is that since there are only four cova-... 

At shorter distances, particularly those characteristic of H-bonded and other charge-transfer complexes, the concepts of partial covalency, resonance, and chemical forces must be extended to intramolecular species. In such cases the distinction between, e.g., the covalent bond and the H-bond may become completely arbitrary. The concept of supramolecular clusters as fundamental chemical units presents challenges both to theory and to standard methods of structural characterization. Fortunately, the quantal theory of donor-acceptor interactions follows parallel lines for intramolecular and intermolecular cases, allowing seamless description of molecular and supramolecular bonding in a unified conceptual framework. In this sense, supramolecular aggregation under ambient thermal conditions should be considered a true chemical phenomenon. 

One of the simplest molecules in which it is customary to invoke outer d-orbital participation in a bonding is the triiodide ion. This ion has been observed with a large number of different cations and X-ray crystal studies have revealed both symmetrical and unsymmetrical species, although in both forms it is essentially linear. If the bonding involves only the valence p-orbitals then the Hiickel orbitals for the symmetric species are those shown in Fig. 12. This description is exactly equivalent to the covalent-ionic resonance formulation VII. 

Instead of using this description of the bond as involving resonance between an extreme covalent bond H C1 and an extreme ionic bond H+Cl-, we may describe the bond as a covalent bond with partial ionic character, and make use of the valence line, writing H—Cl (or H—Cl )... 

The theory of resonance has been applied to many problems in chemistry. In addition to its use in the discussion of the normal covalent bond (involving the interchange of two electrons, with opposed spins, between two atoms) and to the structure of molecules for which a single valence-bond structure does not provide a satisfactory description, it has rendered service to chemistry by leading to the discovery of several... 

In extreme cases a multiple-scattering, sharp resonant structure can result in which the electron is in a quasi-bound state (155). One example is the white line, which is among the most spectacular features in X-ray absorption and is seen in spectra of covalently bonded materials as sharp ( 2eV wide) peaks in absorption immediately above threshold (i.e., the near continuum). The cause of white lines has qualitatively been understood as being due to a high density of final states or due to exciton effects (56, 203). Their description depends upon the physical approach to the problem for example, the LiUii white lines of the transition metals are interpreted as a density-of-states effect in band-structure calculations but as a matrix-element effect in scattering language. 

Initially, hypervalent molecules like PCI5 and SFe were described in terms of sp d" hybrid orbitals, in an extension of the sp" orbital description used to characterize molecules that obey the octet rule. However, ab initio calculations have shown that d orbitals play only a minor role in the bonding of hypervalent molecules. An alternative approach proposed by Pauling invoked combinations of resonance stmctures involving four covalent bonds and one or more additional ionic bonds, as shown in Scheme 1 for PF5. 

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. 

Resonance attempts to correct a fundamental defect in Lewis formulas. Lewis formulas show electrons as being localized they either are shared between two atoms in a covalent bond or are unshared electrons belonging to a single atom. In reality, electrons distribute themselves in the way that leads to their most stable arrangement. This sometimes means that a pair of electrons is delocalized, or shared by several nuclei. What we try to show by the resonance description of ozone is the delocalization of the lone-pair electrons of one oxygen and the electrons in the double bond over the three atoms of the molecule. Organic chemists often use curved arrows to show this electron... 

Notice that l.lla describes the bonding in H2 in terms of a localized two-centre two-electron, 2c-2e, covalent bond. A particular resonance structure will always indicate a localized bonding picture, although the combination of several resonance structures may result in the description of the bonding in the species as a whole as being delocalized (see Section 4.3). 

The hydrogen molecule has provided an example of covalent-ionic resonance in a particular bond. Because structures (3-IVb) and (3-IVc) are of importance in an accurate description of the bond from the VB point of view, we say that the bond has some ionic character. However, the polarity that (3-Vb) introduces is exactly balanced by the polarity that (3-Vc) introduces, so that the bond has no net polarity. It is therefore called a nonpolar covalent bond. It is important not to confuse polarity and ionic character, although, unfortunately, the literature contains many instances of such confusion. When we turn to a heteronuclear diatomic molecule, we necessarily have bonds that have both ionic and polar character. Even for the pure covalent canonical structure of HC1 (3-Ia) there is bond polarity... 

A description of covalent bonding mainly through overlap-dependent resonance integrals... 

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