Lewis structure theory

Although the most naive form of valence-bond and Lewis-structure theory would not predict the paramagnetism of O2, the VB-like NBO donor-acceptor perspective allows us to develop an alternative localized picture of general wavefunctions, including those of MO type. Let us therefore seek to develop a general NBO-based configurational picture of homonuclear diatomics to complement the usual MO description. 

Resonance theory tells us that molecules which cannot be adequately represented in terms of a single Lewis structure are likely to be unusually stable. What the simple theory does not tell us is the magnitude of the effect, the so-called resonance energy. This can be assessed via molecular modeling. 

Gilbert Newton Lewis (1875-1946 was born in Weymouth, Massachusetts, and received his Ph.D. at Harvard in 1899. After a short time as professor of chemistry at the Massachusetts Institute of Technology (1905-1912), he spent the rest of his career at the University of California at Berkeley (1912-1946). In addition to his work on structural theory, Lewis was the first to prepare heavy water," D20, in which the two hydrogens of water are the 2H isotope, ceuterium. 

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. 

The Lewis structures encountered in Chapter 2 are two-dimensional representations of the links between atoms—their connectivity—and except in the simplest cases do not depict the arrangement of atoms in space. The valence-shell electron-pair repulsion model (VSEPR model) extends Lewis s theory of bonding to account for molecular shapes by adding rules that account for bond angles. The model starts from the idea that because electrons repel one another, the shapes of simple molecules correspond to arrangements in which pairs of bonding electrons lie as far apart as possible. Specifically ... 

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

Lewis s theory also fails to account for the compound diborane, B2H6, a colorless gas that bursts into flame on contact with air. The problem is that diborane has only 12 valence electrons (three from each B atom, one from each H atom) but, for a Lewis structure, it needs at least seven bonds, and therefore 14 electrons, to bind the eight atoms together Diborane is an example of an electron-deficient compound, a compound with too few valence electrons to be assigned a valid Lewis structure. Valence-bond theory can account for the structures of electron-deficient compounds in terms of resonance, but the explanation is not straightforward. 

The boranes are electron-deficient compounds (Section 3.8) we cannot write valid Lewis structures for them, because too few electrons are available. For instance, there are 8 atoms in diborane, so we need at least 7 bonds however, there are only 12 valence electrons, and so we can form at most 6 electron-pair bonds. In molecular orbital theory, these electron pairs are regarded as delocalized over the entire molecule, and their bonding power is shared by several atoms. In diborane, for instance, a single electron pair is delocalized over a B—H—B unit. It binds all three atoms together with bond order of 4 for each of the B—H bridging bonds. The molecule has two such bridging three-center bonds (9). 

Long before their theories were supported by computations, organic chemists found a way to use resonance structures to explain the product distribution in electrophilic substitution. Thus, the Lewis structure for phenol is regarded as a resonance hybrid of the following structures ... 

Lewis, G. N., 60, 64, 398 Lewis acid, 398, 473, 671 Lewis base, 398, 473 Lewis structure, 65 writing, 67 Lewis symbol, 60 Lewis theory, limitations of, 115... 

The ab initio calculations of various three-electron hemibonded systems [122, 123] indicated that the inclusion of electron correlation corrections is extremely important for the calculation of three-electron bond energies. The Hartree-Fock (HF) error is found to be nonsystematic and always large, sometimes of the same order of magnitude as the bond energy. According to valence bond (VB) and MO theories, the three-electron bond is attributed to a resonance between the two Lewis structures... 

The many higher boranes such as B5H9 and BgH 2 are similarly electron deficient and cannot be described by a single Lewis structure. They can often be described in terms of a combination of two- and three-center bonds. Alternatively, their structures can be rationalized by electron-counting schemes such as those proposed by Wade. Analysis of the electron density of these molecules by the AIM method shows that there are bond paths between all adjacent pairs of atoms. So from the point of view of the AIM theory there are bonds between each adjacent pair of atoms, but these cannot all be regarded as Lewis two-center, two-electron bonds as is the case in B2H6. 

There are several ways in which valence bond theory is superior to Lewis structures in... 

However, we should emphasize that the NBO/NRT concepts of hybridization, Lewis structure, and resonance differ in important respects from previous empirical usage of these terms. In earlier phases of valence theory it was seldom possible to determine, e.g., the atomic hybridization by independent theoretical or experimental procedures, and instead this term became a loosely coded synonym for the molecular topology. For example, a trigonally coordinated atom might be categorized as sp2-hybridized or an octahedrally coordinated atom as d2sp3-hybridized, with no supporting evidence for the accuracy of these labels as descriptors of actual... 

As predicted by elementary hybrid bonding theory, the multiple bonds of the chemist s Lewis-structure diagram are usually found to correspond to two distinct types of NBOs (1) sigma-type, having exact or approximate cylindrical symmetry about the bond axis (as discussed in Sections 3.2.5-3.2.7), and (2) pi-type, having a nodal mirror plane passing through the nuclei 44... 

Let us first inquire whether basic criteria for the validity of low-order perturbation theory are actually satisfied in the present case. As described in Section 1.4, the perturbative starting point is an idealized natural Lewis-structure wavefunction (t//,l )) of doubly occupied NBOs. The accuracy of this Lewis-type starting point may be assessed in terms of the percentage accuracy of the variational energy (E) or density (p(l ). as shown for each molecule in Table 3.20. 

Nevertheless, even today, we often discuss the bonding of organic compounds in terms of Lewis structures and valence bond theory. 

Sometimes when writing the Lewis structure of a species, we may draw more than one possible correct Lewis structure for a molecule. The nitrate ion, N03 , is a good example. The structures that we write for this polyatomic anion differ in which oxygen has a double bond to the nitrogen. None of these three truly represents the actual structure of the nitrate ion—it is an average of all three of these Lewis structures. We use resonance theory to describe this situation. Resonance occurs when more than one Lewis structure (without moving atoms) is possible for a molecule. The individual structures are called resonance structures (or forms) and are written with a two-headed arrow (<- ) between them. The three resonance forms of the nitrate ion are ... 

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