Lewis Structures Continued

A shortcut method to determine the number of bonds in a Lewis structure is given in the Student Study Guide/Solutions Manual (page 1-4). 

Formal Charge Observed with Common Bonding Patterns for C, N, and O 

Atom Number of valence electrons +1 Formal charge 0 -1 

The discussion of Lewis structures concludes with the introduction of isomers and exceptions to the octet rule. 

In drawing a Lewis structure for a molecule with several atoms, sometimes more than one arrangement of atoms is possible for a given molecular formula. For example, there are two acceptable arrangements of atoms for the molecular formula C2H6O. 

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In practice, the NBO program labels an electron pair as a lone pair (LP) on center B whenever cb 2 > 0.95, i.e., when more than 95% of the electron density is concentrated on B, with only a weak (<5%) delocalization tail on A. Although this numerical threshold produces an apparent discontinuity in program output for the best single NBO Lewis structure, the multi-resonance NRT description depicts smooth variations of bond order from uF(lon) = 1 (pure ionic one-center) to bu 10n) = 0 (covalent two-center). This properly reflects the fact that the ionic-covalent transition is physically a smooth, continuous variation of electron-density distribution, rather than abrupt hopping from one distinct bond type to another. 

In all the cases presented so far, an adequate picture of bonding was obtained by making the approximation that each MO is localized on only two atoms. Sometimes, however, these localized MOs are not a very good model for certain bonds. The situations in which localized MOs fail are the same ones in which normal Lewis structures are inadequate—situations in which resonance is needed to describe the bonding. In such molecules it is convenient to use delocalized MOs (MOs that include AOs from more than two atoms) to describe the bonding. In fact, resonance and delocalized MOs are just different ways to describe the same type of bonds. It is not necessary to delve deeply into delocalized MOs, but a brief discussion of them will provide a much better understanding of resonance and when it should be used. As we continue, we will find the resonance picture adequate for most of our discussions. 

At this point the dispute continues about which position is correct. For second-row elements where the octet rule is never exceeded it seems clear that the best Lewis structures conform to the following rules ... 

Appendix A Lewis Structures of Common FuncUonal Groups continued)... 

The forms shown place the negative charge on two of the oxygen atoms. Continue with a Lewis structure that places the charge on the third oxygen ... 

Stock continued these studies between 1912 and 1936 and succeeded in isolating diborane (B2Hg). Tbe formulas of diborane and the higher boron hydrides do not fit the comfortable Lewis-Langmuir electron pair structures. For example, BH3 (the monomer of diborane) has a total of only 6 valence shell electrons. Diborane is electron deficient by two electrons with no nice Lewis structure. It would be some decades... 

We continue to practice writing Lewis structures for molecules and ions and use formal charges to study the distribution of electrons in these species. (9.6 and 9.7)... 

Each Kekule structure makes an equal contribution to the hybrid thus, the C—C bonds are neither single nor double bonds, but something intermediate. We recognize that neither of these contributing structures exists (they are merely alternative ways to pair 2p orbitals with no reason to prefer one over the other) and that the actual structure is a superposition of both. Nevertheless, chemists continue to use a single contributing structure to represent this molecule because it is as close as we can come to an accurate structure within the limitations of classical Lewis structures and the tetravalence of carbon. 

Dissociation (Ionization) Constants (Kj,) of Selected Amine Bases (continued) Name and Formula Lewis Structure ... 

Although molecular orbital theory is in many ways the most powerful of the bonding models, it is also the most complex, so we continue to use the other models when they do an adequate job of explaining or predicting the properties of a molecule. For example, if you need to predict the three-dimensional shape of an AB molecule on an exam, you should draw its Lewis structure and apply the VSEPR model. Don t try to draw its molecular orbital diagram. On the other hand, if you need to determine the bond order of a diatomic molecule or ion, you should draw a molecular orbital diagram. In general chemistry, it is best to use the simplest theory that can answer a particular question. 

We recall that the length of carbon-carbon O bonds depends on the hybridization of both carbon atoms. The bond length between the sp -hybridized methyl carbon atom and the sp -, sp -, and sp-hybridized central carbon atoms of propane, propene, and propyne, show this effect. The bond length for sp -sp bonded atoms, 151 pm, is 3 pm shorter than that of sp -sp bonded atoms. If there were no double bond character between C-2 and C-3 in butadiene, we might expect a bond length of 148 pm. However, the bond length of 1,3-butadiene is 146 pm. The molecular orbital model predicts a shorter bond length than would be expected from a Lewis structure because of the continuous overlap... 

Age-old questions concerning the nature of the bonds between atoms in molecules culminated in the remarkable Lewis structure model of G. N. Lewis (1916). The notion that such bonds were formed from directed hybrids was subsequently developed by Linus Pauling (1932), shortly after the discovery of quantum mechanics. Although many theoretical advances have ensued, it is fair to say that the underlying concepts of valence-shell hybridization, shared-electron pair bonds, and Lewis structural dot diagrams continue to dominate chemical thinking and pedagogy to this day. 

In this and many cases, we shall not be surprised to discover that the electronic propensities of highly distorted excited-state species exhibit strong NBO/NRT analogies to those studied in earlier chapters. The simple example of acrolein suggests how Lewis structural concepts continue to yield rich explanatory dividends as NBO-based tools are employed to penetrate ever deeper into the excited-state domain. 

We could continue applying ideas introduced in this section, but our ability to write plausible Lewis structures will be greatly aided by a couple of new ideas that we introduce in Section 10-3. 

The octet rule has been our mainstay in writing Lewis structures, and it will continue to be one. Yet at times, we must depart from the octet rule, as we will see in this section. 

The controversy as to how best to write Lewis structures will no doubt continue in the chemical literature, but you should not be too dismayed by this situation. Our approach to depicting the electronic structure of a molecule is based on the simplest Lewis structure and its concomitant use in determining the shape of a molecule through VSEPR theory. In order to probe more deeply into the nature of a chemical bond—for example, to understand experiment results, such as bond enthalpy values— we must analyze a computed electron density map for that molecule rather than rely just on the Lewis structure. 

A flow battery is a battery in which materials (reactants, products, electrolytes) pass continuously through the battery. The battery is simply a converter of chemical to electrical energy. Formal charge is the charge an atom would acquire if the bonding electrons in each bond were divided equally between the two atoms involved. It is the number of outer-shell (valence) electrons minus the number of electrons assigned to that atom in a Lewis structure. 

The beginning of the twentieth century also marked a continuation of studies of the structure and properties of electrolyte solution and of the electrode-electrolyte interface. In 1907, Gilbert Newton Lewis (1875-1946) introduced the notion of thermodynamic activity, which proved to be extremally valuable for the description of properties of solutions of strong electrolytes. In 1923, Peter Debye (1884-1966 Nobel prize, 1936) and Erich Hiickel (1896-1981) developed their theory of strong electrolyte solutions, which for the first time allowed calculation of a hitherto purely empiric parameter—the mean activity coefficients of ions in solutions. 

Methylpropene can be made to continue the process to yield high polymers—cationic polymerisation—but most simple alkenes will go no further than di- or tri-meric structures. The main alkene monomers used on the large scale are 2-methyIpropene (— butyl rubber ), and vinyl ethers, ROCH=CH2 (— adhesives). Cationic polymerisation is often initiated by Lewis acid catalysts, e.g. BF3, plus a source of initial protons, the co-catalyst, e.g. traces of HzO etc. polymerisation occurs readily at low temperatures and is usually very rapid. Many more alkenes are polymerised by a radical induced pathway, however (p. 320). 

Because a review of previous work revealed that C02 could act as a monomer in some polymerizations catalyzed by Lewis acids [141], for our systems it was important to demonstrate that the supercritical C02 being employed as the continuous phase was not being incorporated into the backbone of the polymer chain. Spectral analysis consisting of H and 13C NMR as well as infrared spectroscopy demonstrated that no differences existed in the structure of the polymers prepared in hexane and those prepared in C02, proving that the C02 was acting as an inert solvent in these polymerizations and was not acting as a monomer [139],... 

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