Lewis structures delocalized electrons

Instead, structures I and II are the closest we can get to a structure for benzene within the limitations of its molecular formula, the classic rules of valence, and the fact that the six hydrogen atoms are chemically equivalent. The problem with the Kekule structures is that they are Lewis structures, and Lewis structures portray electrons in localized distributions. (With benzene, as we shall see, the electrons are delocalized.) Resonance theory, fortunately, does not stop with telling us when to expect this kind of trouble it also gives us a way out. 

Constructed from Lewis Structures Delocalization of a Lone-Pair Electron into a Vacant Bonding Orbital... 

Electron delocalization stabilizes a molecule A molecule in which electrons are delocalized is more stable than implied by any of the individual Lewis structures that may be written for it The de gree of stabilization is greatest when the contrib uting Lewis structures are of equal stability... 

Electron Delocalization in the Conjugate Base With a of —1 4 nitnc acid is almost completely ionized m water If we look at the Lewis structure of nitric acid m light of what we have said about inductive effects we can see why The N atom m nitric acid IS not only electronegative m its own right but bears a formal charge of +1 which enhances its ability to attract electrons away from the —OH group... 

Nitrate ion is stabilized by electron delocalization which we can represent m terms of resonance between three equivalent Lewis structures... 

Electron delocalization m the conjugate base usually expressed via resonance between Lewis structures increases acidity... 

The resonance effect of the carbonyl group Electron delocalization expressed by resonance between the following Lewis structures causes the negative charge in acetate to be shared equally by both oxygens Electron delocalization of this type IS not available to ethoxide ion... 

Resonance (Section 1.9) Method by which electron delocalization may be shown using Lewis structures. The true electron distribution in a molecule is regarded as a hybrid of the various Lewis structures that can be written for a molecule. 

Examine the geometry and atomic charges of 2-pyridone to see if it is localized as indicated in the drawing above, or delocalized (as in 2-hydroxypyridine). If you need to, write alternative Lewis structures to that provided above. How many n electrons does 2-pyridone possess Is 2-pyridone aromatic ... 

The process is exothermic, suggesting that the phenoxy radical is particularly stable. Display the spin density surface for phenoxy radical. Is the unpaired electron localized or delocalized over several centers Is the unpaired electron in the a or 7t system Draw appropriate Lewis structures that account for your data. 

Use geometries, electrostatic potential maps and spin densities to help you draw Lewis structures for butanal radical cation, the transition state and product. Where is the positive charge and the unpaired electron in each Is the positive charge (the unpaired electron) more or less delocalized in the transition state than in the reactant In the product ... 

In molecular orbital theory, electrons occupy orbitals called molecular orbitals that spread throughout the entire molecule. In other words, whereas in the Lewis and valence-bond models of molecular structure the electrons are localized on atoms or between pairs of atoms, in molecular orbital theory all valence electrons are delocalized over the whole molecule, not confined to individual bonds. 

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). 

Lewis structure of polvacetyl-ene can be drawn showing that the electrons may be delocalized along the carbon chain. 

Values for c in each method are obtained by solving the equation for various values of each c and choosing the solution of lowest energy. In practice, both methods give similar solutions for molecules that contain only localized electrons, and these are in agreement with the Lewis structures long familiar to the organic chemist. Delocalized systems are considered in Chapter 2. 

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. 

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