Bonding resonance Lewis structures

Our treatment of O2 shows that the extra complexity of the molecular orbital approach explains features that a simpler description of bonding cannot explain. The Lewis structure of O2 does not reveal its two unpaired electrons, but an MO approach does. The simple (t-tt description of the double bond in O2 does not predict that the bond in 2 is stronger than that in O2, but an MO approach does. As we show in the following sections, the molecular orbital model has even greater advantages in explaining bonding when Lewis structures show the presence of resonance. 

IN the past twenty years the electronic structures of many organic molecules, particularly benzene and related compounds, have been discussed in toms of the molecular orbital and valence bond methods.1 During the same period the structures of inorganic ions have been inferred from the bond distances f a bond distance shorter than the sum of the conventional radii has been attributed to the resonance of double bonded structures with the single bonded or Lewis structure. 

Alongside electron-pair bonds, odd-electron bonds play an important role in chemistry, and constitute therefore a compulsory test case for any computational method. Odd-electron bonds can be represented as two resonating Lewis structures that are mutually related by charge transfer, as shown in (13) for two-center, one-electron (2c,le) bonds and in (14) and (15) for typical two-center, three-electron (2c,3e) bonds. 

Acetonitrile (CH3CN) has a pKg of 25, making it more acidic than many other compounds having only C - H bonds. Draw Lewis structures for acetonitrile and its conjugate base. Use resonance structures to account for the acidity of acetonitrile. 

Blood Pressure and the Sodium lon/Potassium Ion Ratio Lewis Structures of Polyatomic Ions Lewis Structure, Stability, Multiple Bonds, and Bond Energies Lewis Structures and Resonance Lewis Structures and Exceptions to the Octet Rule... 

When electrons become delocalized, it s important to remember that all the resonant Lewis structures are just partial pictures of how the molecule may appear. Just as the ozone molecule can be drawn with the double bond left or right, it can more correctly be drawn with a partial double bond to both oxygen atoms as shown on the right. 

Characteristics of Lewis Structures In addition to covalent bonds, a Lewis structure also shows lone pairs, which are pairs of electrons not involved in bonding, on atoms and formal charges, which are the result of bookkeeping of electrons used in bonding. A resonance structure is one of two or more Lewis structures for a single molecule that cannot be described fiiUy with only one Lewis structure. 

Bond Energy and Enthalpy The Localized Electron Bonding Model Lewis Structures Exceptions to the Octet Rule Resonance... 

If one wishes to use only Heitler-London wave-functions for all electron-pair bonds, as indicated above when considering Coulson-Fischer A-B orbitals, we can still speak of the delocalization of a Y electron of structure (4) into the antibonding A-B orbital. When this is done, we obtain Lewis structures (14) and (15), with configurations (y) (a) (b) and (y) (a) (b) that involve Heitler-London formulations of the wave-functions for the Y-A and long Y-B bonds. Resonance between structures (14) and (15) is equivalent to the utilization of increased-valence structure (1) (with a Heitler-London type wave-function for the fractional Y-A bond). 

To deal with circumstances such as the bonding in ozone, the notion of resonance between Lewis structures was developed. According to the resonance concept, when more than one Lewis structure may be written for a molecule, a single structure is insufficient to describe it. Rather, the true structure has an electron distribution that is a hybrid of all the possible Lewis structures that can be written for the molecule. In the case of ozone, two equivalent Lewis structures may be written. We use a double-headed aiTow to represent resonance between these two Lewis structures. 

Aeeording to your Lewis structure(s) and to the actual geometry of the molecule, is the bonding in planar corannulene fully delocalized (as in benzene), or are some CC bonds long and some short Do your results support the notion that planar corrannulene is resonance stabilized Explain. 

Mechanistically the 1,3-dipolar cycloaddition reaction very likely is a concerted one-step process via a cyclic transition state. The transition state is less symmetric and more polar as for a Diels-Alder reaction however the symmetry of the frontier orbitals is similar. In order to describe the bonding of the 1,3-dipolar compound, e.g. diazomethane 4, several Lewis structures can be drawn that are resonance structures ... 

Strategy Write a Lewis structure for the N02 ion, following the usual steps. Then write the other resonance form by changing the position of the multiple bond. Do not change the skeleton structure. 

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. 

Because all three bonds are identical, a better model of the nitrate ion is a blend of all three Lewis structures with each bond intermediate in properties between a single and a double bond. This blending of structures, which is called resonance, is depicted in (9) by double-headed arrows. The blended structure is a resonance hybrid of the contributing Lewis structures. A molecule does not flicker between different structures a resonance hybrid is a blend of structures, just as a mule is a blend of a horse and a donkey, not a creature that flickers between the two. 

I Stratospheric ozone, O, protects life on Farth from harmful ultraviolet radiation from the Sun. Suggest two Lewis structures that contribute to the resonance structure for the 02 molecule. Experimental data show that the two bond lengths are the same. 

STRATEGY Write a Lewis structure for the molecule by using the method outlined in Toolbox 2.1. Decide whether there is another equivalent structure that results from the interchange of a single bond and a double or triple bond. Write the actual structure as a resonance hybrid of these Lewis structures. 

Self-Test 2.7A Write Lewis structures contributing to the resonance hybrid for the acetate ion, CH ,CC)2. The structure of CH COOH is described in Example 2.4 the acetate ion has a similar structure, except that it has lost the final H atom while keeping both electrons from the O -H bond. 

Write the Lewis structure, including typical contributions to the resonance structure (where appropriate, allow for the possibility of octet expansion, including double bonds in different positions), for (a) sulfite ion (b) hydrogen sulfite ion (c) perchlorate ion (d) nitrite ion. 

Draw the most important Lewis structure for each of the following ring molecules (which have been drawn without showing the locations of the double bonds). Show all lone pairs and nonzero formal charges. If there are equivalent resonance... 

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 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... 

Which of these options is the best Lewis structure Actually, no single Lewis structure by itself is an accurate representation of NO3. Any single structure of the anion shows nitrate with one NDO double bond and two N— O single bonds. In Section 9 1, we show that single and double bonds between the same types of atoms have different lengths and different energies. In contrast, experiments show that the three nitrate N—O bonds are identical. To show that the nitrate N—O bonds are all alike, we use a composite of the three equivalent Lewis structures. These are traditionally called resonance structures. Resonance stmctures are connected by double-headed arrows to emphasize that a complete depiction requires all of them. 

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