Lewis theory double bonds

Lewis acids can greatly accelerate the cycloaddition. Instructive examples are the AlQs-catalyzed reaction of cycloalkenones with 1,3-butadienes [12]. The catalytic effect is explained by FMO theory considering that the coordination of the carbonyl oxygen by Lewis acid increases the electron-withdrawing effect of the carbonyl group on the carbon-carbon double bond and lowers the LUMO dienophile energy. 

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 ... 

Pure activated aluminas are also capable of catalyzing the skeletal isomerization of olefins (104, 105), but at considerably higher temperatures (350°-400°C) than those required for double-bond isomerization. The results obtained by Pines and Haag (105,106) leave little doubt that this type of isomerization is acid catalyzed. They found that (a) skeletal isomerization of cyclohexane or 3,3-dimethylbutene-l over pure alumina was poisoned upon ammonia addition and (b) the order of appearance of products from 3,3-dimethylbutene-l isomerization as contact time is increased was that predicted from carbonium ion theory. They also used indicator tests to show that the seat of acid activity in -y-alumina consists of Lewis, not Br0nsted, acidity. Independent infrared studies of pyridine chemisorbed on pure alumina have verified the existence of Lewis acidity and the absence of Brpnsted acidity in pure alumina (23, 107). 

Molecular orbital theory predicts that O2 is paramagnetic, in agreement with experiment. Note that the Lewis structure of O2 does not indicate that it has two unpaired electrons, even through it does imply the presence of a double bond. In fact, the prediction/confirmation of paramagnetism in O2 was one of the early successes of molecular orbital theory. Also, the ions 0+ (dioxygen cation), Oj (superoxide anion), and 0 (peroxide anion) have bond orders 2V2, U/2, and 1, respectively. The experimental energy levels of the molecular orbital for the O2 molecule are shown in Fig. 3.3.3(b). 

Particularly relevant to the present crmtext is the fact that the olefinic double bond is considered as a soft base in Pearson s theory, while many Lewis acids used in cationic polymerisation (BF3, BCI3, AICI3, etc.) are classed as hard acids. Obviously, n-acceptors like chloranil or tetracyanoethylene are considered as soft acids. Thus, the interactions between Lewis acids and olefins must be considered as very weak in the context of the HSAB theory. This prediction is well substantiated by the tenuous character of the complexes observed in experimental studies (see Chap. IV). On the other hand, carbenium ions are usually placed at the borderline between hard and soft acids and are definitely softer than the Lewis acids mentioned above. Consequently, their interactions with olefins must be rather strong, which suggests that that propagation in cationic polymerisations promoted by Lewis acids should be faster than initiation. 

In the LCAO MO description, the H2 molecnle in its ground state has a pair of electrons in a bonding MO, and thus a single bond (that is, its bond order is 1). Later in this chapter, as we describe more complex diatomic molecules in the LCAO approximation, bond orders greater than 1 are discussed. This quantum mechanical definition of bond order generalizes the concept first developed in the Lewis theory of chemical bonding—a shared pair of electrons corresponds to a single bond, two shared pairs to a double bond, and so forth. 

Figure 9-11 Representations of the bonding in the benzene molecule, CgHg. (a) Lewis formulas of the two valence bond resonance structures, (b) The six p orbitals of the benzene ring, shown overlapping to form the (hypothetical) double bonds of the two resonance forms of valence bond theory, (c) In the MO description the six electrons in the pi-bonded region are dehcalized, meaning they occupy an extended pi-bonding region above and below the plane of the six C atoms.

From the Lewis structure of the carbon dioxide molecule (Figure 2- 18a) it is seen that the carbon atom is surrounded by two electron groups (two double bonds). Two electron groups mean that there is a need for two identical orbitals 180° apart according to the VSEPR theory and Table 2- 1 on page 70. The carbon atom solves this problem by forming two identical so-called sp-hybrid orbital. As the name sp indicated these orbitals are made from one s-orbital and one p-orbital. 

With O2, we really see the power of MO theory compared to theories based on localized orbitals. For years, it seemed impossible to reconcile bonding theories with the bond strength and magnetic behavior of O2. On the one hand, the data show a double-bonded molecule that is paramagnetic. On the other hand, we can write two possible Lewis structures for O2, but neither gives such a molecule. One has a double bond and all electrons paired, the other a single bond and two electrons unpaired ... 

Two types of bonds, namely, ionic and nonionic (covalent) bonds, were recognized early on in the formulation of the electronic theory of valency. Lewis made clear distinctions between ionic and covalent bonds. The first are formed by transfer of electrons and production of separate charged ions (as explained by Kossel, 1916), the second, according to Lewis, by sharing of electrons in pairs, a single bond consisting of one shared pair, a double bond of two, and a triple bond of three. 

Ans. The Lewis structure for ethylene is H—C=C—H. The VSEPR theory treats double and triple bonds as though they were single bonds. Therefore we can conclude that there are three bonding pairs around each carbon atom to be accommodated in some symmetrical structure. That structure is the equilateral triangle. VSEPR theory predicts that ethylene consists of two equilateral triangles connected at one corner. The carbon atoms are connected by a single line which we know to be the double bond. The angles between all atoms are equal, and are 120 °. Because of the symmetry of the molecule, it has no dipole moment. 

Houk s model (Fig. 3 [206,207]) assiunes that the factors that are responsible for the non-planarity of the norbornene double bond intervene in the transition states of the cycloadditions of 5-cis-butadiene moieties grafted at C(2), C(3) of norbornane and 7-oxanorbornane systems. The tighter the transition states, the higher the endo face selectivity. The less synchronous the cycloaddition (e.g., with non-symmetrical dienophiles coordinated to Lewis acids), the lower the endo face preference, steric factors competing in favor of the exo mode of addition. This theory is verified for a large number of cycloadditions including those we reported for deuterated 2,3-dimethylidenebicyclo [2.2.1] heptane [208]. 

Resonance theory not only explained structure, but stability as well. If a molecule (or ion) is represented by two (as in benzene) or more equivalent resonance contributors, extra stability is the result. Pauling rationalized benzene s amazing resistance to addition reactions to its C=C bonds by invoking resonance stabilization. The carbonate ion (C03 ) is commonly found in minerals. Despite a Lewis strucmre that suggests one short double bond and two longer single bonds, the ion has threefold symmetry... 

This theory of resonance was particularly useful in establishing the structure of benzene, which had been puzzling in Kekule s day (see page 100) and which had retained questionable points ever since. As usually drawn, the structure of benzene is that of a hexagon with alternating single bonds and double bonds. By the Lewis-Langmuir system, two-electron pools and four-electron pools alternated. Benzene, however, lacked almost completely the characteristic properties of other compounds which contained double bonds, or four-electron pools. 

Bonding in O2 provides an interesting test case for molecular orbital theory. The Lewis structure for this molecule shows a double bond and complete pairing of electrons ... 

The short O — O bond distance (1.21 A) and relatively high bond enthalpy (495 kj/mol) are in agreement with the presence of a double bond. However, Figure 9.43 tells us that the molecule contains two unpaired electrons and should therefore be paramagnetic, a detail not discernible in the Lewis structure. The paramagnetism of O2 is demonstrated in Figure 9.45, which confirms the prediction from MO theory. The MO description also correctly predicts a bond order of 2 as did the Lewis structure. 

Following is another example of adding a proton. Here, the proton is added across the pi bond of the C — C double bond. The compounds below are labeled as proton donor and proton acceptor, terms used to describe Br0nsted acids and bases. They can also be labeled according to Lewis acid-base theory as electrophile and nucleophile. ... 

The theories of HLPS might be called electron-pairing theories if Lewis is called an electron-pair theory. It should also be pointed out that the HLPS electron pair differs considerably from Lewis conception of the electron-pair bond, in that the electrons are much less closely associated in this respect it approaches the truth much more closely than does Lewis conception - Pauling and Slater consider a double bond to be merely two ordinary single bonds sticking out from each atom in different directions, and treat the triple bond in a similar way. In this way they do not agree very well with Lewis, nor do they agree with the results obtained from molecular orbital theory. 

Because the 2p orbitals have a node at the atomic nucleus, the electron distribution of the VB wavefunction formed by their side-on overlap does not have cylindrical symmetry, but instead changes sign when rotated by 180° about the bond axis. Bonds that have this property are called pi (tt) bonds. Thus, the double bond in the oxygen molecule consists of a sigma bond and a pi bond, which are not equivalent. The existence of two different types of bonds in double bond formation is something that is not predicted by simple Lewis theory. A triple bond such as that in N2 is described within the VB approach as a sigma bond formed from the head-on overlap of 2p orbitals and two pi bonds formed by the overlap of the 2py and 2p orbitals on one N atom with their counterparts on the other atom (Figure 3.7). 

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