Lewis structure electron-deficient molecules

Lewis acids are thus electron-deficient molecules or ions such as BF-, or carbo-cations, whereasTewis bases are molecules or ions containing available electrons. such as amines, ethers, alkoxide ions, and so forth.,A Lewis acid-base reaction is the combination of an acid and a base to form a complex, or adduct. The stabilities of these adducts depend on the structures of the constituent acid and base and vary over a wide range. Some examples of Lewis acid—base reactions are given in Table 3.19. Lewis acid-base reactions abound in organic chemistry ... 

A stepwise process is used to convert a molecular formula into a Lewis structure, a two-dimensional representation of a molecule (or ion) that shows the relative placement of atoms and distribution of valence electrons among bonding and lone pairs. When two or more Lewis structures can be drawn for the same relative placement of atoms, the actual structure is a hybrid of those resonance forms. Formal charges are often useful for determining the most important contributor to the hybrid. Electron-deficient molecules (central Be or B) and odd-electron species (free radicals) have less than an octet around the central atom but often attain an octet in reactions. In a molecule (or ion) with a central atom from Period 3 or higher, the atom can hold more than eight electrons by using d orbitals to expand its valence shell. 

Strategy In Lewis acid-base reactions, the acid is usnaUy a cation or an electron-deficient molecule, whereas the base is an anion or a molecule containing an atom with lone pairs, (a) Draw the molecular structure for C2H5OC2H5. What is the hybridization state of A1 in AICI3 (b) Which ion is likely to be an electron acceptor An electron donor ... 

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). 

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. 

Borane has the same structure as a carbocation. The boron is sjr hybridized, with trigonal planar geometry, and has an empty p orbital. Although neutral, it is electron deficient because there are only six electrons around the boron. It is a strong Lewis acid. An electron-deficient compound often employs unusual bonding to alleviate somewhat its instability. In the case of borane, two molecules combine to form one molecule of diborane ... 

The splitting of a Cl2 molecule is an initiation step that produces two highly reactive chlorine atoms. A chlorine atom is an example of a reactive intermediate, a short-lived species that is never present in high concentration because it reacts as quickly as it is formed. Each Cl- atom has an odd number of valence electrons (seven), one of which is unpaired. The unpaired electron is called the odd electron or the radical electron. Species with unpaired electrons are called radicals or free radicals. Radicals are electron-deficient because they lack an octet. The odd electron readily combines with an electron in another atom to complete an octet and form a bond. Figure 4-1 shows the Lewis structures of some free radicals. Radicals are often represented by a structure with a single dot representing the unpaired odd electron. 

The LE model is a simple but very successful model, and the rules we have used for Lewis structures apply to most molecules. To implement this model we have relied heavily on the octet rule. So far we have treated molecules for which this rule is easily applied. However, inevitably, cases arise where the importance of an octet of electrons is called into question. Boron, for example, tends to form compounds in which the boron atom has fewer than eight electrons around it—it does not have a complete octet. Boron trifluoride (BF3), a gas at normal temperatures and pressures, reacts very energetically with molecules such as water and ammonia that have available lone pairs. The violent reactivity of BF3 with electron-rich molecules occurs because the boron atom is electron-deficient. Boron trifluoride has 24 valence electrons. The Lewis structure that seems most consistent with the properties of BF3 is... 

The common feature in such reactions is the attack of a hydride ion on a relatively electron-deficient carbon atom. The hydride may migrate from another carbon atom in the same molecule, from a carbon atom on another molecule, or even from a Lewis-type site (—Al—H ) on the catalyst. Whatever the source, the structural environment to which the hydride adds becomes more hydrogen-rich, and that from which it derives, more hydrogen-deficient. In Scheme 6, some of these reactions are tabulated. 

Not all molecules involves Lewis structures with complete octets of electrons for example BeH2 and BF3 involve four and six electron shells, respectively, a configuration that makes them reactive towards electron donors. The Working Method still applies to such electron deficient systems. 

The positively charged carbon atom in a carbocation is an extremely electron-dehcieni (electrophilic) carbon. As such, its behavior is dominated by a need to obtain an electron pair from any available source. The Sn I reaction illustrates the most obvir>"s fate of a < nr .- -t on c a.bination with an external Lewis base, forming a new bond to carbon. However, the electron deficiency of cationic carbon is so great that even under typical SnI solvolysis conditions, surrounded by nucleophilic solvent molecules, some of the cations won t wait to combine with external electron-pair sources. Instead, they will seek available electron pairs within their own molecular structures. The most available of the.se are electrons in carbon-hydrogen bonds one carbon removed from the cationic center (at llic so-called carbon) ... 

In solid Bep2, a complex network is formed with a Be atom coordination number of 4 (see Figure 3.7). BeCl2 dimerizes to a 3-coordinate structure in the vapor phase, but the linear monomer is formed at high temperatures. This monomeric structure is unstable due to the electronic deficiency at Be in the dimer and the network formed in the solid-state, the halogen atoms share lone pairs with the Be atom in an attempt to fill beryllium s valence shell. The monomer is still frequently drawn as a singly bonded structure, with only four electrons around the beryllium and the ability to accept lone pairs of other molecules to relieve its electronic deficiency (Lewis acid behavior, discussed in Chapter 6). 

The violent reactivity of BF3 with electron-rich molecules arises because the boron atom is electron-deficient. The Lewis structure that seems most consistent with the properties of BF3 (twenty-four valence electrons) is... 

Fig. 12 The extension of the Lewis ideas for hypo-valent and h3rper-valent molecules. In the former, the electron deficiency is relieved by the formation of dative bonds either from lone pairs or B-H bonds. In the latter, the hyper-valency is brought into line with the EAN rule by using Lewis bond structures where the central atom bears positive charges which compensate the negative charges on the fluorines. The results in Table 1 suggest that x is associated with the number of three-centre four-electron Xe-F bonds in these structures...

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