Lewis pairs

Where FCl is the solute gas-liquid partition coefficient, r is the tendency of the solvent to interact through k- and n-electron pairs (Lewis basicity), s the contribution from dipole-dipole and dipole-induced dipole interactions (in molecular solvents), a is the hydrogen bond basicity of the solvent, b is its hydrogen bond acidity and I is how well the solvent will separate members of a homologous series, with contributions from solvent cavity formation and dispersion interactions. 

Electrophiles molecules or ions that can accept an electron pair => Lewis acids. 

Geometry Number of lone pairs Number of bonding pairs Lewis dot structure Shape ... 

Almost simultaneous with the publication of Kossel s paper there appeared a rival electronic theory. The American chemist Lewis introduced the idea of the covalent electron-pair bond. Like Kossel, he was impressed by the apparent stability of the noble gas configuration. He was also impressed by the fact that, apart from many compounds of the transition elements, most compounds when rendered as molecules have even numbers of electrons, suggesting that electrons are usually found in pairs. Lewis devised the familiar representations of molecules and polyatomic ions (Lewis structures, or Lewis diagrams) in which electrons are shown as dots (or as noughts and crosses) to show how atoms can attain noble gas configurations by the sharing of electrons in pairs, as opposed to complete transfer as in Kossel s theory. It was soon apparent from the earliest X-ray studies that Kossel s theory was more appropriate... 

When atoms possess an incomplete outer shell (e.g., nonpaired electrons), yet their net charge is zero, attraction between such atoms takes place because of their strong tendency to complete their outer electron orbital shell by sharing their unpaired electrons. This gives rise to a covalent bond. One example of a covalent bond is the bimolecular chlorine gas (Cl2) (Fig. 1.1). Covalent bonding is a characteristic of some nonmetals or metalloids (bimolecular molecules), but may also arise between any two atoms when one of the atoms shares its outer-shell electron pair (Lewis base) with a second atom that has an empty outer shell (Lewis acid). Such bonds are known as coordinated covalent bonds or polar covalent bonds. They are commonly weaker than the covalent bond of two atoms which share each other s unpaired outer-shell electrons (e.g., F2 and 02). Coordinated covalent bonds often involve organometallic complexes. 

Brensted-Lowry Acids and Bases and Conjugate Pairs Lewis Acids and Bases Titration and Neutralization Hydrolysis... 

Thus, the bonding in metal compounds and complexes has traditionally been viewed as ionic, with a positive metal center interacting with negative ions, such as HO , 0 , Cl , AcO , and coordinate donor, Lewis acid-base interactions with a positive metal center interacting with negative ions and electron-pair Lewis bases, such as iNHs, iPPhs, HOH. Examples of ionic versus covalent bonding illustrate the tradition H+CH versus H-Cl (ls-3p), C +(C1 )4 versus C(-C1)4 [2sp -(3p)4], Fe +(C1 )3 versus Fe(-Cl)3 [3d sp2-(3p)3], H+-OH versus H-OH (ls-2p), C +(Q2-)2 versus 0=C=0 [2sp -(2p )2], and Fe +(0 ) versus Fe=0 [3d sp-(2p )]. Such ionic formulations for these molecules in an inert matrix are not consistent with their physical... 

The Lewis concept of acids and bases (G. N. Lewis, 1923) interprets the combination of acids with bases in terms of the formation of a coordinate covalent bond. A Lewis acid can accept and share a lone pair of electPDns donated by a Lewis base. Because protons readily attach themselves to lone electron pairs, Lewis bases are also Biyinsted bases. Lewis acids, however, include a large number of substances in addition to proton donors for examjjle, metal ions, acidic oxides, or atoms. 

Table 2.11 lists the principal types of solid base catalysts. We should remember, however, that base catalyst is a relative definition and thus the materials listed in Table 2.11 do not necessarily function as a base in all cases. Some of these materials may act as an acid if the reactants are strongly basic. The terms, acid and base, should be used according to the function. The materials may be called solid base catalysts only if acting as a base toward the reactants by abstraction of a proton (Bronsted base) or by donation of an electron pair (Lewis base) to form anionic intermediates that undergo catalytic cycles. 

Use an example to illustrate each of the following terms lone pairs, Lewis stmcture, the octet rule, bond length. 

Plan We examine the formulas to see which species accepts the electron pair (Lewis acid) and which donates it (Lewis base) in forming the adduct. 

According to Lewis s octet rule, each atom should be surrounded by four pairs of electrons, either shared or free pairs. Lewis derived stmctures for halogen molecules, the ammonium ion, and oxy acids, inexplicable according to previous valence theories. He viewed polar bonds as unequally shared electron pairs. Because the complete transfer of electrons was only an extreme case of polarity, he abandoned his earlier dualistic view the polar theory was just a special case of his more general theory. 

An L-function ligand is one which interacts with a metal center via a dative covalent bond (i.e. a coordinate bond), in which both electrons are donated by the L ligand. As such, an L-function ligand donates two electrons to a metal center. Since the metal uses no electrons in forming the M-L bond, an L-function ligand does not influence the valence of a metal center. Simple examples of L-type ligands include R3P, R2O, and CO, i.e. donor molecules that have lone pairs (Lewis bases). 

Lewis acid-base theory (Jensen, 1980) is an outgrowth of the Lewis model of chemical bonds. A Lewis acid is a chemical species that can accept an electron pair. Lewis acids can be cations like Fe or Cu or they can be species with empty or partially empty valence orbitals such as CO2 or SO2. Lewis bases can donate an electron pair. Lewis bases are anions like OH" or S or they can be species with lone pairs such as HjO or NHj. The transfer of cations from a solid, such as szomolnokite (FeS04 H20), to form a hydrated ferrous ion in solution is a typical Lewis acid-base reaction. 

Sections 4.1.2 and 4.1.3 compared complex oxides and halides with the binary (parent) compounds. In this section we attempt to show on a semiquantitative basis how complex oxides and halides favor high oxidation states, primarily because they represent acid-base reactions in which the presence of a basic oxides (e.g., BaO) or halide (e.g., CsCl) donates electron pairs (Lewis basicity) to the acidic (high charge density) f-element ion. 

There are some cases in homogeneous anionic polymerization in which the initiator dissociates completely with quantitative transformation into the active ioitic form and the process is also virtually instantaneous stoichiometric polymerization). This is the case, for example, when one uses, as initiators, alkali orgaiuc compounds (e.g., phenyUithium, butylhthium, or sodium naphthalene) in solvents which have unshared electron pairs (Lewis bases). The alkah forms stable positively charged complex ions with the Lewis base (Lenz, 1967), while the organic residue becomes negatively charged (carbanion) and can initiate an ioiuc polymerization [cf. Eqs. (8.11) and (8.12)] ... 

It is known that the lone pair electrons of amines are capable of capturing a carbon radical long enough to convert it to a nonradical species, as is the case with hindered amine light stabilizers (HALS) [70], It stands to reason that species with available lone pairs (Lewis bases) stabilize the TFE biradicaloid, allowing the end opposite to the catalyst to exist as a full radical and subsequently react. Barabanov s [23] observation of polymer formation is evidence for this terminal radical. Other perfluorinated species capable of forming biradicaloids would undergo similar radical stabilization. 

Traditionally, Lewis acids have been employed to facilitate the opening of similar oxabicyclic intermediates and we thought that this case might react similarly. With this in mind, we began to methodically pair Lewis acids with a series of trialkylphosphine reagents. Initial results were poor, with a variety of Lewis acid/phosphine combinations still falling short with no... 

Lewis s definition of an acid is also straightforward. Clearly, it has to be an entity that can attach to a base by sharing the latter s lone pair. Lewis proposed that that ability alone is sufficient for a species to be regarded as an acid. That is, according to Lewis, an acid is any species that can stick to a lone pair of electrons. Notice that the proton is not mentioned any suitably accepting species can be called an acid the proton is only one of a myriad possibilities. 

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