Lewis acids/bases strengths

Table 2-6. Some E and C parameters expressing Lewis acid/base strength according to Drago [217] ) cf Eq. (2-12).

The mechanisms of ionic and coordination polymerizations are more complex and are not as clearly understood as those of free radical polymerization. Here, we will briefly highlight the essential features of these mechanisms, and more details will be given in Chapter 7. Initiation of ionic polymerization usually involves the transfer of an ion or an electron to or from the monomer. Many monomers can polymerize by more than one mechanism, but the most appropriate polymerization mechanism for each monomer is related to the polarity of the monomers and the Lewis acid-base strength of the ion formed. 

No symbol) Softness (used with Lewis acid/ base strengths) Pearson (1963)... 

Cosolvents ana Surfactants Many nonvolatile polar substances cannot be dissolved at moderate temperatures in nonpolar fluids such as CO9. Cosolvents (also called entrainers, modifiers, moderators) such as alcohols and acetone have been added to fluids to raise the solvent strength. The addition of only 2 mol % of the complexing agent tri-/i-butyl phosphate (TBP) to CO9 increases the solubility ofnydro-quinone by a factor of 250 due to Lewis acid-base interactions. Veiy recently, surfac tants have been used to form reverse micelles, microemulsions, and polymeric latexes in SCFs including CO9. These organized molecular assemblies can dissolve hydrophilic solutes and ionic species such as amino acids and even proteins. Examples of surfactant tails which interact favorably with CO9 include fluoroethers, fluoroacrylates, fluoroalkanes, propylene oxides, and siloxanes. 

Many years ago, geochemists recognized that whereas some metallic elements are found as sulfides in the Earth s crust, others are usually encountered as oxides, chlorides, or carbonates. Copper, lead, and mercury are most often found as sulfide ores Na and K are found as their chloride salts Mg and Ca exist as carbonates and Al, Ti, and Fe are all found as oxides. Today chemists understand the causes of this differentiation among metal compounds. The underlying principle is how tightly an atom binds its valence electrons. The strength with which an atom holds its valence electrons also determines the ability of that atom to act as a Lewis base, so we can use the Lewis acid-base model to describe many affinities that exist among elements. This notion not only explains the natural distribution of minerals, but also can be used to predict patterns of chemical reactivity. 

Even acetone shows appreciable shifts following contact with metal halides - the magnitude of which correlates with the expected strength of the Lewis acid-base interaction. 

As we have seen, the Lewis theory of acid-base interactions based on electron pair donation and acceptance applies to many types of species. As a result, the electronic theory of acids and bases pervades the whole of chemistry. Because the formation of metal complexes represents one type of Lewis acid-base interaction, it was in that area that evidence of the principle that species of similar electronic character interact best was first noted. As early as the 1950s, Ahrland, Chatt, and Davies had classified metals as belonging to class A if they formed more stable complexes with the first element in the periodic group or to class B if they formed more stable complexes with the heavier elements in that group. This means that metals are classified as A or B based on the electronic character of the donor atom they prefer to bond to. The donor strength of the ligands is determined by the stability of the complexes they form with metals. This behavior is summarized in the following table. 

Steric interactions between bulky substituents such as t-Bu, leading to larger C-E-C bond angles, obviously affect the Lewis basicity caused by the increased -character of the electron lone pair. However, the strength of the Lewis acid-base interaction within an adduct as expressed by its dissociation enthalpy does not necessarily reflect the Lewis acidity and basicity of the pure fragments, because steric (repulsive) interactions between the substituents bound to both central elements may play a contradictory role. In particular, adducts containing small group 13/15 elements are very sensitive to such interactions as was shown for amine-borane and -alane adducts... 

Two such well studied systems are pyridine chemisorbed on alumina (15) and pyridine chemisorbed on silica-alumina (16). It had been previously shown that alumina contains only sites which adsorb pyridine in a Lewis acid-base fashion whereas silica-alumina has both Lewis and Bronsted acid sites. These two different kinds of sites are distinguishable by the characteristic vibrational bands of pyridine adducts at these sites (see Table I). Photoacoustic and transmission results are compared in Table II. Note that the PA signal strength depends on factors such as sample particle size and volumes of solid sample and transducing... 

An inner-sphere complex is formed between Lewis acids and bases, while an outer-sphere complex involves a water molecule interposed between the acid and the base. A hard Lewis acid is a molecular unit of small size, high oxidation state, high electronegativity, and low polarizability whereas a soft Lewis acid is a molecular unit of relatively large size, characterized by low oxidation state, low electronegativity, and high polarizability. Based on this characterization, hard bases prefer to complex hard acids, and soft bases prefer to complex soft acids, under similar conditions of acid-base strength. 

The thermodynamic tendency of a substance to act as a Lewis acid. The strength of a Lewis acid depends on the nature of the base with which the Lewis acid forms a Lewis adduct. Hence, comparative measures of Lewis acidities are given by equilibrium constants for the formation of the adducts by a common reference base. See Lewis Acid Electrophilicity Hard Acids Soft Acids Acceptor Number... 

In this article, I shall begin by showing the tremendous scope of Lewis acid-base considerations. Although it is not fully reeilized, it is very difficult to find chemical reactions in which these effects are not operative. This will be followed by a discussion of the kind of data that should be obtained and analyzed in order to learn about the strength of bonding. Since data selection is important, a good deal of space is devoted to complications that can arise from improper design of experiments and improper analysis of experimental results. 

It is hoped that the terms donor and acceptor strengths will be reserved for inferences made about Lewis acid-base properties from data in the gas phase or poorly solvating solvents. This is to be contrasted with the more complex phenomena contributing to acidity and basicity. 

These qualitative explanations, whether they be hard-soft or ionic-covalent or Class A-Class B, all suffer from the arbitrary way in which they can be employed. All Lewis acid-base type interactions are composed of some electrostatic and some covalent properties, i.e., hardness and softness are not mutually exclusive properties. Predictions are straightforward when dealing with the extremes, but with other more ambiguous systems, one can be very arbitrary in explaining results and the predictive value is impaired. What is needed is a quantitative assessment of the essential factors which can contribute to donor strength and acceptor strength. Proper combination of these parameters should produce the enthalpy of adduct formation. Until this can be accomplished, one could even question the often made assumption that the strength of the donor-acceptor interaction is a function of the individual properties of a donor or acceptor. 

Owing to the strength of the B—F bond, die BF3 complexes are of widespread use as model compounds, for investigating Lewis acid-base interactions and the nature of the donor-acceptor bond. BF3 is frequently employed as a standard Lewis acid, for the quantitative characterization of the Lewis basicity of donor mojecules.62,63 The gas-phase equilibrium constants for some BF3 complexes are shown in Table 5. 

A quantitative determination of the strength of Lewis acids to establish similar scales (Ho) as discussed in the case of protic (Br0nsted-type) superacids would be most useful. However, to establish such a scale is extremely difficult. Whereas the Brpnsted acid-base interaction invariably involves a proton transfer reaction that allows meaningful comparison, in the Lewis acid-base interaction, involving for example Lewis acids with widely different electronic and steric donating substituents, there is no such common denominator.25,26 Hence despite various attempts, the term strength of Lewis acid has no well-defined meaning. 

More recently, a quantitative scale for Lewis acidity based on fluoride ion affinities was calculated using ab initio calculations at the MP2/B2 level of theory.26 Due to its high basicity and small size, the fluoride ion reacts essentially with all Lewis acids thus the fluoride affinity (or reaction enthalpy) may be considered as a good measure for the strength of a Lewis acid. An abbreviated pF scale is given in Table 1.3. This scale was used recently by Christe and Dixon112 for estimating the stability of salts of complex fluoro anions and cations. The pF value represents the fluoride affinity in kcal mol 1 divided by 10. 

The topic of interactions between Lewis acids and bases could benefit from systematic ab initio quantum chemical calculations of gas phase (two molecule) studies, for which there is a substantial body of experimental data available for comparison. Similar computations could be carried out in the presence of a dielectric medium. In addition, assemblages of molecules, for example a test acid in the presence of many solvent molecules, could be carried out with semiempirical quantum mechanics using, for example, a commercial package. This type of neutral molecule interaction study could then be enlarged in scope to determine the effects of ion-molecule interactions by way of quantum mechanical computations in a dielectric medium in solutions of low ionic strength. This approach could bring considerable order and a more convincing picture of Lewis acid base theory than the mixed spectroscopic (molecular) parameters in interactive media and the purely macroscopic (thermodynamic and kinetic) parameters in different and varied media or perturbation theory applied to the semiempirical molecular orbital or valence bond approach [11 and references therein]. 

Solvents can be classified as EPD or EPA according to their chemical constitution and reaction partners [65]. However, not all solvents come under this classification since e.g. aliphatic hydrocarbons possess neither EPD nor EPA properties. An EPD solvent preferably solvates electron-pair acceptor molecules or ions. The reverse is true for EPA solvents. In this respect, most solute/solvent interactions can be classified as generalized Lewis acid/base reactions. A dipolar solvent molecule will always have an electron-rich or basic site, and an electron-poor or acidic site. Gutmann introduced so-called donor numbers, DN, and acceptor numbers, AN, as quantitative measures of the donor and acceptor strengths [65] cf. Section 2.2.6 and Tables 2-3 and 2-4. Due to their coordinating ability, electron-pair donor and acceptor solvents are, in general, good ionizers cf. Section 2.6. 

The sttuctures of organoaluminum halides R AlXs are dominated by dimeric structures with AI2X2 four-membered rings for X = Cl, Br, and I. Unlike the aluminum alkyls, these hahde bridges are electron-precise and result from intermolecular Lewis acid-base complexation. These dimers do dissociate and readily react with Lewis bases, but the enthalpies of dissociation are higher than those for the aluminum alkyls. In many cases, the strength of the association is sufficient that dimers are observed in the gas phase. 

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