Lewis acid-base equilibria

Thirdly, the intramolecular assodation of a solvent affects the Lewis acid - base equilibrium Upon... 

The rate-determinating step in reaction (45) is second order, illustrating that Lewis acid-base interactions are involved in this process. Eqs. (43)-(45) are similar to reactions (36) and (37) since the tin(II) compounds formed are highly associated (tin(II) chloride and tin(II) sulfide can be isolated as pure and large crystals), the equilibrium is shifted to the right side. 

The general Lewis-acid-base reaction (3.95) exemplifies the two-electron stabilizing donor-acceptor interaction of Fig. 1.3 (namely the nN->-nB interaction for (3.94)), which may be distinguished from the complementary bi-directional donor-acceptor interactions of covalent-bond formation (Section 3.2.1). However, this leaves open the question of whether (or how) the equilibrium bond reflects the formal difference between heterolytic (3.95) and homolytic (3.96) bond formation. 

Figure 3.34 Atomic-charge variations AQ= Q(R) - Q(oo) for boron (circles) and nitrogen (squares) atoms of the

The Lewis acid-base properties of these ionic liquids are determined by the chloroaluminate species. The equilibrium of the chloroaluminate liquid is primarily described by two equilibria at x AICI3 below 0.67 ... 

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. 

Although BH3 exists in the form of Lewis acid-base adducts and as 4 presumable intermediate in reactions of dihorjne. only trace quantities of the free molecule have been detected. The equilibrium constant for dimerization is approximately MV and the enthalpy of dissociation of the dimer to the monomer is about -HSO kJ mol-1 or slightly more. <20... 

A Lewis base transfers an electron pair to a Lewis acid. A Bronsted acid transfers a proton to a Bnansted base. These exist in conjugate pairs at equilibrium. In an Arrhenius base, the proton acceptor (electron pair donor) is OH-. All Arrhenius acids/bases are Bronsted acids/bases and all Bransted acids/bases are Lewis acids/bases. Each definition contains a subset of the one that comes after it. 

Equation 58 defines the equilibrium between free silene and its Lewis acid-base complex with a nucleophilic solvent. Since the complexed form of the silene can clearly be expected to be relatively unreactive toward nucleophiles compared to the free silene, then the result will be a reduction in the overall rate constant for reaction with a nucleophilic reagent (Nu-H) in a complexing solvent relative to a non-complexing one like hexane. This rate reduction is described quantitatively in equation 59. 

The position of this equilibrium depends on the electrophilicity or nucleophilicity of A and B , respectively, as well as the solvation capability of the surrounding medium. The solvent can influence the association as well as the electron-transfer step (or in the reverse reaction the ionization and dissociation step). The position of the Lewis acid/base equilibrium given in Eq. (4-30) will depend mainly on the differential solvation of the ionic and covalent species (a) and (b). 

Analogous results have been obtained for the Lewis acid/base equilibrium between ionic tropylium azide and covalent 7-azidocycloheptatriene [178]. Again, in less polar solvents such as deuterio-trichloromethane and even [Dgjacetone, no ionization to give the tropylium and azide ions could be detected. Dipolar liquid sulfur dioxide, however, induces complete ionization at low temperature (—70 °C). 

An interesting example of a Lewis acid/base reaction between neutral reactants is the formation of tris(n-butyl)phosphonium-dithiocarboxylate, ( -Bu)3P" — 82 , from tris(n-butyl)phosphane and carbon disulfide in solution. As expected, this equilibrium is strongly shifted in favour of the dipolar zwitterion with increasing solvent polarity (diethyl ether dimethyl sulfoxide) [272, 273]. 

Figure 6.3. Titration of H30 and Cu aq with ammonia (a) and with tetramine (trien) (b). Equilibrium diagrams for the distribution of NH3-NH4 (c) of the amino coppeifll) complexes (d) and of Cu ", Cu-trien (e). The similarity of titrating with a base and titrating a metal ion with a base (Lewis acid-base interaction) is obvious. Both neutralization reactions are used analytically for the determination of acids and metal ions. A pH or pMe indicator electrode (glass electrode for and copper electrode for Cu " ) can be used for the end point indication.

This equation has been used in several correlations of solvent effects on solute properties such as reaction rates and equilibrium constants of solvolyses, energy of electronic transitions, solvent-induced shifts in UV/visible, IR, and NMR spectroscopy, fluorescence lifetimes, and formation constants of hydrogen-bonded and Lewis acid/base complexes [Kamlet et al., 1986b]. 

Furthermore, saturated aldehydes are somewhat less basic than saturated ketones or esters, resulting in reversible complexation even with bulky aluminum reagents. However, whether the equilibrium [Lewis acid -i- base Lewis acid-base complex] is reversible or irreversible, the selective functionalization of more labile or sterically less-encumbered aldehydes is facile using bulky or mild Lewis acids. 

Titanium tetrachloride and amine yield an equilibrium mixture of Lewis acid-base complexes. At least one of these complexes has an increased reactivity towards the carbonyl oxygen and forms a ternary ketone—amine—TiCl,j complex [1] which in the presence of excess amine yields a titanium enolate. The reaction then proceeds as... 

The second important solvent effect on Lewis acid-Lewis base equiUbria concerns the interactions with the Lewis base. Since water is also a good electron-pair acccplor, Lewis-type interactions are competitive. This often seriously hampers the efficiency of Lewis acid catalysis in water. Thirdly, the intennolecular association of a solvent affects the Lewis acid-base equilibrium-" -. Upon complexation, one or more solvent molecules that were initially coordinated to the Lewis acid or the Lewis base are liberated into the bulk liquid phase, which is an entropically favourable process. This effect is more pronounced in aprotic than in protic solvents which usually have higher cohesive energy densities. The unfavourable entropy changes in protic solvents are somewhat counterbalanced by the formation of new hydrogen bonds in the bulk liquid. 

E4.43 Equation (b) is better in explaining the observ ations described in the exercise. Most notably it contains the Lewis acid-base adduct [FeCblOPC 1,3)4] in which OPCI3 is a Lewis base and coordinates via its O atom, consistent with vibrational data. It also has the Fe Clb/fFeCy (i.e., red/yellow) equilibrium. The titration would have to start in either case from a concentrated (red) solution, and the equivalence point at 1 1 mole ratio FeClj/Et4NCI can be explained based on the reaction ... 

As I noted in the Preface, my particular scientific background and the pure joy of chemical synthesis played major roles in the discovery. My thesis work concerned metal hydrides, in particular Lewis acid-base complexes of dialkylzinc species with hydride ion, which in some cases were in equilibrium with hydride-bridged com-... 

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