Acid-Base Reactions in the Gas Phase

As discussed in the previous section, solvent effects have an important influence on pfCg values determined in solution. Indeed, there are pairs of acids for which the relative order of acidity can be reversed by a change of solvent. The order of acidity of a pair of compounds may also be a function 

The aqueous pfCa values shown are estimated from measurements using the acidity fimction. Ho, but there are uncertainties with these measurements. 

In principle, one might try to study the ionic dissociation of an acid (equation 7.3) directly in the gas phase, but AH for dissociation of a neutral species to a proton and an anion is usually quite large without solvent stabilization of the ions. For example, the AH for the gas phase dissociation of methane to methyl anion and a proton (AH° jj) was calculated to be - -417kcal/mol. This is much greater than the homolytic C—H bond dissociation energy of methane (-I-104 kcal/mol), so thermolysis of methane in the gas phase leads to radicals instead of ions. The pKg value of an acid can be determined indirectly, however, by measuring the equilibrium for proton transfer from the acid to a base with a known pKg. With a series of measurements, a scale of gas phase acidity values can be established by referencing one compound to another. 

For a discussion of each of these techniques, see Pellerite, M. J. Brauman, J. I. in Buncel, E. Durst, T., Eds. Comprehensive Carbanion Chemistry, Part A Structure and Reactivity Elsevier Scientific Publishing Amsterdam, 1980 p. 55 ff. 

Electron transfer from an anionic base to a neutral acid can also compete with proton transfer from the acid to the base in the gas phase. Han, C.-C. Brauman, J. 1. /. Am. Chem. Soc. 1988,110, 4048. 

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The standard state for reporting enthalpies and free energies of acid-base reactions in the gas phase is 298 K. The spectroscopic parameters must be corrected because homolytic dissociation energies are usually reported at 298 K, but electron affinity and ionization potential values derived from spectroscopic data refer to enthalpies at OK. ... 

Techniques to Study Acid-Base Reactions in the Gas Phase... 

Bronsted s definition acknowledges the existence of acid-base reactions in the gas phase, in which the complicating effects of solvation are absent. The acidity of a substance B is then measured by the proton affinity PA of the molecule, i.e. the variation of enthalpy involved in the reaction ... 

Thermodynamics of complex formation of silver with several ligands such amines,368 hindered pyridine bases,369 nitrogen donor solvents,370 and azoles371 have been carried out. Other studies include the secondary-ion mass spectra of nonvolatile silver complexes,372 the relationship between Lewis acid-base behavior in the gas phase and the aqueous solution,373 or the rates of hydride abstraction from amines via reactions with ground-state Ag+.374... 

It can be seen from the foregoing discussion that the interpretations of the observed acidities leave something to be desired even for such a fundamental series of compounds as the simple hydrides. The matter has been reopened in recent—, years by the development of techniques for measuring acidities in the gas phase.86 The available results reemphasize the fact, already well known from previous work, that solvation factors have a profound influence on the course of acid-base reactions. But the gas-phase experiments do more than this they call into question some of the fundamental assumptions and interpretations that haVe long been used to account for observed acidities in terms of molecular structure. 

The present chapter is concerned primarily with measured molecular structural effects on reactions 2 and 4 in the gas phase. These have been obtained only very recently from direct equilibrium-constant determinations. Work in this area is still in a very active state, so that this chapter serves as a preliminary progress report. Useful comparison can now be made of structural effects on equation 2 with the following related topics (1) proton-transfer equilibria in condensed phases (2) other Lewis acid-base equilibria in the gas phase (3) theoretical calculations of proton-transfer energetics (4) hydrogen-atom transfer equilibria between cation radicals and saturated cations (5) hydrogen-bonded complex formation, in hydrocarbon solvents and (6) gas-phase equilibria for attachment of neutral molecules to cations and anions. Each of these topics is considered at least briefly. 

Proton transfer is one of the prominent representatives of an ion-molecule reaction in the gas phase. It is employed for the determination of GBs and PAs (Chap. 2.11.2) by either method the kinetic method makes use of the dissociation of proton-bound heterodimers, and the thermokinetic method determines the equilibrium constant of the acid-base reaction of gaseous ions. In general, proton transfer plays a crucial role in the formation of protonated molecules, e.g., in positive-ion chemical ionization mass spectrometry (Chap. 7). 

Even more general is the Lewis concept of acids and bases a Lewis base has a lone pair available for formation of a coordinate bond, and a Lewis acid has a vacant acceptor orbital handy. This concept is applicable to reactions in the gas phase or in inert solvents (as discussed in the previous section) as well as to complex formation in solution and the acid/ base phenomena studied by Arrhenius, Br0nsted and Lowry. 

The alternative approach is the attempt to quantify substituent effects, and this has been most successfully done by the Hammett equation and its various extensions (Hammett, 1970). Here one set of free energy data is compared with another set. One set is taken as standard (originally the dissociation constants of benzoic acids) and other rate or equilibrium data are compared (by logarithmic plots). So much has been written about this treatment that discussion here is unnecessary. Absolute values of a (the substituent constant) are not to be expected, in fact one would expect a different a for every reaction (i.e. for every p). In the present context it is important to note that both the Hammett equation and the closely related Taft treatment are based on systems where solvation is known to be important and therefore the application of these treatments using parameters derived from solution phase studies to reactions in the gas phase may be of uncertain value. 

Three new experimental techniques, developed within the past decades, now make it possible to study ionic reactions in the gas phase as well. These are pulsed ion-cyclotron-resonance (ICR) mass spectrometry, pulsed high-pressure mass spectrometry (HPMS), and the flowing afterglow (FA) technique [469-478 see also the references given in Section 4.2.2]. Although their approaches are quite independent, the results obtained for acid/base and other ionic reactions agree within an experimental error of 0.4... 1.3 kJ/mol (0.1... 0.3 kcal/mol) and are considered as reliable as those obtained in solution. 

The gas-phase exothermicity is due to the enhanced resonance stabilization of aniline compared to benzene for the nonzwitterionic amino acids, as found in the gas phase. On the other hand, the aniline resonance stabilization is lost in the zwitterionic amino acids of the solid phase and thus the reaction is essentially thermoneutral. This is, of course, related to the weak basicity of aniline compared to related nonaromatic bases such as cyclohexylamine, as exhibited by the ca 50 kJ mol-1 exothermicity of the formal reaction 29. 

The free-energy change, AG, of such a reaction in the gas phase is simply PA(RO ) — PA(B). When B is OH and ROH is phenol, then AGg is PA(PhO ) — PA(OH ) = —41 kcalmoR, so in this reaction the phenol molecule acts as the Br0nsted acid and the OH anion as the Br0nsted base. Clearly, relative proton-affinity values determine the relative acidity scale of molecules in the gas phase. 

Equations such as (135) contain only vibrational frequencies, which can be derived from spectroscopic observations for species of not too great complexity. It should thus be possible to predict the isotope effect on simple equilibria, and this has been done successfully for a number of reactions in the gas phase. The problem is more difficult for reactions in solution, and few such quantitative predictions can be made for acid-base equilibria, though existing experimental data can be rationalized to some extent. 

The first item in Table 9.6, the reaction between propanoic add (CH3CH2CO2H) and ammonium hydroxide, corresponds to an add-base reaction. Carboxylic acids are, as noted from the pKa s in Table 9.11, acidic and while, in the gas phase, the low molecular weight carboxylic acids contain significant quantities of dimer (Figure 9.16), it is likely that solvated monomer is the major spedes (at least in solvents capable of hydrogen bonding, e.g., H2O). 

We must measure the AH° of these three reactions to derive the acidity of HA in the gas phase. The acidity is reported as a AH° value, not a pK, because it does not reflect the donation of the proton to any solvent. Instead, the value is a heteroly tic bond dissociation energy. In those cases where the base (anionic or neutral) is stable in the gas phase, one can determine the proton affinity of this conjugate base of the acid, which would be the negative of the heteroly tic bond dissociation energy (Eq. 5.26). The most pertinent number for any reaction is AG° not AH°, which would require knowledge of AS°. AS° can often be estimated, but for most gas phase acidities the entropies of the reactions are all similar, such that relative AG° is barely different than relative AH°. Therefore, Table 5.4 lists gas phase acidities solely as AH° values. 

The energy of reaction of an acid HX with a given base B in the gas phase can be estimated by the following dissection of equation (5.113) ... 

First, the simple thermodynamic description of pe (or Eh) and pH are both most directly applicable to the liquid aqueous phase. Redox reactions can and do occur in the gas phase, but the rates of such processes are described by chemical kinetics and not by equilibrium concepts of thermodynamics. For example, the acid-base reaction... 

The first systematic measurements of the reactions of ions with molecules in the gas phase were initiated largely by workers associated with analytical mass spectrometry.4-6 It was the rapidly expanding area of ion-molecule reactions which led to the origin of Gas-Phase Ion Chemistry as a distinct field.7 The discovery that ion-molecule equilibria in the gas phase can be determined by mass spectrometric techniques8 led to an explosion of thermochemical measurements based on determination of equilibria by a variety of techniques.9 Significantly, for the first time, information could be obtained on the thermochemistry of reactions which had solution counterparts of paramount importance such as acidities and basicities. These were obtained from proton transfer equilibria such as,... 

These models refer to reactions with the simplest nucleophile, H, both under neutral conditions and in the protonated form. Chemical reactivity can be strongly altered by catalytic effects acid/base catalysis is of particular importance. We regard the studies on ga phase acidities and on proton affinities discussed in the above sections to bear special significance for quantitative modelling of acid/base catalysis in the future. 

One of the problems encountered when dealing with the interaction of Lewis acids and bases in a quantitative way is in evaluating the role of the solvent. Bond energies in molecules are values based on the molecule in the gas phase. However, it is not possible to study the interaction of many Lewis acids and bases in the gas phase because the adducts formed are not sufficiently stable to exist at the temperature necessary to convert the reactants to gases. For example, the reaction between pyridine and phenol takes place readily in solution as a result of hydrogen bonding ... 

It is the interaction of the Lewis acid and base in the gas phase that gives the enthalpy of the A B bond, AHAB, but what is more often (and conveniently) measured is the enthalpy change when the reaction is carried out in solution, AH AB. 

Chemical/Physical. Begins to polymerize at 80.2 °C (Weast, 1986). Slowly hydrolyzes in water forming methyl alcohol and acrylic acid (Morrison and Boyd, 1971). Based on a hydrolysis rate constant of 0.0779/M-h at pH 9 at 25 °C, an estimated half-life of 2.8 yr at pH 7 was reported (Roy, 1972). The reported rate constant for the reaction of methacrylonitrile with ozone in the gas phase is 2.91 x lO cm moFsec (Munshi et al, 1989a). 

Photolytic. Irradiation of vinyl chloride in the presence of nitrogen dioxide for 160 min produced formic acid, HCl, carbon monoxide, formaldehyde, ozone, and trace amounts of formyl chloride and nitric acid. In the presence of ozone, however, vinyl chloride photooxidized to carbon monoxide, formaldehyde, formic acid, and small amounts of HCl (Gay et al, 1976). Reported photooxidation products in the troposphere include hydrogen chloride and/or formyl chloride (U.S. EPA, 1985). In the presence of moisture, formyl chloride will decompose to carbon monoxide and HCl (Morrison and Boyd, 1971). Vinyl chloride reacts rapidly with OH radicals in the atmosphere. Based on a reaction rate of 6.6 x lO" cmVmolecule-sec, the estimated half-life for this reaction at 299 K is 1.5 d (Perry et al., 1977). Vinyl chloride reacts also with ozone and NO3 in the gas-phase. Sanhueza et al. (1976) reported a rate constant of 6.5 x 10 cmVmolecule-sec for the reaction with OH radicals in air at 295 K. Atkinson et al. (1988) reported a rate constant of 4.45 X 10cmVmolecule-sec for the reaction with NO3 radicals in air at 298 K. 

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