Electron Affinities from Solution Data

Relative electron affinities of organic molecules can be obtained from half-wave reduction potentials in aprotic solvents. The electron affinities are related to the half-wave reduction potentials by 

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The biochemical applications involve the electronic nature of the components of DNA and proteins, especially the charge distributions, electron affinities, and gas phase acidities of purines and pyrimidines and amino acids. The role of electron reactions in diverse areas such as cancer, electron conduction, and sequence recognition all depend on fundamental energetic properties such as electron affinities and solution energies. We explain nonadiabatic experimental data from radiation chemistry by excited anionic states of biological molecules [24],... 

The relativistic coupled cluster method starts from the four-component solutions of the Drrac-Fock or Dirac-Fock-Breit equations, and correlates them by the coupled-cluster approach. The Fock-space coupled-cluster method yields atomic transition energies in good agreement (usually better than 0.1 eV) with known experimental values. This is demonstrated here by the electron affinities of group-13 atoms. Properties of superheavy atoms which are not known experimentally can be predicted. Here we show that the rare gas eka-radon (element 118) will have a positive electron affinity. One-, two-, and four-components methods are described and applied to several states of CdH and its ions. Methods for calculating properties other than energy are discussed, and the electric field gradients of Cl, Br, and I, required to extract nuclear quadrupoles from experimental data, are calculated. 

Basicity in the gas phase is measured by the proton affinity (PA) of the electron donor and in solution by the pAj,. A solution basicity scale for aldehydes and ketones based on hydrogen bond acceptor ability has also been established [186]. Nucleophilicity could be measured in a similar manner, in the gas phase by the affinity for a particular Lewis acid (e.g., BF3) and in solution by the equilibrium constant for the complexation reaction. In Table 8.1 are collected the available data for a number of oxygen systems. It is clear from the data in Table 8.1 that the basicities of ethers and carbonyl compounds, as measured by PA and p , are similar. However, the nucleophilicity, as measured by the BF3 affinity, of ethers is greater than that of carbonyl compounds, the latter values being depressed by steric interactions. 

The lattice enthalpy U at 298.20 K is obtainable by use of the Born—Haber cycle or from theoretical calculations, and q is generally known from experiment. Data used for the derivation of the heat of hydration of pairs of alkali and halide ions using the Born—Haber procedure to obtain lattice enthalpies are shown in Table 3. The various thermochemical values at 298.2° K [standard heat of formation of the crystalline alkali halides AHf°, heat of atomization of halogens D, heat of atomization of alkali metals L, enthalpies of solution (infinite dilution) of the crystalline alkali halides q] were taken from the compilations of Rossini et al. (28) and of Pitzer and Brewer (29), with the exception of values of AHf° for LiF and NaF and q for LiF (31, 32, 33). The ionization potentials of the alkali metal atoms I were taken from Moore (34) and the electron affinities of the halogen atoms E are the results of Berry and Reimann (35)4. 

The most controversial and one of the most important quantities is the bond dissociation energy for the reaction HO2 H + O2. The understanding of autoxidation depends to some extent on this value. Much higher values have been postulated—e.g., 67 kcal. by Walsh (26). The evidence for the lower value of 36 kcal. is, apart from the agreement with kinetic data, based on the peaks of the electron transfer spectra of various ferric ion pair complexes. These allow an estimate of the electron affinity of the HO2 radical in solution and consequently produce a value of 102 kcal. for Dh.. .02H and 36 kcal. for Dh.. .02- This in turn leads to a dissociation constant of 10 for the HO2 radical in aqueous solutions. 

The physical properties determined using the ECD are important to different areas of chemistry. Analytical chemistry deals with how much and what are involved in a chemical reaction. Expressed differently, it establishes what we refer to as the QQQ quantitation, qualitative identification, and the quality of the results. The determination of the electron affinities of the chlorinated biphenyls, dioxins, and phenols and the prediction of the response of the ECD and NIMS are important to qualitative and quantitative analyses of environmental pollutants [21]. Polarographic reduction in solutions likewise gives accurate and precise qualitative and quantitative results. The quality of the analyses is expressed by the random and systematic uncertainties in the reported values. These are obtained from the same principle of weighted least squares used to obtain information from ECD data. Wentworth has described the application of the general least-squares procedure to chemical problems [22, 23]. 

In Figure 6.17 the electron affinities of several substituted quinones determined from E /2 measurements, and/or charge transfer spectra, are plotted versus the current evaluated gas phase values. They are chosen to give a comparison of the values obtained by the two methods and to note the deviations from the unit slope line. The displacement of the unit slope lines by a constant amount is a systematic uncertainty. This indicates that the solution energy differences and constants used to correlate the charge transfer complex data are both a function of the electron affinity and/or the type of molecules [78]. Thus, it will be possible to reduce deviations by classifying the molecules and identifying the functional relationship. 

The electron affinities of halogenated aromatic and aliphatic compounds and nitro compounds have been evaluated. Additional electron affinities for halogenated benzene, freons, heterocyclic compounds, dibenzofuran, and the chloro- and fluoroben-zenes are reported from ECD data. The first positive Ea for the fluorochloroethanes were obtained from published ECD data. The Ea of halogenated aromatic radicals have been estimated from NIMS data. The AEa of all the halobenzenes have been calculated using CURES-EC. The Ea of chlorinated biphenyls and chlorinated napthalenes obtained from reduction potentials have been revised based on variable solution energy differences. 

The feasibility of electron transfer from alkoxides to acceptor species, attractive as it is for its simplicity, has been prospected as a real possibility only in a few instances (44, 45a) and, in some cases, questioned (9, 45b, 46a). Buncel and Menon (45b) estimated that electron transfer from t-BuO- to 4-nitrotoluene in f-BuOH is energetically unfavorable to the extent of about 3 eV. The estimate is based on calculations involving gas-phase electron affinity data for t-BuO (1.93 eV) instead of its oxidation potential in solution, which, as is the case for other alkoxides, is not known. The approximation is necessarily crude, because solvation effects could be of significant magnitude nevertheless, the estimated value has met wide acceptance. 

Sources Data from Ionization energies cited in this chapter are from C. E. Moore, Ionization Potentials and Ionization Limits Derived fwm the Analyses of Optical Spectra, National Standard Reference Data Series, U.S. National Bureau of Standards, NSRDS-NBS 34, Washington, DC, 1970, unless noted otherwise. Electron affinity values listed in this chapter are from H. Hotop and W. C. Lineberger, J. Phys. Chem. Ref Data, 1985,14, 731. Standard electrode potentials listed in this chapter are from A. J. Bard, R. Parsons, and J. Jordan, eds., Standard Potentials in Aqueous Solutions, Marcel Dekker (for lUPAC), New York, 1985. Electronegativities cited in fiiis chapter are from J. B. Mann, T. L. Meek, and L. C. Allen, J. Am. Chem. Soc., 2000,122, 2780, Table 2. Other data are from N. N. Greenwood and A. Earnshaw, Chemistry ofthe Elements, Pergamon Press, Elmsford, NY, 1984, except where noted. J. Emsley, The Elements, Oxford University Press, New York, 1989. S. G. Bratsch, J. Chem. Educ., 1988,65, 34. 

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