Proton affinities experimental values

Computed quantities are from ref. (50). Experimentally determined enthalpies for 1.1 phenol-base complexation in apolar solvents (51). Values in brackets are predicted from eq. 10. Experimentally determined OH frequency shifts for methanol-base complexes in carbon tetrachloride. These values were obtained by Berthelot and co-workers (52-54). Values in brackets are predicted from eq. 11. Experimentally determined gas phase proton affinities (55). Values in parentheses were not included in the correlation with computed quantities. Values in brackets are predicted from eq. 12. ... 

Data on proton affinities (gasphase) ofmany different compounds (see Table 2) demonstrate the high level of accuracy possible in determining energies of related species. In this report by Dewar and Dieter, the enthalpy of formation of is the experimental value (367.2 kcal/mol). The calculated value for is unreliable. 

Molecular orbital calculations predict that oxirane forms the cyclic conjugate acid (39), which is 30 kJ moF stabler than the open carbocation (40) and must surmount a barrier of 105kJmoF to isomerize to (40) (78MI50500). The proton affinity of oxirane was calculated (78JA1398) to be 807 kJ mol (cf. the experimental values of 773 kJ moF for oxirane and 777-823 kJ moF for dimethyl ether (80MI50503)). The basicity of cyclic ethers is discussed in (B-67MI50504). 

The original paper defining the Gaussian-2 method by Curtiss, Raghavachari, Trucks and Pople tested the method s effectiveness by comparing its results to experimental thermochemical data for a set of 125 calculations 55 atomization energies, 38 ionization potentials, 25 electron affinities and 7 proton affinities. All compounds included only first and second-row heavy atoms. The specific calculations chosen were selected because of the availability of high accuracy experimental values for these thermochemical quantities. 

At the DFT level the preferred protonation site was found to be O4 with a proton affinity of 208.8 kcal/mol at 298 K [98JCC989]. This value agrees nicely with the experimental value of 209.0 kcal/mol. 

Proton affinities of ethene (684 121) and 680129) kJ mol-1) measured experimentally correspond with results from ab initio calculations (698 kJ mol-1 130)). MINDO/3 calculations (with AHf(H+) = 1528 kJ mol-1 91)) also deliver a result of comparable value (714 kJ mol 1) when the formation of a classical carbocation during the protonation is assumed. 

This approach may be possible for small molecules because the gas-phase proton affinity can be obtained quantum mechanically with an accuracy of 1-2 kcal/mol [19]. However, the solvation free energy of H+ cannot be calculated and the experimental value is only known approximately, from 259.5 to 262.5 kcal/mol [60]. Also, because the proton affinity and solvation free energies in Eq. (10-7) are on the order of hundreds of kcal/mol, small percentage errors in the calculation can give rise to large error in AGaqP and pKa. Thus, this method for calculation of absolute pifa s remains impractical at the present time [6],... 

The zero-point corrected MP3/6-31G 76-31G value of 174.7 kcal mol-1 for H2C=0 agrees well with the experimental value of 171.7 kcal mol-1. The proton affinities increase in the order 3 (174.7 kcal mol-1) <1 (190.5 kcal mol-1) <2 (208.3 kcal mol-1). This is explained in terms of the predominance of the electrostatic overcharge transfer interactions, because the charge separations in the double bonds increase in the order H2C+0 2-0-a4 < H2Si+0 7-S-0 4 < H2Si+1 0-O-a7 while the frontier n orbital levels rise in the order 3 (—11.8 eV) 2 ( — 1 T9 eV) <1 (—9.8 eV). 

To verify the mechanism presented, the quantum-chemical calculations of proton affinity, Aa, were carried out for modifiers, since the corresponding experimental data are quite rare. The calculations were performed for isolated molecules, since the properties of species in the interlayer space are probably closer to the gas phase rather than the liquid. The values of Ah were calculated as a difference in the total energy between the initial and protonated forms of the modifier. Energies were calculated using the TZV(2df, 2p) basis and MP2 electron correlation correction. Preliminarily, geometries were fully optimized in the framework of the MP2/6-31G(d, p) calculation. The GAMESS suite of ah initio programs was employed [10]. Comparison between the calculated at 0 K proton affinities for water (7.46 eV) and dioxane (8.50 eV) and the experimental data 7.50 eV and 8.42 eV at 298 K, respectively (see [11]), demonstrates a sufficient accuracy of the calculation. 

By means of appropriate thermochemical cycles, it is possible to calculate proton affinities for species for which experimental values are not available. For example, using the procedure illustrated by the two foregoing examples, the proton affinities ofions such as HC03-(g) (1318 k J mol-1) and C032-(g) (2261 kj mol-1) have been evaluated. Studies of this type show that lattice energies are important in determining other chemical data and that the Kapustinskii equation is a very useful tool. 

A linear relationship between calculated An0+ values and experimental values of proton affinity (PA) for N-heteroaromatic compounds was also established (Figure 3). the points 1-4 (compounds 32h, 46, 32d and 31d respectively) have been excluded from the statistical treatment, r = 0.969]. 

For proton affinities, W1 theory can basically be considered converged [26]. The W2 computed values are barely different from their W1 counterparts, and the latter s MAD of 0.43 kcal/mol is well below the about 1 kcal/mol uncertainty in the experimental values. W1 theory would appear to be the tool of choice for the generation of benchmark proton affinity data for calibration of more approximate approaches. 

In most papers the experimental proton affinity difference between water and ammonia is taken as being equal to 37.5 kcalmol-1. Various calculated values of this gap have appeared. The PA values of ammonia and water were 205.6 and 168 or 168.6 kcalmol-1 respectively (APA = 37 or 37.6 kcalmol-1) according to Dixon and Lias26 the work of Defrees and McLean27 led to similar calculations of P/HN1I3) = 204.0 and PA (H30) = 165.1, with A PA = 38.9 kcalmol-1, although Pople and coworkers28 predicted a A PA = 39.4 kcalmol-1. 

These authors conclude that the problem of internal solvation is still an experimental and theoretical challenge GB measurements for this type of molecules of low volatility are not always in good agreement194. Molecular orbital calculations may help to solve the difficult experimental problems, but they have to take into account conformational isomerisms and the prototropic tautomerisms of the amidine and guanidine moieties. In light of the above discussion, the proton affinities deduced from the experimental GB values should be based on accurate estimations of the entropy of cyclization 86. 

Methyl cation affinities of benzene and some substituted benzenes have been calculated. These follow a simple additivity rule and the value for benzene shows good agreement with the experimental estimate. Conclusive evidence is presented that these values are linearly related to the corresponding proton affinities. The competition between deuteriation and alkylation in the reaction of radiolytically formed perdeuterio ethyl cations with iV-methylpyrrole and with thiophene has been studied. Deuteriation, the Brpnsted acid pathway, predominates and intramolecular selectivities have been determined for each reaction. ... 

The experimental germane PAq of 162.1 to 164.0 kcal mor is somewhat higher than our best estimate. The assumption that the experimental PAg is correct would lead to the conclusion that theory does a somewhat better job for germane than for protonated germane. However, the very recent study of Smith and Radom showed that the usual standard (PAq of isobutene) for the experimental determination of proton affinities may have to be corrected downwards by 2.4 to 4.8 kcal mol", which brings the theoretical PA<, close to the experimental value. ... 

Froelicher and coworkers215 and Schleyer and coworkers216 calculated the proton affinity of several carbanions (Table 25). Especially, the HF/4-31+G results are in good agreement with experimental values. Calculated proton affinities are compared to that of CH3" with the help of the isodesmic reaction given in equation 20. 

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