A method is presented to determine the absolute hydration enthalpy of the proton, ∆H aq °[H + ], from a set of cluster-ion solvation data without the use of extra thermodynamic assumptions. The absolute proton hydration enthalpy has been found to be ∼50 kJ/mol different than traditional values and has been more precisely determined (by about an order of magnitude). Conventional ion solvation properties, based on the standard heat of formation of H + (aq) set to zero, have been devised that may be confusing to the uninitiated but are useful in thermochemical evaluations because they avoid the unnecessary introduction of the larger uncertainties in our knowledge of absolute values. In a similar strategy, we have motivated the need for a reassessment of ∆H aq °[H + ] by the trends with increased clustering in conventional cluster-ion solvation enthalpy differences for pairs of oppositely charged cluster ions. The consequences of particular preferred values for ∆H aq °[H + ] may be evaluated with regard to cluster-ion properties and how they connect to the bulk. While this approach defines the problem and is strongly suggestive of the currently determined proton value, it requires extra thermodynamic assumptions for a definitive determination. Instead, a unique reassessment has been accomplished without extra thermodynamic assumptions, based on the known fraction of bulk absolute solvation enthalpies obtained by pairs of oppositely charged cluster ions at particular cluster sizes. This approach, called the cluster-pair-based approximation for ∆H aq °[H + ], becomes exact for the idealized pair of ions that have obtained the same fraction of their bulk values at the same cluster size. The true value of ∆H aq °[H + ] is revealed by the linear deviations of real pairs of ions from this idealized behavior. Since the approximation becomes exact for a specific pair of oppositely charged ions, the true value of ∆H aq °[H + ] is expected to be commonly shared on plots of the approximation vs the difference in cluster-ion solvation enthalpy for pairs of ions sharing the same number of solvating waters. The common points on such plots determine values of -1150.1 ( 0.9 kJ/mol (esd) for ∆H aq °[H + ] and -1104.5 ( 0.3 kJ/mol (esd) for ∆G aq °[H + ]. The uncertainties (representing only the random errors of the procedure) are smaller than expected because the cluster data of 20 different pairings of oppositely charged ions are folded into the determination.

No abstract

A new method of determining standard absolute solvation energies (enthalpies and Gibbs free energies) for individual ions is presented. This method originated from the cluster pair based approximation used in earlier work [Tissandier et al., J. Phys. Chem. A 1998, 102, 7787] and is called the cluster pair correlation scheme. Unlike the earlier approximation, the new scheme makes analysis possible in solvent systems for which bulk ion-solvation data are complemented by cluster ion data, either experimental or theoretical, for clusters containing only a single solvent molecule. The correlation scheme features linear plots of half the difference between standard conventional cluster solvation energies for anions and cations containing equal numbers of solvent molecules in the cluster vs their bulk counterparts. All of the correlation lines should intersect at a point whose ordinate defines the unknown shift between the conventional and absolute energy scales, the standard absolute energy of formation of the proton. The results obtained using this correlation scheme to analyze the available cluster ion data for water are in substantial agreement with the results reported earlier using the cluster pair based approximation. Preliminary values for the standard absolute enthalpy and Gibbs free energy of ammonation of the proton are obtained when the cluster pair correlation scheme is applied to ammonia, and a clear indication is obtained of what additional data are needed to improve these preliminary estimates.

No abstract

Urea has a resonance energy of approximately 41 kcal.,28 and at least in the solid state is a completely planar molecule.29 This planarity is destroyed when the activated complex is formed. Thus the activation process should require enough energy to compensate for the loss of resonance energy involved in activated complex formation. This resonance energy loss will be given by the difference in resonance energy between reactant and complex. If it is assumed that the activated complex involved in the urea decomposition possesses resonance energy approximately equal to that of an alkyl isocyanate, i.e., about 8 kcal.28; then the loss in resonance energy should be 41 -8 = 33 kcal.30

To calculate the ratio, we have used intensity formulas for Zeeman components in a Russell-Saunders scheme.We have seen no resonance of the 3250A and 3536A lines of Cd 11. Possibly, the polarization of the ^D^2 state of Cd II is too low.With our detection methods we have not been able to observe the 6215A and 7479A lines of Znll.We believe that this is the first time that magnetic dipole resonance of free ions has been observed.The observation of the magnetic resonance of atomic ions might, possibly, be of use for the determination of local magnetic fields in a plasma containing these ions as impurities.

A fundamental many-particle theory of temperature-dependent spectral moments is developed for the enhanced optical absorption bands attributed to solvated electrons in various polar solvents. Several new results are obtained (expressed in atomic units) : (1) nof = 1, where no is a mean index of refraction of the solvent and f i s the empirical oscillator strength of the band; (2) (lAre~2> = 3(1 /w)av, where is an equilibrium-averaged dispersion-in-position of the solvated electron and ( 1 1 ~) ~~ is the mean reciprocal absorption frequency of the band ; (3) pe > -$Wav, where pe is the standard chemical potential of the solvated electron and way is the mean absorption frequency of the band; (4) wth < $way, where W t h is the (vertical) photoejection threshold frequency of the solvated electron.For solvated-electron spectra in ammonia, water and a number of other solvents, no more than about 25 % of the pertinent absorption band can be ascribed to bound-bound transitions involving excited states with energies less than that of the photoejection threshold of the solvated electron.

Solubilities of alkali metal chlorides except lithium were determined in ammonia, EDA, methyl-and ethylamines, DME, and THF at three different temperatures. From these data, standard thermodynamic functions of solution and solvation with respect to ion formation were calculated. The standard free energies of solvation increased linearly as a function of D_1. However, the requirements of the Bom equation were not satisfied because the slopes and intercepts of these plots were not equal. Values of the entropies of solvation are roughly accounted for on the basis of the freezing out of a fixed number of solvent molecules. The superiority of ammonia as an agent for solvating ions is attributed to its relatively low molecular entropy compared to other solvents. A criterion for salt-like behavior is introduced and used to suggest that saturated alkali metal solutions are salt like and are probably dominated by the +• ~ion pairs. On this basis solubilities of metals in ammonia are too large to be accounted for without introducing additional species.