MARGARET FRENCH and PAUL KEBARLE. Can. J. Chem. 53,2668Chem. 53, (1975. The major reactions in methane containing traces of ethane were studied with a pulsed electron beam high ion source pressure mass spectrometer. The CH4: C2H6 ratios were changed from 50: 1 to 100 000: 1 at reaction temperatures between 28-210 "C. The reaction 1 : CH5 + + CzHs = CH4 + C,H7+ still proceeded from left to right even at the highest dilution ratios.Measurement of the C,H7+/CH5+ ratio under these conditions leads to a lower limit of the proton a f i i t y difference PA(C2H6) -PA(CH4) > 10 kcal/mol. This result is in agreement with measurements by Bohrne.It was observed that C,H7+ decomposes slowly at 30°C. The decomposition becomes more rapid at higher temperature. Measurements of the temperature dependence of the thermal reaction 2: CzH7+ + CH4 = CZH5+ + HZ + CH4, and an Arrhenius plot of k, led to the activation energy Ez = 10.5 kcal/mol and preexponential factor A, = 8.3 x (cm3 molecule-' s-I). Assuming E, = AH, one obtains AHI(C2H7+) = 208.5 f 2 kcal/mol and PA(C2H7+) = 137.4 f 2 kcal/mol. This is close to the proton affinity PA(C,H7+) = 139 2 kcal/mol that can be deduced from Bohme's results.At higher dilution ratios the ion C2H5+ was observed to react not only with ethane but also with methane by reaction 6: C2H5+ + CH4 = C3H7+ + Hz, k6 NN 1 x 10-14cm3 molecule-' s-' at 86 "C. The reaction has positive temperature dependence.MARGARET FRENCH et PAUL KEBARLE. Can. J. Chem. 53,2668Chem. 53, (1975. A des taux de dilution plus elevb on a remarque que l'ion C2H5 + reagissait non seulement avec l'ethane mais aussi avec le methane selon la reaction 6: CzH5+ + CH4 = C3H7+ + HZ, k6 NN 1 x 10-l4 cm3 mol6cules-' s-' a 86 "C. La reaction a une dkpendance thermique positive.[Traduit par le journal]
Measurement of the gas phase ion equilibria between ions M and solvent molecules S1 provide binding energies of the complexes M *(SO, (for n = 1 to -6). Comparison of these data with single ion energies of solvation shows that differences in ion solvation in solution are reflected in the binding energies of ion-molecule complexes in the gas phase. The weaker solvation of negative ions (relative to positive ions) observed in liquid aprotic solvents is reflected in the binding energies of negative ion aprotic molecule complexes, a weaker binding being found for the first and subsequent few aprotic molecules. An analysis of the bonding in CI-(CH,CN) and K+(CH3CN) shows that the weaker bonding to C1-is due to the very diffuse distribution of the positive pole of the dipole in acetonitrile. In effect the dipole can not come close to the negative ion. Analysis shows a similar picture also for acetone. Experimental results for the bonding between CI-HR are given for a variety of compounds HR. These show that for RH = protic compounds, like oxygen acids, the hydrogen bond in Cl-HR increases with the acidity of HR. For aprotic compounds, i.e., carbon acids, no relationship between the bond in the complex and the acidity of HR is found. An examination of the solvation of substituted phenoxide ions by protic and aprotic solvents shows that solvation by protic solvents is adversely affected by charge dispersal in the ion, while aprotic solvents are much less sensitive to charge dispersal. The reasons for this important difference in behaviour are examined.The present work utilizes data obtained from gas phase ion equilibria measure-ments1S2 to provide insights into ion solvation by protic and aprotic solvents. The equilibria (n, n -1) can be measured for various anions Alike F-, C1-, PhO-and cations C + like Na+, K+, H30+, PhNHf and diverse solvent molecules S1 like HzO, NH3, CH3CN etc., A-(Sl),The apparatus used employs mass spectrometric detection of ions which have reached equilibrium in a reaction chamber containing known pressures of the neutral reactants S1. The temperature dependence of the equilibrium constant Knsn-provides, via van't Hoff plots, the desired binding energies AHan,n-l and Early observations that organic ions are very reactive in dipolar aprotic solvents like dimethylsulphoxide (DMSO) and dimethylf~rmamide~ led to the s ~g g e s t i o n ~* ~ that these solvents solvate negative ions less well than positive ions. Studies of the enthalpies of tran~fer,~ activity coefficients and Volta potentials' have confirmed this supposition. When searching for an explanation of this behaviour one encounters the question whether the effect is due to the interactions of the ion with the " closest " solvent molecules or to bulk properties of the solvent. As shown in the Results and Discussion section, gas phase measurements of (n, n -1) equilibria involving positive
Far-infrared spectroscopy is shown to provide a powerful new method for studying ion aggregates in solution. The interionic vibration frequency of Na+BP14-is at 175 cm-1 in pyridine 1 4-dioxane piperidine and tetrahydrofuran, indicating that bare, rather than solvated ions, produce the' absorption. The isotop~ shift of the 198-cm-1 band in ~+BP14to 183 cm-lin ND4+ BP14-also fits a bare-ion model. The Li+BP14salt gives a band at 410 cm-1 in THF, which is probably due to a Li+ THF vibration, while the K+BP14band at 133 cm-1 arises from interionic vibration. The Na+, NH4+, ND4+, and K+ salt stretching force constants, based on a bare ion-pair model, are nearly identical, 3.85±0.11 X 10 4 dyn cm-1 . These force constants differ markedly from those predicted from the Pettit and Bruckenstein ion-pair potential used to interpret conductivity data of strong electrolyte solutions. Present observations indicate that the dominating terms in the ion-pair potential are -(ill R) + (BI Rn), where n»7 (Pettit and Bruckenstein) and may approach 30 for a pair of polyatomic ions. The dependence of the band intensities on solvent: concentration, and the nature of the cation appear to be primarily due to variations in the proportion of absorbing systems, rather than in the intrinsic absorption coefficient, and support the assignment of the bands to ion pairs, rather than higher aggregates. Intensity measurements in a mixed benzene-pyridine solvent demonstrate the phenomenon of preferential solvation by pyridine. 4
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