2226
NOTES Approximate values were 2.0 X and 8.5 X 10-8 moll/z cc - 1 / 1 sec-'/* at the respective pressures. This strongly suggests that an adsorbed phase reaction between CFs radicals and TFA to produce CFaH does (CH&OCH3)ad, + not occur. The reaction CHa CH4 (CHaCOCH2)ad, has recently been established.laJ4 The reader is referred to Konstantatos and Quinn's paperla for a detailed discussion of the variwith ation in the comparable function RCH,/RC~H~~/' varying surface to volume ratios. We see from Figure 1 that PI/,is close to 1 mm (-0.8 f 0.2 mm) a t room temperature and 120'. This value correlates well with published data obtained by other methods" and is close to the RRKM calculated values in the temperature range, based on a "loose" activated c ~ m p l e x . ~In their studies below 150°, Grotewold, Lissi, et U Z . , ~ + ~ find an increasing discrepancy with Rabinovitch and Setser's calc~lation,~ and a t 15" report that Pl12= mm from their cross-combination s t ~ d y . The ~ data are interpreted in terms of a more "rigid" activated complex, where the free rotations of the methyl groups are replaced by a low-frequency bending. Our data would seem to dispose of misgivings concerning the validity of the "loose" complex model for CH, recombination, a conclusion arrived at by Hole and Mulcahy.ll
+
Pressure, Torr. Figure 1. Semilog plot of $-z vs. pressure. In the case of mixed bath gases the pressures are taken as additive. Although deactivating efficiencies may vary, as long as some of the excess vibrational energy of C2Ha*is transferred on the initial HFA mixtures: 0, collision, it will not redissociate. Ac room temperature. All data at 70 mm, work of Giles and Whittle;% 0,157"; e, overlapping points from runs at 23, 63, and 103'. TFA: 0, -120', numbers refer to overlapping points. Point a t 30 mm overlapping points of runs a t 90, 130, and 150'; 0, reduced light beam: C), -150'; 8, wall 1 mm c-CeFlz. 0, two points reactor; C), 0.07 mm TFA at 220'; B, 330".
+
+
tion is due solely to analytical errors, and thus the difference is real, at least below 150". A few experiments were carried out at higher temperatures but the possibility of increasing olefin loss due to radical addition reactions1*2cannot be discounted. is found experimentally1 to be 2 at 300" and 30 mm, presumably, in part at least, due to this cause. Also propane becomes an increasing minor product, which renders the mechanism less clear-cut. A trend toward larger values of PI,, is apparent, in agreement with experiment and t h e ~ r y . ~ , ~ ) ~ ~ Some experiments were also carried out with an incident light beam reduced in diameter to 1.9 cm, so that only the central portion of the reactor was illuminated, and it was found that J, was unaffected. I n addition, a reaction vessel was constructed with three quartz tubes of different diameters mounted axially and running nearly the length of the reactor; the volume was 892 cma, with a surface to volume ratio of 3.1 cm-l. With this vessel the value of J, obtained at 0.07 mm agreed reasonably well with the other values, as seen in Figure 1. However, between 1 and 1.5 mm, $ values of -2.2 were obtained at 120"; some unassessed heterogeneous effect may be involved. l 2 The homogeneous pressure-independent value of J, is clearly >2.5 in Figure 1. Interestingly enough, the at 120" was indevalue of the function RCFaH/RC2Ba1'2 pendent of the reactor employed and the diameter of the incident light beam at both 0.07 and 1.0 mm. The Journal of Physical Chemistry, Vol. 76,N o . 14, 1971
+
Acknowledgments. G. 0. Pritchard thanks The National Science Foundation for support and S. Toby thanks Rutgers University for a Faculty Fellowship. (11) K. J. Hole and M. F. R. Mulcahy, J . Phys. Chem., 73, 177 (1969). (12) K. M. Maloney, ibid., 74, 4177 (1970). (13) J. Konstantatos and C. P. Quinn, Trans. Faraday Soc., 65, 2693 (1969). (14) H.Shaw and 5. Toby, J . Phvs. Chem., 72, 2337 (1968).
The Gas-Phase Acidities of Alcohols'
by Mary Jane McAdams and Larry I. Bone* Department of Chemistry, East Texas State University, Commerce, Texas 76@8 (Received January 16, 1971) Publication coats assisted bv the Robert A. Welch Foundation
Brauman and Blair2,*have reported a scale of the relative acidities of alcohols in the gas phase. Using ion cyclotron resonance techniques, they Rere able to show that as the size of the alkyl group increases on an (1) This work is supported by a Faculty Research Grant, East Texas State University, and the Robert A. Welch Foundation. (2) J. I. Brauman and L. K. Blair, J. Amer. Chem. &c., 92, 6987 (1970). (3) J. I. Brauman and L. K. Blair, ibid., 90,6561 (1968).
2227
NOTES alcohol, so does the acidity. Tiernan and Hughes4 more recently reported the same scale of relative gasphase acidities using a tandem mass spectrometer. These results indicate gas-phase acidities are opposite those known for the liquid phase. This phenomenon is of interest as it suggests that relative acidities are changed by removing the molecule to the gas phase. For several years Kebarle5 and coworkers have studied ion-solvent interactions in the gas phase. These studies are particularly informative in that they investigate ions surrounded by a cluster of solvent molecules without interference of the bulk of a solvent. I n a study of competitive solvation of hydrogen ion by water and methanol, Kebarle5 showed that the proton is not associated with any particular solvating molecule and that for small clusters, methanol is taken up preferentially to water. This preference for methanol is because methanol is more polarizable than water and thus provides enhanced charge stabilization. I n larger clusters water is favored because of its larger dipole moment. We have developed a novel technique, which to our knowledge has not been tried before, for studying ionsolvent interactions and how they relate to gas-phase acidities of alcohols. Equimolar mixtures of two alcohols were irradiated with 50-keV X-rays from a Siemens unit operated a t 26 mA. During irradiation an electric field of 500 to 1000 V/cm was applied between two 1-in. diameter copper electrodes. At a pressure from 50 to 100 Torr these field strengths were insufficient to contribute substantially to the kinetic energy of the ions but, as ion current measurements indicate, were sufficient to collect SO-lOO% of the ions produced. Alkoxy ions formed by dissociative electron capture migrate to an anode covered with a mossy silver deposit which binds the surviving ion. Since the ion collides many times in transient to the plate, equilibrium should be well established. After irradiation, the vessel was thoroughly evacuated and the anode was placed in a measured amount of distilled water, where the alkoxy ions were hydrolyzed. This mixture was analyzed on a Hewlett-Packard Model 5750 flame ionization gas chromatograph using a 12-ft 10% Carbowax 2031 on Chromosorb W column. The chromatograph was calibrated such that the yield of each alkoxy ion could be determined. Two sets of control experiments were performed. I n both sets the vessel was filled in an identical manner. One set was not irradiated although the electric field was applied while the other set was irradiated without applying the electric field. Analysis showed no alcohol present in the hydrolyzed sample of either set. If we had been able to carry out these experiments at very low pressures (below a few microns) the following equilibrium would be established. RO-
+ R’OH
ROH
+ R’O-
Table I : Equilibrium Constants for Alkoxide-Alcohol Proton Transfer Reactions4 Ke 9
Reaction
OHCH30CzH50CaH,OCaHg0CsHii0-
+ CH3OH + C2H50H + n-CaH,OH + n-CdHgOH
F2 CHaO-
+
+
+
+ HzO + CHaOH + CnHsOH + CaH7OH + CaH90H
C&OCaHrO& CaHg0n-CsHi1OH F! CsHiiOn-CaHiaOH C~Hiao-
*
+ CsHiiOH
80.0 10.9 2.6 1.4
1.4
1.7
The equilibrium constant for this reaction is a measure of the acidity of the alcohol since the most abundant alkoxy ion is the conjugate base of the most acidic acid. We were unable to do experiments in this pressure range since even experiments at 50 Torr required some 20 hr radiation time to obtain reliable quantities of product. Tiernan and Hughes4 have reported equilibrium constants for this process for a series of alcohols and water. Their equilibrium constants which are determined mass spectrometrically are calculated from the rate constant for the forward and reverse reactions. Table I reproduces their data. As can be seen from Table I the trends observed by Tiernan and Hughes agree with those reported by Brauman and Blair.213 As the polarizability of the alcohol increases so does the acidity. Since our experiments were at pressures between 50 and 100 Torr, the equilibrium studied involves solvated ions. Thus the actual process taking place is more complicated than a simple proton-transfer equilibrium between alkoxy ions. The ion-molecule collisions taking place actually establish the equilibrium concentration of methanol and ethanol in the ion clusters reaching the anode. By comparing ion current measurements with product collected at the anode, we find that there are 4 j = 2 alcohol molecules in the hydrolyzed sample per charge collected. The uncertainty reported is a standard deviation. Assuming this is representative of a gasphase cluster, we estimate that the cluster size in our experiments averages 4 including what is normally considered the central ion. Kebarle has pointed out that for small clusters incorporation of the most polarizable molecule is favored. In our experiments where the cluster size is small, ethanol should be more abundant in the cluster than methanol. Furthermore, since by analogy with Kebarle’s results the charge is not associated with any particular molecule, the most polarizable molecule should be favored for cluster sizes of one, i.e., single ions. For methanol-ethanol mixtures a t low pressure where the ions are unsolvated, CzHb0- should be favored over CHsO-. This is consistent with the results reported on gas-phase (4) T . 0. Tiernan and B. M. Hughes, Seventeenth Annual Conference on Mass Spectrometry and Allied Topics, 18-23,Dallas, Tex., May 1969. ( 5 ) (a) P. Kebarle, R. N. Haynes, and J. G. Collins, J . A m e r . Chem. SOC.,89, 5763 (1967); (b) P. Kebarle, Advan. Chem. Ser., No. 72 (1967). The Journal of Physical Chemistry, Vol. ‘76,N o . 14, 1971
NOTES
2228
acidity. As the number of molecules in the cluster increases the probability of the most polarizable molecule being in the cluster should decrease. From a series of 14 experiments with methanol and ethanol, the equilibrium constant which can probably best be thought of as the equilibrium for the replacement of a methanol with an ethanol in the solvation sphere was found to be 4.3 f 1.4. The reported uncertainty is a standard deviation. Since our cluster size averages 4, the preference for ethanol in the cluster should result in an equilibrium constant lower than the value of 10.9 observed by Tiernan and Hughes for isolated ions. In fact, as the cluster size continues to increase, one would expect that methanol would appear to be more acidic than ethanol. This is exactly the case at complete solvation or the liquid state. Using values of heats of formation of positive and negative ions reported by Krimler and B ~ t t r i l l we ,~ calculate that the enthalpy for the reaction CHDO- CzHsOH CH30H C2HsOassuming unsolvated species is -4.8 f 6 kcal. Since only a small change in entropy is expected for the reactionj5 the free-energy change should be close to the enthalpy change. From our value of the equilibrium constant, the enthalpy is -830 f 200 cal. Even though this is within experimental error of the literature value, the comparison is probably not valid because our determination includes solvation enthalpies. Although the data are not sufficiently accurate to warrant such a calculation, the difference in our values and the literature values represents the difference in solvation enthalpies between methanol and ethanol. As Kebarle and Yarndagn? have shown, this value is a function of cluster size. We have also obtained qualitative results for ethanol-isopropyl alcohol mixtures. However, the large uncertainty in the methanol-ethanol results were compounded by the lower vapor pressure of the higher homologs. For this reason, the project has been at least temporarily discontinued.
+
+
(6) P. Krimler and S. E. Buttrill, Jr., J . A m e r . Chem. Soc., 92, 1142 (1970). (7) R. Yamdagni and P. Kebarle, private communication.
Electron Spin Resonance Spectra of Copper Acetate in Acetic Acid Solutions
by Graeme Nyberg Physical Chemistry Diuiswn, L a Trobe University,' Bundoora, Victoria, Australia, SO83 (RecriPed January $9, 1971) Publication costs assisted by L a Trobe University
The properties of copper acetate monohydrate in the crystal are well studied. Structurally2 there are two T h e Journal of Physical Chemistry, "01. 76, N o . 1.6, 1971
adjacent copper ions bridged by four acetate groups, with the water molecules coordinating in the axial positions of the dimer. The magnetic behavior* above 20°K is that of a triplet, thermally excited from the singlet ground state. I n water solution, on the other hand, the spectrum resembles that of any other simple hydrated (monomer doublet) cupric ion. The question then arises as to the nature of the spectra in acetic acid solvents and the complex ions present. That such solutions do exhibit interesting electron resonance spectra has previously been m e n t i ~ n e d but , ~ without any elaboration.
Results and Discussion In pure glacial acetic acid there is no observable spectrum. As soon as water is added a signal begins to appear, and by 2% (by volume) its shape is unambiguously established. Continued dilution increases the signal intensity with little effect on the shape until the solvent is about 25% water, after which the intensity remains constant but the shape steadily approaches that of a pure water solution. Spectra for the two quoted dilutions are displayed in Figure 1, along with computed reconstructions. From the glacial acetic acid result we can immediately conclude that the copper species is a dimer, whose triplet resonance is so broad (in solution) that it is undetectable. While the addition of water must generate the monomer, dissociation of the dimer cannot be complete until about 25% dilution. Since the appearance of the absorption is rather unusual (more like a triplet half-field line than a solution doublet) this raises the possibility that the line shape might be determined by the chemical exchange between dimer triplet and monomer doublet species which are in equilibrium. Before pursuing this further, however, it was considered advisable to attempt a reconstruction with lorentzian components whose widths are given by the formula W(MI,TZ)= ( A
+ B M I + C M I ' ) ~+~ E/rz
This method, which has been previously described,s is easiest to apply when there is a sequence of spectra (over a range of temperatures) which encompasses one - as this immediately deterwith minimum line widths, mines r 2 = r o = d E / A . Here there is no such minimum, so the fitting process is substantially more difficult. The results displayed in Figure 1 are, however, (1) Work carried out at the Department of Theoretical Chemistry, University of Cambridge, Cambridge, England. (2) J. N. Van Niekerk and F. R . L. Schoening, Acta Crystallogr., 6, 227 (1953). iS'er. A , 214,451 (3) B. Bleaney and K. D. Bowers, Proc. R o y . SOC. (1952). (4) G. Nyberg, Mol. Phys., 12, 69 (1967). (5) G . Nyberg, ibid., 17, 87 (1969).
