J. Phys. Chem. 1981,85,488-492
488
appears that the Cl-Cz stretch does not couple significantly with the other two C-C single bonds. These latter bonds in 1-butene were considered to couple together to give a symmetric stretch at 854 cm-l and an asymmetric stretch a t 1020 cm-l. The C-C bond with the highest frequency stretch, at 1223 cm-' in cis-2-pentenenitrile, has a bond order greater than unity because of the extended conju-
gation between the unsaturated linkages already mentioned, and so may be considered to be a different type of bond from the other two C-C bonds.
Acknowledgment. We thank D. P. Braun of DuPont (Intermediates Marketing Division) for supplying the sample of 2-pentenenitrile.
Thermochemistry and Proton Bond Energies of Gas-Phase Proton-Bound Dimers of Aliphatic Alcohols D. S. Bomse' and J. L. Beauchamp" Arthur Amos Noyes Laboratory of Chemical Physics, Callfornia Institute of Technology, Pasadena, California 9 1125 (ReceiveO: September 3, 1980; I n Final Form: November 12, 1980)
Ion cyclotron resonance (ICR) spectroscopy is used to measure equilibrium constants for exchange reactions B' + BH+B * B'H+B + B where B' and B are the oxygen-containing bases ethanol, 2-propanol, 2-butanol, tert-butyl alcohol, and dimethyl ether. Thermochemical results indicate that the dimer dissociation energies D(BH+-B) all lie between 30 and 33 kcal/mol. This agrees very well with Kebarle's measurements of the D(BH+-B) value for oxygen-containing bases B = dimethyl ether, water, methanol, acetone, and dimethyl sulfoxide. Results are analyzed by comparison with a qualitative potential-energy surface for proton bond formation. Comparison is also made with energetics of the solvated negative ion species RO-(HOR). New estimates for proton affinities of 2-propanol and 2-butanol are reported.
Introduction A prominent application of gas-phase ion chemistry is the determination of intrinsic acidities and basicities of molecules free from complicating solvent effect^.^-^ Now that a substantial body of thermodynamic information on simple acid-base interactions has been acquired, the scope of this work has been expanded to studies of cluster formation.&' Quantitative measurements of clustering reactions, generalized in eq 1,provide a detailed description M*B, + B * M*B,+1 (1) of solvation energetics and help bridge the gap between gas and condensed phases. Typically Mi is a simple monotomic ion such as H+, Li+, Na+, K+, F-,Cl-, Br-, and Iand B = H20, NH3, et^.^-' One surprising result obtained by Kebarle is that, for all oxygen-containing bases B studied, enthalpy changes for the association reaction BH+ + B B2H+lie in the range AHo = -31 f 2 k~al/mol.~,"'~Only one alcohol, B = CH,OH, is among the compounds studieds9 Previous work in our laboratory has focused on the chemical consequences of strong hydrogen bonding in the gas phase.l'-'* In particular, ion-molecule chemistry of
-
(1)Josephine de Kirmin Fellow and Monsanto Fellow, 1979-1980. (2)J. I. Brauman and L. K. Blair, J.Am. Chem. SOC.,92,5986(1970); 90,6561 (1968). (3)J. F.Wolf, R. H. Staley, I. Koppel, M. Taagepera, R. T. McIver, Jr., J. L. Beauchamp, and R. W. Taft, J.Am. Chem. SOC.,99,5417(1977). (4)J. E. Bartmess and R. T. McIver, Jr. in "Gas Phase Ion Chemistry", Vol. 2.,M. T. Bowers, Ed., Academic Press, New York, 1979. (5)P. Kebarle, Annu. Reo. Phys. Chem., 28,445(1977),and references contained therein. (6)A. W. Castleman, Jr., P. M. Holland, D. M. Lindsay, and K. I. Peterson, J. Am. Chem. Soc., 100,6039 (1978). (7)R. L. Woodin and J. L. Beauchamp, Chem. Phys., 41, 1 (1979). (8)Y.K. Lau, P. P. S. Saluja, and P. Kebarle, J. Am. Chem. SOC.,in press. (9)E. P. Grumsrud and P. Kebarle, J. Am. Chem. SOC.,95, 7939 (1973). (10)A. J. Cunningham, J. D. Payzant, and P. Kebarle, J.Am. Chem. SOC.,94,7627 (1972). 0022-3654/81/2085-0488$01.25/0
secondary and tertiary alcohols is dominated by reactions which proceed through proton-bound intermediates."J2J4 Exchange reactions of proton-bound alcohol dimers were studied to verify the reversibility of cluster formation,14 and mechanisms for cluster formation at low pressures have been postulated.12J4 Recently in our laboratory we have investigated photophysical processes involving ions irradiated with CW infrared 1a~ers.l~ These studies have been extended to include proton-bound dimers of aliphatic alcohols.le To understand better the reaction dynamics and thermochemistry of these species, and to test the generalization implied by Kebarle's results, we have measured heats of formation and dissociation energies of proton-bound dimers of several aliphatic alcohols. These values are obtained by using the technique of ion cyclotron resonance (ICR) spectroscopy from equilibrium constants of exchange reactions, process 2, where B and B' represent B'
+ BHB+ * B/H+B + B
(2)
Lewis bases. McIverl' has studied energetics of the analogous hydrogen-bonded negative ions, RO-(HOR). Results of this study are compared with his reports of alkoxide cluster (11)J. L. Beauchamp, J. Am. Chem. SOC.,91,5925(1969). (12)M.C.Caserio and J. L. Beauchamp,J.Am. Chem. SOC.,94,2638 (1972). (13)D.P.Ridge and J. L. Beauchamp, J.Am. Chem. SOC.,93,5925 (1971). ,- - . -,. (14)J. L. Beauchamp, M. C. Caserio, and T. B. McMahon, J. Am. Chem. SOC.,96,6243 (1974). (15)R. L. Woodin, D. S. Bomse, and J. L. Beauchamp, J. Am. Chem. SOC.,100,3428 (1978);D.S.Bomse, R. L. Woodin, and J. L. Beauchamp, ibid., 101,5503(1979);R. L. Woodin, D. S. Bomse, and J. L. Beauchamp, Chem. Phys. Lett., 63,630 (1979);D.S.Bomse and J. L. Beauchamp, J. Am. Chem. SOC.,102, 3967 (1980). (16)D.S.Bomse and J. L. Beauchamp, J. Am. Chem. SOC.,submitted for publication. (17)R. T. McIver, Jr., J. A. Scott, and J. M. Riveros, J. Am. Chem. SOC.,95,2706 (1973).
@ 1981 American Chemical Society
Proton-Bound Dimers of Aliphatic Alcohols
The Journal of Physical Chemistv, Vol. 85, No. 5, 198 1 489 I
thermochemistry. Also, new estimates for proton affinities of some secondary alcohols are presented. Experimental Section The theory, techniques, and instrumentation of trapped-ion ICR spectroscopyhave been described in detail.18J9 The spectrometer used in this study was built at the California Institute of Technology and incorporates a 15-in. electromagnet capable of 23.4 kG. A three-section flat ICR cell equipped for ion trapping is used. All ICR experiments were carried out in the range lO-'-lO" torr, corresponding to neutral particle densities of 3 X log-3 X loll molecules ~ m - ~ Pressure . is measured with a Schulz-Phelps-type ionization gauge calibrated against an MKS Instruments Baratron Model SOHI-E capacitance manometer. It is expected that absolute pressure determinations are within *20% by using this method, with pressure ratios being somewhat more accurate. Sample mixtures are prepared directly in the instrument by using three sample inlets and the SchulzPhelps gauge. Typically, alcohol pressure ratios were adjusted so that two of three proton-bound dimers (i-e.,(R10H)H+(R20H) and (RIOH),H+) were in equilibrium while minimizing concentration of the third dimer ion, (R20H)2H+.Equilibrium constants were obtained both by determination of product-to-reactant ratios as well as by direct measurement of forward and reverse reaction rates. All reagents were obtained from commercial suppliers and were used without further purification except for freeze-pumpthaw cycles to remove noncondensable gases. Mass spectrometry revealed no detectable impurities. Some dehydration of tert-butyl alcohol occurs in the stainless-steel inlet system yielding 2-methyl-1-propene. This problem was minimized by conditioning the inlet system overnight with the alcohol. All experiments were performed at ambient instrument temperature, -30 OC. Results and Discussion Proton-bound dimer formation occurs by one or more sequences of ion-molecule reactions following ionization of secondary or tertiary alcohols. These reactions are described in detail elsewhere.11J2J4Briefly, ions which contain labile protons (such as protonated alcohols or a-cleavage fragment ions) induce ionic dehydration of the neutral alcohol to form hydrated ions, such as (ROH)H+(OH,). Displacement of water by alcohol yields proton-bound alcohol dimers. Proton-bound dimers of ethanol (which does not undergo ionic dehydration) were formed by exchange reactions (i.e., eq 2) in mixtures with 2-propanol. Equilibrium constants for exchange reactions (eq 2) were measured for Lewis bases ethanol, 2-propanol, 2-butanol, tert-butyl alcohol, and dimethyl ether. Results for one equilibrium, reaction 3, are presented in Figures 1and 2. k $ EtOH + (i-PrOH)H+(EtOH)
i-PrOH + (EtOH)2H+
(3) EtOH and i-PrOH represent ethanol and 2-propanol, respectively. Figure 1 shows ICR trapped-ion data for the two proton-bound species in reaction 2. After -0.5-s trapping time, the ratio of ion intensities reaches the constant value [ (i-PrOH)H+(EtOH)]/[ (EtOH)2H+]= 1.55 (18) T. A. Lehman and M. M. Bursey, "Ion Cyclotron Resonance Spectrometry",Wiley-Interscience, New York, 1976; J. L. Beauchamp, Annu. Rev. Phys. Chem., 22,527 (1971). (19)T. B. McMahon and J. L. Beauchamp, Rev. Sci. Instrum., 43,509 (1972).
I
I
I
I
I
J
I
I
I
I
[C~HSOH] = 2 . 6 IO-' ~
Torr
= 1.3 x IO-'
[i-C,H,OH]
Torr
0 '
I 0
0.2
0.4
0.6
1.0
0.8
Time (sec.)
Flgure 1. Ion intensities as a function of trapping time for (EtOH)*H+ and (i-PrOH)H+(EtOH formed by ion-molecule reactions in a mlxture containing 2.0 X 10- torr of EtOH and 1.2 X lo-' torr of i-PrOH. Ion intensities must be divided by ion mass to correct for differences in ICR sensitivity. Ionization Is by a 10-ms pulse of 70-eV electrons.
b
I
I
I
I
I
I
-
d e t e c t IEtOHIH'li-PrOH)
eject (EtOH)H+(i-PrOHl
05
07
eject (EIOHI~H+
09
05
07
09
Time ( s e c )
Flgure 2. Data for direct measurement of forward and reverse rates of reaction 2. Plot at left shows exponential decay of normalized (EtOH),H+ intensity following the start of double-resonance ejection of (EtOH)H+(/-PrOH)at 500 ms (vertical arrow). Slope of the semilog plot indlcates that the forward reaction rate is 4.2 s-'. Data at right show decay of normalized (EtOH)H+(i-PrOH) signal resulting from doubleresonance ejectlon of (EtOH),H+. Reverse reaction rate is 3.4 s-'. Reagent pressures and ionization conditions are identical wlth those of Figure 1.
f 0.05 as equilibrium is established. The ratio of neutral alcohol pressures is [EtOH]/[i-PrOH] = 20, yielding an equilibrium constant for reaction 3 of K = 31 f 6 and AGO = -2.0 f 0.2 kcal/mol. In Figure 1,the decrease in proton-bound dimer intensities is due primarily to ion loss from the trapping cell. At the low i-PrOH pressure used, formation of (iPrOH)2H+is very slow and does not reach equilibrium concentration during the time scale of the experiment shown in Figure 1. The intensity of (i-PrOH),H+ does not exceed 30% of the total proton-bound alcohol dimer population. Figure 2 shows data for a direct measure of both forward and reverse reaction rates of process 2. ICR double-reson a n ~ e ' ~ejection J~ of (i-PrOH)H+(EtOH) results in an exponential decay of the (EtOH)2H+population as the latter ion reacts to form (i-PrOH)H+(EtOH)and is ejected from the cell. Double-resonance ejection occurs during
490
The Journal of Physical Chemistry, Vol. 85, No. 5, 1981
Bomse and Beauchamp
TABLE I: Energetics of Proton-Bound Dimer Exchange Reactions AGO; T A S oc a d , " AH" ex? PA(B'),~ B' B kcal/mol kcal/mol kcal/mol kcal/mol C,H,OH C,H,OH 2.3 0 2.3 191.4' C,H,OH i-C,H,OH 0.5 0.4 0.9 i-C,H,OH i-C,H,OH -0.3 0 - 0.3 194.4 f 0.5d i-C, H, OH sec-C,H,OH - 1.6 0.4 - 1.2 sec-C,H,OH sec-C,H,OH -1.8 0 - 1.8 194.9 f 0.5d i-C,H,OH t-C,H,OH - 1.8 0.4 - 1.4 t-C,H,OH t-C,H,OH -2.5 0 -2.5 196.6e i-C,H,OH -0.7 0.4 - 0.3 (CHAO 0 194.8' (CH3 1 2 0 0 (CH,),O " Relative to [(CH,),O],H+ as indicated by reaction 5. Uncertainties are t 0 . 3 kcal/mol. All thermochemical data are at Assumes PA(",) = 207.0 kcal/mol (see Appendix). ' Reference 3. Estimated, see Appendix. e Reference 298 K. 23.
TABLE 11: Hydrogen-Bond Strengths in Proton-Bound Dimers of Oxygen Basesa isectH,O CH,OH C,H,OH (CH,),O C,H,OH C,H,OH (CH,),CO C,H,OH (CH,),SO H*O 32b 3' 6 40.5d 42.5' 42' 42.5' 44.5' 42' 5 3' CH OH 25' 33.1b C,H,OH 24d 31.7e 3 6e 22.5' 30.7b 31.4e (CH3)20 i-C,H,OH 22.5' 32.0e 31.0e 31.4e 32.4e 30.7e 22.5' 32.ge 32.4e sec-C,H,OH 21.0' 30f (CH3)2C0 20.5' 28.5e 29.6e t-C,H,OH (CH3)2S0 15' 31f a Bond energies D(B'H+-B)in kcal/mol. Reference 9. Estimated by using empirical method of ref 25. Reference 5. e This work. Uncertainties are 20.3 kcal/mol plus uncertainty in proton-affinity measurement (see Table I). f Refere n 2 8.
B'
B
,
alternate ICR trapping cycles allowing measurement of (EtOH)2H+ intensity with and without ejection of (EtOH)H+(i-PrOH). Dividing these two signals yields a normalized (EtOH)2H+intensity. The left-hand portion of Figure 2 shows the normalized population of (EtOH),H+ as a function of trapping time when double-resonance ejection of (i-PrOH)H+(EtOH) begins at 0.5 s. The ordinate scale is logarithmic, and the negative of the slope of the line is the forward reaction rate, 4.2 s-l. Combined with a measured 2-propanol pressure of 1.3 X lo-' torr, this result yields kf = 9.97 X cm3 mol-'s-l. Data shown in the right-hand portion of Figure 2 yield a rate of 3.4 s-l for the reverse process in reaction 3 and k, = 4.04 X lo-" cm3mol-l s-l. Combined resulta in Figure 2 give K = kJk, = 23.5 & 5,corresponding to AGO = -1.9 f 0.2 kcal/mol for reaction 3. The two methods (Figures 1 and 2) for measuring the free-energy change of the exchange reaction agree within experimentaluncertainty. For the experiment illustrated in Figure 1, a total pressure of 3 X lo4 torr corresponds to an ion-neutral collision rate of 102 s-l. Thus each ion suffers 100 collisions with the major component and 10 collisions with the minor component, and reaction 3 "cycles" at most five times during the 1-strapping period. The ratio of ion abundances does not change at even longer times, however, and it appears that the reaction has come to equilibrium under the indicated conditions. A different mixture of EtOH and i-PrOH ([i-PrOH] = 5.3) was used to evaluate the equilibrium constant for the exchange process involving (i-PrOH)2H+, reaction 4.
-
Scheme I 2(CH3)20
+
A%.l
[(CH3)20]2H+
+
B'
-
D IB' H+- B
B'H'B
+
B
+30 7 kcol/mol
+
~ ( c H ~ ) ~ oB ' H +
+
B
-
lPA(CH30CH3) PA(B')
(CHJ20H+
+ (CH3),0 + 8' + B
It is assumed that proton-bound species undergo free internal rotation about the oxygen-proton bond. Thus the symmetry number of a given Lewis base is invariant whether or not the molecule is bound in a complex. Symmetrically substituted proton-bound dimers such as (EtOH)2H+ possess a 2-fold rotational symmetry not present in asymmetric dimers such as (EtOH)H+(i-PrOH). Thus for reaction 3, TASocdd = RT In 2 = 0.41 kcal/mol. In Table I, AGO and AHo, values are computed relative to reaction with [(CH3)20]2H+,process 5. Also listed are B + B' + [(CH,),O],H+ BH+B' + 2(CH3)20 (5)
absolute proton affinities of the bases used in this study. Note that proton affinities of some alcohols are not experimentally determined and must be estimated. These estimates are discussed in the Appendix. Cluster dissociation energies, D(B'H+-B), are calculated from the thermodynamic cycle shown in Scheme I. Kebarle's value of 30.7 kcal/mol is used for the dissociation energy of the (CH3)20proton-bound dimer? Experimentally determined values of AH",, are taken from Table I. The data obtained for the alcohol proton-bound dimers have been augmented by cluster dissociation energies re(i-PrOH)H+(EtOH) + i-PrOH ported for other oxygen-containing ba~es,l*~JO and the (i-PrOH)2H+ EtOH (4) combined set of D(B'H+-B) values are presented in matrix form in Table 11. For asymmetric ions, B'H+B, dissociUnder these conditions the concentration of (EtOH)2H+ ation is assumed to give B'H+ + B and not B' + BH+. is nearly zero once equilibrium is established. A striking result is observed by reading the entries along Resulta of all measured equilibria are presented in Table the main diagonal of Table 11. ,All measured dissociation I. Listed are AGO, ASodcd and AHo,, for the exchange energies of symmetrically substituted proton-bound dimers reactions. Entropy changes, ASodcd, are calculated from ratios of symmetry numbers of reactants and p r o d u ~ t s . ~ lie between 30 and 33 kcal/mol for oxygen-containing
+
Proton,-BoundDimers of Aliphatic Alcohols
The Journal of Physical Chemistry, Vol. 85, No. 5, 1981 491
one assumes this distance to be correct for oxygen-containing ions B’H+B, it is apparent that steric interactions are not very important in determining energetics of proton-bound dimer formation for the species listed in Table 11. The one possible exception is the bulkiest base studied, tert-butyl alcohol. D(BH+-B) = 29.6 kcal/mol for B = tert-butyl alcohol is the lowest measured dimer dissociation energy (Table 11). However, the steric effect, if any, is very slight, and no discernible change in bond strength with variation in molecular bulk exists among the other alcohols. Attempts were made to include in this ICR study exchange reactions of other oxygen-containingLewis bases. I 1 Among these compounds are several dialkyl ethers. 0-H+.,.O Coordinate However, ambient temperature equilibrium constants for Figure 3. Qualitative potential-energy surfaces for proton motion reaction 1 are beyond the ICR dynamic range when choice between the two oxygen atoms in the summetrlc proton-bound dimer of one base was restricted to those compounds listed in BH+B where B represents an oxygencontaining Lewis base. Surface Table I. These negative results are consistent with mea1 is for ”infinite” oxygen-oxygen separation. Surface 2 represents oxygen-oxygen separation in the lowestenergy configuration of BH’B. sured proton-affinity data and the inference that DVertical energy scales of curves 1 and 2 are not directly comparable. (BH+-B) = 31 f 2 kcal/mol for all oxygen-containing bases. bases. The 3 kcal/mol variation among measured energies In a study of solvated alkoxide ions (ROH-OR-), McIver D(BH+-B) is minute compared to the 38 kcal/mol range observed that methoxide (a strong base) forms stronger in proton affinities encompassed by the Lewis bases listed. hydrogen bonds than does ethoxide (a weaker base) to a In fact, differences in proton-bound alcohol dimer dissosingle neutral methanol m01ecule.l~ Thus the solvation ciation energies reflect the uncertainties in alcohol proton energetics of proton-bound alkoxide anions is qualitatively affinities more than any systematic variation with physical similar to those of protonated alcohols. Although absolute properties of the alcohols. The similarity in proton bond values for D(R0--HOR) are not reported, the difference strengths can be understood by examining Scheme I.8*20 in energies between the case R = CH, and the case R = Consider the case B’ = B. Increasing the proton affinity CzH5is only 0.3 kcal/mol with CH30-(HOCH3)possessing of B results in a large endothermic AHHo,(see Table I) and the stronger bond. In comparison, CH30- is more basic an equally large but exothermic proton-affinity difference than CzHbO-by 1.9 kcal/mol. Thus, the reported results relative to (CHJ2O. The effects nearly cancel to yield a are not inconsistent with the possibility of similar dissodimer dissociation energy of 31 f 2 kcal/mol. Similar ciation energies for all complexes RO-(HOR). Definitive results have been reported for proton-bound dimers of conclusions cannot be drawn from this limited amount of aliphatic amines in which the symmetric species yield information. D(BH+-B) values between 21 and 25 kcal/mol for B = NH3, CH3NH2, (CH3l2NH,and (CH3)3N.20 Conclusions A qualitative potential-energysurface for proton motion Dissociation energies, D(BH+-B), of proton-bound dibetween the two oxygen centers in BH+B is shown in mers are nearly identical for all oxygen-containing bases Figure 3. Curve 1 depicts the system when the two oxyexamined to date. When one combines the results of this gen-containingbases are at “infinite” separation. The wells study with Kebarle’s6*9J0~20~21 data, the dimer dissociation represent BH+ formation with the well depth equal to the energies all lie between 30 and 33 kcal/mol for bases proton affinity of B. No interaction between bases occurs. containing cr-bonded or r-bonded oxygen atoms. The Increasing the base strength of B results in greater delothermodynamic cycle shown in Scheme I illustrates that calization of charge and increased electron density on the increasing the base strength of B stabilizes both ions BH+ proton. As the second base approaches BH+, a new proton and BH+B by nearly the same amount, leaving D(BH+-B) bond is formed. However at short 0-0 separation, ininvariant. This result is understood in terms of a potenteractions across the proton between basic lone pairs on tial-energy diagram, Figure 3, and consideration of oxythe two oxygen atoms are repulsive and result in a weakgen-oxygen repulsive interactions as well as attractive ening of the first proton bond. Curve 2 in Figure 3 shows oxygen-proton bonding. the potential-energy curve for the proton-bound dimer. Nearly identical bond strengths for all symmetrically The two oxygen atoms generate a symmetric bonding well substituted dimers listed in Table I1 provide evidence for for the proton (there may or may not be a double well) with a simple hydrogen bond structure, 1,common to all ions the 0-0 distance determined by opposing effects of the H R u + acid-base interaction and repulsive orbital overlap. In asymmetrically substituted dimers, bonding strength of the second base decreases as the difference in proton affinities increases.6~20~21 A limiting case is reached where 3 1 2 the second bond is purely electrostatic with virtually no electron donation by the weaker base to the proton. This BH+B. Complex interactions,such as the four-center bond explains the decrease in D(B’H+-OH2)values as one reads shown in 2, are not possible for species without extra, labile down the first column in Table 11. hydrogens (Le., (CH3)2O and acetone). Hence only Typical experimental and theoretical values for oxystructure 1 is available to all bases. Furthermore, calcugen-oxygen separations in (H20)2H+are 2.4-2.6 A.22 If lations have shown that 1 (R = H) is the most stable geometry for (Hz0)2H+with species 2 being higher in energy.22 Similar arguments can be used to exclude 3 as (20) It. Yamdagni and P. Kerbarle, J.Am. Chem. SOC.,95,3504(1973). (21) J. D. Payzant, R. Yamdagni, and P. Kebarle, Can. J. Chem., 49, 3308 (1971);R.Yamdagni and P. Kebarle, J. Am. Chem. SOC.,95,3504 (1973);W.R.Davidson, J. Sunner, and P. Kebarle, J. Am. Chem. SOC., 101,1675 (1979).
(22) P. A. Kollman and L. C. Allen, J. Am. Chem. SOC.,92,6101(1970);
W.P. Kraemer and G. H. F. Diercksen,Chem. Phys. Lett., 5,463(1970).
492
The Journal of Physical Chemlstry, Vol. 85, No. 5, 1981
175 10 0
II 0
12 0
13 0
I o n i z a t i o n Potential ( e V )
Figure 4. Plot of proton affinities vs. ionization potentials for water and alcohols used In this study. Experimentally determined proton afflnities are represented by filled circles, whereas estimated values are open circles. Proton-affinity data are from ref 3 and assume PA(NH3) = 207.0 kcal/mol. Ionlzatlon potentials are from ref 26.
geometry of the acetone proton-bound dimer. Acknowledgment. This work was supported by the U.S. Department of Energy. We are grateful to Professor Kebarle for communicating his results before publication. Appendix Proton transfer equilibria are not observed among ionmolecule reactions of most secondary and tertiary alcohols. Instead, the protonated alcohols react rapidly with neutral alcohols yielding dehydration p r o d ~ c t s . ' ~ J ~Thus, J ~ accurate proton affinities (usually determined by proton transfer equilibria) have not been determined for these alcohols. One exception is tert-butyl alcohol, for which the heat of formation of the protonated species, AH,O(tC4HgOH2+) = 93.9 kcal/mol, is obtained from equilibrium A1.= This predicts PA(t-C4HgOH)= 196.6 kcal/mol. The
(CH3)3C++ H20
(CH&COH2+
(AI)
(23)K. Hiraoka and P. Kebarle, J. Am. Chem. SOC.,99,360 (1970). (24)J. Long and B. Munson, J. Am. Chem. SOC.,96, 2427 (1973). (25)W.R.Davidson, J. Sunner, and P. Kebarle, J.Am. Chem. SOC., 101,1675 (1979). (26)H. M.Rosenstock, K. Draxl, B. W. Shiner, and J. T. Herron, J. Phys. Chem. Ref. Data, Suppl., 6,1 (1977).
Bomse and Beauchamp
proton affinity of tert-butyl alcohol is tied to the standard proton-affinity scale (relative to PA(",)) by the heat of formation of (CH3),C+which is calculated from the proton affinity of 2-methyl-1-propene. Proton affinities of 2-propanol and 2-butanol are estimated by comparison with substituent effects shown by other Lewis bases. For example, differences in proton affinities between aliphatic alcohols ROH and the corresponding amines RNH2approach a value of 26.7 kcal/mol as the size of the unbranched alkyl chain increases? Using the reported value of PA(i-C3H7NH2)= 221.2 kcal/mol (relative to the proton affinity of NH3) predicts P A ( t C3H70H) = 194.5 kcal/mol. Similarly, proton-affinity differences between RCHO and ROH are -1.0 kcal/mol for R groups larger than CHa3 A value of 195.3 kcal/mol is reported for PA(i-C3H7CHO)which yields an estimate of PA(i-C3H70H) = 194.3 kcal/mol. The two estimates agree well, and their average is listed in Table I. The error limits on the estimate fall within the limits placed on the proton affinity of 2-propanol by the experimental study of Long and M ~ n s o n . ~ ~ There are no reference compounds available which contain sec-butyl groups. Instead, the 2-butanol proton affinity is estimated by assuming comparable increases in basicity in going from 1-propanolto 2-propanol as between 1-butanol and 2-butanol. Thus an estimate of PA(secC4H90H) = 194.9 kcal/mol is obtained. There is, in general, an inverse correlation between the adiabatic ionization potential of a compound and ita proton affinity.14 Figure 4 is a plot of proton affinities vs. ionization potentials for water and alcohols pertinent to this study. Both experimentally determined (filled circles) and estimated (open circles) proton affinities are grouped close to the straight line (which is a least-squares fit excluding H20) drawn in the figure. The slope of the line equals -0.43 and is substantially different from unity, which is the value normally observed in such correlations. Apparently alkyl substituents stabilize or delocalize charge more effectively in the radical cations than in the conjugate acids of aliphatic alcohols. All proton-affinity data in the present work are given relative to PA(NHJ = 207 kcal/mol. This number remains controversial but certainly lies between 204 and 208 kcal/rn01.~~In the absence of a resolution of this important problem, we have somewhat arbitrarily picked the value 207 kcal/mol since this number has been used in related studies. (27)F. A. Houle and J. L. Beauchamp, J.Am. Chem. SOC.,101,4067 (1979).
