nichol, sunner 1988

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International Journal of Mass Spectrometv and Zon Processes, 84 (1988) 135-155Elsevier Science Publishers B.V., Amsterdam - Printed in The Netherlands

135

KINETICS AND THERMODYNAMICS OF PROTONATIONREACTIONS: H,O +(H ,O), + B = BH+(H ,O), + (h - 6 + 1) H,O ,WHERE B IS A NITROGEN, OXYGEN OR CARBON BASE

GORDON NICOL, JAN SUNNER * and P. REBARLE * *Chemistty Department, University of Alberta, Edmonton T6G 2G2 (Canada)(First received 19 November 1987; in final form 20 January 1988)

ABSTRACTThe extent of proton transfer from H90+ (H20),, to compounds B due to proton transfer

reactions (PI)Hs0+(H20)h+B=BH-(HzO)b+(h-b+1)H,0was studied with a pulsed-electron high-pressure mass spectrometer (PHPMS). TheH30+ (H,O),,were equilibrium hydrate populations in a third gas (air or methane) containingwater pressure in the torr range and B in the sub-millitorr range. When a limited reactiontime was available (t -z 0.5 ms), the compounds B fell into three groups regarding theobserved extent of PT. Compounds with gas-phase basic&y (GB) higher than 200 kcal mol-experienced maximum proton transfer. For these B, which are mostly nitrogen bases, theextent of proton transfer was proportional to the rate constant k, (kinetic control) and thekm were found close to the ion/molecule collision-rates. Compounds B with GB -z 200 kcalmol- were found to reach PT equilibria (thermodynamic control). The oxygen bases in thisgroup showed an extent of PT that decreased fairly regularly as GB(B) decreased. The extentof PT observed for a group of B under thermodynamic control was much less than for oxygenbases of the same GB. This group of compounds, which included the carbon bases pyrrole,furan, and thiophene, corresponds to BH+ which form hydrates of very low stability.Hydration equilibria were measured for BH+ of pyrrole, furan, and thiophene. The rateconstants for PT involving the carbon base pyrrole were also measured and found to decreaseas the PT exothermicity decreased.

INTRODUCTIONThe impetus for the present work arose out of an experimental study ofthe sensitivitiesof various compounds, B, in a commercial atmospheric-pres-

* Present address: Department of Chemistry, Montana State University, Bozeman, MT59717, U.S.A.* * To whom correspondence should be addressed.0168-1176/88/$03.50 8 1988 Elsevier Science Publishers B.V.


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136sure chemical-ionization (APCI) analytical mass spectrometer [l]. In theAPCI source, the analytes B, trace organic compounds present in atmo-spheric air, are protonated by the reagent ions H30+(HzO)h. At the prevail-ing relatively high water pressure due to the natural humidity of the air ( - 5torr) and low analyte pressures P, < 10S6 torr, the protonation reaction isrelatively slow compared with reactions hydrating the H,O+ and the productBH+. Thus, the reagent ions H,O+(H,O), are equilibrium hydrate popula-tions and the protonation of B is followed by a very rapid achievement ofequilibrium BH+(H20)b populations.- The method is very well suited for the detection of trace compounds(pollutants) in air; however, it suffers from one problem. The sensitivities ofdifferent classes of compounds can be very different. In particular, someenvironmentally important compounds like mercaptans, thioethers,thiophenes, and halobenzenes are detected with relatively very low sensitivi-ties.The factors responsible for the observed very different sensitivities wereexamined on the basis of experiments performed with the APCI instrument[l] and available literature information. Although the major causes could beelucidated in a general way, more direct experimental information on thekinetics and thermochemistry of the proton transfer from H,O+(H,O), to Bwas needed in order to put the findings on a firm base. The work describedin the present publication deals with these additional experiments andinterpretation. The present experiments were performed with one of ourpulsed-electron high-pressure mass spectrometers (PHPMS). This methodpermits the time resolved detection of the ion concentrations in a high-pressure ion source ( - 3-5 torr) after a short electron pulse [2].While the work to be described was prompted by the desire to understandthe factors governing the APCI sensitivities, the findings are of more generalinterest. Thus, the changes of reactivity of ion clusters like H30+(HzO),,with degree of hydration, or generally solvation, are of fundamental signifi-cance (see previous work by Bohme et al. [3,4] and Henchman and co-workersIS]). Another factor that turns out to govern the sensitivities is the stabilityof the hy&ztes BH+(H,O), and this fact leads one into the general area ofgas-phase ion solvation [6].The efficiency of protonation by H30+(HzO),, of compounds B presentas traces in the atmosphere is not only of interest in the analytical APCItechnique. The H30+(HzO)I, species are the major naturally occurring ionsin the troposphere and stratosphere as well as in the D region of theionosphere [7]. The BH+(H20)b ions resulting from the protonation reac-tions of traces of B can be detected with air- or rocket-borne mass spec-trometers and can be used as probes for the presence of traces of B [7].Examples of possible stratospheric B are CH,O, CH,OH, HNO,, and NH,.


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In order to relate the observed abundance of BH+(H,O)b to the unknownconcentration of B, one needs to understand the kinetics of the protonationreactions. A laboratory study of the protonation of B by H,O+(H,O),undertaken with the purpose of identifying the possible core ions BH+ andthus the precursors B of the observed [7] stratospheric ions, was reported byFerguson and co-workers [7b]. The investigation dealt only with B = CH,OHand CH,O and H30+(H20),, with h I 3, since these were relevant reactantsin the stratosphere. The present work provides data for a larger variety ofcompounds B and H,O+(H,O),, with higher values of h. Thus, it extends thelaboratory data to conditions present in the lower stratosphere and particu-larly the troposphere.EXPERIMENTAL

The experiments were performed with one of our pulsed-electron high-pressure ion source mass spectrometers [2], which utilizes a quadrupole formass analysis and an ion counting and multiscalar system for real timeresolution of the ion signals. The mass-dependent transmission of thequadrupole was determined by comparing low-pressure ion source electron-impact spectra of various compounds with those obtained with a magneticsector instrument.The experiments were generally performed at 3 torr total pressure.Methane was used as the major gas. The dominant ions resulting fromelectron ionization and subsequent ion/molecule reactions in methane areCHf and C,Hl. These protonate water leading to H,O+. At the prevailinghigh concentration of H,O, hydration equilibria establish rapidly and theequilibrated H30+(HzO),, are the only ions observed in the absence of Band impurities.Approximately 4 X 10T5 torr Ccl, was added to all reaction mixtures.The capture of electrons, e + Ccl, = Cl- + Ccl,, leads to ambipolar posi-tive ion-negative ion diffusion which is slower and improves the conditionsfor the measurements [2].RESULTS AND DISCUSSION

Factors responsible for the very different proton transfer efficiencies fromH,O +(H,O), to compounds B

The sensitivities for some representative compounds B relative to that forthe high-sensitivity compound pyridine observed in the APCI measurements[l] are shown in Table 1. The sensitivities are proportional to the extent of


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138TABLE 1Total rate constant, k,, and equilibrium ratio, R,, for proton transfer reactions *B GBWb Rqd AG&(PT) i k,(PQc k,(FQe SENS fMeOH 174 340 -6 s 1.7 3x1O-3 TEtOH 180 2730 -7 1.9 1.7 3~10-~ TMeCN 181 8500 3.0 6~10-~ TFtUan 185


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0 0 0 . 2 0 . 4 0 . 6 0 . 8 1 . 0 1 . 2 1 . 4 1 . 6 1 . 8 2 . 0Time (ma@

Fig. 1. Observed time dependence of ions at 32C, 2 torr CH,, 1.1 torr HzO, and 10e4 torrpyridine. The H,O+ (H,O), maintain hydration equilibrium as they proton transfer topyridine. The py-ridine H+ hydrates are converted to the dimer (pyridine),H+. This reactionwould be suppressed at much lower pyridine concentrations.

The pressure conditions in the APCI apparatus are: p(air) = 700 torr,p(H,O) - 5 torr; p(B) < 10s6 torr, ion reaction time - 300 ps at roomtemperature. These conditions can not be reproduced in the PHPMS ap-paratus. However, experiments under conditions that are reasonably close,i.e. p(H,O) up to a few torr and p(B) down to 10V5 torr, can be performedand were found to provide direct insights into the factors responsible for thedifferent sensitivities of the three groups of compounds.The ion-time profile for a typical K group compound (pyridine) observedwith PHPMS is shown in Fig. 1. The H,O l (H,O) ,, ions, which are in rapidclustering equihbrium [Eq. (l)] and thus at constant ratios are not shownindividually in Fig. 1. AII the H30+ hydrates decrease at the same rate dueto proton transfer to pyridine [see Eq. (2)]. The protonated products ofreaction (2) rapidly reach equilibria via Eq. (3).H30+(H20)h_1 + H,O = H@+(H20)h (1)H30+(H,0), + B = BH+(H,O), + (h - x + l)H,O (0


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B is formed which dissociates to give the products shown in reaction (2).Thus, the reaction could also be called an H,O displacement by B. We usethe expression proton transfer since it is simpler and the mechanisticdistinction is unimportant in the present context.An additional reaction, the formation of the proton-bound pyridine dimercorresponding to &H+, is also evident from the results in Fig. 1. It is alsoobserved that, while the H,O+(H,O),, disappear completely, the BH+(H,O)band &H+ reach equilibrium at longer reaction times. For the higher p(H,O)and much lower p(B) prevailing under the APCI conditions, one expectsthat BH+(H,O), will be the major and &H+ a very minor product.Plots of the logarithm of the H30+(H,0),, ion intensities from normal-ized plots like that in Fig. 1 versus time t lead to the average pseudo-first-order rate constant v, = k,[B] for the proton transfer reaction (2), from

which k , can be evaluated since [B] is known. Thus, k , = 1.7 x 10d9molecules-1 cm3 s- for the conditions in Fig. 1. This is close to the orbitingcollision limit, i.e. k A m , that can be calculated with the average dipoleorientation theory [ 81.At the much lower B concentrations (PB = 10T6 torr, [B] = 3 X lOlomolecules cmm3) and shorter reaction times (At = 300 ps) prevailing for theAPCI conditions, the proton transfer reaction (2) will be very incomplete,i.e. only a small fraction of the H,O+(H,O),, ions will have had time toreact. This fraction can be evaluated from

(4)Substitution of the numerical values leads to a fraction of 1.4 x 10m2, whichmeans that less than 2% of the H,O+(H,O)* react under these conditions.Thus, the extent of proton transfer is kinetically limited, i.e. under kineticcontrol for pyridine and for the other compounds of group K. The observedsimilar sensitivities for the group K compounds (see Table 1 and ref. 1)derive from the rather similar rate constants expected for collision-limitedrates [8].The time dependence of ion intensities for compound B = acetone fromthe T group is shown in Fig. 2. Proton transfer from H30+(HzO),, leads torapid decrease of the hydronium hydrates and increase of BH+(H,O)bfollowed by formation of &H+ and B,H+(H,O),. This is similar to thesituation observed above for pyridine; however, in the present case, theH,O+(H,O), ions do not disappear completely but achieve equilibrium withthe BH+(H,O), and B,H+(H,O), ions after some 500 ps. From the initialdecrease of the H30+(HzO)h, an average rate constant k , = 3.2 X 10m9molecules- cm3 s-r can be obtained. This is in the expected collision limit


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2I2 -ii5 w-I- f.k,OH+W,O),6dr _

0 n 0 0 . 2 0 . 4 0 . 6 0 . 6 1 . 0 1 . 2 1 . 4 1 . 6 1 . 6 2 . 0Time (msec)Fig. 2. Observed time dependence of ions at 32C, 2 torr CH,, 1.1 torr HzO, and 6~10~~tom acetone. The equilibratedH30+ (H,O), proton transfer to acetone, however the reactionreaches equil ibriumat - 0.5 ms.

k range (see Table 1). Rate constants k, obtained for other compoundsB*?the T group are also shown in Table 1.For the higher p (H,O) and lower p(B) pressures prevalent under atmo-spheric conditions, the proton transfer equilibria will be shifted muchfarther towards the H,O+(H,O), than is the case for conditions leading tothe results shown in Fig. 2. Under these conditions, the attainment ofequilibrium is faster than in Fig. 2 so that the extent of proton transfer tocompounds B of group T is thermodynamically controlled, i.e. the protontransfer for equilibrium is attained within the available reaction time ( - 300Ers).The ion-time dependence for a low-sensitivity group L compound, B =furan, is shown in Fig. 3(a). The proton transfer equilibrium 5 is reachedrapidly and the BH+(H,O), ions at equilibrium are much less abundantthan was the case for acetone (Fig. 2) even though, in the present case, P, is5 times higher and PHzO ten times lower. Thus, the proton transfer Eq. (5)at equilibrium is far less efficient for furan than for acetone while the twocompounds have essentially the same gas-phase basicity (see Table 1).Similar results were obtained with the PHPMS for other compounds of thelow-sensitivity group L, which are thus characterized by thermodynamiccontrol but very unfavorable position of the observed proton transferequilibria.A decisive clue to the reasons for the unfavorable equilibrium protonationof the L group of compounds was the observation that the hydrates


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FH+U-$O),,,ILII

' 0 ' I. 2 0 . 4 0 . 6 0 . 6 1 . 0 1 . 2 1 . 4 1 . 6 1 . 6 2 . 0

Time (msec)1 0 0

\ ( b )

Time (msec)Fig. 3. (a) Observed time dependence of ions at 32O torr CH4,

is achievedrapidly. of the equilibrium due to inefficient hydration ofprotonated furan (see in percent h 2, 0.1; h = 3,40.9; h = 4, h 5, 5. of ions at higher of198OC, 2 1.3X10- torr furan. Proton transfer equilibriumshifted in favor of furan due to lower h in H30+(H20),, reactants: h =l, 3.3; h = 2, 82;h = 3, 14.7%.

BH+(H20)b of these compounds had very much lower b, generally b = 0 orb = 1, than T group compounds with similar GB. The consequences of themuch lower stability of the hydrates of the L group are explored in the nextsection.


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Thermodynam ics of the proton transferThe extent of proton transfer from H,O+(H,O),, to B occurring within a

limited time depends on the kinetics and thermodynamics of the reactions(5). The thermodynamics will be discussed first since they are also ofsignificance to the kinetics.H,O+(H,O), + B = BH+(H,O), + (h + 1 - b)H,O (5)AGS GB(H,O) - GB(B) + AG,O,O(H,O+)- AG&(BH+) (6)M+(H,O) = M+ + n H,O AG; = AG;,,(M+) (7)M+(HzO)._i + H,O = M+(H20), AGS AG,O_,,,(M+) (8)AGnp(M+) = AG,q,_i(M+) + . . . AG,q,(M+) (9)AG:-,,, = -AG,q,_i

The thermodynamic data required for evaluation of the BH+(H,O),concentrations relative to the H,O+(H,O), concentration at proton-transferequilibrium, Eq. (5), are given in Eq. (6). Comprehensive information on theGB(B) is available in the literature [9]. The AG& for H30+ and a number ofBH+ have also been determined previously, predominantly with the PHPMSmethod [2,6,10,11] and a recent compilation of such data is available [12].The data are generally given as AG,_,,,(M+), AH,__,,,(M+), andASz_,,,(M+) for the stepwise solvation [see Eq. (S)]. The required AG,$ canbe then obtained from a summation of the (n - 1,n) values as shown in Eq.(9).Under the present experimental and the APCI conditions, water is inlarge excess over B and the H30+(HzO),, hydrates are in equilibriumamongst themselves. This is also true for the BH+(H20), hydrates eventhough the proton transfer reaction may not be in equilibrium. The relativeH30+(H20),, concentrations can be evaluated from the AG~_i,,(H,O)data. Similarly, the BH+(H,O)b concentrations are obtained from theAGf_, (BH+), when these are available. Finally, the relative H,O+(H,O),to BI-I+(H20)b concentrations at proton transfer equilibrium can be ob-tained by fixing the relative concentrations of one given pair, most conveni-H,O+ + B = H,O + BH+ (10)AG,o,= GB(H,O) - GB(B)ently the bare ion pair: H30+ and BH+ on the basis of the proton transferequilibrium, Eq. 10, i.e. AC:,-, where - AG& = RT ln K,,. A computerprogram written in Microsoft Quick BASIC and an IBMXT computer wereused to obtain the H30+(HzO)h and BH+(H,O)b populations at hydration


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144TABLE 2Hydration thermochemical data for compounds BB GB(B) - AG,O_r.,b

HOH 159MeOH 174EtOH 180MeCN 181Furan 185Me&O 189Thiophene 190Et,0 192Pyrrole 200Pyridine 213

5.4 3.8 2.74.9 3.5.6 3.3 (Z) *4.6 3.3 2.3

w 12 2,3 394 4,5 59624.2 12.8 9.218.4 12.1 6.916.1 10.7 6.316.1 9.8 8.04.4 c

12.6 6.5 6.03.9 ,c11.86.9 7.2

(6.1) *3.1=3.6

(5.6) *(1.0) *2.3

4.2 3.1 (2.1) *(3.9) * (2.9) * (1.9) *(0) * (-1.0) * (-1.5) *(1.3) * (0.7) * (0.0) *

* From gas-phase basic&y compilation of Lias et al. [9] in kcal mol- at 300 K.b Free energy for hydration of BH+ [see Eq. (8)] in kcal mol-. Standard state 1 atm. 300 K.From compilation of Keesee and Castleman [12]. AH,_,,, and AS~_rPn, allowing evaluationof AG;-, ,, at other temperatures, available in ref. 12. De&n&rations from present work. Also determined were: furan, AH& = -10.4 kcalmol-, AS& = - 20 cal deg- mol-; thiophene, AH& = - 10.0 kcal mol-, A$, = -20.6cal deg-; pyrrole, AH& = 14.0, AH& =cal deg- mol-.

- 10.4 kcal mol-, AS& = - 23.3, AS& = - 22.6* Rough estimate on the basis of available lower AGZ-,,, and fall of AG,$, with increasingn observed for other similar compounds in Table 2.

and proton transfer equilibrium for given PHzO, PB , and temperature Tconditions.The thermodynamic data used for some representative compounds aregiven in Table 2.No AH:_,,, and AS:_,, hydration data were available for the BH+ ionsof B = pyrrole, furan, and thiophene, see structures I-III, which are ofspecial interest since they are typical low-sensitivity group L bases (seeTable 1) whose BH+ hydrate very inefficiently.


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2.0 2.2 2.4 2.6(loow

Fig. 4. vant Hoff plots for n - 1,n equil ibriaof hydrationBH+(H20),_,+HzO=BH+(H,0),for B = pyrrole, furan, and thiophene. Resulting, AGi_l,,, AH&,, and ASj _,, given inTable 2.

The hydration equilibriafor these BH+ were studied in the present workin separate experiments. Equilibrium constants K,,_l,n were obtained atdifferent temperatures. The resulting vant Hoff plots are shown in Fig. 4and AH,_,,, and AS&+ obtained from these plots are given in Table 2.As expected, the hydration energies of these ions are much lower. Forexample, - AG& at 300 K for protonated water and ethanol is 24 and 16kcal mol-, respectively, while that for protonated furan is only 4 kcalmol- (see Table 2). The low hydration energies for the BH+ of furan,pyrrole, and thiophene are expected, since these compounds protonate noton the lone pair of the basic heteroatom (0, N, or S , respectively) but onthe ring [13,14].The occurrence of ring protonation predominantly in the (Yposition (seestructure IV for furan) was demonstrated by Schwarz and coworkers [13] onthe basis of experimental (ICR) studies of the protonation of these com-pounds and theoretical calculations (MMDO and MNDO) [13]. Thus, thesecompounds are not n donor but s donor, and thus carbon bases. As aconsequence of the low electronegativityof carbon and the charge delocali-


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zation (see resonance structures IV), there is only little positive charge on thehydrogen(s) and this leads to a weak hydrogen bonding in BH+ - - - OH,.Other examples of such weak hydration due to carbon protonation ofsubstituted benzenes have been pointed out in earlier work from thislaboratory [15].It is interesting to note that pyrrole - H+ hydrates relatively much morestrongly than furan and thiophene (see Fig. 4 and Table 2). The reasons forthis are not entirely clear. Theoretical calculations have shown that inprotonated benzene, C,H; , the water interacts (hydrogen bonds) with oneof the two equivalent hydrogens on the ipso carbon and that these hydro-gens carry the largest positive net atomic charge [16]. It is possible that, forthe protonated pyrrole, a relatively strong hydrogen bond to water is formedthrough the nitrogen atom hydrogen which might be expected to carry ahigher net positive charge due to the greater electronegativity of N relativeto C. It is also possible that, for pyrrole - H+, hydration promotes a protonmigration to the nitrogen. Evidence for such H+ migrations was given inearlier work for substituted anilines [15].Given in Table 1 are also equilibrium ratios R,. These are defined byR = ~[BH+(H,O),l [Wleq C[H,O+(H,O),] [Bl (11)and were measured experimentally with the PHPMS. The sums over thehydrates were obtained by summing over the corresponding ion intensitiesat equilibrium. It can be shown [l] that R, is a constant at constanttemperature and water pressure. The R, in Table 1 are for 300 K and 1 torrH,O. The R, express the extent of proton transfer to B at equilibrium forvariable B pressures. They are proportional to the APCI sensitivities for thegroup T and L compounds. The R, can lsoe calculated via Eqs. (5)-(10)when the thermochemical data are available. A comparison between R,obtained by APCI, PHPMS, and calculation is available in the previouswork [l(b)].

The R, for the poorly hydrating compounds of the low-sensitivity groupL are seen to be particularly low (see Table 1 and Fig. 5). The dependence ofR, on GB(B) shown in Fig. 5 follows the same trends as the resultsobtained with the APCI apparatus [l] which were observed at higher waterpressure [ p(H,O) = 5 torr].Since the hydration exothermicity of H,O+(H,O), exceeds the ex-othermicity of BH+(H20), at all values of n (see Table 2 for AGf_,,,values), the same results obtain also for AH:_,,, values [12]; it is obviousthat the proton transfer equilibria (5) will shift to be more in favor of theBH+ hydrates when low H,O+(H,O)h are present as reactants. This can be


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tBoOH* l eZCo. Et20

*M&Nl EtOH

GWWFig. 5. plots of experimentallydetermined R,(B) [see Eq. (ll)] versus the GB(B). Data fromTable 1 for 32C and 1.1 torr HzO.

achieved either by working at low water pressures or at high temperatureswhere only the lower H30+ hydrates are stable. The effect of temperature isshown in Fig. 3 for B = furan where a dramatic shift in favor of BH+(H,O),relative to H30+(HzO)h is observed for a temperature change from 32 [Fig.3(a)] to 198 C [Fig. 3(b)]. Dramatic increases of sensitivities for the group Lcompounds could be achieved in the APCI method by performing theanalysis of atmospheric air at elevated temperatures [l(b)].Kinetics of the proton transfer to nitrogen, oxygen, and carbon bases

As mentioned earlier, the proton transfer kinetic measurements involved ahydration equilibrium assembly of H30+(Hz0),. Plots of ln[H30+(H20),,]vs. time for any h gave the same slope, -v,, where v, is the averagepseudo-first-order rate constant v, = k,[B]. Rate constants k, obtained withthis procedure are given in Table 1.The two major H,O+(H,O), at 1 torr H,O are h = 4 and 5, each at- 45%. For hydrate distributions under different conditions, see Table 3.Bohme et al. [3,4] have measured rate constants for proton transfer from


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148TABLE 3Rate constants, k,, for proton transfer to pyrrole under different conditions *4110 b 0.1 0.5 1.1k, = 2.7 1.2 0.3n %H30+ (H,O), * AGO n,n0 0 0 01 0 0 02 0.7 0.1 - 0.01 - 14.03 47.5 14.9 7.9 -6.14 47.5 61.1 50.3 -0.75 4.3 22.7 37.1 +4.16 0 1.2 4.4 +7.8* For proton transfer to pyrrole at 32 o C.b Pressure in torr. In (molecules- cm3 s-l) X lo+* Measured abundances of H30+(HzO),, at hydration equil ibrium. Free energy for proton transfer reaction with equal hydrates [see Eq. (12)] in kcal mol- at- 32C.

H30+(H,0)h to B at low H,O and high B pressures where the hydrationequilibria were slow and the proton transfer fast. This allowed rate constantdeterminations for proton transfer from H,O+(H,O), with a given h anddata for h = 0 to h = 3 were obtained. The rate constants for several oxygenbases [CH,O, HCOOH, CH,OH, C,H,OH, CH,COOH, HCOOCH,,(CH,),O, (CH,),CO] were measured. It was found that these were close tothe collision rate constants, IcAno, and decreased very little ( - 20%) as hw as increased, see Table 2 in ref. 3. Bohmes data for h = 3 are given in thepresent Table 1 wherever compounds examined in both investigations occur.It can be seen that Bohmes and the present rate constant values are close.Since, for the present measurements, H,O+ with h = 4 and h = 5 are themajor reactants, the present results extend the validity of the observedcollision rates to these higher hydrates.

Considering the general equation for proton transfer to B [see Eq. (2)], amultiplicity of rate constants k,,, exist depending on the reactantH,O+(H,O), and product BH+(H,O),. Neither Bohmes previous work[3,4] nor the present work allow for independent observation of the effect ofX. In the absence of such direct information, a meaningful examination ofthe proton transfer rates can be obtained through a consideration of theequal hydrate reactionH,O+(H,O). + B $ BH+(H,O),, + H,Or 02)


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The forward and reverse rates of this reaction can be expected to be able tomaintain the proton transfer equilibrium and therefore one can relatedirectly the kinetics to the thermodynamics viaAG,q,(PT) = -RT ln K,,, 03)

Bohme et al. [3] found that k, for proton transfer to oxygen basesremains close to the collision rate constant k, = k,, right down to nearzero exothermicity, i.e. AG,&(FT) < 0, although only a few cases wereexamined and n was less than 3. The present results for EtOH, MeCN,Me&O, Et ,O, and the nitrogen bases are in accordance with this result.AG&(FT) is given in Table 1 for the reactions. The sixth hydrate is probablythe relevant one at 1 torr H,O since H30+(Hz0), is the last hydrate ofsignificant concentration (see Table 3). In all of the above cases, - AG&(FT)is negative and the k, are near k,. Unfortunately, examples where- AG:,,(FT) is close to zero were not obtained.The carbon bases, furan, and thiophene, have large positive AG&(FT)values. The proton transfer equilibria were shifted far to the left, i.e. in favorof H30+(H,0), (see small R, in Table 1) so that the rate constants k,could not be measured for these carbon bases. A rate constant k, = 3 X 10-lmolecules-1 cm3 s-l could be obtained only for pyrrole (see Table 1). Whenk, becomes considerably less than k,, proton transfer only from the low nof H,O+(H,O), is indicated, i.e. n for which the proton transfer AG&(FT)is still exothermic.Kinetic measurements for the proton transfer to pyrrole were also madeat lower water pressures. The rate constant determination plots are shown inFig. 6. The resulting k, are given in Table 3 together with the AGO,, forall the relevant H,O+(H,O), present under these conditions. k, = k,, forpyrrole can be calculated as k,, = 1.6 X low9 molecules- cm3 s-l. Themeasured k, increases by a factor of - 10 as the water pressure is decreasedbut still remains far below k, = k,, = 1.6 x 10m9 molecules- cm3 s-l.Examining the k , and the corresponding H30+(H20)h distributions, onefinds that the data can be fitted if one assumes that proton transfer fromH30+(Hz0)3 occurs at - 0.35 of collision rate while that from h = 4 isnegligible. This result also suggest that proton transfer from h = 2 is atcollision rates. Measurements of k , at 1 torr H,O and 198C, whereH30+(H20)2 = 84% of the total hydrates, led to k, = 1.4 x 10m9 mole-cules-lcm3 s-l which is near collision rates, thus supporting the assumptionthat proton transfer for h Q 2 occurs at collision rates.Proton-transfer from the h = 3 corresponds to a AGg3(FT) = -6.1 kcalmol- (see Table 3). This is a substantial exothermicity and it might appearsurprising that the reaction is slow, even though the exothermicity is


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0 . 1 0 . 2 0 . 3 0 . 4 0 . 5 0 . 6 0 . 7Time (msec)Fig. 6. Rate constant determinations for the reaction~H30+(H20),+pyrrole=CpyrroleH+(H,0),+~(h-p+l)H,OAll runs at 32 C, p,,,, = 3 torr, and 1.5 x 10m4 torr pyrrole. Variable water pressures: (a) 1.1;(b) 0.5; (c) 0.1; (d) 0.05 tom Resulting k, are given in Table 3.

substantial. The low rate is most likely due to the fact that a carbon base isinvolved. Fameth and Brauman [17] have shown that the proton transfer eq.(14) involving delocalized negatively charged carbon bases B-AH+B-=A-+BH 04)proceeds at rates slower than collision rates. Experiments described in thenext section examine whether exothermic proton transfer from AH+ to B(pyrrole) slows down as the exothermicity is decreased.Kinetics of proton transfer, AH + + B = A + BH , to carbon bases B

It was established in the preceding section that proton transfer fromH30+(Hz0), to the carbon base B (pyrrole) falls below the collision limit atrelatively large exothermicity, while proton transfer to oxygen bases remainsnear collision rates, even for exothermicities that are close to zero. In thissection, we describe proton transfer measurements [see eq. (15)] whereAH++B=A+BH+ (1%different protonated compounds, A, are involved in proton transfer topyrrole. Use of A with differing gas-phase basicities leads to proton transferequation (15) with different exothermicities. The experimentally determinedrate constants are summarized in Table 4. The measurements were per-formed at near constant temperature (232-238 o C). The elevated tempera-


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TABLE 4Rate constants * for proton transfer: AH+ + pyrrole = A + pyrrole.H+A - AG:: k,, x 1010

(kcal mol- ) (molecules- cm3 s-l)Oxygen bases(iso-Pr) ,O 3.0 5.4MeCOPh b 3.1 4.8(n-B1&0 5.0 9.2(Et) ,O a.2 12.0PhCHO b 8.2 9.5(MeOH), 8.7 18.5Carbon basesMeOPh b 8.1 0.46m-Xylene 12.5 4.2Toluene 18.6 8.0* Measurements at temperatures between 232 and 238 C.b Earlier work [18] established that MeCOPh and PhCHO are oxygen-protonated whileMeOPh (ardsole) is ring-protonated.

tures were required in order to suppress the formation of the proton-helddimers [see reaction (16)]. Even at these temperatures, for some compounds,AH++A=AzH+ 06)A, significant concentrations of A,H+ were present. Since the concentra-tions of A were sufficiently high, AH+ and A,H+ remained in equilibriumwhile the slower proton transfer to B via Eq. (15) proceeded. This was thecase for A = diethyl ether and methanol. For diethyl ether, the protonationof pyrrole by the dimer A2H+ is endothermic. Therefore, the rate constantwas determined by assuming that only the AH+ reacts while AzH+ andAH+ remain in equilibrium. For methanol, the dimer equilibrium, eq. (16)was shifted almost completely to the AzH+ side and the rate of protonationwas evaluated by assuming that it occurred only from the dimer A,H+. Theproton transfer from A2H+ is exothermic for A = methanol (see Table 4).The compounds A for which - AG& was small led to the achievement ofproton transfer equilibrium [Eq. (15)]. In these cases (diisopropyl ether,acetophenone, and di-n-butyl ether), the rate constant was obtained from theion concentration changes on approach to the equilibrium, i.e. with inclusionof the reverse rate. An examples of such a determination is given in Fig. 7.The rate constants of Table 4 are plotted in Fig. 8 versus the free energychange AG,o of the proton transfer [reaction (15)]. The rates are seen to fallinto two groups depending on whether the AH+ are oxygen- or carbon-pro-tonated. The oxygen-protonated AH+ lead to higher rates for the same


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0 . 5 5Time (msec)Fig. 7. Rate constant determination for proton transferAH++B=A+BH+where A = acetophenone and B = pyrrole. Since this reaction reaches equilibrium, rateconstants were derived from concentration changes assuming reversible reaction. I(AH+ ) =100% at i = 0, I(AH+), is the equilibrium intensity in !%%.uns l-3, p(pyrrole) = (0.59;0.30; 0.15) mtorr. Deduced kls = (0.48; 0.43; 0.50)~10-~ molecules-1 cd s-l.

exothermicity. In both groups, the rate constants are found to increase withincreasing exothermicity. These results can be rationalized on the basis ofthe double minimum reaction coordinate model [11,19,20] shown in Fig. 9.Since the observed rates of reaction (15) were not exceedingly slow, Ediff isexpected to be negative. Therefore, the rate constants will increase as theabsolute value 1 & 1 increases. For the cases where AH+ is an oxygen acid,1Edifr 1 is expected to be larger and the rate faster for two reasons. The welldepth E, = DE(AH+ - B) will be deeper and the barrier AE * should besmaller than is the case for carbon acids AH+. The expected poorer bondingof carbon acids AH+ to water due to the relative lack of positive charge onthe hydrogen(s) was discussed in the previous section and the same reasonswill lead to weak bonding of AH+ to B. In general, the barriers AE * forproton transfer are low when both A and B are oxygen or nitrogen bases;however, larger barriers are expected when either A or B is a carbon base[17] and even larger when both A and B are carbon bases.The increase of k,, with exothermicity, i.e. increase of - AEL, for bothoxygen and carbon acid AH+ is also expected. It is generally observedwhenever ion/molecule reactions proceed at rates considerably below colli-sion rates [21]. The decrease of AE * with increasing exothermicity isincorporated into theoretical treatments such as the Marcus equation [22,23].In light of the results in Table 4 and Fig. 8, the slow-down of protontransfer from H30+(H,0)h to pyrrole, and presumably also to the othercarbon acids, with decreasing exothermicity observed in the preceding


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1 0 %1i

1 0 - 1 0LI , I 40 5 1 0 15-AG;, (kcal/mol)

D

Fig. 8. Plot of rate constant k15 for the proton transferAH++B=A+BH+versus AG,4 of reaction. B = pyrrole, o, A = oxygen bases; l A = carbon bases.

section, is seen to be part of a general trend valid when oxygen acids AH+engage in proton transfer to carbon bases.The above results show that proton transfer from H30+(H20)h to carbon,bases like pyrrole, furan, and thiophene is inefficient on two counts. First,due to the poor exothermicity of hydration of BH+(H,O),, the proton

Fig. 9. Double miuimum reaction coordinate for proton transfer reaction (15). Notation usedis that of Dodd and Brauman [23].


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154transfer equilibria are shifted in favor of H,O+(H,O)h. Second, the kinetics.of proton transfer equation (12) slow down as the exothermicity decreases.Under atmospheric (APCI) conditions only the first factor is of importance.Due to the unfavorable thermodynamics, the proton transfer equilibria areshifted almost completely towards H,O+(H,O),. Even though the rateconstant for proton transfer in the forward direction may be slow, theproton transfer equilibria are achieved in a very short time such that theproton transfer is thermodynamically controlled. It should be recalled thatthe relaxation time, 7, for achievement of the equilibrium depends on thesum of the forw ard and reverse rates.+ = k,[B] + k,[H,O] (16)Even though both k, and k, are smaller than is the case for the oxygenbases, the reverse rate is very fast due to the relatively very high waterconcentration.REFERENCES

5

6I

89101112

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(a) J. Sunner, G. Nicol and P. Kebarle, Anal. Chem., in press. (b) J. Sunner, M.Ikonomou and P. Kebarle, Anal Chem., in press.P. Kebarle, in W. Saunders, Jr. and J.M. Farrar @is.), Techniques of Chemistry.Techniques for the Study of Ion Molecule Reactions, Wiley-Interscience, New York, 1988.D.K. Bohme, G.I. Mackay and S.D. Tanner, J. Am. Chem. Sot., 101 (1979) 3729.A.B. Raksit and D.K. Bohme, Can. J. Chem., 61 (1983) 1683. D.K. Bohme and A.B.Raksit, J. Am. Chem. Sot., 106 (1984) 3447. G.I. Mackay, A.B. Raksit and D.K. Bohme,Can. J. Chem., 60 (1982) 2594. D.K. Bohme, A.B. Raksit and G.I. Mackay, J. Am. Chem.Sot., 104 (1982) 1100.D. Smith, N.G. Adams and M.J. Henchman, J. Chem. Phys., 72 (1980) 4951. P.M. Hierl,A.F. Ahrens, M. Henchman, A.A. Viggiano, J. Paulson and D.C. Clary, J. Am. Chem.Sot., 108 (1986) 3140, 3142.P. Kebarle, Ann. Rev. Phys. Chem., 28 (1977) 445.(a) F. Arnold, 0. Krankowsky and K.M. Marien, Nature (London), 267 (1977) 30. (b)F.C. Fehsenfeld, I. Dotan, D.L. Albritton, C.J. Howard and E.E. Ferguson, J. Geophys.Res., 83 (1978) 1333.T. Su and M.T. Bowers, in M.T. Bowers (Ed.), Gas Phase Ion Chemistry, Vol. 1,Academic Press, New York, 1979.S.G. Lias, J.F. Liebman and RD. Levin, J. Phys. Chem. Ref. Data, 13 (1984) 695.Y.K. Lau, S. Ikuta and P. Kebarle, J. Am. Chem. Sot., 104 (1982) 1462.M. Meat-Ner, J. Am. Chem. Sot., 106 (1984) 1265.R.G. Keesee and A.W. Castleman, Jr., J. Phys. Chem. Ref. Data, 15 (1986) 1011.R. Honriet, H. Schwarz, W. Zummack, J.G. Andrade and P.v.R. Schleyer, Nouv. J. Chim.,5 (1981) 505.J. Catalan and M. Yanez, J. Am. Chem. Sot., 106 (1984) 421.Y.K. Lau, K. Nishizawa, A. Tse, R.S. Brown and P. Kebarle, J. Am. Chem. Sot., 103(1981) 6291.


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16 S. Ikuta an d P. Kebarle, unpu blished work.17 W.E. F am eth an d J.I. Brau man , J . Am. Chem. Sot., 98 (1976) 7891.18 Y.K. L au a nd P. Kebar le, J . Am. Ch em. Sot., 98 (1976) 7452.19 W.N. Olmstea d a nd J .I. Brau ma n, J. Am. Chem . Sot., 99 (1977) 4219.20 K. Hira oka a nd P. Kebarle, J . Am. Chem. Sot., 98 (1976) 6119.21 M. Meot-Ner (Maut ner ) an d F.H. Field, J. Am. Chem . Sot., 100 (1978) 1356. D.K. Sen

Sha rm a and P. Kebarle, J . Am. Chem. S ot., 104 (1982) 19. E.P. Grimsru d, S. Chowdhur yand P. Kebarle, J. Chem. Phys., 83 (1985) 1059.

22 D.E . Magnoli an d J R. Mur doch, J . Am. Chem . Sot., 103 (1981) 7465.23 J .A. Dodd an d J.I. Bra um an , J . Ph ys. Chem., 90 (1986) 3559.

nichol, sunner 1988

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