J . Phys. Chem. 1987, 91, 2551-2556
2551
Bonding of Charge Delocalized Anions to Protic and Dipolar Aprotic Solvent Molecules S. Chowdhurry, E. P. Grimsrud, and P. Kebarle* Department of Chemistry, University of Alberta, Edmonton, Alberta, Canada T6G 2G2 (Received: June 17, 1986)
The equilibria (2) M- + S1 = M-81, where 27 different M which are substituted nitrobenzenes and quinones and SI are the solvent molecules CH30H, CH3CN, (CH3)2S0,(CH3)2NCOH,and tetrahydrofuran, were measured in a pulsed electron high-pressuremass spectrometer. The bond free energies and enthalpies,-AGO2 and - M 2 , that are deduced from the equilibria are found to decrease with increasing electron affinity of M. Since for the compounds M high EA(M) corresponds to Min which the charge is highly delocalized, bonding decreases with increasing charge delocalization in M-, an expected effect. The decrease of bond energy is more rapid for the protic S1 (CH30H) than for the dipolar aprotic S1 (CH3CN). This is attributed to the different location of the dipoles on the protic and aprotic molecules. The changes for the “one molecule solvation” of M- are compared with the solvation changes for M- in the corresponding liquid solvents SI.
Introduction The study of ion-molecule clusters can provide information on ion solvation, ion-ligand complexes, and ion-molecule reaction complexes, i.e. systems of importance in the chemistry of ions in the condensed phase. It is this aspect of ion cluster studies that has held the greatest fascination for us and which was emphasized from the very beginning in studies reported from this The present work represents a continuation of this effort. Negative ions represent one-half of the ion population in condensed phases and are obviously important. Furthermore, in organic reactions carried by ions, negative ions turn out to be much more important than positive ions, a fact well known to synthetic organic chemists. A particularly advantageous combination for synthetic work in the condensed phase involves negative ions and dipolar aprotic solvents like acetonitrile, dimethylformamide, dimethyl sulfoxide, etc. The advantage of this combination can be illustrated by examining reaction 1. The activation energy E* is smaller CI- t CH3Br
-
r I-CI-C-Br
1:
ClCH3
+
Br-
(1)
in dipolar aprotic solvents and the rates of (1) many orders of magnitudes faster than is the case for protic solvent^.^ Determinations of the reaction coordinate of (1) in the gas phase5q6show that the energy of the transition state, shown in brackets in eq 1, is slightly lower than that of the reactants and this means that E* in solution is entirely due to the less favorable solvation of the transition state relative to the solvation of the reactants. Since the solvation terms of the ions are dominant, E* in solution is composed largely by the difference of the solvation energy of the larger transition-state negative ion relative to the smaller ionic reactant (Cl-). The smaller Et observed in dipolar aprotic solvents thus show that the solvation in these solvents decreases less with an increase of negative ion size than is the case for protic ~olvents.~ Measured solvation energies of the ions X- = C1-, Br-, I- in liquid protic (HOH, methanol = MeOH, acetic acid) and dipolar aprotic solvents (CH3CN = MeCN, ( C H J 2 S 0 = Me2S0, dimethylformamide = DMF) are found to decrease faster from C1(1) Hogg, A. M.; Kebarle, P. J . Chem. Phys. 1965,43,449. Hogg, A. M.; Haynes, R. N.; Kebarle, P.J . Am. Chem. SOC.1966, 88, 28. (2) Kebarle, P. “Ions and Ion Pairs in the Gas Phase” in Ions and Ion Pairs in Orgunic Reacrions, Szware, M.,Ed.; Wiley-Interscience: New York, 1972. (3) Kebarle, P. Annu. Rev. Phys. Chem. 1977, 28, 445. (4) Bathgate, R.H.; Moelwyn-Hughes, E. A. J. Chem. SOC.1959, 2642. Parker, A. J. Chem. Rev. 1964,69, 1. (5) Farneth, W. E.; Brauman, J. I. J. Am. Chem. SOC.1976, 98, 5546. Olmstead, W. N.; Brauman, J. I. Ibid. 1977, 99, 4219. (6) Caldwell, G.; Magnera, T. F.;Kebarle, P. J . Am. Chem. Soc. 1984, 106, 959.
(7) Magnera, T. F.; Caldwell, G.; Sunner, J.; Ikuta, S.;Kebarle, P. J . Am. Chem. Soc. 1984,106,6140. Kebarle, P.; Caldwell, G.; Magnera, T.; Sunner, J. Pure Appl. Chem. 1985.57, 339.
0022-3654/87/2091-2551$01.50/0
to I-, i.e. with increasing ion radius, for the protic solvents relative to the dipolar arotic solvents (see Figure 8, of ref 7), in agreement with the expectations based on the reaction rates observations. Gas-phase ion-cluster bond energies obtained from measurements of the ion equilibria between X- = C1-, Br-, I- and protic SI = HOH and aprotic solvent molecules S1 = MeCN, Me2C0, Me2S0 showed7-*that the bond energies between X- and one to four solvent molecules, X-61 or X-.(Sl),, decrease faster with increasing ion size, i.e. from C1- to I- for protic relative to the aprotic SI. The different behavior for the two types of S1 toward negative ions with increasing radius was attributed7 to the different location of the dipole in the protic and dipolar aprotic S1. In aprotic SI the dipole is located on one end of the molecules, viz. H3C-+CEN-, such that the attractive approach of a positive ion is favored, while a negative ion cannot come close to the dipole due to steric interference of the alkyl group. These relationships were also confirmed by theoretical calculation^.^^^ More commonly the reactant and product ions are not spherical monoatomic ions like C1- and Br- in reaction 1 but organic anions and particularly organic anions with .rr conjugation leading to extensive delocalization of the negative charge over a large part of the molecule. The present work, which presents measurements of binding energies between some 20 radical anions, two phenoxide anions, and protic and dipolar aprotic molecules, thus extends the analysis to delocalized ions and protic and dipolar aprotic solvent molecules.
Experimental Section The present and earlier measurements were performed with the same i n ~ t r u m e n t . ~A, ~mixture of gases with suitable composition flows slowly through the thermostated ion source-reaction chamber. A short pulse (10 to 50 ps) of 2OOO-V electrons produces positive ions and secondary electrons. At the ---Torr gas preshre used, carrier gas CH4 -4 Torr, thermalization of the secondary electrons is rapid. Near thermal electrons are captured by compound M which has a positive electron affinity and is present a t mTorr partial pressures in the ion source. Collisional deactivation of the resulting (M-)* by the major gas leads to thermal Ma-. These engage in adduct-forming reactions2 with protic or dipolar aprotic solvent molecules S1 present in the ion source at mTorr partial pressure. The number densities of the ions are low, such that the dominant loss of charge is due to diffusion to the wall. Ambipolar positivenegative ion diffusion to the wall establishes -100 ps after the electron pulse. Within a similar time the fundamental diffusion mode is established also. The ion intensities in the ion source are determined by measuring the wall current, i.e. bleeding the gas mixture through a narrow slit into an (8) Davidson, W. R.; Kebarle, P. J. Am. Chem. Soc. 1976, 98, 6125. Yamdagni, R.; Kebarle, P. J. Am. Chem. Soc. 1972, 94, 2940. (9) Yamabe, S.; Hirao, K. Chem. Phys. Lett. 1981, 84, 598. Hirao, K.; Yamabe, S.;Sano, M. J . Phys. Chem. 1982, 86, 2626.
0 1987 American Chemical Society
2552 The Journal of Physical Chemistry, Vol. 91, No. 10, 1987
Chowdhurry et al. 6.0
o-N02NB-.NCCH3
-
----------- - _ _ _ _ _ _ _ _ ................................... - N0, NB-.C H30H 4.c
I
L
x" P,
0
Figure 1. Ion intensities observed after electron pulse. Ion source pressures 2.8 Torr of CH,; 0.02 mTorr of 2-nitrobenzene;93 mTorr of MeOH, 36 mTorr of MeCN at 70 'C. Adduct equilibria M- SI = M-41 are achieved rapidly; the constant ratio of [M--SI]/[M-] appears as constant vertical distance between respective ion intensities.
J
+
-
-
2.6
" > -
d
2s
3.0
2.8
1000/T(" K-' )
CH3CN = SO
0
Figure 3. van't Hoff plots of some of equilibria (2) M- + SI = M-81. The numbers beside the plots correspond to the numbered reactions given in Table 11.
5 .o 0
-I
1
n,
"
C H ~ O H= sc
5 -
2.0
e
"
4.0
3.0
5.0
6.0
I
1'O.O
+
Figure 2. Plot showing invariance of equilibrium constant K2for MSI = M-.SI with increase of total ion source pressure at constant partial
8.0
pressure ratios. evacuated region where the ions are mass analyzed, detected, and time analyzed with a multiscaler. The conditions are selected in such a manner that the ion-molecule reactions of interest have rates that are much faster than the first-order ion loss to the wall.
Results and Discussion Measurement of the Bond Energies between Charge Delocalized Anions M- and Solvent Molecules SI. The bond energy of M-.SI was determined by measuring equilibria 2. The time dependence M-
+ SI = M--Sl
(2)
I
-0
h
E
\
3 0 25
02 0 I
I
20
40
60
of the ion intensities in a typical experiment used for this purpose is shown in Figure 1. Ions like M-and M-431 after having reached equilibrium and thus achieved an intensity ratio (M-.SI)/M- that is constant with time maintain, in the logarithmic plot used, a constant vertical distance. Inspection of the figure shows that equilibria 2 are achieved very early in the measurement, which means that the forward and reverse rates of reaction 2 are fast under the conditions of the experiment. The equilibrium constants K2 are evaluated from the observed ion intensities and lead to AGO, values as shown in eq 3. The results in the figure also demonstrate
EA (M) kcal/mol Figure 4. Plot of AGO2 at 70 OC and No2 for reaction M- + SI = M;Sl vs. electron affinity of M. The three plots are for SI = dimethyl sulfoxide (m), acetonitrile (A),and methanol ( 0 ) . The actual M- are identified by the number. All M- used in the plot were substituted nitrobenzenes XNB; substituent X is 2,3-diCH3 (l), 3-OCH3 (2); 4-OCH, (3); 2-CH3
(3)
nitrobenzenes, XNB-, some substituted quinones and naphthoquinones, azulene, and maleic anhydride, and with the five solvent molecules, methanol, acetonitrile, dimethyl sulfoxide, dimethylformamide, and tetrahydrofuran, are given in Table I. The temperature dependence of the equilibrium constants K2 for some of the equilibria were determined also and AHo2 and AS0,values obtained from the van't Hoff plots are shown in Figure 3. These values are summarized in Table 11. A plot of the binding energies in M--SI vs. the electron affinity of M is given in Figure 4. The electron affinities used were
AGO, = -RT In K2
AGO = AHo - TAS"
that for M = 3-dinitrobenzene, the M-.MeCN adduct has much greater stability than M-aMeOH. In all equilibria measurements it was established that K2 remains invariant as the partial pressure of M and partial pressure of S1 is changed, by a factor of 2 or 3, see Figure 2. The AGO2 values obtained with -30 different M- ions, mostly substituted
(4); 3-CH3 ( 5 ) ; 4-CH3 (6); H (7); 2-F (8); 3-F (9); 4-F (10); 2-C1 (1 1); 3-C1 (12); 4 4 1 (13); 3-CF3 (14); 2-CN (15); 3-CN (16); 4-CN (17); 2-N02 (18); 3-NO2 (19); 4-N02 (20). The AHo, values are approximate and were obtained by assuming that ASo,= -25 cal/(mol deg), which is an average value for So2, see Table 11.
The Journal of Physical Chemistry, Vol. 91, No. 10, 1987
Bonding of Charge Delocalized Anions TABLE I: Thermochemical Data for Reaction 2, M-
2553
+ SI = M-61, at 70 OC -AG,O(SI), kcal/mol
EA: kcal/mol 19.4 23.4 20.3 20.7 22.1 21.2 22.8 24.2 27.7 25.0 25.7 28.8 28.2 31.6 27.8 35.8 34.8 38.7 36.9 36.9 44.3 40.2 42.4 49.2 60.9 -80.4 -68.3
Mb 2,3-(CH,)2NB 3-OCH3NB 4-OCH3NB 2-CH3NB 3-CHSNB 4-CHjNB NB 2-FNB 3-FNB 4-FNB 2-CINB 3-CINB 4-CINB 3-CF3NB MaAn 2-CNNB 3-CNNB 4-CNNB 2-NO2NB 3-NO2NB 4-NO2NB BQ NPQ 2,3-C1,NpQ F4BQ % N o 2 phenoxided 2 - N 0 2 phenoxide
-
CHBOH ( p = 1.4 D) 6.3 6.1 6.5 6.5 6.3 6.4 6.3 5.8 5.6 6.0 5.4 5.4 5.6 4.9 4.7 4.5 4.7 4.1 3.3 3.8 2.3 5.5 4.6 2-N02NB > 4-N02NB. The same order is observed also for the cyano-substituted nitrobenzenes. On the other hand, for the aprotic acetonitrile and dimethyl sulfoxide the order is 2-N02NB > 3-N02NB > 4-N02NB. We assume that the adduct molecule S1 bonds to the NO2 group with geometries approximately as shown in Scheme 11. As pointed out in earlier work,I6 the hydrogen-bonding adduct (MeOH) is much more affected by losses of electron density from the negative center. The withdrawing effect of the second NO2 group (or the C N group) is strongest for 4-N02NB (or 4-CNNB) and weakest for 3-N02NB (or 3-CNNB). Therefore one observes the strongest MeOH adduct to be to the 3-N02NB anion (or 3-CNNB). The dipolar aprotic molecules having their dipole located some distance away from the negative center are less sensitive to the detailed negative charge distribution; they respond rather to more global charges, so that for them the allignment of the two negative charge carrying NO2 groups (NO2, CN), which occurs in the 2-substituted XNB-, is the most favorable and leads to the strongest adduct. Presumably, in this adduct the aprotic S1 interacts with the negative charge of both substituents. The adduct binding energies for the nitrobenzene anions substituted by the a-donating and u-withdrawing substituents CH30 and F decrease in the order 4-XNB > 2-XNB > 3-XNB for both the protic MeOH as well as the aprotic MeCN, Me2S0, and DMF. This indicates that the increasing electron density on the NO2 oxygens by the a donation from C H 3 0 and F which is strongest in 4-XNB- has a favorable effect for both protic and aprotic adducts. Significantly however the difference between 4- and 3-substituted NB- is larger for the protic C H 3 0 H than for the aprotic S1 (see Table I). This difference is in line with the expected greater sensitivity of the protic adduct to the electron density on the negative center to which it bonds. The bond free energies AGO2 for M- = 2-nitrophenoxide and 4-nitrophenoxide anion and S1 = CH,OH, CH3CN, and (CH3)2SO are also given in Table I. The correlation observed between the electron affinities of M- (radical anions) and the bond energy M--SI obviously does not include the phenoxide anions. Thus, while the phenoxides have the highest electron affinities and are thus expected to bond S1 least strongly, the bond energies N02PhO--S1 are relatively high (see Table I). This effect is expected. The phenoxide anions are even electron systems where the extra electron occupies a bonding orbital which is less diffuse than the a* orbital of the radical anions. Also, bonding in the (15) Birch, A. J.; Hinde, A. L.; Radom, L. J . Am. Chem. SOC.1980,102, 3370. (16) McMahon, T. B.; Kebarle, P. J . Am. Chem. SOC.1977, 99,2222.
2556
J . Phys. Chem. 1987, 91, 2556-2562 of the phenoxides XPhO- and the electron affinities of XPhO. This relationship has the same origin as the one observed for M-, i.e. the electron affinity of XPhO increases with increasing charge delocalization in XPhO- and increasing charge delocalization decreases the bonding in XPhO-.SI and the solvation energies of XPhO-.
phenoxide--Sl occurs to the phenoxy oxygen and, other things being equal, one expects a larger negative charge density on this single charged atom than for the two oxygens of the NO2 group of the nitrobenzene (radical) anions. Measurements of the other bond energies for the substituted phenoxides XPhO-61 would have been of considerable interest. The nitrophenoxide anions were produced in the ion source by a reaction of NOz- or 02-with 2-ClNB (or 4-ClNB) described in earlier work." Unfortunately, the preparation of the other phenoxide anions, so far, is possible in our apparatus only by deprotonation of the corresponding phenols,16 and the presence of the strongly hydrogen-bonding XPhOH leads to the corresponding proton held dimers XPhO-.HOPhX as the completely dominant species, precluding thus the observation of the XPhO-41. While the phenoxide ions do not fall on the lines observed for the radical anions M- in Figure 4, it is quite certain that the bond energies for XPhO-81 as a group by themselves will fall on an approximate straight line when plotted vs. EA(XPh0) or vs. the gas-phase acidity of XPhOH. An approximately linear inverse relationship is known to e x i ~ t ~ ~between . ' ~ . ' ~ the solvation energies
Acknowledgment. This work was supported by the Canadian Natural Sciences Research Council. Registry No. DMF, 68-12-2; THF, 109-99-9; 2,3-Me2NB", 3596333-8; 3-MeONB-', 42206-54-2; 4-MeONB-', 34473-10-4; 2-MeNB-', 34505-30-1; 3-MeNB-', 34505-29-8; 4-MeNB-', 34509-96-1; NB-', 12169-65-2; 2-FNB-', 34467-51-1; 3-FNB-', 34470-17-2; 4-FNB-', 34467-53-3; 2-CINB-', 34470-27-4; 3-C1NB-', 34467-54-4; 4-CINB-', 34473-09-1; 3-CF3NB-', 34526-71-1; MaAn-', 51978-3 1-5; 2-CNNB-', 12402-45-8; 3-CNNB-', 12402-46-9; 4-CNNB-', 12402-47-0; 2NOZNB-', 34505-38-9; 3-NOzNB-', 34509-56-3; 4-NOzNB-', 3450533-4; BQ-', 3225-29-4; NpQ-', 20261-01-2; 2,3-CI,NpQ-', 22062-59-5; F,BQ-', 42439-31-6; MeOH, 67-56-1; MeCN, 75-05-8; Me2S03, 6768-5; 4 - N 0 2 phenoxide, 14609-74-6; 2 - N 0 2 phenoxide, 16554-53-3. ~
~
~~
~
~
(19) The relationship is actually between the gas-phase acidities and the solution acidities of the substituted phenols. From this relationship since the bond energies D(XPhO-H) are expected to change very little one can deduce a linear inverse relationship between the electron affinities EA(XPh0) and the solvation energies of XPhO' ions.
(17) See reactions 6, 10, and 11 in Grimsrud, E. P.; Chowdhury, S.; Kebarle, P. Int. J. Mass Spectrom. Ion Processes 1986, 68, 57. (18) Fujio, M.; McIver, R.T.; Taft, R. W. J. Am. Chem. Soc. 1981, 103, 4017.
Thermal Energy Charge-Transfer Reactlons of Ar+ and Ar,+ R. J. Shul, B. L. Upschulte, R. Passarella, R. G. Keesee, and A. W. Castleman, Jr.* Department of Chemistry, The Pennsylvania State University, University Park, Pennsylvania I6802 (Received: June 17, 1986)
The rate coefficients for a number of thermal energy charge-transfer reactions are obtained with a recently completed selected ion flow tube (SIFT). The ion-molecule reactions studied involve Ar+ and Ar2+with a variety of neutral molecules including 02,CS2, C02, SO2,H 2 0 , H2S, NH3, NO, SF6, CH,, N20, NO2, and CO. The relative magnitudes of the observed rate coefficients are not in accord with an energy resonance model which requires favorable Franck-Condon factors. Furthermore, we find that the dimer ion reaction rate constant is not always greater than that of the monomer with a specific neutral although there is greater phase space in the case of the dimer where the dissociativechannel leads to a three-body final state. However, the proximity of the recombination energy of Ar+ and Ar2+to a band in the photoelectron spectra of the neutral appear to explain the relative rates of the monomer and the dimer reactions with a specific neutral.
coefficients for atomic (Ar+) and molecular (Arz+) ions with a variety of polyatomic neutrals. Two proposed mechanisms, the long-range electron jump mechanism and the collision complex formation mechanism, have dominated the literature over the past decade. The long-range electron jump mechanism is considered to occur over a relatively large distance by a nonorbiting collision8 on more than one potential surface. Early evidence for this mechanism at thermal energies included the observation of rate c ~ e f f i c i e n t swhich ~ ~ ' ~ were greater than the calculated collision rates from either the Langevin" or average dipole orientation (AD0)I2 theories, both of which imply orbiting collisions. Such long-range, nonadiabatic reactions were considered to occur by vertical transitions where Franck-Condon factors are significant.I0 Laudenslager et al.I3 have examined a number of charge-transfer reactions and have concluded that, in general, favorable Franck-Condon factors are necessary for a fast thermal energy charge-transfer reaction (koW 3 O.lOkcald).Exceptions have been observed, notably reactions
Introduction Thermal energy charge-transfer reactions have been of longstanding interest in the field of gas-phase ion-molecule chemistry. Work has been prompted by both practical and fundamental considerations. For example, charge-transfer reactions are known to be one of the dominating processes in the ionosphere' and in interstellar chemistry.2 Also, thermal energy charge-transfer reactions have been considered as possible sources of chemical laser^.^^^ The mechanism involved in a thermal energy chargetransfer reaction has yet to be definitively established, and much debate has appeared in the literature over this issue. The distribution of excess energy in the reaction into internal modes, translational energy, or dissociative channels has also been disc u s ~ e d . ~The ~ ~ -subject ~ of the present paper is to compare rate (1 ) Man, R.In Interactions Between Ions and Molecules, Ausloos, P.. Ed.; Plenum: NewYork, 1975. (2) Utterbach, N. G. Phys. Rev. Lett. 1965, 15, 608. (3) Bowers, M. T. In High Power Gas Lasers, Bekefi, G.,Ed.; Wiley: New York, 1976. (4) Wilson, L. C.; Suchard, S . N.; Steinfeld, J. I. In Electronic Transition Lasers; MIT: Cambridge, MA, 1977. (5) Hunton, D. E.; Hofmann, M.; Lindeman, T. G.; Castleman, Jr., A. W. J. Chem. Phys. 1985, 82, 134. (6) Levanthal, J. J.; Earl, J. D.; Harris, H. H. Phys. Reu. Left. 1975, 35, 719. (7) Fehsenfeld, F. C.; Ferguson, E. E. J . Geophys. Res. 1971, 76, 8453.
0022-3654/87/2091-2556$01.50/0 , , I
(8) Bowers, M. T.; Su,T. Ado. Electron Phys. 1973, 34, 223. (9) Gauglhofer, J.; Kevan, L. Chem. Phys. Lett. 1972, 16, 492. (IO) Bowers, M. T.; Elleman, D. D. Chem. Phys. Lett. 1972, 16, 486. (1 1) Langevin, M. P. Ann. Chim. Phys. 1905, 5, 245. (12) Su,T.; Bowers, M. T. J . Chem. Phys. 1973, 58, 3-27. (13) Laudenslager, J. B.; Huntress, Jr., W. T.; Bowers, M. T. J . Chem. Phys. 1971.61, 4600.
0 1987 American Chemical Societv -
