J. Phys. Chem. 1996, 100, 8875-8881
8875
The Dimethylhydroxysulfuranyl Radical Ga´ bor Mere´ nyi* and Johan Lind Nuclear Chemistry Laboratory, Department of Chemistry, Royal Institute of Technology, Stockholm, Sweden
Lars Engman Department of Organic Chemistry, Institute of Chemistry, UniVersity of Uppsala, Uppsala, Sweden ReceiVed: December 6, 1995; In Final Form: February 15, 1996X
By use of pulse radiolysis the one-electron reduction potentials, E0(DMS•+/DMS) and E0((DMS)2•+/2DMS) (dimethyl sulfide, DMS) were determined to be 1.66 ( 0.03 and 1.40 ( 0.02 V vs NHE, respectively. DMS•+ was found to be in equilibrium with DMSOH• with a pKa ) 10.2. The conditional equilibrium constant for the reaction DMSOH• + DMS h (DMS)2•+ + OH- was found strongly dependent on both ionic strength and DMS concentration. In the thermodynamic limit this equilibrium constant is ≈1.3. The dimerization reaction DMS•+ + DMS h (DMS)2•+ was shown to have its equilibrium constant between 104 and 5 × 104 M-1. DMSOH• reacts with O2 with a rate constant of 2 × 108 M-1 s-1, independent of pH (11-14). From this and other observations, we estimate the pKa for deprotonation of DMSOH• to exceed 17. From spectral and kinetic data, the maximum lifetime of the radical (DMS)2OH• was predicted to be 10 ns. The stability of DMS-X• (X ) OH, I, Br, Cl) in aqueous solution was shown to correlate with the one-electron reduction potential of X•. Comparison of gaseous and aqueous behavior of DMS-OH• reveals that aqueous solvation strongly stabilizes the S-O bond against dissociation into DMS and OH•. The Gibbs free energy of solvation of DMSOH• was calculated to be -12 ( 3 kcal/mol, an unusually large value for a neutral species.
Introduction Dimethyl sulfide (DMS) is an important molecule, being the principal organosulfur compound to be released into the atmosphere. Its destruction there is mainly effected by OH• radicals and to a lesser extent by Cl• and NO3•. In the gas phase the reactivity of OH• toward DMS has been the subject of a number of investigations.1 In these studies the disappearance of the OH• radical was measured and thus the occurrence and properties of other intermediates could only indirectly be assessed. By monitoring the time course of OH• radicals in the absence and presence of O2 and by varying the amount of DMS, the following reaction scheme was arrived at, where M denotes the third body.
DMS + OH• h DMSOH•
(3)
DMS + OH• f •CH2SCH3 + H2O
(1)
DMS + OH• f •CH2SCH3 + H2O
(4)
DMS + OH• + M f DMSOH• + M
(2)
DMSOH• h DMS•+ + OH-
(5)
DMS•+ + H2O h DMSOH• + H+
(6)
DMSOH• + DMS h (DMS)2•+ + OH-
(7)
At all temperatures studied the main reaction mode was hydrogen abstraction by OH• to form the (methylthio)methyl radical (•CH2SCH3). However, the lower the temperature, the more important the second channel was found to become, namely, formation of the dimethylhydroxysulfuranyl radical (DMSOH•) in equilibrium with DMS + OH•. While Hynes et al.2 estimate the S-O bond strength in DMSOH• at ca 13 kcal/ mol, Gu and Turecek3 believe that DMSOH• does not exist at all except as a transition state. More recently, a high-level ab initio quantum chemical calculation4 has substantiated the existence of DMSOH•, but it also revealed that the S-O bond is somewhat elusive and rather different from the S-X bonds in DMS-halogen atom (X) adducts. The existence of the latter species was readily predicted even at a relatively simple computational level and they were found to be well described in terms of soft acid-soft base complexes. On the other hand, X
the S-O bond in DMSOH• appeared to be a relatively nonpolar two-center three-electron bond. In contrast to the elusive nature of DMSOH• in the gas phase, the occurrence of this species is well documented in water during OH•-induced oxidation of DMS. By means of optical,5-7 conductivity,5,8 and ESR9 monitoring, a whole spectrum of transient species has been detected. In the course of kinetic investigations, mostly by means of pulse radiolysis, reactions 3-11 have been identified. In particular, the properties of the two-center three-electron bond in (DMS)2•+ and similar species have been extensively covered.10-12
Abstract published in AdVance ACS Abstracts, May 1, 1996.
S0022-3654(95)03613-6 CCC: $12.00
DMS+ + H2O(B-) f •CH2SCH3 + H3O+(BH) (8, 8′) DMS•+ + DMS h (DMS)2•+
(9)
DMSOH• h DMSO•- + H+
(10)
DMSO + eaq- f DMSO•-
(11) OH•
While the qualitative picture of DMS degradation by is clear, not all parameters appear sufficiently well established to allow for a consistent description of the system over the whole concentration and pH range. In particular, some ambiguities persist regarding
© 1996 American Chemical Society
8876 J. Phys. Chem., Vol. 100, No. 21, 1996 the acid-base equilibria and reactivity of the DMSOH• species. In their pioneering work, Meissner et al.13 observed an acid-base equilibrium of some transient formed subsequent to both reactions 3 and 11. The above authors ascribed this equilibrium, characterized by a pKa ) 10.2 ( 0.3, to eq 10, the deprotonation of DMSOH•, a suggestion adopted and retained ever since. Recently,14 however, this claim was challenged on the following grounds. In ref 7 k5 was measured to be (1.3 ( 0.2) × 106 s-1. On the reasonable assumption that k-5 is probably close to the diffusioncontrolled limit, i.e., ca. 1010 M-1 s-1, and certainly cannot exceed this value, the observed pKa ) 10.2 was ascribed to eq 6. As a result, pKa(10) was deduced to be larger than 10 and probably even larger than 14. Clearly, if pKa(6) and pKa(10) were close to each other, one would observe a steeper change with pH of the transient signal around pH 10 than was reported in ref 13. In addition, the yield of (DMS)2•+ should decrease with increasing pH at a constant ratio of [OH-]/[DMS]. In fact, we shall, in what follows, show that this is not the case. In the present work we shall firmly establish the acid-base properties of DMSOH• and discuss the differences between its properties in water and the gas phase.
Experimental Section Pulse radiolysis was performed at room temperature utilizing doses of 2-15 Gy/pulse corresponding to 1.2 × 10-6-9 × 10-6 M radicals. The 4.6-MeV linear accelerator is characterized by a fwhm of 8 ns and a beam current of 1.8 A. The computerized optical detection system15 has been described elsewhere. In all experiments a monochromator bandwith of 2.5 nm was used. When recording spectra each point represents 5-10 superimposed traces. During titration as many as 30 superpositions were utilized for each point. The relative dose/ pulse was monitored by a secondary emission chamber, the latter having been calibrated against an aerated 10-2 M KSCN solution taking16 G ) 2.23 × 104 100 eV-1 M-1 cm-1 at 500 nm. The solutions were made up in Millipore-deionized water and purged for 10 min by the appropriate gas. DMS was added, pure or from stock solutions, after purging. In order to keep the O2 concentration well below 5 × 10-6 M, the added volume was always 11. We note that above pH 11 identical dynamics were observed, independently of whether DMSOH• was produced in reaction 3 or subsequent to reaction 11. In the presence of DMS, the same dynamics were seen at 360 and 465 nm, an observation attesting to equilibria 5 and 9 being maintained throughout. From the pH behavior of the slow decay rate we obtained the same pK(6) ) 10.2 as from the absorbance measurements. We interpret the fast process to be the attainment of eq 6. The pH behavior of the slow component is then given by kobs ) (k8[H+] + klimK6)/ ([H+] + K6). The pH-independent rate constant, klim, for the disappearance of DMSOH• is interpreted as follows. The main reaction of DMS•+ with OH- is addition to form DMSOH• in reation -5. A minor side reaction is proton transfer between them to yield •CH SCH in reaction 8′. From K ) 1.6 × 10-4 M-1 and k 2 3 5 5 ) 1.3 × 106 s-1, k-5 ) 8 × 109 M-1 s-1 is calculated. Hence the formal rate constant, k′8, with OH- as the base B-, would come out as ca. 3 × 108 M-1 s-1. Thus klim ) k5k′8[OH-]/ (k-5[OH-] + k′8[OH-]). More likely, however, is the following scenario. DMS•+ and OH- diffuse together with a rate constant of ca. 8 × 109 M-1 s-1. Then in the encounter complex ca. 4% of the OH- ions deprotonate a methyl group while the rest adds to sulfur. Evidently, these two cases are mathematically indistinguishable. The combination of equilibria 5 and 9 leads to eq 7. Clearly K7 ) K5K9. K7 can be measured directly by, e.g., monitoring the 465 nm absorbance due to (DMS)2•+ as a function of [OH-] at constant [DMS]. Consequently, we performed a number of such titrations, presented in Figure 3. Several features emerge from the figure. The main observation is that all experiments are described by equilibria, which are first order in [OH-]. Therefore, invoking the arguments in the Introduction, we can safely exclude eq 10; i.e., DMSOH• does not deprotonate below pH 13. As expected, the equilibrium constant K7 is dependent on ionic strength (I), varying by a factor of 2 when I goes from 0 to 0.2 M (compare the curve pertaining to 5 × 10-4 M DMS at I ) 0.2 M with that at I ) 0 M and [DMS] ) 1.1 × 10-3 M). However, even though we have not elaborated it in detail, K7 also seems to increase with increasing DMS concentration. Thus, when I is made up of added NaOH in the presence of 1.1 × 10-2 M DMS, the (DMS)2•+ signal is halved at [OH-] ) 0.1 M, which corresponds to a formal K7 ) 9. On the other hand, K7 is merely ca. 3 at I ) 0.2 M and [DMS] < 1.5 × 10-3 M. Formally, this implies a decrease in the activity coefficient of (DMS)2•+ with increasing [DMS]. On the molecular level, this could be taken to imply complexation between DMS and (DMS)2•+. However, no spectral indication (see Figure 1a), whatsoever, could be had for such a process. The value of K7 at [DMS] ) 1.1 × 10-3 M and negligible ionic strength is calculted to be 1.4. This is almost the same as 1.3, extracted from the data in ref 13 pertaining to zero ionic strength and
8878 J. Phys. Chem., Vol. 100, No. 21, 1996
Mere´nyi et al. SCHEME 1
Figure 3. 3 pH titration of the yield of (DMS)2•+ measured at 465 nm in N2O-saturated solutions at various ionic strengths (I) made up of NaClO4 and NaOH. The yields are normalized to the values observed in neutral solutions at the pertinent DMS concentrations. (×) 1.1 × 10-3 M DMS, I ) 0 M; (b) 5 × 10-4 M DMS, I ) 0.2 M; ([) 1.5 × 10-3 M DMS, I ) 0.2 M; (9) 4.5 × 10-3 M DMS, I ) 0.2 M; (O) 1.1 × 10-2 M DMS, I ) [NaOH] M; (1) 1.1 × 10-2 M DMS, I ) 1 + [NaOH] M (Lines) The lines are calculated using the expression (K7[DMS] + 0.1[OH-])/(K7[DMS] + [OH-]), where K7[DMS] (M) is set to (a) 1.6 × 10-3, (b) 5.6 × 10-3, (c) 3.2 × 10-2, (d) 0.1.
[DMS] ) 4 × 10-4 M. We thus set K7 ) 1.3 as the true thermodynamic equilibrium constant. From this value and K5 ) 1.6 × 10-4 M, we obtain K9 ) 0.8 × 104 M-1, in fair agreement with 2.5 × 104 M-1, the value calculated above from redox potentials. Does (DMS)2OH• Exist? In the literature the possible role of a species, surmised to be an adduct of DMS to DMSOH•, or, what is equivalent, an OH- adduct to (DMS)2•+, has sometimes been invoked.7,24 We therefore thought it of interest to inquire into the nature of this species. Consider spectrum c, recorded in Figure 1 at pH 14 in the presence of 10-2 M DMS. It is readily seen that the experimental points can be fitted to a composite spectrum consisting of 79 ( 2% DMSOH• and 21 ( 2% (DMS)2•+. On the reasonable assumption that (DMS)2OH•, were it to exist, would not absorb (or would absorb considerably weaker than (DMS)2•+) above ca. 500 nm, and granting the error margin to bear on the maximum amount of (DMS)2OH• present in eq 16, we deduce K16 > 10 M.
(DMS)2OH• h (DMS)2•+ + OH-
(16)
Extrapolated to zero ionic strength and low [DMS] (see above) K16 > 1 M is obtained. As K7 ) K16K17 ) 1.3 it follows that K17 < 1 M-1.
DMSOH• + DMS h (DMS)2OH•
(17)
In conditions, where the reverse rate in eq 7 was negligible, the forward rate constant of eq 7 was measured7 to be 1.5 × 108 M-1 s-1. Within the above tentative model, this rate should be k7 ) k17k16/(k16 + k-17). If k16 . k-17, then k7 ) k17, and since K17 < 1 M-1, it follows that k-17 > 108 s-1 and hence, a fortiori, k16 > 108 s-1. If, on the other hand, k16 , k-17, then k7 ) k16K17, whence both k16 and k-17 come out larger than 108 s-1. In conclusion, either way, the lifetime of (DMS)2OH• can be 10 ns at most. This is a tenuous basis for postulating such an intermediate and we rather propose reaction 7 to be a onestep bimolecular homolytic substitution, a SH2 process, albeit catalyzed by the water solvent. In ref 7 the large solvent kinetic deuterium isotope effect on k7 was given a plausible interpretation, namely that water transfers a proton to the oxygen atom
of (DMS)2OH•. We agree, in the main, with the mechanism of the authors, except that we suggest their proposed intermediate to be the transition state depicted in Scheme 1. By the logic of microscopic reversibility k-7 has also to involve a H2O molecule. The necessity for H2O can be rationalized by assuming a H2O molecule to be strongly complexed to (DMS)2•+. We note that in ref 25 a tight complex between DMS•+ and H2O was predicted and this possibility is substantiated by our finding14 that DMS•+ is hydrated much more strongly (by ca. 14 kcal/ mol) than (C6H5)2S•+. In view of these points, the suggestion of a relatively tight (DMS)2•+-H2O complex would not appear too audacious. Deprotonation of DMS•+ by H2O. When in neutral solutions DMSO (2 M) was reduced by eaq- in the presence of (4-OH-C6H4)2S or promethazine, the latter compounds were oxidized to their radicals. At low concentrations of the reductant, [Red], the rate of growth of the radical absorbance was proportional to [Red]. However, when the observed rate, kobs, approached ca. 106 s-1, kobs vs [Red] deflected from straight-line behavior. This was analyzed in terms of a limiting rate corresponding7 to k5 ) 1.3 × 106 s-1. Now in reaction 11 DMSO•- is formed primarily. As DMSO•- is extremely unlikely to expel O2- to produce DMS•-, one has to assume that reaction 5 is preceded by protonation of DMSO•- by water with a rate >1.3 × 106 s-1. Deprotonation of DMSOH• by OH- cannot exceed ca. 1010 M-1 s-1. Consequently, pKa(10) is larger than 10, a conclusion deduced in ref 14 in an alternative way. As a spin-off of these titrations by (4-OH-C6H4)2S or promethazine, the rate of DMS•+ disappearance, k8, was obtained with a value of 1.7 × 105 s-1. The rates were extracted from the intercepts of the observed rate constants vs (4-OH-C6H4)2S or promethazine concentration.26 The competition kinetic plots representing the inverse of the size of the Red-radical signal vs [Red]-1 yielded rate ratios which tallied with the directly determined rates. In agreement with ref 20, k8 was found to increase upon addition of phosphate. In another set of experiments in neutral solutions we measured the entire dynamics of the 465 nm absorbance while varying the DMS concentration from 5 × 10-5 to 3 × 10-4 M during its oxidation by OH• radicals. The observed curves were compared with simulated ones, where the values for k-9 and k8 were varied and k9 ) 3 × 109 M-1 s-1 was taken from ref 6. Fits were obtained for k-9 ) 2 × 105 s-1, k8 ) 6 × 104 s-1 or k8 ) 2 × 105 s-1, k-9 ) 5 × 104 s-1. The second set of data is in agreement with k8 measured in 2 M DMSO solutions. It is conceivable though that k8 may be more rapid in the presence of DMSO than in pure water. In support of this we have observed that even (C6H5)2S•+, which cannot undergo deprotonation, decayed more rapidly in the presence of 2 M DMSO than in its absence. All in all, determination of K9 in neutral solution is not straightforwardsonly an interval 1.5 × 104 M-1 < K9 < 6 × 104 M-1 can be obtained. Even though the true value of k8 and hence of K9 is left somewhat in limbo in the simulations, we believe k8 ) 6 × 104 s-1 to be the more reliable rate constant, it leading to a K9 in better agreement with the independnet measurements described above. At this point it may be worthwhile to digress for a brief comment on the acidity of DMS•+. In ref 27 the C-H bond strength in DMS was reported to be ca. 95 kcal/mol. On the
Dimethylhydroxysulfuranyl Radical
J. Phys. Chem., Vol. 100, No. 21, 1996 8879
Figure 4. Spectra obtained upon pulsed irradiation of N2O-saturated solutions containing 5 × 10-4 M tert-butyl sulfide: (b) pH 5, ([) pH 12. Inset: pH titration of the absorbance at 425 nm. The drawn line is calculated for pKa ) 10.4.
reasonable assumption that S0(•CH2SCH3) - S0(DMS) ≈ S0(•CH2OH) - S0(CH3OH) ) 1.4 eu28 and the free energies of solvation of DMS and •CH2SCH3 are equal we calculate E0(•CH2SCH3,H+/DMS) ) 1.64 V. Then by means of E0(DMS•+/DMS) ) 1.66 V we obtain pKa(DMS•+) ≈ 0. Thus DMS•+ would seem to be an extremely strong acid and its relatively slow deprotonation by water should be ascribed to kinetic stabilization of carbon acids. One-Electron Reduction Potential and Pseudobase Formation of the tert-Butyl Sulfide Radical Cation, (t-Bu)2S•+. In view of the complications present with DMS we undertook measurements on (t-Bu)2S as well. The one-electron oxidized radical of this compound, (t-Bu)2S•+, undergoes neither deprotonation nor dimerization (reactions 8 and 9). Figure 4 shows the spectra taken at pH 7 and 12, respectively, upon oxidation of (t-Bu)2S by the OH• radical. The inset presents a titration curve, which discloses a pKa ) 10.4. This value is very close to pKa(6) and is surely the corresponding pseudobase equilibrium constant of (t-Bu)2S•+. We have also determined the below equilibrium by bringing (t-Bu)2S to react with the anisole couple.
(t-Bu)2S
•+
•+
+ CH3O-C6H5 h (t-Bu)2S + CH3O-C6H5 (18)
The equilibrium constant K18 was found to be 1.5. From this value and E0(CH3OC6H5•+/CH3OC6H5) ) 1.62 V29 we obtain E0((t-Bu)2S•+/(t-Bu)2S) ) 1.63 V. Within experimental error this is the same as E0(DMS•+/DMS), and thus our confidence in the latter value increases. The Properties of DMSOH•. In the above sections it has been shown that the pKa for deprotonation of DMSOH•, pKa(10), is significantly higher than 10 and in ref 14 its value was surmised to exceed even 14. We shall now provide additional evidence for this conjecture. First of all, DMSO•- is expected to be a powerful one-electron reductant, similar to CO2•-. This can be deduced from known thermochemistry as follows: From E0(DMS•+/DMS) ) 1.66 V, pKa(6) ) 10.2, and 2E0(DMSO,2H+/DMS,H2O) ) 0.574 V,30 we calculate E0(19) to be -1.11 V.
DMSO + e- + H+ h DMSOH•
(19)
Were pKa(10) no higher than 10, E0(DMSO/DMSO•-) would still be as negative as -1.7 V and would become progressively more negative with increasing pKa(10).
In order to check the reducing properties of DMSOH• (DMSO•-?) we produced it at pH 14 by O•- oxidation of 5 × 10-3 M DMS in N2O-saturated solutions in the presence of methyl viologen, MV2+. The concentration of MV2+ was varied from 5 × 10-5 to 3 × 10-4 M. We could not observe any buildup of the characteristic absorbance of MV•+.31 This is telling evidence for DMSOH• to be a poor reductant, and by implication, of pKa(10) being above 14. Indeed, as will be demonstrated, DMSOH• rather acts as an oxidant. When produced of pH 14 in the presence of (4-O--C6H4)2S, DMSOH• oxidized the latter compound to its radical with a rather high rate (≈109 M-1 s-1). Formally, this reaction is somewhat similar to reaction 7, except that no sulfide dimer is formed. Even more to the point, DMSOH•, produced at pH 14 in the presence of 10-2-1 M N3-, oxidized the latter rapidly (>108 M-1 s-1) to N3•. While we have not explored the mechanism of this reaction, the overall process is in keeping with known redox properties. Thus, from E0(DMS•+/DMS) and K5 we calculate E0(20) to be 1.43 V, well above E0(N3•/N3-) ) 1.33 V.
DMSOH• + e- h DMS + OH-
(20)
As a final confirmation of the protonation state of DMSOH• we investigated its reactivity toward molecular oxygen. We produced DMSOH• in two ways: by reducing DMSO (2-4 M solutions) with eaq- in the presence of (1.3-6.5) × 10-4 M O2 or by having OH• react with DMS using N2O/O2 mixtures as purging gas. From the decay of the 360 nm absorbance we obtained the rate constant k21 to be (2.0 ( 0.5) × 108 M-1 s-1. The relatively low value of k21 is compatible with the almost quantitative formation of (DMS)2•+, observed in ref 24 at pH values well below pKa(6), in the presence of 10-3 M DMS and up to 5 × 10-4 M O2. It should be mentioned that we found (t-Bu)2SOH• to react with O2 immeasurably slowly, i.e., with a rate constant below ca. 107 M-1 s-1. This is ascribed to steric hindrance and is in keeping with the known reluctance12 of e.g. (t-Bu)2S•+ to form three-electron bonded complexes.
DMSOH• + O2 f products
(21)
k21 was independent of the pH (10-14) and of DMS concentration (0-10-2 M), as long as eq 7 was sufficiently shifted to the left. These findings suggest that, up to at least pH 14, DMSOH• is the only species present. Had DMSO•- formed, it would have been expected to react with O2 by way of a simple outersphere electron transfer, just like CO2•-, with a diffusioncontrolled rate, i.e., ca 3 × 109 M-1 s-1. The fact that the measured value of k21 at pH 14 is, within experimental accuracy (10%), equal to those at lower pH then suggests that pKa(10) is probably 16 or higher. A similar estimate is arrived at upon considering the absence of any significant buildup at pH 14 of the MV•+ absorbance, as described above. As this implies a second-order rate constant 10-3 M DMS the yield of (DMS)2•+ is equal to that of OH• within experimental error. As the free energies of hydration, ∆G0g-aq, of DMS31 and OH•32,33 are 0.4 and -2.6 kcal/mol, respectively, it follows that the corresponding ∆G0g-aq for DMSOH• is as negative as -12 ( 3 kcal/mol. Figure 5 presents the equilibria Khet and Khom as a function of the redox potentials of the X•/X- couples.34 There is a linearity observed for both Khet and Khom. The fully drawn line represents the stability of the DMS-X species irrespective of the mode of its dissociation. As expected, the crossing point of the two lines corresponds to E0 ) 1.66 V, the reduction potential of DMS•+. It then follows that the DMS-X adduct is at its most stable, when E0(X•/X-) ) E0(DMS•+/DMS). Now, other things being similar, it is reasonable to expect that the S-X bond in a DMSX• adduct in the gas phase will be the stronger the closer the ionization potentials, i.e., the smaller ∆Ip, of X- and DMS. Unfortunately, no values on gaseous DMSX• other than DMSOH• exist. Nevertheless, quantum chemical calculations in ref 35 predict a relationship between ∆Ip and the dissociation energies of three-electron bonded species. However, even if such a relationship exists in the gas phase, it is far from obvious why a linear relationship should
Although the linearity in Figure 5 is impressive, E0(X•/X-) is apparently not the sole determinant of adduct equilibria. For instance, we were unable to identify a DMS-N3 adduct under any conditions, which we should have were N3 to follow quantitatively the above relationships. We suspect, however, oxygen-centered radicals such as HO2• and alkylperoxyl radicals, ROO•, to conform, like OH•, to the linear relationship in Figure 5. If so, this would imply a very low stability for a DMSOOR• adduct, E0(ROO•/ROO-) merely being ca. 0.77 V.37 By way of concluding remark we observe that, in view of the dramatic difference in stability of the DMSOH• adduct in water and the gas phase, the mechanism and rate of the atmospheric destruction of DMS may vary considerably, depending on whether or not water droplets are present. Acknowledgment. This work owes much to stimulating discussions with Professor K.-D. Asmus and Dr. C. Scho¨neich. We are grateful to the Swedish Natural Science Research Council for its financial support. References and Notes (1) (a) For critical reviews see: Atkinson, R. J. Phys. Chem. Ref. Data Monograph 1989, 1, 1. (b) Atkinson, R.; Baulch, D. L.; Cox, R. A.; Hampson, Jr., R. F.; Kerr, J. A.; Troe, J. J. Phys. Chem. Ref. Data 1992, 21, 1341-1342. (2) Hynes, A. J.; Wine, P. H.; Semmes, D. H. J. Phys. Chem. 1986, 90, 4148. (3) Gu, M.; Turecek, F. J. Am. Chem. Soc. 1992, 114, 7146. (4) McKee, M. L. J. Phys. Chem. 1993, 97, 10971. (5) Bonifacic, M.; Mo¨ckel, H.; Bahnemann, D.; Asmus, K.-D. J. Chem. Soc., Perkin Trans. 2 1975, 675. (6) Chaudri, S. A.; Go¨bl, M.; Freyholt, T.; Asmus, K.-D. J. Am. Chem. Soc. 1984, 106, 5988. (7) Scho¨neich, C.; Bobrowski, K. J. Am. Chem. Soc. 1993, 115, 6538. (8) Janata, E.; Veltwisch, D.; Asmus, K.-D. Radiat. Phys. Chem. 1980, 16, 43. (9) Gilbert, B. C.; Hodgeman, D. K. C.; Norman, R. O. C. J. Chem. Soc., Perkin Trans. 2 1973, 1748. (10) Asmus, K.-D. Acc. Chem. Res. 1979, 12, 436. (11) Go¨bl, M.; Bonifacic, M.; Asmus, K.-D. J. Am. Chem. Soc. 1984, 106, 5984. (12) Asmus, K.-D. In Sulfur-Centered ReactiVe Intermediates in Chemistry and Biology; Chatgilialoglu, C., Asmus, K.-D., Eds.; Nato Advanced Study Institutes Series A; Plenum: New York, 1990; Vol. 197, pp 155172. (13) Meissner, G.; Henglein, A.; Beck, G. Z. Naturforsch. 1967, 22b, 13. (14) Engman, L.; Lind, J.; Mere´nyi, G. J. Phys. Chem. 1994, 98, 3174. (15) Eriksen, T. E.; Lind, J.; Reitberger, T. Chem. Scr. 1976, 10, 5. (16) Fielden, E. M. In The Study of Fast Processes and Transient Species by Electron Pulse Radiolysis; Baxendale, J. H., Busi, F., Eds.; Nato Advanced Study Institutes Series, Reidel: Dordrecht, Holland, 1982; pp 49-62. (17) Bonifacic, M.; Asmus, K.-D. J. Chem. Soc., Perkin Trans. 2 1980, 758. (18) Laurence, G. S.; Thornton, A. T. J. Chem. Soc., Dalton Trans. 1973, 1637.
Dimethylhydroxysulfuranyl Radical (19) Stanbury, D. M. In AdVances in Inorganic Chemistry; Sykes, A. D., Ed.; Academic Press: San Diego, CA, 1989; Vol. 33, pp 81-83. (20) Mo¨nig, J.; Goslich, R.; Asmus, K.-D. Ber. Bunsenges, Phys. Chem. 1986, 90, 115. (21) From the second-order component the third-order rate constant, i.e., k{(C6H5)S+ + 2DMS} was estimated to be ≈1.5 × 1010 M-2 s-1. (22) Recalculated from the value in ref 14 by use of E0(Br2•-/2Br-) ) 1.62 V. (23) Above pH 12.5 use of low grade (97+%) NaOH caused the rate to increase. (24) Scho¨neich, C.; Aced, A.; Asmus, K.-D. J. Am. Chem. Soc. 1993, 115, 11376. (25) Clark, T. In Sulfur-Centered ReactiVe Intermediates in Chemistry and Biology; Chatgilialoglu, C., Asmus, K.-D., Eds.; Nato Advanced Study Institutes Series A; Plenum: New York, 1990; Vol. 197, pp 13-18. (26) From the slopes of the lines the following rate constants were obtained: k{DMS+ + (4-OH-C6H4)2S} ) 4.2 × 109 M-1 s-1 and k{DMS•+ + promethazine} ) 1.2 × 109 M-1 s-1. (27) Stickel, R. E.; Nicovich, J. M.; Wang, S.; Zhao, Z.; Wine, P. H. J. Phys. Chem. 1992, 96, 9875.
J. Phys. Chem., Vol. 100, No. 21, 1996 8881 (28) Ruscic, B.; Berkowitz, J. J. Phys. Chem. 1993, 97, 11451. (29) Jonsson, M.; Lind, J.; Reitberger, T.; Eriksen, T. E.; Mere´nyi, G. J. Phys. Chem. 1993, 97, 11278. (30) Wood, P. M. FEBS Lett. 1981, 124, 11. (31) A referee suggested that methyl viologen may be unstable at pH 14 and therefore may not be reduced even by CO2•-. However, in the presence of formate at pH 14 we could confirm the rapid reduction of MV2+ by CO2•- to form MV•+. The only effect of the high pH was the diminished lifetime of MV•+ as compared with low pH. (32) Hine, J.; Mookerjee, P. K. J. Org. Chem. 1975, 40, 292. (33) Schwarz, H. A.; Dodson, R. W. J. Phys. Chem. 1984, 88, 3643. (34) Kla¨ning, U. K.; Sehested, K.; Holcman J. J. Phys. Chem. 1985, 89, 760. (35) The necessary data for the DMS-halogen adducts were taken from ref 17. (36) Clark, T. J. Am. Chem. Soc. 1988, 110, 1672. (37) Mere´nyi, G.; Lind, J.; Engman, L. J. Chem. Soc., Perkin Trans. 2. 1994, 2551.
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