3354
J. Phys. Chem. 1992, 96, 3354-3359
TABLE I V Rate Constants for Reaction of BIZ'- with ChlOrDrOnIaZifIe in Water/Acetonitrile Solutions HzO/CHJCN k , M-' S-' H,O/CH,CN 1oo/o 7.6 x 109 20/80 80/20 3.5 x 109 ioj90 50/50 1.6 x 1 0 9
k , M-'S-' 6.2 X IO8 3.4 x 108
caused by the use of CPZ-HCl, the DMSO-OH apparently forms the protonated form, D M S O O O H ~ which + , ~ ~ can oxidize Br- and I-, and thus the yields remain in the same range (f25%). The rate constants measured for the reactions of Br;- and 12*depend strongly on solvent composition. The value of k for Br;-
+
Br2*- CPZ
+
-
2Br-
+ CPZ"
(11)
CPZ decreases from 7.6 X lo9 M-I s-I in 100% H 2 0 (nearly identical with the reported value)26 to 110' M-' s-I in 100% DMSO. The rate constants for Br2'- 3-methylindole and for 12'- CPZ also decrease in the same fashion. Very good linear correlation is found (Figure 3) between log k and the mole fraction of DMSO. The slope of the line is larger for the slower 12'reaction than for the faster Br2'- reaction. The magnitude of the slope appears too high to be accounted for by simple solvent polarity effects. In fact, the same reaction of Br2+ CPZ when examined in various water/acetonitrile mixtures (Table IV) also gives a similar correlation with the mole fraction of acetonitrile, but the slope of that line is much smaller than that found with DMSO despite the fact that the dielectric constants for acetonitrile (39) and DMSO (45) are not too different. For solvents of high dielectric constants such as these, the effect of the medium di-
+
+
+
(25) For rate constants of hydroxyl radicals and hydrated electrons with various compounds, see compilation by: Buxton, G. V.;Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 513. (26) Bahnemann, D.;Asmus, K.-D.; Willson, R.L . J . Chem. Soc., Perkin Trans 2 1983, 1661.
electric constant on the reaction is of minor importance compared with specific solvation effectsz7 Reaction 11 becomes slower when the solvent is changed from water, which is a good solvator for both cations and anions, to acetonitrile, which is a poorer solvator for both ions. However, DMSO has a stronger effect because it is a good solvator for the cations and very poor solvator for the anions,28which lowers the driving force for reaction 11. Conclusion
Deoxygenated DMSO solutions containing small amounts of acetone and triethylamine (1%) are a good medium for clean radiolytic one-electron reduction of many solutes. Aerated DMSO solutions containing small amounts of CC14 (or other halogenated hydrocarbons), with or without LiCl, provide a medium for radiolytic oxidation. The yield of reduction and oxidation depends somewhat on concentration and approaches 0.4 pmol/J, the sum of the primary yields of DMSO'+ and e- produced by the radiolysis. This overall yield of reducing and oxidizing species in DMSO is about one-third lower than the value in aqueous solutions. Thus, a wide variety of compounds and reactions may be studied conveniently in DMSO when they cannot be studied in aqueous solutions. However, DMSO exerts a strong effect on the rate constants of various reactions, partly due to its high solvent polarity as compared with many other organic solvents and, more importantly, due to its ability to effectively solvate cations but not anions. Thus, the reactivity of Br{- decreases tremendously upon going from water to DMSO. Furthermore, DMSO forms complexes with C1 and Br atoms which exhibit reactivities different from Clz*- and Br2*-observed in water and other solvents.
Acknowledgment. This research was supported by the Office of Basic Energy Sciences of the Department of Energy. (27) Hammett, L. P. Physical Organic Chemistry. Reaction Rates, Equilibria, and Mechanisms; McGraw-Hill: New York, 1970; p 235. (28) Parker, A. J. Inf.Sci. Technol. 1965, August, 28. Also, ref 27, p 234.
Rate Constants and Temperature Effects for Reactions of CI2- with Unsaturated Alcohols and Hydrocarbons in Aqueous and Acetonitrlle/Water Solutlons S. Padmaja, P. Neta, and R. E. H i e * Chemical Kinetics and Thermodynamics Division, National Institute of Standards and Technology, Gaithersburg, Maryland 20899 (Received: November 18, 1991; In Final Form: January 8, 1992) Absolute rate constants for reactions of the dichlorine radical anion, C12'-, with unsaturated alcohols and hydrocarbons have been measured at various temperatures. The alcohol reactions were measured in aqueous solutions and the hydrocarbon reactions in 1:l aqueous acetonitrile (ACN) solutions. The rate constants for two alcohols and one hydrocarbon were also examined as a function of solvent composition. The room temperature rate constants varied between lo6 and lo9 M-I s-I. The pre-exponential factors, A, were about (1-5) X lo9 M-'s-' for the alcohols in aqueous solutions and about (0.1-1) X lo9 M-' s-' for the hydrocarbons in aqueous ACN solutions. The activation energies, E,, varied considerably, between 4 and 12 kJ mol-' for the alcohols and between 2 and 8 kJ mol-' for the hydrocarbons. The rate constants, k298,decrease with increasing ionization potential (IP) of the unsaturated compound, in agreement with an electrophilic addition mechanism. The activation energies for the unsaturated alcohols decrease when the IP decreases from 9.7 to 9.1 eV but appear to level off at lower IP. Most alkenes studied had IP < 9.1 eV and showed little change in E,. Upon addition of ACN to the aqueous solution, the values of log kZ9,decreased linearly by more than 1 order of magnitude with increasing ACN mole fraction. This decrease appears to result from a combination of changes in the activation energy and in the pre-exponential factor. The reason for these changes may lie in changes in the solvation shell of the CI2'- radical, which will affect the A factor, in combination with changes in solvation of CI-, which will affect the energetics of the reactions as well. Introduction
Although rate constants have been measured for a large number of reactions of inorganic radicals in aqueous solution,' until recently few measurements have been made at other than room temperature. In studies of the temperature dependence of electron-
transfer reactions of inorganic radicals with organic2 and inorganic3 reactants, the reactivity has been found to correlate more with the pre-exponential factor in the Arrhenius equation than with (2) Alfassi, 2.B.; Huie, R. E.; Neta, P.; Shoute, L.C. T.J . Phys. Chem. 1990. 94, 8800.
(1) Neta, P.; Huie, R.E.; Ross, A. B. J . Chem. Phys. Ref. Data 1988, 17, 1027.
(3) Shoute, L. C. T.; Huie, R.E.; Neta, P.;Alfassi, Z . B. J . Phys. Chem. 1991, 95, 3238.
0022-365419212096-3354$03 .OO/O 0 1992 American Chemical Society
The Journal of Physical Chemistry, Vol. 96,No. 8, 1992 3355
Reactions of C12'- with Alcohols and Hydrocarbons
TABLE I: Rate Constants for Reactions of CIi- with Unsaturated Alcohols in Aqueous Solutions reactant T, O C k, M-' s-l T, 'C k , M-I s-I T, O C propenamide 15 5.3 x 106 27 6.6 X IO6 37 i-piopen- 1-01 20 4.9 x IO' 30 5.2 x 107 42 3-buten-2-01 10 4.0 x 107 24 5.2 x 107 37 4-penten-2-01 11 8.2 x 107 24 9.6 x 107 35 9.9 x 107 30 3-buten- 1-01 11 8.5 X IO' 20 2-cyclohexen-1-01 18 1.7 X IO8 31 1.8 X IO8 43 2-buten-1-01 11 1.6 X IO8 23 1.9 X IO8 35 2-methyl-2-propen-1-01 12 2.6 X IO8 25 3.1 X IO8 35 3-methyl-3-buten-1-01 15 4.4 x 108 29 4.7 x 108 45 3-methyl-2-buten- 1-01 12 6.3 X IO8 25 7.0 X lo8 36
the activation energy. Similar behavior was observed for the sulfate radical, SO4'-, in electron-transfer reactions4, but the reactivity of this radical in hydrogen abstraction reactions was found to be determined by the activation energy.s There have been no systematic temperature effect studies of reactions of simple inorganic radicals with organic compounds where addition to a double bond would be expected to be the predominant mechanism. In the gas phase, addition to the double bond is frequently characterized by activation energies which decrease with increasing substitution around the double bond and by minimal changes in the pre-exponential factorsU6 The activation energies, in turn, are found to correlate with the ionization potential of the reactant. In the present work, we have investigated the reaction of the dichlorine radical anion, C12'-, with a number of alcohols and hydrocarbons containing carbon-carbon double bonds and with propenamide (acrylamide). In addition, we have determined rate constants for the reactions of SO4*-and Br2*-with a few of these compounds.
Experimental Section The rate constants were determined by laser flash photolysis with kinetic spectrophotometric detection. The basic apparatus and experimental technique has been described in detail previ0usly.49~ (During the latter part of this work, the Lambda Physik EMG 201 MSC excimer laser was replaced by a Questek Model 2320 excimer laseras) The SO4*-radical was produced by the excimer laser (248-nm) flash photolysis of a sodium persulfate solution (typically, 1 mM)7 S2OS2- hv 2S04*(1) and then allowed to react with 0.1 M C1-: SO4'- + C1- S042- C1' (2) C1' + c1- c12(3) The rate constant for reaction 2 is 4.7 X lo8 M-l s-I at 0.1 M ionic strength, independent of t e m p e r a t ~ r e .Therefore, ~ in less than 0.1 ~.tsall of the SO,*-is converted to C12'-. The pH of the solution was adjusted to the range 2-3 by addition of HC104. Although the rate constants for the reaction of C12'- with unsaturated alcohols could be measured in aqueous solutions where the production of C1;- from SO4' is well understood, the reactions of the alkenes could not be measured in the same system because of the limited solubility of hydrocarbons in water. Therefore, we examined the possibility of measuring these rate constants in mixed aqueous/organic solvents. The organic solvent that is least likely to interfere with the desired reactions is acetonitrile (ACN), since we expect its reactions with SO4*-and with C1;- to be quite slow9 and its optical absorption at 248 nm is negligible. Indeed,we found that the reactions of these radicals can be monitored in the presence of up to -80% ACN with little loss in sensitivity.
+
-
--
+
~
C. L. J . fhys. Chem. 1990, 91, 8561. ( 5 ) Clifton, C. L.; Huie, R. E. Inr. J . Chem. Kinel. 1989, I t . 677. (4) Huie, R. E.; Clifton,
(6) See,for example: Huie, R. E.; Herron, J. T. Inr. J . Cbem. Kincf., Symp. No. I 1975, 165. (7) Huie, R. E.; Clifton, C. L.;Altstein, N. Radiaf. Phys. Chem. 1989, 33, 361. (8) The mention of commercial equipment or material d m not imply recognition or endorsement by the National Institute of Standards and Technology, nor does it imply that the material or equipment used are necessarily the best available for the purpose (9) Steenken, S.;McClelland, R. A. J. Am. Chem. Soc. 1989, I l l , 4967.
k , M-I s-I 7.6 X IO6 6.5 x 107 6.5 X 1.0 x 1.1 x 1.9 X
IO' 108 108
IO8
2.0 x 108 3.2 X IO8 5.6 X IO8 7.4 x 108
k, M-I s-l 8.7 X IO6 7.4 x 107 6.8 X lo7 1.2 x 108 1.2 x 108
T, OC 48 54 49 47 40 52 45 45 56 46
2.0 x 108
2.1 x 108 3.4 x 108 5.7 x 108 8.2 X IO8
TABLE 11: Arrhenius Parameters and k2% Values for the Reactions of CIi- with Unsaturated Alcohols in Aqueous Solutions
propenamide 2-propen-1-01 3-buten-2-01 4-penten-2-01 3-buten-1-01 2-cyclohexen- 1-01 2-buten-1-01
9.50 11.7 f 0.7 9.67 9.9 f 2.2 9.50 11.2 f 1.3 9.38 9.56
8.92 9.13 2-methyl-2-propen-1-01 9.26 3-methyl-3-buten-1-01 9.11
(corrected) 3-methyl-2-buten-1-01
(corrected)
8.63
7.9 f 0.9 8.9 f 1.0 4.4 f 0.3 4.8 f 0.8 5.8 f 1.2
5.9 & 0.9 5.2 5.2 f 0.5 4.1
8.9
6.3 X IO6 5.0 X IO'
9.4 9.7 9.4 9.6 9.0 9.1 9.5 9.7
1.0 x 108 1.8 X IO8 1.8 X IO8 3.0 X IO8 4.7 x 108
9.6 9.7
6.9
5.1 x 107
9.6 x 107
X
lo8
9.6
Experiments could not be carried out in 100% ACN since Na2S20s is not soluble. The suitability of this system for studies of C12'- reactions also depends upon the rate of reaction of SO4'- with Cl-. Experiments showed that the rate constant decreases upon an increase in the fraction of the organic component. In a 1:l solvent mixture the rate constant was 1.0 X lo7, 1.3 X lo7, 1.0 X lo7,and 9.1 X lo6 M-' s-l at 14,27, 39, and 50 OC, respectively. As in the case of aqueous solutions, there is virtually no temperature dependen~e.~ Although these values are nearly 50 times lower than the rate constants in the absence of ACN, the reaction with 0.1 M C1is complete within several microseconds, and thus the reaction of C12*-with the hydrocarbons can be monitored on the longer time scales. The reactions of C12*- with the organic compounds were monitored by following the optical absorbance of the Cl;- radical at 340 nm. The organic radicals produced in these reactions have significant absorptions only below 300 nm, which are generally too weak to permit accurate determination of the rate constant. The concentrations of the organic compounds were kept sufficiently low to minimize their reaction with SO4'-, which for most compounds has a rate constant higher than that of C1- with SO4'-. The results from multiple runs were averaged before undergoing linear least-squares analysis to extract the first-order rate constant. First-order rate constants were determined at four different substrate concentrations, varying by a factor of about 3-5. Second-order rate constants were derived from a weighted least-squares fit to a plot of the first-order rate constant against the substrate concentration. Rate constants were typically measured at four temperatures. Other details of the experiments and the materials used were as described p r e v i o ~ s l y . ~ ~ ~ ~ ' ~ The second-order rate constants for the reactions of CI2'- with the various unsaturated compounds were fit (Figure 1) to the Arrhenius expression k = Ae-E/RTby a weighted least-squares routine (weighted by the reciproeal of the squares of the standard errors from the sccond-order fits). The error limits reported with the Arrhenius parameters are the standard errors from the least-squares fit.
Results Rate constants were measured as a function of temperature for the reactions of the dichlorine radical anion with unsaturated (10) Nahor, G. S.;Neta, P. Inr. J. Chem. Kiner. 1991, 23, 941
3356 The Journal of Physical Chemistry, Vol. 96,No. 8, 1992
Padmaja et al.
TABLE 111: Rate Constants for Reactions of CIj - in Various Water/Acetonitrile Mixtures % T, k T, k, reactant CH,CN OC M-I s-1 M-I s-l OC 2-propen-1-01 0 20 4.9 x 107 5.2 x 107 30 5
18 14 14
IO 30 50 70 0
3-methyl-2-buten-1-01
IO 16 12 13 12 14 16 22 15 16 13 17
IO 30 50 70 20 30 40 50 70
2.3-dimethy1-2-butene
4.0 x 3.6 X 1.3 X 5.5 x 1.8 X 6.3 X 4.5 X 1.6 X 6.4 X 2.0 X 3.4 x 3.6 X 2.2 X 9.2 X 2.3 X
107
IO7 IO7 106 IO6 IO8 IO8 IO8 IO7 10' 108
IO8
IO8 IO7 IO7
30 27 26 20 30 25 29 25 27 27 34 26 28 22 28
5.3 x 4.0 x 1.4 x 7.2 X 1.9 X 7.0 X 4.5 X 1.8 X 6.6 X 2.1 x 2.7 X 2.5 X 2.1 X 9.3 x 3.8 x
107 107 107
IO6 IO6 IO8 IO8 IO8 IO7 107
lo8 IO8 IO8
io7 107
T, OC 42 41 38 38 30 42 36 41 41 37 38 43 38 39 33 40
k, M-I s-I
6.5 x 5.3 x 4.4 x 1.6 x 8.8 X 2.5 X 7.4 X 4.7 X 1.8 X 6.9 X 2.6 x 2.5 X 2.0 X 2.0 X 1.1 x 4.0 x
107 107 107 107
IO6 IO6 IO8 IO8 10'
lo7 107
IO8 IO8 IO8 io8 107
1 I
3.0
'
I
I
I
"
'
1
1
1
"
3.2
'
k,
T, OC 54 51 49 49
7.4 x 5.4 x 4.8 x 1.8 x
107 107 107 107
46 53 53 51 51 52 49 52 49 52
8.2 X 4.9 X 1.9 X 7.7 X 2.9 x 1.6 X 1.4 X 1.7 X 1.2 x 4.7 x
lo8 lo8 IO8
M-I s-I
IO' 107
IO8 IO8 IO8 io8 107
0 0
I
3.4
3.6
1000/T, K.' Figure 1 . Rate constants for the reactions of C12'- with unsaturated alcohols as a function of temperature: 0,propenamide, a, 2-propen-1-01; 0,3-buten-2-01; A,4-penten-2-01; A, 3-buten-1-01; 0, 2-cyclohexen-1-01; +, 2-buten-1-01;V,2-methyl-2-propen-1-01;V, 3-methyl-3-buten-1-01; 0 , 3-methyl-2-buten- 1-01,
alcohols in aqueous solutions, with alkenes in a 1:l acetonitrile/water mixture and with two unsaturated alcohols and one alkene in a number of acetonitrile/water mixtures. The rate constants for reaction of C12'- with unsaturated alcohols in aqueous solutions are summarized in Table I and the Arrhenius parameters in Table 11. The room temperature rate constants ranged from 5 X IO' to 7 X IO8 M-'s-l. The pre-exponential factors were all in the range (1-5) X IO9 M-' s-I, and the activation energies ranged from 4.4 to 11.2 kJ mol-]. The reactions of C12'- with 3-methyl-3-buten-1-01 and 3-methyl-2buten-1-01 are sufficiently fast that they are partly limited by the rate at which the reactants can diffuse toward each other. We have corrected for diffusion by taking" a radius of 0.16 nm for (21;- and estimating 0.3 nm for the alcohols. (Since the alcohols are not ionized, the calculated diffusion rate is not very sensitive to the value of the radii). The extent of the correction is greater at the lower temperatures, so that even though the absolute corrections are small, at most lo%, they have a significant impact on the derived Arrhenius parameters. After a correction is made for diffusion, the average pre-exponeital factor is 3 X lo9 M-' s-'. The rate constants for the reactions of two alcohols and one hydrocarbon were measured as a function of ACN fraction (Table 111 and Figure 2), and the Arrhenius parameters and k298 values are presented in Table IV. Clearly, as in the case of SO4'- + C1-, the rate constants for the reactions of C12'- with the organic ( 1 1 ) Stanbury,
D.M . Inog. Chem. 1984, 23, 2914.
Figure 2. Rate constants for the reactions of CI2'- with 2,3-dimethyl-2butene at different temperatures and in several water/acetonitrile mixtures: 0 , 2 0 % ACN; V, 30% ACN; 0 , 4 0 % ACN; A, 50% ACN; 0,70% ACN.
TABLE I V Arrhenius Parameters and k m Values for the Reactions of CIj- in Various Water/Acetonitrile Mixtures
2-propen- 1-01
0 5
IO
3-methyl-2buten- 1-01
2,3-dimethyl-2butene
30 50 70 0
IO 30 50 70 20 30 40 50 70
9.9 f 2.2 5.9 f 3.1 6.4 f 0.2 7.2 f 0.8 10.8 f 1.8 9.5 f 1.7 5.2 i 0.5 1.8 f 0.5 2.0 f 1.1 4.4 f 0.8 10.1 i 2.2 -15.4 f 3.8 -21.5 f 1.6 -5.2 f 1.2 6.3 f 0.8 15.1 f 4.3
9.4 8.7 8.7 8.4 8.8 8.0 9.7 9.0 8.6 8.6 9.0 5.8
4.7 7.6 9.1 10.5
5.0 x 4.7 x 3.9 x 1.4 x 7.5 X 2.0 X 6.9 X 4.5 X 1.8 X 6.5 X 2.1 x 3.2 X 2.5 X 2.1 X 9.9 x 3.8 X
107
107 107 107 IO6 IO6
IO8 IO8 IO8 IO' 107 10' lo8 IO8 10' IO7
compounds decrease with increasing fraction of ACN. The decrease in k298amounted to more than a factor of 30 on increasing the fraction of ACN from 0 to 70% for the alcohols and almost a factor of 10 for the alkene upon increasing the amount of ACN from 20 to 70%. The reactions of a series of hydrocarbons were measured in 1:1 aqueous ACN; the rate constants are summarized in Table V and Figure 3 and the Arrhenius parameters in Table VI. The value
The Journal of Physical Chemistry, Vol. 96, No. 8, 1992 3357
Reactions of C12'- with Alcohols and Hydrocarbons
TABLE V: Rate Constants for Reactions of Cli- with Unsaturated Hydrocarbons in 1:l Water/Acetonitrile Solutions
T, reactant
O C
-
1 hexene cycloheptene 2-ethyl-I -butene cyclohexene cyclopentene 3-methyl-2-pentene
12 13 12 11
IO
3-ethyl-2-methyl-2-pentene 3-ethyl-2-pentene 2-methyl-2-butene 2,3-dimethyl-2-butene
16 13 15 17 13
k,
M-1 s-I 4.7 x 106 1.5 x 107 2.7 X 2.1 x 4.2 x 5.7 x 5.2 X 6.1 x 6.7 X 9.2 x
10' 107 107 107
IO7 107 IO7 107
T, OC
26 22 22 25 16 21 25 30 31 22
T,
k,
M-I 5.9 x 1.5 x 2.8 x 3.2 x
s-1
106 107 107 107
5.0 x 107 6.2 x 5.5 x 6.3 x 7.1 X 9.3 x
107 107 107 10' 107
TABLE VI: Arrhenius Parameters and k2%Values for the Reactions of Cli- with Unsaturated Hydrocarbons in 1:l Acetonitrile Aqueous Solutions
I-hexene cycloheptene 2-ethyl- I-butene cyclohexene cyclopentene 3-methyl-2-pentene 3-ethyl-2-methyl-2pentene 3-ethyl-2-pentene 2-methyl-2-butene 2,3-dimethyl-2-butene
IO6 107 107 107
9.44 8.91 9.06 8.95 9.01 8.58 8.17
7.1 f 1.4 5.0 1.0 4.1 f 0.4 5.5 1.4 3.0 1.8 5.5 f 1.6 4.3 f 0.5
**
8.1 8.1 8.2 8.4 8.2 8.7 8.5
5.6 X 1.6 x 2.9 x 3.0 x 5.2 x 5.4 x 5.5 x
8.68 8.27
1.6 0.3 4.1 0.5 6.3 f 0.8
8.1 8.6 9.1
6.6 X IO7 7.6 X IO7 9.9 x 107
*
107 107 107
of k298increased by a factor of 18 going from 1-hexene to 2,3dimethyl-2-butene. We have also measured the rate constants for the reaction of Cl,' with propenamide (Tables I and 11). A lower pre-exponential factor and higher activation energy was found than for any of the alcohols, resulting in a kB8value that is 1 order of magnitude lower than that for the least reactive alcohol, reflecting a strong deactivating effect of the amide group. Rate constants for the reactions of SO4'- and Br;- with some of the unsaturated alcohols have also been measured. The rate constant for the reaction of SO4'- with 2-propen-1-01 was (2.9 f 0.1) X lo9 M-I s-I a t 12 OC, (3.0 f 0.1) X lo9 M-l s-l at 27.5 OC, (3.6 f 0.1) X lo9 M-I s-l at 41 "C, and (3.6 f 0.2) X lo9 M-l s-I at 52 "C. These values result in an Arrhenius expression of k = (2.8 f 2.0) X 1Oloexp((3.5 f 1.8 kJ mol-')/RT). SO4'was found to react with propenamide with rate constants of (1.3 f 0.1) X lo* and (1.4 X 0.1) X lo8 M-ls-' a t 25 and 36 "C. These values are in good agreement with the previous room temperature determinations by pulse radiolysis.I2 The reactivity of Br2+ toward 2-propen-2-01 and 3-methyl-2buten-1-01 was investigated in a solution containing 0.1 M KBr and 1 mM Na2S208at concentrations of alcohol ranging up to 10 mM. No reaction was observed. Therefore, we estimate that the rate constant for these reactions is less than lo6 M-' s-I. There has been a previous pulse radiolysis determination of the rate constant for the reaction of C12'- with 2-propen-1-01 (allyl alcohol) at room temperature.13 The reported value, 5.9 X lo8 M-' s-I, is almost 10 times greater than our 25 OC value. We have also carried out pulse radiolysis studies of this reaction at -20 OC in an N 2 0 saturated solution of 0.1 M KCl at pH 2.3. We obtained a rate constant of (8.5 f 0.5) X lo7 M-' s-l, much closer to the present flash photolysis value. Further, the high value reported previously for 2-propen-1-01 is out of line with the other rate constants reported in that paper for reaction of Clz'- with other unsaturated compounds. Discussion
The rate constants for reaction of C12'- with alkenes increase with increasing alkyl substitution at the double bond, indicating ( I 2) Maruthamuthu, P. Makromol. Chem., Rapid Commun. 1980, I , 23. (13) Hasegawa, K.; Neta, P. J . Phys. Chem. 1978, 82, 854.
IO8
2
M-l
40 30 31 40 26 41 39 43 40 33
7.1 X 1.6 x 3.1 x 3.3 x 5.2 x 6.4 x 6.2 X 6.4 X 7.5 x 1.1 x
L
T,
k,
OC
k,
OC
M-I
IO7 lo7
51 40 40 52 36 55 53 53
6.9 X 1.8 x 3.1 x 3.7 x 5.3 x 7.3 x 6.4 x 6.6 X
107 108
49
1.2 x 108
s-I
IO6 107 107 107 107 107
* - - - - - --e - - - - - * -- - -
4
s-l
IO6 107 107 107 107 107 107
IO7
4 1
_ I
I
t lo6 3.0
3.2
3.4
3.6
1000/T, K-' Figure 3. Rate constants for the reactions of C12*- with alkenes as a function of temperature in 1: 1 water/acetonitrile mixtures: 0, I-hexene; 0 , cycloheptene; V,2-ethyl-1-butene; V, cyclohexene; 0,cyclopentene;
A, 3-ethyl-2-methyl-2-pentene; m, 3-methyl-2-pentene; A, 3-ethyl-2pentene; 0 , 2-methyl-2-butene; 4 , 2,3-dimethyl-2-butene.
an electrophilic addition mechanism, as found before for reactions of perfluorobutylperoxyl radicals with alkenes.'O Indeed, chlorine atom adducts have been observed by ESR following the reactions of C12'- with several such compound^.'^ The higher reactivity of disubstituted alkenes with internal double bonds compared to those with terminal double bonds is consistent with what has been observed for ozone or atomic oxygen in its gas-phase reactions with alkenesU6Furthermore, similar to O3and 03P,the activation energy for the reactions of all these electrophilic reagents decreases with increasing alkyl substitution at the double bond. Generally, the rate constant and the activation energy for the addition of an electrophilic reagent to a double bond are expected to be related to such properties as the ionization potential (IP) or the sum of the TaftI5 polar substituent constants, u*, for the substituents about the double bond. Indeed, we find that log k298 decreases with increasing IPI6 for both the unsaturated alcohols in aqueous solutions and the alkenes in a 1:1 water/ACN solvent (Figure 4), although the correlations have a relatively wide scatter. Correlation of log k298 vs u* gave similar results for the hydrocarbons (not shown) but was not done for the alcohols because of insufficient data on u* values. A reasonable, albeit nonlinear, relationship is found also between the activation energy and the IP for the alcohols (Figure 5 ) . The main outlier is the result for 3-buten-2-01, which has a higher (14) Gilbert, B. C.; Stell, J. K.; Peet, W. J.; Radford, K. J. J . Chem. Soc., Faraday Trans. I 1988, 84, 3319. (IS) Taft, R. W., Jr. In Steric Effects in Organic Chemistry; Newman, M. S.,Ed.; Wiley: New York, 1956; p 619. (16) The IP values were taken from: Lias, S. G.;Bart", J. E.; Liebman, J. F.; Holms, J. L.; Levin, R. D.; Mallard, W. G. J . Phys. Chem. Ref. Data 1988, 17, Supplement No. I . The IP values for some of the alcohols (2cyclohexen- 1-01, 3-methyl-3-buten-1-01,and 3-methyl-2-buten- 1-01) were not available and were estimated to be the average of the values for the alkenes formed by replacing the OH group with a H or a CH,.
3358 The Journal of Physical Chemistry, Vol. 96, No. 8,1992
Padmaja et al.
io9
1
1OB
1o8
-
-k
-i:
v)
v)
9
J
10'
10'
1o6
9
8
0.0
10
0.1
0.2
0.4
0.3
0.5
Mole Fraction CH3CN
I. P., eV Figure 4. Dependence of kZ9*on the ionization potentials for the unsaturated alcohols ( 0 )and for the alkenes (0). 1 2 , , , , , 1 ! , , , [ , , , , ,
c
1
,
e I
,
I
Figure 6. Dependence of
k298 on the mole fraction A C N for 2.3-di3-methyl-2-buten-1-01 (A),and 2-propen-1-01 ( 0 ) . methyl-2-butene (0).
ZOr,
]
-I
I
1
I
I
I I
I
I 1
I
I
I 1
I
I
L
10
1
I
I I
i
12
-1
10
-
T
-
0-
2
6
2
t 2 ~ 8.0
1
A
"
"
~
6.5
"
"
~
1 " 9.0
"
"
"
9.5
~
-20
10.0
activation energy than would be predicted, although it is not found to be an outlier in the plot of log k298 vs IP (Figure 4) because the effect of the high activation energy is compensated by the high pre-exponential factor. In contrast with the correlation observed for the alcohols in aqueous solutions, the alkenes in 1:l water/ ACN solutions exhibit no apparent correlation between the activation energy and the IP. It should be pointed out, however, that the IP values for the hydrocarbons are generally lower than those of the unsaturated alcohols studied and that the curve for the alcohols appears to level off at lower IP values. Addition of acetonitrile to the aqueous solvent affects both k298 and E, very strongly (Table IV). The values of log k298decrease linearly with t h e mole fraction of ACN (Figure 6)for both the alcohols and the hydrocarbon. The values of E, and log A, on the other hand, are found to decrease and then increase with the mole fraction of ACN. For 2,3-dimethyl-2-butene (Figure 7), these changes in E, and log A are quite substantial and almost parallel, indicating the existence of an isokinetic relationship" (inset, Figure 7).18 (17) Hammett, L. P. Physical Organic Chemistry. Reaction Rates, Equilibria, and Mechanisms, 2nd ed.; Mdjraw-Hill: New York, 1970. Exner, E. frog. Phys. Org. Chem. 1973, 10, 41 I . Linert, W.;Jameson, R. F. Chem. SOC.Rev. 1989, 18, 477. (18) The values of log A and E, for the unsaturated alcohols vary over a narrower range than those for the alkene and are, therefore, probably not sufficiently accurate to give a good isokinetic relationship.
6
-
-
I.P., eV Figure 5. Correlation between E, and the ionization potential for the unsaturated alcohols ( 0 )and for the alkenes (A).
E
4 M 0 4
-
"-10
111
0.0
0.1
0.2
0.4
0.3
4
0.5
Mole Fraction CH,CN Figure 7. Correlation of E. ( 0 )and log A (0)with ACN mole fraction for 2,3-dimethyl-2-butene. Inset: the isokinetic relationship between E, and log A.
The nonlinear changes in E, and log A with ACN mole fraction are probably the result of competing effects arising from differences in the solvation of reactants, intermediates, and products. The reactions of C12'- with unsaturated organic compounds can be considered to proceed through an intermediate complex followed by dissociation to products.19 Cl2'-
+ RZC=CR*
-
[R2C'-C(Clz-)R2] C1-
-
+ R$'-C(CI)R2
(4)
C12'- should be strongly solvated in water but less so as the acetonitrile fraction is increased. On the other hand, the alkenes are much better solvated in ACN. The intermediate complex would probably by solvated both by water and by ACN. The chloride ion would be highly solvated in water, while the chlorine adduct of the alkene would be better solvated by the ACN. The ~
(19) The possibility that the reaction involves addition of CI' present in equilibrium with C12*-,to form the adduct directly, can k ruled out since the
magnitude of the rate constants measured for these reactions, when corrected for the fraction of CI' present in the equilibrium, would result in values much greater than the diffusion controlled rate
J. Phys. Chem. 1992,96, 3359-3366
3359
net result of all of this is likely to be a region of greater stability of the intermediate complex, corresponding to the solvent mixture with the lowest E,. The fact that log A is also at a minimum in this range suggests that the process involves structure making in the solvent shell. The overall rate constant is also influenced by the rate of the second step, which should decrease with increasing ACN, and thus k298decreases with increasing ACN fraction. In a recent study,*O we reported on the effect of solvent composition (water/2-propanol (2-PrOH)) on a series of reactions of chlorinated methylperoxyl radicals with organic reductants.
In that series, log kZg8decreases with increasing organic mole fraction. This was due to a decrease in the pre-exponential factor upon increasing the 2-PrOH mole fraction. This, however, was partly compensated for by a concurrent decrease in the activation energy. In that work, we found an isokinetic relationship between E, and log A, as we found here for 2,3-dimethyl-2-butene; both cases point to the complexity of solvent effects on rate constants and especially the influence of temperature on these effects, which may be helpful in distinguishing between polarity effects and specific solvation effects.
(20) Alfaui, Z. B.; Huie, R. E.; Kumar, M.; Neta, P. J . Phys. Chem. 1992, 96, 767.
Acknowledgment. This research was supported by the Office of Basic Energy Sciences of the U. S. Department of Energy.
Ultrafast Studies of Solvent Effects in the Isomerization of cis-Stilbene Jane K. Rice* and A. P. Baronavski Naval Research Laboratory, Washington, D.C. 20375-5000 (Received: November 25, 1991; In Final Form: February 6, 1992)
The ultrafast dynamics of cis-stilbene in solution are studied using a femtosecond pumpprobe laser system. Excited-state lifetimes in n-hexane, tetradecane, methanol, 1,Cdioxane, and cyclohexane are examined following excitation at 3 12 nm. In addition, spectra of excited cis-stilbene within 500 fs of excitation are recorded using an optical multichannel analyzer (OMA). The lifetimes observed in n-hexane and tetradecane are in excellent agreement with literature values and correlate well with solvent bulk viscosity. The measured lifetimes in methanol, 1,4-dioxane, and cyclohexane do not correlate with solvent bulk viscosity. The lifetime in methanol is 0.50 0.03 ps, only slightly longer than the gas-phase value of 0.32 f 0.02 ps (Greene, B. I.; Farrow, R. C. J . Chem. Phys. 1983, 78, 3336-3338). The lifetime in 1,4-dioxane is similar to that of an alcohol of the same bulk viscosity. The lifetime in cyclohexane is much longer than lifetimes in linear alkanes of comparable viscosity. An examination of several probe wavelengths between 490 and 690 nm in several solvents reveals little or no wavelength dependence of the lifetime with the possible exception of cyclohexane. The cis-stilbenelifetime data in methanol and 1,Mioxane suggests that microscopic solventsolute interactions based on polarity are present and are larger than the bulk viscosity dependencies. The relatively slow cis- to trans-stilbene isomerization rate in cyclohexane may indicate a steric hindrance to isomerization.
Introduction Recent advances in ultrafast laser sources have led to direct time-resolved studies of photochemically initiated relaxation processes. Molecules that undergo large-amplitude structural changes in the excited state through radiationless relaxation are of particular interest. The polyenes and diphenylpolyenes have been the focus of many studies to determine the location and assignment of their excited states.I-j Numerous other studies have addressed the solvent dependence of large-amplitude relaxation rates on excited-state s u r f a c e ~ and ~ - ~the development of theoretical models9J0 to explain these rates. Stilbene (1,2diphenylethylene) undergoes isomerization through large-amplitude motion on its excited-state surface. A 1986 review of stilbene photophysics and other molecules which involve radiationless relaxation are presented in a book on ultrafast phenomena by G. R. Fleming." New information on the stilbene excited-state ( I ) Hudson, B. S.; Kohler, B. E.; Schulten, K. Excited States 1982, 6, 1-95. (2) Goldbeck, R. A.; Twarowski, A. J.; Russell, E. L.; Rice, J. K.; Birge, R. R.; Switkes, E.; Kliger, D. S.J . Chem. Phys. 1982, 77, 3319-3328. (3) Kohler, B. E. J . Chem. Phys. 1990, 93, 5838-5842. (4) Keery. K. M.; Fleming, G. R. Chem. Phys. Lett. 1982, 93, 322-326. ( 5 ) Rothcnberger, G.;Negus, D. K.; Hochstrasser, R. M.J . Chem. Phys. 1983, 79. 5360-5367. (6) Hicks, J. M.; Vandenall, M. T.; Sitzmann, E. V.; Eisenthal, K. B. Chem. Phys. Letr. 1987, 135, 413-420. (7) Ben-Amotz, D.; Harris, C. B. J . Chem. Phys. 1987, 86, 4856-4870. (8) Zeglinski. D. M.; Waldeck, D. H. J . Phys. Chem. 1988,92,692-701. (9) Kramers, H. A. Physica 1940, 7, 284-304. (IO) Bagchi, B.; Fleming, G. R.; Oxtoby, D. W. J . Chem. Phys. 1983, 78, 7375-7385.
surface and an improved understanding of the isomerization dynamics continue to be of interest as indicated by several recent reports which examine excited cis-stilbene lifetimes on the picosecond and femtosecond time s ~ a l e . ' ~ - ' ~ A diagram of the potential energy surface of the cis- and trans-stilbene isomers in n-hexane is shown in Figure 1. This is based on early descriptions by Saltiel.'sJ6 There is a large barrier to isomerization on the ground-state potential centered near the twisted configuration at 90'. On the excited-state surface near 90' is a minimum referred to as the phantom state. Relaxation on the excited-state surface from either cis- or transstilbene leads to the phantom state. A second minimum shown in the trans-stilbene excited state is the result of an avoided crossing of the 'B, state with the higher energy (planar geometry) IA state.l7 bollowing excitation of the trans-stilbene ground state to the excited state (planar) in hexane, the molecule crosses a 3.5 kcal barrier1s,'9 to reach the twisted phantom state. The excited ( I I ) Fleming, G.Chemical Applications of Ultrafasr Spectroscopy; Oxford University Press: New York, 1986; pp 179-184. (12) Sension. R. J.; Repinec, S.T.; Hochstrasser. R. M. J . Phys. Chem. 1991,95,2946-294a. (13) Todd, D. C.; Jean, J. M.; Rosenthal, S.J.; Ruggiero, A. J.; Yang, D.; Fleming, G. R. J. Chem. Phys. 1990, 93, 8658-8668. (14) Abrash, S.;Repinec, S.;Hochstrasser, R. M. J. Chem. Phys. 1990, 93, 1041-1053. (15) Saltiel, J. J . Am. Chem. SOC.1967, 89, 1036-1037. (16) Saltiel, J. J. Am. Chem. SOC.1968, 90, 6394-6400. (17) Orlandi, G.; Siebrand, W. Chem. Phys. Lett. 1975, 30, 352-354. (18) Courtney, S.H.; Fleming, G.R. J. Chem. Phys. 1985,83, 215-222. (19) Sundstrom, V.; Gillbro, T. Ber. Bunsen-Ges. Phys. Chem. 1985.89, 222-226.
This article not subject to US.Copyright. Published 1992 by the American Chemical Society
