J. Phys. Chem. 1995,99, 13970-13975
13970
Rate Constant Determination for the Reaction of Sulfhydryl Species with the Hydrated Electron in Aqueous Solution? Stephen P. Mezyks Radiation Laboratory, University of Notre Dame, Notre Dame, Indiana 46556 Received: April 10, 1995@
The techniques of pulse radiolysis and absorption spectroscopy have been used to determine hydrated electron reaction rate constants with the sulfhydryls mercaptoethanol, 2-aminoethanethiol, cysteine, and penicillamine and the disulfides cystine and penicillamine disulfide, over the pH range 5.0-13.0 in aqueous solution. These values were found to be strongly dependent on the acid-base properties of the sulfhydryl compounds, with the rate constants dependent on the state of protonation of both the amino and mercapto groups, as well as the overall charge in these compounds. Rate constants for reaction of individual sulfhydryl and disulfide species have been determined using literature ionization constants. These values are contrasted to previous literature data.
Introduction The role of sulfides (RSH) and disulfides (RSSR) in radiation protection has been known for many years.',2 The redox properties of organic sulfur compounds and, in particular, their radicals are of considerable interest for the understanding of many processes in biological and related model systems. Previous studies have demonstrated that organic disulfides can readily undergo sulfur-sulfur bond cleavage by one electron r e d u c t i ~ n ,and ~ . ~the versatile chemistry of the resulting radicals has resulted in increased attention being paid to the formation and subsequent reactions of such sulfur centered radical specie^.^ In several recent investigations,6.' rate constants for the reactions of thiyl radicals with sulfhydryls in aqueous solution to form the reducing disulfide anion have been determined. The determination of unique values for this process is complicated, as the observed disulfide anion growth has contributions from several processes with similar rate constants.* Individual rate constant elucidation can be performed only if the fundamental radiation chemistry of these species is well established. Most of the numerous investigations reported in the literature have focussed on the formation of radicals formed by oxidation p r o c e s ~ e s . ~ -There '~ is surprisingly little kinetic information available on the reaction of the hydrated electron with simple sulfides and disulfides in aqueous solution,* and the effects of ionic forms on these rate constants have only been sporadically in~estigated.~.'~ In general it has been found that the rate constants decrease with increasing pH, and can be directly related to the pK, values of the -SH and other ionizable groups in the molecules. The absolute values of hydrated electron reaction rate constant have been shown to follow the overall charge of the sulfhydryl, and the effects of multiple deprotonations appear to be compounding.I4 Product studies have shown that hydrated electron reaction with sulfhydryls has two potential pathways:
'The research described herein was supported by the Office of Basic Energy Sciences of the Department of Energy. This is Contribution No. NDRL 3816 from the Notre Dame Radiation Laboratory. Permanent address: Research Chemistry Branch, AECL, Whiteshell Laboratories, Pinawa, Manitoba, ROE 1L0,Canada. Abstract published in Advance ACS Absrrucrs, August 15, 1995. 7
@
0022-3654/95/2099- 13970$09.00/0
eaq-
+ RSH (+Hi')
eaq-
+ RSH (+H+)
-
+ H2S RS' + H, R'
(1)
(2)
with reaction 1 generally being dominant.I4 The corresponding reaction with disulfides3 is by an addition process to produce the radical anion:
+
eaq- RSSR
-
RSSR'-
(3)
In this work the systematic investigation of the measurement of reaction rate constants of the hydrated electron with the sulfhydryls mercaptoethanol, 2-aminoethanethiol (cysteamine), cysteine, and penicillamine and the disulfides cystine and penicillamine disulfide in aqueous solution has been performed. From the measured pH dependence of these values, limiting rate constants for each of the sulfhydryl ionic forms have been calculated.
Experimental Section The setup used for these experiments has been described in detail in several previous publication^,'^-'^ and hence only a brief description shall be given here. The pulse radiolysis system consists of a 8 MeV LINAC, which delivered 10 ns electron pulses to the sample cell. In all cases the pulse radiolysis traces were normalized for small fluctuations in radiolysis dose. Absolute dosimetry was carried out using N20 saturated SCN- solutions (IO-* mol dm-3, 1 = 472 nm, GE= 4.92 x 104). The compounds used in this study (Aldrich) were of the highest purity available and used as received. Solutions were prepared immediately before use, by dissolving known amounts of the sulfhydryls in N2 saturated Millipore Milli-Q filtered water containing 1.0 mol dm-3 tert-butyl alcohol and Baker Analyzed monobasic phosphate or borax (2.0 x mol dm-3). Exact pH values were obtained by using a pH meter with the addition of small amounts of NaOH (Fischer, ACS) or HC104 (Aldrich, ACS 70%) prior to the sulfhydryl addition in order to minimize the oxidation of these compounds. All solutions were flowed through the irradiation cell at a sufficient rate to ensure that a fresh sample was irradiated each time. During the irradiation 0 1995 American Chemical Society
J. Phys. Chem., Vol. 99, No. 38, 1995 13971
Rate Constant Determinations of Sulfhydryl Species
TABLE 1: pH Dependence of the Corrected Rate Constants Determined in This Study for the Reaction of Hydrated Electron with Sulfhvdrvls in Aqueous Solution penicillamine mercaptoethanol cysteamine cysteine penicillamine cystine disulfide pH (dm3mol-' s-I) (dm3 mol-' s-') (dm3 mol-' s-l) (dm3 mol-' s-I) (dm3 mol-' s-I) (dm3 mol-' s-') 5.0 (8.26 f 0.10) x lo9 (1.62 f 0.02) x 1O'O (8.52 f 0.42) x lo9 (5.14 f 0.09) x lo9 (1.04 f 0.03) x 1O'O (6.44 & 0.12) x lo9 6.0 (8.34 f 0.92) x lo9 (1.58 f 0.07) x 1O1O (8.43 f 0.13) x lo9 (4.95 f 0.13) x lo9 (1.03 f 0.03) x 1Olo (6.45 f 0.12) x lo9 7.0 ( 8 . 1 8 f 0 . 1 1 ) x lo9 (1.55f0.06) x 1O1O (8.28 f 0.18) x 109 (4.87 f 0.06) x lo9 (1.03 f 0.04) x 1O'O (5.80 f 0.06) x lo9 8.0 (8.38 f 0.15) x lo9 (1.27 f 0.06) x 1O'O (6.62 f 0.06) x lo9 (3.60 f 0.04) x lo9 (8.37 f 0.06) x lo9 (4.60 f 0.07) x lo9 9.0 (6.32 f 0.09) x lo9 (5.70 f 0.11) x lo9 (3.22 f 0.06) x lo9 (1.14 f 0.02) x lo9 (5.13 & 0.10) x lo9 (2.19 f 0.05) x lo9 10.0 (1.70 f 0.02) x lo9 (1.53 f 0.03) x lo9 (1.05 ~f0.20) x 109 (4.98 f 0.07) x lo8 (3.41 f 0.04) x lo9 (1.79 f 0.04) x lo9 11.0 (2.05 f 0.01) x lo8 (9.21 & 0.05) x los (5.13 f 0.05) x lo8 (2.1 1 f 0.02) x lo8 (2.22 f 0.08) x lo9 (8.26 f 0.09) x lo8" 12.0 (3.63 f 0.05) x lo7" (7.41 f 0.06) x lo8 (8.50 f 0.19) x lo7" (6.96 f 0.08) x lo6" (1.54 f 0.06) x lo9" (7.66 f 0.09) x lo8" 13.0 (2.12 f 0.10) x lo7" (3.55 f 0.03) x lo8" (3.60 f 0.08)x 1070 (2.12 f 0.13) x (1.51 f 0.03) x lo9" (8.30 & 0.08) x lo8"
~~~~
~~~
a
~~~
Corrected for ionic strength according to eq 7.
0
HSCH,CH,OH
0'5
H+
+ -SCH,CH,OH
(6)
Furthermore, the ionic strength of the solutions at the two highest pH's was not negligible and hence would increase the observed rate of reaction. These observed rate constants were thus corrected to zero ionic strength, using the equationla
0.4
-
-
0.3
u)
2 E
'4
0.2
(7)
0.1
"."
0.0
1.o
2.0
3.0
4.0
1 O3 (Mercaptoethanol] mol dm'3
Figure 1. Comparison of the directly measured (0)and ionic strength
corrected (W) pseudo-first-order decay rate constants for hydrated electron reaction with mercaptoethanol at pH 12.0. process, the solution vessels were bubbled with only the minimum amount of N2 required to prevent air ingress, to prevent loss of the volatile sulfhydryls. Typically about 1020 pulses were averaged to obtain a single trace. The rate constants for hydrated electron reaction were determined from the pseudo-fust-order decay of its absorbance at 700 nm, using a Corning 3-71 filter in the analysis beam path to minimize interference from shorter wavelengths. All experiments were performed at room temperature (22 f 2 "C).
Results and Discussion Sulfhydryl reactions. Mercaptoethanol. This compound was chosen for initial study, as it is one of the simplest sulfhydryls, with only a single pKa of 9.50.l' Typical results for a hydrated electron reaction at pH 12.0 are seen in Figure 1, with the scavenging rate constants obtained across the pH range 5.0-13.0 shown in Figure 2a (see individual values in Table 1). The error bars shown for these points corresponds to 1 standard deviation. Over the pH range of this study, the charge on this sulfhydryl changes from 0 at pH 5.0 to -1 at pH 13.0, therefore at intermediate pH values, the overall reaction is by a combination of the two limiting reactions:
+ eaq + -SCH,CH,OH
-
ea¶- HSCH2CH20H
products
(4)
products
(51
where these two species exist in equilibrium:
where Z is the charge on the sulfhydryl, (-1), ,n is the total ionic strength of the solution, and a = R13 where R is the sum of the radii of the two reactants. Calculation of the ionic strength of the different solutions allowed for the sulfhydryl concentration, as well as the different ionic forms of the buffer, using the known pKa values for borate (9.14, 12.74, and 13.80),19 as well as corrections for the small amount of polymerization that occurred.20 It was assumed that the pK, values for these polymer species were the same as for borate. The reaction radius of the hydrated electron is known to be 2.5 5 r 5 3.0 A,,' an average value of 2.75 8, was used in these calculations. However, no equivalent value could be found for mercaptoethanol in the literature. Thus this radius was calculated utilizing the geometry optimization routines of the HyperChem 3.0 program. The inbuilt RHF semiempirical methods Ah41 and PM3, which are based on NDDO (neglect of diatomic differential overlap) approximation, were used for these calculations. Optimized geometries were f i s t determined for the protonated form of only the mercaptoethanol. This structure was then surrounded by a periodic box of water molecules, and then the entire system geometry re-optimized. This calculation was repeated with increasing numbers of water molecules until the change in the sulfhydryl geometry was less than 5%. The reaction radius of the sulfhydryl was taken as half the greatest linear distance between two atoms within the molecule, corresponding to a value of 2.48 8,. The radius for the ionized form was assumed equivalent to the fully protonated value. Values for all the sulfhydryls were determined in this manner, and the calculated values used in the ionic strength corrections are listed in Table 2. Using the above parameters, the mercaptoethanol rate constants obtained at pH 12.0 were corrected for ionic strength, and these values are contrasted to the directly measured rate constants Figure 1. This ionic strength correction to measured rate constants at lower pH's was not possible, as the individual rate constants for reactions 4 and 5 were not known. However, as the ionic strength of the solutions with pH's less than 12 was significantly lower and had a greater fraction of the mercaptoethanol in its fully protonated (neutral) form, such corrections are believed
Mezyk
13972 J. Phys. Chem., Vol. 99, No. 38, 1995 1.4
I
I
I
I
1
I
TABLE 3: Summary of Limiting Rate Constants for Hydrated Electron Reaction with Sulfhydryl Species in Aqueous Solution rate constant (dm-'mo1-I s-')
species
I
4.0
I
I
I
I
1
mercaptoethanol HSCHKH,OH
(8.34 0.1 I ) x 109 (1.75 i 0.04j x 107
cysteamine
(1.53 f 0.05) x 1O'O (1.73 rt 0.11) x 109 (4.11 f 0.30) x IO8
cysteine
(8.53 f 0.14) x IO9 (1.53 f 0.14) x lo9 (4.06 f 0.91) x lo7 (5.06 f 0.04) x lo9 (4.12 f 0.53) x lo8 -4 x 106
penicillamine
c
'a 3.0
-
r
m
2'ol
(1.51 f 0.21) x 109
......*...
6 x-
2
(1.06 f 0.02) x 1O'O (4.95 rt 0.30) x 109
cystine
0
0
penicillamine disulfide
1.0
(6.30 f 0.15) x lo9 (4.71 f 0.49) x IO9
0.0
4.0
6.0
(8.91 f 1.18) x lo8 10.0
8.0
12.0
14.0
PH Figure 2. (a) pH dependence of the scavenging rate constants for hydrated electron reaction with mercaptoethanol (H) in comparison to previous literature data; (0)ref 23, (-) ref 24. Dashed line corresponds to calculated values of eq 8 with values given in Table 3. (b) Rate constants for cysteamine (W) reaction in comparison to previous literature; (0)ref 14 values. Dashed line corresponds to calculated values of eq 14 using values of Table 3.
TABLE 2: Reacting Species Radii Used for Rate Constant Ionic Strength Corrections species radius 8, hydrated electron mercaptoethanol cysteamine cysteine penicillamine cystine penicillamine disulfide
2.75 2.48 2.58 2.74 2.79 4.99 5.24
to be minor. Throughout this study, ionic strength corrections were only made when there was a single sulfhydryl (>95%) species present. The measured pH dependent rate constants shown in Figure 2a are described by reactions 4 and 5, and a general expression for the total measured rate constant can be derived as2,
in Figure 2a. Both are seen to be significantly higher than the values determined in this study. Cysteamine. The effects of multiple ionizations on the hydrated electron reaction rate constants were investigated using the sulfhydryl cysteamine. For this sulfhydryl, there are three species present in the pH range of study, formed by the two ionization equilibria
-
-
+ -SCH,CH,NH,+ H+ + -SCH,CH,NH2
HSCH2CH2NH3+ H+ -SCH,CH2NH3+
(9) (10)
with known pKa values of 8.60 and 10.75.17 Ionic strength corrections, using eq 7, were performed for only the pH 13 rate constant, using the calculated cysteamine reaction radius of 2.58 8, (Table 2). The individual rate constants obtained over the pH range of study are given in Table 1 and also shown in Figure 2b. The total measured pH-dependent values are described by the three reactions
-
+ HSCH2CH2NH3+ eaq-+ -SCH2CH2NH,+ eaq- + -SCH,CH,NH,
eaq-
products
(11)
products
(12)
products
(13)
and by analogy with eq 8, the general expression for the total measured rate constant is where K6 corresponds to the mercaptoethanol deprotonation equilibrium constant. By fitting eq 8 to the data of Figure 2a, with a fixed value of K6 = 3.16 x mol dm-3,'7 gives the limiting rate constants k4 = (8.34 f 0.11) x lo9 dm3 mol-' s-I and k5 = (1.75 f 0.04) x lo7 dm3 mol-' s-' (Table 3). The calculated rate constant values using this model are shown as the dashed line in Figure 2a and are seen to be in excellent agreement with the experimental data. There have been two previous measurements for hydrated electron reaction with m e r c a p t ~ e t h a n o l , and ~ ~ . ~these ~ rate constants are also shown
krx, =
+ k12K9[pl + ki,K9K10 K9K,0 + K9[&1 + [&I2
k11[&12
(14)
Fitting this equation to the data shown in Figure 2b and fixing K9 and Klo at their known pKa valuesi7 gives the limiting rate constants k11 = (1.53 f 0.05) x lolo, k12 = (1.73 f 0.11) x lo9, and kl3 = (4.11 f 0.30) x lo8 dm3 mol-' SKI(Table 3). The predicted rate constant values using this equation are shown as the dashed line in Figure 2b; again excellent agreement with the experimental data is observed.
J. Phys. Chem., Vol. 99, No. 38, 1995 13973
Rate Constant Determinations of Sulfhydryl Species
f 0.14) x lo9, (1.53 f 0.14) x lo9, and (4.06 f 0.91) x lo7 dm3 mol-' s-I, respectively, for the 0, -1, and -2 charged 1.2 0 species. The calculated rate constants using the above values v) 0 in eq 14 are shown as the dashed line in Figure 3a. 1.0 ...... ..... &...=.* 0 Penicillamine. Rate constant measurements for hydrated o , 8 :electron reaction with the cysteine analogue penicillamine were m E also performed in this study. Ionic strength corrections based 4%.0 '0 0.6 V on eq 7 used the calculated reaction radius of 2.79 A. The C 0.4 !! 'Y individual rate constants are given in Table 1 and shown in Y Figure 3b in comparison to the two previous determinations for this ~ u l f h y d r y l . ~Although ~'~ the single value obtained at pH 6S4 is in very good agreement, the other values are much higher than the measurements of this study. I I I I I I The fitting of eq 14 to these data, using the known pKa values 0 1.0 of 7.90 and 10.4217 gives the dashed line shown in Figure 3b, 0.8 corresponding to the rate constants (5.06 f 0.04) x lo9, (4.12 -0 f 0.53) x lo8, and -4 x lo6 dm3 mol-' s-I, respectively, for E 0.6 the neutral, single, and double negatively charged species. For ...... m the last rate constant, which was the slowest value observed E 0.4 for any of the studied sulfhydryls, the fitting procedure gave C an error larger than the calculated value, and thus only an Y E 0.2 approximate value can be given. +*.... .*. ....*. .., 0 z ...).... Despite the lack of absolute agreement of the hydrated 0 0.0 electron rate constant measurements of this study with literature 4.0 6.0 8.0 10.0 12.0 14.0 determinations, the obtained values generally corroborated the trends observed previou~ly.'~The measured rate constants PH generally followed the total charge on the sulfhydryl, with the Figure 3. (a) pH-dependent rate constants for hydrated electron reaction highest rate constant measured for the positively charged form with cysteine (m) in comparison to literature data; (0)ref 14, (A) ref of cysteamine, 1.53 x 1Olo dm3 mol-' s-l. This value was 25, (v)ref 26, and (0)ref 27. Dashed line corresponds to calculated about a factor of 2 higher than for the neutral sulfhydryl rate values of eq 14 using values given in Table 3. (b) Rate constants for penicillamine (m) reaction in comparison to the previous literature constants, which were consistent for mercaptoethanol, cysteine, values; (A) ref 4,and (0)ref 14. Dashed line corresponds to calculated and penicillamine at (5-8) x lo9 dm3 mol-' s-l. The values of eq 14 using values of Table 3. equivalent, uncharged, cysteamine rate constant was approximately an order of magnitude lower, however, this can be There has been one previous determination of this rate explained by the deprotonation of its sulfhydryl group. This constant,I4 where in this study, a full pH dependence of the decrease was also observed in the measured values for the single reaction rate constant was also measured. These data, which negatively charged values for cysteine and penicillamine, and were not corrected for any ionic strength effects, are also shown is indicative of hydrated electron reaction being dominant at in Figure 2b. Reasonable agreement between the two deterthis group. minations is seen for pH values above 8.0; however, at lower The rate constant measured for the deprotonated sulfur form pH's, the literature values are much higher than observed in of mercaptoethanol was much lower than for the other sulfhythis study. These literature values, however, do not follow the dryls, being in better agreement with the doubly ionized forms limiting behavior corresponding to only reaction via eq 11, of cysteine and penicillamine, again an order of magnitude lower expected at lower pH values. than their single negatively charged forms. The above comCysteine. There have been several previous determinations of the hydrated electron rate constant with ~ y s t e i n e , ' ~ and . ~ ~ - ~ ~parisons show that all three factors, deprotonation at the sulfur atom, deprotonation of the amine group and the overall charge these values are summarized in Figure 3a. There is a large of the molecule, have major effects on the hydrated electron scatter observed for the values below pH 8. The single, fully reaction rate constant with sulfhydryls. The effects of multiple pH-dependent study performedI4 again does not give the deprotonations causing double negatively charged species are expected limiting behavior expected at low pH, where reaction seen to be additive. with the fully protonated form only would be expected. The Debye relationship'* for the dependence of diffusionThis sulfhydryl parallels cysteamine, as there are three controlled rate constants on the charges of the reacting species different species reacting over the pH range of study, with two predicts a decrease in rate constant by a factor of ca. 6 when known pKa values of 8.33 and 10.78.'' Eq 7 was used to the hydrated electron reacts with a species that has a change in perform the ionic strength corrections for the pH 12 and 13 charge of -2 units. The much larger changes observed for these data, using a calculated cysteine reaction radius of 2.74 A (Table sulfhydryls indicates that the rate of hydrated electron reaction 2). The importance of these ionic strength corrections was is mainly governed by the nature of the site of attack. evident at basic pH's for this sulfhydryl, where the observed The deprotonation of the sulfur atom directly correlates with rate constants decreased up to pH 11.0 but then increased at a very large decrease in the hydrated electron reaction rate higher pH. The individual rate constants of this study are given constant. The effect of deprotonation of the amino group in in Table 1, and shown in Figure 3a. In general the values of these compounds is not as consistent. For cysteamine there is this study are seen to be lower than the literature data, although very little difference in rate constant found for the second in excellent agreement with the data of bra am^,^^ and again deprotonation, while for cysteine and penicillamine, order of follow a sigmoid curve. magnitude decreases are observed. These analyses, however, Applying eq 14 to this data, with the equilibrium constants are complicated by the compounding effect of the deprotonated again fixed at their known values, gives the rate constants (8.53 I
1.4
0
r
*
T
r
I
I
I
(a)
Mezyk
13974 J. Phys. Chem., Vol. 99, No. 38, 1995
-
1.6
v)
7-
L
c7 -
E
1.2
E 0.8
U 0 x'
0.4
E
1
5
z
0
7
o
F1
o
o
0
......r....r......x . 0 u*.
1
0.8
A
i
0
A
*....z... "'it....
m
x'
A
1
i
0.4
A *-.,
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0.2
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The products from hydrated electron reaction with cystine have been established as the thiyl radical and the simple ~ulfhydryl,~ formed either directly at acidic conditions or by proton addition at higher pH values. Penicillamine Disulfide. As a comparison to the cystine system, reaction rate constants for the hydrated electron with penicillamine disulfide were also measured. This disulfide also has two pKa values over the pH range of this study, 7.60 and 8.80," with a calculated reaction radius of 5.24 A (determined as for cystine). The individual corrected rate constants are given in Table 1 and shown in Figure 4b in comparison to the two previous determinations for this ~ u l f h y d r y l . ~ . ~ The rate constants of this study are in very good agreement with the data of Purdie et al. above pH 9.0: where these previous values were corrected for ionic strength. Their rate constants at lower pH's are a little higher but in much better agreement than the values of the other study.3 By fitting eq 14 to the corrected data of this study, fixing the two equilibrium constants at their literature values," rate constants for the individual species can again be obtained. Values of (6.30 f 0.15) x lo9, (4.71 f 0.49) x lo9 and (8.91 f 1.18) x los dm3 mol-' s-I were calculated for the neutral, single, and double negatively charged disulfide, respectively (Table 3). The hydrated electron reaction with penicillamine disulfide does not produce thiyl radicals: instead the formation of the short-lived sulfenium radical is believed to break a C-S bond to produce the tertiary carbon centered radical. The rate constant changes observed for these two disulfides over the pH range of study are far smaller than seen for the single sulfhydryls. The decrease in the rate constants for penicillamine disulfide from pH 5.0 (neutral) to pH 13.0 (-2) is of the same magnitude as expected from the Debye equation predictions,ls which implies that only the total charge of this molecule needs to be considered. The larger decrease deter/ mined for cystine over this range indicates that the presence of the electron-withdrawing -NH3+ group influences the disulfide S - S bond of cystine, but not the more sterically hindered penicillamine disulfide. This effect has been previously observed for hydrated electron reaction with S-methyl~ysteine,'~ where deprotonation of the amino group reduced the observed rate constant by a similar factor. These hydrated electron reaction rate constants for the different ionic forms of both these disulfides further indicates that the most important consideration for these sulfhydryls is the ionization state of the sulfur atom.
Conclusion The systematic investigation of hydrated electron reaction with selected sulfhydryls and disulfides has been performed over a wide range of pH. Utilizing ionization constants from literature data, specific rate constants for the individual charged species of mercaptoethanol, cysteamine, cysteine, penicillamine, cystine, and penicillamine disulfide have been determined from the measured values. The rate constants have shown that the decrease in the hydrated electron reaction rate constant due to the amino group deprotonation is small and that the major effect is due to the deprotonation of the sulfur atom in these compounds.
Acknowledgment. I would like to thank Dr. David Armstrong for useful discussions throughout the course of these experiments. References and Notes (1) Bacq, Z. M. Chemical Protection against Ionizing Radiation; Thomas, C. C., Ed.; Springfield, IL, 1965.
Rate Constant Determinations of Sulfhydryl Species (2) Hollaender, A.; Doherty, D. G. Radiation Damage and Sulfhydryl Compounds; International Atomic Energy Agency: Vienna, 1969. (3) Hoffman, M. Z.; Hayon, E. J. Am. Chem. Soc. 1972, 94, 7950. (4) Purdie, J. W.; Gillis, H. A.; Klassen, N. V. Can. J. Chem. 1973, 51, 3132. (5) Asmus, K.-D., In Radioprotectors and Anticarcinogens, Nygaard, 0. F.; Simic, M. G., Eds.; Academic Press: New York, 1984, p 23. (6) Mezyk, S. P. Chem. Phys. Lett. 1995, 89, 325. (7) Zhao, R.; Lind, J.; Merenyi, G.; Eriksen, T. E. J. Am. Chem. Soc. 1994, 116, 12010. (8) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ret Data 1988, 17, 513 and references therein. (9) Mockel, H.; BonifaEiC, M.; Asmus, K.-D. J. Phys. Chem. 1974, 78, 282. (10) BonifaEiC, M.: Schafer, K.; Mockel, H.; Asmus, K.-D. J. Phys. Chem. 1975, 79, 1496. (11) Gilbert, B. C.; Laue, H. A.; Norman, R. 0. C.; Sealy, R. C. J. Chem. Soc., Perkin Trans. 1975, 2, 892. (12) BonifaEiC, M.; Asmus, K.-D. lnt. J. Radiat. Biol. 1984, 46, 35. (13) BonifaEiC, M.; Asmus, K.-D. J. Phys. Chem. 1976, 80, 2426. (14) Hoffman, M. Z.; Hayon, E. J. Phys. Chem. 1973, 77, 990. (15) Janata, E.; Schuler, R. H. J. Phys. Chem. 1982, 86, 2078. (16) Ebbesen, T. W. Radiat. Phys. Chem. 1989, 34, 619. (17) Fasman, G . D., Ed. Handbook of Biochemistry and Molecular Biology, Physical and Chemical Data, 3rd ed.; CRC Press Inc.: Boston, 1976; Vol. 1.
J. Phys. Chem., Vol. 99, No. 38, 1995 13975 (18) Weston, R. E. Jr.; Schwarz, H. A. Chemical Kinetics; Prentice Hall: Englewood Cliffs, NJ, 1972. (19) Lide, D. R., Ed. Handbook of Chemistry and Physics, 71st ed.: CRC Press Inc.: Boston, 1990. (20) Mesmer, R. E.; Baes, C. F. Jr.; Sweeton, F. H. lnorg. Chem. 1972, 11, 537. (21) Hart, E. J.; Anbar, M. The Hydrated Electron, Wiley-Interscience: New York, 1970. (22) Mezyk, S. P. Radiat. Res., accepted. (23) Karmann, W.; Granzow, A,; Meissner, G.; Henglein, A. lnt. J. Radiat. Phys. Chem. 1969, 1, 395. (24) Jayson, G. G.; Stirling, D. A,; Swallow, A. J. lnt. J. Radiat. Biol. 1971, 19, 143. (25) Braams, R. Radiat. Res. 1966, 27, 319. (26) Samahy, A. E.; White, H. L.; Trumbore, C. N. J. Am. Chem. Soc. 1964, 86, 3177. (27) Armstrong, D. A,; Wilkening, V. G. Can. J. Chem. 1964,42,2631. (28) Hayon, E. Nature 1972, 238, 76. (29) Hart, E. J.; Gordon, S.; Thomas, J. K. J. Phys. Chem. 1964, 68, 1271. (30) Weast, R. C., Ed. CRC Handbook of Chemistry and Physics, 1st student edition; CRC Press Inc.: Boca Raton, FL, 1991; p C-703. JW5 10047
