J. Phys. Chem. 1996, 100, 8295-8301
8295
LINAC/LASER Determination of the Absolute Rate Constant for Thiyl and Hydroxyl Radical Reaction with Sulfhydryls in Aqueous Solution: Mercaptoethanol, Cysteamine, and N-Acetyl-L-cysteine† Stephen P. Mezyk* Radiation Laboratory, UniVersity of Notre Dame, Notre Dame, Indiana 46556 ReceiVed: October 16, 1995; In Final Form: January 2, 1996X
The technique of combined pulse radiolysis/laser photolysis has been used to investigate the disulfide radical anion formation reactions for mercaptoethanol, cysteamine, and N-acetyl-L-cysteine over the pH range 7.013.0. The photolysis of the transient radical anions, formed by the one electron oxidation of the sulfhydryl, perturbs the disulfide radical anion/thiyl radical equilibrium, allowing rate constants for thiyl radical reaction with the parent sulfhydryl to be uniquely determined from the absorption bleach and subsequent recovery. These pH-dependent values were combined with measured disulfide radical anion equilibrium constants to calculate first-order disulfide radical anion dissociation rate constants. By computer modeling of established mechanisms for the disulfide radical anion growths, rate constants for the reaction of hydroxyl radicals with individual sulfhydryl species were calculated. These values are contrasted with previously reported values determined by using competition kinetics.
Introduction The important role of sulfides (RSH) and disulfides (RSSR) in radiation protection has been known for many years.1,2 The redox properties of these organic sulfur compounds, and especially their radicals, are of considerable interest for the understanding of many processes in biological and related model systems. The discovery of in-situ thiyl (RS•) radicals in biological material3,4 and the establishment of their versatile chemistry as intermediates in the repair reaction of carbon centered radicals5 have prompted much attention being paid to the reactions of such sulfur-centered radicals. These RS• radicals are stronger oxidizing species than the simple aliphatic peroxyl radicals, ROO•,6-8 and have been demonstrated to be of similar reactivity to their oxyanalogues, RO•, toward many biologically available electron donors.9 Thiyl radicals have even been found to initiate lipid peroxidation.10,11 These findings indicate that sulfur-centered radicals play a very important role in freeradical-mediated tissue damage. The reaction of the hydroxyl radical with free sulfhydryls, in neutral or slightly basic solution, results in the formation of the reducing radical disulfide anion, (RSSR•-), via the general mechanism12
RSH + OH- h H2O + RS•
OH + RSH f H2O + RS•
•
(1) (2)
OH + RS- f OH- + RS•
(3)
RS• + RS- h RSSR•-
(4)
RS• + RSH h RSSR•- + H+ (5) The determination of specific rate constants for these reactions from only the disulfide transient absorption is difficult, as all * Present address: Research Chemistry Branch, AECL-Whiteshell Laboratories, Pinawa, Manitoba, R0E 1L0, Canada. † The research described herein was supported by the Office of Basic Energy Sciences of the Department of Energy. This is contribution No. NDRL 3883 from the Notre Dame Radiation Laboratory. X Abstract published in AdVance ACS Abstracts, March 15, 1996.
S0022-3654(95)03067-X CCC: $12.00
the formation rate constants are of the same order of magnitude,13 and the observed growths represent an approach to equilibrium. This difficulty is reflected in the scatter in the available literature data.13 In several recent investigations14-16 the technique of combined LINAC radiolysis/LASER photolysis was used with computer modeling to determine the individual rate constants for these reactions for the sulfhydryls cysteine and glutathione. These experiments involved using electron pulse radiolysis to generate the radical disulfide, and then photolyzing this anion using the laser pulse14,16
RSSR•- + hν f RS• + RS-
(-4)
The photolysis of this transient causes a decrease in the RSSR•absorption, which subsequently recovers by first-order kinetics. The bleach and recovery of the transient absorption is due only to the perturbation of the RS•/RSSR•- equilibrium and thus the exponential recovery gives the unique (pseudo-first-order) rate constant for the disulfide anion formation reaction. By repeating these experiments at different pHs and sulfhydryl concentrations, specific rate constants for the disulfide formation by reactions 4 and 5 were obtained. These rate constants were then combined with pH-dependent disulfide radical anion equilibrium constants, independently determined from concentration-dependent integrated RSSR•yield measurements,15 to give rate constants for the disulfide dissociation reaction. Knowing both component rate constants in these two equilibria, the observed pH-dependent growth profiles of the RSSR•- transient could then be fitted to the general reaction mechanism to determine individual rate constants for reactions 2 and 3. The hydroxyl radical reaction rate constants obtained in this manner for both cysteine15 and glutathione16 were much lower than previously reported literature values,13 which had been determined by using competition kinetics. Most of these competition studies used SCN- as a reference reaction, which complicates the elucidation of individual rate constants as the two intermediates formed, RSSR•- and (SCN)2•-, have overlapping spectra. This means that accurate absorption coefficients © 1996 American Chemical Society
8296 J. Phys. Chem., Vol. 100, No. 20, 1996 are required to calculate the hydroxyl radical rate constant with the sulfhydryls. Furthermore, it has been reported that an interfering reaction occurs in the irradiated glutathione/SCNsystem,17 to give another product that absorbs in the same region. This latter study also used tetramethyl urea as a competitor and obtained a much lower rate constant for hydroxyl radical reaction. Based on these potential complications, the competition-kinetics literature values must be treated with some caution. Another interesting aspect observed for hydroxyl radical reaction with cysteine and glutathione in the LINAC/LASER studies was that the overall rate constant decreased at higher pHs. This behavior was opposite to that determined in previous measurements for glutathione,18,19 although only sporadic pHdependent investigations had been performed, and inconsistent with analogous studies of hydroxyl radical reaction with simple amines, which showed an order of magnitude increase in rate constant upon deprotonation of the amino group.20 In order to further investigate the generality of the behavior observed for cysteine and glutathione, this study utilized the LINAC/LASER technique to determine rate constants for the reactions of hydroxyl radicals with other simple thiols in aqueous solution. The sulfhydryls used were mercaptoethanol, cysteamine (2-aminoethanethiol), and N-acetyl-L-cysteine, due to their relatively intense disulfide anion transient absorption across a large pH range and because the photolysis of this transient recovered fully, implying that there was no additional chemistry such as photoionization of RS- occurring on the experimental time scales. The determined hydroxyl radical rate constants have been compared to the results of previous determinations, and by utilizing known ionization constants, the constituent rate constants for hydroxyl radical reaction with individual sulfhydryl species have also been calculated. Experimental Section The experimental setup used for these measurements has been detailed in several previous publications,21-23 hence only a brief description shall be given here. The pulse radiolysis system consists of a 7-MeV LINAC, which delivered 10-20-ns electron pulses to the sample cell. An excimer laser (Lumonics Hyper EX-400), which could be used to pump a dye laser (Lumonics, Hyperdye-300), had its excitation pulse synchronized to photolyze the irradiated solution after a fixed time period. The pulse radiolysis traces in all cases were normalized for the radiolysis dose; the laser intensity did not vary significantly during an experiment so its normalization was not required. Most data in this study were collected by using the output of the excimer laser at 308 nm, of typical power ca. 140 mJ/pulse. Some experiments used the dye laser output over the range 400430 nm; however, the significantly lower energy (ca. 7 mJ/ pulse) meant that the signals obtained were much smaller and correspondingly more noisy. The few rate constants obtained using the dye laser photolysis were found to be in good agreement with those measured with the 308-nm excimer excitation. The sulfhydryls used in this study (Aldrich) were of the highest purity available and used as received. Care was taken to avoid oxidation of these compounds; all solutions were prepared immediately before irradiation by dissolving known amounts of the sulfhydryl into N2O saturated Millipore Milli-Q filtered water which had been buffered to the appropriate pH by using Baker Analyzed monobasic phosphate or borax at a concentration of 2.0 × 10-2 mol dm-3. Exact pH values were obtained by additions of small amounts of NaOH (Fisher, ACS) or HClO4 (Aldrich, ACS 70%) and using a pH meter. All
Mezyk
Figure 1. (a) Transient absorption change at 410 nm observed in the electron pulse radiolysis (9) of N2O-saturated mercaptoethanol (9.61 × 10-4 mol dm-3) solution at pH 10.0. Second trace (o) shows this transient which has subsequently been photolyzed at 308 nm (130 mJ), resulting in a reduction of absorption followed by recovery. (b) Difference between the above two absorption profiles, inverted only for display purposes. Solid line is an exponential decay fit to the difference data, corresponding to a pseudo-first-order rate constant of (4.76 ( 0.28) × 106 s-1.
solutions were flowed through the irradiation cell at a sufficient rate to ensure that a fresh sample was irradiated each time. During the irradiation process, the solution vessels were bubbled with only the minimum amount of N2O required to prevent air ingress, to prevent loss of the volatile sulfhydryl. Typically 50-80 pulses were averaged to obtain a single trace. Dosimetry was performed using N2O saturated SCN- solutions, (10-2 mol dm-3, λ ) 472 nm, G ) 4.92 × 104).24 All experiments were performed at room temperature (22 ( 2 °C) using initial hydroxyl radical concentrations of 2-5 × 10-6 mol dm-3. Results and Discussion Mercaptoethanol. This compound was chosen for initial study, as mercaptoethanol, HSCH2CH2OH, is one of the simplest sulfhydryls with only a single pKa of 9.50.25 In neutral or basic solution, the mercaptoethanol radical disulfide anion, MSSM•-, has a characteristic absorbance with a maximum at 410-420 nm;26 at lower pHs this transient is replaced by a very weak absorbance due to the thiyl radical MS•. Under the experimental conditions of this study, the best signals were obtained at 410 nm; at this wavelength the interference from the MS• radical absorption was negligibly small. The optical transient obtained from only the pulse radiolysis of a mercaptoethanol solution (9.61 × 10-4 mol dm-3 at pH 10.0) is shown in Figure 1a. The additional photolysis of this transient causes a bleach in the absorbance, and the subtraction of these two traces gives the recovery trace shown in Figure 1b (data inverted only for presentation purposes). This recovery trace shows first-order (decay) kinetics. By performing this experiment at different total (MS ) MSH + MS-) mercaptoethanol concentrations the scavenging curve shown in Figure
2a was obtained. The straight line fit to this data gives the rate constant for mercaptoethanol disulfide anion formation,
MS• + MS f MSSM•-
HSCH2CH2OH h H+ + -SCH2CH2OH
k6 )
k7[H+] + k8K9
[H+] + K9 (6)
at pH 10.0 as k6 ) (3.80 ( 0.27) × 109 dm3 mol-1 s-1. The rate constants obtained in this manner over the pH range 8.012.5 are shown in Figure 2b, the error bars shown for these points correspond to one standard deviation obtained from the straight line fits. These individual values are also listed in Table 1. No comparison values could be found in the literature. Over the pH range of this study, the charge on this sulfhydryl changes from 0 to -1, therefore at intermediate pH values, the overall reaction is by a combination of the two limiting reactions
MS• + MSH f MSSM•- + H+ (7)
MS• + MS- f MSSM•(8)
where these two species exist in equilibrium
(9)
On the basis of these reactions, a general expression for the overall measured rate constant, k6, can be derived as
(10)
By fitting eq 10 to the data of Figure 2b, with the value of K9 fixed at its pKa, the limiting rate constants k7 ) (1.22 ( 0.08) × 109 and k8 ) (4.85 ( 0.19) × 109 dm3 mol-1 s-1 were obtained. The predicted rate constant values using this model 13,
8.1 ( 1.3 7.7 ( 1.2 7.3 ( 1.5 5.65 ( 0.80 5.2 ( 1.3 3.74 ( 0.90 3.46 ( 0.33 3.11 ( 0.24 3.00 ( 0.21 3.57 ( 0.25 3.74 ( 0.24 3.74 ( 0.45 3.78 ( 0.36
s-1
8.2 ( 1.1 8.33 ( 0.78 6.65 ( 0.80 5.83 ( 0.48 4.37 ( 0.47 2.93 ( 0.37 3.12 ( 0.46 3.02 ( 0.39 2.51 ( 0.28
s
10-9k14, 3 -1 -1
dm mol
Calculated from fitted limiting values (see also ref 14).
151.9 ( 13.8 162.0 ( 13.3 178.7 ( 18.5 273.3 ( 21.7 417.5 ( 36.6 805.8 ( 71.6 1099 ( 26 1452 ( 51 1600 ( 84 1349 ( 40 1318 ( 18 1292 ( 39 1284 ( 72
1.23 ( 0.09a 1.25 ( 0.09a 1.31 ( 0.14 1.54 ( 0.10a 2.15 ( 0.34 3.01 ( 0.46 3.80 ( 0.27 4.52 ( 0.19 4.79 ( 0.09 4.82 ( 0.19 4.93 ( 0.25 4.83 ( 0.44 4.85 ( 0.20a
7.0 7.5 8.0 8.5 9.0 9.5 10.0 10.5 11.0 11.5 12.0 12.5 13.0
a
K11, dm3 mol-1
10 6, dm3 mol-1 s-1
10-6k
mercaptoethanol
pH
-9k
0.45 ( 0.02 0.72 ( 0.02 1.34 ( 0.03 2.26 ( 0.05 3.04 ( 0.06 3.43 ( 0.07 3.57 ( 0.07 3.60 ( 0.07 3.64 ( 0.07 3.65 ( 0.07 3.65 ( 0.07 3.65 ( 0.07 3.65 ( 0.07
s
a 17, -1 -1
dm3 mol
10-9k 208.9 ( 13.1 436.3 ( 37.2 1575 ( 92 3300 ( 45 2892 ( 143 1815 ( 110 1157 ( 50 1060 ( 88 1005 ( 42 1058 ( 57 1130 ( 47 1244 ( 65 1309 ( 34
K20, dm3 mol-1
s
2.15 ( 0.25 1.65 ( 0.20 0.85 ( 0.07 0.69 ( 0.03 1.05 ( 0.08 1.89 ( 0.16 3.09 ( 0.20 3.40 ( 0.38 3.62 ( 0.23 3.45 ( 0.27 3.23 ( 0.20 2.93 ( 0.22 2.79 ( 0.13
10-6k21, -1
cysteamine
3.45 ( 0.39 3.51 ( 0.28 3.50 ( 0.34 3.09 ( 0.41 2.92 ( 0.16 2.33 ( 0.19 1.90 ( 0.31 1.61 ( 0.09 1.35 ( 0.10
3.46 ( 0.60
22, dm3 mol-1 s-1
10-9k
Figure 2. (a) Total mercaptoethanol concentration dependence of measured rate constants for radical disulfide anion formation in N2Osaturated solution at pH 10.0. Solid line corresponds to fitted second order rate constant, with value k6 ) (3.80 ( 0.27) × 109 dm3 mol-1 s-1. (b) pH dependence of calculated radical disulfide anion formation rate constants of this study. Error bars are one standard deviation obtained from the linear fits in (a). Individual values are given in Table 1. Solid line corresponds to predicted values using eq 10 and limiting thiyl radical reaction rate constants as given in the text.
TABLE 1: Summary of pH-Dependent Rate and Equilibrium Constants Obtained in This Study
1.31 ( 0.19a 1.21 ( 0.11 1.40 ( 0.14 1.41 ( 0.08 1.84 ( 0.07 3.00 ( 0.33 4.57 ( 0.05 7.12 ( 0.33 10.3 ( 0.6 10.3 ( 0.5 10.6 ( 0.5
29, dm3 mol-1 s-1
10-8k
44.8 ( 12.3 64.6 ( 26.8 93.0 ( 15.2 159.6 ( 34.1 265.7 ( 27.9 431.1 ( 64.2 639.9 ( 21.2 1065 ( 30 1156 ( 36 1015 ( 64 958.5 ( 60.6
K32, dm3 mol-1
2.9 ( 1.3 1.87 ( 0.95 1.51 ( 0.40 0.88 ( 0.24 0.69 ( 0.10 0.70 ( 0.18 0.71 ( 0.03 0.67 ( 0.05 0.89 ( 0.08 1.02 ( 0.11 1.11 ( 0.12
10-6k33, s-1
N-acetyl-L-cysteine
3.34 ( 0.55 3.16 ( 0.38 3.22 ( 0.33 3.05 ( 0.19 2.66 ( 0.14 2.20 ( 0.11 1.84 ( 0.17 1.67 ( 0.13 1.60 ( 0.08
10-9k34, dm3 mol-1 s-1
Determination of Absolute Rate Constant J. Phys. Chem., Vol. 100, No. 20, 1996 8297
8298 J. Phys. Chem., Vol. 100, No. 20, 1996
Mezyk
Figure 4. pH dependence of calculated rate constants for the first order dissociation of mercaptoethanol disulfide radical anion, reaction 13, from measured equilibrium values of Figure 3b and measured (9) or calculated (0) scavenging rate constants of Figure 2b. Error bars correspond to one standard deviation as calculated from the uncertainty in these two measured values. Individual values are given in Table 1.
the mercaptoethanol radical disulfide anion, Figure 3. (a) Equilibrium constant determination from integrated transient yields for mercaptoethanol radical disulfide anion formation in N2O-saturated solution at pH 10.0. From fitted linear slope and intercept values (see text), an equilibrium constant of K11 ) 1099.0 ( 26.3 dm3 mol-1 is calculated. (b) pH dependence of the calculated equilibrium constants for mercaptoethanol radical disulfide anion formation. Individual values are given in Table 1.
are shown as the solid line in Figure 2b and are seen to be in very good agreement with the experimental data. The overall equilibrium constants for mercaptoethanol disulfide radical anion formation
MS• + MS h MSSM•-
(11)
were determined by using integrated yield measurements. These measurements were performed at low initial radical concentrations in order to minimize the disulfide radical decay. The dependence of total mercaptoethanol concentration, [MS], on the total yield (G) of the formed radical anion is given by27,28
G∞ 1 )1+ G K11[MS]
(12)
where G∞ is the limiting absorbance. Thus a plot of (G)-1 against [MS]-1 should give a straight line of slope (G∞K11)-1, intercept (G∞)-1, and ratio K11 ) intercept/slope. Typical data obtained at pH 10.0 are shown in Figure 3a, where from the calculated slope of (2.31 ( 0.03) × 10-8 dm3 mol-1 and intercept (2.54 ( 0.02) × 10-5 values, K11 ) 1099.0 ( 26.3 dm3 mol-1 is obtained. These measurements were repeated over the pH range 7.013.0, with the individual equilibrium constants obtained summarized in Table 1 and shown in Figure 3b. The pH dependence of these values is similar to that observed for cysteine15 and glutathione16 previously, increasing from pH 7.0 to 11.0 (ca. 150 to >1600 dm3 mol-1) and then decreasing to a limiting value at higher pHs. Although in very good agreement with the values determined for cysteine and glutathione, these equilibrium constants are much lower than measured for other, similar, sulfhydryls.12 From the measured, pH-dependent k6 rate and K11 equilibrium constants, the rate constants for the first-order dissociation of
MSSM•- f MS• + MS-
(13)
can readily be calculated. These values are shown in Figure 4 and also listed in Table 1. The obtained pH behavior is complex; showing a minimum value around pH 10.5 and rising with increasing pH to a plateau. A slower rise at lower pHs is also observed. This behavior is consistent with that observed for cysteine15 and glutathione16 previously and also in accord with previous measurements for cystamine and cystine,29 where the pH-dependence of the decay of the disulfide radical anion was directly measured by following the reaction of the hydrated electron with the disulfide. Knowing the individual, pH-dependent rate constants for the component reactions in equilibrium (11), the individual rate constants for the reactions of hydroxyl radicals with mercaptoethanol •
OH + MS f MS•
(14)
can then be determined. By use of the established general reaction mechanism, each mercaptoethanol disulfide radical anion growth was modeled by using the differential equation solving code FACSIMILE30 to optimize the best pseudo-firstorder rate constant for reaction 14. Typical fits to three mercaptoethanol concentrations at pH 10.0 are shown in Figure 5. These fitted rate constants were then plotted against total mercaptoethanol concentration to obtain linear scavenging curves, at pH 10.0 the overall rate constant was k14 ) (4.37 ( 0.47) × 109 dm3 mol-1 s-1. By repeating this analysis over the pH range 8.0-12.0, the values shown in Figure 6a and listed in Table 1 were obtained. This plot again exhibits the characteristic sigmoidal shape obtained previously for cysteine and glutathione,15,16 with the higher rate constants being at lower pHs. Over the pH range 10.5-12.0, the calculated hydroxyl radical reaction rate constants are equivalent within error, which implies that the rate constant for reaction of hydroxyl and oxide radical with mercaptoethanol is essentially the same. Thus, the composite values shown in Figure 6a are well described by only the reactions
Determination of Absolute Rate Constant •
J. Phys. Chem., Vol. 100, No. 20, 1996 8299
OH + HSCH2CH2OH f H2O + •SCH2CH2OH (15)
•
OH + -SCH2CH2OH f -OH + •SCH2CH2OH (16)
and, by analogy to eq 10, limiting rate constants of k15 ) (8.62 ( 0.49) × 109 and k16 ) (2.67 ( 0.18) × 109 dm3 mol-1 s-1 are calculated. The predicted rate constants by using these values in eq 10 are shown as the solid line in Figure 6a and are seen to be in very good agreement with the experimental data. There have been several previous determinations of the rate constant for hydroxyl radical reaction with mercaptoethanol. By monitoring the formation of MSSM•- at 420 nm, a value of (4.0 ( 0.6) × 109 dm3 mol-1 s-1 at pH 11 has been measured;26 this rate constant is in reasonable agreement with the values of this study. However, attempts to measure the rate constant for hydroxyl radical reaction with the fully protonated form of mercaptoethanol, using SCN- competition kinetics (k ) 1.1 × 1010 dm3 mol-1 s-1),13 gave values of 3.2 × 1010 (pH 7)26 and 1.9 × 1010 dm3 mol-1 s-1 (pH 6.5).31 These values are much greater than those of this study and also higher than another competition kinetics determination, using ferrocyanide as a reference, which gave (6.5 ( 0.7) × 109 dm3 mol-1 s-1.31 Based on the conclusion that thiocyanate was not suitable as a reference for mercaptoethanol, the only comparison rate constant shown in Figure 6a at lower pHs is the ferrocyanide value, which is seen to be slightly lower than the rate constants calculated in this study. Cysteamine. The rate constants for cysteamine thiyl radical, CyS•, reaction with cysteamine (CyS), to give the cysteamine disulfide radical anion
CyS• + CyS f CySSCy•-
Figure 5. Typical fits obtained in the determination of pseudo-firstorder rate constants for hydroxyl radical reaction with mercaptoethanol at pH 9.0. Data shown corresponds to mercaptoethanol concentrations of 2.00 × 10-3 (0), 1.26 × 10-3 (4), and 7.44 × 10-4 (]) mol dm-3. Fitted lines correspond to calculated values, using determined pHdependent rate constants for reactions 6 and 13 (Table 1) and hydroxyl radical rate constants of 8.28 × 106, 6.81 × 106, and 4.48 × 106 s-1, respectively.
(17)
have been previously measured by using the LINAC/LASER technique.14 From the determined limiting values for the reactions
CyS• + CyS- f CySSCy•-
(18)
CyS• + CySH f CySSCy•- + H+
(19)
of k18 ) (1.21 ( 0.04) × 109 and k19 ) (3.39 ( 0.31) × 108 dm3 mol-1 s-1, and the known pKa value of 8.3,25 individual rate constants for the formation of this transient could be calculated for the pH range of this study (see Table 1). Equilibrium constants for cysteamine disulfide radical anion formation
CyS• + Cys h CySSCy•-
(20)
(21)
Figure 6. (a) pH-dependent rate constants of this study, for the reaction of hydroxyl and oxide radicals with mercaptoethanol in aqueous solution, in comparison to previous literature data: ref 26 (4) and ref 31 (0). Individual values are given in Table 1. Error bars are one standard deviation obtained from second-order rate constant determinations (not shown). Solid line corresponds to best fit of eq 10 with limiting rate constants as given in text. (b) pH dependence of hydroxyl and oxide radical reaction with cysteamine in aqueous solution (values in Table 1). Fitted solid line corresponds to eq 28, with limiting rate constants as given in text. (c) pH dependence for hydroxyl and oxide radical reaction with N-acetyl-L-cysteine (see also Table 1). Fitted line corresponds to eq 10 with values as given in text.
All these individual values are listed in Table 1. The observed pH dependence again displays the characteristic shape seen for mercaptoethanol and also is in qualitative agreement with a previous pH-dependent study performed for this sulfhydryl.12 There have been several previous measurements for the k21 rate constant, with values of ca. 8 × 105 (pH 8)32 and 1.3 × 106 s-1 (pH 11.8),12 obtained from the hydrated electron reaction
with the disulfide. The lower pH rate constant is in very good agreement with the value of this study, (8.5 ( 0.7) × 105 s-1; however, at the basic pH, the value of this study, (3.23 ( 0.20) × 106 s-1, is considerably higher. Although this discrepancy cannot be explained at this time, it is important to note that the calculated values of this study correspond to the total rate of loss of the disulfide radical anion under the exact experimental
were determined over the pH range 6.0-13.0 as for mercaptoethanol. The concentration-dependent yields at the transient maximum of 410 nm were used for these measurements, where the absorbance due to the CyS• radical was again seen to be negligible. These equilibrium and rate constants were again used to calculate the pH dependence of the equivalent firstorder disulfide anion dissociation rate constants,
CySSCy•- f CyS• + CyS-
8300 J. Phys. Chem., Vol. 100, No. 20, 1996
Mezyk
conditions used to determine hydroxyl radical reaction rate constants, and moreover, it has previously been demonstrated that the hydroxyl radical rate constants are not very sensitive to the dissociation rate constants used in the fitting procedure.16 On the basis of the K20 equilibrium and k17/k21 rate constant values of this study, individual cysteamine disulfide radical anion transient absorptions were fitted as for mercaptoethanol, to obtain the pseudo-first-order hydroxyl radical reaction rate constants for the reaction •
OH + CyS f CyS•
(22)
From the variation in cysteamine concentration, scavenging rate constants over the pH range 7.0-12.0 were obtained, these values are shown in Figure 6b and listed in Table 1. This plot again shows the characteristic pH-dependence obtained for mercaptoethanol, with decreasing rate constants at higher pHs. For this sulfhydryl there are three species present in the pH range of study, formed by the two ionization equilibria
HSCH2CH2NH3+ h H+ + -SCH2CH2NH3+
(23)
SCH2CH2NH3+ h H+ + -SCH2CH2NH2
(24)
-
with known pKa values of 8.6 and 10.75.25 The calculated total pH-dependent hydroxyl radical rate constants are therefore described by the three reactions •
OH + HSCH2CH2NH3+ f H2O + •SCH2CH2NH3+ (25)
•
OH + -SCH2CH2NH3+ f OH- + •SCH2CH2NH3+ (26) •
OH + -SCH2CH2NH2 f OH- + •SCH2CH2NH2
(27)
with oxide radical reaction being treated as equivalent to hydroxyl, as for mercaptoethanol. By analogy with eq 10, the general expression for the total rate constant for this sulfhydryl is
krxn )
k25[H+]2 + k26K23[H+] + k27K23K24 K23K24 + K23[H+] + [H+]2
(28)
Fitting this equation to the data shown in Figure 6b, fixing K23 and K24 at their known pKa values, gives the limiting rate constants k25 ) (3.66 ( 0.31) × 109, k26 ) (3.14 ( 0.16) × 109, and k27 ) (1.29 ( 0.08) × 109 dm3 mol-1 s-1. The predicted rate constant values using these values in eq 28 are shown as the solid line in Figure 6b; again excellent agreement with the experimental data is observed. The hydroxyl radical rate constants of this study are significantly lower than the only previous measurements for this sulfhydryl,32 where over the pH range 6.5-9, a value of 8.5 × 109 dm3 mol-1 s-1 was obtained by the SCN- competition kinetics method. N-Acetyl-L-cysteine. The rate constants for N-acetylcysteine thiyl radical, nCysS•, reaction with N-acetylcysteine (nCysS), to give the disulfide radical anion
nCysS• + nCysS f nCysSSnCys•-
(29)
were determined in this study as for mercaptoethanol. The individual pH-dependent rate constants, calculated from the total N-acetylcysteine, [nCysS], concentration dependence on the laser bleach recovery, are given in Table 1. The pH-dependence of these values follows the expected sigmoidal curve, consistent with its single pKa value. However, the fitting of eq 10 to these
values gave a calculated pKa of 10.27 ( 0.06, higher than the literature value of 9.52.25 Using the calculated pKa value of this study, the rate constants for the two limiting reactions
nCysS• + nCysS- f nCysSSnCys•-
(30)
nCysS• + nCysSH f nCysSSnCys•- + H+
(31)
were determined as k30 ) (1.10 ( 0.03) × 109 and k31 ) (1.31 ( 0.19) × 108 dm3 mol-1 s-1, respectively. The k30 value of this study is in excellent agreement with a previous determination at pH 11.2, of k ) 1.1 × 109 dm3 mol-1 s-1.12 Equilibrium constants for N-acetylcysteine disulfide radical anion formation
nCysS• + nCysS h nCysSSnCys•-
(32)
were determined over the pH range 7.0-12.0 as before. These equilibrium and rate constants were again used to calculate the pH dependence of the equivalent first-order disulfide anion dissociation rate constants,
nCysSSnCys•- f nCysS• + nCysS-
(33)
All these individual values are listed in Table 1. The observed pH dependence again displays the characteristic shape seen for the other two sulfhydryls of this study. Based on the K32 equilibrium and k29/k33 rate constant values of this study, individual N-acetylcysteine disulfide radical anion transient absorptions were again fitted to the general formation mechanism, to obtain the pseudo-first-order hydroxyl radical reaction rate constants for the reaction •
OH + nCysS f nCysS•
(34)
The N-acetylcysteine concentration dependence of these values gave the scavenging rate constants shown in Figure 6c and listed in Table 1. This plot also shows the characteristic pHdependence obtained previously, with decreasing rate constants at higher pHs. As this sulfhydryl has only a single pKa value, the calculated total pH-dependent hydroxyl radical rate constants are described by the two limiting reactions
OH + HSCH2CH(NHCOCH3)CO2- f
•
H2O + •SCH2CH(NHCOCH3)CO2- (35) OH + -SCH2CH(NHCOCH3)CO2- f
•
OH- + •SCH2CH(NHCOCH3)CO2- (36) again equating the hydroxyl and oxide radical reaction rate constants under basic conditions. By fitting eq 10 to this data, using the fitted pKa value of this study, the rate constants k35 ) (3.66 ( 0.31) × 109 and k36 ) (3.14 ( 0.16) × 109 dm3 mol-1 s-1 are obtained. The predicted rate constant values by using these values in eq 10 are shown as the solid line in Figure 6c, again excellent agreement with the experimental data is observed. The hydroxyl radical rate constants of this study are significantly lower than the previous measurements for this sulfhydryl; values of 1.0 × 1010 dm3 mol-1 s-1 at pH 5.2 and 8.4 were obtained by using pulse radiolysis techniques33 (unfortunately no further experimental details were given), and a rate constant
Determination of Absolute Rate Constant
J. Phys. Chem., Vol. 100, No. 20, 1996 8301
of 1.36 × 1010 dm3 mol-1 s-1 at pH 7.034 was obtained by using the SCN- competition method.
•
Summary
being determined as k30 ) (1.10 ( 0.03) × 109, k31 ) (1.31 ( 0.19) × 108, k35 ) (3.66 ( 0.31) × 109, and k36 ) (3.14 ( 0.16) × 109 dm3 mol-1 s-1, respectively. The calculated hydroxyl radical reaction rate constants for these three sulfhydryls are considerably lower than previously reported competition kinetics measurements using SCN-, and all exhibit the same pH dependence, with faster values at lower pHs, in agreement with the behavior seen for cysteine and glutathione previously.
The techniques of electron pulse radiolysis and laser photolysis have been used to determine individual rate constants for the mercaptoethanol disulfide radical anion formation reactions
HOCH2CH2S• + HOCH2CH2SH f HOCH2CH2S∴SCH2CH2OH- + H+ (7) HOCH2CH2S∴SCH2CH2OH- (8) as k7 ) (1.22 ( 0.08) × 109 and k8 ) (4.85 ( 0.19) × 109 dm3 mol-1 s-1, respectively. The values have been combined with pH-dependent equilibrium constant measurements, determined from mercaptoethanol disulfide radical anion transient intensities, to allow computation of rate constants for the hydroxyl radical reactions
OH + HOCH2CH2SH f H2O + HOCH2CH2S• (15)
•
OH + HOCH2CH2S- f OH- + HOCH2CH2S• (16)
as k15 ) (8.62 ( 0.49) × 109 and k16 ) (2.67 ( 0.18) × 109 dm3 mol-1 s-1. Similar measurements were also performed for cysteamine, where from literature rate constants for the disulfide radical anion formation and measured equilibrium constants, limiting hydroxyl radical reaction rate constants for the three separate species, •
OH + HSCH2CH2NH3+ f H2O + •SCH2CH2NH3+ (25)
•
OH + -SCH2CH2NH3+ f OH- + •SCH2CH2NH3+ (26) •
OH + -SCH2CH2NH2 f OH- + •SCH2CH2NH2
(27)
were calculated as k25 ) (3.66 ( 0.31) × 109, k26 ) (3.14 ( 0.16) × 109, and k27 ) (1.29 ( 0.08) × 109 dm3 mol-1 s-1. These measurements were also repeated for N-acetyl-Lcysteine, with limiting rate constants for the reactions •
SCH2CH(NHCOCH3)CO2- + -
•
SCH2CH(NHCOCH3)CO2- f
O2C(CH3CONH)CHCH2S∴SCH2CH(NHCOCH3)CO2(30)
SCH2CH(NHCOCH3)CO2- + HSCH2CH(NHCOCH3)CO2- f H+ + -
•
OH- + •SCH2CH(NHCOCH3)CO2- (36)
References and Notes
HOCH2CH2S• + HOCH2CH2S- f
•
OH + -SCH2CH(NHCOCH3)CO2- f
O2C(CH3CONH)CHCH2S∴SCH2CH(NHCOCH3)CO2(31)
OH + HSCH2CH(NHCOCH3)CO2- f H2O + •SCH2CH(NHCOCH3)CO2- (35)
(1) Bacq, Z. M. In Chemical Protection against Ionizing Radiation; Thomas, C. C., Ed.; Springfield, IL, 1965. (2) Hollaender, A.; Doherty, D. G. Radiation Damage and Sulfhydryl Compounds; International Atomic Energy Agency: Vienna, 1969. (3) Mason, R. P.; Ramakrishna, Rao, D. N. In Sulphur-Centered ReactiVe Intermediates in Chemistry and Biology; Chatgilialoglu, C., Asmus, K.-D., Eds.; NATO-ASI Series A; Life Sciences; Plenum-Press: New York, 1990; Vol. 197, p 401. (4) Mason, R. P.; Maples, K. R. In Sulphur-Centered ReactiVe Intermediates in Chemistry and Biology; Chatgilialoglu, C., Asmus, K.-D., Eds.; NATO-ASI Series A; Life Sciences; Plenum-Press: New York, 1990; Vol. 197, p 429. (5) Von Sonntag, C. The Chemical Basis of Radiation Biology; Taylor and Francis: London, 1987; p 353. (6) Forni, L. G.; Mo¨nig, J.; Mora-Arellano, V. O.; Willson, R. L. J. Chem. Soc., Perkin Trans. 2 1983, 961. (7) Forni, L. G.; Willson, R. L. Biochem. J. 1986, 240, 897. (8) Forni, L. G.; Willson, R. L. Biochem. J. 1986, 240, 905. (9) Erben-Russ, M.; Michel, C.; Bors, W.; Saran, M. J. Phys. Chem. 1987, 91, 2362. (10) Tieu, M.; Bucher, J. R.; Aust, S. D. Biochem. Biophys. Res. Commun. 1982, 107, 279. (11) Searle, A. J. F.; Willson, R. L. Biochem. J. 1983, 212, 549. (12) Hoffman, M. Z.; Hayon, E. J. Phys. Chem. 1973, 77, 990. (13) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 513 and references therein. (14) Mezyk, S. P. Chem. Phys. Lett. 1995, 235, 89. (15) Mezyk, S. P. Radiat. Res. 1996, 145, 102. (16) Mezyk, S. P. J. Phys. Chem., in press. (17) Liphard, M.; Bothe, E.; Schulte-Frohlinde, D. Int. J. Radiat. Biol. 1990, 58, 589. (18) Eriksen, T. E.; Fransson, G. J. Chem. Soc., Perkin Trans. 2 1988, 1117. (19) Quintiliani, M.; Badiello, R.; Tamba, M. Int. J. Radiat. Biol. 1977, 32, 195. (20) Simic, M.; Neta, P.; Hayon, E. Int. J. Radiat. Phys. Chem. 1971, 3, 309. (21) Janata, E.; Schuler, R. H. J. Phys. Chem. 1982, 86, 2078. (22) Ebbesen, T. W. Radiat. Phys. Chem. 1989, 34, 619. (23) Natarajan, P.; Fessenden, R. W. J. Phys. Chem. 1989, 93, 6095. (24) Buxton, G. V.; Stuart, C. R. J. Chem. Soc., Faraday Trans. 1995, 91, 279. (25) Fasman, G. D., Ed. Handbook of Biochemistry and Molecular Biology, Physical and Chemical Data, 3rd ed.; CRC Press Inc.: Boston, 1976; Vol. 1. (26) Karmann, W.; Granzow, A.; Meissner, G.; Henglein, A. Int. J. Radiat. Phys. Chem. 1969, 1, 395. (27) Baxendale, J. H.; Bevan, P. L. T.; Stott, D. A. Trans. Faraday Soc. 1968, 64, 2389. (28) Baxendale, J. H.; Bevan, P. L. T. J. Chem. Soc. A 1969, 2240. (29) Hoffman, M. Z.; Hayon, E. J. Am. Chem. Soc. 1972, 94, 7950. (30) Curtis, A. R.; Sweetenham, W. P. FACSIMILE/CHECKMAT Users Manual; Harwell Research Report AERE-R 12805; 1988. (31) Jayson, G. G.; Stirling, D. A.; Swallow, A. J. Int. J. Radiat. Biol. 1971, 19, 143. (32) Adams, G. E.; McNaughton, G. S.; Michael, B. D. In Chemistry of Ionization and Excitation; Johnson, G. R. A., Scholes, G., Eds.; Taylor and Francis: London, 1967; p 281. (33) Rougee, M.; Bensasson, R. V.; Land, E. J.; Pariente, R. Photochem. Photobiol. 1988, 47, 485. (34) Aruoma, O. I.; Halliwell, B.; Hoey, B. M.; Butler, J. Free Rad. Biol. Med. 1989, 6, 593.
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