J. Phys. Chem. 1996, 100, 14751-14761
14751
One-Electron Oxidation Reactions of Some Purine and Pyrimidine Bases in Aqueous Solutions. Electrochemical and Pulse Radiolysis Studies Moshe Faraggi,*,† Federico Broitman,† Jeffrey B. Trent,‡ and Michael H. Klapper‡ Department of Chemistry, Nuclear Research Centre-NegeV, Beer-SheVa 84190, Israel, and Biological Chemistry DiVision, Department of Chemistry, The Ohio State UniVersity, Columbus, Ohio 43210 ReceiVed: February 27, 1996; In Final Form: May 21, 1996X
The reduction potentials of some purine and pyrimidine bases and the guanine nucleoside and nucleotide at pH values between 7 and 13 were investigated using the techniques of cyclic voltammetry, differential pulse polarography, and pulse radiolysis. The results clearly show that the lowest reduction potentials, in volts vs NHE, at pH 7 are those of xanthine, 0.88 V, and 1-methylguanine, 1.06 V (NHE). The extrapolated value of guanine is ca. 1.0 V. We also studied the one-electron oxidation reaction of the azide radical with the above compounds. We monitored their transient absorption spectra and determined their formation secondorder rate constants. The decay kinetics of these radicals was followed. Radicals derived from the bases decayed via a radical-radical mechanism with a second-order rate constant. However, the guanine nucleosides and nucleotide radicals have shown at all pHs two consecutive processes (first order followed by a second order). The first-order reaction is pH dependent. Formation of a new transient was observed at pH g9 for guanosine, pH g11 for 2′-deoxyguanosine, and at pH 13 for 5′-GMP. The observed new transient spectra were similar to that observed for the oxidized guanine radical. Therefore, we suggest that in these oxidized guanine nucleosides and nucleotide the oxidized guanine radical has been released. As previously suggested our results imply that the radiation-induced base release from nucleosides in alkaline solution is due to the preferred reaction of the nucleoside sugar with O•- Using our results, those of Steenken et al. (J. Am. Chem. Soc. 1992, 114, 4701-4709), and Symons model (Symons, M. C. R. J. Chem. Soc., Faraday Trans. 1, 1987, 83, 1-11) we suggest a proton assisted mechanism for a double-strand break in DNA.
Introduction The interaction of ionizing radiation with living cells leads to their inactivation. This inactivation is believed to be mainly due to DNA lesions partly produced by direct effect where the radiation is absorbed by the DNA molecule itself.1 Since ionizing radiation absorption is not specific to special residue(s) of the molecule, chemical transformation (radical cation formation and solvated electron) can occur, in principle, on any constituent of DNA (nucleobases, sugars, or phosphate). DNA lesions are partly generated also by indirect effect where the radiation is absorbed by the solvent (water) in close proximity to the DNA molecule to form the oxidizing OH• radicals and the reducing hydrated electrons (eaq•-).1a The OH• radical is a powerful oxidant that often does not react via an outer-sphere electron transfer mechanism but via a process that is best described by an addition/elimination mechanism.1 Therefore, the OH• radical can react with many organic hydrogens on the macromolecule. The solvated electron is a powerful reducing species and reacts rapidly with all nucleobases.1 Although the direct and indirect processes are of different nature, there is growing evidence that these processes may lead to the formation of similar DNA lesions.2 Both processes form highly reactive free radicals and through pulse radiolysis experiments of DNA constituents we acquired some knowledge of their properties.1 However, the chemical reactions that follow these primary radicals are still not clear. We believe that one of the reasons for that is due to the nonspecificity of both processes and the different reactions that may follow their initial reaction with DNA.3 If subsequent reactions lead to DNA lesions, then one * To whom correspondence should be addressed. † Nuclear Research Centre-Negev. ‡ The Ohio State University. X Abstract published in AdVance ACS Abstracts, August 1, 1996.
S0022-3654(96)00590-4 CCC: $12.00
may ask what is the nature of these reactions (electron, proton, and hydrogen transfers) and what governs them (reduction potentials, pKa). Recently these problems have attracted considerable attention. These new directions of research include notions developed recently in studies of electron transfer in proteins.3,4 ESR studies of irradiated solid DNA and its constituents at 77 K have suggested the formation of one electron-deficient (a radical cation) and one electron-rich (a radical anion) center on the bases but not on the sugar or phosphate.5-8 There is a general agreement that the positive charge migrates to and is localized in guanine.5-8 Earlier ESR studies have suggested thymine as a final location for the negative charge.5 However, recent studies have questioned this assignment and proposed cytosine instead.6-8 These migration processes change the random nature of the initial events in DNA to a site-specific location of both charges.9 Charge migration processes, either short- or long-range, are governed (among other factors) by the reduction potentials of the donor and acceptor and their pKa. Reduction potentials for the one-electron oxidation reaction of some purine and pyrimidine bases were reported by Jovanovic and Simic.10 However, this study was limited to pH 13 only. Very recently Steenken estimated a value of 1.17 V (NHE) for guanine at pH 7.11 The oxidations of these bases by OH• (primary radical), SO4•-, and Br2•- (secondary radicals) were reported in the literature.1,10,12-19 The first two radicals are strong oxidizing species, with a midpoint potential at pH 7 (E′m) of 1.9 and 2.4 V, respectively.20 These radicals may be unspecific, creating more than one radical of a given nucleobase. The use of Br2•- as oxidant (E′m ) 1.63 V versus NHE) should be more specific.10,13,18 Release of unaltered bases from irradiated nucleosides, nucleotides, and DNA in aqueous solution is a well-known © 1996 American Chemical Society
14752 J. Phys. Chem., Vol. 100, No. 35, 1996 phenomenon.1a,21-23 It was assumed, from product analysis studies, that this is due to OH• radical attack on the sugar moiety, followed by decomposition of the sugar radicals.1a This assumption was based on Dizdaroglu et al. findings that in thymidine the yield of thymine in neutral solution is low.21 This result, the authors argue, agrees with the fact that in neutral solution OH• radicals react mostly with the base and only a small fraction reacts with the sugar. Fujita, however, using OH• and other secondary oxidizing radicals (Br2•-, Cl2•-, (CNS)2•-, and SO4•-), has shown that in alkaline solutions of various nucleosides, base release is high and could be a major process.22 Recently Scholes et al. have suggested that the results of Fujita are due to the O•- radical anion that preferentially reacts with the sugar moiety of the nucleosides.23 We have preferred to study the kinetics of the oxidation of the nucleobases using N3• radical (E′m ) 1.31 V), a milder and a genuine one-electron oxidizing agent.24-26 The azide radical is produced via
OH• + N3- f N3• + OH-
(k ) 1.2 × 1010 M-1 s-1)
and the dependence of its second-order rate constants on pH was reported.24 Moreover, the use of an oxidant with low reduction potential, and therefore with lesser probability to create more than one transient species, is a better approach to the question involving internal electron transfer processes. This paper presents results on two important features of electron transfer reactions, namely, one-electron reduction potentials and oxidation rate constants. Since the difference between the electrochemical potentials of two redox centers is an important factor governing the rate of electron transfer,27 it is essential either to know or to be able to determine the individual potentials of the two reduction centers. Pulse radiolysis is a particularly appropriate technique to generate and study unstable radicals, and the determination of their reduction potentials is well grounded on the procedure of equilibrating the radical under investigation with a reference compound.28 Recently, Alfassi et al. and Harriman have shown that by using cyclic voltammetry reduction potentials of couples such as N3•/N3-, Trp•/TrpH, and TyrO•/TyrOH can be measured in aqueous solutions.25,29 We, in aqueous solution and at different pHs, have measured reduction potentials of peptides containing these amino acids and some of their related compounds using both pulse radiolysis and the electrochemical methods of cyclic voltammetry (CV) and differential pulse polarography (DPP).28c,30 We were pleasantly surprised to observe that the electrochemically determined reduction potentials were in good agreement with those determined by pulse radiolysis. Since CV and DPP are simpler, faster, and cheaper techniques than pulse radiolysis, they are more useful for the determination of reduction potentials for radicals provided they could be compared with some results obtained by pulse radiolysis. These electrochemical studies are based on Saveant et al’s. findings that although cyclic voltammetry cannot be used to determine reduction potentials of unstable radicals, it can be successfully used when the radical undergoes rapid dimerization to form an electrode inactive species.31 Experimental Section Purine and pyrimidine bases and their nucleosides and nucleotides were purchased from Sigma (St. Louis, Mo.), Aldrich (Milwaukee, WI), and Fluka (Buchs, Switzerland). Cytosine was recrystallized. All other compounds were used without further purification. Water was obtained from a
Faraggi et al. Millipore Q apparatus. All other chemicals were of analytical grade and were used as received from standard sources. The electrochemical experiments, cyclic voltammetry (CV) and differential pulse polarography (DPP), were performed using Princeton Applied Research (PAR, Princeton, NJ) equipment. These were a Model 173D potentiostat/galvanostat, a Model 179 digital coulometer, and a Model 175 universal programmer for the CV experiments and a Model 174A polarographic analyzer for the DPP experiments. These experiments were conducted in a three-electrode glass cell (Metrohm), with a highly polished glassy carbon working electrode (Metrohm) of 0.2-cm exposed area, platinum wire as a counter electrode, and a calomel electrode as a reference electrode. Before every run the working electrode was washed with hot nitric acid, rinsed with water, polished by successive abrasions (Buehler Micropolish A and B), polishing cloth, rinsed with water, and air dried. Nonetheless, the best and reproducible results were obtained with brand new electrodes. The observed potentials were vs a saturated calomel electrode (SCE). These potentials were converted to those vs a normal hydrogen electrode (NHE) by the addition of 0.242 V. All the experiments were performed in deaerated solutions (argon) at 25 ( 1 °C. Solutions contained 1-10 mM of the base, 0.1 M KCl, and, unless otherwise stated, 20 mM of the appropriate buffer (pH 7, pH 8, and pH 12 phosphate, pH 9 borate, and pH 10 and pH 11 carbonate), and 0.1 M NaOH for pH 13. Pulse radiolytic experiments were carried out using the LINAC facility at The Ohio State University.32 We initiated radical reactions with a short pulse (ca. 100 ns) of high-energy electrons (ca. 3.5 MeV from a Varian linear accelerator) introduced into a solution containing 5 mM of the appropriate buffer (at the following pH values: 13, 11, 9, and 7) and 0.05 M NaN3. The solution was deaerated with N2O treated to remove contaminating oxygen.33 To ensure pseudo-first-order formation kinetics, the solute to N3• concentration ratio was kept below 5%. N3• concentration was normally kept between 0.5 and 2 µM,34 while that of the nucleobase was between 20 µM and 10 mM. The pseudo-first-order rate constants were extracted from the optical density (OD) vs time profiles with a nonlinear regression analysis for first-order (exponential), two consecutive exponential, and parallel first- and second-order processes. The pseudo-first-order rate constants were plotted against the solute concentration to obtain the second-order rate constant from the best fit linear slope. Determination of the decay kinetics of the oxidized nucleobase was extracted from similar profiles and nonlinear regression analysis including also the analysis of a single second-order process. Spectra were investigated in solutions containing at least 1 mM of the nucleobase between 325 and 750 nm. Extinction coefficients were determined from the OD values at the end of the formation reaction. Results and Discussion Electrochemical Studies. Reduction Potentials. Typical cyclic voltammetric current-potential curves for the B/B+ couples (B for base) were obtained with thoroughly degassed solutions. As an example we present in Figure 1 the results with thymine at pH 8. As expected and similarly to Harriman,29 the oxidation product(s) were unstable during the time scale of the run (seconds), therefore, they survive only during the forward scan with no peak(s) observed during the reverse scan. From the position of the peak and its maximum potential value (Ep), the reduction potential (in fact the midpoint potential, Em) of these one-electron oxidation waves could be calculated. However, these peaks were not always well resolved and were
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J. Phys. Chem., Vol. 100, No. 35, 1996 14753
Figure 1. Cyclic voltammetry (CV) of 5 mM thymine. The deaerated solutions (argon) contained 0.1 M NaCl and 20 mM phosphate buffer at pH 8.0. Scan rates from top to bottom were 0.2, 0.1, 0.05, 0.02, and 0.01 V s-1. Inset: CV of 5 mM thymine with the same solution conditions at pH 6.0 and scan rate of 0.02 V s-1.
Figure 2. Differential pulse polarography (DPP) of 5 mM thymine. The deaerated solutions (argon) contained 0.1 M NaCl and 20 mM phosphate buffer at pH 8.0. Scan rates: top, 0.002 V s-1; bottom, 0.005 V s-1.
dependent on the scan rate (Figure 1 inset). This made them difficult to define with some precision. A much better defined oxidation peak was observed by the DPP method (Figure 2). Since the values of the peak potentials (Ep) observed by CV and DPP were similar for equal scan rates (Figure 3), we used eq I developed by Nicholson for
Ox + e- f Red
(1)
2Red f Products
(2)
to calculate Em from both the CV and DPP experiments:35
Ep ) Em - 0.9(RT/nF) + (RT/3nF) ln(2kCRT/3nFν) (I) or
Ep ) Em - 0.9(RT/nF) + (RT/3nF) ln(2kCRT/3nF) - (RT/3nF) ln(ν) with n ) 1 and at 25 °C, RT/nF ) 0.0591 V,
Ep ) Em - 0.053 + 0.020 log(0.04kC) - 0.020 log(ν) (II) where C is the solute initial concentration, ν is the potential scan rate, and 2k is the second-order rate constant for the bimolecular radical-radical (recombination or dismutation) reaction (reaction 2). Experiments were carried out at several concentrations (1, 3, 5, 7, and 10 mM) and showed the normal linear ip (peak current) vs concentration and ip vs the square root of the scan rate (Figure 3 insets) relationships. The nucleobase diffusion coefficient, D (in cm2 s-1), can be calculated from the slope of the plot ip vs (ν)1/2, eq III.36
ip ) (2.69 × 105)n3/2AD1/2C(ν)1/2
(III)
where n is the number of electrons involved in the redox process (n ) 1) and A is the working electrode area in cm2. In all compounds studies we observed a linear relationship between ip and ν1/2. The diffusion coefficients of these compounds at all pH values varied between 2 × 10-6 cm2 s-1 and 4 × 10-6 cm2 s-1. We varied the scan rates from 200 to 5 mV/s in the CV experiments and 5 and 2 mV/s in the DPP experiments. Bimolecular second-order rate constants of the oxidized radical were obtained from pulse radiolysis experiments where the initial radical was generated from N3• (see later). The second-
Figure 3. Dependence of E(p) on scan rates observed by the CV and DPP methods in thymine at pH 10. Other solutions conditions were those described in Figures 1 and 2. Inset (left): i(p) (peak current) vs thymine concentration at pH 10. Inset (right): Dependence of the peak current i(p) vs the square root of the scan rate. Solutions conditions were those described in Figure 1. The individual i(p) values were obtained from CV profiles such as those in Figure 1. Scan rates were from 200 mV/s to 5 mV/s.
order rate constants (2k) for the decay of the oxidized radicals of all compounds and pH values varied between 1 × 108 and 4 × 108 M-1 s-1. These results are similar to those found in the literature.11-13 This range of rate constants introduces a variation of (0.014 V in the determined reduction potential and is within our experimental standard deviation. Reduction potentials (Em) were extracted from the Ep data by nonlinear squares fit of eq II with the program MINSQ (MicroMath Scientific Software, Salt Lake City). It should be noted that eq II defines a slope value of 0.02 for the typical curve presented in Figure 3. The experimental slopes found were all within less than 10% of that value. Each of the Em collected into Table 1 was calculated from the average of two independent determinations at each concentration. All the values of Table 1 have an estimated standard deviation of (20 mV or better. At the concentration range studied some of the nucleobases were not soluble at lower pH. We therefore estimated the reduction potential at pH 7 by extrapolation taking into account its pKa and assuming a linear increase of 60 mV per pH unit. Nucleosides and nucleotides have been also tested (guanosine, 2′-deoxyguanosine, guanosine 5′-monophosphate, adenosine, adenosine 5′-monophosphate, cytidine 5′-monophosphate, and uridine). The results for the guanine-containing compounds allowed the determination of their one-electron reduction potentials. The values obtained were about 0.15 V higher in 2′-deoxyguanosine (0.91, 1.02, and 1.17 V at pH 13, 10, and 7, respectively) and about 0.25 V higher in 5′-GMP (1.03 and 1.25
14754 J. Phys. Chem., Vol. 100, No. 35, 1996
Faraggi et al.
TABLE 1: One-Electron Reduction Potentials of Some Purine and Pyrimidine Bases in Neutral and Alkaline Solutionsa pH compound
(pKa)b
adenine (4.15; 9.8) guanine (9.2; 12.2) 1-methylguanine (3.2; 10.4) xanthine (1.2; 7.5; 11.0)g 1-methylxanthine (1.3; 7.9; 11.8)g 3-methylxanthine (0.8; 8.5; 11.5)g 7-methylxanthine (0.8; 8.4; Å10.5)g 8-methylxanthine 1,3-dimethylxanthine (0.7; 8.5)g 1,7-dimethylxanthine (0.5; 8.6)g hypoxanthine (8.9; 12.1) cytosine (4.55; 12.2) uracil (9.5; >13) thymine (9.9; >13)
7 1.32 1.04e
8
9
10
11
12
13
13c
1.26 1.21 1.16 1.14 1.16 1.15 0.87c 0.89f 0.83 0.83 0.83 0.75c
1.06
1.00 0.93 0.85 0.81 0.82 0.82
0.93
0.88 0.83 0.77 0.72 0.70 0.71 0.63 0.79 0.74 0.70 0.70 0.83 0.70 0.71 1.20 1.15 1.09 1.04 0.98 0.92 0.84 0.77 0.76 0.71 0.73 1.15 1.13 1.14 1.15 1.13
1.22
1.15 1.09 1.05 1.08 1.06 1.08
1.16h
1.05 1.00 0.99 0.98 0.99 0.86
1.44h
1.39 1.33 1.28 1.23 1.18 0.93
1.34e
1.16i 1.13 1.14 1.00
1.29
1.23 1.16 1.12 1.07 1.07 1.05 0.91
Em values in volts vs. NHE and are within (20 mV. b pKas were taken from: Data for Biochemical Research, 2nd ed.; Dawson, R. M. C., Elliott, D. C., Elliott, W. H., Jones, K. M., Eds.; Calendon Press: Oxford, U.K. c Jovanovic and Simic (ref 10) corrected values. The reference compound (tryptophan) redox potential at pH 13 was changed from 0.57 to 0.69 V. See: Butler, J.; Land, E. J.; Prutz, W. A.; Swallow, A. J. J. Chem. Soc., Chem. Commun. 1986, 348-349 and refs 11 and 28c. d At pH 7.5. e Estimatend value obtained by extapolation, assuming a linear increase of 60 mV per pH unit and the pKa value of the base. f pH 10.2. Due to limited solubility at this pH, the concentration of guanine was 0.1 mM. g pKa values for xanthine and methylated xanthines were taken from: Bergmann et al. J. Chem. Soc. (C), 1971, 1676. h Estimatend value obtained by extapolation, assuming a linear increase of 60 mV per pH unit. i At pH 11.6. a
V at pH 13 and 7, respectively). However, from studies on the behavior of the corresponding oxidized radicals of the nucleosides and nucleotide (see later) we have concluded that their chemistry could not be described by reaction 2 (radical-radical second-order reaction). As will be shown, the mechanism that describes the decay of these oxidized species was a fast firstorder (monomolecular) reaction followed by a radical-radical second-order reaction. Based on these kinetics findings and the time dependent spectra, we are suggesting that these reactions describe: (i) the formation of the oxidized guanine radical and a free sugar or its molecular fragments; (ii) radicalradical recombination reaction of the guanine radical. If this mechanism is correct then we suggest that the use of the Nicholson equation is possible. Attempts to measure reduction potentials of the other nucleosides and nucleotides failed since their potentials (NHE) were above 1.5 V (the maximum limit of our glassy carbon working electrode). Our results at pH 13 are in good agreement with those of Jovanovic and Simic (after the correction made by Steenken et al.;11 see note c in Table 1) only for guanine and xanthine, compounds with relatively low reduction potentials.10 For all other compounds our values are significantly higher (Table 1). The study of Jovanovic and Simic was performed using the technique of pulse radiolysis.10 They used p-methoxyphenol, Trolox, and tryptophan with
reduction potentials at pH 13 of respectively of 0.4 (corrected to 0.44 V), 0.19, and 0.57 V (corrected to 0.69 V). We have shown that a reliable estimate of an unknown reduction potential may not be obtained when the reference and unknown potential difference is not too large.28c The use of tryptophan for the determination of guanine fulfills this requirement. However, for the other purine and pyrimidine, even if we include guanine as a reference, the reduction potential of the reference compound is too low and, as we previously have shown, may lead to low reduction potential values of the unknown.28c Table 1 clearly shows that the reduction potentials of the purine bases are lower than those of the pyrimidine bases. It also shows that between the two purine bases of DNA guanine has the lowest reduction potential. However, because of solubility problems with guanine, its reduction potential at pH 7 is an extrapolated value. We believe that since the reduction potentials of 1-methylguanine (soluble at all pHs) are similar to those of guanine, the guanine extrapolated value could be estimated as reasonable. In adenine the Em decreases ca. 60 mV per pH unit between pH 7 and pH 10. At pHs above 10, Em appears to be constant. This behavior, similar to that observed with tyrosine, is expected from a compound with a pK of 9.7 of the N(9) hydrogen.30 In cytosine, with a pKa of 12.2 of the N(1) hydrogen, a decrease of the Em of ca. 60 mV per pH unit is observed between pH 9 and pH 12. At pH 8 and below cytosine at millimolar concentrations precipitates. For thymine the variation of the Em values with pH fits its pKa values. These pyrimidine bases of DNA have E′m (midpoint potentials at pH 7) above 1.2 V. An explanation for the relatively low reduction potential in guanine could be the presence of the carbonyl group and the formation of an enolate group upon deprotonation. Therefore, and in agreement with the literature, we suggest that of the four nucleobases of DNA guanine are the best electron donors. This conclusion implies that upon random oxidation of a nucleobase (radical cation formation) in DNA, as thermodynamics suggests, the final site for electron deficiency could be guanine, provided there are sufficiently rapid short-range or long-range charge transfer process(es). As expected, based on our previous suggestion, the replacement of an amino group by a second carbonyl group in the purine base (guanine to xanthine) decreases the reduction potential. The elimination of the amino group (electron donating group) from the purine base (guanine to hypoxanthine) increases the reduction potential. Reduction potentials of the monomethylated guanine and xanthines seem to follow their nitrogen hydrogens pKa values. NMR studies on guanine show that protons are attached to the 1- and 9-nitrogen atoms.37 In xanthine these protons are on the 1-, 3-, and 7-nitrogen atoms.37 Studies on deprotonation of xanthine and its 1- and 7-methylated derivatives suggested that the first proton to dissociate is from N(3) (pKa 7.4). Further dissociation occurs at N(7) (pKa 11.2) and at N(1) (pKa > 13).37 Deprotonation of guanine is at N(1) (pKa 9.2) and at N(9) (pKa 12.2).38 Reduction potentials of xanthine and its 1- and 8methyl derivatives are similar. The similarity of the 8-methyl derivative is clear since it does not involve dissociable hydrogens. The similarity of the 1-methyl derivative could be due to the similarity of their monoanion tautomeric structures (Scheme 1). According to Elguero et al. the monoanion of both will be a mixture of form I and II derived by N(3)H and N(7)H ionization.37 In other words, the N(1)H in xanthine is so tightly bound to the nitrogen that it does not matter if we substitute it with a methyl group. Assuming that in guanine the monoanion
Oxidation of Purine and Pyrimidine Bases
J. Phys. Chem., Vol. 100, No. 35, 1996 14755
SCHEME 1: Structures of the Xanthine Monoanion
takes the form of structure II, then the similarity in the reduction potentials of guanine and its 1-methyl derivative is also explained. We suggest that the increase of the reduction potentials in the 3-methyl and 7-methyl substituted xanthines are due to the different proton loss sequence. For the 3-methyl, loss is first at N(7) followed by N(1). For the 7-methyl, it is at N(3) followed by N(1). With the same argument the increase of the reduction potentials of the 1,3- and 1,7-methyl substitute derivatives of xanthine could be explained. If this interpretation is valid, then the measured value of 1.06 V for 1-methyl guanine at pH 7 (where guanine itself could not be measured because of solubility) should not be far from that of guanine. Indeed, this value at pH 7 is only slightly larger than that obtained by extrapolation. It should be noted that the potentials of the two bases at high pH are closely similar. Recently, Steenken et al. have measured the reduction potentials of the T/T•- and the C/C•- couples (T, thymidine; C, cytidine).11 They found that the two have similar values.39 We measured the reduction potentials using the same electrochemical methods but with a dropping mercury as a working electrode. Guanosine, adenosine, thymidine, and cytidine were studied. Preliminary results suggest that at pH 7 the reduction potentials (in volts vs. NHE) are -1.2 V for both thymidine and cytidine, -1.4 V for adenosine, and g-1.5 V for guanosine (the voltage limit of the working electrode). The values obtained for the pyrimidines are in good agreement with those of Steenken et al.11,39 Pulse Radiolysis Studies. Purine and Pyrimidine Bases. We turn now to our second objective, namely, the reactivity of the nucleobases toward the oxidant N3•. The use of this radical or any secondary radical in pulse radiolysis studies required close attention to the conditions of the experiment including not only the concentration of the solutes but also the pH of the solution studied. An example of a mistake in pulse radiolysis of dilute aqueous solutions was an attempt to study of the reactions of Cl2•- in neutral or alkaline solution containing Cl-.22 In contrast to acid media where Cl- reacts very rapidly with OH• radical (4.3 × 109 M-1 s-1), it reacts very poorly in neutral or alkaline solutions (
