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Pulse radiolysis studies of the reactions of eaq− and ·OH with ClO3− ions in aqueous solution
Radiation Physics and Chemistry 59 (2000) 309±312
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Pulse radiolysis studies of the reactions of eÿ aq and OH with ClOÿ 3 ions in aqueous solution
B.G. Ershov a, 1, M. Kelm b, E. Janata a,* a Hahn-Meitner-Institut Berlin, Glienicker Str. 100, 14109 Berlin, Germany Institut fuÈr Nukleare Entsorgungstechnik, Forschungszentrum Karlsruhe, PO Box 3640, 76021 Karlsruhe, Germany
b
Received 07 October 1999; accepted 18 January 2000
Abstract ÿ The rate constants of the reactions of eÿ aq and OH with ClO3 in aqueous solutions are determined by pulse ÿ radiolysis studies. The hydrated electron reacts with ClO3 with a rate constant of 1.6 105 Mÿ1 sÿ1 to yield the radical anion. Its optical absorption exhibits two bands at 245 nm e 900 Mÿ1 cmÿ1) and at 500 nm ClO2ÿ 3 e 500 Mÿ1 cmÿ1). The reaction of the OH radical with ClOÿ 3 is too slow to measure a rate constant; an upper limit is estimated to be <1 105 Mÿ1 sÿ1. The rate constants for the reactions of the ClO2ÿ 3 radical anion with the OH radical and the Oÿ radical ion are determined to be 1 1010 and 1.2 109 Mÿ1 sÿ1, respectively. 7 2000 Elsevier Science Ltd. All rights reserved.
1. Introduction Recent experiments have shown (Kelm and Bohnert, 2000a) that the main products formed during the gradiolysis of concentrated sodium chloride solutions are hydrogen, oxygen and chlorate ions. These products are produced in proportional to the absorbed dose. Consequently, the mathematical simulation of the radiolytic processes could demonstrate the linear production of these species and their accumulation (Kelm and Bohnert, 2000b). For this simulation, the value of the rate constants of the reactions of eÿ aq and OH with chlorate ions is important. Until now, only an upper limit for these constants (<106 Mÿ1 sÿ1) has been published (Buxton and Subhani, 1972). In the
* Corresponding author. Fax: +49-30-8062-2434. E-mail address: [email protected] (E. Janata). 1 On leave of absence from the Institute of Physical Chemistry, Russian Academy of Sciences, Moscow, Russia.
present paper, we report on pulse radiolysis investigations to determine these rate constants.
2. Experimental The experiments were carried out at the van de Graa accelerator facility ELBENA of the Hahn-Meitner-Institut using 3.8 MeV electrons. The experimental apparatus (Janata, 1992) and the computer software (Janata, 1994) have previously been described. The duration of the electron pulse was set between 3 and 200 ns (Janata and Gutsch, 1998). Doses corresponding to radical concentrations of the order of 10ÿ7 M were used to avoid second-order reactions of intermediates. Dosimetry was done by observing the absorption of the hydrated electron at 700 nm and 100 ns after the impingement of the electron pulse, taking into account a molar absorbance of 1.9 104 Mÿ1 cmÿ1 (Hug, 1981) and a radiation chemical yield of 2.6 electrons per 100 eV absorbed radiation energy (Bux-
0969-806X/00/$ - see front matter 7 2000 Elsevier Science Ltd. All rights reserved. PII: S 0 9 6 9 - 8 0 6 X ( 0 0 ) 0 0 1 9 1 - 2
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B.G. Ershov et al. / Radiation Physics and Chemistry 59 (2000) 309±312
ton et al., 1988). In presenting the data, the observed changes in absorbance are divided by the concentration of the radicals and the thickness of the measuring cell. Thus, the changes are given in terms of molar absorbance in Mÿ1 cmÿ1. The sodium chlorate (NaClO3) was purchased from Merck and was further puri®ed by recrystallization. The solutions were freshly prepared and were purged with argon or nitrous oxide (N2O) prior to irradiation.
3. Results and discussion Hydrated electrons react slowly with chlorate ions in NaClO3 solutions. The measurements were carried out in slightly alkaline solutions (pH 8.5) in order to avoid the fast reaction of the hydrated electron with the hydrogen ion. As the absorption of the hydrated electron decays over several tens of microseconds, a new absorption, consisting of two low intensity bands at 245 and 500 nm, builds up. We ascribe this new absorption to the ClO2ÿ radical anion, which is 3 formed according to the reaction ÿ 2ÿ eÿ aq ClO3 )ClO3 :
1
Similar adducts have been observed as the result of the addition of eÿ aq to certain inorganic anions, for ÿ example to NOÿ 2 (GraÈtzel et al., 1969) or to NO3 (GraÈtzel et al., 1970). The rate constant of reaction (1) is determined by
monitoring the decay of the eÿ aq absorption at 700 nm. Fig. 1 shows the plot of the inverse ®rst half-life time versus the NaClO3 concentration. The inverse half-life time increases with the concentration of NaClO3. If a straight line was drawn through the measuring points, all points would be close to the line, and there would be an oset at the vertical axis. The oset is due to the reaction of the hydrated electron with itself, and with other products of the radiolysis of water, hydroxyl radicals, hydrogen atoms and hydrogen peroxide. All these species compete with ClOÿ 3 for the hydrated electron. As the concentration of ClOÿ 3 increases, the hydrated electron begins to react more and more with the chlorate ion. The eect of ionic strength has to be taken into account for reaction (1) and for the reaction of the hydrated electron with itself. A reaction mechanism including reaction (1) and the competition reactions mentioned above was simulated, using known rate constants and yields for the competition reactions (Buxton et al., 1988). On comparing the result of the simulation with experimental data, a rate constant of 1.6 105 Mÿ1 sÿ1 was obtained for the reaction of the hydrated electron with the chlorate ion. The values obtained by the simulation of the inverse half-life time (open squares) and the amount of hydrated electrons reacting with chlorate ions (open circles) have also been included in Fig. 1. The calculated and the experimental values of the inverse half-life time are in excellent agreement. At low NaClO3 concentrations, the decay of the hydrated electron is mainly due to the competition reactions. At high NaClO3 concentrations,
Fig. 1. Inverse half-life time of the eÿ aq decay at 700 nm versus the NaClO3 concentration. Irradiation dose per pulse corresponds to 4.0 10ÿ7 M hydrated electrons. pH = 8.5. Closed squares: experimental values, open squares: values obtained by simulation (both left-hand scale), and open circles: calculated amount of hydrated electrons reacting with chlorate ions (right-hand scale).
B.G. Ershov et al. / Radiation Physics and Chemistry 59 (2000) 309±312
0.5 M for example, the eÿ aq decay is determined solely by reaction (1), since about 90% of the hydrated electrons react with chlorate ions. In order to estimate the reaction of the hydroxyl radical with the ClOÿ 3 ion, the competition between 2ÿ ClOÿ 3 and CO3 for the hydroxyl radical was studied, i.e., the eect of the addition of ClOÿ 3 on the yield of ÿ the COÿ 3 radical ion. The CO3 radical anion was observed at lmax 600 nm in a 3 10ÿ4 M solution of Na2CO3, saturated with N2O at pH 11. By purging the solution with N2O, the hydrated electrons of the radiolysis of water are converted into hydroxyl radicals; the rate constant being 9 109 Mÿ1 sÿ1 (Janata and Schuler, 1982). The hydroxyl radical reacts with CO2ÿ 3 at a rate constant of 3.9 108 Mÿ1 sÿ1 (Buxton et al., 1988) according to
ÿ ÿ OH CO2ÿ 3 )OH CO3 :
2
The addition of 0.25 M NaClO3 had no eect on the absorption of the COÿ 3 radical ion in the actual experiment. The yield of the COÿ 3 radical ion and its kinetic formation were the same, with and without the addition of ClOÿ 3 ions. From the rate constant of reaction (2) and from the concentrations of Na2CO3 and NaClO3, it follows that the rate constant of the reaction of the OH radical with the ClOÿ 3 ion must be lower than 1 105 Mÿ1 sÿ1. This is in agreement with the results reported earlier (Domae et al., 1994) that the OH radical does not react at all with the ClOÿ 3
311
ion, even in concentrated solutions of about 5 M NaClO3. In solutions purged with Ar, the ClO2ÿ radical 3 anions disappear rapidly in a second-order process, apparently as a result of the reaction with the OH radicals still present
ÿ ÿ OH ClO2ÿ 3 )OH ClO3 :
3
The second-order rate constant of this reaction was estimated on the basis of these decay curves to be 1 1010 Mÿ1 sÿ1. The high reactivity of OH with ClO2ÿ 3 makes it dicult to observe its absorption. At the end of the eÿ aq decay, two bands at 245 and 500 nm were observed, which exhibit low intensities. In alkaline solutions, however, the hydroxyl radical dissociates (pK = 11.9) and forms the Oÿ radical ion, which has a much lower oxidizing ability. The life time of the ClO2ÿ 3 radical anion should, therefore, be much longer at a high pH value, and thus allow a more precise determination of its absorption spectrum. The rate constant of reaction (1) remains unchanged when the pH is increased from 8.5 to 12.5. Curve 1 in Fig. 2 depicts the absorption 12 ms after the electron pulse, measured in a 0.5 M solution of NaClO3 at pH 12.5. This spectrum includes the absorption of the Oÿ radical ion lmax 240 nm, e 240 Mÿ1 cmÿ1) (Hug, 1981) as well as radical anion. The spectrum of the that of the ClO2ÿ 3 ClO2ÿ 3 radical anion (Curve 2 in Fig. 2) was obtained
Fig. 2. Optical absorption spectrum obtained 12 ms after 10 ns electron pulses (curve 1) and after correction for the absorption of the Oÿ radical ion (curve 2). Solution 0.5 M NaClO3, pH = 12.5, Ar. Irradiation dose corresponds to 2.0 10ÿ6 M hydrated electrons.
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B.G. Ershov et al. / Radiation Physics and Chemistry 59 (2000) 309±312
after the correction for the Oÿ radical ion. It exhibits two optical bands at 245 nm e 900 Mÿ1 cmÿ1) and at 500 nm e 500 Mÿ1 cmÿ1). As expected, the radical anion disappears in alkaline solution ClO2ÿ 3 much slower than in neutral ones. The decay follows a second-order law according to
ÿ ÿ Oÿ ClO2ÿ 3 H2 O)2OH ClO3 :
4
The rate constant of this reaction is determined to be 1.2 109 Mÿ1 sÿ1. In summary, the hydrated electron reacts slowly with the ClOÿ 3 ion, resulting in the formation of the ÿ ClO2ÿ 3 radical anion. The reaction of OH with ClO3 is too slow to be measured in neutral or alkaline solutions or does not take place at all. The redox potential of the ClO3 radical has been estimated to be E 0 2:1 V (Stanbury, 1989) in aqueous solution, which is consistent with the conclusion that this potential seems to be between 1.66 and 2.41 V (Domae et al, 1994). On the other hand, the potential of the OH radical at pH = 7 is also about 2.2 V (Wardman, 1989). It follows from this comparison that the oxi dation of the ClOÿ 3 ion by the OH radical is a rather unfavorable process.
References Buxton, G., Greenstock, C.L., Helman, W.P, Ross, A.B., 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solution. J. Phys. Chem. Rev. Data 17 (2). Buxton, G., Subhani, M.S., 1972. Radiation chemistry and photochemistry of oxychlorine ions. J. Chem. Soc. Faraday Trans. 1, 947. Domae, M., Katsumara, Y., Jiang, P.Y., Nagaiashi, R., Hasegawa, C., Ishogura, K., Yoshida, Y., 1994.
Observation of ClO3 radical in aqueous chlorate solution by pulse radiolysis. J. Phys. Chem. 98, 190. GraÈtzel, M., Henglein, A., Lilie, J, Beck, G., 1969. Pulsradiolytische Untersuchungen einiger Elementarprozesse der Oxydation und Reduktion des Nitritions. Ber. Bunsenges. Physik. Chem. 73, 646. GraÈtzel, M., Henglein, A., Taniguchi, S., 1970. Pulsradiolytische Beobachtungen uÈber die Reduction des NOÿ 3 -Ions und uÈber Bildung und Zerfall der persalpetrigen SaÈure in waÈssriger Loesung. Ber. Bunsenges. Physik. Chem. 74, 292. Hug, G.L., 1981. Optical spectra of nonmetallic inorganic transient species in aqueous solution. NSRDS Ð NBS 69, 12. Janata, E., 1982. Instrumentation of kinetic spectroscopy-10. A modular data acquisition system for laser ¯ash photolysis and pulse radiolysis experiments. Radiat. Phys. Chem. 40, 437. Janata, E., 1994. Instrumentation of kinetic spectroscopy-12. Software for data acquisition in kinetic experiments. Radiat. Phys. Chem. 44, 449. Janata, E., Gutsch, W., 1998. Instrumentation of kinetic spectroscopy-15. Part III: Injector pulse generators for electron accelerators. Radiat. Phys. Chem. 51, 65. Janata, E., Schuler, R.H., 1982. Rate constant for scavenging eÿ aq in N2O saturated solutions. J. Phys. Chem. 86, 2078. Kelm, M., Bohnert, E., 2000a. Radiation chemical eects in the near-®eld of a ®nal disposal site. Part I: Radiolytic products formed in concentrated NaCl solutions. Nucl. Technol., 129, 118. Kelm, M., Bohnert, E., 2000b. Radiation chemical eects in the near-®eld of a ®nal disposal site. Part II. Simulation of the radiolytic processes in concentrated NaCl solutions. Nucl. Technol., 129, 123. Stanbury, D.M., 1989. Reduction potentials involving inorganic free radicals in aqueous solution. In: Sykes A.G. (Ed.), Advances in Inorganic Chemistry, vol. 33, p. 69. Wardman, P., 1989. Reduction potentials of one-electron couples involving free radicals in aqueous solution. J. Phys. Chem. Rev. Data 18 (4), 1637±1755.
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Pulse radiolysis studies of the reactions of eÿ aq and OH with ClOÿ 3 ions in aqueous solution
B.G. Ershov a, 1, M. Kelm b, E. Janata a,* a Hahn-Meitner-Institut Berlin, Glienicker Str. 100, 14109 Berlin, Germany Institut fuÈr Nukleare Entsorgungstechnik, Forschungszentrum Karlsruhe, PO Box 3640, 76021 Karlsruhe, Germany
b
Received 07 October 1999; accepted 18 January 2000
Abstract ÿ The rate constants of the reactions of eÿ aq and OH with ClO3 in aqueous solutions are determined by pulse ÿ radiolysis studies. The hydrated electron reacts with ClO3 with a rate constant of 1.6 105 Mÿ1 sÿ1 to yield the radical anion. Its optical absorption exhibits two bands at 245 nm e 900 Mÿ1 cmÿ1) and at 500 nm ClO2ÿ 3 e 500 Mÿ1 cmÿ1). The reaction of the OH radical with ClOÿ 3 is too slow to measure a rate constant; an upper limit is estimated to be <1 105 Mÿ1 sÿ1. The rate constants for the reactions of the ClO2ÿ 3 radical anion with the OH radical and the Oÿ radical ion are determined to be 1 1010 and 1.2 109 Mÿ1 sÿ1, respectively. 7 2000 Elsevier Science Ltd. All rights reserved.
1. Introduction Recent experiments have shown (Kelm and Bohnert, 2000a) that the main products formed during the gradiolysis of concentrated sodium chloride solutions are hydrogen, oxygen and chlorate ions. These products are produced in proportional to the absorbed dose. Consequently, the mathematical simulation of the radiolytic processes could demonstrate the linear production of these species and their accumulation (Kelm and Bohnert, 2000b). For this simulation, the value of the rate constants of the reactions of eÿ aq and OH with chlorate ions is important. Until now, only an upper limit for these constants (<106 Mÿ1 sÿ1) has been published (Buxton and Subhani, 1972). In the
* Corresponding author. Fax: +49-30-8062-2434. E-mail address: [email protected] (E. Janata). 1 On leave of absence from the Institute of Physical Chemistry, Russian Academy of Sciences, Moscow, Russia.
present paper, we report on pulse radiolysis investigations to determine these rate constants.
2. Experimental The experiments were carried out at the van de Graa accelerator facility ELBENA of the Hahn-Meitner-Institut using 3.8 MeV electrons. The experimental apparatus (Janata, 1992) and the computer software (Janata, 1994) have previously been described. The duration of the electron pulse was set between 3 and 200 ns (Janata and Gutsch, 1998). Doses corresponding to radical concentrations of the order of 10ÿ7 M were used to avoid second-order reactions of intermediates. Dosimetry was done by observing the absorption of the hydrated electron at 700 nm and 100 ns after the impingement of the electron pulse, taking into account a molar absorbance of 1.9 104 Mÿ1 cmÿ1 (Hug, 1981) and a radiation chemical yield of 2.6 electrons per 100 eV absorbed radiation energy (Bux-
0969-806X/00/$ - see front matter 7 2000 Elsevier Science Ltd. All rights reserved. PII: S 0 9 6 9 - 8 0 6 X ( 0 0 ) 0 0 1 9 1 - 2
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B.G. Ershov et al. / Radiation Physics and Chemistry 59 (2000) 309±312
ton et al., 1988). In presenting the data, the observed changes in absorbance are divided by the concentration of the radicals and the thickness of the measuring cell. Thus, the changes are given in terms of molar absorbance in Mÿ1 cmÿ1. The sodium chlorate (NaClO3) was purchased from Merck and was further puri®ed by recrystallization. The solutions were freshly prepared and were purged with argon or nitrous oxide (N2O) prior to irradiation.
3. Results and discussion Hydrated electrons react slowly with chlorate ions in NaClO3 solutions. The measurements were carried out in slightly alkaline solutions (pH 8.5) in order to avoid the fast reaction of the hydrated electron with the hydrogen ion. As the absorption of the hydrated electron decays over several tens of microseconds, a new absorption, consisting of two low intensity bands at 245 and 500 nm, builds up. We ascribe this new absorption to the ClO2ÿ radical anion, which is 3 formed according to the reaction ÿ 2ÿ eÿ aq ClO3 )ClO3 :
1
Similar adducts have been observed as the result of the addition of eÿ aq to certain inorganic anions, for ÿ example to NOÿ 2 (GraÈtzel et al., 1969) or to NO3 (GraÈtzel et al., 1970). The rate constant of reaction (1) is determined by
monitoring the decay of the eÿ aq absorption at 700 nm. Fig. 1 shows the plot of the inverse ®rst half-life time versus the NaClO3 concentration. The inverse half-life time increases with the concentration of NaClO3. If a straight line was drawn through the measuring points, all points would be close to the line, and there would be an oset at the vertical axis. The oset is due to the reaction of the hydrated electron with itself, and with other products of the radiolysis of water, hydroxyl radicals, hydrogen atoms and hydrogen peroxide. All these species compete with ClOÿ 3 for the hydrated electron. As the concentration of ClOÿ 3 increases, the hydrated electron begins to react more and more with the chlorate ion. The eect of ionic strength has to be taken into account for reaction (1) and for the reaction of the hydrated electron with itself. A reaction mechanism including reaction (1) and the competition reactions mentioned above was simulated, using known rate constants and yields for the competition reactions (Buxton et al., 1988). On comparing the result of the simulation with experimental data, a rate constant of 1.6 105 Mÿ1 sÿ1 was obtained for the reaction of the hydrated electron with the chlorate ion. The values obtained by the simulation of the inverse half-life time (open squares) and the amount of hydrated electrons reacting with chlorate ions (open circles) have also been included in Fig. 1. The calculated and the experimental values of the inverse half-life time are in excellent agreement. At low NaClO3 concentrations, the decay of the hydrated electron is mainly due to the competition reactions. At high NaClO3 concentrations,
Fig. 1. Inverse half-life time of the eÿ aq decay at 700 nm versus the NaClO3 concentration. Irradiation dose per pulse corresponds to 4.0 10ÿ7 M hydrated electrons. pH = 8.5. Closed squares: experimental values, open squares: values obtained by simulation (both left-hand scale), and open circles: calculated amount of hydrated electrons reacting with chlorate ions (right-hand scale).
B.G. Ershov et al. / Radiation Physics and Chemistry 59 (2000) 309±312
0.5 M for example, the eÿ aq decay is determined solely by reaction (1), since about 90% of the hydrated electrons react with chlorate ions. In order to estimate the reaction of the hydroxyl radical with the ClOÿ 3 ion, the competition between 2ÿ ClOÿ 3 and CO3 for the hydroxyl radical was studied, i.e., the eect of the addition of ClOÿ 3 on the yield of ÿ the COÿ 3 radical ion. The CO3 radical anion was observed at lmax 600 nm in a 3 10ÿ4 M solution of Na2CO3, saturated with N2O at pH 11. By purging the solution with N2O, the hydrated electrons of the radiolysis of water are converted into hydroxyl radicals; the rate constant being 9 109 Mÿ1 sÿ1 (Janata and Schuler, 1982). The hydroxyl radical reacts with CO2ÿ 3 at a rate constant of 3.9 108 Mÿ1 sÿ1 (Buxton et al., 1988) according to
ÿ ÿ OH CO2ÿ 3 )OH CO3 :
2
The addition of 0.25 M NaClO3 had no eect on the absorption of the COÿ 3 radical ion in the actual experiment. The yield of the COÿ 3 radical ion and its kinetic formation were the same, with and without the addition of ClOÿ 3 ions. From the rate constant of reaction (2) and from the concentrations of Na2CO3 and NaClO3, it follows that the rate constant of the reaction of the OH radical with the ClOÿ 3 ion must be lower than 1 105 Mÿ1 sÿ1. This is in agreement with the results reported earlier (Domae et al., 1994) that the OH radical does not react at all with the ClOÿ 3
311
ion, even in concentrated solutions of about 5 M NaClO3. In solutions purged with Ar, the ClO2ÿ radical 3 anions disappear rapidly in a second-order process, apparently as a result of the reaction with the OH radicals still present
ÿ ÿ OH ClO2ÿ 3 )OH ClO3 :
3
The second-order rate constant of this reaction was estimated on the basis of these decay curves to be 1 1010 Mÿ1 sÿ1. The high reactivity of OH with ClO2ÿ 3 makes it dicult to observe its absorption. At the end of the eÿ aq decay, two bands at 245 and 500 nm were observed, which exhibit low intensities. In alkaline solutions, however, the hydroxyl radical dissociates (pK = 11.9) and forms the Oÿ radical ion, which has a much lower oxidizing ability. The life time of the ClO2ÿ 3 radical anion should, therefore, be much longer at a high pH value, and thus allow a more precise determination of its absorption spectrum. The rate constant of reaction (1) remains unchanged when the pH is increased from 8.5 to 12.5. Curve 1 in Fig. 2 depicts the absorption 12 ms after the electron pulse, measured in a 0.5 M solution of NaClO3 at pH 12.5. This spectrum includes the absorption of the Oÿ radical ion lmax 240 nm, e 240 Mÿ1 cmÿ1) (Hug, 1981) as well as radical anion. The spectrum of the that of the ClO2ÿ 3 ClO2ÿ 3 radical anion (Curve 2 in Fig. 2) was obtained
Fig. 2. Optical absorption spectrum obtained 12 ms after 10 ns electron pulses (curve 1) and after correction for the absorption of the Oÿ radical ion (curve 2). Solution 0.5 M NaClO3, pH = 12.5, Ar. Irradiation dose corresponds to 2.0 10ÿ6 M hydrated electrons.
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B.G. Ershov et al. / Radiation Physics and Chemistry 59 (2000) 309±312
after the correction for the Oÿ radical ion. It exhibits two optical bands at 245 nm e 900 Mÿ1 cmÿ1) and at 500 nm e 500 Mÿ1 cmÿ1). As expected, the radical anion disappears in alkaline solution ClO2ÿ 3 much slower than in neutral ones. The decay follows a second-order law according to
ÿ ÿ Oÿ ClO2ÿ 3 H2 O)2OH ClO3 :
4
The rate constant of this reaction is determined to be 1.2 109 Mÿ1 sÿ1. In summary, the hydrated electron reacts slowly with the ClOÿ 3 ion, resulting in the formation of the ÿ ClO2ÿ 3 radical anion. The reaction of OH with ClO3 is too slow to be measured in neutral or alkaline solutions or does not take place at all. The redox potential of the ClO3 radical has been estimated to be E 0 2:1 V (Stanbury, 1989) in aqueous solution, which is consistent with the conclusion that this potential seems to be between 1.66 and 2.41 V (Domae et al, 1994). On the other hand, the potential of the OH radical at pH = 7 is also about 2.2 V (Wardman, 1989). It follows from this comparison that the oxi dation of the ClOÿ 3 ion by the OH radical is a rather unfavorable process.
References Buxton, G., Greenstock, C.L., Helman, W.P, Ross, A.B., 1988. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solution. J. Phys. Chem. Rev. Data 17 (2). Buxton, G., Subhani, M.S., 1972. Radiation chemistry and photochemistry of oxychlorine ions. J. Chem. Soc. Faraday Trans. 1, 947. Domae, M., Katsumara, Y., Jiang, P.Y., Nagaiashi, R., Hasegawa, C., Ishogura, K., Yoshida, Y., 1994.
Observation of ClO3 radical in aqueous chlorate solution by pulse radiolysis. J. Phys. Chem. 98, 190. GraÈtzel, M., Henglein, A., Lilie, J, Beck, G., 1969. Pulsradiolytische Untersuchungen einiger Elementarprozesse der Oxydation und Reduktion des Nitritions. Ber. Bunsenges. Physik. Chem. 73, 646. GraÈtzel, M., Henglein, A., Taniguchi, S., 1970. Pulsradiolytische Beobachtungen uÈber die Reduction des NOÿ 3 -Ions und uÈber Bildung und Zerfall der persalpetrigen SaÈure in waÈssriger Loesung. Ber. Bunsenges. Physik. Chem. 74, 292. Hug, G.L., 1981. Optical spectra of nonmetallic inorganic transient species in aqueous solution. NSRDS Ð NBS 69, 12. Janata, E., 1982. Instrumentation of kinetic spectroscopy-10. A modular data acquisition system for laser ¯ash photolysis and pulse radiolysis experiments. Radiat. Phys. Chem. 40, 437. Janata, E., 1994. Instrumentation of kinetic spectroscopy-12. Software for data acquisition in kinetic experiments. Radiat. Phys. Chem. 44, 449. Janata, E., Gutsch, W., 1998. Instrumentation of kinetic spectroscopy-15. Part III: Injector pulse generators for electron accelerators. Radiat. Phys. Chem. 51, 65. Janata, E., Schuler, R.H., 1982. Rate constant for scavenging eÿ aq in N2O saturated solutions. J. Phys. Chem. 86, 2078. Kelm, M., Bohnert, E., 2000a. Radiation chemical eects in the near-®eld of a ®nal disposal site. Part I: Radiolytic products formed in concentrated NaCl solutions. Nucl. Technol., 129, 118. Kelm, M., Bohnert, E., 2000b. Radiation chemical eects in the near-®eld of a ®nal disposal site. Part II. Simulation of the radiolytic processes in concentrated NaCl solutions. Nucl. Technol., 129, 123. Stanbury, D.M., 1989. Reduction potentials involving inorganic free radicals in aqueous solution. In: Sykes A.G. (Ed.), Advances in Inorganic Chemistry, vol. 33, p. 69. Wardman, P., 1989. Reduction potentials of one-electron couples involving free radicals in aqueous solution. J. Phys. Chem. Rev. Data 18 (4), 1637±1755.