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The reaction of hydrogen atoms with hydrogen peroxide as a function of temperature
Radiation Physics and Chemistry 61 (2001) 109–113
The reaction of hydrogen atoms with hydrogen peroxide as a function of temperature T. Lundstro¨ma,*, H. Christensenb, K. Sehestedc a
Department of Physics and Measurement Technology, Linko¨ping University, SE-581 83 Linko¨ping, Sweden b Studsvik Nuclear AB, SE-611 82 Nyko¨ping, Sweden c Risø National Laboratory, DK-4000 Roskilde, Denmark Received 13 September 2000; accepted 23 December 2000
Abstract The temperature dependence for the reaction of H atoms with H2O2 at pH 1 has been determined using pulse radiolysis technique. The reaction was studied in the temperature range 10–1208C. The rate constant at 258C was found to be 5.1 0.5 107 dm3 mol1 s1 and the activation energy was found to be 10.7 0.5 kJ mol1. # 2001 Published by Elsevier Science Ltd. Keywords: Pulse radiolysis; Rate constant; Activation energy; H atom; Hydrogen peroxide
1. Introduction In order to make a reliable prediction of the concentration of radiolytic species in a nuclear reactor, it is necessary to have data on the rate constants of important chemical reactions at the operational temperature of the reactor. Over the past, several rate constants for important reactions have been determined (Christensen and Sehested, 1980; Elliot et al., 1990; Buxton and Mackenzie, 1992; Elliot, 1994; Christensen et al., 1994 and references therein). In the present investigation, the reaction of the primary species H and H2O2 has been studied at temperatures up to 1208C. H þ H2 O2 ! OH þ H2 O:
ð1Þ
Attempts at higher temperatures failed due to decomposition of H2O2 to O2, which reacts very fast with H and thus disturbs the measurement. The rate constant of reaction (1) has been determined earlier at ambient temperature by Sweet and Thomas *Corresponding author. Tel.:+46-155-22-1407; fax: +46155-26-3150. E-mail address: [email protected] (T. Lundstro¨m).
(1964) by measuring the build up of Cl 2 . They found a rate constant of 9 107 dm3 mol1 s1. Elliot (1989) studied the reaction, with an experimental setup similar to Sweet and Thomas, in the temperature range 20– 1008C and determined a rate constant at 258C of 5.0 107 dm3 mol1 s1 and an activation energy of 16 kJ mol1. This is believed to be more accurate than the determination by Sweet and Thomas, as it included, in the model, correction for Cl 2 decay reactions. Mezyk and Bartels (1995) studied the reaction in the temperature range 7.5–84.58C using EPR measuring techniques and determined a rate constant at 258C of 3.6 107 dm3 mol1 s1 and an activation energy of 21 kJ mol1. We used a chemical system similar to that of Sweet and Thomas (1964) and of Elliot (1989), and found a reaction rate of 5.1 0.5 107 dm3 mol1 s1 at 258C and an activation energy of 10.7 0.5 kJ mol1.
2. Experimental The solutions were saturated with Ar before use and pressurised in the experimental cell with 6 MPa Ar. The chemicals used: NaCl and HClO4 were of p.a. quality
0969-806X/01/$ - see front matter # 2001 Published by Elsevier Science Ltd. PII: S 0 9 6 9 - 8 0 6 X ( 0 1 ) 0 0 1 8 9 - X
110
T. Lundstro¨m et al. / Radiation Physics and Chemistry 61 (2001) 109–113
Fig. 1. The HTP cell at Ris National Laboratory.
Fig. 2. Computer fit to experimental data at 295 K.
and the water was triply distilled. All solutions contained 50 mmol dm3 NaCl and the concentration of H2O2 was varied between 1 and 5 mmol dm3. The pH was adjusted with HClO4 to pH 1 at ambient temperature. The experimental procedure has been described in detail previously by Christensen and Sehested (1980) and Sehested and Christensen (1990). The experiments were performed in our high-temperature, high-pressure (HTP) cell of optical path length 2.5 cm (Fig. 1). The HRC linac at Ris National Laboratory, Denmark, delivered 10 MeV electrons in a single pulse of 600– 800 mA, pulse length 0.1–0.5 ms. The dose was varied between 0.65 and 5 Gy per pulse. The optical absorption of Cl 2 , formed in the reaction, was measured at 340 nm using a Perkin Elmer double prisme monochromator and an 1P28 photomultiplier. The signal was processed on a LeCroy 9400A digital storage oscilloscope (175 MHz) and on an on-line PC. Cl 2 has an absorption maximum at 340 nm with e ¼ 8500 mol1 dm3 cm1 (Jayson et al., 1973). These values were used also at higher temperature.
During the pulse, the formation of Cl 2 is very rapid due to the following fast reactions:
OH þ Cl ! ClOH ;
9
3
1 1
k¼ 4:5 10 dm mol
s ; ð2Þ
ClOH þHþ ! Cl þ H2 O; k¼ 2:1 1010 dm3 mol1 s1 ; ð3Þ Cl þ Cl ! Cl 2;
k¼ 2:1 1010 dm3 mol1 s1 : Cl 2
+
H , OH H H2 OH H2O2 H2O
258C
dG=dT
408C
908C
0.1 3.55 0.4 2.83 0.76 4.45
3.27E-003 4.62E-003 5.77E-004 6.54E-003 3.85E-004
0.15 3.61 0.41 2.93 0.75 4.58
0.31 3.85 0.44 3.25 0.74 5.04
to determine the rate constant. The radiation dose was calculated as the height of the initial steep increase. The reaction rate was corrected using the decay of Cl 2 measured in the tail of the decay curve, see Fig. 2. The G-values were calculated as described in Table 1 and the rate constants as shown in Table 2 (Christensen et al., 1996). At higher temperatures (above 908C) the H2O2 underwent a catalytic thermal decomposition during heating, forming O2. This made the measurement of reaction (1) difficult due to the fast reaction (5): H þ O2 ! HO2 ; k¼ 2:1 1010 dm3 mol1 s1 ðat 258CÞ:
3. Results and discussion
Table 1 G-values (molecules/100 eV) and their change with temperature
ð4Þ
After the pulse, is formed by the slow reaction (1), followed by the fast reactions (2)–(4). The rate constants given are applicable at 258C. The build up has been modelled in the computer program MAKSIMA-CHEMIST (Carver et al., 1979)
ð5Þ This reaction competes effectively with reaction (1) (k¼ 5:1 107 dm3 mol1 s1 ) so that already at 1% decomposition, the main part of the H atoms reacts with O2 and no build up of Cl 2 is seen. At temperatures above 908C this effect is noticeable, making the results more uncertain. The decay of H2O2 is thought to take place at the walls of the cell. Because the wall area is rather large compared to the volume of the cell this effect is more pronounced than in a power reactor with large volumes of water. We have made several attempts to prevent this decomposition. In one attempt the surface of the cell was treated with concentrated phosphoric acid
111
T. Lundstro¨m et al. / Radiation Physics and Chemistry 61 (2001) 109–113 Table 2 Rate constants and activation energies Reaction
Rate constant (dm3mol1 s1)
Activation energy (J mol1)
OH+OH=H2O2 OH+E=OH OH+H=H2O OH+HO2=H2O+O2 OH+O 2 =O2+OH
5.38E+009 2.98E+010 6.84E+009 6.83E+009 9.52E+009
7.90E+03 1.47E+04 8.40E+03 1.42E+04 1.76E+04
OH+H2O2=HO2+H2O OH+H2=H+H2O OH+HO 2 =O2 +H2O OH+OH =H2O+O OH+O=HO 2
2.60E+007 3.22E+007 7.21E+009 1.15E+010 1.73E+010
1.42E+04 1.92E+04 1.42E+04 1.42E+04 1.42E+04
O+H2O=OH+OH E+E=H2+OH+OH E+H=H2+OHH2O E+O 2 =HO2+OH H2O E +HO2=HO2
1.50E+008 5.80E+009 2.31E+010 1.25E+010 1.92E+010
naa naa 1.40E+04 1.42E+04 1.42E+04
E+H2O2=OH+OH E+O2=O 2 E+H+=H E+H2O=H+OH E+HO 2 =O +OH
1.15E+010 1.83E+010 2.22E+010 1.83E+001 3.36E+009
1.56E+04 1.36E+04 1.26E+04 1.42E+04 1.42E+04
H+H=H2 H+O 2 =HO2 H+HO2=H2O2 H+H2O2=H2O+OH H+O2=HO2
5.28E+009 1.92E+010 1.92E+010 5.10E+007 2.04E+010
1.46E+04 1.42E+04 1.42E+04 1.07E+04 1.03E+04
H+OH=E+H2O HO2+HO2=H2O2+O2 HO2+O 2 =O2+HO2 + HO2=H +O2 H++O 2 =HO2
1.98E+007 7.93E+005 9.40E+007 3.80E+005 4.81E+010
3.76E+04 2.06E+04 7.60E+04 naa 1.42E+04
H2O2+OH=HO 2 +H2O HO 2 +H2O=H2O2+OH H2O2=H2O+O O+O=O2 H2O=H++OH
4.75E+008 1.20E+007 3.36E008 4.81E+009 1.47E001
1.88E+04 naa 6.30E+04 1.42E+04 naa
H++OH=H2O + O 2 +O2 =HO2 +O2H H+H2O=H2+OH O2 2 +H2O=HO2 +OH Cl+OH=ClOH
1.38E+001 1.42E+009 9.61E+005 9.61E+005 4.48E+009
1.42E+04 8.00E+04 1.42E+04 1.42E+04 1.42E+04
ClOH+H+=Cl+H2O Cl+Cl=Cl 2 Cl+Cl=Cl2 H+Cl=HCL ClOH=Cl+OH
2.08E+010 2.09E+010 1.03E+008 1.03E+008 5.75E+009
1.42E+04 1.42E+04 1.42E+04 1.42E+04 1.42E+04
+ H+Cl 2 =H +Cl +Cl Cl +Cl =Cl +Cl 2 2 3 Cl 2 =Cl+Cl H2O2+Cl 2 =OH+OH +Cl2
1.60E+010 3.70E+009 1.10E+005 4.20E+008
b
a b
na: Not applicable. Fitted to experimental data.
b b b
112
T. Lundstro¨m et al. / Radiation Physics and Chemistry 61 (2001) 109–113
Fig. 4. Arrhenius plot. Each point represents a mean value of several experimental data. At ambient temperature the error is 2% and at 345 K the error is 5%, referring to precision ð2sÞ.
Fig. 3. Computer fit to experimental data at 363 K.
Table 3 Rate contants for the reaction H+H2O2 as a function of temperature 3
1 1
T (8C)
T (K)
k (dm mol
10 22 31 42 50 62 72 82 90 100 111 120
283 295 304 315 323 335 345 355 363 373 384 393
4.40E+007 5.10E+007 6.00E+007 6.90E+007 7.60E+007 9.00E+007 9.50E+007 1.10E+008 1.10E+008 1.20E+008 1.40E+008 1.50E+008
s )
(Satterfield and Stein, 1957) in order to decrease the number of active sites on the wall. However, this showed no effect on the rate of decomposition of H2O2. In Figs. 2 and 3 the absorption of Cl 2 is shown as a function of time during the pulse and several microseconds after the pulse at 22 and 908C, respectively. The experimental values of k1 at various temperatures between 10 and 1208C are shown in Table 3. At 1308C the build-up could hardly be distinguished from the noise. Consequently, the result is not included in Table 3. The rate constant did not vary with the concentration of H2O2. In Fig. 4 the results are shown in an Arrhenius plot. From this an activation energy of 10.7 0.5 kJ mol1 has been calculated. The error limits include both precision ð2sÞ and estimations of accuracy. With uncertainties included the rate constant can thus be given as kT ¼ ð5:1 0:5Þ107 expðð10:7 0:5=RÞ ðð1=TÞ ð1=298ÞÞÞ:
The activation energy is lower than that found in previous studies. However, the lower value could be a result of the correction made for the presence of O2 in our work, caused by thermal decomposition of H2O2 at higher temperatures, see reaction (5). Elliot (1989), who used a mechanism similar to our, apparently did not make this correction and he thus measured higher reaction rates at the higher temperatures resulting in a higher activation energy. A change of G-values with temperature was not mentioned in his publication. Mezyk and Bartels (1995) measured the decay of the H atom directly by EPR detection and concluded that this corresponded to the reaction with H2O2, apparently without correction for reaction (5). This would result in a higher activation energy in the same way as for Elliot. The residence time of hydrogen peroxide in the cell was probably lower in the experiments of Elliot (1989) and of Mezyk and Bartels (1995).
Acknowledgements We would like to thank the Swedish Centre for Nuclear Technology and Swedish Nuclear Fuel Management Co for financial support. Fruitful discussions with Anders Lund are gratefully acknowledged. H. Corfitzen and T. Johansen are gratefully acknowledged for technical assistance.
References Buxton, G.V., Mackenzie, S.R., 1992. Non-linear Arrhenius behaviour of the rate constants of some activation-controlled reactions of the hydrated electron in the temperature range 20–2008C. J. Chem. Soc. Faraday Trans. 88, 2833–2836.
T. Lundstro¨m et al. / Radiation Physics and Chemistry 61 (2001) 109–113 Carver, M.B., Hanley, D.V., Chaplin, K.R., 1979. MAKSIMACHEMIST. A program for mass action kinetics simulation by automatic chemical equation manipulation and integration using stiff techniques. Atomic Energy of Canada Ltd (AECL-6413). Christensen, H., Molander, A., Pettersson, S.-O., Lassing, A., Tomani, H., 1996. Studies of water radiolysis in the INCA loop at 508C and 1508C. Studsvik Material AB (Studsvik/M96/180). Christensen, H., Sehested, K., 1980. Pulse radiolysis at high temperatures and high pressures. Radiat. Phys. Chem. 16, 183–186. Christensen, H., Sehested, K., Lo¨gager, T., 1994. Temperature dependence of the rate constant for reactions of hydrated electrons with H, OH and H2O2. Radiat. Phys. Chem. 43 (6), 527–531. Elliot, A.J., 1989. A pulse radiolysis study of the temperature dependence of reactions involving H, OH and e aq in aqueous solutions. Radiat. Phys. Chem. 34 (5), 753–758. Elliot, A.J., 1994. Rate constants and G-values for the simulation of the radiolysis of light water over the range 0–3008C. AECL-11073.
113
Elliot, A.J., McCracken, D.R., Buxton, G.V., Wood, N.D., 1990. Estimation of the rate constants for near-diffusion controlled reactions in water at high temperature. J. Chem. Soc. Faraday Trans. 86, 1539–1547. Jayson, G.G., Parsons, B.J., Swallow, A.J, 1973. Some simple, highly reactive, inorganic, chlorine derivatives in aqueous solution } their formation using pulses of radiation and their role in the mechanism of Fricke dosimeter. J. Chem. Soc. Faraday Trans. 69, 1597–1607. Mezyk, S.P., Bartels, D.M., 1995. Direct EPR measurement of Arrhenius parameters for the reaction of H atoms with H2O2 and D atoms with D2O2 in aqueous solution. J. Chem. Soc. Faraday Trans. 91 (18), 3127–3132. Satterfield, C.N., Stein, T.W., 1957. Homogenous decomposition of hydrogen peroxide vapour. J. Phys. Chem. 61, 537–540. Sehested, K., Christensen, H., 1990. The rate constant of bimolecular reaction of hydrogen atoms at elevated temperatures. Radiat. Phys. Chem. 36 (3), 499–500. Sweet, J.P., Thomas, J.K., 1964. Absolute rate constant for H atom reactions in aqueous solutions. J. Phys. Chem. 68 (6), 1363–1368.
The reaction of hydrogen atoms with hydrogen peroxide as a function of temperature T. Lundstro¨ma,*, H. Christensenb, K. Sehestedc a
Department of Physics and Measurement Technology, Linko¨ping University, SE-581 83 Linko¨ping, Sweden b Studsvik Nuclear AB, SE-611 82 Nyko¨ping, Sweden c Risø National Laboratory, DK-4000 Roskilde, Denmark Received 13 September 2000; accepted 23 December 2000
Abstract The temperature dependence for the reaction of H atoms with H2O2 at pH 1 has been determined using pulse radiolysis technique. The reaction was studied in the temperature range 10–1208C. The rate constant at 258C was found to be 5.1 0.5 107 dm3 mol1 s1 and the activation energy was found to be 10.7 0.5 kJ mol1. # 2001 Published by Elsevier Science Ltd. Keywords: Pulse radiolysis; Rate constant; Activation energy; H atom; Hydrogen peroxide
1. Introduction In order to make a reliable prediction of the concentration of radiolytic species in a nuclear reactor, it is necessary to have data on the rate constants of important chemical reactions at the operational temperature of the reactor. Over the past, several rate constants for important reactions have been determined (Christensen and Sehested, 1980; Elliot et al., 1990; Buxton and Mackenzie, 1992; Elliot, 1994; Christensen et al., 1994 and references therein). In the present investigation, the reaction of the primary species H and H2O2 has been studied at temperatures up to 1208C. H þ H2 O2 ! OH þ H2 O:
ð1Þ
Attempts at higher temperatures failed due to decomposition of H2O2 to O2, which reacts very fast with H and thus disturbs the measurement. The rate constant of reaction (1) has been determined earlier at ambient temperature by Sweet and Thomas *Corresponding author. Tel.:+46-155-22-1407; fax: +46155-26-3150. E-mail address: [email protected] (T. Lundstro¨m).
(1964) by measuring the build up of Cl 2 . They found a rate constant of 9 107 dm3 mol1 s1. Elliot (1989) studied the reaction, with an experimental setup similar to Sweet and Thomas, in the temperature range 20– 1008C and determined a rate constant at 258C of 5.0 107 dm3 mol1 s1 and an activation energy of 16 kJ mol1. This is believed to be more accurate than the determination by Sweet and Thomas, as it included, in the model, correction for Cl 2 decay reactions. Mezyk and Bartels (1995) studied the reaction in the temperature range 7.5–84.58C using EPR measuring techniques and determined a rate constant at 258C of 3.6 107 dm3 mol1 s1 and an activation energy of 21 kJ mol1. We used a chemical system similar to that of Sweet and Thomas (1964) and of Elliot (1989), and found a reaction rate of 5.1 0.5 107 dm3 mol1 s1 at 258C and an activation energy of 10.7 0.5 kJ mol1.
2. Experimental The solutions were saturated with Ar before use and pressurised in the experimental cell with 6 MPa Ar. The chemicals used: NaCl and HClO4 were of p.a. quality
0969-806X/01/$ - see front matter # 2001 Published by Elsevier Science Ltd. PII: S 0 9 6 9 - 8 0 6 X ( 0 1 ) 0 0 1 8 9 - X
110
T. Lundstro¨m et al. / Radiation Physics and Chemistry 61 (2001) 109–113
Fig. 1. The HTP cell at Ris National Laboratory.
Fig. 2. Computer fit to experimental data at 295 K.
and the water was triply distilled. All solutions contained 50 mmol dm3 NaCl and the concentration of H2O2 was varied between 1 and 5 mmol dm3. The pH was adjusted with HClO4 to pH 1 at ambient temperature. The experimental procedure has been described in detail previously by Christensen and Sehested (1980) and Sehested and Christensen (1990). The experiments were performed in our high-temperature, high-pressure (HTP) cell of optical path length 2.5 cm (Fig. 1). The HRC linac at Ris National Laboratory, Denmark, delivered 10 MeV electrons in a single pulse of 600– 800 mA, pulse length 0.1–0.5 ms. The dose was varied between 0.65 and 5 Gy per pulse. The optical absorption of Cl 2 , formed in the reaction, was measured at 340 nm using a Perkin Elmer double prisme monochromator and an 1P28 photomultiplier. The signal was processed on a LeCroy 9400A digital storage oscilloscope (175 MHz) and on an on-line PC. Cl 2 has an absorption maximum at 340 nm with e ¼ 8500 mol1 dm3 cm1 (Jayson et al., 1973). These values were used also at higher temperature.
During the pulse, the formation of Cl 2 is very rapid due to the following fast reactions:
OH þ Cl ! ClOH ;
9
3
1 1
k¼ 4:5 10 dm mol
s ; ð2Þ
ClOH þHþ ! Cl þ H2 O; k¼ 2:1 1010 dm3 mol1 s1 ; ð3Þ Cl þ Cl ! Cl 2;
k¼ 2:1 1010 dm3 mol1 s1 : Cl 2
+
H , OH H H2 OH H2O2 H2O
258C
dG=dT
408C
908C
0.1 3.55 0.4 2.83 0.76 4.45
3.27E-003 4.62E-003 5.77E-004 6.54E-003 3.85E-004
0.15 3.61 0.41 2.93 0.75 4.58
0.31 3.85 0.44 3.25 0.74 5.04
to determine the rate constant. The radiation dose was calculated as the height of the initial steep increase. The reaction rate was corrected using the decay of Cl 2 measured in the tail of the decay curve, see Fig. 2. The G-values were calculated as described in Table 1 and the rate constants as shown in Table 2 (Christensen et al., 1996). At higher temperatures (above 908C) the H2O2 underwent a catalytic thermal decomposition during heating, forming O2. This made the measurement of reaction (1) difficult due to the fast reaction (5): H þ O2 ! HO2 ; k¼ 2:1 1010 dm3 mol1 s1 ðat 258CÞ:
3. Results and discussion
Table 1 G-values (molecules/100 eV) and their change with temperature
ð4Þ
After the pulse, is formed by the slow reaction (1), followed by the fast reactions (2)–(4). The rate constants given are applicable at 258C. The build up has been modelled in the computer program MAKSIMA-CHEMIST (Carver et al., 1979)
ð5Þ This reaction competes effectively with reaction (1) (k¼ 5:1 107 dm3 mol1 s1 ) so that already at 1% decomposition, the main part of the H atoms reacts with O2 and no build up of Cl 2 is seen. At temperatures above 908C this effect is noticeable, making the results more uncertain. The decay of H2O2 is thought to take place at the walls of the cell. Because the wall area is rather large compared to the volume of the cell this effect is more pronounced than in a power reactor with large volumes of water. We have made several attempts to prevent this decomposition. In one attempt the surface of the cell was treated with concentrated phosphoric acid
111
T. Lundstro¨m et al. / Radiation Physics and Chemistry 61 (2001) 109–113 Table 2 Rate constants and activation energies Reaction
Rate constant (dm3mol1 s1)
Activation energy (J mol1)
OH+OH=H2O2 OH+E=OH OH+H=H2O OH+HO2=H2O+O2 OH+O 2 =O2+OH
5.38E+009 2.98E+010 6.84E+009 6.83E+009 9.52E+009
7.90E+03 1.47E+04 8.40E+03 1.42E+04 1.76E+04
OH+H2O2=HO2+H2O OH+H2=H+H2O OH+HO 2 =O2 +H2O OH+OH =H2O+O OH+O=HO 2
2.60E+007 3.22E+007 7.21E+009 1.15E+010 1.73E+010
1.42E+04 1.92E+04 1.42E+04 1.42E+04 1.42E+04
O+H2O=OH+OH E+E=H2+OH+OH E+H=H2+OHH2O E+O 2 =HO2+OH H2O E +HO2=HO2
1.50E+008 5.80E+009 2.31E+010 1.25E+010 1.92E+010
naa naa 1.40E+04 1.42E+04 1.42E+04
E+H2O2=OH+OH E+O2=O 2 E+H+=H E+H2O=H+OH E+HO 2 =O +OH
1.15E+010 1.83E+010 2.22E+010 1.83E+001 3.36E+009
1.56E+04 1.36E+04 1.26E+04 1.42E+04 1.42E+04
H+H=H2 H+O 2 =HO2 H+HO2=H2O2 H+H2O2=H2O+OH H+O2=HO2
5.28E+009 1.92E+010 1.92E+010 5.10E+007 2.04E+010
1.46E+04 1.42E+04 1.42E+04 1.07E+04 1.03E+04
H+OH=E+H2O HO2+HO2=H2O2+O2 HO2+O 2 =O2+HO2 + HO2=H +O2 H++O 2 =HO2
1.98E+007 7.93E+005 9.40E+007 3.80E+005 4.81E+010
3.76E+04 2.06E+04 7.60E+04 naa 1.42E+04
H2O2+OH=HO 2 +H2O HO 2 +H2O=H2O2+OH H2O2=H2O+O O+O=O2 H2O=H++OH
4.75E+008 1.20E+007 3.36E008 4.81E+009 1.47E001
1.88E+04 naa 6.30E+04 1.42E+04 naa
H++OH=H2O + O 2 +O2 =HO2 +O2H H+H2O=H2+OH O2 2 +H2O=HO2 +OH Cl+OH=ClOH
1.38E+001 1.42E+009 9.61E+005 9.61E+005 4.48E+009
1.42E+04 8.00E+04 1.42E+04 1.42E+04 1.42E+04
ClOH+H+=Cl+H2O Cl+Cl=Cl 2 Cl+Cl=Cl2 H+Cl=HCL ClOH=Cl+OH
2.08E+010 2.09E+010 1.03E+008 1.03E+008 5.75E+009
1.42E+04 1.42E+04 1.42E+04 1.42E+04 1.42E+04
+ H+Cl 2 =H +Cl +Cl Cl +Cl =Cl +Cl 2 2 3 Cl 2 =Cl+Cl H2O2+Cl 2 =OH+OH +Cl2
1.60E+010 3.70E+009 1.10E+005 4.20E+008
b
a b
na: Not applicable. Fitted to experimental data.
b b b
112
T. Lundstro¨m et al. / Radiation Physics and Chemistry 61 (2001) 109–113
Fig. 4. Arrhenius plot. Each point represents a mean value of several experimental data. At ambient temperature the error is 2% and at 345 K the error is 5%, referring to precision ð2sÞ.
Fig. 3. Computer fit to experimental data at 363 K.
Table 3 Rate contants for the reaction H+H2O2 as a function of temperature 3
1 1
T (8C)
T (K)
k (dm mol
10 22 31 42 50 62 72 82 90 100 111 120
283 295 304 315 323 335 345 355 363 373 384 393
4.40E+007 5.10E+007 6.00E+007 6.90E+007 7.60E+007 9.00E+007 9.50E+007 1.10E+008 1.10E+008 1.20E+008 1.40E+008 1.50E+008
s )
(Satterfield and Stein, 1957) in order to decrease the number of active sites on the wall. However, this showed no effect on the rate of decomposition of H2O2. In Figs. 2 and 3 the absorption of Cl 2 is shown as a function of time during the pulse and several microseconds after the pulse at 22 and 908C, respectively. The experimental values of k1 at various temperatures between 10 and 1208C are shown in Table 3. At 1308C the build-up could hardly be distinguished from the noise. Consequently, the result is not included in Table 3. The rate constant did not vary with the concentration of H2O2. In Fig. 4 the results are shown in an Arrhenius plot. From this an activation energy of 10.7 0.5 kJ mol1 has been calculated. The error limits include both precision ð2sÞ and estimations of accuracy. With uncertainties included the rate constant can thus be given as kT ¼ ð5:1 0:5Þ107 expðð10:7 0:5=RÞ ðð1=TÞ ð1=298ÞÞÞ:
The activation energy is lower than that found in previous studies. However, the lower value could be a result of the correction made for the presence of O2 in our work, caused by thermal decomposition of H2O2 at higher temperatures, see reaction (5). Elliot (1989), who used a mechanism similar to our, apparently did not make this correction and he thus measured higher reaction rates at the higher temperatures resulting in a higher activation energy. A change of G-values with temperature was not mentioned in his publication. Mezyk and Bartels (1995) measured the decay of the H atom directly by EPR detection and concluded that this corresponded to the reaction with H2O2, apparently without correction for reaction (5). This would result in a higher activation energy in the same way as for Elliot. The residence time of hydrogen peroxide in the cell was probably lower in the experiments of Elliot (1989) and of Mezyk and Bartels (1995).
Acknowledgements We would like to thank the Swedish Centre for Nuclear Technology and Swedish Nuclear Fuel Management Co for financial support. Fruitful discussions with Anders Lund are gratefully acknowledged. H. Corfitzen and T. Johansen are gratefully acknowledged for technical assistance.
References Buxton, G.V., Mackenzie, S.R., 1992. Non-linear Arrhenius behaviour of the rate constants of some activation-controlled reactions of the hydrated electron in the temperature range 20–2008C. J. Chem. Soc. Faraday Trans. 88, 2833–2836.
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