12594
J. Phys. Chem. 1994,98, 12594-12599
EPR Measurement of the Reaction of Atomic Hydrogen with Periodate and Iodate in Aqueous Solution? Stephen P. Mezyk* and Roy MacFarlane AECL Research, Whiteshell Laboratories, Pinawa, Manitoba, ROE 1u) Canada
David M. Bartels Chemistry Division, Argonne National Laboratory, Argonne, Illinois 60439 Received: March 16, 1994; In Final Form: June 12, 1994@
Electron paramagnetic resonance free induction decay (FID)attenuation measurements have been used to determine Arrhenius parameters for the reaction of hydrogen atoms with IO4- and 103- in aqueous solution at pH 4.6. At 25 "C, scavenging rate constants of (7.31 f 0.12) x lo8 and (1.22 f 0.04) x lo7 dm3 mol-' s-' were obtained for 104- and I03-, respectively, with corresponding activation energies, measured over the temperature range 6-75 "C, of 22.35 and 27.46 kJ mol-'. From kinetic measurements in 0.10 and 1.00 M HC104 over the same temperature range, the rate constants for hydrogen atom reaction with the acid forms of these species, HsI06 and HI03, were calculated to be 3.5 x lo8 and 6.0 x lo6 dm3 mol-' s-l (25 "C), with activation energies of 53.00 and 36.55 kJ mol-', respectively. We suggest that all of these reactions involve oxygen atom transfer to the hydrogen atom, giving 'OH and a reduced iodine complex as product.
Introduction The analysis of hypothetical nuclear reactor accident scenarios has shown that radioactive iodine-13 1 and iodine-132 release from nuclear fuel to the environment could be the main source of radiation dose to the public.' This estimate is based on a combination of inventory, half-life, volatility, and biological activity. Iodine is initially released as C S I , ~ and ~ ~as large quantities of water are also expected to be present in most accidents from water-cooled reactors, a good knowledge of the aqueous chemistry of iodide under reactor containment conditions is necessary to model its behavior. These considerations have been responsible for a resurgence of interest in the radiation-induced chemistry of iodine-containing compound^^-^ and have stimulated the present study. Of particular concern is the minimization of 12 formation, which is volatile and thus harder to contain than the ionic forms of iodine. However, the irradiation of IO3- solutions has been demonstrated to produce 12 and, in the presence of organic impurities, to form volatile, low molecular weight, i~doalkanes.~ Although no formal mechanisms have been established for these systems, multistep pathways have been proposed,8 involving iodate reaction with water radiolysis products to form 1 0 2 - and radiolytic formation of IO4-, a strong oxidant that readily oxidizes iodide to iodine. In a recent study,9 the techniques of pulse radiolysis and absorption spectroscopy were used to reinvestigate the aqueous radiation chemistry of iodate, with rate constants and mechanisms for its reaction with the hydrated electron, the hydroxyl radical, and the oxide radical being determined. Unfortunately, the radiolysis of acidic 103-ltert-butyl alcohol solutions showed no change in the visible or near-UV absorption spectrum, and thus no measurement of the iodate reaction rate constant with the hydrogen atom was possible. A previous investigation' of
* To whom correspondence is to be addressed.
t Work at Argonne performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Science, US-DOE,under Contract W-31-109-ENG-38. Abstract published in Advance ACS Abstracts, October 15, 1994. @
this reaction, using steady state competition techniques, calculated a value of -6 x lo7 dm3 mol-' s-l at room temperature. The immediate product of this reaction was not established. This paper reports our direct measurement of rate constants and activation energies for the reaction of 'H atoms with iodate species in solution. Measurements were also performed for the analogous periodate (IO4-) anion. Because of its great specificity, direct EPR detection of the decay of the 'H atom following pulse radiolysis was the monitoring method of The pulsed EPR-based free induction decay (FTD) attenuation method12-16 has been shown to be particularly advantageous for rate constant measurements because of the simple pseudofirst-order scavenging kinetics generally obtained. Insight into the nature of the products was gained from 'H atom spin exchange and chemically induced dynamic electron polarization (CIDEP) in the presence of these anions.
Experimental Section The experimental procedure has been described in detail in several previous publication^,^^-^^ and thus only a short description shall be given here. Hydrogen atoms were generated in aqueous solution within an EPR cavity by pulse radiolysis with 3 MeV electrons from a Van de Graaff accelerator. Stock solutions were prepared by addition of HC104 (Mallinkrodt A.R. Grade, 69.05%) or phosphate buffer (Baker, Analyzed) to Millipore filtered water. Exact acid concentrations were determined by standardization of the concentrated HC104 against HC1 (Aldrich, 1.029 N volumetric standard). A total aqueous sample of -190 mL was recirculated through a flat cell in the cavity at a high flow rate, and the actual volume irradiated in each pulse was less than 0.10 mL. The approximate average radiation dose to this volume was 1.5 and 3.0 kradpulse for the 12 and 25 ns pulses used, respectively. A 35 ns microwave probe pulse was applied to the sample immediately after irradiation, and the resulting free induction decay of the 'H atom low-field (mr = +l/2) EPR transition was recorded on a digital oscilloscope. Typically, 2000 shots were averaged to measure each FID, at a repetition rate of 120 Hz.
0022-365419412098-12594$04.50/0 Published 1994 by the American Chemical Society
J. Phys. Chem., Vol. 98, No. 48, 1994 12595
Reaction of H with 1 0 4 - and IO3Scavenging experiments were performed by successive addition of weighed amounts of solid potassium periodate (Aldrich, 99.99%) or iodate (Aldrich, 99.5%, A.C.S. reagent) to the initial known volume of solution. Accuracy of these concentrations is estimated at better than 2%. Hydrogen atom longitudinal magnetization was also recorded as a function of time after the radiolysis pulse for several different scavenger concentrations by scanning the microwave pulse delay relative to the radiolysis pulse and integrating the first 100 ns of the FID signal in a gated integrator. All absorption spectra were recorded using temperaturecontrolled 1.0 cm cells in a Hewlett-Packard 8450A spectrophotometer.
C
H,IO,
0.6 .
Y
,
B
e
u.
-* Haloe= H*
u)
+ H.10;
b
v, 28
0.2
32
30
34
36
'
---
0.0 0.0
Results and Discussion The effective damping rate of the is given by12-15
= ZH,O + io;
1 .o
2.0
FTD in these experiments
3.0
4.0
50
6.0
aH+
Figure 1. (a) Literaturez0data for temperature dependence of periodate equilibria. (b) Calculated acidity dependence of fractional speciation
of periodate at 25 "C. H,IO, == H410,where kJS] is the 'H atom scavenging rate and Z,k',,[Ri] represents the spin-dephasing contribution of second-order spin exchange and recombination reactions between 'H and other free radicals. At the radiation doses typically used in our experiments, the latter term may not be negligible. In previous FTD experiments,12-16 a hydroxyl radical scavenger such as methanol or bromide was added to produce radicals that diffuse and react much more slowly with 'H atoms than the *OHradical, thus minimizing the importance of the latter term in eq 1 while having a minimal impact on the kinetics of disappearance of the 'H atom. For periodate and iodate solutions, however, this was not possible. The addition of Br- to 1 0 3 - solutions under acidic conditions gave scattered results, which was attributed to interference from Br2 formation, from pse~do-Dushman~~ type equilibria. IO,-
+ 2Br- + 3H'
HOBr
=+ 2HOBr
+ IOH
(2)
H,IO,-
H4106- == 2H,O
-
-
+ IO4-
(5) (6)
-
(3)
Bromine reacts at almost a diffusion-controlled rate with 'H atoms ( k 1O'O dm3 mol-l s-l),18 which makes measurement of the iodate value impossible. A similar problem was observed when methanol was used as an 'OH scavenger. This problem became worse with increasing acidity, as expected from reactions 2 and 3, but was found to be significant even with solutions containing phosphate buffer (pH 4.6). Hence, all iodate EPR measurements were performed without any 'OH scavenger present. Despite this, the measured 'H atom free induction decays were found to be well described by a pure exponentially-damped cosine. l2-I6 Combined with the experimentally observed linear scavenging plots, whose slopes showed no dose dependence, this indicated that, under these conditions, 'H atom loss by *OHrecombination did not significantly affect the rate constant measurement.16 Periodate is also known to be a powerful oxidant, so again no added hydroxyl radical scavengers could be used. However, the fast reaction rate constant between *OH and IO4-, -5 x lo9 dm3 mol-' s-',19 ensured that recombination of hydrogen atoms and hydroxyl radicals was minimal in this system. Periodate Reaction. The periodate system has been well studied previously,20,21 with various equilibria characterized. For pH below 5, the main equilibria are
+ H+
(4)
The equilibrium constants for these reactions have been determined over a wide range of temperature?O and these literature values are shown in Figure la. Based on this data, fractional populations of the several periodate species under various conditions can be readily calculated. A typical example at 25 "C is shown in Figure lb. Under phosphate buffered conditions (pH 4.6), it is seen that the IO4- species dominates, while at acidic pH's, a mixture of tetrahedral IO4- and octahedral H5106 is obtained. The fraction of H4IO6- under all conditions is seen to be very small. The overall 'H atom scavenging rate constant in pH 4.6 phosphate buffer 'H
+ Br- + H+ =+ Br, + H,O
H,IO;-
+ H+
+ IO4- - products
(7)
was measured, and a value of (7.31 f 0.12) x lo8 dm3 mol-' s-l was determined at 25 "C. This value cannot be attributed to hydrogen atom reaction with the small fraction of H4IO6present (-2.4%, see Table l), as this would require that the hydrogen atom reaction with w O 6 - is diffusion-controlled ( k 3.7 x 1Olo dm3 mol-' s-l) and also that the reaction rate should become faster at lower temperatures. This was not experimentally observed. Therefore, we believe that the hydrogen atom reaction under these conditions is dominated by Io4-. The mechanism of reaction is also of interest. Of the potential reactions
-
+ IO4- - OH- + 'IO, 'H + IO4- - 'OH + IO,'H + IO4- - OHIO4'H
(9) (10) (1 1)
reaction 8 involving an electron transfer reaction to form 'Io42-
Mezyk et al.
12596 J. Phys. Chem., Vol. 98, No. 48, 1994
TABLE 1: Temperature and pH Dependence of Measured Rate Constants for Hydrogen Atom Reaction with Periodatea fraction tHC1041, 10-8km,, K6, 104K4, 10-8kmnFIoa-, 10-8kc~c(Hd06), M T, "C dm3mol-' molZdm-6 mol dm-3 104HJO6HsIO6 dm3 mol-' s-I dm3mol-' s-' 10.lb 43.6 0.001 0.561 14.62 0.438 36.8 440.1 1.oo 67.3 0.992 0.03b 0.67 0.007 0.001 4.90 0.69 10.93 0.10 6.0 0.986 0.07b 1.67 0.001 4.85 0.014 1.72 23.00 0.10 16.5 0.986 0.07b 1.55 0.001 4.85 0.014 1.60 23.00 0.10 16.5 0.976 0.21b 3.49 0.024 0.001 4.81 3.62 40.42 0.10 25.0 0.855 2.47b 14.9 0.001 7.59 0.144 178.7 50.0 15.2 0.10 0.004 4.44 0.078 0.918 11.76 4.91 4.44 4.60' 7.0 0.002 7.09 0.040 0.958 23.79 4.84 4.60' 17.1 7.09 0.001 7.31 0.975 0.024 40.42 4.81 25.0 7.31 4.61' 0.OOO 19.1 0.005 0.995 188.8 7.85 4.67' 51.3 19.1 0.000 28.6 0.998 0.002 19.8 28.6 638.5 4.83' 75.5 a The fractional contributions from each species is determined from the K6 and K4 values of ref 21 to enable direct calculation of the Arrhenius parameters for the reaction of hydrogen atoms with HsIO6 (Figure 2). * Rate constants calculated based on Arrhenius plot for pH 4.6 data. Measured
pH values.
does not occur, as the 'Io42-species, which has an intense absorption maximum at 360 nm,22was not observed in previous pulse radiolysis measurements over the pH range 1-3.22 If reaction 11 were to occur, then one potential pathway for the *HI04- product is its dissociation, to again give periodate and a hydrogen atom, 'HI04-
- 'H +
IO4-
(12)
22
J
20
C
This situation has been observed previously for hydrogen atom reaction with bromide and iodide.16 For these two systems, it was found that when longitudinal magnetization measurements were performed, the 'H atom CIDEP persisted an order of magnitude longer than expected from the FID scavenging lifetime. Based on this observation the rate constant for this dissociation was estimated to be -lo6 s-l. However, when similar experiments were performed in the presence of IO4-, it was found that periodate attenuated the 'H atom CIDEP signal as quickly as predicted from the FID measured reaction rate constant, consistent with an irreversible 'H atom reaction. It should be noted that this experiment is not conclusive, however, as a dissociation reaction with a rate constant an order of magnitude smaller than lo6 s-l would not be detectable in these experiments. Unfortunately, because of the large AGf" value for the hydrogen atom (+221 kJ no distinction between reactions 9- 11 could be made by thermodynamic calculation. Reaction 9 is considered unlikely, as the large value of AGf0('I03) = 190 kJ mol-' 22 suggests that formation of this intermediate will be avoided. From the above observations, we I03-, believe that the overall reaction products are 'OH formed either by direct reaction or through the intermediate species. This is in agreement with previous finding^^^,^^ which suggest that the reduction of Ivl to Iv compounds generally take place by transfer of an hydroxyl radical. By repeating this measurement over the temperature range 7.0-75.5 "C, the Arrhenius plot shown in Figure 2 was obtained. All of the individual rate constants, equilibrium constants, and fractions of the different species are summarized in Table 1. The listed pH values for the phosphate buffered solutions at the different temperatures were independently determined using a pH electrode. The temperature dependence of this reaction is given by
+
+
k7 = 6.71 x 10l2exp(-22350/RT) Based on the scatter of data in Figure 2, the rate constants calculated from this formula should be considered uncertain by 10-1570.
18
16l
'
28
,
30
1
32
34
36
I
lo3 /Temperature (K) Figure 2. Arrhenius plot of In k,, vs 1/Tfor hydrogen atom reaction with IO4-, measured at pH 4.6 (M) and for H5106, calculated from measurements at 0.10 (0)and 1.00 M (A)HC104. From the 1iterature2Oequilibrium constants shown in Figure la, the periodate speciation in 0.10 and 1.00 M HC104 can be readily calculated. By measuring the overall hydrogen atom reaction rate at these acidities, the reaction rate constant due to H5106reaction
'H + H,IO,
-
products
(13)
can be evaluated. This has been done over the temperature range 6.0-67.3 "C, with individual values and conditions again given in Table 1. The calculated Arrhenius plot is also shown in Figure 2 and is seen to be an excellent straight line, with temperature dependence given by
k,, = 5.90 x 1017exp(-53000/RT) With the additional error from the equilibrium constants, the uncertainty in the calculated values based on this formula is believed to be 15-20%. At 25 "C, the calculated rate constant for hydrogen atom reaction with H5106is 3.5 x lo8dm3 mol-' s-l, about a factor of 2 slower than for Io4-. This provides further justification for the assumption that the hydrated hIO6species makes a negligible contribution to the rate at pH 4.6. These rate constants for 'H atom reaction with H5106 have been evaluated using the rate constant data for IO4- collected at pH 4.6. Strictly speaking, however, to extrapolate the pH 4.6 rate constant data for hydrogen atom reaction with periodate
J. Phys. Chem., Vol. 98, No. 48, 1994 12597
Reaction of H with 104- and 103to the more concentrated acid conditions, one requires the inclusion of activity coefficients of the reactants and the transition state24
200,
I
I
0'O9
I
1
1
t
18 0
One can assume that the activities of the 1 0 4 - and transition state will be dominated by electrostatic interactions with the electrolyte solution in very similar ways, so thatflt)/'IO4-) PZ 1. In this case any change in reaction rate will be dominated by the 'H atom activity. The activity coefficients of small neutral molecules are generally weak functions of electrolyte concentration, usually written in the formz5
s
10
.-."
+
E, = d A = kT/h exp{A$/R
+ RT
+ 11
(16)
(dm3 mol-')
(17)
where k is the Boltzmann constant, h is the Planck constant, and we have assumed a transmission coefficient of unity. The kTlh prefactor is 6.3 x 10l2 s-' at 298 K, so that we ascribe the large preexponential to a large positive activation entropy of A 9 = $88 J mol-' K-l. The "analogous" reaction of 'H with IO4- has an activation entropy which is approximately -6.5 J mol-' K-l. Apart from the negative charge, the major difference is the hydration equilibrium (6), which converts the iodine species from tetrahedral to octahedral coordination. Crouthamel et aL20 noted that the dehydration of HJO6-, eq 6, occurs with a very large increase in entropy, A&* = 180 J mol-' K-l, which compensates the unfavorable enthalpy, A H 6 = 45.6 kJ mol-'. Given this fact, it seems almost certain that reaction 13 releases two water molecules as a product, with much of the dehydration occumng prior to the transition state. We therefore envision this reaction as a concerted oxygen transfer/dehydration/remangementconverting the iodo-oxyacid from the octahedral iodine(VI) complex to the pyramidal iodine(V) structure of IO3-, with overall stoichiometry
+
'H
32
3.0
(15)
where C is the electrolyte concentration in molar or molal units. The activity coefficient can be obtained directly from the "salt effect" on solubilities of gaseous species, f = L&, where L is the Ostwald solubility coefficient. In a recent p~blication,'~ it was demonstrated that HZprovides a very good model for the 'H atom solubility due to the similarity of polarizabilities of the two species. The solubility of H2 in various concentrated electrolyte solutions has been reviewed by Clever in 1980.26 Though perchloric acid was not among the systems tabulated, the salt effect parameter, ks, for other univalent electrolytes, including HCl, is typically less than 0.1 dm3 mol-' for H2 at any temperature. We can conclude that the reaction rate for 'H IO4- in 0.10 M perchloric acid is almost equal to that in the dilute solution, and the rate in 1.00 M HC104 probably differs by less than 10%. We have taken some pains to establish the reliability of our measurements and analysis for reaction 13, because the Arrhenius preexponential of 5.90 x 1017 dm3 mol-' s-l for this reaction is extremely large and unusual. As such, it deserves some comment. If we invoke the transition state theory, we can write the Arrhenius parameters in terms of the (temperature independent) entropy A 9 and enthalpy Al? of activation, using13J4
+ H,I06 - 2 H 2 0 + IO3- + *OH+ H+
(18)
40
-
16.0
2.8
-log f = k,C
20 30 I O ' [IO3'] / M
34
3.6
10~/~emperature (K)
Figure 3. (a) Scavenging rate constant determination for reaction of hydrogen atoms with 103- at pH 4.6 and 25.0 "C. (b) hhenius plot of In k,, vs 1/Tfor this reaction over the temperature range 7.0-75.2 "C.
TABLE 2: Temperature and pH Dependence of Measured Rate Constants for Hydrogen Atom Reaction with Iodate [HC1041, temp, km, fraction fraction M "C dm3mo1-ls-l 103HI03 4.75 x 106 0.255 0.745 1.oo 11.5 1.00 x 107 0.188 0.812 1.oo 25.0 9.47 x lo6 0.188 0.812 1.oo 25.0 1.89 x 107 0.142 0.858 1.oo 43.0 3.97 x 107 0.195 0.805 1.oo 61.0 6.68 x lo7 0.286 0.714 1.00 75.0 6.68 x lo6 0.788 0.212 0.10 10.5 1.17 x 107 0.698 0.302 0.10 25.0 1.32 x 107 0.698 0.302 0.10 25.0 2.57 x 107 0.623 0.377 0.10 43.5 2.46 x 107 0.623 0.377 0.10 43.5 4.79 x 107 0.708 0.292 0.10 61.3 0.10 76.3 1.07 x lo* 0.815 0.185 7.27 x lo6 1.Ooo 4.60" 7.0 1.18 x 107 1.Ooo 4.60" 17.5 4.60" 25.0 1.22 x 107 1.Ooo 2.95 x 107 1.Ooo 4.65" 46.0 7.18 x 107 1.om 4.83" 75.0 Measured pH values.
Iodate. The limiting, zero ionic strength, acidity constant for HI03 has been determined at room temperature as 0.157,27 which means that in phosphate buffered solution (pH 4.6) the only reaction would be
-
'H
+ IO,-
-
products
(19)
Figure 3a shows a typical scavenging plot obtained for 25 "C, with the measured scavenging rate constant being km = (1.22 f 0.04) x lo7 dm3mol-' s-l. This measurement was repeated at temperatures over the range 7.0-75.2 "C, and the calculated Arrhenius plot is given in Figure 3b. All the experimental results and conditions are summarized in Table 2. The temperature dependence of this reaction is well characterized by the equation
k,, = 9.17 x 10" exp(-27460/RI*) The speciation of iodate under these conditions is also of importance. The presence of polymerized iodate species has previously been suggested;28however, a thorough review of the literature by Pethybridge and Prue27has found evidence only for the dimerization equilibrium
12598 J. Phys. Chem., Vol. 98, No. 48, 1994
IO,-
+ HIO, == H(IO,),-
(20)
which has an equilibrium constant of K E 4 dm3 mol-' and hence would be negligible under these conditions. Therefore, we believe that we measure only the 1 0 3 - reaction. The potential mechanism of hydrogen atom reaction with iodate are analogous to reactions 8-1 1 for periodate. However, 0- abstraction from IO3- to give the 'IO2 radical does not occur, as this product radical would have been observed in the transient absorption experiments performed previo~sly.~ Likewise, an electron transfer reaction to give '103,- can be discounted. Longitudinal magnetization experiments again showed that the addition of iodate attenuated the signal as quickly as predicted from the FIJI measured rate constant, indicating an irreversible reaction on the microsecond time scale. By analogy with the periodate system, we believe that the overall reaction is
+ IO,-
'H
-
'OH
+ IO,-
(21)
which occurs directly or through an HIO3- intermediate. Both the HI03- and 1 0 2 - species have been proposed previously7 as intermediates in the steady state radiolysis formation of I2 from irradiated iodate/methane solutions. Given a pKa of -0.8, the direct measurement of the rate constant for 'H atom reaction with iodic acid 'H
+ HIO, - products K23
+ IO,-
g
0
0.2 -
m
e
. 0.1 -
$ n Q
0.0
-
I
I
.
I
I
I
220
200
.
l
.
260
240
I
I
I
I
300
280
Wavelength (nm)
Figure 4. Experimental absorption spectrum of 1.011 x M KIO3 at 0.10 M HC104 and theoretical fit based on fractional composite fits of the HI03 and 1 0 3 - spectra.
(22)
was not possible in dilute acid. Moreover, the temperature dependence of the equilibrium
HIO, =H+
-
0.3
(23)
has only been studied over a limited range to date.27.29-32Thu S,
FID attenuation measurements were performed in 0.10 and 1.00 M HC104, and the rate coefficients were calculated using auxiliary measurements of the fractions of the two iodate species present over the temperature range of this study. Absorption spectra of 103- were measured in 5.0 x lop3 M phosphate solutions (pH 4.6) over the temperature range 2.1-70.3 "C. These measurements were then repeated using the stock HC104 solution, [HC104] 11.7 M, where the spectra were assumed to correspond to just HI03 absorbance. Based on the literature Ka value of 0.157,27and its expected ionic strength behavior,28 in the stock HC104 solution at 25 "C the 1 0 3 - fraction should be '2%. Finally, absorption measurements were made in 0.10 and 1.00M HC104, and the relative fractions of the two limiting forms of iodate were fitted to the observed spectra using a linear regression technique. A typical example of the fit, at 23.3 "C, is seen in Figure 4. Although all of the spectra were somewhat featureless, this procedure was found to give remarkably consistent results over the entire temperature range studied. These values are given in Table 2. Evaluation of the exact K23 value is not necessary for the rate constant determination. However, limiting zero ionic strength equilibrium coefficients are currently being determined and will be reported in a later publication. In 0.10 M HC104 the hydrogen atoms can react via both reactions 19 and 22, and thus the overall FID scavenging decay rate constant, k,, (Table 2), under these conditions is given by
-
N
where A19 and E19 are the Arrhenius parameters for IO3reaction. Using the measured temperature-dependent equilib-
::.:I 15 0 2 80
,
,
3 00
,
,
,
3 20
, 3 40
,
, 3 60
I O 3 / Temperature (K)
Figure 5. Arrhenius plot of In k , vs U T for reaction of hydrogen atoms at 0.10 (0)and 1.00 M (M)HC104 with iodate over the temperature range 10.5-76.3 "C. Solid lines are calculated values based on derived Arrhenius parameters for 1 0 3 - and HI03 reaction in eq 24. rium fractions, eq 24 can be globally fitted to all the rate constant data given in Table 2, to obtain best fit parameters for reaction 22 of
k,, = 2.01 x 1013exp(-36550/RT) At 25 "C, this corresponds to a scavenging rate constant of 8.2 x lo6 dm3 mol-' s-l, The curves generated by using these values in eq 24 are shown in Figure 5 , in comparison to the experimental data at 0.10 and 1.00 M HC104. Apart from the highest temperature value at 0.10 M HC104, which was not included in the global fit, excellent agreement is obtained. As the formation of *IO3is thermodynamically unfavorable,22 and kinetically slow? we believe that the mechanism of reaction for hydrogen atoms with H I 0 3 is not hydrogen abstraction but is similar to that for IO3-, i.e. 'H
+ HIO, - 'OH + HIO,
(25)
Conclusion The Arrhenius parameters for the reaction of the hydrogen atom with Io,-, HIO3, IO4-, and H5106 have been evaluated. The direct measurement of the scavenging rate constants for iodate and periodate ions at pH 4.6 has shown that these
Reaction of H with
104-
and IO3-
temperature-dependent values are given by k,, = 9.17 x 10" exp(-27460/RT) and
4 = 6.71 x
lo', exp(-22350/RT)
respectively, with the calculated rate constants in units of dm3 mol-' s-l. By spectrophotometric measurement of the temperature dependence of the acidity constant for HI03, the rate constant for the hydrogen atom reaction with this species was calculated to be k,, = 2.01 x 1013exp(-36550/RT) An analogous treatment of the data for IO4-,using literature values for all of the temperature-dependent equilibrium constants, gave the expression k13
= 5.90 x 1017exp(-53000/RT)
for the H5106 species. This last reaction, characterized by an extremely large preexponential factor of 5.90 x lo1' dm3mol-' s-l, almost certainly involves dehydration and rearrangement of the octahedral 1" complex in the transition state, leading to the pyrimidal Iv product.
Acknowledgment. The authors thank Dr.David Werst for his assistance in operating and maintaining the Van de Graaff accelerator. We also thank Joan Hnatiw for her technical assistance with the HI03/I03- equilibrium measurements. References and Notes (1) Thompson, T. J.; Beckerley, J. G. The Technology of Nuclear Reactor Safety; h4IT Press: Cambridge, 1973. (2) Cambell, D. 0.;Malinauskas, A. P.; Stratton, W. R. Nucl. Technol. 1981,53, 111. (3) Cubicciotti, D.; Sehgal, B. R. Nucl. Technol. 1984,65,266. (4) Proceedings of the 1st CSNl Workshop on Iodine Chemistry in Reactor Safety; Deane, A. M., Potter, P. E., Eds.; Harwell Research Report, AERE-R 11974, 1986.
J. Phys. Chem., Vol. 98, No. 48, 1994 12599 ( 5 ) Proceedings of the Second CSNI Workshop on Iodine Chemistry in Reactor Safety; Vikis, A. C., Ed.; Atomic Energy of Canada Ltd Research Report, AECL-9923, 1989. ( 6 ) Proceedings of the 3rd CSNI Workshop on Iodine Chemistry in Reactor Safety; Ishigure, K., Saeki, M., Soda, K., Sugimoto, J., Eds.; JAE Research Report, JAERI-M 92-012, 1992. (7) Paquette, J.; Ford, B. L. Radiar. Phys. Chem. 1990,36,353. (8) Sellers, R. M. Central Electricity Generating Board, Report No. TPRD/B/0688/R85, 1986. (9) Mezyk, S. P.; Elliot, A. J. J . Chem. SOC.,Faraday Trans. 1 1994, 6,831. (10) Fessenden, R. W.; Verma, N. C. Faraday Discuss. Chem. SOC. 1977, 63,104. (11) Bartels, D. M.; Craw, M. T.; Han, P.; Trifunac, A. D. J. Phys. Chem. 1989,93,2412. (12) Han, P.; Bartels, D. M. Chem. Phys. Lett. 1989,159, 538. (13) Han, P.; Bartels, D. M. J . Phys. Chem. 1990,94,7294. (14)Han, P.; Bartels, D. M. J . Phys. Chem. 1992,96,4899. (15) Roduner, E.; Bartels, D. M. Ber. Bunsen-Ges. Phys. Chem. 1992, 96,1037. (16) Bartels, D. M.; Mezyk, S. P. J. Phys. Chem. 1993,97,4101. (17) Dushman, S. J . Phys. Chem. 1904,8,453. (18) Sutton, H. C.; Adams, G. E.; Boag, J. W.; Michael, B. D. Pulse Radiolysis; Ebert, M., Keene, J. P., Swallow, A. J., Baxendale, J. H., Eds.; Academic Press: New York, 1965; pp 61-81. (19) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J . Phys. Chem. Re$ Data 1988,17,513. (20) Crouthamel, C. E.; Hayes, A. M.; Martin, D. S. J . Am. Chem. SOC. 1951,73,82. (21) Buist, G. J.; Hipperson, C. P.; Lewis, J. D. J. Chem. SOC. A 1969, 307. (22) K l h n g , U. K.; Sehested, K.; Wolff, T. J . Chem. Soc., Faraday Trans. I 1981,77, 1707. (23) Symons, M. C. R. J . Chem. SOC. 1955,2794. (24) Laidler, K.J. Chemical Kinetics; McGraw-Hill: New York, 1965; p 207. (25) Harned, H. S.; Owen, B. B. The Physical Chemistry of Electrolyte Solutions; Rheinhold Publishing: New York, 1950; p 397. (26) Clever, H. L. In IUPAC Solubility Data Series; Battino, R., Ed.; Pergamon: New York, 1981; Vol. 3, p 25. (27) Pethybridge, A. D.; Rue, J. E. Trans. Faraday SOC.1967,63,2019. (28) Sidgwick, N. V. The Chemical Elements and Their Compounds; Oxford University Press: Oxford, 1950; pp 1228-9. (29) Naidich, S . ; Ricci, J. E. J . Am. Chem. SOC.1939,61, 3268. (30) Hood, G. C.; Jones, A. C.; Reilly, C. A. J . Phys. Chem. 1959,63, 101. (31) Li, N. C. C.; Lo, Y. T. J . Am. Chem. SOC.1941,63,397. (32) Abel, E.; Redlich, 0.; Hersch, P. Z. Phys. Chem. 1934,A170, 112.
