Journal of The Electrochemical Society, 154 共4兲 F70-F76 共2007兲
F70
0013-4651/2007/154共4兲/F70/7/$20.00 © The Electrochemical Society
Interaction of Aqueous Iodine Species with Ag2O/Ag Surfaces
X. Zhang,* S. Stewart, D. W. Shoesmith,** and J. C. Wren**,z
Department of Chemistry, University of Western Ontario, London, Ontario, Canada
The chemical conversion of Ag2O films on Ag surfaces to AgI in aqueous iodide solutions has been studied electrochemically.
Ag2O films were grown potentiostatically and then exposed to I− solutions. The chemical conversion process was followed at
open-circuit potential 共EOC兲 using cathodic stripping voltammetry performed after various exposure periods. The EOC showed a
sudden drop at the completion of the conversion of Ag2O to AgI, reaching a steady-state value close to the equilibrium potential
for AgI/Ag and the iodide solution. This sudden drop in EOC allowed easy determination of the total reaction time required for
complete conversion of Ag2O to AgI. Distinctly separated current peaks were observed for the cathodic reduction of Ag2O and
AgI to Ag, and the charges associated with these peaks provided a measure of the amount of Ag2O converted. The conversion
reaction was 100% efficient. The total reaction times from the EOC measurements and the cathodic stripping results were used to
determine the reaction order and rate constant required for the development of nuclear reactor safety assessment codes.
© 2007 The Electrochemical Society. 关DOI: 10.1149/1.2435696兴 All rights reserved.
Manuscript submitted September 19, 2006; revised manuscript received November 27, 2006.
Available electronically February 13, 2007.
One of the safety issues of nuclear power plants is the potential
radiation dose to the public in the unlikely event of a severe accident. From the perspective of public safety, radioiodine is one of the
most important fission products from the uranium fuel because of its
large fuel inventory, high volatility, and radiological hazard. Such an
incident would lead to fuel and fuel channel damage, and it is assumed that a significant fraction of the radioiodine fuel inventory
would be released from the reactor core into the containment
building.1-3 It has been established that most of the released iodine
would quickly dissolve and remain in the water ubiquitous throughout the containment building following an accident.4 However, a
small fraction could be released to the gas phase due to the continuous conversion of nonvolatile to volatile iodine species under the
highly oxidizing conditions prevailing in the presence of ionizing
radiation.5 Because of its mobility, the gaseous iodine concentration
is a critical parameter for safety assessment and postaccident management.
One reaction of interest in assessing iodine volatility is that between aqueous iodine species and silver surfaces. The control rod
共for neutron flux兲 in some nuclear reactors is made of silvercadmium, and silver is assumed to be released from the reactor core
into the containment building environment in some accident scenarios, and its reaction with aqueous iodine to form insoluble silver
iodide could dominate iodine behavior. In fact, a significant reduction in iodine volatility in the presence of silver has been observed
in engineering-scale experiments simulating postaccident containment conditions.6,7
In the highly oxidizing and ionizing radiation conditions expected, Ag could enter the aqueous phase in either the metallic or
oxidized 共AgOH, Ag2O, and AgO兲 forms. The availability of aqueous iodine species, mainly I− and I2, would vary depending on radiation dose rate, pH, temperature, and the presence of impurities.5
Thus, it is important to establish the mechanism and kinetics of
individual iodine reactions with silver to determine iodine volatility.
Early studies on silver-iodine interactions involved measurements of overall iodine uptake on silver as a function of pH and the
extent of initial Ag oxidation, in either the presence or absence of
radiation,6-9 and suggested reaction involves the interaction of I−
with Ag2O
Ag2O + 2I− + 2H+ → 2AgI + H2O
关1兴
and the reaction of I2 with both metallic Ag and Ag2O
2Ag + I2 → 2AgI
* Electrochemical Society Student Member.
** Electrochemical Society Active Member.
z
E-mail: [email protected]
关2兴
1
Ag2O + I2 → 2AgI + 2 O2
关3兴
However, these studies have not unambiguously established detailed
mechanisms and kinetics of the individual reactions.
A complication is the pseudo steady state that exists between I−
and I2 in the aqueous solution6
radiation
I− ↔ I2
关4兴
I2 + H2O HOI + I− + H+
关5兴
I2 + I− I−3
关6兴
Equilibria 5 and 6 are achieved thermally even in the absence of
radiation. Furthermore, the effects of oxide and iodide film morphologies on the reaction kinetics have not been examined.
In this study, Ag2O film growth on Ag was controlled, and the
effect of the film on the reaction of I− with Ag2O 共Reaction 1兲
examined, using electrochemical and surface analytical techniques.
The reduction potentials for AgI/Ag and Ag2O/Ag are well separated and the open-circuit potentials of Ag and Ag2O in I− solutions
very different, allowing the kinetics of the chemical reaction 1 to be
followed by electrochemical methods. To our knowledge, this is the
first time electrochemical techniques have been used to quantitatively monitor the kinetics of an aqueous anion induced chemical
transformation between two insoluble solids. From these measurements, the rates of individual reactions required for the development
of nuclear reactor safety assessment codes can be extracted.
Experimental
Electrochemical cell and electrodes.— A three-electrode system,
consisting of a silver working electrode, a reference electrode, and a
counter electrode, was used for all experiments. The working electrode was a 7 mm 共in diameter兲 Ag disk, set in resin so that only the
flat front face was exposed to the solution. Prior to each experiment,
this electrode was manually polished with 600 and 800 grit silicon
carbide papers, and polishing residue was removed by sonication in
an acetone/methanol mixture for 5 min. The reference electrode was
a saturated calomel electrode 共SCE兲 and the counter electrode was
platinum mesh. All potential measurements were made with, and are
quoted against, an SCE. A Solartron model 1240 potentiostat was
used to control and measure potentials and to record current responses. Corrware and Corrview software 共supplied by Scribner and
Associates兲 was used to control experiments and analyze data.
Solutions.— Experiments were conducted at room temperature in
Ar-purged 0.02 M NaH2PO4 with the pH adjusted to 12 with NaOH.
Phosphate solutions were used to control the local pH within pores
in the growing oxide film. Solutions were prepared with water purified using a NANOpure Diamond UV ultrapure water system from
Journal of The Electrochemical Society, 154 共4兲 F70-F76 共2007兲
F71
共Ee兲Ag2O/Ag = 0.927 − 0.0592 pH 共vs SCE兲 = 0.216 V 共vs SCE兲
关8兴
The vertical dotted line in Fig. 1 indicates this potential value.
The charges associated with oxide formation 共QA兲 and reduction
共QC兲, obtained by integration of the voltammetric currents, are
equal; i.e., the oxide formation is fully reversible, and Ag is not lost
to dissolution during oxide formation. This is consistent with the
fact that the solubility of Ag2O is at a minimum at pH 12.11
The anodic current on the forward scan shows four distinct
stages of behavior, consistent with the published behavior in highly
alkaline solutions.12-16 The small anodic current in stage 1 at potentials less than 共Ee兲Ag2O/Ag, has been attributed to the chemisorption
process
Ag + OH−共ad兲 AgOH + e−
Figure 1. Cyclic voltammogram recorded at 5 mV s−1 on a Ag electrode in
Ar-purged 0.02 M NaH2PO4 solutions at pH 12 to an anodic scan limit of
0.6 V. The vertical broken line indicates the equilibrium potential for
Ag2O/Ag at pH 12. The straight lines on the current–potential curves are the
slopes with a value of 58 ⍀.
Barnstead International to remove organic and inorganic impurities.
Prepared in this manner, the water had a resistivity of 18.2 M⍀ cm.
Experimental procedure.— All experiments were preceded by a
cathodic cleaning of the electrode at a potential of −1.1 V for 300 s.
Cyclic voltammetric experiments were conducted from the cathodic
cleaning potential to various anodic limits at a scan rate of
5 mV s−1.
After cathodic cleaning, an Ag2O film was grown potentiostatically at a potential in the range +0.2 to +0.6 V until the total oxidation charge reached a desired value, generally in the range
0.01–0.2 C 共or 0.026–0.52 C cm−2兲. The Ag2O-covered electrode
was then transferred to a 0.02 M NaH2PO4 solution containing KI at
a concentration in the range 10−4–10−3 mol dm−3 共M兲. The progress
of the reaction between Ag2O and I− was monitored by measuring
the open-circuit potential 共EOC兲 as a function of time. At various
reaction times, the reaction was terminated by transferring the electrode from the KI solution back to the KI-free solution. The potential was then immediately scanned from EOC to −1.1 V and the
currents for the reduction of residual Ag2O and the reaction product,
AgI, measured.
The morphology of silver electrode surfaces at various stages of
the reaction were obtained with a Leo 440 scanning electron microscope and the composition of products analyzed with energydispersive X-ray 共EDX兲 spectroscopy.
Electrochemical properties of Ag2O/Ag and AgI/Ag.— A cyclic
voltammogram recorded on Ag from −0.6 to 0.6 V in Ar-purged
aqueous solutions 共pH 12兲 free of KI is shown in Fig. 1. The form of
the voltammogram is very similar to that observed in borate buffered solutions, confirming that pH control was maintained throughout the film growth process.10 The oxide formation/reduction process can be considered reversible, since the potential for the onset of
oxide formation on the forward scan, and reduction on the reverse
scan, both occur very close to the equilibrium potential 共Ee兲 for the
reaction
given by the Nernst equation 共at pH 12兲
In stage 2, where potentials are slightly more positive than
共Ee兲Ag2O/Ag, the current rises steeply with potential. If the scan was
reversed in this region, the current on the reverse scan is larger than
that on the forward scan, consistent with a rapid oxide nucleation
and growth process.
In stage 3, the anodic current increases more slowly and, for a
short range, indicated by the tangential line, a near linear i − E
relationship is observed at potentials ⬍+0.3 V. A similar linear i
− E relationship is maintained during the cathodic scan. The slope
of this line yields a resistance value of ⬃58 ⍀, which is close to the
resistance of the solution, suggesting film growth is controlled by
ion migration in water-filled pores.17 The eventual occurrence of the
current peak shows that the thickening of the layer eventually impedes its growth. This behavior is similar to that observed in borate
buffered solutions.10 If the voltammogram is recorded in the absence
of the phosphate buffer, then the current continues to increase with
increasing potential, indicating the presence of the film does not
impede its further growth. This is most likely due to pH variations
within the pores of the growing film.
The current recovery for potentials ⱖ0.5 V suggests an open
structure allowing the three-dimensional 共3D兲 growth of an outer
layer, in agreement with previous claims that a dual-layer porous
oxide is produced.12 In this high potential range, the oxidation of
Ag2O to AgO is also possible, but galvanostatic reduction of the
potentiostatically grown film at +0.6 V 共these results are not shown兲
does not show behavior that can be clearly identified as the reduction of AgO. The variability of the current peak around +0.4 V from
one scan to another is also an indication of the change in transport
behavior rather than the oxidation behavior.
Figure 2 shows voltammograms recorded in aqueous solutions
containing 10−4 M and 10−3 M KI with the anodic limit confined to
a value less than 共Ee兲Ag2O/Ag to avoid the possibility of oxide formation. The onset of AgI formation
Ag + I− AgI + e−
关10兴
occurs at potentials immediately above the equilibrium potential,
共Ee兲AgI/Ag, given by
共Ee兲AgI/Ag = −0.396 − 0.0592 log关I−兴 共V vs SCE兲
Results
2Ag + 2 OH− Ag2O + 2e− + H2O
关9兴
关7兴
= −0.159 V 共vs SCE兲 共at 10−4 MI−兲
= −0.218 V 共vs SCE兲 共at 10−3 MI−兲
关11兴
The potential-independent current at positive potentials can be
attributed to control of AgI formation by I− diffusion in solution. No
attempt was made to control convective conditions: the noise shows
the influence of Ar purging. No current peak for I− adsorption18,19
was observed, probably due to the lower I− concentration and higher
pH.
The onset of AgI reduction at potentials very close to 共Ee兲AgI/Ag,
and the equality in anodic 共AgI formation兲 and cathodic 共AgI reduction兲 charges, confirms that the AgI/Ag reaction is as 共more兲 reversible as 共than兲 the Ag2O/Ag reaction. The linear i − E relationship
F72
Journal of The Electrochemical Society, 154 共4兲 F70-F76 共2007兲
Figure 4. SEM images of oxide-covered Ag electrodes. The films were
grown potentiostatically at 0.6 V for 共a兲 700 and 共b兲 1300 s, or to total
anodic charges of 0.11 C 共or 0.28 C cm−2兲 and 0.2 C 共or 0.52 C cm−2兲, respectively.
Figure 2. Cyclic voltammograms recorded at 5 mV s−1 on a Ag electrode in
Ar-purged 0.02 M NaH2PO4 solutions containing I− at pH 12 to an anodic
scan limit of 0.15 V, for 关I−兴 of 10−4 and 10−3 M. The vertical broken lines
indicate the equilibrium potentials for AgI/Ag for the two 关I−兴.
for AgI reduction 共−0.25 V to the reduction peak兲 suggests a reaction involving ion migration. The slope of this relationship is
⬃150 ⍀, which is slightly greater than the solution resistance, suggesting the reaction is partially controlled by the resistance of the
pore structure in the AgI film.
Extension of the forward scan to +0.4 V shows both AgI and
Ag2O formation 共Fig. 3兲. The observation of two cathodic reduction
peaks at the potentials for Ag2O 共Fig. 2兲 and AgI 共Fig. 3兲 reduction
confirms both phases are present. Once the cathodic reduction of
Ag2O is complete, the current returns to the diffusion-limited value
for AgI anodic formation 共in the potential range −0.05 to −0.2 V兲,
indicating that the two film formation processes may occur independently.
Potentiostatic growth of initial Ag2O films.— To provide a starting point for the study of I− with Ag2O, a known amount of oxide
was grown potentiostatically 共at 共EAg2O兲0兲 on Ag in an Ar-purged
KI-free solution. The anodic charge associated with oxide growth
was obtained by integrating the current–time transient recorded at
共EAg2O兲0. Since negligible oxide dissolution occurs during film
growth at pH 12, this anodic charge, referred to as 共QAg2O兲0, is the
initial inventory of oxide present prior to exposure to KI solution.
Figure 3. Cyclic voltammograms recorded at 5 mV s−1 on a Ag electrode in
Ar-purged 0.02 M NaH2PO4 solutions containing I− at pH 12 to an anodic
scan limit of 0.4 V, for 关I−兴 of 10−4 and 10−3 M.
Figure 4 shows scanning electron microscopy 共SEM兲 images of
potentiostatically grown Ag2O films after 700 and 1300 s 共or total
charges of 0.11 and 0.2 C, respectively兲. The film grown for 700 s is
composed of highly structured, tetrahedral crystals. The film grown
for 1300 s is less structured with poorly defined crystals. This
change in morphology is consistent with a transition from the 2D
growth of a compact inner layer to the 3D growth of a more porous
outer layer. Individual crystals ranged in size from 0.3 to 1 ␮m,
depending on the growth potential and thickness. Different morphologies and crystal structures, depending on the potential applied,
have been observed by others.20
EDX analysis and X-ray diffraction analysis 共XRD兲 show that
incorporation of phosphate into the oxide occurs. However, even
after extensive film growth under potentiostatic conditions, the potential established on switching to open circuit 共EOC兲 always relaxes
to 共Ee兲Ag2O/Ag 共see below兲, consistent with exposure of an Ag/Ag2O
interface to the pH 12 solution. Although presently unconfirmed,
this suggests that despite the incorporation of phosphate, the film is
behaving as an oxide.
Open-circuit potential behavior.— After immersion in an aqueous solution containing KI, the open-circuit potential 共EOC兲 of the
Ag/Ag2O electrode was followed as a function of reaction time, trxn.
The dependence of the reaction kinetics on 关I−兴 was studied on an
oxide film grown until 共QAg2O兲0 = 0.2 C at 共EAg2O兲0 = 0.6 V 共Fig.
5a兲, and the influence of the initial oxide inventory, for films grown
at 0.6 V, was studied in 5 ⫻ 10−4 M KI solution 共Fig. 5b兲.
The initial value of EOC was independent of 关I−兴 共Fig. 5a兲, varied
only marginally for different initial film thicknesses 共Fig. 5b兲, and
was very close to 共Ee兲Ag2O/Ag. With time, EOC slowly decreased to a
second plateau at a time dependent on the initial film thickness
before undergoing a final abrupt transition to a much lower steadystate value. The time taken to reach the final transition decreased as
关I−兴 increased and increased as 共QAg2O兲0 increased. The final steadystate value of EOC is close to 共Ee兲AgI/Ag, as shown in Fig. 6 by a
comparison of final EOC values to calculated values of 共Ee兲AgI/Ag.
Clearly, once the final potential transition has occurred, redox conditions on the Ag surface are dominated by the AgI/Ag reaction.
Amounts of reactant Ag2O and product AgI as a function of
reaction time.— There are two possible mechanisms by which oxide
can be converted to iodide: 共i兲 a direct chemical process via Reaction 1, and 共ii兲 a galvanically coupled process in which the cathodic
reduction of oxide to Ag is coupled to the anodic oxidation of Ag to
iodide. Here we concentrate on the overall conversion process, and a
detailed discussion of mechanistic details will be published elsewhere.
Figure 7a shows the times on the open-circuit potential transient
when the reaction was stopped, the electrode transferred to a KI-free
solution, and a cathodic stripping voltammogram 共CSV兲 recorded.
Figure 7b shows CSVs recorded for the reaction times indicated in
Fig. 7a. The progress of the reaction is clearly monitored, the peak
for AgI reduction increasing as that for Ag2O reduction decreases
Journal of The Electrochemical Society, 154 共4兲 F70-F76 共2007兲
F73
Integration of the two reduction peaks in Fig. 7b yielded values
of 共QAg2O兲t and 共QAgI兲t as a function of reaction time, trxn 共Fig. 8兲.
The values are inversely related, and 共QAg2O兲t = 0 at ␶ f , further confirming there is no residual oxide on the Ag surface at t = ␶ f . The
relationship
共QAg2O兲t + 共QAgI兲t = 共QAg2O兲0
关12兴
is maintained throughout the potential transient, confirming the
100% efficiency of the conversion reaction, as well as the absence of
both Ag2O dissolution and the production of AgI by reaction between Ag and I−. The slopes of the two plots can be used to determine reaction rate constants, as described below.
Figure 9 shows SEM images of the surface prior to the transfer
of the electrode to an I− solution 共b兲, after the first short potential
transition 共c兲, and on completion of the reaction 共t = ␶f兲 共d兲. As the
reaction proceeds the surface becomes progressively covered by
small particulates of AgI. The micrograph in c shows the coexistence of oxide and iodide particles, confirming that AgI does not
form a protective layer and that the underlying Ag2O is continuously
exposed.
Reaction Kinetic Analysis
Reaction orders.— The rate equation for the Ag2O reaction with
I− 共Reaction 1兲 can be expressed as
d共mAg2O兲t
d共mAgI兲t
= −2
= k共10−3 ⫻ 关I−兴t兲 p共AAg2O兲tq
dt
dt
Figure 5. Open-circuit potential 共EOC兲 as a function of reaction time and its
dependence on 关I−兴 and 共QAg2O兲0. The reaction conditions were 共a兲 oxides
grown at 0.6 V to a constant anodic charge 共QAg2O兲0 = 0.2 C, with 关I−兴 ranging from 10−4 to 5 ⫻ 10−3 M, and 共b兲 oxides grown at 共EAg2O兲0 = 0.6 V, up
to 共QAg2O兲0 ranging from 0.01 to 0.2 C, with 关I−兴 constant at 5 ⫻ 10−4 M.
with reaction time. Once the abrupt transition to the final steadystate value of EOC occurred, no oxide remained on the electrode
surface and no further AgI was formed, confirming that this transition signaled the total conversion of oxide to iodide. The time of
occurrence of the transition, ␶ f , is a measure of the time to complete
the conversion.
关13兴
where 共mAg2O兲t and 共mAgI兲t are the number of moles of Ag2O and
AgI at reaction time t; 关I−兴t 共in mol dm−3兲 and 共AAg2O兲t 共in cm2兲 are
the concentration of I− and the surface area of Ag2O at time t,
respectively; k is the reaction rate constant; and superscripts, p and
q, represent reaction orders with respect to 关I−兴 and 共AAg2O兲, respectively. The units of k will depend on p and q, and the factor of 10−3
is included to express 关I−兴t in mol dm−3 共or M兲.
Because the Ag2O surface area to solution volume ratio is very
small, 关I−兴t is constant. The number of moles of unreacted Ag2O and
AgI formed are given by the cathodic charges for reduction of Ag2O
to Ag, 共AAg2O兲t, and AgI to Ag, 共QAgI兲, by
共mAg2O兲t =
共mAgI兲t =
共QAg2O兲t
2F
共QAgI兲t
F
关14兴
where F is the Faraday constant, and Eq. 13 becomes
d共QAg2O兲t
d共QAgI兲t
=−
= kF共10−3 ⫻ 关I−兴0兲 p共AAg2O兲tq
dt
dt
关15兴
Figure 8 shows that this rate does not vary appreciably as the
reaction proceeds, suggesting 共AAg2O兲t = 共AAg2O兲0. This constant
surface area is also indicated by the micrographs in Fig. 9. The rate
equation can, therefore, be further reduced to
d共QAgI兲t
= kF共10−3 ⫻ 关I−兴0兲 p共AAg2O兲q0
dt
关16兴
and integrated to yield
Figure 6. Final steady-state open-circuit potential 共symbols兲 observed for
various 关I−兴 compared to calculated equilibrium potentials for the Ag
+ I− AgI + e− reaction 共line兲.
共QAgI兲t = kF共10−3 ⫻ 关I−兴0兲 p共AAg2O兲q0t
The total reaction time, ␶ f , is then determined by
关17兴
F74
Journal of The Electrochemical Society, 154 共4兲 F70-F76 共2007兲
Figure 7. CSV as a function of reaction
time 共trxn兲. The times when the CSVs were
taken are indicated on the open-circuit potential plot in 共a兲, and the corresponding
CSVs are shown in 共b兲. The experimental
conditions
were
共QAg2O兲0 = 0.2 C
共0.52 C cm−2兲 at 0.6 V, and 关I−兴 = 5
⫻ 10−4 M.
␶f =
共QAgI兲 f
−3
kF共10
⫻ 关I 兴0兲
−
p
共AAg2O兲q0
=
共QAg2O兲0
−3
kF共10
⫻ 关I−兴0兲 p共AAg2O兲q0
关18兴
Figure 8. Cathodic charges measured for the reduction of unreacted Ag2O
and product AgI at various reaction times, trxn. The experimental conditions
were 共QAg2O兲0 = 0.2 C 共0.52 C cm−2兲 at 0.6 V, and 关I−兴 = 5 ⫻ 10−4 M.
where 共QAgI兲 f = 共QAg2O兲0, since all the initially available Ag2O is
converted to AgI. The slope of a plot of ␶ f vs log 关I−兴 yields a value
of p = 1 共Fig. 10兲.
Similarly, a plot of log ␶ f vs log共AAg2O兲0 would yield the negative of the reaction order, q. Since the surface area of the initial
oxide film was not determined, a value of q = 0.82 was obtained
from a plot of ␶ f vs log共AAg2O兲0 共Fig. 11兲. If the initial oxide grew
uniformly on the electrode, then the oxide surface area would be
equal to its geometric area. Then, 共AAg2O兲0 would be independent of
共QAg2O兲0, and the film thickness directly proportional to 共QAg2O兲0.
However, the SEM micrograph in Fig. 4 shows three-dimensional
oxide growth, and consequently, the surface area of the initial Ag2O
film is expected to have a small dependence on 共QAg2O兲0
Figure 9. SEM images of the electrode
obtained after various reaction times. The
times when the reaction was terminated,
and the SEM images taken, are shown in
共a兲, and the corresponding micrographs
are shown in 共b–d兲.
Journal of The Electrochemical Society, 154 共4兲 F70-F76 共2007兲
F75
Rate constant.— Since the true surface area of Ag2O has not
been determined, the apparent rate constant, kapp, was obtained using
the geometric surface area of the electrode
d共mAg2O兲t
d共mAgI兲t
=2
= kapp共10−3 ⫻ 关I−兴0兲AElec
dt
dt
where kapp = kC␤
冉
共QAg2O兲0
F
冊
关22兴
␤
关23兴
For small values of ␤, kapp approaches k.
A value for kapp can be extracted from the total reaction time
observed as a function of 关I−兴. From Eq. 18 共p = q = 1兲
␶f =
Figure 10. Log of 关I−兴 vs log of total reaction time, ␶ f ; the initial
共QAg2O兲0 = 0.2 C at 0.6 V.
共AAg2O兲0 = C␤
冉
共QAg2O兲0
F
冊
=
kF共10
关19兴
AElec
q
kC␤q 共10−3 ⫻ 关I−兴0兲 pAElec
冉
冊
关24兴
yielding a value of kapp = 5.4 ⫻ 10−3 cm s−1.
Alternatively, kapp can be obtained from 共QAg2O兲t and 共QAgI兲t as
a function of reaction time. Equation 22 yields
关25兴
Hence, Fig. 7 yields kapp = 3.8 ⫻ 10 cm s , in good agreement
共within 50%兲 with the kapp value obtained from the total reaction
time.
A more rigorous kinetic study to determine the rate constant
more accurately is presently underway. This includes a determination of whether the incorporation of anions into the electrochemically grown films exerts a significant influence on the kinetics of the
film conversion reaction.
−3
−1
Conclusions
⫻ 关I−兴0兲 p共AAg2O兲q0
1
冉
共QAg2O兲0
1
−
⫻ 关I 兴0兲AElec
F
␤
共QAg2O兲0
−3
kapp 共10
−3
共QAgI兲t = 共共QAg2O兲0 − 共QAg2O兲t兲 = kappF共10−3 ⫻ 关I−兴0兲AElect
where C␤ is a constant or a unit conversion factor. The value of ␤
would be zero and C␤关共QAg2O兲0/F兴␤⫽1 if the film grew uniformly
and featurelessly. If the film grew hemispherically the value of ␤
would be 2/3. The SEM micrographs of the oxide films indicate that
the value lies between 0 and 2/3. Substituting Eq. 19 into Eq. 18
yields
␶f =
1
共QAg2O兲0
F
冊
1−q␤
关20兴
and a plot of log ␶ f vs log共QAg2O兲0 would have a slope equal to 1
− q␤.
Despite this small surface area dependence, it is reasonable to
accept a first-order dependence on the surface area, 共AAg2O兲0, and
the rate law 共Eq. 13兲 can be written
d共mAg2O兲t
d共mAgI兲t
=2
= k共10−3 ⫻ 关I−兴t兲共AAg2O兲0
dt
dt
关21兴
Figure 11. Log of total anodic charge used to grow the initial silver oxide
film, 共QAg2O兲0, vs log of total reaction time, ␶ f ; 关I−兴 = 5 ⫻ 10−4 M.
The chemical conversion of Ag2O by I− to AgI in aqueous iodide
solutions 共pH 12兲 has been studied by following the reaction on
open circuit and periodically determining the amounts of oxide and
iodide on the surface using CSV. The reaction was 100% efficient
since both the oxide and the iodide are very insoluble under the
experimental conditions employed.
Complete conversion was clearly indicated by an abrupt transition in the open-circuit potential from a value close to the equilibrium potential for Ag2O/Ag to a value identical to the equilibrium
potential for AgI/Ag. This abrupt change in EOC allowed easy determination of the total reaction time required for complete conversion of Ag2O to AgI. This time to complete the reaction was inversely proportional to the iodide concentration and was nearly
proportional to the amount of oxide originally grown electrochemically on the surface.
The rate of the chemical conversion, determined by the CSV as a
function of time, was constant, indicating that the effective surface
area of Ag2O available for the reaction was constant. This constant
surface area was attributable to the fact that the AgI product did not
form a protective layer on the Ag2O surface.
A kinetic analysis established the relationship between the total
reaction time and the reaction orders. From the total reaction times
as a function of iodide concentration and initial Ag2O inventory, the
rate was found to be first order in iodide concentration and almost
first order in oxide surface area, although the exact dependence on
surface area remains to be determined.
The apparent rate constant, using the geometric surface area instead of the real surface area of Ag2O, was obtained from both the
total reaction time measurements based on the changes in EOC and
the cathodic charges for AgI and Ag2O reduction to Ag using cathodic stripping voltammetry. The rate constants obtained by these
two different procedures agreed to within 50%.
Based on the kinetic data obtained, rate equations were derived
which can be used in models developed to predict the fate of radioiodine in various proposed nuclear reactor accident scenarios.
F76
Journal of The Electrochemical Society, 154 共4兲 F70-F76 共2007兲
Acknowledgments
This research was funded by a grant from the Natural Science
and Engineering Council of Canada 共NSERC兲 and a start-up grant
from the Faculty of Science, the University of Western Ontario. The
electrochemical analysis equipment was purchased by a grant from
the Canada Foundation for Innovation.
University of Western Ontario assisted in meeting the publication costs of
this article.
References
1. G. W. Parker and C. J. Barton, Fission Product Release, The Technology of Nuclear
Reactor Safety, Vol. 2, The M.I.T. Press, Cambridge, MA 共1973兲.
2. F. Garisto, R. J. Lemire, J. Paquette, P. P. S. Saluja, S. Sunder, D. F. Torgerson, A.
C. Vikis, D. J. Wren, and J. C. Wren, in Proceedings Series of an International
Symposium on Source Term Evaluation for Accident Conditions, Columbus, OH, p.
501 共1986兲.
3. L. Soffer, S. B. Burson, C. M. Ferrell, R. Y. Lee, and J. N. Ridgely, in Accident
Source Terms for Light-Water Nuclear Power Plants, NUREG-1465GAR, Vol. 95,
Issue 13, Washington, DC 共1995兲.
4. J. McFarlane, J. C. Wren, and R. J. Lemire, Nucl. Technol., 138, 162 共2002兲.
5. J. C. Wren, J. M. Ball, and G. A. Glowa, Nucl. Technol., 129, 297 共2000兲.
6. D. Jacquemain, N. Hanniet, C. Poletiko, S. Dickinson, C. Wren, D. A. Powers, E.
Krausmann, F. Funke, R. Cripps, and B. Herrero, in Proceedings of the OECD
7.
8.
9.
10.
11.
12.
13.
14.
15.
16.
17.
18.
19.
20.
Workshop on Iodine Aspects of Severe Accident Management, Vantaa, Finland, p.
57 共1999兲.
N. Girault, S. Dickinson, F. Funke, A. Auvinen, L. Herranz, and E. Krausmann,
Nucl. Eng. Des., 236, 1293 共2006兲.
F. Funke, G. U. Greger, A. Bleier, S. Hellman, and W. Morell, in Proceedings of
the 4th CSNI Workshop on the Chemistry of Iodine in Reactor Safety, Wurenlingen,
Switzerland 共1996兲.
E. Krausemann and Y. Drossinos, J. Nucl. Mater., 264, 113 共1999兲.
A. M. Zaky, F. H. Assaf, S. S. Abd El Rehim, and B. M. Mohamed, Appl. Surf.
Sci., 221, 349 共2004兲.
C. F. Baes and R. Mesmer, The Hydrolysis of Cations, p. 278, John Wiley & Sons,
New York 共1976兲.
T. U. Hur and W. S. Chung, J. Electrochem. Soc., 152, A996 共2005兲.
J. G. Becerra, R. C. Salvarezza, and A. J. Arvia, Electrochim. Acta, 33, 1431
共1988兲.
J. M. Drogg, P. T. Arderliesten, and G. A. Bootsma, J. Electroanal. Chem. Interfacial Electrochem., 99, 173 共1979兲.
B. M. Jovic and G. R. Strafford, Electrochem. Commun., 1, 247 共1999兲.
C. Alonso, R. C. Salvarezza, J. M. Vara, and A. J. Arvia, Electrochim. Acta, 35,
489 共1990兲.
V. I. Birss and G. A. Wright, Electrochim. Acta, 27, 1439 共1982兲.
S. Jaya, T. P. Rao, and G. P. Rao, J. Appl. Electrochem., 17, 997 共1987兲.
F. Fourcade, T. Tzedakis, and A. Bergel, Chem. Eng. Sci., 58, 3507 共2003兲.
B. J. Murray, O. Li, J. T. Newberg, E. J. Menke, J. C. Hemminger, and R. M.
Penner, Nano Lett., 5, 2319 共2005兲.