Interaction of aqueous iodine species with Ag2O/Ag surfaces

Article abstract of DOI:10.1149/1.2435696

The chemical conversion of Ag2 O films on Ag surfaces to AgI in aqueous iodide solutions has been studied electrochemically. Ag2 O films were grown potentiostatically and then exposed to I- solutions. The chemical conversion process was followed at open-circuit potential (EOC) using cathodic stripping voltammetry performed after various exposure periods. The EOC showed a sudden drop at the completion of the conversion of Ag2 O to AgI, reaching a steady-state value close to the equilibrium potential for AgIAg and the iodide solution. This sudden drop in EOC allowed easy determination of the total reaction time required for complete conversion of Ag2 O to AgI. Distinctly separated current peaks were observed for the cathodic reduction of Ag2 O and AgI to Ag, and the charges associated with these peaks provided a measure of the amount of Ag2 O 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.

Full text of DOI:10.1149/1.2435696

F70
Journal of The Electrochemical Society, 154 ͑4͒ F70-F76 ͑2007͒
0013-4651/2007/154͑4͒/F70/7/$20.00 © The Electrochemical Society
Interaction of Aqueous Iodine Species with Ag2O/Ag Surfaces
*
**
**,z
X. Zhang, S. Stewart, D. W. Shoesmith, and J. C. Wren
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 ͑EOC͒ 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. ͓DOI: 10.1149/1.2435696͔ All rights reserved.
Manuscript submitted September 19, 2006; revised manuscript received November 27, 2006.
Available electronically February 13, 2007.
Ag2O + I2 → 2AgI + 1 O2
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
͓3͔
One of the safety issues of nuclear power plants is the potential
radiation dose to the public in the unlikely event of a severe acci-
dent. 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 as-
sumed 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 through-
out the containment building following an accident.4 However, a
small fraction could be released to the gas phase due to the continu-
ous 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 man-
agement.
One reaction of interest in assessing iodine volatility is that be-
tween aqueous iodine species and silver surfaces. The control rod
͑for neutron flux͒ in some nuclear reactors is made of silver-
cadmium, and silver is assumed to be released from the reactor core
into the containment building environment in some accident sce-
narios, and its reaction with aqueous iodine to form insoluble silver
iodide could dominate iodine behavior. In fact, a significant reduc-
tion in iodine volatility in the presence of silver has been observed
in engineering-scale experiments simulating postaccident contain-
ment conditions.6,7
2
radiation
I− ↔ I2
I2 + H2O ꢀ HOI + I− + H+
I2 + I− ꢀ I−3
͓4͔
͓5͔
͓6͔
Equilibria 5 and 6 are achieved thermally even in the absence of
radiation. Furthermore, the effects of oxide and iodide film mor-
phologies 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 ͑Reaction 1͒
examined, using electrochemical and surface analytical techniques.
The reduction potentials for AgI/Ag and Ag2O/Ag are well sepa-
rated 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 quantita-
tively monitor the kinetics of an aqueous anion induced chemical
transformation between two insoluble solids. From these measure-
ments, the rates of individual reactions required for the development
of nuclear reactor safety assessment codes can be extracted.
In the highly oxidizing and ionizing radiation conditions ex-
pected, Ag could enter the aqueous phase in either the metallic or
oxidized ͑AgOH, Ag2O, and AgO͒ forms. The availability of aque-
ous iodine species, mainly I− and I2, would vary depending on ra-
diation 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 measure-
ments 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
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 elec-
trode was a 7 mm ͑in diameter͒ 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 ͑SCE͒ 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 re-
sponses. Corrware and Corrview software ͑supplied by Scribner and
Associates͒ was used to control experiments and analyze data.
Ag2O + 2I− + 2H+ → 2AgI + H2O
and the reaction of I2 with both metallic Ag and Ag2O
2Ag + I2 → 2AgI
͓1͔
͓2͔
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 pu-
rified using a NANOpure Diamond UV ultrapure water system from
*
Electrochemical Society Student Member.
Electrochemical Society Active Member.
**
z E-mail: jcwren@uwo.ca
Downloaded on 2014-12-04 to IP 138.251.14.35 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).

Journal of The Electrochemical Society, 154 ͑4͒ F70-F76 ͑2007͒
͑Ee͒
F71
= 0.927 − 0.0592 pH ͑vs SCE͒ = 0.216 V ͑vs SCE͒
Ag O/Ag
2
͓8͔
The vertical dotted line in Fig. 1 indicates this potential value.
The charges associated with oxide formation ͑QA͒ and reduction
͑QC͒, 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 poten-
tials less than ͑Ee͒Ag O/Ag, has been attributed to the chemisorption
2
process
Ag + OH−͑ad͒ ꢀ AgOH + e−
͓9͔
In stage 2, where potentials are slightly more positive than
͑Ee͒Ag O/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.
2
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 ⍀.
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 Ͻ+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 ϳ58 ⍀, 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 im-
pedes 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.
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 potentiostati-
cally at a potential in the range +0.2 to +0.6 V until the total oxi-
dation charge reached a desired value, generally in the range
0.01–0.2 C ͑or 0.026–0.52 C cm−2͒. 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 ͑M͒. The progress
of the reaction between Ag2O and I− was monitored by measuring
the open-circuit potential ͑EOC͒ as a function of time. At various
reaction times, the reaction was terminated by transferring the elec-
trode from the KI solution back to the KI-free solution. The poten-
tial 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 current recovery for potentials Ն0.5 V suggests an open
structure allowing the three-dimensional ͑3D͒ 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 ͑these results are not shown͒
does not show behavior that can be clearly identified as the reduc-
tion 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 ͑Ee͒
to avoid the possibility of oxide for-
Ag O/Ag
mation. The onset of Ag2I formation
Ag + I− ꢀ AgI + e−
The morphology of silver electrode surfaces at various stages of
the reaction were obtained with a Leo 440 scanning electron micro-
scope and the composition of products analyzed with energy-
dispersive X-ray ͑EDX͒ spectroscopy.
͓10͔
occurs at potentials immediately above the equilibrium potential,
͑Ee͒AgI/Ag, given by
͑Ee͒
= −0.396 − 0.0592 log͓I−͔ ͑V vs SCE͒
= −0.159 V ͑vs SCE͒ ͑at 10−4 MI−͒
= −0.218 V ͑vs SCE͒ ͑at 10−3 MI−͒
Results
AgI/Ag
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 ͑pH 12͒ free of KI is shown in Fig. 1. The form of
the voltammogram is very similar to that observed in borate buff-
ered solutions, confirming that pH control was maintained through-
out the film growth process.10 The oxide formation/reduction pro-
cess 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 ͑Ee͒ for the
reaction
͓11͔
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 ͑Ee͒
,
AgI/Ag
and the equality in anodic ͑AgI formation͒ and cathodic ͑AgI reduc-
tion͒ charges, confirms that the AgI/Ag reaction is as ͑more͒ revers-
ible as ͑than͒ the Ag2O/Ag reaction. The linear i − E relationship
2Ag + 2 OH− ꢀ Ag2O + 2e− + H2O
given by the Nernst equation ͑at pH 12͒
͓7͔
Downloaded on 2014-12-04 to IP 138.251.14.35 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).

F72
Journal of The Electrochemical Society, 154 ͑4͒ F70-F76 ͑2007͒
Figure 4. SEM images of oxide-covered Ag electrodes. The films were
grown potentiostatically at 0.6 V for ͑a͒ 700 and ͑b͒ 1300 s, or to total
anodic charges of 0.11 C ͑or 0.28 C cm−2͒ and 0.2 C ͑or 0.52 C cm−2͒, re-
spectively.
Figure 4 shows scanning electron microscopy ͑SEM͒ images of
potentiostatically grown Ag2O films after 700 and 1300 s ͑or total
charges of 0.11 and 0.2 C, respectively͒. 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 mor-
phologies and crystal structures, depending on the potential applied,
have been observed by others.20
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 ͓I−͔ of 10−4 and 10−3 M. The vertical broken lines
indicate the equilibrium potentials for AgI/Ag for the two ͓I−͔.
for AgI reduction ͑−0.25 V to the reduction peak͒ suggests a reac-
tion involving ion migration. The slope of this relationship is
ϳ150 ⍀, which is slightly greater than the solution resistance, sug-
gesting 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 ͑Fig. 3͒. The observation of two cathodic reduction
peaks at the potentials for Ag2O ͑Fig. 2͒ and AgI ͑Fig. 3͒ 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 ͑in the potential range −0.05 to −0.2 V͒,
indicating that the two film formation processes may occur indepen-
dently.
EDX analysis and X-ray diffraction analysis ͑XRD͒ show that
incorporation of phosphate into the oxide occurs. However, even
after extensive film growth under potentiostatic conditions, the po-
tential established on switching to open circuit ͑EOC͒ always relaxes
to ͑Ee͒
͑see below͒, consistent with exposure of an Ag/Ag2O
Ag O/Ag
interface 2to 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 aque-
ous solution containing KI, the open-circuit potential ͑EOC͒ of the
Ag/Ag2O electrode was followed as a function of reaction time, trxn
.
The dependence of the reaction kinetics on ͓I−͔ was studied on an
Potentiostatic growth of initial Ag2O films.— To provide a start-
ing point for the study of I− with Ag2O, a known amount of oxide
oxide film grown until ͑QAg O͒0 = 0.2 C at ͑EAg O͒0 = 0.6 V ͑Fig.
2
was grown potentiostatically ͑at ͑EAg O͒0͒ on Ag in an Ar-purged
2
5a͒, and the influence of the initial oxide inventor2y, for films grown
at 0.6 V, was studied in 5 ϫ 10−4 M KI solution ͑Fig. 5b͒.
The initial value of EOC was independent of ͓I−͔ ͑Fig. 5a͒, varied
only marginally for different initial film thicknesses ͑Fig. 5b͒, and
KI-free solution. The anodic charge associated with oxide growth
was obtained by integrating the current–time transient recorded at
͑EAg O͒0. Since negligible oxide dissolution occurs during film
2
was very close to ͑Ee͒Ag O/Ag. With time, EOC slowly decreased to a
growth at pH 12, this anodic charge, referred to as ͑QAg O͒0, is the
initial inventory of oxide present prior to exposure to KI2 solution.
2
second plateau at a time dependent on the initial film thickness
before undergoing a final abrupt transition to a much lower steady-
state value. The time taken to reach the final transition decreased as
͓I−͔ increased and increased as ͑QAg O͒0 increased. The final steady-
2
state value of EOC is close to ͑Ee͒AgI/Ag, as shown in Fig. 6 by a
comparison of final EOC values to calculated values of ͑Ee͒
.
AgI/Ag
Clearly, once the final potential transition has occurred, redox con-
ditions 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: ͑i͒ a direct chemical process via Reac-
tion 1, and ͑ii͒ 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 else-
where.
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 ͑CSV͒ 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
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 ͓I−͔ of 10−4 and 10−3 M.
Downloaded on 2014-12-04 to IP 138.251.14.35 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).

Journal of The Electrochemical Society, 154 ͑4͒ F70-F76 ͑2007͒
F73
Integration of the two reduction peaks in Fig. 7b yielded values
of ͑QAg O͒t and ͑QAgI͒t as a function of reaction time, trxn ͑Fig. 8͒.
2
The values are inversely related, and ͑QAg O͒t = 0 at ␶f, further con-
firming there is no residual oxide on the 2Ag surface at t = ␶f. The
relationship
͑QAg O͒t + ͑QAgI͒t = ͑QAg O 0
͒
͓12͔
2
2
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 be-
tween Ag and I−. The slopes of the two plots can be used to deter-
mine 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 ͑b͒, after the first short potential
transition ͑c͒, and on completion of the reaction ͑t = ␶ ͒ ͑d͒. As the
f
reaction proceeds the surface becomes progressively covered by
small particulates of AgI. The micrograph in c shows the coexist-
ence 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− ͑Reaction 1͒ can be expressed as
d͑mAg O t
͒
d͑mAgI t
͒
2
q
t
p
= −2
= k͑10−3 ϫ ͓I−͔ ͒ ͑AAg O
͒
͓13͔
t
2
dt
dt
where ͑mAg O͒t and ͑mAgI͒t are the number of moles of Ag2O and
2
AgI at reaction time t; ͓I−͔ ͑in mol dm−3͒ and ͑AAg O͒t ͑in cm2͒ are
t
2
the concentration of I− and the surface area of Ag2O at time t,
respectively; k is the reaction rate constant; and superscripts, p and
Figure 5. Open-circuit potential ͑EOC͒ as a function of reaction time and its
dependence on ͓I−͔ and ͑QAg O͒0. The reaction conditions were ͑a͒ oxides
q, represent reaction orders with respect to ͓I−͔ and ͑AAg O͒, respec-
2
2
grown at 0.6 V to a constant anodic charge ͑QAg O͒ = 0.2 C, with ͓I−͔ rang-
tively. The units of k will depend on p and q, and the factor of 10−3
0
2
ing from 10−4 to 5 ϫ 10−3 M, and ͑b͒ oxides grown at ͑EAg O͒ = 0.6 V, up
is included to express ͓I−͔ in mol dm−3 ͑or M͒.
0
2
t
to ͑QAg O
͒
ranging from 0.01 to 0.2 C, with ͓I−͔ constant at 5 ϫ 10−4 M.
0
2
Because the Ag2O surface area to solution volume ratio is very
small, ͓I−͔t is constant. The number of moles of unreacted Ag2O and
AgI formed are given by the cathodic charges for reduction of Ag2O
with reaction time. Once the abrupt transition to the final steady-
state value of EOC occurred, no oxide remained on the electrode
surface and no further AgI was formed, confirming that this transi-
tion 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.
to Ag, ͑AAg O͒t, and AgI to Ag, ͑QAgI͒, by
2
͑QAg O t
͒
2
͑mAg O͒t =
2
2F
͑QAgI t
F
͒
͑mAgI͒t =
͓14͔
͓15͔
where F is the Faraday constant, and Eq. 13 becomes
d͑QAg O t
͒
d͑QAgI t
͒
2
q
t
p
= −
= kF͑10−3 ϫ ͓I−͔ ͒ ͑AAg O
͒
0
2
dt
dt
Figure 8 shows that this rate does not vary appreciably as the
reaction proceeds, suggesting ͑AAg O͒t = ͑AAg O͒0. This constant
2
2
surface area is also indicated by the micrographs in Fig. 9. The rate
equation can, therefore, be further reduced to
d͑QAgI t
͒
q
0
p
= kF͑10−3 ϫ ͓I−͔ ͒ ͑AAg O
͒
͓16͔
0
2
dt
and integrated to yield
͑QAgI͒t = kF͑10−3 ϫ ͓I−͔ ͒ ͑AAg O͒qt
͓17͔
p
0
0
Figure 6. Final steady-state open-circuit potential ͑symbols͒ observed for
various ͓I−͔ compared to calculated equilibrium potentials for the Ag
+ I− ꢀ AgI + e− reaction ͑line͒.
2
The total reaction time, ␶f, is then determined by
Downloaded on 2014-12-04 to IP 138.251.14.35 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).

F74
Journal of The Electrochemical Society, 154 ͑4͒ F70-F76 ͑2007͒
Figure 7. CSV as a function of reaction
time ͑trxn͒. The times when the CSVs were
taken are indicated on the open-circuit po-
tential plot in ͑a͒, and the corresponding
CSVs are shown in ͑b͒. The experimental
conditions
were
͑QAg O͒ = 0.2 C
0
2
͑0.52 C cm−2
ϫ 10−4 M.
͒
at 0.6 V, and ͓I−͔ = 5
͑QAg O 0
͒
͑QAgI
͒
f
2
␶ =
f
=
q
0
q
0
p
p
kF͑10−3 ϫ ͓I−͔ ͒ ͑AAg O
͒
kF͑10−3 ϫ ͓I−͔ ͒ ͑AAg O
͒
0
0
2
2
͓18͔
where ͑QAgI͒f = ͑QAg O͒0, since all the initially available Ag2O is
2
converted to AgI. The slope of a plot of ␶ vs log ͓I−͔ yields a value
f
of p = 1 ͑Fig. 10͒.
Similarly, a plot of log ␶ vs log͑AAg O͒0 would yield the nega-
f
2
tive 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 ␶ vs log͑AAg O͒0 ͑Fig. 11͒. If the initial oxide grew
f
uniformly on the electrode, 2then the oxide surface area would be
equal to its geometric area. Then, ͑AAg O͒0 would be independent of
2
͑QAg O͒0, and the film thickness directly proportional to ͑QAg O͒0.
2
2
However, the SEM micrograph in Fig. 4 shows three-dimensional
oxide growth, and consequently, the surface area of the initial Ag2O
Figure 8. Cathodic charges measured for the reduction of unreacted Ag2O
and product AgI at various reaction times, trxn. The experimental conditions
film is expected to have a small dependence on ͑QAg O 0
͒
2
were ͑QAg O͒ = 0.2 C ͑0.52 C cm−2͒ at 0.6 V, and ͓I−͔ = 5 ϫ 10−4 M.
0
2
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
͑a͒, and the corresponding micrographs
are shown in ͑b–d͒.
Downloaded on 2014-12-04 to IP 138.251.14.35 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).

Journal of The Electrochemical Society, 154 ͑4͒ F70-F76 ͑2007͒
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
d͑mAg O t
͒
d͑mAgI t
͒
2
= 2
= kapp͑10−3 ϫ ͓I−͔ ͒AElec
͓22͔
͓23͔
0
dt
dt
␤
͑QAg O 0
͒
2
ͩ ͪ
F
where kapp = kC
␤
For small values of ␤, kapp approaches k.
A value for kapp can be extracted from the total reaction time
observed as a function of ͓I−͔. From Eq. 18 ͑p = q = 1͒
͑QAg O 0
͒
1
1
2
ͩ ͪ
F
␶ =
f
͓24͔
kapp ͑10−3 ϫ ͓I−͔ ͒AElec
0
yielding a value of kapp = 5.4 ϫ 10−3 cm s−1
.
Figure 10. Log of ͓I−͔ vs log of total reaction time, ␶f; the initial
͑QAg O͒ = 0.2 C at 0.6 V.
Alternatively, kapp can be obtained from ͑QAg O͒t and ͑QAgI͒t as
0
2
2
a function of reaction time. Equation 22 yields
͑QAgI͒t = ͑͑QAg O͒0 − ͑QAg O͒t͒ = kappF͑10−3 ϫ ͓I−͔ ͒AElec
t
0
2
2
␤
͓25͔
͑QAg O 0
͒
2
ͩ ͪ
F
͑AAg O͒0 = C
A
͓19͔
Hence, Fig. 7 yields kapp = 3.8 ϫ 10−3 cm s−1, in good agreement
͑within 50%͒ 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 determina-
tion of whether the incorporation of anions into the electrochemi-
cally grown films exerts a significant influence on the kinetics of the
film conversion reaction.
␤
Elec
2
where C␤ is a constant or a unit conversion factor. The value of ␤
would be zero and C ͓͑QAg O͒ /F͔␤ϭ1 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
␤
0
2
͑QAg O 0
͒
2
␶ =
f
Conclusions
q
0
p
kF͑10−3 ϫ ͓I−͔ ͒ ͑AAg O
͒
0
2
The chemical conversion of Ag2O by I− to AgI in aqueous iodide
solutions ͑pH 12͒ 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 transi-
tion in the open-circuit potential from a value close to the equilib-
rium potential for Ag2O/Ag to a value identical to the equilibrium
potential for AgI/Ag. This abrupt change in EOC allowed easy de-
termination of the total reaction time required for complete conver-
sion of Ag2O to AgI. This time to complete the reaction was in-
versely proportional to the iodide concentration and was nearly
proportional to the amount of oxide originally grown electrochemi-
cally on the surface.
1−q␤
͑QAg O 0
͒
1
2
=
͓20͔
ͩ ͪ
kC␤q͑10−3 ϫ ͓I−͔ ͒pAq
F
0
Elec
and a plot of log ␶ vs log͑QAg O͒0 would have a slope equal to 1
f
2
− q␤.
Despite this small surface area dependence, it is reasonable to
accept a first-order dependence on the surface area, ͑AAg O͒0, and
2
the rate law ͑Eq. 13͒ can be written
d͑mAg O t
͒
d͑mAgI t
͒
2
= 2
= k͑10−3 ϫ ͓I−͔ ͒͑AAg O 0
͒
͓21͔
t
2
dt
dt
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 in-
stead 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 ca-
thodic 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 radio-
iodine in various proposed nuclear reactor accident scenarios.
Figure 11. Log of total anodic charge used to grow the initial silver oxide
film, ͑QAg O͒0, vs log of total reaction time, ␶f; ͓I−͔ = 5 ϫ 10−4 M.
2
Downloaded on 2014-12-04 to IP 138.251.14.35 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).

F76
Journal of The Electrochemical Society, 154 ͑4͒ F70-F76 ͑2007͒
Workshop on Iodine Aspects of Severe Accident Management, Vantaa, Finland, p.
57 ͑1999͒.
Acknowledgments
This research was funded by a grant from the Natural Science
and Engineering Council of Canada ͑NSERC͒ 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.
7. N. Girault, S. Dickinson, F. Funke, A. Auvinen, L. Herranz, and E. Krausmann,
Nucl. Eng. Des., 236, 1293 ͑2006͒.
8. 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 ͑1996͒.
9. E. Krausemann and Y. Drossinos, J. Nucl. Mater., 264, 113 ͑1999͒.
10. A. M. Zaky, F. H. Assaf, S. S. Abd El Rehim, and B. M. Mohamed, Appl. Surf.
Sci., 221, 349 ͑2004͒.
11. C. F. Baes and R. Mesmer, The Hydrolysis of Cations, p. 278, John Wiley & Sons,
New York ͑1976͒.
12. T. U. Hur and W. S. Chung, J. Electrochem. Soc., 152, A996 ͑2005͒.
13. J. G. Becerra, R. C. Salvarezza, and A. J. Arvia, Electrochim. Acta, 33, 1431
͑1988͒.
14. J. M. Drogg, P. T. Arderliesten, and G. A. Bootsma, J. Electroanal. Chem. Inter-
facial Electrochem., 99, 173 ͑1979͒.
15. B. M. Jovic and G. R. Strafford, Electrochem. Commun., 1, 247 ͑1999͒.
16. C. Alonso, R. C. Salvarezza, J. M. Vara, and A. J. Arvia, Electrochim. Acta, 35,
489 ͑1990͒.
17. V. I. Birss and G. A. Wright, Electrochim. Acta, 27, 1439 ͑1982͒.
18. S. Jaya, T. P. Rao, and G. P. Rao, J. Appl. Electrochem., 17, 997 ͑1987͒.
19. F. Fourcade, T. Tzedakis, and A. Bergel, Chem. Eng. Sci., 58, 3507 ͑2003͒.
20. B. J. Murray, O. Li, J. T. Newberg, E. J. Menke, J. C. Hemminger, and R. M.
Penner, Nano Lett., 5, 2319 ͑2005͒.
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 ͑1973͒.
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 ͑1986͒.
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 ͑1995͒.
4. J. McFarlane, J. C. Wren, and R. J. Lemire, Nucl. Technol., 138, 162 ͑2002͒.
5. J. C. Wren, J. M. Ball, and G. A. Glowa, Nucl. Technol., 129, 297 ͑2000͒.
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
Downloaded on 2014-12-04 to IP 138.251.14.35 address. Redistribution subject to ECS terms of use (see ecsdl.org/site/terms_use) unless CC License in place (see abstract).