Corrosion Science 52 (2010) 4035–4045

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Corrosion Science

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Electrochemical corrosion behaviour of Incoloy 800 in sulphate solutionscontaining hydrogen peroxide

Tirdad Nickchi ⇑, Akram AlfantaziDepartment of Materials Engineering, The University of British Columbia, Vancouver, BC, Canada V6T1Z4

a r t i c l e i n f o

Article history:Received 13 July 2010Accepted 11 August 2010Available online 17 August 2010

Keywords:A. SuperalloysB. PolarizationB. EISC. Kinetic parameters

0010-938X/$ - see front matter � 2010 Elsevier Ltd. Adoi:10.1016/j.corsci.2010.08.022

⇑ Corresponding author. Tel.: +1 6047797544; fax:E-mail address: tirdadn@interchange.ubc.ca (T. Nic

a b s t r a c t

The aim of this study is to evaluate the electrochemical corrosion behaviour of Incoloy 800 in sulphatesolutions containing H2O2 in the temperature range of 25–80 �C. The open circuit potential measure-ments, cathodic and anodic polarization and electrochemical impedance spectroscopy (EIS) were usedto characterize the corrosion behaviour. The results provide kinetic data for reduction of hydrogen per-oxide on Pt surface. The anodic polarization curves for Incoloy at different pH, temperature and H2O2 con-centration are presented. EIS data generally confirm the polarization interpretations about the effects ofvarious parameters. An equivalent circuit was used to fit all the acquired data.

� 2010 Elsevier Ltd. All rights reserved.

Introduction

One of the promising designs for the Gen IV reactors is thesupercritical water reactor, which benefits from high heat capacityof water in the supercritical state. Evaluation of corrosion issues forcandidate materials for this reactor is important because the corro-sion mechanisms in supercritical water are not completely under-stood. Furthermore, the variation of water chemistry induced bygamma irradiation can affect the corrosion mechanisms. The irra-diation leads to generation of radicals in water [1] whose effectson corrosion should be evaluated. One of the most important ofthese species is hydrogen peroxide. In this regard, Glass et al. havereported that gamma irradiation leads to electrochemical potentialshift, which is due to hydrogen peroxide generation [2]. Therefore,evaluation of effect of hydrogen peroxide on corrosion behaviour ofactive–passive alloys (due to their superior corrosion resistance) isessential.

A literature survey shows that previous studies in this regardhave evaluated the corrosion behaviour of alloys of interest inpressurized water reactors (PWR), constructional materials for tri-tiated water plants, and so on. These include evaluation of inter-granular stress corrosion cracking [3], and the crack growth rate[4] of 304 stainless steel in the presence of hydrogen peroxide, acomprehensive evaluation of welded S32550 duplex steel [5] andInconel 690 [6] corrosion mechanism by Bellanger et al., and eval-uation of transpassivity of 304 stainless steel in hydrogen peroxidecontaining alkaline solutions [7]. In addition, recently, fundamen-

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+1 6048223619.kchi).

tal studies regarding the interaction of H2O2 with carbon steel sur-face and iron corrosion products have been carried out [8–10].

The high-temperature high-pressure evaluation of corrosion inthe presence of H2O2 has also been studied. These include theeffect of water chemistry, including H2O2 concentration, on elec-trochemical potential and corrosion rate of 304 stainless steel[11,12] and alloy 182 [12], and evaluation of electrochemicalpotential [13], polarization behaviour [14] and oxide film charac-terization [15,16] of stainless steels in the presence of H2O2.

One of the possible candidate alloys to be used in Gen IV reac-tors is Incoloy 800 [17], due to its high temperature corrosionresistance. Some studies have evaluated its corrosion behaviourin high temperature aqueous solutions in terms of long-termimmersion [18] and electrochemical studies [19], as well as theoxidation behaviour in supercritical water simulated environments[20]. The aim of our research is to evaluate the corrosion behaviourof Incoloy 800 in the presence of hydrogen peroxide, due to its sig-nificance in establishing the corrosion mechanism. As a first step, itis necessary to have a fundamental understanding of the effect ofhydrogen peroxide on corrosion. Therefore, in this paper, we pres-ent the low temperature results as they provide a base line and afundamental understanding of corrosion in this medium.

2. Experimental

2.1. Electrochemical methods

A Princeton Applied Research (PAR) Versastat 4 potentiostat/galvanostat was used for all the electrochemical measurements.The open circuit potential (OCP) measurements were carried out

Fig. 1. Optical microstructure of Incoloy 800 polished and etched in Marblessolution for 15 s.

4036 T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045

for 1 ks or until a steady state potential with changes no more than2 mV per 100 s was recorded. The cathodic polarization tests onplatinum substrate were performed from OCP potential down to�2 V vs. reference electrode (Ag/AgCl, [Cl�] = 4 M) with a scan rateof 1 mV s�1. The anodic potentiodynamic experiments on alloy 800were performed from �1 V up to 2 V vs. Ag/AgCl with scan rate of1 mV s�1. Some of the polarization experiments were repeated toensure the reproducibility of the results. The electrochemicalimpedance spectroscopy (EIS) tests on alloy 800 substrate wereperformed after OCP stabilization with amplitude of 10 mV atOCP in the frequency range of 0.01–10,000 Hz with sampling rateof 10 points per decade. All the EIS experiments were performedin the forward and backward direction to ensure the validity ofthe results.

2.2. Electrochemical apparatus configuration

A conventional water-jacketed three-electrode cell, connectedto a heater (Cole-Parmer) to maintain the temperature at 25, 40,60 and 80 �C, was used. This cell consists of graphite as counterelectrode (CE), silver/silver chloride as reference electrode (RE)and platinum or alloy 800 as the working electrode. All the poten-tials in this work are reported vs. Ag/AgCl reference electrode atthe specific temperature. The working electrode was connectedto a copper wire by conductive silver paste, mounted in epoxy re-sin and positioned in the cell. The solution was purged with argonprior the experiments until complete deoxygenating, and contin-ued during the experiment. The argon flow rate was monitoredusing an area changing flow meter with an accuracy of 2% (Cole-Parmer) during the experiment and was maintained at 300 ml/min. This enabled us in maintaining the oxygen-free environmentand a fixed mass transfer state in the cell. The cell off-gas, passedthrough a condenser, was introduced into another containerincluding the same solution, but at 25 �C. The oxygen content ofthis solution was monitored with a dissolved oxygen sensor (OME-GA DOB21 accuracy = 12 ppb). This design was firstly because ofthe absence of an extra hole in the cell for positioning the dissolvedoxygen probe; however, this also enabled us to monitor the prob-able hydrogen peroxide decomposition, as the off-gas in this casewould contain oxygen.

2.3. Materials and preparation methods

Incoloy 800 HT with nominal composition (in wt.%) of C 0.08%,Mn 1.02%, Fe 45.57%, S 0.001%, Si 0.31%, Cu 0.16%, Ni 32.16%, Cr19.59%, Al 0.46%, Ti 0.08%, and heat treatment consisting of anneal-ing at 1176 �C for 9.1 min and air-cooling, was used in this study.Fig. 1 shows the optical microstructure of this alloy after beingetched in Marble’s agent (10 g copper sulphate, 50 ml HCl and50 ml water) for 15 s. It shows a fully annealed equiaxed-grainmicrostructure.

The test solution contained 25 mM sodium sulphate andvarying concentration of hydrogen peroxide. The concentrationswere selected in a range to simulate the actual concentration inthe presence of gamma irradiation [21]. Studies regarding H2O2

electrochemical reduction usually use NaClO4 as the supportingelectrolyte, since the ClO�4 ion does not specifically adsorb onthe substrate [22]. However, the ClO�4 ion itself may affect thecorrosion mechanisms, and therefore, a conventional sulphatebased supporting electrolyte was selected for this study. The con-centration range of hydrogen peroxide in the reactor will be lessthan 10�4 M; however, in order to understand the effects ofhydrogen peroxide concentration on corrosion behaviour, sur-veyed concentration region should be sufficiently higher thanthe target concentration. Hence, the concentrations of up to 1 Mwas studied. All the chemical agents were procured from Fisher

Scientific and solutions were prepared with deionized water.Some of the hydrogen peroxide containing solutions were titratedwith KMnO4 in the presence of sulphuric acid to confirm the H2O2

concentration. The pH of the solution was adjusted to pH = 5, 7, 9by addition of NaOH or H2SO4 at 25 �C. The actual pH at highertemperatures were calculated and not measured.

Surface preparation of platinum consists of grinding up to1200 SiC emery paper, degreasing using ethyl alcohol in ultrasoundfor 5 min and thorough rinsing in deionized water. The alloy 800was mechanically ground up to 1200 SiC emery paper, degreasedwith ethyl alcohol in ultrasound for 5 min, rinsed with deionizedwater followed by cathodic conditioning at �2 V vs. Ag/AgCl for15 min in the test solution.

3. Results and discussion

3.1. Kinetics of hydrogen peroxide reduction

The main purpose of this study is to elucidate the effect ofhydrogen peroxide as a reducing agent on corrosion behaviour ofIncoloy 800. Therefore, in the first step, studying the electrochem-ical behaviour of H2O2 at low concentration range in the solution ofinterest is essential, as it provides an understanding of the electro-chemical behaviour of total cathodic species in the solution. Previ-ously, electrochemistry of H2O2 reduction has been extensivelystudied since it is an important chemical in processes such aswaste treatment, oxygen reduction, corrosion and water radiolysis.Some of the research in this regard include reduction of H2O2 oncopper [23,24], gold [25], passive forming alloys [26], passive ironin alkaline solutions [22] and corrosion of uranium dioxide [27]and titanium [28] in hydrogen peroxide. Nonetheless, it is neces-sary to conduct experiments on electrochemistry of hydrogen per-oxide reduction in this specific solution to have an understandingof the overall cathodic processes.

3.1.1. Open circuit potential (OCP) measurementsThe evolution of open circuit potential as affected by hydrogen

peroxide concentration, pH, and temperature was studied. Thechange of OCP vs. time as a function of H2O2 concentration at25 �C and pH = 7 is presented in Fig. 2. It shows that as the concen-tration of H2O2 is increased from 10�6 to 10�4 M, the OCP isincreased from 174 mV to 232 mV vs. Ag/AgCl. However, furtherincrease in the concentration does not substantially increase theOCP potential. It was also observed that at 10�2 and 1 M, the

Fig. 2. The evolution of open circuit potential vs. time at 25 �C, [Na2SO4] = 25 mM,pH = 7 and different hydrogen peroxide concentrations.

Fig. 3. comparison of measured OCP and calculated equilibrium potentials forelectrochemical reduction and oxidation of hydrogen peroxide, the detailedexperimental conditions are presented in Table 1.

T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045 4037

solution was not fully deoxygenated after addition of H2O2. Inaddition, decomposition of H2O2 and evolution of oxygen at theplatinum surface was evident at 1 M. It was thus concluded thatthe surface concentration of H2O2 is significantly lower than thebulk concentration. This phenomenon necessitates a more thor-ough analysis of possible reactions on the surface at differentconditions.

Reactions concerning the reduction and the surface catalyzeddecomposition of H2O2 are taken into account. The correspondingreactions are as follows:

H2O2 þ 2Hþ þ 2e� ) 2H2O ð1Þ

H2O2 ) O2 þ 2Hþ þ 2e� ð2Þ

The calculated equilibrium potentials for these two reactions atvarious conditions of pH, H2O2 concentration and temperature arepresented in Table 1 as well as the measured final open circuit po-tential. The equilibrium potentials for these reactions are calcu-lated using the Nernst equation partly based on temperaturedependant Gibbs free energy data for H2O, H+ and O2 [29]. The tem-perature dependant thermodynamic data for aqueous H2O2 wasnot available and was thus estimated based on Gibbs free energydata on gaseous H2O2 [30] along with the free energy of dissolutionof gaseous hydrogen peroxide into water [31]. The corrections forAg/AgCl temperature coefficients were extracted from Ref. [32].The contribution of Kw and sulphuric acid second dissociationconstant (HSO�4 ) SO2�

4 þHþ) [33] and hydrogen peroxide dissoci-ation constant (H2O2 ) Hþ þHO�2 ) [34] variation with tempera-ture was taken into account for calculation of actual pH atelevated temperatures. In the case of H2O2 acidic decompositionconstant, the data is only available up to 35 �C and, assuming a

Table 1Summary of OCP and kinetic data on reduction of hydrogen peroxide on platinum in [Na2

Experimentnumber

Temperature(�C)

pH H2O2

concentration(M)

Reaction 1equilibriumpotential (V vs.AgAgCl)

ReactionequilibripotentiaAgCl)

1 25 7.00 10�6 0.967 �0.0992 25 7.00 10�4 1.026 �0.0403 25 7.00 10�2 1.085 0.0194 25 7.00 1 1.144 0.0785 25 5.00 10�4 1.145 0.0796 25 9.00 10�4 0.908 �0.1587 40 6.84 10�4 1.036 �0.0698 60 6.58 10�4 1.051 �0.0969 80 6.32 10�4 1.068 �0.121

constant heat of dissociation, it was extrapolated to higher temper-atures. Comparison of measured OCP values and calculated equi-librium potentials for reactions 1 and 2, as shown in Fig. 3,demonstrates that all measured values fall between the two theo-retical potentials, indicating a mixed potential on the platinumsurface. In addition, the OCP is, for all cases, closer to electrochem-ical oxidation of hydrogen peroxide reaction. This shows the moresignificant efficacy of reaction 2 on total mixed potential. Astemperature is increased from 25 �C to 80 �C, the reaction 1equilibrium potential increases by 42 mV, while that of reaction2 decreases by 80 mV, and therefore, the measured potential doesnot substantially change with temperature. In addition, the varia-tion of the measured potential vs. pH and H2O2 concentration isas expected for contribution of H+ and H2O2 concentration to theNernst potential.

3.1.2. Cathodic polarizationThe kinetics of H2O2 reduction was studied by cathodic polari-

zation method. The cathodic polarization curves of hydrogen per-oxide reduction on Pt at pH = 7, 25 �C and different H2O2

concentration are presented in Fig. 4. The general features of thecurves are the presence of a restricted region of activation-con-trolled reduction, followed by mixed control region and purelymass transfer-controlled region for hydrogen peroxide reduction.The limiting current density is increased as the bulk concentrationof hydrogen peroxide is increased. On further cathodic sweep, the

SO4] = 25 mM and various conditions of pH, temperature and [H2O2].

2uml (V vs. Ag/

MeasuredOCP (V)

i0 (A/cm2) bc

(V/dec)iL (A/cm2) D (cm2 s�1)

0.174 1.30 � 10�6 0.56 5.10 � 10�6 6.73 � 10�6

0.232 1.27 � 10�6 0.10 3.82 � 10�5 6.60 � 10�8

0.222 4.03 � 10�5 0.18 3.82 � 10�3 2.09 � 10�8

0.242 4.57 � 10�4 0.26 1.57 � 10�2 2.37 � 10�9

0.402 – – 3.82 � 10�5 3.96 � 10�7

0.183 1.20 � 10�6 0.09 1.91 � 10�5 6.19 � 10�8

0.221 2.85 � 10�6 0.06 7.64 � 10�5 1.48 � 10�7

0.232 1.96 � 10�6 0.08 7.64 � 10�5 1.02 � 10�7

0.231 4.13 � 10�6 0.09 1.02 � 10�5 2.14 � 10�7

Fig. 4. The cathodic polarization curves for hydrogen peroxide on Pt at 25 �C,[Na2SO4] = 25 mM, pH = 7 and different hydrogen peroxide concentration.

4038 T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045

electrochemical reduction of H+ would occur. By increasing theconcentration of hydrogen peroxide as a strong oxidant, the onsetof H+ is delayed and finally at 1 M concentration, the mass transfer-controlled reduction of hydrogen peroxide completely masks theproton electrochemical reduction.

The same experiments were performed to investigate the effectof pH and temperature on cathodic polarization behaviour. Fig. 5shows the polarization curves for solutions containing 10�4 MH2O2, at 25 �C, and various pH. The limiting current density atpH = 5 and 7 are identical, as the electro active species (H2O2) con-centration is the same, however, the limiting current density forpH = 9 is lower. The reason for this observation is not completelyclear to us; however, it may be explained by considering the aciddissociation of H2O2 as a weak acid: as pH increases, the acid–baseequilibrium is shifting so that H2O2 is consumed and HO�2 is gener-ated. Consequently, the overall available hydrogen peroxide as theelectro active species is declined.

The other interesting observation in Fig. 5 is the presence of twocathodic waves at pH = 5, indicating the occurrence of two differ-ent reactions at this region. Prabhu et al. [35] stated that at lowpH values, oxygen is firstly reduced to hydrogen peroxide. Thisoxygen is previously formed on the platinum surface by catalyticdecomposition of H2O2. The second wave is then assumed to be

Fig. 5. cathodic polarization curves for hydrogen peroxide on Pt surface at 25 �C,[Na2SO4] = 25 mM, 10�4 M H2O2 concentration, and different pH.

related to the direct reduction of H2O2 to water. The polarizationcurve features at pH = 5 is consistent with the hypothesis men-tioned above. The first wave continuing down to +150 mV vs.Ag/AgCl is probably due to reduction of oxygen which is generatedby decomposition of hydrogen peroxide, and the second wavestarting from +150 mV vs. Ag/AgCl down to �700 mV Ag/AgCl isassociated with direct reduction of H2O2.

The general features of polarization curves at different temper-atures, not shown here, are activation-controlled reduction fol-lowed by mixed and pure mass transfer-controlled reduction ofhydrogen peroxide. This behaviour is similar at different tempera-tures except a slight increase in limiting current density byincreasing the temperature. The summary of these values is pre-sented in Table 1.

The summary of kinetic parameters, such as steady state cur-rent density (the corrosion current density at open circuit poten-tial), Tafel slopes and limiting current density, for all theexperiments is also provided in Table 1. The onset of mixed controlregion appears at low overpotentials, due to low concentration ofH2O2; and therefore, Tafel extrapolation method cannot beperformed on the results. Hence, these kinetic parameters wereextracted by simulating the total curve using extended Butler–Volmer equation (Eq. (3)), taking into account both the anodicand cathodic currents at low overpotentials and the effect of masstransfer at higher overpotentials. Prabhu et al., have reported thatthe charge transfer coefficient for anodic and cathodic reactions ofH2O2 on platinum were identical [35]. Therefore, equal anodic andcathodic Tafel slopes were assumed in these calculations. As seenTable 1, kinetic data for pH = 5 are not presented, because theappearance of two cathodic waves prevents an accurate simulationusing Eq. (3).

i ¼i0 exp �acnF

RT g� �

� exp aanFRT g

� �� �

1þ i0 exp �ac nFRT gð Þ

iL;H2O2

ð3Þ

This is the extended Butler–Volmer equation in which the con-tribution of mass transfer as well as the reaction kinetics is takeninto account. In this equation, i0 represents the exchange currentdensity, ac and aa are the transfer coefficients for cathodic and ano-dic reactions, g is the overpotential and F, R and T have their con-ventional meaning.

The exchange current density values are increasing as the bulkconcentration of hydrogen peroxide is increased. However, thisincrease is not as significant as is expected for concentratedsolutions. This is in accordance with Section 3.1.1 findings thatprobably a mixed potential is established at the surface of the elec-trode and that the actual surface concentration is less than the bulkconcentration due to decomposition. Increasing the pH from 7 to 9leads to a slight decrease in i0. The effect of increasing temperatureis to increase the exchange current density. This trend was furtheranalyzed by considering the Arrhenius behaviour: The linearregression of ln(i0) � T�1 data showed that the thermal activationenergy for H2O2 reduction is equal to 16 kJ/mol, with a regressionerror of R2 = 0.88. The R2 value, the coefficient of determination,varies between 0 and 1, and the nearer it is to 1, the better is theregression [36].

The Tafel slope values, as shown in Table 1, are in a very widerange: from 560 mV/decade for 10�6 M H2O2 concentration to60 mV/decade at elevated temperatures. Large values for Tafelslopes are usually attributed to either an electrochemical reactiontaking place at the surface of oxide layers or coupled chemical–electrochemical (CE) mechanism [24].

Sun et al. mentioned that the oxide layers significantly alter theredox reaction rates on the noble metal surfaces [37] and therefore,the high value for Tafel slope can be attributed to the difficult elec-

T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045 4039

tron transfer within the platinum oxide layer. However, formationof oxide or hydroxide compounds are not feasible at pH = 7 andmeasured OCP values in these experiments. For example, the Pour-baix diagram for platinum indicates that formation of Pt(OH)2 isfavored at 0.395 V vs. Ag/AgCl at pH = 7. Still the possibility offormation of oxides prior to immersing in the solution cannot beexcluded, because the Pt electrode was not cathodically condi-tioned in the solution before the experiments.

If the latter is taking place, the high values for Tafel slope at10�6 M H2O2 concentration will be probably associated with achemical reaction occurring at the surface preceding the electrontransfer step. One possible mechanism would be the adsorptionof hydrogen peroxide on the platinum surface followed by electro-chemical reduction of the adsorbed species:

2PtþH2O2 ) 2Pt-OH ð4Þ

Pt-OHþHþ þ e� ) PtþH2O ð5Þ

Reactions (4) and (5) demonstrate a possible CE mechanism forreduction of H2O2 in dilute solutions. This may explain the high Ta-fel slope values at very dilute hydrogen peroxide solutions. None-theless, as the H2O2 concentration is increased, all availablereaction sites on platinum surface would be occupied and a directpathway for reduction would be activated. The Memming’s classi-cal mechanism for H2O2 [38] would explain this direct reductionpathway in the following step reactions:

H2O2 þ e� ) OH�ads: þ OH� ð6Þ

OH�ads: þ e� ) OH� ð7Þ

Therefore, it seems that as the concentration of hydrogen per-oxide is increased, the reduction mechanism is shifted from a CEmechanism, in which the platinum substrate does participate, toa direct two-step reduction of H2O2 independent of reaction siteson Pt surface. It is also observed that the pH and temperatureeffects on the Tafel slop are not significant. Hence, it can be con-cluded that the temperature and pH do not substantially alterthe reaction pathways.

The limiting current density data in Table 1 was used to esti-mate the diffusion coefficient at different conditions. The limitingcurrent density as calculated by interpretations of Fick’s first lawis:

iL;H2O2 ¼ nFDH2O2

dCb;H2O2

ð8Þ

In this equation, iL;H2O2 is the limiting current density for hydro-gen peroxide, DH2O2 and Cb;H2O2 are the diffusion coefficient andbulk concentration of hydrogen peroxide, d is the boundary layer

Table 2Summary of anodic polarization tests performed on Incoloy 800 in [Na2SO4] = 25 mM and

Exp no. Temperature (�C) pH H2O2 concentration (M) Ecorr

(mV vs. Ag/Ag

1 25 5 0 �5432 25 7 0 �7423 25 9 0 �7304 40 7 0 �6715 60 7 0 �6296 80 7 0 �6247 25 7 10�6 �7178 25 7 10�4 �2199 25 7 10�2 11410 25 5 10�4 �24711 25 9 10�4 �20912 40 7 10�4 �21013 60 7 10�4 �18214 80 7 10�4 �167

thickness, and n is number of electrons involved. Therefore, the dif-fusion coefficient can be calculated by assuming the diffusionboundary layer. The boundary layer thickness is typically around10 lm for the forced convection systems, and this value is assumedin the diffusion coefficient calculations. The boundary layer thick-ness is believed to be fixed for all experiments, because the geom-etry of the cell and the argon flow rate were maintained. Thecalculated diffusion coefficient values are presented in Table 1.The values are almost an order of magnitude lower than previousreported values [23,39]. This difference is probably because of un-der-estimating the boundary layer thickness. The other possibilityis the contribution of sulphate ions specific adsorption to the reac-tion rate, which would affect the calculations. These calculationsshow that as the concentration of hydrogen peroxide is increased,the coefficient of diffusion is decreased. This is most likely becauseof the difference in actual and theoretical concentration of hydro-gen peroxide, as well as the evolution of oxygen, which affect theboundary layer thickness. Hence, the calculated diffusion coeffi-cients at high concentration values of hydrogen peroxide are notreliable. In addition, data show that the diffusion coefficient isincreasing with temperature as expected from Arrhenius law.

This part of the study showed the general kinetic behaviour ofhydrogen peroxide in dilute sodium sulphate solutions. The opencircuit potential values provide an insight into the reactions occur-ring at the surface of an inert electrode at the rest potential. Thekinetic parameters and mass limiting currents extracted fromcathodic polarizations curve would help in the interpretation ofthe effect of hydrogen peroxide on corrosion behaviour of Incoloy800.

3.2. Electrochemical corrosion behaviour of Incoloy 800

Electrochemical corrosion behaviour of Incoloy 800 was evalu-ated by comparing the anodic polarization curves as well as theelectrochemical impedance spectroscopy data in the absence andpresence of hydrogen peroxide in the sulphate solution.

3.2.1. Anodic polarization in the absence of hydrogen peroxideAnodic polarization tests were performed on Incoloy 800 at var-

ious conditions as presented in Table 2. The typical behaviour ofpotentiodynamic curves in sulphate solutions without H2O2, aspresented in Fig. 6 for 25 �C and pH = 7, consists of activation-controlled hydrogen reduction, active dissolution, followed byactive–passive transition region and passivity. It is noteworthy thatin the preliminary experiments, the active–passive region was notobserved when the sample was not cathodically conditioned. Thisis consistent with Carranza and Alvarez results that alloy 800

various conditions of pH, temperature and [H2O2].

Cl)Icorr

(lA cm�2)bc

(mV/dec)ba

(mV/dec)Ipassive

(lA cm�2)Etrans

(V vs. Ag/AgCl)

2.9 816 96 11 0.6342.6 91 337 14 0.6593.1 107 474 10 0.512.1 91 180 13 0.5346.8 119 255 14 0.4899.1 82 207 17 0.4181.5 53 177 10 0.5896.4 202 356 8 –2.1 157 406 – –6.3 220 470 10 0.5645.9 191 385 8 0.5607.1 251 247 8 –6.9 194 196 7 –7.2 197 221 6 –

Fig. 6. Anodic polarization curve of Incoloy 800, in solution containing 25 mMNa2SO4, 25 �C, pH = 7 and no H2O2.

4040 T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045

remains in a passive state at the open circuit potential [40]. Thealloy undergoes transpassivity at higher potentials; however, thetranspassive dissolution is not dominant as potential is sweptand another state of passivity is immediately established. This isshown in Fig. 6 as a decrease in current at potentials higher than900 mV vs. Ag/AgCl. The final increase of current in the polariza-tion curves is associated with evolution of oxygen. It is noteworthythat no evidence of severe localized corrosion was observed on thesample after the potentiodynamic test, indicating that the final in-crease of current was solely due to oxygen evolution.

The effect of pH on the polarization curves was also studied. Thegeneral behaviour of the anodic portion of the curve is not affectedby pH variation. The main impact of increasing pH is to decreasethe corrosion potential and to alter the active region of cathodicwave of the curve. This was expected because the only availablereduction reaction in this medium would be the hydrogen evolu-tion, whose equilibrium potential would decrease as pH increasesand consequently reduces the corrosion potential. The comparisonof the potentiodynamic tests at 25 �C and different pH values arepresented in Fig. 7. In addition, the corrosion current does not sig-nificantly change by altering the pH.

Figs. 8a and b show the effect of temperature on the anodicpolarization behaviour of alloy 800. The general trend was con-

Fig. 7. Anodic polarization of Incoloy 800 in 25 �C, [Na2SO4] = 25 mM, and variouspH.

cluded to be the same as temperature was increased from 25 �Cto 80 �C. Nonetheless, the curves illustrate that the transpassivepotential is significantly decreased by increasing the temperature;the values of transpassive potentials, presented in Table 2, show adecrease from 0.659 to 0.417 V vs. Ag/AgCl. The passive layerformed on Incoloy 800 has been studied in acidic [41,42] and alka-line conditions [43]. In both cases, these studies have showed thatCr2O3 is the main constituent of the passive layer at the onset ofpassivity. The passive layer formed on Incoloy 800 is thus believedto mainly consist of chromium oxide. Therefore, the decrease intranspassive potential can be attributed to the effect of tempera-ture on transpassive dissolution of Cr2O3. The transpassive dissolu-tion reaction is: [44]

2CrO2�4 þ 10Hþ þ 6e� ! Cr2O3 þ 5H2O ð9Þ

The temperature dependant thermodynamic data for water[29], chromate ion [29], and Cr2O3 [45] were used to calculatethe equilibrium potential of reaction (9) as a function of tempera-ture. These values of potentials, referenced to Ag/AgCl electrode atthe specific temperature, are marked in Fig. 8b, which shows thepassive–transpassive region of the potentiodynamic curves at dif-ferent temperatures. In Fig. 8b, the critical current/potential are de-fined by the region between the passive and transpassive states.Fig. 8b and Table 2 show that the critical potential is decreasingand the critical current is increasing whereas the passive currentis slightly increasing as the temperature is increased. Furthermore,the temperature facilitated the hydrogen evolution reaction. This isshown in Fig. 8a as shifting the cathodic portion of the curve to theright, which has ultimately increased the corrosion current densityfrom 2.6 to 9.1 lA cm�2.

3.2.2. Anodic polarization in the presence of hydrogen peroxideThe effect of hydrogen peroxide on polarization curves is pre-

sented in Fig. 9. This shows the evolution of potentiodynamiccurves as H2O2 is added to the solution up to a concentration of10�2 M. It reveals that hydrogen peroxide at 10�6 M level, doesnot affect the overall behaviour; nonetheless, the passive currenthas decreased in the presence of H2O2 from 14 lA cm�2 to10 lA cm�2. The passive layer thickening due to addition of hydro-gen peroxide was observed earlier in the case of Inconel 690 [6],and in the present case, it can probably be associated with thatas well. The thicker passive layer increases the charge transferresistance at a specific potential, which is observed in the polariza-tion curves as a decreased passive current at a specific potential.The potentiodynamic curve in the solution containing 10�4 MH2O2 shows the regions of diffusion controlled reduction of hydro-gen peroxide from �700 to �400 mV vs. Ag/AgCl, activation-con-trolled hydrogen peroxide reduction, anodic dissolution portionleading to passive region from 350 to 500 mV vs. Ag/AgCl, andtranspassivity and oxygen evolution. Subsequent increase in theH2O2 concentration up to 10�2 M leads to the disappearance ofthe active–passive transition region by the hydrogen peroxidecathodic wave in the polarization curve. Therefore, the corrosionpotential (114.5 mV vs. Ag/AgCl) relies in the passive state. Fur-thermore, no distinct region of passivity can be defined in thecurve. It seems that at higher potentials, the anodic oxidation ofhydrogen peroxide is participating in the overall anodic curve,since the reaction is favored at potentials higher than 175 mV vs.Ag/AgCl, which finally masks the expected passive behaviour. Itis noteworthy that the partial currents for hydrogen peroxide ano-dic oxidation at concentration of 10�6 and 10�4 M were so low thatthe higher potential region of the curves in the presence of H2O2 isoverlaid with the curve in the absence of hydrogen peroxide. Theevolution of potentiodynamic curves by increasing the hydrogenperoxide concentration showed that hydrogen peroxide leads to

a b

Fig. 8. Effect of temperature on anodic dissolution behaviour of Incoloy 800 at pH = 7, [Na2SO4] = 25 mM, and [H2O2] = 0, the arrows referring to the calculated redoxpotential for Cr2O3/CrO2

4�. The inset shows the whole polarization curve.

Fig. 9. Effect of H2O2 concentration in anodic polarization behaviour of Incoloy 800at 25 �C, [Na2SO4] = 25 mM, and pH = 7.

T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045 4041

passivation of Incoloy 800 when it is in the concentration rangehigher than 10�4 M. This was consistent with the Bellanger andRameau [6] results regarding the polarization behaviour of Inconel690 in tritiated hydrogen peroxide containing sulphate solutions.They showed that the polarization behaviour was similar at H2O2

concentration lower than 10�3 M. Zhang et al. also reported similarresults in which corrosion potential after addition of 10�3 Mhydrogen peroxide was reproducible regardless of the initial sur-face chemistry state of carbon steel [10].

The effect of pH on polarization curves in the presence of hydro-gen peroxide in the level of 10�4 M was also studied. The resultswere similar to what were obtained in sulphate solution in the ab-sence of H2O2; consisting of slight decrease in corrosion potentialand cathodic wave of the curve, and thus are not presented here.

Studying the effect of temperature shows that the corrosioncurrent densities are less sensitive to temperature in the presenceof than in the absence of hydrogen peroxide; in the former case theIcorr increased from 2.6 to 9.1 lA cm�2, while in the latter case, itfluctuated between 6.4 and 7.1 lA cm�2 (Table 2). This is probablyrelated to the temperature dependent H2O2 kinetic data (Table 1).This illustrates that H2O2 reduction is mass transfer-controlled inthe range of observed corrosion potentials and therefore is therate-determining step. Because it is not practically temperature

dependant, the corrosion current density in the presence of hydro-gen peroxide is not temperature sensitive. Table 2 does not showthe transpassive potential, because the onset of hydrogen peroxideanodic oxidation starts at about 220 mV vs. Ag/AgCl, which pre-vents the measurement of the real transpassive potential.

The cathodic polarization curves presented in the first part ofthe current paper were used to examine how accurately the kineticdata for electrochemical reduction of H2O2 can be used to predictthe electrochemical state of Incoloy 800 in the presence of hydro-gen peroxide. The previously presented curves for cathodic polari-zation of H2O2 on the platinum surface (CP) and anodicpolarization curve for Incoloy 800 in the presence of H2O2 (APH),as well as the anodic polarization curve in the absence of hydrogenperoxide (AP) at the same conditions are presented in Fig. 10a andb. It shows that at [H2O2] = 10�4 M, the curve CP intersects with thecurve AP at two points, one before and one after the active–passivetransition region. Since the sample is first cathodically conditionedbefore the polarization test, intersection before the transitionregion defines the corrosion potential as seen in the APH curve.On the other hand, the corrosion potential lies in the passive regionat H2O2 concentration of 10�2 M, as the CP curve intersects with APcurve in the passive region. Therefore, it seems that the corrosionpotential in the presence of hydrogen peroxide can be explainedby Wagner’s mixed potential theory [46] considering the AP andthe CP curve as overall anodic and cathodic curves. An attemptwas made to indirectly calculate the corrosion potential. Usingthe kinetic data for hydrogen peroxide reduction reaction andthe anodic polarization data for Incoloy 800 in the absence ofH2O2, the corrosion potentials were calculated and compared withthe actual observed Ecorr as shown in Fig. 11. It shows that the cor-rosion potentials are better estimated when they lie in the passiveregion. This is probably because of the contribution of hydrogenevolution reaction at lower potentials. In this case, hydrogen per-oxide reaction will not be the only cathodic reaction participatingin the process and introduces errors in the calculations.

3.2.3. Electrochemical impedance spectroscopyEIS measurements were performed on all the polarization test

conditions, after cathodic conditioning and subsequent OCP mea-surement until a stable potential was achieved. Firstly, the effectof pH and temperature on impedance spectrum was studied. Thecomplex plane graph of specimens subjected to different pH valuesin the absence of H2O2 and 25 �C are presented in Fig. 12. The gen-eral trend by increasing the pH is to increase the impedance value;

Fig. 10. A combined presentation of anodic polarization of Incoloy 800 in presence and absence of hydrogen peroxide and cathodic polarization of hydrogen peroxide on Ptsurface, at 25 �C, [Na2SO4] = 25 mM, and H2O2 concentration of (a) 10�4 M and (b) 10�2 M.

Fig. 11. comparison of calculated and measured corrosion potentials in thepresence of H2O2 at various conditions.

Fig. 12. The complex plane impedance data for Incoloy 800, in 25 �C,[Na2SO4] = 25 mM, [H2O2] = 0, and various pH values.

Fig. 13. The complex plane impedance data for Incoloy 800, in pH = 7,[Na2SO4] = 25 mM, [H2O2] = 0, and various temperatures.

4042 T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045

this may be seen as increasing the radius of a capacitive loop. Thisconfirms the decrease of corrosion rate observed in polarizationmeasurements. Furthermore, at pH = 5, the linear region following

the initial capacitive loop is obvious. This can, as Asselin indicated[47], only be attributed to solution phase mass transfer effects andWarburg type behaviour when the frequency is lower than 1 mHz(not the case here). It is thus related to oxygen vacancy transportwithin the passive layer. This linear region is diminishing as pH in-creases. Probably, the resistance of the passive layer is increased athigher pH and masks the linear region. This point will be discussedmore thoroughly in a follow-up of the present paper, where equiv-alent circuits are used to explain the physical significance of thedata. The trend in experiments at various temperatures (Fig. 13)also shows a general decrease in total impedance value, which isconsistent with corrosion rate data in Table 2.

The impedance data are compared at different hydrogen perox-ide levels in Fig. 14. It shows the focused low impedance region todifferentiate the data groups and the complete set of data is pre-sented as an inset. It is evident that H2O2 has different effects atdifferent concentrations. However, this varied behaviour can beexplained considering the polarization curves discussed in Section3.2.2. The curves are essentially the same at 10�6 M; while thedepression in capacitive loop at 10�4 M could be related to polari-zation data indicating that, the alloy is in the active–passive tran-sition region at this level of concentration. On the other hand, theimpedance magnitude is the greatest at 10�2 M, which shows thatthe alloy is effectively passivated. Furthermore, the low frequency

Fig. 14. Comparison of impedance data at different levels of hydrogen peroxideconcentration at 25 �C, [Na2SO4] = 25 mM, and pH = 7.

a

b

c

Fig. 15. (a) Equivalent circuit used to model the EIS results. Measured andsimulated Bode and Bode-phase diagrams for 25 �C, [Na2SO4] = 25 mM, pH = 7 and(b) [H2O2] = 0 M and (c) [H2O2] = 10�2 M.

T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045 4043

linear relationship disappears at 10�2 M. This is probably becauseof a parallel electron transfer taking place at the passive film/solu-tion interface. The most probable reaction would be reduction/oxi-dation reactions of hydrogen peroxide, as was previously discussedin Section 3.1. The effect of temperature and pH in H2O2 containingsolutions were similar to the trends in solutions without H2O2 andare not presented here.

In addition to the qualitative comparison of data presentedabove, EIS can provide an insight into the mechanisms involvedin the corrosion process. A detailed study of the mechanisms of ac-tive–passive alloys corrosion as affected by hydrogen peroxideusing EIS is ongoing and will be presented in a future publication.An attempt to provide equivalent circuits for the impedance spec-tra, to have a better understanding of the effect of H2O2 on the cor-rosion mechanisms, will be presented in the following.

Carranza and Alvarez [40] used a {R(RQ)} equivalent circuit (Qrepresents a constant phase element) for modelling the impedancespectra of Incoloy 800 in sodium sulphate solutions in the tempera-ture range of 60–280 �C. However, this circuit did not satisfactorilyfit the data and does not explain the action of hydrogen peroxideas redox species and the way it alters the passive film behaviour.Another possible equivalent circuit is the {R(C(R(CR)))(CR)} circuitdeveloped by Bellanger and Rameau [5], which was developed fora duplex stainless steel in solutions containing carbonate and chlo-ride ions as well as the tritiated hydrogen peroxide. This model did

Table 3EIS fitting data for alloy 800 in [Na2SO4] = 25 mM and various conditions of pH, temperature and [H2O2].

Exp no. Temperature (�C) pH H2O2 conc. (M) EOCP

(mV vs. Ag/AgCl)Rsol (X) Cm/f (F) Rm/f (X) Cf (F) Rf (X) QY0 Qn Rf/s (X) Chi-Squared

S � sn

1 25 5 0 �500 36.89 0.0002859 168.4 0.0003432 458.6 0.000716 0.6956 5961 5.12 � 10�5

2 25 7 0 �672 23.49 0.0002645 1.248 0.0007983 7470 0.0002627 0.8443 2135 4.54 � 10�5

3 25 9 0 �704 49.95 0.0004065 1.202 0.0009415 4498 0.0002627 0.841 2826 4.04 � 10�5

4 40 7 0 �617 39.37 0.0005728 214.4 0.001171 7130 0.0005307 0.7747 1449 8.72 � 10�5

5 60 7 0 �591 12.83 0.0004028 85.68 0.001547 3746 0.001065 0.7189 474.6 1.11 � 10�4

6 80 7 0 �422 91.44 0.001538 110.9 0.0007891 2759 0.0008811 0.7738 123.2 3.47 � 10�3

7 25 7 10�6 �659 37.38 0.0005711 280.1 0.001386 11200 0.000494 0.7688 1941 6.30 � 10�5

8 25 7 10�4 �296 33.77 0.0003957 1017 0.0006541 2557 0.0004847 0.7634 664.6 1.26 � 10�4

9 25 7 10�2 121 34 0.000308 12500 0.0001275 2.352 6.76 � 10�5 0.9053 1.54 � 104 4.70 � 10�5

10 25 5 10�4 �355 44.85 0.0002429 1.181 0.002051 9537 0.0002025 0.8568 4488 1.80 � 10�5

11 25 9 10�4 �832 32.78 5.28 � 10�5 0.6229 0.001115 2698 0.0001834 0.8369 4948 8.12 � 10�5

12 40 7 10�4 �226 24.14 0.0003404 2121 0.0007456 6644 0.0004395 0.7669 1065 5.16 � 10�5

13 60 7 10�4 �125 18.15 0.0007948 802.9 0.0002752 8446 0.0003826 0.7738 1723 4.11 � 10�5

14 80 7 10�4 �111 14.83 0.0003638 1179 0.0001703 14300 0.0005421 0.7396 419.7 6.67 � 10�5

4044 T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045

not fit our data, probably because of the effects of carbonate andchloride ions on the electrochemical interface. Furthermore, they as-sumed that two oxide layers (barrier and porous) are established onthe surface, which is not the case here.

The features of EIS spectra are similar to what was observed inthe case of Inconel 625 in ammonia solutions [48] and nickel inphosphate buffer solutions [49]. Priyantha et al. have presentedan equivalent circuit for this situation [50]. It consists of a{Rsol(Rm/fCm/f)(RfCf)(Rf/sQf/s)} circuit, in which contribution of solu-tion resistance, metal/barrier layer electron transfer impedance,bulk passive layer ionic transport properties and film/solutionreactions are considered. This circuit was fitted to the present re-sults with a chi-squared value in the order of 10�5, and therefore,it validates the selection of this circuit. The equivalent circuit ispresented in Fig. 15a. The Bode and Bode-phase diagrams of themeasured and simulated data are presented in Fig. 15b and c.The detailed results of fitting are presented in Table 3. It shows thatas the concentration of hydrogen peroxide is increased, the passivelayer resistance is increased from 280 X cm2 to 12,500 X cm2, indi-cating a passive layer growth. In addition, the exponent in the con-stant phase element is approaching unity. This can be referred to asa pure red/ox reaction on the film/solution interface, which is theanodic decomposition and cathodic reduction of H2O2.

4. Conclusion

This study showed that hydrogen peroxide establishes a mixedpotential on the Pt surface, and that all the mixed potentials lie be-tween the anodic decomposition and cathodic reduction reactionpotential at various conditions. Cathodic polarization data enableus to have an overall understanding of electrochemical behaviourof H2O2 in the sulphate solution. Comparison of anodic polarizationcurves in the presence and absence of hydrogen peroxide allowsthe calculation of corrosion potential based on kinetic data forhydrogen peroxide and Incoloy 800. The impedance data showedthat the resistance and thickness of the passive layer would be in-creased as the concentration of hydrogen peroxide increases. EISresults confirmed the data from polarization curves that an in-crease in temperature and a lowering in the pH would lead to ahigher corrosion rate. The next phase of our study will be to inves-tigate the corrosion at elevated temperature regions.

Acknowledgements

Financial support for this work was provided by Natural Sci-ences and Engineering Research Council of Canada (NSERC) andthe Natural Resources Canada (NRCan) Office of Energy Researchand Development, in collaboration with Atomic Energy CanadaLimited (AECL) through the NSERC/NRCan/AECL Generation IVEnergy Technologies Program.

References

[1] J.W.T. Spinks, R.J. Woods, An Introduction to Radiation Chemistry, John Wileyand Sins, 1990.

[2] R.S. Glass, G.E. Overturf, R.A. Van Konynenburg, R.D. McCright, Gammaradiation effects on corrosion – I. electrochemical mechanisms for theaqueous corrosion processes of austenitic stainless steels relevant to nuclearwaste disposal in the tuff, Corros. Sci. 26 (1986) 577–590.

[3] P.L. Andresen, Effects of transients in water chemistry, temperature, andloading on intergranular stress corrosion cracking of AISI 304 stainless steel,Corrosion 42 (1986) 169–180.

[4] H. Anzai, K. Nakata, J. Kuniya, S. Hattori, Effect of hydrogen peroxide on thestress corrosion cracking of 304 stainless steel in high temperature water,Corros. Sci. 36 (1994) 1201–1211.

[5] G. Bellanger, J.J. Rameau, Effects of tritiated hydrogen peroxide in tritiatedwater containing Cl� and CO2�

3 on behavior of welded type S32550 duplexstainless steel, Corros. Sci. 40 (1998) 1985–2022.

[6] G. Bellanger, J.J. Rameau, The effects of tritiated hydrogen peroxide on Inconel690 passivity, Corros. Sci. 39 (1997) 209–230.

[7] N.J. Laycock, R.C. Newman, J. Stewart, Transpassive corrosion of stainless steelin stabilized alkaline peroxide solution, Corros. Sci. 37 (1995) 1637–1642.

[8] K. Daub, X. Zhang, J.J. Noel, J.C. Wren, Effects of gamma-radiation versus H2O2

on carbon steel corrosion, Electrochim. Acta. 55 (2010) 2767–2776.[9] D. Fu, X. Zhang, P.G. Keech, D.W. Shoesmith, J.C. Wren, An electrochemical

study of H2O2 decomposition on single-phase c-FeOOH films, Electrochim.Acta. 55 (2010) 3787–3796.

[10] X. Zhang, W. Xu, D.W. Shoesmith, J.C. Wren, Kinetics of H2O2 reaction withoxide films on carbon steel, Corros. Sci. 49 (2007) 4553–4567.

[11] Y.-J. Kim, Effect of variations in boiling water reactor water chemistry on type304 stainless steel electrochemical corrosion potential, Corrosion 58 (2002)208–215.

[12] D.D. Macdonald, H. Song, K. Makela, K. Yoshida, Corrosion potentialmeasurements on type 304 SS and alloy 182 in simulated BWRenvironments, Corrosion 49 (1993) 8–16.

[13] J. Sugama, S. Uchida, N. Yamashiro, Y. Morishima, T. Hirose, T. Miyazawa, et al.,Effects of hydrogen peroxide on corrosion of stainless steel, (II) evaluation ofoxide film properties by complex impedance measurement, J. Nucl. Sci.Technol. 41 (2004) 880–889.

[14] S. Uchida, Y. Morishima, T. Hirose, T. Miyazawa, T. Satoh, Y. Satoh, et al., Effectsof hydrogen peroxide on corrosion of stainless steel (VI) – effects of hydrogenperoxide and oxygen on anodic polarization properties of stainless steel inhigh temperature pure water, J. Nucl. Sci. Technol. 44 (2007) 758–766.

[15] T. Miyazawa, S. Uchida, T. Satoh, Y. Morishima, T. Hirose, Y. Satoh, et al., Effectsof hydrogen peroxide on corrosion of stainless steel, (IV) determination ofoxide film properties with multilateral surface analyses, J. Nucl. Sci. Technol.42 (2005) 233–241.

[16] T. Miyazawa, T. Terachi, S. Uchida, T. Satoh, T. Tsukada, Y. Satoh, et al., Effectsof hydrogen peroxide on corrosion of stainless steel, (V) characterization ofoxide film with multilateral surface analyses, J. Nucl. Sci. Technol. 43 (2006)884–895.

[17] L. Tan, X. Ren, K. Sridharan, T.R. Allen, Effect of shot-peening on the oxidationof alloy 800H exposed to supercritical water and cyclic oxidation, Corros. Sci.50 (2008) 2040–2046.

[18] M.G. Alvarez, A.M. Olmedo, M. Villegas, Corrosion behaviour of Alloy 800 inhigh temperature aqueous solutions: long term autoclave studies, J. Nucl.Mater. 229 (1996) 93–101.

[19] A.M. Olmedo, M. Villegas, M.G. Alvarez, Corrosion behaviour of Alloy 800 inhigh temperature aqueous solutions: electrochemical studies, J. Nucl. Mater.229 (1996) 102–114.

[20] M. Fulger, D. Ohai, M. Mihalache, M. Pantiru, V. Malinovschi, Oxidationbehavior of Incoloy 800 under simulated supercritical water conditions, J.Nucl. Mater. 385 (2009) 288–293.

[21] J.M. Joseph, B.S. Choi, P. Yakabuskie, J.C. Wren, A combined experimental andmodel analysis on the effect of pH and O2(aq) on c-radiolytically produced H2

and H2O2, Radiat. Phys. Chem. 77 (2008) 1009–1020.[22] E.J. Calvo, D.J. Schiffrin, The reduction of hydrogen peroxide on passive iron in

alkaline solutions, J. Electroanal. Chem. 163 (1984) 257–275.[23] K.L. Stewart, A.A. Gewirth, Mechanism of electrochemical reduction of

hydrogen peroxide on copper in acidic sulfate solutions, Langmuir 23 (2007)9911–9918.

[24] M.V. Vazquez, S.R. de Sanchez, E.J. Calvo, D.J. Schiffrin, The electrochemicalreduction of hydrogen peroxide on polycrystalline copper in borax buffer, J.Electroanal. Chem. 374 (1994) 179–187.

[25] N.M. Markovic, I.M. Tidswell, P.N. Ross, Oxygen and hydrogen peroxidereduction on the Au(100) surface in alkaline electrolyte. The roles of surfacestructure and hydroxide adsorption, Langmuir 10 (1994) 1–4.

[26] G. Bellanger, Corrosion Induced by Low-energy Radionuclides: Modeling ofTritium and its Radiolytic and Decay Products Formed in Nuclear Installations,Elsevier Science, 2005.

[27] S. Sunder, N.H. Miller, D.W. Shoesmith, Corrosion of uranium dioxide inhydrogen peroxide solutions, Corros. Sci. 46 (2004) 1095–1111.

[28] C. Fonseca, M.A. Barbosa, Corrosion behaviour of titanium in biofluidscontaining H2O2 studied by electrochemical impedance spectroscopy, Corros.Sci. 43 (2001) 547–559.

[29] E.H. Oelkers, H.C. Helgeson, E.L. Shock, D.A. Sverjensky, J.W. Johnson, V.A.Pokrovskii, Summary of the apparent standard partial molal Gibbs freeenergies of formation of aqueous species, minerals, and gases at pressures 1to 5000 Bars and temperatures 25 to 1000 [degree]C, J. Phys. Chem. Ref. Data.24 (1995) 1401–1560.

[30] J. Chao, R.C. Wilhoit, B.J. Zwolinski, Internal rotation and thermodynamicproperties of hydrogen peroxide and its deuterated species, J. Chem.Thermodyn. 3 (1971) 497–505.

[31] H. Hwang, P.K. Dasgupta, Thermodynamics of the hydrogen peroxide–watersystem, Environ. Sci. Technol. 19 (1985) 255–258.

[32] D.T. Sawyer, Electrochemistry for Chemists, John Wiley and Sons Inc., 1995.[33] A.G. Dickson, D.J. Wesolowski, D.A. Palmer, R.E. Mesmer, Dissociation constant

of bisulfate ion in aqueous sodium chloride solutions to 250 degree C, J. Phys.Chem. 94 (1990) 7978–7985.

[34] M.G. Evans, N. Uri, The dissociation constant of hydrogen peroxide and theelectron affinity of the HO2 radical, Trans. Faraday Soc. 45 (1949) 224–230.

[35] V.G. Prabhu, L.R. Zarapkar, R.G. Dhaneshwar, Electrochemical studies ofhydrogen peroxide at a platinum disc electrode, Electrochim. Acta. 26 (1981)725–729.

T. Nickchi, A. Alfantazi / Corrosion Science 52 (2010) 4035–4045 4045

[36] J.T. McClave, F.H. Dietrich, T. Sincich, Statistics, seventh ed., Prentice Hall, NewJersy, 1997.

[37] A. Sun, D.D. Macdonald, Growth and properties of oxide films on platinum II.pH dependence in alkaline solutions, ECS Trans. 2 (2007) 1–10.

[38] R. Memming, Mechanism of the electrochemical reduction of persulfates andhydrogen peroxide, J. Electrochem. Soc. 116 (1969) 785–790.

[39] D.M.H. Kern, The polarography and standard potential of the oxygen–hydrogen peroxide couple, J. Am. Chem. Soc. 76 (1954) 4208–4214.

[40] R.M. Carranza, M.G. Alvarez, The effect of temperature on the passive filmproperties and pitting behavior of a Fe–Cr–Ni alloy, Corros. Sci. 38 (1996) 909–925.

[41] S. Bera, S. Rangarajan, S.V. Narasimhan, Electrochemical passivation of ironalloys and the film characterisation by XPS, Corros. Sci. 42 (2000) 1709–1724.

[42] R.S. Dutta, R. Purandare, A. Lobo, S.K. Kulkarni, G.K. Dey, Microstructuralaspects of the corrosion of Alloy 800, Corros. Sci. 46 (2004) 2937–2953.

[43] K. Hashimoto, K. Asami, XPS study of surface film on nickel alloys in hotconcentrated NaOH, Corros. Sci. 19 (1979) 427–435.

[44] D.D. Macdonald, On the tenuous nature of passivity and its role in the isolationof HLNW, J. Nucl. Mater. 379 (2008) 24–32.

[45] Knovel Critical Tables, second ed. Knovel Publications, 2003.[46] T. Traud, C. Wagner, On the interpretation of corrosion processes through

superposition of electrochemical partial processes and on the potential ofmixed electrodes, Z. Elektrochem. Ang. Physik. Chemie. 44 (1938) 391.

[47] E. Asselin, PhD Thesis, High Temperature and High Pressure Corrosion of Ni-based Alloys and Stainless Steels in Ammoniacal Sulfate Solution, University ofBritish Columbia, 2007.

[48] E. Asselin, A. Alfantazi, S. Rogak, Effect of oxygen on the corrosion behavior ofalloy 625 from 25 to 200 C, J. Electrochem. Soc. 154 (2007) C215–C229.

[49] C.Y. Chao, L.F. Lin, D.D. Macdonald, Point defect model for anodic passive films,J. Electrochem. Soc. 129 (1982) 1874–1879.

[50] D.D. Macdonald, A. Sun, N. Priyantha, P. Jayaweera, An electrochemicalimpedance study of Alloy-22 in NaCl brine at elevated temperature: II.Reaction mechanism analysis, J. Electroanal. Chem. 572 (2004) 421–431.