Influence of pitting and iron oxide formation during corrosion of carbon steel in unbuffered NaCl solutions

Corrosion Science 51 (2009) 971–978

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Influence of pitting and iron oxide formation during corrosion of carbon steel in unbuffered NaCl solutions Luis Cáceres a,*, Tomás Vargas b, Leandro Herrera b a

CICITEM, Centro de Investigación Científico y Tecnológico para la Minería, Departamento de Ingeniería Química, Universidad de Antofagasta, Av Angamos 601, II Region, Antofagasta, Chile b Departamento de Ingeniería Química y Biotecnología, Universidad de Chile, Beauchef 861, Santiago, Chile

a r t i c l e

i n f o

Article history: Received 4 March 2008 Accepted 24 February 2009 Available online 5 March 2009 Keywords: A. Mild steel B. Polarisation C. Oxidation C. Oxide coatings C. Pitting corrosion

a b s t r a c t The corrosion evolution over time of a carbon steel rotating disk immersed in aerated NaCl solutions was analyzed using a superposition model. Using this approach, partial polarization curves for iron oxidation and oxygen reduction were synthesized from experimental current–potential data at different corrosion time in order to determine the kinetics parameters, corrosion potential and current density of the underlying anodic and cathodic subprocesses. The distinctive features of the polarization curves are well described in terms of the simplifying assumptions of the model. In particular, the time evolution of the corrosion current density was linked to the morphology of the corroding surface under different NaCl concentrations. Ó 2009 Elsevier Ltd. All rights reserved.

1. Introduction Iron as carbon steel has been one of the most extensively studied metals in the environment. Under neutral conditions, oxide or hydroxide layers from its corrosion remain on the surface forming distinctive layers. These layers have a significant structure that tends to be determined by the anions present in the solution [1]. Moreover, the corrosion kinetics becomes independent of pH, and hydrogen ion reduction is no longer an important reaction [2–4]. The major reaction governing corrosion in most practical applications is the reduction of oxygen present in solution [1,2]. The rate of corrosion of steels in neutral NaCl solutions is initiated through two main mechanisms: formation and build-up of a passivating iron oxide layer and partial destruction of this layer by pitting [5–7]. It is generally accepted that pitting corrosion is preceded by the appearance of tiny corrosion seeds on the metal surface, which is naturally protected by an oxide layer [8]. In pure dry air at normal temperatures a thin protective oxide film forms on the surface of mild steel. Unlike that formed on stainless steels, it is not protective in the presence of electrolytes and it usually breaks down [5]. Several researchers examined early pit development in mild steel immersed in chloride solutions in the presence of various passivating agents [6,7,9–12]. An increasing tendency toward stable pitting of carbon steel with increasing Cl concentration in neutral NaCl solution was demonstrated by Chen et al., 2000 [7]. It was also shown * Corresponding author. Tel.: +56 55 637342; fax: +56 55 240152. E-mail address: [email protected] (L. Cáceres). 0010-938X/$ - see front matter Ó 2009 Elsevier Ltd. All rights reserved. doi:10.1016/j.corsci.2009.02.021

that pits in carbon steel may tend to occur around the damaged area, leading to accumulated local dissolution [9,10,12,13]. Iron oxide scales deposited on a corroding iron surface are usually a mixture of different oxides [14]. Layer-type rust arises as a result of potential and chemical gradients across the film which varies with film thickness [15,16]. Oxygen diffusion across the oxide scales and subsequent reduction have been studied by Stratmann and Muller, 1994 [17]; their main finding was that oxygen is predominantly reduced within the rust scale and not at the metal/ electrolyte phase boundary. Recently, the analysis of current–potential curves obtained from linear potential sweep technique with a superposition model has been successfully applied to study corrosion of carbon steel in aerated unbuffered NaCl solutions [18]. This methodology made it possible to separately evaluate the influence of the cathodic subprocess, oxygen reduction, and the anodic sub process, iron dissolution, and was successfully applied to characterize the corrosion behaviour of a rust-free carbon steel surface. The present work extends the application of this method to characterize the time evolution of carbon steel corrosion when a homogeneous oxide layer builds up on the metal surface. It will be shown that this approach allows the determination of how the oxide layer formation and pitting process affect corrosion rate evolution. 2. Experimental Experiments were conducted in a conventional three-electrode cell in which the working rotating electrode was made of carbon

L. Cáceres et al. / Corrosion Science 51 (2009) 971–978

40

5h 6h

3h

8h

2h

20

i (A/m2)

steel, the counter electrode was a platinum wire, and the reference electrode was Ag/AgCl (sat. KCl). The working electrode was made of a 4-mm diameter  5-mm length SAE 1010 carbon steel rod. The composition (in wt%) of the studied carbon steel was Fe–98.5, Mn–0.6, C–0.2, and traces of P, S, Si, Sn, Cu, Ni, Cr, and Mo. It was prepared according to a previously described procedure which prevented the influence of crevice corrosion on the measurements [18]. The working electrode was abraded in a rotating plate with wet SiC paper (initially with 400 grade and thereafter with 1200 grade), then degreased by immersion in acetone, cleaned in an ultrasonic bath for 2 min, rinsed with distilled water, and immediately inserted in the cell. The cell operated at 22 ± 0.5 °C using a circulating water bath. A Bioanalytical Systems model BAS100 B/W potentiostat with a BAS/RDE-1 rotating electrode were used for the voltammetry measurements. In each run the abraded electrode was immersed in the air-bubbling electrolyte once the oxygen-saturation condition was achieved. The oxygen concentration in the electrolyte was measured using a Wissenschaftlich-Technische Werkstätten GmbH & Co. KG (WTW) handheld model 340 oxygen meter. The immersed rotating electrode was kept at open-circuit for a period of up to 12 h. During this period current–potential polarization curves were obtained from periodical application of linear potential sweeps at a rate of 1 mV s1, a sweep rate reported to guarantee obtaining steady-state current–potential curves [19]. Potential sweeps started from the cathodic limit and covered the 700 to 100 mV/SHE range. Measurements were made using a working electrode rotating at 170 rad s1 in 0.02, 0.1, 0.5, and 1 M NaCl electrolytes. The electrolytes were prepared by dissolving laboratory grade NaCl in distilled water. In some complementary experiments carbon steel electrodes were removed at different time during the corrosion period in order to conduct morphological characterizations of their surface. In each case, immediately after removal from the cell, the working electrode was washed in distilled water, dried, and stored in a nitrogen atmosphere. The oxidised surface of the working electrode was inspected using an LEO Electron Microscopy Ltd. Model Leo 1420p equipped with a model EDS 7424 elemental analysis accessory.

1h 0h

0

-20

-0.7

-0.5

E (V/SHE)

-0.3

-0.1

Fig. 2. Experimental polarization curves of carbon steel rotating at 170 rad s1 in aerated 0.5 M NaCl solution obtained at different immersion time.

40 4h 9h 12 h

20

2h 6h

i (A/m2)

972

1h 0h

0

-20

-0.7

-0.5

E (V/SHE)

-0.3

-0.1

Fig. 3. Experimental polarization curves of carbon steel rotating at 170 rad s1 in aerated 0.1 M NaCl solution obtained at different immersion time.

40 6h

4h 3h 2h

i (A/m2)

20

3. Results and discussion 3.1. Polarization curves

1h 0h

0

Polarization curves for the carbon steel electrodes obtained at different immersion time are shown in Figs. 1–4. Reproducible polarization curves with well-defined segments were obtained for immersion time not exceeding 12 h in 1, 0.5, and 0.1 M NaCl

-20

-0.7

-0.5

-0.3

-0.1

E (V/SHE) 40

Fig. 4. Experimental polarization curves of carbon steel rotating at 170 rad s1 in aerated 0.02 M NaCl solution obtained at different immersion time.

4h 3h

20

i (A/m2)

2h

1h

0h

0

-20

-0.7

-0.5

E (V/SHE)

-0.3

-0.1

Fig. 1. Experimental polarization curves of carbon steel rotating at 170 rad s1 in aerated 1 M NaCl solution obtained at different immersion time.

electrolytes. For longer time the curves were distorted due to random noise of varying intensity. In 0.02 M NaCl electrolytes the curves showed noise after 6 h of immersion. Iron oxide particle detachment from the corroding carbon steel surface was noticeable as a gradual reddish colour appearing in the electrolyte, which evolved to particle coagulation and settling at the bottom of the cell. The iron oxide fraction remaining attached to the iron surface formed two well-defined layers on the surface: a friable outer layer that could be easily removed by gently scraping or rubbing with a solid object or by induced random turbulence in the solution, and a thin inner layer strongly attached to the surface. These two layers are responsible for the decrease of total current

L. Cáceres et al. / Corrosion Science 51 (2009) 971–978

-0.1

0

2

4

6

8

10

Time h

Ecorr (V/SHE)

-0.2 0.02 M NaCl 0.1 M -0.3 0.5 M -0.4

1M

-0.5 Fig. 5. Variation of corrosion potential with immersion time of carbon steel electrode rotating at 170 rad s1 in aerated NaCl solutions.

density with time observed at every potential value. It was found that after several hours of immersion the current density could be partially restored to the initial value of rust-free surface condition by gently rubbing the outer oxide layer with a paintbrush. The presence of certain elements such as chromium induces a compact inner oxide layer formation which improves the long-term corrosion protection [20]. From the results illustrated in Figs. 1–5, three distinctive trends can be observed to occur consistently during carbon steel corrosion at all NaCl concentrations. First, the anodic current branch becomes steeper at longer corrosion time, which in principle is indicative of an increase in the rate of iron oxidation with time. Second, the cathodic current plateau systematically decreases with increasing corrosion time. A decrease in the cathodic current plateau corresponds to a decrease in the limiting current density for oxygen reduction on the carbon steel surface, which is a result of

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increasing resistance to oxygen diffusion exerted by oxide layer build-up (see Section 3.2). Third, potential values at which the net current becomes zero (corrosion potential) are displaced to more cathodic values with increasing corrosion time and NaCl concentration. The graph of corrosion potential variation with time and NaCl concentration reported in Fig. 5, illustrates this feature. 3.2. Surface morphology of corroded carbon steel electrodes A surface morphology examination of carbon steel electrodes kept at open-circuit in an aerated electrolyte was conducted using SEM. Fig. 6 shows the morphology of the carbon steel surface corroded in 0.02 M NaCl electrolyte. After 2 min of immersion time (Fig. 6, top) some pits become distinguishable, with trails that suggest a wake of fluid emerging from them, in agreement with the description given in former reports [12,13]. The spiral direction of the wake should be a result of the centrifugal action caused by the electrode’s rotation. After 30 min of immersion (Fig. 6, bottom) the size of the pits had increased, but without any significant increase in the number of initially formed pits. In this case there is no oxide propagation towards bare sectors of the metal surface and corrosion remains focused on these separate entities. Fig. 7 shows SEM micrographs of carbon steel surface exposed to 0.5 M NaCl electrolyte at two different immersion times. At 2 min (Fig. 7, top), there is a large number of pits surrounded by a thin iron oxide layer which almost fully covers the electrode surface. Both the number density of cavities and their size are much larger than those seen on the surface exposed to 0.02 M NaCl electrolyte (Fig. 6). At 30 min (Fig. 7, bottom) the accumulated iron oxide has spread over the surface covering all the cavities. Removal of the superficial oxide layer in an ultrasonic bath revealed a number of distinctive pits formed on the corroded surface in 0.5 M NaCl. Fig. 8 shows SEM microphotographs of such rustfree samples after 30 min (Fig. 8, top) and 6 h (Fig. 8, bottom). These micrographs show that at both times there is a wide range of pits that increase in number and size with time. This indicates

Fig. 6. Images of pitting cavities produced on the surface of a carbon steel electrode in 0.02 M NaCl solution. Top figure: 2-min immersion. Bottom figure: 30-min immersion.

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Fig. 7. Surface images of a carbon steel electrode in 0.5 M NaCl solution with an oxide layer. Top figure: 2-min immersion. Bottom figure: 30-min immersion.

that pit formation under these conditions occurs continuously during the experiment while iron oxide builds up over the surface. 3.3. The model In this section an iron corrosion model for a surface covered by an iron oxide layer immersed in neutral aerated NaCl solution is presented. Consider an initial oxide-free polished metal surface immersed in NaCl solution where oxygen reduction is the main oxidation reaction. In the potential range of interest it is assumed that: (a) The oxide layer formed on the metal surface during immersion in NaCl solution is uniform, continuous, and of the single-phase type; (b) hydrogen evolution does not play any significant role in the corrosion current; (c) redox reactions are irreversible; and (d) the corrosion kinetics is independent of pH [3,4]. The major reaction governing corrosion in most practical applications is the reduction of oxygen present in solution. Thus the expression for the total current density i is,

i ¼ iO2 þ iFe

ð1Þ

where iO2 is the current density associated with oxygen reduction, the predominant cathodic reaction, and iFe is the current density associated with iron oxidation. As the oxide layer composed essentially of goethite and lepidocrocite [21,22] builds up, it coats the surface producing rust layer porosity [23]. This porosity may be the result of disordered aggregation of iron oxide particles formed onto the corroding iron surface immersed in neutral NaCl solutions. In contrast, in the presence of a passivating agent the iron oxide layer can become highly ordered and compact [24]. If a connection with the underlying metal exists, then iron oxides on the surface of the pores can act as a cathodic area [17], and oxygen can then be reduced on that surface. The cathodic and anodic areas are decoupled, with the oxidation of iron taking place on

the metal in contact with electrolyte at the bottom of the pores, and the reduction reaction on the large cathodic area formed by the iron oxides. The actual mechanism of oxygen reduction in the presence of iron oxides has not yet been clearly resolved [23]. Assuming that the cathodic reaction of oxygen reduction is under mixed diffusion and charge-transfer control, then for reaction under a surface condition covered by an oxide layer, a model similar to that of a bare iron surface under neutral conditions is proposed [19,25]:

 m   gO iO iO2 ¼ i0O2 1  2 exp 2:3 2 ilO2 tc

ð2Þ

where, ilO2 (Am2) is the limiting current density for oxygen reduction; i0O2 (Am2) is the exchange current density for oxygen reduction; m is the order of the oxygen reduction reaction on a solid surface; gO2 ¼ E  EeqO2 (V) is the oxygen reduction overpotential; E (V) is the applied potential; EeqO2 (V) is the equilibrium potential for oxygen reduction; t c ¼ 2:3RT anF is the cathodic Tafel slope expressed as V dec1; R = 8.314 J K1 mol1; T (K) is the temperature; a is the transfer coefficient; n is the number of electrons involved in the cathodic reaction, and F is the Faraday´s constant (96,500 C mol1). Eq. (2) considers a linear oxygen concentration profile between the bulk electrolyte phase and a fixed plane within the oxide film where a charge transfer reaction takes place. This assumption implies that the oxide layer has to be uniform, continuous, and of the single-phase type. Strictly speaking, even for pure metals, this assumption is rarely valid. The well-defined steady flow pattern in a rotating electrode may contribute to uniform layer formation. At certain negative potential (where both, the iron oxidation and hydrogen evolution are negligible), the diffusive transport of oxygen through both, the hydrodynamic boundary layer and corrosion product layers is the rate-limiting step of the corrosion pro-

L. Cáceres et al. / Corrosion Science 51 (2009) 971–978

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condition [28]. In view of this situation, two alternative values for m, 0.5 and 1, were adopted for parameter determination. These values have been previously used for the kinetics expressions of oxygen reduction on the surface of various metals [19,26]. Rearranging the partial current density, iO2 , in Eq. (2) in order to obtain an explicit equation, and using Eq. (3), a convenient expression of the total current density (Eq. (1)) for each m value is: m = 0.5

i¼

  i0O2 e exp 2:3gO2 =tc h

  i0O2 exp 2:3gO2 =t c 2ilO2 rffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi #   2  2 þ i0O2 exp 2:3gO2 =t c þ 4 ilO2 þ i0Fe exp

  2:3gFe ta

ð4Þ

m=1

    i0O2 exp 2:3gO2 =tc 2:3gFe   þ i i¼ exp 0Fe   ta 1 þ i0O2 =ilO2 exp 2:3gO2 =tc

ð5Þ

Equivalent forms of Eq. (5) have been used in former reports for mathematical description and modelling of iron corrosion processes [26,29]. The determination of parameters i0O2 ; i0Fe ; tc ; t a , and ilO2 were carried out from experimental current–potential data according to the numerical method reported previously [18]. These values are shown in Table 1. The quality of the fit for a given set of experimental current–potential data was evaluated using the root mean square error (RMSE) [18]. 3.4. Fitted parameters and model application Fig. 8. Surface images of a carbon steel electrode in 0.5 M NaCl solution. Top figure: after 30 min immersion; Bottom figure: after 6 h immersion. The oxide layer was removed by an ultrasonic bath.

cess. This value is represented by the intermediate plateau in a polarization curve. The maximum value of the limiting reduction current ilO2 applies at the initial stage of corrosion where oxide layer is absent. As thickness of the oxide layer film influences the inward diffusion of oxygen, therefore, as it grows, the ilO2 value should become attenuated. A decrease of ilO2 at increasing immersion time is observed in Figs. 1–4. A gradual current density increase in the range of the intermediate plateau at more negative potentials may be due to a combination of hydrogen evolution and iron oxide reduction. The iron kinetic expression for iFe at the metal surface is assumed as [19,26]:

  g iFe ¼ i0Fe exp 2:3 Fe ta

ð3Þ

where, i0Fe (Am2) is the exchange current density for oxidation; gFe = E  EeqFe (V) is the iron oxidation overpotential; EeqFe (V) is 2:3RT the equilibrium potential for iron oxidation; and ta ¼ ð1 aÞnF is the anodic Tafel slope expressed as V dec1. A ferrous iron concentration value of 1106 M was arbitrarily assumed for the calculation of EeqFe following the suggestion of Bockris [27]. The corrosion potential Ecorr and corrosion current density icorr were calculated from expressions (3)–(5) under a condition of null-current density (i = 0). Due to the lack of convergence, exponent m of Eq. (2) could not be obtained using the numerical method for parameter determination. This value, which is related to the oxygen reduction mechanism, could vary during corrosion according to the surface

Electrochemical kinetics parameters for all experimental polarization runs are listed in Table 1. Referring to 0.5 M NaCl experiments at different immersion time, Fig. 9 shows the fitted polarization curves as dashed lines, which were calculated from the parameters listed in Table 1 using Eqs. (4), (5) with values for EeqFe and EeqO2 equal to 0.614 and 0.762 (V/SHE) respectively. For comparison, the corresponding experimental current–potential curves are also plotted as continuous lines in the same graph. The initial polarization curve under a rust-free surface condition had a best fit when using Eq. (4) (m = 0.5), while for curves with immersion time larger than 1 h, best fits were obtained using Eq. (5) (m = 1). In some cases the quality of the fit obtained from Eqs. (4) and (5) was quite similar. An example of this situation for carbon steel in 1 M NaCl after 3 h immersion is shown in Fig. 10. In the range of coincidence O–A the best quality of the fit cannot be visually resolved for any of the fitted curves which showed different RMSE values. As seen in Fig. 9, the early experimental polarization curve shows at point C, located in the anodic branch, a noticeable deviation from Tafel’s law. The pitting activity at more positive values than a limit known as ‘pitting potential’ has been characterized by noting the potential at which a sudden increase of the anodic current takes place [30]. Thus, this deviation between the experimental and the fitted values in the anodic branch could be identified as a superimposed pitting current corrosion density. It is interesting to note that at an immersion time of 2 h or longer such deviation between the modelled and the experimental curve is not observed. In this case a large fraction of the pits on the carbon steel surface that have become covered with rust are already established and the variation of the dissolution area during the sweep is apparently negligible. In the case of the simulated cathodic branch, modelled assuming a mixed charge transfer–mass transfer controlled reaction, fitting is good only up to point ‘A’. After that point the

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L. Cáceres et al. / Corrosion Science 51 (2009) 971–978

Table 1 Electrochemical parameters calculated from experimental current–potential data by using a numerical method reported in a previous work [18]. NaCl (M)

i0Fe (Am2)

Immersion time (h)

ta (mV dec1)

i0O2 (Am2)

tc (mV dec1)

ilO2 (Am2)

Ecorr (mV)

icorr (Am2)

0.02

0 1 2 3 4 6

0.0134 0.0415 0.0584 0.0273 0.0005 0.0003

206 217 197 163 94 82

6.0E-06 1.1E-05 8.2E-12 2.8E-17 2.5E-04 1.7E-11

177 168 71 51 221 90

14.5 12.0 7.1 6.7 7.7 7.4

198 194 206 225 243 283

1.4 3.6 6.9 6.6 4.0 3.7

0.1

0 1 2 4 6 9 12

0.0019 0.0200 0.0359 0.0019 0.0015 0.0041 0.0042

142 171 169 104 98 104 105

3.2E-07 5.9E-07 1.1E-13 2.9E-24 4.9E-23 2.7E-30 5.4E-35

140 140 70 42 45 35 31

12.7 9.1 5.9 3.3 2.7 2.0 1.4

191 226 250 280 299 336 350

1.8 3.7 5.1 3.1 2.5 2.0 1.4

0.5

0 1 2 3 4 5 6

0.0405 0.0029 0.0002 0.00004 0.0007 0.0027 0.0051

177 107 74 62 77 85 90

2.4E-03 1.4E-12 2.4E-12 9.8E-18 4.7E-19 4.4E-07 5.2E-23

278 75 82 62 58 139 48

12.4 5.2 3.2 2.1 1.7 1.2 1.2

238 269 301 328 360 388 403

5.4 4.9 2.8 1.4 1.5 1.2 1.1

1

0 1 2 3 4

248 81 72 71 84

3.5E-06 2.6E-01 2.3E-02 1.6E-01 3.8E-02

141 886 658 1151 677

8.0 3.2 4.0 2.3 1.7

271 306 372 383 415

7.2 1.8 0.9 0.9 0.9

0.2996 0.0003 0.0004 0.0005 0.0039

interference of a secondary cathodic reaction, hydrogen evolution and/or oxide reduction become significant.

20 6h

i (A/m2)

10

4h

2h

0

3.5. Corrosion current and its dependence on anodic and cathodic subprocesses

C 0h

A' S'

-10

S

A

-20 -0.7

-0.5

E (V/SHE)

-0.3

-0.1

Fig. 9. Experimental (continuous) and fitted (dashed) polarization curves in aerated 0.5 M NaCl solution at different immersion time.

2

Values of corrosion current density as a function of time calculated using the model for different NaCl concentrations are plotted in Fig. 11. Two patterns of dependence of corrosion current density with time can be seen. In the high NaCl concentration range, 0.5 and 1.0 M, corrosion current density steadily decays with time during the whole experiment. Even though corrosion current density in 1 M NaCl solutions decreases faster than in 0.5 M NaCl solutions, after 4 h the corrosion current density in both cases is in the same range, about 1 Am2. In the low NaCl concentration range, 0.02–0.1 M, the corrosion current density initially increases, reaches a maximum at about 2–4 h, and then starts dropping. The variation of corrosion current density with time shown in Fig. 11 can now be explained in terms of the variation of the kinetics of the anodic subprocesses and the morphological observations

8

O

6

icorr (A/m2)

i (A/m2)

1

0 Fitted curve m = 0.5

A

-1

0.02 M NaCl

4 0.1 M

2 Fitted curve m = 1

-2 -06.0

0.5 M

Experimental curve

-05.0

1M

-04.0

E (V/SHE) Fig. 10. Experimental (dotted) and fitted polarization curves using Eqs. (4) and (5) for aerated 1 M NaCl solution at 3 h immersion time.

0 0

2

4

Time h

6

8

10

Fig. 11. Variation of corrosion current density with immersion time of carbon steel electrode in aerated NaCl solutions determined from the superposition model.

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L. Cáceres et al. / Corrosion Science 51 (2009) 971–978

i Fe

iFe, IO2 (A/m2)

i O2

3h

2h

5 6h

1h

4h

-0.5

-0.4

-0.3

-0.2

-0.1

E (V/SHE)

0h

Fig. 15. Simulated current–potential curves for anodic and cathodic subprocesses in 0.02 M NaCl solutions at different corrosion time.

5

4h

3h 1h 2h

0 -0.6

-0.5

-0.4

-0.3

-0.2

E (V/SHE)

-0.1

Fig. 12. Simulated current–potential curves for anodic and cathodic subprocesses in 1 M NaCl solutions at different corrosion time.

15 i Fe

10

iFe, iO2 (A/m2)

10

0 -0.

i O2

i O2 0h

5

1h 2h 8h

0 -0.6

-0.5

6h

4 h

-0.4

3h

-0.3

-0.2

-0.1

E (V/SHE) Fig. 13. Simulated current–potential curves for anodic and cathodic subprocesses in 0.5 M NaCl solutions at different corrosion time.

15

i Fe i O2

iFe, iO2 (A/m2)

i Fe

0h

10

10

2h 1h

5 9h

6h

12 h

0 -0.6

15

iFe, iO2 (A/m2)

made on the corroded carbon steel electrodes. Figs. 12–15, for 1, 0.5, 0.1, and 0.02 M NaCl, respectively, show current–potential curves calculated for the anodic and cathodic subprocesses at different corrosion times. These curves were calculated using the parameters from Table 1 for each case. Each dot in Figs. 12–15 indicates the intersecting points of the anodic and cathodic curves at a given corrosion time, and define the corresponding corrosion current density and corrosion potential values.

-0.5

-0.4

4h 0h

-0.3

-0.2

-0.1

E (V/SHE) Fig. 14. Simulated current–potential curves for anodic and cathodic subprocesses in 0.1 M NaCl solutions at different corrosion time.

The curves for the anodic subprocess, iron dissolution, were calculated according to Eq. (3), which corresponds to a reaction with pure charge transfer control. It is seen that the rate of the anodic subprocess increases steadily with corrosion time at the four NaCl concentrations studied. From morphological observations of the corroded steel surface, this behaviour can be related to an increase of the iron dissolution area resulting from an increase in the number and size of pits formed with time (see Figs. 6–8). This trend is shown as a significant decrease of the ta value at longer corrosion times (see Table 1). Curves for the cathodic subprocess, oxygen reduction, were calculated according to Eq. (2), which corresponds to a reaction with mixed charge transfer and mass transfer control. It can be seen that the rate of the cathodic subprocess follows two main trends with corrosion time increase. First, at high NaCl concentrations, 0.5 and 1 M, the cathodic current is predominantly mass-transfer controlled, and steadily decreases with corrosion time due to the rapid decay of the limiting current density. From morphological observations of the corroded carbon steel, this behaviour can be explained by the rapid formation of a continuous surface layer of iron oxide which establishes a regime of planar diffusion control for oxygen reduction. In this case the rapid drop of the limiting current determines the steady drop of the corrosion current density (see Figs. 12 and 13). This implies that the inhibitory influence of the formation of a continuous oxide layer predominates over the increase of the iron dissolution area due to the increase in the number and size of the pits (see Fig. 8). Finally, at low NaCl concentrations, 0.02 and 0.1 M, there is a significant potential range (0.2 to 0.3 V/SHE) where the cathodic current is predominantly charge-transfer controlled and increases with corrosion time. This trend can be related to the increase in the size of the pits formed initially, which here grew as separate entities far apart from each other, without merging into a continuous oxide layer. This initial increase of the cathodic current results in the increase of the corrosion current observed during the first 2–3 h at 0.1 and 0.02 M NaCl (see Fig. 11). Once there is build-up of an isolated oxide cap on each pit (see Fig. 6, bottom), oxygen reduction becomes mainly mass-transfer controlled. The cathodic limiting current then starts decreasing steadily, a factor which predominates and triggers a steady drop of the corrosion current (see Figs. 14 and 15). Experimental observations that suggest similar effects on corrosion have been described in former investigations. Frangini and De Cristofaro (2003) [31] described a galvanostatic experiment conducted with stainless steel in a pit promoting electrolyte under deaerated conditions, the resulting chronopotentiometric curve shows a maximum potential value; such behaviour was linked to the sum of two competing processes: passivation and pit nucleation. Lalvani and Zhang (1995) [32] observed a maximum in a current density–time experimental data

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from a carbon steel electrode immersed in 3.33% NaCl solution. This system was under nitrogen purge and a constant peak voltage of a positive half-cycle rectified alternative voltage signal. In the present research the described behaviour, can be explained discarding cathodic curves for oxygen reduction in Figs. 12–15. In fact, a maximum potential value in time emerges in Figs. 12 and 13 when a horizontal line representing a constant current density value (lower than 5 Am2) intercepts anodic curves at various immersion times. Also a maximum current density emerges when a vertical line representing a constant potential intercepts anodic curves in Figs. 14 and 15. These considerations imply that both, partial polarization curves for iron oxidation and hydrogen evolution are not significantly affected by dissolved oxygen in the NaCl solution. 4. Conclusions The analysis of the evolution in time of carbon steel corrosion in NaCl solutions using a superposition model for periodic polarization curves allows a kinetics characterization of the system influenced by iron oxide build-up. The results of this analysis showed that the anodic subprocess, dissolution of iron, is well described in terms of a pure chargetransfer controlled kinetics. On the other hand, the cathodic subprocess, oxygen reduction on iron, is well described in terms of a mixed mass-transfer and charge-transfer controlled kinetics. Model analysis indicated that progressive pitting results in an increase of the rate of anodic dissolution of iron, while iron oxide formation results in a decrease of the oxygen reduction rate due to a rapid decrease of the limiting current. These two factors determine corrosion rate as follows: In the high NaCl concentration range, 0.5–1 M, the rate of carbon steel corrosion continuously drops with time, a behaviour determined by the inhibitory influence of the formation of a continuous surface oxide layer which reduces the limiting current density for oxygen reduction. This factor predominates over the increase of the iron dissolution area related to the steady increase in the number and size of the pits. In the low NaCl concentration range, 0.02–0.1 M, the rate of carbon steel corrosion initially increases, as determined by the size increase of the initially formed pits. However, corrosion rate reaches a maximum at about 3 h and starts to drop when the initially formed pits, which keep growing as separate entities, become covered by an iron oxide cap that inhibits oxygen diffusion. Acknowledgment Financial support under FONDECYT project 1070930 carried out at the Facultad de Ingeniería of the Universidad de Antofagasta, Antofagasta, Chile, is greatly appreciated. References [1] M. Fischer, Possible models of the reaction mechanism of the function of passivating additions during the chemical passivation of iron in weakly acid, neutral and basic pH ranges, Werkst. Korros. 29 (1978) 188. [2] B.J. Andrzejaczek, Influence of dissolved oxygen in the iron corrosion kinetics in water, Korrosion 15 (1984) 239.

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