Experimental Demonstration of Corrosion Phenomena The Corrosion, Passivation, and Pitting of Iron in Aqueous Media Omar Solma and Lucrecia Olivares CINVESTAV-IPN, Ap. Postal 14-470. 07000 MBxlco. D. F., MBxlco Jorge 0. Ibaiiez' Unlvenldad Iberoamerlcana, Depto. de ICQ, Prol. Paseo de la Reforma 880, 01210 MBxlw, D. F., MBxlco Metal corrosion is a problem of primary concern to our contemoorarv societv. According to estimates (1-4, the economic losses due to corrosion may represent aa much as 3 or 4% of the moss national ~ r o d u c(GNP) t of several countries. Not onlyUis the economic aspect of importance, but also several other aspeds are frequently affected; for example, from an industrial perspective, it is necessary to considerihe safety of the personnel and of the properties, the reliability of the equipment, the energy waste, the preservation of materials, and the environmental impact (5). A considerable of such . . - ~fraction ~ ~ ~ oroblems ~ ~ mav be avoided if we educate our present-day add future prdfessionals to understand the factors involved in corrosion Drocesses and their prevention. The theoretical discussions in textbooks have been relatively scarce as have the design and implementation of didactic lab experiments (6-9). With the followine exoeriments, students taking inorganic chemistry, electrochemistry, corrosion science,physical chemistry or materials science will be able to ohsewe and analyze the effect that electric potential, the pH, and the concentration of aggressive ions (e.g., C1-) have upon the corrosion, passivation, and pitting of iron electrode in aqueous media.
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The electrochemical nature of metal corrosion in aqueous media has been madeevident and has been widelvdiscussed in this Journal (see, for example, refs 3,6,10, l l j . In eeneral, a metal (M) ex~osedto an aqueous medium an electric potential may: (1) formsoluble species and (e.g., Mn+, M(OH),("-'I+, M(OH),b-n)-), (2) form insoluble species (e.g., M,O,,M(OH),), or (3) remain unaltered (M). These reactions give rise to the E-pH (Pourbaix) diagrams that present the thermodynamic equilibria in aqueous solutions in terms of the predominant species (12) and which have been widely analyzed, exemplified, and discussed in this Journal (see, for example, refs 13-17). In terms of corrosion. case (1) . . corresoonds to reeions of active corrosion. case (2) of passivation, and case (3) of immunity. The simplified E-pH diagram (adapted from ref 17) in t e r n s of the behavior of Fe a t room temperature, is shown in Figure 1. Thus, when a voltage scan is initiated (anodic direction) on an iron electrode a t constant pH, the Fe is oxidized firstlv to Fez+ in the active dissolution region (corrosion). As the sweep continues, an oxide layer is formed that prevents any further corrosion under these conditions (passivity). Finally, a potential where either this layer breaks down or else the solvent decomposes (or both) is reached. The anodic polarization curve thus obtained will show one of the forms indicated in Firmre 2 (18). .deoendine . -uoon . the solution composition. In the present experiment, several of
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Figure 1. Simplitled potential-pH (Pourbaix) dlagram showing me conasion behavior of Fe at room temperahre (adapted from ref 17). Dotted lines: experimentel points for iron in 1.25 M KNOa showing the effect of the pH solution upon: (a) h mixad mosion potential (h).(b) the passivation In the pH range potential (Ed,end (c)Ute water dewmporition potentla1 (6). 0.5-6.2.
these curves are produced with the aid of a potentiostat in order to obtain the variation of the mixed corrosion poten, of tial (EM),of the passivation (or Flade) potential ( E F )and the solvent decomposition (ED)as functions of pH, and to understand such variations with the aid of the E-pH diagram. The last series of experiments shows that the passivated metals may undergo a localized breakdown of their passivating layer, due (among other possible factors) to the action of aggressive ions (e.g., halides). Such ions preferentially attack the heterogeneities or disslocations of the passivating layer (oxide or hydroxide) and finally reach the metal, thus initiating the process of pitting (I). ExpMlmental Several solutions of supporting electrolyte (1.25 M RN03 were prepared; the pH was adjusted with dil. HN03. Analytical-grade reagents were used without further purification. In the pitting exVolume 68 Number 2
February 1991
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Figure 3. Sample anodic polarllation curves for iron in 1.25 M KNOS.(a) pH = 1.5. (b) pH = 2.6. (c) pH = 6.2 ( i X 10).
Figwe 2. Differentfarmsof an anodic poladration m e : 1, active dissolution: 2, passivatlon: 3, passive region: 4, pitting: 5, transpassivation: 6. solvent oxidation: 7. oxide growth.
periments the required amounts of a NaCl solution where added to produce the desired concentrations. Once inside the cell, the soluiions were deaerated with Nz before eaeh run. A conventional fourelectrode cell (working, reference, and two auxiliary electrodes), with an approximate capacity of 25 mL, was used. A typical experimental setup was used (see, for example, the following references in this Journal: 19,20,21). Thin disks of high-purity Fe (cut from an iron foil) with an approximate area of 0.05 em2 were pressuremounted on a cylindrical Teflon matrix and forced to make contact with a copper wire inside the matrix; such disks were mirror-polished with very fine sandpaper and washed with deionized water before each run. Two graphite rods were used as auxiliary electrodes, with an area of 5 cm2eaeh, in order to obtain a good distribution of the current lines. The potential was measured versus a saturated calomel electrode (SCE), controlled with a Tacussel 202X potentiostat and measured with a homemade current-to-voltage converter (22);the i vs. E signal was plotted with a Houston 2000 XY recorder. However, practically any simpler equipment could he used as well. The anodic polarization curves were obtained by sweeping the potential at a rate of 1 mV 8-1, from the equilibrium potential to the oxygen production potential for each solution. Results and Dlscuscrion T h e anodic polarization curves obtained for Fe (in 1.25 M KNO.,)at different valuesof pH are shown in Figure 3. It can be noticed t h a t these curves show an active dissolution zone, followed by t h e formation of a passivating layer, a passive zone, and water oxidation. Notice t h a t the potential values plotted directly from the potentiostat are given in volts vs. SCE, whereas t h e values in t h e E-pH diagram are generally given vs. S H E (standard hydrogen electrode). T o convert E vs. S C E t o E vs. SHE, simply add 0.244 V (T = 25 'C).
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Figure 4. Anodio polartration curves for iron in 1.25 M HNOl at pH = 2.8. wlth differentCI- concentration: (a) 0 ppm, (b) 2500. (c)2750, (d) 3000.
According t o the E-pH diagram (Fig. I), the Fe corrosion potential should not he pH-dependent; however, the true process occurs simultaneously with the production of Hz, thus yielding a mixed corrosion potential (EM), which is cathodically displaced with a pH increase. This general tendency can he observed in the dotted line, inFigure 1.TheEpH diagram also predicts a cathodic displacement for both the passivation potential (EF)and the anodic decomposition potential of water (ED);both tendencies are confirmed by the results ohtained from the potentiodynamic curves (see dottedlines b a n d c in Fig. 1). The discrepancies between the predicted and the obtained potentials are mainly due t o the fact that theE-pH diagrams are based upon thermodynamics data, whereas the results ohtained herein were obtained from kinetic curves ( I ) ; to illustrate such phenomena better, i t is sufficient to recall that the thermodynamic potentials for the decomposition of water are frequently very different from the experimental ones due to the different overpotentials required for the oxygen production on each metal; in addition, the thermodynamically favored products may he kinetically controlled and yield different products (e.g., oxides instead of hydroxides) (1) resulting in the "displacement" of the equilibrium lines. As for the Fe pitting, the curves obtained upon increasing the [Cl-] are shown in Figure 4. As can he noticed, the presence of C1- produces a peak located a t 0.3-0.4 V vs. SCE, which does not exist in the absence of C1-. The magnitude of such a peak increases with the increase of C1-, thus proving that this peak is due to the presence of the C1- ions. The students may verify whether the area under peak is proportional to the C1- concentration. Conclusions
The potentiodynamic curves may he good tools in helping the students understand the phenomena involved in corrosion, passivation, and pitting of metals in aqueous media, as well as the decomposition of the solvent. Even though these curves are not ohtained under rigorous equilihrium conditions, there is a general and illustrative (although not exact)
concordance with the results that can be predicted from the E-pH diagrams, which involve thermodynamic equilibria.
Acknowledgment
We wish to thank Jose A. GonzHlez (CENIM-Madrid) for helpful comments. This paper is dedicated to Wayne Wentworth (of the U. of Houston) on the occasion of his 60th birthday. One of us (JGI) acknowledges partial support from project INQ-050 of the U. Iberoamericana and from Dow Quimica Mexicana.
Literature cned L. Cufler,A.J.B.:Tuek,C.D.S.;~eld,S.P.:Williams.D.E.Ch~m.Br.1986,22,1109. 2. GonzBez, .I. A. Teorio y P16rtico de lo Lucho contra lo Corroaibn; CS.1.C.C.E.N.I.M.: Madrid. 1984;Chapferl. 3. Laurm.P. M. J. Chem.Educ. 1978,55,319. 4. Uh1ig.H. H.; Revie. R. W. Corrosion and Corrmion Control; Wiley: New York, 1 9 8 5 ; ~
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