the influence of hydrogen sulfide on corrosion of iron under different conditions

The in¯uence of hydrogen sul®de on corrosionof iron under di�erent conditions

Houyi Maa, Xiaoliang Chenga, Guiqiu Lib, Shenhao Chena,*,Zhenlan Quana, Shiyong Zhaoa, Lin Niua

aDepartment of Chemistry, Shandong University, Jinan, 250100, People's Republic of ChinabDepartment of Photo-Electronic Information Engineering, Shandong University, Jinan, 250100, People's

Republic of China

Received 8 January 1999; accepted 20 November 1999

Abstract

Hydrogen sul®de (H2S) can either accelerate or inhibit corrosion of iron under di�erentexperimental conditions. What H2S has done to both the anodic iron dissolution and

cathodic hydrogen evolution, in most cases, is to have a strong acceleration e�ect, causingiron to be seriously corroded in acidic medium, but H2S can also have a strong inhibitionon the iron corrosion under certain special conditions where H2S concentration is below0.04 mmol dmÿ3, pH value of electrolyte solution is within 3±5 and the immersion time of

the clectrode is over 2 h. The inhibition e�ect of H2S on the iron corrosion is attributed toformation of ferrous sul®de (FeS) protective ®lm on the electrode surface. Moreover, thestructure and composition of the protective ®lm is closely related to H2S concentration, pH

of solutions and the immersion time of iron. Accordingly, the in¯uence of the three factorson the inhibition e�ect is investigated in this paper by means of AC impedance technologytogether with the potentiostatic steady-state polarization. A probable reaction mechanism is

proposed to interpret theoretically how H2S inhibits the corrosion of iron. 7 2000 ElsevierScience Ltd. All rights reserved.

Keywords: Iron; Acid corrosion; Acid inhibition

0010-938X/00/$ - see front matter 7 2000 Elsevier Science Ltd. All rights reserved.

PII: S0010 -938X(00)00003 -2

Corrosion Science 42 (2000) 1669±1683

www.elsevier.com/locate/corsci

* Corresponding author. Fax: +86-531-8565167.

E-mail address: [email protected] (S. Chen).

1. Introduction

Recently, we have investigated the electrochemical behaviours of iron,chromium, nickel in the acidic solutions containing H2S, by means of thepotentiostatic steady-state technology, and AC impedance method and found thatH2S can strongly accelerate corrosion of these metals [1±6]. Among the metals, weare particularly interested in corrosion behaviour of iron in the acidic solutionswith H2S. Results have shown that H2S exhibited an acceleration to both theanodic iron dissolution and the cathodic hydrogen evolution in most cases [1±3,7,8], but under the certain special conditions where the lower H2S concentration(R0.04 mmol dmÿ3), pH value of 3±5, and the longer immersion time (r2 h) aremet simultaneously, what H2S does to the corrosion of iron is not an accelerationbut a strong inhibition role [1,9±11]. This kind of inhibition e�ect is related to theformation of FeS with di�erent crystal shapes, such as pyrite, troilite andmackinawite [12]. It is well known that the dissolution of iron in acidic mediumhas been extensively studied for a long time. However, the systematic reportsabout the inhibiting e�ect of H2S on iron corrosion in acidic solutions have notbeen encountered so far. On the basis of previous studies [1±3], how H2S inhibitsthe iron corrosion is interpreted in terms of a probable reaction mechanism.Additionally, the in¯uence of pH value, H2S concentration and immersion time ofthe electrode on this role is also investigated by using AC impedance technique,together with the potentiostatic steady-state polarization technology.

2. Experimental

A cylindrical iron electrode was used in experiments, with 3.14 cm2 cross-sectional area, directly made from 99.99% pure iron rods of 5 mm diameter(Johnson Matthey). It was embedded in the epoxy resin mold and only its cross-section was allowed to contact the electrolyte. Prior to each experiment, theelectrode surface was polished with #600 and #1200 emery papers in proper order,then rinsed with alcohol and triply with distilled water.

Experiments were conducted in the potentiostatic mode at 222 0.28C using athree-electrode cell. The counter electrode was a large piece of platinum blackwith a surface area of over 4 cm2 and reference electrode was a saturated calomelelectrode (SCE). All potentials reported in this paper were measured with respectto an SCE.

All concentrated solutions, including 1.0 mol dmÿ3 H2SO4, 1.0 mol dmÿ3

Na2SO4 and 0.4 mol dmÿ3 Na2S were prepared from analytical reagent gradechemicals and triply distilled water. The dilute solutions of di�erent pH wereprepared by mixing various proportions of the H2SO4 and Na2SO4 concentratedsolution. The H2S-containing electrolyte solutions were prepared from theconcentrated Na2S solution and the H2SO4 solutions in the following way: theH2SO4 solution was at ®rst purged of oxygen by degassing with high-purity

H. Ma et al. / Corrosion Science 42 (2000) 1669±16831670

nitrogen for 30 min, and then the suitable volume of the Na2S solution was addedto the solution [1].

Excited by a sinusoidal perturbation signal of 5 mV amplitude, impedancemeasurements were made at the potential of interest in the frequency range from65 kHz to 10 MHz at 5 points per frequency decade. The impedancemeasurements system consists of a Solartron 1286 Electrochemical Interface and1250 Frequency Response Analyzer controlled by a Hewlett Packard 9122computer, and the measured impedance spectra were drawn by a 7475A plotter.

3. Results and discussion

3.1. Acceleration e�ect of hydrogen sul®de

3.1.1. Characteristics of the steady-state polarization curves of iron in H2S-containing acidic solutions

In recent years, the corrosion behaviour of iron in acidic solutions with H2Shave been studied in detail in our laboratory [1±3]. The experimental results haveshown that addition of H2S into electrolytes greatly accelerates both anodicdissolution current and cathodic hydrogen evolution current occurring on ironsurface, making the corrosion potential move strongly towards negative direction.The phenomenon seems to be more remarkable in the low pH solutions. Forexample, the corrosion potential of iron in pH 0.75 H2S-containing solution (0.2mol dmÿ3 Na2SO4/H2SO4 and 0.4 mmol dmÿ3 H2S) is close to ÿ620 mV (vs.SCE), which is 80 mV more negative than that measured in the H2S-free solutionwith the same pH and SO 2ÿ

4 concentration (0.2 mol dmÿ3 Na2SO4/H2SO4) [1].What is more, compared with the polarization behaviour in the solutions withoutH2S, the potentiostatic steady-state polarization curves for iron in the solutionswith H2S exhibit two distinct characteristics: (i) although the anodic dissolutioncurrent is much greater at a given potential in the presence of H2S than in itsabsence, it increases more slowly with the positive potential, and (ii) consideringthat the corrosion current of iron in strong acidic solutions with H2S (pH R 2) isvery high even at potentials near corrosion potential, the pure anodic polarizationcurve can not be obtained in the lower positive potential region, which isparticularly di�erent from that in the H2S-free acidic solutions [1].

3.1.2. Impedance characteristics of iron in H2S-containing acidic solutionsOne of the AC impedance method's advantage is that many parameters

involved in corrosion processes occurring on the electrode surface, such as thecharge-transfer resistance, the polarization resistance, the double-layer capacitanceand pseudo-capacitance caused by relaxation processes of adsorbed intermediates,can be obtained in the same measurement. In this sense, the electrochemicalbehaviour of iron in the H2S-containing solution at corrosion potential can becharacterized better by AC impedance technology.

Impedance tests were conducted in a group of H2S-containing solutions of

H. Ma et al. / Corrosion Science 42 (2000) 1669±1683 1671

di�erent pH at the respective corrosion potential. The corrosion potentials of ironin these solutions were described in the caption of Fig. 1. It was found that thecorrosion potential of iron in H2S-containing acidic solutions decreased with pHincrease under otherwise identical conditions.

Fig. 1 shows a set of Nyquist impedance diagrams measured in the solutionsmentioned above. The shape and size of these impedance spectra strongly depend

Fig. 1. (a) Nyquist impedance diagrams for iron in the four H2S solutions of di�erent pH values at the

respective corrosion potential. (b) Magni®cation of the impedance plot of iron in the solution of pH

0.75. (w) 0.2 mol dmÿ3 Na2SO4/H2SO4 and 0.4 mmol dmÿ3 H2S, pH 0.75, Ecorr = ÿ620 mV (vs.

SCE); (r) 0.2 mol dmÿ3 Na2SO4/H2SO4 and 0.4 mmol dmÿ3 H2S, pH 2, Ecorr = ÿ680 mV (vs. SCE);

(+) 0.5 mol dmÿ3 Na2SO4/H2SO4 and 0.4 mmol dmÿ3 H2S, pH 2, Ecorr = ÿ672 mV (vs. SCE); (.) 0.5mol dmÿ3 Na2SO4/H2SO4 and 0.4 mmol dmÿ3 H2S, pH 3.5, Ecorr = ÿ700 mV (vs. SCE).


on the pH value of solutions. As can be found in Fig. 1, the impedance diagramsmeasured in the solutions of low pH, for example, in electrolytes of pH 0.75 and2, displayed two well-separated capacitive loops, while those diagrams measuredin solutions of high pH (3 and 3.5) seemed to be made up of two loops poorlyseparated in the frequency domain. It was also observed from Fig. 1 that, theimpedance diagrams measured in pH 0.75 and 2 solutions respectively displayedthe high-frequency capacitive arcs with the smaller diameter and higher topfrequency in comparison with those measured in pH 3 and 3.5 solutions,indicating low pH value is favourable for the corrosion of iron, that is to say, thelower the pH is, the more seriously iron was corroded.

3.1.3. Analysis of impedance dataAll impedance plots may be analyzed with the equivalent circuit in Fig. 2. In

order to give more accurate ®t results, the constant phase elements (CPE) weresubstituted for capacitors. The so-called CPE is an element whose admittance orimpedance value is a function of the frequency and whose phase is independent ofthe frequency. Its admittance and impedance are, respectively, de®ned as:

YQ � Y0�jo�n �1�and

ZQ � 1

Y0�jo�ÿn �2�

where, Q represents a CPE, Y0 is the modulus, o the angular frequency and n thephase [13]. The reason why CPEs are used as substitutes for capacitors in analysisof impedance spectra is that most impedance loops measured in experiments arenot ideal semi-circles but the depressed ones [14,15]. The depression degreedepends on the phase of the CPE [13].

The circuit element values required for ®tting the impedance behaviors are givenin Table 1. Moreover, a CPE can be treated as a parallel combination of a purecapacitor and a resistor being inversely proportional to the angular frequency,according to our recent theoretical research about analysis of impedance data [16].For a parallel circuit composed of a CPE �Qi� and a resistor �Ri), �QiRi), therelaxation time constant of this circuit, ti, is

Fig. 2. The equivalent circuit to ®t the Nyquist impedance diagrams with two capacitive loops.


ti � CiRi �3�where, Ci is capacitance of the pure capacitor comprising the CPE, Qi: Seeing thatthere exists a relation among Y0i, ni, ti and Ri:

Y0i � tnii =Ri �4�where, Y0i and ni are, respectively, the modulus and the phase of Qi: Associationof Eqs. (3) and (4) leads to

Ci ��Y0iR

1ÿnii

�1=ni �5�

On the basis of Eq. (5), the double-layer capacitance (Cdl) and pseudo-capacitance(Ca) involved in two CPEs, CPEdl and CPEa, of equivalent circuit shown in Fig. 2can be calculated (see Table 2).

Usually the double layer charging±discharging is a very rapid process. Thus it isnormal to associate the experimental ®rst capacitive loop seen in high frequencyrange with the double layer relaxation. Of course, this hypothesis is not only onepossibility and some rapid Faradaic pseudo-capacitive processes can also becoupled with the double layer relaxation [17]. Some authors usually judge whetherthe coupling process occurs only based on a depressed form of the higherfrequency capacitive loop. This viewpoint is not correct, according to Barcia and

Table 1

Element values of equivalent circuit to ®t the impedance diagrams in Fig. 1

pH Rt (O cm2) CPEdl Ra (O cm2) CPEa

Y0 (Oÿ1 cmÿ2 sn) n (0±1) Y0 (O

ÿ1 cmÿ2 sn) n (0±1)

0.75 13.2 3.76� 10ÿ5 0.95 5.8 7.98� 10ÿ3 0.69

2 45.9 4.0� 10ÿ5 0.94 91.7 1.79� 10ÿ3 0.64

3 62.1 7.37� 10ÿ5 0.94 1.65� 102 9.52� 10ÿ4 0.69

3.5 72.4 2.18� 10ÿ4 0.90 2.74� 102 6.29� 10ÿ4 0.79

Table 2

Values of the double-layer capacitance and pseudo-capacitance for iron in the H2S-containing acidic

solutions of di�erent pH value

pH Cdl (F cmÿ2) Ca (F cmÿ2)

0.75 2.52� 10ÿ5 2.01� 10ÿ3

2 2.68� 10ÿ5 6.48� 10ÿ4

3 5.23� 10ÿ5 4.14� 10ÿ4

3.5 1.37� 10ÿ4 3.94� 10ÿ4


Matto [17], since almost all loops obtained experimentally with electrochemicalpolycrystalline solid interface are scarcely regular semi-circles but depressed onesto some extent [17,18]. We know, the appearance of depressed capacitive loops isalso related to the surface roughness of solid electrode [19], the inherent physicaland chemical heterogeneous nature of the solid surface [20], the nonuniformdistribution of current density on the surface [21] and so on, in addition to thecoupling e�ect. We do not want to make too much comment on origin ofdepressed impedance loop here. On the contrary, we are more interested incapacitance value of the electrical double layer. In general, the double-layercapacitance for pure iron electrode in most cases should be less than 100 mF cmÿ2

[17]. Therefore, the double-layer capacitance listed in Table 2, determined by®tting the impedance spectra in Fig. 1, is reasonable.

The time constant caused by relaxation process of an adsorbed intermediatespecies on electrode surface is much greater than that of the electric double layer,and the low frequency impedance loops are therefore attributed to the relaxationof the adsorbed species by most authors. Comparison of Cdl and Ca in Table 2shows that Ca values, measured in the three solutions except the one with pH 3.5are higher 1±2 order of magnitude than the corresponding Cdl value. It is wellknown that H2S molecules easily adsorb on iron surface. Based on the dissolutionmechanism of iron in acidic solutions with H2S as proposed by us [2,3], we thinkthat the low frequency capacitive loop observed in the H2S-containing solutions atcorrosion potential originates from the relaxation of the adsorbed sul®de species,�FeSH�ÿads:

3.2. Inhibiting e�ect of hydrogen sul®de

3.2.1. Impedance measurements at corrosion potentialWe have reported that H2S can also exhibit a strong inhibition e�ect on iron

corrosion under certain special conditions in Ref. [1]. This kind of role of H2S willbe investigated systematically in the following sections of this paper.

Three representative experiments were carried out. One experiment was tomeasure the impedance diagram for iron in the pH 3 H2S-free acidic solution (0.5mol dmÿ3 H2SO4/Na2SO4). Before experiments, the iron electrode is at ®rstimmersed in the solution for 5 min. The experiment was used as the referencepoint in order to better understand the role of H2S on the corrosion of iron underdi�erent conditions. The other two experiments were performed in 0.5 mol dmÿ3

H2SO4/Na2SO4 and 0.02 mmol dmÿ3 H2S solutions (pH = 3), but the ironelectrode were immersed 5 min and 3 h respectively, before the onset ofexperiments. The purpose to conduct the latter two experiments was to study howimmersion time a�ects what H2S have done with iron corrosion. All impedancemeasurements were performed at corrosion potential.

The measured open-circuit potential or corrosion potential was about ÿ615 mVwhen the iron electrode was just immersed in the H2S-free solution. Then thecorrosion potential gradually moved towards negative direction. This potentialapproached a relatively stable value (ÿ653 mV) about 5 min later. Di�erent from


variation of the corrosion potential with the immersion time in the H2S-freesolution, the change rate of corrosion potential with time in the acidic solutionwith H2S looked more rapid, decreasing from original value (ÿ632 mV) by ÿ690mV within 5 min, after the iron electrode was immersed in the solution. It is ofinterest to note that the corrosion potential no longer decreased about 2 h later;on the contrary, it began to increase, reaching ÿ620 mV in 3 h, which was evenmore positive than ÿ653 mV in H2S-free solution. This unusual phenomenon isprobably related to formation of sul®de protective ®lm on the iron surface.

Fig. 3 gives the impedance spectra for iron under di�erent experimentalconditions. The Nyquist diagram of iron in H2S-free acidic solution (described by``t'' in Fig. 3) was a depressed capacitive loop with 407 O cm2 diameter and 2.5Hz top frequency. The double-layer capacitance calculated from the ®tting resultsin Table 3 is 1.03 � 10ÿ4 F cmÿ2 (see Table 4), slightly higher than the normalvalue. In the presence of H2S, the impedance behaviour of iron was quitedi�erent. Moreover, the immersion time of iron strongly a�ected its impedancedisplay. When the immersion time was short, the complex plane plot (shown by ``� '' in Fig. 3) presented two smaller capacitive loops, as would be expected [1,3].It was determined from the high frequency loop that the charge-transfer is about64.4 O cm2 and double-layer capacitance 7.76 � 10ÿ5 F cmÿ2. At the same time,the corrosion potential also shifted from ÿ653 mV in the absence of H2S to ÿ690mV in the presence of H2S. These two changes show the remarkable accelerationof hydrogen sul®de to the iron corrosion. On the contrary, when the electrode wasimmersed for 3 h, it is found that, the corrosion potential became about ÿ620

Fig. 3. The impedance spectra for the iron electrode in the three solutions at corrosion potential. (t)

in the pH 3 solution of 0.5 mol dmÿ3 H2SO4/Na2SO4 (Ecorr =ÿ653 mV); (�) after 5 min immersion in

the pH 3 solution of 0.5 mol dmÿ3 H2SO4/Na2SO4 and 0.02 mmol dmÿ3 H2S (Ecorr = ÿ690 mV); (q):

after 3 h immersion in the pH 3 solution of 0.5 mol dmÿ3 H2SO4/Na2SO4 and 0.02 mmol dmÿ3 H2S

(Ecorr = ÿ620 mV).


mV; on the other hand, the original characteristic capacitive loop in low frequencydisappeared and the impedance plot (hollow square in Fig. 3) only gave acapacitive loop. Analysis of the impedance diagram shows that the charge-transferresistance increased by 584 O cm2, which is even greater than that of iron in H2S-free solution, and the double capacitance (2.75� 10ÿ5 F cmÿ2) is also much lowerthan the value (1.03 � 10ÿ4 F cmÿ2) in the absence of H2S. This indicates thatwhen the iron experienced a long time immersion in solutions containing a lowerH2S concentration, what H2S has done to the iron corrosion is inhibition e�ectrather than acceleration e�ect.

3.2.2. Characterization of sul®de protective ®lmThe potentiostatic steady-state polarization measurements also con®rm the

inhibition e�ect of H2S on corrosion of iron, provided iron is immersed for 3 h inthe solutions with the lower H2S concentration. As can be found in Fig. 4, thesteady-state polarization curves are substantially a�ected by the immersion timebefore experiments. Comparing the anodic polarization curves (1 and 2) and thecathodic polarization curves (1 ' and 2 '), one will observe that, after a longerimmersion time, anodic iron dissolution current is signi®cantly reduced. However,the behaviour of cathodic current is not remarkable. It is lowered because thepotential is decreased as a result of the changes in the anodic polarizationbehaviour. Additionally, comparison of the anodic polarization curves 1 and 2show that, the Tafel slope of curve 2 was greatly smaller than that of curve 1, but

Table 3

Element values of equivalent circuit to ®t the impedance diagrams in Fig. 3

Experimental condition Rt (Ocm2)

CPEdl Ra (Ocm2)

CPEa

Y0 (Oÿ1 cmÿ2

sn)

n (0±

1)

Y0 (Oÿ1 cmÿ2

sn)

n (0±

1)

In H2S-free solution 4.07�102

1.55� 10ÿ4 0.87 ± ±

Immersed 3 h in the solution with

H2S

5.84�102

3.94� 10ÿ5 0.91 ± ±

Immersed 5 min in the solution

with H2S

64.4 1.39� 10ÿ4 0.89 2.03�102

1.63� 10ÿ3 0.63

Table 4

Values of the double-layer capacitance and pseudo-capacitance determined from impedance spectra in

Fig. 3

Experimental condition Cdl (F cmÿ2) Ca (F cmÿ2)

In H2S-free solution 1.03� 10ÿ4 ±

Immersed 3 h in the solution with H2S 2.71� 10ÿ5 ±

Immersed 5 min in the solution with H2S 7.76� 10ÿ5 8.51� 10ÿ4


the current indicated by curve 2 increased more rapidly with anodic potential andit could have reached the current value of curve 1 at the higher anodic potential.

The occurrence of inhibition e�ect is related to formation of a protective layerwith ferrous sul®de grain at the electrode surface. Some other author's studieshave shown that a layer of FeS ®lm will form on the surface of iron after the ironexperiences a long time immersion [12,13]. By comparing the di�erence betweencurves 1 and 2 in Fig. 4, we think at the lower anodic potential, the protective®lm can inhibit the anodic current increase to some extent; while at the higheranodic potential, this ®lm will be destroyed and it can not continue to prevent theinside of the iron from anodic dissolution. In this case, the anodic current willsurely approach to the value of curve 1.

In order to verify the assumption, we have done another interesting experiment.The whole process is as follows: (i) the iron electrode was at ®rst immersed in thesolution (0.5 mol dmÿ3 H2SO4/Na2SO4 and 0.02 mmol dmÿ3 H2S, pH = 3) for 3h, (ii) the electrode was polarized at a higher anodic potential (see A point inFig. 4) for 1 min, and (iii) the anodic steady-state polarization curve (see curve 3in Fig. 4) was obtained by adjusting electrode potential gradually from highpotential to low potential. The results show that this curve is very similar to curve1, indicating that once FeS protective ®lm was destroyed, H2S molecules willadsorb on the electrode surface again and accelerate the iron dissolution.

Fig. 4. The polarization curves of iron in 0.5 mol dmÿ3 Na2SO4/H2SO4 + 0.02 mmol dmÿ3 H2S (pH

= 3): 1. immersed for 5 min, 2. immersed for 3 h, 3. after immersing for 3 h, the electrode was

polarized at point A for 1 min, then the potential decreased steeply.


In a sense, the FeS protective ®lm is similar to those formed by inhibitor'sadsorption at electrode surface. This protective ®lm will gradually be damagedwith the increasing potential. AC impedance measurement results are in a goodagreement with the anodic polarization in Fig. 4 and have further con®rmed ourconclusion. Fig. 5(a) is the Nyquist impedance diagrams for iron measured atcertain anodic potential (ÿ620 mV) after the electrode was immersed for 5 min in0.5 mol dmÿ3 Na2SO4 + 0.02 mmol dmÿ3 H2S solution, which displays 3 timeconstants. When the immersion time increased, as observed in Fig. 5(b), theimpedance diagram measured at this potential will only give a capacitive loopwith a larger diameter. Obviously, the inhibiting e�ect of H2S on the irondissolution has occurred because of the formation of sul®de protective ®lm.Provided the iron electrode with the FeS ®lm is polarized at ®rst at ÿ500 mV andthe impedance data measured at the same potential will give a plot almost thesame as Fig. 5(a) (see Fig. 5(c)). The reason is that the FeS ®lm on the electrodesurface is destroyed and loses the ability to protect the inside iron.

3.3. Factors in¯uencing the inhibiting e�ect

3.3.1. Immersion time and pHThe protective ®lm introduced here is quite di�erent from the common

passivation ®lms in structure or composition. The latter belongs to a kind ofdissipation system, which has to be maintained by the applied electric ®eld,whereas the protective ®lm is composed of di�erent ferrous sul®de crystal grain,such as pyrite, troilite, kansite, pyrrhotite and mackinawite [12,16], whoseprotective property depends on the H2S concentration, pH of solution andimmersion time of electrode. To date, there is considerable controversy about thecomponents of sul®de ®lm, but most authors are inclined to believe that ametastable species, mackinawite, forms at electrode surface at ®rst [12], and thespecies will further transform into the more stable species, troilite and pyrite[12,22]. Based on our recent researches [1±3], a probable mechanism of iron in theacidic solutions containing H2S is described as follows:

Fe� H2S� H2O,K1

Kÿ1FeSHÿads � H3O� �6�

FeSHÿads,K2

Kÿ2Fe�SH�ads � eÿ �7�

Fe�SH�adsÿÿÿ4K3FeSH� � eÿ �8�

According to Shoesmith [12], the species FeSH+ at electrode surface may beincorporated directly into a growing layer of mackinate via Eq. (9)


Fig. 5. The Nyquist impedance diagrams for iron measured at ÿ620 mV in the pH 3 solution of 0.5

mol dmÿ3 H2SO4/Na2SO4 + 0.02 mmol dmÿ3 H2S: (a) immersed for 5 min at the open-circuit

corrosion potential before experiment; (b) immersed for 3 h at the open-circuit corrosion potential

before experiment; (c) immersed 3 h at the open-circuit corrosion potential before experiment and then

was polarized at ÿ500 mV.


FeSH�ÿÿÿ4K1FeS1ÿx � xSHÿ � �1ÿ x�H� �9�

or it may be hydrolyzed to yield Fe2+ via Eq. (10)

FeSH� � H3O�,K5

Kÿ5Fe2� � H2S� H2O �10�

If reaction (9) leads to local supersaturation at the electrode surface, thennucleation and growth of one or more of the iron sul®des, mackinate, cubicferrous sul®de and troilite occurs. H2S will exhibit the di�erent role on anodicprocess of iron depending on the pH value. At the lower pH values (<2) irondissolves via reaction (10) and little iron sul®de forms because of the relativelygreater solubility of the iron sul®de phases. In this case, H2S only exhibits theacceleration e�ect on the dissolution of iron. However, at the pH values of 3±5,H2S begins to exhibit its inhibiting e�ect as FeSH+ species may form partiallymackinawite via reaction (9). Furthermore, the mackinawite can convert intotroilite with the greater stability and more protective property. At the much higherpH (r5), mackinate was the only corrosion product observed. As the protectiveability of mackinawite is worse that that of troilite, the inhibiting e�ect of H2Sdecreases.

As mentioned above, the FeS protective ®lm involves a growing layer ofmackinawite through Eq. (9) at ®rst. Consequently, even if at lower H2Sconcentration and pH range within 3±5, what H2S exhibits at the beginning is theacceleration e�ect corrosion rather than the inhibiting e�ect on iron. Only whenthe electrode is immersed for a long period of time, its inhibiting e�ect can bedisplayed due to the formation of pyrite and troilite.

3.3.2. H2S concentrationH2S concentration has great in¯uence on the protective ability of the sul®de

®lm. As H2S concentration increases, the amount of kansite in the ®lm alsoincreases [12]. As kansite has some defects in its solid structure, it can note�ectively prevent the iron from corrosion like pyrite and troilite. In addition,when the H2S concentration is higher, appreciable amount of mackinawitedeposits on electrode surface at 3±5 pH, under open-circuit condition, forms aloose sul®de ®lm, which is not contributory to the inhibiting e�ect. SEMmeasurement showed that the electrode surface was more greatly corroded in thesolution of high H2S concentration than in that of the low concentration andscratch marks partially disappeared (see Fig. 6(a) and (b)).

4. Summary

H2S displays the acceleration role on both the anodic iron dissolution and thecathodic hydrogen evolution in most of the cases. Only under certain specialconditions, such as the lower H2S concentration (R0.04 mmol dmÿ3), pH value of


3±5 and the longer immersion time (r2 h), can H2S exhibit the inhibiting e�ect onthe iron corrosion. This role is related to the formation of FeS protective ®lm atelectrode surface. Among FeS with di�erent crystal shapes, a metastablemackinawite form at ®rst via Eq. (6), then converts into troilite and pyrite withthe greater stability and the more protective property. Therefore, this inhibitinge�ect can occur only in the case of the lower H2S concentration, pH 3±5 and witha long immersion time.

Acknowledgements

This project is supported by the Chinese Natural Science Fund and the StateKey Laboratory for Corrosion and Protection, Academia Sinica.

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the influence of hydrogen sulfide on corrosion of iron under different conditions

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