CORROSION OF CARBON STEEL BY H2S IN CO2 CONTAINING OILFIELD ENVIRONMENTS

Stephen N. Smith ExxonMobil Production Company

800 Bell, CORP-EMB-2003H Houston, Texas 77381

Michael W. Joosten

ConocoPhillips Bartlesville Technology Center Bartlesville, Oklahoma 74004

ABSTRACT

The effect that even small concentrations of H2S can have upon CO2 corrosion has been recognized since at least the 1940's. Early studies showed that the FeS corrosion products that were formed had an impact, but disagreed whether the impact was beneficial or not. Although H2S corrosion has not received the level of attention given to CO2 corrosion, the literature has shown that there are a number of different forms of FeS that can form as corrosion products, depending upon the exposure conditions. Between the corrosion, geochemical and thermodynamics literature, a great deal is known about the corrosion chemistry involved with the formation of the various FeS species as well as the impact that each has upon further corrosion.

However, there is still a great deal that is not known. For example, there are currently no generally accepted prediction algorithms for any form of H2S corrosion. There are also still a number of unknowns about the corrosion reactions that lead to pitting, which is the most common mode of sour service equipment failure. This paper reviews a sampling of the H2S corrosion literature over the past 60 years and describes some of the areas of research that remain.

INTRODUCTION

A great deal has been written over the years about the effects of CO2 on corrosion and the various changes in environmental parameters that impact the corrosion rate due to CO2. Factors such as

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temperature, pH, bicarbonate, velocity and a host of others had been studied extensively. By comparison, only a small amount of effort has gone into the study of the impact of H2S upon carbon steel corrosion rates. This is most likely a result of a concentrated focus on H2S cracking of various metallic materials. This does not mean that there is no information available about the effects of H2S upon corrosion mechanisms and reaction rates. Research and field work in this area dates back to the 1940's. Research relevant to iron sulfide formation chemistry has also been conducted by geochemists, microbiologists, thermodynamicists and geophysicists. Much of this work has direct relevance to the study of the mechanism of CO2/H2S corrosion.

Although there is over 60 years of H2S related corrosion research work, much of this literature is somewhat confusing and often seemingly contradictory. Iron sulfide chemistry is very complex and seemingly minor changes in test conditions can often lead to dramatically different results. The mineralogists and thermodynamicists still cannot completely agree upon the number of types of iron sulfide that actually exist, even though they have been studying the materials for far longer than corrosion engineers.

This paper reviews some of the history relative to corrosive oilfield environments that contain CO2 and H2S. We will attempt to examine the effects of H2S on the variety of different environments where the various iron sulfide corrosion products are formed. Finally, we will discuss the areas of further research needed to more fully understand the impact of H2S on CO2 corrosion in oilfield environments. This understanding will subsequently improve the accuracy of corrosion rate predictions.

HISTORY OF OILFIELD CORROSION BY H2S

Oilfield corrosion engineers have recognized since the 1940's that the presence of H2S changes the corrosivity of produced fluids as compared to sweet production with only CO2. Many investigators working with H2S have concentrated on the cracking of carbon steels. Others have studied various aspects of the role of H2S upon corrosion and the role of the various FeS corrosion products. Some have concluded that H2S reduces corrosion as compared to CO2 and others have concluded that H2S increases corrosion. To understand how both of these positions can be correct and how our current understanding of the complexities of H2S corrosion has evolved, we must start from the work conducted in the 1950's and move forward in time. 1950's

In 1951, H.R. Copson1 published a paper in Corrosion that was a literature review of field

experience in unaerated oil well brines. In the paper, Copson concluded that "in fields which produce large quantities of hydrogen sulfide bearing brine, there is little or no corrosion irrespective of hydrogen sulfide concentration."

Several years later, Walter Rogers and J.A. Rowe, Jr.2 conducted laboratory electrochemical studies of corrosion in oil field brines with either CO2 or H2S. Their tests were run at a temperature of 100oF (38oC) and at atmospheric pressure. The measured pH varied from 4.5 to 9 depending upon the concentration of CO2 or H2S in the test and the composition of the field brine being evaluated. Tests were run for periods of as long as 480 days. They observed a cathodic effect due to the formation of FeS corrosion deposits in H2S corrosion. Based upon their observation of the very positive potential of an FeS cathode, they developed a theory of sulfide corrosion whose principle factors were:

1. The high potential of the anode due to low iron solubility

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2. The low potential of the sulfide cathode, which they said approximated a platinum cathode, and 3. The ability of the iron sulfide cathode to receive large current densities without polarizing.

Their theory went on to address the low early corrosion rates in H2S systems that increase with time and also result in local pitting. The theory basically involved increasing cathode surfaces with time due to the continued deposition of more FeS cathodes, which drives more anode areas to develop due to the potential difference between iron and FeS.

Rogers and Rowe also developed equations to predict corrosion current in H2S and CO2 corrosion that were functions of pH and activity of sulfide or carbonate. The H2S equation was:

( )( )[ ]222 log0295.01495.11 +−−−= aHaS

RI SH (1)

At the same time that Rogers and Rowe were doing their research, Scott Ewing3 was conducting

similar electrochemical studies. Ewing's contributions involved a discussion of the role of pH upon the solubility and/or precipitation tendency of FeS. Ewing also addresses the role of increased total pressure upon the partial pressure of H2S as an explanation for increasing corrosivity observed at higher total pressures. He also observed that the corrosivity in H2S drops sharply at pH levels greater than7, which was attributed to the protectiveness of the FeS deposit.

In 1956, NACE Task Group T-5B-2 published NACE Technical Committee Report 5B156, "Collection and Correlation of High Temperature Hydrogen Sulfide Corrosion Data4." This report compiled tables and graphs of laboratory and field data for 1 atm partial pressure of H2S at a wide range of temperatures and for a number of different materials. The alloys included ranged from various carbon steels through 13Cr and 316 stainless steels to the Cr-Ni alloys.

In 1958, Meyer, Riggs, McGlasson and Sudbury5 conducted extensive studies of the corrosion products that form on mild steel in H2S environments. In addition to the normal weight loss corrosion tests, they performed XRD analysis of corrosion products formed in H2S-H2O, H2S+NaCl+H2O and H2S+CO2+H2O for various lengths of time up to 127 days. They observed sequential formation of various FeS species starting with Kansite (also known as Mackinawite) moving on to Pyrrhotite and finally ending with Pyrite. They also introduce the concept of solid state diffusion of Fe2+ ions through the scale as the mechanism for Kansite scale growth. The also list the various forms of FeS as Kansite, Troilite, Pyrrhotite, Smythite, Marcasite and Pyrite. Finally, they concluded that "contaminants of H2S environments such as brine, carbon dioxide, free sulfur and hydrogen influence the forms taken by the corrosion products”.

Andrew Dravnieks and Carl Samans6 studied the kinetics of H2S reactions on steel as a function of temperature in the range of 250 to 500oC. Although this is well beyond the range of temperatures found in oil field corrosion, it is a useful reference because they found that sulfidation is a three step process comprised of H2S adsorption, formation of the diffusing species, and diffusion through the FeS scale.

In 1959, Donald Shannon and James Boggs7 investigated a variety of parameters such as type of liquid hydrocarbon, NaCl concentration, H2S concentration and time in order to devise a screening test procedure for evaluating oil field corrosion inhibitors. They concluded that "the initial rate of corrosion of steel coupons by oil-water-H2S mixtures is dependent upon the H2S concentration. After an exposure of one day, the rate becomes diffusion controlled and independent of H2S concentration over a rather wide range. At high concentrations, the corrosion rate is actually lowered because of the formation of a

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denser, more crystalline FeS coating through which diffusion is slower." It is also important to note that they recognized that in their low H2S tests, they did not have adequate H2S at end of test for the tests to be considered valid.

In 1959, Newman8 proposed conducting inhibitor performance tests for secondary recovery systems using a flow loop. The loop was used to determine the critical corrosion inhibitor concentration for a variety of conditions, including those with H2S. The H2S tests were run at 105oF (40oC) and found that corrosivity did not increase substantially in range from 50 to 400 ppm H2S. However, he concluded that at concentrations in excess of approximately 600ppm H2S, inhibitor would be required. 1960's

In 1963, Hughes and Stromberg9 conducted weight gain/loss corrosion tests to determine the

influence the addition of various chemicals would have upon the adherence of FeS and its ability to be an effective barrier to corrosion. They conducted their studies over a wide pH range, from 4.5 to 9.5. The concluded that "Hydrogen sulfide may not, under all circumstances, be an undesirable component from the stand point of corrosion; in fact, its presence may be advantageous in some instances. It is also concluded that inhibitors enter into and modify the sulfide scale to some extent."

Sardisco, Wright and Greco10 evaluated the effect of H2S partial pressure upon the type of FeS formed. Their tests included H2S partial pressures from 0.001 to 4 psi (0.006 to 27 kPa) with no NaCl. They found the same variety of FeS species that had been reported by Meyer, et al.5, five years earlier. In their weight loss corrosion rate tests, they found higher corrosion rates when significant levels of Kansite (Mackinawite) were present as compared to the other FeS species. This was concluded to be due to the formation of a non-protective, porous scale. They also evaluated the crystal size of the FeS as a function of H2S partial pressure and found a parabolic increase.

In 1965, Sardisco and Pitts11-12 presented two papers on the corrosion of iron in an H2S-CO2-H2O system. The first paper discussed the mechanism and kinetics of the sulfide film formation. The second paper addressed the protectiveness of the sulfide film as a function of pH. In the mechanism paper, weight loss/gain tests were run and were found to reach steady state corrosion rates between 5 and 15 hrs. Tests were run with H2S concentrations of 0.00958 to 3.25 psia (0.066 to 22 kPa) in water at 24oC. The FeS films formed consisted of Troilite, Pyrrhotite, Marcasite and Kansite. No attempt was made to determine if the staged scale formation observed by Meyer, et al., was duplicated since the scale surfaces were analyzed as a function of exposure time. They concluded that "during liquid phase corrosion of iron by H2S-CO2-H2O, over-all reaction is controlled partially by interface reaction and partially by passage (diffusion) of ions and electrons across film. At low H2S concentrations, reaction mechanism approaches complete diffusion control and at high H2S concentrations mechanism approaches complete interface control."

In their subsequent paper on protectiveness of FeS as a function of pH, Sardisco and Pitts found the least level of protection in the pH range of 6.5 to 8. Kansite (Mackinawite) was identified as the primary scale in this pH range. Troilite and Pyrite were found to predominate in more acidic range of 4.0 to 6.5.

In 1966, Charles Milton13, a geologist with the U.S. Geological Survey published a Technical Note in Corrosion titled "Kansite = Mackinawite, FeS" The publication explained that the geologists had observed a mineral for over a century that had been confused with a similar mineral, Vallerite, which is CuFeS2. The geologists did not clear up their confusion until 1963, well after Fred Prange had

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named the corrosion product that he had observed forming in wells in Kansas as Kansite. It is also perhaps ironic that a key step in clarifying the confusion within the geological circles was a Mackinawite synthesis test conducted in 1962 that consisted of exposing an iron wire to an aqueous solution of H2S in the absence of air.

In 1968, Edward Greco and William Wright14 reported on corrosion studies conducted on H2S-CO2-H2O systems under static and dynamic conditions. The static tests involved room autoclave temperature tests in solutions with 400 ppm NaCl and H2S partial pressures from 6x10-5 to 80 psia (0.0004 to 550 kPa). The dynamic tests were conducted at a low flow rates simply to maintain constant H2S levels at the lower concentration levels for exposure times up to 120 hrs. They found that the corrosion rate drops with increasing H2S at the lower concentration with a minimum corrosion rate reached at 0.001 psia (7 Pa). The corrosion rate then holds fairly constant up to an H2S partial pressure of about 1 psia (7 kPa). Further increases in H2S partial pressure resulted in increasing corrosion rate. The type of FeS scale formed in these tests was not reported. 1970's

In 1972, Hausler, Goeller, Zimmerman and Rosenwald15 reported on experiments run at 150oC with 100 psi (670 kPa) H2S. They varied the pH and time to study film formation and diffusion kinetics through the FeS layer. They found significant variations in film weight versus pH. The film weights and corrosion rates were both low values at pH 7, the film weight increased sharply but there was a lot of scatter when plotted versus corrosion rate at pH 6. At pH 4.2, the film weight produced a hyperbolic function of decreasing film weight with increasing corrosion rate. They concluded that this was evidence of film diffusion control and that changes to the corrosive media could change the film. The paper also discusses the concepts of the ion exchange properties of FeS film and presents a concept of a coherent FeS film that provides protectiveness as a means to explain the differences in FeS protection as a function of pH.

Hausler, Goeller and Rosenwald16 expanded upon the earlier work of Hausler, Goeller, Zimmerman and Rosenwald by studying the effects of corrosion inhibitors. The paper discusses the concept of adsorbed inhibitors and how they can become ineffective through growth of the FeS layer such that the inhibitor becomes buried in FeS.

Yamaguchi and Moori17 conducted electrochemical corrosion tests on steel to prepare Greigite films. This corrosion product was analyzed by electron diffraction to define the Greigite crystal type (Fd3m type spinel), lattice constant (9.875A) and crystal structure with a composition Fe3S4.

Zitter18 reported on the results of his examination of production tubing from a well with production that contains 15% CO2 and 1.7% H2S at 4 MPa and temperatures that range from 30 to 150oC. Analysis of the multilayer corrosion product scale found Troilite and Magnetite near the metal surface and Marcasite, Pyrite and Troilite closer to the flow stream. Siderite (FeCO3) was also observed, but only in areas of heavy CaCl2 accumulation. An explanation for the presence of Siderite was offered that is based upon the formation of HCl through a reaction involving the precipitation of CaCO3 from the mixture of CO2 and CaCl2.

In 1975, Smith and Miller19 published a review paper that describes the nature of iron sulfides and their corrosive effect. The paper provides thermodynamic and crystallographic data for a variety of FeS phases, including Mackinawite, Cubic FeS, Pyrrhotite, Greigite, Smythite, Marcasite and Pyrite, with a discussion of each phase. The paper also reproduced a formation/progression interrelationship

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diagram from a geochemistry reference that relates the various reaction paths between the various forms of FeS (and Siderite). This figure appears below as Figure 1.

Tewari, Wallace and Campbell20 is the first of a series of corrosion reports and papers discussed herein that was produced by researchers at Whiteshell Nuclear Research in Canada. This document is an extensive report on the formation chemistry and kinetics of formation for Mackinawite, Pyrrhotite and Pyrite in Girdler-Sulfide heavy water extraction plants. The experimental work that was reported involved the determination of dissolution kinetics for Mackinawite, Troilite, monoclinic and hexagonal Pyrrhotite, and Pyrite as a function of temperature (25 to 125oC) and pH (1.0 to 4.0). Dissolution rate constants were determined for each form of FeS at each of the temperature studied. A relationship between the solubility of the various FeS morphologies and pH at 25oC and an H2S partial pressure of 0.1 MPa can be seen in Figure 2.

In 1978, William Thomason21 reported on laboratory tests that exposed steel weight loss coupons to 1 atm H2S for 2 to 7 days at temperatures from 30 to 90oC. These tests were conducted to see information about "film formation at higher temperatures" based upon speculation that the reduction in sulfide stress corrosion cracking (SSC) at temperatures greater than 65oC was due to the formation of a passive film. The tests consisted of coupon exposures as well as hydrogen permeation rate measurements and the analysis of the films formed. The two to seven day tests only produced Mackinawite as corrosion product, but a change in the appearance of the film at higher temperatures was observed. The paper concluded that "a protective film alone would not be a reliable means of protecting high strength steels in a SSC environment."

Bruce Craig22 evaluated the change in Mackinawite corrosion products that occurs upon subsequent exposure to atmospheric oxygen. He found that Mackinawite can transform to Lepidocrocite [γ-FeO(OH)]) and elemental sulfur. Upon further exposure, the Lepidocrocite transformed to Magnetite [Fe3O4]. Craig also measured the oxidation rate and found that it varied as a function of the degree of crystallization of the corrosion product and increasing oxygen concentrations. The timeframe for the oxidation reactions studied was on the order of days, not seconds or minutes.

Fryt, Smeltzer and Kirkaldy23 present a high temperature sulfidation study that was performed at temperatures of 600 to 1000oC. Normally this study would not have relevance in a discussion of aqueous corrosion. However, Fryt et. al., measured the chemical diffusivity and concentration of Fe ion vacancies in Pyrrhotite as a function of sulfur activity and temperature where the sulfur fugacity ranged from 10-11 to 10-3 atm. The diffusivity values may not be applicable at lower temperatures, especially considering the phase transformations that occur in Pyrrhotite at 308oC, but perhaps the data could be useful as a starting point for future research.

Murata, Matsuhashi, Taniguchi and Yamamoto24 evaluated the corrosion rate as a function of H2S and CO2 partial pressure and temperature. Their data is presented as 3D plots of corrosion rate versus gas composition and temperature. An example is reproduced as Figure 3. The plots clearly show that under some conditions, the corrosion rate in high H2S gas mixtures is lower than for high CO2 and that the relationship is reversed for other test conditions. These plots provide a clear illustration of the basis for confusion in the early literature that cited FeS films as both increasing and decreasing corrosivity when H2S was added to CO2.

Milliams and Kroese25 studied the effect of NaCl concentration upon corrosion in H2S and CO2 as a function of temperature. They concluded that, "at 25oC high salt concentrations prevent the formation of protective surface layers that could stifle further attack" and "at higher temperatures (80oC)

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better protective layers are formed, even in the case of high salt concentrations." Milliams and Kroese also attempted to study pitting due to chlorides and H2S. They found that the pits that they generated electrochemically lost their activity in a few days as a result of buildup of protective corrosion product layers. Since pitting is a practical concern in the field, they concluded that more research in this area was required.

Hamby26 reported field corrosion problems for sour gas production wells producing 28 to 46% H2S and 3 to 8% CO2 for wells with bottom hole pressures (BHP) of 17500 to over 22000 psi (120 - 152 MPa) and bottom hole temperatures (BHT) of 365 to 385oF (185 to 195oC). The wells experienced downhole corrosion failures in less than six months. The cause was reported to be due to vaporization of the corrosion inhibitor solvent into the dry gas that left the carbon steel tubing string with no inhibition.

Tewari and Campbell27 reports on a laboratory rotating disk study performed to determine Mackinawite solubility and dissolution rates as a function of velocity at standard temperature and pressure conditions. The Mackinawite dissolution rate was found to be 20 times faster than Troilite and 1000 times faster than Pyrrhotite. The research also measured the diffusion coefficient of the FeSH+ ion at 22oC.

Tewari, Bailey and Campbell28 extended the Mackinawite dissolution rate work of Tewari and Campbell to conditions of 120oC and 1.6 MPa pressure. Flow effects upon dissolution using a rotating disc that was pre-scaled with Pyrrhotite and Pyrite. The study found that at high velocities, Mackinawite remained as the predominant surface film, while stagnant and low velocities allowed the Mackinawite to transform to the higher FeS phases were iron ion transport from the surface is not as rapid.

Hausler29 expanded upon the discussion by Hausler, Goeller and Rosenwald of the interaction between corrosion inhibitors with FeS films. The paper discusses the role of the diffusion of iron vacancies in FeS scale upon corrosion and the interaction between the inhibitor, flow and the crystal defects on the scale surface. 1980's

Shoesmith, Taylor, Bailey and Owen30 studied the formation of FeS at 21oC and 1 atm H2S at pH values from 2 to 7 and for times ranging from minutes up to 96 hrs. They found that Mackinawite is formed by both solid state and precipitation processes. Cubic FeS and Troilite were found to occur as precipitates between pH 3 and pH 5 by growth that resulted from cracking of the initial Mackinawite layer. As the tests proceeded, Cubic FeS was later transformed via solid state to Mackinawite over a period of about 80 hours. The interrelationship between the scales is shown in Figure 4. They found that the scale formation rates are controlled by pH, the applied electrochemical current and the degree of convection. The corrosion rate increases with decreasing pH. The quantity of produced FeS scale peaks at pH 4 below which scale dissolution rates become the predominant process. Passivation was only observed at pH 7.

Morris, Sampaleanu and Veysey31 conducted polarization studies on rotating disks and in a flow loop at 25oC and 1 atm H2S with pH values from 3 to 5. They concluded that H2S does not change the Tafel slopes of the anodic and cathodic processes within this pH range. The cathodic process remained reversible, but H+ ion diffusion control gradually disappears with increasing H2S concentration. The Mackinawite corrosion product observed in their testing was non-adherent, particularly at flow velocities in excess of 2 m/sec.

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Foroulis32 presented a theoretical review of H2S corrosion of iron that looks at the chemistry of

both the dissolution/precipitation and direct film formation sulfidic corrosion reactions, the potential-pH diagram for the Fe-H2O-H2S system at 25oC and the physical nature of the corrosion products scales formed in testing performed at 25oC.

Wikjord, Rummery, Doern and Owen33 studied the complex, multiphase FeS films that formed on rotating disks exposed to H2S solutions. Their test conditions included exposures for up to 840 hrs to 1.5 MPa H2S at temperatures of 308, 373 and 433oK. They found an evolution sequence of corrosion products with progression from Mackinawite to Cubic FeS to Troilite to Pyrrhotite and finally to Pyrite. All phases except for Mackinawite were found as crystals of regular geometry, indicating relatively slow growth. Higher temperatures were found to accelerate the transformation while increasing rotation speeds slowed the transformation. Additional oxidants were also found to speed the formation of Pyrite.

Martin and Annand34 conducted an electrochemical study of the impact of small O2 contamination upon the corrosion of steel by suspended iron sulfides. Tests were run in 3.5% NaCl solutions at 23oC and sparged with H2S at 1 atm. They concluded that the increase in corrosion rate due to the presence of suspended FeS was due to an increase in the cathodic reaction through hydrogen adsorption by the suspended FeS particles. They also concluded that inhibition of the oxidized suspended sulfides required different compounds than those that were shown effective for simple H2S corrosion.

Shoesmith35 presented an extensive review of film formation, transformation and dissolution processes on surfaces as the Lash Miller Award address to the Electrochemical Society in 1981. The presentation included special attention to sulfide films involved in the Girdler-Sulfide process that had earlier been reported by others. The paper does not present new information, but provides an excellent overall review of the H2S corrosion process in the Girdler-Sulfide system.

Lichti, Soylemezoglu and Cunliffe36 reviews experience with geothermal wells in New Zealand that produce brines containing CO2 and H2S. They found the formation of corrosion products of Mackinawite, Troilite and Pyrrhotite. They concluded that sulfide films reduced corrosion rates, even when present in minute amounts. Pitting was also observed with all FeS scales, even when the FeS scales were adherent with pitting depths often greater than 10 times the average wall loss due to corrosion.

Pound, Wright and Sharp37 performed cyclic voltammetry tests for New Zealand geothermal wells with test conditions very similar to Lichti, et al. They reported that the cathodic process occurs by the reduction of H2S to H2 through a process that involves adsorbed hydrogen intermediates.

Rogne38 reports on a study that evaluates the influence of temperature (25, 60 and 80oC) and H2S concentration (0.03 to 4%) on corrosion of steel in 5% NaCl solution. The work also compares corrosion in tests performed with either H2S/CO2 or H2S/N2. The study found that at 60 and 80oC and high H2S, the corrosion rate was reduced to 1/10 of the initial values due to the formation of protective FeS films. At 25oC, a reduction in corrosion rate was observed with a less protective film, but the reduction in corrosion rate was not as pronounced. They also found that H2S/CO2 mixtures are more corrosive than H2S/N2, but that the addition of even small amounts of H2S had "inhibiting" effect on CO2 corrosion.

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Narayan, Anderegg and Chen39 Conducted an ESCA study of the surface of iron exposed to H2S at partial pressures of 10-6 to 10 torr (10-4 to 103 Pa) at temperatures of 100 to 773oK. They found that the mode of sulfidation of the iron surface occurred by dissociative adsportion of H2S at ambient temperatures and the direct formation of FeS at temperatures in excess of 423oK. Sulfidation was also found to occur at higher rates on oxidized iron than on sputter cleaned surfaces.

Dunlop, Hassell and Rhodes40 is a paper that primarily discusses CO2 corrosion. However, the paper proposes the use of a CO2/H2S ratio of 500 at 25oC to determine whether the corrosion product will be FeCO3 or FeS. For values greater than 500, the product will be FeCO3 and less than 500 the product will be FeS.

Craig41 offered an explanation why steel alloys with small Cu additions can either enhance or reduce the rate of corrosion upon exposure in H2S environments. Previous work found that small FeS grain sizes with Cu, at pH values greater than 5.0, resulted in reduced corrosion and hydrogen permeation rates. This was explained by changes in the semiconductive properties of the FeS that makes it more insulating. For pH in the range of 4.6 to 4.8, the corrosion rate was increased. The proposed explanation for this is that the Cu+1 ions increase the n-type semiconductive properties of Mackinawite. For pH values less than 4.0, Cu additions to the steel had no impact due to the increased solubility of FeS.

Ikeda Ueda and Mukai42 examined the effects of H2S and O2 on CO2 corrosion of pure iron. 3.3 ppm H2S was found to accelerate the cathodic reaction while 33 ppm H2S suppressed the corrosion rate due to formation of a temporary FeS film. Above 150°C, FeCO3 was found to dominate the corrosion process.

Ogundele and White43 conducted potentiodynamic testing in simulated aqueous sour gas production environments at temperatures and pressure up to 95oC and 4.2 MPa. The results were used to produce potential-pH diagrams and to evaluate the mechanisms of corrosion in H2S and CO2 environments. They concluded that corrosion in oilfield environments was very complex with multiple potential oxide and sulfide reaction products and complex ion intermediates.

Ramanarayanan and Smith44 studied the growth kinetics of FeS at 218oC with 10% H2S at 2000 psi (14 MPa). The FeS film produced was Pyrrhotite with small crystals of Pyrite on outer surface. They proposed a mechanism of Fe+2 ion growth through the semiconductive Pyrrhotite and with H2S breakdown at the liquid/sulfide interface to form FeS2. Pyrrhotite growth and dissolution kinetics are also discussed with respect to the practical limitations to the parabolically growing iron sulfide scale.

Pound, Wright and Sharp45 performed cyclic voltamagrams in NaCl, NaHCO3 and Na2SO4 solutions with H2S at pH 5.8. They concluded that Mackinawite "film formation in chloride solutions does not fit a conventional diffusion or pore-resistance model for multilayer films. In contrast, the film in sulfate solutions can be represented in terms of a pore-resistance model.

Criaud and Fouillac46 studied scales that formed in geothermal wells in France. They concluded that the scale formed in the wells was FeS corrosion product caused by reaction of dissolved sulfides with steel and not a FeS mineral deposit from the produced brine. A comparison with similar produced fluids from other regions concluded that chlorides in the brine played an important role in the corrosion mechanism.

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1990's

Ramanarayanan and Smith47 conducted laboratory corrosion scale formation and thermogravimetric tests at 400oF (205oC) with 10% H2S in gaseous and aqueous environments with and without CO2. The scale formed was Pyrrhotite. In the gas phase tests, the rate limiting step of the reaction was found to be the dissociation of H2S on scale surface. In the aqueous testing, short term corrosion rates were limited by Fe2+ diffusion through the scale. At longer times, scale growth was found to be limited by Pyrrhotite dissolution.

Panov, Getmanskii, Enikeev and Fokin48 used photoelectric spectroscopy to examine the surface of steel exposed to a H2S solution both with and without a corrosion inhibitor. Without the corrosion inhibitor, the initial surface was always Mackinawite. When the octadecyl amine inhibitor was included, the FeS surface was more amorphous.

Marcus and Protopopoff49 produced potential-pH (Pourbaix) diagrams starting with Gibbs Free Energies of the Fe-S-H2O system at 25 and 300oC. The sulfur ion activity used was 10-4 and the iron ion activity was 10-6.

Asperger50 investigated mechanisms of pit initiation, propagation and inhibition in high pH, sour environments and compares the proposed mechanism to two pits observed in the field. Mechanisms were proposed to explain the various pit morphologies observed.

Vedage, Ramarayanan, Mumford and Smith51 conducted electrochemical impedance spectroscopy tests of corrosion at various H2S and acidic pH levels at 22 to 95oC. Analysis of the results found two processes involved in the corrosion reaction, a charge transfer process between the FeS and the metal and a transport process through the FeS film. Analysis of the impedance curves also indicates the latter step is rate limiting.

Smith52 proposed an oilfield corrosion mechanism for low H2S environments that is based upon the formation and dissolution of Mackinawite corrosion products. The discussion includes methods for determining the equilibrium constants for Mackinawite and Pyrrhotite at temperature as well as the impact of Fe2+ complexes upon the FeS formation and dissolution equilibriums.

Smith and Wright53 presents an equation for the calculation of the minimum H2S required to form Mackinawite that is based in equilibrium thermodynamics. The paper also discusses the effects of temperature, total pressure, HCO3

- ion concentration, the partial pressure of CO2, and non-ideality of both the gaseous and liquid phases upon the minimum H2S calculation.

Lyle and Schutt54 studied CO2 / H2S corrosion under wet gas pipeline conditions in the presence of alkalinity, chloride, and oxygen. Weight loss and LPR measurements were used over a range of 0 to 2 psi (0 to 14 kPa) H2S and 0 to 20 psi (0 to 140 kPa) CO2 pressures at 60°F (15oC) to establish operational guidelines for safe H2S concentrations. Tests were conducted with both fully and partially immersed samples for 14 days at a pH of 3.8 to 7.2. Pyrrhotite was detected in the vapor phase corrosion product where pitting occurred. The authors suggest that a sulfur-containing species such as polysulfide caused the pitting found under the vapor phase corrosion product. A predictive equation was developed for the fully immersed results, but the partially immersed results could not be modeled.

Cheng, Ma, Zhang, Chen, Chen and Yang55 conducted potentiostatic polarization and electrochemical impedance spectroscopy tests in Na2SO4 solutions with pH levels of 1 to 5 at 20oC. The

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H2S concentrations ranged from 0.04 mmol/l to 30 mmol/l. Nyquist plots of the results showed two capacitive loops corresponding to two adsorbed intermediates. A mechanism including adsorbed FeSH- and FeSH species was postulated.

Kvarekval56 measured corrosion rates on AISI 1010 carbon steel in a glass/titanium flow loop. The flow loop contained a single-phase test fluid of distilled water with 3% NaCl saturated with CO2 or a CO2/H2S (1000:1) gas mixture. The tests were run at temperatures from 25 to 90°C, pH of 3.85 to 4.15, and flow velocities ranged from 0.25 to 2.5 m/sec. Steady-state corrosion currents were compared to cathodic limiting diffusion currents obtained under corresponding conditions. The limiting currents were calculated from potentiodynamic sweeps carried out on platinum. Limiting current plateaus were easily seen at pH values below 5, but vanished with increasing pH.

Anderko and Young57 present a comprehensive thermodynamic/kinetic computational model for H2S/CO2/brine corrosion of carbon steel. The predicted results were compared with data presented by Greco14, et al. and were found to be in good agreement. 2000's

Anderko58 is a further discussion of his earlier paper. This paper compares the results with corrosion rate versus pH data from Shoesmith35 and data from Ikeda42. Once again, good agreement was obtained.

Smith and Pacheco59 presents the results of lab tests that were run to evaluate the minimum H2S levels required to form Mackinawite. The paper also presents a "simplified" equation for the calculation of the minimum H2S required for Mackinawite formation and for the determination of the critical H2S/CO2 ratio. A method to determine the conditions required for Mackinawite corrosion products to transition to Pyrrhotite is also presented.

γP*10SppH

pH*2T*4.575

∆G

2

0T −

= (2)

where

++−−−+=2

7520

TT

10*2.147T

10*4.196T*.135ln(T)*T*11.297T*93.473961.1∆G

T

ln(T)*10*7.2ln(T)*T*0.0196 42 + (3)

ppH2S is the minimum partial pressure of H2S required to form mackinawite γ is the fugacity coefficient for H2S is the gas phase, T is temperature in oK, and P is the system pressure.

Dougherty60 provides a succinct review of the factors that influence protective films, including iron sulfide films. The factors include galvanic effects, chlorides, pH, and velocity.

11

Smith and Pakalapati61 reports on 35 years of field experience with corrosion in a sour gas field producing 42% CO2 and 19% H2S. Downhole corrosion was a major issue early in the field life despite corrosion inhibition. With time, the corrosion rates reduced and the downhole tubing life was greatly extended. This was tied to a reduction in reservoir pressure. The explanation for this reduction was the elimination of minute levels of elemental sulfur from flow stream as pressure dropped below the hydrocarbon dew point and the produced condensate began to serve as in in-situ sulfur solvent.

GEOSCIENCE AND THERMODYNAMICS LITERATURE

In addition to the corrosion literature, valuable information about the chemistry, structure and formation of iron sulfide scales can be found in the Geoscience and Thermodynamics literature. Often, geochemists will react iron or iron compounds with sulfides to synthesize minerals for study. The data produced by these studies can provide valuable insights into the corrosion reaction as well as information that is required to predict the formation and transformation of iron sulfides. Amorphous FeS

The literature on amorphous iron sulfides is limited, partially because they are so unstable and either dissolve or transform into Mackinawite so quickly. Bagander and Carman62 provides a solubility constant for amorphous FeS. The experimental conditions used to determine the Ksp included a sea salt solution at 15.8oC. The amorphous FeS composition was estimated to be FeS0.86. Mackinawite

Clark63 reviews of the available data that is used to define the upper stability temperature for Mackinawite. This value was said to be near 135oC, but the upper temperature limit is difficult to define precisely due to the reaction times involved and the effect that low level impurities of ions such as nickel and cobalt can have upon Mackinawite stability. Clark also proposes a composition for Mackinawite that contains between 50.9 and 51.6 atomic percent Fe.

Berner64 determined the solubility constants and standard free energy of formation values for Mackinawite, Greigite and amorphous FeS. The reactions as FeS transitions from amorphous FeS to Mackinawite and then on to Greigite, Pyrrhotite or Pyrite also are discussed.

Berner65 is also credited with the initial "proof" that demonstrated that Kansite and Mackinawite are the same compound. This led to the crossover paper by Milton discussed earlier and the eventual replacement of the term Kansite with Mackinawite.

Rickard66 describes the kinetics of FeS formation from solution including intermediate complexes and phases and the kinetic rate coefficients that are required to predict Mackinawite formation rates. The paper concludes that the formation of Mackinawite is very rapid, on the order of seconds or fractions of seconds.

Ritvo, White and Dixon67 very recently identified a new FeS intermediate between amorphous FeS and Mackinawite with the proposed (and unapproved) name of Dorite. Dorite is formed in seawater with trace concentrations of H2S. The structure of the new phase is very similar to Mackinawite, but contains two crystallographic sheets, one is similar to Mackinawite and the other contains far more disorder than Mackinawite.

12

Mullet, Boursiquot, Abdelmoula, Genin and Ehrhardt68 synthesized Mackinawite by reacting

H2S with iron in pH 4.6 deionized water. They studied the resulting solids by XRD, TEM, Transmission Mossbauer spectroscopy and XPS. The study results address non-stoichiometry of FeS both in the bulk and on the crystal surface and propose the presence of both Fe3+ as well as polysulfide and elemental sulfur in the Mackinawite lattice.

Wolthers, Van der Gaast and Rickard69 studied synthesized Mackinawite and found that there are two forms, a disordered and highly reactive form and the more stable, more ordered form. The study defined difference between two crystal structures as a difference in the scale of the long range order of the crystal defect structure.

Pankow and Morgan70 evaluated the dissolution rates of Mackinawite as a function of pH, temperature and ionic strength. Their study measured the dissolution rate coefficients and estimated the Arrhenius activation energy for the reaction in the temperature range of 5 to 35oC.

Evans, Milton, Chao, Adler, Mead, Ingram and Berner71 is the initial geochemical paper introducing the mineral Mackinawite. Previously, Mackinawite had been confused with a rare mineral, Valleriite, which had the composition CuFeS2. Mackinawite does not contain copper and is named for the Mackinaw mine in Snohomish County, Washington, USA.

Taylor, Rummery and Owen72 studied the reaction sequence for the formation and transition of various FeS compounds at temperatures up to 160oC. They found that Mackinawite usually forms first and can then transforms on to Greigite then to Pyrrhotite and finally to Pyrite. They also found that even when conditions for transformation are optimal, the transformation sequence to reach Pyrite can require a few days for temperatures as low as 100oC.

Taylor73 discusses the formation chemistry that leads to the precipitation of Mackinawite from bulk solution. The reaction sequence starts with Fe2+ ions, which react with H2S to form FeSH+ ions. Two FeSH+ ions then combine to form a Fe2S2 molecule, which starts to combine with other Fe2S2 molecules. Once a critical number of Fe2S2 molecules have joined, the solid begins the transformation to Mackinawite. Cubic FeS

Takeno, Zoka and Niihara74 exposed iron plates to H2S saturated distilled water adjusted to a variety of pH values. The exposures ranged in time from 6 to 384 hours. In the tests, Mackinawite, Troilite, and hexagonal Pyrrhotite were formed. Cubic FeS seemed to form most readily in the neutral pH range. After formation, Cubic FeS was also found to transforms gradually to Mackinawite at room temp with change complete within a month and transformed into hexagonal Pyrrhotite when heated to 200oC.

De Medicis75 defined the crystal structure and x-ray diffraction pattern for Cubic FeS. De Medicis formed Cubic FeS on iron in a 2 to 3 day exposure to a 0.01M H2S solution at pH 4 to 4.5. The testing also found that Cubic FeS did not form in presence of O2 or chlorides.

Murowchick and Barnes76 produced Cubic FeS at temperatures below 92oC and pH in the range of 2 to 6 in 4 to 85 hours. They also found the need to avoid chlorides. Above pH 6, Mackinawite was

13

formed in preference to Cubic FeS. Above 92oC, the FeS phase formed was either Troilite or Mackinawite. Smythite

Erd, Evans and Richter77 described the Smythite crystal structure. Smythite deposits tend to form in hexagonal flakes, much like Pyrrhotite, and it is strongly ferromagnetic. It has a rhombohedral crystal structure. It is similar in overall crystal structure to Pyrrhotite along a axis, but stacked differently along the c axis. Erd described the ideal composition of Smythite as Fe3S4. Erd was not able to synthesize Smythite.

Almost a decade later, Rickard78 was able to synthesize Smythite by reacting Siderite (FeCO3) with H2S at pH of 6.5 to 8 and 25oC. In these tests, Mackinawite was always found as a minor phase in the acidic pH range. At higher pH values, unreacted Siderite was frequently found. Greigite

Lennie, Redfern, Champness, Stoddart, Schofield and Vaughan79 studied the transformation reaction of Mackinawite to Greigite. They found that the transformation occurs rapidly during heating at temperatures above 373oK. They also found that the general structure of the Mackinawite was retained in the new Greigite. In their tests, Mackinawite most readily transformed to Greigite upon the exposure to oxidizing species such as dissolved O2. Since Greigite is generally believed to be Fe3S4, the oxidation of Fe2+ to Fe3+ may play a role in the transformation. This may also explain why Greigite is often an intermediary between Mackinawite and Pyrite.

Skinner, Erd and Grimaldi80 described the structure of Greigite as Fe3S4 and defined the crystal structure and x-ray diffraction pattern. In their work, they concluded that Greigite is a polymorph of Smythite since both have the nominal composition of Fe3S4. Pyrrhotite

Pyrrhotite is a non-stoichiometric form of FeS that forms hexagonal close packed (HCP) crystals. Since Pyrrhotite is not stoichiometric, ordering of the Fe2+ ion vacancies in the HCP crystal lattice under equilibrium conditions give rise to a variety of super cell crystal structures. Troilite is the form of Pyrrhotite when the Fe2+ ion vacancy concentration is essentially zero.

Desborough and Carpenter81 describe the phase relationships between the various forms of Pyrrhotite crystal structures as a function of temperature. Subtypes of Pyrrhotite include monoclinic and a variety of hexagonal supercells that differ only in the length of the c axis between repeating Fe2+ vacancies. Fe2+ content in Pyrrhotite was reported to range from 45.5 to 50 percent at temperatures below 315oC.

Clark82 defined the stability field of monoclinic Pyrrhotite as a function of temperature and Fe2+ content (46.4 to 46.8 atomic %). The maximum temperature for monoclinic Pyrrhotite was found to be 308oC, which is the β transformation temperature for a number of Pyrrhotite crystalline structures into a single hexagonal crystalline form.

14

Carpenter and Desborough83 discuss the crystallographic relationships between Troilite and the various Pyrrhotite forms. Troilite is defined as stoichiometric FeS whereas Pyrrhotite generally contained 46.5 to 48 percent Fe in the samples studied.

Arnold84 extends the work of Desborough and Carpenter by discussing the two phase Pyrrhotite regions where one Pyrrhotite phase exist as mixed FeS lamellae within another.

Arnold and Reichen85 provide a method for using x-ray diffraction to determine the Fe2+ content of Pyrrhotite. The procedure uses the shift in the d(102) spacing as a function of Fe2+ vacancy concentration for this correlation.

Gronvold, Westrum and Chou86 measured the thermodynamic heat capacity, the enthalpy and the entropy of Troilite and Pyrrhotite in the temperature range from 5 to 350oK. At 298.15oK, the heat capacity, entropy and enthalpy were reported respectively as 12.08 cal/mole oK, 14.415 cal/mole oK and 7.496 cal/mole for Troilite and 11.92 cal/mole oK, 14.529 cal/mole oK and 7.396 cal/mole oK for Pyrrhotite.

Tokonami, Nishiguchi and Morimoto87 discuss the defect structure of monoclinic Pyrrhotite with the vacancies located in every other HCP layer and with a long range repeating pattern of vacancy alignment along the c axis. Pyrite

Rickard88 studied the kinetics of Pyrite formation from the reaction FeS + S → FeS2. The investigation concluded that there must be an oxidizing agent present for Pyrite to form. The rate constants are very slow in comparison to Mackinawite formation, even at 125oC. The reaction pathway was found to require the presence of polysulfide species and possibly requires the involvement of a Greigite intermediary.

Rickard and Luther89 also studied the complex multi-step reaction mechanism required to convert FeS to Pyrite. The proposed reaction requires H2S adsorbed to the FeS surface to decompose to adsorbed So and H2 in order for form the FeS2.

Gronvold and Westrum90 report on the heat capacity, enthalpy and entropy of Pyrite from 5 to 350oK. At 298.15oK, the heat capacity, entropy and enthalpy was reported respectively as 14.86 cal/mole oK, 12.65 cal/mole oK and 2302 cal/mole

Wilkin and Barnes91 discuss Pyrite formation by the reaction of Mackinawite and Greigite with various sulfur species at 70oC and pH in the range of 6 to 8.

Schoonen and Barnes92 describe a detailed reaction sequence for the formation of Pyrite from amorphous FeS, Mackinawite and Greigite. They also conclude that H2S and/or HS- are not adequate oxidizers to form Pyrite at 65oC.

Rickard93 presents a study of rate constants for Pyrite formation from 20 to 50oC.

15

CHEMISTRY/THERMODYNAMICS

In order to be able to develop corrosion predictions for environments that contain H2S and CO2, it is necessary to have a detailed knowledge and reliable thermodynamic for the various ionic and solid species that could potentially form. The following series of papers either present procedures for developing potential-pH (Pourbaix) diagrams for the Fe-H2S-CO2-H2O system or documents the thermodynamic data for the solids or ion complexes that could be encountered.

Criaud, Fouillac and Marty94 developed potential-pH diagrams for the Fe/S/H2O/CO2 system at 60 and 80oC with H2S, Fe and CO2 activities respectively at 10-5, 10-5 and 5x10-3

Similarly, Horvath andNovak95 developed numerous potential-pH diagrams for a variety of metal sulfides, including iron, and discussed the relevance of each diagram to corrosion control.

Dyrssen96 calculates the ionic speciation of metal-sulfide complexes in seawater. Log K data are provided for FeHS+, Fe(HS)2, Fe(HS2)-, and FeS2

- at 25oC.

Landing and Westerlund97 used Dyrssen's data to review the solubility of Mackinawite and Greigite in seawater.

Luther and Ferdelman98 evaluated the reported the stability constants for FeSH+, Fe2(SH)3+,

Fe(H2S)2+, and FeS(aq) in marine waters against acid-base titration data for bay water samples from the Chesapeake Bay. They found that the complex ions formed in the laboratory were inconsistent with the data whereas field samples were much more stable and had log K values consistent with values reported in the literature. The proposed explanation was the presence of organic chelating agents in the field samples that stabilize the Fe-S complexes.

McDonald and Syrett99 present potential-pH diagrams for geothermal brines with H2S at temperatures of 25 and 250oC.

Ribbe100 provides a general summary of the phase behavior, crystal structure, thermal stability and thermochemical relationships between the various Fe-S compounds.

Pound, Wright and Sharp101 present both potential-pH and potential-pS diagrams for geothermal brines at 298-573oK.

Berner102 produced potential-pS stability diagrams for the Fe-S-CO2-H2O system at 25oC and for various pH and PCO2 levels. The diagrams shows the interrelationship between FeSx, FeCO3 and FexOy phases as the partial pressure of H2S and the EH increase.

Bouet and Brenet103 published a theoretical potential-pH diagram at 25oC for the Fe-S system with 10-1 and 10-4 molar S. Diagrams based upon the Fe(OH)3 and Fe2O3 type oxide phases were provided.

And finally, Morse, Millero, Cornwell and Rickard104 published an extensive review of the chemistry of H2S and FeS in natural water systems. The various aqueous phases, FeS transformation sequences, crystal structures, reaction rate kinetics and even the effects of some cationic impurities are all discussed.

16

CURRENT SITUATION

The current state-of-the-art for H2S corrosion recognizes that there are a variety of FeS corrosion products that form depending upon the conditions. Figure 1 illustrates the transformation relationships between the primary forms of FeS as well as Fe2+ ions and FeCO3. Figure 5 presents a different type of view depicting the effects of temperature and H2S partial pressure upon the type of FeS corrosion product that forms.

From a practical standpoint, knowing which form(s) of FeS are formed in the field can now provide insights into the corrosion conditions in the facilities. Knowledge of the specific type(s) of FeS formed in the field is also necessary for the construction of a realistic laboratory simulation. Obviously, if the Smythite is being formed in the field, a laboratory test that produces Mackinawite or Pyrrhotite will not accurately reflect field conditions.

Generally accepted corrosion rate prediction algorithms similar to deWaard and Milliams do not yet exist for any of the FeS species. However, the geochemical and thermodynamic literature provides a great deal of information about the reactions that are required to form and retain each of the species. The literature also provides guidance for the estimation of the concentrations of Fe-H2S complex species in solution. Any algorithms developed will probably need to account for these compounds to accurately predict FeS formation, especially at the lower end of the H2S activity spectrum where there is a question whether the corrosion product will be FeS, FeCO3 or FexOy.

The existing corrosion knowledge for each of the species can be summarized as follows: Mackinawite - Tewari21, 27, 28 and Rickard61, 64, 99 have identified the corrosion mechanism that leads to Mackinawite formation as well as the metal surface reactions, the kinetics of the reactions. Others have provided definition to the conditions that define the stability regions for Mackinawite. Mackinawite is generally viewed as the initial form of FeS corrosion product to form over a period that may be less than one second. Cubic FeS - DeMedicis70 and Murowchick71 defined the conditions required to form Cubic FeS. The conditions are similar to those required to form Mackinawite except that the environment cannot contain chlorides. Cubic FeS is also stable to higher temperatures, but will decompose to Mackinawite or Pyrrhotite over a period of a day or two. Pyrrhotite - Ramanarayan43, 46 and Vedage49 defined the corrosion mechanism for aqueous environments and identified the rate limiting step. However, there are currently practical limitations to converting the Fe2+ ion diffusion rate through the Pyrrhotite scale into a practical algorithm for use in predicting corrosion rates. There is also a major gap in the absence of a defined mechanism for pit initiation and growth since equipment failures in Pyrrhotite forming environments are generally due to pitting and not general weight loss corrosion. Pyrrhotite formation generally occurs over a period of few days, but can occur much more quickly at temperatures where Mackinawite and Cubic FeS are not stable. Pyrite and Marcasite - The geochemical literature has established the need for elemental sulfur to be present for the formation of either type of FeS2. Rickard83, 84, 88 defined the kinetics for Pyrite formation, which indicates that Pyrite will generally form over a period of about a week. Ramanarayan found that even when Pyrite did form, the kinetics often resulted in the formation of Pyrite cubes buried in a matrix of Pyrrhotite.

17

Greigite - Only a limited about of work has been done to define the formation and/or transition requirements for Greigite. It appears that there is a need for an oxidation reaction step that forms Fe3+ to produce Greigite from Mackinawite as an alternative to the direct transformation of Mackinawite to Pyrrhotite. The difficulty with past work may be the question of inadvertent oxygen contamination of test solutions that could have resulted in Fe3+ generation. The observation by Lennie74 that Greigite is often an intermediary between Mackinawite and Pyrite in laboratory synthesized scales may also be an indicator of the requirement for an oxidizer. Smythite - Smythite is definitely the least studied form of FeS by the geochemists and is almost completely ignored as a corrosion product. Rickard73 was able to demonstrate that Smythite appears to the transformation product of Siderite that has been exposed to H2S.

AREAS FOR FURTHER STUDY • Flow effects on Mackinawite and Pyrrhotite mechanisms • Confirmation of CO2/H2S transition levels of H2S reported in C/2003 paper • Determination of role of chloride in pitting mechanism • Confirmation of relationship between H2S activity and Pyrrhotite defect concentration at lower

temperatures • Definition of roles of Greigite and Cubic FeS • Determination of effect of Smythite upon further corrosion • Information of Fe-S ion pairs and complex species to enhance predictions • Top of Line Corrosion in the presence of H2S + CO2 • Influence of CO2 partial pressure on H2S corrosion • Effect of solids (produced sand) deposition on H2S corrosion. • Further information on role of oxidizing agents on H2S corrosion and power of agents required to

effect any changes. • Measure the chemical diffusivity and concentration of Fe ion vacancies in Pyrrhotite as a function of

sulfur activity and temperature for production conditions. • The influence of rogue H2S on corrosion in typical packer fluids.

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Figure 1 Corrosion product reaction pathways from Smith and Miller19 and Morse104

-7

-6

-5

-4

-3

-2

-1

1 2 3 4 5

pH

log(

Ferr

ous

mol

al)

Pyrite Hexagon Pyrrhotite Monoclinic PyrrhotiteTroilite Mackinawite

Figure 2

Solubility of FeS as a Function of pH at 25oC from Tewari20

Fe++

FeCO3

S0

(S--)

Pyrite

(S--n)

Marcasite

Greigite

Mackinawite

Pyrrhotite Smythite equilibrium

equilibrium

equilibrium

equilibrium

(S--)

(S--)

(S--n)

24

0

50

100

150

200

250

0 0.2 0.4 0.6 0.8 1

Percent H2S

Cor

rosi

on R

ate,

mpy

25C 60C

Figure 3

Corrosion Rate as a function of CO2/H2S and temperature from Murata24

Figure 4 Corrosion Product formation transitions from Shoesmith30

Carbon Steel + H2S + H2O

film rupture & precipitation

solid state Mackinawite Cubic FeS

FeS1-x Troilite

MackinawiteFeS1-x

25

Pyrrhotite

Mackinawite Pyr

ite

FeCO3

Cubic FeS

Fe+2

Oilfield Corrosion Products

log H2S activity

Tem

pera

ture

Figure 5 Corrosion product formation as a function of temperature and H2S

from Smith and Pacheco59

26