06115 Corrosion of Carbon Steel by h2s in Co2 Containing Oilfield Environments (51300-06115-Sg)

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CORROSION OF CARBON STEEL BY H2S IN CO2 CONTAINING OILFIELD

ENVIRONMENTS

Stephen N. Smith

ExxonMobil Production Company

800 Bell, CORP-EMB-2003HHouston, Texas 77381

Michael W. Joosten

ConocoPhillips

Bartlesville Technology CenterBartlesville, Oklahoma 74004

ABSTRACT

The effect that even small concentrations of H2S can have upon CO2 corrosion has beenrecognized 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 exposureconditions. 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 theimpact 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 ofunknowns 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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06115

Paper No.

©2006 NACE International. Requests for permission to publish this manuscript in any form, in part or in whole must be in writing to NACE International,Conferences Division, 1440 South Creek Drive, Houston, Texas 77084. The material presented and the views expressed in this paper are solely those ofthe author(s) and are not necessarily endorsed by the Association. Printed in the U.S.A.

Copyright



temperature, pH, bicarbonate, velocity and a host of others had been studied extensively. Bycomparison, 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 variousmetallic 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. Themineralogists 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 containCO2 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 oilfieldenvironments. 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 investigatorsworking 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 haveconcluded 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 understandingof 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 of100oF (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. Testswere 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, and3. 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 andalso 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:

( )( )[ ]22

2 log0295.01495.11 +−−−= aH aS R

I S H

(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 corrosiontests, 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 andfinally 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 ofliquid 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 exposureof 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 thatthey 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 concentrationfor 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 concludedthat 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 bean 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 alsoconcluded 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 FeSformed. 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 ofKansite (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 15hrs. 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 todetermine 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 phasecorrosion 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 mechanismapproaches complete interface control."

In their subsequent paper on protectiveness of FeS as a function of pH, Sardisco and Pitts foundthe 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 of4.0 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 geologistshad 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 aqueoussolution 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 constantH2S 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 pressureof 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 weightsand 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 wasevidence 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 asa 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 layersuch 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.

Zitter 18

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 alsoobserved, 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 ofCaCO3 from the mixture of CO2 and CaCl2.

In 1975, Smith and Miller 19 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 ofFeS 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

5



diagram from a geochemistry reference that relates the various reaction paths between the various formsof 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, Pyrrhotiteand 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 hexagonalPyrrhotite, and Pyrite as a function of temperature (25 to 125

oC) 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 couponsto 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 wasobserved. 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 ofcrystallization of the corrosion product and increasing oxygen concentrations. The timeframe for theoxidation 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 attemperatures 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 rangedfrom 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 beuseful as a starting point for future research.

Murata, Matsuhashi, Taniguchi and Yamamoto24 evaluated the corrosion rate as a function ofH2S 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)

6



better protective layers are formed, even in the case of high salt concentrations." Milliams and Kroesealso 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 productlayers. 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 - 152MPa) and bottom hole temperatures (BHT) of 365 to 385

oF (185 to 195

oC). 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 determineMackinawite 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+ ionat 22oC.

Tewari, Bailey and Campbell28 extended the Mackinawite dissolution rate work of Tewari andCampbell 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 totransform to the higher FeS phases were iron ion transport from the surface is not as rapid.

Hausler 29

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 ironvacancies 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 foundthat 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 wasonly observed at pH 7.

Morris, Sampaleanu and Veysey31

conducted polarization studies on rotating disks and in a flowloop 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-pHdiagram 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 to1.5 MPa H2S at temperatures of 308, 373 and 433

oK. 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 rotationspeeds 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 dueto 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 H2Scorrosion.

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 excellentoverall 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, evenwhen 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 tocorrosion.

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 comparescorrosion 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 protectiveFeS films. At 25

oC, 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 H2Sat 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 ambienttemperatures 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 25o

C to determine whether the corrosion productwill 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 orreduce 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 ofMackinawite. 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 solutionsdoes 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 concludedthat 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 producedfluids 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 andthermogravimetric 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 termcorrosion 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 wasincluded, the FeS surface was more amorphous.

Marcus and Protopopoff 49 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.

Asperger 50 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. Mechanismswere proposed to explain the various pit morphologies observed.

Vedage, Ramarayanan, Mumford and Smith51

conducted electrochemical impedancespectroscopy 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 indicatesthe 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 theimpact 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 oftemperature, 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 phasecorrosion 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 wasdeveloped for the fully immersed results, but the partially immersed results could not be modeled.

Cheng, Ma, Zhang, Chen, Chen and Yang55

conducted potentiostatic polarization andelectrochemical impedance spectroscopy tests in Na2SO4 solutions with pH levels of 1 to 5 at 20oC. The

10



H2S concentrations ranged from 0.04 mmol/l to 30 mmol/l. Nyquist plots of the results showed twocapacitive 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 comparedto cathodic limiting diffusion currents obtained under corresponding conditions. The limiting currents

were calculated from potentiodynamic sweeps carried out on platinum. Limiting current plateaus wereeasily seen at pH values below 5, but vanished with increasing pH.

Anderko and Young57

present a comprehensive thermodynamic/kinetic computational model forH2S/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 wasobtained.

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 calculationof 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.

γ

P*10

S ppH

pH*2T*4.575

∆G

2

0

T −

= (2)

where

++−−−+=2

7520

TT

10*2.147

T

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, andP is the system pressure.

Dougherty60

provides a succinct review of the factors that influence protective films, includingiron 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 greatlyextended. 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 andeither dissolve or transform into Mackinawite so quickly. Bagander and Carman62 provides a solubility

constant for amorphous FeS. The experimental conditions used to determine the K sp included a sea saltsolution at 15.8oC. The amorphous FeS composition was estimated to be FeS0.86.

Mackinawite

Clark 63 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 nickeland cobalt can have upon Mackinawite stability. Clark also proposes a composition for Mackinawite

that contains between 50.9 and 51.6 atomic percent Fe.

Berner 64 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.

Berner 65 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 eventualreplacement of the term Kansite with Mackinawite.

Rickard66

describes the kinetics of FeS formation from solution including intermediatecomplexes 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 amorphousFeS 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 moredisorder 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, TransmissionMossbauer 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 aretwo forms, a disordered and highly reactive form and the more stable, more ordered form. The studydefined 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 theArrhenius activation energy for the reaction in the temperature range of 5 to 35oC.

Evans, Milton, Chao, Adler, Mead, Ingram and Berner 71 is the initial geochemical paperintroducing 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.

Taylor 73 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 transformationto 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. DeMedicis formed Cubic FeS on iron in a 2 to 3 day exposure to a 0.01M H2S solution at pH 4 to 4.5. Thetesting 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 Richter 77

described the Smythite crystal structure. Smythite deposits tend toform 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 stackeddifferently along the c axis. Erd described the ideal composition of Smythite as Fe3S4. Erd was not ableto 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 wasretained 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 isoften an intermediary between Mackinawite and Pyrite.

Skinner, Erd and Grimaldi80

described the structure of Greigite as Fe3S4 and defined the crystalstructure and x-ray diffraction pattern. In their work, they concluded that Greigite is a polymorph ofSmythite 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 underequilibrium 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 Carpenter 81 describe the phase relationships between the various forms of

Pyrrhotite crystal structures as a function of temperature. Subtypes of Pyrrhotite include monoclinic anda 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.

Clark 82 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 asingle hexagonal crystalline form.

14



Carpenter and Desborough83

discuss the crystallographic relationships between Troilite and thevarious 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+

contentof Pyrrhotite. The procedure uses the shift in the d(102) spacing as a function of Fe

2+ vacancy

concentration for this correlation.

Gronvold, Westrum and Chou86

measured the thermodynamic heat capacity, the enthalpy and theentropy 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 and7.496 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 aGreigite intermediary.

Rickard and Luther 89

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 toadsorbed 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 14.86cal/mole oK, 12.65 cal/mole oK and 2302 cal/mole

Wilkin and Barnes91

discuss Pyrite formation by the reaction of Mackinawite and Greigite withvarious 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 fromamorphous 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 solidspecies 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 andNovak 95 developed numerous potential-pH diagrams for a variety of metalsulfides, 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 andGreigite 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 thedata 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 attemperatures of 25 and 250oC.

Ribbe100

provides a general summary of the phase behavior, crystal structure, thermal stabilityand 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.

Berner 102 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 areall 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 FeSformed 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 Pyrrhotitewill not accurately reflect field conditions.

Generally accepted corrosion rate prediction algorithms similar to deWaard and Milliams do notyet 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 speciesin 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 thanone second.

Cubic FeS - DeMedicis70

and Murowchick 71

defined the conditions required to form Cubic FeS. Theconditions 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 orPyrrhotite over a period of a day or two.

Pyrrhotite - Ramanarayan43, 46

and Vedage49

defined the corrosion mechanism for aqueous environmentsand 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 andgrowth 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 matrixof Pyrrhotite.

17



Greigite - Only a limited about of work has been done to define the formation and/or transitionrequirements 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 toPyrrhotite. 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 anindicator of the requirement for an oxidizer.

Smythite - Smythite is definitely the least studied form of FeS by the geochemists and is almostcompletely 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 lowertemperatures

• 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 toeffect 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.

REFERENCES

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in Oil Production," Proceedings of the Fourth World Petroleum Congress, Rome, 479-499, 1955.3. Ewing, Scott P., "Electrochemical Studies of the Hydrogen Sulfide Corrosion Mechanism,"

Corrosion, Vol. 11, pp. 497t-501t, 1955.

4. Sorell, G., and W.B. Hoyt, "Collection and Correlation of High Temperature Hydrogen SulfideCorrosion Data," NACE Publication 5B156.

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Sulfide/Hydrogen Mixtures," J. Electrochemical Society, Vol. 105, pp. 183-191, 1958.

18



7. Shannon, Donald W., and James E. Boggs, "Factors Affecting the Corrosion of Steel byOil-Brine-Hydrogen Sulfide Mixtures," Corrosion, Vol. 15, pp. 299-302, 1959.

8. Newman, T.R., "A Laboratory Method for Evaluating Corrosion Inhibitors for Secondary Recovery,"Corrosion, Vol. 15, pp. 307-310, 1959.

9. Hughes, W.B. and V.L. Stromberg, "Effect of Inhibitors on Protective Properties of Iron Sulfide

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18. Zitter, H., "Corrosion Phenomena in Sour Gas Wells," Erdoel-Erdgas Zeitschrift, Vol. 89, No. 3, pp.

101-106, 1973.19. Smith, J.S., and J.D.A. Miller, "Nature of Sulfides and Their Corrosive Effect on Ferrous Metals: A

Review," British Corrosion Journal, Vol. 10(3), pp. 136-143, 1975.

20. Tewari, P.H., G. Wallace and A.B. Campbell, "The Solubility of Iron Sulfides and Their Role inMass Transport in Girdler-Sulfide Heavy Water Plants," AEC of Canada Report AECL-5960, 1978.

21. Thomason, William H., "Formation Rates of Protective Iron Sulfide Films on Mild Steel in H2S

Saturated Brine as a Function of Temperature", Corrosion/78, paper 41, 1978.

22. Craig, Bruce D., "The Nature of Iron Sulfides Formed on Steel in an H2S-O2 Environment,"

Corrosion, Vol. 35, pp. 136-138, 1979.23. Fryt, E.M., W.W. Smeltzer and J.S. Kirkaldy, "Chemical Diffusion and Point Defect Properties of

Iron Sulfide at Temperatures 600-1000 C," J. Electrochemical Society, Vol. 126, pp. 673-682,1979.

24. Murata, T., R. Matsuhashi, T. Taniguchi and K. Yamamoto, "The Evaluation of H2S ContainingEnvironments from the Viewpoint of OCTG and Linepipe for Sour Gas Applications," Offshore

Technology Conference, paper 3507, 1979.

25. Milliams, D.E. and C.J. Kroese, "Aqueous Corrosion of Steel by H2S and H2S/CO2," ThirdInternational Conference on the Internal and External Protection of Pipes, 1979

26. Hamby, Tyler W., "Development of High Pressure Sour Gas Technology," SPE paper 8309, 1979.

27. Tewari, Param H. and Allan B. Campbell, "Dissolution of Iron During the Initial Corrosion ofCarbon Steel in Aqueous H2S Solutions," Canadian Journal of Chemistry, Vol. 57, pp. 188-196,

1979.28. Tewari, P.H., M.G. Bailey and A.B. Campbell, "The Erosion-Corrosion of Carbon Steel in Aqueous

H2S Solutions up to 120oC and 1.6MPa Pressure," Corrosion Science, Vol. 19, pp. 573-585, 1979.

29. Hausler, R.H., "Corrosion Inhibition in the Presence of Corrosion Product Layers," 6th European

Conference on Corrosion Inhibition, Ferrara, Italy, Sept. 1985.

19



30. Shoesmith, David W., Peter Taylor, M. Grant Bailey and Derrek G. Owen, "The Formation ofFerrous Monosulfide Polymorphs During the Corrosion of Iron by Aqueous Hydrogen Sulfide at

21oC," J. Electrochemical Society, Vol. 127(5), pp. 1007-1015, 1980.31. Morris, D.R., L.P. Sampaleanu and D.N. Veysey, "The corrosion of steel by aqueous solutions of

hydrogen sulfide," J. Electrochemical Society, 127(6) 1228-35, 1980.

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23



Figure 1

Corrosion product reaction pathways from Smith and Miller 19 and Morse104

-7

-6

-5

-4

-3

-2

-1

1 2 3 4 5

pH

l o g ( F e r r o u s m o l a l )

Pyrite Hexagon Pyrrhotite Monoclinic Pyrrhotite

Troilite 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

C o r r

o s i o n R a t e , m p y

25C 60C

Figure 3Corrosion Rate as a function of CO2/H2S and temperature from Murata24

Figure 4Corrosion Product formation transitions from Shoesmith30

Carbon Steel + H2S + H2O

film rupture & precipitation

solidstate

MackinawiteCubic FeS

FeS1-xTroilite

MackinawiteFeS1-x

25



Pyrrhotite

Mackinawite P y r i t e

FeCO3

Cubic FeS

Fe+2

Oilfield Corrosion Products

log H2S activity

T e m p e r a t u r

e

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

from Smith and Pacheco59

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06115 Corrosion of Carbon Steel by h2s in Co2 Containing Oilfield Environments (51300-06115-Sg)