the role of anion contaminants on corrosion in refinery amine units

Technical Article

of 17 Form No. 170-00276

The Role of Anion Contaminants on Corrosion in Refinery Amine Units L.E. Hakka Union Carbide Canada Inc. P.O. Box 700 Pointe-aux-Trembles, QC HIB 5K8 Sidney F. Bosen Union Carbide Corporation 335 Pennbright, Suite 120 Houston, TX 77090 H.J. Liu Union Carbide Corporation 39 Old Ridgebury Road Danbury, CT 06817-0001 Prepared for Presentation at AIChE 1995 Spring National Meeting Session 55: Advances in Gas Treating March 21-23, 1995 • Houston, TX UNPUBLISHED AIChE shall not be responsible for statements or opinions contained in papers or printed in its publications.

of 17 *Trademark of The Dow Chemical Company Form No. 170-00276

Using the definition

Abstract

Measurement of corrosion rates in aqueous MDEA solutions containing various heat stable amine salts (HSAS) commonly encountered in refinery amine treaters has shown that those formed from medium strength acids, such as formic, acetic and oxalic, are indeed corrosive. Neutralization with strong base reduced the observed corrosion rates by two orders of magnitude. HSAS from strong acids, such as thiocyanic and thiosulfuric, are not corrosive. The observed corrosion rates correlated well with the calculated concentration of undissociated acid present in the solutions, suggesting strongly that the observed corrosion process is under cathodic control, even in the case of the strongly chelating oxalate anion. Hydrogen sulfide and, to a lesser extent, thiocyanate were found to give false positive corrosion rate results, likely due to a competing electrochemical reaction. The role played by heat stable amine salts (HSAS) in the corrosion of refinery amine units is not well understood, in spite of its very significant economic impact. In particular, the presumed relative corrosiveness of the various anions comprising the HSAS and suitable means of mitigation are largely based on experience and anecdotal evidence rather than scientific studies. A recent literature review article on refinery amine system corrosion by Nielson et al.(1), lists only one reference on the corrosiveness of HSAS. A recent series of papers (2,3,4), of which this is one, has started reporting in this are of technology.

Introduction

Refinery amine units utilize aqueous solutions of alkanolamines, such as N-methyldiethanolamine (MDEA), diethanolamine (DEA), and di-isopropanolamine (DIPA), to remove hydrogen sulfide (H2S), carbon dioxide (CO2), carbonyl sulfide (COS) and mercaptans from various product and gas streams. Sulfur recovery unit tail gas clean-up units also use amines, such as MDEA. The amine solutions selectively absorb the contaminants, which are acidic, by means of an acid-based chemical reaction: Amine + HAcid ⇔ AmineH+ + Acid- (1)

Reaction (1) describes the reaction of the amine in the free base form with a proton donor (an acid by the Bronsted-Lowry definition) in equilibrium with the protonated amine (the conjugate acid of the amine) and the anion produced by the ionization of the acid. In the usual way, we can define an equilibrium constant for Reaction (1) as

[AmineH+] [Acid-]

K1 = (2) [Amine] [HAcid]

Also, the acid dissociation constant for the acid and the conjugate acid of the amine are given in the usual form as

[Acid-] [H+] Ka, acid = (3)

[HAcid]

Amine] [H+] Ka, amine = (4)

[AmineH+] Substituting Equations 3 and 4 into Equation 2 and taking logarithms, we get log K1 = log Ka, acid – Ka, amine (5)


PKa = -log Ka (6)

quation 5 becomes log K1 = pKa, amine – pKa, acid (7)

quation 7 clearly illustrates that the position of equilibrium in (1) is determined by the

4.

he e

efinery streams usually contain minor quantities of other acids, most of them stronger 3,

s

ven

ielson et al.(1) provide a good recent overview of the types and location of corrosion e

the

listing of common HSAS anions encountered in refinery systems and their possible he

ugs

he accumulation of high levels of HSAS is undesirable from a process performance point

ed by

ase

- + Na+OH- ⇒Amine + Na+ Acid- + H2O (8)

E Erelative strengths of the acid and base. Table 1 gives pKa values for various acids and bases of interest. The dilute solution pKa’s of MDEA H2S, and CO2 at 25°C are 8.52, 7.0and 6.37 respectively. Thus, for MDEA solutions, using the preceding values, K1 is 30.2 for H2S and 141 for CO2, indicating that absorption is favored at low temperatures. At 120°C however, the pKa’s are 7.03, 6.46, and 6.50 in the same order as above, yielding K1 valuesof 3.7 for H2S and 3.4 for CO2. This shift of equilibrium in Equation 1 towards the left is utilized in a multistage countercurrent steam-driven stripper to regenerate the amine absorbent into a lean amine, which after cooling is ready for absorption once again. Tstripper overhead is an acid gas stream containing the desorbed H2S and CO2. The choicof amine, and thereby the pKa, determines a suitable balance of absorption and desorption tendencies for efficient and economical operation of the scrubbing process. Racids than H2S and CO2. To pick an example, SO2 upon hydration to sulfurous acid, H2SOhas a pKa for the first ionization of 1.82 at 25°C. Substituting this value into Equation 7 gives an equilibrium constant K1 of 5.1 x 106, indicating that the absorption of strong acidinto the amine solvent results in salt formation that is essentially irreversible even at the higher stripper temperature. Thus, the salts of amine with strong acids are called heat stable amine salts (HSAS). Since the amine unit very efficiently captures strong acids, esmall concentrations in the feed stream will eventually lead to a high steady state HSAS level, providing that mechanical losses are not excessive. The role of the various anions that accumulate in the system in the corrosion process is the focus of this study. Nobserved in amine treaters. Liquid phase corrosion attributable to the corrosivity of thamine solution is most often seen in the areas where hot amine is in contact with metal—lean-rich heat exchanger, rich amine piping to the stripper, the stripping tower, the reboiler and the hot lean amine line to the L/R exchanger. This paper will not discuss vapor phase corrosion, erosion-corrosion, hydrogen-induced cracking, or other corrosion phenomena attributable to process, equipment or metallurgical deficiencies. Asources is given in Table 2. Anions, such as formate, acetate and chloride, can enter ttreater with gaseous feed streams because their corresponding acids are sufficiently volatile. The others must enter either 1) from liquid treaters, 2) be carried in by liquid slthrough poorly functioning inlet separators, 3) with poor quality makeup water or 4) be formed in situ through chemical reaction. Tof view. The pronated amine cation of the HSAS is unavailable for absorbing acid gases, thereby decreasing capacity. Secondly, conventional wisdom states that HSAS are corrosive(4,5,6). The acid gas scrubbing capacity of the amine absorbent can be restor“neutralization” with a strong base, such as caustic (sodium hydroxide). The reaction involved is really a displacement of a moderately strong base (the amine) by a strong b(a hydroxide ion), from the salt: AmineH+ Acid


he unprotonated or free base amine so produced is again available for acid gas absorption

ost authors in the literature believe that the inorganic salts produced by neutralization are

, care

g

iu et al.(2) advanced two plausible hypotheses to explain the decrease in corrosion rate

acid

[HAcid] [AmineH+] (9)

earranging,

Ka amine [AmineH+] cid-] (10)

rom Equation 10 it can be seen that with a given amine, varying concentrations of

ble at

reement in

Tand the sodium salt is essentially inert with respect to scrubbing. Mless corrosive than the corresponding amine salts(1,2,3,4,7,8,9), but this position is not unanimous(6). The present data address this question. If neutralization is practicedmust be taken not to exceed the solubility of the alkali metal salts in the amine solution. When solid salts precipitate, scaling of heat transfer and other surfaces will occur, resultinin poorer heat transfer and possible under-deposit corrosion. Foaming is also promoted by suspended solids. Lobserved when MDEA formate HSAS is neutralized. These were 1) the formation of a protective iron oxide film as the pH is raised and 2) the decrease in the concentration ofundissociated acid in equilibrium with the formate anion. The “undissociated acid” hypothesis best fits the experimental data, and it requires that it is the undissociatedthat is the active species promoting corrosion. The concentration of undissociated acid is calculated from the equilibria involved in an amine solution as given in Equations 3 and 4. Since both equilibria are in the same amine solution, the hydrogen ion concentrations in both equations are the same: [H+] = Ka, acid = Ka, amine [Acid-] [Amine] R [HAcid] = [A Ka acid [Amine] Fundissociated acid [HAcid] can be obtained at constant pH by choosing an acid with a different dissociation constant. The ratio of pronated amine over free base amine is a surrogate for the pH. Conversely, equal undissociated acid concentrations are obtainadifferent pH values by changing the strength of the acid. Thus, we have a method of testing the undissociated acid hypothesis. If the observed corrosion rate correlates with [HAcid] for various acids, that would support the hypothesis, on the assumption that the various acids are equally active in promoting corrosion. This assumption is probably not valid quantitatively even for homologous acids, such as formic and acetic acids, but agtrend is reasonable to expect. Furthermore, since activity coefficient data is not available for the systems of interest, except by estimation methods, concentrations rather than activities will be used, further detracting from quantitative interpretation.


Experimental Procedure

Test Solutions Test solutions were prepared gravimetrically using deionized water and commercial-grade MDEA from Union Carbide Corporation. Reagent-grade formic, acetic and oxalic acids (from Aldrich Chemical Company, Inc.) were added to aqueous MDEA to prepare test solutions with the corresponding HSAS. The thiocyanate and thiosulfate HSAS solutions were prepared from the corresponding ammonium salts (purchased from Aldrich Chemical Company, Inc.) by 1) mixing the ammonium salt with MDEA and deionized water in the mole ratio 1: 2: 3.3, 2) heating to about 40°C, 3) stripping off ammoinia a ta pressure of about 150 mm Hg with a slow nitrogen purge and 4) adding gravimetrically the HSAS stock solution to the required amount of water and MDEA to give the desired MDEA and HSAS concentration. The concentration of HSAS is expressed as the weight percent of MDEA in HSAS form, based on the total solution. Experiments requiring neutralized or partially neutralized HSAS solutions were prepared by adding the calculated amount of UCARSOL® Neutralizer DHM to solutions prepared as above. The HSAS content of the test solutions was determined by titration with 0.5N KOH to a pH of 11.2 as measured with an Orion combination glass electrode and an Orion Model 610 pH meter. The initial pH, prior to the start of titration, was recorded and reported. The test solution containing H2S was prepared by bubbling a gas mixture of 2.08% H2S in nitrogen through the test solution until the desired loading was achieved. The H2S concentration was determined by iodimetric titration, and reported as moles H2S/mole MDEA. The concentration of MDEA in all cases was determined by titration with 0.5N HCl to a pH of 4.5. Anion concentrations were determined with a Dionex ion chromatograph. Corrosion Tests Corrosion rates were measured by the Linear Polarization Resistance (LPR) technique using a “Gamry” potentiostat(10). The experimental technique is reported in detail by Asperger and Dean(3). All measurements were performed at 120 °C in a borosilicate glass-lined autoclave, using a stirrer speed of 200 rpm. Where used, SAE 1018 corrosion coupons were mounted on the LPR electrode posts with “Teflon” tape so as to be totally immersed in the test solution. Prior to the test, the coupons were degreased and weighed. After the test, the coupons were cleaned with inhibited 15% HCl to a constant weight and the weight loss was reported. Coupon weight loss was converted to a corrosion rate in micrometers per year (μm/y), using the surface area of the coupon (1” x 1/2” x 1/8”, 3.375 in2 or 21.77 cm2), test duration, and a density of 7.86 g/cc for the steel.

Results As discussed in the Introduction, a critical test of the undissociated acid hypothesis is to demonstrate that the observed corrosion rate is proportional to the undissociated acid concentration independently of the pH at which the observation is made. Liu et al.(2) reported that with formate HSAS corrosion increased rapidly as the pH of the amine solution fell below about 9.4. Acetic acid is convenient for this test, both because its pKa is one unit higher than that of formic acid and because it is commonly found in refinery treaters. Other commonly found anions were also included in this study, both to provide further examples and counterexamples for the hypothesis as well as generating data useful for operating refinery amine treaters. The corrosion test results are given in Table 3. Interference With LPR Measurement Experiment 31 was performed in order to determine whether systems containing H2S, which are of interest to refinery applications of amine scrubbing, could be studied by LPR. The test liquid was a nominal 50% MDEA solution with an initial 14% formate HSAS neutralized to 0.85% remaining HSAS and with an H2S loading of 0.04 moles/mole of MDEA.


Results Cont. Liu et al.(2) have shown this system to be non-corroding by coupon studies. The LPR curve is shown in Figure 1. The noisy and non-monotonic indicated corrosion rate is quite different from runs in other systems, with the possible exception of Experiment 41 with thiocyanate HSAS (Figure 2). The average corrosion rate indicated in Figure 1 by LPR is about 3600 μm/y (140 mpy) over the 60-hour test. The coupon weight losses were 3.8 and 6.1 mg, for an average of 5 mg. This translates to a corrosion rate of 43 μm/y (1.7 mpy). The excess electrochemical reaction recorded by the LPR instrument is most likely the oxidation of sulfur species in the -2 oxidation state(3). Pourbaix diagrams generated with the HSC computer program(11) indicate that at 120°C the thermodynamic equilibrium potential for iron-magnetite is about -0.85 V while that for HS-/S2O3= is about -0.50 V, both at a pH of 9.0. Since the interference in LPR measurements by H2S completely precludes obtaining valid data, sour system studies must be performed with other experimental methods. Also, in order to avoid confounding the results, no CO2 was used in the test fluids of the present study. Since lean amine in general has low concentrations of these acid gases, excluding them from the experimental program is not expected to reduce the validity of the conclusions. A less serious but still significant interference was also noted with thiocyanate HSAS. Experiment 41. Again, a very noisy LPR trace was obtained (Figure 2) which at the end of 90 hours indicated a corrosion rate of 330 μm/y (13mpy). Both corrosion coupons had weight losses of less than 1 mg, corresponding to a corrosion rate of <10 μm/y, again indicating an interfering reaction. Thus, care must be taken when studying corrosion by the LPR method to show that no interference is present. Formate(3) and acetate anions do not seem to interfere, since low corrosion rates are obtained upon neutralization. Acetic acid is known to be oxidatively stable. A general method of preparing carboxylic acids is the oxidation of the corresponding alcohol with various strong oxidizing agents such as potassium permanganate, hot concentrated nitric acid or chromic acid(12). Time to Reach Steady State Attainment of steady state corrosion rate values in the LPR experiments was slow, particularly in the case of corrosive media. This is reasonable in view of the fact that in noncorrosive systems the rates of the chemical reactions involved are low to begin with, so that little change over time is possible on an absolute scale. Corrosive media, on the other hand, involve rapid reaction and chemically aggressive species which can hinder the formation of a passivating deposit on the metal. Figure 3 shows the LPR curve for Experiment 42 with 5.1% thiosulfate HSAS in a nominal 50 wt % MDEA solution. A rise in corrosion rate from time zero to about 2 hours, during which time the autoclave is heating up from room temperature to the test temperature of 120°C, was observed in all runs. However, within about 3 hours, the measured corrosion rate in Experiment 42 dropped to about 50 μm/y (2 mpy), compared to 30 μm/y at 60 hours. In strong contrast, Experiment 34 with a solution containing 14.8% acetate HSAS exhibited a slow, steady decrease in corrosion rate from about 8000 μm/y (320 mpy) to 5700 μm/y over a period of 135 hours, still with no attainment of steady state (Figure 4). Since steady state was not reached, we cannot say with certainty whether this medium is corrosive or not, under the test conditions.


Results cont. In a real refinery amine treater, however, if any conditions, such as erosion or the presence of chelating species, exist that tend to remove the passivating layer, then high rates of corrosion would likely occur. The LPR curve of the same 14.8% acetate HSAS solution, neutralized to 4.7% HSAS remaining (Experiment 37), the balance of the acetate being present as an alkali metal salt, is given in Figures 5 and 6. After a high initial corrosion rate maximum of 6600 μm/y (260 mpy) during the heatup period, the corrosion rate drops to less than 300 μm/y within 2 hours. At the end of 42 hours, the measured corrosion rate is about 25 μm/y and still decreasing. Comparing Figures 4 and 5, it seems reasonable to conclude that at the lower HSAS level the corrosivity of the medium is lower, since passivation is more rapid and effective. Oxalate HSAS containing MDEA solutions exhibit LPR curves qualitatively similar to those with acetate HSAS except that the absolute corrosion rates are higher and attainment of passivation (in the neutralized solution) is slower. These LPR curves are shown in Figure 7 to 9. Due to the limited solubility of the oxalate alkali metal salt, precipitation occurred upon neutralization. In order to do a worst case experiment, the precipitated solid was left in the test solution. It is not known whether the solids dissolved at the 120°C test temperature. If not, then mild erosion due to the solids would produce an increased corrosion rate. Nevertheless, the observed LPR curves make it reasonable to assume that the unneutralized solution is more corrosive than the neutralized one. LPR Versus Coupons The LPR Experiment 40 with oxalate HSAS neutralized to 1.2% was run with corrosion coupons immersed in the liquid. The weight losses were 161 and 164 mg, which converts to an average corrosion rate of 740 μm/y (29 mpy) over the 112-hour test. This is in semi-quantitative agreement with the value of 2100 μm/y obtained by “integrating” the LPR corrosion rate curve by summing time segments of the curve. As the Tafel constants utilized by the Gamry instrument are not specific for the present solutions, quantitative agreement would be fortuitous. Also, the average LPR-measured corrosion rate is heavily weighted by the very high rate during the initial 15 hours of the run. The system during this period is far from steady state. Therefore, as pointed out by Asperger et al. (3), the corrosion rate values measured cannot validly be related to an actual corrosion rate. It is interesting to note that short duration coupon tests can overestimate the corrosion rate, due to rapid initial corrosion prior to passivation. Experiment 42, which had a corrosion rate indication of about 30 μm/y (1.25 mpy) by LPR, had coupon weight losses of less than 1 mg, corresponding to a corrosion rate of < 10 μm/y. This tends to indicate that LPR is not susceptible to giving false negative results. Correlation of Corrosion Rate to pH Figure 10 is a plot of LPR-measured corrosion rate versus room temperature pH of the test solution. While a given anion shows an increase in corrosivity as the pH of the medium becomes more acidic, the correlation of corrosion rate with pH alone is poor when all the anions are considered. This observation weakens the hypothesis(2) that a passivating layer, whose protectiveness is a function of pH, is responsible for the variation in corrosion rate. Correlation of Corrosion Rate to Undissociated Acid The concentration of undissociated acid in the various test solutions shown in Table 3 were calculated using Equation 10. Figure 11 shows the plot of corrosion rate versus the acid concentration in molal units.


Results cont. Allowing for experimental error, a good correlation is evident: the corrosion rate increases with increasing acid concentration and anions which provide low acid levels (thiocyanate and thiosulfate) have low corrosivities. Not unexpectedly, the slopes of the plots for acetate and oxalate are different. While the pKa of acetic acid is virtually identical to the pKa2 of oxalic acid, oxalate co-ordinates much more strongly with FeII than does acetate(13). Thus, in the case of oxalic acid, the corrosivity due to the acid concentration may be co-operating with its chelating ability to produce a high total corrosiveness. A possible mechanism for the corrosiveness of the acids is that they are functioning as very efficient cathodic depolarizers. In the absence of oxygen, the likely cathodic reaction is (14)

H+ + e- ⇒ H• (11) At 120 °C, the concentration of H+ is of the order of 10-8 to 10-9 molar in the amine solutions of this study. The boundary layer adjacent to the corroding surface further decreases the H+ concentration available for the cathodic reaction, thus limiting the overall corrosion rate. Asperger et al.(13) demonstrated that the corrosion rate is indeed diffusion rate controlled. Systems which generate significant concentrations of undissociated acid according to Equation 10 have an alternative source of protons for the cathodic reaction from the undissociated acid: HAcid + e- ⇒ Acid- + H• (12) The ease of removing the proton is of course related to the strength of the acid, which would make it reasonable to expect that the stronger the acid, the greater the slope of the corrosion rate vs. acid concentration plot. A further consideration which can explain the effectiveness of even small concentrations of undissociated acids producing large increases in corrosion rate is the fact that as acid is consumed by Reaction 12, it is immediately regenerated by reestablishment of equilibrium: Acid- + H2O ⇔ HAcid + OH- (13) Proton transfer reactions of this nature are usually considered to be instantaneous. The generation of hydroxide ions of course tends to depress the undissociated acid concentration but outward diffusion through the boundary layer into the bulk solution will limit the extent of this effect. The above concept of cathodic depolarization by undissociated acid can also be used to explain why neutralization of the HSAS with a strong base is effective in decreasing corrosion rate by two orders of magnitude or more. Equation 8 illustrates the fact that the strong base OH- substitutes for the weaker base MDEA in the process of neutralization. The pKa of OH- is about 14, i.e., the hydroxide ion is more than 5 orders of magnitude stronger than MDEA. According to Equation 10, the corresponding undissociated acid concentration is then 5 decades less. Cathodic polarization then again becomes the rate limiting factor. Even in the case of oxalic acid where the chelating ability of the oxalate anion would be expected to speed up the anodic reaction Fe + C2O4= ⇒ FeC2O4 + 2e- (14) neutralization brings the overall corrosion rate under cathodic control due to lack of undissociated acid.


Conclusions Liquid phase corrosion in lean MDEA solutions, in the absence of HSAS, is slow and likely under cathodic control due to the low concentration of hydrogen ions available for the reduction reaction to hydrogen. HSAS formed from medium-strength acids, such as formic, acetic, and oxalic, increase corrosion rates by two orders of magnitude or more by providing an alternate source of protons. These HSAS are in equilibrium with unprotonated MDEA and the undissociated acid corresponding to the anion. The undissociated acid seems to function as a very effective cathodic depolarizer, explaining the observed increase in corrosion rate with these HSAS. Neutralization is an effective means of reducing the corrosion rate in these cases by reducing the concentration of undissociated acid by several orders of magnitude. HSAS produced from strong acids, such as thiocyanic and thiosulfuric acids, were found to be non-corrosive, which is consistent with the undissociated acid hypothesis since they do not produce significant quantities of undissociated acid by equilibraiton. Even the oxalate HSAS corrosion rate is greatly reduced by neutralization, in spite of the oxalate anion’s ability to strongly chelate with Fe(II), probably due again to cathodic polarization. Caution must be exercised in the use of LPR in corrosion studies. Hydrogen sulfide seriously interferes with measuring the corrosion reaction rate by providing an alternative electrochemical reaction. While LPR is an excellent method for studying the course of corrosion reactions, prudent experimentation requires demonstrating the absence of interferences.

Acknowledgments The authors thank M.G. Pavelchak for performing the experimental work and Dr. R.G. Asperger for helpful discussions

References 1. R.B. Nielsen, K.R. Lewis, J.G. McCullough and D.A. Hansen, NACE Corrosion/95. Orlando, FL.

2. H.J. Liu, J.W. Dean and S.F. Bosen, NACE Corrosion/95, Orlando, FL. 3. R.G. Asperger, H.J. Liu and J.W. Dean, AlChE Spring 1995 Meeting, Houston, TX. 4. D. Burns and R.A. Gregory, 1995 Laurance Reid Gas Conditioning Conference,

University of Oklahoma, Norman, OK. 5. M.S. Dupart, T.R. Bacon and D.J. Edwards, Hydrocarbon Processing, May (1993): p.

89. 6. A.E. Keller, R.M. Kammiller, F.C. Veatch, A.L. Cummings and J.C. Thompsen, 1992.

Laurance Reid Gas Conditioning Conference, University of Oklahoma, Norman, OK. 7. R.F. Smith and A.M. Younger, Hydrocarbon Processing, July (1972): p. 98. 8. K.F. Butwell, D.J. Kubek and P.W. Sigmund, Hydrocarbon Processing, March (1982):

p.108. 9. H.J. Liu and R.A. Gregory, NACE Corrosion/94, Baltimore, MD. 10. Gamry Instruments Inc., Willow Grove, PA. 11. HSC Chemistry for Windows, Version 2.0, Outokumpu Research Oy. 12. R.J. Fessenden and J.S. Fessende 1, “Organic Chemistry”, 4th edition, pg. 285,

Brooks/Cole Publishing Company, Pacific Grove, CA., 1990. 13. L.G. Sillen and A.E. Martell, “Stability Constants of Metal-Ion Complexes”, The

Chemical Society, London, Special Publication No. 17, 1964. 14. J.C. Scully, “The Fundamentals of Corrosion”, 3rd Edition, p.55, Pergamon Press, 1990. 15. G. Kortum, W. Vogel and W. Andrussow, “Dissociation Constants of Organic Acids in

Aqueous Solution”, Butterworths, London, 1961. 16. D.M. Austgen, Ph.D


Table 1 PKa Values

pKa Species 25 °C 120 °C* Reference MDEA 8.52 7.03 16

H2S 7.04 6.46 16 CO2 6.37 6.50 16

Formic Acid 3.75 4.02 15 Acetic Acid 4.75 4.99 15 Oxalic Acid 4.27 5.00 15

Glycolic Acid 3.83 15 Thiosulfuric Acid 1.56 13 Thiocyanic Acid 0.85 13 Sulfurous Acid 1.82 13

Sodium Hydroxide 14.0 13 *Obtained by graphical extrapolation of referenced data Table 2. Sources of Refinery Anions

Anion Source Acetate Feed gas from FCCU and coker, LPG Formate Feed gas from FCCU and coker, LPG Oxalate Oxygen oxidation of amine Glycolate Oxygen oxidation of amine Thiocyanate Reaction of cyanide with sulfur compounds Cyanide Feed gas from FCCU and coker Thiosulfate Oxidation of H2S Sulfate Oxidation of sulfur compounds Chloride Make-up water

Table 3 Corrosion Test Results

Rate(1) [HAcid] Time(2)

Run # Anion % HSAS

pH LPR, μm/y

Coupon, μm/y

Molal x 104

Hours

31 Formic(3) 0.85 10.54 2030(4) 43 0.4 60 33 Formic 11.68 9.31 1320 3.1 138 34 Acetic 14.78 9.38 5715(4) 52.7 135 35 Acetic 5.27 10.06 1524(4) 14.7 44 37 Acetic 4.71 10.25 25(4) 12.9 42 36 Acetic 2.60 10.37 14 6.8 120 38 Oxalate 8.94 9.72 7620(4) 7.5 44 39 Oxalate 4.53 9.89 1020(4) 3.7 92 40 Oxalate 1.16 10.48 130 740 0.9 112 41 SCN- 4.84 9.93 330(4) <10 10-4 90 42 S2O3= 5.05 9.67 32 <10 10-4 60

(1) Indicated rate at end of test period (2) From start of heating (3) With 0.4 moles H2S/mole of amine (4) Steady state not attained


Figure 1 Experiment 31



Figure 9 Experiment 40 (cont’d)

Figure 10 Corrosion VS pH


Figure 11 Corrosion Rate VS [HAcid]


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