Environ. Sci. Techno/. 1995,29, 2936-2945
Anaerobic Corrosion of Granular Iron: Measurement and Interpretation of Hydrogen Evolution Rates ERIC J. REARDON Department of Earth Sciences, University of Waterloo, Waterloo, Ontario, Canada N2L 3G1
corrodes by the oxidative action of water itselk Fe(s)
+ 2H20(1)- Fez++ 20H- + H,(g)
In this reaction, 1 mol of hydrogen gas is generated for every mole of iron corroded. The accompanying production of Fez+and OH- ions is of considerable importance to the geochemical design of a potential groundwater treatment system. Corrosion-produced OH- ions will increase pH and react with dissolved carbonic acid and bicarbonate species in the groundwater to produce carbonate ions:
+ 20H- - C0,'- + 2H,0(1) HC0,- + OH- - C0:- + H,O(1)
H,CO,'
Anaerobic corrosion of iron metal produces Fez+, OH-, and H2(g). Growing interest in the use of granular iron in groundwater remediation demands accurate corrosion rates to assess impacts on groundwater chemical composition. In this study, corrosion rates are measured by monitoring the hydrogen pressure increase in sealed cells containing iron granules and water. The principal interference is hydrogen entry and entrapment by the iron. The entry rate is described by Sievert's law ( R = k P ~ , 0 . ~and ) , the rate constant, k, is evaluated by reducing the cell pressure once during a test. For the 10-32 mesh iron used in this study, k initially was 0.015 but decreased to 0.009 mmol kg-' d-l kPa-o.5 in 150 d. The corrosion rate in a saline groundwater was 0.7 f.0.05 mmol of Fe kg-' d-' a t 25 "C-identical under water-saturated or fully-drained conditions. The rate decreased by 50% in 150 d due to alteration product buildup. The first 40-200 h of a corrosion test are characterized by progressively increasing rates of pressure increase. The time before steady-state rates develop depends on the solution composition. Data from this period should be discarded in calculating corrosion rates. Tests on pure sodium salt solutions a t identical equivalent concentrations (0.02 equiv/L) show the following anion effect on corrosion rate: HC03 > S042- > CI-. For NaCl solutions, corrosion rates decrease from 0.02 to 3.0 m.
Introduction Interest is growingin the use of granulated iron to remediate groundwater contaminated with halogenated hydrocarbons, such as trichloromethane (TCM),tetrachloroethene (PCE), and trichloroethene (TCE) (1-6). Although the chemical mechanisms are still a subject of research, granulated iron promotes rapid dechlorination of these hydrocarbons in groundwater. Any future applications of this technology to groundwater remediation would be designed to operate under anaerobic conditions so as to prevent or inhibit corrosion of the iron by oxygen. Nevertheless, even in the absence of oxygen, iron still *
F a : (519) 746-0183.
2936 ENVIRONMENTAL SCIENCE &TECHNOLOGY
/ VOL. 29, NO. 12, 1995
(1)
(2)
(3)
Buildup of carbonate ions eventually will result in the precipitation of carbonate solid phases. Depending on the chemical composition of the groundwater, these may be calcites,siderites, or magnesium hydroxycarbonates. Since considerable ion substitution occurs in the CaC03-FeC03MgC03 system, the solids that form will likely be complex carbonate solid solutions rather than pure phases (7, 8). If these solids precipitate as surface coatings, they may reduce the effectiveness of iron to degrade halogenated hydrocarbons and reduce flow rates by porosity and permeability reductions. If groundwater carbonate is exhausted through precipitation of carbonate minerals, continued corrosion of iron leads to saturation and precipitation of Fe(OHI2,itself a relativelyinsoluble solid phase. In the absence of oxygen and at low temperatures, Fe(OH), is stable but is predicted thermodynamically to convert to magnetite:
The disproportionation of Fe(OHIzis known as the Schikorr reaction and occurs at temperatures =- 373 K (9). At lower temperatures, it proceeds slowly because of kinetic restraints, and Fe(OHI2may be a metastable phase in lowtemperature anaerobic groundwater environments. However, there are known inorganic and organic compounds that can catalyze this mineralogic transformation (10, 11). None of these identified catalysts, however, are major constituents of groundwater. In this study, measurements of hydrogen production in sealed cells containing iron and water are interpreted as a measure of corrosion, and so the possible conversion of Fe(OH)2to Fe304contributes uncertainty to the calculated corrosion rates. If the conversion does not occur, there will be 1 mol of Hz(g) produced for every mole of iron corroded. However, if all of the produced Fe(OH)*converts to FesO4, there will be an additional 0.33 mol of Hz(g) produced for every mole of iron corroded. This represents a sizable, but for many purposes acceptable, uncertainty in relating hydrogen evolution rates to iron corrosion rates. A knowledge of iron corrosion rates under anaerobic conditions is necessary to predict the flux of OH- and Fezions to the aqueous phase over time. With this information, geochemical models can be used to predict the impact of corrosion on the chemical composition of ground-water flowing through granular iron material as well as the
0013-936X/95/0929-2936$09.00/0
C 1995 American Chemical Society
quantity of carbonate precipitate that is likely to be produced. The purpose of this paper is to present the methodologyto monitor hydrogen evolution from granular iron material under strictly oxygen-free conditions; to examine the processes that may interfere with deriving corrosion rates from hydrogen evolution measurements; to present corrosion rate results for several common salt solutions; and to show results of a long-term corrosion experiment on a saline groundwater, which indicate the impact of reaction product buildup on corrosion rates.
Experimental Section Corrosion Cell Apparatus. Four corrosion test cells following a design by Dewaele et al. (12) were assembled for this study. Each cell consisted of a 250-cm3 stainless steel canister with a bolt-down cap outfitted with a Vitor O-ring. Avalve was installed in the base of the canister for injection and removal of solution during the experiments. A network of tubing and Swage-Lok fittings led from the canister cap to another valve and an Omega PX-302-015 absolute vacuum pressure transducer. The upper valve was used for sample evacuation or reducing hydrogen gas pressure produced from the corrosion reaction. Stainless steel was used for all material: canister, fittings, valves, and tubing. Analytical-grade, oxygen-free nitrogen was used in all gas flushing operations. The pressure transducers used in this study were energized using a 9.920 & 0.002Vpower supply and yielded approximatelya 100-mvresponsefor a 0- 100kPa pressure change. The stated accuracyofthe transducers was f0.25% BFSL (best fit straight line) over this pressure range. The transducers were individually calibrated by measuring their millivolt response at zero vacuum and atmosphericpressure and assuming linearity over the operating pressure range, which was from 0 to 140kPa. Millivolt readingswere logged at 30-s intervals using a 16-bit high-resolution data acquisition board and a 386 PC. The millivolt readings of the acquisition board were calibrated using two different multimeters. Prolonged exposure of the PX-302-015 pressure transducers to a moist hydrogen atmosphere usually resulted in a progressive decrease in their sensitivity and sometimes produced spurious responses. It is believed this was due to embrittlement of the stainless steel membranes of the transducers as a result of hydrogen contamination. Placement of the affected transducers under vacuum for several days usually restored their original response characteristics, but some transducers were permanently damaged after extended use. Solid and Solution Samples. The granular iron used in the corrosion tests was Master Builders Inc. (Cleveland, OH) Blend A, a 50/50 mixture of 10-18 and 18-32 mesh size iron granules (C, 3.2%;Mn, 0.65%; S, 0.09%; surface area, 1.5 m2g-'). The material is produced for use in concretes by rotary kiln processing of iron borings obtained from the iron and steel borings market and is being examined in several laboratories across North America for its effectiveness in dehalogenation of hydrocarbons. A variety of solutions were used in the corrosion tests: IR-gw, a saline groundwater from Inrine, (2%IR-sim, a laboratory-prepared solution designed to simulate the major chemical composition of IR-gw; deionized water; and pure salt solutions of NaC1, NaHC03, and Na2S04. Triplicate runs using different solid/water/gas phase combinations were made on one solution (0.02 m NaC1) to
TABLE 1
Composition of Water Samples Used in Corrosion Tests water constituent PH
Ca
Mg Na
K CI so4
alkalinity as HCO3
IR-gw (mgR)
8.0 200 1500 11 000
18 9 000 26 000 730
IR-sim (mg/L) 7.5
2008 1 485 10 900 19 8 900 16 500 600a
Actual Ca2+and HCO3- concentrations may be less than indicated. The target values were generated by adding 0.5 g of CaC03 to 1 L of solution and bubbling with 5% C02(g) after all other salts had been added. The solution, however, remained slightly turbid after this procedure, indicating that not all the CaC03 dissolved. a
ensure reproducibility of the determined rate measurements. The chemical compositions of the actual and simulated groundwater are recorded in Table 1. Solutionssampled for iron analyseswere extracted from the reaction cell by syringe, immediately filtered through a 0.22-pm filter, and then acidified with 1:1 HN03. Iron analyses were made on a Varian atomic absorption spectrophotometer. Uncertainties are on the order of 5%,and the detection limit was 0.05 mg L-l. PH measurements were made immediately on unfiltered aliquots. Iron oxidation occurs quickly at high pH, and a stabilized reading will progressively drift downward. If this effect was observed before a clear stabilization, fresh solution was extracted and the measurement was repeated. Corrosionrate measurements on specimens of iron and steel are usually reported in units of thickness of Fe per unit time (such as mm/yr or nm/d). This approach normalizes the rate measurement for the surface area of the specimen tested and facilitates direct comparisons with the results of other studies. In this paper, rates are reported in units of mmol of Fe kg-' d-], uncorrected for surface area. This is done because the high measured surface area of the iron material used in this study (1.5m2g-l) is usually associated with material much finer than 10-32 mesh. Although SEM evidence reveals considerable intragranular porosity, which may account for this U. Warren, personal communication),a question remains whether the measured surface area overestimatesthat of the iron granules because of the large amount of intimately-admixed, fine-grained carbonaceous material (3.2mass %). Because this question has not been resolved, rates in this study are reported with respect to mass rather than surface area. Test Procedure. In the corrosiontests, weighed samples of the iron were placed in the test cell and evacuated using a Vac Torr Model 25 pump, with a 0.01 Pa vacuum rating. The void volume of the cell was measured, usually several times, by injecting nitrogen gas and recording the volume required to return the internal pressure to atmospheric conditions. The accuracyof this volume measurement was within &2%. A test was initiated by first placing the cell and its contents under vacuum for several hours. Next, a known volume of a test solution was injected through the bottom valve of the corrosion cell. The cell was then placed in a constant temperature bath maintained at 25 "C, and the data acquisition system was activated. The headspace volume of the cell was calculated as the difference between VOL. 29. NO. 12, 1995 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
2937
the total void volume and the volume of solution injected. Subsequent tests on the same sample were performed by draining and flushing the cell with fresh solution using the bottom valve. The rate of hydrogen production, Le., corrosion rate, measured for the iron material in this study ranged from 0.1 to 0.7 mmol kg-I d-I depending on the solution composition. Exhaustive tests were conducted on the corrosion cell, without iron or solution present, to ensure that its leak rate, under both vacuum and positive pressure conditions, was unimportant compared to the hydrogen generation rate from corrosion. The positive pressure leak tests were conducted using both nitrogen and hydrogen gases. The measured leak rates in these tests were less than 0.001 mmol d-I or less than 1%of the lowest measured rate of hydrogen production from corrosion. Solutions were first placed under vacuum for less than 1 min to remove dissolved atmospheric gases before injection into the corrosion cell. Because of the variable quantities of dissolved gases that may be present in natural waters and because gas may be produced through bacterial activity, a blank test must be run where the water sample is injected into an evacuated corrosion cell, without iron present, and the pressure monitored with time. This test simultaneously evaluates the possibility of hydrogen gas generation as a result of the corrosion of the stainless steel cell itself. Corrosion Rate Calculations. Hydrogen gas produced as a result of anaerobic corrosion at the surfaces of iron granules can enter three phases: the aqueous phase, where it is limited by its solubility (7.515 x mol kg-l for pure water at 100 kPa HZat 25 "C) (13); the gas phase, where it is monitored with the pressure transducer; and the solid phase. The possible diffusion and entrapment of hydrogen into the solid phase, Le., within the iron granules themselves, is not considered in this section but will be evaluated in the discussion of the Long-Term Corrosion Experiment. The measured pressure (Pmea,) in the cell at each point during a corrosion test is interpreted as the sum of the water vapor pressure (PH~o) and the hydrogen pressure ( P H ~ ) . For pure water at 25 "C, P H ~ isO 3.167 kPa (13). Therefore, the hydrogen partial pressure in the cell is
no effect on the corrosion rate calculated from the pressure changes occurring in the cell over time; ( 2 ) the hydrogen dissolved in the water phase applying Henry's law without salt correction is small compared to that in the gas phase (always less than 5 mass W),and the reduction in Hz(g) solubilitywith salt concentration is modest, dropping to at most 3 or 4 mass % at concentrations of 1.0 m for a variety of electrolytes(14). Thus, ignoring the salt effect contributes an uncertainty in the calculated corrosion rates of less than 2%. The total moles of hydrogen in the cell (gas plus aqueous phase) for each recorded pressure reading is then
The corrosion rate (expressed in units of mmol of Fe kg-' d-l) can then be calculated over a selected time interval (tz - t l ) as
where M F is ~ the mass of the iron in kilograms. Usually, corrosion rates were calculated from the recorded pressure measurements over intervals of approximately 1 h.
Results and Discussion
Interference Tests. Two tests were performed with saline groundwater (IR-gw)but without iron present to evaluate the possible generation ofhydrogen gas due to the corrosion of the stainless steel reaction cell. These tests simultaneously evaluated the possible interference of the degassing of volatiles from the water sample. A third test was conducted to determine if hydrogen gas generated from the corrosion reaction would react chemically or biochemicallywith dissolved constituents in the water sample during the test period or diffuse into the stainless steel reaction vessel. In the first test, 90 g of IR-gw was placed in an evacuated stainless steel reaction cell, and the 28-cm3headspace was flushed with nitrogen for 10 min. The cell was then sealed with the nitrogen headspace at atmospheric pressure, and the internal pressure was monitored for 92 h. There was PH,= Pmea, - 3.167 (5) no pressure increase over this period, only a slight pressure decrease of 0.7 kPa, corresponding to a loss of 0.2 cm? of and the moles of hydrogen in the gas phase of the cell is nitrogen from the 28-cm3headspace. In the second test, 90 g of IR-gw was injected directly into an empty evacuated cell. The pressure increased by 0.5 kPa over the first 10 h, corresponding to 0.14 cm3of gas exsolution, and there was no detectable change over the next 35 h. where vis is the headspace volume in cm3, Tis temperature After completion of these tests, the nitrogen headspace in degrees K, and P H is ~ the pressure in Wa. Assuming of the cell used in the first test was removed by repeatedly equilibrium between the gas and aqueous phase, the moles flushing with hydrogen gas, and then the cell was sealed. of hydrogen dissolved in the aqueous phase is This initiated a test for the reaction of hydrogen with the saline groundwater and for the possible loss of hydrogen H,(aq) = 7.515 x PH,MHz0 (7) by diffusion into the stainless steel reaction cell. The cell was periodically shaken to facilitate the contact of the where M H ~isOthe mass of the test solution in kilograms. hydrogen gas with the water phase. The total pressure Salt concentration corrections on both the equilibrium P H ~ O decrease over 85 h was 0.8 @a, corresponding to a loss of in the gas phase and on hydrogen solubility in the water 0.22 cm3 of hydrogen from the gas phase. phase were not made in this study. These are reasonable In comparison with the results of these three interference approximations because (1) the equilibrium P H ~ in O the tests, 24.4 cm3of hydrogen is generated in atypical corrosion cell [lo% lower than above pure water for the highest test over an 80-h period (assuming a 0.5-kg sample with a concentration solution used in this study (3.0 m NaCl)] corrosion rate of 0.6 mmol kg-' d-l). Thus, all these would not change during a test and, therefore, would have potential sources of error, such as corrosion of the stainless 2938 ENVIRONMENTAL SCIENCE &TECHNOLOGY / VOL. 29, NO. 12, 1995
120
c
80 -
a Y
0
6 0 -
30 -
0
I
1
97
06
95
3
a
1 0.-
2Y
I
1
1
I
10
:I
lngasalng Detalls 9 -
l(7ZZ==Degasslng Detalls
I
steel cell or hydrogen entry into the cell contribute negligible uncertainty to the corrosion rate measurements. The only effect of concern is when the headspace pressure is reduced by evacuation at the beginning or during a test. This would cause dissolved hydrogen or other gases to exsolve from the water phase. Even though the pressure changes associated with this gaslwater solubility readjustment will generally be small, they occur over short periods and contribute to apparent enhancements in the calculated corrosion rate immediately following a pressure decrease. An experiment was conducted to determine the time dependency of water phase ingassing and degassing following pressure changes to a drained sample of iron granules wetted with deionized water. Under drained conditions, a continuous gas phase exists throughout the material. The experiment was conducted within 3 h of adding deionized water to dry iron granules. During this time, pressure increases from hydrogen entry into the gas phase due to corrosion would not be important. The iron granules were first placed in a cell, and its total mass was determined. The cell was placed under vacuum for 2 h and then flushed three times with deionized water. In each flushing, the water entered the cell under vacuum and drained under nitrogen atmosphere. After the final draining, the cell was evacuated, its mass was redetermined (to establish the net water content by difference with the dry mass), and then the cell was placed back in the 25 "C temperature bath. The data acquisition system was activated to log cell pressures at 5-s intervals. After 10 min, the internal pressure was stable at 3.1 H a , corresponding to the establishment of waterhapor equilibria at this temperature. Nitrogen gas was then injected to ambient atmospheric pressure conditions (96.6 kPa), and the cell
8 .
1
I
was sealed to initiate monitoring of the ingassing of nitrogen to the water phase. After 1.3 h, the cell was evacuated to 8.4 kPa and sealed again to initiate monitoring of nitrogen degassing from the water phase. The results of this experiment are presented with both coarse and expanded pressure scales in Figure 1. From the solubility of nitrogen in water and the measured amounts of iron granules, liquid water, and void space, the predicted pressure changes associatedwith the nitrogen ingassing and degassing(curve segments 2 and 3, respectively) were calculated and are shown as vertical bars in the diagram. Figure 1 shows that degassing occurs more rapidlythan ingassingbut that both will be 90%completewithin 1h following apressure change. Although the results of this experiment pertain to nitrogen gas, gadwater equilibration times for hydrogen should be even shorter because of the greater diffusivity of hydrogen. For water-saturated samples of iron granules, the period of solubility readjustment will be longer, perhaps 10 h or more, as indicated by the second of the above interference tests. Short-TermCorrosion hperiments. Three short-term corrosion tests were conducted on Master Builders Blend A iron filings. A 544.4-g sample of the material was placed in a cell, sealed, and evacuated. The void space (144 i. 2 mL) was determined by recording the volume of nitrogen required to return the internal pressure to the atmospheric value. The cell and its contents then were placed under vacuum for several hours. IR-gwgroundwater (100 mL) was next introduced through a septum and valve arrangement at the bottom of the reaction cell. The pressure was then monitored for 3 d; however, the first several hours were not logged. This completed run A l , whereupon much of the solution phase was expelled through the bottom valve VOL. 29, NO. 12, 1995 / ENVIRONMENTAL SCIENCE &TECHNOLOGY
2939
i
1.2
0 ' 0
I 20
40
60
Time (h) FIGURE 2. Corrosion rate versus time for three sequential tests (Al,
A2,and A3) on a 544.4-8 iron sample in contact with 100 mL of IR-gw groundwater (see Table 1). Void space during runs was 44 cm3, Individual points represent rates determined on differences in cell pressure at 50-min intervals.
and septum by injecting nitrogen gas into the headspace through another valve at the top of the cell. The cell was then evacuated, and a known volume of fresh groundwater solution was introduced through the bottom valve. This flushingprocedure was repeated three times, which reduced the original water content to less than 20% of the final. The headspace was then reevacuated and sealed, which commenced run B2. After 3 d of pressure monitoring, the cell was flushed in a similar manner to start run A3. Net corrosion rates were calculated over each 50-min interval for all three runs from the pressure data, knowing the quantity of iron and the volumes of solution and headspace present in each run. Henry's law corrections were made for the solubility of hydrogen gas in the solution phase as outlined in the Experimental Section. The results are presented in Figure 2 and reveal sizable differences in the calculated corrosion rates for all three runs at early time. After about 40 h of reaction, all rates are in good agreement, averaging 0.6 i 0.05 mmol of Fe kg-' d-l. The reasons for the variability in calculated corrosion rates at early times is not completely understood. However, it depends on whether the sample is initially dry or wet. For initially dry samples, e.g., run A l , a lag time always occurs before a steady-state rate of hydrogen entry into the gas phase is attained. One explanation is that rapid injection of the solution into a dry iron sample could result in a substantial number of closed-off evacuated pores. Over time, these pores would gradually fill with water, which would contribute to an overall decrease in headspace pressure and counteract the pressure increase due to hydrogen production. A second explanation is that some surface-held oxygen is not removed by the sample evacuation step and undergoes cathodic reaction with earlyformed H2to form water. Lag times for hydrogen evolution from iron metal immersed in water have been observed by Seo et al. (15) and attributed to this mechanism. A third explanation for the time lag in the attainment of steadystate hydrogen production to the gas phase is that produced hydrogen, upon initial contact of iron and water, would tend to accumulate on the surface of iron granules until any sorption capacity for hydrogen gas molecules is filled. In addition, accumulating bubbles of hydrogen would remain adhered to the surface of iron granules until their sizes were large enough (Le.,internal pressures and densities low enough) to overcome surface tensional forces to enable 2940 rn ENViRONMENTAL SCIENCE & TECHNOLOGY / VOL. 29, NO. 12.1995
their release through the water phase. Likely, all three effects contribute to the low apparent hydrogen production at early time for initially dry samples, and tests results on pure salt solutions, presented later in this paper, reveal that this nascent period can range from 40 to 200 h and is dependent on both solution composition and concentration. In runs on previously reacted material, where the solution was drained and refilled with fresh solution (runs A2 and A3 in Figure 21, low rates of pressure increase at early time are not observed. Instead, anomalously highpressure changes are recorded at early time. This may reflect corrosion rate changes as a result of chemical changes to the water or it may reflect the effect of minor sample degassing. (Note that the less than 1-h degassing equilibration time does not apply in these tests because the iron granules were water-saturated.) Determining the possible causes for early low rates of hydrogen entry into the gas phase for initially dry material and high rates for previously wetted material will require considerable experimentation. However, for the purposes of applying the hydrogen evolution monitoring technique to determine corrosion rates for water in long contact with iron granules, data collected during this early period of reaction should be discarded in determining corrosion rates. Water Saturated vs Unsaturated Corrosion Test Results. A corrosion test can be conducted on granular iron under water-saturated or fully-drained (unsaturated) conditions. An advantage to conducting the test underwatersaturated conditions is that the headspace volume can be kept small. This increases the sensitivity of the corrosion rate measurements because, for a given rate, the pressure change is inversely proportional to the headspace volume. However, there is the associated problem of the produced hydrogen gas migrating upward through the water phase and into the headspace during the test. This could contribute to sporadic pressure fluctuations in the headspace. Conversely, if the test is conducted on a fully-drained sample, the gas phase will be continuous throughout the sample. This avoids the problem of gas bubble migration and facilitates hydrogen transfer to the gas phase because of the increased surface area of the gadwater interface. The disadvantage is that less control over the void space volume is possible. A minimum volume is dictated by the porosity of the granular iron sample in its drained condition. Thus, there may be reduced sensitivity in the corrosion rate measurements compared to performing the test under water-saturated conditions. An experiment was conducted to determine if corrosion rates are similar under saturated and unsaturated water content conditions. In the test, 551.4 g of dry granular iron was added to a cell; the cell was placed under vacuum overnight; and 110 mL of IR-sim, a laboratory-prepared simulation of IR-gw (see Table 11, was injected into the evacuated cell to fully water saturate the iron material. The headspace volume was determined to be 33 cm3. The corrosion rate was monitored for 3 d, and the results are presented in Figure 3 as curve B1. After this experiment was completed, nitrogen was injected into the top of the cell to expel the drainable portion of the solution phase. A total of 42 mL of solution was collected from the bottom of the cell. Thus, the new headspace volume was 75 cm3. The sample was then evacuated and sealed, and the pressure was monitored for
9
"* 0.9
1 c 1
0 ' 0
82
- Unsaturated
I
I
60
40
20
Time (h) FIGURE 3. Corrosion rate versus time for two sequential tests on a 551.4-9 iron sample in contact with IR-sim groundwater (see Table 1). In run B1, the iron granules were initially dry and then water saturated with 110 mL of solution, void space 33 c d . Following completion of this test, 42 mL of solution was drained and the cell was evacuated to initiate run 82.
0.02 mNaHCO3
0.01 mNaZSO4
p0 E E
0.4
aJ
2+
indicate iron and steel are more susceptible to pitting corrosion under aerobic conditions with increasing NaCl concentration. Pitting corrosion, however, is most commonly observed in the presence of strong oxidants such as oxygen and is a form of localized corrosion, which will generally be a small component of the overall or uniform corrosion that occurs over the entire surface of the iron material. Studies of the effect of NaCl on corrosion rates under aerobic corrosion reveal that the overall rate reaches a maximum at 3 mass % NaCl (0.5 m)and systematically decreases at higher concentrations (16). One study was found that examined the effects of NaCl concentration on corrosion rates under strictly anaerobic conditions (17). These authors report results at 80 "C that show the anaerobic corrosion rate of cast steel at 0.23 m NaCl is only one-sixth that at a concentration of 0.023 m NaC1, thus supporting this study's findings. They also present data to show that the corrosion rate decreases below a NaCl concentration of 0.023 m, which indicates there must be a maximum rate near this concentration. Results for the three different sodium salts at identical equivalent concentrations (0.02 equivll), showed the following anion effect on corrosion rate: HC03 > S042- > C1-. The results also indicate substantial variation in the initial 'nascent' period, i.e., the time before steady-state rates of hydrogen entry into the gas phase are attained upon solution addition to initially dry iron material. The solution phase inside the sealed reactions cells could not be conveniently sampled for pH measurement during a test because it would disrupt the monitoring of pressure required to determine the corrosion rate. However, a parallel set of experiments was mounted to enable pH measurements for four of the salt solutions shown in Figure 4 (0.02 NaC1,O.Ol Na2S04,0.02 m NaHC03, and deionized water) and the Irvine, CA, groundwater. Specifically, 135 g of iron granuleswere added to five glass syringes, outfitted with rubber septa at the top and bottom. Atmospheric oxygen was removed by purging each syringe with nitrogen for 1 h. A total of 35 mL of the five solutions was then injected into the syringes, which were then wrapped in several sheets of black plastic to prevent photochemical reactions and totally immersed in a 25 "C water bath. pH measurements were made on 3-mL solution extracts expelled from the bottom septum by injecting nitrogen gas into the top septum. Measurements were made after 72, 120,200,and 400h. On the last sampling, solution samples were taken for Fe analysis. The results are presented in Table 2. Although there appear to be slight systematic changes in pH for some of the solutions, they are generally constant with time and, with the exception of the Irvine saline groundwater, are similarly of high pH ('9.0). For these neutral to alkaline conditions, pH does not appear to be a factor in governing anaerobiccorrosionrates because the solution with the highest pH (0.02 m NaHC03) and the solution with the lowest pH (Irvinegroundwater)both show similar corrosion rates (see Figures 4 and 3). Whitman et al. (18) were the first to show that pH does not directly increase iron corrosion rates until below pH 4. Using HC1 and NaOH to adjust pH, they observed constant rates of iron corrosion over the pH range 4-10 and ascribed this to the stability of the protective oxide coating present on the iron. Solution composition is a more important factor in determining corrosion rates over this pH range, but given the wide range of solution composition and concentration considered in this study, it is of note that the corrosion
0.2
0.0
'
0
,
1
1
I
100
200
300
400
Time (h) FIGURE 4. Corrosion rate versus time for granular iron in contact with various pure salt solutions. Data set labeled ' D W refers to deionized water run.
another 3 d. The results of this run are shown by curve B2 in Figure 3. Aside from the early time differences between the two runs, the reasons which have alreadybeen discussed, Figure 3 shows that differing water content conditions have no effect on the corrosion rate. This is an important finding because it allows an experimenter greater flexibility in designing a corrosiontest for a particular irodwater system. All factors considered, it is recommended to conduct corrosion tests on drained samples, principally because gas solubility readjustments following any pressure change that may be imposed on the cell during the experiment will be mostly complete within 1 h. Effect of Solution Composition on Corrosion Rates. Corrosion tests were conducted using deionizedwater (DW) and several pure salt solutions (0.02, 0.2, 0.8, and 3.0 m NaCl; 0.01 Na2S04;and 0.02 m NaHC03). The results are presented in Figure 4. The time required for the attainment of steady-state hydrogen production in these solutions appears longer, and the corrosion rates appear lower than for the saline groundwater. For the four corrosion tests using NaCl solutions over the concentration range 0.02-3.0 molal, Figure 4 shows that corrosion rates decrease with increasing NaCl concentration. The results for the deionized water run are only slightly lower than the 0.02 m NaCl. These results are seemingly at odds with extensive published data that
VOL. 29, NO. 12, 1995 / E N V I R O N M E N T A L SCIENCE & TECHNOLOGY
2941
TABLE 2
Results of pH Measurements at Various Times and Fe Analyses (in mg 1-l) at Completion of Corrosion Tests for Several Solutions in Contact with Master Buildets Blend A Iron GranulesB solution deionized water 0.01 m N a 2 S 0 4 0.02 m NaCl 0.02 m NaHC03 lrvine groundwater
PH Fe PH Fe PH Fe PH Fe PH Fe
120 h
200 h
400 h
modeled values
9.04
9.18
9.36
6.51
10.50
10.40
10.31
6.35
9.64
9.46
9.38
8.35
10.75
10.71
10.74
9.24 [Fe(OH)21 0.23 9.34 [Fe(OH)21 0.10 9.29 [Fe(OH)21 1.33 10.91 [FeCOdFe(OH)21
8.01
7.75
7.70
7.74
9.40 0.60 10.12 0.85 9.36 0.75 10.81 0.03 7.70 5.72
Oh
12 h
7.O6
0.08
"The final column records the pH and dissolved Fe concentrations predicted by the chemical equilibrium program PHREEQE with the modelpredicted saturated solid phase(s) shown in brackets. pH 7.0 assumed.
rates for the same iron material vary only by a factor of 5. The chemical modeling program PHREEQE (19) was used to simulate the effect of corrosion on several of the solution compositions used in this study (see Table 2). The Irvine groundwater was not modeled because its composition is too complex for an unequivocal simulation (several solid phases are predicted to precipitate) and because its total ion concentration is too high for PHREEQE's DebyeHuckel-based approach for the calculation of ion activity coefficients. For the solution compositions considered, the model predicts that an invariant solution composition develops upon attainment of saturation with respect to Fe(OH)2 for deionized water (pH = 9.24), 0.01 m Na2S04(pH = 9.34), and 0.02 m NaCl (pH = 9.29) and with respect to both FeC03 and Fe(OH)*for 0.02 m NaHC03 (pH = 10.91).The predicted pH and dissolved Fe concentrations are in reasonable agreement with the final measured values for allsolutions. The importance of secondary iron precipitates in controlling the low iron levels in solution is underscored by a calculation of the rates of iron entering the solution from anaerobic corrosion. For the deionized water run, for example, iron enters the solution at a rate of 100 mg L-l d-], yet at the time of sampling, the solution concentration was only 0.23 mg L-]. In other words, virtually all the iron that corrodes precipitates. The 0.01 m Na2S04solution shows the greatest difference between modeled and actual pH and iron concentrations. There is also a marked pH decrease over time for this solution. Arecent study of iron corrosion in Na2S04solution using Raman spectroscopy (21)revealed that sulfate ion is incorporated into the iron hydroxide phase and may be present as a separate phase known as green rust (22,23). Green rust is a mixed ferrous-ferric phase and may require oxygen to form. If this phase formed initiallywhen surfacesorbed oxygen might still be present and later redissolved as reducing conditions become established, it would account for an initially high and progressively decreasing pH until all sulfate is released and the solution is saturated with respect to Fe(OH)Z(pH 9.34). Although this seems a reasonable explanation, it is not clear why the Irvine groundwater, which is essentially a concentrated Na2S04 solution, would not have behaved in a similar manner. The fact that pH's did not rise above the initial pH for this saline groundwater with time is perplexing. It is not due to the buffercapacity of the water itself. Base titration of a sample of the water revealed that, given the measured corrosion rate, pH should have risen to over 9.0 in 24 h. 2942
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 29. NO. 12,1995
These limited and preliminary results on the effects of salts and their concentration indicate that anaerobic corrosion rates and lag times for steady-state hydrogen production into the gas phase are complex functions of the types of solutes present and their total concentrations. The results in Figure 4 are presented merely as a demonstration of the application of the corrosion measurement technique, and no explanations are offered here for the specific effects of differingsalts on corrosion rate. However, the effects of salts and the categorization of corrosion products under anaerobic conditions is an important part of our current research program. Long-Term Corrosion Experiment. Run B2, the corrosion experiment conducted on iron material that was first saturated with IR-sim water and then allowed to drain (see Figure 31, was allowed to react for 5 months. The cell pressure was intermittently monitored over this period, and the results are shown in Figure 5. At various times duringthe 5-month test, it was necessary to relieve the hydrogen pressure buildup by opening the valve at the top of the cell to the atmosphere for several seconds. On two occasions, a vacuum line was attached to the cell to reduce the internal pressure to near vaccum ( ~kPa). 5 For the data shown in Figure 5, pressure release to the atmosphere was performed at the 15,33,42,46,and 159 day mark and evacuation was performed on days 17 and 54. Figure 5 shows that each time the pressure was reduced, there was an apparent enhancement in the hydrogen production rate. These appear as vertical discontinuities in the rate versus time curves. The vertical discontinuity on day 31 was not due to a pressure decrease. On this day, the bath temperature was increased from 23 to 25 "C when it was realized that the electronic temperature monitor was reading low. This produced an abrupt ( ~ 0 . 0 6mmol of Fe kg-' d-l) increase in the corrosion rate. The results of the long-term corrosion test showed a progressive decrease in hydrogen production (see Figure 5). However, each time the internal cell pressure was reduced, the apparent hydrogen production rate increased. In addition, the greater the pressure reduction, the greater was the rate increase. Part of this effect can be attributed to hydrogen degassing of the aqueous phase due to the release in pressure. As reported earlier, however, hydrogen degassing should be 90% complete within 1 h after a pressure decrease. Thus, degassing could not account for the sustained enhancement in hydrogen production ob-
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FIGURE 5. Diagrams a-c are corrosion rates versus time for three monitoring periods during a 160-day corrosion test, uncorrected for hydrogen entry into the iron granules. Run conditions are identical to test B2 (see caption of Figure 3 for details). Vertical discontinuities correspond to times when the cell pressure was reduced and one instance when the bath temperature was increased by 2 "C. Diagrams al, bl, and c l are the corresponding rate measurements corrected for hydrogen entry into the iron granules (see text for details).
served over days of reaction. In fact, the effects of hydrogen degassing can be seen as the very short, steeply dipping front ends of the various curve segments in Figure 5. There are two possible explanations for the sustained increase in rate. Because hydrogen is a product of corrosion, the rate of the backward reaction may become important as the PH*in the cell increases with time. The other possibility is that hydrogen produced at the corroding iron surfaces not only enters the aqueous and gas phases but the solid phase as well and that the rate of uptake in the solid increases with increasing hydrogen pressure. It appears the first explanation can be ruled out. Thermodynamic calculations indicate that the equilibrium PHz for the iron corrosion reaction under anaerobic conditions is on the order of 4000 kPa (24).Consequently, the Fe(s)
+ 2H20(1)- Fe(OH),(s) + H,(g)
(10)
rate of the backward reaction is unlikely to be significant at the pressure conditions of the corrosion tests (5-140 kPa). Abundant evidence, however, exists for entry of hydrogen into the iron particles. Although the solubility of hydrogen in pure iron is very low ( e cm3 kg-' at a P H of ~ 100 H a ) , hydrogen can bond at dislocation sites within the iron lattice and possibly within microfractures at the surface of iron particles. One theory proposes that
dislocations bearing hydrogen migrate through the lattice and accumulate at sites of foreign inclusions and grain boundaries where the hydrogen can proceed no further (25). The resulting internal pressure buildup at these locations is believed to be the principal cause for embrittlement and blistering of iron and steel exposed to anaerobic environments. The rate of uptake at the surface is reaction-controlled but is proportional to the square root of the hydrogen partial pressure (Sievert's law):
indicating that H2 molecules disproportionate at the iron surface and enter the iron as individual atoms rather than as diatomic molecules (26). The uptake capacity can be considerable. For example, over 1 L of HZ can be induced into 1 kg of a-iron by electrolytic charging (27). Notably, diffusion into the solid does not limit the rate of hydrogen uptake. This is due to the high diffusion coefficient of hydrogen in iron. Kedzierzawski (28)reports measured diffusion coefficients for hydrogen in a-iron from eight studies. These values average 8.0 & 1.0 x m2 s-l at 298 K. Thus, hydrogen diffuses faster through a-iron than it does through water ( D x 2.0 x 10-9 m2s-1). Diffusion coefficientsthrough cold-worked VOL. 29, NO. 12, 1996 / E N V I R O N M E N T A L SCIENCE &TECHNOLOGY
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2943
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Time (h) FIGURE 6. Rate of HZuptake versus time for 500-9 dry samples of a coarse (10-18 mesh) and a fine (64-32 mesh) sample of Master Builders granular iron.
iron and steels,however, can be severalorders of magnitude lower. Many factors influence the rate constant for the hydrogen uptake reaction. These include the nature and concentration of impurities in the iron, dislocation site density, degree of cold working, thermal history, and the composition of the solution in contact with the iron. Nevertheless, the basic construct of the reaction rate expression with its PH;.~ dependence can be evaluated in terms of how well it describes the discontinuities observed in the long-term corrosion experiment (see Figure 5). Corrosion rates were recalculated for the long-term experiment by modifymg eq 9 to account for the uptake of hydrogen in the solid:
where &Z is the average hydrogen partial pressure in the cell over the time interval 4 - fl. All discontinuities observed during the first 60 days of the reaction (Figure 5a,b) were removed with a reaction rate constant (k)of 0.015 mmol kg-l d-' kPa-O.j. The recalculated corrosion rates are presented in Figure 5al and b l . The one vertical discontinuity that remains at the 31-day mark was due to an increase of the bath temperature by 2 "C as described previously. In Figure 5c, the discontinuity on day 159,which was produced as a result of a pressure decrease, was removed using a rate constant of 0.009 mmol kg-l d-' kPa-0,5 (see Figure 5cl). This indicates a decrease in the rate of hydrogen uptake by the solid with time. Based on the uptake rates and the recorded hydrogen partial pressures, the total quantity of hydrogen taken up by the iron over 160 d is estimated at 0.01 mol/kg. From the average corrosionrate (=0.5mmolkg-l d-l),only2.5gofthe551.4-g iron sample had corroded during this period. To provide direct evidence for hydrogen uptake by solid iron, two hydrogen exposure tests were conducted on 500-g samples of coarse and fine mesh size samples of Master Builders iron granules. These tests were conductedwithout water to eliminate the confounding effect of hydrogen production due to corrosion. Both dry samples were exposed to 100 kPa of H2, and the pressure decrease was monitored for 750 h. The calculated rates of uptake are plotted versus time in Figure 6. The results are similar and indicate uptake rates that are about a factor of 2 lower than those required to remove the early time discontinuities in 2944
ENVIRONMENTAL SCIENCE & TECHNOLOGY
i
VOL. 29, NO. 12, 1995
q
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401 240
260
280
300
Time (h)
FIGURE 7. Extension of the deionized waterliron corrosion run presented in Figure 4 during which the cell's pressure was reduced by evacuation after 261 h of reaction. Corrosion rates are shown corrected (a) and uncorrected (b) for hydrogen entry into the solid. Diagram c displays the raw pressure data used to calculate the rates.
Figure 5a,b. This difference is probably due to the absence of water. With water present, the effective hydrogen pressure at the surface of a corroding iron particle will be higher than the bulk pressure measured in the headspace. This is because, in wet iron material, Hz is produced at the solidlwater interface and only enters the gas phase by advection and diffusion through the water film. For dry iron that is simply exposed to hydrogen gas, hydrogen pressure at the surface of the solid is identical to the gas phase. Thus, HZuptake rates for dry granular iron should be low compared to wetted material at the same bulk gas phase hydrogen partial pressures. A final test was conducted to determine ifthe 0.015 mmol kg-l d-l kPa-o.5rate constant for solid phase uptake of hydrogen, which was derived for the saline groundwater, is significantly influenced by salt concentration. In an extension of the deionized water run shown in Figure 4, the cell pressure was reduced from 52 to 7 KPa after 261 h of reaction. The bottom and middle diagrams in Figure 7 show the recorded pressure readings versus time and the calculated corrosion rate without correction for solid phase uptake of hydrogen, respectively. The middle diagram shows the characteristic discontinuity between rate data collected prior and following a pressure decrease as well as the short-term leading 'tail' of high apparent rates due to hydrogen degassing from the water phase. To remove the discontinuity between the calculated rates for these two periods, a rate constant of 0.013 i0.002 mmol kg-' d-l kPa-o.5is required, which is only slightly lower than that obtained using the saline groundwater. The criterion used to arrive at this value is that it yields identical average corrosion rates for the pre-degassing and post-degassing periods excluding, of course, data immediately following the pressure decrease because they incorporate the anomalous effects of degassing. The upper diagram in the figure presents the corrected corrosion rates using eq 12. The
average rate for both the pre-degassing (240-261 h) and post-degassing period (265-290 h) is (0.457 & 0.016 mmol of Fe kg-’ d-l). Other factors may influence the rate of hydrogen uptake by the solid, such as the absolute value of the corrosion rate itself. This has not yet been investigated, but it appears from this limited data that total salt concentration does not play a major role in determining solid phase uptake rates. In summary, pressure transducers afford avery sensitive means to continuously monitor in situ corrosion rates of granular iron material. The principal effects that must be evaluatedto correct and interpret these rate measurements, however, are hydrogen entry into the iron granules and the lag time before steady-statehydrogen production rates are realized. It is also noted that the stainless steel sensing membranes of the particle transducers used in this study underwent embrittlement and loss of some sensitivity upon long exposure to hydrogen gas.
Acknowledgments The author thanks R. Gillham, M. Odziemkowski, S. O’Hannesin, J. Vogan, J. Warren, and many graduate students at Waterloo for stimulating discussions on aspects of ironlgroundwater interactions over the past 2 years. Support for iron corrosion research at Waterloo is provided by Envirometal Technologies, Inc., the National Science and EngineeringResearch Council of Canada (NSERC),and the Waterloo Center for Groundwater Research (WCGR).
literature Cited (1) Senzaki, T.; Yasuo, K. Kogyo Yosui 1989, 369, 19-25. (2) Reynolds, G. W.; Hoff, J. T.; Gillham, R. W. Environ. Sci. Technol. 1990, 24, 135-142. (3) Senzaki, T. Koeyo Yosui 1991, 369, 29-35. (4) Gillham, R. W.; O’Hannesin, S. F. Presented at the IAH Conference on Modern Trends in Hydrogeology, Hamilton, Ontario, 1992. (5) Gillham, R. W.; O’Hannesin, S. F. Ground Water 1994,32, 958967. (6) Matheson, L. J,; Tratnyek, P. G. Environ. Sci. Technol. 1994,24, 2045-2053.
(7) Stumm, W. J. Am. Water Works Assoc. 1956, 48, 300-310. (8) Woods, T. L.; Garrels, R. M. Geochim. Cosmochim.Acta 1992,56, 3031-3043. (9) Schikorr, G. Z. Elektrochim. 1929, 35, 62-65. (10) Shipko, F. J.; Douglas, D. L. 1.Phys. Chem. 1956,60, 1519-1523. (11) Joshi, P. S.; Venkateswaran, G.; Venkateswarlu, K. S.; Rao, K. A. Corros. Sci. 1993, 49, 300-309. (12) Dewaele, P. J.; Reardon, E. J.; Dayal, R. Cem. Concr. Res. 1991, 21, 441-454. (13) Dean, J. A. Lunge’s Handbook of Chemistry, 14th ed; McGrawHill Inc.: New York, 1992. (14) Randall, M.; Failey, C. F. J. Am. Chem. SOC. 1927, 49, 270-284. (15) Seo, M.; Hultquist, G.; Grasjo, L.; Sato, N. 10th International Conference on Metal Corrosion, Madras, India; Trans Tech Publishers: Brookfeld, VT, 1987; Vol. 1, p p 481-488. (16) Jones, D. A. Principles annd Prevention of Corrosion, MacMillan Publishing Co.: New York, 1992. (17) Simpson, J. P.; Schenk, R.; Knecht, B. Mater. Res. Symp. Ser. 1985, 50, 429-436. (18) Whitman, W.; Russel, R.; Akieri, V. Ind. Eng. Chem. 1924, 16, 665-679. (19) Parkhurst, D. L.; Thorstenson, D. C.; Plummer, L. N. PHREEQE-A computer programforgeochemical calculations;Water Resources Investigations 80-96; U.S. Geological Survey: Reston, VA, 1980. (20) Pitzer, K. S. Theory: Ion interaction approach. In Activity Coefficients in Electrolyte Solutions; Pytkowicz, R. M., Ed.; CRC Press Cleveland, 1979. (21) Gui, J.; Devine, T. M. Corros. Sci. 1994, 36, 441-462. (22) Hansen, H. C. B. Clay Miner. 1994, 24, 663-669. (23) Refait, P.; Gknin, J-M. Corros. Sci. 1994, 36, 55-65. (24) Naish, C. C.; Balkwill, P. H.; O’Brien, T. M.; Taylor, K. H.; Marsh, G. P. Final Report to the Commission of European Communities on Characterization of Radioactive Waste Forms. Task 3, No. 33, 1991, 10 pp. (25) Gourmelon, A. Mem. Sci. Rev. Metall. 1975, 72, 475-489. (26) Kumnick, A. J.; Johnson, H. H. Metall. Trans. A 1975,6A, 10871091. (27) Raczynkski, W. In Corrosion of Metals and Hydrogen-Related Phenomena; Flis, J., Ed.; Elsevier: Amsterdam, 1991; p p 267292. (28) Kedzierzawski, P. 1.Electrochem. SOC. 1985, 132, 1866-1871.
Received for review November 29, 1994. Revised manuscript received J u n e 12, 1995.Accepted J u l y 28, 1995.@
ES940726T @Abstractpublished in Advance ACS Abstracts, October 1, 1995.
VOL. 29, NO. 12, 1995 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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