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Iron (II) Activated Persulfate Oxidation of MGP Contaminated Soil

November 15, 2007

By Killian, Paul F Bruell, Clifford J; Liang, Chenju; Marley, Michael C

The persulfate anion (S^sub 2^O^sup 2-^^sub 8^) is a strong oxidant with a redox potential of 2.01 V. However, when mixed with iron (II), it is capable of forming the sulfate radical (SO^sup – ^^sub 4^sup .^^) that has an even higher redox potential (E[degrees] = 2.6 V). In these studies the sulfate radical was investigated to determine if it was a feasible oxidant for the destruction of BTEX and PAH compounds found in MGP contaminated soil. The sulfate radical was generated by either the sequential addition of iron (II) solutions or by a single addition of a citric acid chelated iron (II) solution. The sequentially added iron destroyed 86% of the total BTEX concentration and 56% of the total PAH concentration in the soil. The citric acid chelated iron destroyed 95% of the total BTEX concentration and 85% of the total PAH concentration. A second dose of persulfate and citric acid chelated iron (II) resulted in the destruction of 99% of the total BTEX concentration and 92% of the total PAH concentration. In both the sequential and chelated iron studies the lower molecular weight BTEX compounds were oxidized to a greater extent than the higher molecular weight BTEX compounds, whereas the oxidation of PAH compounds showed no preference to molecular weight. Keywords ISCO, BTEX, PAHs, Chemox

1. Introduction

Manufactured gas plants (MGPs) converted coal to gas, which was used to light the city streets, homes, and businesses from the early 1800s until the 1950s. The plants were often centrally located in cities or towns to assist with the distribution of the gas and were referred to as “town gasworks.” The process also created several undesirable byproducts such as coaltars, ash, purifier waste, blue billy, and lampblack (Hatheway, 2000). These byproducts contained hazardous chemicals such as metals, cyanides, BTEX compounds (benzene, toluene, ethylbenzene, and xylenes), and a class of semivolatile organic compounds known as polycyclic aromatic hydrocarbons (PAHs). If the MGP waste could not be sold it was often dumped at the plant or into nearby waterways (Murarka et al., 1992). Many of these contaminants are very resistant to degradation, and even though they were disposed of more than 50 years ago, remain essentially intact. Eventually, electricity and natural gas made the manufactured gas plants obsolete; however, the byproducts remain where they were disposed. There were an estimated 5,000 former MGP sites throughout the United States, few of which have been cleaned up (Lee et al., 2001).

In situ chemical oxidation (ISCO) is an innovative treatment technology in which chemicals are injected into the aquifer to oxidize the organic contaminants. The benefits of ISCO include: (i) the ability to oxidize DNAPLs; (ii) the reduction in overall treatment time; (iii) the ability to treat contaminated areas without disturbing above-ground structures; and (iv) the elimination of the cost of excavating and handling contaminated soil (Amarante, 2000). Typical oxidants used include ozone, potassium permanganate, and hydrogen peroxide. Each of these oxidants has its limitations, such as effectiveness and persistence.

Persulfate is used as a microetchant for printed circuit boards, in cosmetics to increase hair bleaching performance, and in oil and gas production in “down hole” gel formation and breaking (FMC, 2001). It is often used in the determination of total organic carbon (TOC) in wet chemical oxidation methods, where the organic carbon is oxidized to carbon dioxide (McKenna and Doering, 1995). It is also used as an initiator in polymerization reactions (Lin, 2001). In this study the sodium salt of persulfate was investigated as apossible oxidant for the destruction of MGP contaminated soil by in situ chemical oxidation.

2. Background and Theory

Liang et al. (2004) showed that adding the iron (II) in small sequential doses was much more efficient than adding the iron (II) in a large single dose. Therefore, in this study the iron (II) was added sequentially or chelated to citric acid before the addition. Citric acid is a non-toxic tetradentate chelating agent (Muller et al., 1997), which complexes with iron in a 1:1 ratio (Hamm et al., 1954).

3. Experimental Section

3.1. Materials

Reagent grade sodium persulfate, ferrous sulfate, and citric acid were purchased from VWR International. Aqueous solutions were made with reverse osmosis (RO) water generated from a Barnstead ROpure LP system located in the laboratory. The contaminated soil was obtained from a former MGP site located in the continental United States. The soil was a “well sorted” soil consisting of 2% medium sands, 70% fine sands, and 27% fines, with a f oc of 3.02%. [The foe was determined using a modified Mebius procedure (Nelson and Sommers 1982).] The soil contained MGP waste consisting of approximately 660 mg/kg BTEX and 6,400 mg/kg PAHs. Test bottles consisted of 4-oz (125- mL) wide mouth jars and 16-oz (500-mL) wide mouth jars, both contained Teflon liner caps and were obtained from Environmental Sampling Supplies (ESS). The experiments were performed using a constant temperature chamber (VWR BOD Incubator model 2020) set at 20[degrees]C.

3.2. Experimental Procedures

3.2.1. Study 1-sequentially added iron (II). Sample bottles were prepared by placing 27 g of contaminated soil in a 125-mL jar. Next, 95 mL of either a 1.05 M sodium persulfate solution (treated sample) or RO water (control sample) was added to the jar. Then 8 mL of a 0.9 M ferrous sulfate solution was added to each jar. (This left approximately 10 mL of headspace in the jar. The headspace was necessary because gases are generated during the reaction. Control samples contained the same headspace to account for any losses due to volatilization. The headspace was relieved during the daily addition of iron. No attempt was made to collect and analyze the headspace.) The samples were placed in a constant temperature chamber set at 20[degrees]C. On a daily basis for the next 5 days, 5 mL of the supernatant were removed and replaced with 5 mL of the 0.9 M ferrous sulfate solution. The solutions were mixed by inverting the jars once, then placing them back in the constant temperature chamber. On Day 6, the samples were removed, the supernatant decanted and analyzed for residual persulfate and iron (II). The soil phase was submitted to Consumer Energy Laboratory, Jackson, MI, for BTEX and PAH analyses in accordance with EPA methods 8260 and 8270, respectively.

3.2.2. Study 2-citric acid chelated iron (II). Sample bottles were prepared by placing 75 g of contaminated soil in a 500-mL jar. Next, 150 mL of either a 2.1 M sodium persulfate solution or RO water was added to the sample jar. (Two sample jars were prepared for the treated sample and one sample jar was prepared for the control sample.) Then 150 mL of a 1.0 M citric acid/ 0.20 M ferrous sulfate solution was added to each jar. The jars were placed in a constant temperature chamber set at 20[degrees]C. The lids of the jars were loosely attached to allow any gas generated during the reaction to escape. After 3 days the samples were removed and the aqueous phase decanted. An aliquot of the aqueous phase was analyzed in-house for residual persulfate concentration; the remaining portion of the aqueous phase was submitted to Consumer Energy Laboratory, for BTEX and PAH analyses. The soil phase was submitted to Consumer Energy Laboratory for BTEX and PAH analyses.

3.2.3. Study 3-effect of citric acid concentration and iron concentration. Sample bottles were prepared by placing 75 g of contaminated soil in a 500-mL jar. Next, 150 mL of either a 2.0 M sodium persulfate solution or RO water was added to the jar. (Two sample jars were prepared for each treated sample and one sample jar was prepared for the control sample.) Then 150 mL of a 1.0 M citric acid/ 0.04 M ferrous sulfate solution was added to the control sample and the low citric acid samples, and 150 mL of a 2.0 M citric acid/ 0.04 M ferrous sulfate solution was added to the high citric acid samples. The samples were placed in a constant temperature chamber set at 20[degrees]C. On a daily basis, the samples were opened and an aliquot of the aqueous phase was removed to measure residual persulfate concentration. After 7 days the samples were removed from the constant temperature chamber and the aqueous phase decanted and analyzed in-house for residual persulfate concentration. The soil phase was submitted to Resource Laboratory, Inc., Portsmouth, NH, for BTEX and PAH analyses.

3.2.4. Study 4-effect of multiple doses. Sample bottles were prepared by placing 75 g of contaminated soil in a 500-mL jar. Next, 150 mL of either a 2.0 M sodium persulfate solution or RO water was added to the jar. (Two sample jars were prepared for the treated sample and one sample jar was prepared for the control sample.) Then 150 mL of a 1.0 M citric acid/ 0.04 M ferrous sulfate solution was added to each jar. The samples were placed in a constant temperature chamber set at 20[degrees]C. On a daily basis the samples were opened and an aliquot of the aqueous phase removed to measure the residual persulfate concentration. After 7 days one set of samples (single dose sample) were removed from the chamber. The aqueous phase decanted and analyzed in-house for residual persulfate concentration. The soil phase was submitted to Resource Laboratory, Inc., for BTEX and PAH analyses. The other set of samples (double dose samples) were decanted, and a second 150 mL dose of either the 2.0 M sodium persulfate solution or RO water was added to the jar, then a second 150 mL dose of a 1.0 M citric acid/ 0.04 M ferrous sulfate solution was added to each jar. The samples were returned to the constant temperature chamber set at 20[degrees]C. Again on a daily basis the samples were opened and an aliquot of the aqueous phase removed to measure residual persulfate concentration. After a total of 19 days (12 days with the second dose) the samples were removed from the chamber. The aqueous phase decanted and analyzed in- house for residual persulfate concentration. The soil phase was submitted to Resource Laboratory, Inc., for BTEX and PAH analyses. 3.3. Sample Analysis

3.3.1. Persulfate and iron (II) measurements. Persulfate concentrations were determined by iodiometric titration with thiosulfate (Kolthoff and Stenger, 1947). Iron (II) concentrations were determined using a Hach DR/2000 spectrophotometer set at 510 nm following Hach method 8146 (binding to 1,10-phenanthroline).

3.3.2. BTEX and PAH measurements. The soil BTEX samples were preserved with methanol (1 mL methanol per 1 g of soil). The aqueous BTEX samples were prepared by filling a 40-mL Volatile Organic Analysis (VOA) vial such that there was no headspace. The VOA vial was pre-preserved with hydrochloric acid (HCl), such that the final pH was less than 2. The samples were analyzed by EPA method 8260 for the BTEX compounds. The soil PAH samples and the aqueous PAH samples were submitted to the laboratory unpreserved. The samples were analyzed by EPA method 8270 for the PAH compounds. In the first two studies the samples were analyzed by Consumer Energy Laboratory, Jackson, MI. In the last two studies the samples were analyzed by Resource Laboratory, Inc., Portsmouth, NH.

4. Results and Discussion

4.1. Study 1-Sequentially Added Iron (II)

As can be seen in Table 1, the persulfate destroyed 87% of the total BTEX concentration. The compound that had the greatest reduction in concentration was toluene (88.4 mg/kg), then total xylene (85 mg/kg), ethylbenzene (64.6 mg/kg), benzene (34 mg/kg), 1,2,4-trimethylbenzene (15 mg/kg), and finally 1,3,5- trimtheylbenzene (4.7 mg/kg). However, when the data is examined by percent reduction in concentration, then the order becomes: benzene (100%), toluene (99%), ethylbenzene (88%), total xylene (85%), and finally the trimethylbenzenes (50%). This order is the same as their aqueous solubilities (i.e., benzene has the highest relative aqueous solubility and the trimethylbenzenes have the lowest relative aqueous solubility). Figure 1 shows this shift, expressed as mole fraction within the mixture, from a greater percentage of the lighter molecular weight BTEX compounds (higher relative aqueous solubility) present in the untreated control sample to a greater percentage of the heavier molecular weight BTEX compounds (lower relative aqueous solubility) in the treated samples. As seen in Figure 1, the benzene and toluene make up approximately 42% of all the BTEX compounds found in the control sample but only 2% of all the BTEX compounds found in the treated sample. Likewise, the trimethylbenzenes make up approximately 10% of all the BTEX compounds found in the control sample, but account for 42% of all the BTEX compounds found in the treated sample.

Table 1

Residual Soil Concentrations (mg/kg) Study 1-Sequentially Added Iron (II)

Table 1 also shows that the persulfate destroyed 57% of the total PAH concentration. The compound that had the greatest reduction in concentration was naphthalene (1,200 mg/kg), then phenanthrene (400 mg/kg), and pyrene (383 mg/kg). However, when the data is examined by percent reduction in concentration, no similarity to relative aqueous solubility is observed (unlike the BTEX compounds). The percent reduction in concentration ranged from a low of 45% for phenanthrene to 100% for acenaphthene. If the data is examined based on the mole fraction of the individual PAH compounds, there is little difference between the control sample and the treated sample (Figure 2). (The compounds in Figure 2 are arranged by molecular weight, which is similar to the order of their aqueous solubility.) This indicates that the PAHs are oxidized as a group, instead of individually.

Figure 1. Plot of mole fraction of the individual BTEX compounds remaining in the soil from Study 1, sequentially added iron (II).

Figure 2. Plot of mole fraction of the individual PAH compounds remaining in the soil from Study 1, sequentially added iron (II).

The supernatant in both the treated sample and the control sample were analyzed to determine the residual persulfate concentrations and the residual iron (II) concentrations. The treated sample had over 50% of the persulfate still remaining and less than 1% of the iron (II) remaining, whereas the control sample had over 88% of the iron (II) remaining (no persulfate was added to the control sample). This indicates that there was sufficient persulfate present, but insufficient iron (II).

4.2. Study 2-Citric Acid Chelated Iron (II)

Though the sequential addition of iron worked fairly well, the logistics involved with adding multiple additions of iron may be impractical for field applications. Therefore, the use of citric acid as a chelate to regulate the iron was investigated. Presented in Table 2 are the residual concentrations of Study 2. (The results for the treated sample are the averages of the two sample jars.) The table shows similar results to those in the previous study. Though the total BTEX percent concentration reduction (78%) was lower than in Study 1 (86%), the concentration reduction (1,034 mg/kg) was much greater than in Study 1 (291 mg/kg). This was because Study 2 had a shorter duration than Study 1 and because the sample jars were not opened on a daily basis to add the ferrous sulfate solution. Both of these reasons allowed less BTEX compounds to volatilize. As in Study 1, this study also showed that the BTEX compounds that had the greatest percent reduction in concentration were the lower molecular weight BTEX compounds, and this again resulted in a shift in BTEX composition from the lighter molecular weight compounds to the heavier molecular weight compounds following treatment. Table 2 also shows that the PAH results were similar to those in Study 1. Again there was no relationship between percent reduction in concentration and relative aqueous solubility for the PAHs. As with the previous study, the composition (mole fraction) of the PAH compounds in both the control sample and treated sample were similar.

Table 2

Residual Soil Concentrations (mg/kg) Study 2-Citric Acid Chelated Iron (II)

In this study the aqueous phase was also analyzed for residual BTEX concentrations and residual PAH concentrations. The results are presented in Table 3. (The treated sample results are the averages of the results from the two sample jars.) The aqueous phase of the control sample contained 13,130 [mu]g/L total BTEX and 5,663 [mu]g/ L total PAH. The aqueous phase of the treated sample contained 70 [mu]g/L total BTEX and 315 [mu]g/L total PAH. This indicates that the persulfate oxidized over 99% of the total BTEX and over 94% of the total PAH leached from the soil.

The supernatant was also analyzed for residual persulfate concentrations. The two treated samples had 66.6 mmoles and 70.3 mmoles of persulfate remaining of the 315 mmoles added. This indicates that the chelated iron reaction consumed 78% of the persulfate over three days.

Table 3

Residual Aqueous Phase Concentrations ([mu]g/L) Study 2-Citric Acid Chelated Iron (II)

43. Study 3-Effect of Citric Acid Concentration and Iron Concentration

Study 2 demonstrated that citric acid chelated iron persulfate was partially successful at oxidizing the BTEX and PAH from MGP waste. However, at a citric acid to iron (II) molar ratio of 5:1, the persulfate was consumed quickly. Therefore a lower iron concentration was used at two different citric acid concentrations to slow down the consumption of persulfate and to increase the overall degradation of the MGP waste.

Presented in Figures 3a, 3b, and 3c, respectively, are the total BTEX and total PAH concentrations remaining in the soil, and milimoles of persulfate remaining in the aqueous phase. (For the purpose of comparison the results from Study 2 are also included.) As shown in Figure 3a, both concentrations of citric acid oxidized over 90% of the total BTEX, compared to only 78% for Study 2. (The control sample for Study 3 had a lower total BTEX concentration than the control sample in Study 2 because Study 3 had a longer duration and the sample jars were opened on a daily basis to measure residual persulfate. Both of these reasons allowed more BTEX compounds to volatilize.) Both citric acid concentrations also reduced the total PAH by over 85%, compared to 47% for Study 2. (Though Study 2 was 4 days shorter in length than Study 3, it is estimated that if Study 2 had run for 7 days then all the persulfate would have been exhausted and the reduction of total PAHs would not have been much greater than 47%), The reason for the increase in oxidation was due to the slower rate of persulfate consumption. As shown in Figure 3c, Study 2 had approximately 22% of the persulfate remaining after 3 days, whereas both samples in Study 3 had approximately 60% of the persulfate remaining after 3 days, and approximately 40% of the persulfate remaining after 7 days. The sample with a higher concentration of citric acid did consume more of the persulfate than the low citric acid sample, but as can be seen in Figures 3a and 3b, this neither increased nor decreased the degradation of total BTEX or total PAHs in the soil. Figure 3a. Plot of residual Total BTEX concentration remaining in the soil after Study 2 (after 3 days) and Study 3 (after 7 days). (Error bar for the Treated Sample represents the difference between sample concentration and average concentration.)

4.4. Study 4-Effect of Multiple Doses

Study 3 demonstrated that over 85% of the BTEX and PAHs were destroyed in a slow consumption of persulfate; however, 15% still remained. Therefore, Study 4 was performed to determine if a second dose and an extended reaction period would increase the destruction to over 90%. In this study, the soil was treated with a dose of persulfate-citric acid-iron for 7 days, then the aqueous phase was decanted, and a second dose of persulfate-citric acid-iron was added and allowed to react for an additional 12 days. Presented in Table 4 are the residual total BTEX concentrations and residual total PAH concentrations in the soil. The sample treated with a second dose of oxidant had less than 1% of the total BTEX concentration found in the control sample, and the only compounds remaining were 1,2,4- trimethylbenzene and 1,3,5-trimethylbenzene. The total PAH concentration was reduced to less than 8% of the concentration found in the control sample. As shown in Table 4, the second dose significantly reduced lower molecular weight compounds such as naphthalene (270 mg/kg versus 61.5 mg/kg.), 2-methylnaphthalene (68.5 mg/kg versus 22 mg/kg), and phenanthrene (130 mg/kg versus 105 mg/kg). Whereas the higher molecular weight compounds were reduced only slightly such as chrysene (20.5 mg/kg versus 16.5 mg/kg), benzo(b)fluoranthene (17.5 mg/kg versus 17.5 mg/kg), and benzo(k)fluoranthene (7 mg/kg versus 7 mg/kg). This shift from a greater percentage of the lower molecular weight PAH compounds to a greater percentage of the higher molecular weight PAH compounds is reflected in Figure 4. Figure 4a shows that in the single dose samples there is little difference in composition between the treated sample and the control sample (even though there was a 85% reduction in concentration). However, Figure 4b shows that in the two dose samples, the treated sample contained only 25% naphthalene compared to 50% found in the control sample. Conversely, the phenanthrene made up 30% of the treated sample compared to only 10% in the control sample. This indicates that in the treated sample a greater number of the lower molecular weight PAH compounds were oxidized compared to the higher molecular weight PAH compounds.

Figure 3b. Plot of residual Total PAH concentration remaining in the soil after Study 2 (after 3 days) and Study 3 (after 7 days). (Error bar for the Treated Sample represents the difference between sample concentration and average concentration.)

Figure 3c. Plot of mmoles of persulfate remaining in the aqueous phase after Study 2 and Study 3. (Error bar represents the difference between sample concentration and average concentration.)

Table 4

Residual Soil Concentrations (mg/kg) Study 4 – Effect of Multiple Doses

Figure 4a. Plot of mole fraction of the individual PAH compounds remaining in the soil after single doses of oxidant, Study 4. (The error bar for the Treated Sample represents the difference between sample concentration and the average concentration.)

Figure 4b. Plot of mole fraction of the individual PAH compounds remaining in the soil after two doses of oxidant, Study 4. (The error bar for the Treated Sample represents the difference between sample concentration and the average concentration.)

4.5. Influence of Solubility

In all four studies, the BTEX compounds were selectively oxidized, as the more aqueous soluble compounds (benzene and toluene) were oxidized to a greater percentage than the less aqueous soluble compounds (trimethylbenzenes). This was not an unexpected observation as the sulfate radical reaction is an aqueous phase reaction. Therefore, the greater extent to which the BTEX compounds dissolve into the aqueous phase, the greater the extent of their oxidation. This is supported by the observation that the longer the duration of the study (thereby allowing more to dissolve), the greater the percentage of BTEX oxidized. In Study 2, which lasted only 3 days, only 78% of the BTEX was oxidized, whereas in Study 4, which lasted 19 days, over 99% of the BTEX was oxidized.

By contrast, in most of the studies the PAHs compounds were observed to be oxidized as a group rather than selectively. This indicates that the PAHs, with their lower aqueous solubilities, were not oxidized in the bulk aqueous phase as dissolved components. Rather, the sulfate radical oxidized the non-aqueous phase liquid (NAPL) as a whole, with the reaction most likely occurring at the non-aqueous phase/aqueous phase interface. However, when sufficient time was allowed for more of the PAHs to dissolve (as in Study 4), the PAHs were observed to be more selectively oxidized, with the more soluble compounds being oxidized to a greater extent.

5. Conclusion

These studies demonstrated the ability of persulfate to oxidize both the BTEX and PAH compounds found in MGP waste. The use of citric acid chelated iron (II) as a persulfate activator was superior to the use of sequentially added iron (II). The persulfate oxidation of the BTEX compounds was more extensive for the lighter molecular weight compounds (which have higher relative aqueous solubility) than the heavier molecular weight compounds. The persulfate oxidation of the PAH compounds initially showed no such preference, as all compounds were oxidized to about the same degree, and the composition of the PAHs remained unchanged. However, when a second dose of oxidant was added, the lower molecular weight PAHs were oxidized to a greater extent than the higher molecular weight PAHs. The persulfate was also very successful at oxidizing the compounds that leached into the aqueous phase, as 99% of the BTEX compounds and 94% of the PAH compounds were destroyed.

Soil and contaminant conditions vary at different MGP sites. Therefore, site-specific laboratory treatability studies are required. However, it is postulated that the use of optimized citric acid chelated iron (II) activated persulfate could be used in the field for the destruction of MGP waste in soils.

References

Amarante, D. 2000. Applying in situ chemical oxidation. Pollution Engineering February 2000, 40-42.

Eberson, L. 1987. Electron Transfer Reactions in Organic Chemistry, Springer-Verlag, Berlin, pp. 88-91.

FMC Corporation. 2001. Persulfates Technical Information, Ammonium, Potassium, Sodium Salts, FMC9487-2500, pp. 1-16, www.fmcchemicals.com/Content/CPG/Images/AOD-Brochure-Persulfate.PDF

Hach Company. 1993. DR/2000 Spectrophotometer Procedures Manual, Hach Company, Loveland, CO.

Hamm, R., Shull, C., and Grant, D. 1954. Citrate complexes with iron (II) and iron (III). J. Am. Chem Soc. 76, 2111-2114.

Hatheway, A. 2000. Former Manufactured Gas Plants. Site and Waste Characterization and Remedial Engineering of Former Manufactured Gas Plants and Other Coal-Tar Sites. www.hatheway.net Allen Hatheway, Rolla, MO.

House, D. 1962. Kinetics and mechanism of oxidation by peroxydisulfate. Chem. Rev. 62, 185-203.

Kolthoff, I. and Stenger, V. 1947. Volumetric Analysis, Second Revised Edition, Volume II Titration Methods: Acid-base, Precipitation, and Complex Reactions, Interscience Publishers, Inc., New York, pp. 287-290.

Kolthoff, I, Medalia, A., and Raaen, H. 1951. The reaction between ferrous iron and peroxides. IV. Reaction with potassium persulfate. J. Am. Chem. Soc. 73, 1733-1739.

Latimer, W.M. 1952. The Oxidation States of the Elements and their Potentials in Aqueous Solutions. Prentice-Hall, New York, p. 78.

Lee, P., Ong, S, Golchin, J., and Nelson, G. 2001. Use of solvents to enhance PAH biodegradation of coal tar-contaminated soils. Water Resources 35, 3941-3949.

Liang, C. 2002. Chemical Oxidation of Chlorinated Solvents by Sodium Persulfate in Aqueous and Soil Slurry Systems, Doctoral Dissertation, University of Massachusetts, Lowell.

Liang, C., Bruell, C., Marley, M., and Sperry, K. 2004. Persulfate oxidation for in situ remediation of TCE. I. Activated by ferrous ion with and without a persulfate-thiosulfate redox couple. Chem. 55, 1213-1223.

Lin, H. 2001. Solution polymerization of acrylamide using potassium persulfate as an initiator: kinetic studies, temperature and pH dependence. European Polymer Journal 37, 1507-1510.

McKenna, J., and Doering, P. 1995. Measurement of dissolved organic carbon by wet chemical oxidation with persulfate: influence of chloride concentration and reagent volume. Marine Chemistry 48, 109-114.

Muller, B., Klager, W., and Kubitzki, G. 1997. Metal chelates of citric acid as corrosion inhibitors for zinc pigment. Corrosion Science 39, 1481-1485.

Murarka, L, Neuhauser, E., Sherman, M, Taylor, B., Mauro, D, Ripp, J., and Taylor, T. 1992. Organic substances in the subsurface: delineation, migration, and remediation. J. of Hazardous Materials 32, 245-261.

Nelson, D. and Sommers, L. 1982. Total carbon, organic carbon, and organic matter. Methods of Soil Analysis-Part 2 Chemical and Microbiological Properties, Second Edition. American Society of Agronomy, Inc., Soil Science Society of America, Inc., Madison, WI, pp. 571-573.

Neta, P., Madhavan, V., Zemel, H., and Fessenden, R. 1977. Rate constants and mechanism of reaction of SO^sup -^^sub 4^ with aromatic compounds. J. Am. Chem Soc. 99, 163-164.

Pugh, J. 1999. In situ remediation of soils containing organic contaminants using the electromigration of peroxysulfate ions, United States Patent No. 5,967,348.

Walling, C., Camaioni, D., and Kim, S. 1978. Aromatic hydroxylation by peroxydisulfate. J. Am. Chem. Soc. 100, 4814-4818.

PAUL F. KILLIAN,1 CLIFFORD J. BRUELL,2 CHENJU LIANG,3 AND MICHAEL C. MARLEY4

1 Ambient Engineering, Inc., Concord, MA, USA

2 Civil and Environmental Engineering Department, University of Massachusetts, Lowell, MA, USA 3 Department of Environmental Engineering, National Chung-Hsing University, Taichung City, Taiwan

4 Xpert Design & Diagnostic, Inc., Stratham, NH, USA

Address correspondence to Paul F. Killian, Ambient Engineering, Inc., 100 Main Street, Concord, MA 01742, USA. E-mail: pkillian@ambient-engineering.com

Copyright Taylor & Francis Ltd. 2007

(c) 2007 Soil & Sediment Contamination. Provided by ProQuest Information and Learning. All rights Reserved.




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