Perchlorate Chemistry: Implications for Analysis and Remediation Edward T. Urbansky U.S. Environmental Protection Agency, National Risk Management Research Laboratory, Water Supply and Water Resources Division, Treatment Technology Evaluation Branch, Cincinnati, Ohio 45268 Abstract: Since the discovery of perchlorate in the ground and surface waters of several western states, there has been increasing interest in the health effects resulting from chronic exposure to low (parts per billion [ppb]) levels. With this concern has come a need to investigate technologies that might be used to remediate contaminated sites or to treat contaminated water to make it safe for drinking. Possible technologies include physical separation (precipitation, anion exchange, reverse osmosis, and electrodialysis), chemical and electrochemical reduction, and biological or biochemical reduction. A fairly unique combination of chemical and physical properties of perchlo- rate poses challenges to its analysis and reduction in the environment or in drinking water. The implications of these properties are discussed in terms of remediative or treatment strategies. Recent developments are also covered. Keywords: perchlorate, bioremediation, reductase, thyroid, anion exchange, electrochemical reduction, kinetic barrier, oxidant, reverse osmosis, electrodialysis, ion chromatography, capillary electrophoresis, analysis. Introduction Several factors have brought about the current interest in perchlorate (ClOj), which, because of its chemical and physical nature, presents challenges for analysis and remediation. Perchlorate has been found in ground- water and in surface waters in several western states, including the Colorado River. Concentrations ranging from 8 ng mL"1 to 3.7 mg mlr1 have been measured. The extensive use of Colorado River water in this region and the proximity of some of these sites to the river have heightened the concern. Other local water supplies also are affected. Perchlorate targets the thy- roid, bone marrow, and muscle tissue at sufficiently high concentrations; however, it is unknown what ef- fects, if any, occur at the levels currently encountered in the contaminated water sources. Although addi- tional toxicological studies are ongoing (Table 1), an action level of 18 ng mlr1 has been adopted by Cali- fornia and informally by other affected states. The fundamental physical and chemical nature of perchlorate make it difficult to uniquely analyze for and to remediate, especially at the low concentrations typically encountered (i.e., <500 u.g mL"1). Although ion chromatography is capable of determining very low levels (e.g., 5 ng mL"1), retention time is not considered a unique identifier, and known confirma- tory tests have much higher detection limits. Perchlo- rate ion is unreactive as a ligand and its salts are extremely soluble, even in organic solvents. Despite its strength as an oxidizing agent, perchlorate is nonlabile, that is, very slow to react. This kinetic bar- rier is well known and widely made use of in chemical studies. Common reducing agents do not reduce per- chlorate, and common cations do not precipitate it. Consequently, standard practices of water treatment 1058-8337/98/$.50 © 1998 by Battelle Memorial Institute Bioremediation Journal 2(2):81-95 (1998) 81 ------- Table 1. Perchlorate facts at a glance. Date of discovery Discovered in 1997 in Western U.S. ground and surface waters Physiological implications Known to target the thyroid, bone marrow, and muscle at high concentrations Medical use Once used medicinally to treat Graves' disease (hyperthyroidism) Remediation implications Difficult to quantitate and remediate because of its high solubility and low kinetic lability towards reduction will neither remove it physically nor destroy it chemi- cally. Sources and Nature of the Contamination Ammonium perchlorate (NH4C1O4) has been used as an energetics booster in rocket fuels, and it appears most perchlorate contamination is the result of dis- charge from rocket fuel manufacturing plants or from the demilitarization of weaponry (missiles). It is im- portant to emphasize that ammonium perchlorate is not rocket fuel; it is an additive. Potassium perchlorate (KC1O4) can be used as a solid oxidant for rocket propulsion, and it was the original source for a fraction of the contamination. However, most of the contami- nation appears to have come from the legal discharge decades ago of then unregulated waste effluents con- taining high levels of ammonium perchlorate. Although ammonium perchlorate was released initially, the salt is highly soluble and dissociates completely to ammo- nium and perchlorate ions upon dissolving in water (equation 1): NH4C104(s) (1) It is likely that most of the ammonium has been biode- graded and the cation is now best viewed as mostly sodium (Na+) or possibly hydrogen (H+), especially where levels are below 100 |ig mL"1; nevertheless, those regions with high concentrations of perchlorate ion probably retain at least some ammonium ion. At those sites where contamination dates back decades, very little (if any) ammonium ion has been found. To date, there has been no quantitative determination of the cations responsible for the charge balance. Three states are known to be substantially af- fected: Utah, California, and Nevada. Arizona also may be affected since it too draws water from the Colorado River. Perchlorate concentrations in Utah range from 4 to 200 ng mL"1 in groundwater wells on the property of rocket motor manufacturer Alliant Techsystems. A level of 13 ng mL"1 at the Kennecott Utah Copper mines in Magna, Utah has led the com- pany to supply its miners with bottled water for drink- ing. In Henderson, Nevada, water samples taken 1000 ft (300 m) from the site of the former Pacific Engineer- ing & Production Company of Nevada (PEPCON) rocket fuel plant, which exploded in 1988, contain as much as 630 )J.g mL"1. Wells near the site show con- centrations ranging from 51.4 to 630 |J.g mL'1. Samples drawn from 50 wells near ammonium perchlorate manufacturer Kerr-McGee Chemical Corporation, lo- cated about 1 mile (1.6 km) from the abandoned PEPCON site, also showed significant perchlorate contamination. The Kerr-McGee samples showed per- chlorate levels as high as 3.7 mg mL"1 in the ground- water. Surface water samples taken in August 1997 from the Las Vegas Wash, which feeds into Lake Mead, had perchlorate concentrations between 1.50 and 1.68 (J.g mL"1. Lake Mead is formed by the Hoover Dam on the Colorado River and thus affects' the water supply of southern California, including Los Angeles. Testing by the Los Angeles Metropolitan Water District found 8 ng mL"1 perchlorate at its intake anc$VLake Mead at the Hoover Dam. Lake Mead lies at the southern tip of Nevada, straddling the Arizona border. This prompted the Southern Nevada Water Authority (SNWA) to begin testing its water; the SNWA found 11 ng mL"1 perchlorate in tap water. The California Department of Health Services (DHS) began testing wells in early 1997, and has closed more than 20 wells for exceeding the action level of 18 ng mL"1. Some wells could not be closed because of the high water demand of the regions served by the utility companies. Because the utilities often blend water from several sources, it is possible to dilute water from some of these wells. The DHS cer- tifies commercial testing laboratories to perform per- chlorate determinations, and has therefore stopped its own testing program (Cal DHS, 1997a). DHS reports that some regions of Lake Mead showed levels up to 165 ng mL"1. 82 Urbansky ------- Physiological and Health Effects In 1992, the U.S. Environmental Protection Agency (EPA) reviewed and assessed the health effects of perchlorate administered chemotherapeutically to pa- tients with hyperthyroidism (Dollarhide, 1992, 1995; Stanbury, 1952). This study showed a no observable adverse effects level (NOAEL) of 0.14 mg kg'1 day"1. Doses of 6 mg kg"1 day"1 or more for periods of at least 2 months led to fatal bone marrow changes. The EPA study recommended the following safety/error factors: 10 (nonchronic study), 10 (sensitive persons), 10 or 3 (database error margin) and allowed for two possible uncertainty factors, 1000 and 300. Using the somewhat arbitrary, but relatively ac- cepted, uncertainty factor of 300, the California DHS established 1.8 ng mL"1 as the action level for initiating remediation and stopping water usage (Cal DHS, 1997b). This cut-off assumes a 70-kg person consum- ing 0.5 |4.g perchlorate for each kilogram body mass who drinks 2 L of water daily (18 ng mL"1 = 0.016 mg mL"1 = 0.14 mg kg"1 day"1 x 70 kg x 1 day/2 L •*• 300). The 0.5 [J.g number introduces a rounding error that was carried through (Cal DHS, 1997b). This 18 ng mL"1 action level has been adopted informally by other governmental agencies in the region as well. Using the same assumptions, we would calculate that harmful thyroid effects begin to occur at 49 jig mL"1, and fatalities occur at 210 to 490 (ig mL"1. Meanwhile, the European Communities (1982) set a maximum admis- sible guide level of 20 jig NaClO4 mL"1 for drinking water. This corresponds to 16 |0.g C1OJ mL"1. Perchlorate exerts its most commonly observed physiological effects on or through the thyroid gland. The primary effect is a decrease in thyroid hormone output. The thyroid gland takes up iodide ion from the bloodstream and converts it to organic iodide in the form of hormones that regulate metabolism. The mecha- nism responsible for this process, the cellular iodide pump, preferentially selects for anions on the basis of ionic volume: I"=SCN" < CIO;, TcO; (Chiovato et al., 1997; Cooper, 1991; Foye, 1989; Orgiassi, 1990). Consequently, the presence of any large anion in the serum reduces thyroid hormone production. This phenomenon was once used pharmaceuti- cally to treat hyperthyroidism, which is known as Graves' disease (Foye, 1989; Chiovato et al., 1997; Cooper, 1991; Orgiassi, 1990). Chemotherapeutic use of perchlorate was reduced substantially in the United States after several instances of aplastic anemia and renal damage were observed (Foye, 1989; Hobson, 1961). Domestic perchlorate use now is restricted al- most exclusively to use as a diagnostic tool for the evaluation of thyroid hormone production. As a diag- nostic tool, perchlorate is still the standard for evalu- ating thyroid activity; the protocol at the University of California, Los Angeles (UCLA) requires a dose of 0.6 g (pediatric) or 1 g (adult) (UCLA, 1997). Following the administration of radiolabeled iodide, perchlorate is used to displace iodide anion in the iodide pump. When thyroid function is low, most of the radioiodine remains as inorganic iodide (rather than being con- verted to an organic iodide) and is lost; therefore, very little intrathyroid iodine shows the radiolabel. Although perchlorate has been used as a treatment for hyperthyroidism, under the right circumstances it also can act as goitrogen in rodents and prevent thyroid hormone formation by interfering with iodide uptake (Capen and Martin, 1989). The low level of hormones is recognized by the pituitary gland which then stimu- lates the thyroid gland to work harder, eventually lead- ing to goiter. A recent study of thyroid hormone levels in the Sprague-Dawley rat supported the EPA refer- ence dose of 0.14 mg kg"1 day"1. Male rats exhibited a thyroid NOAEL of 0.44 mg kg"1 day"1, but females exhibited a thyroid NOAEL of only 0.124 mg kg"1 day"1 (King, 1995). Potassium perchlorate has been used to treat thyrotoxicosis without toxicity at doses ranging from 40 to 120 mg day"1 (Cooper, 1996). If we assume a daily intake of 3 L of water, this would correspond to 13 to 40 (ig KC1O4 mL"1, or about 9 to 12 fig CIO; mL"1. This is a factor of about 1000 times the California DHS action level, but close in line with the European Communities level. It is unknown whether secondary effects resulting from decreased thyroid function, indirectly caused by perchlorate, will be con- sequential. No studies link perchlorate to any second- ary adverse health effects at this time. Perchlorate can directly affect organs and tissues in addition to the thyroid gland. The mouse mammary gland has a mechanism similar to the thyroid iodide pump that is inhibited by perchlorate (Rillema and Rowady, 1997); however, it is unclear whether this has any significance for human health. At high (millimo- lar) concentrations, perchlorate is known to potentiate excitation-contraction (E-C) coupling and charge move- ment in muscle cells (Bruton et al., 1995; Gonzalez and Rios, 1993; Jong et al., 1997; Khammari et al., 1996; Ma et al., 1993; Pereon et al., 1996). At this level, part of the E-C effect is due to activation of calcium ion release from the sarcoplasmic reticulum (Fruen et al., 1994; Percival et al., 1994; Yano et al., 1995). In fact, this property of perchlorate often is exploited to study muscle physiology in animals. Much of what is known about perchlorate's effects on living organisms is derived from studies of acute toxicity over relatively short periods of time rather than chronic exposure to very low concentrations over a lifetime. Perchlorate Chemistry: Implications for Analysis and Remediation 83 ------- The U.S. Air Force Research Laboratories (AFRL) are conducting animal toxicology studies under the guidance of the EPA's National Center for Environ- mental Assessment (NCEA) and in consultation with state agencies. These studies are intended to refine the NOAEL and to set a standard to replace the current action level; preliminary results are scheduled for re- lease in September 1998. EPA's Office of Water has added perchlorate to the candidate contaminant list (CCL), but it is unclear whether this will eventually culminate in the establishment of a maximum contami- nant level (MCL) for perchlorate in drinking water. Chemical and Physical Properties Perchlorate as a Noncomplexing Anion Perchlorate is widely known to be a very poor complexing agent (Cotton et al., 1987) and is used extensively as a counter ion in studies of metal cation chemistry, especially in nonaqueous solution. In this use, it is comparable with other noncomplexing or weakly ligating anions, e.g., trifluoromethanesulfonate (triflate, CF3SOJ), tetrafluoroborate (BFj), and to a lesser extent nitrate (NOj). Some exceptions are known, but these are rare. All of these anions have a highly delocalized (NOJ, ClOj, CF3SOj) or sterically blocked (BFJ) monovalent anionic charge and large volume; the low charge density reduces their affinity for cat- ions and their extent of aquation (see Table 2). This low association with cations is responsible for the extremely high solubilities of perchlorate salts in aque- ous and nonaqueous media. It is important to point out that the solubility is not due to association with the solvent While perchlorate is often described as strongly retained on anion exchange resins, the truth is that the Table 2. Gibbs free energies of formation for selected anions in aqueous solution.3 Anion AG,°, kJ moM BF<- so|- HCOj OH- ci- NOj Br cio<- ClOj -1490" -1019 -744 -587 -157 -131 -109 -104 -8.5 -8.0 From Barrow (1988), except BF;. From Dean (1985). ion in the starting form of the resin (e.g., chloride or hydroxide) is much more hydrophilic than perchlorate (Table 2). Perchlorate Salts as Supporting Electrolytes or Ionic Strength Adjustors In addition to its use in synthesizing transition metal compounds where competition for coordination is un- desirable, sodium perchlorate is used extensively as a means of adjusting ionic strength for equilibrium, ki- netics, and electrochemical studies where potassium nitrate cannot be used. For example, many halogen species undergo multiple simultaneous equilibria in . which a central halogen atom expands its octet and forms a hypervalent species; perchlorate does not act as a ligand in these situations (Urbansky et al., 1997). Inorganic perchlorate salts are generally extremely soluble, with potassium perchlorate the notable excep- tion (-17 g L"1 = 0.12 M). The solubility of sodium perchlorate in water is extremely high, just under 8 M; only the mineral acids and the alkali metal hydroxides surpass it in solubility. Kinetics and Thermodynamics of Perchlorate Reduction Besides its weak ligating ability, perchlorate often is used to adjust ionic strength because of its low reactiv- ity as an oxidant. At first glance, this seems surprising given that it contains a highly oxidized central halogen atom, chlorine(VII). However, the low reactivity is a matter of kinetic lability rather than thermodynamic stability. The standard reduction potentials (Emsley, 1989) for the half-reactions (equations 2 and 3) clearly indicate that reductions to chloride or chlorate are very favorable processes from a thermodynamic standpoint: 8H+ + 8e^Cl- + 4H2O E° = 1.287V (2) C1O4- + 2 H+ + 2 e ;= C1O3- + H2O E° = 1.201 V (3) Given thermodynamics alone, we would expect per- chloric acid to oxidize water to oxygen because the water-oxygen couple has an oxidation potential of -1.229 V (Emsley, 1989). 2 H20 ,=* 4 H+ + O2 -E° = -1.229 V (4) Consequently, we can confidently state that the ob- served behavior of perchlorate is dominated largely by 84 Urbansky ------- its kinetics. In fact, perchloric acid is quite unreactive toward most reducing agents when cold and dilute. For example, digestion of organic material (wet ashing) with perchloric acid requires heating the material with the concentrated acid (Harris, 1991; Schilt, 1979). The kinetic barrier to perchlorate reduction is very useful in many oxidation-reduction investigations because it allows control of the ionic strength with a non- complexing anion, even at moderately acidic condi- tions (e.g., 1 M) where nitrate would be reduced. It is fortunate, for the sake of remediation, that the behavior of perchlorate is due to a kinetic activation barrier rather than a thermodynamic barrier, because all kinetic problems can be reduced. It comes down to the matter of "where there's a will, there's a way." Also fortunate for us, the way is fairly well defined for perchlorate. Depending on the reductant, perchlorate may be reduced to either chlorate or chloride (Cotton and Wilkinson, 1988). Ruthenium(H) reduces perchlo- rate to chlorate, whereas vanadium(II), vanadium(ni), molybdenum(in), dimolybdenum(in), chromium(n), and titanium(in) all reduce perchlorate to chloride. The first work to show reduction of perchlorate by a metal cation was done by Rothmund (1909). He showed that Ti(lH), V(H), and Cr(H) all reduce per- chlorate to chloride in acidic aqueous solution at am- bient temperature. Bredig and Michel (1922) refined Rothmund's Ti(Tfl) work, and they showed that Mo(ffl) also reduces perchlorate to chloride. With the right catalyst, other reductants will react with perchlorate. In the presence of ruthenium(ni,IV) (Crowell et al., 1929) or osmium(TV) (Crowell et al., 1940), bromide will reduce perchlorate. Tin(II) will reduce perchlorate in the presence of molybdate (Haight and Sager, 1952). While these studies were significant and substantial at the time of publication, the treatment of the data was insufficiently rigorous to apply it here directly. Nevertheless, these papers laid the ground- work for many of the later investigations and still supply directions for future study. King and Garner (1954) published the results of the first thorough kinetic investigation of the reaction of vanadium(II) and vanadium(in) with perchlorate. Reactions 5 through 7 summarize the behavior they observed. The oxidation-reduction reactions of per- chlorate with V(II) or V(III) occur on comparable time scales (equations 5 and 6). The comproportionation in equation 7 is much faster, essentially instantaneous. 8 V2+ + C1OJ + 8 H+ -» 8 V3+ + Cl- + 4 H2O (5) 8 V3+ + C1O4- + 4 H2O -> 8 VO2+ + Cl~ + 8 H* (6) Kallen and Barley (1971) published a detailed investigation on the reaction between hexaaquo- ruthenium(n) and perchlorate (equation 8). They also discussed the factors that control reaction rates of per- chlorate reduction by metal cations. 2 Ru + C1OJ + 2 H+ -> 2 Ru3+ + ClOj + H2O (8) Duke and Quinney (1954) published the first rig- orous study on the reaction of titanium(IH) and per- chlorate (equation 9). They found that the reaction proceeds through an initial complexation, after which Ti(IH) is oxidized to a titanyl ion, TiO2+(equation 10). 8 Ti3+ + C104- + 8 H+ -H> 8 Ti(IV) + CT + 4 H2O (9) Ti3+ + ClOj ^= TiO2+ + C1O3° (10) Their postulation of the radical chlorine trioxide as the first product is still accepted today. Cope et al. (1967) studied the Tim-ClO4-reaction in the absence of chloride. They obtained the differen- tial rate expression given by equation 11: rate = -d[Tini]/df = (ifc + JfcfH+f^TPjaO;] (11) where* = 1.9 x 1Q-4 M'1 s-1 and kf = 1.25 x W4 M~2 Possibly the most significant paper with regard to chemical reduction deals with the redox reaction of perchlorate with Af-(hydroxyethyl)ethylenediamine- NJf ,Ar-triacetatopentaaquotitanium(in) ion (Liu et al., 1984). The net reaction is shown in equation 12. 8 Ti(Hedta) + CIO; + 8 H+ -> 8 Ti(IV) + CT + Hedta3~ + 4H20 (12) This Ti(in) chelate is reasonably stable in air. The Ti(IV) produced begins to form hydrous oxides over a matter of an hour or so. Over the course of hours to days, fine suspended crystallites of TiO2 develop. They found the rate expressible as rate = -d[lim J/dr = + £ j [Ti(Hedta)] (13) VO2+ + V2+ + 2 H* -» 2 V3* + H2O (7) where k = 2 x lO"3 M~2 s~' and kf = 2 x lO'8 s~l. Based on this study by Liu et al. (1984), we might propose to treat perchlorate-contaminated waters with Ti(m) chelates under anaerobic conditions. The chlo- Perchlorate Chemistry: Implications for Analysis and Remediation 85 ------- ride produced is harmless, and the TiO2 may be re- moved by agglutination and sedimentation or by filtra- tion. TiO2 is very insoluble and quite nontoxic. A number of stable titanium(H[) chelates and complexes have been prepared; however, in general, precautions have been taken to exclude oxygen (Chaudhuri and Diebler, 1986). It is unclear if any fairly air-stable titaniumCEH) chelate may be synthesized that will still react quickly with perchlorate. Even if other factors can be overcome, this reaction is still too slow to be useful. If we could lower the pH to 4, for 1 mM ClO^ and 800 mM Ti(Hedta), the rate of the perchlorate oxidation of Ti(J3I) would be 1.6 x 10~10 M s'1, while the rate of the decomposition would be 1.6 x 10"8 M s'1, i.e., 100 times faster! Even if we stopped the decomposition, the half-life for this reaction would be 50 days, far too long to be practicable in a water treatment plant. Hills et al. (1986) demonstrated that molyb- denum(ni) and dimolybdenum(m) are capable of re- ducing perchlorate in acidic solution. This is notewor- thy in light of the fact that molybdenum is not known to have any stable "-yl" ions in aqueous solution, the significance of which is described below. Taube (1982) has speculated that the relative sta- bility of the resulting "-yl," i.e., M=O"+, ions is largely responsible for the behavior observed for titanium, vanadium, and other metals, whereas others (Kallen and Barley, 1971; Liu et al., 1984) have put forth an alternative explanation that they feel is superior relat- ing to the polarizability of the metal d-orbitals and the interaction of these orbitals with the lowest unoccu- pied molecular orbital (LUMO) of the perchlorate. Although Taube's explanation is more well known, the Kallen-Earley postulate is reasonable and fits bet- ter with some of the experimental evidence. Very recently, methylrhenium dioxide (CH3ReC>2) has been shown to abstract oxygen atoms from halates and perhalates (Cl, Br, and I) (Abu-Omar and Espenson, 1995; Abu-Omar et al., 1996). The net reaction is given by equation 14: 4 CH3Re02 + CIO; -H> 4 CH3ReO3 + Cl~ (14) The first step in the process is a reduction to chlorate: CH3ReO2 + C1OJ -> CH3ReO3 + C1OJ (15) with a reaction rate given by rate = -d[ CH3ReO2 ]/dt = 4fc[CH3ReO2 ][C1O;] (16) where 4 it = 29 M"1 s"1; the subsequent chlorate reduc- tion is about 1000 times faster. The reaction rate for equation 15 is not dependent on acid concentration, which is unusual. The rate is unusually fast, several orders of magnitude faster than for any other transition metal compound. Abu-Omar and Espenson (1995) provide a very convenient and useful table of rate constants for the reduction of perchlorate and chlorate by transition metal complexes. There are two serious problems with chemical reductants: (1) they tend to suffer from oxidation by atmospheric oxygen, and (2) they are too slow under normal conditions (i.e., pH and concentration). Conse- quently, anaerobic conditions would be required for their delivery as well as for the duration of their reac- tion time. The electrochemical reduction of perchlorate ion has been reported for a wide variety of cathodes, in- cluding platinum (Hordnyi and Vertes, 1975b; Vasina and Petrii, 1969), tungsten carbide (Hordnyi and Vertes, 1974, 1975a; Vertes and Horanyi, 1974), ruthenium (Gonzales Tejera and Colom Polo, 1984), carbon doped with chromium(III) oxide or aluminum oxide (Mouhandess et al., 1980), aluminum (Kiss et al., 1973), and titanium (Mathieu and Landolt, 1978; Mathieu et al., 1978; Tsinman et al., 1975). Brown (1986) showed that chlorate and perchlorate may be reduced success- fully to chloride by active titanium, and he discusses the possibility of passivation of the titanium electrode as titanium(IV) forms, presumably from the deposition of titanium dioxide. Given the exceptionally fast reaction between methylrhenium dioxide and perchlorate, it seems pos- sible that alkyl rhenium oxides might be used catalyti- cally if not chemically. Moreover, this suggests an entire area of research into the use of organometallic compounds as chemical reductants for perchlorate. Besides the possible application to remediation, this also suggests possibilities for use in kinetic methods of analysis since anyone with an ultraviolet-visible spec- trophotometer could follow this reaction in as low as submillimolar levels of perchlorate based on the molar absorptivity of CH3ReO3. Quantitative Analytical Chemistry Gravimetry The strongly dissociative property of perchlorate, which makes it so highly desirable in the uses described, makes it correspondingly difficult to quantitate by gravimetry or to remove by precipitation. However, some insoluble or sparingly soluble perchlorate com- pounds are known. The first methods reported for 86 Urbansky ------- quantitation of perchlorate were based on gravimetry with nitron (DOC, 1996; Harris, 1991; Welcher, 1947), C20HI6N4, Mr = 312.36 g moH. Other names for nitron include 4,5-dihydro-2,4-diphenyl-5-(phenyl-imino)- lH-l,2,4-triazolium hydroxide, inner salt; 1,4-diphe- nyl-3-(phenylimino)-1,2,4-triazolidine, mesoionic didehydroderivative; 1,4-diphenyl-3,5-endanilodihy- drotriazol; and 3,5,6-triphenyl-2,3,5,6-tetraaza- bicyclo[2.2.1]hex-l-ene. As might be expected, nitron also quantitatively precipitates BFj, WO^", ReOj, and NOj as well as a few other anions. In fact, nitron takes its name from its ability to precipitate nitrate anion. It is important to note that the precipitates contain the acidic hydrogen ion, which protonates the positive nitrogen of the inner ring; however, there is no significant association be- tween the proton and the inorganic anion. The anion is simply a large counterion to balance the charge. Table 3 gives the solubilities reported for some nitron salts. In addition to the anions in Table 3, other large anions should be expected to interfere; interference is docu- mented for ferrocyanide, ferricyanide, picrate, and oxalate (Welcher, 1947). Perchlorate may be assayed gravimetrically using tetraphenylarsonium chloride, AsPh4Cl (Ph = C6H5) (Harris, 1991). While this material may be very useful in the analysis of perchlorate, it is doubtful that it will establish itself in any way for remediation because of the arsenic. The low solubilities of Nit • HC1O4 (Nit = nitron) and AsPh4ClO4 immediately suggest several opportu- nities for both analysis and remediation. Besides tradi- tional gravimetric analyses, this property may be ex- ploited for electrochemical analyses. As might be expected, potentiometric titrimetry with nitron has been used for perchlorate and other anions (Selig, 1988). An indirect method has been used to determine as little perchlorate as 0.05 to 0.15 nmol using nitron followed Table 3. species.3 Solubilities of some nitron • HX HX Solubility, g L~1 HCI04 HN03 HI HSCN HCrO4H HCIO3 MONO HBr 0.08 0.099 0.17 0.4 0.6 1.2 1.9 6.1 From Welcher, 1947. by an iodimetric back-titration (Shahine and Ismael, 1976). Potentiometry and Ion-Selective Electrodes (ISEs) One technique that holds excellent promise as a rou- tine monitoring device is potentiometric measurement via an ion-selective electrode. The perchlorate ion- selective electrode (ISE) has an extensive history and has been under development for about 20 years. Many designs and components abound, some with outstand- ing response. A liquid-membrane ISE has been shown to have nearly Nernstian response to perchlorate in the range of 10~5 to 10"2 M with a lower limit of detection of -10-6 M (Hassan and Elsaied, 1986). A polyvinyl chloride (PVC) membrane impregnated with HNitSCN has been used to determine perchlorate down to 2.5 x 10"5 M, and one impregnated with As(C6H5)4SCN has been constructed, but it was not tested for response to C1O4 (Elmosalamy et al., 1987). Perchlorate, along with several other anions, has been determined using flow injection analysis with a carbon electrode and &M(diphenylphosphino)propane-copper complex as an ion exchanger (Wang and Kamata, 1992). Making use of the low solubility of potassium perchlorate, a potas- sium cation ISE was used to study the migration of perchlorate into a PVC membrane (Verpoorte and Harrison, 1992). Perchlorate ISEs based on a barium complex with a macrocyclic Schiff base have been developed (Masuda et al., 1991). The newer perchlo- rate ISEs are based on large, inert, metal-ligand com- plexes that do not undergo complexation with small, hard Lewis bases and have no open coordination sites. The perchlorate ISE has already established itself as a research tool, and it has been used to monitor perchlorate concentration in a variety of investigations (Alegret et al., 1986; Baczuk and Dubois, 1968; Cakrt et al., 1976; Ciavatta et al., 1989; Efstathiou and Hadjiioannou, 1977a, 1977b; Fogg et al., 1977; Hiiro et al., 1979; Hopirtean and Stefaniga, 1978; Hopirtean et al., 1976, 1977; Hseu and Rechnitz, 1968; Lnato et al., 1980; Ishibashi and Kohara, 1971; Ishibashi et al., 1973; Jain et al., 1987; James et al., 1972; Jasim, 1979; Jyo et al., 1977, 1983; Kataoka and Kambara, 1976; Manahan et al., 1970; Matei et al., 1986; Nikolskii et al., 1977; Pathan and Fogg, 1974; Rohn and Guilbault, 1974; Sharp, 1972; Silber and Zhang, 1991; Sykut et al., 1979; Tateda et al., 1970; Vosta and Havel, 1973a, 1973b; Wilson, 1979; Wilson and Pool, 1976). Based on the significant advances in perchlorate ISEs, the time is ripe for exploration of perchlorate ISEs as a technique both for first-line assay in the field and for Perchlorate Chemistry: Implications for Analysis and Remediation 87 ------- continuous monitoring within a treatment facility, es- pecially in regions with known contamination. Ion Chromatography and Capillary Electrophoresis Perchlorate ion is commonly described as strongly retained on anion exchange resins. What this really means is that other ions are more strongly attracted to the aqueous mobile phase. It does not mean that the resin has a high affinity for perchlorate. Chloride and hydroxide have much higher charge densities than perchlorate and therefore associate more strongly with water than they do with a fairly diffuse quaternary ammonium site. If eluent components are not chosen wisely, perchlorate elution times can run over an hour, allowing substantial diffusion and thus peak broaden- ing. The California Department of Health Services has established an ion chromatography (1C) method that usesp-cyanophenoxide to displace the perchlorate from the resin (Cal DHS, 1997b). Recently, Maurino and Minero (1997) showed that hydrogen cyanurate ions, HzA" and HA2" (cyanuric acid = H3A), can be used to effect excellent separation and peak shape for perchlo- rate. Biesaga et al. (1997) showed that phthalate can be used, but retention times are long (40 min) and some degradation of peak shape is observed. Jandik et al. (1990) used acetonitrile to modify the eluent dielectric constant and a solid-phase reagent to obtain retention times under 20 min. In a study of retention time and polarizability, Daignault et al. (1990) used 1.7 mM NaHCO3-1.8mMNa2CO3 eluent for perchlorate; how- ever, they used analyte concentrations of 2 mg mL"1. Buchberger and Haider (1997) used 1C with particle beam mass spectrometry to detect perchlorate, provid- ing a definitive identification of the ion. An earlier study used micropacked alumina columns to separate anions and cations simultaneously, making use of the amphoteric nature of aluminum oxide (Takeuchi et al., 1988). Wirt et al. (1998) reported a new 1C method using only NaOH as the eluent with a retention time of <10 min. One of the advantages of capillary electrophoresis (CE) over 1C is readily apparent when considering strongly retained ions, such as perchlorate. In CE sepa- ration, the electrophoretic (ionic) mobility is the most important factor, unlike 1C where interactions with the stationary and mobile phase are important. In CE, associations with the wall or a modifier are generally unintended, undesirable, and, most importantly, avoid- able. Furthermore, a slow eluter when using 1C may be a very fast one when using CE. Avdalovic et al. (1993) showed that perchlorate can be eluted by CE at 10 min; meanwhile, bromide, chloride, and iodide all take more than 15 min by their method. Hauser et al. (1995) used CE with a micro-ISE to quantitate as little as 10 |OM perchlorate. Corr and Anacleto (1996) used CE coupled with mass spectrometry with ion spray introduction on a wide variety of cations and anions. Although it is common practice to use quaternary ammonium salts, such as cetyltrimethylammonium chloride, to reverse the electroosmotic flow, Krokhin et al. (1997) used water-soluble polymers with quaternary ammonium moieties to promote the elution of perchlorate by CE. AFRL and EPA's National Exposure Research Laboratory (NERL) have begun the process of interlaboratory validation for the California DHS 1C method (Tsui and Pia, 1998). There currently is no EPA-approved method for the quantitation of perchlo- rate in drinking water; however, Cal DHS has estab- lished its own approval process for laboratories that seek to analyze for perchlorate. AFRL continues to work on methods for determining perchlorate in other media, e.g., soils and plant tissues. Other Techniques In addition to gravimetry and electrochemistry, other techniques also offer promise when coupled with these reagents, particularly spectrophotometry. Methylene blue forms an insoluble complex with perchlorate; the loss of methylene blue from the supernatant is determined spectrophotometrically (Nabar and Ramachandran, 1959). The neocuproine-cuprous ion complex also has been used to extract perchlorate into ethyl acetate; the fcz'.s(neocuproine)cuprous perchlorate species has a visible absorption spectrum (Xmax = 456 nm) or, alternatively, the copper content may be deter- mined by atomic absorption photometry (A, = 324.7 nm) (Collinson and Boltz, 1968). A similar method designed for biological fluid samples makes use of an anion exchange resin (Amberlite IR-45) to pre- concentrate the perchlorate and thereby lower (im- prove) the quantitation limit (Weiss and Stanbury, 1972). Both tetraphenylarsonium chloride and nitron can be used to determine perchlorate concentration spec- trophotometrically by difference in samples contain- ing nitrate (Shahine and Khamis, 1979). An aliquot of nitron solution is added to the sample, and the excess nitron is determined photometrically as (HNit)2 [Co(NCS)4] at 625 nm; perchlorate alone is deter- mined by precipitation with AsPh^, and the excess is determined as (AsPh4)2[Co(NCS)4] at 620 nm. Al- though not reported, it seems reasonable that laser- 88 Urbansky ------- induced fluorescence would be be a sound technique, especially if the HNitClO4 or AsPh4ClO4 were ex- tracted into an organic solvent. Reverse-phase high-performance liquid chroma- tography also could be used on such an extract to separate perchlorate from interferents such as nitrate or bromide with detection by ultraviolet (UV) absor- bance. The use of a photodiode array detector may provide a UV absorption spectrum sufficiently distinct to ensure definitive identification of perchlorate, elimi- nating the retention time problem. It is known that AsPh4ClO4 has a unique infrared absorption spectrum that permits it to be distinguished from other tetraphenylarsonium salts. Numerous methods using wet chemical or instrumental techniques or a combina- tion thereof are known, and these have been reviewed extensively for the literature prior to 1979 (Schilt, 1979). Although mulls in Nujol or perfluorosilicone grease may be safe, compressing AsPh4ClO4 in a KBr pellet seems to be flirting with disaster. Even if KBr pellets can be made safely, there is no guarantee that some decomposition of the perchlorate will not occur or that a reaction with bromide will not take place during pressurization. Walter Selig at Lawrence Livermore Laboratory has investigated a number of potentiometric precipita- tion titrations to determine perchlorate using quater- nary ammonium compounds as titrants (Selig, 1977, 1979, 1980a, 1980b). A carbon paste electrode that uses thallium for catalysis has been used to voltammetrically quantitate perchlorate in drinking water samples down to 50 ng mL~!; however, it suffers from a number of significant direct and indirect interferences (Neuhold et al., 1996). Of these, bromide, chlorate, and nitrate are most likely to be found in a drinking water matrix. Remediation and Treatment Overview It is helpful to keep in mind the following criteria for any drinking water treatment technology. The treat- ment must not (1) adversely affect other treatment technologies for regulatory compliance, (2) produce water that corrodes the distribution system, (3) pro- duce water that is unpalatable, (4) suffer from degra- dation by other components in the water, (5) fail to perform reliably, (6) produce excessive waste, or (7) fail to meet time and expense constraints. The best choice for any situation will require a careful evalua- tion of options and probably some combination of techniques. We must remember that the potential for success of any technology is dependent on two factors: the establishment of a safe level of perchlorate for drinking water and a quantitative chemical analysis that ensures this safe level is in fact achieved. Remediation and Treatment by Physical Processes Membrane-Based Techniques. Membrane-based techniques can be effective, but they suffer from sev- eral drawbacks. While reverse osmosis (RO) would effect sufficient remediation, it can be impractical for a municipal treatment system because of the fouling of membranes and the associated cost. RO-treated water has to be remineralized with sodium chloride, sodium bicarbonate, and other innocuous salts to prevent deg- radation of the distribution system and to make the water palatable, since deionized water generally is considered to have an unpleasant taste. Therefore, as long as sufficient salts are taken in from food and other sources, consumption of deionized water is not likely to pose a threat to the normal electrolyte balance. As with RO, electrodialysis also might be used in this fashion. These two techniques are probably best suited for point-of-use or small systems. Anion Exchange. Although perchlorate ion is strongly retained by quaternary ammonium resins, the crux of the matter is its initially low concentration in most cases. For example, it might be necessary to reduce perchlorate concentration from 1 jig mL-1 to 20 ng mL-1. However, consider that bicarbonate, carbon- ate, chloride, and a host of other anions are all likely to be present at much higher concentrations. Assuming that a chloride-form resin is used, the presence of phosphates, carbonates, and sulfate remains an issue. Although it may be possible to produce a resin salt that matches the proportions of the major anions in the influent water, to do so would be extremely inconve- nient. In addition, the low concentration of perchlorate in the raw water substantially reduces the driving force for its removal. In other words, to adequately remove the perchlorate may require essentially demineralizing and remineralizing the water, depending on its anion content. It is possible to modify resins so as to improve their selectivity for particular anions. Kawasaki et al. (1993) have used Dowex 1X-8 to selectively precon- centrate perchlorate; the selectivity of the resin for perchlorate is about 100 times that for chloride and 10 times that for nitrate. In addition to selectivity in a thermodynamic sense, there is the matter of rapid equili- bration and anion exchange. If the rate of exchange is too slow, a resin will not be usable no matter how high its selectivity. The U.S. Department of Energy has Perchlorate Chemistry: Implications for Analysis and Remediation 89 ------- developed a mixed triethylammonium-trihexylammo- nium resin that is capable of removing pertechnetate down to the parts-per-trillion level (Brown, 1997). Tethered triphenylarsonium or phosphonium moi- eties or a tethered (through a phenyl group) nitron might work in an anion-exchange resin to selectively preconcentrate perchlorate as a step in remediation. The disadvantage of the tethered triphenylarsonium group is that normal degradation of the resin would lead to the release of arsenic into the treated water. Although the health effects of nitron are unknown, it would be expected to undergo biodegradation; further- more, it would be destroyed readily by UV irradiation (A,£ 185 nm), whereas arsenic would remain as an inorganic oxyanion even if the organic portion of the species were destroyed. Precipitation. The low solubility of the HNitClO4 ion pair reveals a strong association between the pro- tonated nitron cation and the perchlorate anion. All insoluble ion pairs and complexes exist at some level in solution. It may be possible to exploit this pairing for purposes of remediation. If the addition of nitron to perchlorate-containing waters results in formation of the soluble ion pair, it may be possible to subsequently induce an intramolecular reaction in which both the perchlorate and the nitron are destroyed. Photoactivation of the perchlorate by UV or laser irradiation may promote an intramolecular redox reaction (probably by oxygen atom transfer). The proximity of the HNif- ClOj ion pair within a solvent (water) cage means that it is not necessary to form an encounter complex. In addition, the local concentration of the two species is very high within the solvent cage. This should help to reduce the effects of the perchlorate kinetic barrier (discussed below). Of course, irradiation with UV light also will promote destruction of the nitron by hydroxyl radical formation. Ideally, the largest possible wave- length (lowest frequency and energy) light would be used to reduce side reactions that would destroy the nitron. Unfortunately, nitron has potential to remediate only those sites with very high perchlorate concentra- tions unless it can be synthesized more cheaply. At present, nitron is about 52 times more expensive than an equal mass of reagent-grade sodium chloride. How- ever, should a method involving nitron prove effec- tive, bulk synthesis of the material would likely drop the cost by 40 to 50%, and use of a less refined tech- nical (rather than reagent) grade would probably re- duce it by another 10 to 25%. At some of the sites where the perchorate concen- tration is 0.037 M, nitron could readily be used as a precipitant since the nitron-hydrogen perchlorate salt has a solubility of only 0.19 mM. Although the action level of 18 ng mL"1 corresponds to 0.18 M, a level of 0.19 mM is certainly preferable to 37 mM. Of course, one drawback is that a source of acid (usually 5% acetic acid) must be present. On the other hand, vin- egar is probably preferable to 0.037 M ClOj. More- over, such post-remediation acetic acid and acetate would be biodegradable. In addition to cost, all physical separation pro- cesses have one major problem: waste disposal. Pre- sumably, the regenerant from ion exchange and the concentrate from RO or electrodialysis would contain perchlorate at concentrations too high to be released into a sewage system. This waste presents a problem in terms of cost and post-treatment needs. Although these techniques take the perchlorate out, they concentrate it somewhere else where it must be dealt with later. Remediation and Treatment by Chemical Processes Chemical and Electrochemical Reduction. Here we refer to reduction specifically in the redox sense of adding electrons. From the description of the oxida- tion-reduction reactions of perchlorate above, it is clear that chemical reduction will play no role in drinking water treatment in the near future. Chemical reduction is simply too. slow. Unless safe new catalysts become available, this appears unlikely to change. Common- place reductants (e.g., iron metal; thiosulfate, sulfite, iodide, and ferrous ions) do not react at any observable rate, and the more reactive species are too toxic (and still too sluggish). In addition, any reductant will nec- essarily have oxidized by-products. The toxicity of the by-products must be considered. There is more hope for electrochemical reduction. A decided advantage of electrochemical reduction is the large amount of control over kinetics that results from control of the operating potential. Electrode re- duction kinetics reasonably can be viewed as being limited by three factors: (1) diffusion of the ions to the electrode surface, (2) association with the electrode surface, and (3) activation past the overpotential re- quired to reach the transition state. Although overpotential usually is the greatest barrier, it also is the one that can be dealt with best. Because we are not concerned with other reductions (including reducing water to hydrogen), the only barrier is the limit of a negative potential that is practical and safe to apply. Fortunately, most of the materials in raw water are reducing agents. Although some may be affected by electroreduction, this probably does not present a sig- nificant obstacle. To date, this option has not been explored for low-concentration treatment at anything 90 Urbansky ------- approaching pilot scale. Although electrochemical tech- nologies are well established for other industries (e.g., electroplating of metals, electrolysis of brine), they have not yet found a place in drinking water treatment. In this category, it appears that the most success- ful strategies for remediating perchlorate contamina- tion will utilize metal cation-catalyzed reduction by either chemical or electrochemical means. Several metal chelates have potential at this point, especially if em- bedded in an electrode for use in electrochemical re- duction. Biological and Biochemical Techniques. Bio- remediation is another matter entirely, and it may prove to be the most practical approach. A number of bacte- ria that contain nitrate reductases (Payne, 1973) are capable of reducing perchlorate (Schilt, 1979). Staphy- lococcus epidermidis is capable of reducing perchlor- ate in the absence of nitrate. Cell-free extracts of ni- trate-adapted Bacillus cereus also reduce perchlorate (and chlorate) (Hackenthal, 1965). As would be ex- pected, sodium perchlorate, especially in higher con- centrations, has been shown to be toxic to several species of bacteria. Unfortunately, S. epidermidis is also pathogenic; it increasingly is encountered as a source of nosocomial infections, especially opportu- nistic infection with in-dwelling intravenous or uri- nary catheters (Archer, 1995). It is encountered with other medical apparatus such as prosthetic joints, pace- makers, heart valves, and breast implants (Archer, 1995). Like S. epidermidis, B. cereus is pathogenic. B. cereus is known for food poisoning, ocular infections, and pneumonia with other sites sometimes affected (Tuazon, 1995). Rikken et al. (1996) reported that perchlorate and chlorate are reduced to chloride by Proteobacteria with acetate as a nutrient (reductant) at near-neutral pH. While they did show loss of perchlorate and chlo- rate, their mechanisms failed to include contributions from uncatalyzed reactions. Specifically, they con- cluded that a dismutase is responsible for all elimina- tion of toxic chlorite from the cell, catalyzing its dis- proportionation to dioxygen and chloride. However, the uncatalyzed disproportionation of chlorite to chlo- ride and chlorate is not necessarily negligible. Korenkov et al. (1976) patented Vibrio dechloraticans Cuzensove B-1168 for perchlorate reduction; V. dechloraticans is nonsporulating, motile, and gram negative. Malmqvist et al. (1994) showed that Ideonella dechloratans can reduce chlorate, but they did not test for perchlorate reduction. Over the past 8 years, work at AFRL has shown that perchlorate is metabolized by Wolinella succinogenes, strain HAP-1 (U.S. Air Force, 1994; Wallace and Attaway, 1994). W. succinogenes is ca- pable of using either chlorate or perchlorate to oxidize Brewer's yeast. Pilot-scale systems at Tyndall AFB, Florida showed that perchlorate levels could be re- duced from 3000 |J,g mL-1 to below 0.5 |j,g mlr1 (Hurley et al., 1997). HAP-1 was first isolated from a munici- pal anaerobic digester. The bacterium is an antibiotic- resistant, nonsporulating, motile, Gram-negative, obli- gately anaerobic bacillus (Wallace et al., 1996). This sort of remediation may be effective at a site where perchlorate concentrations in water are high, but it very likely would be impractical for the treatment of drinking water unless it can be demonstrated to reach even lower perchlorate concentrations. AFRL's efforts have led to the implementation of a production-scale bioreactor in Utah for meeting perchlorate discharge requirements (Hurley, 1998). Very little research has been done on perchlorate reductases. It may be possible to isolate these from bacteria and use them directly as reagents without the parent organisms. The mechanisms of these catalysts are not well understood, and the reductases themselves have not been well characterized. It may be possible to synthesize an analogous catalyst based on the reduc- tase, but only if the fundamental bioinorganic chemis- try is understood. Although nitrate reductases are based on molybdenum (Coughlan, 1980), it has not been verified whether this is also true for the perchlorate reductases. Several projects are ongoing in the affected areas of EPA's Region 9 and were described at a recent meeting. Catts (1998) reported that a pilot-scale bioreactor has been constructed for the Baldwin Park Operable Unit in California using microbes derived from the food-processing industry. Operation of this pilot unit over a period of several months showed that perchlorate and nitrate could be reduced to undetect- able levels, i.e., [ClOj] < 4 ng mL-1. Ethanol was used as a food source and minerals were added to the sys- tem. The perchlorate-reducing microbes were not iso- lated or characterized. Sase (1998) reported on a spe- cially designed anion exchange system with alternately regenerating columns that is undergoing testing by the Main San Gabriel Basin Watermaster. 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