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
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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.
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Urbansky
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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
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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.
Conclusions
Bioremediation and biological or biochemical treat-
ment appear to be the most economically feasible,
fastest, and easiest means of dealing with perchlorate-
laden waters at all concentrations. Although other tech-
niques may find application to select systems, e.g.,
point-of-use or small utilities, it appears that biological
and biochemical approaches will play the greatest role
Perchlorate Chemistry: Implications lor Analysis and Remediation
91
-------
in solving the perchlorate problem. Some situations
may require a combination of technologies to best
meet unique needs. This is a complex problem, and
many of the standard technologies that have domi-
nated the drinking water industry for the past several
decades will not work for this contaminant when used
in the conventional ways. Many of the possibly effec-
tive technologies have not been applied to drinking
water specifically, and the interplay with other treat-
ment technologies required for regulation compliance
must be assessed. In addition to rapid implementation
of effective and workable technologies, ongoing de-
velopment will be required to find new technologies
and to make them affordable and assimilable into the
industry.
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