Reduction of haloacetic acids by Fe0: Implications for treatment and fate

Article abstract of DOI:10.1021/es001785b

To predict the fate of haloacetic acids (HAAs) in natural or engineered systems, information is needed concerning the types of reactions that these compounds undergo, the rates of those reactions, and the products that are formed. Given that many drinking water distribution systems consist of unlined cast iron pipe, reactions of HAAs with elemental iron (Fe⁰) may play a role in determining the fate of HAAs in these systems. In addition, zerovalent iron may prove to be an effective treatment technology for the removal of HAAs from chlorinated drinking water and wastewater. Thus, batch experiments were used to investigate reactions of four trihaloacetic acids, trichloroacetic acid (TCAA), tribromoacetic acid (TBAA), chlorodibromoacetic acid (CDBAA), and bromodichloroacetic acid (BDCAA), with Fe⁰. All compounds readily reacted with Fe⁰, and investigation of product formation and subsequent disappearance revealed that the reactions proceeded via sequential hydrogenolysis. Bromine was preferentially removed over chlorine, and TBAA was the only compound completely dehalogenated to acetic acid. In compounds containing chlorine, the final product of reactions with Fe⁰ was monochloroacetic acid. Halogen mass balances were 95-112%, and carbon mass balances were 62.6-112%. The pseudo-first-order rate constants for trihaloacetic acid degradation were as follows: BDCAA (10.6 ± 3.1 h^-1) > CDBAA (1.43 ± 0.32 h^-1) ≈ TBAA (1.41 ± 0.28 h-^-1) ? TCAA (0.08 ± 0.02 h^-1).

Full text of DOI:10.1021/es001785b

Environ. Sci. Technol. 2001, 35, 2258-2263
Reduction of Haloacetic Acids by
TABLE 1. List of the Nine HAAs of Concern in Treated Drinking
Water and the MDL and RSD Values Obtained for the CE
0
Fe : Implications for Treatment and Analysis Method
MDLa RSDb
Fate
name
formula
abbreviation (µM) (%)
m onochloroacetic acid
CH
CH
2
ClCOOH
2
BrCOOH
MCAA
MBAA
DCAA
11.0 1.74
5.76 2.91
5.36 2.06
3.94 0.49
5.77 0.21
6.52 1.08
4.93 3.20
7.25 2.42
4.27 1.45
R A YM O N D M . H O Z A L S K I , *
m onobrom oacetic acid
L I Z H A N G , A N D W I L L I A M A . A R N O L D
dichloroacetic acid
2
CHCl COOH
CHBrClCOOH BCAA
CHBr COOH
3
CCl COOH
Department of Civil Engineering, University of Minnesota,
Minneapolis, Minnesota 55455-0116
brom ochloroacetic acid
dibrom oacetic acid
trichloroacetic acid
chlorodibrom oacetic acid CClBr
brom odichloroacetic acid CBrCl
2
DBAA
TCAA
CDBAA
BDCAA
TBAA
2
COOH
COOH
2
tribrom oacetic acid
3
CBr COOH
To predict the fate of haloacetic acids (HAAs) in natural
or engineered systems, information is needed concerning
the types of reactions that these compounds undergo,
the rates of those reactions, and the products that are
formed. Given that many drinking water distribution systems
consist of unlined cast iron pipe, reactions of HAAs with
a
MDL: m ethod detection lim it. b RSD: relative standard deviation
determ ined from four m easurem ents at
proxim ately 40 ppm .
a concentration of ap-
In addition to their occurrence in chlorinated drinking
water supplies, in wastewaters, and in swimming pool water
0
elemental iron (Fe ) may play a role in determining the fate
(
5), chloroacetic acids (TCAA, DCAA, MCAA) are widespread
of HAAs in these systems. In addition, zerovalent iron
may prove to be an effective treatment technology for the
removal of HAAs from chlorinated drinking water and
wastewater. Thus, batch experiments were used to investigate
reactions of four trihaloacetic acids, trichloroacetic acid
in the environment (6-8). Chloroacetic acids are plant growth
inhibitors that were used as herbicides until the late 1980s
(
6). Chlorinated HAAs are believed to be formed in the
atmosphere during the photochemical degradation of chlo-
rinated solvents (9) and have been detected as byproducts
of the oxidation of chlorinated ethylenes using the photo-
catalyst titanium dioxide (10). It appears that the chlorinated
HAAs also occur naturally (11). Other HAAs, such as the
trifluoroacetic acid and chlorodifluoroacetic acid produced
upon atmospheric degradation of hydrochlorofluorocarbons
and hydrofluorocarbons (6, 12), are also of environmental
concern. Because of their widespread occurrence, phyto-
toxicity, and suspected human carcinogencity, information
regarding the reactivity of these compounds is needed to
develop treatment methods and to predict their fate in
engineered and natural environments.
(TCAA), tribromoacetic acid (TBAA), chlorodibromoacetic
acid (CDBAA), and bromodichloroacetic acid (BDCAA), with
0
0
Fe . All compounds readily reacted with Fe , and
investigation of product formation and subsequent
disappearance revealed that the reactions proceeded via
sequential hydrogenolysis. Bromine was preferentially
removed over chlorine, and TBAA was the only compound
completely dehalogenated to acetic acid. In compounds
0
containing chlorine, the final product of reactions with Fe
was monochloroacetic acid. Halogen mass balances
were 95-112%, and carbon mass balances were 62.6-
Unfortunately, published research concerning the biotic
and abiotic reactivity of HAAs in aquatic environments is
very limited. A few studies report the biodegradation of TCAA
(13, 14). Yu and Welander (14) discovered an aerobic
bacterium that was able to grow on TCAA as the sole carbon
and energy source. De Wever et al. (13) isolated an anaerobic
bacterium that was capable of growth via reductive deha-
logenation of TCAA with acetate as the electron donor. The
organism, named Trichlorobacter thiogenes, was unable to
dehalogenate DCAA or MCAA, and DCAA was the ultimate
product of the biotransformation of TCAA. In addition,
Boethling and Alexander (15) demonstrated that MCAA can
1
12%. The pseudo-first-order rate constants for trihaloacetic
acid degradation were as follows: BDCAA (10.6 ( 3.1
-
1
-1
-1
h ) > CDBAA (1.43 ( 0.32 h ) ≈ TBAA (1.41 ( 0.28 h )
-
1
.
TCAA (0.08 ( 0.02 h ).
Introduction
Haloacetic acids (HAAs) are important byproducts of the
chlorination of water (1) and wastewater (2). Chlorinated
HAAs are formed when hypochlorous acid (HOCl) reacts with
organic matter that is ubiquitous in water supplies and
wastewaters. In addition, brominated HAAs are formed when
free chlorine reacts with bromide to form hypobromous acid
be mineralized to CO
streamwater.
2
by natural microbial communities in
There is little or no evidence of abiotic transformations
of HAAs in the literature. In particular, there are no reported
investigations of the reactivity of HAAs with zerovalent iron
(
HOBr), which then reacts with the organic matter in a fashion
similar to HOCl. The nine possible chlorinated, brominated,
and combined chlorinated-brominated HAAs are listed in
Table 1. Some of these compounds are suspected carcinogens
and pose a potential health risk to water consumers (3). The
United States Environmental Protection Agency began
regulating HAAs in drinking water as part of the Disinfectants/
Disinfection By-Products (D/ DBP) Rule promulgated on
December 16, 1998 (4). This rule established a maximum
contaminant level (MCL) for the sum of five HAAs (HAA5 )
MCAA + DCAA + TCAA + MBAA + DBAA) of 60 µg/ L (4).
0
0
(
Fe ). Reactions of HAAs with Fe may have important
implications for understanding the fate of HAAs in water
distribution systems, which primarily consist of cast iron
piping, and for development of HAA treatment systems. This
research was performed to evaluate the reactivity of HAAs
0
with Fe .
Numerous halogenated solvents are susceptible to re-
duction by zerovalent iron including chlorinated alkanes (e.g.,
methanes and ethanes) (16-29), alkenes (19, 30-49), alkynes
(
48), and aromatic compounds, such as chlorinated phenols
*
To whom correspondence should be addressed. E-mail:
hozal001@tc.umn.edu; phone: (612) 626-9650; fax: (612) 626-7750.
(50) and polychlorinated biphenyls (51). Brominated com-
2
2 5 8
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ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 35, NO. 11, 2001
10.1021/es001785b CCC: $20.00

2001 Am erican Chem ical Society
Published on Web 04/27/2001

pounds, such as 1,2-dibromoethane (52) and 1,2-dibromo-
0
3
-chloropropane (53), are also reduced by Fe . For the alkanes
and alkenes, it has been demonstrated that the reduction
occurs via competing pathways: hydrogenolysis and reduc-
tive elimination (18, 39, 40, 48). Hydrogenolysis is the
replacement of a halogen by hydrogen. Reductive â-elimina-
tion may occur when halogens are present on adjacent
carbons (39, 40, 48). When a carbon is multiply halogenated,
reductive R-elimination may occur, resulting in a carbene-
(
oid) intermediate (18, 48).
Previous work with 1,1,1-trichloroethane has shown that
0
FIGURE 1. CE electropherogram from the analysis of a sample
containing 10 ppm of each of the nine possible chlorinated and
brominated HAAs and the internal standard trifluoroacetic acid
reduction by Fe occurs by a combination of hydrogenolysis
and reductive R-elimination (18). Given the similar structure
of the trihaloacetic acids to 1,1,1-trichloroethane, we hy-
pothesized that similar pathways may be important for the
HAAs. The specific research objectives of this study included
the following: (1) identification of the predominant reaction
pathways of the trihalogenated acetic acids (i.e., TCAA, TBAA,
(
TFAA) in DI water. Peak identification is as follows: 1, TFAA; 2,
MCAA; 3, MBAA; 4, DCAA; 5, BCAA; 6, DBAA + TCAA; 7, BDCAA;
, CDBAA; and 9, TBAA.
8
0
process. The 1-methyl-3-nitro-1-nitrosoguanidine (MNNG)
used to prepare the diazomethane is a potent carcinogen
CDBAA, and BDCAA) with Fe , and (2) preliminary deter-
mination of the rates of reaction of these species.
(
55). The diazomethane itself also represents a potential
explosion hazard. There is a great deal of concern, therefore,
about working with MNNG and diazomethane (56). We
speculate that it is largely due to the time and difficulty of
the current HAA analysis methods that little research on the
fate of HAAs in aquatic systems has been performed.
Experimental Section
Materials. Aqueous HAA stock solutions were prepared by
diluting neat chemical in distilled and deionized (DI) water
(
Milli-RO 12, Millipore). TCAA and TBAA were obtained
from Sigma-Aldrich, and BDCAA and CDBAA were obtained
from Supelco. The purity of the neat chemicals used in
these experiments is as follows: TCAA (99+%), BDCAA (99%),
CDBAA (99%), and TBAA (99%). The chemicals were used
as received. To prepare calibration standards for all nine of
the HAAs (Table 1), 1000 ppm standard solutions of HAA in
methyl tert-butyl ether (MTBE) obtained from Supelco
were dissolved in DI water and diluted to the desired
concentration.
Batch Experim ents. The batch experiments were per-
formed by diluting a freshly prepared 100 mg/ L trihaloge-
nated acetic acid stock solution with deoxygenated DI water
to the desired concentration (100-200 µM) and adding 36
mL of the resulting solution to a 38 mL glass serum bottle
To overcome the limitations of current HAA analysis
methods, a new method of analyzing for HAAs using CE was
developed in our laboratory. This new method is safe and
rapid, so that experiments on HAA fate could be readily
performed with frequent sampling and adequate replicates.
Our method is similar to other CE methods recently described
in the peer reviewed literature (5, 57).
HAA Analysis Using CE. Analyses of HAAs, chloride, and
bromide were performed using a Hewlett-Packard capillary
electrophoresis ion analyzer (model HP 3DCE) controlled with
HP Chemstation software. A 75 µm inner diameter by 56 cm
(
effective length) fused silica capillary was used for the
separation. The separation buffer consisted of phosphate
(6.25 mM NaH PO , 6.25 mM Na HPO ) with 5.0 ppm of the
0
along with 0.5 g of electrolytic Fe (100 mesh size; Fisher
2
4
2
4
0
cationic surfactant hexadimethrine hydroxide (HDMOH) as
electroosmotic flow modifier. The buffer pH was adjusted to
Scientific). The Fe was used as received. Although the
electrolytic iron may not be representative of cast iron
distribution system piping, it has been demonstrated that
this iron type allows detailed analyses of reaction pathways
and kinetics to be performed (18, 39, 40, 48). The control
reactors were prepared in a similar manner except that no
5
.7 with phosphoric acid. The HDMOH was prepared by
dissolving hexadimethrine bromide (Sigma-Aldrich) in DI
water and passing the solution through an anion exchange
cartridge (Maxi-clean ICOH plus, Alltech Associates). The
buffer solution was filtered and degassed before use with a
liquid chromatography mobile phase filter/ degasser (Kontes).
Pressure injection (200 mbar‚s) was used to introduce the
sample onto the capillary. The separation voltage was -20
kV (reverse polarity), and the temperature was held constant
at 25 °C. Analytes were detected by direct ultraviolet (UV)
light absorbance at 191 nm with the capillary itself being the
detector cell. Finally, measurements of solution pH were
performed using a pH meter (Fisher Scientific).
Fe0 was added to the bottles. The reactors with Fe0 were
prepared in duplicate along with a single control bottle. All
batch bottles were filled in an anaerobic glovebag (Coy
Laboratory Products), topped with PTFE-faced butyl rubber
septa, and sealed with aluminum crimp caps. The bottles
were then removed from the glovebag and placed on a rotator
(
3
Glas-Col) such that the bottles were mixed end-over-end at
8 rpm at room temperature. Experiments were run for
periods of up to 94 h. No pH buffer was employed, leading
to a rise in pH over the course of the experiments. Initial pH
values ranged from 3.62 to 4.14 and rose to 5.60 to 6.23 by
the end of the experiments. Experiments in which pH was
monitored as a function of time revealed that much of the
pH change occurred in the first hour of incubation. The
bottles were periodically sampled by withdrawing 100 µL
from each bottle using a syringe and then filtering the solution
through a 0.45 µm pore size filter capsule (Gelman Acrodisc)
into a polyethylene autosampler vial for subsequent analysis
by capillary electrophoresis (CE).
The CE method developed in our laboratory was capable
of separating and detecting the HAAs listed in Table 1, except
for TCAA and DBAA, which comigrated and were unable to
be resolved (Figure 1). The inability to resolve TCAA and
DBAA did not create problems for the experiments in this
research because these compounds were never simulta-
neously present in a batch bottle. In addition to analysis of
the HAAs, the method was also capable of detecting the
anticipated degradation products bromide, chloride, and
acetate in the same run. It is also important to note that the
calibration curves for all nine HAAs, acetate, bromide, and
Analytical Procedures. The most widely used HAA
method, 6251B, involves liquid-liquid extraction with MTBE,
methylation with diazomethane, and analysis by GC-ECD
2
chloride were linear (R g 0.9966) over the concentration
ranges used in the experiments (e1 mM). The method
detection limits and relative standard deviation values
obtained for the nine HAAs are given in Table 1.
(
54). The entire process is time-consuming and presents
several safety concerns that arise from the methylation
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FIGURE 3. Reduction of TCAA (O) by 0.5 g of Fe(0) in 36 mL of DI
-
water. Major products are DCAA (3), MCAA (0), and Cl (]). The
dashed line represents the observed carbon mass balance, and the
solid lines are pseudo-first-order model fits.
FIGURE 2. Reduction of TBAA (O) by 0.5 g of Fe(0) in 36 mL of DI
water. Major products are DBAA (3), MBAA (0), acetate (]), and
Br (4). The dashed line represents the observed carbon mass
-
balance, and the solid lines are pseudo-first-order model fits. The
inset figure shows the disappearance of the parent compound and
the appearance of the breakdown products DBAA and MBAA during
the first 2 h of incubation.
Results and Discussion
Reaction Pathways. A plot of the degradation of TBAA in the
0
presence of Fe is shown in Figure 2. The TBAA was rapidly
consumed within the first 60 min of the experiment with the
concomitant rise of DBAA as the product. Similarly, as DBAA
was reduced, MBAA appeared, and the disappearance of
MBAA was accompanied by the production of acetate. These
results are consistent with the degradation proceeding via
sequential hydrogenolysis from TBAA to acetate. If reductive
R-elimination occurred, MBAA or acetate would have been
detected at early times during the disappearance of TBAA.
The total dehalogenation of TBAA was confirmed by the
complete bromide mass balance (98.5%) (Figure 2). The low
carbon mass balance (62.6%) in the TBAA experiment may
have been due to complexation of the terminal product
acetate with dissolved or solid-phase iron species, or both
(
22). In addition, there was no loss of TBAA in the control
bottle (result not shown), suggesting that the TBAA degrada-
tion in the iron-containing bottles was due to dehalogenation
0
by the Fe . A proposed pathway for the degradation of TBAA
in the presence of Fe is as follows:
FIGURE 4. Reduction of CDBAA (O) by 0.5 g of Fe(0) in 36 mL of DI
water. Major products are CBAA (3), MCAA (0), and Br (]).The
0
-
dashed line represents the observed carbon mass balance, and the
solid lines are pseudo-first-order model fits. The inset figure shows
the disappearance of the parent compound and the appearance of
the breakdown products CBAA and MCAA during the first 2 h of
incubation.
Similar to the reactions of TBAA, the degradation of TCAA
proceeded via sequential hydrogenolysis (Figure 3). The
reductive dehalogenation of TCAA, however, was much
slower than TBAA and also incomplete as evidenced by the
accumulation of MCAA. The incomplete dehalogenation is
also supported by the inability to detect acetate in the batch
bottles and the release of only 2 mol of chloride per mole of
TCAA. The carbon mass balance was 97.8%.
For the mixed HAAs, the first step in the degradation was
the loss of a bromide ion (Figures 4 and 5). Reduction of
CDBAA proceeded via hydrogenolysis to form BCAA, which
was subsequently reduced to MCAA (Figure 4). A plot of CE
eletropherograms as a function of time clearly illustrates the
destruction of the parent compound CDBAA, the production
of the reaction intermediates, and accumulation of final
products (Figure 6). BDCAA lost a bromide to form DCAA.
A chloride was then removed to again give MCAA as the final
product (Figure 5). Both carbon and halogen mass balances
for the mixed HAAs ranged from 93.3 to 112%. There was no
loss of either TCAA, BDCAA, or CDBAA in the control bottles
0
without Fe . The results from the control bottles were not
plotted in Figures 3-5 for clarity. Proposed pathways for the
degradation of TCAA, BDCAA, and CDBAA, which all lead to
2
2 6 0
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 35, NO. 11, 2001

FIGURE 5. Reduction of BDCAA (O) by 0.5 g of Fe(0) in 36 mL of DI
-
-
water. Major products are DCAA (3), MCAA (0), Br (]), and Cl
4). The dashed line represents the observed carbon mass balance,
(
and the solid lines are pseudo-first-order model fits. The inset figure
shows the disappearance of the parent compound and the
appearance of the first breakdown product DCAA during the first
0.25 h of incubation.
0
MCAA as the final product, in the presence of Fe are given
in Figure 7.
Unlike the structurally similar 1,1,1-trichloroethane, there
is no evidence that the trihaloacetic acids undergo reductive
R-elimination. The work of Fennelly and Roberts (18)
demonstrated that the reactions pathways of 1,1,1-trichlo-
roethane can be altered by plating a catalytic metal (Cu, Ni)
onto the iron. Reductive R-elimination was also increased
0
0
by using Zn instead of Fe . The reaction pathways of HAAs
mediated by other reductants merit further investigation.
Degradation Rates. For each trihaloacetic acid, the
concentration versus time data for the degradation of the
parent compound, the formation and decay of daughter acetic
acid species, and the production of the halide ions were
simultaneously fit using a pseudo-first-order kinetic model
for reactions in series. Rate constants were determined by
a least-squares fit of numerically integrated solutions of the
differential equations to the experimental data. In Figures
FIGURE 6. CE electropherograms as a function of time for the
reduction of CDBAA in the presence of Fe(0). The electropherograms
show the disappearance of the parent compound CDBAA, the
subsequent appearance and disappearance of the intermediate
BCAA, and the accumulation of the final products MCAA and bromide.
trihalomethanes by Cr(II) (60) and by an iron porphyrin (61).
For each of these reductants, the degradation rates for the
bromine-containing compounds were much greater than for
the trichlorinated compound. The results observed for both
the trihaloacetic acids and the trihalomethanes are consistent
with the preferential removal of bromine over chlorine.
Although we have reported fitted rate constants for the
HAAs appearing as intermediates in our experimental
systems, we cannot currently make any claims about the
relative reactivity of these species. The initial conditions for
the trihaloacetic acids were similar (e.g., pH, iron surface
oxide coating), and this allows a comparison to be made.
The intermediates were produced over different time scales
and thus were exposed to different solution and iron surface
conditions, which may lead to changes in reactivity. The
obvious example of this are the different pseudo-first-order
rate constants obtained from the model fits for the inter-
mediate DCAA produced from TCAA and BDCAA. BDCAA is
completely consumed within 15 min, leading to rapid
production of DCAA. The DCAA produced, therefore, is
degraded under conditions very similar to those at the
2
-5, the solid lines represent these model fits. In general,
the model captures the major features of the data well. The
concentrations of intermediates (e.g., DBAA from TBAA or
DCAA from TCAA) are slightly underpredicted by the model,
and the predicted decay of these species is slightly slower
than the experimentally observed decay rates. The overall
success of this reaction-in-series model, however, provides
support for the assertion that sequential hydrogenolysis is
the pathway by which reactions are occurring.
The rate constants calculated for all the starting materials
and intermediates that were detected are presented in Table
2
. Based on these pseudo-first-order degradation rate
constants, the ranking of degradation rates from fastest to
slowest for the trihaloacetic acids was as follows: BDCAA
-
1
-1
(
0
10.6 ( 3.1 h ) > CDBAA (1.43 ( 0.32 h ) ≈ TBAA (1.41 (
-
1
-1
.28 h ) . TCAA (0.08 ( 0.02 h ). The degradation of
BDCAA is so rapid that it is likely to be mass transfer limited
based on a calculation for mass transfer to suspended
particles (58, 59). The trend in these rate constants is
consistent with the trends observed for the reduction of
VOL. 35, NO. 11, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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these potentially carcinogenic compounds. Drinking water
distribution systems primarily consist of unlined cast iron
pipe. The abundance of iron metal in distribution systems
may account for some of the loss of HAAs observed in full-
scale distribution systems (64). It is expected, however, that
the corroded and passivated iron surfaces in a distribution
system would be much less reactive than the laboratory grade
electrolytic iron used in this research. Thus, reaction rates
of HAAs in the presence of cast iron filings and cast iron
piping from actual distribution systems merit investigation.
The roles of pH, temperature, and potential complexing
agents also need to be established. In addition to impacting
0
HAA fate in distribution systems, destruction of HAAs by Fe
suggests a potential approach for in-home water treatment
to eliminate these compounds from tap water. Although the
observed persistence of MCAA may be of concern, we will
investigate possible approaches for degrading this compound
such as the use of a small amount of a catalytic metal (e.g.,
palladium, copper, or nickel) in combination with iron.
Acknowledgments
The authors acknowledge the funding provided by the
American Water Works Association Research Foundation
FIGURE 7. Reaction pathways of trichloroacetic acid (TCAA),
bromodichloroacetic acid (BDCAA), and chlorodibromoacetic acid
(
Project 2644) to support the development of a method for
(CDBAA) with Fe(0). For each species, reduction occurs via
analyzing HAAs using capillary electrophoresis. We also thank
the three anonymous reviewers, whose comments and
suggestions significantly improved the manuscript.
hydrogenolysis with chloroacetic acid (MCAA) as the terminal
product.
TABLE 2. Pseudo-First-Order Rate Constants for Haloacetic
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beginning of the experiment. DCAA is produced much more
slowly from TCAA and therefore encounters different iron
surface and pH conditions than the DCAA produced from
BDCAA. These differences may explain why the rate constant
for DCAA produced from TCAA is only 40% of that for DCAA
formed via the degradation of BDCAA.
Even though the pseudo-first-order model provides a
reasonable fit to our experimental data and has often been
used to fit data obtained from zerovalent metal experiments,
the degradation rates of TBAA, CDBAA, and TCAA appeared
to undergo a subtle acceleration as the reaction progressed.
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Received for review October 18, 2000. Revised manuscript
received February 15, 2001. Accepted February 19, 2001.
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