Environ. Sci. Technol. 2009, 43, 5307–5313
Reduction of Hg(II) to Hg(0) by
Magnetite
H E A T H E R A . W I A T R O W S K I , †,|
S O U M Y A D A S , ‡,⊥ R A V I K U K K A D A P U , §
EUGENE S. ILTON,§ TAMAR BARKAY,†
A N D N A T H A N Y E E * ,‡
Department of Biochemistry and Microbiology, Rutgers
University, New Brunswick, New Jersey, Department of
Environmental Sciences, Rutgers University, New Brunswick,
New Jersey, and Pacific Northwest National Laboratory,
Richland, Washington
Received February 4, 2009. Revised manuscript received
May 8, 2009. Accepted May 20, 2009.
Mercury (Hg) is a highly toxic element, and its contamination
of groundwater presents a significant threat to terrestrial
ecosystems. Understanding the geochemical processes that
mediate mercury transformations in the subsurface is necessary
to predict its fate and transport. In this study, we investigated
the redox transformation of mercuric Hg (Hg[II]) in the
presence of the Fe(II)/Fe(III) mixed valence iron oxide mineral
magnetite. Kinetic and spectroscopic experiments were
performed to elucidate reaction rates and mechanisms. The
experimental data demonstrated that reaction of Hg(II)
with magnetite resulted in the loss of Hg(II) and the formation
of volatile elemental Hg (Hg[0]). Kinetic experiments showed
that Hg(II) reduction occurred within minutes, with reaction rates
increasing with increasing magnetite surface area (0.5 to 2 m2/
L) and solution pH (4.8 to 6.7), and decreasing with increasing
chloride concentration (10-6 to 10-2 mol/L). Mössbauer
spectroscopic analysis of reacted magnetite samples revealed
a decrease in Fe(II) content, corresponding to the oxidation
of Fe(II) to Fe(III) in the magnetite structure. X-ray photoelectron
spectroscopy detected the presence of Hg(II) on magnetite
surfaces, implying that adsorption is involved in the electron
transfer process. These results suggest that Hg(II) reaction with
solid-phase Fe(II) is a kinetically favorable pathway for Hg(II)
reduction in magnetite-bearing environmental systems.
Introduction
In the United States, mercury (Hg) associated with mixed
waste generated by nuclear weapons manufacturing has
contaminated vast areas of soil and groundwater (1, 2).
Mercury released from spills and waste disposal typically
enters the subsurface as inorganic mercury. In anoxic
sediments, mercuric Hg (Hg[II]) can be subsequently converted into the neurotoxic substance methylmercury (MeHg)
by anaerobic bacteria (3-5). Elevated levels of MeHg have
been shown to accumulate in fish inhabiting surface waters
* Corresponding author e-mail: [email protected].
†
Department of Biochemistry and Microbiology, Rutgers University.
‡
Department of Environmental Sciences, Rutgers University.
§
Pacific Northwest National Laboratory.
|
Department of Biology, Clark University, Worcester, Massachusetts (present address).
⊥
Department of Geological Sciences, University of Saskatchewan, Canada (present address).
10.1021/es9003608 CCC: $40.75
Published on Web 06/12/2009
 2009 American Chemical Society
receiving hydrologic inputs from mercury-contaminated sites
(6-8). Critical to understanding the formation of methylmercury is an accurate knowledge of the chemical reactions
that proceed as inorganic mercury moves from contaminant
sources to anoxic methylation zones.
The fate of Hg(II) in soil and groundwater is strongly
influenced by complexation and redox reactions (9). Hg(II)
complexation with organic matter and mineral surfaces can
retard its subsurface migration (10, 11). Oxide minerals in
particular have been found to be efficient sorbents of Hg(II)
(12-15). The adsorption of Hg(II) onto iron and aluminum
oxide surfaces has been studied extensively, with the extent
of adsorption varying as a function of mineral surface area,
pH, and chloride concentrations. Iron oxides such as
ferrihydrite and goethite are known to adsorb Hg(II) ions
above pH 4 (16-18). For example, in ferrihydrite suspensions,
Hg(II) adsorption increases from near zero at pH 3 to greater
than 90% at pH 5 (17). The mechanism for this process is
attributed to surface complexation of Hg(II) ions with surface
hydroxyl functional groups at the mineral-water interface
(19). X-ray absorption spectroscopy has shown that the
dominant mode of Hg(II) adsorption to goethite occurs by
the formation of monodentate and bidentate inner sphere
surface complexes (18). At high pH, the extent of adsorption
onto iron oxide surfaces decreases due to pH dependence
of Hg(II) hydrolysis (16). The presence of chloride also inhibits
adsorption due to the formation of nonsorbing mercury
chloride complexes (13, 17, 20).
The redox transformation of Hg(II) to Hg(0) significantly
alters the fate of mercury in soil and groundwater. Due to
its low solubility in water and high volatility, Hg(0) readily
partitions to the gas phase in the vadose zone (21, 22). Loss
of gaseous Hg(0) to the atmosphere decreases the amount
of mercury remaining for groundwater transport and limits
the concentration of Hg(II) available for methylation. In
groundwater aquifers where gas exchange is restricted, Hg(0)
may become supersaturated and mobilized to drinking water
sources (23). Important chemical reductants of Hg(II) include
dissolved organic carbon and sorbed/precipitated ferrous
iron. Alberts et al. (24) and Allard and Arsenie (25) demonstrated that natural organic matter such as humic and fulvic
acids can reduce Hg(II). Bacteria can promote mercury
reduction by catalyzing electron transfer from an electron
donor to Hg(II) (26-28). Mineral-associated ferrous iron has
been identified as another possible reductant for Hg(II).
Charlet et al. (29) reported the reduction of Hg(II) to Hg(0)
by Fe(II) adsorbed onto phlogopite surfaces, and O’Loughlin
et al. (30) observed reduction of Hg(II) by Fe(II)-containing
mineral hydroxysulfate green rust.
In anoxic groundwater, ferrous iron often accumulates in
soils and sediments as the iron oxide mineral magnetite
(Fe3O4). Magnetite is a mixed-valence iron oxide that contains
both Fe(II) and Fe(III) ions in an inverse spinel structure
with oxygen atoms in a cubic closest packing array. White
and Peterson showed that the Fe(II) in magnetite can reduce
a wide range of metal ions, including ferric iron, copper(II),
vanadate, and chromate (31). Previous experimental studies
have found that magnetite can also reduce Np(V) to Np(IV)
(32), Pu(V) to Pu(IV) (33), U(VI) to U(IV) (34), and Se(IV) to
Se(0) (35). However, despite its affinity to reduce metal
contaminants (31-35), and its common occurrence in soils
and sediments (36), the extent to which magnetite can act
as a chemical reductant for the reduction of Hg(II) is currently
unknown.
In this study, we conducted laboratory experiments to
investigate the interaction of Hg(II) with magnetite in
VOL. 43, NO. 14, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
5307
deoxygenated water. We present macroscopic and spectroscopic evidence that the ferrous iron in magnetite reduces
Hg(II) to Hg(0). Hg(II) reduction experiments were conducted
as a function of magnetite surface area, pH, and chloride
concentration to quantify the rates of reaction. Reacted
magnetite samples were analyzed using Mössbauer and X-ray
photoelectron spectroscopy to examine solid-phase changes
in Fe speciation, and to identify the Hg charge state on the
magnetite surface. The results of this work suggest that
surface-catalyzed Hg(II) reduction is a kinetically favorable
pathway for the formation of Hg(0) in magnetite-bearing
soils and sediments.
Experimental Section
Synthesis and Characterization of Iron Oxides. Magnetite,
goethite, and ferrihydrite were synthesized according to the
methods described by Cornell and Schwertmann (36). Briefly,
magnetite was prepared in an anaerobic glovebox (Coy
Laboratories, Grass Lake, MI) by titrating a FeSO4 · 7H2O
solution with KOH and KNO3 at 90 °C. Ferrihydrite was
synthesized by titrating a FeCl3 solution with NaOH to pH
7. Goethite was prepared by reacting ferrihydrite in a KOH
solution at 70 °C for 60 h. All precipitates were washed with
deoxygenated distilled deionized water until the supernatant
exhibited a constant pH approximately equal to the pHzpc of
the iron oxide mineral. Aliquots of the suspension were
filtered, dried under N2 atmosphere, and characterized using
X-ray powder diffraction (XRD) (Philips X’Pert diffractometer).
The identities of the minerals were confirmed by comparing
the X-ray diffraction patterns to standards in the Joint
Committee on Powder Diffraction Standards database.
Surface area was determined with an 11-pt BET-Nitrogen
isotherm (Micromeritics Gemini 2375). BET measurements
were conducted on samples outgassed at 80 °C for 24 h.
Trapping of Hg(0). Magnetite (0.2 g/L) suspended in 20
mL of deoxygenated water was reacted with 48.4 ( 2.0 nM
HgCl2, corresponding to a total of 0.63 ( 0.04 µg of Hg in the
reactor. Experiments were conducted in foil-wrapped sealed
serum bottles, and Hg(0) gas was collected continuously by
purging N2 through the reactor for 1 h into midget bubblers
(Ace Glass, Vineland, NJ, catalog 75320-20) containing an
Hg(0) trapping solution (0.6% potassium permanganate, 2.5%
sulfuric acid, 2.5% nitric acid). Samples were collected from
the reaction vessel and trapping solution at the beginning
and at the end of the experiment. Additionally, at the end of
the experiment, the walls of the bubblers and serum bottles
were washed with concentrated acid to remove any mercury
sorbed onto the glassware. Mercury was digested using 1:1
concentrated sulfuric and nitric acids (trace metal grade)
and 10 MΩ Milli-Q water, according to a variation of EPA
method 245.1 (37). These samples were heated at 65 °C for
2 h, incubated with 250 µL of 5% potassium permanganate
at room temperature overnight, and reacted with 100 µL of
12% hydroxylamine hydrochloride. Finally, samples were
diluted with 2% HCl and analyzed by cold vapor atomic
absorbance spectroscopy (CVAAS) using a Leeman Laboratories Hydra AA (Hudson, NH). The detection limit of our
instrument, defined as 3 times the standard deviation of 10
blank samples, was 0.4 nM.
Hg(II) Reduction Kinetic Experiments. For the kinetic
experiments, Hg was provided as HgCl2, with 203HgCl2 as a
radioactive tracer (provided by Prof. D. Barfuss). Reactions
were performed in sealed serum bottles containing deoxygenated water. Experiments with magnetite were conducted
as a function of magnetite surface area (0.5-2 m2/L), pH
(4.8-6.7), and chloride concentration (10-6 to 10-2 mol/L).
Samples (0.5 mL) were removed from the sealed serum bottles
using a needle and syringe, every 35 s for approximately 20
min. The radioactivity of the unfiltered mineral suspension
was measured to determine the amount of total Hg remaining
5308
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 43, NO. 14, 2009
(e.g., Hg in solution and adsorbed to the mineral particles).
203
Hg analysis was performed by liquid scintillation counting
using a Beckman LS-6500 Counter (Beckman Instruments,
Fullerton, CA), with EcoLume Scintillation Cocktail (ICN
Radiochemicals, Irvine, CA). Counts per minute (CPM) were
determined for 3 min using a wide channel. Mercury
reduction experiments were also performed with goethite
and ferrihydrite suspensions (0.2 g/L) as Fe(II)-free controls.
Additional control experiments were conducted with deoxygenated water in absence of iron oxide minerals. As initial
CPM readings in mineral suspensions were indistinguishable
from those from the water controls, it was assumed that
quench due to minerals in suspension was not complicating
data collection.
57
Fe Mössbauer Spectroscopy. 57Fe Mössbauer spectroscopy was performed on magnetite samples reacted with 1
mM of HgCl2 for 14 days. Samples from a control experiment
conducted with magnetite suspended in deoxygenated water
without Hg(II) for 14 days were also analyzed. All samples
were dried in an anoxic chamber and care was taken to
prevent magnetite oxidation by sample handling (38). The
sample holder was filled with the magnetite sample and
sealed with transparent tape and an oxygen-impermeable
polymer film (aluminized Mylar stable to 4 K). The tape and
polymer were snapped into the holder with carbonized
polyethyletherketone (PEEK) polymer rings to ensure tightness. Mössbauer spectra were collected according to the
procedure given in Kukkadapu et al. (38). A closed-cycle
cryostat was used for the 125 K measurements. The Mössbauer data were modeled with the Recoil software using a
Voight-based spectral fitting routine (39). Additional details
of the Mössbauer analysis are included in the Supporting
Information.
X-ray Photoelectron Spectroscopy. XPS analysis was
performed on magnetite samples reacted with 0.1 mM and
1 mM HgCl2. After reaction for 14 days, magnetite suspensions
were pipetted from the glass serum bottle in a glovebox at
<0.1 ppm O2, and centrifuged/filtered at 4500 rpm using 30K
molecular weight cut off Whatman VectaSpin centrifuge
filters. The resulting pastes were smeared onto tantalum
coupons with a stainless steel spatula and allowed to dry.
Samples were then placed in a dry seal desiccator and
transferred to a glovebag (∼35 ppm O2) attached to the XPS
entry port. Exposure to atmosphere was minimized.
The XPS measurements were performed using a Physical
Electronics Quantum 2000 Scanning ESCA Microprobe.
Details of the XPS analysis are included in the Supporting
Information. Scans of the Hg4f and Cl2p regions were
recorded and the energy scale was referenced to adventitious
carbon 1s at 285.0 eV. Because the Hg4f region is strongly
overlapped with Fe3s, the Fe3s region was characterized for
unreacted magnetite and used in the fit for the experimental
spectra. Silicon (Si) was detected and the Si2p line was also
included in the fit. The Hg4f lines are characterized by two
simple spin-orbit split peaks, where the Hg4f5/2 peak is clearly
visible but the Hg7/2 peak is buried under the Fe3s and Si2p
peaks. The Hg7/2 peak was generated using known values for
the spin orbit splitting and relative intensities of the two
peaks. The spectra were best fit by nonlinear least-squares
using the CasaXPS curve resolution software. A spin-orbit
splitting of 4.0 eV was used. The Hg4f5/2:Hg4f7/2 intensity
ratio was set at 0.75, which is the ideal branching ratio. Each
spin-orbit peak was modeled using only one component,
with variable but equal fwhm (full width at half-maximum).
This is consistent with the closed shell electronic structure
of Hg(II) and resulting lack of multiplet structure. Elemental
ratios were semiquantified using Scofield photoionization
cross sections.
TABLE 1. Loss of Hg from Iron Oxide Mineral Suspensions
percent Hg lost from suspension
magnetitea
ferrihydrite
goethite
water
2h
24 h
surface area (m2/g)
79.5 ( 0.8b
9.0 ( 1.2
4.2 ( 1.5
1.0 ( 1.8
82.5 ( 0.2
16.2 ( 1.2
9.1 ( 4.2
5.7 ( 1.8
10.4
46.4
13.6
NAc
a
All iron oxide minerals were suspended at a
concentration of 0.2 g/L in deoxygenated water. b Values
represent means of triplicate experiments, and errors
represent the standard deviation of the mean. c Not
applicable.
FIGURE 1. Formation of gaseous Hg(0) during Hg(II) reaction
with magnetite. Grey bars represent mercury remaining in the
reaction vessel, and white bars represent mercury in the
trapping solution. Reactions were performed under anoxic
conditions in sealed 100 mL serum bottles with 70 mL of water
or magnetite suspension (0.2 g/L). The initial concentration of
Hg(II) added to the reaction vessel was 48.3 ( 2.0 nM.
Reactions were performed in triplicate, and error bars
represent standard deviation.
Results and Discussion
Reduction of Hg(II) to Hg(0). The reduction of Hg(II) by
magnetite was investigated by reacting 0.2 g/L of magnetite
with 100 nM Hg(II) in deoxygenated water, using 203Hg as a
tracer (Table 1). After 2 h, radioactivity measurements
indicated that 79.5 ( 0.8% of the total Hg was lost from the
magnetite suspension, compared to 1.0 ( 1.8% Hg loss in
the reaction vessel containing only deoxygenated water. The
reaction with magnetite was nearly complete in 2 h, as
reaction for 24 h resulted in only a marginal increase in Hg
loss (82.5 ( 0.2%). In the control experiments with the Fe(III)
oxide minerals goethite and ferrihydrite, the amounts of Hg
lost from the suspensions after 2 h were 4.2 ( 1.5% and 9.0
( 1.2%, respectively. These measurements indicate that
reaction of Hg(II) with magnetite results in loss of Hg that
does not occur with goethite or ferrihydrite.
A mercury trapping experiment was conducted to determine if the loss of mercury in the magnetite suspension
was due to formation of volatile Hg gas. After 1 h of reaction
time, 30.8 ( 5.2% of the Hg remained in suspension with the
magnetite, and 70.9 ( 16.3% was recovered in a potassium
permanganate trap (Figure 1). For the control experiment
performed in deoxygenated water without magnetite, 110.2
( 2.8% of the Hg remained in the reaction vessel and 2.5 (
1.5% was recovered in the trap. These results indicate that
gaseous Hg is a product of Hg(II) reaction with magnetite.
We interpret this gaseous Hg to be strong evidence for the
formation of Hg(0).
Kinetic Experiments. Experiments were conducted to
measure Hg(II) reduction rates, and to determine the effect
of magnetite surface area, pH, and chloride concentration
on the rates of reaction. The data show that the rate of
reduction increases with increasing magnetite surface area
FIGURE 2. Kinetics of Hg(II) reduction by magnetite.
Experiments were conducted in deoxygenated water with 100
nM of HgCl2, using 203Hg as a tracer. Symbols indicate
experimentally determined data, and lines indicate the pseudo
first-order kinetic model fit. Percent Hg(II) remaining in solution
is plotted as a function of (A) magnetite surface area (0.5-2 m2/
L), (B) pH (4.8-6.7), and (C) chloride concentration (10-6 to 10-2
mol/L).
(0.5 to 2 m2/L) (Figure 2A) and pH (4.8 to 6.7) (Figure 2B),
and decreases with increasing chloride concentration (10-6
to 10-2 mol/L) (Figure 2C). At a magnetite surface area
concentration of 2 m2/L, over 80% of the Hg(II) loss occurred
in less than 15 min of reaction (Figure 2A). This Hg(II)
reduction rate is approximately 10 times faster than Hg(II)
reduction by Fe(II) sorbed onto phlogopite (29).
As the magnetite surface site density is in large excess
compared to the Hg(II) concentration, we can assume that
the magnetite concentration remains constant during the
experiment. Accordingly, the Hg(II) reduction kinetics can
be described using a pseudo first-order kinetic model:
d[Hg(II)]
) -krxn[Hg(II)]
dt
(1)
where krxn is pseudo first-order rate constant, and [Hg(II)] is
the concentration of total Hg(II) remaining in the system.
The reaction rate constants determined for each experimental
system are given in Table 2. The reaction rates predicted by
each pseudo first-order rate constant are plotted in Figure
VOL. 43, NO. 14, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
5309
TABLE 2. Pseudo First-Order Reaction Rate Constants: Hg(II)
Reduction by Magnetite at Varying pH, Magnetite Surface
Area, and Chloride Concentrations
pH
[Fe3O4] (m2/L)
[Cl-] (mol/L)
rate constant (
(2.0 × 10-4) (s-1)
6.73 ( 0.08
6.73 ( 0.08
6.73 ( 0.08
4.77 ( 0.23
6.05 ( 0.15
6.63 ( 0.11
6.63 ( 0.11
6.63 ( 0.11
2.08 ( 0.11
1.04 ( 0.05
0.52 ( 0.04
2.08 ( 0.11
2.08 ( 0.11
2.08 ( 0.11
2.08 ( 0.11
2.08 ( 0.11
1 × 10-2
1 × 10-4
1 × 10-6
1.6 × 10-3
9.0 × 10-4
4.0 × 10-4
3.0 × 10-4
9.0 × 10-4
1.0 × 10-4
5.0 × 10-4
9.0 × 10-4
2. The rate constants describe the overall reaction rates, which
include the rates of adsorption, electron transfer, and
volatilization.
Comparison between the experimental measurements
and model fits indicate that the pseudo first-order kinetic
model provides an excellent description of the Hg(II)
reduction data. The magnetite, [H+], and [Cl-] reaction order
terms are reported in the Supporting Information (Figure
S1).
Spectroscopic Analysis. To examine the solid-phase
changes in Fe speciation by 57Fe-specific Mössbauer spectroscopy and the Hg surface state by surface-sensitive X-ray
photoelectron spectroscopy, magnetite samples were treated
with higher concentrations of Hg(II) (100 µM to 1.4 mM).
Mössbauer spectra were collected at room temperature and
125 K to examine the purity of the product and to follow
oxidation of Fe(II) magnetite by Hg(II). Figure 3 shows 125
K spectra of a sample that was suspended in deoxygenated
water (Figure 3A) and a sample reacted with Hg(II) (Figure
3B). The spectrum of the water-treated sample exhibited two
sextet peaks with relative areas and Mössbauer parameters
that are similar to stoichiometric magnetite [Fe(II)/Fe-total
of 0.33] (40). The spectral features are also in agreement with
the absence of any impurity phases, e.g., ferrihydrite,
hematite, or goethite. The outer sextet (36% area) is due to
Fe(III) in tetrahedral sites of the inverse spinel structure. The
inner sextet (64% area), on the other hand, is due to an average
of Fe(II) and Fe(III) or “Fe2.5+” in the octahedral sublattice,
which is the result of fast electron hopping between the Fe(II)
and Fe(III) sites, at temperatures above 120 K (40). Because
half of the Fe in the octahedral sublattice is occupied by
Fe(II), the fit-derived Fe(II)oct/[Fe(II,III)oct + Fe(III)tet] of 0.36
is approximately a third of the total Fe, in agreement with
Fe charge state distribution in stoichiometric magnetite.
The redox reaction with Hg(II) resulted in significant
changes in the 57Fe Mössbauer spectrum of the magnetite
sample. This is evident from different relative areas of the
sextets in the water-treated and Hg(II)-treated spectra (Figure
3A and B). Increase in the outer sextet area to 46% from 36%
in the water-treated sample implied partial oxidation of the
octahedral Fe(II). Oxidized Fe(II) exhibits parameters similar
to the tetrahedral Fe(III), hence their peaks are unresolved
from each other. Based on the spectral area of the Fe2.5+
contribution (54%), the Fe(II) content of the Hg(II)-treated
sample was estimated to be 27% or half of the inner sextet
contribution. This change represents oxidation of ∼18% of
the Fe(II) in the magnetite sample by ionic mercury. Features
characteristic of ferrihydrite or goethite were absent in the
oxidized sample indicating that secondary Fe(III) oxide
phases did not form.
Results of the XPS analysis are summarized in the Table
3 (also see Figure S2 in Supporting Information). Adsorbed
Hg was detected on all samples that were reacted with HgCl2,
where Hg4f binding energies (BE) are consistent with Hg(II).
As expected, the Hg/Fe ratio was about 10-fold smaller for
5310
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 43, NO. 14, 2009
the 0.1 mM compared to the 1 mM HgCl2 magnetite
experiments, indicating that the amount of adsorbed Hg(II)
was proportional to the amount of initial HgCl2(aq). In
contrast, the same decrease in HgCl2(aq) only achieved a
2-fold decrease in surface Cl/Fe ratio. The Hg/Cl ratios for
the magnetite experiments are all far less than 0.5, the ratio
for HgCl2. This is consistent with preferential loss of total Hg
from the system.
Hg(II) Interaction with Magnetite Surfaces. Electron
transfer at the magnetite-water interface involves direct
interactions between the Hg(II) and structural ferrous iron
as free electrons are not readily transferred into aqueous
solution. Adsorption sites are one location where reduction
of metal ions can occur (41). The XPS data indicate that Hg(II)
ions can adsorb onto magnetite surfaces. Furthermore, the
Mössbauer results suggest that the adsorbed Hg(II) interacts
with magnetite surfaces by accepting electrons from Fe(II)
in the magnetite structure. On the magnetite surface, the
half cell potential for the solid-state oxidation reaction Fe(II)
f Fe(III) + e- is approximately -0.34 to -0.65 V (31). For
mercury reduction, the standard potential for the half reaction
Hg(II) + 2e- f Hg(0) is +0.85 V. The sum of the half cell
potentials yields a positive value, indicating that electrons
can spontaneously flow from Fe(II) on magnetite surfaces to
adsorbed Hg(II) ions. The reactivity of solid phase Fe(II) on
magnetite surface can be attributed to the shift in electron
density from surface hydroxyl groups, which increases the
reducing power of the Fe(II) ion (42). In comparison, the
reduction of Hg(II) by aqueous Fe(II) (-0.77 V) is energetically
less favorable, and kinetically inhibited (29).
Our experimental data demonstrate that the kinetics of
Hg(II) reduction by magnetite systematically varies as a
function of magnetite concentration, pH, and chloride
concentration. We propose that the rate of reaction is
controlled by the chemical speciation of Hg(II) ions and
magnetite surfaces. First, the effect of magnetite concentration on overall reaction rates is attributed to the increase of
mineral surface area available for interaction with Hg(II).
The surface area of the magnetite particles used in our
experiments was 10.4 m2/g (Table 1), with a surface site
density of approximately 3.62 × 10-5 mol/m2 (43). At the
concentration range of 0.5 to 2 m2/L, the concentration of
surface sites increased from 1.88 × 10-5 to 7.52 × 10-5 mol/L.
Second, the effect of pH on Hg(II) reduction rates can be
explained by the pH-dependent adsorption of Hg(II) onto
magnetite. Hg(II) complexation with surface hydroxyl groups
onto iron oxide surfaces occurs in the pH range of 4-7, with
adsorption increasing at pH > 4 and decreasing at pH > 7
(16, 17). This adsorption behavior is controlled by the
deprotonation of surface hydroxyl groups above pH 4, and
the formation of neutral Hg(OH)2 aqueous complexes above
pH 7. Between pH 4 and 7, deprotonated surface hydroxyl
groups generate negative surface charge and electrostatically
attract Hg(II) cations to adsorption sites at magnetitewater interface. Finally, the effect of chloride concentration
on Hg(II) reduction is attributed to mercury-chloride
aqueous complexation. At high chloride concentrations,
Hg(II) predominately exists as mercuric chloride complexes
(13, 19). The formation of stable nonsorbing aqueous mercury
complexes limits direct contact of mercuric Hg with magnetite
surfaces, thereby hindering the electron transfer reaction.
Environmental Significance. The results presented in this
study show that magnetite can rapidly reduce Hg(II). The
evidence for this reaction includes the loss of Hg(II) from
magnetite suspensions, the formation of gaseous Hg(0), and
the oxidation of solid-state Fe(II). These experimental
observations may help us understand the natural redox
transformations of Hg in magnetite-bearing soils and sediments. In oxic surface water, biologic and photoreduction
may dominate. However, because magnetite is common in
FIGURE 3. 125 K 57Fe Mossbauer spectra of magnetite. (A) magnetite samples suspended in deoxygenated water for 14 days; and (B)
magnetite samples reacted with 1.37 ( 0.07 mM of Hg(II) for 14 days.
TABLE 3. Elemental Ratios and Binding Energies from X-ray
Photoelectron Spectroscopy
sample
Hg/Fe
H2O+mag.
1 mM Hg+mag
area 2
1 mM Hg+mag
area 4
0.1 mM Hg+mag
area 5
0.1 mM Hg mag
area 6
a
(0.1 eV.
b
Cl/Fe
b
Hg/Cl
binding energiesa
Hg4f7/2
Hg4f5/2
Hg ND
0.0095
Hg ND Hg ND
Cl NAc Cl NA
102.0
106.0
0.012
0.25
0.049
102.0
106.0
0.0008
0.13
0.0066
102.0
106.0
0.0015
0.11
0.014
102.1
106.1
not detected. c not analyzed.
iron-reducing sediments and waterlogged soils, the reduction
of Hg(II) by solid-phase Fe(II) may occur in anoxic groundwater. This process may limit the discharge of MeHg from
hyporheic zones to surface waters (44) by controlling the
concentration of Hg(II), the substrate for mercury methylation. Conversely, Hg(II) reduction in saturated sediments
may promote the mobilization of Hg as Hg(0). The formation
of mobile Hg(0) is of particular concern as groundwater can
be a source of Hg to surface water (45) and water distribution
systems where Hg(0) is the major form of Hg (23, 46).
Intriguingly, elevated Hg concentrations in groundwater have
been correlated with high levels of Fe(II) and low redox (47).
Based on our laboratory experiments, Hg(II) reduction
by magnetite in soils and sediments is expected to be facile,
however the reaction rates would be strongly dependent on
groundwater composition and reactive surface area. For the
reaction to proceed, the pH of the groundwater must be high
enough for Hg(II) adsorption on the magnetite surface to
occur. In the presence of elevated chloride concentrations,
Hg(II) adsorption is inhibited and surface-mediated reduction
by magnetite is expected to be slow. Furthermore, Hg(II)
reduction by magnetite is likely to be most favorable under
anoxic conditions, or in soil environments where oxygen
diffusion is limited. White and Peterson showed that natural
magnetite sands weathered under reducing conditions are
VOL. 43, NO. 14, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
5311
electrochemically active, where as magnetite particles weathered under oxidizing vadose conditions show minimal
reactivity toward metal reduction (31). Oxidative weathering
of magnetite results in the formation of surface oxidation
products that may impede electron transfer. Therefore, the
oxidation of magnetite surfaces by oxygen is expected to
pacify its reactivity toward Hg(II). Other factors such as
organic matter and sulfide complexation with Hg(II) may
also inhibit mercury reduction by magnetite through the
formation of nonsorbing Hg chemical species. Additional
field and experimental studies are required to further
elucidate the competing pathways of Hg(II) transformation
in contaminated subsurface environments.
Acknowledgments
This research was supported by the Office of Science (BER),
U.S. Department of Energy Grant DE-FG02-08ER64544. XPS
and Mössbauer experiments were performed at EMSL, a
national scientific user facility sponsored by the Department
of Energy’s Office of Biological and Environmental Research
located at Pacific Northwest National Laboratory.
Supporting Information Available
This material is available free of charge via the Internet at
http://pubs.acs.org.
Literature Cited
(1) Riley, R. G.; Zachara, J. M.; Wobber, F. J. Chemical contaminants
on DOE lands and selection of contaminant mixtures for
subsurface science research; DOE/ER-0547; U.S. DOE, 1992.
(2) Campbell, K. R.; Ford, C. J.; Levine, D. A. Mercury distribution
in Poplar Creek, Oak Ridge, Tennessee, USA. Environ. Toxicol.
Chem. 1998, 17 (7), 1191–1198.
(3) Compeau, G. C.; Bartha, R. Sulfate-reducing bacteria: principle
methylators of mercury in anoxic estuarine sediment. Appl.
Environ. Microbiol. 1985, 50, 498–502.
(4) King, J. K.; Kostka, J. E.; Frischer, M. E.; Saunders, F. M. Sufatereducing bacteria methylate mercury at variable rates in pure
culture and in marine sediments. Appl. Environ. Microbiol. 2000,
66 (6), 2430–2437.
(5) Kerin, E. J.; Gilmour, C. C.; Roden, E.; Suzuki, M. T.; Coates,
J. D.; Mason, R. P. Mercury methylation by dissimilatory ironreducing bacteria. Appl. Environ. Microbiol. 2006, 72 (12), 7919–
21.
(6) Southworth, G. R.; Turner, R. R.; Peterson, M. J. Form of mercury
in stream fish exposed to high-concentrations of dissolved
inorganic mercury. Chemosphere 1995, 30 (4), 779–787.
(7) Southworth, G. R.; Peterson, M. J.; Ryon, M. G. Long-term
increased bioaccumulation of mercury in largemouth bass
follows reduction of waterborne selenium. Chemosphere 2000,
41 (7), 1101–5.
(8) Santoro, E. D.; Koepp, S. J. Mercury levels in organisms in
proximity to an old chemical site (Berry’s Creek, Hackensack
Meadowlands, New Jersey, USA). Mar. Pollut. Bull. 1986, 17
(5), 219–224.
(9) Gabriel, M. C.; Williamson, D. G. Principal biogeochemical
factors affecting the speciation and transport of mercury through
the terrestrial environment. Environ. Geochem. Health 2004, 26
(4), 421–34.
(10) Melamed, R.; Boas, R. C. V. Phosphate-background electrolyte
interaction affecting the transport of mercury through a Brazilian
Oxisol. Sci. Total Environ. 1998, 213 (1-3), 151–156.
(11) Miretzky, P.; Bisonti, M. C.; Jardim, W. F. Sorption of mercury(II)
in Amazon soils from column studies. Chemosphere 2005, 60
(11), 1583–1589.
(12) Lockwood, R. A.; Chen, K. Y. Adsorption of Hg (II) by ferric
hydroxide. Environ. Lett. 1974, (6), 151–166.
(13) Mac Naughton, M. G.; James, R. O. Adsorption of aqueous
mercury (II) complexes on the oxide/water interface. J. Colloid
Interface Sci. 1974, 46 (2), 431–440.
(14) Gunneriusson, L.; Sjorberg, S. Surface Complexation in the
H+ - Goethite (a-FeOOH)-Hg(II)-Chloride System. J. Colloid
Interface Sci. 1993, 156 (1), 121–128.
(15) Sarkar, D.; Essington, M. E.; Misra, K. C. Adsorption of Mercury(II)
by Variable Charge Surfaces of Quartz and Gibbsite. Soil Sci.
Soc. Am. J. 1999, 63 (6), 1626–1636.
5312
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 43, NO. 14, 2009
(16) Barrow, N. J.; Cox, V. C. The effects of pH and chloride
concentration on mercury sorption. I. By goethite. Eur. J. Soil
Sci. 1992, 43 (2), 295–304.
(17) Kinniburgh, D. G.; Jackson, M. L. Adsorption of Mercury(II) by
Iron Hydrous Oxide Gel. Soil Sci. Soc. Am. J. 1978, 42 (1), 45–47.
(18) Kim, C. S.; Rytuba, J. J.; Brown, G. E. EXAFS study of mercury(II)
sorption to Fe- and Al-(hydr)oxides: I. Effects of pH. J. Colloid
Interface Sci. 2004, 271 (1), 1–15.
(19) Tiffreau, C.; Lutzenkirchen, J.; Behra, P. Modeling the adsoprtion
of mercury(II) on (hydr)oxides 0.1. amorphous iron-oxide and
alpha-quartz. J. Colloid Interface Sci. 1995, 172 (1), 82–93.
(20) Kim, C. S.; Rytuba, J. J.; Brown, G. E. EXAFS study of mercury(II)
sorption to Fe- and Al-(hydr)oxides: II. Effects of chloride and
sulfate. J. Colloid Interface Sci. 2004, 270 (1), 9–20.
(21) Schlüter, K. Review: evaporation of mercury from soils. An
integration and synthesis of current knowledge. Environ. Geol.
2000, 39 (3), 249–271.
(22) Zhang, H.; Lindberg, S. Processes influencing the emission of
mercury from soils: a conceptual model. J. Geophys. Res. 1999,
104 (D17), 21889–21896.
(23) Murphy, E. A.; Dooley, J.; Windom, H. L.; Smith, R. G. Mercury
species in potable ground water in southern New Jersey. Water
Air Soil Pollut. 1994, 78, 61–72.
(24) Alberts, J. J.; Schindler, J. E.; Miller, R. W.; Nutter, D. E. Elemental
mercury evolution mediated by humic acid. Science 1974, 184,
895–897.
(25) Allard, B.; Arsenie, I. Abiotic reduction of mercury by humic
substances in aquatic system s an important process for the
mercury cycle. Water Air Soil Pollut. 1991, 56 (1), 457–464.
(26) Barkay, T.; Miller, S. M.; Summers, A. O. Bacterial mercury
resistance from atoms to ecosystems. FEMS Microbiol. Rev. 2003,
27, 355–384.
(27) Takeuchi, F.; Iwahori, K.; Kamimura, K.; Negishi, A.; Maeda, T.;
Sugio, T. Volatilization of Mercury under Acidic Conditions from
Mercury-polluted soil by a Mercury-resistant Acidothiobacillus
ferooxidans SUG2-2. Biosci. Biotechnol. Biochem. 2001, 65 (9),
1981–1986.
(28) Wiatrowski, H. A.; Ward, P. M.; Barkay, T. Novel reduction of
mercury(II) by mercury-sensitive dissimilatory metal reducing
bacteria. Environ. Sci. Technol. 2006, 40 (21), 6690–6696.
(29) Charlet, L.; Bosbach, D.; Peretyashko, T. Natural attenuation of
TCE, As, Hg linked to the heterogeneous oxidation of Fe(II): an
AFM study. Chem. Geol. 2002, 190 (1-4), 303–319.
(30) O’Loughlin, E. J.; Kelly, S. D.; Kemner, K. M.; Csencsits, R.; Cook,
R. E. Reduction of Ag(I), Au(III), Cu(II), and Hg(II) by Fe(II)/
Fe(III) hydroxysulfate green rust. Chemosphere 2003, 53 (5),
437–46.
(31) White, A. F.; Peterson, M. L. Reduction of aqueous transition
metal species on the surfaces of Fe(II) - containing oxides.
Geochim. Cosmochim. Acta 1996, 60 (20), 3799–3814.
(32) Nakata, K.; Nagasaki, S.; Tanaka, S.; Sakamoto, Y.; Tanaka, T.;
Ogawa, H. Sorption and reduction of neptunium(V) on the
surface of iron oxides. Radiochim. Acta 2002, 90 (9-11), 665–
669.
(33) Powell, B. A.; Fjeld, R. A.; Kaplan, D. I.; Coates, J. T.; Serkiz, S. M.
Pu(V)O2+ Adsorption and Reduction by Synthetic Magnetite
(Fe3O4). Environ. Sci. Technol. 2004, 38 (22), 6016–6024.
(34) Scott, T. B.; Allen, G. C.; Heard, P. J.; Randell, M. G. Reduction
of U(VI) to U(IV) on the surface of magnetite. Geochim.
Cosmochim. Acta 2005, 69 (24), 5639–5646.
(35) Scheinost, A. C.; Charlet, L. Selenite Reduction by Mackinawite,
Magnetite and Siderite: XAS Characterization of Nanosized
Redox Products. Environ. Sci. Technol. 2008, 42 (6), 1984–1989.
(36) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure,
Properties, Reactions, Occurences, and Uses; Wiley-VCH: New
York, 2003.
(37) O’Dell, J. W.; Potter, B. B.; Lobring, L. B.; Martin, T. D.
Determination of mercury in water by cold vapor atomic
absorption spectroscopy; Method 245.1; U.S. Environmental
Protection Agency, 1994.
(38) Kukkadapu, R. K.; Zachara, J. M.; Fredrickson, J. K. Biotransformation of two-line silica-ferrihydrite by a dissimilatory Fe(III)reducing bacterium: Formation of a carbonate green rust in the
presence of phosphate. Geochim. Cosmochim. Acta 2004, 68
(13), 2799–2814.
(39) Rancourt, D. J.; Ping, J. Y. Voigt-based methods for arbitraryshape static hyperfine parameter distribution in Mossbauer
spectroscopy. Nucl. Instrum. Meth. Phys. Res., B 1991, 58 (1),
85–97.
(40) Cashion, J.; Murad, E. Mössbauer Spectroscopy of Environmental
Materials and their Industrial Utilization; Kluwer Academic
Publishers: Boston, MA, 2004.
(41) Stumm, W.; Hering, J. G. Oxidative and reductive Dissolution
of Minerals. In Mineral-Water Interface Geochemistry; Hochella,
M. F., Jr., White, A. F., Eds.; Mineralogical Society of America:
Chantilly, VA, 1990; pp 427-465.
(42) Luther, G. W. The frontier-molecular-orbital theory approach
in geochemical processes. In Aquatic Chemical Kinetics; Stumm,
W., Ed.; Interscience: New York, 1990; pp 173-198.
(43) Wesolowski, D. J.; Machesky, M. L.; Palmer, D. A. Magnetite
surface change studies from 190 degrees C from in situ pH
titrations. Chem. Geol. 2000, 167 (1-2), 193–229.
(44) Stoor, R. W.; Hurley, J. P.; Babiarz, C. L.; Armstrong, D. E.
Subsurface sources of methyl mercury to Lake Superior from
a wetland-forested watershed. Sci. Total Environ. 2006, 368,
99–110.
(45) Bone, S. E.; Charette, M. A.; Lamborg, C. H.; Gonneea, M. E. Has
Submarine Groundwater Discharge Been Overlooked as a Source
of Mercury to Coastal Waters? Environ. Sci. Technol. 2007, 41
(9), 3090–3095.
(46) Barringer, J. L.; Szabo, Z. Overview of investigations into mercury
in ground water, soils, and septage, New Jersey Coastal Plain.
Water Air Soil Pollut. 2006, (175), 193–221.
(47) Barringer, J. L.; Szabo, Z.; Kauffman, L. J.; Barringer, T. H.;
Stackelberg, P. E.; Ivahnenko, T.; Rajagopalan, S.; Krabbenhoft,
D. P. Mercury concentrations in water from an unconfined
aquifer system, New Jersey coastal plain. Sci. Total Environ.
2005, 346 (1-3), 169–83.
ES9003608
VOL. 43, NO. 14, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
5313