Experimental study and steady-state simulation of ...web.engr.oregonstate.edu/~istokj/Microbial_Growth... · Experimental study and steady-state simulation of biogeochemical processes

Experimental study and steady-state simulation of

biogeochemical processes in laboratory columns

with aquifer material

Aria Amirbahmana,*, Rene Schonenbergerb, Gerhard Furrerc,Jurg Zobristb

aDepartment of Civil and Environmental Engineering, University of Maine, 5711 Boardman Hall, Orono, Maine,

ME 04469 5711, USAbSwiss Federal Institute of Environmental Science and Technology (EAWAG), Switzerland

c Institute of Terrestrial Ecology (ITO), Swiss Federal Institute of Technology (ETH), Zurich, Switzerland

Received 18 June 2001; received in revised form 5 August 2002; accepted 23 August 2002

Abstract

Packed bed laboratory column experiments were performed to simulate the biogeochemical

processes resulting from microbially catalyzed oxidation of organic matter. These included aerobic

respiration, denitrification, and Mn(IV), Fe(III) and SO4 reduction processes. The effects of these

reactions on the aqueous- and solid-phase geochemistry of the aquifer material were closely

examined. The data were used to model the development of alkalinity and pH along the column. To

study the independent development of Fe(III)- and SO4-reducing environments, two columns were

used. One of the columns (column 1) contained small enough concentrations of SO4 in the influent to

render the reduction of this species unimportant to the geochemical processes in the column.

The rate of microbially catalyzed reduction of Mn(IV) changed with time as evidenced by the

variations in the initial rate of Mn(II) production at the head of the column. The concentration of Mn

in both columns was controlled by the solubility of rhodochrosite (MnCO3(S)).

In the column where significant SO4 reduction took place (column 2), the concentration of

dissolved Fe(II) was controlled by the solubility of FeS. In column 1, where SO4 reduction was not

important, maximum dissolved Fe(II) concentrations were controlled by the solubility of siderite

(FeCO3(S)). Comparison of solid-phase and aqueous-phase data suggests that nearly 20% of the

produced Fe(II) precipitates as siderite in column 1. The solid-phase analysis also indicates that

during the course of experiment, approximately 20% of the total Fe(III) hydroxides and more than

70% of the amorphous Fe(III) hydroxides were reduced by dissimilatory iron reduction.

The most important sink for dissolved S(-II) produced by the enzymatic reduction of SO4 was its

direct reaction with solid-phase Fe(III) hydroxides leading initially to the formation of FeS.

0169-7722/03/$ - see front matter D 2003 Elsevier Science B.V. All rights reserved.

doi:10.1016/S0169-7722(02)00151-1

* Corresponding author. Tel.: +1-207-581-1277; fax: +1-207-581-3888.

E-mail address: [email protected] (A. Amirbahman).

www.elsevier.com/locate/jconhyd

Journal of Contaminant Hydrology 64 (2003) 169–190

Compared to this pathway, precipitation as FeS did not constitute an important sink for S(-II) in

column 2. In this column, the total reacted S(-II) estimated from the concentration of dissolved sulfur

species was in good agreement with the produced Cr(II)-reducible sulfur in the solid phase. Solid-

phase analysis of the sulfur species indicated that up to half of the originally produced FeS may have

possibly transformed to FeS2.

D 2003 Elsevier Science B.V. All rights reserved.

Keywords: Aquifer; Groundwater; Iron; Redox processes; Sulfur

1. Introduction

The presence of high concentrations of oxidizable organic matter in groundwater can

lead to the development of reducing conditions due to microbial activities. Geochemical

transformations brought about by the oxidation of organic matter can have important

implications for the quality of groundwater (Jacobs et al., 1988). These transformations

include reduction of electron acceptors such as dissolved O2, NO3, and SO4, reductive

dissolution of Fe- and Mn-(hydr)oxides, and precipitation or dissolution of various

carbonate and sulfide minerals. Alkalinity and pH of the groundwater also change due

to these transformations (von Gunten et al., 1994). The presence of a reduced zone can

have significant consequences for the mobility of toxic organic and inorganic chemicals.

We have simulated the development of Fe(III)- and SO4-reducing environments

downgradient from a municipal solid waste landfill (Riet, Switzerland) under well-

controlled laboratory conditions. This landfill has been the subject of previous studies

(Amirbahman et al., 1998; Abbaspour et al., 1999; Hoehn et al., 2000). Solid material

collected from the oxic part of the aquifer and inoculated with bacteria from the anaerobic

leachate was packed in two flow-through column reactors. Anaerobic conditions were

established by running solutions of known concentrations of organic carbon and electron

acceptors through the columns. Geochemical changes caused by anaerobic conditions in

aqueous and solid phases along the columns were studied. Of special interest was to study

the contributions of the enzymatic Fe(III)-reducing conditions and the dissolved S(-II)

species generated by the SO4-reducing bacteria to the oxidation capacity (OXC) of the

aquifer as defined by Heron and Christensen (1994, 1995) and Heron et al. (1994a). This

approach allows us to quantify the rates of geochemical transformations brought about by

the redox processes, and to evaluate the capacity of the aquifer solid phase to buffer

against the development of highly reducing conditions.

2. Experimental

2.1. Column design and operation

Columns were made of Plexiglas with an inner diameter of 5 cm and contained side

ports with Teflon-lined septa at 0.5, 1, 1.5, 2, 3, 4, 5, 6, 10, 14, 18, 22 and 26 cm. They

were filled with aquifer solid collected from the oxic zone of well 5 in the vicinity of

A. Amirbahman et al. / Journal of Contaminant Hydrology 64 (2003) 169–190170

landfill Riet, Winterthur, Switzerland (Amirbahman et al., 1998). The aquifer solid was

first dried at 60 jC and then fractionated by nylon sieves. The fine sand fraction between

150 to 465 Am was filled into the columns up to a depth of 29 cm according to the

procedure of Stauffer and Dracos (1986) to insure homogeneous packing. The columns

were then filled with degassed water at ionic strength of 10 mM NaCl.

Two columns were operated separately in this study. Analysis of the effluent break-

through curves with tritium indicated that column 1 had a pore volume of 175 ml and a

porosity of 0.31. The pore volume and porosity for column 2 were 192 ml and 0.34,

respectively. Columns were operated at a flow rate of 0.45 ml min� 1 in the dark by

covering them with aluminum foil at room temperature.

The background influent solution was prepared twice a week by bubbling 100% CO2(g)

for 10–30 min and then 1% CO2(g) for 60 min through a suspension of calcite. This

procedure produced an alkalinity of 3.72F 0.65 meq l� 1. The average influent concen-

trations of sodium propionate (NaC3O2H5) and NaNO3 in both columns were 0.61 and

0.13 mM, respectively. The concentration of propionate in column 2 was increased to 1.0

mM after day 78 of operation. To suppress the dissimilatory SO4 reduction in column 1, no

SO4 was added to this column. However, the influent of column 1 contained a small

amount of SO4 due to impurities in CaCO3 and other salts used. This SO4 did not produce

any appreciable concentration of S(-II) that influenced the dissolved Fe concentration in

column 1. The influent solution in column 2 consisted of an average of 0.52 mM Na2SO4.

The influent solutions in both columns also contained 10–100 nM concentrations of

Co(II), Zn(II), Cu(II), Ni(II) and Mo(VI), 0.6 AM PO4 as nutrient and 10 mM NaCl as

ionic strength buffer.

Both columns were inoculated at the inlet with anaerobic bacteria which were present in

groundwater samples collected under anoxic conditions from well 1 of the landfill Riet

(Amirbahman et al., 1998). Anoxic condition was established in both columns within the

first 7 days of operation. All chemicals used were of reagent grade or better and the water

employed was double distilled.

2.2. Analysis of the aqueous phase

Dissolved oxygen was measured using a combined LF-196 instrument (WTW, Well-

heim, Germany) in a closed flow-through cell. A combination pH electrode (Metrohm)

and a pH meter (Metrohm 654) were used to measure the pH. Alkalinity was measured by

standardized 10 mM HCl to pH 4.5. Alkalinity and pH were measured immediately after

the samples were collected. The concentration of the dissolved inorganic carbon (DIC)

was calculated using pH and alkalinity measurements and by considering the contributions

of propionate, acetate and dissolved S(-II) to alkalinity. Propionate, acetate, SO4 and NO3

were measured simultaneously by ion chromatography (Dionex) with an AS11 anion

exchange column using the gradient technique (Ammann and Ruttimann, 1995). The

dissolved organic carbon (DOC) analysis of the effluent for a selected few samples

indicated that more than 95% of the effluent DOC was due to propionate and acetate.

Effluent samples for metal and S(-II) analyses were collected in 1 M HNO3 and 0.1 M zinc

acetate solutions, respectively, such that oxidation or volatilization would be minimized.

Dissolved Fe, Mn, K, Na, Mg, and Ca were measured using ICP-OES (Spectro Analytical

A. Amirbahman et al. / Journal of Contaminant Hydrology 64 (2003) 169–190 171

Instruments). Dissolved S(-II) concentrations were measured using the photometric

method with methylene blue (Zhabina and Volkov, 1978).

2.3. Analysis of the solid phase

Chemical extraction techniques were used to analyze the solid samples for Fe and

reduced sulfur. The solid-phase Fe(III) was extracted by the Ti(III)–EDTA method, which

represents the total reducible Fe(III) phases (Heron et al., 1994a). This technique is

capable of completely reducing amorphous synthetic hydrous ferric oxide (HFO) and

lepidocrocite, and reducing more than 90% of synthetic hematite and goethite, but extracts

only 5% of magnetite and less than 5% of siderite (Heron et al., 1994b).

The concentration of the amorphous Fe(III) hydroxide was determined by chemical

extraction using ascorbate at pH 8 (Kostka and Luther, 1994). Extraction of synthetic

Fe(III) hydroxides showed that nearly all of amorphous HFO and less than 2% of

lepidocrocite and goethite are extracted by this technique (Amirbahman et al., 1998).

The ion exchangeable Fe(II) was extracted by mixing a suspension of sand in a 1 M CaCl2solution at pH 7 for 24 h (Heron et al., 1994b).

Acid volatile sulfide (AVS), which primarily reflects the amorphous FeS, was measured

using the technique developed by Zhabina and Volkov (1978). Cr(II)-reducible sulfide

(CRS) technique was used to measure the total reducible S (Canfield et al., 1986). In

general, CRS is a measure of FeS, FeS2, and elemental and organic S. Our previous CRS

measurements of similar soil under a landfill indicate that the elemental and organic S are

very small fractions of total reducible S (Amirbahman et al., 1998). A detailed description

of the extraction methods used is given previously (Amirbahman et al., 1998).

Quantitative analysis of the solid before packing was performed with a scanning

electron microscope (JEOL JSM-840) using an energy dispersive X-ray spectrometer

TRACOR 5402. The results were in agreement with X-ray diffraction patterns and

indicated that the main constituents are Si (25.3F 2.0% wt.) from quartz and silicates

(primarily feldspar and clay minerals) and Ca (23.5F 2.9% wt.) from calcite. Metals such

as Fe (3.0F 0.7% wt.) and Al (4.5F 0.4% wt.) were detected in smaller amounts. Due to

their relatively small amounts, which is typical of most natural samples, crystalline and

amorphous Fe(III) hydroxides were not distinguished.

3. Steady-state geochemical modeling

We have used the computer program STEADYQL (Furrer et al., 1989) to perform

geochemical calculations and sensitivity analysis of the results of our column experiments.

This program considers chemical speciation and reaction kinetics in a continuous-flow

stirred tank reactor. The flow domain is subdivided into several box reactors in series, where

the outflow from the first reactor constitutes the inflow for the second reactor. Therefore, the

column system is modeled as several continuously stirred tank reactors in series. Reaction

rates for geochemical processes are obtained by combining the concentration gradient of a

given component with the flow velocity. Reaction rates are assumed to be at steady-state

conditions during the course of sample collection. STEADYQL has been previously used in


other studies involving biogeochemical changes in flow-through systems (Furrer et al.,

1996; Amirbahman et al., 1998). The compartment between two consecutive sampling

ports were considered one box reactor, where the concentrations of components in the

inflow and the outflow are known, and are used to determine the reaction rate coefficients.

The processes in the columns were divided into ‘‘fast’’ and ‘‘slow’’ categories. The

‘‘fast’’ processes consist of equilibrium acid–base and aqueous complexation reactions,

and are reported elsewhere (Amirbahman et al., 1998). The ‘‘slow’’ processes are kineti-

cally controlled and consist of all of the microbially catalyzed redox reactions (Table 1), and

reaction of dissolved S(-II) and solid-phase Fe(III) hydroxides. The reaction rate coef-

ficients were determined based on concentration of the independent components, such as

dissolved O2, NO3, and SO4. The dependent species are those species whose turnover rates

and concentrations are affected by all of the biogeochemical processes involving the

independent species. The dependent species here are propionate, acetate, DIC and pH. The

modeled values of the dependent species can be compared to the actual experimental

measurements to determine the validity of the modeling approach and assumptions.

Zero-order rate coefficients were used to describe the geochemical processes. For

microbially catalyzed processes, the rate coefficients were calculated by dividing the

change in the concentration of the corresponding independent components in a given

reactor by the mean residence time in that reactor. The computer program STEADYQL

performs sensitivity analysis with respect to the rate coefficients that control the

concentrations of the dependent species. This provides us with useful information on

changes in concentrations of the dependent species brought about by changes in

concentrations of the independent components. The reported sensitivity coefficients are

normalized with respect to the rate coefficients, which are directly calculated from

concentrations of the independent components.

4. Results and discussion

4.1. Reduction of dissolved oxygen and nitrate

The microbially catalyzed redox processes for both columns are listed in Table 1. At the

end of the first week, all of dissolved O2 (0.25 mM) was consumed within the first 0.5 cm

Table 1

Stoichiometry of microbially catalyzed kinetic processes

Processes Stoichiometry

Aerobic respiration 2C3H5O2� + 7O2 + 2H

+! 6CO2 + 6H2O

Denitrification 5C3H5O2� + 14NO3

� + 4H+! 15HCO3� + 7N2 + 7H2O

Mn(IV) reduction C2H3O2� + 4MnO2 + 7H

+! 4Mn2 + + 2HCO3� + 4H2O

Fe(III) reduction C2H3O2� + 8FeOOH+ 15H+! 8Fe2 + + 2HCO3

� + 12H2O

Sulfate reduction 4C3H5O2� + 3SO4

2�! 4C2H3O2� +H+ + 4HCO3

� + 3HS�

Fermentation of propionate to acetate C3H5O2� + 3H2O!C2H3O2

� +H+ +HCO3� + 3H2

C3H5O2�: propionate

C2H3O2�: acetate


of the columns. Since O2 is likely to have been depleted at a shorter distance than the first

sampling port at 0.5 cm, the estimated zero-order rate coefficient of 9 nM s� 1 constitutes a

minimum value.

Nearly all added NO3 (average inlet concentration of 0.13 mM) was reduced in the

first 0.5 cm of both columns, even though at different initial rates in each column. The

onset of denitrification in column 1 was between days 7 and 14 after inoculation. By

day 14, nearly 95% of NO3, and by day 35, more than 98% of NO3 was reduced within

the first 0.5 cm. In column 2, a total of approximately 35% of NO3 was reduced by day

7. This reduction was almost all limited to the first 0.5 cm. By day 21, more than 90%

of NO3 was reduced in the first 0.5 cm with the rest being reduced in the following 1.5

cm. By day 35, nearly 97% of NO3 was reduced in the first 0.5 cm of column 2. The

increase in the rate of denitrification with time is due to the establishment and

acclimation of the NO3-reducing bacteria. In a similar system, von Gunten and Zobrist

(1993) have observed the onset of denitrification between days 1 and 6 after the

inoculation. They observed a virtually complete NO3 reduction within the first 1 cm of

the column.

4.2. Manganese reduction

Fig. 1 shows the concentration distribution of dissolved Mn(II) on different days in

both columns. Solid-phase Mn(IV) undergoes early reduction following the establishment

of anoxic environment in both columns. The onset of Mn(IV) reduction in both columns

was before day 7 of operation. For the first 14 days, when still negligible concentrations of

SO4 were reduced in column 2, nearly equivalent amounts of dissolved Mn(II) were

observed in both columns. In column 1, the initial zero-order rate coefficients of dissolved

Mn(II) production on days 7 and 14 were 0.28 and 1.02 nM s� 1, respectively.

Corresponding values of 0.27 and 1.11 nM s� 1 were observed in column 2. By day 35,

the initial dissolved Mn(IV) reduction rate coefficients diminished to 0.32 and 0.48 nM

s� 1 in columns 1 and 2, respectively. These rates were observed close to the inlet of the

columns, where initially the bulk of the bacterial biomass was located, the solution was

still undersaturated with respect to rhodochrosite (MnCO3(S)), and the adsorption sites at

the soil surface were likely to be saturated with respect to Mn(II). Saturation of surface

sites with respect to reduced metals has been substantiated by independent adsorption

experiments for the same aquifer material (results not shown). Therefore, we believe that

these rates are those of the overall Mn(II) production. The initial dissolved Mn(II)

production rate coefficients diminished with time to below 0.05 nM s� 1 by day 219 in

both columns.

4.3. Iron reduction

Dissolved Fe(II) was observed in both columns before day 7. It is assumed that the

Fe(II) release is due to the dissimilatory reduction of solid-phase Fe(III) hydroxides in the

presence and absence of SO4 (Jakobsen and Postma, 1999). First, we present the results

from column 1, where precipitation of FeS did not take place due to the presence of small

concentrations of SO4 in the influent.


Fig. 2 exhibits the distribution of dissolved Fe(II) along the length of column 1 at

different times. This distribution is characterized by relatively higher Fe(II) release rates

close to the inlet than at the outlet of the column. Dissolved Fe(II) concentrations along the

length also increased up to day 85, after which they decreased.

The observed release rate of dissolved Fe(II) reflects not only the reduction rate of

Fe(III) hydroxides, but also adsorption, desorption and precipitation of Fe(II). To study

variations in the reduction rate of Fe(III), we have considered the dissolved Fe(II)

Fig. 1. Distribution of dissolved Mn(II) in (a) column 1, and (b) column 2. The numbers refer to the days samples

were collected for measurement.


distribution at the first 2 cm of column 1. Since at this location the adsorption sites are

likely to be saturated with respect to Fe(II) and the alkalinity and Fe(II) concentrations are

not high enough to induce precipitation of siderite, the measured initial production rates of

dissolved Fe(II) may be the best approximation for the dissimilatory Fe(III) reduction

rates. Close to the inlet of column 1, a zero-order rate of 0.1 nM s� 1 for Fe(III) reduction

is observed on day 7 (data not shown). However, the concentration of dissolved Fe(II)

decreases after the first 2 cm due to precipitation of siderite (Fig. 3a). The initial Fe(II)

release rate increases to 4.5 nM s� 1 by day 70, after which it decreases perhaps due to the

depletion of the more reactive (amorphous) solid Fe(III) hydroxide species. By day 85 of

operation, the initial rate of Fe(II) production in column 1 is 2.5 nM s� 1 and continues at

an initial rate of 2.1 nM s� 1 until the end of column operation on day 219. Even though

the rate of dissimilatory Fe(III) reduction cannot be directly determined for downgradient

Fig. 2. Distribution of dissolved Fe(II) in column 1. The numbers refer to the days samples were collected for

measurement.


Fig. 3. Saturation index with respect to (a) siderite in column 1, (b) rhodochrocite in column 1, and (c)

rhodochrocite in column 2.


locations in column 1, we expect that time variation of this rate would be similar to that

measured close to the inlet of this column.

4.3.1. Precipitation of siderite

Solubility calculations indicate that siderite should be considered as a sink for the

dissolved Fe(II) in column 1. Fig. 3a shows the saturation index of siderite as a function

of distance at different times in this column. Precipitation of siderite in not expected close

to the inlet of this column at any time. Therefore, at a distance no longer than 1.5–2 cm,

concentration of dissolved Fe(II) is controlled only by dissimilatory Fe(III) reduction.

Initially, the solution is undersaturated with respect to siderite throughout the column, due

to low concentrations of dissolved Fe(II) and DIC. By day 41, however, precipitation may

occur toward the end of the column (Fig. 3a). This increase in saturation index is due to

the increasing production of DIC and dissolved Fe(II) during this time. The pore water

remains close to saturation along most of the length of the column until the end of the

experiment. Toward the end of the operation of column 1, where high concentrations of

DIC are produced, precipitation of siderite is expected at distances as short as 1.5 cm

from the inlet. Unlike field conditions, complexation of Fe with organic ligands is

negligible in this study. Therefore, in chemical speciation modeling, all measured Fe

species in solution have been considered as either free Fe(II) or Fe(II) complexed with

carbonate species.

4.3.2. Solid-phase analysis of column 1

Solid-phase extractions of both columns were performed before and after the experi-

ment. Fig. 4a illustrates the distribution of the solid-phase Fe(III) hydroxides in column 1

obtained using Ti(III)–EDTA extraction. These are the residual Fe(III) hydroxides that

were measured after the completion of the experiment. The data in Fig. 4 have been

corrected for the Fe(II) adsorbed to aquifer matrix by subtracting the concentration of

Fe(II) extracted by CaCl2 (results not shown here).

Ti(III)–EDTA extraction results indicate that initially the solid phase contained a total

concentration of 21 Amol Fe(III) g� 1. The total residual concentration of Ti(III)–EDTA

extractable Fe(III) averaged over the entire column length is 17 Amol g� 1. Therefore,

according to these measurements, a total of 4 Amol Fe(III) g� 1, or approximately 19% of

the total Ti(III)–EDTA extractable Fe(III) hydroxides, has been reduced. The total

adsorbed Fe(II) concentration, as estimated by CaCl2 extraction of the solid phase and

averaged over the entire length of column 1, is only 0.06 Amol g� 1, which is a very small

part of the total produced Fe(II).

Fig. 4b shows the distribution of the residual Fe(III) hydroxides extracted using

ascorbate along the length of column 1. This fraction, which consists largely of the

amorphous Fe(III) phases, had an initial concentration of 1.1 Amol g� 1. The residual

concentration of the ascorbate extractable Fe(III) hydroxides averaged over the length of

the column is 0.3 Amol g� 1. This concentration may even be lower due to the possible

oxidation of small amounts of adsorbed Fe(II) during sample handling and measurements.

Previous studies of biotic and abiotic reductive dissolution of solid Fe(III) hydroxides

have shown that amorphous phases dissolve more readily than the crystalline phases

(Lovley et al., 1991; Amirbahman et al., 1997). The same behavior is also observed in this


study. Comparison of the initial and residual masses of ascorbate extractable Fe(III)

indicates that more than 70% of the amorphous Fe(III) phase has been reductively

dissolved in column 1. Even though only 5% of the total extractable Fe(III) hydroxides

are amorphous, they constitute nearly 20% of the total reduced Fe(III). During the course

of the experiment, higher concentrations of the ascorbate-extractable Fe(III) phases have

been dissolved closer to the inlet of the column.

Table 2 summarizes solid- and liquid-phase measurements of total Fe in column 1.

Total initial and residual Ti(III)–EDTA extractable amounts of Fe(III) hydroxides are 22.0

and 17.8 mmol, respectively. The total mass of reduced Fe(III) hydroxides estimated from

solid-phase data is, therefore, 4.2 mmol. The total mass of aqueous phase Fe(II) that has

been released from column 1 during the course of experiment is 3.3 mmol. Therefore, the

total mass of Fe(II) reacted in the column, which is the difference between the total

reduced Fe(III) hydroxides and the total aqueous Fe(II) output, is 0.9 mmol. Accounting

for the total adsorbed mass of Fe(II) as estimated by CaCl2 extractions, the total mass of

Fig. 4. Concentration of solid-phase Fe(III) hydroxides in column 1 extracted by (a) Ti(III)–EDTA, and (b)

ascorbate. A background concentration of 3.8F 0.7 Amol Fe/g solid before the experiment has been subtracted

from the Ti(III) –EDTA extractable Fe data.


Fe(II) precipitated as siderite is approximately 0.8 mmol (Table 2). Adsorbed Fe(II)

constitutes only a small part of the total Fe(II) budget.

Given that the reduced Fe(III) in column 1 is only a small fraction of the total Ti(III)–

EDTA extractable Fe(III) hydroxides, any uncertainties in analytical measurements could

lead to significant errors in the estimation of precipitated siderite. According to the data

presented in Table 2, nearly 20% (F 10%) of the total produced Fe(II) precipitates as

siderite, while the rest leaves the column.

4.3.3. Average rates of Fe(III) reduction

Analysis of the rates of release of dissolved Fe(II) close to the inlet of column 1

discussed above suggests that these rates change with time due to the preferential depletion

of the amorphous compared to the more crystalline Fe(III) hydroxides. We have used the

initial and residual concentrations of Ti(III)–EDTA extractable Fe(III) hydroxides to

estimate the zero-order rate coefficients for the reduction of Fe(III) at different locations

along the column, by taking into account the flow rate and the mean residence time. These

reduction rate coefficients are averaged over the entire duration of operation of column 1

and are shown in Fig. 5. According to Fig. 5, the average zero-order Fe(III) reduction rate

decreases approximately three-fold from the inlet to the outlet of the column. Given the

uniformly available surface area of Fe(III) hydroxides along the column, the most likely

explanation for the decrease in the Fe(III) reduction rate in the direction of flow would be

the non-uniform distribution of the density of the Fe(III)-reducing bacteria.

The average rate of reduction of Fe(III) hydroxides at 1.5 cm estimated from the solid-

phase analysis is 1.9 nM s� 1 . Using the aqueous phase data, this rate is 2.4 nM s� 1 at the

first 2 cm of column 1. The latter rate was estimated by integrating and averaging the

initial zero-order rates with respect to time. The agreement between the data estimated

Table 2

Mass balance of iron species in column 1

Measurements and procedures Mass (mmol)

Total initial Ti(III) –EDTA extractable Fe(III)a 22.0

Total residual Ti(III) –EDTA extractable Fe(III)a,b 17.8

Total reduced Fe(III) hydroxidesc 4.2

Initial ascorbate extractable Fe(III)a 1.2

Residual ascorbate extractable Fe(III)a,b 0.3

Total reduced amorphous Fe(III) hydroxides 0.9

Residual CaCl2 extractable Fe(II)d 0.06

Total dissolved Fe(II) outpute 3.3

Total precipitated Fe(II)f 0.8

a Total reducible Fe(III) in the solid phase extracted directly and integrated over the entire length of column.b These values have been corrected for the CaCl2 extractable Fe(II).c The difference between total initial Fe(III) and total residual Fe(III) obtained by Ti(III) –EDTA extraction.d The ion exchangeable Fe(II) adsorbed to the solid surface.e Total dissolved Fe(II) released from the column over the course of experiment.f Total precipitated Fe(II) = Total reduced Fe(III) hydroxides�Total dissolved Fe(II) output�Total residual

CaCl2 extractable Fe(II).


from the solid and aqueous phases is reasonable considering that sensitivity analysis

indicates an approximately 40% variation in the values of the average reduction rates with

a 20% change in the concentration of residual Ti(III)–EDTA extractable Fe(III).

4.4. Comparison of Mn and Fe turnover

During the first 3 weeks of operation of the columns, the dissolved Mn(II) concen-

trations were higher than those of the dissolved Fe(II), indicating a higher rate of reduction

of Mn(IV) oxides. By day 35, however, higher concentrations of dissolved Fe(II) were

observed in column 1, where there was no FeS precipitation. The decrease in the initial

rate of Mn(II) production with time is likely due to the depletion of the more available

fractions of Mn(IV) oxides by the Mn-reducing bacteria, as suggested previously (von

Gunten and Zobrist, 1993). The removal of dissolved Mn(II) can be explained by two

possible mechanisms: precipitation as rhodochrosite and/or adsorption to FeS and other

solids, with adsorption to FeS possible only in column 2. As indicated by equilibrium

calculations, after day 14, precipitation of Mn(II) as rhodochrosite took place (Fig. 3b and

c). Equilibrium with respect to rhodochrosite was maintained throughout the duration of

this study in both columns. The overall decrease in the dissolved Mn(II) concentration

with time is, therefore, explained by the gradual increase in alkalinity, which was in turn

generated by the microbially catalyzed reactions.

Other likely mechanisms for the removal of Mn(II) are its adsorption to and

incorporation into FeS and other solids. Morse and Luther (1999) have argued that

Fig. 5. The average rate of dissimilatory Fe(III) hydroxide reduction in column 1 estimated from the solid-phase

measurements.


despite the higher rate of water exchange for Mn2 + than for Fe2 +, Mn2 + does not

readily form a MnS phase. Instead, they have proposed incorporation of Mn(II) into FeS

and eventually into pyrite. However, the latter mechanism may be kinetically con-

strained, as observed for the dissolution of sulfide minerals (Morse and Luther, 1999).

Kornicker and Morse (1991) and Arakaki and Morse (1993) have shown that dissolved

Mn(II) can also adsorb onto FeS phases and be pyritized. From our data, we cannot

estimate the fraction of Mn(II) removal by each mechanism. However, the relatively

constant concentrations of this species at any given time throughout the length of both

columns suggest that precipitation as rhodochrosite would be the dominant removal

mechanism.

4.5. Sulfate reduction

This section discusses the geochemical changes brought about due to the reduction of

SO4 in column 2. Production of S(-II) from SO4 (average influent concentration of 0.52

mM in column 2) was completely catalyzed by the SO4-reducing bacteria, with propionate

as the sole electron source for the reduction of SO4 (Table 1).

The diagenetic sulfur cycle has been described in detail elsewhere (Davison, 1991;

Luther et al., 1992). Reaction of S(-II) with iron species generally follows two major

reaction pathways. The first pathway involves the direct attack on the solid-phase Fe(III)

by aqueous S(-II) to readily form FeS at the surface. This is the main sink for the dissolved

S(-II) in most natural systems, where a sufficient pool of available solid-phase Fe(III)

hydroxides exists. The FeS mineral formed via this pathway undergoes further reaction to

form pyrite (FeS2) (Berner, 1984). The stoichiometry of direct S(-II) attack is (von Gunten

and Zobrist, 1993)

2FeOOHþ 3HS� ! 2FeSþ ^ Sð0Þ]]þ 3OH� þ H2O ð1Þ

FeSþ ^ Sð0Þ]] ! FeS2 ð2Þ

The second pathway is the reaction between dissolved S(-II) and Fe(II) and precipitation as

FeS, provided that the solubility limit of the mineral is reached:

Fe2þ þ HS� ! FeSþ Hþ ð3Þ

Typical profiles for the time-dependent distribution of aqueous S(-II) along the length

of column 2 are shown in Fig. 6. On day 7, the measured concentrations of S(-II) were

below the detection limit (0.15 AM) along the column. Dissolved S(-II) was detected at 0.5

cm from the inlet on day 14. The total concentration of dissolved S(-II) along the column

increased in the downgradient direction with time. After day 85, S(-II) reached the end of

the column and was detected in the effluent. By the end of operation of column 2,

dissolved S(-II) traveled nearly conservatively (Fig. 6).

Fig. 6 also shows the change in the concentration of SO4 for day 85. A similar pattern

was observed throughout the duration of the experiment, where SO4 reduction initiated

close to the column inlet because of the high concentration of the organic substrate and

density of the SO4-reducing bacteria. For example, on day 85, more than 98% of SO4 was


reduced within the first 4 cm of column 2. Kinetics of SO4 reduction process in this zone

may be characterized by a pseudo-first-order reaction with a rate coefficient of 7.4� 10� 4

s� 1 . The limiting factor in the dissimilatory reduction of SO4 was the availability of

propionate. During the course of this experiment, the bacteria did not use acetate as an

electron acceptor to reduce SO4, despite its presence in column 2.

Lack of stoichiometric correspondence between the distribution of SO4 and dissolved

S(-II) is evident (Fig. 6) and is attributed to the reactions of S(-II) with Fe(II) and Fe(III) as

depicted in Eqs. (1)–(3). The kinetics of production of dissolved S(-II) may be followed

by changes in SO4 concentration.

4.5.1. Solid-phase analysis of column 2

Solid-phase extraction of sulfur species in column 2 was performed before and after the

experiment. These extractions included AVS and CRS as described in the Experimental

section. Fig. 7 shows the results from AVS and CRS extractions of solids after the

experiment in column 2. Solid-phase sulfur forms initially as amorphous FeS. Further

transformation to pyrite takes place only later in the solid phase (Eq. (2)).

Fig. 6. Distribution of dissolved sulfur species (S(-II) and SO4) in column 2. The numbers refer to the days

samples were collected for measurement. The empty symbols represent S(-II) and the filled circle represents SO4

on day 85.


Integration of concentrations of the solid-phase sulfur species along the length of the

column indicates that a total of 30.4 mmol of CRS and a total of 15.2 mmol of AVS were

generated in column 2 during the course of the experiment. This means that half of the

originally formed FeS has been transformed into pyrite. Given our analytical procedure for

determination of AVS, however, it may be difficult to resolve the exact extent of

transformation of FeS to pyrite. This is due to the possible enhanced oxidation of H2S

by dissolved Fe(III) under acidic condition and at an elevated temperature during the AVS

extractions (von Gunten and Zobrist, 1993). Such artifact, if indeed effective, would

underestimate the concentration of AVS. Therefore, the measured AVS reported here

should be considered as the lower limit. Likewise, the extent of transformation into pyrite

would perhaps be less than half of the originally formed FeS.

Table 3 presents the mass balance of sulfur species in column 2. During the 219 days of

experiment, the total SO4 input into column 2 was 72 mmol, and the measured total SO4

output from the end of this column was 16 mmol. The latter value was estimated by

integrating and averaging the total SO4 output with respect to time. The difference between

total SO4 input and output (56 mmol) is equivalent to the total S(-II) produced in column 2.

A total dissolved S(-II) output mass of 23 mmol was estimated from our measurements. The

total reacted S(-II) of 33 mmol during the course of the experiment is simply the difference

between the total S(-II) produced and the S(-II) released from the column. This is in good

agreement with the total mass of solid-phase sulfur species produced from the dissimilatory

Fig. 7. Concentration of solid-phase reduced sulfur in column 2 extracted by AVS and CRS techniques. The

background concentrations of AVS and CRS in oxic solid phase are 0.3F 0.1 and 16.0F 7.1 Amol S/g solid,

respectively.


SO4 reduction as determined by the CRS method (30 mmol). The two values are

highlighted in bold in Table 3.

Heron et al. (1994a) defined the OXC of an aquifer as its ability to restrict the

migration of a reduced plume. Operationally, they quantified the OXC as the equivalent

concentration of Ti(III)–EDTA reducible Fe(III). Assuming that all of the AVS (15.18

mmol) consists of sulfur in FeS and all of the CRS (30 mmol) consists of sulfur in FeS

and FeS2, a total mass of 22.5 mmol of available Fe(III) is estimated in column 2. This is

in very good agreement with the initial mass of Ti(III)–EDTA reducible Fe(III) of 22.0

mmol (Table 2). This observation suggests that the Ti(III)–EDTA reducible Fe(III) is

indeed a good indicator for the OXC of an aquifer where the principal reductant is

dissolved S(-II).

4.6. Geochemical modeling of column 2

The geochemical processes in column 2, where significant SO4 reduction takes place,

have been quantified using the computer program STEADYQL. One of the aims of this

modeling effort is to evaluate the sensitivity of the modeled dependent species (DIC, pH,

propionate and acetate) with respect to the kinetics of the biogeochemical processes. Of

special interest here is to evaluate the effect of sulfur cycling processes such as reduction

of SO4 and direct reaction of S(-II) with surface Fe(III) hydroxides (Eq. (1)) on the

development of pH and alkalinity. Comparison between measured and simulated results

also allows us to validate our assumptions.

We present here the modeling results of the aqueous phase data of column 2 collected

on day 85. The concentrations of O2, NO3, SO4 and propionate in the influent were 0.25,

0.10, 0.55 and 1.0 mM, respectively. The influent DIC concentration was 4.7 mM and pH

was 7.7. Equilibrium reactions reported elsewhere have been used for speciation

calculations (Amirbahman et al., 1998). The microbially catalyzed kinetic processes

relevant to this data are listed in Table 1.

Zero-order rate coefficients for the microbially catalyzed redox processes were obtained

from the changes in concentrations of the relevant independent species and by taking into

account the mean residence time in each box reactor as described previously. Rates of

Table 3

Mass balance of sulfur species in column 2

Measurements and procedures Mass (mmol)

Sulfate inputa 72

Sulfate outputa 16

S(-II) producedb 56

S(-II) outputa 23

S(-II) reactedc 33

Acid volatile sulfide producedd 15

Cr(II)-reducible sulfur producedd 30

a Total dissolved mass over the course of experiment.b The difference between sulfate input and output.c The difference between S(-II) produced and S(-II) released from the column, initially forming FeS.d Total sulfur in the solid phase measured directly and integrated over the entire length of column.


change of concentrations of O2, NO3 and SO4 were used to predict the concentrations of

propionate and acetate. Aerobic respiration, denitrification and SO4 reduction all use

propionate as the carbon source (Table 1). Aerobic respiration was modeled only in the

first 0.5 cm of column 2, where complete reduction of 0.25 mM of O2 was observed.

Denitrification process was observed up to 2 cm from the inlet, where the entire added 0.1

mM NO3 was reduced. Sulfate reduction was considered for the entire length of the

column, even though SO4 was reduced mainly close to the inlet. This process produces 1

mol of acetate for every mole of oxidized propionate (Table 1).

Dissimilatory Fe and Mn reduction processes use acetate as the organic carbon source

in this study. However, no decrease in acetate concentration was observed due to Fe and

Mn reduction processes, since the concentrations of produced metals are nearly two orders

of magnitude less than the concentration of available acetate. Rates of Fe and Mn

reduction processes measured in column 1 in the absence of significant SO4 reduction

were used for column 2.

The measured and simulated concentrations of the two organic acids are shown in Fig.

8. Sulfate reduction was originally considered as the only dissimilatory process here that

produces acetate by oxidation of propionate. However, changes in SO4 concentration

could not account for the entire production of acetate. To model the distribution of acetate,

fermentation of propionate to acetate and hydrogen gas was considered (Table 1; Thauer et

al., 1977). This process was considered only between 1 and 1.5 cm from the inlet and

accounts for the production of nearly 15% of acetate on day 85. Fermentation of

propionate to acetate in column 2 started sometime after the 10th week of operation

and increased in rate with time. By the end of operation on day 219, this process was

Fig. 8. Concentrations of propionate and acetate in column 2 on day 85. Circles are experimental measurements

and lines are STEADYQL simulations.


complete at the end of column 2 with all of propionate converted to acetate (data not

shown).

The modeled aerobic respiration, denitrification, SO4 reduction and propionate

fermentation processes together account for 92% of consumption of propionate, whereas

the measured data indicate a 98% consumption. This difference may be due to the

consumption of propionate by the bacteria to generate biomass. Such a process is expected

to take place close to the inlet of the column, due to higher density of bacteria.

Even though the saturation index of solution with respect to calcite varies from 0.5

close to the inlet to 0.9 at the end of the column, precipitation of CaCO3 is not observed.

Using typical rate coefficients for precipitation of CaCO3(S) estimated under field

conditions in a similar porous media (Amirbahman et al., 1998), the expected decrease

in the concentration of dissolved Ca from the inlet to the outlet of column 2 would be in

the order of 1 AM. This is more than three orders of magnitude smaller than the measured

Ca concentrations at any location along the columns.

The simultaneous reduction of SO4 and Fe in the same zone has been previously

reported and may be attributed to the low availability of solid-phase Fe(III) hydroxides

(Canfield, 1989). In column 2, the low dissolved concentrations of Fe(II) indicate that the

distribution of this species is controlled by the solubility of amorphous FeS. This low

concentration of Fe(II) keeps the solution well below the solubility of siderite.

Reaction of dissolved S(-II) with the surface Fe(III) hydroxides was considered as a

kinetically controlled process. Due to the small concentrations of dissolved Fe(II) (0.1–1

AM), precipitation of FeS was considered as an equilibrium reaction (Eq. (3)). von Gunten

and Furrer (2000) have observed kinetically controlled precipitation of FeS under similar

conditions. As discussed below, however, inclusion of such a process has a very small

effect on the overall S(-II) balance in column 2 and development of pH and alkalinity.

The rate of reaction of dissolved S(-II) with the surface Fe(III) hydroxides was

calculated for each box by taking into account the total reduced SO4 concentration, the

change in the dissolved S(-II) concentration and the concentration of S(-II) that is

precipitated as FeS (Furrer et al., 1996). To calculate the latter concentration in column

2, a rate coefficient of 2.5 nM s� 1 for the dissimilatory Fe reduction was considered. This

rate was observed close to the inlet of column 1 on day 85, where SO4 reduction was

negligible, and was applied uniformly along the length column 2. Given the data presented

in the previous section, this rate represents a high end estimate for the rate of production of

dissolved Fe(II) in both columns. The difference in the estimated total and measured Fe(II)

concentrations is the precipitated concentration of Fe(II), which is also equivalent to the

concentration of precipitated S(-II).

Direct reaction with surface Fe(III) hydroxides constitutes a more important sink for the

dissolved S(-II) than precipitation of FeS in column 2. Experimental measurements

indicate that a total of 0.53 mM of SO4 was reduced to S(-II) on day 85. Approximately

99% of the produced S(-II) reacts in the column on day 85. Considering a rate coefficient

of 2.5 nM s� 1 for the dissimilatory Fe(III) reduction in column 2, a total maximum

concentration of 63 AM of Fe(II) would be generated in this column by this mechanism.

Given that 21 AM of Fe(II) leaves the column, a total of 42 AM of Fe(II) has been

precipitated as FeS. The precipitated FeS is therefore, more than an order of magnitude

smaller than the total reacted S(-II).


Results of STEADYQL simulations of column 2 are shown in Fig. 9. This figure

exhibits the measured and modeled values of pH and DIC for day 85. The modeled values

are in reasonable agreement with the measured values, indicating that our basic assump-

tions are correct. Sensitivity analysis shows that close to the inlet of the column, aerobic

respiration, denitrification and SO4 reduction processes largely control the evolution of pH

and DIC. Propionate fermentation to acetate, where considered, has a large effect on

lowering the pH. Toward the end of column2, where the direct reaction of dissolved S(-II)

with surface Fe(III) hydroxides is important, this process controls almost all of the increase

in pH. In column 1, the Mn and Fe reduction processes do not contribute to the evolution

of pH and DIC significantly. In this column, pH and DIC evolution are controlled by

aerobic respiration and denitrification processes.

Fig. 9. (a) pH, and (b) dissolved inorganic carbon concentrations in column 2 on day 85. Circles are experimental

measurements and lines are STEADYQL simulations.


5. Conclusions

We have studied the evolution of reducing environments in two separate flow-through

packed bed reactors. All geochemical transformations were directly or indirectly related to

the microbial oxidation of organic matter by various electron acceptors. The rates of

enzymatic reduction of Mn(IV) and Fe(III) exhibited an initial increase and a gradual

decrease in time, due to the preferential dissolution of the more reactive mineral phases.

Regardless of the presence of dissolved S(-II), dissolved Mn(II) concentration was

controlled by the solubility of rhodochrosite. Dissolved Fe(II) concentration was con-

trolled by the solubility of siderite or amorphous FeS, depending on the concentration of

dissolved S(-II). However, in the column where Fe(II) concentration was controlled by

dissolved S(-II), the most important mechanism for the reduction of Fe(III) hydroxide was

the direct attack by the dissolved S(-II). For the duration of our experiment, enzymatic

Fe(III) reduction contributed only to a small extent to the total mass of reduced Fe (19%)

due to the slower kinetics of this process than the enzymatic sulfate reduction process and

the reactions depicted in Eqs. (1)–(3).

Acknowledgements

The authors wish to thank Dr. A. Ammann and T. Ruttimann of EAWAG for their help

in analyzing the samples. Professor R. Giovanoli of the University of Bern performed the

SEM and XRD analysis of the soil samples. Critical comments of three reviewers are

greatly acknowledged.

References

Abbaspour, K., Matta, V., Huggenberger, P., Johnson, C.A., 1999. A contaminated site investigation: comparison

of information gained from geophysical measurements and hydrogeological modeling. J. Contam. Hydrol. 40,

365–380.

Amirbahman, A., Sigg, L., von Gunten, U., 1997. Reductive dissolution of Fe(III) (hydr)oxides by cysteine:

mechanism and kinetics. J. Colloid Interface Sci. 194, 194–206.

Amirbahman, A., Schonenberger, R., Johnson, C.A., Sigg, L., 1998. Aqueous- and solid-phase biogeochemistry

of a calcareous aquifer system downgradient from a municipal solid waste landfill (Winterthur, Switzerland).

Environ. Sci. Technol. 32, 1933–1940.

Ammann, A.A., Ruttimann, T.B., 1995. Simultaneous determination of small organic and inorganic anions in

environmental water samples by ion-exchange chromatography. J. Chromatogr., A 706, 259–269.

Arakaki, T., Morse, J.W., 1993. Coprecipitation and adsorption of Mn2 + with mackinawite (FeS) under con-

ditions similar to those found in anoxic sediments. Geochim. Cosmochim. Acta 57, 9–15.

Berner, R.A., 1984. Sedimentary pyrite formation: an update. Geochim. Cosmochim. Acta 48, 605–615.

Canfield, D.E., 1989. Reactive iron in marine sediments. Geochim. Cosmochim. Acta 53, 619–632.

Canfield, D.E., Raiswell, R., Westrich, J.T., Reaves, C.M., Berner, R.A., 1986. The use of chromium reduction in

the analysis of reduced inorganic sulfur in sediments and shales. Chem. Geol. 54, 149–155.

Davison, W., 1991. The solubility of iron sulfides in synthetic and natural waters at ambient temperature. Aquat.

Sci. 53, 309–329.

Furrer, G., Westall, J., Sollins, P., 1989. The study of soil chemistry through quasi-steady-state models: I.

Mathematical definition of model. Geochim. Cosmochim. Acta 53, 595–601.


Furrer, G., von Gunten, U., Zobrist, J., 1996. Steady-state modelling of biogeochemical processes in columns

with aquifer material: 1. Speciation and mass balances. Chem. Geol. 133, 15–28.

Heron, G., Christensen, T.H., 1994. The role of aquifer sediment in controlling redox conditions in polluted

groundwater. In: Dracos, J., Stauffer, F. (Eds.), Transport and Reactive Processes in Aquifers. Balkema,

Rotterdam, The Netherlands, pp. 73–78.

Heron, G., Christensen, T.H., 1995. Impact of sediment-bound iron on redox buffering capacity in a landfill

leachate polluted aquifer (Vejen, Denmark). Environ. Sci. Technol. 29, 187–192.

Heron, G., Christensen, T.H., Tjell, J.C., 1994a. Oxidation capacity of aquifer sediments. Environ. Sci. Technol.

28, 153–158.

Heron, G., Crouzet, C., Bourg, A.C.M., Christensen, T.H., 1994b. Speciation of Fe(II) and Fe(III) in contami-

nated aquifer sediments using chemical extraction techniques. Environ. Sci. Technol. 28, 1698–1705.

Hoehn, E., Johnson, C.A., Huggenberger, P., Amirbahman, A., Peter, A., Zweifel, H.R., 2000. Investigative

strategies and risk assessment of old unlined municipal solid waste landfills. Waste Manag. Res. 18, 577–589.

Jacobs, L.A., von Gunten, H.R., Keil, R., Kuslys, M., 1988. Geochemical changes along a river-groundwater

infiltration flow path: Glattfelden, Switzerland. Geochim. Cosmochim. Acta 52, 2693–2706.

Jakobsen, R., Postma, D., 1999. Redox zoning, rates of sulfate reduction and interactions with Fe reduction and

methanogenesis in a shallow sandy aquifer, Romo, Denmark. Geochim. Cosmochim. Acta 63, 137–151.

Kornicker, W.A., Morse, J.W., 1991. The interactions of divalent cations with the surface of pyrite. Geochim.

Cosmochim. Acta 55, 2159–2172.

Kostka, J.E., Luther, G.W., 1994. Partitioning and speciation of solid phase iron in saltmarsh sediments. Geo-

chim. Cosmochim. Acta 58, 1701–1710.

Lovley, D.R., Phillips, E.J.P., Lonergan, D.J., 1991. Enzymatic versus nonenzymatic mechanisms for Fe(III)

reduction in aquatic sediments. Environ. Sci. Technol. 25, 1062–1067.

Luther, G.W., Church, T.M., Kostka, J.E., Sulzberger, B., Stumm, W., 1992. Seasonal iron cycling in the marine

environment: the importance of ligand complexes with Fe(II) and Fe(III) in the dissolution of Fe(III) minerals

and pyrite, respectively. Mar. Chem. 40, 81–103.

Morse, J.W., Luther, G.W., 1999. Chemical influences of trace metal sulfide interactions in anoxic sediments.

Geochim. Cosmochim. Acta 63, 3373–3378.

Stauffer, F., Dracos, T., 1986. Experimental and numerical study of water and solute infiltration in layered porous

media. J. Hydrol. 84, 9–34.

Thauer, R.K., Jungermann, K., Decker, K., 1977. Energy conservation in chemotrophic anaerobic bacteria.

Bacteriol. Rev. 41, 100–180.

von Gunten, U., Furrer, G., 2000. Steady-state modelling of biogeochemical processes in columns with aquifer

material: dynamics of iron–sulfur interactions. Chem. Geol. 167, 271–284.

von Gunten, U., Zobrist, J., 1993. Biogeochemical changes in groundwater-infiltration systems: column studies.

Geochim. Cosmochim. Acta 57, 3895–3906.

von Gunten, H.R., Karametaxas, G., Keil, R., 1994. Chemical processes in infiltrated riverbed sediment. Environ.

Sci. Technol. 28, 2087–2093.

Zhabina, N.N., Volkov, I.I., 1978. A method of determination of various sulfur compounds in sea sediments and

rocks. In: Krumbein, W.E. (Ed.), Environmental Biogeochemistry; Methods, Metals and Assessments, vol. 3.

Ann Arbor Science Publishers, Ann Arbor, MI, pp. 735–745.


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