Pollutant Reduction in Heterogeneous Fe(II)-Fe(III) - American

Chapter 17

Pollutant Reduction in Heterogeneous Fe(II)-Fe(III) Systems 1,3

2

Downloaded by NORTH CAROLINA STATE UNIV on September 26, 2012 | http://pubs.acs.org Publication Date: January 15, 1999 | doi: 10.1021/bk-1998-0715.ch017

Stefan B. Haderlein and Klaus Pecher 1

Ralph M. Parsons Laboratory, Department of Civil and Environmental Engineering, Massachusetts Institute of Technology, Cambridge, MA 02139 Department of Hydrology, University of Bayreuth, D-95440 Bayreuth, Germany

2

Several classes of priority pollutants can undergo abiotic reductive transformation reactions in anoxic aqueous environments. Such reactions may be environmentally benign since they often lead to products which are more bioavailable and more easily to degrade. In certain cases, however, the products of such reactions are more toxic than the parent compounds. Thus, knowledge of the reductants and processes involved in the reductive transformation of pollutants is essential to evaluate the risk of pollution and the need for remediation of contaminated sites. In recent years, increasing evidence for the participation of ferrous iron in reductive transformation of pollutants became available through laboratory and field studies. The pollutants studied include halogenated solvents (e.g., (1-3)), nitroaromatic compounds (4-6), nitrite (7, 8), nitrate (9), chromate (10, 11), selenate (12), and pertechnetate (13). A common feature of most of these laboratory studies is that ferrous iron species present at mineral surfaces were much more reactive than dissolved Fe(II) present in the background electrolyte. These findings can be rationalized by inspection of Figure 1. Shown are the reduction potentials in aqueous solution, E °(w), of some environmentally relevant iron redox couples as well as of selected organic pollutants (for details see (14, 15)). As can be seen, iron is an extremely versatile redox sensitive element. The redox potentials of geochemically important forms of iron vary by more than 1000 mV, depending on which species are involved. Species which stabilize Fe(III) formed in the reaction tend to exhibit the lowest redox potentials. Aqueous Fe(II) is a poor reductant compared to minerals containing structural iron such as magnetite or iron silicates which can reduce polyhalogenated alkanes, nitroaromatic compounds , or aromatic azo compounds under standard environmental conditions. Since oxidation of iron(II) to iron(III) involves the transfer of only one electron, the reduction potential of a given iron(III)/iron(II) couple can be directly applied to evaluate the relative reactivities of iron(II) species in various chemical environments. However, this consideration is only applicable if the actual electron transfer from Fe(II) to a pollutant is at least in part rate limiting. This can not always be assumed, particularly for heterogeneous reactions, where in the course of the reaction the properties of the reductant (i.e., the surface) is changed (see below). Although minerals containing structural Fe(II), e.g., phyllosilicates, were found to h

1

Current address: EAWAG/ETH Zurich, CH-8600 Dubendorf, Switzerland (Fax: +41-1-823 5471; e-mail: [email protected]).

342

© 1999 American Chemical Society

In Mineral-Water Interfacial Reactions; Sparks, D., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1999.


343

Eh(w) (volt)

Red

Ox

CI C-CCI

CCI3-CCI3

Red

Ox

2

Eh(w)

(volt)

2

1.0

1.0 CBr

4

— CHBr

3

Fe (aq)

Fe (aq)

2+

3+

CCI4

0.5

— CHCI3

CHCI3

CH CI 2

2

0.5 Fe""OH2+ Fe»«sal

— Fe OH n

Fe sal

+

n

Fe porph m

0.0 2^-NH

+

Fe(OH) am(s)

Fe porph M

0.0

Fe (10-5M) Fe 0 (s) oc-FeOOH(s) — cx-FeOOH(s) Fe2+ (10-5M) a - F e 0 ( s ) — — Fe2+ (10-5M) a-FeOOH(s)"" - " F e C 0 ( s ) 2+

3

2

3

4

%

2

3

3

Fe 0 (s) 3

4

— Fe Si0 (s) 2

4

-0.5

4 -0.5

Figure 1. Representative redox couples of various organic pollutants and iron species. Given are standard potentials at pH = 7 and molar concentrations of the reactants but at standard environmental concentrations of the major anions involved: [HCO3-] = [C1-] = ΙΟ" M ; [Br~] = 10" M ; (s) = solid. Adapted from (44). 3

5


344

reduce organic substances such as CCI4 (76), the rate constants for such reactions typically are very low. Obviously, the reactivity of the various Fe(II) species cannot be predicted from their £ ° ( ) values alone. Other processes than the electron-trans fer, such as sorption, precipitation or dissolution reactions, may be rate-limiting and must be considered as well to evaluate the reactivity of Fe(II) species. The purpose of this chapter is to give an overview of the chemical and biologi cal processes that control the reactivity of Fe(II) in heterogeneous aqueous systems with respect to pollutant transformation. To this end, we will evaluate data collected in various laboratory systems as well as field studies. Two classes of model com pounds with complementary properties will be used to monitor the reactivity of Fe(II) species in the various systems. Nitroaromatic compounds (NACs) primarily served to characterize the systems in terms of mass and electron balances. Reduction of NACs by Fe(II) species results in only a few major products (aromatic amines and hydroxylamines) which can be easily quantified by standard HPLC-UV methods in the low μΜ range. Polyhalogenated aliphatic compounds (PHAs) were used if little perturba tion of the systems in terms of electron transfer to the organic substrates was crucial. Reduction of PHAs requires fewer electrons than nitro reduction and PHAs can be quantified by standard GC-ECD methods in the low ppb range. w


n

Factors C o n t r o l l i n g the Reactivity o f Ferrous I r o n i n Heterogeneous Aqueous Systems

Figure 2a shows a typical example of the reactivity of ferrous iron at ambient pH in (anoxic) aqueous solutions or suspensions containing Fe(M) bearing minerals, respec tively. Significant reactivity of aqueous Fe(II) was observed neither for NACs (17) nor for PHAs (18). Addition of Fe(III)(hydr)oxides or magnetite to anoxic solutions of Fe(II) resulted in rapid reduction of the organic substrates. Apparently, uptake of Fe(II) by such surfaces created very reactive Fe(II) species. Systems containing dissolved Fe(II) and iron oxy-hydroxide minerals showed a high reactivity as long as sufficient aqueous Fe(II) was available to replace Fe(II) sur face species that had reacted with NACs (Figure 2b). During the entire course of the reaction, the number of electrons transferred to NACs matched the consumption of aqueous Fe(II). Thus, reactive Fe(II) species formed at these surfaces can be regener ated rapidly by uptake of Fe(II) from solution, once they have been consumed by oxi dation of pollutants. Although the reduction of NACs with most Fe(II)-containing minerals thermodynamically is highly favorable at neutral pH (Figure 1), virtually no transformation was observed in suspensions of magnetite without added aqueous Fe(II). Long-term reac tion rates of iron sulfides or phyllosilicates such as biotite and vermiculite for CCI4 or NACs in aqueous suspension were also found to be very slow at near-neutral pH val ues (16, 19-22). Because the oxidative dissolution of such minerals by organic pollu tants under environmental conditions is a relatively slow process, it seems rather unlikely that such reactions can compete with the significant rates of formation and regeneration of reactive, surface-bound Fe(II). In the following, the various factors that control the reactivity of Fe(II) at iron oxy-hydroxide surfaces will be reviewed. p H Value. The pH of the system is a master variable that is related to several factors affecting the reactivity of heterogeneous Fe(II)/Fe(III) systems. With increasing pH values, observed pseudo-first-order reaction rates with respect to pollutant transfor mation increased for both NACs (Figure 3a) and PHAs (18). Apparently, uptake of Fe(II) from aqueous solution by iron oxy-hydroxide surfaces, which is a prerequisite for formation of reactive Fe(II) surface species, is strongly pH dependent (Figure 3b and (23)). However, other important system variables also change with pH, namely the redox potential and the speciation of both aqueous Fe(II) and surface sites of the



345

Ο

15

0

60

30 Time (min)

120

45

2500

180

480

Time (min)

Figure 2. a) Reduction of 50 μΜ 4-chloro nitrobenzene (4-C1-NB) in the presence of 17 m L magnetite and an initial concentration of 2.3 m M Fe(II) at pH 7.0 and 25 °C (•). The rate law deviates from pseudo-first-order behavior for longer obser vation times. 4-C1-NB was not reduced significantly in suspensions of magnetite without dissolved Fe(II) (V) or in solutions of Fe(II) without magnetite (Δ); (adap ted from (77)) b) Electron balance (x) for the reduction of 4-C1-NB repeatedly added at times t = 0, 6, 17 min to a suspension containing 11.2 m L magnetite and an initial concentration of dissolved Fe(II) of 1.6 m M at pH = 7.75 and Τ = 25 °C; (adapted from (6)). 2

_ 1

2

_ 1



346

3.0-

PH

Figure 3. a) Pseudo-first-order rate constants, fc^as a function of pH for the reduction of 50 μΜ nitrobenzene in suspensions of 11.2 m L magnetite and an initial concentration of 1.5 mM Fe(II). b) Adsorption of 0.2 mM Fe(II) onto 0.55 g L magnetite as a function of pH (adapted from (17)). 2

_ 1

_1



347 minerals. Moreover, different types of reactive Fe(II) species may be formed at surfaces, depending on the amount of Fe(II) uptake. Thus, we will try to discuss the reactivity of the system by considering variations in only one of these variables at a time. During reduction of nitrobenzenes in heterogeneous Fe(II)/Fe(III) systems, a pseudo-first-order rate law was obtained only for the initial phase of the experiments, followed by a decrease in reaction rates (Figure 2a). Such behavior may indicate that the most reactive Fe(II) species were depleted and less reactive species and/or the changing properties of the system controlled the reaction rates at later stages of the experiments. The further evaluation of the reactivity of heterogeneous Fe(II)/Fe(III) systems will be based on the reduction of polyhalogenated alkanes. In contrast to NACs, which take up six electrons per nitro group during the reduction to the respective amines, reduction of hexachloroethane and dibromodichloromethane to tetrachloroethene and monobromdichloromethane, respectively, consumes only two electrons per molecule. Furthermore, reduction of these PHAs can be studied at much lower concentrations. The reduction of PHAs by surface-associated Fe(II) caused sufficiently low perturbation of the system properties to result in pseudo-first-order reduction kinetics during the entire course of the reaction (18). Remodeling of F e ( I I ) at Iron(hydr)oxides w i t h T i m e . The reactivity of Fe(II) in

suspensions of goethite strongly depended on how long the suspended minerals had been exposed to dissolved Fe(II) before the organic substrate was spiked to the systems. Figure 4 shows normalized reaction rates, k' rm, for two PHAs as a function of this equilibration time, t . Reaction rates increased with increasing t and somehow levelled off after about 20 h. This behavior was independent of the pollutant studied since it was observed for both PHAs and NACs. It demonstrates that Fe(II) bound to the oxide surface alters its reactivity with time and that during the course of interaction of ferrous iron with the goethite surface various reactive Fe(II) species are formed. Such remodeling of surface-bound Fe(II) species is corroborated by results of experiments where Fe(II) was desorbed from goethite as a function of equilibration time of Fe(II) and goethite. Figure 5 shows results of such experiments where phenanthroline was used to desorb Fe(II) (24)). The non desorbable fraction of surface-bound ferrous iron increased with equilibration time, t q. until a plateau was reached after about 20 h. Thus, under appropriate reaction conditions (pH, surface density of Fe(II)), more strongly bound Fe(II) species are being formed with increasing equilibration time, consistent with a remodeling of the surface in terms of formation of surface precipitates or surface clusters. Coughlin & Stone (23) also reported considerable uptake of Fe(II) by goethite after 18 h equilibration time even under acidic conditions (pH 4 - 5.5). Similar to our study they found that a significant fraction of surface-bound ferrous iron was not desorbable, even after treatment with H N 0 for 28 h. n0

eq

eq

e

3

Surface Density of Fe(II)-Species. Figure 6 shows the rate constants for the reduction of dibromodichloromethane in suspensions containing goethite and Fe(II) as a function of total ferrous iron present and pre-equilibration time of Fe(II) with the surface. A strong dependence of pseudo-first-order reaction rates on total ferrous iron concentrations was observed for long pre-equilibration times (te > 30 h) which provides further evidence that surface species of Fe(II) formed after prolonged contact of ferrous iron with iron(hydr)oxide surfaces are most reactive. Experiments such as shown in Figure 6 do not allow one to calculate second-order rate constants as it is remains unclear which species or fraction(s) of surface-bound Fe(II) is involved in the reaction. The increase of reduction rates with higher concentrations of Fe(II) t may also be due to a decreasing reduction potential of the system. Assuming that the redox potential of the Fe(II)/Fe(III) couple in homogeneous solution represents the redox q

to



ο' • • • • Ο

10

20

30

40

50

'

Uh] Figure 4. Dependence of normalized pseudo-first-order rate constants (k^orm) on pre-equilibration time of F e with goethite (te ( pH 7.2, Fe(II) t=l mM, 1=20 mM, T=25°C, 25 m L"* goethite, ( • ) : C (C Cl6)=0.4 μΜ, (•): C (CCl Br )=5 μΜ); adapted from (7). 2 +

q

to

2

0

2

0

2

2

Figure 5. Non-desorbable surface bound Fe(II) as a function of contact time of Fe(II) with suspended goethite (pH 7.2, Fe(II) t=l mM, 1=20 mM, T=25°C, 25 m L goethite); adapted from(24). to

2

_ 1


349

potential of the system reasonably well and that the ratios of aqueous Fe(II)/Fe(III) for the two sets of experiments shown in Figure 6 are similar at the respective Fe(II)tot concentrations, the results suggest that changes in the redox potential were not a major factor for the increase in reaction rates with Fe(II) .


tot

Interaction of Fe(II) with Iron(hydr)oxides - Formation and Characterization of Reactive Sites. Adsorption kinetics of Fe(II) to iron oxide minerals was studied under strictly anoxic conditions (glove box). Kinetics of adsorption at Fe(II) surface saturation > 60% exhibited two distinct features (shown in Figure 7 for several iron oxides at pH = 7.2). In an almost immediate first step, a certain fraction of the ferrous iron added was bound to the oxides. The different amounts of initially adsorbed Fe(II) for the various minerals may be due to differences in electrostatic forces (different pH of the minerals) as well as morphologic features (differences in crystallographic planes, kinks, steps, edges, etc.). In a second step, the fraction of bound ferrous iron slowly increased and levelled off after about 30 h. At given Fe(II)tot» this behavior was significant for goethite at pH values > 6.0 (Figure 8). The suspensions showed a distinct change in color from bright yellow to blue-green at pH > 7.6, which may indicate the formation of green rust phases during the experiments (8, 25). Two-step adsorption kinetics of metal cations at iron oxides has been observed for several metals (26-28) and mainly attributed to interaction of the metal cations with different sites of the metal oxide surface (29). Diffusion of cations into micro- or macropores can be ruled out for synthetic goethites (30) due to an essential lack of porosity as determined by t-plot analysis (31). However, diffusion processes may be important for other oxides, especially amorphous iron oxides or oxides with certain coatings and oxide surfaces which have been subject to substantial dissolution processes. Thus, other processes than diffusion must have caused the two different kinetic domains of Fe(II) sorption to iron oxide surfaces. Several authors proposed slow electron transfer reactions, induced by adsorption of Fe(II) (32-36), at the interface of spinel-like iron oxides containing Fe(II) and Fe(III) and also provided spectroscopic evidence for product formation (37). However, these results must be carefully evaluated since contamination of the systems with molecular oxygen cannot be excluded, especially in the early studies. Thus, the experimental conditions and systems reported in this literature are not directly comparable to our study where predominantly non mixed-valence iron oxides were investigated under strict exclusion of molecular oxygen. Direct electron transfer between adsorbed Fe(II) and structural Fe(III) at the goethite surface seems unlikely as Fe(III) within the crystal structure is stabilized by neighbouring oxygen ligands. An alternative explanation for the observed slow kinetics of Fe(II) uptake is the formation of surface precipitates. The greenish color developed in goethite suspensions at pH > 7.6 may indicate Fe(II)/Fe(III)-precipitates, possibly green rust phases. Traces of Fe(III), a prerequisite for the formation of such mixed valence iron phases, may either have been present as impurity in the Fe(II) spike solutions or due to dissolution of the iron(hydr)oxides. Incorporation of divalent metal cations into goethite or hematite has recently been shown during aging of hydrous iron oxides (38). The hypothesized Fe(II)/Fe(IH) surface precipitates, which seem to be responsible for the high reactivity of systems studied at high pH values, are currently analyzed by surface spectroscopic techniques. Despite the missing spectroscopic characterization of Fe(II) surface precipitates in such systems, analysis of adsorption isotherms of ferrous iron on several iron oxides (24) supports this hypothesis. Modeling of Fe(II) adsorption isotherms considering surface precipitation reactions (39) gave excellent fits to adsorption data (Figure 9). The model used allows for a transition from surface complexation to surface precipitation, which becomes relevant at surface saturations exceeding 20% (26). Fitted surface site density data of the studied iron oxides (Table I) agree with experimental data. Furthermore, the solubility product fitted for a hypothetical Fe(OH)2(s) p z c



350

1

2

3

Fe(II) [mM] lol

Figure 6. Pseudo-first-order rate constant of transformation of C C l B r in suspen sions of goethite (25 m L ) as a function of concentrations of Fe(II) t (pH=7.2, 1=20 mM, T=25°C, C (CCl Br )=0.5 μΜ, (•): t >30 h, (A):teq=0 h); adapted from (7). 2

2

2

- 1

to

0

2

2

eq

Fe(n) = 1 mM w

A-A-A

2 °Β

A

A

A

1

A — A

wt

0.8

8

0.6

0.6 0.4 0.2

Fe(II) = 1 mM

0.4 •-• Α-A ·-· •-•

Goethite Hematite Lepidocrocite Magnetite 20

40

0.2

60

80

100

t [min]

10

20

30

40

50

60

t M

eq

eq

Figure 7. Uptake kinetics of dissolved Fe(II) by different iron oxides in anoxic aqueous supensions (pH 7.2, 25 m L" ,1=20 mM, T=25°C, Fe(II) i= dissolved concentration of ferrous iron). The difference between Fe(n) t and Fe(II) i gives surface-bound ferrous iron. Adapted from (7). 2

1

so

to


so


8.91

100

50

time [h] Figure 8. Uptake kinetics of dissolved Fe(II) by goethite in anoxic aqueous supensions at different pH values (25 m L" ,1=20 mM, T=25°C, Fe(II) i= dissolved concentration of ferrous iron). The difference between Fe(U) t and Fe(II) i gives surface-bound ferrous iron. Adapted from (24). 2

1

so

to

so

c** [M] Figure 9. Isotherm fits of the surface precipitation model for the sorption of ferrous iron on iron oxides: ( • ) magnetite, (•) goethite, ( · ) lepidocrocite, (A) hematite (for parameter values see Table I, Fe(II) i = concentration of dissolved ferrous iron, Γρ = concentration of surface-bound ferrous iron per total concentration of iron oxide, pH 7.2,1=20 mM, T=25°C, 25 m L " , t =15 min); adapted from (1). so

β

2

1

eq


352 precipitate closely matched the experimental value reported by Baes & Mesmer (40), independent of the iron oxide used. However, we are aware that this model is a first approximation and needs to be refined in terms of pH-dependent isotherms and inclusion of measured stoichiometrics of the surface precipitates. Surface precipita tion may be neglected at lower total ferrous iron concentrations due to low surface coverage of sorbed Fe(II) (41). Table I . Calculated Parameter Values o f Isotherm Fits o f the Surface Precipitation M o d e l f o r Uptake o f F e ( I I ) b y Various Iron(hydr)oxides. a


b

Mineral

Kads

magnetite

0.014

4.7 10-

12

3.9 ΙΟ"

goethite

0.013

1.5 10-

13

1.3 10-

hematite

0.0026

2.5 10-

13

4.7 10-

e

S

c

T 4

4

5

lepidocrocite 1.06 100.096 1.4 10shown in Figure 9; for details see Ref. (39); K ds=equilibrium coefficient for adsorption; K M=solubility product of the metal hydroxide formed; ST=surface site density [M]. 4

13

a

b

c

a

d

e

sp

Significance o f Surface-Bound Fe(II) f o r Pollutant Reduction i n N a t u r a l Systems

In natural anoxic matrices, several biogeochemical redox processes may take place in parallel, giving rise to several potential reductants for subsurface contaminants. Such species may be present in aqueous solution or at solid surfaces, which complicates chemical analyses. Under such complex redox conditions, the traditional approach to predicting the fate of reactive pollutants often fails due to inadequate system characterization. The use of reactive tracers as in situ probes may help to identify the predominant reductants in anoxic natural systems (42). Riigge et al. (43) demonstrated how in situ information on reductive transfor mation processes in anaerobic aquifers can be obtained using NACs as reactive tracers. NACs can be used to identify the predominant reductants since the effects of N A C substituents on reduction rates are characteristic for various important subsurface reductants (44). Furthermore, depending on the pattern of substituents, reduction rates of NACs by a given species may vary by orders of magnitude. This allows one to study reactions at very different time scales by choosing an appropriate set of substituted NACs as probe compounds. The reactivity pattern of five nitroaromatic probe compounds was studied in an anaerobic leachate plume of a landfill at Grindsted, Denmark, by means of an injection experiment in the plume area, by in situ microcosm experiments installed along a flow line in the plume, and by laboratory batch experiments containing groundwater and aquifer sediment. The in situ reactivity of the compounds in the aquifer was compared to their reactivity in well defined model systems to delineate processes and reductants that were active in the leachate plume. The redox conditions in the plume were typical for complex redox environments with several processes occurring simultaneously (Figure 10a). Closest to the landfill, methane production and sulfate reduction prevailed, followed by sulfate- and iron reduction. Generally, the groundwater was rich in dissolved or colloidal organic matter (DOM = 50-120 mg C rg L ) and in dissolved inorganic Fe(II) (>1.4 mM). Since reduction of NACs by dissolved H S , Fe(II),(45) and Fe(II) bound to D O M relatively slow, the major _ 1

0

2



353

reductants for NACs expected under these conditions include microorganisms, reduced D O M , and Fe(n) associated with iron(hydr)oxide surfaces. In the aquifer, dinitrobenzenes generally were reduced at very high rates and without any noticeable lag phase. These compounds were always reduced first within the mixture of the NACs studied. These results and the very similar pattern of Ν A C reduction under unchanged and microbially deactivated conditions suggest that N A C reduction was dominated by abiotic reactions rather than biotransformation. The presence of aquifer sediment strongly enhanced the reduction of all NACs compared to filtrated groundwater. Throughout the plume, a very similar pattern of N A C re duction was found, indicating that the same type of reactions occurred in the entire anaerobic part of the aquifer. The reactivity pattern of the NACs within the aquifer resembled very much the pattern observed in model systems where surface-bound Fe(II) species were the reductants (Figure 10b). In both systems, the range of relative reactivities and the sequence of N A C reduction were similar. Also, similar competition effects among NACs present in mixtures occurred in the two systems which indicates that the reactive sites present in the aquifer had a similar affinity spectrum for NACs than reactive, surface-bound Fe(II) species. If D O M was the predominant reductant in the leachate plume, the N A C mixture was expected to be reduced without competition effects. Also, dinitrobenzenes were expected to react 2-3 orders of magnitude faster than mono-NACs, and 4-CH3-NB should have been reduced at the lowest rate (Figure 10b). The results suggest that ferrous iron adsorbed to Fe(III)(hydr)oxides was the dominating reductant of the NACs throughout the plume. Formation and Regeneration of Reactive Fe(II) Surface Sites in Natural Systems. As has been addressed in the discussion of Figure 2b, reactive Fe(II) surface sites at Fe(III)-containing minerals are formed and can be regenerated by adsorption of dissolved Fe(II) from solution. However, such sites can also be formed by microbial activity. Heijman et al. (4) and Heijman (46) investigated the formation and reactivity of Fe(II) surface sites in microbially active laboratory columns containing sandy aquifer material and a model sediment consisting of pure quartz sand coated with amorphous FeOOH, respectively. The columns were operated under ferrogenic conditions and reduced NACs quantitatively to the corresponding anilines. It was shown that reduction of NACs occurred exclusively by abiotic reaction with surfacebound iron(II) species as discussed above. The formation of reactive iron(II) species was coupled to oxidation of organic material by dissimilatory iron reducing bacteria and mediated the transfer electrons from the organic substrate of the bacteria to the pollutants (Figure 11). In contrast to sterile Fe(II)/Fe(III) batch systems, under the prevailing conditions, the NACs studied were all reduced at the same rates. Within the columns virtually no ferrous iron appeared in aqueous solution. Under such conditions, the regeneration of the reactive sites (i.e., the surface-bound iron(II) species) was not possible by re-adsorption of Fe(II) from a reservoir in solution and thus was entirely controlled by the activity of the iron reducing bacteria. Factors that stimulated the activity of iron-reducing microorganisms (e.g., addition of electron donors such as acetate) increased the rates of pollutant reduction. The conditions prevailing in the these column experiments may be regarded as an extreme case since reducing conditions in natural porous media often involve the presence of significant amounts of dissolved ferrous iron due to chemical and micro bial reduction of iron(hydr)oxide minerals or dissolution of Fe(II)-containing minerals. Thus, in natural anoxic media, re-adsorption of Fe(II) from solution will compete with direct microbial regeneration of reactive Fe(II) surface sites. Throughout the studied part of the Grindsted aquifer, the concentrations of aqueous Fe(II) were very high (about 1.5 mM), providing a large reservoir of Fe(II) bulk reductants. Under these conditions, the (re)generation of reactive Fe(II) species at the minerals proceeded predominately by adsorption of aqueous Fe(II) to mineral



354

b) Χ II χ ω c φ

4 " ρ

- 3-Ν0

Ν C 1C Φ

2 -

ft

1 -

c ο φ >

2-Ν0

2

3 -

ο

-

-1 -

2

• Η • 4-CH

2

3-N0

2

i

3

- 2-CH

-|-2-N0

Η 2-CH 4-CH3 3

3

Φ

-2reduced DOM

suspensions >Fe(lll) / Fe(ll)

Grinsted aquifer

Figure 10. a) Schematic representation of the field injection experiment of Riigge et al. (43) conducted in the anaerobic part of a landfill leachate plume. The reactivity patterns of nitroaromatic tracers injected in the aquifer were used to characterize the active reductants present, b) Comparison of the reactivity of the tracers within the aquifer and in model systems exhibiting only N O M or surface-bound Fe(II) as reductants.

"OX

"^v

"surface bound' Fe(ll) adsorption of Fe(ll) f r o m solution

abiotic reaction

red

• iron reducing bacteria

Fe(lll) "surface bound'

Figure 11. General scheme for reduction of pollutants by reactive Fe(II) surface sites and processes that regenerate such reductants.



355 surfaces as reflected by a similar reactivity of untreated and microbially inhibited samples. Nielsen et al. (47) studied the transformation of various xenobiotic organic compounds including NACs and CCI4 in the anaerobic plume the Vejen landfill, Denmark. In contrast to the Grindsted aquifer, the Vejen aquifer exhibited considerable variability with respect to the concentrations of aqueous ferrous iron throughout. In the presence of high concentrations of dissolved Fe(II) their results were similar to those obtained in the Grindsted aquifer. However, further out in the plume, i.e., at lower concentrations of dissolved Fe(II), inhibition of microbial activity by adding biocides to in situ microcosms significantly slowed down the reduction rate constants. Their conclusion was, that biological processes or chemical processes coupled to biological activity were important in the reduction of nitrobenzene under these conditions, while chemical processes were responsible for the reduction in the more reduced part of the plume where Fe(II) was more abundant. This is in good accordance with our discussion on the interplay of microbial iron reduction and adsorption of aqueous Fe(II) in regenerating reactive Fe(II) surface sites at Fe(III)minerals. Conclusions Ferrous iron associated with iron(III)(hydr)oxide surfaces is a versatile and powerful reductant for many reducible organic pollutants of environmental concern. Evidence is presented for the importance of such Fe(II)/Fe(III) species in the overall transformation rates of reducible pollutants in anoxic aquifers. In contrast to the reactivity of minerals containing structural Fe(II), which is often very low once the reactive surface sites are exhausted, the high reactivity of surface-bound Fe(II) can be maintained over long periods of time since such species may continuously be regenerated by uptake of Fe(II) from solution or by the activity of iron reducing microorganisms. As dissolved Fe(II) and iron(III)(hydr)oxides are almost ubiquitously present in anoxic environments, abiotic reduction of pollutants by such species generally may be significant in reducing subsurface systems. The type and the reactivity of Fe(II) species formed at Fe(III)(hydr)oxide surfaces are subject to various environmental factors. The most important variables include solution pH, surface charge and morphologic properties of the supporting Fe(III) mineral, and degree of surface saturation with respect to Fe(II). Generally, reaction conditions favouring surface precipitation, i.e., high surface saturation of oxides with Fe(II) and high pH values, speed up kinetics of reduction reactions and extent the range of reducible substrates. Various macroscopic evidence suggests that surface clusters of Fe(II) or mixed valence iron precipitates such as green rusts may be responsible for the high reactivity of these systems. This hypothesis currently is further investigated by surface sensitive spectroscopic techniques. Groundwater remediation using metal iron as a bulk reductant for pollutants has become an emerging technique in recent years. However, the processes and the nature of the actual reductants that lead to the observed high reduction rates and low substrate specificity of metal iron systems are not yet fully resolved. In such systems, reaction conditions prevail that favor the surface precipitation of Fe(II) species (i.e., high pH values, high surface saturation of oxides with Fe(II)). In fact, spectroscopic studies on corrosion the of steel recently showed that metastabile mixed valence iron coatings are formed as precursors to stabile Fe(III)oxide coatings of the metal (4850). Thus, Fe(II) present at Fe(III)(hydr)oxides may play a pivotal role for the abiotic reduction of pollutants in natural reducing matrices as well as in technical groundwater remediation systems.


356 Acknowledgments Financial support was provided by the Swiss National Science Foundation (SNF grant No. 8220-46517) and by the German Research Council (DFG grant No. Pe 581/1-1).


Literature Cited (1) Pecher, K.; Haderlein, S.B.; Schwarzenbach, R.P. In: 213th ACS National Meeting, Symposium on Redox Reactions in Natural and Engineered Aqueous Systems; ACS, Division of Environmental Chemistry, Ed.; Preprints of Papers 37; 1997, 185-187. (2) Sivavec, T.M.; Horney, D.P. In: 213th ACS National Meeting, Symposium on Redox Reactions in Natural and Engineered Aqueous Systems; ACS, Division of Environmental Chemistry, Ed.; Preprints of Papers 37; 1997, 115-117. (3) Rügge, K.; Bjerg, P.L.; Christensen, T.H. Water Resources Research 1997, submited. (4) Heijman, C.G.; Grieder, E.; Holliger, C.; Schwarzenbach, R.P. Environ. Sci. Technol. 1995, 29, 775-783. (5) Heijman, C.G.; Holliger, C.; Glaus, M.A.; Schwarzenbach, R.P.; Zeyer, J. Appl. Environ. Microbiol. 1993, 59, 4350-4353. (6) Klausen, J. Ph.D. Thesis, ΕΤΗ Zurich, 1995. (7) Sørensen, J.; Thorling, L. Geochchim. Cosmochim. Acta 1991, 55, 1289-1294. (8) Hansen, H.C.B.; Borgaard, O.K.; Sørensen, J. Geochim. Cosmochim. Acta 1994, 58, 2599-2608. (9) Ottley, C.J.; Davison, W.; Edmunds, W.M. Geochim. Cosmochim. Acta 1997, 67, 1819-1828. (10) Buerge, I.J.; Hug, S.J. Environ. Sci. Technol. 1997, 31, 1426-1432. (11) Sedlak, D.L.; Chan, P.G. Geochim. Cosmochim. Acta 1997, 61, 2185-2192. (12) Myreni, S.C.B. et al. Science 1997, 278, 1106ff. (13) Cui, D.; Eriksen, T.E. Environ. Sci. Technol. 1996, 30, 2259-2262. (14) Schwarzenbach, R.P.; Gschwend, P.M.; Imboden, D.M. Environmental Organic Chemistry; John Wiley & Sons: New York, 1993. (15) Stumm, W. Chemistry of the Solid-Water Interface; Wiley: New York, 1992. (16) Kriegmann-King, M.R.; Reinhard, M. Environ. Sci. Technol. 1992, 2198-2206. (17) Klausen, J.; Tröber, S.P.; Haderlein, S.B.; Schwarzenbach, R.P. Environ. Sci. Technol. 1995, 29, 2396-2404. (18) Pecher, K.; Haderlein, S.B.; Schwarzenbach, R.P. Environ. Sci. Technol. 1998, in preparation. (19) Lanz, K.; Schwarzenbach, R.P. 1988. unpublished results. EAWAG (CH)., (20) Yu, S.Y.; Bailey, G.W. J. Environ. Qual. 1992, 21, 86 - 94. (21) Kriegmann-King, M.R.; Reinhard, M. Environ. Sci. Technol. 1994, 28, 692-700. (22) Assaf-Anid, N.; Lin, K.-Y.; Mahony, J. In: 213th ACS National Meeting, Symposium on Redox Reactions in Natural and Engineered Aqueous Systems; ACS, Division of Environmental Chemistry, Ed.; Preprints of Papers 37; 1997, 194-195. (23) Coughlin, B.R.; Stone, A.T. Environ. Sci. Technol. 1995, 29, 2445-2455. (24) Pecher, K.; Schwarzenbach, R.P. Geochim. Cosmochim. Acta 1998, in preparation. (25) Hansen, H.C.B.; Koch, C.B.; Nancke-Kroh, H.; Borggaard, O.K.; Sørensen, J. Environ. Sci. Technol. 1995, 30, 2053-2056. (26) Dzombak, D.A.; Morel, F.M.M. Surface complexation modeling. Hydrous ferric oxide.; John Wiley & Sons: New York, 1990.



357 (27) Kinniburgh, D.G.; Jackson, M.L.; Anderson, Α.Α.; Rubin, A.J. Cation adsorption by hydrous metal oxides and clays/Adsorption of inorganics at solid -liquid interfaces.; Ann Arbor Science: Ann Arbor (MI), 1981. (28) Sposito, G. In Geochemical Processes at Mineral Surfaces; ACS Symposium Series, 323, Davis, J. Α.; K. F. Hayes, Ed.; American Chemical Society: Washington, 1986. 217-228. (29) Hachiya, K.; Sasaki, M.; Ikeda, T.; Mikami, N.; Yasunaga, T. J. Phys. Chem. 1984, 88, 27-31. (30) Cornell, R.M.; Schwertmann, U. The Iron Oxides; VCH: Weinheim, 1996. (31) De Boer, J.; Lippens, B.; Linsen, B.; Broekhiff, J.; van den Heuvel, Α.; Osinga, T. J. Colloid Interf. Sci. 1966, 21, 405-414. (32) Jolivet, J.-P.; Tronc, Ε. J. Coll. Interf. Sci. 1988, 125, 688-701. (33) Tronc, Ε.; Jolivet, J.-P. Ads. Sci. Technol. 1984, 1, (34) Tronc, Ε.; Jolivet, J.-P.; Lefebvre, J.; Massart, R. J. Chem. Soc. Faraday Trans. 1984, 80, 2619-2629. (35) Tronc, Ε.; Belleville, P.; Jolivet, J.-P.; Livage, J. Langmuir 1992, 8, 313-319. (36) Tamaura, Y.; Ito, K.; Katsura, T. J. Chem. Soc. Dalton Trans. 1983, 189-194. (37) Tronc, Ε.; Jolivet, J.P.; Belleville, P.; Livage, J. Hyperfine Interact. 1989, 46, 637-643. (38) Ford, R.G.; Bertsch, P.M.; Farley, K.J. Environ. Sci. Technol. 1997, 31, 20282033. (39) Farley, K.J.; Dzombak, D.A.; Morel, F.M.M. J. Coll. Interf. Sci. 1985, 106, 226242. (40) Baes Jr., C.F.; Mesmer, R.E. The Hydrolysis of Cations.; John Wiley & Sons: New York, 1976. (41) Zhang, Y.; Charlet, L.; Schindler, P.W. Colloids Surf. 1992, 63, 259-268. (42) Tratnyek, P.G.; Smolen, J.M.; Weber, E.J. In: 213th ACS National Meeting, Symposium on Redox Reactions in Natural and Engineered Aqueous Systems; ACS, Division of Environmental Chemistry, Ed.; Preprints of Papers 37; 1997, 119-121. (43) Rügge, K.; Hofstetter, T.; Haderlein, S.B., et al. Environ. Sci. Technol. 1998, 32(1), in press. (44) Haderlein, S.B.; Schwarzenbach, R.P. In Biodegradation of Nitroaromatic Compounds; Spain, J., Ed.; Plenum: New York, 1995. Chapter 12, 199-225. (45) Colon, D.; Weber, E.J.; Anderson, J.L. In: 213th ACS National Meeting, Symposium on Redox Reactions in Natural and Engineered Aqueous Systems; ACS, Division of Environmental Chemistry, Ed.; Preprints of Papers 37; 1997, 193-194. (46) Heijman, C.G. Ph.D. Thesis, ΕΤΗ Zurich, 1995. (47) Nielsen, P.H.; Bjarnadottir, H.; Winter, P.L. J. Contam. Hydrol. 1995, 20, 51-66. (48) Abdelmoula, M.; Refait, P.; Drissi, S.H.; Mihe, J.P.; Génin, J.-M.R. Corrosion Science 1996, 38, 626-633. (49) Novakova, A.A.; Gendler, T.S.; Manyurova, N.D.; Turishcheva, A. Corrosion Science 1997, 39, 1585-1594. (50) Réfait, P.H.; Drissi, S.H.; Pytkiewicz, J.; Génin, J.M.R. Corrosion Science 1997, 39, 1699-1709.


Pollutant Reduction in Heterogeneous Fe(II)-Fe(III) - American

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