ISSN 10642293, Eurasian Soil Science, 2013, Vol. 46, No. 12, pp. 1139–1149. © Pleiades Publishing, Ltd., 2013.
Original Russian Text © Yu.N. Vodyanitskii, 2013, published in Pochvovedenie, 2013, No. 12, pp. 1437–1448.
SOIL
CHEMISTRY
Determination of the Oxidation States of Metals and Metalloids:
An Analytical Review
Yu. N. Vodyanitskii
Faculty of Soil Science, Moscow State University, Moscow, 119992 Russia
Email: [email protected]
Received March 10, 2011
Abstract—The hazard of many heavy metals/metalloids in the soil depends on their oxidation state. The
problem of determining the oxidation state has been solved due to the use of synchrotron radiation methods
with the analysis of the Xray absorption nearedge structure (XANES). The determination of the oxidation
state is of special importance for some hazardous heavy elements (arsenic, antimony, selenium, chromium,
uranium, and vanadium). The mobility and hazard of each of these elements depend on its oxidation state.
The mobilities are higher at lower oxidation states of As, Cr, V, and Se and at higher oxidation states of Sb and U.
The determination of the oxidation state of arsenic has allowed revealing its fixation features in the rhizo
sphere of hydrophytes. The known oxidation states of chromium and uranium are used for the retention of
these elements on geochemical barriers. Different oxidation states have been established for vanadium dis
placing iron in goethite. The determination of the oxidation state of manganese in the rhizosphere and the
photosynthetic apparatus of plants is of special importance for agricultural chemists.
Keywords: arsenic, antimony, selenium, chromium, uranium, vanadium, mobility of elements, geochemical
barriers
DOI: 10.1134/S1064229313120077
INTRODUCTION
The hazard of many heavy metals/metalloids in the
soil depends on their oxidation state. However, the
determination of this parameter faces some problems.
Suffice to say that the prevalent methods for the chem
ical fractionation of heavy metals/metalloids provide
no such information [14, 84]. The problem was solved
with the development of synchrotron radiation meth
ods based on the use of accelerators. These methods
are utilized in different disciplines, including soil sci
ence [36, 46, 77, 83]. The thirdgeneration synchro
trons are most efficient; in them, elementary particles
are accelerated in a magnetic field and form powerful
Xray radiation of very high brightness and purity.
Since the 1980s, abundant information has emerged
on the oxidation states of different elements in the soil
from Xray absorption analysis [42, 54, 55].
The synchrotron radiation methods currently allow
studying the solidphase composition in a microvol
ume, the oxidation states of elements with variable
valences, the distribution of heavy metals and metal
loids in undisturbed soil samples, and the revelation of
their linking to carrier phases. Xray microfluores
cence spectroscopy (µXRF), Xray microdiffraction
(µXRD), Xray absorption nearedge spectroscopy
(XANES), and extended Xray absorption fine struc
ture analysis (EXAFS) are used for this purpose. These
structural tools are sufficiently specialized: they are
sensitive to lowordered particles and have sufficient
identification limits for heavy metals if their contents
exceed 100 mg/kg. Xray synchrotron analysis is espe
cially efficient in the study of lowordered compounds
and elements with low clarke values, when the Xray
and electron diffraction methods are useless. Another
important advantage is the study of samples with natu
ral water content [19]. This research area is actively
being developed. In 2010, Elsevier published Synchro
tronBased Techniques in Soils and Sediments (Series:
Advances in Soil Science), where the advances of
chemistry and mineralogy in the studies of soils and
bottom sediments by synchrotron radiation methods
are summarized [36, 77].
The Xray absorption analysis allows determining
the oxidation states of many chemical elements with
variable valences [74]. The determination of the oxi
dation state is of special importance for the hazardous
heavy elements whose contents increase with soil con
tamination. Their retention on natural and artificial
geochemical barriers depends on their oxidation state.
The aim of this review is to generalize the data on
the oxidation states of arsenic, antimony, selenium,
uranium, vanadium, chromium, and manganese in
soils derived using the XANES technique.
1139
Absorbance
1142
VODYANITSKII
OR1
SR1
HF1
U(VI)
U(IV)
17140 17160 17180 17200 17220 17240
Energy, eV
Fig. 3. XANES absorption spectra of uranium in UO2 and
UO2(NO3)2 ⋅ H2O standards and the spectra of uranium in
contaminated soils from Oak Ridge in Tennessee (OR), near
the Hanford nuclear power plant in Washington (HF1),
and the Savannah River in South Caroline (SR). (Modi
fied from [25]).
only Sb(0) or Sb(V) bound to iron hydroxides (pre
dominantly goethite) [71].
The study of soils contaminated with Sb and Sb2O3
at the emission from smelters also showed the presence
of oxidative processes. The technogenic forms of
reduced antimony are converted into the most oxi
dized and mobile form of Sb(V) [80]. This complicates
the remediation of soils contaminated with antimony.
Uranium and its oxidation states. Uranium occupies
the 47th position among the elements in the Earth’s
crust (2.3 mg/kg) [8]. The content of uranium varies
among the soils of the world from 0.7 to 10.7 mg/kg
[12]. The content of uranium in the soils of uranium
provinces is appreciably higher than in the soils
depleted of the element [13]. It is 5.8–10.7 mg of U/kg in
the IssykKul Depression and only 0.5–0.8 mg U/kg in
the Kursk Depression. The study of the uranium con
tent in soils of the United States revealed that the dif
ferences are related to the particlesize distribution
rather than to the soil type: the content of uranium
decreases to 0.3 mg/kg in light soils and increases to
10.7 mg/kg in heavy soils [3]. Very high concentrations
(up to 100 mg U/kg) are related to the technogenic
contamination of soils. The mean content of uranium
is 2.6, 1.2, 0.79, and 11 mg U/kg in the soils of Great
Britain, Canada, Poland, and India, respectively. In
the soils of countries of the temperate zone, the mean
uranium content is 2 ± 1.5 mg/kg [3].
Uranium has a variable valence; the main degrees
of oxidation are +4 and +6. This determines the sen
sitivity of uranium to redox conditions. Under oxida
tive conditions, uranyl UO 22+ forms highly mobile
compounds [17]. Under reducing conditions, U4+ is
oxidized to the stable oxide (uraninite) UO2. This fact
determines the different behavior of uranium in soils.
The behavior of uranium radically differs depend
ing on the oxidation state; therefore, the determina
tion of its valence in an undisturbed sample is an
important task. The XANES spectra show signals of
U(VI) and U(IV) obtained at the shift of the energy
(Е) relative to Е0 = 17166 eV, the standard energy of
the uranium LIII line. Due to the high sensitivity of
synchrotron equipment, the difference between the
positions of the U(VI) and U(IV) spectra is significant
and easily identified. The XANES spectra of U(IV) in
UO2 and U(VI) in UO2(NO3)2 ⋅ H2O are shown in Fig. 3.
They are compared to the spectra of uranium in con
taminated soils of the United States [25]. It can be
seen that the positions of the U(VI) and U(IV) maxi
mums on the energy scale are clearly different; the
absorption energy of U(VI) is higher than that of
U(IV). In soil samples, uranium occurs in the danger
ous oxidized state.
A pathway for the fixation of uranium in watersat
urated soil layers is the chemical reduction of U(VI) by
Fe(II) at the participation of dissimilatory metal
reducing bacteria [34, 38]. These bacteria reduce iron
(hydr)oxides and saturate the water solution with
Fe(II), which reduces U(VI) and favors its precipita
tion. This complex process cannot be studied without
the XANES technique. It was found that the reaction
of Fe(II) with U(VI) in water is catalyzed on the sur
face of solidphase particles [34, 43]. Soil particles are
covered with organic compounds; therefore, searchers
are focused on studying the role of organic ligands in
the reduction of U(VI) due to the oxidation of Fe(II).
Synthetic colloidal microspheres with surface car
boxyl functional groups were used in model experi
ments [27]. The use of XANES revealed the following
peculiarities. In the U + Fe + carboxyl system, the
reaction strongly depends on the pH. At pH 7.5, car
boxyls inhibit the capacity of Fe(II) to reduce U(VI).
However, the reaction is accelerated at pH 8.4. Thus,
XANES is an efficient tool for studying the oxidation
state of uranium and the reactions with its participa
tion.
Vanadium and its oxidation states. Vanadium occu
pies the 19th position among the elements of the
Earth’s crust (136 mg/kg) [8]. In soils, it is associated
with iron and titanium oxides, which are usually
inherited from the parent rock. Its content depends on
the composition of the parent rocks; shales and clays
contain much vanadium. Contaminated soils contain
a high portion of technogenic vanadium. In the ash of
highmoor peat contaminated with oil in Western
Siberia, the content of vanadium reaches 2000 mg/kg,
which exceeds its background values by 20 times [7].
Vanadium complexes are usually of anionic nature,
but they are electroneutral and cationic under acid
conditions. Vanadium has a variable degree of oxida
tion, from +2 to +5 with the main degree of oxidation
being +5. In automorphic soils under oxidative condi
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DETERMINATION OF THE OXIDATION STATES OF METALS AND METALLOIDS
5. Yu. N. Vodyanitskii, A. A. Vasil’ev, S. N. Lesovaya,
E. F. Sataev, A. V. Sivtsov, “Formation of Manganese
Oxides in Soils,” Eur. Soil Sci. 37 (6), 572–584 (2004).
6. Yu. N. Vodyanitskii, A. I. Gorshkov, and A. V. Sivtsov,
“Peculiarities of Manganese Oxides in Soils of the Rus
sian Plain,” Eur. Soil Sci. 35 (10), 1037–1045 (2002).
7. Yu. N. Vodyanitskii, A. T. Savichev, S. Ya. Trofimov, and
E. A. Shishkonakova, “Accumulation of Heavy Metals
in OilContaminated Peat Soils,” Eur. Soil Sci. 45 (10),
977–982 (2012).
8. N. Greenwood and A. Earnshaw, Chemistry of the Ele
ments (2nd ed.) (Elsevier, 1997).
9. V. V. Dobrovol’skii, Geography of Microelements. Global
Scattering (Mysl’, Moscow, 1983) [in Russian].
10. V. V. Dobrovol’skii, Fundamentals of Biogeochemistry
(ACADEMIA, Moscow, 2003) [in Russian].
11. V. V. Ivanov, Environmental Geochemistry of the Ele
ments (NedraEkologiya, Moscow, 1994) [in Russian].
12. A. KabataPendias and H. Pendias, Trace Elements in
Soils and Plants, (CRS Press, Boca Raton, USA, 1985).
13. V. V. Koval’skii, Geochemical Ecology (Nauka, Moscow,
1974) [in Russian].
14. D. V. Ladonin, “Heavy Metal Compounds in Soils:
Problems and Methods of Study,” Eur. Soil Sci. 35 (6),
605–613 (2002).
15. The Encyclopedia of Mineralogy, Ed. by K. Frye
(Hutchinson Ross Publ. Company, Pa, 1981).
16. D. S. Orlov, L. K. Sadovnikova, and N. I. Sukhanova,
Soil Chemistry (Vysshaya shkola, Moscow, 2005) [in
Russian].
17. A. I. Perel’man and N. S. Kasimov, Geochemistry of
Landscape (Astreya, Moscow, 1999) [in Russian].
18. F. V. Chukhrov, A. I. Gorshkov, and V. A. Drits, Superg
enous Manganese Oxides (Nauka, Moscow, 1989) [in
Russian].
19. J. R. Bargar, B. M. Tebo, and J. E. Villinski, “In situ
characterization of Mn(II) oxidation by spores of the
marine Bacillus sp. strain Sg1,” Geochim. Cosmo
chim. Acta 64, 2775–2778 (2000).
20. S. Beauchemin and Y. T. J. Kwong, “Impact of redox
conditions on arsenic mobilization from tailings in a
wetland with neutral drainage,” Environ. Sci. Technol.
40, 6297–6303 (2006).
21. N. Belzile, Y. W. Chen, and Z. J. Wang, “Oxidation of
antimony(III) by amorphous and manganese oxyhy
droxides,” Chem. Geol. 174, 379–387 (2001).
22. D. W. Blowes, C. J. Ptacek, S. G. Benner, C. W. T. McRae,
T. A. Bennet, R. W. J. Puls, “Treatment of inorganic
contaminants using permeable reactive barriers,” Con
tam. Hydrol 45, 123–137 (2000).
23. D. W. Blowes, C. J. Ptacek, and J. L. Jambor, “In situ
remediation of Cr(VI)contaminated groundwater
using permeable reactive walls: laboratory studies,”
Environ. Sci. Technol. 31, 3348–3357 (1997).
24. N. K. Blute, D. J. Branander, H. F. Hemond, S. R. Sut
ton, M. G. Nrwville, M. L. Rivers, “Arsenic sequestra
tion by ferric iron plaque on cattail roots,” Environ.
Sci. Technol. 38, 6074–6077 (2004).
25. B. C. Bostic, S. Fendorf, M. O. Barnett, P. M. Jardine,
S. C. Brooks, “Uranyl surface complexes formed on
EURASIAN SOIL SCIENCE
Vol. 46
No. 12
2013
26.
27.
28.
29.
30.
31.
32.
33.
34.
35.
36.
37.
38.
39.
40.
41.
1147
subsurface media from DOE facilities,” Soil Sci. Soc
Am. J. 66, 99–108 (2002).
H. J. M. Bowen, Environmental Chemistry of the Ele
ments (Academic, London, 1979).
M. I. Boyanow, E. J. O’Loughlin,, E. E. Roden, J. B. Fein,
and K. M. Kemner, “Adsorption of Fe(II) and U(VI) to
carboxylfunctionalized microspheres: the influrnce of
speciation on uranyl reduction studied by titration and
XAFS,” Geochim. Cosmochim. Acta 71, 1898–1912
(2007).
L. Charlet and A. Manceau, “In situ characterization of
heavy metal surface reactions: the chromium case,”
Int. J. Environ. Anal. Chem. 46, 97–108 (1992).
W. R. Cullen and K. J. Reomer, “Arsenic speciation in
the environment,” Chem. Rev. 89, 713–764 (1989).
S. H. R. Davies and J. J. Morgan, “Manganese (II) oxi
dation kinetics on metaloxide surfaces,” J. Coll.
Interf. Sci. 129, 63–77 (1989).
K. S. Dhillon and S. K. Dhillon, “Distribution and
management of seleniferous soils,” Adv. Agron. 79,
119–184 (2003).
M. Filliela, N. Belzile, and Y. W. Chen, “Antimony in
the environment: a review focused on natural waters. i.
occurrence,” Earth Sci. Rev. 57, 125–176 (2002).
M. Filliela, N. Belzile, and Y. W. Chen, “Antimony in
the environment: a review focused on natural waters. ii.
relevant solution chemistry,” Earth Sci. Rev. 59, 265–
285 (2002).
G. K. Fredrickson, J. M. Zachara, D. W. Kennedy,
M. C. Duff, Y. A. Gorby, S.M. Li, K. N. Krupka,
“Reduction of U(VI) in goethite (αFeOOH) suspen
sions by dissimilatory metalreusing bacterium,”
Geochim. Cosmochim. Acta 64, 3085–3098 (2000).
G. F. Friedl, B. Wehrli, and A. Manceau, “Solid phases
in the cycling of manganese in eutrophic lakes: new
insights from EXAFS spectroscopy,” Geochim. Cos
mochim. Acta 61, 275–290 (1997).
M. GinderVogel and D. L. Sparks, “The impacts of
Xray absorption spectroscopy on understanding soil
processes and reaction mechanisms,” in Synchrotron
Based Techniques in Soils and Sediments, Ed. by
B. Singh and M. Grafe, Developments in Soil Sci. 34,
1–26 (2010).
S. Goldberg and R. A. Glaubig, “Anion sorption on a
calcareous, montmorillonitic soil—arsenic,” Soil Sci.
Soc. Am. J. 52, 1297–1300 (1988).
Y. A. Gorby and D. R. Lovley, “Enzymatic uranium
precipitation,” Environ. Sci. Technol. 26, 205–207
(1992).
C. M. Hansel, S. G. Benner, J. Neiss, A. Dohnalkova,
R. K. Kukkadapu, S. Fendorf, “Secondary mineraliza
tion pathways induced by dissimilatory iron reduction
of ferrihydrite under advective flow,” Geochim. Cos
mochim. Acta 67, 2977–2992 (2003).
C. M. Hansel, S. Fendorf, M. J. LaForce, and S. Sut
ton, “Spatial and temporal associations of As and Fe
species on aquatic plant roots,” Environ. Sci. Technol.
36, 1988–1994 (2002).
B. P. Jackson and W. P. Miller, “Effectiveness of phos
phate and hydroxide for desorption of arsenic and sele
nium species from iron oxides,” Soil Sci. Soc. Am. J.
64, 1616–1622 (2000).
Article
pubs.acs.org/est
Abiotic Reductive Immobilization of U(VI) by Biogenic Mackinawite
Harish Veeramani,*,† Andreas C. Scheinost,‡ Niven Monsegue,§ Nikolla P. Qafoku,∥ Ravi Kukkadapu,⊥
Matt Newville,¶ Antonio Lanzirotti,¶ Amy Pruden,# Mitsuhiro Murayama,§,#
and Michael F. Hochella, Jr.†,∥
†
Department of Geosciences, Virginia Tech, Blacksburg, Virginia, United States
Institute of Radiochemistry, FZD and Rossendorf Beamline, European Synchrotron Radiation Lab, Grenoble, France
§
Department of Materials Science and Engineering, Virginia Tech, Blacksburg Virginia, United States
∥
Geosciences Group, Pacific Northwest National Laboratory, Richland, Washington, United States
⊥
Environmental Molecular Sciences Laboratory, Pacific Northwest National Laboratory, Richland, Washington, United States
¶
Center for Advanced Radiation Sources, Advanced Photon Source (APS), Argonne, Illinois, United States
#
Institute of Critical Technology and Applied Sciences, Virginia Tech, Blacksburg, Virginia, United States
‡
S Supporting Information
*
ABSTRACT: During subsurface bioremediation of uranium-contaminated sites,
indigenous metal and sulfate-reducing bacteria may utilize a variety of electron
acceptors, including ferric iron and sulfate that could lead to the formation of various
biogenic minerals in situ. Sulfides, as well as structural and adsorbed Fe(II)
associated with biogenic Fe(II)-sulfide phases, can potentially catalyze abiotic U(VI)
reduction via direct electron transfer processes. In the present work, the propensity
of biogenic mackinawite (Fe1+xS, x = 0 to 0.11) to reduce U(VI) abiotically was
investigated. The biogenic mackinawite produced by Shewanella putrefaciens strain
CN32 was characterized by employing a suite of analytical techniques including
TEM, SEM, XAS, and Mö ssbauer analyses. Nanoscale and bulk analyses
(microscopic and spectroscopic techniques, respectively) of biogenic mackinawite
after exposure to U(VI) indicate the formation of nanoparticulate UO2. This study
suggests the relevance of sulfide-bearing biogenic minerals in mediating abiotic
U(VI) reduction, an alternative pathway in addition to direct enzymatic U(VI)
reduction.
iron [Fe2+], sorbed Fe(II) species,12 and the formation of
secondary mineralization products in situ including reactive
Fe(II)-bearing biogenic minerals.13−23 Biogenic Fe(II)-bearing
minerals can provide a reservoir of reducing capacity where
reduction of U(VI) may occur due to abiotic interactions17,22
and potentially compete with direct enzymatic reduction24 of
U(VI). Abiotic U(VI) reduction is a thermodynamically
favorable but often kinetically limited process and has been
reported to be mediated by adsorbed Fe(II) species,23−31
structural Fe(II) present in Fe(II)-bearing17,22,32−34 and
ferrous-sulfide bearing minerals such as pyrite (FeS2),35−37
mackinawite (Fe1+xS),38−40 and amorphous iron-sulfide.41
Mackinawite is an environmentally relevant biogenic mineral42
and is the initial ferrous sulfide solid phase that forms under
sulfate reducing conditions, both in column42−44 and field-scale
studies.45 It plays a critical role in serving as a precursor to the
formation of most other stable iron sulfide phases46,47 among
1. INTRODUCTION
Microbially mediated reduction of aqueous hexavalent uranium
U(VI) to promote the formation of the sparingly soluble
mineral uraninite [UO2] represents a promising strategy for the
in situ immobilization of uranium in subsurface sediments and
groundwater at contaminated sites. In compositionally
heterogeneous subsurface environments such as sediments,
indigenous microbes including dissimilatory metal reducing
(DMRB) and dissimilatory sulfate reducing (DSRB) bacteria
can encounter multiple electron acceptors including Fe(III),
Mn(IV), sulfate, and nitrate. Although the utilization of
terminal electron acceptors is often assumed to be sequential
from the highest to the lowest energy yield,1 iron and sulfate
reduction have been observed to occur either concurrently or
sequentially in several field studies.2−5 While preferential or
competitive terminal electron accepting processes reported in
most laboratory studies do not necessarily represent natural
events in the subsurface, their potential occurrences cannot be
excluded during biostimulation trials for uranium remediation.6
Due to the abundance of Fe(III) in the subsurface,7−9 the
biostimulation of DMRB will likely lead to biological Fe(III)
reduction3,10,11 resulting in the formation of aqueous ferrous
© 2013 American Chemical Society
Received:
Revised:
Accepted:
Published:
2361
October 3, 2012
January 27, 2013
February 4, 2013
February 4, 2013
dx.doi.org/10.1021/es304025x | Environ. Sci. Technol. 2013, 47, 2361−2369
Environmental Science & Technology
Article
(lepidocrocite): Comparison of several Shewanella species. Geomicrobiol. J. 2007, 211−230.
(17) O’Loughlin, E. J.; Kelly, S. D.; Kemner, K. M. XAFS
Investigation of the Interactions of UVI with Secondary Mineralization
Products from the Bioreduction of FeIII Oxides. Environ. Sci. Technol.
2010, 44 (5), 1656−1661.
(18) Komlos, J.; Moon, H.; Jaffe, P. Effect of Sulfate on the
Simultaneous Bioreduction of Iron and Uranium. J. Environ. Qual.
2008, 2058−2062.
(19) Roden, E. E.; Zachara, J. M. Microbial Reduction of Crystalline
Iron(III) Oxides: Influence of Oxide Surface Area and Potential for
Cell Growth. Environ. Sci. Technol. 1996, 30 (5), 1618−1628.
(20) Zachara, J. M.; Fredrickson, J. K.; Li, S. M.; Kennedy, D. W.;
Smith, S. C.; Gassman, P. L. Bacterial reduction of crystalline Fe3+
oxides in single phase suspensions and subsurface materials. Am.
Mineral. 1998, 83 (11-12, Part 2), 1426−1443.
(21) Roh, Y.; Zhang, C.-L.; Vali, H.; Lauf, R. J.; Zhou, J.; Phelps, T. J.
Biogeochemical and environmental factors in Fe biomineralization:
Magnetite and siderite formation. Clays Clay Miner. 2003, 51 (1), 83−
95.
(22) Veeramani, H.; Alessi, D. S.; Suvorova, E. I.; Lezama-Pacheco, J.
S.; Stubbs, J. E.; Sharp, J. O.; Dippon, U.; Kappler, A.; Bargar, J. R.;
Bernier-Latmani, R. Products of abiotic U(VI) reduction by biogenic
magnetite and vivianite. Geochim. Cosmochim. Acta 2011, 75 (9),
2512−2528.
(23) Behrends, T.; Van Cappellen, P. Competition between
enzymatic and abiotic reduction of uranium(VI) under iron reducing
conditions. Chem. Geol. 2005, 220 (3−4), 315−327.
(24) Fredrickson, J.; Zachara, J.; Kennedy, D.; Duff, M.; Gorby, Y.;
Li, S.; Krupka, K. Reduction of U(VI) in goethite (alpha-FeOOH)
suspensions by a dissimilatory metal-reducing bacterium. Geochim.
Cosmochim. Acta 2000, 3085−3098.
(25) Charlet, L.; Liger, E.; Gerasimo, P. Decontamination of TCEand U-Rich Waters by Granular Iron: Role of Sorbed Fe(II). J. Environ.
Eng. 1998, 124 (1), 25−30.
(26) Jeon, B.-H.; Kelly, S. D.; Kemner, K. M.; Barnett, M. O.; Burgos,
W. D.; Dempsey, B. A.; Roden, E. E. Microbial Reduction of U(VI) at
the Solid−Water Interface. Environ. Sci. Technol. 2004, 38 (21), 5649−
5655.
(27) Liger, E.; Charlet, L.; Van Cappellen, P. Surface catalysis of
uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999,
63 (19−20), 2939−2955.
(28) Regenspurg, S.; Schild, D.; Schäfer, T.; Huber, F.; Malmström,
M. E. Removal of uranium(VI) from the aqueous phase by iron(II)
minerals in presence of bicarbonate. Appl. Geochem. 2009, 24 (9),
1617−1625.
(29) Chakraborty, S.; Favre, F.; Banerjee, D.; Scheinost, A. C.;
Mullet, M.; Ehrhardt, J.-J.; Brendle, J.; Vidal, L. c.; Charlet, L. U(VI)
Sorption and Reduction by Fe(II) Sorbed on Montmorillonite.
Environ. Sci. Technol. 2010, 44 (10), 3779−3785.
(30) Jeon, B.-H.; Dempsey, B. A.; Burgos, W. D.; Barnett, M. O.;
Roden, E. E. Chemical Reduction of U(VI) by Fe(II) at the Solid−
Water Interface Using Natural and Synthetic Fe(III) Oxides. Environ.
Sci. Technol. 2005, 39 (15), 5642−5649.
(31) Boyanov, M. I.; O’Loughlin, E. J.; Roden, E. E.; Fein, J. B.;
Kemner, K. M. Adsorption of Fe(II) and U(VI) to carboxylfunctionalized microspheres: The influence of speciation on uranyl
reduction studied by titration and XAFS. Geochim. Cosmochim. Acta
2007, 71 (8), 1898−1912.
(32) O’Loughlin, E.; Kelly, S.; Cook, R.; Csencsits, R.; Kemner, K.
Reduction of Uranium(VI) by mixed iron(II/iron(III) hydroxide
(green rust): Formation of UO2 manoparticies. Environ. Sci. Technol.
2003, 721−727.
(33) Ilton, E. S.; Haiduc, A.; Moses, C. O.; Heald, S. M.; Elbert, D.
C.; Veblen, D. R. Heterogeneous reduction of uranyl by micas: Crystal
chemical and solution controls. Geochim. Cosmochim. Acta 2004, 68
(11), 2417−2435.
(34) Ilton, E. S.; Heald, S. M.; Smith, S. C.; Elbert, D.; Liu, C.
Reduction of Uranyl in the Interlayer Region of Low Iron Micas under
Anoxic and Aerobic Conditions. Environ. Sci. Technol. 2006, 40 (16),
5003−5009.
(35) Wersin, P.; Hochella, M. F., Jr.; Persson, P.; Redden, G.; Leckie,
J. O.; Harris, D. W. Interaction between aqueous uranium (VI) and
sulfide minerals: Spectroscopic evidence for sorption and reduction.
Geochim. Cosmochim. Acta 1994, 58 (13), 2829−2843.
(36) Bruggeman, C.; Maes, N. Uptake of Uranium(VI) by Pyrite
under Boom Clay Conditions: Influence of Dissolved Organic Carbon.
Environ. Sci. Technol. 2010, 44 (11), 4210−4216.
(37) Qafoku, N. P.; Kukkadapu, R. K.; McKinley, J. P.; Arey, B. W.;
Kelly, S. D.; Wang, C.; Resch, C. T.; Long, P. E. Uranium in
Framboidal Pyrite from a Naturally Bioreduced Alluvial Sediment.
Environ. Sci. Technol. 2009, 43 (22), 8528−8534.
(38) Hyun, S. P.; Davis, J. A.; Sun, K.; Hayes, K. F. Uranium(VI)
Reduction by Iron(II) Monosulfide Mackinawite. Environ. Sci. Technol.
2012, 46 (6), 3369−3376.
(39) Moyes, L. N.; Parkman, R. H.; Charnock, J. M.; Vaughan, D. J.;
Livens, F. R.; Hughes, C. R.; Braithwaite, A. Uranium Uptake from
Aqueous Solution by Interaction with Goethite, Lepidocrocite,
Muscovite, and Mackinawite: An X-ray Absorption Spectroscopy
Study. Environ. Sci. Technol. 2000, 34 (6), 1062−1068.
(40) Livens, F. R.; Jones, M. J.; Hynes, A. J.; Charnock, J. M.;
Mosselmans, J. F. W.; Hennig, C.; Steele, H.; Collison, D.; Vaughan,
D. J.; Pattrick, R. A. D.; Reed, W. A.; Moyes, L. N. X-ray absorption
spectroscopy studies of reactions of technetium, uranium and
neptunium with mackinawite. J. Environ. Radioact. 2004, 74 (1−3),
211−219.
(41) Hua, B.; Deng, B. Reductive Immobilization of Uranium(VI) by
Amorphous Iron Sulfide. Environ. Sci. Technol. 2008, 42 (23), 8703−
8708.
(42) Abdelouas, A.; Lutze, W.; Nuttall, H. E. Oxidative dissolution of
uraninite precipitated on Navajo sandstone. J. Contam. Hydrol. 1999,
36 (3−4), 353−375.
(43) Personna, Y. R.; Ntarlagiannis, D.; Slater, L.; Yee, N.; O’Brien,
M.; Hubbard, S. Spectral induced polarization and electrodic potential
monitoring of microbially mediated iron sulfide transformations. J.
Geophys. Res. 2008, 113 (G2), G02020.
(44) Williams, K. H.; Ntarlagiannis, D.; Slater, L. D.; Dohnalkova, A.;
Hubbard, S. S.; Banfield, J. F. Geophysical Imaging of Stimulated
Microbial Biomineralization. Environ. Sci. Technol. 2005, 39 (19),
7592−7600.
(45) Williams, K. H.; Kemna, A.; Wilkins, M. J.; Druhan, J.; Arntzen,
E.; N’Guessan, A. L.; Long, P. E.; Hubbard, S. S.; Banfield, J. F.
Geophysical Monitoring of Coupled Microbial and Geochemical
Processes During Stimulated Subsurface Bioremediation. Environ. Sci.
Technol. 2009, 43 (17), 6717−6723.
(46) Butler, I. B.; Rickard, D. Framboidal pyrite formation via the
oxidation of iron (II) monosulfide by hydrogen sulphide. Geochim.
Cosmochim. Acta 2000, 64 (15), 2665−2672.
(47) Sweeney, R. E.; Kaplan, I. R. Pyrite Framboid Formation;
Laboratory Synthesis and Marine Sediments. Econ. Geol. 1973, 68 (5),
618−634.
(48) Berner, R. A. Sedimentary pyrite formation. Am. J. Sci. 1970, 268
(1), 1−23.
(49) Patterson, R. R.; Fendorf, S.; Fendorf, M. Reduction of
Hexavalent Chromium by Amorphous Iron Sulfide. Environ. Sci.
Technol. 1997, 31 (7), 2039−2044.
(50) Scheinost, A. C.; Charlet, L. Selenite Reduction by Mackinawite,
Magnetite and Siderite: XAS Characterization of Nanosized Redox
Products. Environ. Sci. Technol. 2008, 42 (6), 1984−1989.
(51) Stookey, L. L. FerrozineA new spectrophotometric reagent
for iron. Anal. Chem. 1970, 42 (7), 779−781.
(52) Integrated X-ray powder diffraction software. Rigaku J. 2010, 26,
(1), 23−27.
(53) Peretyazhko, T. S.; Zachara, J. M.; Kukkadapu, R. K.; Heald, S.
M.; Kutnyakov, I. V.; Resch, C. T.; Arey, B. W.; Wang, C. M.; Kovarik,
L.; Phillips, J. L.; Moore, D. A. Pertechnetate (TcO4−) reduction by
reactive ferrous iron forms in naturally anoxic, redox transition zone
2368
dx.doi.org/10.1021/es304025x | Environ. Sci. Technol. 2013, 47, 2361−2369
SCHINDLER & ILTON
CHAPTER 7: URANIUM MINERALOGY AND GEOCHEMISTRY ON THE NANO- TO
MICROMETRE SCALE: REDOX, DISSOLUTION AND PRECIPITATION
PROCESSES AT THE MINERAL-WATER INTERFACE
Michael Schindler,
Department of Earth Sciences, Laurentian University,
Sudbury, Ontario, Canada P3E 2C6
e-mail: [email protected]
and
Eugene S. Ilton,
Pacific Northwest National Laboratory,
Richland, Washington, USA 99352 
e-mail: [email protected]
hundred μg/g or less (Schindler et al. 2013).
Nonetheless, in a small discharge site at the Key
Lake mine site, Saskatchewan, a limited area
contains numerous uranyl minerals formed through
neutralization of mill process solutions (pH range of
2–6) by tailings material containing an excess of
unreacted slaked lime (calcium hydroxide). The
formation of uranyl minerals in this environment is
strongly controlled by dissolution and precipitation
processes at the tailings discharge interface resulting
in complex intergrowths involving U minerals
(Schindler et al. 2013). Here, unraveling interfacial
processes and complex mineralogical relationships
involving U would be beneficial for better
understanding the long term stability of U within
these tailings (Schindler et al. 2013).
The United States nuclear weapons program
has left a legacy of U contamination at various U.S.
Department of Energy sites including those at
Hanford WA, Oak Ridge TN, Rifle CO, and
Savannah River GA (Riley & Zachara 1992). The
mineralogy and hydrology of some of these sites has
been covered elsewhere in this volume (e.g., see
Zachara et al. 2013). Suffice it to say here that the
accidental release of caustic or acidic solutions
containing U into sediments induced dissolution–
precipitation reactions creating complicated textural
relationships between U-bearing minerals and the
surrounding host lithology. Unraveling the
mineralogy and petrology of U is key to
understanding the long term fate and transport of U
at these DOE legacy sites.
Likewise, the purposeful disposal of Ucontaining waste has to consider a host of issues
including dissolution–precipitation processes,
whether of vitrified waste or spent fuel. For
INTRODUCTION
The nuclear fuel cycle includes mining and
milling of U-ore, conversion of UO2+x phases into
uranium-hexafluouride, fuel fabrication of UO2
pellets, power generation, storage of used fuel, and
the purposeful disposal or accidental release of
materials during weapons production. Consequently,
there is substantial public concern with respect to
the long term fate of U and other radionuclides in
the environment.
The contamination of soils by metals and
radionuclides poses a threat to the biosphere
(including the food chain) and groundwater
resources. The fate and transport of metals and
radionuclides in porous media such as soil, tailings
and aquifers is strongly influenced by interactions
occurring at mineral–water interfaces. As outlined
below, a comprehensive understanding of surface
structures and U valences as well as chemical
reactions occurring at mineral–water interfaces is
essential to the management of tailings,
contaminated soil and groundwater systems.
Relevance of processes at the mineral–water
interface involving uranium
The first step in the nuclear fuel cycle is the
mining of U ore which leaves a legacy of mine
tailings. From an environmental perspective, U in
these tailings is usually not a great concern.
However, despite the generally low concentration of
U left in tailings, hot spots can persist. For example,
in the Athabasca Basin, northern Saskatchewan,
Canada, the efficiency of the U extraction process is
generally around 99% or greater. As a result, the
discharged tailings solids have U concentrations
which are generally only on the order of a few
Mineralogical Association of Canada Short Course 43, Winnipeg MB, May 2013, p. 203-253.
203
SCHINDLER & ILTON
wide range of environmentally relevant conditions,
but that the homogeneous reaction may be initiated
under certain conditions. What we do know,
regardless of the specifics, is that in order to reduce
one U6+ to U4+ a minimum of two electrons are
required and there needs to be a thermodynamic
driving force for the reaction to proceed irreversibly.
One pathway is 2Fe2+ reduces 1U6+ to 1U4+ yielding
in parallel 2Fe3+. Alternatively, the reaction of 1Fe2+
with 1U6+ could yield 1U5+. Mechanistic concerns
could then be raised, such as whether the concerted
two electron reaction proceeds directly or whether
2U6+ → 2U5+ which then disproportionate to form
1U4+ and 1U6+. Of course, favorable thermodynamics have to be in place for such reactions to
occur, but reaction pathways could be important as
well.
In the following we discuss three potential
mechanisms that may account for surfaces
facilitating U–Fe electron transfer. The first
concerns the electronic structure of adsorbed Fe2+
(Wehrli et al. 1989), the second regards surfaces
acting as structural templates for the nucleation and
growth of stable Fe3+ reaction products (Felmy et al.
2011), and the third considers surfaces as a means to
concentrate the reactants (Amonette et al. 2000,
Boyanov et al. 2007). The last mechanism may also
be dependent on whether the substrate is conductive
or non-conductive. Aside from the concentration
effect of surfaces, the emphasis is on Fe more so
than U. This is a reasonable simplification given
that surfaces strongly increase the rates of reaction
between Fe2+ and a wide range of oxidants with
very different properties. However, in the section on
structural incorporation of U, the emphasis shifts to
the bonding environment of U.
In this regard, we begin with the pioneering
study of Liger et al. (1999) who observed that
introduction of hematite particles into a solution
containing U6+(aq) and Fe2+(aq) appeared to induce
the redox reaction at pH 7.5, whereas the reaction
did not occur in the absence of an initial solid
despite their calculations indicating a favorable
driving force. We pause to mention that a recent
paper by Du et al. (2011) raises issues concerning
some of the protocols used in Liger et al. and
provides evidence that the redox reaction between
Fe2+aq and U6+(aq) could be initiated in the absence
of a solid, even at slightly acidic pH, in accord with
their thermodynamic calculations. We note that Du
et al. used higher concentrations of both reactants
(i.e., 1 mM Fe2+ and 0.2 mM U) than seen in most
environments as well as higher than used in most
experimental studies. Indeed, Zeng & Giammar
(2011) appeared to confirm that hematite does
indeed facilitate reduction of U6+ by Fe2+ for more
environmentally relevant concentrations of U6+(aq)
where no reduction occurred for the homogeneous
case, in broad agreement with Liger et al. (1999).
Further, thermodynamic calculations by Du et al.
(2011) predict that the redox reaction should not
occur for the pH 6 samples in Liger et al. (1999),
contrary to assertions made in that earlier study. As
discussed later, the thermodynamic calculations in
Felmy et al. (2011) clearly agree with Du et al.
(2011) if the Fe3+ reaction product is ferrihydrite,
but allow the redox reaction as low as pH ~6.2 if
(bulk) hematite or goethite (not shown) forms (Fig.
7-2; reproduced with kind permission of the
Mineralogical Society from a paper by Felmy et al.
2011). Although pH 6.2 is not exactly pH 6, this
illustrates the importance of the Fe(III) reaction
product in determining whether the redox reaction is
viable and we return to this point later in the
discussion. With these caveats, we believe that the
weight of the evidence indicates that surfaces do
facilitate reduction of U6+ by Fe2+, although there
might be conditions that are conducive to the rapid
initiation of the pure homogeneous redox reaction
(i.e., Du et al. 2011).
Returning to Liger et al. (1999), based on
surface complexation modeling, the authors
concluded that the formation of a dominant inner
spherically bound surface O > Fe2+(OH)n complex at
pH 7.5 was the key reactive species and that once
formed readily reduced U6+. Thus, it was proposed
that the surface of hematite functioned to activate
Fe2+, in a manner akin to that suggested by Wehrli
et al. (1989). In essence, the argument is that the
combination of inner sphere sorption and hydrolysis
increases the electron density around the Fe2+
surface species, lowering the activation energy for
oxidation of Fe2+ and consequently electron transfer
to the oxidant. In contrast, Zeng & Giammar (2011),
although largely confirming the basic tenet of Liger
et al. (1999), see above, cited a decrease in
reduction rate from pH 7.5 to 9 as inconsistent with
a dominant role for a ferrous hydroxide surface
species. (Note that the experiments were performed
in carbonate-free systems and adsorption did not
decrease with pH.)
Two recent studies (Felmy et al. 2011,
Boyanov et al. 2007) provide further insight on the
role of surfaces in facilitating reduction of U6+ by
adsorbed Fe2+. Felmy et al. (2011) used thermodynamics to make a straightforward argument
208
SCHINDLER & ILTON
ABD EL-NABY, A.H.H. (2009): Role of argillic
alteration in uranophane precipitation along shear
zones of the Gattar Granites, Eastern Desert,
Egypt. JKAU: Earth Sciences 20, 45-69.
ABDELOUAS, A. (2006): Uranium mill tailings:
Geochemistry, mineralogy, and environmental
impact. Elements 2, 335-341.
ABRAMOWSKI, M., GRIMES, R. W. & OWENS, S.
(1999): Morphology of UO2. J. Nucl. Mater. 275,
12-18.
ALESSI, D.S., USTER, B., VEERAMANI, H.,
SUVOROVA, E.I., LEZAMA-PACHECO, J.S.,
STUBBS, J.E., BARGAR, J.R. & BERNIER-LATMANI,
R. (2012): Quantitative separation of monomeric
U(IV) from UO2 in products of U(VI) reduction.
Env. Sci. Tech. 46, 6150-6157.
ALLARD, T., ILDEFONSE, P., BEAUCAIRE, C. &
CALAS, G. (1999): Structural chemistry of
uranium associated with Si, Al, Fe gels in a
granitic uranium mine. Chem. Geol. 158, 81-103.
AMONETTE, J., ISMAIL, F. T. & SCOTT, A. D. (1985):
Oxidation of iron in biotite by different oxidizing
solutions at room-temperature. Soil Sci. Soc.
Amer. J. 49, 772-777.
AMONETTE, J. E. & SCOTT, A. D. (1995): Oxidative
weathering of trioctahedral micas buffered by
H2O2 solutions. In: Churchman, G. J. and al., E.
Eds.) Clays Controlling the Environment. CSIRO
Publishing, Melbourne, Australia.
AMONETTE, J. E., WORKMAN, D. J., KENNEDY, D.
W., FRUCHTER, J. S. & GORBY, Y. A. (2000):
Dechlorination of carbon tetrachloride by Fe(II)
associated with goethite. Env. Sci. Tech. 34,
4606-4613.
AREY, J. S., SEAMAN, J. C. & BERTSCH, P. M.
(1999): Immobilization of uranium in
contaminated sediments by hydroxyapatite
addition. Env. Sci. Tech. 33, 337-342.
BEAZLEY, M. J., MARTINEZ, R. J., SOBECKY, P. A.,
WEBB, S. M. & TAILLEFERT, M. (2007): Uranium
biomineralization as a result of bacterial
phosphatase activity: Insights from bacterial
isolates from a contaminated subsurface. Env. Sci.
Tech. 41, 5701-5707.
BEAZLEY, M. J., MARTINEZ, R. J., SOBECKY, P. A.,
WEBB, S.M. & TAILLEFERT, M. (2009):
Nonreductive Biomineralization of Uranium(VI)
Phosphate via microbial phosphatase activity in
anaerobic conditions. Geomicrobiol. J. 26, 431441.
BEHRENDS, T. & VAN CAPPELLEN, P. (2005):
Competition between enzymatic and abiotic
reduction of uranium(VI) under iron reducing
conditions. Chem. Geol. 220, 315-327.
BELAI, N., FRISCH, M., ILTON, E. S., RAVEL, B. &
CAHILL, C. L. (2008): Pentavalent Uranium Oxide
via Reduction of [UO2](2+) Under Hydrothermal
Reaction Conditions. Inorg. Chem. 47, 1013510140.
BERNIER-LATMANI, R., VEERAMANI, H., VECCHIA,
E. D., JUNIER, P., LEZAMA-PACHECO, J. S.,
SUVOROVA, E. I., SHARP, J. O., WIGGINTON, N. S.
& BARGAR, J. R. (2010): Non-uraninite products
of microbial U(VI) reduction. Env. Sci. Tech. 44,
9456-9462.
BOLAND, D. D., COLLINS, R. N., PAYNE, T. E. &
WAITE, T. D. (2011): Effect of amorphous Fe(III)
oxide transformation on the Fe(II)-mediated
reduction of U(VI). Env. Sci. Tech. 45, 13271333.
BOYANOV, M. I., O'LOUGHLIN, E. J., RODEN, E. E.,
FEIN, J. B. & KEMNER, K. M. (2007): Adsorption
of Fe(II) and U(VI) to carboxyl-functionalized
microspheres: The influence of speciation on
uranyl reduction studied by titration and XAFS.
Geochim. Cosmochim. Acta. 71, 1898-1912.
BOYARCHENKOV, A. S., POTASHNIKOV, S. I.,
NEKRASOV, K. A. & KUPRYAZHKIN, A. Y. (2012):
Molecular dynamics simulation of UO2
nanocrystals surface. J. Nucl. Mater. 421, 1-8.
BRANTLEY, S. L. (2005): Reaction kinetics of
primary rock-forming minerals under ambient
conditions. In: Dreyer, J. W. (Ed.), Treat. Geoch.
5. Elsevier.
BRUGGEMAN, C. & MAES, N. (2010): Uptake of
uranium(VI) by pyrite under boom clay
conditions: influence of dissolved organic carbon.
Env. Sci. Tech. 44, 4210-4216.
BRUGGER, J., BURNS, P. C. & MEISSER, N. (2003):
Contribution to the mineralogy of acid drainage
of uranium minerals: Marecottite and the zippeitegroup. Amer. Min. 88, 676-685.
BRUNO, J., DEPABLO, J., DURO, L. & FIGUEROLA, E.
(1995): Experimental-study and modeling of the
U(VI)-Fe(OH)3 surface precipitation coprecipitation equilibria. Geochim. Cosmochim. Acta. 59,
4113-4123.
BUCK, E. C., BROWN, N. R. & DIETZ, N. L. (1996):
Contaminant uranium phases and leaching at the
Fernald site in Ohio. Env. Sci. Tech. 30, 81-88.
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Available online at www.sciencedirect.com
Geochimica et Cosmochimica Acta 117 (2013) 266–282
www.elsevier.com/locate/gca
Abiotic U(VI) reduction by sorbed Fe(II) on natural sediments
Patricia M. Fox a,⇑, James A. Davis a, Ravi Kukkadapu b, David M. Singer a,1,
John Bargar c, Kenneth H. Williams a
b
a
Lawrence Berkeley National Laboratory, Berkeley, CA 94720, USA
Environmental and Molecular Sciences Laboratory, Pacific Northwest National Laboratory, Richland, WA 99354, USA
c
Stanford Synchrotron Radiation Lightsource, SLAC, Menlo Park, CA 94025, USA
Received 31 October 2012; accepted in revised form 1 May 2013; available online 10 May 2013
Abstract
Laboratory experiments were performed as a function of aqueous Fe(II) concentration to determine the uptake and oxidation of Fe(II), and Fe(II)-mediated abiotic reduction of U(VI) by aquifer sediments from the DOE Rifle field research site in
Colorado, USA. Mössbauer analysis of the sediments spiked with aqueous 57Fe(II) showed that 57Fe(II) was oxidized on the
mineral surfaces to 57Fe(III) and most likely formed a nano-particulate Fe(III)-oxide or ferrihydrite-like phase. The extent of
57
Fe oxidation decreased with increasing 57Fe(II) uptake, such that 98% was oxidized at 7.3 lmol/g Fe and 41% at 39.6 lmol/g
Fe, indicating that the sediments had a limited capacity for oxidation of Fe(II). Abiotic U(VI) reduction was observed by
XANES spectroscopy only when the Fe(II) uptake was greater than approximately 20 lmol/g and surface-bound Fe(II)
was present, possibly as oligomeric Fe(II) surface species. The degree of U(VI) reduction increased with increasing Fe(II)
-loading above this level to a maximum of 18% and 36% U(IV) at pH 7.2 (40.7 lmol/g Fe) and 8.3 (56.1 lmol/g Fe), respectively in the presence of 400 ppm CO2. Greater U(VI) reduction was observed in CO2-free systems [up to 44% and 54% at pH
7.2 (17.3 lmol/g Fe) and 8.3 (54.8 lmol/g Fe), respectively] compared to 400 ppm CO2 systems, presumably due to differences
in aqueous U(VI) speciation. While pH affects the amount of Fe(II) uptake onto the solid phase, with greater Fe(II) uptake at
higher pH, similar amounts of U(VI) reduction were observed at pH 7.2 and 8.3 for a similar Fe(II) uptake. Thus, it appears
that abiotic U(VI) reduction is controlled primarily by sorbed Fe(II) concentration and aqueous U(VI) speciation. The range of
Fe(II) loadings tested in this study are within the range observed in biostimulation experiments at the Rifle site, suggesting that
Fe(II)-mediated abiotic U(VI) reduction could play a significant role in field settings.
Published by Elsevier Ltd.
1. INTRODUCTION
Because Fe-oxides are strong adsorbents for U(VI), redox cycling of Fe in sediments is an important process controlling the mobility of U(VI) in the environment (Waite
et al., 1994; Bargar et al., 2000). Fe(III)-bearing minerals
in sediments may undergo reductive dissolution, primarily
through biologically-mediated reactions, resulting in the
production of aqueous Fe(II), the formation of new Fe
⇑ Corresponding author. Address: Lawrence Berkeley National
Laboratory, 1 Cyclotron Road MS 74R316C, Berkeley, CA 94720,
USA.
E-mail address: [email protected] (P.M. Fox).
1
Current address: Kent State University, Kent, OH 44242, USA.
0016-7037/$ - see front matter Published by Elsevier Ltd.
http://dx.doi.org/10.1016/j.gca.2013.05.003
mineral phases, release of sorbed U(VI), and its reduction
to U(IV) (Behrends and Van Cappellen, 2005). Several
studies have demonstrated that U(VI) can be abiotically reduced by Fe-oxides and clay minerals in the presence of
Fe(II) (Liger et al., 1999; Behrends and Van Cappellen,
2005; Jeon et al., 2005; Nico et al., 2009; Chakraborty
et al., 2010; Boland et al., 2011), and by Fe(II)-bearing minerals, such as magnetite, Fe(II)-micas, and green rust
(O’Loughlin et al., 2003; Ilton et al., 2004, 2005, 2006,
2010; Jeon et al., 2005; Jang et al., 2008; Latta et al.,
2011; Singer et al., 2012a,b). In addition, Boyanov et al.
(2007) demonstrated U(VI) reduction by Fe(II) adsorbed
onto carboxyl-functionalized polystyrene microspheres at
pH 8.4, but no U(VI) reduction occurred at pH 7.5. While
reaction between aqueous Fe(II) and U(VI) (i.e. in
274
P.M. Fox et al. / Geochimica et Cosmochimica Acta 117 (2013) 266–282
There are a number of mineral phases present in this
sediment which may oxidize added Fe(II). Such minerals include crystalline Fe-oxides, ferrihydrite, and Fe-clays, that
are evident from Fig. 2 and XRD data, and Mn-oxides
(based on Table 1). Studies on the reaction of Fe(II) with
crystalline Fe oxides such as goethite, hematite, and nonstoichiometric magnetite have found that the oxidized
product resembled the structure of the sorbent mineral
phase (Williams and Scherer, 2004; Larese-Casanova and
Scherer, 2007; Gorski and Scherer, 2009; Handler et al.,
2009). However, based on the spectral areas, this doesn’t
appear to be the case in this study. Ferrihydrite, which
may represent 5–7% of the sediment Fe (see discussion in
Section 3.2.1), may react with Fe(II) to produce both goethite and magnetite (Hansel et al., 2005; Pedersen et al.,
2005; Nico et al., 2009; Boland et al., 2011), with goethite
produced at lower Fe(II) concentrations and magnetite produced at higher Fe(II) concentrations (>1 mmol Fe(II)/g)
(Hansel et al., 2005). Alternatively, 57Fe oxidation could
be due to transformation of the ferrihydrite-like mineral
to a more crystalline ferrihydrite form, as noted by Boland
et al. (2011) in a study on silicate-ferrihydrite. Reaction of
Fe(II) with Mn-oxides may also account for a fraction of
the Fe(II) oxidation. Villinski et al. (2001) noted precipitation of ferrihydrite and jacobsite (MnFe2O4) in a MnO2 and
Fe(II) system at pH 3. The sediment contains 5.5 lmol/g
Mn (or 300 lg/g Mn; Table 1), which if we assume is present entirely as Mn(IV), may oxidize up to 11 lmol/g Fe(II)
to ferrihydrite (with or without Mn substitutions) or jacobsite. While this is a large fraction of the added Fe(II) at lower Fe-loadings (e.g. 100% at 7.3 lmol/g), it represents a
smaller fraction at higher Fe(II) loadings (e.g. 28% at
39.6 lmol/g). Furthermore, it is unlikely that all of the sediment Mn can react with Fe(II) due to both physical constraints (some Mn may be physically separated from the
dissolved phase) and poisoning of Mn-oxide surfaces with
Fe(III) as the reaction proceeds (Villinski et al., 2001). A
Mn-containing ferrihydrite may display a doublet feature
at >77 K. Thus, the oxidized product could be predominantly a Fe(II)- or Fe(II)–Mn(II)-rich ferrihydrite-like or
nanogoethite mineral. Fe(II)-rich ferrihydrite was predominant oxidation product of aqueous Fe(II) and Tc(VII)
(Zachara et al., 2007). Oxidation of Fe(II), coupled to
structural Fe(III) reduction on surfaces of smectite and illite
may also occur. Schaefer et al. (2011) recently has shown
oxidation of Fe(II) by structural Fe(III) in nontronite, a
Fe-rich and predominantly Fe(III)-containing smectite.
Oxidation of Fe(II) on surfaces of clays in our samples,
however, may be limited or absent. This assessment was
based on: (a) Fe(II) content of the clay fraction, which is
>50% of the total clay Fe (Fig. 2c), and (b) lack of sorbed
Fe(II) oxidation in nontronite beyond 15% (Schaefer
et al., 2011).
Because of complex mineralogy of the sediment, we are
unable to determine the exact mechanism of Fe(II) oxidation or nature of the oxidized product and it is likely that
several reaction mechanisms are occurring simultaneously.
However, the similarity between the sextet features in the
7.3 and 39.6 lmol/g samples (Fig. 3h) and presence of intense sextets below 12 K suggest that predominant oxida-
tion could be due to transformation of the intrinsic
ferrihydrite-like phase by spiked 57Fe(II) to Fe(II)-rich ferrihydrite-like or nanogoethite phase, with or without Mn
substitutions.
3.2.4. Nature of the unoxidized 57Fe(II)
Closer examination of the low temperature (<RT)
Mössbauer spectra of the 39.6 lmol/g amended sample
(Fig. 3f–h) provides some clues as to the nature of the unoxidized 57Fe(II). Unlike in the 7.3 lmol/g sample, where
nearly 100% of the added 57Fe(II) was oxidized, in the
39.6 lmol/g amended sample a fraction of the Fe(II) is
magnetically ordered at 4.5 K. This is evident from the
broad feature at 4.25–5.25 mm/s that is a mix of the highenergy peak of magnetically ordered Fe(II) ( in Fig. 3h)
and 5th peak of the Fe(III)-sextet (Fig. 3h). To our knowledge, this is the first report of a magnetically ordered adsorbed Fe(II) on sediment surfaces. Such magnetic
ordering is characteristic of Fe(II)-rich domains with established Fe(II)–O–Fe(II) networks. Fe(OH)2 precipitated in a
Fe(II) and hematite system displayed a magnetically ordered feature at 13 K (Larese-Casanova and Scherer,
2007), while Genin et al. (1986) showed that Fe(OH)2 magnetically orders at 24 K. A detailed analysis of the
39.6 lmol/g sample in the temperature range of 40 and
4.5 K further indicated that bulk of the magnetic ordering
occurred below 8 K (Fig. EA2), which implies that the
Fe(II)–Fe(II) associations are weaker than in Fe(OH)2.
This assessment is in agreement with the chemical composition of the medium which is undersaturated with respect to
Fe(OH)2 (Table EA3). Boyanov et al. (2007) observed a
transition from monomeric to oligomeric Fe(II) surface
species from pH 7.5 to pH 8.4 on carboxyl-functionalized
microspheres. These oligomeric species would also be expected to become more common with increasing Fe(II) concentration at a given pH (Benjamin, 2002) and may play an
important role in the reactivity of the surface-bound Fe(II).
Thus, it is possible that the fraction of Fe(II) that underwent magnetic ordering at 4.5 K in the 39.6 lmol/g sample
is due to oligomeric Fe(II) surface species. The presence of a
somewhat less intense but similar feature in the 23.6 lmol/g
sample provides further support for this hypothesis
(Fig. EA5).
Larese-Casanova and Scherer (2007) detected adsorbed
Fe(II) on hematite only when Fe(II) loadings exceeded a
monolayer surface coverage. We can estimate a monolayer
surface coverage of 13.8 lmol/g for the whole soil in our
study using the surface area (3.59 m2/g, Table 1) and
assuming a site density of 3.84 lmol/m2 (Davis and Kent,
1990). This value is close to the Fe-loading at which we first
detected sorbed Fe(II) (13.2 lmol/g) and the maximum Fe
oxidation observed in this study (16.3 lmol/g). While the
concept of a monolayer surface coverage is much more
complex in a soil where Fe(II) may react with and adsorb
to a number of different mineral phases with different reactivities, this value may serve as a general guide to conditions
where adsorbed Fe(II) would be expected to exist. Thus, it
is possible that the 4.5 K Fe(II) doublet (not magnetically
ordered feature) could be due to 57Fe(II) adsorbed [e.g.,
as Fe(III)–O–Fe(II) (OH)0 and/or Fe(III)–O–Fe(II)+], on
276
P.M. Fox et al. / Geochimica et Cosmochimica Acta 117 (2013) 266–282
Bioreduced field sediments collected from the Rifle site
had similar levels of solid-associated Fe(II) to those used
in the batch experiments (Table EA10). HCl-extractable
Fe(II) concentrations ranged from 16.0 to 39.8 lmol/g for
sediments collected at various depths and distances from
the point of injection following acetate amendment and
prolonged (>100 day) sustenance of microbial iron and sulfate reduction. The highest Fe(II) concentration
(39.8 lmol/g) was observed in a sample collected nearest
to the point of acetate injection (P104, 6 m below ground
surface) of which it was estimated that no more than 7–
8 lmol/g of this Fe(II) was derived from solubilization of
FeS associated with HCl addition (Williams et al., 2011).
The Fe(II) concentrations observed in these field samples
are in the same range as those investigated in this laboratory study. Therefore, we consider the results of this study
to be very relevant to field conditions.
3.4. Abiotic reduction of U(VI)
The effect of Fe(II) addition on U(VI) uptake onto sediment was studied in batch systems in pH 7.2 and pH 8.3
buffered solutions in the presence of 400 ppm CO2
(Fig. 5). At pH 7.2 there was no observable effect of Fe(II)
on U(VI) uptake, as there was >99% U(VI) uptake occurring even in the absence of Fe(II). However, at pH 8.3
U(VI) uptake onto the sediment increased with increasing
Fe(II) addition. At the highest Fe(II) concentration
(5.6 mM) U(VI) uptake was equal at pH 7.2 and 8.3. It
should be noted that despite the use of a pH buffer, the
pH drifted downward by 0.1–0.3 pH units in the pH 7.2 system and by 0.1–0.4 pH units in the pH 8.3 system, with
greater drift occurring at higher Fe(II) uptake (Table 3).
The pH drift may account for some of the observed enhanced U(VI) uptake at higher Fe(II) concentrations in
the pH 8.3 system. However, it is clear from the U XANES
analysis (Table 3) that some of the enhanced U(VI) uptake
was also due to abiotic U(VI) reduction. The enhanced
U(VI) uptake, especially at low Fe(II) concentrations,
may also be explained by incorporation of U(VI) into newly
formed Fe-oxides as observed by Nico et al. (2009) and Boland et al. (2011).
3.4.1. XANES results
At both pH values, U(VI) reduction to U(IV) was observed by XANES, with the extent of reduction increasing
with increasing Fe(II) concentration (Table 3 and Fig. 6).
While greater levels of U reduction were observed at pH
8.3 compared to pH 7.2, it appears that this effect may be
primarily due to the greater uptake of Fe(II) onto sediments
at higher pH. For example, 18% of the solid phase U was
present as U(IV) at both pH 7.2 and 8.3 for similar Fe-loadings (40.7 and 37.6 lmol/g, respectively). In addition, a
slightly greater fraction of unoxidized or “sorbed” Fe(II)
was observed in the Mössbauer data at pH 8.3 compared
to pH 7.2 (Table 2), possibly leading to greater U(VI)
reduction at pH 8.3. To our knowledge there are no other
papers which examine abiotic U(VI) reduction by sorbed
Fe(II) as a function of Fe(II) concentration. The formation
of oligomeric Fe(II) surface species is expected to increase
Fig. 5. The effect of Fe(II) addition on U(VI) uptake onto
carbonate-free LRC. Experiments were performed in pH buffered
solutions in the presence of 400 ppm CO2 and an initial U(VI)
concentration of 50 lM. Samples were reacted for 7 days under
sterile conditions. Error bars are standard deviations of multiple
ICPMS measurements; errors for duplicate samples were equal to
or less than the measurement error.
with increasing Fe(II) concentration (Benjamin, 2002),
and the presence of these surface species may explain the
enhanced U(VI) reduction at higher Fe(II) loadings. Boyanov et al. (2007) found rapid and complete reduction of
U(VI) to U(IV) at pH 8.4, but no reduction at pH 7.5 on
carboxyl-functionalized microspheres in the presence of
Fe(II). They observed a transition from monomeric to oligomeric Fe(II) surface species from pH 7.5 to 8.4 which the
authors hypothesized to account for the reduction of U(VI)
at the higher pH through enhancement of electron transfer.
An alternative explanation for the enhanced U reduction
observed at higher pH involves the thermodynamics of
the reaction between Fe(II) and U(VI), which becomes
more favorable at higher pH (Ginder-Vogel et al., 2006;
Du et al., 2011). This explanation was invoked by Du
et al. (2011) to explain the onset of U(VI) reduction at
about pH 5.4–5.5 in a mixture of U(VI) and Fe(II) solutions. However, in our study the fact that a similar degree
of U(VI) reduction was observed at pH 7.2 and 8.3 for similar Fe(II) uptake suggests that reduction is controlled primarily by the concentrations of surface-bound or sorbed
Fe(II), particularly the presence of oligomeric Fe(II) surface
species observed in Mössbauer spectra. While pH may have
had a small effect on the extent of U(VI) reduction, it is
278
P.M. Fox et al. / Geochimica et Cosmochimica Acta 117 (2013) 266–282
in the presence of 400 ppm CO2 were undersaturated with
respect to all U(VI) mineral phases, and U(VI) is assumed
to be present solely as an adsorbed species under these conditions. In the presence of 400 ppm CO2, U(VI) aqueous
speciation is dominated by uranyl–carbonate complexes
[(UO2)2CO3(OH)3 and UO2(CO3)34 at pH 7.2 and 8.3,
respectively] that are more stable than the uranyl hydroxyl
species that dominate U(VI) aqueous speciation in the absence of CO2 (Ginder-Vogel et al., 2006). Other researchers
have noted an effect of U(VI) speciation on reduction in
homogenous solution (Du et al., 2011), in pure mineral systems (Singer et al., 2012a), and during microbial Fe(III)
reduction (Neiss et al., 2007; Ulrich et al., 2011).
3.4.2. l-XRF results
To characterize the reduced U(IV) phase, thin sections
of U(IV) and Fe(II)-reacted sediments were examined using
l-XRF. Fig. 7 shows selected elemental maps for samples
exposed to 5.6 mM Fe(II) at pH 7.2 (CO2-free) and pH
8.3 (400 ppm CO2 and CO2-free). Additional maps are
shown in the SI (Figs. EA6–8). The data show that there
are two apparent populations of U in the samples: (1) dispersed U which is spatially correlated with mineral surfaces
and (2) discrete hot spots of U. The discrete hot spots of U
likely represent precipitated phases, which may be either
U(VI) or U(IV) precipitates. As discussed above, the
samples reacted in the CO2-free system were oversaturated
with respect to the U(VI) mineral phases schoepite and
b-UO2(OH)2, however, the sample reacted with 400 ppm
CO2 was undersaturated with respect to U(VI) minerals
and thus any discrete U hot spots are likely due to U(IV)
precipitates, such as uraninite. A number of studies have
identified nanoparticulate uraninite as the reduced U phase
during abiotic reduction by Fe(II) (O’Loughlin et al., 2003;
Boyanov et al., 2007; Nico et al., 2009; Singer et al.,
2012a,b). However, others have observed the formation of
sorbed U(IV) (Chakraborty et al., 2010), U(V) incorpo-
rated into an Fe-mineral structure (Ilton et al., 2005,
2010; Boland et al., 2011), or non-uraninite U(IV) or
U(V) phases (Latta et al., 2012b). Mineral phase identification and determination of the dominant U redox state of
the discrete U particles was attempted using l-XRD and
l-XANES spectroscopy as part of the l-XRF experiment,
but ultimately conclusive identifications could not be made
within experimental error. The l-XRD patterns indicated
that these particles were X-ray amorphous, and the local
U concentration was too low to collect usable l-XANES
spectra under the experimental conditions. Ultimately, the
l-XRF results are consistent with the bulk XANES analyses; samples exposed to Fe(II) result in partial reduction of
the U(VI), possibly to uraninite.
3.5. Conclusions and implications
In this study, abiotic U(VI) reduction was controlled
primarily by concentrations of surface-bound (“sorbed”)
Fe(II) and aqueous U(VI) speciation. The extent of U(VI)
reduction was greater in CO2-free experiments than in the
presence of 400 ppm CO2. Under a CO2-free atmosphere
the solution phase was oversaturated with respect to
U(VI) mineral phases and aqueous U(VI) speciation was
dominated by uranyl hydroxyl species. These species are
more readily reduced than the uranyl-carbonate species that
dominate in the presence of CO2 [(UO2)2CO3(OH)3 and
UO2(CO3)34 at pH 7.2 and 8.3, respectively] (Ginder-Vogel et al., 2006; Ulrich et al., 2011). While not investigated
in this study, the presence of Ca would likely further decrease the extent of U(VI) reduction due to the formation
of the highly stable aqueous Ca–uranyl–carbonate complexes. In fact, in studies involving pure Fe minerals, other
researchers have noted that abiotic reduction of U(VI) by
Fe(II)-bearing minerals only occurs in the absence of Ca
(Singer et al., 2009b, 2012a,b). During bioremediation of
the Rifle aquifer, U(VI) aqueous speciation was dominated
Fig. 7. Representative maps showing distributions of U, Fe, and Si from l-XRF data of thin sections prepared from CO2-free samples reacted
with 5.7 mM Fe(II) and 0.1 mM U(VI) at pH 7.2 and 8.3 and a 400 ppm CO2 sample reacted with 5.6 mM Fe(II) and 0.05 mM U(VI) at pH
8.3. Other elemental data (K, Ca, Ti, Br, and S), pictures of each location investigated, and maps from other locations in each thin section are
available in the SI.
P.M. Fox et al. / Geochimica et Cosmochimica Acta 117 (2013) 266–282
by Ca–uranyl–carbonate complexes (Williams et al., 2011),
which likely constrained abiotic U(VI) reduction in this
environment. Competition between Fe(II) and Ca for sorption sites may further decrease the extent of abiotic U(VI)
reduction by decreasing the amount of sorbed Fe(II). Mössbauer data indicates that concentrations of surface-bound
or sorbed Fe(II) were greater at pH 8.3 than at pH 7.2
and increased with increasing Fe(II) loading. Even under
conditions where the Fe(II) mineral phases Fe(OH)2 and
siderite (FeCO3) were oversaturated, these species were
not observed in Mössbauer spectra, indicating that Fe(II)
sorption and oxidation outcompetes Fe(II) precipitation
under these conditions. Under conditions where nearly
100% of the added Fe(II) was oxidized, no abiotic U(VI)
reduction was observed. In fact, U(VI) reduction was only
detected when total Fe(II) uptake was greater than 20–
30 lmol/g, which corresponds to a sorbed Fe(II) concentration of about 7.3 lmol/g (Tables 2 and 3). The formation of
oligomeric Fe(II) surface species at higher Fe(II) uptake
(e.g., at 39.6 lmol/g) may further enhance U(VI) reduction.
Our results suggest that in field settings, Fe(II) mediated
abiotic U(VI) reduction may occur in the presence of high
Fe(II) concentrations. Solid-phase Fe(II) concentrations
(measured by extraction with 0.5 M HCl) of 16–40 lmol/g
were observed in bioreduced sediments collected from the
Rifle aquifer after a field biostimulation experiment. A similar range of HCl-extractable Fe(II) concentrations (16–
58 lmol/g) was observed in a zone of natural bioreduction
in the Rifle aquifer (Campbell et al., 2012) and in bioreduced sediments produced during column experiments
(20–90 lmol/g) performed under low sulfate concentrations
(i.e., in the absence of sulfate reduction) (Komlos et al.,
2008). While Fe(II)-mediated abiotic U(VI) reduction is
not expected to occur to a large extent in sediments at the
lower end of this range (<20–30 lmol/g), it may be an
important process in sediments at the higher Fe(II) loadings, such as may exist closer to the point of injection during field bioremediation or after extensive Fe-reduction, or
after acetate addition to the aquifer has ceased (when biological activity has decreased). The abiotic removal process
may thus partially account for repeated observations at the
Rifle site of prolonged removal of U from groundwater far
beyond the period where injected acetate levels fall below
detection (Williams et al., 2011). Such a pathway may operate in conjunction with other postulated post-stimulation
removal pathways tied to biological pathways (N’Guessan
et al., 2008) and promote prolonged uranium accumulation
[as U(IV)] within Fe(II)-rich, geochemically reduced regions
of the subsurface.
ACKNOWLEDGMENTS
Portions of this work were carried out at the Stanford Synchrotron Radiation Lightsource (SSRL) and the Environmental Molecular Sciences Laboratory (EMSL). SSRL is a Directorate of SLAC
National Accelerator Laboratory and an Office of Science User
Facility operated for the U.S. Department of Energy Office of Science by Stanford University. EMSL is a national scientific user
facility sponsored by the Department of Energy’s Office of Biological and Environmental Research and located at Pacific Northwest
National Laboratory. This research was supported by the U.S.
279
Department of Energy (DOE), Office of Science, Biological and
Environmental Research, Subsurface Biogeochemical Research
Program and was conducted as part of the Rifle IFRC project.
The Rifle IFRC project is a multidisciplinary, multi-institutional
project initially managed by the Pacific Northwest National Laboratory (PNNL) and now by Lawrence Berkeley National Laboratory. The use of trade names does not constitute endorsement by
the US government. The authors gratefully acknowledge the assistance of Christopher Fuller at the U.S. Geological Survey for assistance with the gamma-spectroscopy analysis.
APPENDIX A. SUPPLEMENTARY DATA
Supplementary data associated with this article can be
found, in the online version, at http://dx.doi.org/10.1016/
j.gca.2013.05.003.
REFERENCES
Amirbahman A., Kent D. B., Curtis G. P. and Davis J. A. (2006)
Kinetics of sorption and abiotic oxidation of arsenic(III) by
aquifer materials. Geochim. Cosmochim. Acta 70, 533–547.
Anderson R. T., Vrionis H. A., Ortiz-Bernad I., Resch C. T., Long
P. E., Dayvault R., Karp K., Marutzky S., Metzler D. R.,
Peacock A., White D. C., Lowe M. and Lovley D. R. (2003)
Stimulating the in situ activity of geobacter species to remove
uranium from the groundwater of a uranium-contaminated
aquifer. Appl. Environ. Microbiol. 69, 5884–5891.
Bargar J. R., Brown, Jr., G. E., Evans I., Rabedeau T., Rowen M.,
and Rogers J. (2002). A new hard X-ray XAFS spectroscopy
facility for environmental samples, including actinides, at the
Stanford Synchrotron Radiation Laboratory. In Proceedings of
the Euroconference and NEA Workshop on Speciation, Techniques, and Facilities for Radioactive Materials at Synchrotron
Light Sources, Grenoble, France, Sept. 10–12, 2000, Nuclear
Energy Agency/Organization for Economic Cooperation and
Development. AEN/NEA, Paris.
Bargar J. R., Reitmeyer R., Lenhart J. J. and Davis J. A. (2000)
Characterization of U(VI)-carbonato ternary complexes on
hematite: EXAFS and electrophoretic mobility measurements.
Geochim. Cosmochim. Acta 64, 2737–2749.
Behrends T. and Van Cappellen P. (2005) Competition between
enzymatic and abiotic reduction of uranium(VI) under iron
reducing conditions. Chem. Geol. 220, 315–327.
Benjamin M. M. (2002) Modeling the mass-action expression for
bidentate adsorption. Environ. Sci. Technol. 36, 307–313.
Bernier-Latmani R., Veeramani H., Vecchia E. D., Junier P.,
Lezama-Pacheco J. S., Suvorova E. I., Sharp J. O., Wigginton
N. S. and Bargar J. R. (2010) Non-uraninite products of
microbial U(VI) reduction. Environ. Sci. Technol. 44, 9456–
9462.
Boland D. D., Collins R. N., Payne T. E. and Waite T. D. (2011)
Effect of amorphous Fe(III) oxide transformation on the Fe(II)mediated reduction of U(VI). Environ. Sci. Technol. 45, 1327–
1333.
Boyanov M. I., O’Loughlin E. J., Roden E. E., Fein J. B. and
Kemner K. M. (2007) Adsorption of Fe(II) and U(VI) to
carboxyl-functionalized microspheres: the influence of speciation on uranyl reduction by titration and XAFS. Geochim.
Cosmochim. Acta 71, 1898–1912.
Campbell K. M., Kukkadapu R. K., Qafoku N. P., Peacock A. D.,
Lesher E., Williams K. H., Bargar J. R., Wilkins M. J.,
Figueroa L., Ranville J., Davis J. A. and Long P. E. (2012)
Geochemical, mineralogical and microbiological characteristics
Applied Radiation and Isotopes 70 (2012) 872–881
Contents lists available at SciVerse ScienceDirect
Applied Radiation and Isotopes
journal homepage: www.elsevier.com/locate/apradiso
Uranium removal from water using cellulose triacetate membranes added
with activated carbon
R. Villalobos-Rodrı́guez a, M.E. Montero-Cabrera a,n, H.E. Esparza-Ponce a,
E.F. Herrera-Peraza a, M.L. Ballinas-Casarrubias b
a
b
Centro de Investigación en Materiales Avanzados, Miguel de Cervantes 120, Compl. Ind. Chihuahua, CP 31109, Chihuahua, Chih., Mexico
Facultad de Ciencias Quı́micas, Universidad Autónoma de Chihuahua, Nuevo Campus s/n, Chihuahua, Chih., Mexico
a r t i c l e i n f o
a b s t r a c t
Article history:
Received 7 June 2011
Received in revised form
10 January 2012
Accepted 20 January 2012
Available online 8 February 2012
Ultrafiltration removal of uranium from water, with composite activated carbon cellulose triacetate
membranes (AC-CTA), was investigated. The filtrate was provided by uraninite dissolution with pH ¼
6–8. Removal efficiencies were calculated measuring solutions’ radioactivities. Membranes were
mainly characterized by microscopy analysis, revealing iron after permeation. Uranyl removal was
357 7%. Chemical speciation indicates the presence of (UO2)2CO3(OH)3 , UO2CO3, UO2(CO3)22 and
Fe2O3(s) as main compounds in the dissolution, suggesting co-adsorption of uranium and iron by the AC
during filtration, as the leading rejection path.
& 2012 Elsevier Ltd. All rights reserved.
Keywords:
Uranium removal
Membranes
Activated carbon
Adsorption
1. Introduction
The problem of uranium contamination in water has become
more important because of the shortage of this liquid resource for
human consumption. This is a particular concern in those areas
where, in addition to this, minerals containing uranium are
significant components of the ground.
In average, uranium content in fresh water can range from
0.01 ppb to 50 ppb, in surface water, and up to 2000 ppb in
groundwater. This amount depends on factors such as water flow,
leaching contact time with uranium sources, evaporation, partial
carbon dioxide pressure, presence of oxygen, redox conditions
and pH. Availability of complex ions such as carbonates, phosphates, vanadates, fluorides, sulfates and silicates, as well as the
interaction among these elements also play a role in uranium
content in natural water (Gómez et al., 2006; Ivanovich and
Harmon, 1992; Reyes-Cortes et al., 2007; Rossiter et al., 2010).
Uranium appears in nature with oxidation states of þ2, þ3, þ4,
þ5 and þ6, with the U(IV) and U(VI) states being the most
common. Uranium(IV) is not soluble in water and usually precipitates, while uranium(VI) forms soluble ions and is diluted in
water, and thus can be ingested.
It has been verified that uranium has toxic effects, particularly in
the urinary system (Domingo, 2001; Kurttio et al., 2006; Rossiter
n
Corresponding author. Tel.: þ52 614 4391123; fax: þ52 614 4391170.
E-mail address: [email protected] (M.E. Montero-Cabrera).
0969-8043/$ - see front matter & 2012 Elsevier Ltd. All rights reserved.
doi:10.1016/j.apradiso.2012.01.017
et al., 2010). However, uranium can be transformed into other
radioactive substances, which can cause cancer as a stochastic effect
(Bosshard et al., 1992; Kurttio et al., 2006). The USEPA (USEPA, 1974,
1986), based on a large number of studies, has established the
appropriate limit for drinking water at 30 mg L 1 for uranium and at
0.56 Bq L 1 for gross alpha counting rate. In Mexico, norm NOM127-SSA1-1994 (SSA, 2000) establishes the allowable limit for gross
alpha radioactivity at 0.56 Bq L 1.
The problem of uranium contamination in Mexico has been
studied in the north of the country. The total activity concentration of uranium present in some wells in the state of Chihuahua
ranges from 0.03 to 1.34 Bq L 1 (Villalba et al., 2006). The northern area of the city of Chihuahua is served by the Sacramento
River, linked to the San Marcos dam. In this place, the presence of
uranium-rich minerals, such as pitchblende, uraninite, uranophane, tyuyamunite and becquerelite, have been verified
(Reyes-Cortés et al., 2010). Therefore, the leaching of uranium
from geological subtract might be the main path to explain the
presence of uranium in underground and surface water. Enhanced
activity concentrations of uranium in drinking water could
represent a deep risk for the population; therefore, an efficient
and not expensive procedure for the removal of uranium from
water could be necessary.
There are several techniques for uranium removal from
groundwater. Some methods use reactive materials as barriers,
such as hydroxyapatite (Krestou et al., 2004; Simon et al., 2008),
activated carbon (Mellah et al., 2006) carbon nanotubes (Schierz
and Zanker, 2009) and elemental iron (Noubactep et al., 2006).
R. Villalobos-Rodrı́guez et al. / Applied Radiation and Isotopes 70 (2012) 872–881
2.4. Speciation analysis
The speciation analyses of the pitchblende and uranyl solution
were performed using the HYDRA/MEDUSA software, 32-bit vers,
Aug. 26, 2009 (Puigdomenech, 2009). MEDUSA means: make
equilibrium diagrams using sophisticated algorithms. It is a
Windows interface to MS-DOS versions of INPUT, SED and PREDOM which are a collection of FORTRAN programs which can
draw chemical equilibrium diagrams. Concentrations of concomitant elements in the uranium solution from Table 1 were
provided to the code as explained later further.
2.5. SEM of activated carbon
Morphological and elemental analysis has been made by SEM
JSM-5800 as it was previously indicated in 2.2.2, but AC has not
been covered with gold. In a sample of carbon taken as ‘‘blank’’, C,
O, S and Si were observed.
2.6. Activated carbon adsorption
Activated carbon (Carbochem LQ1000; 1058 m2/g surface area,
5.75 Å average pore size) previously milled and sieved, was used
for adsorption. The AC particles were of 1.6 mm average diameter
size (Ballinas-Casarrubias et al., 2006). A preliminary treatment
was performed as it is reported elsewhere (Ballinas-Casarrubias
et al., 2006) for AC solvation. For adsorption experiments on AC,
a solution with 1200 ppm uranium concentration was prepared;
0.63 g of uranyl nitrate were dissolved in 250 mL of tridistilled
water. Based on this liquid, solutions at uranium concentrations
of 120, 12 and 1.2 ppm were prepared. From each of the above
solutions, aliquots of 20 mL were poured in a conical (Erlenmeyer) flask and stirred for 12 h with 0.1 g of activated carbon
LQ1000. The adsorption of uranium was determined by liquid
scintillation alpha spectrometry using the same detector type
Triathler Hidex 425-034. For this determination, 1 mL aliquot of
the solution was taken before mixing with AC, and another, after
completing the experiment. Each aliquot was added to 19 mL of
Ultima Gold AB. It allows calculating the activity concentration
of the extracted uranium. The removal efficiencies of AC from
the uranium input solution were determined with Appendix B
Eqs. (B1) and (B2), respectively.
875
was tested in the range of 12–600 ppm. Equations are presented
in Appendix B. Removal efficiencies are shown in Table 3. AC
adsorption capacity, calculated by Eq. (B3) is evidenced in Fig. 2.
These results agree with other published works, using activated carbon of vegetal origin (Kutahyali and Eral, 2004), where
adsorption capacity was of 57.33 mg g 1.
In order to elucidate how the adsorption process occurs, a
speciation analysis (of compounds present in aqueous phase) is
made for uranium.
Speciation analysis was made for the uranyl solution, considering the presence of carbon dioxide (Green, 2008). Results are
shown in Fig. 3.
Activated carbon is mainly adsorbing U(VI) as a uranyl complex with carbonate, into the functionalized structure. Uranyl
carbonates are very stable compounds in aqueous solutions
(Langmuir, 1978). Thus, there is a different mechanism of interaction, which establishes the adequate conditions for uranium to
be adsorbed into the solid structure of carbon. It could be trapped
into the AC as was observed by SEM-EDAX. For instance, it is
proposed that the iron present within the carbon nanoporosity, is
playing an important role as a co-adsorbing agent. Iron has been
reported to interact with uranyl by surface interaction forming
mononuclear inner sphere complexes (Waite et al., 1994). Moreover, carbon functionalization also could affect U(VI) adsorption.
Carboxyl groups could form surface complexes with U(VI) as it
has been studied in other substrates (Boyanov et al., 2007).
Further studies should be done to elucidate the exact mechanism
of interaction.
Table 3
Mean values for removal efficiencies obtained
in adsorption experiments. Uncertainties are
expressed in 1s.
U Conc. (ppm)
RE (7 r)
600
450
225
120
96
48
24
12
0.107 0.11
0.22 7 0.08
0.32 7 0.06
0.56 7 0.12
0.607 0.12
0.607 0.05
0.47 7 0.06
0.33 7 0.10
3. Results and discussion
3.1. Uranium removal efficiency at different pH levels with Ac-CTA
membranes
As a preliminary study, activated carbon was tested in batch
adsorption experiments, using uranyl dissolution at several initial
concentrations, as it was mentioned in Section 2.6.
In this experiment it was found that the adsorption of uranium
from 1200 ppm solution saturates the AC. Conversely, for the
solution of 1.2 ppm, adsorption is not observed. In solutions with
concentrations of 12 and 120 ppm, the removal efficiencies (RE)
were 0.43 and 0.65, respectively. After adsorption, SEM-EDX was
made to verify the presence of uranium within the carbon.
In the SEM-EDX analysis, AC samples provided by the adsorption of 120 and 1200 ppm uranium solution, showed the presence
of C, O, S, Al, Fe, Si, K, Ca and U. Some of these elements (Al, Si, K
and Ca) could come from carbon ash as it was previously reported
for this material, but in lower proportion than 8% w/w (Rueda
Ramı́rez, 2005).
AC adsorption capacity of uranium is defined as the uranium
quantity that could be adsorbed by the activated carbon, and it
Fig. 2. Activated carbon adsorption capacity (Initial pH: 5; shaking time: 12 h,
temperature: 25 1C, adsorbent amount: 0.1 g).
880
R. Villalobos-Rodrı́guez et al. / Applied Radiation and Isotopes 70 (2012) 872–881
where A(Utotal) is the Sample activity concentration, A238U is the
activity concentration of the U-238 standard, V238U is the volume
of the U-238 standard, cpssample the Sample counting rate, cpsblank
is the blank counting rate, cpsstd is the standard counting rate, sA
is the relative uncertainty in sample activity concentration, s2e is
the squared efficiency uncertainty, countssample is the sample
counts, Vsample is the sample volume, sAesp is the absolute
uncertainty in sample activity concentration
Appendix B. Removal efficiency and calculation of Ac
adsorption capacity
The determination of the relative filtration, the removal and
the adsorption efficiencies was performed applying the procedure
for activity concentration determination by liquid scintillation
detection, described in Appendix A. Therefore, the following
expressions were applied:
FE ¼
Aconc f iltration
Aconc input
RE ¼ 1FE
ðB1Þ
ðB2Þ
where FE is the filtration efficiency, RE is the removal efficiency,
Aconcfiltration is the activity concentration of filtered solution,
Aconcinput is the Activity concentration of input solution.
The AC capacity of U(VI) adsorption, or the amount of U(VI)
adsorbed at equilibrium per unit mass of AC, qe (mg g 1), was
calculated using the following mass balance equation:
qe ¼
ðC 0 C e ÞV
V
¼ C0
RE
W
W
ðB3Þ
where C0 and Ce are the initial and equilibrium uranyl concentrations (mg L 1), V is the uranyl solution volume (L) and W is the
amount of adsorbent (g). Three measurements were made for
each sample and the results were averaged. Uncertainties were
calculated by propagation from the counting rate at the scintillation detector.
References
Anson, M., Marchese, J., Garis, E., Ochoa, N., Pagliero, C., 2004. ABS copolymeractivated carbon mixed matrix membranes for CO2/CH4 separation. J. Membr.
Sci. 243 (1-2), 19–28.
Baeza, A., Salas, A., Legarda, F., 2008. Determining factors in the elimination of
uranium and radium from groundwaters during a standard potabilization
process. Sci. Total Environ. 406 (1-2), 24–34.
Ballinas-Casarrubias, L., Terrazas-Bandala, L.P., Ibarra-Gómez, R., Mendoza-Duarte,
M., Manjarrez-Nevárez, L., González-Sánchez, G., 2006. Structural and performance variation of activated carbon-polymer films. Polym. Adv. Technol. 17
(11-12), 991–999.
Barton, C.S., Stewart, D.I., Morris, K., Bryant, D.E., 2004. Performance of three resinbased materials for treating uranium-contaminated groundwater within a
PRB. J. Hazard. Mater. 116 (3), 191–204.
Bosshard, E., Zimmerli, B., Schlatter, C., 1992. Uranium in the diet: risk assessment
of its nephro- and radiotoxicity. Chemosphere 24 (3), 309–322.
Boyanov, M.I., O’Loughlin, E.J., Roden, E.E., Fein, J.B., Kemner, K.M., 2007. Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: the influence
of speciation on uranyl reduction studied by titration and XAFS. Geochim.
Cosmochim. Acta 71 (8), 1898–1912.
Domingo, J.L., 2001. Reproductive and developmental toxicity of natural and
depleted uranium: a review. Reprod. Toxicol. 15 (6), 603–609.
Donat, R., 2009. The removal of uranium (VI) from aqueous solutions onto natural
sepiolite. JChTh 41 (7), 829–835.
Duff, M.C., Amrhein, C., 1996. Method for the separation of uranium(IV) and (VI)
oxidation states in natural waters. J. Chromatogr. A 743 (2), 335–340.
El-Sayed, A.A., 2008. Kinetics and thermodynamics of adsorption of trace amount
of uranium on activated carbon. Radiochim. Acta 96 (8), 481–486.
Favre-Réguillon, A., Lebuzit, G., Murat, D., Foos, J., Mansour, C., Draye, M., 2008.
Selective removal of dissolved uranium in drinking water by nanofiltration.
Water Res. 42 (4-5), 1160–1166.
Gäfvert, T., Ellmark, C., Holm, E., 2002. Removal of radionuclides at a waterworks.
J. Environ. Radioact. 63 (2), 105–115.
Geise, G.M., Park, H.B., Sagle, A.C., Freeman, B.D., McGrath, J.E., 2011. Water
permeability and water/salt selectivity tradeoff in polymers for desalination.
J. Membr. Sci. 369 (1-2), 130–138.
Gómez Escobar, V., Vera Tomé, F., Lozano, J.C., Martı́n Sánchez, A., 1998. Extractive
procedure for uranium determination in water samples by liquid scintillation
counting. Appl. Radiat. Isot. 49 (7), 875–883.
Gómez, P., Garralón, A., Buil, B., Turrero, M.J., Sánchez, L., de la Cruz, B., 2006.
Modeling of geochemical processes related to uranium mobilization in the
groundwater of a uranium mine. Sci. Total Environ. 366 (1), 295–309.
Green, D.W., 2008. Perry’s chemical engineers’ handbook, 8th ed. McGraw-Hill,
New York.
Grenthe, I., Lagerman, B., 1991. Studies on metal carbonate equilibria. 22.
A coulometric study of the uranium (VI)-carbonate system, the composition
of the mixed hydroxide carbonate species. Acta Chem. Scand., 45122–45128.
Huikuri, P., Salonen, L., Raff, O., 1998. Removal of natural radionuclides from drinking
water by point of entry reverse osmosis. Desalination 119 (1-3), 235–239.
Ivanovich, M., Harmon, R.S., 1992. Uranium-series disequilibrium: applications to
earth, marine, and environmental sciences. Clarendon Press.
Jang, J.-H., Dempsey, B.A., Burgos, W.D., 2007. A model-based evaluation of
sorptive reactivities of hydrous ferric oxide and hematite for U(VI). Environ.
Sci. Technol. 41 (12), 4305–4310.
Krestou, A., Xenidis, A., Panias, D., 2004. Mechanism of aqueous uranium(VI)
uptake by hydroxyapatite. Miner. Eng. 17 (3), 373–381.
Kryvoruchko, A.P., Yurlova, L.Y., Atamanenko, I.D., Kornilovich, B.Y., 2004. Ultrafiltration removal of U(VI) from contaminated water. Desalination, 162229–162236.
Kurttio, P., Salonen, L., Ilus, T., Pekkanen, J., Pukkala, E., Auvinen, A., 2006. Well
water radioactivity and risk of cancers of the urinary organs. Environ. Res. 102
(3), 333–338.
Kutahyali, C., Eral, M., 2004. Selective adsorption of uranium from aqueous
solutions using activated carbon prepared from charcoal by chemical activation. Sep. Purif. Technol. 40 (2), 109–114.
Kütahyali, C., Eral, M., 2010. Sorption studies of uranium and thorium on activated
carbon prepared from olive stones: kinetic and thermodynamic aspects.
J. Nucl. Mater. 396 (2-3), 251–256.
L’Annunciata, M.F., Kessler, M.J., 2003. Liquid Scintillation Analysis: Principles and
Practices. In: L’Annunciata, M.F. (Ed.), Handbook of Radioactivity Analysis, 2nd
ed. Elsevier Science, New York, USA.
Langmuir, D., 1978. Uranium solution-mineral equilibria at low temperatures
with applications to sedimentary ore deposits. Geochim. Cosmochim. Acta,
42547–42569.
Langmuir, D., 1997. Aqueous Environmental Geochemistry. Prentice Hall, Upper
Saddle River, NJ, USA.
Manceau, A., Gates, W.P., 1997. Surface structural model for ferrihydrite. Clays
Clay Miner. 45 (3), 448–460.
Marchese, J., Anson, M., Ochoa, N.A., Prádanos, P., Palacio, L., Hernández, A., 2006.
Morphology and structure of ABS membranes filled with two different
activated carbons. Chem. Eng. Sci. 61 (16), 5448–5454.
Maya, L., 1982. Hydrolysis and carbonate complexation of dioxouranium(VI) in the
neutral-pH range at 25 1C. Inorg. Chem. 21 (7), 2895–2898.
Mellah, A., Chegrouche, S., Barkat, M., 2006. The removal of uranium(VI) from
aqueous solutions onto activated carbon: Kinetic and thermodynamic investigations. J. Colloid Interface Sci. 296 (2), 434–441.
Noubactep, C., Schöner, A., Meinrath, G., 2006. Mechanism of uranium removal from
the aqueous solution by elemental iron. J. Hazard. Mater. 132 (2-3), 202–212.
Osmond, J.K., Cowart, J.B., 1992. Ground Water. In: Ivanovich, M., Harmon, R.S.
(Eds.), Uranium series disequilibrium. applications to earth, marine, and
environmental science. Clarendon Press, Oxford, UK, pp. 290–333.
Phillips, D.H., Gu, B., Watson, D.B., Parmele, C.S., 2008. Uranium removal
from contaminated groundwater by synthetic resins. Water Res. 42 (1-2),
260–268.
Puigdomenech, I., 2009. Hydra/Medusa Chemical Equilibrium Database and
Plotting Software, Inorganic Chemistry. KTH Royal Institute of Technology,
Stockholm, Sweeden.
Qadeer, R., Hanif, J., Saleem, M., Afzal, M., 1992. Effect of alkali metals, alkaline
earth metals and lanthanides on the adsorption of uranium on activated
charcoal from aqueous solutions. J. Radioanal. Nucl. Chem. 165 (4), 243–253.
Raff, O., Wilken, R.-D., 1999. Removal of dissolved uranium by nanofiltration.
Desalination 122 (2-3), 147–150.
Regenspurg, S., Schild, D., Schäfer, T., Huber, F., Malmström, M.E., 2009. Removal of
uranium(VI) from the aqueous phase by iron(II) minerals in presence of
bicarbonate. Appl. Geochem. 24 (9), 1617–1625.
Reyes-Cortés, M., Fuentes-Cobas, L., Torres-Moye, E., Esparza-Ponce, H., MonteroCabrera, M., 2010. Uranium minerals from the San Marcos District, Chihuahua,
Mexico. MinPe 99 (1), 121–132.
Reyes-Cortes, M., Montero-Cabrera, M.E., Renteria-Villalobos, M., Fuentes-Montero, L., Fuentes-Cobas, L., Herrera-Peraza, E.F., Esparza-Ponce, H., RodriguezPineda, A., 2007. Radioactive mineral samples from the northwest of Chihuahua City, Mexico. RMxF S53 (3), 23–28.
Reyes Baeza, J., 2008. Cuarto Informe de Gobierno del estado de Chihuahua,
Administración 2004–2010, Chihuahua, Chih., Mexico.
Rossiter, H.M.A., Graham, M.C., Schäfer, A.I., 2010. Impact of speciation on
behaviour of uranium in a solar powered membrane system for treatment of
brackish groundwater. Sep. Purif. Technol. 71 (1), 89–96.
Rueda Ramı́rez, A., 2005. Métodos de solvatación quı́mica para la reducción del
tamaño de partı́cula de carbón activado, Facultad de Ciencias Quı́micas.
Universidad Autónoma de Chihuahua, Chihuahua, Mexico.
Journal of Membrane Science 389 (2012) 499–508
Contents lists available at SciVerse ScienceDirect
Journal of Membrane Science
journal homepage: www.elsevier.com/locate/memsci
Direct quantification of negatively charged functional groups on membrane
surfaces
Alberto Tiraferri, Menachem Elimelech ∗
Department of Chemical and Environmental Engineering, Yale University, P.O. Box 208286, New Haven, CT 06520-8286, USA
a r t i c l e
i n f o
Article history:
Received 30 July 2011
Received in revised form 9 November 2011
Accepted 9 November 2011
Available online 20 November 2011
Keywords:
Surface charge
Thin-film composite membranes
Carboxylic groups
Titration
Uranyl
Uranyl cation binding
Charge density
Polyamide
Water purification
Toluidine blue O
Charge quantification
a b s t r a c t
Surface charge plays an important role in membrane-based separations of particulates, macromolecules,
and dissolved ionic species. In this study, we present two experimental methods to determine the concentration of negatively charged functional groups at the surface of dense polymeric membranes. Both
techniques consist of associating the membrane surface moieties with chemical probes, followed by
quantification of the bound probes. Uranyl acetate and toluidine blue O dye, which interact with the
membrane functional groups via complexation and electrostatic interaction, respectively, were used as
probes. The amount of associated probes was quantified using liquid scintillation counting for uranium
atoms and visible light spectroscopy for the toluidine blue dye. The techniques were validated using selfassembled monolayers of alkanethiols with known amounts of charged moieties. The surface density of
negatively charged functional groups of hand-cast thin-film composite polyamide membranes, as well
as commercial cellulose triacetate and polyamide membranes, was quantified under various conditions.
Using both techniques, we measured a negatively charged functional group density of 20–30 nm−2 for the
hand-cast thin-film composite membranes. The ionization behavior of the membrane functional groups,
determined from measurements with toluidine blue at varying pH, was consistent with published data
for thin-film composite polyamide membranes. Similarly, the measured charge densities on commercial membranes were in general agreement with previous investigations. The relative simplicity of the
two methods makes them a useful tool for quantifying the surface charge concentration of a variety of
surfaces, including separation membranes.
© 2011 Elsevier B.V. All rights reserved.
1. Introduction
Liquid separation by polymeric membranes is used routinely in
a variety of applications, including water and wastewater treatment [1,2], seawater desalination [3], liquid food processing [4],
industrial separation processes [5,6], and more recently in energy
production and storage systems [7]. Polymeric membranes often
possess surface moieties and consequently acquire surface charge
when in contact with an aqueous solution. The charged functional
groups at the surface affect the interactions of solutes with the
membrane surface, thus impacting the membrane performance.
In particular, most commercial reverse osmosis (RO) and forward
osmosis (FO) membranes have a thin-film composite (TFC) structure, whereby a thin, selective polyamide layer is cast on top of
a polysulfone support [8]. The polyamide layer possesses innate
carboxyl- and amino-groups when immersed in aqueous solution
due to incomplete cross-linking of the polymer during fabrication.
∗ Corresponding author. Tel.: +1 203 432 2789; fax: +1 203 432 4387.
E-mail address: [email protected] (M. Elimelech).
0376-7388/$ – see front matter © 2011 Elsevier B.V. All rights reserved.
doi:10.1016/j.memsci.2011.11.018
For “loose” nanofiltration (NF) and ultrafiltration membranes,
the separation of charged solutes by electrostatic (Donnan) exclusion is directly related to the density of surface charges [9–12]. In
addition, the dissociation of charged groups also affects the “openness” of the pores and therefore, the separation by sieving or size
exclusion [10]. In “tighter” NF, RO, and FO membranes, membrane
separation is governed by a solution-diffusion mechanism [13].
Generally, the presence of functional groups in these membranes
is an indication of a lower extent of polymer cross-linking due
to incomplete reaction of the monomers during interfacial polymerization. Membrane surface charge also plays a role during the
fouling of tight and loose membranes by charged macromolecules
and colloidal matter by governing the electrostatic interactions
between the foulants and the membrane surface or pore walls
[14]. Furthermore, surface moieties can be exploited as reactive
sites for binding of surface coatings [8,15] or nanomaterials [16–18]
for membrane surface modification. For these reasons, the development of simple methods for direct quantification of membrane
surface charge density is of paramount importance.
Direct quantitative measurement of membrane surface charge
has remained challenging due to limitations associated with the
characterization techniques. Titration methods cannot be readily
502
A. Tiraferri, M. Elimelech / Journal of Membrane Science 389 (2012) 499–508
Fig. 2. Calibration curve for the concentration of toluidine blue dye using optical
density at 630 nm wavelength and 1-cm path length. The light absorbance of 0.3-mL
solutions containing a known concentration of dye was measured using a 96-plate
well analyzer. Linear fit is shown for the linear region of the experimental data.
Fig. 1. Calibration curves for the number of uranium atoms using a liquid scintillation counter. Known amounts of uranyl acetate were diluted in 15 mL of scintillation
fluor in the (A) absence and (B) presence of a gold surface used for self-assembled
monolayers. Data points represent results for the 70–250 keV channel window,
after subtraction of the blank (no uranyl acetate molecules). Detection limit is
around 1 count per minute (CPM), corresponding roughly to (A) 1–3 × 10−8 and
(B) 2–5 × 10−11 moles of uranium atoms in solution.
3. Results and discussion
3.1. Calibration curves and detection limit
The calibration plots in Fig. 1A show the counts per minute
(CPM) obtained by liquid scintillation (70–250 keV energy range)
of solutions containing a known amount of uranium atoms. CPM
increased linearly with the concentration of uranyl acetate, allowing the use of a linear equation to relate CPM raw data to the
concentration of uranium atoms. The lower detection limit was
approximately 1 CPM, corresponding to around 1–3 × 10−8 moles
of uranium atoms.
Fig. 1B presents the CPM data when a clean gold sheet (used for
self-assembled monolayer to be discussed later) is immersed in the
calibration solutions. The resulting photoelectric effect in the presence of gold [47] significantly enhanced the instrument response
and altered the relationship between CPM and the uranium concentration. The enhanced sensitivity in the presence of gold resulted in
a lower detection limit of approximately 2–5 × 10−11 moles of uranium atoms in solution. The measured data in Fig. 1B are described
by a power law, which was used to convert CPM to moles of uranium
atoms.
The UCB technique relies on a number of assumptions, some
of which can result in underestimation of the density of functional groups, while others in their overestimation. First, we assume
that the complexation of uranyl ions with nearly all carboxylic
groups occurs fast, meaning that the process is diffusion-limited.
This assumption is corroborated by observations showing that the
UVI -carboxyl stability constants for surface carboxylates are high
and significantly higher than those determined in bulk solutions
under the same conditions [34,48–50]. Second, we assume that
bidentate complex formation is more probable than monodentate
complexation, giving rise to a 1:1 stoichiometric ratio of uranyl
ion to carboxylic group. Bidentate complexation has been observed
to be favorable when uranyl ions form complexes with carboxylic
groups associated with a solid surface [49,51–53]. Third, we assume
that complexation is maximized at pH 4.5. Protonation of carboxylates at lower pH can decrease the complexation kinetics [51,54],
while at high pH, association of uranyl ions with hydroxyl groups
and subsequent precipitation can occur [54,55].
All the assumptions discussed above can lead to underestimation of functional groups. Overestimation can result from the
following assumptions. First, the constructed calibration curves
allow for the quantification of uranium atoms from scintillation
data, even without ␣/␤ emission discrimination. Second, nonspecific adsorption of uranyl ions in the thin film is negligible.
Third, the support side of the membrane comes in contact with
no or a negligible amount of uranyl acetate; radioactivity of the
rinsing solutions was undetectable, thereby corroborating this
assumption. Fourth, we assume that a statistically minor amount
of alpha particles are completely emitted into the polymer and
thus, are undetectable by the scintillation counter. The fraction
of alpha particles emitted into the polymer is a complex function of the surface roughness and of the morphology of surface
features. Understanding this process would require a thorough
statistical analysis that is beyond the scope of this paper. We
note, however, that this fraction of radioactive particles could be
significant, which results in underestimation of the total number of uranyl ions at the surface. Finally, we assume complete
dissolution of uranyl acetate in solution. Filtering the solution
through a 0.1-␮m filter prior to use is recommended to avert
the presence of unbound crystals at the surface of the materials
characterized.
Fig. 2 shows the light absorbance of solutions containing a
known amount of TBO at a 630 nm wavelength and a path length
of 1 cm. The linear portion of the graph was confined in the region
between approximately 0.3 and 100 ␮M TBO, which represent the
lower and upper detection limits in our study. All experimental
data obtained in this study were within the linear response region.
The sensitivity of the instrument corresponded to variations in
absorbance on the order of 10−4 units, equivalent to 0.24 nM of
TBO in solution.
508
A. Tiraferri, M. Elimelech / Journal of Membrane Science 389 (2012) 499–508
[42] K. Bower, A. Angel, R. Gibson, T. Robinson, D. Knobeloch, B. Smith, Alphabeta discrimination liquid scintillation-counting for uranium and its daughters,
Journal of Radioanalytical and Nuclear Chemistry – Articles 181 (1994) 97–107.
[43] R. Blackburn, M.S. Almasri, Determination of uranium by liquid scintillation
and cerenkov counting, Analyst 119 (1994) 465–472.
[44] D.L. Horrocks, Measurement of low-levels of normal uranium in water and
urine by liquid scintillation alpha-counting, Nuclear Instruments & Methods
117 (1974) 589–595.
[45] S.K. Alpat, O. Ozbayrak, S. Alpat, H. Akcay, The adsorption kinetics and removal
of cationic dye, Toluidine Blue O, from aqueous solution with Turkish zeolite,
Journal of Hazardous Materials 151 (2008) 213–220.
[46] J. Jebaramya, M. Ilanchelian, S. Prabahar, Spectral studies of toluidine blue O in
the presence of sodium dodecyl sulfate, digest, Journal of Nanomaterials and
Biostructures 4 (2009) 789–797.
[47] R.H. Pratt, A. Ron, H.K. Tseng, Atomic photoelectric effect above 10 keV, Reviews
of Modern Physics 45 (1973) 273–325.
[48] K.S. Rajan, A.E. Martell, Equilibrium studies of uranyl complexes.4. Reactions
with carboxylic acids, Journal of Inorganic & Nuclear Chemistry 29 (1967)
523–529.
[49] M.I. Boyanov, E.J. O’Loughlin, E.E. Roden, J.B. Fein, K.M. Kemner, Adsorption
of Fe(II) and U(VI) to carboxyl-functionalized microspheres: the influence of
speciation on uranyl reduction studied by titration and XAFS, Geochimica et
Cosmochimica Acta 71 (2007) 1898–1912.
[50] D.A. Fowle, J.B. Fein, A.M. Martin, Experimental study of uranyl adsorption onto Bacillus subtilis, Environmental Science & Technology 34 (2000)
3737–3741.
[51] S.D. Kelly, K.M. Kemner, J.B. Fein, D.A. Fowle, M.I. Boyanov, B.A. Bunker,
N. Yee, X-ray absorption fine structure determination of pH-dependent Ubacterial cell wall interactions, Geochimica et Cosmochimica Acta 66 (2002)
3855–3871.
[52] M.A. Denecke, T. Reich, M. Bubner, S. Pompe, K.H. Heise, H. Nitsche, P.G. Allen,
J.J. Bucher, N.M. Edelstein, D.K. Shuh, Determination of structural parameters
of uranyl ions complexed with organic acids using EXAFS, Journal of Alloys and
Compounds 271 (1998) 123–127.
[53] S. Tsushima, S. Nagasaki, S. Tanaka, A. Suzuki, Investigation of the interaction between organic compounds and uranyl ions by Raman spectroscopy,
Czechoslovak Journal of Physics 49 (1999) 783–787.
[54] P. Selvam, K. Vidya, N.M. Gupta, Influence of pH on the sorption behaviour of
uranyl ions in mesoporous MCM-41 and MCM-48 molecular sieves, Materials
Research Bulletin 39 (2004) 2035–2048.
[55] P.K. Jha, G.P. Halada, The catalytic role of uranyl in formation of polycatechol
complexes, Chemistry Central Journal 5 (2011).
[56] B.S. Day, J.R. Morris, Packing density and structure effects on energy-transfer
dynamics in argon collisions with organic monolayers, Journal of Chemical
Physics 122 (2005).
[57] P.E. Laibinis, G.M. Whitesides, D.L. Allara, Y.T. Tao, A.N. Parikh, R.G. Nuzzo, Comparison of the structures and wetting properties of self-assembled monolayers
of normal-alkanethiols on the coinage metal-surfaces, Cu, Ag, Au, Journal of the
American Chemical Society 113 (1991) 7152–7167.
[58] S.N. Patole, C.J. Baddeley, D. O’Hagan, N.V. Richardson, F. Zerbetto, L.A. Zotti,
G. Teobaldi, W.A. Hofer, Self-assembly of semifluorinated n-alkanethiols on
{1 1 1}-oriented Au investigated with scanning tunneling microscopy experiment and theory, Journal of Chemical Physics 127 (2007).
[59] M.L. Lind, D.E. Suk, T.V. Nguyen, E.M.V. Hoek, Tailoring the structure of thin film
nanocomposite membranes to achieve seawater RD membrane performance,
Environmental Science & Technology 44 (2010) 8230–8235.
[60] C.Y.Y. Tang, Y.N. Kwon, J.O. Leckie, Effect of membrane chemistry and coating
layer on physiochemical properties of thin film composite polyamide RO and
NF membranes II. Membrane physiochemical properties and their dependence
on polyamide and coating layers, Desalination 242 (2009) 168–182.
[61] B.X. Mi, O. Coronell, B.J. Marinas, F. Watanabe, D.G. Cahill, I. Petrov, Physicochemical characterization of NF/RO membrane active layers by Rutherford
backscattering spectrometry, Journal of Membrane Science 282 (2006) 71–81.
[62] F.A. Pacheco, I. Pinnau, M. Reinhard, J.O. Leckie, Characterization of isolated
polyamide thin films of RO and NF membranes using novel TEM techniques,
Journal of Membrane Science 358 (2010) 51–59.
[63] S.H. Choi, Y.C. Nho, Adsorption of UO22+ by polyethylene adsorbents with
amidoxime, carboxyl, and amidoxime/carboxyl group, Radiation Physics and
Chemistry 57 (2000) 187–193.
[64] E.L. Cussler, Diffusion: Mass Transfer in Fluid Systems, 3rd ed., Cambridge University Press, 2009.
[65] C.Y.Y. Tang, Y.N. Kwon, J.O. Leckie, Effect of membrane chemistry and coating
layer on physiochemical properties of thin film composite polyamide RO and
NF membranes I. FTIR and XPS characterization of polyamide and coating layer
chemistry, Desalination 242 (2009) 149–167.
[66] H.U. Demisch, W. Pusch, Ion-exchange capacity of cellulose–acetate membranes, Journal of the Electrochemical Society 123 (1976) 370–374.
[67] H.U. Demisch, W. Pusch, Electric and electrokinetic transport properties of
homogeneous weak ion-exchange membranes, Journal of Colloid and Interface
Science 69 (1979) 247–270.
[68] C.P. Minning, K.S. Spiegler, Polarization at membrane – solution interfaces in
reverse osmosis (hyperfiltration), in: E. Selegny (Ed.), Charged Gels and Membranes I, Reidel, Dordrecht, Holland, 1976, p. p. 277.
[69] A.E. Childress, M. Elimelech, Effect of solution chemistry on the surface charge
of polymeric reverse osmosis and nanofiltration membranes, Journal of Membrane Science 119 (1996) 253–268.
[70] M. Elimelech, W.H. Chen, J.J. Waypa, Measuring the zeta (electrokinetic)
potential of reverse-osmosis membranes by a streaming potential analyzer,
Desalination 95 (1994) 269–286.
BIOMINERALIZATION AND BIOSORPTION INVOLVING BACTERIA:
METAL PHOSPHATE PRECIPITATION AND MERCURY ADSORPTION EXPERIMENTS
A Dissertation
Submitted to the Graduate School
of the University of Notre Dame
in Partial Fulfillment of the Requirements
for the Degree of
Doctor of Philosophy
by
Sarrah M. Dunham-Cheatham
Jeremy B. Fein, Director
Graduate Program in Civil and Environmental Engineering and Earth Sciences
Notre Dame, Indiana
August, 2012
Figure 6: XRD patterns from analysis of run products from U system experiments.
(Hennig et al., 2001) is a characteristic feature of the U(VI) valence state (Boyanov et al,, 2007),
and is present in the spectra of all of the samples. Both lines of evidence indicate that the vast
majority of the uranium in the biotic and abiotic samples remained as U(VI) during the
experiments, with no measureable reduction to U(IV).
EXAFS spectra at the U L3-edge show that at low saturation state conditions (biotic
sample U5), uranyl ions are present in the biotic sample dominantly as adsorbed species, bound
to carboxyl and phosphoryl groups on the bacterial cell walls. The signal strength of the
phosphorous peak (located at 3.0 Å) is weak compared to the HUP reference spectrum (Figure
8), and in general, the biotic U5 sample exhibits a markedly different spectrum than does the
HUP standard. The second oxygen peak is more distinguishable from the other samples, and the
36
Beveridge T. J. and Murray R. G. E. (1980) Sites of metal-deposition in the cell-wall of Bacillus
subtilis. Journal of Bacteriology. 141. 876-887.
Beveridge T. J. (1989) Role of cellular design in bacterial metal accumulation and mineralization.
Annual Review of Microbiology. 43. 147-171.
Billen G., Servais P. and Becquevort S. (1990) Dynamics of bacterioplankton in oligotrophic and
eutrophic aquatic environments: Bottom-up or top-down control. Hydrobiologia. 207.
37-42.
Blum and Bartha R. (1980) Effect of salinity on methylation of mercury. Bulletin of
Environmental Contamination and Toxicology. 25. 404-408.
Bonnissel-Gissinger P., Alnot M., Lickes J.-P., Ehrhardt J.-J., Behra P. (1999) Modeling the
Adsorption of Mercury(II) on (Hydr)oxides II: α-FeOOH (Goethite) and Amorphous Silica.
215. 313-322.
Bonny S., and Jones B. (2003) Microbes and mineral precipitation, Miette Hot Springs, Jasper
National Park, Alberta, Canada. Canadian Journal of Earth Sciences. 40. 1483-1500.
Borrok D and Fein J. B. (2004) Distribution of protons and Cd between bacterial surfaces and
dissolved humic substances determined through chemical equilibrium modeling.
Geochimica et Cosmochimica Acta. 68. 3043-3052.
Borrok D., Kulpa C. F., Fein J. B. (2004) Proton and Cd adsorption onto natural bacterial
consortia: testing universal adsorption behavior. Geochimica and Cosmochimica Acta.
68. 3231.
Borrok D., Turner B. F., and Fein J. B. (2005) A universal surface complexation framework for
modeling proton binding onto bacterial surfaces in geologic settings. American Journal
of Science. 305. 826-853.
Borrok D., Aumend K., Fein J. B. (2007) Significance of ternary bacteria-metal-natural organic
matter complexes determined through experimentation and chemical equilibrium
modeling. Chemical Geology. 238. 44-62.
Bosak T. and Newman D. K. (2005) Microbial kinetic controls on calcite morphology in
supersaturated solutions. Journal of Sedimentary Research. 75. 190-199.
Boyanov M. I., O’Loughlin E. J., Roden E. E., Fein J. B., and Kemner K. M. (2007) Adsorption of
Fe(II) and U(VI) to carboxyl-functionalized microspheres: The influence of speciation on
uranyl reduction studied by titration and XAFS. Geochimica et Cosmochimica Acta. 71.
1898-1912.
Braissant O., Cailleau G., Dupraz C., and Verrecchia E. P. (2003) Bacterially induced
mineralization of calcium carbonate in terrestrial environments: The role of
exopolysaccharides and amino acids. Journal of Sedimentary Research. 73. 485-490.
111
Applied Geochemistry 27 (2012) 1499–1511
Contents lists available at SciVerse ScienceDirect
Applied Geochemistry
journal homepage: www.elsevier.com/locate/apgeochem
Geochemical, mineralogical and microbiological characteristics of sediment from
a naturally reduced zone in a uranium-contaminated aquifer
K.M. Campbell a,b,⇑, R.K. Kukkadapu c, N.P. Qafoku c, A.D. Peacock d, E. Lesher e, K.H. Williams f, J.R. Bargar g,
M.J. Wilkins h,c, L. Figueroa e, J. Ranville e, J.A. Davis b,f, P.E. Long c,f
a
U.S. Geological Survey, Boulder, CO, United States
U.S. Geological Survey, Menlo Park, CA, United States
Pacific Northwest National Laboratory, Richland, WA, United States
d
Haley and Aldrich, Oak Ridge, TN, United States
e
Colorado School of Mines, Golden, CO, United States
f
Lawrence Berkeley National Laboratory, Berkeley, CA, United States
g
Stanford Synchrotron Radiation Laboratory, Menlo Park, CA, United States
h
University of California, Berkeley, CA, United States
b
c
a r t i c l e
i n f o
Article history:
Received 8 December 2011
Accepted 30 April 2012
Available online 23 May 2012
Editorial handling by D. Fortin
a b s t r a c t
Localized zones or lenses of naturally reduced sediments have the potential to play a significant role in
the fate and transport of redox-sensitive metals and metalloids in aquifers. To assess the mineralogy,
microbiology and redox processes that occur in these zones, several cores from a region of naturally
occurring reducing conditions in a U-contaminated aquifer (Rifle, CO) were examined. Sediment samples
from a transect of cores ranging from oxic/suboxic Rifle aquifer sediment to naturally reduced sediment
were analyzed for U and Fe content, oxidation state, and mineralogy; reduced S phases; and solid-phase
organic C content using a suite of analytical and spectroscopic techniques on bulk sediment and size fractions. Solid-phase U concentrations were higher in the naturally reduced zone, with a high proportion of
the U present as U(IV). The sediments were also elevated in reduced S phases and Fe(II), indicating it is
very likely that U(VI), Fe(III), and SO4 reduction has occurred or is occurring in the sediment. The microbial community was assessed using lipid- and DNA-based techniques, and statistical redundancy analysis
was performed to determine correlations between the microbial community and the geochemistry.
Increased concentrations of solid-phase organic C and biomass in the naturally reduced sediment suggests that natural bioreduction is stimulated by a zone of increased organic C concentration associated
with fine-grained material and lower permeability to groundwater flow. Characterization of the naturally
bioreduced sediment provides an understanding of the natural processes that occur in the sediment
under reducing conditions and how they may impact natural attenuation of radionuclides and other
redox sensitive materials. Results also suggest the importance of recalcitrant organic C for maintaining
reducing conditions and U immobilization.
Published by Elsevier Ltd.
1. Introduction
Uranium-contaminated groundwater is a long-term environmental problem resulting from the legacy of U mining, ore processing, and radioactive waste disposal. Even after extensive clean-up
efforts, groundwater U concentrations can still exceed levels
acceptable for site closure. The need for additional remediation
at these sites often stems from the inability of natural attenuation
to decrease dissolved U concentrations within a reasonable timeframe (Curtis et al., 2006). One such site is the Old Rifle Mill Processing site (Rifle, CO, Fig. 1), where in situ bioreduction is being
⇑ Corresponding author. Address: USGS, 3215 Marine Street, Suite E127, Boulder,
CO 80303, United States. Tel.: +1 303 541 3035; fax: +1 303 541 3084.
E-mail address: [email protected] (K.M. Campbell).
0883-2927/$ - see front matter Published by Elsevier Ltd.
http://dx.doi.org/10.1016/j.apgeochem.2012.04.013
explored as a potential strategy for the long-term remediation of
low-level U contamination in the groundwater as part of the U.S.
Department of Energy’s Integrated Field Research Challenge (IFRC)
Site in Rifle, CO. Although U attenuation in the form of natural
flushing was originally predicted to be sufficient for the site
(DOE, 1999), elevated U concentrations have persisted, making pilot-scale testing of alternative treatment methods necessary. The
processes resulting in residual U in the aquifer are not completely
understood, but may be due in part to zones or lenses of naturally
reduced sediments with elevated concentrations of U.
Naturally reduced sediments may be common in alluvial aquifers (Bargar et al., 2011) and because of their elevated concentrations of trace elements and redox-active phases, are likely to be
important for accurate estimation of natural attenuation capacity.
Although the reduced sediments may be a relatively small fraction
1506
K.M. Campbell et al. / Applied Geochemistry 27 (2012) 1499–1511
Table 1
Nitric acid extractable U, bicarbonate/carbonate extractable U, and selected organic C contents in size fractions in cores D05–D08 and one background sample (RABS). Weight
percent of each size fraction is reported as a fraction of the <2 mm size fraction.
Sediment
sample
Particle size fraction
(lm)
Weight % of the <2000 lm
fraction
Acid extractable U
(lg g1)
Bi/carbonate extractable U(VI)
(lg g1)
Organic carbon
(%)
D-08-160
<53
<106 > 53
<149 > 106
<250 > 149
<500 > 250
<1000 > 500
<2000 > 1000
<2000
4.1
14.6
7.6
14.8
23.0
26.3
9.7
–
19.0
14.5
10.4
9.2
7.5
3.9
4.3
9.3
18.9
11.9
8.5
7.8
5.6
2.9
5.2
–
0.9
–
0.4
0.3
–
0.1
0.1
0.4
D-05-160
<53
<106 > 53
<149 > 106
<250 > 149
<500 > 250
<1000 > 500
<2000 > 1000
<2000
3.0
5.6
4.1
11.6
29.9
36.1
9.9
–
3.3
2.0
1.4
1.0
0.8
0.6
0.8
1.0
2.2
1.1
0.8
0.6
0.4
0.3
0.4
–
–
–
–
–
–
–
–
–
RABS
<53
<106 > 53
<149 > 106
<250 > 149
<500 > 250
<1000 > 500
<2000 > 1000
<2000
3.1
8.3
4.4
10.5
27.5
34.1
12.2
–
4.9
3.7
2.9
2.7
1.7
1.2
0.9
1.7
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
–
Table 2
Specific surface area measured by BET analysis of <2 mm samples from D05–D08 and
background samples RABS and BKG-A.
Sediment
Specific surface area (m2 g1)
BKG-A
RABS
D-05-130
D-05-160
D-05-190
D-06-130
D-06-160
D-06-190
D-07-130
D-07-160
D-07-190
D-08-130
D-08-160
D-08-190
4.1
4.9
2.7
3.8
3.5
4.6
2.5
3.2
2.6
3.8
3.2
6.9
5.9
3.8
the sediment are likely detrital in origin; biogenic and/or authigenic magnetites that are products of biotransformation of
Fe(III)-oxides by dissimilatory Fe-reducing bacteria are small-particle and exhibit Mössbauer signatures that are distinct from those
in the samples examined in this study (Kukkadapu et al., 2005).
EXAFS and Fourier transform (FT) of sample P103-10 are shown
in Fig. 4. Initial fits were performed with up to 10% of U as uranyl,
as constrained by the XANES results. Fits to the EXAFS were performed over the data range 3–7 Å1. Addition of 10% uranyl using
parameters typical of uranyl to the fits resulted in a significant decrease of the statistical R factor. EXAFS fits to the first coordination
shell using O atoms indicated the presence of 7 ± 1 atoms at 2.32 Å,
consistent with values expected for U(IV), which is often (pseudo-)
cubically coordinated to O (Bernier-Latmani et al., 2010; Schofield
et al., 2008). A small but clear second shell frequency can be seen in
the FT at ca. 2.8 Å (Fig. 4B). Fits to this shell were attempted using a
single shell of atoms that could be present in the sediments,
including phosphate or carboxylate functional groups in biomass
or refractory organic material derived by biomass (Bernier-Latmani
Fig. 4. EXAFS (A) and corresponding Fourier transform (FT) (B) from P103-10 bulk
sample. Solid lines are data, and dotted lines are fits.
et al., 2010; Fletcher et al., 2010), and Fe or Si functional groups on
mineral surfaces. The simple EXAFS model used here assumed that
U was coordinated to a single phosphate group, or to a single Fe or
Si surface functional group. For carboxylate groups, it was assumed
that two such groups were bonded to each U(IV). Addition of 2nd
frequency shell to O-shell-only fits produced a decrease in the statistical R significant at the 90% confidence interval, as judged by
the Hamilton test (Hamilton, 1965). Fit-derived interatomic distances to U were: 3.60 Å (U–P), 3.41 Å (U–C), 3.41 Å (U–Fe), and
3.73 Å (U–Si). Whereas P or C atoms provided the lowest R factors
(0.0129 and 0.0118, respectively), the R factors for the fits using Fe
and Si (0.0145 and 0.0135, respectively) were not significantly different from those obtained using P/C. It is, therefore, not possible to
distinguish the identity of the neighboring atoms from the EXAFS
fits. Fits using P to model the 2nd shell are presented in Table S6.
The EXAFS data can be used to exclude several U(IV) phases
from consideration. The lack of U–U pair correlations at 3.85 Å allows concluding that uraninite and coffinite are at most minor
phases, and could be present only below the detection limit
K.M. Campbell et al. / Applied Geochemistry 27 (2012) 1499–1511
1507
Table 3
Descriptions of PLFA and qPCR parameters used in Fig. 5.
Abbreviation
Name
qPCR or
PLFA
Description
Geo
Monos
Cells (pmol/g)
Polys
MGN
NSats
Ebac
TbSats
Geobacter
Monoenoic
Cells in pmol/g
Polyenoic
Methyl coenzyme M reductase
Normal saturated
Bacterial biomass
Terminally branched saturated
qPCR
PLFA
PLFA
PLFA
qPCR
PLFA
qPCR
PLFA
T/C
MBSats
Hydroxy
Cy/mono
Dioic
Proteo
Trans to cis ratio
Mid-chain branched saturated
Hydroxy
Cyclopropyl to monounsaturate ratio
Dioic
Includes iron and sulfate reducing
bacteria
Dissimilatory sulfite reductase
PLFA
PLFA
PLFA
PLFA
PLFA
qPCR
Assay for Geobacter-type bacteria
Abundant in gram-negative bacteria
Biomass
Found in eukaryotes (fungi and protozoa)
Assay for archeal methanogens
High proportions often indicate less diverse populations
Biomass
Characteristic of firmicutes and gram-negative bacteria; indicates presence of anaerobic
fermenting bacteria
The higher the ratio, the less fluid in bacterial membrane; a measure of microbial stress
Often associated with anaerobic sulfate and iron reducing bacteria
Indicative of iron reducing bacteria
High ratio representative of an older population; indicates how active microbial population is
Marker for iron reducing bacteria
Cells/g delta proteobacteria
qPCR
Assay for sulfate reducing bacteria
DSR
(15–20% of total U). The absence of a 3.13 Å Si shell further suggests that coffinite is not present. The absence of P atoms at interatomic distances of 3.1–3.2 Å argues against the presence of U(IV)
phosphate minerals containing U–P pairs at these distances, such
as ningyoite (UCa(PO4)2(H2O)x), but the presence of other U(IV)
phosphate minerals cannot be excluded. Uranium–P/C interatomic
distances similar to those observed for the P103-10 sample also
have been observed for monomeric complexes of U(IV) bound to
biomass (Bernier-Latmani et al., 2010; Fletcher et al., 2010). It is,
therefore, possible that some of the U(IV) in the sediment could
be complexed by organic matter.
In a previous study focused on framboidal pyrites from P10310, <53 lm fraction, U was found to be associated with Fe and S
by electron microscopy and electron microprobe analysis (Qafoku
et al., 2009). Although it is not possible to quantify how much of
the U is associated with reduced Fe–S phases from the present
analyses, the direct evidence of U–Fe–S co-occurrence in Qafoku
et al. (2009) is consistent with bulk data presented in this study.
The absence of uraninite suggests that the mechanism of U(VI)
reduction is not similar to that of metal-reducing bacteria cultured
in a laboratory batch environment (e.g., Lovley et al., 1991; Suzuki
et al., 2002; Sharp et al., 2009). Although it is still possible that
U(VI) reduction may be enzymatic, it has also been shown in laboratory studies that adsorbed Fe(II) on Fe oxide surfaces and Fe(II)containing mineral phases (e.g., green rust, magnetite, Fe(II)-sulfides) can also reduce adsorbed U(VI) (Wersin et al., 1994; Liger
et al., 1999; Missana et al., 2003; O’Loughlin et al., 2003; Jeon
et al., 2005; Boyanov et al., 2007; Hua and Deng, 2008). Iron(II)containing clays are another possible reducing phase for U(VI);
when U diffuses into clay interlayers, electron transfer with Fe(II)
can occur (Ilton et al., 2006). Several possible reactive phases for
U are present in the natural reduced zone, and include pyrite, magnetite and Fe(II)-containing clays. However, the relative contribution of enzymatic and various possible abiotic reductants is
currently not known.
3.2. Correlating microbial community analysis to sediment
geochemistry
The microbial community varied across the sediment transect
(D05–D08), correlating to changes in the sediment composition,
particularly the abundance of biomass, Fe- and SO4-reducing bacteria, and indicators of diversity (Table 3). The results of the statis-
tical redundancy analysis calculations are presented in Fig. 5, with
the gradient from oxidizing to reducing conditions marked by an
arrow superimposed on the data from the upper left to lower right
of the plot. The advantage of RDA for this system is in the ability to
simultaneously and quantitatively compare multiple microbial and
geochemical parameters along a redox gradient. In the RDA plot,
the AVS, U, and to a lesser degree the organic C and Fe(II), are correlated, defining the most reducing sediments, located in the lower
right of Fig. 5. The HA-extractable Fe(III) is negatively correlated to
AVS and U, as expected for the more oxidizing sediments. Nitric
acid-extractable Fe(T)HNO3 does not lie on the gradient because it
was relatively constant in samples taken from cores D05–D08.
The total amount of biomass (pmol PLFA/g sediment) is highly correlated to organic C. PLFA is a better estimate of live/active bacteria
than total DNA (Ebac) since lipids are readily degraded in the environment after cell death, Whereas DNA may be relatively stable to
degradation resulting in an overestimation of the active microbial
population. The normal saturated and terminally branched saturated PLFA (NSats and TBSats) were correlated to AVS and U, suggesting that a less diverse, possibly slower growing, gram positive
bacterial community exists in the reduced sediments. Moreover,
the methyl coenzyme-M reductase (mcrA, MGN in Fig. 5) gene,
indicative of methanogens, was highly correlated with the reduced
sediments. Although methanogens probably do not directly influence the redox state of U, Fe or S, the data suggest that they be
one of the important groups of organisms in the naturally reduced
zone. It is possible that methanogens are important in metabolism
for the overall community structure given the relative scarcity of
Fe(III) and SO4 as electron acceptors. The presence of proteobacteria (Proteo), which includes Fe- and SO4-reducing bacteria, and the
dissimilatory sulfite reductase (DSR) gene targets are correlated to
HA-extractable Fe(III) in the oxidizing portion of the redox gradient. The correlation of dioic PLFA, indicative of Fe-reducing bacteria, supports this observation, although other indicators of Fe- and
SO4-reducing bacteria are uncorrelated (mid-chain branched saturated and hydroxy PLFAs, MBSats and Hydroxy in Fig. 5). In addition, the total cyclopropyl to monounsaturated precursor ratio
(Cy/mono in Fig. 5) in the more oxidized region of the RDA plot
suggests that the microbial community may be relatively active
compared to the natural reduced zone. Iron/metal-reducing bacteria are dominant in the more oxidizing sediments possibly because
there are more abundant terminal electron acceptors available for
metabolism. Sulfate-reducing bacteria have a weaker correlation
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Anderson, R.T., Vrionis, H.A., Ortiz-Bernad, I., Resch, C.T., Long, P.E., Dayvault, R.,
Karp, K., Marutzky, S., Metzler, D.R., Peacock, A., White, D.C., Lowe, M., Lovley,
D.R., 2003. Stimulating the in situ activity of Geobacter species to remove
uranium from the groundwater of a uranium-contaminated aquifer. Appl.
Environ. Microbiol. 69, 5884–5891.
Bargar, J.R., Campbell, K.M., Stubbs, J.E., Suvorova, E., Williams, K., Lezama-Pacheco,
J.S., Blue, L.Y., Cerrato, J., Bernier-Latmani, R., Giammar, D. E., Long, P.E., 2011.
Speciation and Dynamics of Biologically Reduced Uranium(IV) in the Old Rifle
Aquifer. Paper Presented at American Chemical Society Annual Meeting,
Denver, CO.
Bernier-Latmani, R., Veeramani, H., Vecchia, E.D., Junier, P., Lezama-Pacheco, J.S.,
Suvorova, E., Sharp, J.O., Wigginton, N.S., Bargar, J.R., 2010. Non-uraninite
products of microbial U(VI) reduction. Environ. Sci. Technol. 44, 5104–5111.
Boyanov, M.I., O’Loughlin, E.J., Roden, E.E., Fein, J.B., Kemner, K.M., 2007. Adsorption
of Fe(II) and U(VI) to carboxyl-functionalized microspheres: the influence of
speciation on uranyl reduction studied by titration and XAFS. Geochim.
Cosmochim. Acta 71, 1898–1912.
Campbell, K.M., Veeramani, H., Ulrich, K.U., Blue, L.Y., Giammar, D.E., BernierLatmani, R., Stubbs, J.E., Suvorova, E., Yabusaki, S., Lezama-Pacheco, J.S., Mehta,
A., Long, P.E., Bargar, J.R., 2011. Oxidative dissolution of biogenic uraninite in
ground water at Old Rifle, CO. Environ. Sci. Technol. 45, 8748–8754.
Curtis, G.P., Davis, J.A., Naftz, D.L., 2006. Simulation of reactive transport of
uranium(VI) in groundwater with variable chemical conditions. Water Resour.
Res. 42, 1–15.
Cutshall, N.H., Larsen, I.L., Olsen, C.R., 1983. Direct analysis of 210Pb in sediment
samples: self-absorption corrections. Nucl. Instrum. Methods 306, 309–312.
Davis, J.A., Meece, D.E., Kohler, M., Curtis, G.P., 2004. Approaches to surface
complexation modeling of uranium(VI) adsorption on aquifer sediments.
Geochim. Cosmochim. Acta 68, 3621–3641.
Davis, J.A., Curtis, G.P., Wilkins, M.J., Kohler, M., Fox, P., Naftz, D.L., Lloyd, J.R., 2006.
Processes affecting transport of uranium in a suboxic aquifer. Phys. Chem. Earth
31, 548–555.
Dhillon, A., Lever, M., Lloyd, K.G., Albert, D.B., Sogin, M.L., Teske, A., 2005.
Methanogen diversity evidenced by molecular characterization of methyl
coenzyme M reductase A (mcrA) genes in hydrothermal sediments of the
Guaymas Basin. Appl. Environ. Microbiol. 71, 4592–4601.
DOE, 1999, Final Site Observational Work Plan for the UMTRA Project Old Rifle Site,
U0042501, DOE, Grand Junction, 122 p.
Dong, W., Brooks, S.C., 2006. Determination of the formation constants of ternary
complexes of uranyl and carbonate with alkaline earth metals (Mg2+, Ca2+, Sr 2+,
and Ba2+) using anion exchange method. Environ. Sci. Technol. 40, 4689–4695.
Dong, W., Xie, G., Miller, T.R., Franklin, M.P., Oxenberg, T.P., Bouwer, E.J., Ball, W.P.,
Halden, R.U., 2006. Sorption and bioreduction of hexavalent uranium at a
military facility by the Chesapeake Bay. Environ. Pollut. 142, 132–142.
Fletcher, K.E., Boyanov, M.I., Thomas, S.H., Wu, Q., Kemner, K.M., Löffler, F.E., 2010.
U(VI) reduction to mononuclear U(IV) by Desulfitobacterium species. Environ.
Sci. Technol. 44, 4705–4709.
Fredrickson, J.K., Zachara, J.M., Kennedy, D.W., Duff, M.C., Gorby, Y.A., Li, S.M.W.,
Krupka, K.M., 2000. Reduction of U(VI) in goethite (a-FeOOH) suspensions by a
dissimilatory metal-reducing bacterium. Geochim. Cosmochim. Acta 64, 3085–
3098.
Fuller, C.C., van Geen, A., Baskaran, M., Anima, R., 1999. Sediment chronology in San
Francisco Bay defined by 210Pb, 234Th, 137Cs, and 239,240Pu. Mar. Chem. 64, 7–27.
Fysh, S.A., Clark, P.E., 1982. Aluminous goethite: a Mössbauer study. Phys Chem.
Miner. 8, 180–187.
Ginder-Vogel, M., Criddle, C.S., Fendorf, S., 2006. Thermodynamic constraints on the
oxidation of biogenic UO2 by Fe(III) (hydr)oxides. Environ. Sci. Technol. 40,
3544–3550.
Granger, H.C., Warren, C.G., 1969. Unstable sulfur compounds and the origin of rolltype uranium deposits. Econ. Geol. 64, 160–171.
Granger, H.C., Warren, C.G., 1974. Zoning in the Altered Tongue Associated with
Roll-Type Uranium Deposits. International Atomic Energy Agency, Vienna, pp.
185–199.
Gu, B., Yan, H., Zhou, P., Watson, D.B., Park, M., Istok, J., 2005. Natural humics impact
uranium bioreduction and oxidation. Environ. Sci. Technol. 39, 5268–5275.
Guckert, J.B., Antworth, C.P., Nichols, P.D., White, D.C., 1985. Phospholipid, esterlinked fatty acid profiles as reproducible assays for changes in prokaryotic
community structure of estuarine sediments. FEMS Microbiol. Lett. 31, 147–
158.
Hamadeh, H.H., Barghout, K., Ho, J.C., Stand, P.M., Miller, L.L., 1999. A Mössbauer
evaluation of cation distribution in titanomagnetites. J. Magn. Magn. Mater.
191, 72–78.
Hamilton, W.C., 1965. Significance tests on the crystallographic R factor. Acta
Crystallogr. 18, 502–510.
Hee, S.M., Komlos, J., Jaffé, P.R., 2007. Uranium reoxidation in previously
bioreduced sediment by dissolved oxygen and nitrate. Environ. Sci. Technol.
41, 4587–4592.
Hua, B., Deng, B., 2008. Reductive immobilization of uranium(VI) by amorphous
iron sulfide. Environ. Sci. Technol. 42, 8703–8708.
Hyun, S.P., Fox, P.M., Davis, J.A., Campbell, K.M., Hayes, K.F., Long, P.E., 2009. Surface
complexation modeling of U(VI) adsorption by aquifer sediments from a former
mill tailings site at Rifle, Colorado. Environ. Sci. Technol. 43, 9368–9373.
Ilton, E.S., Heald, S.M., Smith, S.C., Elbert, D., Liu, C., 2006. Reduction of uranyl in the
interlayer region of low iron micas under anoxic and aerobic conditions.
Environ. Sci. Technol. 40, 5003–5009.
Jeon, B.H., Dempsey, B.A., Burgos, W.D., Barnett, M.O., Roden, E.E., 2005. Chemical
reduction of U(VI) by Fe(II) at the solid-water interface using natural and
synthetic Fe(III) oxides. Environ. Sci. Technol. 39, 5642–5649.
Juottonen, H., Galand, P.E., Yrjala, K., 2006. Detection of methanogenic Archaea in
Peat: comparison of PCR primers. Res. Microbiol. 157, 914–921.
Kohler, M., Curtis, G.P., Meece, D.E., Davis, J.A., 2004. Methods for estimating
adsorbed uranium(VI) and distribution coefficients of contaminated sediments.
Environ. Sci. Technol. 38, 240–247.
Komlos, J., Peacock, A., Kukkadapu, R.K., Jaffé, P.R., 2008. Long-term dynamics of
uranium reduction/reoxidation under low sulfate conditions. Geochim.
Cosmochim. Acta 72, 3603–3615.
Křepelová, A., Sachs, S., Bernhard, G., 2006. Uranium(VI) sorption onto kaolinite in
the presence and absence of humic acid. Radiochim. Acta 94, 825–833.
Kukkadapu, R.K., Zachara, J.M., Fredrickson, J.K., Kennedy, D.W., Dohnalkova, A.C.,
McCready, D.E., 2005. Ferrous hydroxy carbonate is a stable transformation
product of biogenic magnetite. Am. Mineral. 90, 510–515.
Liger, E., Charlet, L., Van Cappellen, P., 1999. Surface catalysis of uranium(VI)
reduction by iron(II). Geochim. Cosmochim. Acta 63, 2939–2955.
Lovley, D.R., Phillips, E.J.P., Gorby, Y.A., Landa, E.R., 1991. Microbial reduction of
uranium. Nature 350, 413–416.
Marcelino, T., Suzuki, Y., Taylor, L.T., DeLong, E.F., 2000. Quantitative analysis of
small subunit rRNA genes in mixed microbial populations via 50 -nuclease
assays. Appl. Environ. Microbiol. 66, 4605–4614.
Mason, C.F.V., Turney, W., Thomson, B.M., Lu, N., Longmire, P.A., Chisholm-Brause,
C.J., 1997. Carbonate leaching of uranium from contaminated soils. Environ. Sci.
Technol. 31, 2707–2711.
Mayberry, W.R., Lane, J.R., 1993. Sequential alkaline saponification/acid hydrolysis/
esterification: a one-tube method with enhanced recovery of both cyclopropane
and hydroxylated fatty acids. J. Microbiol. Methods 18, 21–32.
Michalsen, M.M., Peacock, A.D., Spain, A.M., Smithgal, A.N., White, D.C., SanchezRosario, Y., Krumholz, L.R., Istok, J.D., 2007. Changes in microbial community
composition
and
geochemistry
during
uranium
and
technetium
bioimmobilization. Appl. Environ. Microbiol. 73, 5885–5896.
Missana, T., García-Gutiérrez, M., Fernńdez, V., 2003. Uranium (VI) sorption on
colloidal magnetite under anoxic environment: experimental study and surface
complexation modelling. Geochim. Cosmochim. Acta 67, 2543–2550.
Moon, H.S., Komlos, J., Jaffe, P.R., 2007. Uranium reoxidation in previously
bioreduced sediment by dissolved oxygen and nitrate. Environ. Sci. Technol.
41, 4587–4592.
N’Guessan, A.L., Vrionis, H.A., Resch, C.T., Long, P.E., Lovley, D.R., 2008. Sustained
removal of uranium from contaminated groundwater following stimulation of
dissimilatory metal reduction. Environ. Sci. Technol. 42, 2999–3004.
O’Loughlin, E.J., Kelly, S.D., Cook, R.E., Csencsits, R., Kemner, K.M., 2003. Reduction of
U(VI) by mixed iron(II/III) hydroxide (green rust): formation of UO2
nanoparticles. Environ. Sci. Technol. 37, 721–727.
Owen, D.E., Otton, J.K., 1995. Mountain wetlands: efficient uranium filters –
potential impacts. Ecol. Eng. 5, 77–93.
Plant, J.A., Simpson, P.R., Smith, B., Windley, B.F., 1999. Uranium ore deposits –
products of the radioactive earth. Rev. Miner. Geochem. 38, 254–319.
Qafoku, N.P., Kukkadapu, R.K., McKinley, J.P., Arey, B.W., Kelly, S.D., Wang, C., Resch,
C.T., Long, P.E., 2009. Uranium in framboidal pyrite from a naturally bioreduced
alluvial sediment. Environ. Sci. Technol. 43, 8528–8534.
Rancourt, D.G., Christie, I.A.D., Royer, M., Kodama, H., Robert, J.-L., Lalonde, A.E.,
Murad, E., 1994. Determination of accurate [4]Fe3+, [6]Fe3+, and [6]Fe2+ site
populations in synthetic annite by Mössbauer spectroscopy. Am. Miner. 79, 51–
62.
Read, D., Bennett, D.G., Hooker, P.J., Ivanovich, M., Longworth, G., Milodowski, A.E.,
Noy, D.J., 1993. The migration of uranium into peat-rich soils at Broubster,
Caithness, Scotland, UK. J. Contam. Hydrol. 13, 291–308.
Regenspurg, S., Margot-Roquier, C., Harfouche, M., Froidevaux, P., Steinmann, P.,
Bernier-Latmani, P.J.R., 2010. Speciation of naturally-accumulated uranium in
an organic-rich soil of an alpine region (Switzerland). Geochim. Cosmochim.
Acta 74, 2082–2098.
Rehr, J.J., Zabinsky, S.I., 1992. High-order multiple scattering calculations of X-ray
absorption fine structure. Phys. Rev. Lett. 69, 3397–3400.
Reynolds, R.L., Goldhaber, M.B., 1983. Iron disulfide minerals and the genesis of rolltype uranium deposits. Econ. Geol. 78, 105–120.
Schofield, E.J., Veeramani, H., Sharp, J.O., Suvorova, E., Bernier-Latmani, R., Mehta, A.,
Stahlman, J., Webb, S.M., Clark, D.L., Conradson, S.D., Ilton, E.S., Bargar, J.R., 2008.
Structure of biogenic uraninite produced by Shewanella oneidensis strain MR-1.
Environ. Sci. Technol. 42, 7898–7904.
Senko, J.M., Kelly, S.D., Dohnalkova, A.C., McDonough, J.T., Kemner, K.M., Burgos,
W.D., 2007. The effect of U(VI) bioreduction kinetics on subsequent reoxidation
of biogenic U(IV). Geochim. Cosmochim. Acta 71, 4644–4654.
Sharp, J.O., Schofield, E.J., Veeramani, H., Suvorova, E.I., Kennedy, D.W., Marshall,
M.J., Mehta, A., Bargar, J.R., Bernier-Latmani, R., 2009. Structural similarities
between biogenic uraninites produced by phylogenetically and metabolically
diverse bacteria. Environ. Sci. Technol. 43, 8295–8301.
Stevenson, F.J., 1994. Humus Chemistry: Genesis, Composition, Reactions, second
ed. John Wiley and Sons, New York.
Stookey, L.L., 1970. Ferrozine – a new spectophotometric reagent for iron. Anal.
Chem. 42, 779–781.
Stucki, J.W., Lee, K., Goodman, B.A., Kostka, J.E., 2007. Effects of in situ
biostimulation on iron mineral speciation in a sub-surface soil. Geochim.
Cosmochim. Acta 71, 835–843.
J Nanopart Res (2011) 13:3741–3754
DOI 10.1007/s11051-011-0296-0
RESEARCH PAPER
U(VI) reduction by Fe(II) on hematite nanoparticles
Hui Zeng • Daniel E. Giammar
Received: 21 May 2010 / Accepted: 14 February 2011 / Published online: 26 February 2011
Ó Springer Science+Business Media B.V. 2011
Abstract Nanoscale size effects on U(VI) reduction
by Fe(II) on hematite were investigated with four
aerosol-synthesized hematite nanoparticles (12, 30,
50, 125 nm) and one aqueous-synthesized hematite
(70 nm). Batch experiments were conducted at loadings of 0.01 mM U(VI) and 5 mM Fe(II) at pH 7.5
and 9.0. Rate constants for reduction of U(VI) to
U(IV) were determined using a pseudo-first order
reaction rate law. Reduction was faster at pH 7.5 than
at pH 9.0. Rate constants were higher for aerosolsynthesized hematite than for aqueous-synthesized
hematite. Rate constants were not significantly
different for the 30, 50, and 125 nm particles.
However, reduction was two orders of magnitude
faster for the 12 nm hematite particles. Possible
explanations for the dramatically faster reduction
with the 12 nm hematite include the formation of a
more reactive solid such as magnetite, effects on
H. Zeng D. E. Giammar (&)
Department of Energy, Environmental and Chemical
Engineering, Washington University, St. Louis,
MO 63130, USA
e-mail: [email protected]
D. E. Giammar
Center for Materials Innovation, Washington University,
St. Louis, MO 63130, USA
Present Address:
H. Zeng
TLC EnviroTech, Dallas, TX 75423, USA
electron conduction through hematite, and quantum
confinement effects.
Keywords Nanoparticles Hematite Uranium reduction Adsorption Environmental remediation Aerosols
Introduction
Iron oxides have reactive surfaces for adsorption
(Benjamin et al. 1996; Madden and Hochella 2005;
Madden et al. 2006; Payne et al. 1998; Waychunas
et al. 2005) and surface-mediated oxidation–reduction reactions (Jeon et al. 2005; Liger et al. 1999;
Williams and Scherer 2004). The structure and
energetics of the surfaces of nanoparticles can be
substantially different from those of larger particles.
Nanoscale size may affect iron oxide reactivity
towards adsorbates through the large specific surface
areas of nanoparticles and through specific size
effects. Size effects include greater proportions of
surface sites at edges or corners, distorted coordination environments of adsorbate atoms or sorbent
surface atoms, and quantum confinement effects
(Brus 1983; Chen et al. 2002; Chernyshova et al.
2007; Korgel and Monbouquette 1997; Madden et al.
2006; Rossetti et al. 1983; Toyoda and Tsuboya
2003; Waychunas et al. 2005).
Iron oxide structures are influenced by nanoscale
size effects. An X-ray absorption near-edge structure
123
3752
reduction of adsorbed U(VI) (Liger et al. 1999). In
agreement with this observation, other researchers
reported slightly higher pseudo-first order reduction
rate constants of 4-chloronitrobenzene by hydrolyzed
Fe(II) surface complexes on TiO2 at pH 9.0 than at
pH 7.5 (Nano and Strathmann 2008). However, they
also found that the link between Fe(II) speciation and
the rates of redox reactions was partially dependent
on the identity of the oxidizing species. For oxamyl,
reduction rates were similar at pH 7.5 and 9.0 and
were most strongly correlated with the volumetric
total adsorbed Fe(II) concentration. For U(VI) reduction by surface-bound Fe(II) on carboxyl-functionalized microspheres, reduction was limited at pH 7.5
but extensive at pH 8.4 (Boyanov et al. 2007), a
finding attributed to the formation of Fe(II) oligomers
(e.g., adsorbed Fe(II) polymers and surface precipitates) at higher pH values that were responsible for
the enhanced reactivity.
The results of the present study are not consistent
with those of these previous investigations. According to the interpretation of Liger et al., a similar or
higher solid-bound Fe(II) concentration at pH 9.0
would correspond to a higher concentration of
hydroxylated Fe(II) surface complexes and a higher
U(VI) reduction rate, but for both 30 and 70 nm
hematite a lower U(VI) reduction rate was found at
pH 9.0 than at 7.5 (Liger et al. 1999). The hydroxylated Fe(II) surface complexes are probably not the
dominant reactive surface species. Reduction rates
also do not correlate with total Fe(II) surface
complexes. In our hematite-free control experiment
at pH 9.0, nearly all Fe(II) precipitated, and U(VI)
adsorbed to the Fe(II) precipitate (presumably
Fe(OH)2(s)) and was slightly reduced after 24 h of
reaction (Fig. 3). The reduction rate constant with the
Fe(II) precipitate was similar to that for solid-bound
Fe(II) on the 70 nm hematite at pH 9.0 (Table 2).
This suggests that Fe(OH)2(s) precipitates may contribute to U(VI) reduction. Therefore, the decrease in
U(VI) reduction at pH 9.0 relative to pH 7.5 may be
primarily caused by the change in distribution of
Fe(II) among hematite-bound and Fe(OH)2(s) species.
Further investigation of structural differences in
adsorbed Fe(II) species (e.g., coordination environment) at pH 7.5 and 9.0 is needed to more firmly
determine the mechanisms for the observed decreases
in the rates and extents of U(VI) reduction with
increasing pH.
123
J Nanopart Res (2011) 13:3741–3754
Implications
The reduction of U(VI) by solid-bound Fe(II) on
hematite is a potentially important pathway for immobilization of uranium in subsurface environments.
Hematite exerts important effects on mineral–water
reactions such as adsorption and surface-mediated
redox reactions. Nanoparticles have advantages of
high surface area to mass ratios, high adsorption
affinity, and capacity to dissolved ions, and also faster
oxidation–reduction rates. Hematite nanoparticles
have potential applications in environmental remediation of uranium contamination. Delivery of hematite
nanoparticles to uranium-contaminated soil and
groundwater may immobilize U(VI) by adsorption
and by reduction to less mobile U(IV). Further
understanding of the reactivity of hematite water
interfaces and their reactivity with respect to U(VI)
will benefit the development of uranium remediation
technologies.
Acknowledgments This research was supported by the
National Science Foundation (BES 0608749). The Center for
Materials Innovation at Washington University provided
supplemental support. Dr. Soubir Basak and Manoranjan Sahu
in Dr. Pratim Biswas’ Aerosol and Air Quality Research
Laboratory at Washington University provided the aerosolsynthesized hematite nanoparticles. Zimeng Wang performed
selected control experiments. The valuable comments of four
anonymous reviewers were helpful in revising this manuscript.
References
Anschutz AJ, Penn RL (2005) Reduction of crystalline iron(III)
oxyhydroxides using hydroquine: influence of phase and
particle size. Geochem Trans 6:60–66
Behrends T, Van Cappellen P (2005) Competition between
enzymatic and abiotic reduction of uranium(VI) under
iron reducing conditions. Chem Geol 220:315–327
Benjamin MM, Sletten RS, Bailey RP, Bennett T (1996)
Sorption and filtration of metals using iron-oxide-coated
sand. Water Res 30:2609–2620
Boyanov MI, O’Loughlin EJ, Roden EE, Fein JB, Kemner KM
(2007) Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: the influence of speciation on
uranyl reduction studied by titration and XAFS. Geochim
Cosmochim Acta 71:1898–1912
Brus LE (1983) A simple model for the ionization potential,
electron affinity, and aqueous redox potential of small
semiconductor crystallites. J Chem Phys 79:5566–5571
Catalano JG, Fenter P, Park C, Zhang Z, Rosso KM (2010)
Structure and oxidation state of hematite surfaces reacted
with aqueous Fe(II) at acidic and neutral pH. Geochim
Cosmochim Acta 74:1498–1512
Available online at www.sciencedirect.com
Geochimica et Cosmochimica Acta 75 (2011) 2512–2528
www.elsevier.com/locate/gca
Products of abiotic U(VI) reduction by biogenic magnetite
and vivianite
Harish Veeramani a,1,⇑, Daniel S. Alessi a, Elena I. Suvorova a,
Juan S. Lezama-Pacheco b, Joanne E. Stubbs b, Jonathan O. Sharp a,2, Urs Dippon c,
Andreas Kappler c, John R. Bargar b, Rizlan Bernier-Latmani a
a
Environmental Microbiology Laboratory, Ecole Polytechnique Fédérale de Lausanne (EPFL), Switzerland
b
Stanford Synchrotron Radiation Lightsource (SSRL), USA
c
Geomicrobiology, Center for Applied Geosciences (ZAG), Universität Tübingen, Germany
Received 20 September 2010; accepted in revised form 14 February 2011; available online 21 February 2011
Abstract
Reductive immobilization of uranium by the stimulation of dissimilatory metal-reducing bacteria (DMRB) has been investigated as a remediation strategy for subsurface U(VI) contamination. In those environments, DMRB may utilize a variety of
electron acceptors, such as ferric iron which can lead to the formation of reactive biogenic Fe(II) phases. These biogenic
phases could potentially mediate abiotic U(VI) reduction. In this work, the DMRB Shewanella putrefaciens strain CN32
was used to synthesize two biogenic Fe(II)-bearing minerals: magnetite (a mixed Fe(II)–Fe(III) oxide) and vivianite (an
Fe(II)-phosphate). Analysis of abiotic redox interactions between these biogenic minerals and U(VI) showed that both
biogenic minerals reduced U(VI) completely. XAS analysis indicates significant differences in speciation of the reduced uranium after reaction with the two biogenic Fe(II)-bearing minerals. While biogenic magnetite favored the formation of structurally ordered, crystalline UO2, biogenic vivianite led to the formation of a monomeric U(IV) species lacking U–U
associations in the corresponding EXAFS spectrum. To investigate the role of phosphate in the formation of monomeric
U(IV) such as sorbed U(IV) species complexed by mineral surfaces, versus a U(IV) mineral, uranium was reduced by biogenic
magnetite that was pre-sorbed with phosphate. XAS analysis of this sample also revealed the formation of monomeric U(IV)
species suggesting that the presence of phosphate hinders formation of UO2. This work shows that U(VI) reduction products
formed during in situ biostimulation can be influenced by the mineralogical and geochemical composition of the surrounding
environment, as well as by the interfacial solute–solid chemistry of the solid-phase reductant.
Ó 2011 Elsevier Ltd. All rights reserved.
1. INTRODUCTION
Uranium mining and processing for nuclear weapons
production has led to extensive uranium contamination of
soil and groundwater at US Department of Energy
(DOE) sites. Among options for remediating uranium⇑ Corresponding author. Tel.: +1 540 922 5540.
E-mail address: [email protected] (H. Veeramani).
Current address: Department of Geosciences, Virginia Polytechnic Institute and State University, USA.
2
Current address: Environmental Science & Engineering, Colorado School of Mines, USA.
1
0016-7037/$ - see front matter Ó 2011 Elsevier Ltd. All rights reserved.
doi:10.1016/j.gca.2011.02.024
contaminated sites, in situ reductive bioremediation has
received appreciable attention due to its perceived costeffectiveness when compared to pump-and-treat methods,
and because it obviates the need for off-site handling of hazardous materials (Palmisano and Hazen, 2003). Hexavalent
uranium [U(VI)], the valence state of contaminant uranium
at most sites, is stable in oxic environments and typically
occurs as aqueous carbonate complexes in oxic groundwater at circumneutral pH. In contrast, tetravalent uranium
[U(IV)], produced by biological or abiotic processes, is
stable in anoxic environments and often occurs as the
sparingly soluble mineral, uraninite (UO2) (Langmuir,
1978).
Abiotic U(VI) reduction by biogenic magnetite and vivianite
Subsurface uranium can be reduced by a number of abiotic (Behrends and Van Cappellen, 2005) and microbiallymediated processes (Abdelouas et al., 1998; Elias et al.,
2003) including reductive immobilization of uranium by
dissimilatory metal-reducing bacteria (DMRB) (Lovley,
1993). These bacteria catalyze U(VI) reduction using organic acids, alcohols or H2 as electron donors and utilize
Fe(III) as growth-supporting electron acceptors.
Fe-oxides and iron-bearing clay minerals, which are
widely distributed in soils and sediments, represent a large
reserve of Fe(III) for DMRB (Kostka et al., 2002; Zachara
et al., 2002; Kappler and Straub, 2005). This includes contaminated US-DOE sites where cleanup efforts are underway to immobilize uranium as uraninite (N’Guessan
et al., 2008). Evidence from both field and laboratory studies also suggest a nexus between iron redox cycling and uranium redox processes (Galloway, 1978; Posey-Dowty et al.,
1987). The biostimulation of DMRB will likely lead to biological Fe(III) reduction (Wielinga et al., 2000; Finneran
et al., 2002; Anderson et al., 2003; Elias et al., 2004;) and
production of sorbed Fe(II) or Fe(II)-bearing minerals as
metabolic products. The Fe(II)-bearing phases found include magnetite, siderite, vivianite, ferruginous smectite,
and green rust (Bell et al., 1987; Roden and Zachara,
1996; Fredrickson et al., 1998; Zachara et al., 1998; Dong
et al., 2000; Roh et al., 2003; O’Loughlin et al., 2007;
Komlos et al., 2008; O’Loughlin et al., 2010). Sorbed Fe(II)
and the Fe(II)-bearing biogenic phases can provide a reservoir of reducing capacity where reduction of U(VI) may
occur due to abiotic interactions (O’Loughlin et al.,
2010). This process may compete with direct enzymatic
microbial reduction of U(VI) (Fredrickson et al., 2000).
Although reduction of U(VI) by aqueous Fe(II) is thermodynamically favorable, it can be kinetically limited, often
necessitating an appropriate adsorbent to react with aqueous Fe(II) and catalyze the reaction. Research thus far has
demonstrated U(VI) reduction by Fe(II) sorbed onto a variety of iron oxides/oxyhydroxides (Charlet et al., 1998; Liger
et al., 1999; Fredrickson et al., 2000; Jeon et al., 2004),
Fe(II)-containing natural sediments (Behrends and Van
Cappellen, 2005; Jeon et al., 2005), Fe(II)-containing
carboxyl-functionalized microspheres (Boyanov et al.,
2007), Fe(II) sorbed on corundum (Regenspurg et al.,
2009) and Fe(II) sorbed on montmorillonite (Chakraborty
et al., 2010). These studies primarily consider surface
catalyzed processes that involved either concomitant or
sequential adsorption of aqueous Fe(II) and U(VI) species
onto a solid phase adsorbent or mineral to mediate abiotic
U(VI) reduction.
Likewise, U(VI) can adsorb directly onto Fe(II)-bearing
minerals and undergo reduction by structurally bound
Fe(II). For instance, chemogenic green rust and silicates
including various micas as well as ferrous-bearing sulfide
minerals such as galena and pyrite have been shown to adsorb and reduce U(VI) (Wersin et al., 1994; O’Loughlin
et al., 2003; Ilton et al., 2004; Ilton et al., 2005; Ilton
et al., 2006; Bruggeman and Maes, 2010).
Biogenic Fe(II)-bearing minerals are of interest in the
context of uranium redox cycling and bioremediation
because they are formed under Fe-reducing conditions
2513
(Behrends and Van Cappellen, 2005). Previous studies that
focused on chemogenic analogs may not have accounted
for important properties characteristic of biogenic minerals
such as their nano-size and associated enhanced reactivity
(O’Loughlin et al., 2003; Regenspurg et al., 2009). Two
such minerals are biogenic magnetite and vivianite both
of which have shown to be produced as an end product
of microbial Fe(III) reduction and environmentally pertinent under Fe(III) reducing conditions (Fredrickson
et al., 1998; Kostka et al., 2002).
Interactions between U(VI) and magnetite have received
appreciable attention because magnetite is a ubiquitous,
environmentally relevant ferrous-bearing oxide, a metabolic byproduct of bacterial respiration, and a corrosion
product of steel with ramifications for nuclear waste repositories (Ishikawa et al., 1998; Dodge et al., 2002; Ilton et al.,
2010). Microbial reduction of amorphous ferric oxyhydroxide (Fe(OH)3) has been reported to induce the formation of
magnetite (Bell et al., 1987; Lovley et al., 1987; Moskowitz
et al., 1989; Zhang et al., 1997; Konhauser, 1998). Similarly,
magnetite formation from the reduction of aqueous Fe(III)
precursors catalyzed by sulfate-reducing microorganisms
such as Desulfovibrio spp. has been reported (Sakaguchi
et al., 1993; Sakaguchi et al., 2002). Magnetite formation
has also been reported during biooxidation of Fe(II) coupled to denitrification (Chaudhuri et al., 2001). A number
of studies have investigated the role of magnetite in uranium reduction and the findings varied greatly ranging
from no observable reduction (Dodge et al., 2002) to clear
evidence of reduction (Scott et al., 2005; Aamrani et al.,
2007; O’Loughlin et al., 2010) to the formation of a
mixed-valence U(IV)–U(VI) phase (Missana et al., 2003;
Aamrani et al., 2007; Regenspurg et al., 2009) or the formation of U(V) (Ilton et al., 2010). The variation in findings is
presumably linked to variability in morphology, specific
surface area and phase stoichiometry (Gorski and Scherer,
2009, Gorski et al., 2010) of the magnetite used as well as
differences in experimental conditions.
In phosphate-rich reducing environments, vivianite
(Fe3(PO4)28H2O) is an important sink for dissolved Fe(II)
and is considered a stable mineral due to its low solubility
at neutral pH (Nriagu and Dell, 1974; Buffle et al., 1989;
Manning et al., 1991; Al-Borno and Tomson, 1994; Viollier
et al., 1997; Sapota et al., 2006). Under anoxic conditions,
vivianite is very stable (Ksp = 10 36; (Nriagu, 1972)) and
can exert significant control over the geochemical cycles
of Fe and P. Vivianite has also been reported as an end
product of bacterial Fe(III) reduction (Fredrickson et al.,
1998; Zachara et al., 1998; Roh et al., 2007; Peretyazhko
et al., 2010). To our knowledge, the role of biogenic vivianite in abiotic uranium reduction and subsequent immobilization has not been investigated.
Redox processes linking biogenic magnetite or vivianite
and uranium were systematically investigated and the factors controlling the product of U(VI) reduction probed in
the present study. Biogenic magnetite and vivianite were
produced by Shewanella putrefaciens CN32 and characterized by scanning electron microscopy (SEM), transmission
electron microscopy (TEM), X-ray powder diffraction
(XRD) and Mössbauer spectroscopy. Their propensity to
Abiotic U(VI) reduction by biogenic magnetite and vivianite
including monomeric U(IV) species. Importantly, the presence of structural or sorbed phosphate inhibits uraninite
formation. While the precise mechanism of this inhibition
is unknown, it appears that monomeric U(IV) is associated
with the phosphate groups that are either adsorbed and/or
structurally bound to Fe(II)-bearing minerals.
While the reactivity of biogenic uraninite has been studied and documented (Ulrich et al., 2008; Ulrich et al., 2009),
the reactivity and stability of monomeric U(IV) in the environment is unknown. The results presented in this paper
suggest that there is a wealth of U(IV) chemistry not fully
understood in these systems, and that there may be complex
mixtures of U(IV) products in the field. For accurate predictions of the stability of reduced U in the subsurface, it
will be critical to consider the stability of these species in future hydrogeochemical models. A thorough understanding
of the structure, composition, occurrence, and stability of
these species is crucial to assess the feasibility of in situ
reductive bioremediation.
ACKNOWLEDGEMENTS
We thank Dorothy Parker, Anca Haiduc, Dan Giammar and
Brad Tebo for helpful discussion and feedback in preparing this
manuscript. Funding for this project was provided by a DOEOBER Grant to S.L.A.C. (work package number 2009-SLAC10006), and Grant No. DE-FG02-06ER64227 to E.P.F.L. and
Swiss NSF Grants No. 20021-113784 and No. 200020-126821/1.
Portions of this research were carried out at the Stanford Synchrotron Radiation Lightsource, a national user facility operated by
Stanford University on behalf of the US DOE, Office of Basic Energy Sciences. D.S.A. was partially funded by a Marie Curie International Incoming Fellowship (FP7-PEOPLE-2009-IIF-254143)
from the European Commission. We also thank CIME (Interdisciplinary Centre for Electron Microscopy) at EPFL for use of the
electron microscope facility, Chris Fuller (USGS) for providing adsorbed U(IV) standards and Takuya Echigo (Virginia Tech) for
providing XRD reference spectra.
APPENDIX A. SUPPLEMENTARY DATA
Supplementary data associated with this article can be
found, in the online version, at doi:10.1016/j.gca.2011.02.
024.
REFERENCES
Aamrani S. E., Giménez J., Rovira M., Seco F., Grivé M., Bruno
J., Duro L. and De Pablo J. (2007) A spectroscopic study of
uranium(VI) interaction with magnetite. Appl. Surf. Sci. 253,
8794–8797.
Abdelouas A., Lu Y., Lutze W. and Nuttall H. E. (1998) Reduction
of U(VI) to U(IV) by indigenous bacteria in contaminated
ground water. J. Contam. Hydrol. 35, 217–233.
Al-Borno A. and Tomson M. B. (1994) The temperature dependence of the solubility product constant of vivianite. Geochim.
Cosmochim. Acta 58, 5373–5378.
Anderson R., Vrionis H., Ortiz-Bernad I., Resch C., Long P.,
Dayvault R., Karp K., Marutzky S., Metzler D., Peacock A.,
White D., Lowe M., and Lovley D. (2003). Stimulating the
in situ activity of geobacter species to remove uranium from the
2525
groundwater of a uranium-contaminated aquifer. Appl. Environ. Microbiol. 5884–5891.
Bargar J. R., Bernier-Latmani R., Giammar D. E. and Tebo B. M.
(2008) Biogenic uraninite nanoparticles and their importance
for uranium remediation. Elements 4, 407–412.
Behrends T. and Van Cappellen P. (2005) Competition between
enzymatic and abiotic reduction of uranium(VI) under iron
reducing conditions. Chem. Geol. 220, 315–327.
Bell P. E., Mills A. L. and Herman J. S. (1987) Biogeochemical
conditions favoring magnetite formation during anaerobic iron
reduction. Appl. Environ. Microbiol. 53, 2610–2616.
Bernier-Latmani R., Veeramani H., Della Vecchia E., Junier P.,
Lezama-Pacheco J. S., Suvorova E. I., Sharp J. O., Wigginton
N. S. and Bargar J. R. (2010) Non-uraninite products of
microbial. Environ. Sci. Technol. 44, 9456–9462.
Boyanov M., Latta D., O’loughlin E., Gorski C., Scherer M., and
Kemner K. (2009) Distinct uranium(IV) products result from
uranyl reduction in different ferrous-ferric oxyhydroxide systems. Geochim. Cosmochim. Acta A151–A151.
Boyanov M., O’Loughlin E., Roden E., Fein J., and Kemner K.
(2007) Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The influence of speciation on uranyl
reduction studied by titration and XAFS. Geochim. Cosmochim.
Acta 1898–1912.
Bruggeman C. and Maes N. (2010) Uptake of uranium(VI) by
pyrite under boom clay conditions: influence of dissolved
organic carbon. Environ. Sci. Technol. 44, 4210–4216.
Buffle J., De Vitre R. R., Perret D. and Leppard G. G. (1989)
Physico-chemical characteristics of a colloidal iron phosphate
species formed at the oxic–anoxic interface of a eutrophic lake.
Geochim. Cosmochim. Acta 53, 399–408.
Chakraborty S., Boivin F., Banerjee D., Scheinost A., Mullet M.,
Jacques Ehrhardt Jea, Brendle J., Vidal L. and Charlet L.
(2010) U(VI) Sorption and reduction by Fe(II) sorbed on
montmorillonite. Environ. Sci. Technol. 44, 3779–3785.
Charlet L., Liger E. and Gerasimo P. (1998) Decontamination of
TCE- and U-rich waters by granular iron: role of sorbed Fe(II).
J. Environ. Eng. 124, 25–30.
Chaudhuri S. K., Lack J. G. and Coates J. D. (2001) Biogenic
magnetite formation through anaerobic biooxidation of Fe(II).
Appl. Environ. Microbiol. 67, 2844–2848.
Conradson S. D., Manara D., Wastin F., Clark D. L., Lander G.
H., Morales L. A., Rebizant J. and Rondinella V. V. (2004)
Local structure and charge distribution in the UO2–U4O9
system. Inorg. Chem. 43, 6922–6935.
Cornell R., Giovanoli R., and Schneider W. (1989) Review of the
hydrolysis of iron(III) and the crystallization of amorphous
iron(III) hydroxide hydrate. J. Chem. Technol. Biotechnol. 115–
134.
Corr S. A., Gun’ko Y. K., Douvalis A. P., Venkatesan M. and
Gunning R. D. (2004) Magnetite nanocrystals from a single
source metallorganic precursor: metallorganic chemistry vs
biogeneric bacteria. J. Mater. Chem. 14, 944–946.
Daou T. J., Begin-Colin S., Grenèche J. M., Thomas F., Derory A.,
Bernhardt P., Legaré P. and Pourroy G. (2007) Phosphate
adsorption properties of magnetite-based nanoparticles. Chem.
Mater. 19, 4494–4505.
Dodge C. J., Francis A. J., Gillow J. B., Halada G. P., Eng C. and
Clayton C. R. (2002) Association of uranium with iron oxides
typically formed on corroding steel surfaces. Environ. Sci.
Technol. 36, 3504–3511.
Dong H., Fredrickson J. K., Kennedy D. W., Zachara J. M.,
Kukkadapu R. K. and Onstott T. C. (2000) Mineral transformations associated with the microbial reduction of magnetite.
Chem. Geol. 169, 299–318.
Available online at www.sciencedirect.com
Geochimica et Cosmochimica Acta 75 (2011) 7277–7290
www.elsevier.com/locate/gca
Heterogeneous reduction of U6+ by structural Fe2+
from theory and experiment
F.N. Skomurski, E.S. Ilton, M.H. Engelhard, B.W. Arey, K.M. Rosso ⇑
Pacific Northwest National Laboratory, Richland, WA 99352, United States
Received 14 March 2011; accepted in revised form 2 August 2011; available online 9 August 2011
Abstract
Computational and experimental studies were performed to explore heterogeneous reduction of U6+ by structural Fe2+ at
magnetite (Fe3O4) surfaces. Molecular Fe–Fe–U models representing a uranyl species adsorbed in a biatomic bidentate
fashion to an iron surface group were constructed. Various possible charge distributions in this model surface complex were
evaluated in terms of their relative stabilities and electron exchange rates using ab initio molecular orbital methods. Freshlycleaved, single crystals of magnetite with different initial Fe2+/Fe3+ ratios were exposed to uranyl-nitrate solution (pH 4) for
90 h. X-ray photoelectron spectroscopy and electron microscopy indicated the presence of a mixed U6+/U5+ precipitate
heterogeneously nucleated and grown on stoichiometric magnetite surfaces, but only the presence of sorbed U6+ and no precipitate on sub-stoichiometric magnetite surfaces. Calculated electron transfer rates indicate that sequential multi-electron
uranium reduction is not kinetically limited by conductive electron resupply to the adsorption site. Both theory and experiment point to structural Fe2+ density, taken as a measure of thermodynamic reducing potential, and sterically accessible
uranium coordination environments as key controls on uranium reduction extent and rate. Uranium incorporation in solid
phases where its coordination is constrained to the uranate type should widen the stability field of U5+ relative to U6+. If
uranium cannot acquire 8-fold coordination then reduction may proceed to U5+ but not necessarily U4+.
Ó 2011 Elsevier Ltd. All rights reserved.
1. INTRODUCTION
Interest in the role that iron oxides play in limiting
contaminant uranium mobility stems from their ubiquity
in natural and man-made environments (e.g., corrosion
products of steel), their strong sorptive properties, and
the reductive potential of Fe2+-bearing oxides. Contaminant uranium is important because of its long half-life
(4.0 Ga), radioactivity (alpha-radiation emitter), and potential high mobility under oxidizing conditions (Bruno and
Ewing, 2006). Oxidized uranium exists as the uranyl cation
ðUO2þ
2 Þ which forms strong complexes with carbonate and
other ligands in aqueous solution that contribute to its high
⇑ Corresponding author. Address: Pacific Northwest National
Laboratory, P.O. Box 999, MSIN K8-96, Richland, WA 99352,
United States. Tel.: +1 509 371 6357; fax: +1 509 371 6354.
E-mail address: [email protected] (K.M. Rosso).
0016-7037/$ - see front matter Ó 2011 Elsevier Ltd. All rights reserved.
doi:10.1016/j.gca.2011.08.006
mobility (Bargar et al., 1999). Reduction to U4+, on the
other hand, yields sparingly soluble U4+ compounds (e.g.,
UO2, uraninite), thus limiting transport (Shoesmith,
2000). It is therefore of interest to know the rate and extent
of U6+ reduction by common reductants in the environment, such as Fe2+ in its various chemical and mineralogic
forms.
Due to the well known fact that homogeneous reduction
of U6+ by aqueous Fe2+ is kinetically hindered (e.g., Liger
et al., 1999), many studies have investigated relatively rapid, coupled sorption–reduction of uranyl on various
Fe2+-containing oxides and silicates, including magnetite
(El Aamrani et al., 1999, 2007; Missana et al., 2003a,b;
Scott et al., 2005a; Ilton et al., 2010), steel corrosion products (Moyes et al., 2000; Dodge et al., 2002; Eng et al.,
2003; O’Loughlin et al., 2003; Scott et al., 2005b; Rovira
et al., 2007; Duro et al., 2008; Ferriss et al., 2009), siderite
(Ithurbide et al., 2009, 2010), and micas (Ilton et al., 2004,
2005). Further, considerable work has focused on reduction
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F.N. Skomurski et al. / Geochimica et Cosmochimica Acta 75 (2011) 7277–7290
Fig. 4. Back-scattered SEM images of the large magnetite single crystal following exposure to U-bearing solution show sub-micron-sized
crystallites covering the surface at 20,000 (a) and 100,000 (b) magnification. Images of the small magnetite crystal following exposure to Ubearing solution show the absence of a detectable surface precipitate at similar 25,000 (c) and 100,000 (d) magnifications.
at the surface of the sub-stoichiometric magnetite crystal
(Fig. 4c and d) which, when combined with the XPS results,
suggests that U6+ sorbed to this surface in a finer-scale, possibly molecular adsorbate form that did not undergo reduction, ostensibly because of the relative depletion of Fe2+ at
the initial surface.
3.3. Mechanistic implications from combined modeling and
experiment
The modeling highlights the importance of both
uranium coordination and Fe2+ density for reduction of
adsorbed U6+. The Fe2+ density can be considered a reasonable proxy quantity for the thermodynamic reducing
potential at the surface, with higher Fe2+ density equating
to a more reducing surface. For the ideal Fe2+–Fe2+ case,
reduction from U6+ to U5+ is predicted to be facile, with
no change in coordination required, just a lengthening of
the axial oxygen bonds. In contrast, for the depleted
Fe2+–Fe3+ case, reduction was inhibited ( 2 kT) unless
uranium began to acquire uranate-like coordination,
whereupon U5+ was strongly stabilized relative to U6+.
Interestingly, stabilization of U4+ relative to U6+ required
both a comprehensive change to 8-fold U coordination
and a locally excess Fe2+ electron supply, such as might
be supplied from the magnetite bulk. This prediction is in
accord with that fact that U4+ is most often found in 8-fold
coordination with nearly equidistant U–O bonds either in
the solid state, such as in UO2 (Wyckoff, 1963) and coffinite
(USiO4; Fuchs and Gebert, 1958), or in sorbed molecular
form (Wu et al., 2007; Kelly et al., 2008; Bernier-Latmani
et al., 2010; Fletcher et al., 2010), although the latter is often associated with bacterially-mediated reduction of U6+.
There are rare exceptions such as octahedrally coordinated
U4+ in ianthinite (Burns et al., 1997), but its rarity correlates with its lower relative stability compared to 8-fold
coordination. In contrast, U6+ generally maintains the
two short U–O axial bonds and coordination with an additional 4, 5, or 6 equatorial ligands (Burns, 1999).
The apparent role of Fe2+ density in our model relates to
a study by Boyanov et al. (2007), where uranium and Fe2+
were co-adsorbed to non-conducting carboxyl-functionalized latex spheres in the presence of relative excess Fe2+
but fixed U adsorption density. No documented uranium
reduction occurred until the pH was raised to 8.4. At pH
8.4, adsorbed U6+ was reduced to U4+ although not necessarily UO2. EXAFS indicated that U6+ reduction was coincident with polymerization of sorbed Fe2+ (i.e., formation
of dimers and higher order edge-sharing surface structures),
and presumably closer approach of Fe2+ and U6+ with
increasing Fe2+ sorption density. In particular, our model
suggests that Fe dimers might not be sufficient to yield
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F.N. Skomurski et al. / Geochimica et Cosmochimica Acta 75 (2011) 7277–7290
Waste Management Graduate Fellowship Program. Helpful comments of two anonymous reviewers and Associate Editor Mike
Machesky are gratefully acknowledged.
REFERENCES
Allison J. D., Brown D. S. and Novo-Gradac K. J. (1991)
MINTEQA2/PRODEFA2, A geochemical assessment model for
environmental systems: Version 3.0. Environmental Protection
Agency, Athens, GA.
Bargar J. R., Reitmeyer R. and Davis J. A. (1999) Spectroscopic
confirmation of uranium(VI)-carbonato adsorption complexes
on hematite. Environ. Sci. Technol. 33(14), 2481–2484.
Belai N., Frisch M., Ilton E. S., Ravel B. and Cahill C. L. (2008)
Pentavalent uranium oxide via reduction of [UO2]2+ under
hydrothermal reaction conditions. Inorg. Chem. 47, 10135–
10140.
Bernier-Latmani R., Veeramani H., Dalla Vecchia E., Junier P.,
Lezama-Pacheco J., Suvorova Buffat E., Sharp J. O., Wigginton
N. S. and Bargar J. R. (2010) Non-uraninite products of
microbial U(VI) reduction. Environ. Sci. Technol. 44(24), 9456–
9462.
Boland D. D., Collins R. N., Payne T. E. and Waite T. D. (2011)
Effect of amorphous Fe(III) oxide transformation on the Fe(II)mediated reduction of U(VI). Environ. Sci. Technol. 45, 1327–
1333.
Boyanov M. I., O’Loughlin E. J., Roden E. E., Fein J. B. and
Kemner K. M. (2007) Adsorption of Fe(II) and U(VI) to
carboxyl-functionalized microspheres: The influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim. Cosmochim. Acta 71, 1898–1912.
Bruno J. and Ewing R. C. (2006) Spent nuclear fuel. Elements 2(6),
343–349.
Burns P. C., Finch R. J., Hawthorne F. C., Miller M. L. and Ewing
R. C. (1997) The crystal structure of ianthinite,
[U24+(UO2)4O6(OH)4(H2O)4](H2O)5: A possible phase for
Pu4+ incorporation during the oxidation of spent nuclear fuel.
J. Nucl. Mater. 249(2–3), 199–206.
Burns P. C. and Finch R. J. (1999) Wyartite: crystallographic
evidence for the first pentavalent-uranium mineral. Am. Mineral. 84(9), 1456–1460.
Burns P. C. (1999) The crystal chemistry of uranium. In Uranium:
Mineralogy, Geochemistry, and the Environment. Reviews in
Mineralogy and Geochemistry (eds. R. Finch and P. C. Burns).
Mineralogical Society of America and the Geochemical Society,
Chantilly, VA, 38, pp. 23–86.
Bylaska E. J., de Jong W. A., Kowalski K., Straatsma T. P., Valiev
M., Wang D., Apra E., Windus T. L., Hirata S., Hackler M. T.,
Zhao Y., Fan P.-D., Harrison R. J., Dupuis M., Smith D. M.
A., Nieplocha J., Tipparaju V., Krishnan M., Auer A. A.,
Nooijen M., Brown E., Cisneros G., Fann G. I., Fruchtl H.,
Garza J., Hirao K., Kendall R., Nichols J. A., Tsemekhman K.,
Wolinski K., Anchell J., Bernholdt D., Borowski P., Clark T.,
Clerc D., Dachsel H., Deegan M., Dyall K., Elwood D.,
Glendening E., Gutowski M., Hess A., Jaffe J., Johnson B., Ju
J., Kobayashi R., Kutteh R., Lin Z., Littlefield R., Long X.,
Meng B., Nakajima T., Niu S., Pollack L., Rosing M.,
Sandrone G., Stave M., Taylor H., Thomas G., van Lenthe
J., Wong A. and Zhang Z. (2006) NWChem, A Computational
Chemistry Package for Parallel Computers, Version 5.0. Pacific
Northwest National Laboratory, Richland, Washington 993520999, USA.
Catalano J. G., Trainor T. P., Eng P. J., Waychunas G. A. and
Brown, Jr., G. E. (2005) CTR diffraction and grazing-incidence
EXAFS study of U(VI) adsorption onto a-Al2O3 and a-Fe2O3
(1 1 0 2) surfaces. Geochim. Cosmochim. Acta 69(14), 3555–
3572.
Chakraborty S., Favre F., Banerjee D., Scheinost A. C., Mullet M.,
Ehrhardt J.-J., Brendle J., Vidal L. and Charlet L. (2010) U(VI)
sorption and reduction by Fe(II) sorbed on montmorillonite.
Environ. Sci. Technol. 44, 3779–3785.
Chamberlain S. C., Robinson G. W., Lupulescu M., Morgan T. C.,
Johnson J. T. and deLorraine W. B. (2008) Cubic and
tetrahexahedral magnetite. Rocks Mineral. 83, 224–239.
Charlet L., Silvester E. and Liger E. (1998) N-compound reduction
and actinide immobilization in surficial fluids by Fe(II): the
surface „FeIIIOFeIIOH species, as major reductant. Chem.
Geol. 151, 85–93.
Dodge C. J., Francis A. J., Gillow J. B., Halada G. P., Eng C. and
Clayton C. R. (2002) Association of uranium with iron oxides
typically formed on corroding steel surfaces. Environ. Sci.
Technol. 36, 3504–3511.
Duff M. C., Coughlin J. U. and Hunter D. B. (2002) Uranium coprecipitation with iron oxide minerals. Geochim. Cosmochim.
Acta 66(20), 3533–3547.
Duro L., El Aamrani S., Rovira M., de Pablo J. and Bruno J.
(2008) Study of the interaction between U(VI) and the anoxic
corrosion products of carbon steel. Appl. Geochem. 23, 1094–
1100.
El Aamrani F., Casas I., de Pablo J., Duro L., Grivé M. and Bruno
J. (1999) Experimental and modeling study of the interaction
between uranium(IV) and magnetite. Technical Report TR-9921, SKB, 7–30.
El Aamrani S., Giménez J., Rovira M., Seco F., Grivé M., Bruno
J., Duro L. and de Pablo J. (2007) A spectroscopic study of
uranium(VI) interaction with magnetite. Appl. Surf. Sci. 253,
8794–8797.
Eng C. W., Halada G. P., Francis A. J., Dodge C. J. and Gillow J.
B. (2003) Uranium association with corroding carbon steel
surfaces. Surf. Interface Anal. 35, 525–535.
Ferriss E. D. A., Helean K. B., Bryan C. R., Brady P. V. and Ewing
R. C. (2009) UO2 corrosion in an iron waste package. J. Nucl.
Mater. 384, 130–139.
Fleet M. E. (1981) The structure of magnetite. Acta Cryst. B37,
917–920.
Fletcher K. E., Boyanov M. I., Thomas S. H., Wu Q., Kemner K.
M. and Löffler F. E. (2010) U(VI) reduction to mononuclear
U(IV) by Desulfitobacterium species. Environ. Sci. Technol. 44,
4705–4709.
Fuchs L. H. and Gebert E. (1958) X-ray studies of synthetic
coffinite, thorite and uranothorites. Am. Mineral. 43, 243–
248.
Hay P. J. (1977) Gaussian basis sets for molecular calculations. The
representation of 3d orbitals in transition-metal atoms. J.
Chem. Phys. 66(10), 4377–4384.
Ikeda-Ohno A., Hennig C., Tsushima S., Scheinost A. C.,
Bernhard G. and Yaita T. (2009) Speciation and structural
study of U(IV) and-(VI) in perchloric and nitric acid solutions.
Inorg. Chem. 48, 7201–7210.
Ilton E. S., Haiduc A., Moses C. O., Heald S. M., Elbert D. C. and
Veblen D. R. (2004) Heterogeneous reduction of uranyl by
micas: Crystal chemical and solution controls. Geochim. Cosmochim. Acta 68(11), 2417–2435.
Ilton E. S., Haiduc A., Cahill C. L. and Felmy A. R. (2005) Mica
surfaces stabilize pentavalent uranium. Inorg. Chem. 44, 2986–
2988.
Ilton E. S., Boily J.-F. and Bagus P. S. (2007) Beam induced
reduction of U(VI) during X-ray photoelectron spectroscopy:
The utility of the U4f satellite structure for identifying uranium
oxidation states in mixed valence uranium oxides. Surf. Sci.
601(4), 908–916.
Environ. Sci. Technol. 2011, 45, 951–957
Competitive Reduction of
Pertechnetate (99TcO4-) by
Dissimilatory Metal Reducing
Bacteria and Biogenic Fe(II)
ANDREW E. PLYMALE,†
J A M E S K . F R E D R I C K S O N , * ,†
JOHN M. ZACHARA,†
ALICE C. DOHNALKOVA,†
STEVE M. HEALD,‡ DEAN A. MOORE,†
DAVID W. KENNEDY,†
MATTHEW J. MARSHALL,†
CHONGMIN WANG,† CHARLES T. RESCH,†
AND PONNUSAMY NACHIMUTHU†
Pacific Northwest National Laboratory, P.O. Box 999,
Richland, Washington 99352, United States, and Argonne
National Laboratory, Argonne, Illinois 60439, United States
Received August 12, 2010. Revised manuscript received
November 16, 2010. Accepted December 3, 2010.
The fate of pertechnetate (99Tc(VII)O4-) during bioreduction
was investigated in the presence of 2-line ferrihydrite (Fh) and
various dissimilatory metal reducing bacteria (DMRB)
(Geobacter, Anaeromyxobacter, Shewanella) in comparison
with TcO4- bioreduction in the absence of Fh. In the presence
of Fh, Tc was present primarily as a fine-grained Tc(IV)/Fe
precipitate that was distinct from the Tc(IV)O2 · nH2O solids
produced by direct biological Tc(VII) reduction. Aqueous Tc
concentrations (<0.2 µm) in the bioreduced Fh suspensions (1.7
to 3.2 × 10-9 mol L-1) were over 1 order of magnitude lower
than when TcO4- was biologically reduced in the absence of Fh
(4.0 × 10-8 to 1.0 × 10-7 mol L-1). EXAFS analyses of the
bioreduced Fh-Tc products were consistent with variable chain
length Tc-O octahedra bonded to Fe-O octahedra associated
with the surface of the residual or secondary Fe(III) oxide.
In contrast, biogenic TcO2 · nH2O had significantly more Tc-Tc
second neighbors and a distinct long-range order consistent
with small particle polymers of TcO2. In Fe-rich subsurface
sediments, the reduction of Tc(VII) by Fe(II) may predominate
over direct microbial pathways, potentially leading to lower
concentrations of aqueous 99Tc(IV).
Introduction
Technetium-99 (99Tc) is a long-lived (t1/2 ) 2.13 × 105 y)
fission product of nuclear production and nuclear fuel
reprocessing that is an environmental contaminant (1 and
references therein). Environmental contamination by 99Tc is
of particular concern at the U.S. Department of Energy’s
Hanford Site, where subsurface Tc exists as the pertechnetate
oxyanion, Tc(VII)O4-, which is weakly adsorbed by mineral
phases and consequently mobile in vadose-zone water and
groundwater (2 and references therein]). However, under
anoxic conditions, Tc(VII)O4- can be reduced to Tc(IV),
* Corresponding author phone (509)371-6943; fax (509) 371-6946;
e-mail: [email protected]; mailing address: MS J4-16.
†
Pacific Northwest National Laboratory.
‡
Argonne National Laboratory.
10.1021/es1027647
 2011 American Chemical Society
Published on Web 01/06/2011
forming sparingly soluble Tc(IV)O2 · nH2O at circumneutral
pH and in the absence of strong complexing ligands (3-5).
Microbial processes can contribute to Tc(VII) reduction
directly, by enzymatic reduction (6-12) or through redoxactive organic molecules, such as quinones (8, 13). The
terminal reductases for direct enzymatic Tc(VII) reduction
by dissimilatory metal reducing bacteria (DMRB) and sulfate
reducing bacteria (SRB) include periplasmic hydrogenases
(10, 14), outer membrane multiheme c-type cytochromes
(10), or, conceivably, periplasmic c-type cytochromes.
Iron oxides and Fe-bearing clay minerals are widespread
in the terrestrial subsurface, and ferrous iron (Fe(II)) can be
a strong reductant of Tc(VII) when in the sorbed or mineral
structural state (15-20). Although reduction of Tc(VII) by
aqueous Fe(II) (i.e., homogeneous reduction) is kinetically
slow and pH dependent (15, 19), Tc reduction in a system
without an initial solid phase can be accelerated by Fe(II)
sorption to the insoluble Fe/Tc(IV) redox product resulting
from homogeneous reduction of Tc(VII) by Fe(II) (19).
However, Fe(II) reactivity toward Tc(VII) depends on the
chemical environment and distribution of Fe(II), as some
forms of Fe(II) appear to be less reactive than others
(1, 15, 17, 21). Biogenic Fe(II) is similarly reactive toward
Tc(VII), whether the Fe(II) is associated with magnetite
(8, 9, 12), mineral surface complexes (1, 21), or fine-grained
phyllosilicates (20, 22). In natural sediments that contain
reactive Fe(III), direct and indirect bioreduction pathways
of Tc(VII) by dissimilatory metal reducing bacteria (DMRB)
may compete, with the predominant pathway being controlled by the fastest reaction rate.
The intent of the current investigation was to assess the
nature and distribution of the end products resulting from
concurrent bioreduction of 2-line ferrihydrite (Fh) and
Tc(VII)O4- by DMRB, as a means to determine the relative
contributions of direct and indirect pathways under conditions where both are potentially operational. Final aqueous
Tc solution concentrations were carefully measured and
related to Fe mineralogy, and the nature of Fe-Tc solids was
determined by electron microscopy and X-ray absorption
spectroscopy.
Experimental Section
Washed late-log-phase cultures (1 × 108 cells mL-1) of four
DMRB (Shewanella oneidensis MR-1, S. putrefaciens CN-32,
Anaeromyxobacter dehalogenans 2CP-C, and Geobacter sulfurreducens PCA) were incubated with H2 (80 mL added to
80 mL of N2 headspace, at overpressure, to give ∼4 mmol L-1
H2 in solution), 30 mmol L-1 2-line ferrihydrite, and 0.3 mmol
L-1 ammonium pertechnetate, (NH499TcO4-). To prevent the
complexation and solubilization of Tc(IV) (11), PIPES (30
mmol L-1, pH 7), which has been shown in our laboratory
to be noncomplexing toward Tc(IV) (9, 10), was used.
The DMRB-Fh-Tc samples were incubated under anoxic
conditions (50:50 N2:H2 headspace) for 4 days at 50 rpm in
the dark. Additionally, a series of biogenic Tc(IV) solids were
prepared from incubations of S. oneidensis, A. dehalogenans,
and G. sulfurreducens with Tc(VII)O4- and H2 for X-ray
absorption spectroscopy (XAS). The G. sulfurreducens treatment without Fh was also examined by transmission electron
microscopy (TEM). Thin sections prepared from subsamples
of the various DMRB-Fh-Tc suspensions were analyzed by
TEM and X-ray energy dispersive spectroscopy (EDX) to
examine the nature and distribution of Tc in relation to Fe
solids and bacterial cells.
Soluble 99Tc was measured by filtering (0.2 µm) and
assaying the filtrates by liquid scintillation counting (detecVOL. 45, NO. 3, 2011 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
951
FIGURE 1. Electron micrographs of thin sections from Tc-Fh suspensions incubated with H2 and S. oneidensis MR-1 (A, B) and A.
dehalogenans 2CP-C (C, D), uninoculated Fh incubated under identical conditions (E), and G. sulfurreducens incubated with TcO4and H2 in the absence of Fh (F).
X-ray Absorption Spectroscopy. Iron XANES of the
various Tc-Fh-DMRB suspensions revealed variable amounts
of Fe(II) (Figure S12), consistent with measured 0.5 N HClextractable Fe(II) concentrations (Table 1). The Fe-edge
positions in the suspensions with S. oneidensis and A.
dehalogenans were similar to those for Fh, Fh + Tc(IV), and
hematite (Fe2O3) standards. The edge positions for the S.
putrefaciens and G. sulfurreducens suspensions were downshifted more closely to the position of the magnetite standard,
again consistent with the presence of magnetite, in addition
to goethite, in these suspensions (Table 1, Figure S4).
The valence of Tc in all bioreduced Fh suspensions,
regardless of the organism, was confirmed to be Tc(IV) by
XANES (Figure S10). Tc X-ray absorption fine structure
(EXAFS) (χ (k) data) (Figure S13) and radial transforms for
the bioreduced Fh suspensions were similar to a Tc(IV) + Fh
standard (made by mixing Fh with Tc(IV) in 2 N HCl and
adjusting to pH 7) (19), and were distinct from a
Tc(IV)O2 · nH2O standard generated by dithionite reduction
(Figure 2). These spectra, in turn, were almost identical to
those resulting from Tc(VII) reaction with (i) sorbed Fe(II)
on goethite and hematite (17), (ii) fine-grained biomagnetite,
and (iii) aqueous Fe(II) (19). The EXAFS spectra for these
different Tc(IV)-Fe(III) oxide associations are indistinguishable from one another, and they all can be closely described
with a model where variable chain-length Tc-O octahedra
(n ) 1-3) are bonded in an edge-sharing fashion to Fe-O
octahedra associated with the Fe oxide (17, 19). In this regard,
the Tc(IV) associations with Fh, goethite, and magnetite are
indistinguishable from one another. The peak at ∼2 Å is
diagnostic of these particular molecular associations. Given
the high levels of Tc associated with the poorly crystalline
nanoparticle clusters (Figure 1, Figures S5-S7), we assume
that the bulk EXAFS signature is representative of this specific
redox product.
The EXAFS spectra for biogenic TcO2 · nH2O produced by
three different organisms (Figure 3) were quite similar to
each other. They exhibited spectral features comparable to
the Tc(IV) standard with a distinctive long-range order. The
biogenic phases, however, had fewer Tc-Tc second neighbors
at 2.2 Å, consistent with their nanometer size (by analogy to
biogenic UO2 26, 27).
VOL. 45, NO. 3, 2011 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
953
Acknowledgments
We thank Oleg Geydebrekht for assistance with culturing A.
dehalogenans; Eric Roden for advice on culturing G. sulfurreducens; Yuanxian Xia, Nancy Hess, and Ken Krupka for
helpful discussions of Tc chemistry; Yuanxian Xia for
preparing XAS standards; Tetyana Peretyazhko and Carolyn
Pearce for discussing our results and reviewing the manuscript; and Gailann Thomas-Black and Sonia Enloe for
assistance with manuscript preparation. This research was
supported by the Subsurface Biogeochemical Research
Program (SBR), Office of Biological and Environmental
Research (OBER), U.S. Department of Energy (DOE), and is
a contribution of the PNNL Scientific Focus Area. Transmission electron microscopy and micro-XRD measurements
were performed in the William R. Wiley Environmental
Molecular Sciences Laboratory, a national scientific user
facility sponsored by OBER and located at Pacific Northwest
National Laboratory (PNNL). PNNL is operated for the DOE
by Battelle. Use of the Advanced Photon Source for XANES
and EXAFS measurements was supported by the DOE’s Office
of Science under contract DE-AC02-06CH11357.
Supporting Information Available
Additional data as referenced in this article. This information
is available free of charge via the Internet at http://
pubs.acs.org.
Literature Cited
(1) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Kukkadapu,
R. K.; McKinley, J. P.; Heald, S. M.; Liu, C.; Plymale, A. E.
Reduction of TcO4- by sediment-associated biogenic Fe(II).
Geochim. Cosmochim. Acta 2004, 68 (15), 3171–3187.
(2) Zachara, J. M.; Serne, J.; Freshley, M.; Mann, F.; Anderson, F.;
Wood, M.; Jones, T.; Myers, D. Geochemical processes controlling migration of tank wastes in Hanford’s vadose zone. Vadose
Zone J. 2007, 6 (4), 985–1003.
(3) Hess, N. J.; Xia, Y. X.; Rai, D.; Conradson, S. D. Thermodynamic
model for the solubility of TcO2 · xH2O(am) in the aqueous Tc(IV)Na+-Cl--H+-OH--H2O system. J. Solution Chem. 2004, 33 (2),
199–226.
(4) Meyer, R. E.; Arnold, W. D.; Case, F. I.; Okelley, G. D. Solubilities
of Tc(IV) oxides. Radiochim. Acta 1991, 55 (1), 11–18.
(5) Rard, J. A.; Rand, M. H.; Anderegg, G.; Wanner, H. Chemical
Thermodynamics of Technetium; Elsevier: Amsterdam, 1999.
(6) Lloyd, J. R.; Macaskie, L. E. A novel phosphorimager-based
technique for monitoring the microbial reduction of technetium.
Appl. Environ. Microbiol. 1996, 62 (2), 578–582.
(7) Lloyd, J. R.; Ridley, J.; Khizniak, T.; Lyalikova, N. N.; Macaskie,
L. E. Reduction of technetium by Desulfovibrio desulfuricans:
Biocatalyst characterization and use in a flowthrough bioreactor.
Appl. Environ. Microbiol. 1999, 65 (6), 2691–2696.
(8) Lloyd, J. R.; Sole, V. A.; Van Praagh, C. V. G.; Lovley, D. R. Direct
and Fe(II)-mediated reduction of technetium by Fe(III)- reducing bacteria. Appl. Environ. Microbiol. 2000, 66 (9), 3743–3749.
(9) Marshall, M. J.; Dohnalkova, A.; Kennedy, D. W.; Plymale, A. E.;
Thomas, S. H.; Loffler, F. E.; Sanford, R. A.; Zachara, J. M.;
Fredrickson, J.; Beliaev, A. S. Electron donor-dependent radionuclide reduction and nanoparticle formation by Anaeromyxobacter dehalogenans strain 2CP-C. Environ. Microbiol.
2009, 11 (2), 534–543.
(10) Marshall, M. J.; Plymale, A. E.; Kennedy, D. W.; Wang, Z.; Reed,
S. B.; Dohnalkova, A. C.; Simonson, C. J.; Liu, C.; Saffarini, D. A.;
Romine, M. F.; Zachara, J. M.; Beliaev, A. S.; Fredrickson, J. K.
Hydrogenase- and outer membrane c-type cytochromefacilitated reduction of technetium(VII) by Shewanella oneidensis MR-1. Environ. Microbiol. 2008, 10 (1), 125–136.
(11) Wildung, R. E.; Gorby, Y. A.; Krupka, K. M.; Hess, N. J.; Li, S. W.;
Plymale, A. E.; McKinley, J. P.; Fredrickson, J. K. Effect of electron
donor and solution chemistry on products of dissimilatory
reduction of technetium by Shewanella putrefaciens. Appl.
Environ. Microbiol. 2000, 66 (6), 2451–2460.
(12) Lloyd, J. R.; Chesnes, J.; Glasauer, S.; Bunker, D. J.; Livens, F. R.;
Lovley, D. R. Reduction of actinides and fission products by
Fe(III)-reducing bacteria. Geomicrobiol. J. 2002, 19 (1), 103–
120.
(13) Fredrickson, J. K.; Kostandarithes, H. M.; Li, S. W.; Plymale,
A. E.; Daly, M. J. Reduction of Fe(III), Cr(V), U(VI), and Tc(VII)
956
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 45, NO. 3, 2011
(14)
(15)
(16)
(17)
(18)
(19)
(20)
(21)
(22)
(23)
(24)
(25)
(26)
(27)
(28)
(29)
(30)
(31)
(32)
by Deinococcus radiodurans R1. Appl. Environ. Microbiol. 2000,
66, 2006–2011.
De Luca, G.; De Philip, P.; Dermoun, Z.; Rousset, M.; Vermeglio,
A. Reduction of technetium(VII) by Desulfovibrio fructosovorans
is mediated by the nickel-iron hydrogenase. Appl. Environ.
Microbiol. 2001, 67 (10), 4583–4587.
Cui, D.; Eriksen, T. E. Reduction of pertechnetate by ferrous
iron in solution: Influence of sorbed and precipitated Fe(II).
Environ. Sci. Technol. 1996, 30 (7), 2259–2262.
Cui, D.; Eriksen, T. E. Reduction of pertechnetate in solution
by heterogeneous electron transfer from Fe(II)-containing
geological material. Environ. Sci. Technol. 1996, 30 (7), 2263–
2269.
Peretyazhko, T.; Zachara, J. M.; Heald, S. M.; Jeon, B. H.;
Kukkadapu, R. K.; Liu, C.; Moore, D.; Resch, C. T. Heterogeneous
reduction of Tc(VII) by Fe(II) at the solid-water interface.
Geochim. Cosmochim. Acta 2008, 72 (6), 1521–1539.
Peretyazhko, T.; Zachara, J. M.; Heald, S. M.; Kukkadapu, R. K.;
Liu, C.; Plymale, A. E.; Resch, C. T. Reduction of Tc(VII) by Fe(II)
Sorbed on Al (hydr)oxides. Environ. Sci. Technol. 2008, 42 (15),
5499–5506.
Zachara, J. M.; Heald, S. M.; Jeon, B. H.; Kukkadapu, R. K.; Liu,
C. X.; McKinley, J. P.; Dohnalkova, A. C.; Moore, D. A. Reduction
of pertechnetate [Tc(VII)] by aqueous Fe(II) and the nature of
solid phase redox products. Geochim. Cosmochim. Acta 2007,
71 (9), 2137–2157.
Jaisi, D. P.; Dong, H. L.; Plymale, A. E.; Fredrickson, J.; Zachara,
J. M.; Heald, S. C.; Liu, C. Reduction and long-term immobilization of technetium by Fe(II) associated with clay mineral
nontronite. Chem. Geol. 2009, 264, 127–138.
Fredrickson, J.; Zachara, J. M.; Plymale, A. E.; Heald, S. C.;
McKinley, J. P.; Kennedy, D. W.; Liu, C.; Nachimuthu, P. Oxidative
dissolution potential of biogenic and abiogenic TcO2 in subsurface sediments. Geochim. Cosmochim. Acta 2009, 73 (8),
2299–2313.
Wildung, R. E.; Li, S. W.; Murray, C. J.; Krupka, K. M.; Xie, Y.;
Hess, N. J.; Roden, E. E. Technetium reduction in sediments of
a shallow aquifer exhibiting dissimilatory iron reduction
potential. FEMS Microbiol. Ecol. 2004, 49 (1), 151–162.
Currie, L. A. Limits for qualitative detection and quantitative
determination - application to radiochemistry. Anal. Chem.
1968, 40 (3), 586–593.
Hansel, C. M.; Benner, S. G.; Fendorf, S. Competing Fe(II)induced mineralization pathways of ferrihydrite. Environ. Sci.
Technol. 2005, 39 (18), 7147–7153.
Borch, T.; Masue, Y.; Kukkadapu, R. K.; Fendorf, S. Phosphate
imposed limitations on biological reduction and alteration of
ferrihydrite. Environ. Sci. Technol. 2007, 41 (1), 166–172.
Boyanov, M. I.; O’Loughlin, E, J.; Roden, E. E.; Fein, J. B.; Kemner,
K. Adsorption of Fe(II) and U(VI) to carboxyl-functionalized
microspheres: The influence of speciation on uranyl reduction
studied by titration and XAFS. Geochim. Cosmochim. Acta 2007,
71, 1898–1912.
O’Loughlin, E, J.; Kelly, S. D.; Cook, R. E.; Csencsits, R.; Kemner,
K. Reduction of uranium(VI) by mixed iron(II)/iron(III) hydroxide (green rust): Formation of UO2 nanoparticles. Environ.
Sci. Technol. 2003, 37, 721–727.
Zachara, J. M.; Kukkadapu, R. K.; Fredrickson, J. K.; Gorby, Y. A.;
Smith, S. C. Biomineralization of poorly crystalline Fe(III) oxides
by dissimilatory metal reducing bacteria (DMRB). Geomicrobiol.
J. 2002, 19 (2), 179–207.
Kukkadapu, R. K.; Zachara, J. M.; Fredrickson, J. K.; Kennedy,
D. W. Biotransformation of two-line silica-ferrihydrite by a
dissimilatory Fe(III)-reducing bacterium: Formation of carbonate green rust in the presence of phosphate. Geochim. Cosmochim. Acta 2004, 68 (13), 2799–2814.
Zachara, J. M.; Fredrickson, J. K.; Li, S. W.; Kennedy, D. W.;
Smith, S. C.; Gassman, P. L. Bacterial reduction of crystalline
Fe3+ oxides in single phase suspensions and subsurface materials. Am. Mineral. 1998, 83, 1426–1443.
Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Dong, H.;
Onstott, T. C.; Hinman, N. W.; Li, S. W. Biogenic iron
mineralization accompanying the dissimilatory reduction of
hydrous ferric oxide by a groundwater bacterium. Geochim.
Cosmochim. Acta 1998, 62 (19/20), 3239–3257.
Coker, V. S.; Bell, A. M. T.; Pearce, C. I.; Pattrick, R. A. D.; van
der Laan, G.; Lloyd, J. R. Time-resolved synchrotron powder
X-ray diffraction study of magnetite formation by the Fe(III)reducing bacterium Geobacter sulfurreducens. Am. Mineral.
2008, 93 (4), 540–547.
Geomicrobiology Journal, 28:497–506, 2011
Copyright © Taylor & Francis Group, LLC
ISSN: 0149-0451 print / 1521-0529 online
DOI: 10.1080/01490451.2010.512033
Uranium Redox Cycling in Sediment and Biomineral Systems
Gareth T. W. Law,1 Andrea Geissler,1 Ian T. Burke,2 Francis R. Livens,3
Jonathan R. Lloyd,1 Joyce M. McBeth,1 and Katherine Morris1
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1
Research Centre for Radwaste and Decommissioning and Williamson Research Centre for Molecular
Environmental Science, School of Earth, Atmospheric and Environmental Sciences, The University
of Manchester, Manchester, United Kingdom
2
Earth System Science Institute, School of Earth and Environment, University of Leeds,
Leeds, United Kingdom
3
Centre for Radiochemistry Research, School of Chemistry, The University of Manchester, Manchester,
United Kingdom
Under anaerobic conditions, uranium solubility is significantly
controlled by the microbially mediated reduction of relatively soluble U(VI) to poorly soluble U(IV). However, the reaction mechanism(s) for bioreduction are complex with prior sorption of U(VI)
to sediments significant in many systems, and both enzymatic
and abiotic U(VI) reduction pathways potentially possible. Here,
we describe results from sediment microcosm and Fe(II)-bearing
biomineral experiments designed to assess the relative importance
of enzymatic vs. abiotic U(VI) reduction mechanisms and the longterm fate of U(IV). In oxic sediments representative of the UK Sellafield reprocessing site, U(VI) was rapidly and significantly sorbed
to surfaces and during microbially-mediated bioreduction, XAS
analysis showed that sorbed U(VI) was reduced to U(IV) commensurate with Fe(III)-reduction. Additional control experiments with
Fe(III)-reducing sediments that were sterilized after bioreduction
and then exposed to U(VI), indicated that U(VI) reduction was
inhibited, implying that enzymatic as opposed to abiotic mechanisms dominated in these systems. Further experiments with model
Fe(II)-bearing biomineral phases (magnetite and vivianite) showed
that significant U(VI) reduction occurred in co-precipitation systems, where U(VI) was spiked into the biomineral precursor phases
prior to inoculation with Geobacter sulfurreducens. In contrast,
when U(VI) was exposed to pre-formed, washed biominerals, XAS
analysis indicated that U(VI) was recalcitrant to reduction. Reoxidation experiments examined the long-term fate of U(IV). In
sediments, air exposure resulted in Fe(II) oxidation and significant
U(IV) oxidative remobilization. By contrast, only partial oxidation
Received 21 April 2010; accepted 27 July 2010.
Current affiliation for A. Geissler: Forschungszentrum DresdenRossendorf, Dresden Germany
Current affiliation for J. M. McBeth: Bigelow Laboratory for Ocean
Sciences, West Boothbay Harbour, Maine, USA
Address correspondence to Katherine Morris, Research Centre for
Radwaste and Decommissioning and Williamson Research Centre for
Molecular Environmental Science, School of Earth, Atmospheric and
Environmental Sciences, The University of Manchester, Manchester,
M13 9PL, United Kingdom. E-mail: [email protected]
of U(IV) and no remobilization to solution occurred with nitrate
mediated bio-oxidation of sediments. Magnetite was resistant to
biooxidation with nitrate. On exposure to air, magnetite changed
from black to brown in colour, yet there was limited mobilization
of uranium to solution and XAS confirmed that U(IV) remained
dominant in the oxidized mineral phase. Overall these results highlight the complexity of uranium biogeochemistry and highlight the
importance of mechanistic insights into these reactions if optimal
management of the global nuclear legacy is to occur.
Keywords
uranium, sediment, bioreduction, biomineral, redox
INTRODUCTION
Uranium-238 is a long-lived (238U = 4.5 × 109 years) alphaemitting-radionuclide that is present as a subsurface contaminant at nuclear legacy sites (Morris et al. 2002; Istok et al.
2004). In oxic environments, U(VI) dominates as the uranyl
cation (UO2+
2 ), which displays a range of environmental behaviors, ranging from being highly soluble in acidic or carbonate
dominated environments (Lovley et al. 1992; Clark et al. 1995)
to being extensively sorbed to geomedia in the absence of complexants (Sylwester et al. 2000; Barnett et al. 2002; Ortiz-Bernad
et al. 2004; Jeon et al. 2005; Dong et al. 2006; Begg et al. 2010).
Under anoxic conditions, highly insoluble U(IV)O2 dominates
speciation (Lovley et al. 1991; Lloyd and Renshaw 2005).
In axenic culture, microcosm, and in situ studies, microbiallymediated reduction has been shown to facilitate formation of
insoluble U(IV) from both soluble and sorbed U(VI) (e.g., Lovley et al. 1991; Fredrickson et al. 2000; Finneran et al. 2002;
Istok et al. 2004; Wilkins et al. 2007; Begg et al. 2010). Here,
U(VI) is reduced to U(IV) commensurate with the development of Fe(III)- and/or sulfate-reducing conditions, with reduction facilitated via enzymatic processes and/or abiotic reaction
with the reduced by-products of microbial metabolism (e.g.,
Fe(II)-biominerals). Indeed, in systems where U(VI) is partially
497
503
URANIUM REDOX CYCLING
TABLE 2
U(VI) in solution and solid-phase uranium LIII edge XANES linear combination fitting results for biomineral reduction and
reoxidation experiments
XANES linear combination modelling
Sample
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Ferric gel
Magnetite co-precipitation
Magnetite co-precipitation air (20 days)
Magnetite sorption (10 days exposure)
Vivianite co-precipitation
Vivianite sorption (10 days exposure)
% U(VI)(aq)
% spectrum 1
% spectrum 2
∼0
∼0
2
∼0
∼0
∼0
52†
10†
5‡
48†
90†
95‡
Linear combination fitting (LCF) errors were estimated to be +/− ∼15%. End-member spectra used in linear combination modelling (denoted
by (-) symbol) were (spectrum 1) U(IV) sorbed to “co-precipitated” magnetite† or vivianite‡, and (spectrum 2) U(VI) sorbed to ferric gel.
(Figure 3, Table 1). To assess the mechanism of bioreduction
(i.e. enzymatic vs. abiotic), a control experiment was conducted.
Here, sterile Fe(III)-reduced sediment was spiked with U(VI)
and left to equilibrate for 10 days (Table 1). The XANES spectra
and linear combination fitting of the sample indicated a predominantly U(VI)-like environment (Figure 3; Table 1), indicating
that U(VI) reduction in this system may be primarily facilitated
via enzymatic processes. Alternatively, changes in the physicochemical conditions of the sediments after autoclaving may have
altered the U(VI) reduction potential. Regardless, these results
are similar to findings for sediments from a range of nuclear
legacy sites and suggest that under certain environmental conditions, U(VI) reduction is dominated by enzymatic pathways
(Liu et al. 2005; Fox et al. 2006; Wilkins et al. 2007; Begg et al.
2010).
Uranium Interaction with Fe(II)-Bearing Biominerals
Whilst enzymatic reduction appears to dominate U(VI) reduction in Sellafield sediments, reduction in other systems has
been attributed to U(VI) reaction with Fe(II)-bearing mineral
phases and/or Fe(II) sorbed to surfaces (i.e., abiotic U(VI) reduction) (Moyes et al. 2000; Fredrickson et al. 2000; Misanna et
al. 2003; Scott et al. 2005; Behrends et al. 2005; Jeon et al. 2005;
O’Loughlin et al. 2003, 2010; Sharp et al. 2008; Ithurbide et al.
2009). This apparent inconsistency may reflect U(VI) specificity
for differing Fe(II) phases, variation in the reactivity of different
Fe(II) minerals, or differences in the reactivity of synthetic vs.
biogenic Fe(II) minerals due to, for example, surface area or
pH/surface speciation effects (Boyanov et al. 2007).
To further assess whether environmentally relevant Fe(II)bearing biominerals can reduce U(VI), uranium interactions
with model biogenic Fe(II)-bearing biominerals (magnetite and
vivianite) were investigated. Two experimental treatments were
undertaken: (i) co-precipitation, where U(VI) was added to magnetite and vivianite precursors prior to inoculation with Geobacter sulfurreducens and biomineral formation; and (ii) sorption,
where U(VI) was spiked into the preformed, washed biomineral
phases. In the co-precipitation treatments, uranium remained
soluble in the Fe(III)-citrate medium used to prepare vivianite,
but was completely sorbed to ferric gel (which was bioreduced
to magnetite) as U(VI) (Figure 4). After magnetite and vivianite
formation, all of the added uranium was sorbed to the mineral
phases, with uranium XANES spectra reflecting a predominantly U(IV)-like environment (Figure 4).
In the sorption treatments, U(VI) removal was also marked,
with ∼90% of the added uranium sorbed to each mineral phase
after 1 h, and ∼100% sorbed after 2 days (Table 2). However,
after 10 days equilibration, the XANES spectra were predominantly U(VI), with linear combination fitting suggesting that
≤10% of the uranium was present as U(IV) in the magnetite
and vivianite samples (Figure 4; Table 2). When considered
alongside the co-precipitation mineral and sediment data (Figures 3 and 4), these results highlight that U(VI) reduction is
dominated by enzymatic pathways in these systems and further imply that biogenic magnetite and vivianite are ineffectual
U(VI) reductants under the conditions of study.
These results are similar to past work (Moyes et al. 2000;
Jeon et al. 2005; Ithurbide et al. 2009; O’Loughlin et al. 2010;
Finneran et al. 2002) but contrast with the observations of several
workers who have reported significant U(VI) reduction on exposure to synthetic and biogenic Fe(II)-bearing mineral phases
(Misanna et al. 2003; Scott et al. 2005; Behrends et al. 2005;
Boyanov et al. 2007; O’Loughlin et al. 2003, 2010; Sharp et
al. 2008). Overall, it is clear that the key factors that control
whether electron transfer to U(VI) can occur in the presence of
Fe(II)-bearing mineral phases are highly specific to the conditions of study, and that under the ambient conditions studied
here in both sediments and model mineral phases, enzymatic
processes appear to enhance the extent of U(VI) reduction.
Uranium Reoxidation Behavior in Sediment Systems
To understand the long-term fate of bioreduced U(IV), we
examined uranium behavior during air and nitrate reoxidation
of Fe(III)-reducing, U(IV) labelled sediments. Air reoxidation
resulted in rapid Fe(II) oxidation and almost complete uranium
remobilization to solution as U(VI) within 24 hours (Table 1).
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URANIUM REDOX CYCLING
Behrends T, Van Cappellen P. 2005. Competition between enzymatic and abiotic reduction of uranium(VI) under iron reducing conditions. Chem Geol
220:315–327.
Boyanov M, O’Loughlin EJ, Roden E, Fein J, Kemner K. 2007. Adsorption
of iron(II) and uranium(VI) to carboxyl-functionalized microspheres: the
influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim Cosmochim Acta 71:1898–1912.
Brewer PG, Spencer DW. 1971. Colourimetric determination of Mn in anoxic
waters. Limnol Oceanogr 16:107–110.
Burns PC, Finch R. 1999. Uranium: mineralogy, geochemistry and the environment. Rev Mineral 38:1–679.
Caccavo JR, Lonergan DJ, Lovley DR, Davis M, Stolz JF, McInerney
MJ. 1994. Geobacter sulfureducens sp. nov., a hydrogen- and acetateoxidizing dissimilatory metal-reducing microorganism. Appl Environ Microbiol 60:3752–3759.
Cheng T, Barnett MO, Roden EE, Zhuang J. 2006. Effects of solid-solution
ratio on uranium(VI) adsorption and its implications. Environ Sci Technol
40:3243–3247.
Clark DL, Hobart DE, Neu MP. 1995. Actinide carbonate complexes and their
importance in actinide environmental chemistry. Chem Rev 95:25–48.
Coppi MV, Leang C, Sandler SJ, Lovley DR. 2001. Development of a
genetic system for Geobacter sulfurreducens. Appl Environ Microbiol
67:3180–3187.
Dong W, Xie G, Miller TR, Franklin MP, Oxenberg TP, Bouwer EJ, Ball WP,
Halden RU. 2006. Sorption and bioreduction of hexavalent uranium at a
military facility by the Chesapeake Bay. Environ Pollut 142:132–142.
Finneran KT, Anderson RT, Nevin KP, Lovley DR. 2002. Potential for bioremediation of uranium-contaminated aquifers with microbial U(VI) reduction.
Soil Sed Contam 11:339–357.
Fox JR, Mortimer RJG, Lear G, Lloyd JR, Beadle I, Morris K. (2006) The
biogeochemical behavior of U(VI) in the simulated near-field of a low-level
radioactive waste repository, Appl Geochem 21:1539–1550.
Fredrickson JK, Zachara JM, Kennedy DW, Duff MC, Gorby YA, Li SW,
Krupka M. 2000. Reduction of U(VI) in goethite ([alpha]-FeOOH) suspensions by a dissimilatory metal-reducing bacterium. Geochim Cosmochim
Acta 64:3085–3098.
Geissler A, Law GTW, Morris K, Boothman C, Burke ITB, Livens FR, Lloyd
JR. 2011. (this issue). Microbial communities associated with the oxidation
of iron and technetium in bioreduced sediments.
Harris SJ, Mortimer RJG. 2002. Determination of nitrate in small water samples
(100 µL) by the cadmium-copper reduction method: A manual technique with
application to the interstitial waters of marine sediments. Int J Environ Anal
Chem 82:369–376.
Istok JD, Senko JM, Krumholz LR, Watson D, Bogle MA, Peacock A, Chang
YJ, White DC. 2004. In situ bioreduction of technetium and uranium in a
nitrate-contaminated aquifer. Environ Sci Technol 38:468–475.
Ithurbide A, Peulon S, Miserque F, Beaucaire C, Chausse C. 2009. Interaction
between uranium (VI) and siderite in carbonate solutions. Radiochim Acta
97:177–180.
Jeon BH, Dempsey BA, Burgos WD, Barnett MO, Roden EE. 2005. Chemical
reduction of U(VI) by Fe(II) at the solid-water interface using natural and
synthetic Fe(III) oxides. Environ Sci Technol 39:5642–5649.
Johnson DA, Florence TM. 1971. Spectrophotometric determination of uranium(VI) with 2-(5-bromo-2-pyridylazo)-5diethylaminophenol. Anal Chim
Acta 53:73–79.
Komlos J, Peacock A, Kukkadapu RK, Jaffé PR. 2008. Long-term dynamics of
uranium reduction/reoxidation under low sulfate conditions. Geochim Cosmochim Acta 72:3603–3615.
Law GTW, Geissler A, Boothman C, Burke IT, Livens FR, Lloyd JR, Morris
K. 2010. Role of nitrate in conditioning aquifer sediments for technetium
bioreduction. Environ Sci Technol 44:150–155.
Liu C, Zachara JM, Zhong L, Kukkadupa R, Szecsody JE, Kennedy DW. 2005.
Influence of sediment bioreduction and reoxidation on uranium sorption.
Environ Sci Technol 39:4125–4133.
505
Lloyd JR, Renshaw JC. 2005. Bioremediation of radioactive waste:
radionuclide-microbe interactions in laboratory and field-scale studies. Curr
Opin Biotech 16:254–260.
Lovley DR, Phillips EJP. 1986. Availability of ferric iron for microbial reduction
in bottom sediments of the freshwater tidal Potomac River. Appl Environ
Microbiol 52:751–757.
Lovley DR, Phillips EJP. 1987. Rapid assay for microbially reducible ferric iron
in aquatic sediments. Appl Environ Microbiol 53:1536–1540.
Lovley DR, Phillips EJP, Gorby YA, Landa ER. 1991. Microbial reduction of
uranium. Nature 350:413–416.
Lovley DR, Philips EJP. 1992. Reduction of uranium by Desulfovibrio desulfuricans. Appl Environ Micro 58:850–856.
Missana T, Maffiotte C, Garcı́a-Gutiérrez M. 2003, Surface reactions kinetics
between nanocrystalline magnetite and uranyl. J Coll Interf Sci 261:154–160.
Moon HS, Komlos J, Jaffé PR. 2007. Uranium reoxidation in previously
bioreduced sediment by dissolved oxygen and nitrate. Environ Sci Technol
41:4587–4592.
Morris K, Raiswell R. 2002. Biogeochemical cycles and the remobilization of
the actinide elements. In: Interactions of Microorganisms with radionuclides,
Keith-Roach MJ, Livens FR, editors. Elsevier, London.
Morris K, Livens FR, Charnock JM, Burke IT, McBeth JM, Begg JDC, Boothman C, Lloyd JR. 2008. An X-ray absorption study of the fate of technetium
in reduced and reoxidized sediments and mineral phases. Appl Geochem
23:603–617.
Moyes LN, Parkman RH, Charnock JM, Vaughan DJ, Livens FR, Hughes CR,
Braithwaite A. 2000. Uranium uptake from aqueous solution by interaction
with goethite, lepidocrocite, muscovite, and mackinawite: an X-ray absorption spectroscopy study. Environ Sci Technol. 34:1062–1068.
O’Loughlin EJ, Kelly SD, Cook RE, Csencsits R, Kemner KM. 2003. Reduction
of uranium(VI) by mixed iron(II)/iron(III) hydroxide (green rust): formation
of UO2 nanoparticles. Environ Sci Technol 37:721–727.
O’Loughlin EJ, Kelly SD, Kemner KM. 2010. XAFS investigation of the interactions of U(VI) with secondary mineralization products from the bioreduction
of Fe(III) oxides. Environ Sci Technol 44:1656–1661.
Ortiz-Bernad I, Anderson RT, Vrionis HA, Lovley DR. 2004. Resistance
of solid-phase U(VI) to microbial reduction during in situ bioremediation
of uranium-contaminated groundwater. Appl Environ Microbiol 70:7558–
7560.
Ravel B, Newville M. 2005. ATHENA, ARTEMIS, HEPHAESTUS: Data analysis for X-ray absorption spectroscopy using IFEFFIT. J Synchrotron Rad
12:537–541.
Scott TB, Allen GC, Heard PJ, Randell MG. 2005. Reduction of U(VI) to U(IV)
on the surface of magnetite. Geochim Cosmochim Acta 69:5639–5646.
Senko JM, Istok JD, Suflita JM, Krumholz LR. 2002. In situ evidence for
uranium immobilization and remobilization. Environ Sci Technol 36:1491–
1496.
Sharp JO, Schofield E, Veeramani H, Suvorova E, Junier P, Bargar JR, BernierLatmani R. 2008. Uranyl reduction by Geobacter sulfurreducens in the presence or absence of iron. In: Merkel BJ, Hasche-Berger A, editors. Uranium,
Mining and Hydrogeology. Springer: Berlin. P 725–732.
Stewart BD, Nico PS, Fendorf S. 2009. Stability of uranium incorporated into
Fe (hydr)oxides under fluctuating redox conditions. Environ Sci Technol
4:4922–4927.
Stookey LL. 1970. Ferrozine—A new spectrophotometric reagent for iron. Anal
Chem 42:779–781.
Sylwester ER, Hudson EA, Allen PG. 2000. The structure of uranium (VI)
sorption complexes on silica, alumina, and montmorillonite. Geochim Cosmochim Acta 64:2431–2438.
Thomas GW. 1996. Soil pH and acidity. In: Sparks D, editor. Methods of Soil
Analysis, Part 3—Chemical Methods. Madison, WI: Soil Science Society of
America.
Um W, Serne RJ, Krupka KM. 2007. Surface complexation modeling of U(VI)
sorption to Hanford sediment with varying geochemical conditions. Environ
Sci Technol 41:3587–3592.
Applied Geochemistry 26 (2011) S167–S169
Contents lists available at ScienceDirect
Applied Geochemistry
journal homepage: www.elsevier.com/locate/apgeochem
Composition, stability, and measurement of reduced uranium phases
for groundwater bioremediation at Old Rifle, CO
K.M. Campbell a,b,⇑, J.A. Davis a,c, J. Bargar d, D. Giammar e, R. Bernier-Latmani f, R. Kukkadapu g,
K.H. Williams c, H. Veramani e, K.-U. Ulrich e,h, J. Stubbs d, S. Yabusaki g, L. Figueroa i, E. Lesher i
M.J. Wilkins g, A. Peacock j, P.E. Long g
a
U.S. Geological Survey, 345 Middlefield Road, Menlo Park, CA 94025, USA
U.S. Geological Survey, 3215 Marine St., Boulder, CO 80303, USA
Lawrence Berkeley National Laboratory, 1 Cyclotron Rd., MS 90-1116, Berkeley, CA 94720, USA
d
Stanford Synchrotron Radiation Lightsource, 2575 Sand Hill Road, Menlo Park, CA 94025, USA
e
Washington University in Saint Louis, One Brookings Drive, St. Louis, MO 63130, USA
f
Environmental Microbiology Laboratory, École Polytechnique Fédérale de Lausanne, Lausanne, CH 1015, Switzerland
g
Pacific Northwest National Laboratory, Richland, WA 99352, USA
h
BGD Boden- und Grundwasserlabor GmbH Dresden, Tiergartenstraße 48, 01219 Dresden, Germany
i
Colorado School of Mines, 1500 Illinois St., Golden, CO 80401, USA
j
Haley and Aldrich, Oak Ridge, TN 37830, USA
b
c
a r t i c l e
i n f o
Article history:
Available online 26 March 2011
a b s t r a c t
Reductive biostimulation is currently being explored as a possible remediation strategy for U-contaminated groundwater, and is being investigated at a field site in Rifle, CO, USA. The long-term stability of
the resulting U(IV) phases is a key component of the overall performance of the remediation approach
and depends upon a variety of factors, including rate and mechanism of reduction, mineral associations
in the subsurface, and propensity for oxidation. To address these factors, several approaches were used to
evaluate the redox sensitivity of U: (1) measurement of the rate of oxidative dissolution of biogenic uraninite (UO2(s)) deployed in groundwater at Rifle, (2) characterization of a zone of natural bioreduction
exhibiting relevant reduced mineral phases, and (3) laboratory studies of the oxidative capacity of Fe(III)
and reductive capacity of Fe(II) with regard to U(IV) and U(VI), respectively.
Published by Elsevier Ltd.
1. Introduction
The legacy of U ore milling and processing has left many sites,
particularly in the western USA, with impacted groundwater even
after extensive reclamation projects. Although the concentrations
of U at these sites are substantially lower after the mill tailings are
removed and the site is remediated, groundwater concentrations
still often exceed the maximum contaminant level required for site
closure. Since conventional remediation technologies (e.g., pump
and treat) are costly for this type of scenario, alternate strategies
are currently being investigated. One promising strategy is reductive bioremediation, where dissolved U(VI) is reduced to relatively
insoluble U(IV) by stimulating a native metal-reducing microbial
community with an organic C substrate such as ethanol, acetate or
molasses. This process has been shown to significantly decrease
dissolved U(VI) concentrations (e.g., Anderson et al., 2003).
⇑ Corresponding author at: USGS, 345 Middlefield Road, Menlo Park, CA 94025,
USA. Tel.: +1 303 541 3035.
E-mail address: [email protected] (K.M. Campbell).
0883-2927/$ - see front matter Published by Elsevier Ltd.
doi:10.1016/j.apgeochem.2011.03.094
Since solid phase U(IV) is the desired product of reductive bioremediation, the long term efficacy of treatment will depend on
the stability of these phases. Therefore, it is important to identify
the mechanisms of formation, characterize the phases, and assess
their stability in the subsurface under oxidizing conditions. The
objectives of this work are to synthesize the results of several studies evaluating (1) the stability of biogenic uraninite (UO2(s)) deployed in groundwater at Rifle, (2) a zone of natural bioreduction
exhibiting similar processes observed in artificially stimulated bioreduction, (3) the possible role of abiotic oxidation of U(IV) by
Fe(III) as well as U(VI) reduction by adsorbed Fe(II).
2. Results and discussion
2.1. In situ uraninite stability in Rifle groundwater
Nanoparticulate biogenic uraninite (UO2(s)) is a well-characterized product of enzymatic U(VI) reduction by several species of
metal-reducing bacteria (e.g., Bargar et al., 2008, and references
therein). Although other forms of U(IV) may be produced during
reduction, such as U(IV) adsorbed to biomass and minerals (Ber-
S168
K.M. Campbell et al. / Applied Geochemistry 26 (2011) S167–S169
nier-Latmani et al., 2010; Fletcher et al., 2010), biogenic uraninite
provides a proxy for various U(IV) phases and can be used to constrain the upper end of U(IV) stability in sediments. Of the possible
oxidants in groundwater, dissolved O2 (DO) is particularly important because it is ubiquitous and may be present in relatively high
concentrations in upgradient groundwater. Rates of oxidative dissolution by DO were measured in situ by deploying biogenic uraninite in two wells at the Rifle site with differing DO concentrations
using a novel membrane-walled cell (Campbell et al., submitted for
publication). After 104 days of incubation in the groundwater,
approximately 50% of the uraninite was dissolved with no accumulation of corrosion products. Compared to laboratory-derived rates,
rates of dissolution in the field are 50–100 times lower. The presence of biomass in the deployment cell additionally retarded the
oxidative dissolution in the field. Molecular diffusion and surface
passivation by groundwater solutes are likely to be key processes
decreasing oxidation rates in the field.
2.2. Characterization of natural bioreduction zone
Several cores were drilled in a zone of natural bioreduction at
the Rifle site in an area that had never been subject to acetate
amendment. Sediment samples from a transect of samples ranging
from typical Rifle sediments to naturally bioreduced sediments
were analyzed to determine U and Fe oxidation state, Fe mineralogy, reduced S phases, solid phase organic C content, and to characterize the microbial community. Solid phase U concentrations
were substantially higher (2–10 times) in the naturally bioreduced
sediments, with significant amounts of U(IV) present. The U(IV)
was found to be in an adsorbed phase, rather than as nanocrystalline uraninite. Elevated concentrations of reduced Fe and S phases
as well as organic C were also measured. Biomass was correlated to
organic C, suggesting that natural bioreduction was stimulated by
a zone of increased organic C, resulting in Fe, U and S reduction.
The zone of natural bioreduction appears to be stabilized with
respect to oxidation, possibly through maintenance of locally
reducing conditions by microbial activity and the presence of
redox-buffering mineral phases (reduced Fe and S phases)
(Campbell et al., in preparation).
2.3. Chemical extraction for determination of labile U(VI) and
oxidizable U(IV) content in sediment
In natural sediments where solid phase U concentrations are
relatively low, as is the case with Rifle, direct spectroscopic measurement of U oxidation state is often beyond the capability of current technology and/or often not feasible for a large number of
samples. Since dissolved inorganic C is a strong ligand for U(VI),
HCO3/CO3 chemical extractions can serve as an alternate method
for measuring solid phase U oxidation state. Anoxic sediment extracted with a HCO3/CO3 solution liberates labile (adsorbed)
U(VI), while a subset of the same sample extracted under oxic conditions releases total oxidizable/labile U(VI); the difference is the
oxidizable U(IV) content of the sediment. Conventionally, the anoxic extraction is performed at pH 9.4 under a 5% CO2 atmosphere
in an anaerobic chamber (Kohler et al., 2004). However, a comparison of oxidation state estimates obtained on a naturally-bioreduced Rifle sediment using the anoxic extraction method and
X-ray absorption spectroscopy showed that substantial oxidation
of U(IV) occurred during anoxic extraction. Subsequent experiments with biogenic uraninite and ferrihydrite and additional thermodynamic calculations demonstrated that Fe(III) can oxidize
U(IV) under anoxic extraction conditions. A new extraction method
was shown to prevent anaerobic oxidation in Rifle sediments by
increasing the pH to 10.5 and decreasing the CO2 atmosphere to
400 ppm. In addition, the experiments and calculations extend
the range of pH and CO2 conditions reported by Ginder-Vogel
et al. (2006), and suggest that U(IV) oxidation by Fe(III) is a potentially relevant abiotic process in natural sediments.
2.4. Abiotic reduction of U(VI) by Fe(II)
Although dissolved Fe(II) is relatively unreactive toward U(VI)
at circumneutral pH, Fe(II) adsorbed to Fe oxides has been shown
to reduce adsorbed U(VI) (Liger et al., 1999). Proposed mechanisms
of this interaction include direct electron transfer between adsorbed species and/or electron migration through a conductive
mineral. To elucidate the former mechanism, Fe(II) and U(VI)
adsorption onto a non-conductive mineral (0.5 g/L c-Al2O3) was
investigated at pH 7 and 8.2 at several different surface loadings
and concentrations of CO2 in an anaerobic chamber. To understand
the effects of competitive adsorption between Fe(II) and U(VI) on
surface loading in the system, Ni(II) was used as a proxy for Fe(II)
in a separate set of adsorption experiments. Nickel(II) was also
used as a non-reactive control condition. At pH 7, no reaction between Fe(II) and U(VI) was observed under any conditions, but
reduction of U(VI) did occur when approximately equal amounts
of adsorbed Fe(II) and U(VI) were present on the alumina surface
at pH 8.2. This suggests that at appropriate surface loadings and
high pH, a direct electron transfer between Fe(II) and U(VI) can occur. This is consistent with an Fe(II) oligomer formation mechanism proposed for this reaction (Boyanov et al., 2007). The
results suggest that the direct reaction of adsorbed Fe(II) with adsorbed U(VI) is unlikely to proceed at the pH of Rifle groundwater.
3. Conclusions
With the goal of producing stable, insoluble reduced U(IV)
phases, biostimulation is currently being investigated as a remediation strategy for U-contaminated groundwater. The stability of
products in the field depends upon a variety of factors, including
rate and mechanism of reduction, mineral associations in the subsurface, and propensity for oxidation. A zone of natural bioreduction suggests that long-term stability of adsorbed U(IV) phases
may be possible, potentially by sustaining locally reduced conditions, precipitating redox-buffering minerals, and even maintaining the presence of biomass. Biogenic uraninite was found to be
more stable to oxidation by DO under aquifer conditions than predicted in laboratory studies, and its nanoparticulate nature does
not appear to make it more susceptible to oxidation on a surface
area normalized basis. Kinetic limitations of chemical diffusion
may extend the lifetime of U(IV) in the subsurface. However, other
oxidants in the subsurface, such as Fe(III) oxides, may be important. Although further research is necessary to determine the redox
balance in the field, laboratory results indicate that the favorability
of Fe(III) as an oxidant over Fe(II) as a reductant is very sensitive to
geochemical conditions, and may be an important consideration
during and after active remediation.
References
Anderson, R.T., Vrionis, H.A., Ortiz-Bernad, I., Resch, C.T., Long, P.E., Dayvault, R.,
Karp, K., Marutzky, S., Metzler, D.R., Peacock, A., White, D.C., Lowe, M., Lovley,
D.R., 2003. Stimulating the in situ activity of Geobacter species to remove
uranium from the groundwater of a uranium-contaminated aquifer. Appl.
Environ. Microbiol. 69, 5884–5891.
Bargar, J.R., Bernier-Latmani, R., Giammar, D.E., Tebo, B.M., 2008. Biogenic uraninite
nanoparticles and their importance for uranium remediation. Elements 4, 407–
412.
Bernier-Latmani, R., Veeramani, H., Vecchia, E.D., Junier, P., Lezama-Pacheco, J.S.,
Suvorova, E., Sharp, J.O., Wigginton, N.S., Bargar, J.R., 2010. Non-uraninite
products of microbial U(VI) reduction. Environ. Sci. Technol. 44, 5104–5111.
Boyanov, M.I., O’Loughlin, E.J., Roden, E.E., Fein, J.B., Kemner, K.M., 2007. Adsorption
of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The influence of
K.M. Campbell et al. / Applied Geochemistry 26 (2011) S167–S169
speciation on uranyl reduction studied by titration and XAFS. Geochim.
Cosmochim. Acta 71, 1898–1912.
Campbell, K.M, Kukkapdapu, R.K., Davis, J.A., Peacock, A.D., Qafoku, N., Lesher, E.,
Figueroa, L., Ranville, J., Williams, K.H., Wilkins, M.J., Resch, C.T., Icenhower, J.P.,
Long, P.E., in preparation. Characterizing the extent and role of natural
subsurface bioreduction in a uranium-contaminated aquifer. Chem. Geol., in
preparation.
Campbell, K.M., Veeramani, H., Ulrich, K.U., Blue, L., Giammar, D., Bernier-Latmani,
R., Stubbs, J.E., Surorova, E., Yabusaki, S., Mehta, A., Long, P. E., Bargar, J.R.,
submitted. Rates and mechanisms of oxidative dissolution of biogenic uraninite
in groundwater at Old Rifle, CO. Eniron. Sci. Technol., submitted for publication.
S169
Fletcher, K.E., Boyanov, M.I., Thomas, S.H., Wu, Q., Kemner, K.M., Löffler, F.E., 2010.
U(VI) reduction to mononuclear U(IV) by desulfitobacterium species. Environ.
Sci. Technol. 44, 4705–4709.
Ginder-Vogel, M., Criddle, C.S., Fendorf, S., 2006. Thermodynamic constraints on the
oxidation of biogenic UO2 by Fe(III) (Hydr)oxides. Environ. Sci. Technol. 40,
3544–3550.
Kohler, M., Curtis, G.P., Meece, D.E., Davis, J.A., 2004. Methods for estimating
adsorbed uranium(VI) and distribution coefficients of contaminated sediments.
Environ. Sci. Technol. 38, 240–247.
Liger, E., Charlet, L., Van Cappellen, P., 1999. Surface catalysis of uranium(VI)
reduction by iron(II). Geochim. Cosmochim. Acta 63, 2939–2955.
Geomicrobiology Journal, 28:160–171, 2011
Copyright © Taylor & Francis Group, LLC
ISSN: 0149-0451 print / 1521-0529 online
DOI: 10.1080/01490451003761137
Bioreduction Behavior of U(VI) Sorbed to Sediments
James D.C. Begg,1 Ian T. Burke,2 Jonathan R. Lloyd,2 Chris Boothman,2
Samual Shaw,1 John M. Charnock,2 and Katherine Morris1
1
Downloaded By: [University of Notre Dame] At: 00:40 5 May 2011
Earth Surface Science Institute, School of Earth and Environment, University of Leeds, Leeds,
United Kingdom
2
Research Centre for Radwaste and Decommissioning, and Williamson Centre for Molecular
Environmental Science, School of Earth, Atmospheric and Environmental Sciences, The University
of Manchester, Manchester, United Kingdom
It is well known that microbially mediated reduction can result
in the removal of U(VI)(aq) from solution by forming poorly soluble
U(IV) oxides; however, the fate of U(VI) already associated with
mineral surfaces is less clear. Here we describe results from both
oxic adsorption and anaerobic microcosm experiments to examine
the fate of sorbed U(VI) during microbially mediated bioreduction.
The microcosm experiments contained sediment representative of
the nuclear facility at Dounreay, UK. In oxic adsorption experiments, uptake of U(VI) was rapid and complete from artificial
groundwater and where groundwater was amended with 0.2 mmol
l−1 ethylenediaminetetraacetic acid (EDTA) a complexing ligand
used in nuclear fuel cycle operations. By contrast, uptake of U(VI)
was incomplete in groundwaters amended with 10 mmol l−1 bicarbonate. Analysis of sediments using X-ray adsorption spectroscopy
showed that in these oxic samples, U was present as U(VI). After
anaerobic incubation of U(VI) labelled sediments for 120 days, microbially mediated Fe(III)- and SO4 2−- reducing conditions had developed and XAS data showed uranium was reduced to U(IV). Further investigation of the unamended groundwater systems, where
oxic systems were dominated by U(VI) sorption, showed that reduction of sorbed U(VI) required an active microbial population
and occurred after robust iron- and sulfate- reducing conditions
had developed. Microbial community analysis of the bioreduced
sediment showed a community shift compared to the oxic sediment with close relatives of Geobacter and Clostridium species,
which are known to facilitate U(VI) reduction, dominating. Overall, efficient U(VI) removal from solution by adsorption under oxic
conditions dominated in unamended and EDTA amended systems.
In all systems bioreduction resulted in the formation of U(IV) in
solids.
Keywords
uranium, speciation, bioreduction, sediments, sorbtion
Received 27 November 2009; accepted 8 March 2010.
Current affiliation for James D. C. Begg: Glenn T. Seaborg Institute, Lawrence Livermore National Laboratory, Livermore, CA, USA.
Address correspondence to Katherine Morris, Research Centre for
Radwaste Disposal, School of Earth, Atmospheric and Environmental
Sciences, The University of Manchester, Manchester M13 9PL, United
Kingdom. E-mail: [email protected]
INTRODUCTION
Uranium is considered a problematic contaminant due to its
expected mobility under oxic environmental conditions, toxicity
to humans and widespread occurrence throughout the world at
sites where nuclear fuel cycle operations have occurred. Thus,
there are strong incentives to understand and control the behavior of uranium in contaminated environments. Under environmental conditions, uranium is present in two chemically stable
forms; under oxic conditions the uranyl cation (U(VI)O2 2+)
dominates and under reducing conditions, insoluble U(IV)O2
uraninite (solubility ca 10−17 M at pH > 4) dominates (Dozol
and Hagemen 1993; Lovley et al. 1991). Under oxic conditions,
the behavior of the uranyl cation is complex and dependent
on a number of factors. It may interact strongly with sediment
components and become sorbed to surfaces (Barnett et al. 2002;
Dong et al. 2006; Jeon et al. 2005; Ortiz-Bernad et al. 2004). Alternatively, at circumneutral pH and where carbonate is present,
it is likely to form relatively soluble anionic species such as
[UO2 (CO3 )2 ]2− (Clark et al. 1995).
The microbially mediated development of anoxia in sediments has a significant effect on aqueous U(VI) behavior with
U(VI) reduction producing poorly soluble UO2 that is retained
on a wide variety of environmental materials (Gu et al. 2005;
Lovley et al. 1991; Wilkins et al. 2007). These observations
have led to the development of bioremediation as a treatment
for subsurface uranium groundwater contamination. Here, an
electron donor is added to the subsurface to promote bioreduction resulting in precipitation of UO2 on sediments (Anderson
et al. 2003; Wu et al. 2006). However, there is a paucity of
information on the biogeochemical behavior of U(VI) when
it is already adsorbed to sediments. Under certain conditions
sediment associated U(VI) is reportedly recalcitrant to bioreduction (Jeon et al. 2005; Ortiz-Bernad et al. 2004) whilst under different conditions, bioreduction may occur (Dong et al.
2006; Kelly et al. 2008). Further complicating the fate of uranium in contaminated environments is its ability to form a range
160
BIOREDUCTION BEHAVIOR OF U(VI)
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FIG. 1. Time dependent U(VI) removal from solution in lower U(VI) concentration (53 µmol l−1) oxic adsorption experiments: oxic unamended (+);
oxic EDTA amended (); oxic carbonate amended (). Inset shows removal of
uranium from solution in oxic unamended systems over 60 min. Error bars are
1σ of three replicates.
confirming that U(VI) was not oversaturated in the groundwater
system. Eh measurements during oxic sorption typically ranged
from +150 mV to +200 mV, indicating oxidizing conditions.
The extremely fast sorption of U(VI) in EDTA amended experiments confirms that sorption occurred despite the presence of an
excess of the complexant EDTA. In the oxic carbonate amended
experiment, removal from solution was much slower with only
30.0 ± 4.0% removal from solution observed after 2 h and a
slow but steady removal occurring over time with 89.5 ± 0.4%
sorption in oxic experiments occurring after 60 days (Figure
1). Presumably, incomplete uptake of UO2 2+ in these high carbonate systems is due to the formation of neutral or negatively
charged U(VI)-carbonate complexes (Clark et al. 1995).
In additional oxic adsorption experiments for XAS analysis, both XAS oxic unamended and XAS oxic EDTA amended
systems showed slower uptake kinetics than the lower uranium
concentration experiments (Figure 2). Nonetheless, complete (>
97.8 ± 0.6%) sorption of U(VI) occurred in both experiments
by 7 days and these samples were taken for XAS analysis. Fast
initial removal of U(VI) followed by slower removal is likely
due to rapid adsorption of U(VI) to surface sites followed by
slower uptake due to structural arrangement on the solid surface
(Cheng et al. 2006; Um et al. 2007; Waite et al. 1994). In the
XAS oxic carbonate amended systems, 25 ± 1.0% of the initial uranium spike was removed after 7 days and 35.6 ± 0.6%
was removed at 120 days. This is a reduced percentage uptake
compared with the lower uranium concentration oxic carbonate amended experiment. Only the 7-day sediment sample was
taken for XAS analysis.
XAS Analysis of Oxic Samples
To assess the speciation of uranium in the oxic sorption experiments, XAS analyses were undertaken on 7-day time point
samples from XAS oxic unamended, carbonate amended and
EDTA amended experiments. These samples were challenging
to measure due to the relatively low concentrations of U sorbed
163
FIG. 2. Time dependent uranium(VI) removal from solution in XAS U(VI)
oxic adsorption experiments: XAS oxic unamended (+); XAS oxic EDTA
amended (); XAS oxic carbonate amended (). XAS oxic carbonate amended
showed a plateau in removal of U(VI) from solution. Error bars are 1σ of three
replicates.
to sediments (an average of several hundred ppm U on solids).
We were able to obtain XANES data for all samples and where
feasible we also collected EXAFS data. In all 3 oxic 7-day samples the XANES spectra were characterized by their similarity
to the shape of the U(VI) standard and to U(VI) standards reported in the literature (Figure 3A; Boyanov et al. 2007). When
compared to the U(VI) and U(IV) standards, linear combination
fitting for these 7-day samples showed an ∼ 80% contribution from U(VI) which strongly supports the observed similarity between spectra from sediments, our U(VI) standards and
published work showing U(VI) spectra (Boyanov et al. 2007).
Thus, as expected in oxic conditions, XAS shows that U(VI)
is predominantly found sorbed to the sediment surface. This
observation was further supported by EXAFS analysis of the 7
day samples where sediment associated U was characterized by
a large peak in the Fourier transforms at ca. 1.80 Å which is
diagnostic for the axial U = O bond length in U(VI) (Figure 4,
Catalano et al. 2004).
Bioreduction
To simulate the behavior of U(VI) in anoxic subsurface environments, a series of bioreduction microcosm experiments were
performed at the lower uranium concentration. Overall bioreduction, indicated by lower Eh values (Figure 5A) occurred over
several weeks and at similar rates to those observed in previous
work (Begg et al. 2007). Development of Fe(III)-reducing conditions, indicated by increases in 0.5 N HCl extractable Fe(II),
was rapid in all systems and was measured at 41 ± 6%, 73 ± 2%
and 24 ± 4% respectively in bioreduction unamended, bioreduction carbonate amended and bioreduction EDTA amended
experiments at 15 days (Figure 5B). As expected, a significant
increase in total Fe in porewaters was also observed at 15 days
due to production of soluble Fe(II) as Fe(III)-reduction progressed (Figure 5C). In all microcosms there was a rise in pH as
bioreduction proceeded, consistent with production of OH− and
HCO3 − from oxidation of carbon-based electron donors (Figure 5D; Chang et al. 2005). Sulfate reduction, as indicated by
170
J. D. C. BEGG ET AL.
TABLE 4
Phylogenetic affiliation in reduced sediment of distinct RFLP types detected in a 16S rDNA clone library obtained by PCR
amplification using broad-specificity primers. Amplification was from the XAS bioreduction unamended 120 d sample.
Downloaded By: [University of Notre Dame] At: 00:40 5 May 2011
Clone
RFLP
Type
JBT120-1
JBT120-16
JBT120-2
1, 12
JBT120-4
JBT120-5
JBT120-6
JBT120-8
JBT120-9
JBT120-11
JBT120-13
JBT120-14
JBT120-15
JBT120-18
JBT120-19
JBT120-20
JBT120-22
JBT120-27
JBT120-28
JBT120-30
JBT120-33
JBT120-34
JBT120-45
3
4
5
6
7
8
9
10
11
13
14
15
16
17
18
19
20
21
22
JBT120-53
23
2
Closest Matching Micro Organism
(accession number)
Identities
(% Match)
%
Present
Clostridium lituseburense
448/464 (96%)
7.2
Uncultured Xiphinematobacteriaceae bacterium
clone EB1116
Uncultured soil bacterium clone CWT SM03 B11
Uncultured bacterium clone ORS25C b04
Clostridium puniceum
Uncultured bacterium clone Amb 16S 1261
Clostridium tunisiense
Candidatus Magnetobacterium bavaricum
Acetivibrio cellulolyticus
Methylocella palustris
Uncultured soil bacterium clone HSB NT53 H06
Uncultured bacterium clone ORSFES f09
Bacillus longiquaesitum
Bacillus litoralis
Unidentified bacterium clone FI-2M B10
Geobacter psychrophilus strain P35
Uncultured bacterium clone aab39a02
Beijerinckia sp. TB13
Methylosinus sporium strain NR3K
Bacillus longiquaesitum
Uncultured alpha proteobacterium clone
Fl-1F C12
Uncultured bacterium clone CON4 C02
458/462 (99%)
10.6
Spartobacteria
409/422 (96%)
459/482 (95%)
466/470 (99%)
495/504 (98%)
383/397 (96%)
232/277 (83%)
432/477 (90%)
422/452 (93%)
426/463 (92%)
435/446 (97%)
483/501 (96%)
497/506 (98%)
469/493 (95%)
483/522 (92%)
499/500 (99%)
344/379 (90%)
310/338 (91%)
189/216 (87%)
282/302 (93%)
8.8
1.8
12.5
1.8
1.8
3.6
3.6
3.6
7.1
1.8
7.1
3.6
3.6
7.1
3.6
1.8
1.8
1.8
3.6
Unknown
Unknown
Clostridiales
Unknown
Clostridiales
Nitrospirales
Clostridiales
Alpha-proteobacteria
Unknown
Unknown
Bacillales
Bacillales
Unknown
Deltaproteobacteria
Unknown
Alpha-proteobacteria
Alpha-proteobacteria
Bacillales
Alpha-proteobacteria
282/309 (91%)
1.8
and highlights that it is essential to maintain reducing conditions
in environments where bioremediation technologies are used to
treat U contaminated sites.
ACKNOWLEDGMENTS
Thanks to Bob Bilsborrow, for invaluable help in XAS data
acquisition, to Gareth Law, University of Leeds, for help in
XANES data acquisition and analysis and to Miranda KeithRoach, Stephanie Handley, Rachael Spraggs and Doug McAllister. This work was supported by grants NE/D00473X/1 and
NE/D005361/1 from the UK Natural Environment Research
Council, NERC studentship NER/S/A/2004/13005 to JDCB and
by Daresbury SRS beamtime allocation from the UK Science
and Technology Facilities Council.
REFERENCES
Anderson RT, Vrionis HA, Ortiz-Bernad I, Resch CT, Long PE, Dayvault R,
Karp K, Marutzky S, Metzler DR, Peacock A, White DC, Lowe M, Lovley
DR. 2003. Stimulating the in situ activity of Geobacter species to remove
Phylogenetic Division
Clostridiales
Unknown
uranium from the groundwater of a uranium-contaminated aquifer. Appl
Environ Microbiol 69:5884–5891.
Barnett MO, Jardine PM, Brooks SC. 2002. U(VI) Adsorption to heterogeneous
subsurface media: application of a surface complexation model. Environ Sci
Technol 36:937–942.
Begg JDC, Burke IT, Charnock JM, Morris K. 2008. Technetium reduction and
reoxidation behavior in Dounreay soils. Radiochim Acta 96:631–636.
Begg JDC, Burke IT, Morris K. 2007. The behaviour of technetium during
microbial reduction in amended soil from Dounreay. U.K. Sci Total Environ
373:297–304.
Binsted N. 1998. CLRC Daresbury Laboratory EXCURV98 program. CLRC
Daresbury Laboratory: Warrington, UK.
Boyanov MI, O’Loughlin EJ, Roden EE, Fein JB, Kemner KM. 2007. Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The
influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim Cosmochim Acta 71:1898–1912.
Burke IT, Boothman C, Lloyd JR, Livens FR, Charnock JM, McBeth JM,
Mortimer RJG, Morris K. 2006. Reoxidation behavior of technetium, iron,
and sulfur in estuarine sediments. Environ Sci Technol 40:3529–3535.
Catalano JG, Heald SM, Zachara JM, Brown Jr. GE. 2004. Spectroscopic and
diffraction study of uranium speciation in contaminated sediments from the
Hanford site, Washington State. Environ Sci Technol 38:2822–2828.
Chang Y-J, Long PE, Geyer R, Peacock AD, Resch CT, Sublette K, Pfiffner
S, Smithgall A, Anderson RT, Vrionis HA, Stephen JR, Dayvault R,
Available online at www.sciencedirect.com
Geochimica et Cosmochimica Acta 75 (2011) 5648–5663
www.elsevier.com/locate/gca
The effect of pH and natural microbial phosphatase activity
on the speciation of uranium in subsurface soils
Melanie J. Beazley a, Robert J. Martinez b, Samuel M. Webb c,
Patricia A. Sobecky b, Martial Taillefert a,⇑
a
School of Earth & Atmospheric Sciences, Georgia Institute of Technology, Atlanta, GA 30332-0340, USA
b
Department of Biological Sciences, University of Alabama, Tuscaloosa, AL 35487, USA
c
Stanford Synchrotron Radiation Lightsource, Menlo Park, CA 94025, USA
Received 19 October 2010; accepted in revised form 6 July 2011; available online 19 July 2011
Abstract
The biomineralization of U(VI) phosphate as a result of microbial phosphatase activity is a promising new bioremediation
approach to immobilize uranium in both aerobic and anaerobic conditions. In contrast to reduced uranium minerals such as
uraninite, uranium phosphate precipitates are not susceptible to changes in oxidation conditions and may represent a longterm sink for uranium in contaminated environments. So far, the biomineralization of U(VI) phosphate has been demonstrated with pure cultures only. In this study, two uranium contaminated soils from the Department of Energy Oak Ridge
Field Research Center (ORFRC) were amended with glycerol phosphate as model organophosphate source in small flowthrough columns under aerobic conditions to determine whether natural phosphatase activity of indigenous soil bacteria
was able to promote the precipitation of uranium(VI) at pH 5.5 and 7.0. High concentrations of phosphate (1–3 mM) were
detected in the effluent of these columns at both pH compared to control columns amended with U(VI) only, suggesting that
phosphatase-liberating microorganisms were readily stimulated by the organophosphate substrate. Net phosphate production
rates were higher in the low pH soil (0.73 ± 0.17 mM d1) compared to the circumneutral pH soil (0.43 ± 0.31 mM d1), suggesting that non-specific acid phosphatase activity was expressed constitutively in these soils. A sequential solid-phase extraction scheme and X-ray absorption spectroscopy measurements were combined to demonstrate that U(VI) was primarily
precipitated as uranyl phosphate minerals at low pH, whereas it was mainly adsorbed to iron oxides and partially precipitated
as uranyl phosphate at circumneutral pH. These findings suggest that, in the presence of organophosphates, microbial phosphatase activity can contribute to uranium immobilization in both low and circumneutral pH soils through the formation of
stable uranyl phosphate minerals.
Ó 2011 Elsevier Ltd. All rights reserved.
1. INTRODUCTION
Uranium contamination represents a major environmental concern at Department of Energy (DOE) sites
across the United States. At the Oak Ridge Field Research
Center (ORFRC) of the Oak Ridge National Laboratory
⇑ Corresponding author. Address: School of Earth & Atmo-
spheric Sciences, 311 Ferst Drive, Atlanta, GA 30332-0340, USA.
Tel.: +1 404 894 6043; fax: +1 404 894 5638.
E-mail address: [email protected] (M. Taillefert).
0016-7037/$ - see front matter Ó 2011 Elsevier Ltd. All rights reserved.
doi:10.1016/j.gca.2011.07.006
Reservation (Tennessee), soils and groundwater are heavily
contaminated with depleted uranium (up to 252 lM) and
nitrate (up to 645 mM) as a result of more than 30 years
of uranium enrichment at the facility (Brooks, 2001).
Remediation of uranium in subsurface environments is difficult mainly because the speciation and mobility of uranium are controlled by its oxidation state (Cotton et al.,
1999) and a complex web of biogeochemical reactions,
including adsorption and desorption (Hsi and Langmuir,
1985; Langmuir, 1997), precipitation and dissolution of
minerals (Langmuir, 1978), and complexation (Barnett
Effect of microbial phosphatase on uranium speciation
5655
Table 1
Parameters derived from fitting of U LIII-edge EXAFS of flow-through column soil samples conducted at pH 7.
0
0
Column
Shell
N
R (Å )
r2 (Å 2)
U amended control
0–1 cm
Oax
2
1.80 (0.006)
0.003 (0.0009)
Oeq
Oeq
C
Fe
1.52
1.48
1.27
0.14
2.33
2.46
2.94
3.46
0.003
0.003
0.003
0.003
Oax
2
Oeq
Oeq
C
Fe
P
POU MSa
OPOU MSa
PO distb
2.47
0.14
1.38
0.25
0.27
0.54
0.27
1.56
Oax
2
1.80 (0.009)
0.005 (0.001)
Oeq
Oeq
C
Fe
P
POU MSa
OPOU MSa
PO distb
2.17 (0.57)
0.17 (0.35)
1.37 (1.5)
0.34 (0.11)
0.21 (0.16)
0.42
0.21
1.7 (0.06)
2.29
2.41
2.92
3.43
3.55
3.62
3.99
0.003
0.003
0.003
0.003
0.005
0.005
0.005
Oax
2
1.80 (0.01)
0.004 (0.001)
Oeq
Oeq
C
Fe
P
POU MSa
OPOU MSa
PO distb
2.68
0.42
1.67
0.40
2.56
5.11
2.56
1.68
2.30
2.45
2.90
3.44
3.79
3.89
3.98
0.003
0.003
0.003
0.003
0.005
0.005
0.005
Column 1
0–1 cm
Column 2
0–1 cm
Column 3
0–1 cm
(0.34)
(0.48)
(0.38)
(0.1)
(0.6)
(0.4)
(1.6)
(0.17)
(0.45)
(0.02)
(0.01)
(0.02)
(0.04)
1.80 (0.009)
0.004 (0.001)
2.29
2.39
2.90
3.42
3.68
3.77
3.85
0.003
0.003
0.003
0.003
0.005
0.005
0.005
(0.01)
(0.25)
(0.06)
(0.03)
(0.11)
(0.06)
(0.05)
v2v
R factor
0.203
0.0112
10.9 (2)
0.227
0.0141
10.4 (1.8)
1.629
0.0134
10.9 (2.1)
0.275
0.0110
DE0 (eV)
8.68 (1.3)
(0.05)
(0.7)
(0.5)
(2.2)
(0.11)
(1.9)
(0.02)
(0.2)
(0.08)
(0.02)
(0.43)
(0.17)
(0.06)
(0.01)
(0.06)
(0.05)
(0.02)
(0.7)
(0.05)
(0.03)
(0.03)
Errors are given in parenthesis (no error means the value was fixed, or calculated from other parameters.
a
MS denoted multiple scattering paths.
b
PO dist is the distance between the P–O in phosphate coordination (used for the MS paths).
immediately after U(VI) addition, possibly due to a combination of U–P precipitation and U(VI) toxicity on NSAPexpressing microbial populations. Concomitantly with the
U(VI) decrease in the effluent, however, organophosphate
hydrolysis rebounded to produce similar phosphate concentrations at low pH and even ca. 50% higher at circumneutral
pH, suggesting that phosphatase activity was restored by the
removal of U(VI) from solution. Previous incubations with
pure cultures of NSAP-carrying Rahnella sp. demonstrated
a similar behavior (Beazley et al., 2007) that was attributed
to the toxicity of U(VI) initially adsorbed to cell membranes
which was desorbed during the precipitation of U–P minerals
(Beazley et al., 2009).
4.2. Speciation of uranium in the solid phase
Despite the depletion of oxygen and nitrate in incubations with both soils, U XANES analysis of the soils iden-
tified U(VI) as the primary oxidation state (Fig. 3a and 6a)
within the first few centimeters of the cores where the
majority of the uranium was located (Fig. 2a and 5a). Unless uraninite was in a colloidal (Suzuki et al., 2002) or
molecular (Fletcher et al., 2010) form that could have escaped from the columns, these results suggest that the
majority of uranium remained oxidized in these incubations. The small concentrations of U(VI) detected in the
effluents of both soils amended with organophosphate compared to the U-controls (Figs. 1e and 4e) indicate that
U(VI) was removed from solution by adsorption onto the
solid phase or precipitation as uranium phosphate minerals.
As adsorption is likely the main process of removal of
U(VI) in the U-control columns, the negligible fraction of
exchangeable uranium and the large release of uranium in
the presence of AcOH (Fig. 2b and 5b) suggest that
U(VI) adsorbs strongly at both pHs, though more notably
at low pH. Interestingly, most uranium was removed from
Effect of microbial phosphatase on uranium speciation
Biological and Environmental Research, and by the National
Institutes of Health, National Center for Research Resources,
Biomedical Technology Program. We thank Dave Watson of
Oak Ridge National Laboratory for providing ORFRC soil cores
and two anonymous reviewers for their valuable comments on an
earlier version of this paper.
REFERENCES
Akob D. M., Mills H. J. and Kostka J. E. (2007) Metabolically
active microbial communities in uranium-contaminated subsurface sediments. FEMS Microbiol. Ecol. 59, 95–107.
Ankudinov A. L., Ravel B., Rehr J. J. and Conradson S. D. (1998)
Real space multiple scattering calculation of XANES. Phys.
Rev. B 58, 7565.
Appukuttan D., Rao A. and Apte S. (2006) Engineering of
Deinococcus radiodurans R1 for bioprecipitation of uranium
from dilute nuclear waste. Appl. Environ. Microbiol. 72, 7873–
7878.
Bargar J. R., Reitmeyer R., Lenhart J. J. and Davis J. A. (2000)
Characterization of U(VI)-carbonato ternary complexes on
hematite: EXAFS and electrophoretic mobility measurements.
Geochim. Cosmochim. Acta 64, 2737–2749.
Barnett M. O., Jardine P. M. and Brooks S. C. (2002) U(VI)
adsorption to heterogeneous subsurface media: application of a
surface complexation model. Environ. Sci. Technol. 36, 937–
942.
Beazley M. J., Martinez R. J., Sobecky P. A., Webb S. M. and
Taillefert M. (2007) Uranium biomineralization as a result of
bacterial phosphatase activity: insights from bacterial isolates
from a contaminated subsurface. Environ. Sci. Technol. 41,
5701–5707.
Beazley M. J., Martinez R. J., Sobecky P. A., Webb S. M. and
Taillefert M. (2009) Nonreductive biomineralization of uranium(VI) phosphate via microbial phosphatase activity in
anaerobic conditions. Geomicrobiol. J. 26, 431–441.
Bostick B. C., Fendorf S., Barnett M. O., Jardine P. M. and Brooks
S. C. (2002) Uranyl surface complexes formed on subsurface
media from DOE facilities. Soil Sci. Soc. Am. J. 66, 99–108.
Boyanov M. I., O’Loughlin E. J., Roden E. E., Fein J. B. and
Kemner K. M. (2007) Adsorption of Fe(II) and U(VI) to
carboxyl-functionalized microspheres: the influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim. Cosmochim. Acta 71, 1898–1912.
Brendel P. J. and Luther, III, G. W. (1995) Development of a gold
amalgam voltammetric microelectrode for the determination of
dissolved Fe, Mn, O2, and S(-II) in porewaters of marine and
freshwater sediments. Environ. Sci. Technol. 29, 751–761.
Bristow G. and Taillefert M. (2008) VOLTINT: a Matlab (R)based program for semi-automated processing of geochemical
data acquired by voltammetry. Comput. Geosci. 34, 153–162.
Brooks S. C. (2001) Waste Characteristics of the Former S-3 Ponds
and Outline of Uranium Chemistry Relevant to NABIR Field
Research Center Studies. NABIR FRC, Technical Report.
Carey E. A. and Taillefert M. (2005) The role of soluble Fe(III) in
the cycling of iron and sulfur in coastal marine sediments.
Limnol. Oceanogr. 50, 1129–1141.
Catalano J. G. and Brown, Jr., G. E. (2004) Analysis of uranylbearing phases by EXAFS spectroscopy: interferences, multiple
scattering, accuracy of structural parameters, and spectral
differences. Am. Mineral. 89, 1004–1021.
Cheng T., Barnett M. O., Roden E. E. and Zhuang J. (2004) Effects
of phosphate on uranium(VI) adsorption to goethite-coated
sand. Environ. Sci. Technol. 38, 6059–6065.
5661
Cheng T., Barnett M. O., Roden E. E. and Zhuang J. (2006) Effects
of solid-to-solution ratio on uranium(VI) adsorption and its
implications. Environ. Sci. Technol. 40, 3243–3247.
Cheng T., Barnett M. O., Roden E. E. and Zhuang J. (2007)
Reactive transport of uranium(VI) and phoshate in a goethitecoated sand column: an experimental study. Chemosphere 68,
1218–1223.
Chinni S., Anderson C., Ulrich K. and Giammar D. (2008) Indirect
UO2 oxidation by Mn(II)-oxidizing spores of Bacillus sp strain
SG-1 and the effect of U and Mn concentrations. Environ. Sci.
Technol. 42, 8709–8714.
Cotton F. A., Wilkinson G., Murillo C. A. and Bochmann M.
(1999) Advanced Inorganic Chemistry. John Wiley & Sons, Inc.,
New York.
De Pablo J., Casas I., Giménez J., Molera M., Rovira M., Duro L.
and Bruno J. (1999) The oxidative dissolution mechanism of
uranium dioxide. I. The effect of temperature in hydrogen
carbonate medium. Geochim. Cosmochim. Acta 63, 3097–
3103.
Diress A. and Lucy C. A. (2005) Study of the selectivity of
inorganic anions in hydro-organic solvents using indirect
capillary electrophoresis. J. Chromatogr. A 1085, 155–
163.
Finneran K. T., Housewright M. E. and Lovley D. R. (2002)
Multiple influences of nitrate on uranium solubility during
bioremediation of uranium-contaminated subsurface sediments. Environ. Microbiol. 4, 510–516.
Fletcher K. E., Boyanov M. I., Thomas S. H., Wu Q., Kemner K.
M. and Löffler F. E. (2010) U(VI) reduction to mononuclear
U(IV) by Desulfitobacterium species. Environ. Sci. Technol. 44,
4705–4709.
Fredrickson J. K., Zachara J. M., Kennedy D. W., Duff M. C.,
Gorby Y. A., Li S.-M. W. and Krupka K. M. (2000) Reduction
of U(VI) in goethite (a-FeOOH) suspensions by a dissimilatory
metal-reducing bacterium. Geochim. Cosmochim. Acta 64,
3085–3098.
Fredrickson J. K., Zachara J. M., Kennedy D. W., Liu C., Duff M.
C., Hunter D. B. and Dohnalkova A. (2002) Influence of Mn
oxides on the reduction of uranium (VI) by the metal-reducing
bacterium Shewanella putrefaciens. Geochim. Cosmochim. Acta
66, 3247–3262.
Fuller C. C., Bargar J. R. and Davis J. A. (2003) Molecular-scale
characterization of uranium sorption by bone apatite materials
for a permeable reactive barrier demonstration. Environ. Sci.
Technol. 37, 4642–4649.
Fuller C. C., Bargar J. R., Davis J. A. and Piana M. J. (2002)
Mechanisms of uranium interactions with hydroxyapatite:
implications for groundwater remediation. Environ. Sci. Technol. 36, 158–165.
Ginder-Vogel M., Stewart B. and Fendorf S. (2010) Kinetic and
mechanistic constraints on the oxidation of biogenic uraninite
by ferrihydrite. Environ. Sci. Technol. 44, 163–169.
Grasshoff K. (1983) Determination of nitrite. In Methods of
Seawater Analysis (eds. K. Grasshoff, M. Ehrhardt and K.
Kremling). Verlag Chemie GmbH, Weinheim.
Hsi C.-K. D. and Langmuir D. (1985) Adsorption of uranyl
onto ferric oxyhydroxides: application of the surface complexation site-binding model. Geochim. Cosmochim. Acta 49,
1931–1941.
Hudson E. A., Allen P. G., Terminello L. J., Denecke M. A. and
Reich T. (1996) Polarized x-ray-absorption spectroscopy of the
uranyl ion: comparison of experiment and theory. Phys. Rev. B
54, 156–165.
Hudson E. A., Terminello L. J., Viani B. E., Denecke M., Reich T.,
Allen P. G., Bucher J. J., Shuh D. K. and Edelstein N. M.
Available online at www.sciencedirect.com
Geochimica et Cosmochimica Acta 75 (2011) 5648–5663
www.elsevier.com/locate/gca
The effect of pH and natural microbial phosphatase activity
on the speciation of uranium in subsurface soils
Melanie J. Beazley a, Robert J. Martinez b, Samuel M. Webb c,
Patricia A. Sobecky b, Martial Taillefert a,⇑
a
School of Earth & Atmospheric Sciences, Georgia Institute of Technology, Atlanta, GA 30332-0340, USA
b
Department of Biological Sciences, University of Alabama, Tuscaloosa, AL 35487, USA
c
Stanford Synchrotron Radiation Lightsource, Menlo Park, CA 94025, USA
Received 19 October 2010; accepted in revised form 6 July 2011; available online 19 July 2011
Abstract
The biomineralization of U(VI) phosphate as a result of microbial phosphatase activity is a promising new bioremediation
approach to immobilize uranium in both aerobic and anaerobic conditions. In contrast to reduced uranium minerals such as
uraninite, uranium phosphate precipitates are not susceptible to changes in oxidation conditions and may represent a longterm sink for uranium in contaminated environments. So far, the biomineralization of U(VI) phosphate has been demonstrated with pure cultures only. In this study, two uranium contaminated soils from the Department of Energy Oak Ridge
Field Research Center (ORFRC) were amended with glycerol phosphate as model organophosphate source in small flowthrough columns under aerobic conditions to determine whether natural phosphatase activity of indigenous soil bacteria
was able to promote the precipitation of uranium(VI) at pH 5.5 and 7.0. High concentrations of phosphate (1–3 mM) were
detected in the effluent of these columns at both pH compared to control columns amended with U(VI) only, suggesting that
phosphatase-liberating microorganisms were readily stimulated by the organophosphate substrate. Net phosphate production
rates were higher in the low pH soil (0.73 ± 0.17 mM d1) compared to the circumneutral pH soil (0.43 ± 0.31 mM d1), suggesting that non-specific acid phosphatase activity was expressed constitutively in these soils. A sequential solid-phase extraction scheme and X-ray absorption spectroscopy measurements were combined to demonstrate that U(VI) was primarily
precipitated as uranyl phosphate minerals at low pH, whereas it was mainly adsorbed to iron oxides and partially precipitated
as uranyl phosphate at circumneutral pH. These findings suggest that, in the presence of organophosphates, microbial phosphatase activity can contribute to uranium immobilization in both low and circumneutral pH soils through the formation of
stable uranyl phosphate minerals.
Ó 2011 Elsevier Ltd. All rights reserved.
1. INTRODUCTION
Uranium contamination represents a major environmental concern at Department of Energy (DOE) sites
across the United States. At the Oak Ridge Field Research
Center (ORFRC) of the Oak Ridge National Laboratory
⇑ Corresponding author. Address: School of Earth & Atmo-
spheric Sciences, 311 Ferst Drive, Atlanta, GA 30332-0340, USA.
Tel.: +1 404 894 6043; fax: +1 404 894 5638.
E-mail address: [email protected] (M. Taillefert).
0016-7037/$ - see front matter Ó 2011 Elsevier Ltd. All rights reserved.
doi:10.1016/j.gca.2011.07.006
Reservation (Tennessee), soils and groundwater are heavily
contaminated with depleted uranium (up to 252 lM) and
nitrate (up to 645 mM) as a result of more than 30 years
of uranium enrichment at the facility (Brooks, 2001).
Remediation of uranium in subsurface environments is difficult mainly because the speciation and mobility of uranium are controlled by its oxidation state (Cotton et al.,
1999) and a complex web of biogeochemical reactions,
including adsorption and desorption (Hsi and Langmuir,
1985; Langmuir, 1997), precipitation and dissolution of
minerals (Langmuir, 1978), and complexation (Barnett
Effect of microbial phosphatase on uranium speciation
5655
Table 1
Parameters derived from fitting of U LIII-edge EXAFS of flow-through column soil samples conducted at pH 7.
0
0
Column
Shell
N
R (Å )
r2 (Å 2)
U amended control
0–1 cm
Oax
2
1.80 (0.006)
0.003 (0.0009)
Oeq
Oeq
C
Fe
1.52
1.48
1.27
0.14
2.33
2.46
2.94
3.46
0.003
0.003
0.003
0.003
Oax
2
Oeq
Oeq
C
Fe
P
POU MSa
OPOU MSa
PO distb
2.47
0.14
1.38
0.25
0.27
0.54
0.27
1.56
Oax
2
1.80 (0.009)
0.005 (0.001)
Oeq
Oeq
C
Fe
P
POU MSa
OPOU MSa
PO distb
2.17 (0.57)
0.17 (0.35)
1.37 (1.5)
0.34 (0.11)
0.21 (0.16)
0.42
0.21
1.7 (0.06)
2.29
2.41
2.92
3.43
3.55
3.62
3.99
0.003
0.003
0.003
0.003
0.005
0.005
0.005
Oax
2
1.80 (0.01)
0.004 (0.001)
Oeq
Oeq
C
Fe
P
POU MSa
OPOU MSa
PO distb
2.68
0.42
1.67
0.40
2.56
5.11
2.56
1.68
2.30
2.45
2.90
3.44
3.79
3.89
3.98
0.003
0.003
0.003
0.003
0.005
0.005
0.005
Column 1
0–1 cm
Column 2
0–1 cm
Column 3
0–1 cm
(0.34)
(0.48)
(0.38)
(0.1)
(0.6)
(0.4)
(1.6)
(0.17)
(0.45)
(0.02)
(0.01)
(0.02)
(0.04)
1.80 (0.009)
0.004 (0.001)
2.29
2.39
2.90
3.42
3.68
3.77
3.85
0.003
0.003
0.003
0.003
0.005
0.005
0.005
(0.01)
(0.25)
(0.06)
(0.03)
(0.11)
(0.06)
(0.05)
v2v
R factor
0.203
0.0112
10.9 (2)
0.227
0.0141
10.4 (1.8)
1.629
0.0134
10.9 (2.1)
0.275
0.0110
DE0 (eV)
8.68 (1.3)
(0.05)
(0.7)
(0.5)
(2.2)
(0.11)
(1.9)
(0.02)
(0.2)
(0.08)
(0.02)
(0.43)
(0.17)
(0.06)
(0.01)
(0.06)
(0.05)
(0.02)
(0.7)
(0.05)
(0.03)
(0.03)
Errors are given in parenthesis (no error means the value was fixed, or calculated from other parameters.
a
MS denoted multiple scattering paths.
b
PO dist is the distance between the P–O in phosphate coordination (used for the MS paths).
immediately after U(VI) addition, possibly due to a combination of U–P precipitation and U(VI) toxicity on NSAPexpressing microbial populations. Concomitantly with the
U(VI) decrease in the effluent, however, organophosphate
hydrolysis rebounded to produce similar phosphate concentrations at low pH and even ca. 50% higher at circumneutral
pH, suggesting that phosphatase activity was restored by the
removal of U(VI) from solution. Previous incubations with
pure cultures of NSAP-carrying Rahnella sp. demonstrated
a similar behavior (Beazley et al., 2007) that was attributed
to the toxicity of U(VI) initially adsorbed to cell membranes
which was desorbed during the precipitation of U–P minerals
(Beazley et al., 2009).
4.2. Speciation of uranium in the solid phase
Despite the depletion of oxygen and nitrate in incubations with both soils, U XANES analysis of the soils iden-
tified U(VI) as the primary oxidation state (Fig. 3a and 6a)
within the first few centimeters of the cores where the
majority of the uranium was located (Fig. 2a and 5a). Unless uraninite was in a colloidal (Suzuki et al., 2002) or
molecular (Fletcher et al., 2010) form that could have escaped from the columns, these results suggest that the
majority of uranium remained oxidized in these incubations. The small concentrations of U(VI) detected in the
effluents of both soils amended with organophosphate compared to the U-controls (Figs. 1e and 4e) indicate that
U(VI) was removed from solution by adsorption onto the
solid phase or precipitation as uranium phosphate minerals.
As adsorption is likely the main process of removal of
U(VI) in the U-control columns, the negligible fraction of
exchangeable uranium and the large release of uranium in
the presence of AcOH (Fig. 2b and 5b) suggest that
U(VI) adsorbs strongly at both pHs, though more notably
at low pH. Interestingly, most uranium was removed from
Effect of microbial phosphatase on uranium speciation
Biological and Environmental Research, and by the National
Institutes of Health, National Center for Research Resources,
Biomedical Technology Program. We thank Dave Watson of
Oak Ridge National Laboratory for providing ORFRC soil cores
and two anonymous reviewers for their valuable comments on an
earlier version of this paper.
REFERENCES
Akob D. M., Mills H. J. and Kostka J. E. (2007) Metabolically
active microbial communities in uranium-contaminated subsurface sediments. FEMS Microbiol. Ecol. 59, 95–107.
Ankudinov A. L., Ravel B., Rehr J. J. and Conradson S. D. (1998)
Real space multiple scattering calculation of XANES. Phys.
Rev. B 58, 7565.
Appukuttan D., Rao A. and Apte S. (2006) Engineering of
Deinococcus radiodurans R1 for bioprecipitation of uranium
from dilute nuclear waste. Appl. Environ. Microbiol. 72, 7873–
7878.
Bargar J. R., Reitmeyer R., Lenhart J. J. and Davis J. A. (2000)
Characterization of U(VI)-carbonato ternary complexes on
hematite: EXAFS and electrophoretic mobility measurements.
Geochim. Cosmochim. Acta 64, 2737–2749.
Barnett M. O., Jardine P. M. and Brooks S. C. (2002) U(VI)
adsorption to heterogeneous subsurface media: application of a
surface complexation model. Environ. Sci. Technol. 36, 937–
942.
Beazley M. J., Martinez R. J., Sobecky P. A., Webb S. M. and
Taillefert M. (2007) Uranium biomineralization as a result of
bacterial phosphatase activity: insights from bacterial isolates
from a contaminated subsurface. Environ. Sci. Technol. 41,
5701–5707.
Beazley M. J., Martinez R. J., Sobecky P. A., Webb S. M. and
Taillefert M. (2009) Nonreductive biomineralization of uranium(VI) phosphate via microbial phosphatase activity in
anaerobic conditions. Geomicrobiol. J. 26, 431–441.
Bostick B. C., Fendorf S., Barnett M. O., Jardine P. M. and Brooks
S. C. (2002) Uranyl surface complexes formed on subsurface
media from DOE facilities. Soil Sci. Soc. Am. J. 66, 99–108.
Boyanov M. I., O’Loughlin E. J., Roden E. E., Fein J. B. and
Kemner K. M. (2007) Adsorption of Fe(II) and U(VI) to
carboxyl-functionalized microspheres: the influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim. Cosmochim. Acta 71, 1898–1912.
Brendel P. J. and Luther, III, G. W. (1995) Development of a gold
amalgam voltammetric microelectrode for the determination of
dissolved Fe, Mn, O2, and S(-II) in porewaters of marine and
freshwater sediments. Environ. Sci. Technol. 29, 751–761.
Bristow G. and Taillefert M. (2008) VOLTINT: a Matlab (R)based program for semi-automated processing of geochemical
data acquired by voltammetry. Comput. Geosci. 34, 153–162.
Brooks S. C. (2001) Waste Characteristics of the Former S-3 Ponds
and Outline of Uranium Chemistry Relevant to NABIR Field
Research Center Studies. NABIR FRC, Technical Report.
Carey E. A. and Taillefert M. (2005) The role of soluble Fe(III) in
the cycling of iron and sulfur in coastal marine sediments.
Limnol. Oceanogr. 50, 1129–1141.
Catalano J. G. and Brown, Jr., G. E. (2004) Analysis of uranylbearing phases by EXAFS spectroscopy: interferences, multiple
scattering, accuracy of structural parameters, and spectral
differences. Am. Mineral. 89, 1004–1021.
Cheng T., Barnett M. O., Roden E. E. and Zhuang J. (2004) Effects
of phosphate on uranium(VI) adsorption to goethite-coated
sand. Environ. Sci. Technol. 38, 6059–6065.
5661
Cheng T., Barnett M. O., Roden E. E. and Zhuang J. (2006) Effects
of solid-to-solution ratio on uranium(VI) adsorption and its
implications. Environ. Sci. Technol. 40, 3243–3247.
Cheng T., Barnett M. O., Roden E. E. and Zhuang J. (2007)
Reactive transport of uranium(VI) and phoshate in a goethitecoated sand column: an experimental study. Chemosphere 68,
1218–1223.
Chinni S., Anderson C., Ulrich K. and Giammar D. (2008) Indirect
UO2 oxidation by Mn(II)-oxidizing spores of Bacillus sp strain
SG-1 and the effect of U and Mn concentrations. Environ. Sci.
Technol. 42, 8709–8714.
Cotton F. A., Wilkinson G., Murillo C. A. and Bochmann M.
(1999) Advanced Inorganic Chemistry. John Wiley & Sons, Inc.,
New York.
De Pablo J., Casas I., Giménez J., Molera M., Rovira M., Duro L.
and Bruno J. (1999) The oxidative dissolution mechanism of
uranium dioxide. I. The effect of temperature in hydrogen
carbonate medium. Geochim. Cosmochim. Acta 63, 3097–
3103.
Diress A. and Lucy C. A. (2005) Study of the selectivity of
inorganic anions in hydro-organic solvents using indirect
capillary electrophoresis. J. Chromatogr. A 1085, 155–
163.
Finneran K. T., Housewright M. E. and Lovley D. R. (2002)
Multiple influences of nitrate on uranium solubility during
bioremediation of uranium-contaminated subsurface sediments. Environ. Microbiol. 4, 510–516.
Fletcher K. E., Boyanov M. I., Thomas S. H., Wu Q., Kemner K.
M. and Löffler F. E. (2010) U(VI) reduction to mononuclear
U(IV) by Desulfitobacterium species. Environ. Sci. Technol. 44,
4705–4709.
Fredrickson J. K., Zachara J. M., Kennedy D. W., Duff M. C.,
Gorby Y. A., Li S.-M. W. and Krupka K. M. (2000) Reduction
of U(VI) in goethite (a-FeOOH) suspensions by a dissimilatory
metal-reducing bacterium. Geochim. Cosmochim. Acta 64,
3085–3098.
Fredrickson J. K., Zachara J. M., Kennedy D. W., Liu C., Duff M.
C., Hunter D. B. and Dohnalkova A. (2002) Influence of Mn
oxides on the reduction of uranium (VI) by the metal-reducing
bacterium Shewanella putrefaciens. Geochim. Cosmochim. Acta
66, 3247–3262.
Fuller C. C., Bargar J. R. and Davis J. A. (2003) Molecular-scale
characterization of uranium sorption by bone apatite materials
for a permeable reactive barrier demonstration. Environ. Sci.
Technol. 37, 4642–4649.
Fuller C. C., Bargar J. R., Davis J. A. and Piana M. J. (2002)
Mechanisms of uranium interactions with hydroxyapatite:
implications for groundwater remediation. Environ. Sci. Technol. 36, 158–165.
Ginder-Vogel M., Stewart B. and Fendorf S. (2010) Kinetic and
mechanistic constraints on the oxidation of biogenic uraninite
by ferrihydrite. Environ. Sci. Technol. 44, 163–169.
Grasshoff K. (1983) Determination of nitrite. In Methods of
Seawater Analysis (eds. K. Grasshoff, M. Ehrhardt and K.
Kremling). Verlag Chemie GmbH, Weinheim.
Hsi C.-K. D. and Langmuir D. (1985) Adsorption of uranyl
onto ferric oxyhydroxides: application of the surface complexation site-binding model. Geochim. Cosmochim. Acta 49,
1931–1941.
Hudson E. A., Allen P. G., Terminello L. J., Denecke M. A. and
Reich T. (1996) Polarized x-ray-absorption spectroscopy of the
uranyl ion: comparison of experiment and theory. Phys. Rev. B
54, 156–165.
Hudson E. A., Terminello L. J., Viani B. E., Denecke M., Reich T.,
Allen P. G., Bucher J. J., Shuh D. K. and Edelstein N. M.
Environ. Sci. Technol. 2010, 44, 8409–8414
Biogenic Formation and Growth of
Uraninite (UO2)
SEUNG YEOP LEE,* MIN HOON BAIK,
AND JONG WON CHOI
Korea Atomic Energy Research Institute, 1045 Daedeok-daero,
Yuseong-gu, Daejeon, Korea
Received June 4, 2010. Revised manuscript received
October 12, 2010. Accepted October 18, 2010.
Biogenic UO2 (uraninite) nanocrystals may be formed as a
product of a microbial reduction process in uranium-enriched
environments near the Earth’s surface. We investigated the
size, nanometer-scale structure, and aggregation state of UO2
formed by iron-reducing bacterium, Shewanella putrefaciens
CN32, from a uranium-rich solution. Characterization of biogenic
UO2 precipitates by high-resolution transmission electron
microscopy (HRTEM) revealed that the UO2 nanoparticles formed
were highly aggregated by organic polymers. Nearly all of
the nanocrystals were networked in more or less 100 nm diameter
spherical aggregates that displayed some concentric UO2
accumulation with heterogeneity. Interestingly, pure UO2
nanocrystals were piled on one another at several positions
via UO2-UO2 interactions, which seem to be intimately related
to a specific step in the process of growing large single
crystals. In the process, calcium that was easily complexed
with aqueous uranium(VI) appeared not to be combined with
bioreduced uranium(IV), probably due to its lower binding energy.
However, when phosphate was added to the system, calcium
was found to be easily associated with uranium(IV), forming
a new uranium phase, ningyoite. These results will extend the
limited knowledge of microbial uraniferous mineralization
and may provide new insights into the fate of aqueous uranium
complexes.
Introduction
Dissimilatory metal-reducing bacteria (DMRB) can couple
the oxidation of organic matter or H2 to the reduction of
oxidized radionuclides. Usually, oxidized uranium(VI) is
much more soluble than the reduced form, uranium(IV),
and typically exists in groundwater as uranyl carbonate
complexes (1–3). Oxidized uranium is readily reduced by
DMRB under anoxic conditions, resulting in the precipitation
of UO2 nanoparticles (4, 5). The rapid rate of oxidized uranium
reduction and the low solubility of the reduced form make
bioremediation an attractive option for removing uranium
from contaminated groundwaters (6–9).
Biogenic UO2 is a fascinating and important nanoscale
biogeological material. Its long-term structural stability is
crucial to the viability of microbial bioremediation strategies (10) that seek to mitigate subsurface uranium contamination. UO2 nanoparticles are potentially highly
mobile because of their small size and can redissolve
quickly if conditions change (11, 12). Size, shape, structure,
degree of crystallinity, and polymer associations all affect
* Corresponding author phone: +82 42 868 4735; fax: +82 42 868
8850; e-mail: [email protected].
10.1021/es101905m
 2010 American Chemical Society
Published on Web 10/27/2010
UO2 solubility, transport in growndwater, and potential
for deposition by sedimentation.
The fate of uranium in natural systems is of great
environmental importance. We report the detailed structure
of aggregated nanobiogenic uraninite produced by Shewanella putrefaciens CN32 in the presence of major cations
of groundwater. A detailed structural observation of biogenic
UO2 with organic polymers has been rarely performed using
direct probes (13–16). The work reported here reveals directly
observed UO2 nanoparticles and their growth within aggregates to provide insight into the fate of biomineralization
products of uranium over longer time scales.
Some uranium ore deposits are believed to involve a direct
microbial reduction process for uranium(VI) (4, 17), as
opposed to an abiotic reduction by reduced species such as
sulfide (18), magnetite (19), and green rust (20). We document
the aggregation of nanoparticles to form submicrometerscale aggregates by organic polymers, and crystal growth
pathways that can lead to morphologies similar to those found
in sedimentary environments. Studies of extracellular nanoparticle growth have been rarely reported for radioactive
chemical-containing minerals. Here, we present in situ the
high-resolution transmission electron microscopy (HRTEM)
characterization for the mineralogy and ultrastructure of
biogenic UO2 formed by an iron-reducing bacterium and
direct evidence that UO2 nanocrytals are grown to larger
sizes by crystallographic attachment.
Experimental Section
S. putrefaciens strain CN32 (ATCC BAA-1097) was obtained
from the American Type Culture Collection (ATCC), U.S. The
S. putrefaciens CN32 was routinely cultured aerobically in a
30 g/L tryptic soy broth (TSB) (Difco Laboratories, Detroit,
MI), and stock cultures were maintained by freezing them
in 40% glycerol at -80 °C.
The aerobically cultured S. putrefaciens CN32 cells were
harvested at mid to late log phase by centrifugating them
from 30 g/L TSB cultures. The cells were centrifuged at 4,000
rpm for 15 min. The supernatant was discarded and the cell
pellets were suspended in a 30 mmol/L NaHCO3 (pH 7) buffer
solution and purged with N2 gas. This process was repeated
four times and washed cells (>4 × 108 cells mL-1) were used
as inoculum.
The NaHCO3 buffer solution (30 mM) was extensively
flushed with N2 to remove dissolved O2. 100 mL of the buffer
solution with lactate (10 mM) as an electron donor were
dispensed into 120 mL serum bottles under N2 condition.
The headspace of the serum bottles was pressurized with
ultrapure nitrogen, then capped with butyl rubber septa and
crimped with an aluminum seal. The bottle and solution
were sterilized by autoclaving at 121 °C and 15 lb/sq for 20
min.
To use background electrolytes similar to groundwaters,
several filtered (0.2 µm, Advantec cellulose acetate) stock
solutions of major cations were aseptically added by syringe
and needle to the serum bottles as soluble forms as follows
(mM): calcium chloride, 1.0; potassium chloride, 1.0; magnesium chloride, 1.0. In addition, P was separately injected
into some of those bottles as a form of sodium hydrogen
phosphate (0.3 mM).
Uranium(VI) stock solutions were prepared by dissolving
a known amount of UO2(NO3)2 · 6H2O (Aldrich) in a previously
acidified HClO4 solution to prevent cation hydrolysis. The
stock solution concentrations were about 1 × 10-3 M, and
uranium(VI) (5 × 10-5 M) was aseptically added using a
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FIGURE 6. The microbial UO2 formation process from uranyl
carbonate complexes and the phosphate effect.
FIGURE 5. A schematic illustration of the UO2 aggregate based
on oriented attachment.
(6, 33) to individual UO2 nanoparticles ranging in size from
0.9 to 5 nm (6, 34, 35) to large aggregates of more than 100
nm. Interaction between nanoparticles and organic molecules is an important step in the stage of initial UO2 growth
because nanometer-size particles can remain firmly attached
without dispersing in aqueous flow environments. Organic
polymers appear to play an important role as fundamental
bases for organizing the three-dimensional framework of UO2.
As shown in Figure 5, the aggregated nanoparticles remain
well defined, in random but oriented attachment, with
bonding at the interfaces. Aggregation ensures small distances
between the nanoparticles, providing the possibility of
oriented attachment of nanoparticles. The illustration depicts
a group of several attached particles with clear grain edges,
showing interfaces between primary particles. Although the
UO2 particle-particle junctions are not initially perfect, their
very short-range UO2 interactions will be crucial for enhancing the intimate binding of their interfaces. Following this
fundamental particle cohesion, pure UO2 may further
outgrow via subsequent attachment and stepwise stacking
of discrete particles.
Biomineralization of Uranyl Carbonate Complexes.
Incorporation of groundwater-dissolved cations into the uranium phase will be critical to predicting uranium-bearing
nanoparticle stability and growth in the environment. Most
equilibrium speciation models predict that the dominant
uranium aqueous species in groundwater will be uranyl
carbonate complexes (36, 37). Generally, Ca-U-CO3 complexes
(CaUO2(CO3)32-, Ca2UO2(CO3)3) have been proposed to play an
important role in the environmental chemistry of uranium (8).
Calcium is usually a common ion in groundwater that can easily
complex with uranium(VI) in bicarbonate solutions.
Unfortunately, for the above reasons, the bioreduction
rate of aqueous uranium(VI) with calcium was slower than
that of uranium(VI) without calcium under the same cell
concentration (1). The calcium caused a significant decrease
in the rate and extent of bacterial uranium(VI) reduction
(2, 3, 8). Interestingly, in our study when uranium(VI) with
calcium was reduced to uranium(IV), the calcium did not
consistently combine with uranium(IV) in the formation of
a uranium phase. The calcium appears to be neglected from
the UO2 nucleation process. This means that calcium is no
longer complexed with uranium(IV) when uranium(VI) is
bioreduced to uranium(IV), probably due to the lower binding
energy of calcium for uranium(IV). This was confirmed by
the result that the uranium solid-phase precipitated from
aqueous uranium complexes had little calcium within its
structure (Figure 2A). However, there have been some reports
that calcium was frequently found in some UO2 ore deposits
as impurities. In such cases, the calcium seems to exist
together with uranium(VI), not with the reduced form(IV).
Uranium(VI) is frequently incorporated into the UO2 mineral
formation as considerable amount in natural condition (38).
Actually, pure UO2.0 has been rarely observed in nature, and
it has been recognized that uraninites formed in the field are
likely to contain much of uranium(VI) and calcium.
In spite of the above result, when phosphate was added
to the system, calcium was found to be easily associated
with uranium(IV) forming a new uranium phase (ningyoite)
(Figure 6). The intimate relationship between uranium and
calcium appears to be maintained by phosphate, even though
uranium(VI) was changed to uranium(IV) during microbial
reduction. We suppose that the separation of calcium from
uranium(IV) is retarded by the adhesion of phosphate, which
can promote their coprecipitation regardless of the uranium
oxidation state. However, if aqueous uranium(VI) or phosphate concentrations were elevated to mmole/L level, an
abiotic precipitation of uranium(VI) phosphate phases could
occur due to their rapid coprecipitation (21) without affordable bioreduction. Nevertheless, a microbial reduction
(bio-transformation) for the solid-phase U(VI) to U(IV) can
occur, but it may need much more time and energies (21, 39).
As amendments of backfill material for radioactive waste
storage, phosphate is currently considered one of the most
important candidate chemicals (40). Considering this situation, the microbial phosphatic uranium(IV) phase, due to
its extremely low solubility under circumneutral pH conditions (21), might play an important role in lowering uranium
mobility in uranium-rich environments.
Acknowledgments
We thank Dr. Wooyong Um for his helpful discussions and
reviews on this manuscript. We appreciate the comments
from three anonymous reviewers. This work was supported
by the Nuclear Research and Development Program of
National Research Foundation of Korea (NRF) funded by the
Ministry of Education, Science and Technology (MEST).
Literature Cited
(1) Liu, C.; Jeon, B. H.; Zachara, J. M.; Wang, Z. Influence of calcium
on microbial reduction of solid phase uranium(VI). Biotechnol.
Bioeng. 2007, 97, 1415–1422.
(2) Neiss, J.; Stewart, B. D.; Nico, P. S.; Fendorf, S. Speciationdependent microbial reduction of uranium within iron-coated
sands. Environ. Sci. Technol. 2007, 41, 7343–7348.
VOL. 44, NO. 22, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
8413
(3) Stewart, B. D.; Neiss, J.; Fendorf, S. Quantifying constraints
imposed by calcium and iron on bacterial reduction of
uranium(VI). J. Environ. Qual. 2007, 36, 363–372.
(4) Lovley, D. R.; Phillips, E. J. P.; Gorby, Y. A.; Landa, E. R. Microbial
reduction of uranium. Nature 1991, 350, 413–416.
(5) Gorby, Y. A.; Lovley, D. R. Enzymatic uranium precipitation.
Environ. Sci. Technol. 1992, 26, 205–207.
(6) Suzuki, Y.; Kelly, S. D.; Kemner, K. M.; Banfield, J. F. Nanometresize products of uranium bioreduction. Nature 2002, 419, 134.
(7) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Liu, C.; Duff,
M. C.; Hunter, D. B.; Dohnalkova, A. Influence of Mn oxides on
the reduction of uranium(VI) by the metal-reducing bacterium
Shewanella putrefaciens. Geochim. Cosmochim. Acta 2002, 66,
3247–3262.
(8) Brooks, S. C.; Fredrickson, J. K.; Carroll, S. L.; Kennedy, D. W.;
Zachara, J. M.; Plymale, A. E.; Kelly, S. D.; Kemner, K. M.; Fendorf,
S. Inhibition of Bacterial U(VI) reduction by calcium. Environ.
Sci. Technol. 2003, 37, 1850–1858.
(9) Burgos, W. D.; McDonough, J. T.; Senko, J. M.; Zhang, G.;
Dohnalkova, A. C.; Kelly, S. D.; Gorby, Y.; Kemner, K. M.
Characterization of uraninite nanoparticles produced by Shewanella oneidensis MR-1. Geochim. Cosmochim. Acta 2008,
72, 4901–4915.
(10) Wielinga, B.; Bostick, B.; Hansel, C. M.; Rosenzweig, R. F.;
Fendorf, S. Inhibition of bacterially promoted uranium reduction: Ferric (hydr)oxides as competitive electron acceptors.
Environ. Sci. Technol. 2000, 34, 2190–2195.
(11) Ulrich, K. A.; Singh, A.; Schofield, E. J.; Bargar, J. R.; Veeramani,
H.; Sharp, J. O.; Bernier-Latmani, R.; Giammar, D. E. Dissolution
of biogenic and synthetic UO2 under varied reducing conditions.
Environ. Sci. Technol. 2008, 42, 5600–5606.
(12) Ulrich, K. A.; Ilton, E. S.; Veeramani, H.; Sharp, J. O.; BernierLatmani, R.; Schofield, E. J.; Bargar, J. R.; Giammar, D. E.
Comparative dissolution kinetics of biogenic and chemogenic
uraninite under oxidizing conditions in the presence of carbonate. Geochim. Cosmochim. Acta 2009, 73, 6065–6083.
(13) Fayek, M.; Utsunomiya, S.; Pfiffner, S. M.; White, D. C.; Riciputi,
L. R.; Ewing, R. C.; Anovitz, L. M.; Stadermann, F. J. The
application of HRTEM techniques and nanosims to chemically
and isotopically characterize Geobacter sulfurreducens surfaces.
Can. Miner. 2005, 43, 1631–1641.
(14) Schofield, E. J.; Veeramani, H.; Sharp, J. O.; Suvorova, E.; BernierLatmani, R.; Mehta, A.; Stahlman, J.; Webb, S. M.; Clark, D. L.;
Conradson, S. D.; Ilton, E. S.; Bargar, J. R. Structure of biogenic
uraninite produced by Shewanella oneidensis strain MR-1.
Environ. Sci. Technol. 2008, 42, 7898–7904.
(15) Marshall, M. J.; Dohnalkova, A. C.; Kennedy, D. W.; Plymale,
A. E.; Thomas, S. H.; Löffler, F. E.; Sanford, R. A.; Zachara, J. M.;
Fredrickson, J. K.; Beliaev, A. S. Electron donor-dependent
radionuclide reduction and nanoparticle formation by Anaeromyxobacter dehalogenans strain 2CP-C. Environ. Microbiol.
2009, 11, 534–543.
(16) Singer, D. M.; Farges, F.; Brown Jr, G. E. Biogenic nanoparticulate
UO2: Synthesis, characterization, and factors affecting surface
reactivity. Geochim. Cosmochim. Acta 2009, 73, 3593–3611.
(17) Mohagheghi, A.; Updegraff, D. M.; Goldhaber, M. B. The role
of sulfate-reducing bacteria in the deposition of sedimentary
uranium ores. Geomicrobiol. J. 1985, 4, 153–173.
(18) Hua, B.; Xu, H. F.; Terry, J.; Deng, B. L. Kinetics of uranium(VI)
reduction by hydrogen sulfide in anoxic aqueous systems.
Environ. Sci. Technol. 2006, 40, 4666–4671.
(19) Scott, T. B.; Allen, G. C.; Heard, P. J.; Randell, M. G. Reduction
of U(VI) to U(IV) on the surface of magnetite. Geochim.
Cosmochim. Acta 2005, 69, 5639–5646.
(20) O’Loughlin, E. J.; Kelly, S. D.; Cook, R. E.; Csencsits, R.; Kemner,
K. M. Reduction of uranium(VI) by mixed iron(II)/iron(III)
hydroxide (green rust): Formation of UO2 nanoparticles. Environ.
Sci. Technol. 2003, 37, 721–727.
(21) Khijniak, T. V.; Slobodkin, A. I.; Coker, V.; Renshaw, J. C.; Livens,
F. R.; Bonch-Osmolovskaya, E. A.; Birkeland, N. K.; MedvedevaLyalikova, N. N.; Lloyd, J. R. Reduction of uranium(VI) phosphate
during growth of the thermophilic bacterium Thermoterrabac-
8414
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 22, 2010
(22)
(23)
(24)
(25)
(26)
(27)
(28)
(29)
(30)
(31)
(32)
(33)
(34)
(35)
(36)
(37)
(38)
(39)
(40)
terium ferrireducens. Appl. Environ. Microbiol. 2005, 71, 6423–
6426.
Marshall, M. J.; Beliaev, A. S.; Dohnalkova, A. C.; Kennedy, D. W.;
Shi, L.; Wang, Z.; Boyanov, M. I.; Lai, B.; Kemner, K. M.; McLean,
J. S.; Reed, S. B.; Culley, D. E.; Bailey, V. L.; Simonson, C. J.;
Saffarini, D. A.; Romine, M. F.; Zachara, J. M.; Fredrickson, J. K.
c-Type cytochrome-dependent formation of U(IV) nanoparticles
by Shewanella oneidensis. PLoS Biol. 2006, 4, 1324–1333.
Shelobolina, E. S.; Coppi, M. V.; Korenevsky, A. A.; DiDonato,
L. N.; Sullivan, S. A.; Konishi, H.; Xu, H.; Leang, C.; Butler, J. E.;
Kim, B. C.; Lovley, D. R. Importance of c-Type cytochromes for
U(VI) reduction by Geobacter sulfurreducens. BMC Microbiol.
2007, 7, 16.
Sogaard, E. G.; Medenwaldt, R.; Abraham-Peskir, J. V. Conditions
and rates of biotic and abiotic iron precipitation in selected
Danish freshwater plants and microscopic analysis of precipitate
morphology. Water Res. 2000, 34, 2675–2682.
Mavrocordatos, D.; Fortin, D. Quantitative characterization of
biotic iron oxides by analytical electron microscopy. Am.
Mineral. 2002, 87, 940–946.
Allain, C.; Cloitre, M.; Parisse, F. Settling by cluster deposition
in aggregating colloidal suspensions. J. Colloid Interface Sci.
1996, 178, 411–416.
Bradford, S. A.; Yates, S. R.; Bettehar, M.; Simunek, J. Physical
factors affecting the transport and fate of colloids in saturated
porous media. Water Resour. Res. 2002, 38, 1327–1340.
Stumm, W.; Morgan, J. J. Aquatic Chemistry: Chemical Equilibria
and Rates in Natural Waters; Wiley: New York, 1996; pp 400414.
Moreau, J. W.; Weber, P. K.; Martin, M. C.; Gilbert, B.; Hutcheon,
I. D.; Banfield, J. F. Extracellular proteins limit the dispersal of
biogenic nanoparticles. Science 2007, 316, 1600–1603.
Penn, R. L.; Banfield, J. F. Oriented attachment and growth,
twinning, polytypism, and formation of metastable phases:
insights from nanocrystalline TiO2. Am. Mineral. 1998, 83, 1077–
1082.
Huang, F.; Zhang, H.; Banfield, J. F. Two-stage crystal-growth
kinetics observed during hydrothermal coarsening of nanocrystalline ZnS. Nano Lett. 2003, 3, 373–378.
Huang, F.; Gilbert, B.; Zhang, H.; Banfield, J. F. Reversible,
surface-controlled structure transformation in nanoparticles
induced by an aggregation state. Phys. Rev. Lett. 2004, 92, 155501.
Boyanov, M. I.; O’Loughlin, E. J.; Roden, E. E.; Fein, J. B.; Kemner,
K. M. Adsorption of Fe(II) and U(VI) to carboxyl-functionalized
microspheres: the influence of speciation on uranyl reduction
studies by titration and XAFS. Geochim. Cosmochim. Acta 2007,
71, 1898–1912.
Sani, R. K.; Peyton, B. M.; Dohnalkova, A.; Amonette, J. E.
Reoxidation of reduced uranium with iron(III) (hydr)oxides
under sulfate-reducing conditions. Environ. Sci. Technol. 2005,
39, 2059–2066.
Senko, J. M.; Kelly, S. D.; Dohnalkova, A. C.; McDonough, J. T.;
Kemner, K. M.; Burgos, W. D. The effect of U(VI) bioreduction
kinetics on subsequent reoxidation of biogenic U(IV). Geochim.
Cosmochim. Acta 2007, 71, 4644–4654.
Clark, D. L.; Hobart, D. E.; Neu, M. P. Actinide carbonte
complexes and their importance in actinide environmental
chemistry. Chem. Rev. 1995, 95, 25–48.
Abdelouas, A.; Lu, Y.; Lutze, W.; Nuttall, H. E. Reduction of
U(VI) to U(IV) by indigenous bacteria in contaminated ground
water. J. Contam. Hydrol. 1998, 35, 217–233.
Janeczek, J.; Ewing, R. C. Structural formula of uraninite. J. Nucl.
Mater. 1992, 190, 128–132.
Ortiz-Bernad, I.; Anderson, R. T.; Vrionis, H. A.; Lovley, D. R.
Resistance of solid-phase U(VI) to microbial reduction during
in situ bioremediation of uranium-contaminated groundwater.
Appl. Environ. Microbiol. 2004, 70, 7558–7560.
Oelkers, E. H.; Montel, J. M. Phosphates and nuclear waste
storage. Elements 2008, 4, 113–116.
ES101905M
A GEOCHEMICAL INVESTIGATION OF HETEROGENEOUS REDOX
REACTIONS BETWEEN FE(II), FE(III), AND URANIUM
by
Drew Eric Latta
A thesis submitted in partial fulfillment
of the requirements for the Doctor of
Philosophy degree in Civil and Environmental Engineering
in the Graduate College of
The University of Iowa
December 2010
Thesis Supervisor: Professor Michelle M. Scherer
159
that of the UVI standard, and contains a post-edge feature indicative of UVI in the uranyl
(UO22+) geometry (vertical arrow), indicating that all the U associated with the solid
phase remained oxidized as UVI (Figure 6.4). The edge position of the U XANES spectra
of all the samples containing Fe(II) is near the UIV standard, and lacks the post-edge
feature indicative of UVI, suggesting that nearly all of the added U has been reduced to
UIV. The EXAFS spectra of the goethite and Al-goethite reacted with U in the presence of
Fe(II) indicate the reduced U product is consistent nanoparticulate uraninite (UIVO2) with
a spectrum close to that of biogenically produced nanoparticulate uraninite (Figure 6.5)
(47).
Formation of stable Fe(II) species on goethite (see discussion in Chapter 5) may
have promoted UVI reduction in the goethite + Fe(II) system. We note that in a study
using insulating beads functionalized with carboxylate groups capable of binding U and
Fe(II), that formation of Fe(II) polymers was required for the reduction of UVI to UIV
(112). In contrast, some UVI reduction has been noted in systems were Fe(II)
concentrations were less than what would be expected to cause surface saturation of
Fe(II). In addition, total Fe(II) loading in that system was less than that required to reduce
all the added UVI (57). Currently, we cannot conclude whether UVI reduction might also
be mediated by electron conduction through the bulk of goethite. Further study of this
mechanism is warranted.
Isotope Exchange between Fe(II) and Goethite
We have begun to develop a method to measure isotope exchange between
goethite and aqueous Fe(II) using a quadrupole-ICP-MS and highly enriched 57Fe(II)
solutions exposed to goethite with natural isotopic composition. We have started with
determining whether 56Fe and 57Fe could be determined in mixtures using the q-ICP-MS
using highly enriched isotope solutions (Figure 6.6). We have used spiked 2 g L-1
goethite suspensions with a natural abundance of Fe isotopes (5.8% 54Fe, 91.8% 56Fe,
183
107. Hansen, H. C. B., Composition, stabilization, and light-absorption of Fe(II)Fe(III)
hydroxy-carbonate (green rust). Clay Minerals 1989, 24, (4), 663-669.
108.
Gustafsson, J. P. Visual MINTEQ, 2.51; 2006.
109. Teixeira, L. S. G.; Costa, A. C. S.; Ferreira, S. L. C.; Freitas, M. D.; de Carvalho,
M. S., Spectrophotometric determination of uranium using 2-(2-thiazolylazo)-p-cresol
(TAC) in the presence of surfactants. Journal of the Brazilian Chemical Society 1999, 10,
(6), 519-522.
110. Hua, B.; Xu, H. F.; Terry, J.; Deng, B. L., Kinetics of uranium(VI) reduction by
hydrogen sulfide in anoxic aqueous systems. Environmental Science & Technology 2006,
40, (15), 4666-4671.
111. Debeer, H.; Coetzee, P. P., Ion Chromatographic Separation and
Spectrophotometric Determination of U(Iv) and U(Vi). Radiochimica Acta 1992, 57, (23), 113-117.
112. Boyanov, M. I.; O'Loughlin, E. J.; Roden, E. E.; Fein, J. B.; Kemner, K. M.,
Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The influence
of speciation on uranyl reduction studied by titration and XAFS. Geochimica Et
Cosmochimica Acta 2007, 71, (8), 1898-1912.
113. Kelly, S. D.; Kemner, K. M.; Fein, J. B.; Fowle, D. A.; Boyanov, M. I.; Bunker,
B. A.; Yee, N., X-ray absorption fine structure determination of pH-dependent Ubacterial cell wall interactions. Geochimica Et Cosmochimica Acta 2002, 66, (22), 38553871.
114. Webb, S. M., SIXpack: a graphical user interface for XAS analysis using
IFEFFIT. Physica Scripta 2005, T115, 1011-1014.
115. Hansen, H. C. B.; Guldberg, S.; Erbs, M.; Koch, C. B., Kinetics of nitrate
reduction by green rusts - effects of interlayer anion and Fe(II):Fe(III) ratio. Applied Clay
Science 2001, 18, 81-91.
116. Lazaridis, N. K.; Asouhidou, D. D., Kinetics of sorptive removal of
chromium(VI) from aqueous solutions by calcined Mg-Al-CO3 hydrotalcite. Water
Research 2003, 37, (12), 2875-2882.
117. Litten, G. R., Quality of ground water used for selected municipal water supplies
in Iowa, 1997-2002 water years: USGS Open File Report 2004-1048. In USGS, Ed.
2004; p 36.
Available online at www.sciencedirect.com
Geochimica et Cosmochimica Acta 74 (2010) 1–15
www.elsevier.com/locate/gca
Role of extracellular polymeric substances in metal
ion complexation on Shewanella oneidensis: Batch
uptake, thermodynamic modeling, ATR-FTIR,
and EXAFS study
Juyoung Ha a,*, Alexandre Gélabert a,1, Alfred M. Spormann b,
Gordon E. Brown Jr. a,b,c
a
Surface & Aqueous Geochemistry Group, Department of Geological & Environmental Sciences, Stanford University, Stanford,
CA 94305-2115, USA
b
Department of Chemical Engineering, Stanford University, Stanford, CA 94305, USA
c
Stanford Synchrotron Radiation Lightsource, SLAC National Accelerator Laboratory, MS 69, 2575 Sand Hill Road, Menlo Park,
CA, 94025, USA
Received 10 September 2008; accepted in revised form 25 June 2009; available online 10 July 2009
Abstract
The effect of cell wall-associated extracellular polymeric substances (EPS) of the Gram-negative bacterium Shewanella
oneidensis strain MR-1 on proton, Zn(II), and Pb(II) adsorption was investigated using a combination of titration/batch
uptake studies, surface complexation modeling, attenuated total reflectance – Fourier transform infrared (ATR-FTIR) spectroscopy, and Zn K-edge extended X-ray absorption fine structure (EXAFS) spectroscopy. Both unmodified (wild-type (WT)
strain) and genetically modified cells with inhibited production of EPS (DEPS strain) were used. Three major types of functional groups (carboxyl, phosphoryl, and amide groups) were identified in both strains using ATR-FITR spectroscopy. Potentiometric titration data were fit using a constant capacitance model (FITEQL) that included these three functional groups.
The fit results indicate less interaction of Zn(II) and Pb(II) with carboxyl and amide groups and a greater interaction with
phosphoryl groups in the DEPS strain than in the WT strain. Results from Zn(II) and Pb(II) batch adsorption studies and
surface complexation modeling, assuming carboxyl and phosphoryl functional groups, also indicate significantly lower Zn(II)
and Pb(II) uptake and binding affinities for the DEPS strain. Results from Zn K-edge EXAFS spectroscopy show that Zn(II)
bonds to phosphoryl and carboxyl ligands in both strains. Based on batch uptake and modeling results and EXAFS spectral
analysis, we conclude that the greater amount of EPS in the WT strain enhances Zn(II) and Pb(II) uptake and hinders diffusion of Zn(II) to the cell walls relative to the DEPS strain.
Ó 2009 Published by Elsevier Ltd.
1. INTRODUCTION
*
Corresponding author. Tel.: +1 650 723 7513; fax: +1 650 725
2199.
E-mail address: [email protected] (J. Ha).
1
Present address: Department of Earth Sciences, University of
Paris 7, IMPMC, IPGP, CNRS, UMR 7590, F-75015 Paris,
France.
0016-7037/$ - see front matter Ó 2009 Published by Elsevier Ltd.
doi:10.1016/j.gca.2009.06.031
With an estimated biomass close to the total amount of
carbon in plants (Newman, 2001), bacteria play a major
role in the sequestration, transformation, and cycling of
contaminant metal ions in soils and aqueous environments
(Daughney and Fein, 1998; Yee and Fein, 2001). Bacterial
sequestration of metal and metalloid ions is mainly the result of sorption and biomineralization processes (Beveridge
2
J. Ha et al. / Geochimica et Cosmochimica Acta 74 (2010) 1–15
and Fyfe, 1985; Urrutia and Beveridge, 1993; Fein et al.,
2001), and various thermodynamic models of metal ion
interactions with bacteria in planktonic forms have been
proposed (e.g., Daughney and Fein, 1998; Fein, 2000; Haas
et al., 2001; Claessens et al., 2004; Claessens and Van Cappellen, 2007; Fein, 2007). Such models are very useful, but
they are inherently macroscopic (Sposito, 1986) and are not
capable of providing molecular-level information on metal
ion reaction products associated with bacterial cell walls.
Information on adsorption site identities, surface complex
geometries and stoichiometries, and types of biominerals
associated with bacteria can, however, be derived from
spectroscopic or scattering studies (e.g., Sarret et al.,
1998a; Webb et al., 2001; Kelly et al., 2002; Panak et al.,
2002; Boyanov et al., 2003; Jiang et al., 2004; Guine
et al., 2006).
Among the various reactive components associated with
bacterial cell walls, extracellular polymeric substance (EPS)
is of particular importance (Beveridge and Fyfe, 1985; Slaveykova and Wilkinson, 2002). During cell growth, the outer-membrane composition evolves with respect to protein
and EPS concentrations, and such changes often result in
different microenvironments around the cells relative to
the bulk environment (Boyanov et al., 2007). It is well
known that EPS affects biofilm formation (e.g., Matsukawa
and Greenberg, 2004; Thormann et al., 2006), cell adhesion
to solid substrates (e.g., Bruinsma et al., 2001; Walker and
Chen, 2006), and local reactive site densities in cells, resulting in high antibiotic resistance in biofilms (e.g., Costerton
et al., 1995). In addition, a number of recent studies have
reported enhanced reactivity of planktonic cells due to the
presence of EPS based on proton and metal ion sorption
using either isolated bacterial EPS (Guibaud et al., 2003;
Comte et al., 2006; Lamelas et al., 2006) or chemically
and/or physically treated cells that are free of EPS
(Merroun et al., 2003; Toner et al., 2005; Tourney et al.,
2008). Liu et al. (2007) used genetically mutated EPS-free
cells to study the transport behavior of cells in porous media containing metal ions, but the role of EPS in metal ion or
proton sorption was not determined in detail. To the best of
our knowledge, there have been no quantitative studies of
the role of EPS in proton and metal ion sorption reactions
using genetically modified cells.
In the present study, Shewanella oneidensis MR-1 wildtype (WT) strain and a genetically modified strain of S.
oneidensis MR-1 in which EPS production has been inhibited (designated here as DEPS; see Thormann et al., 2006)
were used to characterize the functional role of EPS in
metal ion and proton uptake on S. oneidensis. Among
Gram-negative bacteria, S. oneidensis MR-1 is of particular
interest because it lives under both anaerobic and aerobic
conditions, utilizes a vast array of electron acceptors, plays
an important biogeochemical role in metal reduction, and
affects the global cycling of metals (Zachara et al., 1998; Zachara et al., 2001; Neal et al., 2003). The divalent cations
chosen for study, Zn(II) and Pb(II), are important environmental contaminants that are not redox sensitive. In addition, there have been a number of extended X-ray
absorption fine structure (EXAFS) spectroscopic studies
of the interactions of Zn(II) and Pb(II) with mineral sur-
faces (see Brown and Sturchio, 2002 for a review; Juillot
et al., 2003; Ha et al., 2009), natural organic matter (e.g.,
Sarret et al., 1997; Xia et al., 1997a,b; Karlsson and Skyllberg, 2007), plants (e.g., Sarret et al., 1998a; Sarret et al.,
2002), cell walls of bacteria in planktonic form (e.g., Webb
et al., 2001; Guine et al., 2006; Guine et al., 2007; Claessens
and Van Cappellen, 2007), microbial biofilms on mineral
surfaces (e.g., Templeton et al., 2001; Templeton et al.,
2003; Toner et al., 2005), and diatoms (e.g., Gélabert
et al., 2004; Pokrovsky et al., 2005), which provide a basis
for the present study.
In this study, we have investigated the types of surface
functional groups in cell walls and associated EPS and
the nature of H+, Zn(II), and Pb(II) binding on S. oneidensis MR-1 WT and DEPS surfaces at the macroscopic level
using sorption isotherms and a constant capacitance model
and at the molecular-level using attenuated total reflectance
– Fourier transform infrared (ATR-FTIR) and Zn K-edge
EXAFS spectroscopy. Our objectives in this study are to (1)
quantify the concentrations and protonation/deprotonation constants of the major cell membrane functional
groups present in S. oneidensis MR-1 WT and DEPS
strains; (2) characterize the interaction of Zn(II) and Pb(II)
with S. oneidensis at the molecular-level, and (3) quantify
their binding constants for the two different strains of S.
oneidensis, with the overall aim of understanding the effect(s) of reduced amounts of EPS on Zn(II) and Pb(II)
uptake.
2. MATERIALS AND METHODS
2.1. Bacterial cell growth and stock suspension preparation
Both WT and DEPS strains of S. oneidensis MR-1 were
cultivated under aerobic conditions in Luria broth (LB) at
pH 7.4 and 30 °C by shaking at 200 rpm until the late exponential growth phase was reached (22 h). The cells were
then centrifuged at 5000 rpm (3000g) and washed three
times with 0.01 M NaNO3 solution in order to remove excess ions from the cells and LB solutions. After the final
wash, the wet bacterial pellet was weighed and immediately
used for potentiometric titration and electrophoretic mobility measurements and metal ion adsorption experiments.
Stock bacterial solutions were prepared for S. oneidensis
WT and DEPS strains using 3 g (wet weight) of bacteria,
which is equivalent to 3 109 cells mL1, per 1 L of 1,
0.1, or 0.01 M NaNO3 depending on the experimental
objective, and bubbled with N2 for 30 min prior to further
treatments to remove CO2 from the solution.
2.2. ATR-FTIR spectroscopic characterization of cells
After centrifugation to remove the cells from the growth
medium and washing three times with 0.01 M NaNO3, the
wet pastes of cells were deposited on a germanium ATR
crystal in a Nicolet FT-IR spectrometer (NEXUS470),
equipped with a mercury cadmium telluride (MCT) detector and a horizontal attenuated total reflectance attachment
(germanium crystal). The sample holding region was sealed
with a lid to prevent evaporation during measurements. A
Role of extracellular polymeric substances in metal ion complexation on Shewanella oneidensis
Boyanov M. I., O’Loughlin E. J., Roden E. E., Fein J. B. and
Kemner K. M. (2007) Adsorption of Fe(II) and U(VI) to
carboxyl-functionalized microspheres: the influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim. Cosmochim. Acta 71, 1898–1912.
Brown, Jr., G. E. and Sturchio N. C. (2002) An overview of
synchrotron radiation applications to low temperature geochemistry and environmental science. Rev. Mineral. Geochem.
49, 1–115.
Bruinsma G. M., Rustema-Abbing M., van der Mei H. C. and
Busscher H. J. (2001) Effects of cell surface damage on surface
properties and adhesion of Pseudomonas aeruginosa. J. Microbiol. Methods 45, 95–101.
Burnett P.-G. G., Heinrich H., Peak D., Bremer P. J., McQuillan J.
and Daughney C. J. (2006) The effect of pH and ionic strength
on proton adsorption by the thermophilic bacterium Anoxybacillus flavithermus. Geochim. Cosmochim. Acta 70, 1914–1927.
Claessens J., Behrends T. and Van Cappellen P. (2004) What do
acid–base titrations of live bacteria tell us? A preliminary
assessment. Aquat. Sci. 66, 19–26.
Claessens J. and Van Cappellen P. (2007) Competitive binding of
Cu2+ and Zn2+ to live cells of Shewanella putrefaciens. Environ.
Sci. Technol. 41, 909–914.
Comte S., Guibaud G. and Baudu M. (2006) Biosorption properties of extracellular polymeric substances (EPS) resulting from
activated sludge according to their type: soluble or bound.
Process Biochem. 41, 815–823.
Costerton J. W., Lewandowski Z., Caldwell D. E., Korber D. R.
and Lappin-Scott H. M. (1995) Microbial biofilms. Ann. Rev.
Microbiol. 49, 711–745.
Daughney C. J. and Fein J. B. (1998) The effect of ionic strength on
the adsorption of H+, Cd2+, Pb2+, and Cu2+ by Bacillus subtilis
and Bacillus licheniformis: a surface complexation model. J.
Colloid Interface Sci. 198, 53–77.
Daughney C. J., Fein J. B. and Yee N. (1998) A comparison of the
thermodynamics of metal adsorption onto two common bacteria. Chem. Geol. 144, 161–176.
Daughney C. J., Fowle D. A. and Fortin D. (2001) The effect of
growth phase on proton and metal adsorption by Bacillus
subtilis. Geochim. Cosmochim. Acta 65, 1025–1035.
Doonan B. B., Crang R. E., Jensen T. E. and Baxter M. (1979) In
situ X-ray energy dispersive microanalysis of polyphosphate
bodies in Aureobasidium pullulans. J. Ultrastruc. Res. 69, 232–
238.
Effenberger H., Mereiter K. and Zemann J. (1982) Crystal structure
refinements of magnesite, calcite, rhodochrosite, siderite, smithonite, and dolomite, with the discussion of some aspects of the
stereochemistry of calcite type carbonates. Z. Kristallogr. 156,
233–243.
Fein J. B. (2000) Quantifying the effects of bacteria on adsorption
reactions in water-rock systems. Chem. Geol. 169, 265–280.
Fein J. B. (2007) Thermodynamic modeling of metal adsorption
onto bacterial cell walls: current challenges. Adv. Agron. 90,
179–202.
Fein J. B., Daughney C. J., Yee N. and Davis T. A. (1997) A
chemical equilibrium model for metal adsorption onto bacterial
surfaces. Geochim. Cosmochim. Acta 61, 3319–3328.
Fein J. B., Martin A. M. and Wightman P. G. (2001) Metal
adsorption onto bacterial surfaces: development of a predictive
approach. Geochim. Cosmochim. Acta 65, 4267–4273.
Felmy A. R., Girvin D. C. and Jenne E. A. (1984) MINTEQ—A
Computer Program for Calculating Aqueous Geochemical Equilibria. US Environmental Protection Agency, Athens, GA.
Gélabert A., Pokrovsky O. S., Schott J., Boudou A., Feurtet-Mazel
A., Mielczarski J., Mielczarski E., Mesmer-Dudons N. and
Spalla O. (2004) Study of diatom/aqueous solution interface. I.
13
Acid–base equilibria and spectroscopic observation of freshwater and marine species. Geochim. Cosmochim. Acta 68, 4039–
4058.
Gonzalez-Davila M. (1995) The role of phytoplankton cell on the
control of heavy metal concentrations in seawater. Marine
Chem. 48, 215–236.
Guibaud G., Tixier N., Bouju A. and Baudu M. (2003) Relation
between extracellular polymers’ composition and its ability to
complex Cd, Cu and Pb. Chemosphere 52, 1701–1710.
Guine V., Spadini L., Sarret G., Muris M., Delolme C., Gaudet J.
P. and Martins J. M. F. (2006) Zinc sorption to three Gramnegative bacteria: combined titration, modeling, and EXAFS
study. Environ. Sci. Technol. 40, 1806–1813.
Guine V., Martins J. M. F., Causse B., Durand A., Gaudet J. P.
and Spadini L. (2007) Effect of cultivation and experimental
conditions on the surface reactivity of the metal-resistant
bacteria Cupriavidus metallidurans CH34 to protons, cadmium
and zinc. Chem. Geol. 236, 266–280.
Ha J., Trainor T. P., Farges F. and Brown, Jr., G. E. (2009)
Interaction of aqueous Zn(II) with hematite nanoparticles and
microparticles: part 1. EXAFS study of Zn(II) adsorption and
precipitation. Langmuir 25, 5574–5585.
Haas J. R., Dichristina T. J. and Wade, Jr., R. (2001) Thermodynamics of U(VI) sorption onto Shewanella putrefaciens.
Chem. Geol. 180, 33–54.
Harding M. M. (1999) The geometry of metal–ligand interactions
relevant to proteins. Acta Crystallogr., Sect. D 55, 1432–1443.
Herbelin A. L. and Westall J. C. (1996) FITEQL Version 3.2, A
Computer Program for Determination of Chemical Equilibrium
Constants from Experimental Data. Dept. of Chemistry, Oregon
State University, Corvallis, OR.
Heinrich H. T. M., Bremer P. J., McQuillan A. J. and Daughney C.
J. (2008) Modeling of the acid–base properties of two thermophilic bacteria at different growth times. Geochim. Cosmochim.
Acta 72, 4185–4200.
Jiang W., Saxena A., Song B., Ward B. B., Beveridge T. J. and
Myneni S. C. B. (2004) Elucidation of functional groups on
Gram-positive and Gram-negative bacterial surfaces using
infrared spectroscopy. Langmuir 20, 11433–11442.
Juillot F., Morin G., Ildefonse P., Trainor T. P., Benedetti M.,
Galoisy L., Calas G. and Brown, Jr., G. E. (2003) Occurrence
of Zn/Al hydrotalcite in smelter-impacted soils from northern
France:eEvidence from EXAFS spectroscopy and chemical
extractions. Am. Mineral. 88, 509–526.
Kang S. Y., Bremer P. J., Kim K. W. and McQuillan A. J. (2006)
Monitoring metal ion binding in single-layer Pseudomonas
aeruginosa biofilms using ATR-IR spectroscopy. Langmuir 22,
286–291.
Karlsson T. and Skyllberg U. (2007) Complexation of zinc in
organic soils – EXAFS evidence for sulfur associations.
Environ. Sci. Technol. 41, 119–124.
Kawahara A. and Gan K. (1972) The syntheses, thermal and
structural studies of hopeite. J. Mineral. 7, 289–297.
Kelly S. D., Kemner K. M., Fein J. B., Fowle D. A., Boyanov M.
I., Bunker B. A. and Yee N. (2002) X-ray absorption fine
structure determination of pH-dependent U-bacterial cell wall
interactions. Geochim. Cosmochim. Acta 66, 3855–3871.
Kramer U., Cotter-Howells J. D., Charnock J. M., Baker A. J. M.
and Smith J. A. C. (1996) Free histidine as a metal chelator in
plants that accumulate nickel. Nature 379, 635–638.
Lamelas C., Benedetti M., Wilkinson K. J. and Slaveykova V. I.
(2006) Characterization of H+ and Cd2+ binding properties of
the bacterial exopolysaccharides. Chemosphere 65, 1362–1370.
Leone L., Ferri D., Manfredi C., Persson P., Shchukarev A.,
Sjoberg S. and Loring J. (2007) Modeling the acid–base
properties of bacterial surfaces: a combined spectroscopic and
NEW INSIGHTS INTO REDUCTIVE DETOXIFICATION
OF CHLORINATED SOLVENTS AND RADIONUCLIDES
A Dissertation
Presented to
The Academic Faculty
By
Kelly Elizabeth Fletcher
In Partial Fulfillment
Of the Requirements for the Degree
Doctor of Philosophy in Environmental Engineering in the
School of Civil and Environmental Engineering
Georgia Institute of Technology
December 2010
to U(VI) (23, 25-26).
After shaking, these samples were filtered through 0.2 μm
membrane syringe filters. Exposure of the samples to air resulted in U(IV) oxidation, and
subsequent U(VI) measurements yielded total uranium concentrations. Soluble U(VI)
was quantified by laser excitation spectrofluorescence with a luminescence spectrometer
as previously described (27). Briefly, 0.1 mL aliquots from samples were diluted with
0.9 mL filtered, deionized water and amended with 30 µL of a 40 mM sodium
hypophosphite and 80 mM sodium pyrophosphate solution.
Nominal U(IV)
concentrations were calculated by subtracting the concentration of soluble U(VI) from
the nominal concentration of total uranium.
8.3.3 Characterization
Spectroscopy
of
Uranium Precipitates
using
X-Ray
Absorption
Uranium LIII-edge X-ray absorption near-edge structure (XANES) and extended
X-ray absorption fine structure (EXAFS) analyses were performed to determine the
valence state and the average local environment of uranium in the hydrated solid phase.
Measurements were carried out at the MRCAT/EnviroCAT sector 10-ID (28), Advanced
Photon Source (APS), Argonne National Laboratory, Illinois.
Samples for XAFS
analysis were mounted by filtering the suspensions through 0.22 µm membranes in an
anoxic glove box. The membrane and solids were sealed in Kapton™ film (Dupont,
Circleville, OH). Samples prepared in this manner have shown no oxidation changes
under ambient atmosphere for at least 8 hours (29). The sealed sample holders were
exposed to air only for about 30 seconds while being transferred from an O2-free
transport container to the N2-purged detector housing. Beamline parameters have been
published previously (30-31).
Briefly, the beamline undulator was tapered and the
184
PCE1 are not spore formers, indicating that vegetative cells were responsible for U(VI)
reduction.
In contrast to most U(VI)-reducing organisms, including gram-negative model
organisms such as Anaeromyxobacter, Geobacter, Desulfovibrio, and Shewanella (8, 11,
25, 42, 44), Desulfitobacterium spp. did not produce UO2 but generated mononuclear
U(IV).
Biotic factors (e.g., electron transport machineries, cellular components,
extracellular features) and abiotic factors (e.g., solution composition) can influence the
nature of the reduced product.
For example, a U(IV) phase different from UO2 is
produced by the chemical reduction of U(VI) by Fe(II) (30). Microbial U(VI) reduction
yielding UO2 has been observed in a variety of media with diverse solution compositions
including bicarbonate-buffered groundwater (25) piperazine-N,N‟-bis-(2-ethanesulfonic
acid)-buffered artificial groundwater (26), 30 mM bicarbonate buffer (7-8, 26, 41, 43),
and unbuffered water (16). The medium used in our Desulfitobacterium experiments was
similar in composition to aqueous systems used in previous work that determined UO2 as
the reduced product, suggesting biological factors are involved; however, mononuclear
U(IV) formation may be controlled by a complex interplay between biotic and abiotic
(e.g., medium composition) factors, which future studies should explore in more detail.
Similarly to UO2, the mononuclear U(IV) phase produced in Desulfitobacterium
cultures is readily oxidized upon oxygen exposure (Table 8.1), but further
characterization is needed to describe the stability and mobility of mononuclear U(IV)
(e.g., the potential for complexation with organic ligands and colloidal transport).
Comprehensive knowledge of the different processes and mechanisms involved in U(VI)
reduction are crucial for making meaningful predictions about the mobility and fate of
194
(29)
O'Loughlin, E. J.; Kelly, S. D.; Cook, R. E.; Csencsits, R.; Kemner, K. M.
Reduction of uranium(VI) by mixed iron(II)/iron(III) hydroxide (green rust):
Formation of UO2 nanoparticies. Environ. Sci. Technol. 2003, 37, 721-727.
(30)
Boyanov, M. I.; O'Loughlin, E. J.; Roden, E. E.; Fein, J. B.; Kemner, K. M.
Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The
influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim. Cosmochim. Acta 2007, 71, 1898-1912.
(31)
Kemner, K. M.; Kelly, S. D. Synchrotron-based techniques for monitoring metal
transformations. In Manual of Environmental Microbiology, Third ed., Hurst, C.
J., Ed.; ASM Press: Washington, D.C., 2007; pp 1183-1194.
(32)
Kemner, K. M.; Kropf, J.; Bunker, B. A. A low-temperature total electron yield
detector for X-ray absortion fine-structure spectra. Rev. Sci. Instrum. 1994, 65,
3667-3669.
(33)
Newville, M.; Livins, P.; Yacoby, Y.; Rehr, J. J.; Stern, E. A. Near-edge X-ray
absorption fine structure of Pb - A comparison of theory and experiment. Phys.
Rev. B 1993, 47, 14126-14131.
(34)
Newville, M.; Ravel, B.; Haskel, D.; Rehr, J. J.; Stern, E. A.; Yacoby, Y. Analysis
of multiple-scattering XAFS data using theoretical standards. Physica B 1995,
208-209, 154-156.
(35)
Kelly, S. D.; Kemner, K. M.; Fein, J. B.; Fowle, D. A.; Boyanov, M. I.; Bunker,
B. A.; Yee, N. X-ray absorption fine structure determination of pH-dependent Ubacterial cell wall interactions. Geochim. Cosmochim. Acta 2002, 66, 3855-3871.
(36)
Dusausoy, Y.; Ghermani, N. E.; Podor, R.; Cuney, M. Low-temperature ordered
phase of CaU(PO4)2: Synthesis and crystal structure. Eur. J. Mineral. 1996, 8,
667-673.
(37)
Lanthier, M.; Villemur, R.; Lépine, F.; Bisaillon, J. G.; Beaudet, R. Geographic
distribution of Desulfitobacterium frappieri PCP-1 and Desulfitobacterium spp. in
soils from the province of Quebec, Canada. FEMS Microbiol. Ecol. 2001, 36,
185-191.
(38)
Villemur, R.; Lanthier, M.; Beaudet, R.; Lépine, F. The Desulfitobacterium
genus. FEMS Microbiol. Rev. 2006, 30, 706-733.
(39)
Finneran, K. T.; Forbush, H. M.; VanPraagh, C. V. G.; Lovley, D. R.
Desulfitobacterium metallireducens sp. nov., an anaerobic bacterium that couples
growth to the reduction of metals and humic acids as well as chlorinated
compounds. Int. J. Syst. Evol. Microbiol. 2002, 52, 1929-1935.
199
The Pennsylvania State University
The Graduate School
Department of Civil and Environmental Engineering
KINETIC AND MECHANISTIC STUDY FOR THE ABIOTIC OXIDATION OF Fe(II)
CATALYZED AT THE FERRIC (OXYHYDR)OXIDE AND SOLUTION INTERFACE
A Dissertation in
Environmental Engineering
by
Yuan-Liang Tai
© 2009 Yuan-Liang Tai
Submitted in Partial Fulfillment
of the Requirements
for the Degree of
Doctor of Philosophy
December 2009
31
kobs = koverall · [Fe(II)solid-bound]…..…………………………………………………….…(6)
The dashed line in Figure 3-3 represents the linear regression line determined for the reduction of
NO2- by Fe(II)/HFO. When compared with pseudo-second-order rate constant for U(VI)
reduction, the trendline for nitrite reduction has a much flatter slope than the trendline for U(VI)
reduction. Previous research (25) and this study showed instantaneous removal of over 97% U(VI)
from carbonate-free solution at circumneutral pH. Since practically all of the U(VI) is associated
with solid phase, the reduction of U(VI) must have occurred at the solid-water interface (20). Our
study implies that electron transfer through a sorbed oxidant (U(VI)) is more efficient than
through a dissolved oxidant (NO2-). Previously we demonstrated the conservation of solid-bound
Fe(II) during the first redox reaction stage and evidence for the delocalization and migration of
electrons from surface-bound Fe(II) to bulk Fe(III) (hydr)oxide(22, 29, 30). Electron transfer
from this pool of electrons to the surface-bound oxidant could be a kinetically shorter pathway for
redox reactions involving Fe(II)/HFO. This pathway is also consistent with the anode/cathode
mechanism proposed by Park and Dempsey (26) for Fe(II)/O2 redox reaction with reduction and
oxidation occurring on separate sorption sites of Fe(III) (hydr)oxides.
3.3.4 Effect of Fe(II) surface coverage on reaction rate
Park and Dempsey (26) reported that Fe(II)/O2 heterogeneous reaction rate decreased at
high Fe(II) surface coverage on HFO. Jang et al. (25) demonstrated inhibition of uranium
reduction at high uranium surface coverage on HFO. Surface sorption density could affect surface
coordination between sorbates and sorbents and efficiency of subsequent electron transfer
processes (68). Previously we reported a decrease in the Fe(II)/NO2- reaction rate when Fe(II)
coverage of HFO was above the breakpoint of 0.026 mol Fe(II)/mol Fe(III)(22). Park and
Dempsey also reported a breakpoint of 0.022 mol Fe(II)/mol Fe(III) above which the Fe(II)/O2
55
substantial U(VI) residual remained un-reacted in experiments 2, 3 and 4 even though there was
stoichiometric excess of Fe(II).
∆Fe(II)/ ∆U(VI)phosphate ratio was about 2 in experiment 1. At first glance, this was the
stoichiometry of two-electron transfer from Fe(II) to U(VI). However, for U(VI) to acquire two
electrons from two Fe(II) atoms at the same time requires either a appropriate coordination of
three monomeric species or direct contact between Fe(II) oligomer and U(VI) atom (68). The
possibility of these formations for direct two-electron transfer could be extremely small either in
solution or at solid-water interface. Therefore, several researchers have suggested the formation
of U(V) species as the intermediate product for U(VI) reduction (2, 82, 84-86, 92). First U(VI) is
reduced to U(V) by Fe(II) (2):
(≡FeO)2UO20 + ≡FeO FeOH0 + 3 H2O → 3≡FeOH + Fe(OH)3(s) + UO2 (OH)0aq ..........(2)
U(VI)
U(V)
It has been reported that U(V) species are highly unstable in circumneutral pH and
readily disproportionate into U(VI) and U(IV) (2, 85-89):
2UO2 (OH)0aq + 2 ≡FeOH → (≡FeO)2UO20 + UO2(s) + 2H2O .........................................(3)
U(V)
U(VI)
U(IV)
By using analytical techniques (84, 85) and computational methods (92), previous
researchers suggested that second electron transfer in U(VI) reduction to U(IV) is the rate limiting
step. Along with other researcher (2, 58, 90), their results strongly suggested that one electron
transfer and the subsequent disproportionation is the main pathway for U(VI) reduction to U(IV).
The ∆Fe(II)/ ∆U(VI) stoichiometric ratio for this one electron transfer and the subsequent
disproportionation pathway of uranium reduction would also be 2. Lower ∆Fe(II)/ ∆U(VI) ratios
were observed with higher HFO concentrations. The increasing uranium uptake could be due to
the occlusion and bonding of uranium by the Fe(III) oxyhydroxide structure.
100
57. Wazne, M.; Korfiatis, G. P.; Meng, X. G., Carbonate effects on hexavalent uranium
adsorption by iron oxyhydroxide. Environmental Science & Technology 2003, 37, (16), 36193624.
58. Renshaw, J. C.; Butchins, L. J. C.; Livens, F. R.; May, I.; Charnock, J. M.; Lloyd, J. R.,
Bioreduction of uranium: Environmental implications of a pentavalent intermediate.
Environmental Science & Technology 2005, 39, (15), 5657-5660.
59. Charlet, L.; Silvester, E.; Liger, E., N-compound reduction and actinide immobilisation in
surficial fluids by Fe(II): the surface [6-point triple bond; length half of mdash]FeIIIOFeIIOH[deg] species, as major reductant. Chemical Geology 1998, 151, (1-4),
85-93.
60. Elsner, M.; Schwarzenbach, R. P.; Haderlein, S. B., Reactivity of Fe(II)-bearing minerals
toward reductive transformation of organic contaminants. Environ. Sci. Technol. 2004, 38,
(3), 799-807.
61. Klausen, J.; Trober, S. P.; Haderlein, S. B.; Schwarzenbach, R. P., Reduction of substituted
nitrobenzenes by Fe(II) in aqueous mineral suspensio. Environ. Sci. Technol. 1995, 29, (9),
2396-2404.
62. Jeon, B. H.; Kelly, S. D.; Kemner, K. M.; Barnett, M. O.; Burgos, W. D.; Dempsey, B. A.;
Roden, E. E., Microbial reduction of U(VI) at the solid-water interface. Environmental
Science & Technology 2004, 38, (21), 5649-5655.
63. Senko, J. M.; Kelly, S. D.; Dohnalkova, A. C.; McDonough, J. T.; Kemner, K. M.; Burgos,
W. D., The effect of U(VI) bioreduction kinetics on subsequent reoxidation of biogenic
U(IV). Geochimica et Cosmochimica Acta 2007, 71, (19), 4644-4654.
64. Alowitz, M. J.; Scherer, M. M., Kinetics of nitrate, nitrite, and Cr(VI) reduction by iron
metal. Environmental Science & Technology 2002, 36, (3), 299-306.
65. Sowder, A. G.; Clark, S. B.; Fjeld, R. A., The effect of sample matrix quenching on the
measurement of trace uranium concentrations in aqueous solutions using kinetic
phosphorimetry. Journal of Radioanalytical and Nuclear Chemistry 1998, 234, (1-2), 257260.
66. Elias, D. A.; Senko, J. M.; Krumholz, L. R., A procedure for quantitation of total oxidized
uranium for bioremediation studies. Journal of Microbiological Methods 2003, 53, (3), 343353.
67. Wehrli, B.; Sulzberger, B.; Stumm, W., Redox processes catalyzed by hydrous oxide
surfaces. Chemical Geology 1989, 78, (3-4), 167-179.
68. Boyanov, M. I.; O'Loughlin, E. J.; Roden, E. E.; Fein, J. B.; Kemner, K. M., Adsorption of
Fe(II) and U(VI) to carboxyl-functionalized microspheres: The influence of speciation on
uranyl reduction studied by titration and XAFS. Geochimica et Cosmochimica Acta 2007, 71,
(8), 1898-1912.
69. Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Duff, M. C.; Gorby, Y. A.; Li, S.-m. W.;
Krupka, K. M., Reduction of U(VI) in goethite ([alpha]-FeOOH) suspensions by a
dissimilatory metal-reducing bacterium. Geochimica et Cosmochimica Acta 2000, 64, (18),
3085-3098.
70. Jenne, E. A., Trace element sorption by sediments and soils—Sites and processes. In
Molybdenum in the Environment K.K., C. W. R. a. P., Ed. Marcel Dekker Inc.: New York,
1977; pp 425-553.
71. Bruno, J.; De Pablo, J.; Duro, L.; Figuerola, E., Experimental study and modeling of the
U(VI)-Fe(OH)3 surface precipitation/coprecipitation equilibria. Geochimica et
Cosmochimica Acta 1995, 59, (20), 4113-4123.
72. Duff, M. C.; Coughlin, J. U.; Hunter, D. B., Uranium co-precipitation with iron oxide
minerals. Geochimica et Cosmochimica Acta 2002, 66, (20), 3533-3547.
Chemistry Central Journal
Open Access
Research article
Quantum mechanical calculation of aqueuous uranium complexes:
carbonate, phosphate, organic and biomolecular species
James D Kubicki*1,2, Gary P Halada3, Prashant Jha3 and Brian L Phillips4
Address: 1Department of Geosciences, The Pennsylvania State University, University Park, PA 16802, USA, 2The Earth & Environmental Systems
Institute, The Pennsylvania State University, University Park, PA 16802, USA, 3Department of Materials Science and Engineering, Stony Brook
University, Stony brook, New York 11794-2275, USA and 4Dept. of Geological Sciences, Stony Brook University, Stony brook, New York 117942275, USA
Email: James D Kubicki* - [email protected]; Gary P Halada - [email protected]; Prashant Jha - [email protected];
Brian L Phillips - [email protected]
* Corresponding author
Published: 18 August 2009
Chemistry Central Journal 2009, 3:10
doi:10.1186/1752-153X-3-10
Received: 23 September 2008
Accepted: 18 August 2009
This article is available from: http://journal.chemistrycentral.com/content/3/1/10
© 2009 Kubicki et al
Abstract
Background: Quantum mechanical calculations were performed on a variety of uranium species
representing U(VI), U(V), U(IV), U-carbonates, U-phosphates, U-oxalates, U-catecholates, Uphosphodiesters, U-phosphorylated N-acetyl-glucosamine (NAG), and U-2-Keto-3-doxyoctanoate
(KDO) with explicit solvation by H2O molecules. These models represent major U species in
natural waters and complexes on bacterial surfaces. The model results are compared to observed
EXAFS, IR, Raman and NMR spectra.
Results: Agreement between experiment and theory is acceptable in most cases, and the reasons
for discrepancies are discussed. Calculated Gibbs free energies are used to constrain which
configurations are most likely to be stable under circumneutral pH conditions. Reduction of U(VI)
to U(IV) is examined for the U-carbonate and U-catechol complexes.
Conclusion: Results on the potential energy differences between U(V)- and U(IV)-carbonate
complexes suggest that the cause of slower disproportionation in this system is electrostatic
repulsion between UO2 [CO3]35- ions that must approach one another to form U(VI) and U(IV)
rather than a change in thermodynamic stability. Calculations on U-catechol species are consistent
with the observation that UO22+ can oxidize catechol and form quinone-like species. In addition,
outer-sphere complexation is predicted to be the most stable for U-catechol interactions based on
calculated energies and comparison to 13C NMR spectra. Outer-sphere complexes (i.e., ion pairs
bridged by water molecules) are predicted to be comparable in Gibbs free energy to inner-sphere
complexes for a model carboxylic acid. Complexation of uranyl to phosphorus-containing groups
in extracellular polymeric substances is predicted to favor phosphonate groups, such as that found
in phosphorylated NAG, rather than phosphodiesters, such as those in nucleic acids.
Background
The toxicity and radioactivity of U makes it a potentially
hazardous element in the environment. In areas of high U
concentrations, understanding U chemistry is imperative
in order to predict its fate, transport, and risk. Uranium is
capable of forming a wide variety of aqueous and surface
complexes. Furthermore, redox reactions, mainly between
U(VI) and U(IV), are common in subsurface environments (e.g., [1]).
Page 1 of 29
(page number not for citation purposes)
Chemistry Central Journal 2009, 3:10
Research has focused on the environmental chemistry of
U with the goal of managing and remediating U-contaminated sites in the most effective manner ([2-4] and references therein). Recent studies have probed the molecularlevel structures and processes that influence the overall
behavior of U in the environment (e.g., [5]). Both analytical and theoretical studies have discussed complexation
with numerous ligands [6-9] and the redox reactions
between U(VI) and U(IV) (e.g., [10-15]). Computational
chemistry is an important complement to experimental
studies of U chemistry because this methodology can provide information that is not available via experiment,
especially for transient species and those with short
kinetic lifetimes. In order for the molecular modeling to
be useful however, one must demonstrate that the computational methodology produces accurate results compared to known experimental data.
Before one can simulate structures, thermodynamics and
kinetics with confidence, a computational methodology
must be tested against observation. Environmental chemists are interested in U complexation and redox reactions,
so this study focused on evaluating the ability of quantum
mechanical calculations to reproduce experimental data
on aqueous U complexes and redox chemistry. Specifically, models of aqueous U(VI), U(V), and U(IV) were
generated and compared with experiment and previous
calculations. Uranium complexes with inorganic (carbonate and phosphate), organic (oxalate and catechol), and
biological
(phosphodiester,
phosphorylated
glucosamine, and the 2-Keto-3-deoxyoctanoate) ligands were
modeled and analyzed in light of previous observations.
The model results are compared to interatomic distances
from EXAFS, observed vibrational frequencies, and 13C
and 17O NMR chemical shifts. Calculations on the
observed oxidation of catechol by U(VI) are also presented.
Experimental
Computational
Hybrid density functional theory calculations were performed on all model systems using the program Gaussian
03 [16]. The basis set 6-31G(d,p) [17-20] was used for H,
C, and O and the Stuttgart pseudopotential ECP60MWB
and the corresponding ECP60ANO basis set [21,22] were
used for U. This relativistic pseudopotential uses 60 electrons as the "core" electrons and 32 as the valence electrons. The Becke 3-parameter exchange [23,24] and Lee,
Yang and Parr [25] correlation functionals were used for
energy minimizations, frequency analyses and Gibbs free
energy calculations. The Hartree-Fock method was used
for NMR chemical shielding calculations. Excellent results
were obtained by de Jong et al. [6] using a similar method.
All atoms were allowed to relax during energy minimizations, and no symmetry constraints were applied.
http://journal.chemistrycentral.com/content/3/1/10
The models were created including explicit H2O molecules around the complex to account for H-bonding by
aqueous solutions. In this paper, a H-bond is considered
to exist if the H---O distance is less than or equal to 2.0 Å
and if the O-H---O angle is greater than 120°. These criteria are similar to those used by others (e.g., [26]) and are
useful for identifying significant shifts in the calculated OH stretching frequencies [27]. In general, initial models of
solvation were created by positioning H2O molecules
with either their H or O atom at approximately 1.8 Å from
a O or H atom on the solute model with a O-H---O angle
between 120 and 180°.
Previous work [28,29] has shown that including the H2O
molecules in the primary solvation shell of UO22+ is
important for obtaining accurate structures, vibrational
frequencies and energetics. This study (as in [29]) investigates the effects of adding a second solvation shell to the
hydrated UO22+ cation. The number of H2O molecules
was chosen to be at least the minimal number necessary
to form one H-bond to each of the possible H-bonding
atoms in the U coordination sphere (e.g., 2 H2O molecules for each U-OH2 group). In some cases (e.g., UO2oxalate), an increasing number of H2O molecules were
included in the model to assess the effects of explicit solvation on the predicted interatomic distances, vibrational
frequencies, and NMR chemical shifts. Energy minimizations were generally carried out with the default criteria in
Gaussian 03. When imaginary frequencies were calculated
from an energy minimized structure, a re-optimization of
the structure was performed with the "Opt = Tight" option
until a structure with no imaginary frequencies was found.
Although we have obtained potential energy minima,
there is no guarantee that each configuration is the global
minimum because the potential energy surface of these
models will be complicated due to many possible Hbonding configurations. Any energy minimization
scheme is unlikely to find the global potential energy minimum, so molecular dynamics simulations would be useful in the future to investigate configuration space at the
temperature of interest and determine average configurations for these models.
Calculated results were compared to observed EXAFS, IR,
Raman and NMR spectra. Model interatomic distances
were directly compared to the values extracted via analysis
of EXAFS data and vibrational spectra. Theoretical vibrational frequencies were compared to observed values for
uranyl without a scaling factor applied because the appropriate value is not known for this computational methodology. For the vibrations of the ligands, a scaling factor of
0.96 was applied as determined by Wong [30] for B3LYP/
6-31G(d) with the assumption that the p-functions added
to the H atoms do not significantly affect the vibrational
frequencies of the C-C and C-O bonds. This assumption is
Page 2 of 29
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Chemistry Central Journal 2009, 3:10
Model results are shown to be consistent with spectroscopic results on uranyl-organic complexes as well provided the first solvation shell around these complexes is
included in the model. In particular, the NMR spectra collected in this study are consistent with an outer-sphere
uranyl-catechol complex. The oxidation of catechol by
UO22+(aq) was shown to occur through a H-radical mechanism as two phenolic H atoms are transferred in sequence
to the axial O atoms of the UO22+. This results in a U(IV)aq
and quinone. The intermediate quinone radical species
can explain the observation of catechol oxidation and
polymerization in the presence U(VI) in aqueous solutions [101].
For uranyl-cell surface complexation, uranyl is predicted
to favor binding at phosphonate groups rather than phosphodiester groups. Although the inner-sphere bidentate
configuration is predicted to have the lowest Gibbs free
energy in these models, the differences between these configurations and outer-sphere associations is relatively
small suggesting that a significant portion of the observed
complexation could involve outer-sphere binding.
Competing interests
The authors declare that they have no competing interests.
Authors' contributions
http://journal.chemistrycentral.com/content/3/1/10
4.
5.
6.
7.
8.
9.
10.
11.
12.
JDK carried out the quantum mechanical calculations and
wrote these portions of the paper. PJ collected the Raman
spectra. GPH wrote the Raman methods and results sections. BLP collected NMR spectra and wrote these sections
of the paper.
13.
Acknowledgements
15.
This research was funded by the NSF grants "Stony Brook-BNL collaboration to establish a Center for Environmental Molecular Sciences (CEMS)"
and Grant No. CHE-0431328 "Center for Environmental Kinetics Analysis"
(CEKA) at The Pennsylvania State University. Computations were supported by the Materials Simulation Center, a Penn State MRSEC and MRI
facility, and by CEKA, an NSF/DOE Environmental Molecular Science Institute. JDK also thanks AJ Francis and A Clark for commenting on the manuscript before submission, and A Clark for running a modified NPA on the
uranyl-catechol complexes. The detailed and constructive criticisms of
three anonymous reviewers are also appreciated for their suggested
improvements.
14.
16.
17.
18.
19.
20.
References
1.
2.
3.
Abdelouas A, Lutze W, Gong WL, Nuttal EH, Strietelmeier BA, Travis
BJ: Biological reduction of uranium in groundwater and subsurface soil. Sci Total Env 2000, 250:21-35.
Porcelli D, Swarzenski PW: The behavior of U- and Th-series
nuclides in groundwater. In Reviews in Mineralogy & Geochemistry,
Uranium-Series Geochemistry Volume 52. Edited by: Bourdon B, Henderson GM, Lundstrom CC, Turner SP. Washington DC: Mineralogical
Society of America; 2003:317-362.
Swarzenski PW, Porcelli D, Andersson PS, Smoak JM: The behavior
of U- and Th-series nuclides in the estuarine environment. In
Reviews in Mineralogy & Geochemistry, Uranium-Series Geochemistry Volume 52. Edited by: Bourdon B, Henderson GM, Lundstrom CC,
21.
22.
23.
24.
Turner SP. Washington DC: Mineralogical Society of America;
2003:577-606.
Suzuki Y, Suko T: Geomicrobiological factors that control uranium mobility in the environment: update on recent
advances in the bioremediation of uranium-contaminated
sites. J Mineralogical Petrological Sciences 2006, 101:299-307.
Catalano JG, McKinley JP, Zachara JM, Heald SM, Smith SC, Brown
GE: Changes in uranium speciation through a depth
sequence of contaminated Hanford sediments. Env Sci Tech
2006, 40:2517-2524.
de Jong WA, Apra E, Windus TL, Nichols JA, Harrison RJ, Gutowski
KE, Dixon DA: Complexation of the carbonate, nitrate, and
acetate anions with the uranyl dication: density functional
studies with relativistic effective core potentials. J Phys Chem
A 2005, 109:11568-11577.
Dong WM, Brooks SC: Determination of the formation constants of ternary complexes of uranyl and carbonate with
alkaline earth metals (Mg2+, Ca2+, Sr2+, and Ba2+) using
anion exchange method. Env Sci Tech 2006, 40:4689-4695.
Kelly SD, Kemner KM, Brooks SC: X-ray absorption spectroscopy identifies calcium-uranyl-carbonate complexes at environmental concentrations. Geochim et Cosmochim Acta 2007,
71:821-834.
Boyanov MI, O'Loughlin EJ, Roden EE, Fein JB, Kemner KM: Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The influence of speciation on uranyl reduction
studied by titration and XAFS. Geochim Cosmochim Acta 2007,
71:1898-1912.
Vallet V, Schimmelpfennig B, Maron L, Teichteil C, Leininger T, Gropen O, Grenthe I, Wahlgren U: Reduction of uranyl by hydrogen:
an ab initio study. Chem Phys 1999, 244:185-193.
Moskaleva LV, Krüger S, Spörl A, Rösch N: Role of solvation in the
reduction of the uranyl dication by water: A density functional study. Inorg Chem 2004, 43:4080-4090.
Hua B, Xu HF, Terry J, Deng BL: Kinetics of uranium(VI) reduction by hydrogen sulfide in anoxic aqueous systems. Env Sci
Tech 2006, 40:4666-4671.
Marshall MJ, Beliaev AS, Dohnalkova AC, Kennedy DW, Shi L, Wang
Z, Boyanov MI, Lai B, Kemner KM, McLean JS, Reed SB, Culley DE,
Bailey VL, Simonson CJ, Saffarini DA, Romine MF, Zachara JM, Frederickson JK: c-Type cytochrome-dependent formation of
U(IV) nanoparticles by Shewanella oneidensis. Plos Biology 2006,
4:1324-1333.
Stewart BD, Neiss J, Fendorf S: Quantifying constraints imposed
by calcium and iron on bacterial reduction of uranium(VI). J
Env Quality 2007, 36:363-372.
Wander MCF, Kerisit S, Rosso KM, Schoonen MAA: Kinetics of
triscarbonato uranyl reduction by aqueous ferrous iron: a
theoretical study. J Phys Chem A 2006, 110:9691-9701.
Frisch MJT, et al.: Gaussian 03, Revision C.02. Wallingford 2004.
Ditchfield R, Hehre WJ, Pople JA: Self-consistent molecular
orbital methods. 9. Extended Gaussian-type basis for molecular orbital studies of organic molecules. J Chem Phys 1971,
54:724-728.
Hehre W, Ditchfield R, Pople J: Self-consistent molecular-orbital
methods 12. Further extensions of gaussian-type basis sets
for use in molecular-orbital studies of organic-molecules. J
Chem Phys 1972, 56:2257-2261.
Rassolov VA, Pople JA, Ratner MA, Windus TL: 6-31G* basis set for
atoms K through Zn. J Chem Phys 1998, 109:1223-1229.
Rassolov VA, Ratner MA, Pople JA, Redfern PC, Curtiss LA: 631G*basis set for third-row atoms. J Comp Chem 2001,
22:976-984.
Küchle W, Dolg M, Stoll H, Preuss H: Energy-adjusted pseudopotentials for the actinides – parameter sets and test calculations for thorium and thorium monoxide. J Chem Phys 1994,
100:7535-7542.
Cao X, Dolg M, Stoll H: Valence basis sets for relativistic
energy-consistent small-core actinide pseudopotentials. J
Chem Phys 2003, 118:487-496.
Becke AD: Density-functional thermochemistry 5. Systematic
optimization of exchange-correlation functionals. J Chem Phys
1997, 107:8554-8560.
Stephens PJ, Devlin FJ, Chabalowski CF, Frisch MJ: Ab-initio calculation of bibrational absorption and circular-dichroism spec-
Page 26 of 29
(page number not for citation purposes)
REDOX BEHAVIOR OF MAGNETITE IN THE ENVIRONMENT: MOVING
TOWARDS A SEMICONDUCTOR MODEL
by
Christopher Aaron Gorski
An Abstract
Of a thesis submitted in partial fulfillment
of the requirements for the Doctor of
Philosophy degree in Civil and Environmental Engineering
in the Graduate College of
The University of Iowa
December 2009
Thesis Supervisor: Associate Professor Michelle M. Scherer
115
and the underlying solid phase (e.g., Fe3+ oxides, clay minerals) (23, 24, 29, 30, 74). This
work focuses only on redox inactive substrates (i.e., a stable, sorbed Fe2+ species) and our
ability to spectroscopically characterize these phases.
Sorbed Fe2+ is a critical groundwater substituent to several processes including
redox buffering, contaminant fate, microbial respiration, and microbial metabolism (166169). Previous studies have shown that sorbed Fe2+ on clay minerals as well as Al and Ti
oxides is capable of reducing and degrading environmental contaminants which are not
reactive with dissolved Fe2+ alone, including nitroaromatics, Se6+, and Tc7+ (165, 169172). Two recent works have confirmed that Fe2+ sorbed on Al and Ti oxides is a
stronger reductant (lower Eh) than dissolved Fe2+ (172, 173). For dissimilatory iron
reducing bacteria (DIRB), Fe2+ sorbed on cells was shown to diminish the ability of the
cells to respire on Fe3+ oxides and nitrate (166, 167), and other works have argued that
sorbed Fe2+ on cells significantly influences the mineralogy of Fe in the environment
(e.g., 174).
Despite the importance of Fe2+ sorption to biogeochemistry, sorbed Fe2+ is largely
uncharacterized using spectroscopic methods, which is likely due to its amorphous
structure and low abundance within the sample. For example, sorbed Fe2+ is completely
invisible to powder X-ray diffraction. In most studies, Fe2+ sorption is examined only
indirectly, by measuring the amount of dissolved Fe2+ removed from solution after
addition of the solid phase (21, 164, 172, 175). Spectroscopic techniques have been used
in some systems, including XAFS and Mössbauer spectroscopy, but these techniques are
so rarely used that it is difficult to discern if the conclusions drawn in each study can be
applied to broader scopes (23, 176, 177). This lack of spectroscopic data presents a gap in
our understanding of Fe2+ sorption and its environmental implications.
Instead of spectroscopy, surface complexation modeling (SCM) is used more
commonly to describe Fe2+ sorption behavior on environmental surfaces (21, 164, 172,
175). SCM is a valuable technique, and can aid in understanding what is occurring at the
211
176.
Boyanov, M. I.; O'Loughlin, E. J.; Roden, E. E.; Fein, J. B.; Kemner, K. M.,
Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The
influence of speciation on uranyl reduction studied by titration and XAFS.
Geochimica et Cosmochimica Acta 2007, 71, 1898-1912.
177.
Dong, H. L.; Kukkadapu, R. K.; Zachara, J. M.; Kennedy, D. W.; Kostandarithes,
H. M., Microbial reduction of structural Fe(III) in illite and goethite.
Environmental Science and Technology 2003, 37, 1268-1276.
178.
Appelo, C. A. J.; Van der Weiden, M. J. J.; Tournassat, C.; Charlet, L., Surface
complexation of ferrous iron and carbonate on ferrihydrite and the mobilization of
arsenic. Environmental Science and Technology 2002, 36, 3096-3103.
179.
Limousin, G.; Gaudet, J.-P.; Charlet, L.; Szenknect, S.; Barthes, V.; Krimissa, M.,
Sorption isotherms. A review on physical bases, modeling and measurement.
Applied Geochemistry 2007, 22, 249-275.
180.
Griffen, D. T.; Nelson, W. R., Mossbauer spectroscopy of Zn-poor and Zn-rich
rhodonite. American Mineralogist 2007, 92, 1486-1491.
181.
Burns, R. G., Mineral Mössbauer-Spectroscopy - Correlations between chemicalshift and quadrupole splitting parameters. Hyperfine Interactions 1994, 91, 739745.
182.
Chatellier, X.; Fortin, D., Adsorption of ferrous ions onto Bacillus subtilis cells.
Chemical Geology 2004, 212, 209-228.
183.
Martinez, R. E.; Smith, D. S.; Pedersen, K.; Ferris, F. G., Surface chemical
heterogeneity of bacteriogenic iron oxides from a subterranean environment.
Environmental Science and Technology 2003, 37, 5671-5677.
184.
Beveridge, T. J.; Murray, R. G. E., Sites of Metal-Deposition in the Cell-Wall of
Bacillus-Subtilis. Journal of Bacteriology 1980, 141, 876-887.
185.
Amthauer, G.; Groczicki, M.; Lottermoser, W.; Redhammer, G., Mössbauer
spectroscopy: Basic Principles. In Spectroscopic Methods in Mineralogy, Beran,
A.; Libowitzky, E., Eds. 2004; Vol. 6, pp 345-367.
186.
Pettibone, J. M.; Cwiertny, D. M.; Scherer, M.; Grassian, V. H., Adsorption of
organic acids on TiO2 nanoparticles: Effects of pH, nanoparticle size, and
nanoparticle aggregation. Langmuir 2008, 24, 6659-6667.
187.
Campbell, S. J.; Aubertin, F., Evaluation of Distributed Hyperfine Parameters. In
Mössbauer spectroscopy applied to inorganic chemistry, Long, G. J.; Grandjean,
F., Eds. 1984; Vol. 3, pp 183-242.
188.
Vrajmasu, V.; Munck, E.; Bominaar, E. L., Density functional study of the
electric hyperfine interactions and the redox-structural correlations in the cofactor
of nitrogenase. Analysis of general trends in Fe-57 isomer shifts. Inorganic
Chemistry 2003, 42, 5974-5988.
189.
Shenoy, G. K., Mössbauer-Effect Isomer Shifts. In Mössbauer Spectroscopy
Applied to Inorganic Chemistry, Long, G. J.; Grandjean, F., Eds. Plenum Press:
New York, 1989; Vol. 1, pp 57-74.
Geomicrobiology Journal, 26:431–441, 2009
Copyright © Taylor & Francis Group, LLC
ISSN: 0149-0451 print / 1521-0529 online
DOI: 10.1080/01490450903060780
Nonreductive Biomineralization of Uranium(VI) Phosphate
Via Microbial Phosphatase Activity in Anaerobic Conditions
Melanie J. Beazley,1 Robert J. Martinez,2 Patricia A. Sobecky,2 Samuel M. Webb,3
and Martial Taillefert1
1
Downloaded by [University of Notre Dame] at 10:04 16 July 2011
School of Earth & Atmospheric Sciences, Georgia Institute of Technology, Atlanta,
Georgia 30332-0340, USA
2
School of Biology, Georgia Institute of Technology, Atlanta, Georgia 30332-0230, USA
3
Stanford Synchrotron Radiation Laboratory, Menlo Park, California 94025, USA
The remediation of uranium from soils and groundwater at
Department of Energy (DOE) sites across the United States represents a major environmental issue, and bioremediation has
exhibited great potential as a strategy to immobilize U in the subsurface. The bioreduction of U(VI) to insoluble U(IV) uraninite has
been proposed to be an effective bioremediation process in anaerobic conditions. However, high concentrations of nitrate and low
pH found in some contaminated areas have been shown to limit
the efficiency of microbial reduction of uranium. In the present
study, nonreductive uranium biomineralization promoted by microbial phosphatase activity was investigated in anaerobic conditions in the presence of high nitrate and low pH as an alternative
approach to the bioreduction of U(VI). A facultative anaerobe,
Rahnella sp. Y9602, isolated from soils at DOE’s Oak Ridge Field
Research Center (ORFRC), was able to respire anaerobically on
nitrate as a terminal electron acceptor in the presence of glycerol3-phosphate (G3P) as the sole carbon and phosphorus source and
hydrolyzed sufficient phosphate to precipitate 95% total uranium
after 120 hours in synthetic groundwater at pH 5.5. Synchrotron
X-ray diffraction and X-ray absorption spectroscopy identified the
Received 29 December 2009; accepted 20 April 2009.
This research was supported by the Office of Science (BER), U. S.
Department of Energy Grant No. DE-FG02-04ER63906. Portions of
this research were carried out at the Stanford Synchrotron Radiation
Lightsource, a national user facility operated by Stanford University on
behalf of the U.S. Department of Energy, Office of Basic Energy Sciences. The SSRL Structural Molecular Biology Program is supported
by the Department of Energy, Office of Biological and Environmental
Research, and by the National Institutes of Health, National Center
for Research Resources, Biomedical Technology Program. We thank
Hong Yi at the Emory School of Medicine Electron Microscopy Core
for transmission electron microscopy sample preparation and imaging
and Dr. Joan S. Hudson at the Clemson University Electron Microscope
Facility for variable-pressure scanning electron microscopy imaging,
transmission electron microscopy imaging, and EDX analyses. We also
thank the two anonymous reviewers for their valuable comments to improve the manuscript.
Address correspondence to Martial Taillefert, School of Earth &
Atmospheric Sciences, 311 Ferst Drive, Atlanta, GA 30332-0340.
E-mail: [email protected]
mineral formed as chernikovite, a U(VI) autunite-type mineral.
The results of this study suggest that in contaminated subsurfaces,
such as at the ORFRC, where high concentrations of nitrate and
low pH may limit uranium bioreduction, the biomineralization of
U(VI) phosphate minerals may be a more attractive approach for
in situ remediation providing that a source of organophosphate is
supplied for bioremediation.
Keywords
biomineralization, bioremediation,
phatase, uranium (VI)
microbial
phos-
INTRODUCTION
Over 30 years of uranium enrichment at the DOE Oak Ridge,
Tennessee site left a legacy of uranium contamination in soils
and groundwater (Brooks 2001). These contaminated systems
are characterized by high concentrations of uranium and other
toxic metals, as well as high nitrate and low pH (Brooks 2001;
Wu et al. 2006a). Complete removal of uranium from groundwater and soils through methods such as extraction and pump and
treat is infeasible on large spatial scales, and research in recent
years has focused on ways to lower the solubility of uranium in
situ, thereby stopping and/or slowing its migration through the
subsurface (e.g., Arey et al. 1999; Istok et al. 2004; Wu et al.
2006a, 2007).
The solubility of uranium in the subsurface is influenced by
the redox conditions, pH, soil matrix, and the presence/absence
of organic ligands, phosphate, and carbonate (Langmuir 1997).
Uranium exists in the environment in two primary oxidation
states, the soluble uranyl ion (U(VI)) and the insoluble uraninite
mineral (U(IV)). Uranium(VI) reduction to U(IV) occurs either
chemically by Fe(II) adsorbed onto mineral surfaces (Boyanov
et al. 2007; Jeon et al. 2005; Liger et al. 1999; O’Loughlin et al.
2003) or biologically by dissimilatory metal-reducing bacteria
(DMRB) and sulfate-reducing bacteria (SRB) (Fredrickson et al.
2000; Lovley and Phillips 1992; Lovley et al. 1991; North et al.
2004; Wade Jr. and DiChristina 2000).
Unfortunately, uraninite is rapidly oxidized to the more
mobile and reactive uranyl ion (UVI O2+
2 ) in oxic conditions
431
440
M. J. BEAZLEY ET AL.
Downloaded by [University of Notre Dame] at 10:04 16 July 2011
Future experiments should examine the competition between
bioreduction and non-reductive biomineralization of U in natural systems and the effect of the presence of carbonates on the
stability of the minerals formed.
CONCLUSIONS
The results of this investigation illustrate the potential for
controlling the solubility of uranium through phosphatase activity by subsurface soil microorganisms in contaminated waste
sites in both aerobic and anaerobic conditions. The facultative gram-negative anaerobe, Rahnella sp. Y9602, isolated from
the ORFRC subsurface demonstrates strong phosphatase activity in both aerobic and anaerobic conditions and in the presence of high nitrate and low pH. The phosphate hydrolyzed
from an organophosphate substrate is sufficient to precipitate
95% total U(VI) as the uranyl phosphate mineral chernikovite
very rapidly. The precipitation of chernikovite appears to be a
pure chemical process, whose kinetics is governed by pH, the
concentration of phosphate generated by the microbial hydrolysis of organophosphate by phosphatase enzymes, and probably
the adsorption of U(VI) to the cell surfaces. Uranyl phosphates
are stable in a wide range of pH for long periods of time and may
be preferable to the more reactive and easily oxidized uraninite
produced during U(VI) bioreduction.
REFERENCES
Ankudinov AL, Ravel B, Rehr JJ, Conradson SD. 1998. Real space multiple
scattering calculation of XANES. Phys Rev B 58:7565.
Arey JS, Seaman JC, Bertsch PM. 1999. Immobilization of uranium in contaminated sediments by hydroxyapatite addition. Environ Sci Technol 33:337–
342.
Beazley MJ, Martinez RJ, Sobecky PA, Webb SM, Taillefert M. 2007. Uranium biomineralization as a result of bacterial phosphatase activity: Insights
from bacterial isolates from a contaminated subsurface. Environ Sci Technol
41:5701–5707.
Beller HR. 2005. Anaerobic, nitrate-dependent oxidation of U(IV) oxide minerals by the chemolithoautotrophic bacterium Thiobacillus denitrificans. Appl
Environ Microbiol 71:2170–2174.
Betlach MR, Tiedje JM. 1981. Kinetic explanation for accumulation of nitrite,
nitric oxide, and nitrous oxide during bacterial denitrification. Appl Environ
Microbiol 42:1074–1084.
Bollag J-M, Henninger NM. 1978. Effects of nitrite toxicity on soil bacteria
under aerobic and anaerobic conditions. Soil Biol Biochem 10:377–381.
Bothe H, Jost G, Schloter M, Ward BB, Witzel K-P. 2000. Molecular analysis
of ammonia oxidation and denitrification in natural environments. FEMS
Microbiol Rev 24:673–690.
Boyanov MI, O’Loughlin EJ, Roden EE, Fein JB, Kemner KM. 2007. Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The
influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim Cosmochim Acta 71:1898–1912.
Brooks SC. 2001. Waste Characteristics of the Former S-3 Ponds and Outline
of Uranium Chemistry Relevant to NABIR Field Research Center Studies.
Technical Report; NABIR FRC.
Catalano JG, Brown Jr. GE. 2004. Analysis of uranyl-bearing phases by EXAFS spectroscopy: Interferences, multiple scattering, accuracy of structural
parameters, and spectral differences. Am Mineral 89:1004–1021.
Cullity BD. 1978. Elements of X-Ray Diffraction. Menlo Park, CA: AddisonWesley Publishing Company, Inc. 555 p.
De Pablo J, Casas I, Giménez J, Molera M, Rovira M, Duro L, Bruno J. 1999. The
oxidative dissolution mechanism of uranium dioxide. I. The effect of temperature in hydrogen carbonate medium. Geochim Cosmochim Acta 63:3097–
3103.
De Yoreo JJ, Vekilov PG. 2003. Principles of crystal nucleation and growth. In:
Dove PM, De Yoreo JJ, Weiner S, editors. Biomineralization. Washington,
DC: Mineralogical Society of America. P 57–90.
Drysdale GD, Kasan HC, Bux F. 1999. Denitrification by heterotrophic bacteria
during activated sludge treatment. Water SA 25:357–362.
Edwards L, Küsel K, Drake H, Kostka JE. 2007. Electron flow in acidic subsurface sediments co-contaminated with nitrate and uranium. Geochim Cosmochim Acta 71:643–654.
Fields MW, Yan T, Rhee S-K, Carroll SL, Jardine PM, Watson DB, Criddle
CS, Zhou J. 2005. Impacts on microbial communities and cultivable isolates
from groundwater contaminated with high levels of nitric acid-uranium waste.
FEMS Microbiol Ecol 53:417–428.
Finneran KT, Housewright ME, Lovley DR. 2002. Multiple influenences of
nitrate on uranium solubility during bioremediation of uranium-contaminated
subsurface sediments. Environ Microbiol 4:510–516.
Fredrickson JK, Zachara JM, Kennedy DW, Duff MC, Gorby YA, Li S-MW,
Krupka KM. 2000. Reduction of U(VI) in goethite (α-FeOOH) suspensions
by a dissimilatory metal-reducing bacterium. Geochim Cosmochim Acta
64:3085–3098.
Fredrickson JK, Zachara JM, Kennedy DW, Liu C, Duff MC, Hunter DB,
Dohnalkova A. 2002. Influence of Mn oxides on the reduction of uranium
(VI) by the metal-reducing bacterium Shewanella putrefaciens. Geochim
Cosmochim Acta 66:3247–3262.
Fuller CC, Bargar JR, Davis JA, Piana MJ. 2002. Mechanisms of uranium
interactions with hydroxyapatite: Implications for groundwater remediation.
Environ Sci Technol 36:158–165.
Grasshoff K. 1983. Determination of nitrite. In: Grasshoff K, Ehrhardt M, Kremling K, editors. Methods of Seawater Analysis. Weinheim: Verlag Chemie
GmbH. P 139–142.
Guillaumont R, Fanghänel T, Fuger J, Grenthe I, Neck V, Palmer DA, Rand
MH. 2003. Chemical Thermodynamics 5. Update on the Chemical Thermodynamics of Uranium, Neptunium, Plutonium, Americium and Technetium.
Amsterdam: Elsevier. 919 p.
Hsi C-KD, Langmuir D. 1985. Adsorption of uranyl onto ferric oxyhydroxides: Application of the surface complexation site-binding model. Geochim
Cosmochim Acta 49:1931–1941.
Hudson EA, Allen PG, Terminello LJ, Denecke MA, Reich T. 1996. Polarized
x-ray-absorption spectroscopy of the uranyl ion: Comparison of experiment
and theory. Phys Rev B 54:156–165.
Istok JD, Senko JM, Krumholz LR, Watson D, Bogle MA, Peacock A, Chang
Y-J, White DC. 2004. In situ bioreduction of technetium and uranium in a
nitrate-contaminated aquifer. Environ Sci Technol 38:468–475.
Jeon B-H, Dempsey BA, Burgos WD, Barnett MO, Roden EE. 2005. Chemical
reduction of U(VI) by Fe(II) at the solid-water interface using natural and
synthetic Fe(III) oxides. Environ Sci Technol 39:5642–5649.
Jerden Jr. JL, Sinha AK. 2003. Phosphate based immobilization of uranium in
an oxidizing bedrock aquifer. Appl Geochem 18:823–843.
Kelly SD, Kemner KM, Fein JB, Fowle DA, Boyanov MI, Bunker BA, Yee
N. 2002. X-ray absorption fine structure determination of pH-dependent Ubacterial cell wall interactions. Geochim Cosmochim Acta 66:3855–3871.
Knowles R. 1982. Denitrification. Microbiol Rev 46:43–70.
Krieg NR, Holt JG, 1984. Bergey’s Manual of Systematic Bacteriology. In:
Krieg NR, Holt JG Eds.). Baltimore, MD: Williams & Wilkins.
Langmuir D. 1978. Uranium solution-mineral equilibria at low temperatures
with applications to sedimentary ore deposits. Geochim Cosmochim Acta
42:547–569.
Langmuir D. 1997. Aqueous Environmental Geochemistry. Upper Saddle River,
NJ: Prentice Hall. 600 p.
Liger E, Charlet L, Van Cappellen P. 1999. Surface catalysis of uranium(VI)
reduction by iron(II). Geochim Cosmochim Acta 63:2939–2955.
Interfacial and Long-Range Electron Transfer at the MineralMicrobe Interface
Nicholas Scott Wigginton
Dissertation submitted to the faculty of the Virginia Polytechnic
Institute and State University in partial fulfillment of the
requirements for the degree of
Doctor of Philosophy
In
Geosciences
Committee
Michael F. Hochella Jr., Chair
James R. Heflin
Kevin M. Rosso
Brian H. Lower
April 21, 2008
Blacksburg, VA
Keywords: biogeochemistry, geomicrobiology, Shewanella,
scanning tunneling microscopy, hematite
Copyright 2008, Nicholas S. Wigginton
of uraninite (UO2) have been shown to form abiotically when U(VI) is reduced by Fe(II)oxides (Boyanov et al., 2007; O'Loughlin et al., 2003).
Disregarding reoxidation of the solid-phase products (Wan et al., 2005) for the
moment, one issue of great concern for the stability of the product is how unreduced
aqueous U(VI) interacts with the precipitating nanoparticles. One field study showed that
a solid-phase U(IV) precipitate actually contained a large fraction of unreduced U(VI)
(Ortiz-Bernad et al., 2004).
The complexities behind the fate of metal and radionuclide contaminants during
nanoparticle formation make predicting the final products very difficult.
Several
examples in the literature of different end-products for various metal/metalloid
contaminants highlight the fact that we do not yet fully understand the fate of the
contaminant with respect to the precipitating nanoparticle, but a lot of progress is being
made. For example, a recent study by Zachara and colleagues showed that the abiotic
reduction of Tc(VII) by Fe(II) formed relatively stable iron oxide nanoparticles with
homogeneously distributed Tc(IV) in the crystal structure, effectively stabilizing the Tc
(Zachara et al., 2007). The opposite can be true with the metalloid As. Tadanier and
colleagues showed that the microbial-reduction of Fe-oxide aggregates with adsorbed
As(V)
caused
the
deflocculation
of
As-bearing
ferrihydrite
(Fe10O14(OH)2)
nanoparticles, subsequently increasing the mobility of As (Tadanier et al., 2005).
When relying on the reductive transformation of metals and subsequent growth of
nanoparticles as a means for remediation, the problem of nanoparticle stability must be a
chief concern. However, as we will see later, nanoparticles can be incredibly mobile in
the environment; understanding the transport mechanisms of nanoparticles in subsurface
contamination plumes is also very important for analyzing the practicality of such
remediation efforts. One recent study, however, circumvented the need for understanding
transport because they simply removed metal- and nanoparticle-bearing water from a
contaminated mining site and used it to create additional nanoparticles ex-situ (Wei and
Viadero, 2007). Magnetite nanoparticles were grown using ferric iron in the acid-mine
drainage waters that also contained trace amounts of other metals including Zn, Ni, and
Cu. This presents an intruiging remediation case where contamination species are
removed from the site and transformed ex-situ.
128
Bose, S., Hochella, M. F., Lower, B. H., Gorby, Y. A., Kennedy, D. W., McCready, D.
E., and Madden, A. S., submitted. Bioreduction of hematite nanoparticles by
Shewanella oneidensis MR-1. Am. J. Sci.
Boyanov, M. I., O'Loughlin, E. J., Roden, E. E., Fein, J. B., and Kemner, K. M., 2007.
Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: The
influence of speciation on uranyl reduction studied by titration and XAFS.
Geochim. Cosmochim. Acta 71, 1898-1912.
Brantley, S. L., White, T. S., White, A. F., Sparks, D. L., Richter, D. D., Pregitzer, K.,
Derry, L. A., Chorover, J., Chadwick, O. A., April, R. H., Anderson, S. P., and
Amundson, R., 2006. Fronteirs in exploration of the critical zone. Report of a
workshop sponsored by the National Science Foundation, Newark, DE.
Buffle, J. and Leppard, G. G., 1995. Characterization of aquatic colloids and
macromolecules. 1. Structure and behavior of colloidal material. Environ. Sci.
Technol. 29, 2169-2175.
Burleson, D. J., Driessen, M. D., and Penn, R. L., 2004. On the characterization of
environmental nanoparticles. J. Environ. Sci. Health., A 39, 2707-2753.
Chan, C. S., De Stasio, G., Welch, S. A., Girasole, M., Frazer, B. H., Nesterova, M. V.,
Fakra, S., and Banfield, J. F., 2004. Microbial polysaccharides template assembly
of nanocrystal fibers. Science 303, 1656-1658.
Chanudet, V. and Filella, M., 2006. A non-perturbing scheme for the mineralogical
characterization and quantification of inorganic colloids in natural waters.
Environ. Sci. Technol. 40, 5045-5051.
Chen, K. L., Mylon, S. E., and Elimelech, M., 2006. Aggregation kinetics of alginatecoated hematite nanoparticles in monovalent and divalent electrolytes. Environ.
Sci. Technol. 40, 1516-1523.
Chen, K. L., Mylon, S. E., and Elimelech, M., 2007. Enhanced aggregation of alginatecoated iron xxide (hematite) nanoparticles in the presence of calcium, strontium,
and barium cations. Langmuir 23, 5920 -5928.
Chernyshova, I. V., Hochella, M. F., and Madden, A. S., 2007. Size-dependent structural
transformations of hematite nanoparticles. 1. Phase transition. Phys. Chem. Chem.
Phys. 9, 1736-1750.
Ciglenecki, I., Krznaric, D., and Helz, G. R., 2005. Voltammetry of copper sulfide
particles and nanoparticles: Investigation of the cluster hypothesis. Environ. Sci.
Technol. 39, 7492-7498.
Contado, C., Blo, G., Conato, C., Dondi, F., and Beckett, R., 2003. Experimental
approaches for size-based metal speciation in rivers. J. Environ. Monit. 5, 845851.
Council, N. R., 2001. Basic Research Opportunities in the Earth Sciences. National
Academies Press, Washington, D.C.
Daulton, T. L., Little, B. J., Lowe, K., and Jones-Meehan, J., 2001. In Situ Environmental
Cell-Transmission Electron Microscopy Study of Microbial Reduction of
Chromium(VI) Using Electron Energy Loss Spectroscopy. Microsc. Microanal. 7.
Daulton, T. L., Little, B. J., Lowe, K., and Jones-Meehan, J., 2002. Electron energy loss
spectroscopy techniques for the study of microbial chromium(VI) reduction. J.
Microbiol. Methods 50, 39-54.
146
Rev Environ Sci Biotechnol (2008) 7:355–380
DOI 10.1007/s11157-008-9137-8
REVIEW PAPER
Interactions of aqueous U(VI) with soil minerals in slightly
alkaline natural systems
Nikolla P. Qafoku Æ Jonathan P. Icenhower
Published online: 22 August 2008
Battelle Memorial Institute 2008
Abstract Uranium (U) is a common contaminant at
numerous surface and subsurface sites in proximity to
areas involved with weapons manufacturing and
atomic energy related activities. This paper covers
some important aspects of the aqueous hexavalent
uranium [U(VI)] interactions with soil minerals that
are present in contaminated soils and sediments. The
retention of U via interactions with soil minerals has
significant consequences for the prediction of its
short- and long-term behavior in soils and geological
systems. Studies of the nature and type of these
interactions have provided the necessary evidence for
assessing the geochemical behavior of U in natural
systems under different physical, biogeochemical,
hydrological, and reducing or oxidizing conditions.
Over the last 20 years, aqueous U(VI): soil mineral
interactions have been studied by geochemists, soil
chemists, clay and soil mineralogists, and the
progress in some areas is remarkable. Although a
mechanistic description and understanding of the
complex interactions involving U and soil minerals in
natural systems is currently difficult, results from
carefully designed and executed field and laboratory
experiments with these materials have improved our
understanding of the heterogeneous system’s behavior and U contaminant mobility and transport. There
N. P. Qafoku (&) J. P. Icenhower
Pacific Northwest National Laboratory, P.O. Box 999,
MSIN: K3-61, Richland, WA 99352, USA
e-mail: [email protected]
are, however, areas that warrant further exploration
and study. Numerous research publications were
reviewed in this paper to present recent important
findings to reveal the current level of the understanding of the U(VI) interactions with soil minerals, and
to provide ideas for future needs and research
directions.
Keywords Uranium U(VI) U(IV) Adsorption Desorption Redox reactions Soils Sediments Heterogeneous natural media Soil minerals Fe oxides Phyllosilicates Calcite
1 Introduction
1.1 The extent of U contamination
Mainly because of its essential role in the production
of nuclear weapons, uranium (U) is a common
contaminant at numerous sites throughout the world.
For example, U is a common contaminant at sites in
the United States of America (USA), where production of nuclear weapons and handling of U in various
forms has occurred (Riley et al. 1992). Anthropologic
sources of U contamination belong to three categories: (i) U from weapons production; (ii) U from
nuclear energy activities; and (iii) U from scientific
and other uses (Todorov and Ilieva 2006). Elevated
123
372
leads to the formation of soluble Ca2UO2(CO3)3
which inhibits microbial reduction (Brooks et al.
2003). While it is not known if this latter species is
susceptible to heterogeneous reduction, it is likely that
lacking a net charge it has a higher propensity to
interact with particle surfaces than its Ca-free, anionic
counterparts. The presence of aqueous Ca2UO2(CO3)3
in the pore-waters of many contaminated sites warrants an assessment of the redox reactivity of this
molecule.
4.2 Soil mineral role in U(VI) redox reactions
While the redox energetics and kinetics of uranyl
coordination complexes in the aqueous phase are
relatively well studied (Morris 2002), this appears to
be not true for the soil mineral mediated U(VI)
reduction reaction. A slow homogeneous reaction
may be accelerated in the presence of solids because
reactants may be concentrated on surfaces allowing
for longer lives of the encounter complexes. Surfaces
may also increase the reaction driving force.
Redox processes are by nature a series of coupled
reactions, which may in turn be coupled with other
reactions (adsorption/desorption, dissolution/precipitation) and processes (hydrologic, physical and
chemical) that occur in soils, sediments and vadose
zones during the transport of contaminants. This
coupling will affect the extent of these reactions in
geochemical systems, and the overall mobility of U.
The role of soil minerals in the redox reactions that
occur in soils, sediments and aquifers is multifarious,
and their chemical (sorption capacity, surface charge
and composition), and physical properties (with or
without expandable layers) are important determinants
of the extent of their participation in transport controlled, coupled adsorption/desorption, dissolution/
precipitation and redox reactions of contaminants,
although the extent of their involvement is not well
understood or studied.
Stumm et al. have emphasized the importance of
coupled geochemical processes and reactions with the
Fe(II) and Fe(III) redox transformations (Stumm
1992; Stumm and Morgan 1996), and numerous
studies have shown that sorbed Fe(II) is involved in
redox reactions with carbon tetrachloride (Elsner
et al. 2004b), pentachloronitrobenzene (Klupinski
et al. 2004), polyhalogenated methanes (Pecher et al.
2002), oxime carbamate pesticides (Strathmann and
123
Rev Environ Sci Biotechnol (2008) 7:355–380
Stone 2003), and organic contaminants such as 4chloronitrobenzene and hexachloroethane (Elsner
et al. 2004a). The catalytic role of soil minerals in
the redox reactions that occur in soils, sediments and
aquifers is, therefore, clearly demonstrated in the
recent literature (Liger et al. 1999; Strathmann and
Stone 2003; Elsner et al. 2004a, 2004b; Fredrickson
et al. 2004; Ilton et al. 2004; Klupinski et al. 2004).
Potentially important U(VI) reductants in lowtemperature geochemical systems are aqueous and
structural Fe(II), sulfides, and organic matter (Liger
et al. 1999). Fe(II) is abundant in many suboxic and
anoxic soils and sediments (Anderson et al. 1994)
and numerous observations indicate that reduction of
redox sensitive elements can occur in soils and
sediments in the presence of inorganic Fe(II).
Aqueous U(VI) is not involved in homogeneous
redox reactions with aqueous Fe(II) in near neutral
and basic pH (Liger et al. 1999). However, with
Fe(II) sorbed to a surface (Charlet et al. 1998; Liger
et al. 1999; Boyanov et al. 2007) or as a constituent
in a smectite clay (Giaquinta et al. 1997) and biotite
(Ilton et al. 2004), U(VI) reduction proceeds faster.
Sorbed or structural Fe(II) may be also present or
formed when strong reductants (such as dithionite,
H2S gas and Ca-polysulfide liquid) react with Fe(III)
oxides or other soil minerals, such as phyllosilicates.
Heterogeneous, abiotic reduction of U(VI) by
Fe(II) has been observed in zero valent iron (Noubactep et al. 2003), mixed Fe(II)/Fe(III) green rust
(O’Loughlin et al. 2003), nano-magnetite (Missana
et al. 2003) and biotite (Ilton et al. 2004) in acid to
circumneutral pH, and in anoxic-CO2-free systems.
Little information is available on U(VI) reduction by
these Fe(II)-containing phases in alkaline fluids were
anionic uranyl-carbonates and neutral Ca-uranylcarbonate are the dominant aqueous U(VI) species.
Unlike the kinetically inert Cr(III), oxidation of
reduced U(IV) with the intrusion of O2 may occur in
sediments. However, the oxidation of the reduced
species appears to be diffusion and residence time
controlled. U(IV) oxidation from reduced ISRM
sediments shows a slow process that may be
confounded by other reduced species (Szecsody et al.
1998). In a more recent study (Moon et al. 2007),
reoxidation of microbially reduced U with either O2
or nitrate supplied as the oxidant, was investigated.
They found that U reduction occurred simultaneously
with Fe reduction as the dominant electron accepting
Rev Environ Sci Biotechnol (2008) 7:355–380
process. Both O2 and nitrate remobilized the majority
(88 and 97%, respectively) of the U precipitated
during bioreduction within 54 days. Although O2 is
more thermodynamically favorable an oxidant than
nitrate, U oxidation by nitrate occurred significantly
faster at the beginning of the experiments, due to O2
reacting more strongly with other reduced compounds (Moon et al. 2007).
Regardless of whether the redox process is homogeneous or heterogeneous, near-field or far-field, or
whether the reactive surfaces are neoformed or a
natural part of the sediment mineral assemblage,
hydrologic flow characteristics will affect reactions,
reaction rates, and extent of reaction via the transport
of either oxidants or reductants to or away from an
Fe(II) source. The processes, reactions and conditions
that affect the rate of the surface (soil mineral)
mediated redox reactions under advection, mass
transfer and/or diffusion limiting condition, are not
well studied in natural mixtures of soil minerals.
4.3 Recent findings and future trends
There are interesting developments in the area of U
redox reactions in the presence of solid phases. We
previously mentioned that in the presence of soil
mineral catalysts, the redox reaction may occur on the
Fe(II) exchanged surface of a solid phase, or on the
surface of a Fe(II)-bearing mineral. However, the
explanation for the enhanced reactivity of sorbed
Fe(II) remains ambiguous, although recent studies
have shed some light on this subject (Boyanov et al.
2007). These authors conducted experiments to gain
further insights into the U–Fe redox process at a
complexing, non-conducting surface such as carboxyl-functionalized polystyrene. It was reported that
in the Fe ? surface carboxyl system, a transition
from monomeric to oligomeric Fe(II) surface species
was observed between pH 7.5 and pH 8.4 (Boyanov
et al. 2007). In the U ? surface carboxyl system, the
U(VI) cation was adsorbed as a mononuclear uranylcarboxyl complex at both pH 7.5 and 8.4. In the
ternary U ? Fe ? surface carboxyl system, U(VI)
was not reduced by the solvated or adsorbed Fe(II) at
pH 7.5 over a 4-month period, whereas complete and
rapid reduction to U(IV) nanoparticles occurred at pH
8.4 (Boyanov et al. 2007).
Boyanov et al. (2007) also reported that the U(IV)
product reoxidized rapidly upon exposure to air, but it
373
was stable over a 4-month period under anoxic
conditions. The U(IV)–Fe coordination was consistent
with an inner-sphere electron transfer mechanism
between the redox centers and involvement of Fe(II)
atoms in both steps of the reduction from U(VI) to
U(IV). The inability of Fe(II) to reduce U(VI) in
solution and at pH 7.5 in the U ? Fe ? carboxyl
system was explained by the formation of a transient,
‘‘dead-end’’ U(V)–Fe(III) complex that blocked the
U(V) disproportionation pathway after the first electron transfer. These authors suggested that the
increased reactivity at pH 8.4 relative to pH 7.5 could
be attributed to the reaction of U(VI) with an Fe(II)
oligomer, whereby the bonds between Fe atoms
facilitated the transfer of a second electron to the
hypothetical U(V)–Fe(III) intermediate, which may
explain the commonly observed higher efficiency of
uranyl reduction by adsorbed or structural Fe(II)
relative to aqueous Fe(II) (Boyanov et al. 2007).
However, recent results have shown that at low
Fe(II) concentrations, sorbed Fe(II) species on hematite are transient and quickly undergo interfacial
electron transfer with structural Fe(III) (LareseCasanova and Scherrer 2007), forming a Fe(III)
surface coating layer. The formation of a stable,
sorbed Fe(II) phase in hematite was observed only at
higher Fe(II) concentrations and coincided with the
macroscopically observed change in isotherm slope,
and with the estimated surface site saturation,
suggesting that the finite capacity for interfacial
electron transfer is influenced by surface properties
(Larese-Casanova and Scherrer 2007).
Definitely, the heterogeneous reduction of U(VI)
by sorbed Fe(II) is an important pathway for immobilization of aqueous U(VI) species in subsurface
environments, where the homogeneous redox reaction of the aqueous U(VI):Fe(II) couple is slow.
However, many questions remain unanswered:
(i)
What is the extent of sorbed U(VI) reduction by
sorbed Fe(II) on hematite or other Fe(III) oxides
and hydroxide surfaces?
(ii) Is sorbed U(VI) more competitive than structural Fe(III) for the electrons of sorbed Fe(II) on
hematite?
(iii) What are the conditions that may control the
competition between sorbed U(VI) and structural Fe(III) for sorbed Fe(II) electrons?
(iv) How do other Fe(III) oxides behave?
123
Rev Environ Sci Biotechnol (2008) 7:355–380
Bostick BC, Fendorf S, Barnett MO, Jardine PM, Brooks SC
(2002) Uranyl surface complexes formed on subsurface
media from DOE facilities. Soil Sci Soc Am J 66:99–108
Boyanov MI, O’Loughlin EJ, Roden EE, Fein JB, Kemner KM
(2007) Adsorption of Fe(II) and U(VI) to carboxyl-functionalized microspheres: the influence of speciation on
uranyl reduction studied by titration and XAFS. Geochim
Cosmochim Acta 71:1898–1912. doi:10.1016/j.gca.2007.
01.025
Braithwaite A, Livens FR, Richardson S, Howe MT, Goulding
KWT (1997) Kinetically controlled release of uranium
from soils. Eur J Soil Sci 48:661–673. doi:10.1046/j.13652389.1997.00121.x
Braithwaite A, Richardson S, Moyes LN, Livens FR, Bunker
DJ, Hughes CR (2000) Sorption kinetics of uranium-238
and neptunium-237 on glacial sediment. Czech J Phys
50:265–269. doi:10.1023/A:1022842808820
Brooks SC, Fredrickson JK, Carroll SL, Kennedy DW, Zachara
JM, Plymale AE et al (2003) Inhibition of bacterial U(VI)
reduction by calcium. Environ Sci Technol 37:1850–
1858. doi:10.1021/es0210042
Carroll SA, Bruno J (1991) Mineral-solution interactions in the
U(VI)-CO2-H2O system. Radiochim Acta 52(53):187–193
Carroll SA, Bruno J, Petit J-C, Dran J-C (1992) Interactions of
U(VI), Nd, and Th(IV) at the calcite-solution interface.
Radiochim Acta 58(59):245–252
Catalano JG, Brown GE (2004) Analysis of uranyl-bearing
phases by EXAFS spectroscopy: interferences, multiple
scattering, accuracy of structural parameters, and spectral
differences. Am Min 89:1004–1021
Catalano JG, Brown GE (2005) Uranyl adsorption onto
montmorillonite: evaluation of binding sites and carbonate complexation. Geochim Cosmochim Acta 69:2995–
3005. doi:10.1016/j.gca.2005.01.025
Catalano JG, McKinley JP, Zachara JM, Heald SM, Smith SC,
Brown GE (2006) Changes in uranium speciation through a
depth sequence of contaminated Hanford sediments.
Environ Sci Technol 40:2517–2524. doi:10.1021/
es0520969
Chadwick O, Nettleton WD (1990) Micromorphologic evidence
of adhesive and cohesive forces in soil cementation. Soil Sci
19:207–212. doi:10.1016/S0166-2481(08)70332-5
Charlet L, Silvester E, Liger E (1998) N-compound reduction
and actinide immobilization in surfacial fluids by Fe(II):the
surface =FeIIOFeIIOH0 species, as major reductant. Chem
Geol 151:85–93. doi:10.1016/S0009-2541(98)00072-2
Cheng T, Barnett MO, Roden EE, Zhuang JL (2006) Effects of
solid-to-solution ratio on uranium(VI) adsorption and its
implications. Environ Sci Technol 40:3243–3247. doi:
10.1021/es051771b
Clark DL, Hobart DE, Neu MP (1995) Actinide carbonate
complexes and their importance in actinide environmental
chemistry. Chem Rev 95:25–48. doi:10.1021/cr00033a002
Cornell RM, Schwertmann U (1996) The iron oxides. VCH
Publ, Weinheim
Cornell RM, Schwertmann U (2003) The iron oxides. Structure, properties, reactions, occurrences and uses, 2nd edn.
WILEY-VCH
Culver TB, Hallisey SP, Sahoo D, Deitsch JJ, Smith JA (1997)
Modeling the desorption of organic contaminants from
long-term contaminated soil using distributed mass
375
transfer rates. Environ Sci Technol 31:1581–1588. doi:
10.1021/es9600946
Curtis GP, Fox P, Kohler M, Davis JA (2004) Comparison of in
situ uranium Kd values with a laboratory determined
surface complexation model. Appl Geochem 19:1643–
1653. doi:10.1016/j.apgeochem.2004.03.004
Curtis GP, Davis JA, Naftz DL (2006) Simulation of reactive
transport of uranium(VI) in groundwater with variable
chemical conditions. Water Resour Res 42:W04404. doi:
10.1029/2005WR003979
Davis JA (2001) Surface complexation modeling of uranium(VI) adsorption on natural mineral assemblages. U.S.
Geological Survey, NUREG/CR–6708
Davis JA, Curtis GP (2003) Application of surface complexation modeling to describe uranium(VI) adsorption and
retardation at the uranium mill tailings site at Naturita,
Colorado. NUREG Report CR-6820. U.S. Nuclear Regulatory Commission, Rockville
Davis JA, Meece DE, Kohler M, Curtis GP (2004) Approaches
to surface complexation modeling of Uranium(VI)
adsorption on aquifer sediments. Geochim Cosmochim
Acta 68:3621–3641. doi:10.1016/j.gca.2004.03.003
Davis JA, Curtis GP, Wilkins MJ, Kohler M, Fox P, Naftz DL
et al (2006) Processes affecting transport of uranium in a
suboxic aquifer. Phys Chem Earth 31:548–555
de Jong WA, Apra E, Windus TL, Nichols JA, Harrison RJ,
Gutowski KE et al (2005) Complexation of the carbonate,
nitrate, and acetate anions with the uranyl dication: density
functional studies with relativistic effective core potentials.
J Phys Chem A 109:11568–11577. doi:10.1021/jp0541462
Dodge CJ, Francis AJ, Gillow JB, Halada GP, Eng C, Clayton
CR (2002) Association of uranium with iron oxides typically formed on corroding steel surfaces. Environ Sci
Technol 36:3504–3511. doi:10.1021/es011450?
Dong WM, Brooks SC (2006) Determination of the formation
constants of ternary complexes of uranyl and carbonate
with alkaline earth metals (Mg2?, Ca2?, Sr2?, and Ba2?)
using anion exchange method. Environ Sci Technol
40:4689–4695. doi:10.1021/es0606327
Dong W, Ball WP, Liu C, Wang Z, Stone AT, Bai J et al (2005)
Influence of calcite and dissolved calcium on uranium(VI)
sorption to a Hanford subsurface sediment. Environ Sci
Technol 39:7949–7955. doi:10.1021/es0505088
Dong WM, Xie GB, Miller TR, Franklin MP, Oxenberg TP,
Bouwer EJ et al (2006) Sorption and bioreduction of
hexavalent uranium at a military facility by the Chesapeake Bay. Environ Pollut 142:132–142. doi:10.1016/j.
envpol.2005.09.008
Duff MC, Amrhein C (1996) Uranium(VI) adsorption on
goethite and soil in carbonate solutions. Soil Sci Soc Am J
60:1393–1400
Duff MC, Hunter DB, Bertsch PM, Amrhein C (1999) Factors
influencing uranium reduction and solubility in evaporation pond sediments. Biogeochem 45:95–114
Duff MC, Coughlin JU, Hunter DB (2002) Uranium co-precipitation with iron oxide minerals. Geochim Cosmochim
Acta 66:3533–3547. doi:10.1016/S0016-7037(02)00953-5
Elsner M, Schwarzenbach RP, Haderlein SB (2004a) Reactivity of Fe(II)-bearing minerals toward reductive
transformation of organic contaminants. Environ Sci
Technol 38:799–807. doi:10.1021/es0345569
123
Real–Time Speciation of Uranium during Active
Bioremediation and U„IV… Reoxidation
John Komlos1; Bhoopesh Mishra2; Antonio Lanzirotti3; Satish C. B. Myneni4; and Peter R. Jaffé5
Abstract: The biological reduction of uranium from soluble U共VI兲 to insoluble U共IV兲 has shown potential to prevent uranium migration
in groundwater. To gain insight into the extent of uranium reduction that can occur during biostimulation and to what degree U共IV兲
reoxidation will occur under field relevant conditions after biostimulation is terminated, X-ray absorption near edge structure 共XANES兲
spectroscopy was used to monitor: 共1兲 uranium speciation in situ in a flowing column while active reduction was occurring; and 共2兲 in situ
postbiostimulation uranium stability and speciation when exposed to incoming oxic water. Results show that after 70 days of bioreduction
in a high 共30 mM兲 bicarbonate solution, the majority 共⬎90% 兲 of the uranium in the column was immobilized as U共IV兲. After acetate
addition was terminated and oxic water entered the column, in situ real-time XANES analysis showed that U共IV兲 reoxidation to U共VI兲
共and subsequent remobilization兲 occurred rapidly 共on the order of minutes兲 within the reach of the oxygen front and the spatial and
temporal XANES spectra captured during reoxidation allowed for real-time uranium reoxidation rates to be calculated.
DOI: 10.1061/共ASCE兲0733-9372共2008兲134:2共78兲
CE Database subject headings: Uranium; Biodegradation; Oxidation; Iron; Rates.
Introduction
Uranium contamination is a concern at numerous U.S. Department of Energy facilities throughout the United States. Uranium
exists in nature as either U共VI兲 or U共IV兲. The oxidized form,
U共VI兲, tends to be soluble and may exist as different ions depend2−
ing on the alkalinity and pH 关e.g. UO2+
2 , UO2 共CO3兲2 兴. It typically transports in flowing groundwater, whereas the reduced
form of uranium, U共IV兲, forms insoluble minerals such as UO2
共uraninite兲 that precipitates out of solution. The bioreduction of
U共VI兲 to U共IV兲 is an anaerobic process that has been shown to
occur after nitrate is consumed 共Finneran et al. 2002; Senko et al.
2002兲 and during either iron and/or sulfate reducing conditions
共Abdelouas et al. 1999; Anderson et al. 2003; Lovley and Phillips
1992兲. The precipitation of uranium from groundwater through
the addition of an electron donor to stimulate the uranium reducing microbial population has shown potential to prevent uranium
1
Research Staff, Dept. of Civil and Environmental Engineering,
Princeton Univ., Princeton, NJ 08544. E-mail: [email protected]
2
Research Associate, Dept. of Geosciences, Princeton Univ.,
Princeton, NJ 08544. E-mail: [email protected]
3
Senior Research Associate, The Univ. of Chicago—Center for
Advanced Radiation Sources at the National Synchrotron Light Source,
Brookhaven National Laboratory, Upton, NY 11973. E-mail:
[email protected]
4
Associate Professor, Dept. of Geosciences, Princeton Univ.,
Princeton, NJ 08544. E-mail: [email protected]
5
Professor, Dept. of Civil and Environmental Engineering, Princeton
Univ., Princeton, NJ 08544 共corresponding author兲. E-mail: jaffe@
princeton.edu
Note. Discussion open until July 1, 2008. Separate discussions must
be submitted for individual papers. To extend the closing date by one
month, a written request must be filed with the ASCE Managing Editor.
The manuscript for this paper was submitted for review and possible
publication on April 11, 2007; approved on August 3, 2007. This paper is
part of the Journal of Environmental Engineering, Vol. 134, No. 2,
February 1, 2008. ©ASCE, ISSN 0733-9372/2008/2-78–86/$25.00.
migration from contaminated sites 共Anderson et al. 2003; Chang
et al. 2005; Istok et al. 2004兲.
The nature of U共VI兲 reduction, however, in anoxic sediments
is poorly understood. Some studies have shown that the majority
of uranium in sediments under biologically reducing conditions
was present as U共IV兲 共Michalsen et al. 2006; Sani et al. 2005兲,
although other studies have shown that not all of the uranium
measured on mineral surfaces under reducing conditions was reduced 共Gu et al. 2005a; Jeon et al. 2004; Ortiz-Bernad et al. 2004;
Wan et al. 2005兲 and the reasons for the presence of U共VI兲 under
reducing conditions are inconclusive. Possible explanations are
that U共VI兲 was adsorbed and unavailable for microbial reduction.
U共VI兲 can adsorb to Fe共III兲–共hydr兲oxides 共Giammar and Hering
2001; Jeon et al. 2004兲 or form relatively insoluble complexes
with PO3−
4 共Cheng et al. 2006; Langmuir 1978兲, and research have
shown that U共VI兲 sorption can limit the rate and extent of microbial U共VI兲 reduction 共Jeon et al. 2004兲. Calcium can suppress
U共VI兲 sorption 共Zheng et al. 2003兲 but has been shown to also
strongly inhibit U共VI兲 reduction 共Brooks et al. 2003兲 and to
increase abiotic ferrihydrite-dependent U共IV兲 oxidation 共GinderVogel et al. 2006兲. In carbonate containing groundwater at
circumneutral pH, U共VI兲 forms strong soluble complexes with
2−
4−
CO2−
3 共e.g., UO2CO3, UO2共CO3兲2 , UO2共CO3兲3 兲 共Fredrickson
et al. 2000兲 that absorb poorly with mineral surfaces such as
Fe共III兲 共hydr兲oxides 共Duff and Amrhein 1996; Hsi and Langmuir
1985兲 and clays due to the neutral or anionic charge 共Fredrickson
et al. 2000兲. In 共bi兲carbonate containing waters under reducing
conditions, the majority of the uranium has been shown to be
U共IV兲 共Sani et al. 2005; Wan et al. 2005兲 or as a combination of
U共VI兲 and U共IV兲 共Gu et al. 2005a; Wan et al. 2005兲 and the
reasons for the discrepancy are not fully understood. The U共VI兲
complexes mentioned above could decrease U共VI兲 bioavailability
in the 共bi兲carbonate solution. In addition, U共IV兲 can be anaerobically oxidized by denitrification byproducts 共Senko et al. 2002兲
and, under electron donor limitation, U共IV兲 can be oxidized by
Fe共III兲 共Ginder-Vogel et al. 2006; Sani et al. 2005兲 and Mn–
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then restarted on Day 65 at the NSLS facility and operated under
the same biostimulation conditions described above for 4 days to
allow for the system to come back to pretransfer conditions in
case some changes occurred during transportation. XANES spectroscopic analysis was performed at Beamline X26A on Day 69
while the column was maintained under flowing conditions.
In order to mimic the cessation of electron donor addition after
biostimulation in the field, the column reoxidation was initiated
on Day 70 by stopping the electron donor addition and substituting the CO2 / N2 共20:80兲 gas supplied to the influent media with a
gas containing O2 / CO2 / N2 共20:20:60兲. The influent media was
the same used for bioreduction minus NH4Cl and the vitamin
solution 共to prevent dissolved oxygen consumption from ammonia oxidation兲. CO2 共20%兲 and bicarbonate 共30 mM兲 addition was
continued during reoxidation to maintain a pH of 7. The column
remained at NSLS during reoxidation and was also analyzed on
Days 71 and 99. The column was operated at 22–25°C until it was
shipped on ice overnight back to Princeton University where it
was destructively sampled in an anaerobic glove box 共3:97
H2 : N2兲.
Ex Situ XANES Sample Preparation
Sediment used for the ex situ XANES spectroscopy analysis was
from a column bioreduced under the same conditions as described
above except for undergoing reducing conditions for an additional
34 days 共and no reoxidation兲. The column was taken apart and the
sediment removed from the column using a long spatula in an
anaerobic glove box 共3:97 H2 : N2兲, placed in a polycarbonate
sample holder and sealed on both sides with two layers of Kapton
tape. The samples were placed in a pressurized chamber filled
with N2 gas and transported to Brookhaven National Laboratory
for XANES analysis. The XANES spectra of these samples were
collected in air.
Batch Uranium Reduction Experiment
One g of RABS sediment and 9.3 mL of the influent media described above 共except that the uranyl acetate concentration was
1 mM兲 was added to 15 mL plastic centrifuge tubes and purged
for 30 min with a 20% CO2 / 80% N2 gas mixture prior to being
capped with a thick rubber stopper. At the start of the experiment,
U共VI兲 bioreduction was facilitated by the addition of 0.2 mL of
1 M sodium acetate 共resulting in 20 mM acetate after mixing兲 and
0.5 mL of G. metallireducens growth culture 共prepared and rinsed
as described above兲. The samples were stored in the dark at
22–23°C until analyzed after 35 days using extended X-ray absorption fine structure 共EXAFS兲 spectroscopy as described below.
Plastic centrifuge tubes were used to allow for EXAFS analysis of
the sediment in situ through the plastic.
Analytical Measurements
Effluent Fe共II兲 concentrations were measured by adding 0.5 mL
of effluent solution to 0.5 mL of 1 M HCl and analyzing after 1 h
extraction using ferrozine 共Lovley and Phillips 1987兲. Dissolved
oxygen was measured using a Corning 317 dissolved oxygen
共DO兲 meter fitted to an in-line sampling device attached to the
effluent of the column. Anions 共bromide, acetate, sulfate, phosphate兲 were analyzed using a Dionex DX500 ion chromatograph
equipped with a CD25 conductivity detector and a Dionex IonPac
AS14-4 mm column. Influent and effluent U共VI兲 concentrations
were analyzed using reversed phased chromatography coupled to
postcolumn derivatization with the dye Arsenazo III 共SigmaAldrich兲 as described by Lack et al. 共2002兲. All samples were
filtered 共0.2 ␮m兲 and stored at 4°C until analyzed. The total uranium concentration 关U共VI兲 plus U共IV兲兴 in the sediment was quantified by adding 2 – 3 g of sediment to 5 mL of 0.2 M NaHCO3.
The samples were extracted under aerobic conditions to oxidize
U共IV兲 to U共VI兲 for 24 h, filtered 共0.2 ␮m兲, and stored at 4°C
until U共VI兲 was analyzed as described above.
XANES Spectroscopy Measurements
XANES spectroscopy was used to provide information about
the oxidation state of uranium. Uranium L3 edge 共17,166 eV兲
XANES spectroscopy measurements were performed at X26A at
the NSLS 共Brookhaven National Laboratory兲. X26A is a hard
X-ray microprobe bending magnet 共BM兲 beamline. The energy of
the incident X-rays was scanned by using a Si共111兲 reflection
plane of a channel-cut monochromator cooled to 11°C using a
Neslab chiller. The X-ray spot size used for these measurements
was set to 5 ⫻ 5 ␮m. The fluorescence signal of the soil column
was measured using a Canberra 9-element Ge array detector. The
scans were aligned by collecting uranyl acetate solution data after
every 3–4 XANES scans. Scan to scan variation in the energy
calibration of the monochromator was within 0.2 eV even after
several hours. However, a bigger difference was usually seen after
each beam refill or beam dump.
Step scans 共energy scans with 0.5 eV step size, near the edge
and 5.0 eV far below and above the edge兲 were used with an
integration time of 5 – 15 s per point depending on the signal to
noise ratio of the spectra. The bioreduced samples were scanned
from −200 to +300 eV relative to edge position to ensure proper
normalization and background removal of the data. However, a
faster scanning setup was required to monitor the in situ reoxidation profile of the column. Hence the data collected for 0 – 2 h
reoxidation was scanned from −50 to +150 eV relative to the
edge position 共resulting in a scan time of ⬃20 min兲. This energy
range was sufficient for linear combination fitting 共LCF兲 of
XANES data. All XANES data reported in this study were normalized and fit in this data range for consistency.
XANES Spectroscopy Data Processing and Fitting
Interactive data language 共IDL兲 software associated with Beamline X26A was used to collect data. The data were analyzed using
ATHENA 共Ravel and Newville 2005兲 which is based on AUTOBK 共Newville et al. 1993兲 to remove the background. ATHENA was also used for LCF of the uranium data to quantify the
relative amount of U共IV兲 compared to U共VI兲 in a given spectra.
The fitting was done in the normalized ␮ 共E兲 space. A fitting
range of −50– + 150 eV was used for proper normalization of the
XANES spectra. Since uranium is known to be stable in +IV and
+VI oxidation states only, powdered UO2 and UO3 XANES spectra were used as U共IV兲 and U共VI兲 standards for the LCF of the
sediment samples. The sum of their contribution in the unknown
samples was forced to sum to 1. A lower R factor and ␹2v values
were used as the criteria for the goodness of fit. The accuracy of
the valance state determination of uranium from the XANES data
was estimated to be 10–15%, which is similar to the accuracies
previously reported for this analysis 共Boyanov et al. 2007; Jeon
et al. 2004兲. Hence 90 and 100% bioreduction should be considered roughly the same for the treatment presented in this study.
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Fig. 3. Effluent Fe共II兲 concentrations during bioreduction
231 ␮moles of U共VI兲 was removed between the influent and effluent of the column. The effluent pH remained constant at 7.0
during biostimulation. Fe共III兲 reduction 关and subsequent Fe共II兲
production, Fig. 3兴 occurred simultaneously with U共VI兲 reduction
共Fig. 2兲 although neither process reached steady-state conditions
by the start of reoxidation on Day 70, indicating that the overall
biological activity was still increasing. The effluent Fe 共II兲 concentration was 165 ␮M after 70 days of biostimulation. Removal
of sulfate between the influent and effluent of the column was first
detected on Day 16 with 70–96% of the influent 9 ␮M concentration removed between Day 26 and the end of biostimulation
共data not shown兲. Effluent acetate concentrations remained above
1 mM throughout biostimulation 共acetate was not limiting兲. Phosphate 共14 ␮M兲 present in the influent media was not detected at
the column effluent prior to acetate addition and remained below
detection at the effluent throughout biostimulation 共data not
shown兲. The lack of phosphate at the effluent is in contrast to
U共VI兲, whose effluent concentration equaled the influent concentration prior to biostimulation and slowly decreased with time of
biostimulation 共Fig. 2兲. The discrepancy between the trends of
phosphate and U共VI兲 removal indicates that the U共VI兲 removal in
these experiments was not dependent on complexation with phosphate which corresponds to previous work 共Sandino and Bruno
1992兲 showing that U共VI兲 will be associated with aqueous phos2−
phate complexes when the 关PO3−
4 兴T / 关CO3 兴T ratio is greater than
−1
10 共which is higher than the ratio in this study, 0.0004兲. In
addition, the calcium concentration in the feed media 共0.023 mM兲
was lower than that observed to inhibit U共VI兲 reduction 共Brooks
et al. 2003兲. Therefore, for these conditions, phosphate and calcium complexes did not appear to play a role in U共VI兲 reduction.
Fig. 4. XANES spectrum of U共VI兲 and U共IV兲 standard
Fig. 5. XANES spectrum showing U speciation: 共a兲 after 70 days of
bioreduction; 共b兲 during first 2 h of reoxidation; 共c兲 1 day after reoxidation; and 共d兲 29 days after reoxidation. Relative ratio of U共IV兲 to
total uranium is shown 共in percentage兲 in each XANES spectra. Also
time 共t兲 is indicated in min for the first 2 h of reoxidation.
Uranium Speciation during Bioreduction
Fig. 4 compares uranium XANES data from UO2 and UO3 standards. A higher energy position of the absorption edge and a
shoulder at 17,190 eV indicate uranium in the +VI valance state
and uranyl coordination geometry. A lower energy position of the
edge, a lack of the shoulder at 17,190 eV, and higher amplitude of
the peak immediately after the edge indicate uranium in the +IV
valance state 共Boyanov et al. 2007; Ilton et al. 2006; Michalsen
et al. 2006; O’Loughlin et al. 2003; Wan et al. 2005; Wu et al.
2006兲.
Fig. 5共a兲 shows XANES spectra of the uranium sediment at
different heights of the column after 70 days of bioreduction. LCF
indicated that the majority 共⬎ 90%兲 of the uranium was found as
U共IV兲. Further, X-ray fluorescence mapping of the sediment
samples from a column experiment using the same sediment
under identical experimental conditions 共though bioreduced for
slightly longer, 104 days兲 indicated that uranium was homoge-
Fig. 6. XANES spectra of in situ 共solid line兲 and ex situ 共dash line兲
measurements at 5 and 12 cm from influx into column
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Table 1. EXAFS Fitting Parameters
Path
Coordination
number 共N兲
Bond length 共R兲
共Å兲
Debye–Waller
factor 共␴2兲
共10−3 Å2兲
U-O1
7.5± 0.6
2.34± 0.01
12.2± 1.5
U-U1
5.3± 1.4
3.84± 0.01
8.6± 2.8
U-MS
7.5± 0.6
4.68± 0.02
24.4± 3.0
Note: MS denotes two multiple scattering paths: U-O1-U-O1. Passive
electron reduction factor 共S20兲 was set at 0.9, and ⌬E was 3.1± 0.6. Fourier
transform was done over the data range of 2.3– 10.2Å−1, and the fit range
was 1.2– 4.4 Å.
Fig. 7. k2 weighed ␹共k兲 data for bioreduced uranium sediment. Data
range used for Fourier transform was 2.3– 10.2k 共Å−1兲.
neously distributed throughout the sediment 共data not shown兲.
The lack of uranium hotspots in the sediment indicates that the
measured XANES spectra are representative of the sediment and
do not represent any localized feature. Complete reduction of
U共VI兲 to U共IV兲 under bioreduced conditions contradicted
XANES spectroscopy performed on ex situ sediment samples
from the 104 day bioreduction column mentioned above where
not all of the uranium was reduced 共Fig. 6兲. The discrepancy
between the ex situ and in situ XANES analysis was unexpected
and could have been caused by oxygen contamination during
sample preparation, transport, or analysis 共even though efforts
were taken to provide anaerobic conditions兲. The discrepancy
could also have been due to electron donor limitation once acetate
addition was terminated and the sample was removed from the
column, thus allowing U共IV兲 to act as an electron donor for
Fe 共III兲 reduction. Ferrihydrite has been shown to oxidize U共IV兲
under conditions with electron donor limitation 共Ginder-Vogel
et al. 2006兲 and additional research is needed to determine the
impact of available Fe 共III兲 in the RABS sediment on U共IV兲 stability under electron donor limiting conditions.
EXAFS analysis of a sample run in a batch experiment with
similar experimental conditions as the bioreduced column sediment was performed to further investigate the speciation of the
uranium in the column under active bioreduction conditions. The
EXAFS data quality can be seen from the averaged EXAFS
␹共k兲 * k2 spectrum in Fig. 7. Fig. 8 shows the ␹共k兲 * k2 magnitude
and real part of the Fourier transform and fit of the bioreduced
uranium sediment sample. The peak at 1.8 Å is due to the cubical
oxygen shell in UO2, and the double peak between 2.5 and 4.5 Å
is mostly due to the 12-member U shell at 3.87 Å in UO2. The 1:1
ratio of the peak amplitudes between 2.5 and 4.5 Å is consistent
with previously published results for uranium nanoparticulates
共O’Loughlin et al. 2003兲 which is different from the 1:2 ratio
found for a pure UO2 standard 共Boyanov et al. 2007; O’Loughlin
et al. 2003兲. EXAFS modeling of this sample was done using the
model for UO2 crystal structure 共O’Loughlin et al. 2003兲. Best fit
values for the EXAFS analysis are listed in Table 1. The best fit
value for the number of first shell oxygen atoms is 7.5± 0.6 at
2.34 Å, which is consistent with the UO2 structure of eight oxygen atoms in the first shell at the same distance. The drop in
average U – U coordination from 12 in crystalline UO2 to 5.3± 1.4
in the bioreduced sediment could be either due to a thin coating of
UO2 on Fe particles present in the sediment or the formation of
the nanometer sized uraninite particles, both of which indicate the
formation of uraninite nanophases.
Fig. 8. Magnitude 共a兲; real part 共b兲 of Fourier transform of ␹共k兲 * k2 best-fit model and data from bioreduced sediment
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共Contract No. DE-FG02-92ER14244兲, and DOE–Office of Biological and Environmental Research, ERSD. Use of NSLS was
supported by DOE under Contract No. DE-AC02-98CH10886.
References
Abdelouas, A., Lutze, W., and Nuttall, H. E. 共1999兲. “Oxidative dissolution of uraninite precipitated on Navajo sandstone.” J. Contam. Hydrol., 36共3–4兲, 353–375.
Anderson, R. T., et al. 共2003兲. “Stimulating the in situ activity of Geobacter species to remove uranium from the groundwater of a uraniumcontaminated aquifer.” Appl. Environ. Microbiol., 69共10兲, 5884–
5891.
Ankudinov, A. L., Ravel, B., Rehr, J. J., and Conradson, S. D. 共1998兲.
“Real-space multiple-scattering calculation and interpretation of
x-ray-absorption near-edge structure.” Phys. Rev. B, 58共12兲, 7565–
7576.
Boyanov, M. I., O’Loughlin, E. J., Roden, E. E., Fein, J. B., and Kemner,
K. M. 共2007兲. “Adsorption of Fe共II兲 and U共VI兲 to carboxylfunctionalized microspheres: The influence of speciation on uranyl
reduction studied by titration and XAFS.” Geochim. Cosmochim.
Acta, 71共8兲, 1898–1912.
Brooks, S. C., et al. 共2003兲. “Inhibition of bacterial U共VI兲 reduction by
calcium.” Environ. Sci. Technol., 37共9兲, 1850–1858.
Chang, Y. J., et al. 共2005兲. “Microbial incorporation of C-13-labeled acetate at the field scale: Detection of microbes responsible for reduction of U共VI兲.” Environ. Sci. Technol., 39共23兲, 9039–9048.
Cheng, T., Barnett, M. O., Roden, E. E., and Zhuang, J. L. 共2006兲. “Effects of solid-to-solution ratio on uranium共VI兲 adsorption and its implications.” Environ. Sci. Technol., 40共10兲, 3243–3247.
Duff, M. C., and Amrhein, C. 共1996兲. “Uranium共VI兲 adsorption on goethite and soil in carbonate solutions.” Soil Sci. Soc. Am. J., 60共5兲,
1393–1400.
Elias, D. A., Senko, J. M., and Krumholz, L. R. 共2003兲. “A procedure for
quantitation of total oxidized uranium for bioremediation studies.” J.
Offshore Mech. Arct. Eng., 53共3兲, 343–353.
Finneran, K. T., Housewright, M. E., and Lovley, D. R. 共2002兲. “Multiple
influences of nitrate on uranium solubility during bioremediation of
uranium-contaminated subsurface sediments.” Environmental Microbiology, 4共9兲, 510–516.
Fredrickson, J. K., Zachara, J. M., Kennedy, D. W., Duff, M. C., Gorby,
Y. A., Li, S. M. W., and Krupka, K. M. 共2000兲. “Reduction of U共VI兲
in goethite 共alpha-FeOOH兲 suspensions by a dissimilatory metalreducing bacterium.” Geochim. Cosmochim. Acta, 64共18兲, 3085–
3098.
Fredrickson, J. K., Zachara, J. M., Kennedy, D. W., Liu, C. X., Duff, M.
C., Hunter, D. B., and Dohnalkova, A. 共2002兲. “Influence of Mn oxides on the reduction of uranium共VI兲 by the metal-reducing bacterium
Shewanella putrefaciens.” Geochim. Cosmochim. Acta, 66共18兲, 3247–
3262.
Freeze, R. A., and Cherry, J. A. 共1979兲. Groundwater, Prentice-Hall,
Englewood Cliffs, N.J.
Giammar, D. E., and Hering, J. G. 共2001兲. “Time scales for sorptiondesorption and surface precipitation of uranyl on goethite.” Environ.
Sci. Technol., 35共16兲, 3332–3337.
Ginder-Vogel, M., Criddle, C. S., and Fendorf, S. 共2006兲. “Thermodynamic constraints on the oxidation of biogenic UO2 by Fe共III兲 共hydr兲
oxides.” Environ. Sci. Technol., 40共11兲, 3544–3550.
Gu, B. H., et al. 共2005a兲. “Bioreduction of uranium in a contaminated soil
column.” Environ. Sci. Technol., 39共13兲, 4841–4847.
Gu, B. H., Yan, H., Zhou, P., Watson, D. B., Park, M., and Istok, J.
共2005b兲. “Natural humics impact uranium bioreduction and oxidation.” Environ. Sci. Technol., 39共14兲, 5268–5275.
Hsi, C. K. D., and Langmuir, D. 共1985兲. “Adsorption of uranyl onto ferric
oxyhydroxides: Application of surface complexation site-binding
model.” Geochim. Cosmochim. Acta, 49, 1931–1941.
Ilton, E. S., Heald, S. M., Smith, S. C., Elbert, D., and Liu, C. X. 共2006兲.
“Reduction of uranyl in the interlayer region of low iron micas under
anoxic and aerobic conditions.” Environ. Sci. Technol., 40共16兲, 5003–
5009.
Istok, J. D., Senko, J. M., Krumholz, L. R., Watson, D., Bogle, M. A.,
Peacock, A., Chang, Y. J., and White, D. C. 共2004兲. “In situ bioreduction of technetium and uranium in a nitrate-contaminated aquifer.”
Environ. Sci. Technol., 38共2兲, 468–475.
Jeon, B. H., Kelly, S. D., Kemner, K. M., Barnett, M. O., Burgos, W. D.,
Dempsey, B. A., and Roden, E. E. 共2004兲. “Microbial reduction of
U共VI兲 at the solid-water interface.” Environ. Sci. Technol., 38共21兲,
5649–5655.
Kelly, S. D., Kemner, K. M., Fein, J. B., Fowle, D. A., Boyanov, M. I.,
Bunker, B. A., and Yee, N. 共2002兲. “X-ray absorption fine structure
determination of pH-dependent U-bacterial cell wall interactions.”
Geochim. Cosmochim. Acta, 66共22兲, 3855–3871.
Komlos, J., and Jaffe, P. R. 共2004兲. “Effect of iron bioavailability on
dissolved hydrogen concentrations during microbial iron reduction.”
Biodegradation, 15共5兲, 315–325.
Lack, J. G., Chaudhuri, S. K., Kelly, S. D., Kemner, K. M., O’Connor, S.
M., and Coates, J. D. 共2002兲. “Immobilization of radionuclides and
heavy metals through anaerobic bio-oxidation of Fe共II兲.” Appl. Environ. Microbiol., 68共6兲, 2704–2710.
Langmuir, D. 共1978兲. “Uranium solution-mineral equilibria at lowtemperatures with applications to sedimentary ore-deposits.”
Geochim. Cosmochim. Acta, 42共6兲, 547–569.
Lovley, D. R., and Phillips, E. J. P. 共1987兲. “Rapid assay for microbially
reducible ferric iron in aquatic sediments.” Appl. Environ. Microbiol.,
53共7兲, 1536–1540.
Lovley, D. R., and Phillips, E. J. P. 共1988兲. “Novel mode of microbial
energy-metabolism—Organic-carbon oxidation coupled to dissimilatory reduction of iron or manganese.” Appl. Environ. Microbiol.,
54共6兲, 1472–1480.
Lovley, D. R., and Phillips, E. J. P. 共1992兲. “Reduction of uranium by
Desulfovibrio-desulfuricans.” Appl. Environ. Microbiol., 58共3兲, 850–
856.
Michalsen, M. M., Goodman, B. A., Kelly, S. D., Kemner, K. M., McKinley, J. P., Stucki, J. W., and Istok, J. D. 共2006兲. “Uranium and
technetium bio-immobilization in intermediate-scale physical models
of an in situ bio-barrier.” Environ. Sci. Technol., 40共22兲, 7048–7053.
Moon, H. S., Komlos, J., and Jaffe, P. R. 共2007兲. “Uranium reoxidation in
previously bioreduced sediment by dissolved oxygen and nitrate.” Environ. Sci. Technol., 41共13兲, 4587–4592.
Newville, M. 共2001兲. “IFEFFIT: Interactive XAFS analysis and FEFF
fitting.” J. Synchrotron Radiat., 8, 322–324.
Newville, M., Livins, P., Yacoby, Y., Rehr, J. J., and Stern, E. A. 共1993兲.
“Near-edge X-ray-absorption fine-structure of Pb—A comparison of
theory and experiment.” Phys. Rev. B, 47共21兲, 14126–14131.
Newville, M., Ravel, B., Haskel, D., Rehr, J. J., Stern, E. A., and Yacoby,
Y. 共1995兲. “Analysis of multiple-scattering XAFS data using theoretical standards.” Physica B, 209共1–4兲, 154–156.
O’Loughlin, E. J., Kelly, S. D., Cook, R. E., Csencsits, R., and Kemner,
K. M. 共2003兲. “Reduction of uranium共VI兲 by mixed iron共II兲/iron共III兲
hydroxide 共green rust兲: Formation of UO2 nanoparticles.” Environ.
Sci. Technol., 37共4兲, 721–727.
Ortiz-Bernad, I., Anderson, R. T., Vrionis, H. A., and Lovley, D. R.
共2004兲. “Resistence of solid-phase U共VI兲 to microbial reduction during in situ bioremediation of uranium-contaminated groundwater.”
Appl. Environ. Microbiol., 70共12兲, 7558–7560.
Ravel, B., and Newville, M. 共2005兲. “ATHENA, ARTEMIS, HEPHAESTUS: Data analysis for X-ray absorption spectroscopy using IFEFFIT.” J. Synchrotron Radiat., 12, 537–541.
Sandino, A., and Bruno, J. 共1992兲. “The solubility of
共UO2兲3共PO4兲2 * 4H2O共s兲 and the formation of U共VI兲 phosphate
complexes—Their influence in uranium speciation in natural-waters.”
Geochim. Cosmochim. Acta, 56共12兲, 4135–4145.
Sani, R. K., Peyton, B. M., Dohnalkova, A., and Amonette, J. E. 共2005兲.
“Reoxidation of reduced uranium with iron共III兲 共hydr兲oxides under
sulfate-reducing conditions.” Environ. Sci. Technol., 39共7兲, 2059–
2066.
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CRITICAL REVIEW
www.rsc.org/jem | Journal of Environmental Monitoring
Aquatic environmental nanoparticles
Nicholas S. Wigginton, Kelly L. Haus and Michael F. Hochella Jr*
Received 17th August 2007, Accepted 20th September 2007
First published as an Advance Article on the web 4th October 2007
DOI: 10.1039/b712709j
Researchers are now discovering that naturally occurring environmental nanoparticles can play a
key role in important chemical characteristics and the overall quality of natural and engineered
waters. The detection of nanoparticles in virtually all water domains, including the oceans, surface
waters, groundwater, atmospheric water, and even treated drinking water, demonstrates a
distribution near ubiquity. Moreover, aquatic nanoparticles have the ability to influence
environmental and engineered water chemistry and processes in a much different way than similar
materials of larger sizes. This review covers recent advances made in identifying nanoparticles
within water from a variety of sources, and advances in understanding their very interesting
properties and reactivity that affect the chemical characteristics and behaviour of water. In the
future, this science will be important in our vital, continuing efforts in water safety, treatment,
and remediation.
1. Introduction
Environmental nanoparticles are nanometre-sized (B1–100 nm)
crystalline to amorphous solid materials formed in nature.
Scientists in the last 20 years have shown that environmental
nanoparticles are quite literally everywhere in natural environments. They exist stably in nearly all components of
the Earth, including the oceans, atmosphere, and subsurface.
The most important of these occurrences, however, is probably in the Earth’s so-called ‘‘critical zone.’’ The critical zone
of our planet extends from the topmost forest canopy down to
The Center for NanoBioEarth, Department of Geosciences, Virginia
Tech, 4044 Derring Hall, Blacksburg, VA 24061, USA. E-mail:
[email protected]; Fax: +1 540 231 3386; Tel: +1 540 231 6227
the deepest groundwater aquifer.1 It is the portion of the Earth
that provides or strongly influences nearly all of our most vital
resources, including fresh water, air, and soil. Environmental
nanoparticles, in a vast variety of forms, exist in virtually all of
these resources,2 including groundwater, lakes, and rivers.
Although these water resources comprise less than 1% of the
planet’s total water supply, they are the most indispensable
because we are critically reliant on them for drinking water
and agricultural use for a rapidly expanding population.
This article concerns nanoparticles formed by natural
geochemical (abiotic) and biogeochemical (biotic) processes
in water, as well as those formed in natural aqueous environments as an unintended consequence of human activity in
those environments. As an example of the latter, as we will see
Michael Hochella is University Distinguished Professor of Environmental Geochemistry at Virginia Tech, concentrating in the areas of
nanogeoscience and biogeochemistry. He served as President of the
Geochemical Society during 2000 and 2001, received the Alexander
von Humboldt Research Award in 2001, and was awarded the Dana
Medal by the Mineralogical Society of America in 2002. He was
elected Fellow of the American Geophysical Union in 2006, and is a
Fellow of four other professional societies. Nicholas Wigginton and
Kelly Haus are Ph.D. candidates at Virginia Tech. Their research
interests include mineral–microbe interactions and environmental
geochemistry.
Michael F. Hochella Jr, Kelly L. Haus and Nicholas S.
Wigginton (left to right)
1306 | J. Environ. Monit., 2007, 9, 1306–1316
This journal is
c
The Royal Society of Chemistry 2007
directly out of solution must start in the nanoparticle size
range. Certain phases, based on environmental conditions and
growth kinetics, quickly surpass this size region and form
much larger particles. But a large fraction of solid-phase
materials exist at this size range for extended periods of time.
In the simplest systems, many inorganic growth mechanisms
are responsible for nanoparticle formation, including classic
crystal growth,23 aggregation (i.e. ripening),24 and redoxtriggered crystallization based on changes in mineral solubility.25
Examining how nanoparticles are formed and sustained in natural
waters is key to understanding their possible roles in environmental
processes, such as the transport and ultimate fate of contaminants
associated internally or on the surface of the particles.
In many environments, certain microorganisms induce the
formation of nanoparticles. Biogenic nanoparticles are sometimes formed directly by the organism as a metabolic requirement (e.g. magnetite, Fe3O4, produced intracellularly by
magnetotactic bacteria is required for motility).26 Nanoparticles also form as an indirect result of microbial activity. For
example, when a microorganism induces the redox transformation of a metal, the solubility may significantly change
causing the precipitation of a new nanocrystalline mineral
phase (e.g. Fe-oxides27–29 and Mn-oxides14,16,30). With Mnoxidizing bacteria, for example, the final oxidation product is
Mn(IV), which is insoluble and will interact with existing
mineral phases, other aqueous metal species, or the cell wall,
to form nanoparticles. Mineralization can also be promoted
by other metabolites (e.g. electron shuttles) or by microbial
cell surfaces acting as organic templates.31
Understanding the precise growth mechanism of nanoparticles has recently become of high importance because this may
be strongly correlated to particle reactivity. For example,
Fe-oxide nanoparticles grown both abiotically and biotically
show different optical properties,32 and rates of heterogeneous
catalytic efficiency.33 Additionally, defects and phase transitions of abiotically grown hematite (a-Fe2O3) nanoparticles
depends largely on growth kinetics of the particles.34 Phase
transitions on such a scale are often directly correlated to
surface energy and thermodynamics of growth.35
Delineating the origin of nanoparticles from natural samples, however, is often a very challenging task. For example,
when determining the origin/biogenicity of magnetite, criteria
such as oxygen isotope fractionation, magnetic properties,
particle morphology, and crystal size are often too ambiguous
to be used individually.36 A much more rigorous characterization using a combination of such methods can sometimes
allow for the accurate determination of the origin.37 Indeed,
more detailed investigations of the growth mechanisms for
both inorganic and biogenic nanoparticles will undoubtedly
aid in the efforts to understand the origin of nanoparticulate
phases. Thus far, laboratory studies examining the growth
mechanisms of environmentally-relevant nanoparticles have
predominately focused on various sulfide20,38 Fe-oxide,7,29,39
and other metal oxide phases (e.g. TiO2).40
efforts. One system of high interest for countries with a history
of nuclear weapons manufacturing and nuclear power is that
of uranium contaminated soils and groundwater aquifers. For
example, a primary goal of the United States Department of
Energy is to address the nuclear legacy of the weapons
program in the US, including the remediation of uraniumcontaminated subsurface sites. One of the most promising means
of non-invasive clean-up is through the bioremediation of soluble
U(VI) by microorganisms, such as metal-reducing bacteria.41 This
involves microbial-induced redox-transformations from U(VI) to
an insoluble U(IV) phase. Thus, the hope is for the uranium to be
immobilized within the contaminated aquifer. However, one
caveat to this argument is that these precipitates have been shown
to predominately exist as nanoparticles.17,42,43 In fact, nanoparticles of uraninite (UO2) have been shown to form abiotically
when U(VI) is reduced by Fe(II)-oxides.19,44
Disregarding reoxidation of the solid-phase products45 for
the moment, one issue of great concern for the stability of the
product is how unreduced aqueous U(VI) interacts with the
precipitating nanoparticles. One field study showed that a
solid-phase U(IV) precipitate actually contained a large fraction of unreduced U(VI).46
The complexities behind the fate of metal and radionuclide
contaminants during nanoparticle formation make predicting
the final products very difficult. Several examples in the
literature of different end-products for various metal/metalloid
contaminants highlight the fact that we do not yet fully
understand the fate of the contaminant with respect to the
precipitating nanoparticle, but a lot of progress is being made.
A recent study by Zachara and colleagues showed that the
abiotic reduction of Tc(VII) by Fe(II) formed relatively stable
iron oxide nanoparticles with homogeneously distributed
Tc(IV) in the crystal structure, effectively stabilizing the Tc.25
The opposite can be true, however, with the metalloid As.
Tadanier and colleagues showed that the microbial-reduction
of Fe-oxide aggregates with adsorbed As(V) caused the deflocculation of As-bearing ferrihydrite (Fe10O14(OH)2) nanoparticles, subsequently increasing the mobility of As.47
When relying on the reductive transformation of metals and
subsequent growth of nanoparticles as a means for remediation, the problem of nanoparticle stability must be a chief
concern. However, as we will see later, nanoparticles can be
incredibly mobile in the environment so understanding the
transport mechanisms of nanoparticles is also very important
for determining the practicality of remediation efforts. One
recent study, however, circumvented the need for understanding transport because they simply removed metal- and nanoparticle-bearing water from a contaminated mining site.48
Magnetite nanoparticles were grown using ferric iron in the
extracted acid-mine drainage waters which also contained
trace amounts of other metals including Zn, Ni, and Cu. This
presents an intriguing remediation case where contamination
species are removed from the site and transformed ex situ.
Nanoparticle formation associated with metal contaminants
3. Reactivity of nanoparticles: insight from the
laboratory
In systems with heavy metal and/or radionuclide contamination, nanoparticles are often the byproducts of remediation
As alluded to in the Introduction, the principal reason for the
recent trend in the characterization of nanoparticles is that
1308 | J. Environ. Monit., 2007, 9, 1306–1316
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16 M. Villalobos, J. Bargar and G. Sposito, Environ. Sci. Technol.,
2005, 39, 569–576.
17 Y. Suzuki, S. D. Kelly, K. M. Kemner and J. F. Banfield, Nature,
2002, 419, 134–134.
18 (a) Y. Suzuki, S. D. Kelly, K. M. Kemner and J. F. Banfield,
Appl. Environ. Microbiol., 2005, 71, 1790–1797; (b) J. K.
Fredrickson, J. M. Zachara, D. W. Kennedy, M. C. Duff, Y. A.
Gorby, S. M. W. Li and K. M. Krupka, Geochim. Cosmochim.
Acta, 2000, 64, 3085–3098.
19 E. J. O’Loughlin, S. D. Kelly, R. E. Cook, R. Csencsits and K. M.
Kemner, Environ. Sci. Technol., 2003, 37, 721–727.
20 M. Labrenz, G. K. Druschel, T. Thomsen-Ebert, B. Gilbert,
S. A. Welch, K. M. Kemner, G. A. Logan, R. E. Summons,
G. De Stasio, P. L. Bond, B. Lai, S. D. Kelly and J. F. Banfield,
Science, 2000, 290, 1744–1747.
21 T. F. Rozan, M. E. Lassman, D. P. Ridge and G. W. Luther,
Nature, 2000, 406, 879–882.
22 (a) J. H. P. Watson, I. W. Croudace, P. E. Warwick,
P. A. B. James, J. M. Charnock and D. C. Ellwood, Sep. Sci.
Technol., 2001, 36, 2571–2607; (b) J. H. P. Watson, B. A. Cressey,
A. P. Roberts, D. C. Ellwood, J. M. Charnock and A. K.
Soper, J. Magn. Magn. Mater., 2000, 214, 13–30; (c) J. W.
Moreau, R. I. Webb and J. F. Banfield, Am. Mineral., 2004, 89,
950–960.
23 G. A. Waychunas, in Nanoparticles and the Environment,
Mineralogical Society of America, Washington DC, 2001,
vol. 44, pp. 105–166.
24 A. Navrotsky, Proc. Natl. Acad. Sci. U. S. A., 2004, 101,
12096–12101.
25 J. M. Zachara, S. M. Heald, B. Jeon, R. K. Kukkadapu, C. Liu, J.
P. McKinley, A. C. Dohnalkova and D. A. Moore, Geochim.
Cosmochim. Acta, 2007, 71, 2137–2157.
26 D. A. Bazylinski and R. B. Frankel, Nat. Rev. Microbiol., 2004, 2,
217–230.
27 S. Glasauer, S. Langley and T. J. Beveridge, Science, 2002, 295,
117–119.
28 (a) C. M. Hansel, S. G. Benner, P. Nico and S. Fendorf, Geochim.
Cosmochim. Acta, 2004, 68, 3217–3229; (b) J. R. Lloyd, V. A.
Sole, C. V. G. Van Praagh and D. R. Lovley, Appl. Environ.
Microbiol., 2000, 66, 3743–3749; (c) C. M. Hansel, S. G. Benner,
J. Neiss, A. Dohnalkova, R. K. Kukkadapu and S. Fendorf,
Geochim. Cosmochim. Acta, 2003, 67, 2977–2992; (d) R. K.
Kukkadapu, J. M. Zachara, J. K. Fredrickson, D. W. Kennedy,
A. C. Dohnalkova and D. E. McCready, Am. Mineral., 2005, 90,
510–515.
29 J. F. Banfield, S. A. Welch, H. Z. Zhang, T. T. Ebert and R. L.
Penn, Science, 2000, 289, 751–754.
30 (a) B. M. Tebo, J. R. Bargar, B. G. Clement, G. J. Dick, K. J.
Murray, D. Parker, R. Verity and S. M. Webb, Annu. Rev. Earth
Planet. Sci., 2004, 32, 287–328; (b) M. Villalobos, B. Lanson, A.
Manceau, B. Toner and G. Sposito, Am. Mineral., 2006, 91,
489–502.
31 (a) C. S. Chan, G. De Stasio, S. A. Welch, M. Girasole, B. H.
Frazer, M. V. Nesterova, S. Fakra and J. F. Banfield, Science,
2004, 303, 1656–1658; (b) M. Nesterova, J. Moreau and J. F.
Banfield, Geochim. Cosmochim. Acta, 2003, 67, 1177–1187.
32 R. S. Oremland, M. J. Herbel, J. S. Blum, S. Langley, T. J.
Beveridge, P. M. Ajayan, T. Sutto, A. V. Ellis and S. Curran,
Appl. Environ. Microbiol., 2004, 70, 52–60.
33 H. Jung, H. Park, J. Kim, J. Lee, H. Hur, N. V. Myung and H.
Choi, Environ. Sci. Technol., 2007, 41, 4741–4747.
34 I. V. Chernyshova, M. F. Hochella and A. S. Madden, Phys.
Chem. Chem. Phys., 2007, 9, 1736–1750.
35 A. Navrotsky, Geochem. Trans., 2003, 4, 34–37.
36 (a) D. Faivre, P. Agrinier, N. Menguy, P. Zuddas, K. Pachana, A.
Gloter, J. Y. Laval and F. Guyot, Geochim. Cosmochim. Acta,
2004, 68, 4395–4403; (b) D. Faivre, N. Menguy, F. Guyot, O.
Lopez and P. Zuddas, Am. Mineral., 2005, 90, 1793–1800.
37 D. Faivre and P. Zuddas, Earth Planet. Sci. Lett., 2006, 243,
53–60.
38 I. Ciglenecki, D. Krznaric and G. R. Helz, Environ. Sci. Technol.,
2005, 39, 7492–7498.
39 (a) R. L. Penn, J. J. Erbs and D. M. Gulliver, J. Cryst. Growth,
2006, 293, 1–4; (b) Y. Guyodo, A. Mostrom, R. L. Penn and S. K.
Banerjee, Geophys. Res. Lett., 2003, 30.
This journal is
c
The Royal Society of Chemistry 2007
40 (a) R. L. Penn, J. Phys. Chem. B, 2004, 108, 12707–12712; (b) M.
R. Ranade, A. Navrotsky, H. Z. Zhang, J. F. Banfield, S. H.
Elder, A. Zaban, P. H. Borse, S. K. Kulkarni, G. S. Doran and H.
J. Whitfield, Proc. Natl. Acad. Sci. U. S. A., 2002, 99, 6476–6481;
(c) H. Z. Zhang and J. F. Banfield, J. Phys. Chem. B, 2000, 104,
3481–3487; (d) R. L. Penn and J. F. Banfield, Science, 1998, 281,
969–971.
41 (a) J. D. Wall and L. R. Krumholz, Annu. Rev. Microbiol., 2006,
60, 149–166; (b) K. T. Finneran, R. T. Anderson, K. P. Nevin and
D. R. Lovley, Soil Sediment Contam., 2002, 11, 339–357.
42 (a) J. K. Fredrickson, J. M. Zachara, D. W. Kennedy, C. X. Liu,
M. C. Duff, D. B. Hunter and A. Dohnalkova, Geochim.
Cosmochim. Acta, 2002, 66, 3247–3262; (b) M. J. Marshall, A.
S. Beliaev, A. C. Dohnalkova, D. W. Kennedy, L. Shi,
Z. M. Wang, M. I. Boyanov, B. Lai, K. M. Kemner, J. S.
McLean, S. B. Reed, D. E. Culley, V. L. Bailey, C. J. Simonson,
D. A. Saffarini, M. F. Romine, J. M. Zachara and J. K.
Fredrickson, PLoS Biol., 2006, 4, 1324–1333.
43 M. Fayek, S. Utsunomiya, S. M. Pfiffner, D. C. White, L. R.
Riciputi, R. C. Ewing, L. M. Anovitz and F. J. Stadermann, Can.
Mineral., 2005, 43, 1631–1641.
44 M. I. Boyanov, E. J. O’Loughlin, E. E. Roden, J. B. Fein and
K. M. Kemner, Geochim. Cosmochim. Acta, 2007, 71,
1898–1912.
45 J. M. Wan, T. K. Tokunaga, E. Brodie, Z. M. Wang, Z. P. Zheng,
D. Herman, T. C. Hazen, M. K. Firestone and S. R. Sutton,
Environ. Sci. Technol., 2005, 39, 6162–6169.
46 I. Ortiz-Bernad, R. T. Anderson, H. A. Vrionis and D. R. Lovley,
Appl. Environ. Microbiol., 2004, 70, 7558–7560.
47 C. J. Tadanier, M. E. Schreiber and J. W. Roller, Environ. Sci.
Technol., 2005, 39, 3061–3068.
48 X. C. Wei and R. C. Viadero, Colloids Surf., A, 2007, 294,
280–286.
49 F. M. Michel, L. Ehm, S. M. Antao, P. L. Lee, P. J. Chupas, G.
Liu, D. R. Strongin, M. A. A. Schoonen, B. L. Phillips and J. B.
Parise, Science, 2007, 316, 1726–1729.
50 A. S. Madden, M. F. Hochella and T. P. Luxton, Geochim.
Cosmochim. Acta, 2006, 70, 4095–4104.
51 J. R. Rustad and A. R. Felmy, Geochim. Cosmochim. Acta, 2005,
69, 1405–1411.
52 A. J. Anschutz and R. L. Penn, Geochem. Trans., 2005, 6,
60–66.
53 R. L. Poulson, C. M. Johnson and B. L. Beard, Am. Mineral.,
2005, 90, 758–763.
54 O. Larsen and D. Postma, Geochim. Cosmochim. Acta, 2001, 65,
1367–1379.
55 H. D. Pederson, D. Postma, R. Jakobsen and O. Larsen,
Geochim. Cosmochim. Acta, 2005, 69, 3967–3977.
56 S. Bonneville, P. Van Cappellen and T. Behrends, Chem. Geol.,
2004, 212, 255–268.
57 S. Bonneville, T. Behrends, P. Van Cappellen, C. Hyacinthe and
W. F. M. Roling, Geochim. Cosmochim. Acta, 2006, 70,
5842–5854.
58 S. Bose, M. F. Hochella, B. H. Lower, Y. A. Gorby, D. W.
Kennedy and D. E. McCready, Abstr. Paper Am. Chem. Soc.
Natl. Meet., 2006, 232, GEOC 119.
59 S. Glasauer, P. G. Weidler, S. Langley and T. J. Beveridge,
Geochim. Cosmochim. Acta, 2003, 67, 1277–1288.
60 J. T. Mayo, C. Yavuz, S. Yean, L. Cong, H. Shipley, W. Yu, J.
Falkner, A. Kan, M. Tomson and V. L. Colvin, Sci. Technol. Adv.
Mater., 2007, 8, 71–75.
61 (a) K. L. Chen, S. E. Mylon and M. Elimelech, Environ. Sci.
Technol., 2006, 40, 1516–1523; (b) K. L. Chen, S. E. Mylon and
M. Elimelech, Langmuir, 2007, 23, 5920–5928.
62 (a) D. J. Burleson, M. D. Driessen and R. L. Penn, J. Environ. Sci.
Health, Part A, 2004, 39, 2707–2753; (b) C. Contado, G. Blo, C.
Conato, F. Dondi and R. Beckett, J. Environ. Monit., 2003, 5,
845–851; (c) G. G. Leppard, D. Mavrocordatos and D. Perret,
Water Sci. Technol., 2004, 50, 1–8.
63 J. R. Lead and K. J. Wilkinson, in Environmental Colloids and
Particles: Behaviour, Separation and Characterisation, ed. J. R.
Lead and K. J. Wilkinson, John Wiley and Sons, Chichester, UK,
2007, pp. 1–687.
64 J. R. Lead, J. Hamilton-Taylor, W. Davison and M. Harper,
Geochim. Cosmochim. Acta, 1999, 63, 1661–1670.
J. Environ. Monit., 2007, 9, 1306–1316 | 1315
Geobiology (2007), 5, 207–210
DOI: 10.1111/j.1472-4669.2007.00122.x
Introduction to Special Issue
IPerspectives
N T RO DPublishing
U Cfrom
T I Othe
N Ltd
Tmineral–bacteria
O S P E C I A L I Sinterface
SUE
Blackwell
The evolution of geomicrobiology: perspectives from the
mineral–bacteria interface
D. A. F OW LE , 1 J . A . RO B E RT S , 1 D . F O RTIN 2 AN D K. KONHAUSER 3
1
Department of Geology, University of Kansas, Multidisciplinary Research Building, 2030 Becker Dr, Lawrence, KS 66047, USA
of Earth Sciences, University of Ottawa, 140 Louis Pasteur, Ottawa, Ontario, Canada K1N 6N5
3
Department of Earth and Atmospheric Sciences, University of Alberta, Edmonton, Alberta, Canada T6G 2E3
2Department
This issue of Geobiology provides a glimpse into the state
of geomicrobiology with research presented spanning from
molecular-scale cellular metal interactions to field studies of
elemental cycling. The broad link between all of these papers
presented here is the interconnectivity between minerals and
microbial ecology and metabolisms. This issue was organized
and solicited from the session ‘Bacteria–mineral interface’ at
the International Mineralogical Association meeting in Kobe,
Japan, 2006.
From its origins, perhaps some 4 billion years ago, biology
has had a profound effect on shaping our planet. The ‘higher’
organisms, multicellular eukaryotes, are restricted for the most
part to the Earth’s surface, while the ubiquitous nature of
prokaryotic organisms has allowed them to extend from polar
icecaps to the hottest desert, from the most acid acidic mine
waste to salty and highly alkaline lakes, and from atmospheric
dust particles to oceanic trenches, hydrothermal ocean vents
and a myriad of subterranean environments. Indeed, it would
be necessary to penetrate several kilometres into the crust
where temperatures are outside the physiochemical limits for
life to find a sterile environment. Not only are prokaryotes
widespread in the Earth’s crust, but throughout the biosphere,
microbial populations are intimately involved in transforming
both inorganic and organic compounds to meet their metabolic and energetic requirements, and in doing so, they have
modified almost every aspect of the Earth’s biosphere (see
Konhauser, 2007).
Geomicrobiological research has been conducted under
various guises (e.g. microbial ecology, low-temperature
geochemistry, environmental engineering, economic geology,
chemical oceanography) for many years but perhaps the true
blossoming of the science followed the MSA short course and
volume by Banfield & Nealson in (1997). Since this time, the
discipline has grown into a multidisciplinary science that
Corresponding author: David A. Fowle, Tel.: +1 785 864 1955;
fax: +1 785 864 5276; e-mail: [email protected].
© 2007 The Authors
Journal compilation © 2007 Blackwell Publishing Ltd
links microbiologists, genome scientists, geochemists, physicists,
biochemists, and analytical chemists together, and has essentially generated a subfield of molecular geomicrobiology which
has garnered significant attention (e.g. Banfield et al., 2005).
Of course, much of this focus has led to the development of
large environmental or single organism genomic databases,
which are still substantially separated from pairing the genetic
basis to their biogeochemically relevant metabolic pathways.
Molecular geomicrobiology, in this case, also refers to the
increased use and utility of spectroscopic techniques to study
nanoscale processes at the bacteria–mineral interface (e.g.
Jiang et al., 2004). Opportunities abound for continuing on
with molecular geochemistry via further deconstruction using
model organisms (e.g. DiChristina et al., 2005) and the use of
artificial membranes or colloids in spectroscopic studies (e.g.
Boyanov et al., 2007). However, there still remains a need for
system-based science in geomicrobiology. The relevance of
gene expression, and what are considered biogeochemical
relevant genes, will remain unknown, unless they are evaluated
in the context of interdependent communities, chemical
gradients, and typical transport mechanisms in near-surface
geological settings. Here we present a number of papers that,
in their own way, attempt to adjust to the biocomplexity of
these natural settings by: using the microbes’ ecological
settings as models for laboratory-based studies; molecularly
characterizing the surface chemistry of native mineral–
bacteria composites or actively metabolizing bacteria; and
by studying mesoscale and end-member geomicrobiological
settings to isolate microbial influences on, and function within,
near-surface geological settings.
CELL SURFACE REACTIVITY
One of the uniquely characteristic features of microorganisms
is their large surface area to volume ratios. This, coupled
with highly reactive charged surfaces, leads to significant metal
partitioning onto microbial biomass. Many of the metals
bound serve physiological functions, but interestingly, many
207
Perspectives from the mineral–bacteria interface
Rosling et al. also investigated the microbial acquisition of
nutrients from mineral sources, examining the ability of fungi
to release the macronutrient, phosphorus (P), from apatite as
a function of P concentration in solution. Using fungal isolates
from a grassland in northern California, USA, the authors
identified three fungi-mediated modes of dissolution including acidification, moderate acidification, and no acidification.
Acidifying isolates, identified as Zygomycetes in the order of
Mucorales, induce fluorapatite dissolution by producing oxalic
acid while growing in the presence of P. In contrast, the nonacidifying isolate, identified as Ascomycetes belonging to the
family Trichocomaceae, lowered the solution pH and induced
fluroapatite dissolution without the production of low molecular
weight organic acids under P-limited conditions. Results from
this study stress the significance of soil mineralogy as a source
of essential nutrients, such as phosphorus, which is limiting in
many soil environments.
ELEMENTAL CYCLING
All of Earth’s major biogeochemical cycles are also effected
by microbial metabolism. Some cells couple the oxidation
of organic material with the dissolution of mineral phases
(e.g. ferric iron reduction) or dissolved solutes (e.g. sulfate
reduction), whereas other cells have evolved the metabolic
capacity to oxidize inorganic substances for autotrophic
carbon fixation. In either case, microorganisms simply catalyse
reactions that are thermodynamically favoured, yet kinetically
hindered. In sediments, the reductive–oxidative processes
work in tandem, with the by-products of one metabolic guild
the substrate for another. This invariably leads to biochemical
stratification, and although the fundamentals underlying
biogeochemical zonation are established, it is now becoming
apparent that complex recycling in microniches may impart
significant heterogeneity on the overall system. Recent advances
in sampling both pore-water geochemistry and microbiological
populations, using for example, signature lipid biomarkers or
nucleic acid sequence analysis of genes, are providing new
insights into characterizing microbial community structure
and nutritional status.
Here in this issue, two papers focused on the environmental
ramifications of metal reduction. Rowland et al. investigated
the control exerted by organic matter on microbially mediated
Fe(III) reduction and arsenic(III) release in sediments from a
shallow alluvial aquifer in Cambodia. Using natural sediments
and various types of organic carbon, the authors showed
that the rate and magnitude of Fe(III) reduction and As(III)
release under anaerobic conditions were enhanced after
the addition of acetate and AQDS (used in the study as an
analogue for humic substances) when compared to autoclavedcontrol systems. The presence of AQDS, hydrocarbons and
finer grained sediments enhanced As release associated with
the amorphous and crystalline Fe and Al fractions of the
sediments. The native microbial community was initially
© 2007 The Authors
Journal compilation © 2007 Blackwell Publishing Ltd
209
complex and involved in various metabolic processes, but the
addition of acetate and AQDS to the microcosms led to a
predominance of microorganisms closely related to metalreducing Geobacter species. The role of Geobacter in the mobilization of As remains unclear, but the authors proposed
that dissolution of ferric oxides was likely responsible for the
release of As. This study highlights the role of heterogeneity
in sediment geochemistry and carbon source in differentiating
microbial communities resulting in unique geomicrobiological
conditions.
Wilkens et al. provide an intriguing view of the grand
challenges associated with scaling the study of Fe, U, and Tc
reduction in the laboratory to natural systems. The authors
discovered that Fe(III)-reducing microorganisms in the Drigg
site sediments (currently operated by British Nuclear Fuels)
effectively removed both U(VI) and Tc(VII) from the aqueous
phase while continuing to immobilize radium in the sediments.
With the onset of denitrifying conditions, an organism closely
related to Pseudomonas stutzeri, dominated the bacterial community structure and leads to nitrite production. The reoxidation and the introduction of nitrate to the system facilitated
the remobilization of U(VI), whereas Tc remained in an insoluble
form. Ultimately, this work stresses the need for site-specific
information on a variety of scales in order to accurately predict
radionuclide mobility and potential bioremediation outcomes
for the long-term stability of contaminated sites.
All of the studies included in this issue emphasize the broad
implications of geomicrobiology in modern environments,
but they have relevance for the geological past as well. Indeed,
the events that led to the emergence of life and evolution of
the biosphere can only really be elucidated by studying modern ecosystems along with the fossil and stratigraphic records.
Complimentary modern ecosystem studies provide a means
to understand the biogeochemical processes by which life
interacted with its environment through time, and ultimately
how biosignatures are recorded in the geological record.
Furthermore, these geomicrobiological investigations at the
complex laboratory scale and the mesoscale in the field are
perhaps our best opportunities for insight into novel microbial
metabolic pathways and ecosystem function of particular
groups of organisms, thereby providing crucial linkages between
geochemical and microbiological evolution in subsurface
environments.
REFERENCES
Banfield JF, Nealson KH (eds) (1997) Geomicrobiology: interactions
between microbes and minerals. Reviews in Mineralogy and
Geochemistry 35, 448.
Banfield JF, Tyson G, Allen EE, Whitaker RJ (2005) The search for
a molecular-level understanding of the processes that underpin the
Earth’s biogeochemical cycles. Reviews in Mineralogy and
Geochemistry 59, 1–7.
Boyanov MI, O’Loughlin EJ, Roden EE, Fein JB,
Kemner KM (2007) Adsorption of Fe(II) and U(VI)
IOP PUBLISHING
JOURNAL OF GEOPHYSICS AND ENGINEERING
doi:10.1088/1742-2132/4/3/S07
J. Geophys. Eng. 4 (2007) 285–292
Abundances of radioelements (K, U, Th)
in weathered igneous rocks in Hong Kong
L S Chan1, P W Wong1 and Q F Chen2
1
2
Department of Earth Sciences, University of Hong Kong, Hong Kong SAR
Electronics and Geophysical Surveys, Hong Kong SAR
E-mail: [email protected]
Received 2 November 2006
Accepted for publication 5 June 2007
Published 31 August 2007
Online at stacks.iop.org/JGE/4/285
Abstract
Gamma-ray spectrometric measurements and geochemical determinations of major and trace
element contents using ICP-MS and XRF were conducted on 58 igneous rock samples from
Hong Kong. Stripping analyses on the gamma-ray spectra have yielded estimates on the
abundance of K, Th and U. Major element contents were used to compute the Parker
weathering indices for the samples. Only K shows a systematic variation in concentration with
an increasing degree of weathering. The increase in porosity and interconnectivity of
microfractures are probably the cause for the observed nonlinear decline in the K content
during the weathering process. The concentrations of U and Th in the samples do not show
any systematic variations with the weathering index, reflecting the complex mechanisms of
dissolution and deposition of the two radioelements in the weathering profile. A curvilinear
relationship between eK and Wp has been derived from the measurement data, which
possibly provides a quick means of characterizing the extent of alteration in the igneous
rocks.
Keywords: gamma-ray spectrometry, radiometry, radioelements, saprolites, weathered igneous
rock, Hong Kong
Introduction
Gamma-ray spectrometric surveys are often undertaken to
facilitate geological mapping, mineral exploration as well
as assessment of certain environment hazards such as radon.
The method uses a gamma-ray spectrometer to sort detected
gamma rays associated with different radioelements in the
surficial layer by their respective energies. The three most
abundant radioelements in rocks—K, U and Th—follow
very different pathways of evolution during the alteration
of rocks (Pliler and Adams 1962, Dickson et al 1995,
Scott and Dickson 1999, Imaizumi and Ishida 2001, Perrin
et al 2006). Chemical decomposition of K-bearing minerals
such as feldspars and micas results in leaching of potassium
cations with consequential formation of clay minerals depleted
in K. This process of chemical decomposition is generally
accompanied by a reduction of physical strength and
compositional changes. The gradual decline in the potassium
content with an increasing degree of weathering has been
1742-2132/07/030285+08$30.00
described in numerous studies (e.g. Curtis (1976), Wilford
et al (1997), Dickson and Scott (1997), Taboada et al (2006)).
The concentration of U and Th in weathered rock depends
complexly on the processes of dissolution and precipitation
of the two radioelements in the weathered rock. The removal
of U from the host rock is mainly through the dissolution of
uranyl compounds in the presence of groundwater, and to a
lesser extent, by direct alpha recoil of the U-progeny isotopes
into groundwater in the interstitial pores. Many studies have
shown that the mobility of uranyl ions does not depend on
a single soil parameter but is very sensitive to the redox
potential, the alkalinity of the groundwater, and the presence
of complexing agents such as carbonates and phosphates in
the groundwater (e.g. Rosholt et al (1965), Scott et al (1992),
Porcelli and Swarzenski (2003), Curtis et al (2004), Batuk et al
(2006), Vandenhove et al (2007), Boyanov et al (2007)). A
thorough understanding of the distribution and migration of the
radioelements in the weathering process is critically important
for the interpretation of gamma-ray spectrometry data.
© 2007 Nanjing Institute of Geophysical Prospecting
Printed in the UK
285
Abundances of radioelements (K, U, Th) in weathered igneous rocks in Hong Kong
in particular at low ionic strengths (Catalano and Brown 2005).
In a pedogenic profile, U and Th are often leached from the
topsoil and precipitated in the bottom soil horizons (Greenman
et al 1999, Taboada et al 2006). This transfer within the soil
profile depends on the downward percolation of groundwater
of U and Th ions. The presence of vegetation, including
moss and lichens, can further complicate the re-distribution
process of the U and Th compounds in the weathered profile
(Szalay 1964). The adsorption of the U and Th within the
weathered regolith is found to be facilitated by the presence of
certain forms of bacteria. Many works have shown that sulfate
reducing bacteria in soil can lead to reduction of soluble U(VI)
into insoluble uraninite or U(IV) oxides (Abdelouas et al 2000,
Ohnuki et al 2005, Roden and Scheibe 2005), highlighting the
role of microorganisms in the precipitation of U within the
weathered regolith.
The above-mentioned factors and processes have probably
led to the observed complex patterns of variation of U and Th
contents in the weathering rocks in the present study. The
typical groundwater conditions within the weathered regolith
in Hong Kong, however, are not conducive to the U dissolution
process. The groundwater in Hong Kong has an average
pH of about 6.1 and contains hydrochemical complexes of
mostly nitrates, sulphates and chlorides (Leung et al 2005).
The acidity, the relative depletion in carbonate and phosphate
complexes, and generally the abundance of iron oxyhydroxides
probably limit the U mobilization into the groundwater. The
amount of adsorbed or precipitated U and Th, in such cases,
becomes highly localized and controlled by the abundance of
iron oxyhydroxides, the oxygen content of the groundwater,
groundwater percolation rate, and the presence of organic
matter and bacteria within the intergranular pores and fractures
in the altered igneous rocks.
Conclusion
Among the three most common radioelements, the contents
of K have demonstrated a statistically significant correlation
with the weathering index. Both Th and U did not show
any marked trend of variation with an increasing degree
of weathering, corroborating the results from earlier studies
on the complexities of the dissolution and precipitation
mechanisms of the two radioelements in the weathering
process. Although the variations in U and Th concentrations
do not show any distinct pattern with the increasing extent
of weathering, the gamma-ray spectrometric method can still
be used as an indication of the degree of weathering in the
igneous rocks. The gamma-ray intensity integrated over the K
energy window in the gamma-ray spectrum of the weathered
samples is markedly greater than that of U and Th. The
significant correlation between the weathering index and eK
derived from this study allows us to use the eK content as a
proxy index to weathering. Since gamma-ray spectrometry
surveying is fast and inexpensive, the method could become
an efficient tool for characterizing the extent of weathering, as
well as for studying the mechanisms of leaching, sorption and
transporting of radioelements in the weathering process.
Acknowledgments
This study has been funded by a Hong Kong CERG grant
HKU7094/00P. Supplementary funding was also provided by
the University of Hong Kong through two incentive awards
HKU7007/04 and HKU7009/05. The authors are indebted
to two anonymous reviewers for thorough reviews of the
manuscript.
References
Abdelouas A, Lutze W, Gong W, Nuttall E H, Strietelmeier B A and
Travis B J 2000 Biological reduction of uranium in
groundwater and subsurface soil Sci. Total Environ. 250 21–35
Batuk O N, Kalmykov S N, Zakharova E V, Teterin Y A and
Myasoedov B F 2006 Uranium dioxide interaction with
groundwater—leaching behavior and sorption tests J. Nucl.
Mater. 352 241–5
Boyanov M I, O’Louighlin E J, Roden E E, Fein J B and
Kemner K M 2007 Adsorption of Fe(II) and U(VI) to
carboxyl-functionalized microspheres: the influence of
speciation on uranyl reduction studied by titration and XAFS
Geochim. Cosmochim. Acta 71 1898–912
Catalano J G and Brown G E 2005 Uranyl adsorption onto
montmorillonite: evaluation of binding sites and carbonate
complexation Geochim. Cosmochim. Acta 69 2995–3005
Chen Q F 2002 Geophysical and radiometric investigation of
weathered igneous rocks in Hong Kong PhD Thesis
University of Hong Kong
Chen Q and Chan L S 2000 Gamma radioactivity differentiation of
weathered volcanic tuff in Hong Kong Proc. Symp. on the
Application of Geophysics to Engineering and Environmental
Problems (Washington DC, USA) pp 1137–46
Chen M Q F and Chan L S 2002 In-situ gamma-ray spectrometric
study of weathered volcanic rocks in Hong Kong Earth Surf.
Process. Landf. 27 613–25
Chen Q, Chan L S and Ie S 1999 Radiometric study of clay-rich
zones in weathered volcanic tuff in Hong Kong Proc. 2nd
Asian Symposium on Engineering Geology and the
Environment (Malaysia) pp 7-13 to 7-19
Chittleborough D J 1991 Indices of weathering for soils and
paleosols formed on silicate rocks Aust. J. Earth Sci. 38
115–120
Curtis C D 1976 Chemistry of rock weathering: fundamental
reactions and controls Geomorphology and Climate
ed E Derbyshire (London: Wiley) pp 25–57
Curtis G P, Fox P, Kohler M and Davis J A 2004 Comparison of in
situ uranium KD values with a laboratory determined surface
complexation model Appl. Geochem. 19 1643–53
Dickson B L, Fraser S J and Kinsey H A 1995 Interpreting aerial
gamma-ray surveys utilizing geomorphological and weathering
models J. Geochem. Explor. 57 75–88
Dickson B L, Fraser S J and Kinsey-Henderson A 1995 Interpreting
aerial gamma-ray surveys utilising geomorphological and
weathering models J. Geochem. Explor. 57 75–88
Dickson B L and Scott K M 1997 Interpretation of aerial gamma-ray
surveys—adding the geochemical factors AGSO J. Aust. Geol.
Geophys. 17 187–200
Durrance E M 1986 Radioactivity in Geology: Principles and
Applications (Chichester: Ellis Horwood)
Duzgoren-Aydin N S, Aydin A and Malpas J 2002 Re-assessment of
chemical weathering indices: case study of pyroclastic rocks of
Hong Kong Eng. Geol. 63 99–119
Echevarria G, Sheppard M I and Morel J L 2001 Effect of pH on the
sorption of uranium in soils J. Environ. Radioact. 53 257–64
Geotechnical Engineering Office 1988 Guide to Rock and Soil
Descriptions. Geoguide 3: Geotechnical Engineering Office,
Hong Kong Government
291