Environ. Sci. Technol. 1998, 32, 2084-2091
Potential Role of Bicarbonate during
Pyrite Oxidation
V . P . E V A N G E L O U , * ,† A . K . S E T A , ‡ A N D
A. HOLT†
Department of Agronomy, University of Kentucky,
Lexington, Kentucky 40546-0091, and Fakultas Pertanian,
Universitas Bengkulu, Bengkulu 38371-A, Indonesia
According to Frontier molecular orbital (FMO) theory, the
surface-exposed sulfur atom of pyrite possesses an
unshared electron pair which produces a slightly negatively
charged pyrite surface that can attract cations such as
Fe2+. Because of surface electroneutrality and pH considerations, however, the pyrite surface Fe2+ coordinates
OH. We proposed that this surface Fe2+ OH when in the
presence of CO2 is converted to -FeCO3 or -FeHCO3,
depending on pH. In this study, using Fourier transform
infrared spectroscopy (FT-IR) we demonstrated that such
complexes form on the surface of pyrite and continue to
persist even after a significant fraction of the surface
Fe2+ was oxidized to Fe3+. FT-IR spectra also showed the
presence of two carbonyl absorption bands (1682 and
1653 cm-1) on the surface of pyrite upon exposure to CO2
which suggested that pyrite surface carbonate complexes
existed in two different surface chemical environments,
pointing out two potential mechanisms of pyrite surfaceCO2 interactions. One potential mechanism involved formation
of a pyrite surface-Fe(II)HCO3 complex, whereas a second
potential mechanism involved formation of a pyrite
surface-carboxylic acid group complex [-Fe(II)SSCOOFe(II)]. We hypothesized that these pyrite surface-CO2
complexes could promote abiotic oxidation of pyrite by
accelerating the abiotic oxidation of Fe2+. Iron (III) would
oxidize the disulfide (-S2) by accepting its electrons.
Using a miscible displacement technique, oxidation of FeS2
with H2O2 was carried out in the absence or presence
of 10 or 100 mM NaHCO3. The data show that 100 mM
NaHCO3 significantly increased the oxidation rate of FeS2.
Furthermore, the data show that FeS2 oxidation kinetics were
more dependent on HCO3- but were less dependent on
H2O2 for the range of HCO3- and H2O2 concentrations tested.
Introduction
The need to prevent the development of acid mine drainage
(AMD) by oxidation of pyrite has triggered numerous
investigations into the mechanisms of its oxidation (1-9).
Singer and Stumm (1) reported that Fe3+ can oxidize pyrite
at a much higher rate than O2. The role of Fe3+ in the
oxidation of pyrite is demonstrated below
* Corresponding author e-mail: [email protected].
† University of Kentucky.
‡ Universitas Bengkulu.
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ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 32, NO. 14, 1998
FeS2 + 14 Fe3+ + 8H2O
15Fe2++ 2SO42– + 16H+
(1)
O2
iron-oxidizing bacteria
At low pH, Thiobacillus ferrooxidans (an acidophilic,
chemoautotrophic, iron-oxidizing bacterium) catalyzes and
accelerates the oxidation of Fe2+ by a factor larger than 106
(1). Thus, the processes of reduction of Fe3+ to Fe2+ by pyrite
and oxidation of Fe2+ to Fe3+ by atmospheric O2, catalyzed
by T. ferrooxidans, represent an effective continuous pyrite
oxidation cycle. It is believed that pyrite oxidation can be
controlled by adding alkaline material to pyrite, a widely
used field practice. This is due to the fact that the activity
of T. ferrooxidans diminishes and Fe3+ becomes insoluble at
high pH (10). However, recent findings (4) revealed that
nonmicrobial pyrite oxidation rates increased as pH increased. At circumneutral pH, nonmicrobial pyrite oxidation
appeared to be carried out by a surface-catalyzed mechanism
involving Fe2+, O2, and Fe3+ (6, 11).
According to Moses and Herman (4) and Moses et al. (5),
Fe3+ is an effective and direct pyrite oxidant at low pH as well
as at circumneutral pH, and the role played by dissolved O2
is to sustain the reaction by regenerating Fe3+. It is wellknown that the abiotic rate of Fe2+ oxidation increases as pH
increases (1, 12). The role of OH in abiotically oxidizing Fe2+
has been postulated by Luther et al. (9) to be due to the
potential increase in frontier molecular orbital electron
density of Fe2+ upon binding to oxygen by coordinating OH.
Luther et al. (9) demonstrated that an increase in electron
density also increases the potential of Fe2+ to oxidize rapidly
to Fe3+ when in the form of a complex with a ligand containing
oxygen as the ligating atom.
Nicholson et al. (2, 3) reported that pyrite oxidation
kinetics in a bicarbonate (HCO3-) buffered system initially
increased, but hypotheses regarding this behavior were not
given. Hood (11) also reported that the rate of pyrite oxidation
by atmospheric O2 at room temperature under abiotic
conditions increased when HCO3- concentration increased.
The author postulated that the cause for the increase in pyrite
oxidation was formation of a pyrite surface-Fe(II)CO3
complex which facilitated electron transfer from Fe(II) to O2.
However, no evidence of pyrite-Fe(II)CO3 complexes was
presented by Hood (11). The objectives of this study were
to provide molecular evidence of pyrite surface-Fe(II)-CO3
complexes and to evaluate the influence of solution HCO3on pyrite oxidation.
Theoretical Considerations
To explain the potential enhancement of pyrite oxidation by
HCO3- or CO2 gas, we are postulating the following models.
Formation of a ferrous-bicarbonate complex on the surface
of pyrite occurs as follows
Fe-SA-SB:Fe(OH)2 + CO2 f
Fe-SA-SB:FeHCO3+ + OH- (2)
or
Fe-SA-SB:Fe(OH)2 + HCO3- f
Fe-SA-SB:FeHCO3+ + 2OH- (3)
S0013-936X(97)00829-8 CCC: $15.00
 1998 American Chemical Society
Published on Web 05/21/1998
Formation of the ferrous-bicarbonate complex on the surface
of pyrite would increase the basicity of Fe2+ leading to its
rapid oxidation (9, 12)
Fe-SA-SB:FeHCO3+ + 1/4O2 + H+ f
Fe-SA-SB:Fe(OH)2+ + CO2 + 1/2H2O (4)
Reaction 4 spontaneously forms a persulfido bridge after
Fe(II) oxidation followed by decarboxylation and pyrite
oxidation initiates as shown below
Fe–SA–SB:Fe3+
•
Fe–SA–SB :Fe2+
persulfido bridge
(5)
free radical
This leads to a continuous loss of electrons until thiosulfate
is produced as shown below
–
–
–
O
Fe–SA–SB:(Fe3+)5 + 3H2O
–
–
Fe–SA SB O + 6H+ + 5Fe2+ (6)
O
–
thiosulfate
In the above reactions, -S2 undergoes five continuous redox
reactions to produce thiosulfate. Thiosulfate is rapidly
oxidized by H2O2 or Fe3+ to SO4, yielding eight more H+ ions
(details are given in ref 13 and references therein).
It follows that pyrite oxidation by an oxidizer, such as
H2O2, is an autocatalytic process since one of the oxidation
products, Fe3+, can also oxidize pyrite (1, 4, 5)
FeS2 + 7H2O2 f Fe2+ + 2SO42- + 2H+ + 6H2O (7)
Fe2+ + H+ + 0.5H2O2 f Fe3+ + H2O
(8)
and
FeS2 + 14Fe3+ + 8H2O f 15Fe2+ + 2SO42- + 16H+ (9)
The rate law describing production of SO4 through oxidation
of FeS2 by H2O2 (eq 7) and Fe3+ (eq 9) can be written as (13)
d[SO4]/dt ) (k1[H2O2] + k2[Fe3+])S
(10)
where SO4 is directly related to the quantity of decomposed
pyrite at any time t; [H2O2] and [Fe3+] are the concentrations
of H2O2 and Fe3+; k1 and k2 are the rate constants of H2O2 and
Fe3+, respectively; and S is the surface area of pyrite available
to react. Assuming that, during oxidation of a small portion
of the available pyrite, the newly exposed pyrite surface (S)
remains proportional to the number of moles of unreacted
FeS2 in the system, then (14)
S ) K[FeS2]
(11)
where K is a constant. Considering also that during pyrite
oxidation, as t approaches zero S, [H2O2] and Fe3+ (eq 10)
remain relatively constant, the oxidation process can be
expressed as zero-order reaction
SO4produced ) k′t
(12)
where k′ is an empirical rate constant related to the
concentrations of H2O2 and Fe3+; and the corresponding rate
constants, k1 and k2 (eq 10), overall oxidizing solution
composition; and surface area of available pyrite. Since
HCO3- can enhance oxidation of Fe2+, increasing HCO3concentration should increase the rate of pyrite oxidation by
rapidly regenerating Fe3+ (1, 4).
Ionic strength is a second potential reason for the HCO3to affect the oxidation rate of pyrite. Its influence on the rate
of reaction between any two chemical species was described
by Benson (15) as follows:
log(kexp/kid) ) 1.0ZAZB(I)1/2
(13)
where kexp ≈ k′ (eq 12) denotes the experimental rate constant
under a given ionic strength; kid denotes the rate constant
at infinite dilution; ZA and ZB denote the charges of ions A
and B, respectively; and I denotes ionic strength. A plot of
log(kexp/kid) versus (I)1/2 would produce a straight line with
slope ZAZB. Benson (15) pointed out that when one of the
Zi values in eq 13 is 0, the ionic strength would not influence
the reaction rate. On the other hand, when one of the Zi
values is negative and the other is positive, the influence of
ionic strength on the rate of reaction would be negative.
When both Zi values are positive, the influence of ionic
strength on reaction rate would be positive. In the case of
Fe2+ being oxidized to Fe3+ by H2O2, the latter’s charge is 0,
and for this reason the influence of ionic strength on Fe2+
oxidation would be negligible. Assuming that Fe2+ oxidation
is the rate-controlling step in pyrite oxidation (2) under our
experimental conditions, one could then assume that ionic
strength would have negligible effect on pyrite oxidation.
However, any increase in pyrite oxidation due to increasing
HCO3- concentration would imply that -Fe(II)HCO3- complexes were responsible for the rapid regeneration of Fe3+.
Materials and Methods
The pyrite sample used in this study was obtained from the
University of Kentucky Applied Energy Research Laboratory,
Lexington KY. It was separated by gravity, using a water
column, from coal material obtained from seam number 9
in western Kentucky. The sample was pulverized and passed
through a 37 µm sieve, then washed with 4 M hydrofluoric
acid to remove silicate and Fe oxides, rinsed repeatedly with
N2-purged distilled water and acetone, and stored in a
desiccator under vacuum (16, 17). This pyrite sample was
then characterized prior to surface spectroscopic studies.
Characterization of sample crystallinity was carried out by
X-ray diffraction analysis (XRD) using an 1840 Philips Co-Ka
diffractometer, while surface morphologies of individual
pyrite grains were examined using a scanning electron
microscope (SEM) [Hitachi S-800]. Specific surface analysis
was carried by multipoint BET method using a Quantachrome
Autosorb 6 instrument.
Pyrite oxidation experiments were carried out by placing
approximately 50 mg of clean pyrite into small glass vials.
The unstoppered vials were then put into a chamber, and
various gas treatments were introduced. These treatments
were the following: (a) a control sample with a constant
exposure of pyrite to an N2 gas atmosphere, (b) a 14-day
exposure to atmospheric air at 100% relative humidity [After
exposure to 100% relative humidity, a portion of the sample
was rinsed with 10 mL of 4 M HCl through the use of a buchner
filter-funnel under a N2 atmosphere], (c) a 14-day exposure
to atmospheric air at 100% relative humidity followed by
heating at 50, 100, and 249 °C [heating was carried out by
using 25-50 mg of pyrite in a furnace DuPont S/N-X, FTIR-TGA interface] for 4 h under a N2 gas atmosphere, followed
by a 2-h evacuation in a desiccator, and (d) a 14-day exposure
to atmospheric air at 100% relative humidity followed by
extraction with 1 M KCl to determine total iron [Fe3+ plus
Fe2+] and Fe2+ on the surface of pyrite (18) [all samples
representing all treatments were evacuated for 2 h, using a
desiccator, prior to scanning].
For iron analysis, 10 mg of the pyrite samples representing
the control and the 14-day exposure to 100% relative humidity
was placed into 50-mL centrifuge tubes to which 10 mL of
VOL. 32, NO. 14, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2085
1.0 M KCl was added. After equilibration for 2 h using endto-end shaking, the samples were centrifuged and the
supernatants were collected for analysis. Iron was quantified
colorimetrically following procedures outlined by American
Water Works Association (19). For total Fe determination,
150 µL aliquots of samples and standards were pipetted into
a disposable microplate and treated with 60 µL of hydroxylamine/phenanthroline solution and 30 µL of an ammonium
acetate pH 5 buffer solution. The samples were shaken on
a vortex mixer for 10 min before they were scanned at 540
nm (Bio-Tek Instruments microplate autoreader EL311). For
Fe2+ determination, 150 µL aliquots of samples and standards
were pipetted into a disposable microplate to which 60 and
30 µL of the phenanthroline and ammonium acetate buffer
solutions, respectively, were added. The samples were
shaken on a vortex mixer for 10 min before scanning them
at 450 nm.
Infrared (IR) analyses were carried out employing a Nicolet
5XSC Fourier transform infrared spectrophotometer (FT-IR).
The procedure for generating FT-IR spectra involved mixing
100 mg of KBr with 2.5 mg of pyrite and placing the mixture
into the FT-IR sample holder (20). All FT-IR analyses were
performed under a nitrogen gas atmosphere. For background
correction KBr alone was used after the instrument was
purged with N2 gas and 50 scans were collected. This number
of scans was determined by trial and error using as criteria
reproducibility of noice-clean spectra. A similar approach
was used for all pyrite samples scanned. In addition to pyrite
spectra, FT-IR spectra of reagent grade NaHCO3, Na2CO3,
KHCO3, and K2CO3 were also obtained as references for
carbonate and bicarbonate species. All treatments were
duplicated, and a minimum of two subsamples from each
duplicate was used to generate FT-IR spectra.
The kinetics of pyrite oxidation were studied employing
a bed-reactor, consisting of a nuclepore swin-lok filter holder
(25-mm inside diameter) similar to that described by Jardine
and Sparks (21). The segregated pyrite from the coal, as
previously described, was ground and passed through 47and 75-µm sieves. The fraction 47 µm and 75 µm was used
for the present study. Prior to using the pyrite for leaching
studies, the samples were washed with 4 M hydrofluoric acid
and repeatedly rinsed with N2-purged distilled water and
acetone. For each test run, approximately 50 mg of the
separated pyrite was placed in a nuclepore filter holder
between two Whatman no. 42 ashless filters. The column
was first rinsed with 5 mL of 1 N HCl and then 5 mL of 5 mM
NaCl to remove any surface oxidation products prior to
initiating leaching. The leaching solutions were passed
through the pyrite nuclepore filter column at a constant flow
rate of 0.50 mL min-1 at room temperature (22 °C). Aliquots
of the effluent solution were collected every 2 min for the
first 16 min and then at 10 min intervals for a total of 1 h
using an Eldex Universal fraction collector.
Five duplicate tests were run: (1) 10 mM NaHCO3 with
20 mM H2O2; (2) 100 mM NaHCO3 with 20 mM H2O2; (3) 10
mM NaHCO3 with 2 mM H2O2; (4) 100 mM NaHCO3 with 2
mM H2O2; and (5) 10 mM NaHCO3, 20 mM H2O2, and 100
mM NaNO3. The ranges of NaHCO3 and NaNO3 concentration were chosen on the basis of the results reported by
Millero and Izaquirre (12). Their results showed that abiotic
solution Fe2+ oxidation rates were not affected by 10 mM
solutions of NaHCO3 or NaNO3 relative to the control (absence
of NaHCO3 or NaNO3). Furthermore, solution Fe2+ oxidation
rates remained unaffected when in the presence of 100 mM
NaNO3 but increased approximately 10-fold when in the
presence of 100 mM NaHCO3.
Activation energies (Ea) were determined by assessing
pyrite oxidation kinetics at temperatures of 10, 25, 30, and
40 °C using 20 mM H2O2 with 10 or 100 mM NaHCO3.
Constant temperature during each run was maintained by
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ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 32, NO. 14, 1998
a 2095 Bath & Circulator using a pyrite sample holder which
consisted of a Pharmacia XK 26-mm i.d. column equipped
with thermostat jacket. The data were plotted according to
the Arrhenius equation (eq 12) [log k′ vs 1/T, where k′ denotes
experimental rate constant (eq 12) and T denotes temperature
in degrees Kelvin].
Sulfate concentration in all leachates was measured using
turbidimetry with BaCl2 (23). Regression analysis on SO4
production versus time was carried out using the duplicated
data with slope denoting the conditional zero-order rate
constant (eq 12).
Results and Discussion
X-ray diffraction analysis of the iron sulfide sample substantiated that the sample was pyrite (24). The micrograph
in Figure 1A shows the surface morphology of the pyrite
sample, demonstrating a spongy-like appearance and a large
specific surface. The micrograph in Figure 1B reveals the
absolute size and size distribution of the pyrite particles
selected for the wet chemistry oxidation study using H2O2
and HCO3-. It is important to note that the majority of the
particles appear to be fairly uniform in size, although a small
number of extremely small particles are also present. Specific
surface analysis for the less than 75 µm particles from a fourpoint N2 BET isotherm (25) was found to be 9.75 m2 g-1.
Commonly, coarse-grained, massive pyrite specimens in the
range of 38 to 250 µm particles have surface areas that range
from 0.071 to 0.047 m2 g-1 (5, 26). The large specific surface
of the pyrite in the present study may account for the high
acid mine drainage production potential for the geologic
stratum, Kentucky coal seam 9.
Figure 2 shows the carbonate and bicarbonate spectra of
KHCO3/K2CO3 and NaHCO3/Na2CO3, respectively. The data
in Figure 2I reveal a major vibrational band for Na2CO3
occurring around 1406 cm-1 and one for K2CO3 occurring at
around 1450 cm-1 (Figure 2II). These bands represented the
ν3 of the CO3 vibrations (27). The carbonate anion (CO32-)
is highly symmetrical, and ν3 is infrared active (27). When
the symmetry of carbonate is perturbed due to coordination
with cations, the degenerate vibrations (ν3) split into a number
of distinct vibrations which depend on metal-carbonate
coordination. For example, in the case of reagent grade K2CO3
(Figure 2II), the CO3 band split into three distinct bands
signifying that CO3 coordinated in a bidentate fashion (27).
In the case of NaHCO3 (Figure 2I) and KHCO3 (Figure 2II),
the two major vibrational bands were at 1601 and 1631 cm-1,
respectively (28).
There were significant differences in the vibrational
spectra of reagent grade Na2CO3/NaHCO3 and K2CO3/KHCO3
spectra. These differences were most likely due to differences
in the physicochemical properties between the two different
carbonate-associated cations, Na+ and K+. One major
difference between these two cations is their ionic potential,
defined as the ratio between charge and radius (c/r2) (29).
The ionic potential for Na+ is 1.11 and for K+ is 0.56.
Difference in ionic potentials produces differences in the
electron orbital perturbation of CO3 associated with these
two metals; for this reason the vibrational spectra of Na2CO3/
NaHCO3 and K2CO3/KHCO3 differ (Figure 2).
The FT-IR spectra for pyrite before and after oxidation by
exposure to humidified air for 14 days at room temperature
were presented in Figure 3. These spectra show that before
oxidation no absortion bands were apparent. After oxidation,
however, pyrite showed absorption bands at 3539 and 3406
cm-1. These bands represented OH most likely associated
with both Fe on the surface of pyrite and pyrite surface
adsorbed water. The data in Table 1 show the presence of
iron on the surface of the pyrite sample representing the
control. Furthermore, iron concentration on the pyrite
surface increased approximately 4-fold during oxidation,
FIGURE 1. Scanning electron microscope photograph of (A) surface of pyrite and (B) a number of pyrite particles less than 75 µm size
before any oxidation treatment.
although the ratio between Fe2+ and total Fe before and after
oxidation did not differ dramatically (0.38 vs 0.31, respectively)
The difference between total Fe and Fe2+ was assumed to
represent Fe3+.
After exposing the pyrite to humidified air for 14 days, the
sample was heated to various temperatures in a N2 atmosphere and then cooled to room temperature in a desiccator
under vacuum. The spectra of these pyrite samples under
the various heating treatments are presented in Figure 4.
Note, after heating the pyrite to 50 °C, the intensity of the
3539 cm-1 band (Figure 4I) decreased, suggesting that the
OH, coordinated to Fe2+ or Fe3+ on the surface of pyrite, was
consumed perhaps by increasing surface acidity. Pyrite
surface pH may decrease by decreasing water content (30)
on the pyrite surface and/or by increasing electron transfer
between the disulfide and Fe3+ at increasing temperatures.
As temperature increased, the intensity of the H2O band (3406
cm-1) decreased as expected. Finally, when the pyrite was
rinsed with 4 M HCl, neither OH band was apparent (spectra
E in Figure 4I) signifying that most surface iron was removed
from the surface of pyrite (also OH loss via the acid effect
outlined above).
Support for the above observations and conclusions is
also given by the spectra behavior around the 1000 cm-1
frequencies (Figure 4II). The emergence of 1227, 1090, and
1009 cm-1 absorption bands indicates the presence of SO42associated with Fe3+ on the surface of pyrite (31). The sulfate
ion (SO42-) is highly symmetrical, and ν3 (around 1100 cm-1)
is infrared active (27). When the symmetry of sulfate is
perturbed due to coordination with cations, the degenerate
vibrations (ν3) split into a number of distinct vibrations,
depending on strength of the complex, and ν1 is activated
(27). The apparent splitting of ν3 (1227, 1090 cm-1) and the
presence of ν1 (1009 cm-1) on the spectrum of oxidized pyrite
(Figure 4II) suggested that the sulfate may be bonded to Fe3+
in a monodentate fashion (27). When the oxidized pyrite
sample was heated to 249 °C, the 1227 cm-1 band disappeared, which suggested, as in the case of the OH bands
(Figure 4I), that pyrite surface pH decreased through water
loss (30), weakening Fe3+-SO42+ interactions by forming
HSO4-, or through surface acidification by increasing electron
transport between the disulfide and Fe3+ at increasing
temperature; SO4 was most likely coordinated to the newly
produced Fe2+ on the pyrite surface in a weak electrostatic
fashion.
The spectra of oxidized pyrite (Figure 4III) also exhibit
carbonate absorption bands at 1429 cm-1, which are assigned
to the ν3 of CO3 vibrations. According to Nakamoto (27), the
extent of splitting (number of distinguishable bands) of the
1429 cm-1 band, representing the free CO32- ion, could be
directly related to the bond strength between CO32- and a
metal cation or surface. Nonsplitting of the CO3 band in
Figure 4III indicates that pyrite surface-CO3 complexes were
most likely weak electrostatic complexes. One would assume
that the strong acid environment on the surface of pyrite,
created by the oxidation of the disulfide, is expected to prevent
formation of surface carbonate complexes. However, persistence of the carbonate absorption band (1429 cm-1), as
shown in Figure 4III, demonstrated that some of the acid
produced on the pyrite surface was removed through reaction
with pyrite as follows:
FeS2 + H2SO4 ) FeSO4 + H2S + S°
(14)
Equation 14 shows that H2SO4 could be removed from the
pyrite surface through formation of H2S, a weak acid with
pKa’s 7 and 12.9 (32). Therefore, after pyrite exposure to
humidified air and heating, its surface was most likely covered
with ferrous sulfate and S° and most active anodic sites (SB)
were saturated with Fe2+ (4, 32).
The spectra in Figure 4III also show a number of
vibrational bands in the 1600 cm-1 range. The 1621 cm-1
band was most likely due to H2O deformation. The absorption bands at 1653 and 1670 cm-1 most likely represent
carboxylate species in two different chemical environments
(27, 28). When the oxidized pyrite was heated to 50 °C under
a N2 gas atmosphere, the absorption band at 1653 cm-1 was
eliminated perhaps due to pyrite surface acidification inducVOL. 32, NO. 14, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2087
TABLE 1. Total Fe and Fe2+ on the Surface of Pyrite before
and after Oxidation
sample
total Fe (mg g-1)
Fe2+ (mg g-1)
before oxidation
after oxidationa
4.72
22.48
1.82
6.99
a
14-day exposure to humidified air.
1670 cm-1 bands (spectra E, Figure 4III), suggesting that the
pyrite surface carbonate was removed due to surface
acidification, inducing conversion of pyrite surface HCO3to CO2 gas. The absence of the 1621 cm-1 band from spectra
E (Figure 4III) could be due to production of elemental sulfur
on the surface of pyrite (23), a hydrophobic substance.
It follows from above that carbon dioxide may react with
the surface of pyrite via two mechanisms. One mechanism
may involve coordination of HCO3- by Fe2+ on the surface
of pyrite as shown below:
–
–O
–Fe(II)–SA–SB:Fe(II)OH + C
O
O
–Fe(II):SA–SB:Fe(II)OC
–
–
(15)
OH
This pyrite surface-bicarbonate complex is speculated
to be sensitive to acid attack, releasing CO2 at about pH 4.2
[approximately 2 pH units lower than the pKa of H2CO3 (32)].
The second pyrite surface-CO2 complex may involve formation of a sulfur-carbon bond on the pyrite surface as shown
below:
–
–O
–Fe(II)–SA–SB: + C
O
–
FIGURE 2. FT-IR diffuse reflectance spectra of (IA) Na2CO3; (IB)
NaHCO3; (IIA) K2CO3; and (IIB) KHCO3.
FIGURE 3. FT-IR diffuse reflectance spectra of pyrite: (A) before
oxidation (after washing with 4 M hydrofluoric solution) and (B)
after oxidation by exposure to air at 100% relative humidity for 14
days.
ing conversion of pyrite surface HCO3- to CO2 gas (decarboxylation), while the 1670 cm-1 band remained intact. When
the temperature was raised to 100 °C, the intensity of the
1621 cm-1 band decreased due most likely to water removal
from the surface of pyrite; the 1670 cm-1 band was completely
eliminated. This suggested that the 1670 cm-1 band could
be representative of a second carbonyl which was removed
at a higher temperature than the carbonyl represented by
the 1653 cm-1 band. After washing the oxidized pyrite sample
with HCl, its spectrum did not exhibit the 1429, 1653, and
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ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 32, NO. 14, 1998
O
(16)
–Fe(II):SA–SB–C
–
O
This complex is speculated to be more resistant to acid
attack, perhaps because of the covalent interaction between
the carbon and sulfur atoms; its removal from the pyrite
surface takes effect at higher temperature (Figure 4III).
However, the exact reasons for this difference in temperature
resistance between the two carbonyls are not clearly apparent
from the present data.
The data in Figure 5 and Table 2 show that pyrite oxidation
is affected by the presence and concentration of HCO3-. At
a H2O2 concentration of 20 mM, increasing the HCO3concentration from 10 to 100 mM increased oxidation rate
by about 174% (calculated from the difference between
slopes) (Figure 5A). Similarly, when H2O2 concentration was
set at 2 mM, the oxidation rate was about 125% higher at the
100 mM HCO3- concentration than at the 10 mM HCO3concentration (Figure 5B). The concentration of H2O2 in
this experiment also influenced pyrite oxidation. Its influence, however, was found to be smaller than that of HCO3within the range of H2O2 and HCO3- concentrations tested.
At a HCO3- concentration of 10 mM, increasing the H2O2
concentration from 2 to 20 mM increased the oxidation rate
by about 75%. Similarly, when HCO3- concentration was
set at 100 mM, the oxidation rate was about 113% higher at
the 20 mM H2O2 concentration than at the 2 mM H2O2
concentration. The role of H2O2 on pyrite oxidation was
most prominent at the higher HCO3- concentration. The
above results imply that at the higher HCO3- concentration,
the oxidation rate was controlled mainly by the more rapid
regeneration of Fe3+ due to HCO3-. It appears that Fe3+ is
a more effective pyrite oxidant than H2O2 because of its small
ionic radius (steric effect) and large potential to form innersphere complexes (transition-state intermediate) with the
pyrite surface. On the other hand, H2O2 is a molecule with
a relatively large radius, thus limiting its potential to form
FIGURE 5. Influence of sodium bicarbonate concentration on the
oxidation of pyrite using (A) 2 mmol L-1 H2O2 and (B) 20 mmol L-1
H2O2. The two slopes within each graph are significantly different
at the 95% confidence level according to the t-test.
TABLE 2. Summary of Reaction Rates of the Various Systems
Tested
FIGURE 4. FT-IR diffuse reflectance spectra of pyrite after the
following treatments: (A) oxidation with air for 14 days; (B) heating
at 50 °C; (C) heating at 100 °C; (D) heating at 249 °C; and (E) washing
with 4 M HCl (Roman numerals I, II, and III signify different
wavenumber ranges).
pyrite surface complexes with all reactive sites potentially
available to Fe3+, and with a low potential to form pyrite
oxidizing systems
slope
r2
20 mmol L-1 H2O2 plus 10 mmol L-1 HCO320 mmol L-1 H2O2 plus 100 mmol L-1 HCO32 mmol L-1 H2O2 plus 10 mmol L-1 HCO32 mmol L-1 H2O2 plus 100 mmol L-1 HCO320 mmol L-1 H2O2 plus 100 mmol L-1 NaNO3
0.042
0.115
0.024
0.054
0.034
0.999
0.996
0.999
0.996
0.996
surface inner-sphere complexes (16).
The data in Figure 6 show that the increase in pyrite
oxidation when HCO3- concentration was increased from
10 to 100 mM could not be attributed to ionic strength. This
can be deduced from the fact that the addition of NaNO3
(100 mM) in the 10 mM HCO3- solution plus 20 mM H2O2
did not in any way affect the oxidation of pyrite. Activation
energy (Ea) data (Table 3) were also consistent with what one
might expect if indeed -Fe2+HCO3 complexes controlled the
rate of pyrite oxidation. In other words, increasing HCO3concentration in the oxidizing solution should decrease Ea.
The data in Table 3 show that upon increasing HCO3concentration from 10 mM to 100 mM, Ea decreased from
28.96 to 25.92 kJ mol-1. The difference in Ea between the two
systems was small but meaningful. Note, however, the Ea
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9
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TABLE 3. Apparent Activation Energies of Pyrite Oxidation by 20 mmol L-1 H2O2 as Affected by NaHCO3
a
NaHCO3
treatment
(mmol L-1)
Arrhenius regression
equations (log k′ ) log A Ea/RT)
10
100
log k′ ) 3.692 - (1511.5)(1/T)
log k′ ) 3.306 - (1352.4)(1/T)
n
apparent
activation energies
(kJ mol-1)
log k′
vs 1/T
(r 2 )
4
4
28.96 ( 0.045a
25.92 ( 0.023a
0.896
0.977
Confidence interval at 95%.
the abiotic pyrite oxidation rate aided by HCO3- is not
expected to be as rapid as the biotic oxidation rate (1).
Acknowledgments
This research was supported, in part, by funds provided by
the United States Department of Energy and the Pittsburgh
Energy Technology Center under Grant DE-FG22-95PC5226.
This support is gratefully acknowledged. We also wish to
thank Ms. Joni Norris and Mr. Martin Vandiviere for their
help in generating the data. This research was published
with the approval of the Director of Kentucky Agricultural
Experimental Station, Lexington.
Literature Cited
FIGURE 6. Influence of 100 mmol L-1 NaNO3 on the oxidation of
pyrite by 20 mmol L-1 H2O2 and 10 mmol L-1 sodium bicarbonate.
The two slopes are not significantly different at the 95% confidence
level according to the t-test.
values reported represent apparent activation energies, which
include diffusion-controlled reactions on the surface of pyrite
as well as adsorption and precipitation. The overall data
suggest that -Fe2+HCO3 complexes on the surface of pyrite
increased its oxidation rate.
Conclusions
FT-IR spectroscopic evidence is presented which demonstrates that pyrite exposed to humidified CO2 plus O2 formed
pyrite surface-CO2 complexes. Two potential mechanisms
were proposed to account for these complexes. One mechanism involved formation of a weak pyrite surface-Fe(II)HCO3 complex, whereas a second mechanism involved
formation of a pyrite surface-carboxylic acid group complex
[-Fe(II)SSCOOFe2+]. It is hypothesized that both complexes
could promote abiotic pyrite oxidation by accelerating
regeneration of Fe3+, an effective pyrite oxidant. Our
experimental data clearly showed that abiotic pyrite oxidation
increased in the presence of solution HCO3-.
The approach currently used to prevent pyrite oxidation
in the field is mainly based on eliminating Fe3+ from pore
waters (10, 13). This approach includes using limestone to
precipitate Fe3+ as iron hydroxide/oxyhydroxide (13) and
raising pH to diminish activity of T. ferrooxidans (1). Recent
evidence (4), however, showed that Fe3+, associated with the
surface of pyrite, is an effective pyrite oxidant at low pH as
well as at near-neutral pH. The findings in this study suggest
that in pyritic geologic waste the presence of HCO3-, added
in the form of either limestone or Na2CO3, may accelerate
the abiotic oxidation rate of pyrite. Therefore, application
of alkaline materials to pyritic waste may neutralize AMD
but may not inhibit sulfate production. Observations to this
effect have been reported in the literature (13). However,
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