JOURNAL OF COLLOID AND INTERFACE SCIENCE
ARTICLE NO.
178, 628–637 (1996)
0160
Interfacial Electrochemistry of Pyrite Oxidation and Flotation
I: Effect of Borate on Pyrite Surface Oxidation
XIANG-HUAI WANG 1
Department of Mining Engineering and Center for Applied Energy Research, 230 Mining and Mineral Resources Building,
University of Kentucky, Lexington, Kentucky 40506-0107
Received February 7, 1995; accepted September 22, 1995
Sodium tetraborate (Na2B4O7 ) has been widely used as an electrolyte and pH buffer in studying the interfacial electrochemistry
of sulfide minerals in relation to sulfide mineral flotation. In all
the previous studies published so far, borate was regarded as an
inert electrolyte/pH buffer, and its reactions with the sulfide minerals were completely overlooked. In this first part of this series
papers, the complicating effects of borate on the interfacial electrochemistry of pyrite have been studied. It has been demonstrated
that borate is not an inert electrolyte/pH buffer. It strongly reacts
with the surfaces of pyrite. In the borate solutions, the surface
oxidation of pyrite is strongly enhanced. The first and rate-determining step of the reaction between borate and pyrite has been
shown to be the following irreversible reaction:
FeS2 / B(OH) 0
4 c FeS2rrr[B(OH)4 ]ads / e.
This reaction appears in the voltammogram as an anodic oxidation
peak at potentials of more than 0.4 V lower than the commencement of pyrite oxidation in sodium perchlorate or nitrate electrolyte solutions. As the borate concentration increases, the peak current increases linearly, while the peak potential shifts positively
at 240 mV per decade. On a rotating-disc electrode, the peak
becomes a plateau. The limiting current density is a linear function
of the square root of the rotation speed at relatively low rotation
speeds. The Tafel slope, i.e., dE/d{log I}, is close to 240 mV
per decade and is independent of the rotation speed and borate
concentration. The results indicate that charge transfer coefficient
( a ) is 0.25. q 1996 Academic Press, Inc.
Key Words: pyrite; oxidation; flotation; borate; electrochemistry;
adsorption; dissolution; pH buffer, surface complexes.
INTRODUCTION
The interfacial chemistry of pyrite is of great industrial
importance in complex sulfide ore flotation, oxidative
1
To whom correspondence should be addressed. Present address: Research and Development, Betz Paper Chem, Inc., 7510 Baymeadows Way,
Jacksonville, FL 32256.
628
0021-9797/96 $18.00
Copyright q 1996 by Academic Press, Inc.
All rights of reproduction in any form reserved.
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leaching of gold from pyritic concentrate, coal desulfurization, acid mine drainage mitigation, and conversion of
solar energy to electrical or chemical energy. The surface
properties of pyrite, particularly its oxidation behavior,
have been studied extensively in relation to these industrial applications ( e.g., (1 – 21) ) . A number of excellent
review papers on the subject have been published ( 13,
19, 20, 22, 24 ) . However, in spite of the vast volume of
published studies on the interfacial chemistry of pyrite
oxidation, there is little agreement on the details of the
pyrite oxidation processes, particularly the mechanisms
and intermediate products of pyrite oxidation in the alkaline media. For example, using linear sweep voltammetry,
Hamilton and Woods ( 9 ) observed that a monolayer of
elemental sulfur was formed on pyrite when oxidized at
pH 9.2 and 13, while multiple layers of sulfur was produced in acidic media. However, later studies by Buckley
et al. ( 3 – 5 ) and Chander and co-workers ( 6 – 8 ) , using
electrochemical and spectroscopic surface analysis,
showed that the initial oxidation products of pyrite would
be better represented by metal-deficient pyrite surface
( Fe10xS2 ) . On the other hand, Yoon ( 25 ) , Luttrel and
Yoon ( 26 ) , and Zhu et al. ( 27, 28 ) suggested that iron
polysulfide ( FeSx , x ú 2 ) would be the oxidation products
of pyrite. Still other investigators argued that the sulfur
substance is of zero oxidation state and should be represented by elemental or atomic sulfur ( e.g., (13 – 29);
(27) ) . Fuerstenau and Sabacky ( 30 ) showed that pyrite
and most heavy metal sulfide minerals are naturally hydrophobic.
The above issue is fundamentally very interesting and
deserves further extensive studies to obtain more conclusive
results. There are a number of factors that can drastically
change the pyrite surfaces, such as, the source of the pyrite
sample, impurities, crystal defects, and surface preparation
methods. These factors have been examined in several recent
studies (7, 17, 18, 27, 28). However, so far no studies have
ever considered the effects of the electrolyte and the ‘‘impu-
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TABLE 1
Chemical Analysis of the Pyrite Sample
Element
Weight percent
Fe
S
Cu
Pb
Zn
Ca
Si
44.6
53.6
0.31
0.75
1.27
õ0.01
õ0.01
Fe/S Ratio Å 2.09
629
water. The polished electrode was then immediately inserted
into the electrochemical cell and the measurements were
promptly initiated.
Samples for flotation tests. The pyrite samples were first
crushed by hand and then dry-ground in a jar ball mill.
Particles of sized 150 1 200 and 200 1 325 mesh were used
for flotation tests.
Chemicals. All the chemicals used in the present study
were purchased from Aldrich Chemical Company, in the
purest grade available. All the solutions were prepared using
deionized distilled water.
Electrochemical Setup
rities’’ in the electrolyte on the surface oxidation of pyrite.
Our extensive literature review shows that almost all previous studies of pyrite electrochemistry in weakly to medium
alkaline solutions (pH 8–10) were conducted using borate,
carbonate, phosphate, or mixture of them, as the electrolyte
and pH buffer. Unfortunately, in all the previous studies,
these chemicals were treated as inert electrolytes/pH-buffers
and their possible interactions with pyrite were completely
overlooked.
In this paper, the effects of borate on the oxidation of
pyrite surfaces are investigated systematically using electrochemical techniques. To delineate the chemical reactions
occurred at the pyrite/borate solution interface, solution
equilibrium diagrams for the borate–iron–water systems
will also be presented. In the next few papers of the series,
the effects of borate on surface oxidation of lead and copper
sulfides and of pH-buffers/electrolyte on surface oxidation
of pyrites and their flotation will be presented.
A standard three-electrode system was used in the electrochemical measurements. A pyrite electrode prepared as
above described, a platinum electrode, and a saturated calomel electrode were used as the working, counter, and reference electrode, respectively. The potential of a saturated
calomel electrode (SCE) is related to the standard hydrogen
scale by the following equation:
ESCE Å Eh 0 245 (mV).
The electrochemical system used in the present study consists of an EG&G Princeton Applied Research ( PAR )
Model 273 potentiostat / galvanostat and a PAR model
660 rotating-disc electrode apparatus. The system, including data acquisition, is operated by a 386 IBM computer. All the measurements were performed with a usermodified Headstart Creative Software. The rotating-disc
electrode ( RDE ) technique enables one to study the effect
of rotating the electrode on the reaction processes oc-
EXPERIMENTAL
Materials
A number of natural pyrite samples from both ore and
coal sources and a synthetic single pyrite crystal were investigated. Details of the sample sources, their chemical composition and semiconducting properties have been described
elsewhere (18). The emphasis of this paper will be focused
on the ore-pyrite whose chemical composition is shown in
Table 1.
Pyrite electrode. Rectangular (between 5 1 5 and 7 1
7 mm) slabs were cut from the natural single crystals. The
slab specimen was attached to a brass using conductive carbon glue. The electrode was then encapsulated with epoxy
resin. The rectangular face was then ground to expose to the
solution. The measured geometric surface area was used to
calculate the current density. Before each experiment, the
electrode was wet ground with silicon carbide polishing paper (4000 grid), rinsed with ethanol and deionized distilled
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FIG. 1. Cyclic voltammograms of pyrite electrode in 0.1 M KClO4 at
pH 9.4. No pH buffer was used. The solution pH was adjusted with NaOH
and HClO4 . The initial potential was started from the open circuit potential
in the positive going direction. Sweep rate: 20 mV/s. Solid line: stationary.
Dashed line: rotation at 1000 RPM.
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FIG. 2. Cyclic voltammograms of stationary pyrite electrode in 0.1 M
KClO4 solution at pH 9.3 in the absence and presence of 1 1 10 03 and 1
1 10 02 M sodium borate, respectively. Sweep rate: 20 mV/s. Sweeps were
initiated in the positive-going direction from OCP. Upper potential limit:
1000 mV vs SCE.
curring at the electrode surface, and thereby to distinguish
the solution species involved reactions from solid species
reactions. The EG&G Princeton Applied Research manufactured RDE system was modified such that only the
lower portion of the shaft rotates, while the upper part
remains stationary. This enables the upper portion to keep
tight contact with the cell lid. The cell will thus be sealed
from the ambient, which prevent any oxygen / air entrance
into the electrochemical cell. The atmosphere in the cell
can be well controlled.
Procedures
Unless stated elsewhere, the electrochemical measurements were performed in 0.1 KClO4 electrolyte solution at
ambient temperature. The pH value of the solution was adjusted with NaOH or HClO4 . The solutions were purged
with ultrahigh purity nitrogen gas for at least 60 min before
the experiments to remove oxygen from the solution. A nitrogen gas flow was maintained above the solution surface
during the experiments to prevent the entrance of oxygen
from the system.
pendent on the concentration of the electrolyte. The results show that for pyrite only perchlorate and nitrate may
be considered as true inert electrolyte, and that perchlorate and nitrate do not interact with the pyrite electrode
surface. Figure 1 shows typical cyclic voltammograms of
pyrite electrode at pH 9.3 in 0.1 M KClO4 at stationary
and rotation conditions, respectively, in the absence of
any pH buffer. When the potential sweeps are initiated
in the positive-going direction, one anodic current rise
is observed at potential above 0.4 V ( vs SCE ) . This is
attributed to the oxidation of pyrite surfaces. The detailed
mechanisms of the oxidation process are very complex.
It has been observed that when cyclic voltammetric measurements are performed in either of the above electrolyte
solutions, the voltammograms show no pH dependence
in the pH range 5 – 9, if there are no pH-buffers in the
solution. The initial and rate-determining reaction of pyrite oxidation are shown to be as follows:
FeS2 / H2O c FeS2rrrOHads / H / / e.
[1]
This is in agreement with previous studies (1, 17, 18, 46,
47, 48).
In highly alkaline solutions (pH ú 9), an additional anodic oxidation peak appears at the potential above 200 mV
(vs SCE). It has been observed that the peak current is
directly proportional to the concentration of hydroxide ions,
i.e., solution pH (18). On a rotating-disc electrode, the peak
becomes a plateau with the limiting current being a perfect
linear function of the square root of the rotation speed. The
results are in excellent agreement with several previous studies (1, 17, 18, 44). Under such conditions, the rate-determining reaction of this process has been described by the following equation:
RESULTS AND DISCUSSION
Characteristics of Pyrite Oxidation in No-Buffer Solutions
In order to determine whether commonly used electrolytes, such as nitrate, perchlorate, sulfate, chloride, borate, etc, affect the oxidation of pyrite surfaces, systematic experiments have been performed for pyrite electrodes in varying concentrations of the electrolytes. It was
found that only in perchlorate and nitrate solutions, the
cyclic voltammograms are almost identical and are inde-
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FIG. 3. Cyclic voltammograms of stationary pyrite electrode at pH 9.3
in 0.1 M sodium borate and 0.1 M potassium perchlorate electrolyte solutions, respectively. Sweep rate: 20 mV/s.
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FIG. 4. Cyclic voltammograms of stationary pyrite electrode in 0.1 M
NaNO3 electrolyte solution at pH 9.3 in the absence and presence of 1.5
1 10 03 M sodium borate, respectively. Sweep rate: 20 mV/s. Upper potential limit: 800 mV.
FeS2 / OH 0 c FeS2rrrOHads / e.
[2]
More details of the results will be presented in the subsequent papers of this series.
Cyclic Voltammetry of Pyrite in Borate Solutions
Sodium borate ( NaB ( OH ) 4 ) or tetraborate ( Na2B4O7 )
has been widely used as the electrolyte and pH buffer in
many previous studies of sulfide oxidation and collector
adsorption on sulfide surfaces. In all the studies, it was
assumed that borate would not be involved in the reaction
processes. However, as will be shown below, this is not
true.
Figure 2 shows the cyclic voltammograms of the pyrite
electrode in the presence of 0, 1.0, and 10.0 m M of so-
631
dium borate at pH 9.3, respectively. As described above,
in the absence of borate, only one anodic current rise
( note: not a peak ) can be observed. However, in the
presence of borate, an additional anodic peak ( labeled A
in Fig. 2 ) appears at about 250 mV ( vs SCE ) . As the
concentration of borate increases, the peak current density
( Ip ) increases, while the peak potential ( Ep ) shifts positively. At very high borate concentrations, peak A merges
with B into one current rise, as shown in Fig. 3. The
current of the cathodic peak ( peak D ) on the reversing
sweep also increase with increasing the borate concentration. On the second positive-going sweep, the corresponding anodic peak ( E ) currents are increased. Similar results
have been observed when sodium nitrate ( NaNO 3 ) was
used as supporting electrolyte. Figure 4 presents typical
cyclic voltammograms for the pyrite electrode in 0.1 M
NaNO3 in the absence and presence of sodium borate.
Note that in the absence of borate, a significant reduction
peak ( peak C ) can be observed when the upper potential
limit is above 750 mV ( vs SCE ) . The peak current increases as the potential limit in the preceding sweep increases. In the presence of borate, the peak current of this
reduction peak decreases. At very high borate concentration, this peak disappears completely ( see Figs. 2 and 3 ) .
The reaction has been identified to be the reduction of
some soluble ferric ion species to ferrous species. Details
of the analysis for this peak will be presented in a subsequent paper. Here the emphasis is focused on identifying
peak A.
Figure 5 shows the cyclic voltammograms of the pyrite
electrode in 0.1 M KClO4 in the presence of varying concentrations of borate with the upper positive potential
limit being 800 mV ( vs SCE ) . For clarity, the voltammograms are plotted in two diagrams. It can be seen that the
FIG. 5. (a) and (b) Cyclic voltammograms of stationary pyrite electrode in 0.1 M KClO4 electrolyte solution in the presence of 1.0 1 10 03 –1 1
10 02 M sodium borate. Sweep rate: 20 mV/s. Potential sweeps initiated in positive-going direction.
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FIG. 6. Peak current density of peak A in Figs. 2–5 as a function of
borate concentration at pH 9.3 in 0.1 M KClO4 or 0.1 M NaNO3 electrolyte
solution.
peak current increases with increasing the borate concentration. Figure 6 indicates that the peak current density
is a linear function of the borate concentration. Interestingly, both the peak potential ( Ep ) and the half-peak potential ( Ep / 2 ) shifts toward higher positive potentials with
increasing borate concentration. When the borate concentration is greater than 3 1 10 02 M, peak A merges with
B into one current rise, as shown in Fig. 3. Similar results
have been obtained when NaNO3 , instead of KClO4 , was
used as electrolyte.
Figures 7 and 8 show the voltammograms of the rotatingdisc pyrite electrode in the presence of 1 1 10 03 M borate
at pH 9.3 at different rotation speeds. For clarity, in Fig. 8,
only the first potential sweep voltammograms are presented.
It can be seen that on a rotating disc electrode, peak A
becomes a plateau. At lower rotation speed ( õ1000 rpm)
FIG. 7. Cyclic voltammograms of stationary and rotating-disc (2000
rpm) pyrite electrode in 0.1 M NaNO3 electrolyte solution at pH 9.3 in the
presence of 1 1 10 03 M sodium borate. Sweep rate: 20 mV/s. Potential
sweeps were initiated in the positive-going direction.
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FIG. 8. Linear sweep voltammograms of pyrite electrode in 0.1 M
KClO4 electrolyte solution at pH 9.3 in the presence of 1 1 10 03 M sodium
borate at different rotation speeds. Sweep rate: 20 mV/s.
and low borate concentrations, the limiting current can be
well distinguished. Furthermore, the limiting current density
is a linear function of the square root of the rotation speed,
as can be seen from Fig. 9. However, at high rotation speed
or high borate concentrations, no well-defined limiting current can be observed. This is because at such conditions, the
electrode potentials of the mass diffusion superimpose with
that of the subsequent pyrite oxidation reactions (see Fig. 2
and 3).
These results demonstrate that pyrite oxidation in borate
solution is considerably enhanced. Furthermore, the rotation
and borate concentration dependences of the reaction indicate that borate species in the solution must have been involved in the pyrite oxidation process. In order to determine
the precise chemical nature of the reaction, further analysis
of the electrochemical data are required. Let us suppose
FIG. 9. Limiting current density (IL ) of A as a function of the rotation
speed in the presence of 1 1 10 03 M sodium borate at pH 9.3. Sweep rate:
20 mV/s.
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ELECTROCHEMISTRY OF PYRITE OXIDATION AND FLOTATION, I
TABLE 2
Peak Potential and Peak Current Values of Pyrite in Varying Borate Concentrations
Ep 0 Ep/2
Borate conc. (M)
Ep (SCE) (mV)
Ip (mA/cm2)
Ep/2 (SCE) (mV)
Ip/2 (mA/cm2)
(mV)
Note
302
370
384
460
510
566
560
—
0.156
0.207
0.271
0.516
0.734
0.922
1.274
—
131
218
211
283
338
388
358
0.0785
0.103
0.136
0.258
0.366
0.461
0.636
171
162
173
177
172
178
202
KClO4
NaNO3
KClO4
KClO4
KClO4
KClO4
NaNO3
1.0 1 1003
1.5 1 1003
2.5 1 1003
5.0 1 1003
7.5 1 1003
1.0 1 1002
1.2 1 1002
Average
—
that the electrochemical reaction may be expressed by the
following equation:
—
176
log I Å log(nFK 0f)
/ m log[CB ( OH ) 40 ] /
FeS2 / mB(OH) c FeS2rrr[B(OH)4 ]m / ne.
0
4
[3]
Slope Å
Therefore, we need to determine the values of m and n.
The parameters can be evaluated from the kinetics of the
reaction. Assuming that the above reaction is the ratedetermining step, then the reaction kinetics can be expressed by
V Å K f r[C
0
B ( OH ) 4
m
] (1 0 q ).
[4]
In terms of current density, this can be written
I Å nFV Å nFK f [CB ( OH ) 40 ] mr(1 0 q )
Å nFK 0f[CB ( OH ) 40 ] mr(1 0 q )rexp
F
G
anF
E ,
RT
—
mÅ
Z
anF
E
2.303r RT
Ì{log I}
anF
CB ( OH ) 40 Å
ÌE
2.303r RT
Ì{log I}
Ì{log CB ( OH ) 40 }
Z
.
[6]
[7a]
[7b]
E
Therefore, at a given borate concentration, plotting log I
vs E will gives the Tafel slope, from which the value of
( an ) can be calculated; while at a given potential ( E ) ,
plotting log I vs log { borate concentration } will enable
us to evaluate m , the reaction order with respect to borate.
Note that careful inspection of the voltammograms presented above shows two distinguished features: ( a ) at a
stationary electrode, there is the total absence of a reverse
[5]
where
K f Å rate constant in the forward direction;
q Å fraction of surface covered by adsorbed borate;
a Å charge transfer coefficient;
F Å Faraday constant;
T Å absolute temperature in K;
R Å gas constant;
E Å potential (in volt)
Note that if reaction [3] is the rate-determining step, the
subsequent reactions will be very fast compared to reaction
[3]. This means that the surface coverage ( q ) will be fairly
small and constant, that is, 1 0 q c 1. Then, expressing
Eq. [5] in the form of Tafel equation yields
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FIG. 10. Current density (I) of stationary pyrite electrode at electrode
potential E Å 0 mV vs SCE as a function of sodium borate concentration.
Sweep rate: 20 mV/s.
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XIANG-HUAI WANG
FIG. 11. log(concentration) 0 pH distribution diagrams for the borate–water system at total borate concentration (a) 1 1 10 03 M and (b) 0.1 M.
peak on the return sweep ( see Figs. 2 – 5 ) ; ( b ) at the
rotating-disc electrode, the current density is independent
of the rotation speed at low potential and low current
density, while the limiting current is a linear function of
the rotation speed at the mass transport potential region.
These features indicate that the electrode reaction is irreversible ( 31 – 33 ) . This suggests that plotting log ( I ∗ IL /
( IL 0 I ) ) against E should give a linear line and be independent of the rotation speed and borate concentration.
The Tafel plots at different rotation speeds at a given
borate concentration all fall into the same line with a
Tafel slope of 240 mV per decade. Similarly, the Tafel
plots ( log I vs E ) for the stationary electrode at different
borate concentrations give a Tafel slope is close to 240
mV per decade, independent of the borate concentration.
These results mean that the product of and a and n is
0.25. For an irreversible reaction, it has been shown that
theoretically, EP 0 EP / 2 Å 47.7 / ( an ) ( 31 – 33 ) . The results given in Table 2 reveal EP 0 EP / 2 Å 176 mV. This
confirms that the reaction is irreversible and an Å 0.25.
From these analysis we can assume with great confidence
that n Å 1 and a Å 0.25 ( 31 – 34 ) . A more accurate
analysis has been conducted and the results have confirmed this. Details of the method will be described in
another paper of this series ( 45 ). Figure 10 shows the
relationship between the current density at 0.0 V ( vs
SCE ) and the borate concentration. The slope of the regression line is 0.96. This indicates that the reaction order
with respect to borate is one, i.e., m Å 1.
From the above accurate kinetic analysis, the electrode
reaction has been determined, that is,
TABLE 3
Stability Constants of Borate and Ferric–Borate Compounds
The reaction is probably only the rate-determining step
in the overall complicated pyrite oxidations.
Another interesting phenomenon observed in the electrochemical studies is the shift of the peak potential with
the borate concentration. However, classical cyclic voltammetry theories ( 35 ) dealt with only rather simplistic
homogeneous ( electro ) chemical reactions, they ignore
many mechanistic nuances and cannot satisfactorily explain the above phenomenon. From the theories developed by Parker and co-workers ( 35, 39 ) , we may derive
the following equations for the pyrite-borate electrode
reaction [ 8 ] :
Reaction
log K
Ref.
A. Borate–water system
H / B(OH) r HB(OH)4
2H/ / 3B(OH)04 r B3O3(OH)04 / 5H2O
H/ / 3B(OH)04 r B3O3(OH)05 / 4H2O
2H/ / 4B(OH)04 r B4O5(OH)20
4 / 7H2O
4H/ / 5B(OH)04 r B5O6(OH)04 / 10H2O
/
0
4
9.236
20.07
10.40
20.90
38.20
a
1.0)
1.0)
1.0)
1.0)
a
8.58 (I Å 1.0)
15.54 (I Å 1.0)
20.6 (I Å 1.0)
b
(I
(I
(I
(I
Å
Å
Å
Å
a
a
a
B. Ferric–borate–water system
Fe / B(OH) r Fe[B(OH)4]2/
Fe3/ / 2B(OH)04 r Fe[B(OH)4]/2
Fe3/ / 3B(OH)04 r Fe[B(OH)4]3
3/
a
b
0
4
b
b
FeS2 / B ( OH ) 40 c FeS2rrr[ B ( OH )4 ]ads / e .
[8]
ÌEP
Ì log V
Z
CB ( OH ) 40 Å
2.303r RT
Å 240 (mV)
anF
Smith and Martell (40).
Sillen and Martell (41, 42).
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ELECTROCHEMISTRY OF PYRITE OXIDATION AND FLOTATION, I
635
FIG. 12. log(concentration) 0 pH distribution diagrams for the ferric–borate–water system at total ferric concentration of 1 1 10 04 M and total
borate concentration (a) 0 M and (b) 0.05 M.
ÌEP
Ì log CB ( OH ) 40
Z
Å mr
V
2.303r RT
Å 240 (mV)
anF
Fe3/ / 2B(OH) 4 0 r Fe[B(OH)4 ] 2/
Fe
for m Å 1.
The Solution Chemistry of Borate and Borate–Iron
Systems
In order to further understand the chemical principles
of the above electrochemical phenomena, the solution
chemistry of the borate – water and borate – iron systems
has been investigated. It was found that the chemistry of
borate – water system is very complex, much more complicated than it is generally thought. Besides the protonation reactions of borate, borate undergoes polymerization under high concentrations to form a number of polymers, B3O3 ( OH ) 4 0 , B3O3 ( OH ) 5 0 , B4O5 ( OH ) 20
4 , and
B5O6 ( OH ) 4 0 . Table 2 presents the standard reaction constants of the various borate species ( 40 ) . Figure 11 shows
the concentration – pH distribution diagrams for the borate – water system at two typical borate concentrations
used by various previous investigators. It can be seen that
at high concentrations, all the polyborates exist in high
quantities.
More importantly, it has been found that borate is a strong
complexing agent for heavy metal ions such as Fe3/ , Pb 2/ ,
and Cu 2/ (41, 42, 43). In the ferric-borate solution system,
borate reacts with the ferric ions, forming a number of soluble species:
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/ 3B(OH) 4 0 r Fe[B(OH)4 ]3 (aq).
[11]
[12]
[9b]
This predicts that the peak potential of electrode reaction
[8] will shifts 240 mV for a 10-fold change in the borate
concentration, which is in good agreement of the experimental data (Table 2).
Fe3/ / B(OH) 4 0 r Fe[B(OH)4 ] 2/
3/
[10]
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The stability constants of these complexes are given in
Table 3. It can be seen that the stability constants are
relatively large. Thus, it can be expected that in the borate
solution, the iron hydrolysis reactions will be strongly
affected by the formation of such soluble complexes. Figure 12 compares the distribution of ferric species in the
absence and presence of 0.05 M borate. It can be noticed
that in the presence of borate, even ferric hydroxide is
unstable. It dissolves to form ferric-borate species. The
precipitation pH of Fe( OH )3 ( s ) is shifted from pH 3 to
pH 7 – 9, depending on the ratio of the total ferric ion
concentration to borate concentration. Consequently, due
to the formation of the soluble ferric-borate complexes,
the surface oxidation of pyrite will be strongly affected.
The solution chemistry studies provide very valuable information in understanding the effects of borate on oxidation and flotation of pyrite. As demonstrated by the electrochemical studies, borate is chemisorbed on pyrite surfaces according to reaction [ 8 ] . As the borate
concentration increases, the degree of pyrite oxidation
will increase. On the other hand, since borate is bonded
to the surface iron atoms, when the borate concentration
is increased to the appropriate concentration, the surface
ferric-borate complex ( es ) will be detached from surface
into the solution, ( i.e., dissolution ) . This is in good
agreement with the electrochemical results.
It should be pointed out that the detailed chemical
mechanisms of the subsequent pyrite – borate reactions
after the initial reaction ( eqn. ( 8 ) ) cannot be established
from the conventional CV and electrode kinetics studies.
However, from the electrochemical studies presented
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XIANG-HUAI WANG
above, it can be expected that the subsequent reactions
must be very fast. Precise identification of such fast reactions can only be determined by other techniques. More
studies are being conducted. Finally, the present study
suggests that for those previous electrochemical studies
( on the surface oxidation and adsorption of collectors on
pyrite and other sulfide minerals ) that were performed in
borate solutions, the interpretation and conclusions might
have been seriously biased due to that the effects of borate
on the reactions were overlooked. Therefore, more careful
and systematic reinvestigations appear to be highly necessary.
The electrochemical measurements, in conjunction
with flotation and solution chemistry studies, have clearly
demonstrated that borate, a commonly pH buffer and
electrolyte, reacts strongly with pyrite surfaces. Precise
voltammetric analysis show that the initial and rate-determining reaction involves one electron transfer and one
molecule borate:
FeS2 / B(OH) 40 c FeS2rrr[B(OH)4 ]ads / e.
This reaction is completely irreversible, with a Tafel slope
of 240 mV/decade and the charge transfer coefficient of
0.25. At relatively high potentials, the reaction is controlled
by the mass-transfer of borate. The surface hydrophobicity of
pyrite is strongly affected by the presence of borate. Solution
chemistry calculations for the ferric-borate systems indicate
that a number of highly stable ferric-borate complexes can
be formed under the studied conditions. At high borate concentrations, ferric hydroxide precipitate will be dissolved
below pH 9.
ACKNOWLEDGMENTS
The author thanks Dr. Elisabet Ahlberg (Chalmers University of Technology, Sweden) and Dr. Ron Woods (CSIRO, Australia) for their very constructive discussions on this paper. The author also appreciates the help of
the highly competent CAER researchers and staff in conducting the research, in particular to Mr. C. L. Jiang (Department of Mining Engineering,
University of Kentucky), for their assistance with the experiments. This
work was sponsored by the U.S. Department of Energy (DE-FG2290PC90295).
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