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Characterization of surface layers formed during pyrite oxidation
COLLOIDS AND SURFACES
A
Colloids and Surfaces
A: Physicochemical
Characterization
and Engineering
Aspects 93
( 1994)
201-210
of surface layers formed during pyrite oxidation Ximeng Zhu, Jun Li, Milton E. Wadsworth *
Department
of Metallurgical
Engineering,
College of Mines and Earth Sciences,University UT841
It
of Utah, Salt Lake City,
USA
Received 8 September 1993; accepted 30 April 1994
Abstract Surface electrochemical reactions of pyrite have been studied using cyclic voltammetry. in situ laser Raman spectroscopy and potentiostatic measurements. A surface sulfur layer was identified on pyrite surfaces during transpassive oxidation. Sulfur layer formation at high anodic potentials was confirmed by subsequent surface reactions during the cathodic and return anodic cycle in cyclic voltammetry and by in situ Raman spectroscopy. Initially, polysulfides (FeS, of variable X) form as intermediates during anodic oxidation in both acidic and basic solutions. The nature of these intermediates dcpcnds upon pH and the rate of surface film growth. The end product of the film growth is an active form of sulfur, S:,,. which is readily reducible. Raman spectra obtained during an extended period of time indicated the film was mainly St,,. The conditions such as applied potentials and pH that lead lo the build-up of the surface oxidation layer have been determined. Film growth kinetics obey the paralinear rate law. indicating uniform film growth of the sulfur-rich layer to a steady state thickness. The thickness is controlled by the simultaneous rate of dilTusion of cations through the polysulfide layer, with parabolic rate constant k,. and the rate of oxidation of the active-sulfur outer layer, with the linear rate constant k,. to soluble sulfur species in solution. The elTecls of temperature. applied potential and pH on k, and k, were determined using the paralincar rate law. Keywords:
Electrochemistry; Kinetics; Pyrite; Raman spectroscopy; Sulfur
1. Introduction Pyrite is the most abundant metal sulfide in nature. In many cases, the occurrence of pyrite is undesirable. For example, the coexistence of pyrite with other non-ferrous metal sulfide minerals affects the recovery of those valuable metals and its association with coal leads to the production of sulfur dioxide and acid rain. In metallurgical practices, flotation is by far the most widely used method for pyrite separation. Under appropriate
conditions sulfide minerals may be readily floated. Many studies attribute this to the formation of a hydrophobic sulfur-rich layer on the surface of the sulfide minerals [ 11. With potentiodynamic techniques, Hamilton and Woods [2] suggested that the electrochemical oxidation of pyrite results in the formation of elemental sulfur and sulfate ions according to the reactions FeS,+3H,O
= Fc(OH),+2S”+3H++3e-
(1)
FeS,+I1H,O=Fe(OH),+S0~-+19H++15el
Correspondingauthor.
0927-7757;94607.00
Q 1994 Elsevier Science R.V. All rights reserved
SSDI 0927-7757(94)02935-L
(2)
201
X Zhu et al./Culloicis Sur/uws A: Ph_vsicochrm.Eng. Aspects 93 ( 1994) 201-210
Similar results were reported by Ahlberg ef al. [ 11. Janetski et al. [3]. Franklin et al. [4], Chander and Briceno [S] and Chmielewski and Wheelock
CO Many ex situ and in situ studies were carried out to confirm the presence of sulfur on the sulfide surface. Allison and Finkelstein dissolved the mineral surface using carbon disulfide and detected the presence of sulfur using calorimetric techniques [7]. Other techniques that have been used to detect sulfur films successfully on surfaces include X-ray photoelectron spectroscopy (XPS) [S-lo] and Raman spectroscopy [ lO,ll]. The formation of a surface polysultide film has been proposed by Zhu and co-workers [ 12-151, based on cyclic voltammetry and kinetics of film growth. The proposed reactions are sFe!$ +(3.u-6)H,O
= 2 FeS +(.u-2)Fe(OH), +(3s-6)H+
+(3.x-6)e(3)
FeS, +(4.u+3)H,O
= Fe(OHb + xSO:+(8s+3)H+ +(6s+3)c(4)
FcS, + 3H,O = Fc(OH ), + .uS”+ 3H + + 3c-
(5)
with SOi- being the predominant final sulfur product. In situ Raman spectroscopic cxpcrimenls on the elcctrochcmically oxidized pyrite confirmed the presence of clcmcntal sulfur [ 10,l I], polysullidcs [ 10.1I]. sulfate ions [I l] and ferric hydroxide [II].
The present rcscarch concerns the formation and characterization of the surface layer formed during pyrite oxidation. Electrochemistry and Raman spectroscopic techniques were used in this investigation. The conditions for the film formation, the kinetics of the film growth, and the properties of the film were studied. Included in this investigation were various inllucncing factors such as applied potentials. electrode pretrcatmcnt. electrode rotation sprcd. solution pH. and temperature.
2. Esperimcntal
The cxpcrimontal set-up and proccdurcs have been dcscribcd previously [ 133. The working
electrodes, both stationary and rotating disk, were mineral pyrite obtained from Ward’s National Science Establishment. A 12 cm* platinum sheet was used as the counterelectrode and a saturated calomel electrode (SCE) was used as the reference electrode. The electrolyte used was 0.01 M sodium borate in 0.49 M sodium sulfate solution (the resulting pH was 9.2) for electrochemical measurements and 0.5 M sodium chloride for Raman spectroscopic measurements. Higher or lower pH values were obtained by addition of NaOH or HCl solutions. Cyclic voltammetry measurements were conducted at room temperature in a threecompartment Pyrex cell. The scans. unless otherwise specified, were initiated from the most negative potential going positively to the upper potential and then reversed. Data were recorded and plotted by a PAR Model 273 potentiostat/ galvanostat programmed with an IBM PC-XT computer. Potentials were recorded and reported here versus SCE. The spectrometer (Spex Tripmate 1877) consists of a triple monochromater and a 2D CCD array (Photometrics) of 576 by 354 pixels. The laser lint was from an argon ion laser (Spectra Physics) at 514.5 nm, and the incident power was about 200 mW. The spcctroclcctrochcmical cell had a quartz window. It was arranged such that the incident Iascr light was at a grazing angle of approximately 70” to the sample. The laser beam was focused to a spot approximately 150 urn in diamctcr. The electrochemical pretreatment flotation experiments were carried out in a modified Hallimond tube. A platinum plate electrode was used as the working electrode to contact electrically a bed of particulate pyrite. A saturated calomel reference electrode was placed in a separated compartment to measure the potential near the top of the pyrite bed by means of a Luggin capillary tube.
3. Formation of surface sulfur layer during pyrite oxidation
The cyclic voltammograms of pyrite in stirred and quiescent solutions of pH 9.2 using a stationary elcctrodc are shown in Fig. 1. All the peaks but one (peak la) in the voltammograms have been carefully analyzed and assigned to the appro-
X Zhu et al./Colloids
6 -1.00
-0.50
I
I. 0.50
0.00
1.00
Fe(OH)2 + HS- + H+ $ FeS+2H,O
E (V, vs SCE) Fig. I. Cyclic
voltammograms
of pyrite
(pH 9.2,
scan
rate
100 mV s-l).
priate reactions in most of the published literature. Peaks 1 and 6 were related and the corresponding reaction was identified as the conversion between Fe0 and Fe(OH),. Peaks 2 and 5 were conjugate and weakened significantly if peak I did not appear. They represented the transition of ferrous and ferric ions/hydroxides. Peak 3 was the direct oxidation of pyrite in the transpassive region. The passivation of peak 3 under quiescent conditions resulted from the build-up of a loosely held layer of ferric hydroxide on the surface. This film was readily disrupted by the action of stirring, evidenced by the appearance of colloidal ferric hydroxide throughout the electrolyte under conditions of strong agitation. The reactions at different peak positions are listed in Table I. Anodic peak la at about -0.67 V was new and appeared only when certain conditions were met. The conditions for peak la to appear in the cyclic voltammogram were (a) the upper reversal potenTable
Peak
-1.50
-1.00
I and 6
Fe(OH),+ZH+
Fc(OH ), + H + +e-
3
xFeSL +(3x-6)H,O=2FcS, FcS, +(4,+3)H,O=Fc(OH), FcS,+3H,O=
+2e-
Fig. 2. Cyclic
voltammograms
rcvcrse potentials
lrotating
rate lo0 mV s-l).
scans [9.10]
= Fc+ZH,O = Fc(OtI
Fc(Ofl),
): + Hz0 +(3x-6)11’ +(3x-6)e+(.x-- ‘)Fc(OH), +xSO:+(8x+3)1-l +(6s+3)e-
+ sS+3tIc
+3e-
-0.50
0.00
0.50
I.00
E (V. vs SCE)
Reaction
2 and 5
(6)
During the next positive-going voltage scan, the FeS was oxidized to Fe& at about -0.67 V, characterized by peak la. The magnitude of peak la correlates directly with the quantity of sulfur intermediates formed in the transpassive region with
I
Pyrite reactions during cyclic voltammogram
203
201-210
tial exceeded a certain value (Fig. 2); (b) there was enough conditioning time at the potentials above the threshold value (Fig. 3); and (c) the scan passed the threshold potential at least once before going through the potential at peak la (Fig. 4). The products (partially oxidized sulfur intermediates and iron hydroxide) formed in the transpassive region (peak 3) were reduced to Fe(OH), and HS- during the negative-going voltage scan. Fe(OH)t and HS- produced on the negative-going scan would react to form FeS according to the reaction [ 14.IS]
Quiescent
s
I.. -1.50
Surfaces A: Physicochem. Eng. Aspects 93 (1994)
of pyrite
at h&rent
speed 300 rev min-‘.
anodic
pH 9.2. scan
34
X. Zhu et ul./Colloids
IO.OOl6
la
12
I
A&
3 set ,
t....,.. -1.50
Surfuces A: Ph_wicochem. Eng. Aspects 93 ( 1994) 201-210
.1
-0.50
-1.00
1
.‘.‘I
,..‘I
0.00
0.50
E (V, vs SCE) Fig. 3. Cyclic voltammoprams for various times (rotating
of pyrite preconditioned
speed 200 rev min-‘.
at 0.6 V
pH 9.1. scan
rute 50 mV s-‘1.
-1.50
-1.00
-ll.so
0.00
0.50
with peak la increased with increasing electrode precondition time and tended to level off after 30 s. probably indicating saturation of the surface. The maximum charge associated with peak la is about 10 mC, corresponding to about 27 monolayers in thickness [IS]. As mentioned above, peak la was apparent in voltammograms only for scans reversed at an anodic potential higher than a threshold value. This value was 0.6 V for solutions of pH 9.2. A series of voltammograms at various upper potentials in solutions of pH 9.2 was obtained at 50 mV s-l, with the rotating speed of the disk electrode being 200, 700 and 1000 rev mine1 respectively. As the upper potential increased from 0.65 to 0.85 V. the intensity of peak la increased. Further increase of the anodic reversal potential resulted in the decrease in the intensity of peak la. Peak la disappeared for anodic reversal potentials above 1.1 V. This result clearly indicates that, at potentials higher than certain limits, the partially oxidizcd sulfur intermediates oxidized rapidly to sulfate ions. The integrated charge associated with peak la vs. the anodic reversal potential for various clcctrodc rotation speeds is shown in Fig. 5. For each electrode rotation speed, the intensity of peak la increased with increasing anodic reversal potential and rcachcd a maximum WIUC iit about 0.85 -0.9 V. The mcasuromcnts obtained from
I .w
E (V. vs SCE) Fig. 4. Cyclic
voltammograms
of
pyrik
in
various
l
1ooorpnl
0
7OU rpm
I
P?
cycles
(stirred solution, pH 9.2, scim rate IO0 mV se’).
the reaction for peak la being 2FeS+2H20=Fe(OH),_
+ FcS,+2H+
+2e(7)
All the factors that all’cct the transpassivc oxidation of pyrite influence peak la. The faster the clcctrodc rotated. the grcatcr the amount of substances participated in reaction (7). When the clcctrodc was preconditioned at a constant potential for increasing time periods, peak la was continuously cnhanccd. The intcgratcd charge associated
0.4
0.6
0.X
1.0
I.2
I.4
Anodic rcvcrsc potential (V) Fig. 5. Effect of nnodic rcvwsc pokntial scan r:Itc 50 mV s
’ J.
on pc;tk la (pH 9.2.
X. Zhu et aLfColloi& Surfuces A: Physicochem. Eng. Aspects 93 (1994) JOI-210
Raman spectroscopic experiments gave identical results [ 111. The influence of pH on the transpassive oxidation of pyrite was studied with both cyclic voltammetry and Raman spectroscopy. The pH value affected the minimum potential for the appearance either of peak la in the voltammograms or of the sulfur peaks in the Raman spectra. In acid solutions (pH 2.6) there was no peak la in the voltammograms as FeS is stable at low potentials only in the mid-pH range [ 151, although sulfur species could still be detected by Raman spectroscopy at high potentials. In strongly alkaline solutions (pH 12.5), the sulfur species including pyrite did not appear in the Raman spectra. The conditions for sulfur detection by the Raman spectroscopic technique are demonstrated in Fig. 6. It is not surprising that the Raman peaks of sulfur and/or polysulfides did not appear in the strongly alkaline solutions as these species were not stable. However, the overpotentials for sulfur detection seemed quite high in acid, neutral and weakly alkaline solutions. Kinetics were perhaps the main reason for the overpotentials. During the in situ Raman expcrimcnts, the maximum time for the potential application was normally 2 h. Another related factor was the detection limit of the Raman spectrometer set-up. Theoretically Raman spec-
20s
trometry can do monolayer measurements. For the set-up in these experiments. the best result obtained was about 90 monolayers [ 163.
4. Kinetics of surface sulfur layer growth In the transpassive region. aggressive electrochemical oxidation of pyrite occurs. The reaction products in this region are ferric hydroxide, elemental sulfur, sulfate ions and partially oxidized sulfur intermediates, with the majority of the sulfur product being sulfate ions. Sulfur products other than sulfate ion are metastable only. While exposure of pyrite to anodic potentials higher than the transpassive region results in rapid oxidation of sulfur intermediates to sulfate ions, the sulfur intermediates convert to soluble sulfoxy and ultimately sulfate ions at a measurable rate in the transpassive region. A surface layer on top of the oxidized pyrite is formed, based on the dynamic equilibrium between the inner layer growth of sulfur intermediates due to pyrite oxidation and the outer layer diminishmcnt due to its conversion to soluble sulfate ions. A schematic diagram is dcpictcd in Fig. 7 to show the growth mechanism of this sulfur layer. The surface sulfur layer has an iron-deficient structure or consists of a mixture of polysulfidcs FeS,, where ,x> 2. The iron concentration in the layer dccrcascs continuously from the inner to the outer surface while the sulfur concentration increases. Sulfur species may vary from Sg- (pyrite) on the inner W,
0
-1.0 -
Cl9
I
-1.5
.
0
’ . ’ 2 4
’ 6
.
I
a ’ 8
.
1
-w
so:
--_)
Fc(OlI),
F4
’ . ’ IO 12
14
PH Fig. 6. Conditions with
Raman
for sulfur dctcction from oxidation
spectroscopy:
(E[S]-~[Fc]=O.OOl
0.
unobscrvcd;
M in the E,-pH
diagram).
0,
of pyrite
I
observed Fig. 7. Schcmalic diagram
II of surface sulfur layer.
206
X. Zhu et aL/Colloids Surfuces A: Physicochem.
surface to S& (electrochemically reducible elemental sulfur) on the outer surface. Since the outer surface is saturated with sulfur, it appears that the dissolution reactions follow the linear kinetics. Iron lattice ions migrate through vacancies by diffusion, characterized by parabolic kinetics. The kinetics of surface layer growth as pyrite was oxidized in the transpassive region were studied using the potentiostatic technique. Fig. 8 shows the current decay curves obtained at potentials between 0.4 and 0.65 V. For applied potentials higher than 0.65 V, oxygen evolution started. and the current density increased irregularly. During the initial stage, currents decreased quickly, indicating the build-up of a passivation layer. As the layer increased in thickness, the rate of pyrite oxidation decreased and approached the dissolution rate of the outer surface layer. The thickness of the passive layer thus became essentially constant, leading to the steady movement of the whole surface layer toward pyrite. A paralinear rate equation may be derived to describe the kinetics of the transpassive oxidation of pyrite [ 14.153:
where Q is the accumulated charge passed on the electrode, I the time, k, the parabolic rate constant and k, the linear rate constant. Fig. 9 demonstrates ...
I..“,.
‘.I.-.
I 2 3 4
4.0
0.65 V 0.64lv 0.50 v 0.40 v
3.0
0.0
‘. 0
‘. loo0
‘. 2000
’ 3MMl
4MKl
Time (see) Fig. 8. Current
decay
curves of pyrite
(stirred solution, pH 9.2).
at variaus
potentials
Eng. Aspects 93 ( 1994) 201-110
:
-
0
Paralinear Equation
IO00
zoo0
3ooo
4ooo
Time (set) Fig. 9. Correlation equation
of current
decay
curves
with
paralinear
(stirred solution, pH 9.2).
the good fit of the experimental data to the paralinear rate equation. Thus the kinetics results are consistent with the proposed mechanism of surface layer growth: the formation of an intermediate passive layer during pyrite oxidation and the simultancous dissolution of the outer surface layer resulting from sulfur and polysulhdes oxidation to sulfate ions. By fitting the experimental data to Eq. (8). the parabolic and linear rate constants for various applied potentials were obtained. both k, and k, incrcascd with increasing applied potential. Consistent with results from the cyclic voltammetric expcrimcnts, increased anodic oxidation potcntial increased both the amount of sulfur intermediates formed and the amount of intermediates dissolved. The effect of agitation on the reaction rate was examined at dilfcrcnt applied potentials and solution pH values using rotating disk electrodes. For each cast, the initial current increased significantly as the electrode rotation speed jumped from 0 to 100 rev min-‘. However, further increase in the electrode rotation speed up to 7OOrcv min-’ increased the initial current only slightly while the steady state current remained almost constant even when the elcctrodc was stationary. The independcncc of the surface layer movcmcnt on agitation confirmed that the outer layer dissolution was limited by the surface reaction.
X. Zhu et al. JCoIloti
Sur-_aces A: Physicochem. @g. Aspects 93 (1994)
Both the parabolic rate constant k, and the linear rate constant k, increased with increasing solution pH as expected from reactions ( l)-( 5) although the pH influence was much more profound on k,, than on k,. The strong dependence of k, upon pH could be explained on the basis of the following assumption: the surface Fe’+ (lattice) concentration decreased with increasing pH. thus increasing the diffusion gradient and the rate of Fe’+ (lattice) diffusion through the passivation layer. The kinetics of the surface sulfur layer growth was studied in the temperature range 23.4-59S’C. The activation energies were found to be 66.17 kJ mol-r (15.83 kcal mol-‘) for the parabolic rate process and 38.67 kJ mol-’ (9.25 kcal mol- ‘) for the linear rate process. The increase in temperature accelerated both the formation and the dissolution of the surface layer. However, the thickness of the layer will increase with temperature since k, increases more rapidly than k, as temperature goes up. To quantify the influences of temperature, applied potentials and electrolyte pH on k, and k,, a general equation may bc dcvcloped: li=oexp(
-$)exp(;T
JOI-210
207
1
10.0 -
’ -
0
-.-I’.-.-
Panlinur Equation
1000
2000
1
3ooo
4ooa
Time (see) Fig. 10. Comparison estimated
of experimental
rate constants obtained
and calculated data using
from Eq. (9).
lPph
where k could bc k, or k,, E,, is the activation energy, and a, b and c are constants. E,, u, b and c were estimated from k, and k, obtained under different conditions with non-linear curve fitting. The parabolic and linear rate constants k, and k, at given temperature, applied potential and electrolyte pH can be calculated according to Eq. (9). Substituting these k, and k, values into paralinear Eq. (8). a series of charge versus time data was obtained. Fig. 10 illustrates the satisfactory correlation of calculated charge-time data using k, and k, estimated from Eq. (9) with the paralinear equation for various temperatures. Similar results were also obtained for various electrolyte pH values and applied potentials.
5. Properties of surface sulfur layer
Fig. 11 (curve a) shows the Raman spectrum of a sample of pyrite treated potentiostatically at + 1.0 V for 10 min. The sulfur/polysuhidc Raman
I-...@
100
I-...1
200
300
1.
400
500
I
..I
600
700
Frcqucncy shift (cm.‘) Fig. 11. Raman
spectra
of
pyrite
in 0.5 M
NaCl
solution
(pH 9.2): curve a. 10 min at 1.0 V; curve b, 3 min at - I.0 V; curve c. 2 min at + 1.0 V.
peaks can be seen clearly from the figure. The potential was then immediately switched to - 1.0 V. The Raman spectrum in Fig. I1 (curve b) was taken after 3 min at this potential. The sulfur/ polysulfide peaks that appeared at + 1.0 V now disappeared and the spectrum was almost identical to that at the beginning of the experiment. This suggests that cathodic polarization reduces the pyrite surface to a state that is spectroscopically identical to the freshly ground condition. At the end of the experiment the applied potential was
208
X. Zhu et aL/Colloids
Surjices
A: Physicochem. Eng. Aspects 93 ( 1994) 201-210
reversed again to + 1.0 V. Fig. 11 (curve c) demonstrates that the sulfur/polysulfide peaks reappeared after 2 min of oxidation. These results indicate that the sulfur layer formed during pyrite oxidation is readily reducible (S&J. Sulfur produced by electrochemical deposition from sulfide solutions at a platinum electrode was also reducible [ 163. The fact that the surface sulfur layer formed during pyrite oxidation could be reduced electrochemically implies that the sulfur thus formed does not behave as an electrical insulator as does bulk elemental sulfur in its orthorhombic state. The ability of the material to support a photocurrent [ 163 also suggests that the sulfurous film has unusual electronic properties more akin to a semiconductor than to an insulator. The formation of a hydrophobic layer on surfaces of metal sulfide minerals is deemed important to flotation by many researchers [ 11. By nature, the sulfur layer formed on the oxidized pyrite surface has such properties. The interactions between the surface sulfur layer and the flotation agent potassium amyl xanthate wcrc probed by Raman spectroscopy in this study. Raman spectra of sulfur and potassium amyl xanthate taken in air arc given in Figs. 12 and 13 (spectra a and b) as rcfcrcnccs. To increase sensitivity, the surface enhanced Raman spectroscopy (SERS) tcchniquc was employed for pyrite
loo
200
300
400
500
600
700
Frcqucncy shift (cm”) Fig. 12. Raman
spectra of sulfur in air (curve a). xanthatc
in
air (curve b). silver-plated
pyrite in xanthatc
solution (curve c),
oxidized and silver-plated
pyrite in xanthatc solution (curve d).
700
800
900
IO00
1100
1200
Frequency shift (cm”) Fig. 13. Raman
spectra of sulfur in air (curve a). xanthate
in
air (curve b). silver-plated
pyrite in xanthate solution (curve c),
oxidized and silver-plated
pyrite in xanthate solution (curve d).
electrodes. The pyrite electrode, whether in its original form or after being oxidized, was put into the solution containing 0.001 M AgNO, amd 0.5 M KNO, and polarized at -0.60 V for the passage of approximately 5 mC of charge. The silver-plated pyrite elcctrodc was then trnnsfcrrcd to a solution of pH 6.5 containing 0.5 M NaCl and a small amount of potassium amyl xanthate and its Raman spectrum recorded. Figs. 12 and 13 prcscnt the Raman spectra of the pyrite electrode: pyrite unoxidized in spectrum c and pyrite oxidized at 1.0 V for 1 h in spectrum d. A close look revealed that there wcrc several new peaks in the spectrum of the oxidized pyrite electrode. The peaks at about 170, 640 and 1IOOcm-’ were not found in the spectra of either unoxidized pyrite in the presence of amyl xanthate or oxidized pyrite in the absence of amyl xanthate. There are several possible interactions between amyl xanthate and the oxidized pyrite. Physical adsorption and chemical reaction are two cases. If it is the former process then the peak position will not change much (e.g. the peaks at 170 and 1100 cm-‘). However, the latter process could lead to the formation of new species and therefore new peak positions (e.g. the peak at 640cm-‘). Since them arc no new peaks found when unoxidized pyrite elcctrodc was put into a solution of 0.5 M NaCl and pH 6.5 with the addition of a small
X Zhu t-t aL/Colloidr Surfaces A: Phpicochem Eng, Aspects 93 (1994) 201-210
amount of amyl xanthate, it is reasonable to assume that sulfur or polysulfides play an important role in the flotation process. More work, however, needs to be done to identify the newly formed or adsorbed species. The formation of sulfur and polysulfides promotes xanthate adsorption on pyrite and so could enhance its floatability. Fig. 14 illustrates the influence of applied potentials on floatability of coal pyrite. Sodium sulfate and sodium borate were added to the solution as supporting electrolytes. The floatability increased first with the increase of the applied potential up to a maximum value and then decreased with the potential increase. The bell-shaped curve agreed completely with the results of sulfur layer formation obtained with the electrochemical and Raman spectroscopic measurements.
6. Conclusions
Aggressive oxidation of pyrite occurred in the transpnssivc region. The reaction products in this region wcrc Fc(lI1) oxides. sulfate ions, an active form of elcmcntnl sulfur and partially oxidized sulfur intcrmcdiatcs. The surface layer formation at high anodic potentials was confirmed by subscqucnt surbcc reactions during the cathodic and
“g “. l
./
/
\ Potcnfial (V, vs SCE)
Fig. I-8. Eflect of potenkd on tlo&bility of cd pyrik (0.125 M Na,S04, 0.005 M N:I&O,. pH 7.0; pxticlc six. - 100 -325 mesh).
209
return anodic cycle in cyclic voltammetry and by in situ Raman spectroscopy. Polysulfides (FeS, of variable X) form as intermediates during anodic oxidation. The iron concentration in the surface layer decreases continuously from the inner to the outer surface while the sulfur concentration increases. Sulfur species may vary from S:- (pyrite) on the inner surface to So (elemental sulfur) on the outer surface. The conditions such as applied potentials and pH that lead to the build-up of the surface oxidation layer have been determined. The influence of applied potentials on the quantity of sulfur formed can be represented by a bell-shaped curve. The sulfur formed increased with the applied potential up to a maximum value after which less sulfur was produced because of the rapid formation of sulfate ions. Sulfur and its intermediates were detected in both acidic and basic solutions (pH 2.6-10.5). Film growth kinetics obey the paralinear rate law indicating uniform film growth of the sulfurrich layer to a steady state thickness. The thickness is controlled by the simultaneous rate of diffusion of cations through the polysullidc layer, with parabolic rate constant k,, and the rate of oxidation of the active sulfur outer layer, with the linear rate constant k,, to soluble sulfur spccics in solution. The cfTccts of tcmpcrature, applied potential and pH on k, and k, wcrc dctcrmincd using the paralinear rate law. The activation encrgics were found to be 66.17 kJ mot-’ (15.83 kcal mol-‘) for the rate process and 38.67 kJ mot-’ parabolic (9.25 kcal mol-‘) for the linear process. The end product of the lilm growth is an active form of sulfur, SI1,,. which is readily reducible. This indicates that the sulfur formed during pyrite oxidation dots not behave as an electrical insulator as does bulk elemental sulfur in its orthorhombic state. The surface sulfur layer was hydrophobic in nature and its interaction with the flotation agent such as potassium amyl xanthate enhanced the floatability of pyrite as evidenced by Raman spectroscopy and flotation experiments.
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X Zhu et af./CoNoi&
[ 21 I.C.
Surfaces A: Physicochem. Eng. Aspects 93 (1994)
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( 198I ) 327. [3] N.D. Janetski S.I. Woodbum and R. Woods, Int. J. Mineral Process .. 4 ( 1977) 7?7. _[4] T.C. Franklin, R. Nodimele and W.K. Andeniyi, J. Electrochem. Sot.. 134 ( 1987) 2150. [S] S. Chander and A. Briceno, Minerals Metall. Process.. 8 (1987) 171. [6] T. Chmielewski and T.D. Wheelock, J. Min. Metall., 26 (1990) 133. [‘I] S.A. Allison and N.P. Finkelstein. NIM report No. 1495. Johannesburg. 1972. [8] A.N. Buckley, I.C. Hamilton and R. Woods. J. Electroanal. Chem., 216 (1987) 213. [9] A.N. Buckley and R. Woods, Appl. Surf. Sci., 27 ( 1987) 437. [IO] J.R. Mycroft. GM. Bancroft. N.S. McIntyre, J.W. Lorimer and I.R. Hill. J. Electroannl. Chem.. 292 ( 1990) 139. [II] I. Li and M.E. Wadsworth. in J.B. Hiskey and G.W. Warren (Eds.). Fundamentals, Hydrometallurpy:
MI-II0
Technology, and Innovation. Society for Mining, Metallurgy, and Exploration. Inc., Littleton. CO. 1993. p. 127. [ 121 X. Zhu. M.E. Wadsworth. D.M. Bodily and A.M. Riley, in P.R. Dugan. D.R. Quigley and Y.A. Attia (Eds.), Processing and Utilization of High-Sulfur Coals, Vol. IV, Elsevier, Amsterdam. 1991. p. 205. [ 131 X. Zhu, J. Li, D.M. Bodily and M.E. Wadsworth, J. Electrochem. Sot.. 140 (1993) 1927. [ 141 X. Zhu, J. Li and M.E. Wadsworth. in J.P. Hager (Ed.), EPD Congress 1993. The Minerals. Metals % Materials Society. Warrendale. PA. 1993. p. 355. [ 151 M.E. Wadsworth, X. Zhu and J. Li, in J.B. Hiskey and G.W. Warren (Eds.). Hydrometallurgy: Fundamentals, Technology. and Innovation, Society for Mining, Metallurgy. and Exploration, Inc.. Littleton. CO. 1993. p. 85. [ 161 M.E. Wadsworth. Surface Electrochemical Control for Fine Coal and Pyrite Separation. Final Report to DOE (Project No. DE-ACZZ-89PC89758), 1993.
A
Colloids and Surfaces
A: Physicochemical
Characterization
and Engineering
Aspects 93
( 1994)
201-210
of surface layers formed during pyrite oxidation Ximeng Zhu, Jun Li, Milton E. Wadsworth *
Department
of Metallurgical
Engineering,
College of Mines and Earth Sciences,University UT841
It
of Utah, Salt Lake City,
USA
Received 8 September 1993; accepted 30 April 1994
Abstract Surface electrochemical reactions of pyrite have been studied using cyclic voltammetry. in situ laser Raman spectroscopy and potentiostatic measurements. A surface sulfur layer was identified on pyrite surfaces during transpassive oxidation. Sulfur layer formation at high anodic potentials was confirmed by subsequent surface reactions during the cathodic and return anodic cycle in cyclic voltammetry and by in situ Raman spectroscopy. Initially, polysulfides (FeS, of variable X) form as intermediates during anodic oxidation in both acidic and basic solutions. The nature of these intermediates dcpcnds upon pH and the rate of surface film growth. The end product of the film growth is an active form of sulfur, S:,,. which is readily reducible. Raman spectra obtained during an extended period of time indicated the film was mainly St,,. The conditions such as applied potentials and pH that lead lo the build-up of the surface oxidation layer have been determined. Film growth kinetics obey the paralinear rate law. indicating uniform film growth of the sulfur-rich layer to a steady state thickness. The thickness is controlled by the simultaneous rate of dilTusion of cations through the polysulfide layer, with parabolic rate constant k,. and the rate of oxidation of the active-sulfur outer layer, with the linear rate constant k,. to soluble sulfur species in solution. The elTecls of temperature. applied potential and pH on k, and k, were determined using the paralincar rate law. Keywords:
Electrochemistry; Kinetics; Pyrite; Raman spectroscopy; Sulfur
1. Introduction Pyrite is the most abundant metal sulfide in nature. In many cases, the occurrence of pyrite is undesirable. For example, the coexistence of pyrite with other non-ferrous metal sulfide minerals affects the recovery of those valuable metals and its association with coal leads to the production of sulfur dioxide and acid rain. In metallurgical practices, flotation is by far the most widely used method for pyrite separation. Under appropriate
conditions sulfide minerals may be readily floated. Many studies attribute this to the formation of a hydrophobic sulfur-rich layer on the surface of the sulfide minerals [ 11. With potentiodynamic techniques, Hamilton and Woods [2] suggested that the electrochemical oxidation of pyrite results in the formation of elemental sulfur and sulfate ions according to the reactions FeS,+3H,O
= Fc(OH),+2S”+3H++3e-
(1)
FeS,+I1H,O=Fe(OH),+S0~-+19H++15el
Correspondingauthor.
0927-7757;94607.00
Q 1994 Elsevier Science R.V. All rights reserved
SSDI 0927-7757(94)02935-L
(2)
201
X Zhu et al./Culloicis Sur/uws A: Ph_vsicochrm.Eng. Aspects 93 ( 1994) 201-210
Similar results were reported by Ahlberg ef al. [ 11. Janetski et al. [3]. Franklin et al. [4], Chander and Briceno [S] and Chmielewski and Wheelock
CO Many ex situ and in situ studies were carried out to confirm the presence of sulfur on the sulfide surface. Allison and Finkelstein dissolved the mineral surface using carbon disulfide and detected the presence of sulfur using calorimetric techniques [7]. Other techniques that have been used to detect sulfur films successfully on surfaces include X-ray photoelectron spectroscopy (XPS) [S-lo] and Raman spectroscopy [ lO,ll]. The formation of a surface polysultide film has been proposed by Zhu and co-workers [ 12-151, based on cyclic voltammetry and kinetics of film growth. The proposed reactions are sFe!$ +(3.u-6)H,O
= 2 FeS +(.u-2)Fe(OH), +(3s-6)H+
+(3.x-6)e(3)
FeS, +(4.u+3)H,O
= Fe(OHb + xSO:+(8s+3)H+ +(6s+3)c(4)
FcS, + 3H,O = Fc(OH ), + .uS”+ 3H + + 3c-
(5)
with SOi- being the predominant final sulfur product. In situ Raman spectroscopic cxpcrimenls on the elcctrochcmically oxidized pyrite confirmed the presence of clcmcntal sulfur [ 10,l I], polysullidcs [ 10.1I]. sulfate ions [I l] and ferric hydroxide [II].
The present rcscarch concerns the formation and characterization of the surface layer formed during pyrite oxidation. Electrochemistry and Raman spectroscopic techniques were used in this investigation. The conditions for the film formation, the kinetics of the film growth, and the properties of the film were studied. Included in this investigation were various inllucncing factors such as applied potentials. electrode pretrcatmcnt. electrode rotation sprcd. solution pH. and temperature.
2. Esperimcntal
The cxpcrimontal set-up and proccdurcs have been dcscribcd previously [ 133. The working
electrodes, both stationary and rotating disk, were mineral pyrite obtained from Ward’s National Science Establishment. A 12 cm* platinum sheet was used as the counterelectrode and a saturated calomel electrode (SCE) was used as the reference electrode. The electrolyte used was 0.01 M sodium borate in 0.49 M sodium sulfate solution (the resulting pH was 9.2) for electrochemical measurements and 0.5 M sodium chloride for Raman spectroscopic measurements. Higher or lower pH values were obtained by addition of NaOH or HCl solutions. Cyclic voltammetry measurements were conducted at room temperature in a threecompartment Pyrex cell. The scans. unless otherwise specified, were initiated from the most negative potential going positively to the upper potential and then reversed. Data were recorded and plotted by a PAR Model 273 potentiostat/ galvanostat programmed with an IBM PC-XT computer. Potentials were recorded and reported here versus SCE. The spectrometer (Spex Tripmate 1877) consists of a triple monochromater and a 2D CCD array (Photometrics) of 576 by 354 pixels. The laser lint was from an argon ion laser (Spectra Physics) at 514.5 nm, and the incident power was about 200 mW. The spcctroclcctrochcmical cell had a quartz window. It was arranged such that the incident Iascr light was at a grazing angle of approximately 70” to the sample. The laser beam was focused to a spot approximately 150 urn in diamctcr. The electrochemical pretreatment flotation experiments were carried out in a modified Hallimond tube. A platinum plate electrode was used as the working electrode to contact electrically a bed of particulate pyrite. A saturated calomel reference electrode was placed in a separated compartment to measure the potential near the top of the pyrite bed by means of a Luggin capillary tube.
3. Formation of surface sulfur layer during pyrite oxidation
The cyclic voltammograms of pyrite in stirred and quiescent solutions of pH 9.2 using a stationary elcctrodc are shown in Fig. 1. All the peaks but one (peak la) in the voltammograms have been carefully analyzed and assigned to the appro-
X Zhu et al./Colloids
6 -1.00
-0.50
I
I. 0.50
0.00
1.00
Fe(OH)2 + HS- + H+ $ FeS+2H,O
E (V, vs SCE) Fig. I. Cyclic
voltammograms
of pyrite
(pH 9.2,
scan
rate
100 mV s-l).
priate reactions in most of the published literature. Peaks 1 and 6 were related and the corresponding reaction was identified as the conversion between Fe0 and Fe(OH),. Peaks 2 and 5 were conjugate and weakened significantly if peak I did not appear. They represented the transition of ferrous and ferric ions/hydroxides. Peak 3 was the direct oxidation of pyrite in the transpassive region. The passivation of peak 3 under quiescent conditions resulted from the build-up of a loosely held layer of ferric hydroxide on the surface. This film was readily disrupted by the action of stirring, evidenced by the appearance of colloidal ferric hydroxide throughout the electrolyte under conditions of strong agitation. The reactions at different peak positions are listed in Table I. Anodic peak la at about -0.67 V was new and appeared only when certain conditions were met. The conditions for peak la to appear in the cyclic voltammogram were (a) the upper reversal potenTable
Peak
-1.50
-1.00
I and 6
Fe(OH),+ZH+
Fc(OH ), + H + +e-
3
xFeSL +(3x-6)H,O=2FcS, FcS, +(4,+3)H,O=Fc(OH), FcS,+3H,O=
+2e-
Fig. 2. Cyclic
voltammograms
rcvcrse potentials
lrotating
rate lo0 mV s-l).
scans [9.10]
= Fc+ZH,O = Fc(OtI
Fc(Ofl),
): + Hz0 +(3x-6)11’ +(3x-6)e+(.x-- ‘)Fc(OH), +xSO:+(8x+3)1-l +(6s+3)e-
+ sS+3tIc
+3e-
-0.50
0.00
0.50
I.00
E (V. vs SCE)
Reaction
2 and 5
(6)
During the next positive-going voltage scan, the FeS was oxidized to Fe& at about -0.67 V, characterized by peak la. The magnitude of peak la correlates directly with the quantity of sulfur intermediates formed in the transpassive region with
I
Pyrite reactions during cyclic voltammogram
203
201-210
tial exceeded a certain value (Fig. 2); (b) there was enough conditioning time at the potentials above the threshold value (Fig. 3); and (c) the scan passed the threshold potential at least once before going through the potential at peak la (Fig. 4). The products (partially oxidized sulfur intermediates and iron hydroxide) formed in the transpassive region (peak 3) were reduced to Fe(OH), and HS- during the negative-going voltage scan. Fe(OH)t and HS- produced on the negative-going scan would react to form FeS according to the reaction [ 14.IS]
Quiescent
s
I.. -1.50
Surfaces A: Physicochem. Eng. Aspects 93 (1994)
of pyrite
at h&rent
speed 300 rev min-‘.
anodic
pH 9.2. scan
34
X. Zhu et ul./Colloids
IO.OOl6
la
12
I
A&
3 set ,
t....,.. -1.50
Surfuces A: Ph_wicochem. Eng. Aspects 93 ( 1994) 201-210
.1
-0.50
-1.00
1
.‘.‘I
,..‘I
0.00
0.50
E (V, vs SCE) Fig. 3. Cyclic voltammoprams for various times (rotating
of pyrite preconditioned
speed 200 rev min-‘.
at 0.6 V
pH 9.1. scan
rute 50 mV s-‘1.
-1.50
-1.00
-ll.so
0.00
0.50
with peak la increased with increasing electrode precondition time and tended to level off after 30 s. probably indicating saturation of the surface. The maximum charge associated with peak la is about 10 mC, corresponding to about 27 monolayers in thickness [IS]. As mentioned above, peak la was apparent in voltammograms only for scans reversed at an anodic potential higher than a threshold value. This value was 0.6 V for solutions of pH 9.2. A series of voltammograms at various upper potentials in solutions of pH 9.2 was obtained at 50 mV s-l, with the rotating speed of the disk electrode being 200, 700 and 1000 rev mine1 respectively. As the upper potential increased from 0.65 to 0.85 V. the intensity of peak la increased. Further increase of the anodic reversal potential resulted in the decrease in the intensity of peak la. Peak la disappeared for anodic reversal potentials above 1.1 V. This result clearly indicates that, at potentials higher than certain limits, the partially oxidizcd sulfur intermediates oxidized rapidly to sulfate ions. The integrated charge associated with peak la vs. the anodic reversal potential for various clcctrodc rotation speeds is shown in Fig. 5. For each electrode rotation speed, the intensity of peak la increased with increasing anodic reversal potential and rcachcd a maximum WIUC iit about 0.85 -0.9 V. The mcasuromcnts obtained from
I .w
E (V. vs SCE) Fig. 4. Cyclic
voltammograms
of
pyrik
in
various
l
1ooorpnl
0
7OU rpm
I
P?
cycles
(stirred solution, pH 9.2, scim rate IO0 mV se’).
the reaction for peak la being 2FeS+2H20=Fe(OH),_
+ FcS,+2H+
+2e(7)
All the factors that all’cct the transpassivc oxidation of pyrite influence peak la. The faster the clcctrodc rotated. the grcatcr the amount of substances participated in reaction (7). When the clcctrodc was preconditioned at a constant potential for increasing time periods, peak la was continuously cnhanccd. The intcgratcd charge associated
0.4
0.6
0.X
1.0
I.2
I.4
Anodic rcvcrsc potential (V) Fig. 5. Effect of nnodic rcvwsc pokntial scan r:Itc 50 mV s
’ J.
on pc;tk la (pH 9.2.
X. Zhu et aLfColloi& Surfuces A: Physicochem. Eng. Aspects 93 (1994) JOI-210
Raman spectroscopic experiments gave identical results [ 111. The influence of pH on the transpassive oxidation of pyrite was studied with both cyclic voltammetry and Raman spectroscopy. The pH value affected the minimum potential for the appearance either of peak la in the voltammograms or of the sulfur peaks in the Raman spectra. In acid solutions (pH 2.6) there was no peak la in the voltammograms as FeS is stable at low potentials only in the mid-pH range [ 151, although sulfur species could still be detected by Raman spectroscopy at high potentials. In strongly alkaline solutions (pH 12.5), the sulfur species including pyrite did not appear in the Raman spectra. The conditions for sulfur detection by the Raman spectroscopic technique are demonstrated in Fig. 6. It is not surprising that the Raman peaks of sulfur and/or polysulfides did not appear in the strongly alkaline solutions as these species were not stable. However, the overpotentials for sulfur detection seemed quite high in acid, neutral and weakly alkaline solutions. Kinetics were perhaps the main reason for the overpotentials. During the in situ Raman expcrimcnts, the maximum time for the potential application was normally 2 h. Another related factor was the detection limit of the Raman spectrometer set-up. Theoretically Raman spec-
20s
trometry can do monolayer measurements. For the set-up in these experiments. the best result obtained was about 90 monolayers [ 163.
4. Kinetics of surface sulfur layer growth In the transpassive region. aggressive electrochemical oxidation of pyrite occurs. The reaction products in this region are ferric hydroxide, elemental sulfur, sulfate ions and partially oxidized sulfur intermediates, with the majority of the sulfur product being sulfate ions. Sulfur products other than sulfate ion are metastable only. While exposure of pyrite to anodic potentials higher than the transpassive region results in rapid oxidation of sulfur intermediates to sulfate ions, the sulfur intermediates convert to soluble sulfoxy and ultimately sulfate ions at a measurable rate in the transpassive region. A surface layer on top of the oxidized pyrite is formed, based on the dynamic equilibrium between the inner layer growth of sulfur intermediates due to pyrite oxidation and the outer layer diminishmcnt due to its conversion to soluble sulfate ions. A schematic diagram is dcpictcd in Fig. 7 to show the growth mechanism of this sulfur layer. The surface sulfur layer has an iron-deficient structure or consists of a mixture of polysulfidcs FeS,, where ,x> 2. The iron concentration in the layer dccrcascs continuously from the inner to the outer surface while the sulfur concentration increases. Sulfur species may vary from Sg- (pyrite) on the inner W,
0
-1.0 -
Cl9
I
-1.5
.
0
’ . ’ 2 4
’ 6
.
I
a ’ 8
.
1
-w
so:
--_)
Fc(OlI),
F4
’ . ’ IO 12
14
PH Fig. 6. Conditions with
Raman
for sulfur dctcction from oxidation
spectroscopy:
(E[S]-~[Fc]=O.OOl
0.
unobscrvcd;
M in the E,-pH
diagram).
0,
of pyrite
I
observed Fig. 7. Schcmalic diagram
II of surface sulfur layer.
206
X. Zhu et aL/Colloids Surfuces A: Physicochem.
surface to S& (electrochemically reducible elemental sulfur) on the outer surface. Since the outer surface is saturated with sulfur, it appears that the dissolution reactions follow the linear kinetics. Iron lattice ions migrate through vacancies by diffusion, characterized by parabolic kinetics. The kinetics of surface layer growth as pyrite was oxidized in the transpassive region were studied using the potentiostatic technique. Fig. 8 shows the current decay curves obtained at potentials between 0.4 and 0.65 V. For applied potentials higher than 0.65 V, oxygen evolution started. and the current density increased irregularly. During the initial stage, currents decreased quickly, indicating the build-up of a passivation layer. As the layer increased in thickness, the rate of pyrite oxidation decreased and approached the dissolution rate of the outer surface layer. The thickness of the passive layer thus became essentially constant, leading to the steady movement of the whole surface layer toward pyrite. A paralinear rate equation may be derived to describe the kinetics of the transpassive oxidation of pyrite [ 14.153:
where Q is the accumulated charge passed on the electrode, I the time, k, the parabolic rate constant and k, the linear rate constant. Fig. 9 demonstrates ...
I..“,.
‘.I.-.
I 2 3 4
4.0
0.65 V 0.64lv 0.50 v 0.40 v
3.0
0.0
‘. 0
‘. loo0
‘. 2000
’ 3MMl
4MKl
Time (see) Fig. 8. Current
decay
curves of pyrite
(stirred solution, pH 9.2).
at variaus
potentials
Eng. Aspects 93 ( 1994) 201-110
:
-
0
Paralinear Equation
IO00
zoo0
3ooo
4ooo
Time (set) Fig. 9. Correlation equation
of current
decay
curves
with
paralinear
(stirred solution, pH 9.2).
the good fit of the experimental data to the paralinear rate equation. Thus the kinetics results are consistent with the proposed mechanism of surface layer growth: the formation of an intermediate passive layer during pyrite oxidation and the simultancous dissolution of the outer surface layer resulting from sulfur and polysulhdes oxidation to sulfate ions. By fitting the experimental data to Eq. (8). the parabolic and linear rate constants for various applied potentials were obtained. both k, and k, incrcascd with increasing applied potential. Consistent with results from the cyclic voltammetric expcrimcnts, increased anodic oxidation potcntial increased both the amount of sulfur intermediates formed and the amount of intermediates dissolved. The effect of agitation on the reaction rate was examined at dilfcrcnt applied potentials and solution pH values using rotating disk electrodes. For each cast, the initial current increased significantly as the electrode rotation speed jumped from 0 to 100 rev min-‘. However, further increase in the electrode rotation speed up to 7OOrcv min-’ increased the initial current only slightly while the steady state current remained almost constant even when the elcctrodc was stationary. The independcncc of the surface layer movcmcnt on agitation confirmed that the outer layer dissolution was limited by the surface reaction.
X. Zhu et al. JCoIloti
Sur-_aces A: Physicochem. @g. Aspects 93 (1994)
Both the parabolic rate constant k, and the linear rate constant k, increased with increasing solution pH as expected from reactions ( l)-( 5) although the pH influence was much more profound on k,, than on k,. The strong dependence of k, upon pH could be explained on the basis of the following assumption: the surface Fe’+ (lattice) concentration decreased with increasing pH. thus increasing the diffusion gradient and the rate of Fe’+ (lattice) diffusion through the passivation layer. The kinetics of the surface sulfur layer growth was studied in the temperature range 23.4-59S’C. The activation energies were found to be 66.17 kJ mol-r (15.83 kcal mol-‘) for the parabolic rate process and 38.67 kJ mol-’ (9.25 kcal mol- ‘) for the linear rate process. The increase in temperature accelerated both the formation and the dissolution of the surface layer. However, the thickness of the layer will increase with temperature since k, increases more rapidly than k, as temperature goes up. To quantify the influences of temperature, applied potentials and electrolyte pH on k, and k,, a general equation may bc dcvcloped: li=oexp(
-$)exp(;T
JOI-210
207
1
10.0 -
’ -
0
-.-I’.-.-
Panlinur Equation
1000
2000
1
3ooo
4ooa
Time (see) Fig. 10. Comparison estimated
of experimental
rate constants obtained
and calculated data using
from Eq. (9).
lPph
where k could bc k, or k,, E,, is the activation energy, and a, b and c are constants. E,, u, b and c were estimated from k, and k, obtained under different conditions with non-linear curve fitting. The parabolic and linear rate constants k, and k, at given temperature, applied potential and electrolyte pH can be calculated according to Eq. (9). Substituting these k, and k, values into paralinear Eq. (8). a series of charge versus time data was obtained. Fig. 10 illustrates the satisfactory correlation of calculated charge-time data using k, and k, estimated from Eq. (9) with the paralinear equation for various temperatures. Similar results were also obtained for various electrolyte pH values and applied potentials.
5. Properties of surface sulfur layer
Fig. 11 (curve a) shows the Raman spectrum of a sample of pyrite treated potentiostatically at + 1.0 V for 10 min. The sulfur/polysuhidc Raman
I-...@
100
I-...1
200
300
1.
400
500
I
..I
600
700
Frcqucncy shift (cm.‘) Fig. 11. Raman
spectra
of
pyrite
in 0.5 M
NaCl
solution
(pH 9.2): curve a. 10 min at 1.0 V; curve b, 3 min at - I.0 V; curve c. 2 min at + 1.0 V.
peaks can be seen clearly from the figure. The potential was then immediately switched to - 1.0 V. The Raman spectrum in Fig. I1 (curve b) was taken after 3 min at this potential. The sulfur/ polysulfide peaks that appeared at + 1.0 V now disappeared and the spectrum was almost identical to that at the beginning of the experiment. This suggests that cathodic polarization reduces the pyrite surface to a state that is spectroscopically identical to the freshly ground condition. At the end of the experiment the applied potential was
208
X. Zhu et aL/Colloids
Surjices
A: Physicochem. Eng. Aspects 93 ( 1994) 201-210
reversed again to + 1.0 V. Fig. 11 (curve c) demonstrates that the sulfur/polysulfide peaks reappeared after 2 min of oxidation. These results indicate that the sulfur layer formed during pyrite oxidation is readily reducible (S&J. Sulfur produced by electrochemical deposition from sulfide solutions at a platinum electrode was also reducible [ 163. The fact that the surface sulfur layer formed during pyrite oxidation could be reduced electrochemically implies that the sulfur thus formed does not behave as an electrical insulator as does bulk elemental sulfur in its orthorhombic state. The ability of the material to support a photocurrent [ 163 also suggests that the sulfurous film has unusual electronic properties more akin to a semiconductor than to an insulator. The formation of a hydrophobic layer on surfaces of metal sulfide minerals is deemed important to flotation by many researchers [ 11. By nature, the sulfur layer formed on the oxidized pyrite surface has such properties. The interactions between the surface sulfur layer and the flotation agent potassium amyl xanthate wcrc probed by Raman spectroscopy in this study. Raman spectra of sulfur and potassium amyl xanthate taken in air arc given in Figs. 12 and 13 (spectra a and b) as rcfcrcnccs. To increase sensitivity, the surface enhanced Raman spectroscopy (SERS) tcchniquc was employed for pyrite
loo
200
300
400
500
600
700
Frcqucncy shift (cm”) Fig. 12. Raman
spectra of sulfur in air (curve a). xanthatc
in
air (curve b). silver-plated
pyrite in xanthatc
solution (curve c),
oxidized and silver-plated
pyrite in xanthatc solution (curve d).
700
800
900
IO00
1100
1200
Frequency shift (cm”) Fig. 13. Raman
spectra of sulfur in air (curve a). xanthate
in
air (curve b). silver-plated
pyrite in xanthate solution (curve c),
oxidized and silver-plated
pyrite in xanthate solution (curve d).
electrodes. The pyrite electrode, whether in its original form or after being oxidized, was put into the solution containing 0.001 M AgNO, amd 0.5 M KNO, and polarized at -0.60 V for the passage of approximately 5 mC of charge. The silver-plated pyrite elcctrodc was then trnnsfcrrcd to a solution of pH 6.5 containing 0.5 M NaCl and a small amount of potassium amyl xanthate and its Raman spectrum recorded. Figs. 12 and 13 prcscnt the Raman spectra of the pyrite electrode: pyrite unoxidized in spectrum c and pyrite oxidized at 1.0 V for 1 h in spectrum d. A close look revealed that there wcrc several new peaks in the spectrum of the oxidized pyrite electrode. The peaks at about 170, 640 and 1IOOcm-’ were not found in the spectra of either unoxidized pyrite in the presence of amyl xanthate or oxidized pyrite in the absence of amyl xanthate. There are several possible interactions between amyl xanthate and the oxidized pyrite. Physical adsorption and chemical reaction are two cases. If it is the former process then the peak position will not change much (e.g. the peaks at 170 and 1100 cm-‘). However, the latter process could lead to the formation of new species and therefore new peak positions (e.g. the peak at 640cm-‘). Since them arc no new peaks found when unoxidized pyrite elcctrodc was put into a solution of 0.5 M NaCl and pH 6.5 with the addition of a small
X Zhu t-t aL/Colloidr Surfaces A: Phpicochem Eng, Aspects 93 (1994) 201-210
amount of amyl xanthate, it is reasonable to assume that sulfur or polysulfides play an important role in the flotation process. More work, however, needs to be done to identify the newly formed or adsorbed species. The formation of sulfur and polysulfides promotes xanthate adsorption on pyrite and so could enhance its floatability. Fig. 14 illustrates the influence of applied potentials on floatability of coal pyrite. Sodium sulfate and sodium borate were added to the solution as supporting electrolytes. The floatability increased first with the increase of the applied potential up to a maximum value and then decreased with the potential increase. The bell-shaped curve agreed completely with the results of sulfur layer formation obtained with the electrochemical and Raman spectroscopic measurements.
6. Conclusions
Aggressive oxidation of pyrite occurred in the transpnssivc region. The reaction products in this region wcrc Fc(lI1) oxides. sulfate ions, an active form of elcmcntnl sulfur and partially oxidized sulfur intcrmcdiatcs. The surface layer formation at high anodic potentials was confirmed by subscqucnt surbcc reactions during the cathodic and
“g “. l
./
/
\ Potcnfial (V, vs SCE)
Fig. I-8. Eflect of potenkd on tlo&bility of cd pyrik (0.125 M Na,S04, 0.005 M N:I&O,. pH 7.0; pxticlc six. - 100 -325 mesh).
209
return anodic cycle in cyclic voltammetry and by in situ Raman spectroscopy. Polysulfides (FeS, of variable X) form as intermediates during anodic oxidation. The iron concentration in the surface layer decreases continuously from the inner to the outer surface while the sulfur concentration increases. Sulfur species may vary from S:- (pyrite) on the inner surface to So (elemental sulfur) on the outer surface. The conditions such as applied potentials and pH that lead to the build-up of the surface oxidation layer have been determined. The influence of applied potentials on the quantity of sulfur formed can be represented by a bell-shaped curve. The sulfur formed increased with the applied potential up to a maximum value after which less sulfur was produced because of the rapid formation of sulfate ions. Sulfur and its intermediates were detected in both acidic and basic solutions (pH 2.6-10.5). Film growth kinetics obey the paralinear rate law indicating uniform film growth of the sulfurrich layer to a steady state thickness. The thickness is controlled by the simultaneous rate of diffusion of cations through the polysullidc layer, with parabolic rate constant k,, and the rate of oxidation of the active sulfur outer layer, with the linear rate constant k,, to soluble sulfur spccics in solution. The cfTccts of tcmpcrature, applied potential and pH on k, and k, wcrc dctcrmincd using the paralinear rate law. The activation encrgics were found to be 66.17 kJ mot-’ (15.83 kcal mol-‘) for the rate process and 38.67 kJ mot-’ parabolic (9.25 kcal mol-‘) for the linear process. The end product of the lilm growth is an active form of sulfur, SI1,,. which is readily reducible. This indicates that the sulfur formed during pyrite oxidation dots not behave as an electrical insulator as does bulk elemental sulfur in its orthorhombic state. The surface sulfur layer was hydrophobic in nature and its interaction with the flotation agent such as potassium amyl xanthate enhanced the floatability of pyrite as evidenced by Raman spectroscopy and flotation experiments.
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