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An investigation of surface oxidation of pyrite and pyrrhotite by linear potential sweep voltammetry
J. Electroanal. Chem., 118 (1981) 327--343
327
Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
AN INVESTIGATION OF SURFACE OXIDATION OF PYRITE AND PYRRHOTITE BY LINEAR POTENTIAL SWEEP VOLTAMMETRY
I.C. HAMILTON 1 and R. WOODS 2 iChemistry Department, Footscray Institute of Technology, P.O. Box 64, Footscray, Victoria 3011, Australia. 2CSIRO Division of Mineral Chemistry, P.O. Box 124, Port Melbourne, Victoria 3207, Australia.
ABSTRACT The products of surface oxidation of pyrite and pyrrhotite have been determined from analysis of linear potential sweep voltammograms. Pyrite oxidizes to both sulphur and sulphate.
The formation of sulphur is
restricted to the order of a monolayer at pH 9.2 and 13, but significant yield occurs at pH 4.6.
The proportion of sulphate formed increases rapidly with
increase in potential. Sulphur is the major product of pyrrhotite oxidation at pH 4.6, 9.2 and 13. Sulphate is also formed in significant quantities, particularly in the alkaline solutions.
Oxidation of pyrrhotite is strongly inhibited by the surface ferric
oxide produced.
INTRODUCTION Iron forms many compounds with sulphur which have a range of stoichiometries and crystal structures (ref. i).
The iron compounds are the most abundant of
the metal sulphides existing in nature, the predominant species being pyrite (FeS 2) and pyrrhotite (Fel_xS).
Iron sulphides constitute a binary boundary of
many multicomponent base metal systems of economic importance, e.g., copper, nickel and zinc, and they exist widely in sulphide ore bodies.
A knowledge of
their properties is important, therefore, in mining the ore, separating the valuable components from the mineral assemblages and in processing to win the metal values. Fine-grained iron sulphides oxidize rapidly in air; dust explosions (ref. 2) and "hot ground" problems due to spontaneous combustion (refL. 3-5) have occurred during mining.
On the other hand, large crystals of these sulphides, particular-
ly pyrite, retain their natural metallic sheen and, because of this, are used in
0022-0728/81/0000--0000/$02.50, © 1981, Elsevier Sequoia S.A.
328
jewellery.
This contrast in reactivity between the finely divided and massive
forms is similar to that observed for iron metal.
The resemblance between the
iron sulphides and the constituent metal suggests similar involvement of surface films in inhibiting sustained oxidation.
In fact, it has been demonstrated
(ref. 6) that a thin iron oxide layer forms spontaneously when pyrite is exposed to air. Pyrite is a good electrocatalyst for oxygen reduction (refs. 7,8) and hydrogen evolution (ref. 9), being much more active than pyrrhotite and most other metal sulphides (refS. 8,10).
Galvanic coupling between pyrite and other sulphide
minerals can result in an increased rate of oxidation of these minerals because oxygen reduction on the pyrite provides a relatively facile cathodic process (ref. Ii).
Indeed, the ability of pyrite to enhance the oxidation rate of other
sulphides has been known for many years (ref. 12).
Coupling with pyrite leads to
sulphur being formed on the surface of base metal sulphides (ref. ii) and, because sulphur is hydrophobic, this results in difficulties in the selective flotation of complex sulphide ores.
On the other hand, excessive oxidation, in which
sulphur is lost from the surface as soluble oxy-sulphur ions leaving a hydrated metal oxide layer, results in inhibition of flotation (ref. 13). Chalcopyrite (CuFeS2) has been shown to float in the absence of flotation collectors under oxidizing but not reducing conditions (ref~. 14,15). behaviour has been ~ttributed oxidation. 16).
This
(ref. 15) to the formation of sulphur by surface
Other sulphides also float naturally under certain conditions (ref.
Thus the products of surface oxidation can play an important role in
determining the flotation properties of sulphide minerals. The oxidative leaching reactions of pyrite and pyrrhotite have been studied in some detail (ref. 17), these minerals being widely present as impurities in sulphide concentrates.
The conclusions reached from hydrometallurgical studies
are supported by analysis of oxidation reactions at sulphide electrodes (refs. 18-20).
However, these investigations have identified bulk oxidation products
whereas it is the surface reaction products formed initially which are important in flotation. In this communication, we present an investigation of the surface oxidation of pyrite and pyrrhotite employing linear potential sweep voltammetry.
EXPERIMENTAL The pyrite and pyrrhotite electrodes were prepared from natural massive specimens of these minerals.
The pyrite was of Spanish origin and the pyrrhotite
came from Mt Isa, Queensland, Australia.
Pyrite has a stoichiometry which is
always very close to FeS2, while pyrrhotite has a range of stoichiometries with x in Fel_xS varying from 0 to 0.14 (ref. 21).
The ratio of sulphur to iron in
the pyrrhotite employed in the present investigations was found by analysis to
329 be 1.13.
Thus, we have assumed the pyrrhotite
Selected pieces were cut from the mineral epoxy resin as described
previously
to be FeSI.13 , that is, x = 0.115.
specimens and encapsulated
(refs. 7,8).
in an
The sulphide electrodes were
mounted on a Beckman rotated electrode assembly. In order to avoid any oxidation of the electrode
surface before the experiments
were run, all operations were carried out in a nitrogen atmosphere A fresh surface was produced on the mineral
in a glove bag.
surface before each electrochemical
run by wet grinding on 600 grade silicon carbide paper in the glove bag.
The
freshly ground surface was transferred
cell
which was also contained
immediately
to the electrochemical
in the glove bag.
The glove bag and the cell solutions were purged with nitrogen purified by passage over heated copper deposited on kieselguhr. Potentials were measured and converted
against a saturated calomel reference
to the standard hydrogen
electrode
has a potential of 0.245 V against the SHE (ref. 22). are on the SHE scale.
Current-potential
and charge-potential
on a Hewlett-Packard
programmed
and serviced by a current integrator.
and integrator were designed
and constructed
relationships
(SCE)
the SCE
All potentials reported
The potential was controlled by a potentiostat
with a Utah 0151 sweep generator potentiostat
electrode
(SHE) scale, assuming
The
in these laboratories.
were recorded
simultaneously
7046A XYIY 2 recorder.
Experiments were carried out in solutions of (i) pH 4.6; 0.5 M C H 3 C O O H +
0.5 M
CH3COONa , (ii) pH 9.2; 0.05 M Na2B407 and (iii) pH 13; 0.i M NaOH. In all the figures presented
in this paper, positive currents are anodic and
negative currents cathodic.
RESULTS AND DISCUSSION For the investigation
of pyrite and pyrrhotite
techniques used in previous
studies
surfaces,
we have employed
(refs. 15,23) of sulphide surfaces.
the
These
involve oxidation of the surface during a linear potential sweep and analysis of the products of anodic oxidation on a subsequent reverse potential behaviour
from the characteristics
scan.
Interpretation
of their reduction
of the voltammetric
requires a knowledge of the possible reactions which involve the
iron sulphides
and iron-and sulphur-containing
Thermodynamics
of the Fe-S-H20
Eh - pH diagrams
species.
system
for the S-H20 , Fe-H20 , FeS2-H20 and FeI_xS-H20
10 -3 M dissolved species are presented in Fig.
i.
systems
for
These diagrams were computer
generated employing the CSIRO-NPL Thermodata System (ref. 24). The diagram for the S-H20 system, Fig. above pH 4-5.
However,
la, shows that sulphur is not stable
electrode reactions
involving sulphate are irreversible
and sulphur can be formed from oxidation of sulphide species at high pH.
We have
330
1,0
i
i
i
]
i
1,0 :
i
I
I
I
I
i
i
(b)
(o)
0.5
0,5
Eh
H
=
S
SO,~"
~
0
0
-0.5
~
-I.0
0
~
'
~ Fez+
Fe
0.5
HS-
I
J
I
I
I
I
2
4
6
8
I0
12
i
I
I
I
I
14
1.0
J
I
I
I
i
I
2
4
6
8
I0
12
I
I
I
]
I
I
pH
pH 1.0 I ~
1,0
(c)
o L ,:A Eh
~
(
(OHa )
'
(d}
0.5 Fe (OH)s
i
FeZ+
H)3
0
~
__~
-~
(;
-05
• 0.5
.Fe(OH)2
~Fe(CH)= Fe -I.0 0
;
I
I
]
L
]
2
4
6
8
I0
12
pH
14
-I.0
I
I
I
J
I
I
2
4
6
8
I0
12
14
pH
Fig. i. Eh-pH diagrams for the (a) S-H20 , (b) Fe-H20 , (e) pyrite (FeS~)-HoO o m and (d) pyrrhotite (FeSI.143)-H20 systems at 25 C and 10-3M dissolved specles. Diagrams computer generated employing the CSIRO-NPL Thermodata system (ref. 24).
331
included in the S-H20 diagram the boundary between metastable sulphur and HS(dashed line) in addition to the thermodynamically more favourable equilibria. The diagram also indicates that sulphate in solution could be reduced to H2S or HS- at potentials within the range covered by the voltammograms.
However, this
process is highly irreversible and has a negligible rate at these potentials. Peters (ref. 25) points out that zinc electrowinning from sulphate solutions takes place at an overpotential of % 1 V
with respect to H2S formation, but no
trace of this compound is detected in practice. The Fe-H20 diagram, Fig. ib, includes the hydrated iron oxides, since they are more likely to be formed in the present environment than anhydrous oxides. The pyrite diagram, Fig. ic, shows the domain of stability of this mineral. The sulphur-containing products are dissolved sulphide species below, and sulphate ions abovejthe pyrite domain, i.e., in the stability regions of these species in Fig. la.
Oxidation of sulphides to sulphate generally requires high overpotentials
(ref. 9) and occurs at significant rates only at potentials at which sulphur can also be a product.
Hence, the boundary to which pyrite is metastable before
oxidation to Fe 2÷ or Fe(OH) 3 and S is included (dashed lines) in addition to the thermodynamically favoured equilibria. Figure id shows the Eh-pH diagram for pyrrhotite. not been included.
Oxidation to pyrite has
Pyrite can be formed in hydrothermal systems involving H2S ,
e.g., it forms through a series of iron sulphides in the corrosion of iron (refs. 26,27).
However, the oxidation step from pyrrhotite to pyrite is slow even at
100°C and occurs through a dissolution, reprecipitation mechanism rather than a transformation of the solid (ref. 28).
Furthermore, the reaction of pyrrhotite
with dissolved sulphide species is not relevant to the interpretation of the voltammograms presented here, since no sulphide species were present in solution at the commencement of the potential sweeps.
Conversion to pyrite by a reaction
such as,
2Fel_xS
÷
~l-2x)Fe 2+ + FeS 2 + 2(l-2x)e
(i)
has been neglected since there is no evidence of pyrite formation during the leaching of pyrrhotite (ref. 29).
In fact, reactions in nature which give rise
to pyrite are known to be particularly slow (ref. 30).
The anodic reactions
presented are for the formation of sulphate and metastable sulphur (dashed lines). The data in Fig. id correspond to pyrrhotite with composition FESI.143; the diagram for FeSI.13 will not be significantly different.
Oxidation of pyrite Voltammograms for a stationary pyrite electrode are shown in Fig. 2.
The
potential sweeps were commenced at the open circuit potential marked * and each
332 I
I
,
I
,
']
2.0
20 1.5 f
J
1
u
f
f
E
E
"~
1.0
==
g
I0
I
I
0.5
)
8 5
- J~
4
I
-0.2
- 0.4
m 2 I
I
I
I
0
0.2
0.4
0.6
Potential
I 0.8
V vs SHE
Fig. 2. Voltammograms for a stationary pyrite electrode at pH 4.6. Linear potential sweeps at 20 mV s -I reversed at different upper potential limits. Each v o l t a m m o g r a m relates to a freshly ground surface. Dashed curve is the recorded charge on the voltammogram, curve 4.
sweep was carried out on a freshly ground surface. positive
going scan,
It can be seen that,
for a
there is a steady rise in anodic current as the potential
is increased. It is apparent Fig.
FeS 2
2 to either
÷
from Fig.
ic that pyrite can be oxidized at the potentials
in
sulphur,
Fe 2+ + 2S + 2e
(2)
or sulphate,
FeS 2 + 8H20
÷
Fe 2+ + 2S042- + 16H + + 14e
There is a step on the curve at ~0.4 V. (see Fig.
(3)
This can be assigned
to the reaction
ib),
Fe 2+ + 3H20
÷
Fe(OH) 3 + 3H + + e
(4)
o
333 Investigation
of the effect of electrode rotation enables a distinction
to be
made between reactions which result in soluble products and those in which only surface species are produced
(refs.
15,23,31).
For a stationary electrode,
dis-
solved products remain in the vicinity of the surface and can take part in further electrochemical are dispersed
reactions.
When the electrode
is rotated,
the soluble products
as they are formed.
The step assigned This observation
to reaction
(4) is absent on rotating
is consistent with reaction
the pyrite electrode.
involving dissolved
species.
cathodic peak at %0.25 V, which is also absent when the electrode arises from the reverse of reaction
(4).
reduction of iron oxide on chalcopyrite It is to be expected state (Fig.
ib).
A similar peak is observed for the (ref.
(v.i.).
15) and pyrrhotite
that all iron formed above ~0.4 V will be in the ferric
However,
the charge under the cathodic peak at %0.25 V is insuf-
ficient to account for all the iron species arises from dissolution
The
is rotated,
formed anodically.
This probably
of Fe(OH) 3 at this pH to form species such as FeOH 2+ and
Fe(OH)2+ which can diffuse away from the electrode The cathodic peak at lower potentials
surface.
can be assigned
sulphur which had been formed by reaction
(2).
to the reduction of the
In acid solution,
the reduction
will produce H2S.
S + 2H + + 2e
÷
H2S
(5)
When the scan is reversed again at -0.35 V (Fig. 2), the H2S is oxidized back to sulphur by the reverse of reaction 0 V.
(5) and this gives rise to the anodic peak at
The effect of electrode rotation supports
this view.
unaffected but the anodic peak is eliminated because
The cathodic peak is
the H2S produced
cathodically
is dispersed. The charge associated with the lower cathodic peak is a measure of the quantity of sulphur formed anodically by reaction sum of reactions
(2),
(3) and
(4).
(2).
The anodic charge is due to the
If we assume that the iron product
ferrous state below 0.4 V and ferric above this potential, to determine
is in the
then it is possible
the number of moles of pyrite oxidized to sulphur and to sulphate
from integration
of the anodic and cathodic
charges.
An example of the recording of the charge is shown in Fig. 2. anodic charge is given by the maximum on the charge-potential
The total
curve.
The Fe(OH) 3
is given by the step at 0.25 V and the sulphur by the difference between the charge at 0.15 V and the minimum on the second positive going scan. Reduction (Fig.
of pyrite to a lower iron sulphide is expected from the Eh-pH diagram
Ic) and it was not possible
reduction
to set a lower scan limit which allowed complete
of sulphur on the negative-going
reduction of the pyrite itself.
sweep without causing considerable
However a lower limit can be chosen such that
334 reduction of sulphur is completed during the sulphur peak includes cathodic scan.
The sulphur charge was corrected
tion of the mineral by subtracting
the return scan.
Thus integration
charge passed at the beginning
of
of the reverse
for the small amount of background
reduc-
the charge passed during a negative-going
applied to a freshly ground pyrite surface
sweep
(curve i, Fig. 2).
The quantity of pyrite oxidized to sulphur and to sulphate on a potential sweep is presented as a function of the upper potential
limit of the sweep in
Fig. 5. In basic solution,
oxidation of pyrite takes place at potentials
stability domain of hydrated
in the
ferric oxide and hence the oxidation reactions
will be,
FeS 2 + 3H20
+
(6)
Fe(OH) 3 + 2S + 3H + + 3e
and
FeS 2 + IIH20
+
The reduction Figs.
Fe(OH) 3 + 2S042- + 19H + + 15e
(7)
of both Fe(OH) 3 and S occurs in a similar potential range
(see
la and ib) and hence a single cathodic peak appears on the voltammogram
(Figs. 3 and 4).
It is unlikely that HS- would be released
iron oxides are present on the mineral that an iron sulphide will be produced.
surface.
Rather,
to the solution since
it is to be expected
The cathodic peak occurs at potentials
at which pyrite itself is reduced to give a small but significant i, Fig. 3).
Thus, the likely product of sulphur reduction
current
iron oxides is FeS.
This conclusion
product precipitated
from aqueous solutions of iron and sulphide species.
There is negligible 4) when the electrode involve solid species. oxidation
substantiating
at pH 9.2 and 13 (Figs. 3 and
the conclusion
that all reactions
We have taken the reduction of the products of anodic
to be represented
Fe(OH) 3 + S + 3e
is supported by the fact that FeS is the
effect on the voltammograms is rotated,
(curve
in the presence of
by,
*
FeS + 3OH-
(8)
+
Fe(OH) 2 + H20
(9)
and
Fe(OH) 3 + H + + e
If we assume that x moles of pyrite are oxidized sulphate by reactions the potential
(6) and
(7) respectively,
to sulphur and y moles to
then the anodic charge passed on
sweep will be (3x + 15y)F coulombs.
The cathodic charge arising
335 I
I
I
I
I
1
1.5
i%
I.O -
o
6
E >,
•"o
0.5
C
8 6 0
6 I
I
-0.4
-0.2
[
I
I
I
0
0.2
0.4
0.6
Potential
V vs SHE
Fig., 3, Voltammograms for a stationary pyrite electrode at pH 9.2. Linear potential sweeps at 20 mV s -I reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
from reaction
(8) will be 6xF coulombs.
moles of which 2x are consumed
The total Fe(OH) 3 produced
in reaction
(8).
Thus the charge due to reaction
(9) will be (y - x)F and the total cathodic charge
(Sx + y)F.
for the oxidation of pyrite at pH 9.2 and 13, determined are also presented
is (x + y)
Values of x and y
from these equations
in Fig. 5.
It can be seen from Fig. 5 that, in all three solutions, product is sulphur with very little sulphate being formed. proportion of sulphate produced
increases as the potential
the initial oxidation In each case, the is taken to higher
values. At pH 4.6, the sulphur yield also increases with increase in potential,
but
sulphate becomes the dominant product above 0.8 V. At pH 9.2 and 13, the sulphur yield remains constant at ~0.45 x 10 -8 moles cm -2 , while the sulphate produced
increases rapidly.
This suggests that the
first anodic peak (0 V, pH 9.2; -0.2 V, pH 13) constitutes with the steady increase in current at high potentials sulphate.
being oxidation to
The charge passed in the production of the sulphur layer is similar
to that found assigned
sulphur formation
(ref. 23) for a prewave
to monolayer
in the oxidation of galena which was
formation of PbO + S.
It is reasonable
to assume that the
336 [
I
f
I
I
2.0
1.5
0
<~ E
1.0
"tD C L
8
0.5
0
-_iji;5
-0.6
;
I
-0.4
i
-0.2 Potenfiol
J
0
04
0.2
V vs S H E
Fig. 4. Voltammograms for a stationary pyrite electrode at pH 13. Linear potential sweeps at 20 mV s -I reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
sulphur layer on pyrite is also of the order of a monolayer There appears to be a decrease potentials
on the surface.
in the quantity of sulphur present at high
in pH 9.2 and 13 solutions.
However,
this could arise from under-
estimation of the cathodic charge due to either the background reduction
of the
mineral being inhibited by the presence of surface oxides, or failure to reduce all the oxidation products before reversal of the sweep slows the reduction
to a
negligible rate.
Oxidation of pyrrhotite It is well established
(ref. 17) that pyrite oxidizes entirely to sulphate and
sulphur during leaching processes formed. ions.
and that no other oxy-sulphur
On the other hand, pyrrhotite However,
the differential
species are
is known to give rise to lower oxy-sulphur
increase in anodic charge on potential
sweeps to
337 I
%
I
I
I
2.0
o
1.5
/7
1.0
,
f~
0.5
•
. el 2 ' o S '
0 0
0.2 0.4 Potential Limit V vs SHE
0.6
0.8
-i Fig. 5. Quantity of pyrite oxidized to sulphur and sulphate on a 20 mV s potential sweep as a function of the upper potential limit of the sweep: (i), (i') pH 4.6, (2), (2') pH 9.2, (3), (3') pH 13; (i) O , (2) ~, (3) h , sulphur, (i') O, (2') I, (3') A , sulphate.
high potentials in alkaline solution (>0.7 V) is close to 10 times the corresponding increase in cathodic charge, which is the value expected if pyrrhotite is oxidized to sulphate (see reaction 13).
It is possible that the other oxy-
sulphur anions, formed in the leaching of pyrrhotite in ammoniacal and caustic solutions, arise from air oxidation of sulphur which is a product of the anodic oxidation of the mineral at lower potentials
(v.{.).
Voltammograms for a stationary pyrrhotite electrode at pH 4.6 are presented in Fig. 6.
We have analyzed these curves in terms of the anodic reactions,
FeSl.13
÷
Fe2+ + 1.13S + 2e
(i0)
*
Fe(III) + 1.13S + 3e
(li)
or FeSI.13
and
F e S l . 1 3 + 4.52H20
÷
Fe 2+ + 1.13S042- + 9.04H+ + 8.7Se
(12)
338
2.0
--
•
r
I
i
I
i
I0-
?
oE E
o
8 -I0
-
-2.0
-
I
-0.4
-02
I
I
I
I
0
0.2
04
06
Potential
08
V vs SHE
Fig. 6. Voltammograms for alstationary pyrrhotite electrode at pH 4.6. Linear potential sweeps at 20 mV s reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
or
FeSl.13 + 4.52H20
÷
Fe(lll) + 1.13SO42- + 9.04H + + 9.78e
(13)
The peak at 0.43 V in Fig. 6, which is absent when the electrode is equivalent of Fe 2+.
to the step observed with pyrite which is assigned
is rotated,
to the oxidation
The cathodic peak at ~0.2 V arises from reduction of Fe(0H) 3 by the
reverse of reaction
(4).
The charge associated with this peak does not account
for all the iron species formed above 0.4 V so that some of the iron(III) must dissolve.
The peak is not eliminated
diminished
to some extent.
on rotating the electrode,
this peak is formed directly and is not produced As with pyrite,
from a dissolved
the sulphur formed anodically
it is
is reoxidized
second positive going scan unless it is dispersed by rotating to the voltammograms
same manner as for pyrite in order to determine
species.
is reduced at the lower end of
the reverse scan and the H2S produced by this reduction
The charges corresponding
although
This suggests that the oxide which gives rise to
in Fig.
on the
the electrode.
6 were analyzed
in the
the number of moles of pyrrhotite
339 oxidized to sulphur and to sulphate. were assumed to arise from reactions
That is, the anodie currents below 0.4 V (10) and (12) and above this potential,
to
reactions
(ii) and (13), with the cathodic peak at low potentials arising from
reduction
of the product
Voltammograms Fig. 7.
sulphur by reaction
for stationary pyrrhotite
(5).
electrodes
at pH 9.2 are presented
Rotated electrodes produce almost identical curves.
at this pH are due to reactions
(ii) and (13), the Fe(lll)
in
The anodic currents
species being Fe(OH) 3.
The processes which give rise to the cathodic peak on the reverse scan will be the same as those for pyrite,
i.e., reactions
(8) and (9).
The first anodic
wave on the second positive going scan arises from re-oxidation produced on the preeeeding reverse scan by reaction
(9).
of the Fe(OH) 2
The peak at ~0.05 V
we assign to oxidation of the FeS formed on the reverse scan by reaction It is more active than the original pyrrhotite, to be rapidly inhibited by the products,
cathodic peak.
Although the same processes
pyrite and pyrrhotite
I
surfaces,
I
but the oxidation process seems
ferric oxide and sulphur,
total anodic charge to 0.2 V is significantly
(8).
since the
less than that of the preeeeding
have been assumed to take place at
this anodic peak is not clearly discernable
I
I
1
on
I
2.0
4
1.0 E
o 8 -I.0
I -0.4
I -0.2
I 0 Potential
I 0.2
I 0.4
I 0.6
V vs SHE
Fig. 7. Voltammograms for alstationary pyrrhotite electrode at pH 9.2. Linear potential sweeps at 20 mV s- reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
0.8
340
volEammograms
for the former mineral.
This is because the quantity of sulphur,
and hence of FeS, formed on pyrite is much less than on pyrrhotite. The quantity of pyrrhotite
oxidized
in the same manner as for pyrite.
to sulphur and to sulphate can be obtained
Thus, if x moles form sulphur and y sulphate,
the anodic charge will be (3x + 9.78y)F coulombs and the cathodic,
(3.26 + y)F.
The latter value arises because 3.39e are involved in the reduction to FeS of the 1.13x moles of sulphur and the remaining Fe(OH)3,
viz.
(y - 0.13x) moles,
are reduced to Fe(OH) 2. Voltammograms presented
for stationary and rotated pyrrhotite
in Fig. 8.
at t0 V, followed by a steeply increasing
current.
results in inhibition of the anodic process. from the thermodynamic
electrodes at pH 13 are
The anodic oxidation curves are characterized
instability
I
I
Rotation of the electrode
This phenomenon probably results
of pyrrhotite
I
by a peak
at high pH (Fig. id).
Thus,
I
I
I
I
I
- 0.4
- 0.2
I 0
I 0.2
3.0
2.0
?
5 E
==
|.0
--
C
8 0
f -I.0
3j
J
__J I -0.8
i -0.6
Potential
V vs SHE
Fig. 8. Voltammograms for a pyrrhotite electrode at pH 13. Linear potential sweeps at 20 mV s -I with electrode - - stationary and - - - rotated at i00 rps. Each voltammogram relates to a freshly ground surface.
341 part of the reaction could proceed via dissolved
sulphide
dispersed
on stirring and this would lead to a decrease
current.
The observation
rotated electrode
is diminished,
but the following anodic peak due to oxidation constant,
would support a loss of sulphur with-
in Fig, 8, the anodic currents are greater on the
reverse than on the initial positive-going for chalcopyrite
electrodes.
peak for the
loss of iron on stirring.
At the higher potentials
observed
in the overall anodic
(Fig. 8) that the cathodic reduction
of Fe(OH) 2 remains approximately out a corresponding
species which will be
(ref.
scan.
18), pyrite
It can arise from nucleation
(ref.
This behaviour
has also been
19) and galena
(ref. 23)
of active sites or pits which act as
growth centres for further reaction. The voltammograms
at pH 13 were analyzed
in the same manner as at pH 9.2 to
determine the quantity of sulphur and sulphate produced The quantity of pyrrhotite oxidized
to sulphur and to sulphate is presented
in Fig. 9 as a function of the potential
I
1
anodically.
limit of the sweep.
1
It can be seen
I
I
o?
6.0 L) (31.
% x
¢1
4.0
0 O
¢-
0.. (/)
2.0
-
O
/
O 0
0.2
0.4
0.6
I_ 0.8
Potential Limit V vs SHE -i Fig. 9. Quantity of pyrrhotite oxidized to sulphur and sulphate on a 20 mV s potential sweep as a function of the upper potential limit of the sweep: (i), (i') pH 4.6, ( 2 ) , ( 2 ' ) p H 9.2, ( 3 ) , ( 3 ' ) p H 13; (i) O, (2) D , (3) A , sulphur, (i') Q , (2') • , (3') • , sulphate.
342
that sulphur is the major oxidation product in each solution at all potentials investigated and initially accounts for all the pyrrhotite oxidized.
The amount
of sulphate produced increases as the sweep is taken to higher potentials. However, part of the current at the higher potentials in Figs. 7 and 8 still goes to produce sulphur, the sulphur yield continually increasing with potential. We have found that sulphate is the dominant product on sweeps taken to much higher potentials. Oxidation of pyrrhotite is inhibited significantly in alkaline solution, suggesting that iron oxide on the surface retards further reaction.
The anodic
oxidation curves for pyrite are also less steep in the basic media than at pH 4.6 (Figs. 4-6), but the inhibition for this mineral is not large since a regular shift in the anodic wave is observed of ~60 mV per pH unit.
On the other hand,
the oxidation of pyrrhotite at pH 9.2 is actually less than at pH 4.6 at the same potential on the SHE scale, even though the equilibrium potentials of the anodic reactions
(Ii) and (13) (with Fe(OH) 3 as the Fe(III) species) shift to
lower potentials with increase in pH.
The difference between the behaviour of
the two iron sulphides regarding the inhibiting effect of the product iron oxide is also apparent from a comparison of Figs. 5 and 9.
It could arise from different
structures of the hydrated ferric oxide species formed.
Pyrite is known (ref. 17)
to oxidize in alkali to form y-Fe203 which covers the surface but allows further oxidation thromgh a solid-state topotactlc reaction, whereas pyrrhotite becomes coated by a layer of ferric oxide which does not contain any well-defined oxide structure.
REFERENCES 1 P . H . Ribbe, (Ed.), Sulfide Mineralogy, Mineralogical Society of America, Washington, 1974. 2 M.J. Taylor, CIM Bulletin, 70 (1977) 78. 3 G.M. Lukaszewski, in M.J. Jones (Ed.), Mining and Petroleum, I.M.M., London, 1970 pp. 803-19. 4 R.L. Hewett, Proc. Aust. Instn. Min. Met., 226 (1968) 73. 5 D.J.M. Farnsworth, CIM Bulletin, 70 (1977) 65. 6 D. Michel_land R. Woods, Aust. J. Chem., 31 (1978) 27. 7 T. Biegler, D.A.J. Rand and R. Woods, J. Electroanal. Chem., 60 (1975) 151. 8 T. Biegler, D.A.J. Rand and R. Woods, in J.O'M. Bockris, D.A.J. Rand and B.J. Welch (Eds.), Trends in Electrochemistry, Plenum, N.Y., 1977, pp. 291-302. 9 E. Peters, in J.O'M. Bockris, D.A.J. Rand and B.J. Welch (Eds.), Trends in Electrochemistry, Plenum, N.Y., 1977, pp. 267-90. i0 D.A.J. Rand, J. Electroanal. Chem., 83 (1977) 19. ii H. Majima, Canad. Met. Q., 8 (1969) 269. 12 H.A. Buehler and V.H. Gottschalk, Econ. Geol., 5 (1910) 28. 13 N. Arbiter, The Physical Chemistry of mineral-reagent interactions in sulfide flotation, proceeding of symposium held at College Park, Md., 1978 U.S.B.M. ICC 8818, pp. 144-9. 14 G.W. Heyes and W.J. Trahar, Int. J. Miner. Process., 4 (1977) 317. 15 J.R. Gardner and R. Woods, Int. J. Miner. Process., 6 (1979) i. 16 G.W. Heyes and W.J. Trahar, Int. J. Miner. Process., 6 (1979) 229. 17 A.R. Burkin, The Chemistry of Hydrometallurgical Processes, E. & F.N. Spon, London, 1966.
343
18 T. Biegler and D.A. Swift, J. Appl. Electrochem., 9 (1979) 545. 19 T. Biegler and D.A. Swift, Electrochim. Acta, 24 (1979) 415. 20 L.K. Bailey and E. Peters, Can. metall. Q., 15 (1976) 333. 2 1 G . Kullerud, P.H. Abelson (Ed.), Researches in Geochemistry, J. Wiley, Chichester, Vol. 2, 1967, pp. 236-333. 22 R.G. Bates, Determination of pH, Wiley, N.Y., 1964, pp. 458-483. 23 J.R. Gardner and R. Woods, J. Electroanal. Chem., i00 (1979) 447. 24 A.G. Turnbull, Chem. in Aust., 44 (1977) 334. 25 E. Peters, Internal Report, U.B.C. 26 J.S. Smith and J.D.A. Miller, Br. Corros. J., I0 (1975) 136. 27 A.G. Wikjord, T.E. Rummery and F.E. Doern, Can. Mineral., 14 (1976) 571. 28 P. Taylor, T.E. Rummery and D.G. Owen, J. Inorg. Nucl. Chem., 4 (1979) 1683. 29 E. Peters, in D.J.I. Evans and R.S. Shoemaker (Eds.), Intern. Symp. Hydromet., AIME, 1973, pp. 205-28. 30 P.B. Barton and B.S. Skinner, in H.L. Barnes (Ed.), Geochemistry of Hydrothermal Ore Deposits, Holt, Rinehart and Winston, N.Y., 1967 pp. 236-333. 31P.E. Richardson and E. Maust Jr., in M.C. Fuerstenau (Ed.), Flotation, A.M. Gaudin Memorial Volume, AIME, N.Y., 1976, Vol. i, Ch. 12.
327
Elsevier Sequoia S.A., Lausanne -- Printed in The Netherlands
AN INVESTIGATION OF SURFACE OXIDATION OF PYRITE AND PYRRHOTITE BY LINEAR POTENTIAL SWEEP VOLTAMMETRY
I.C. HAMILTON 1 and R. WOODS 2 iChemistry Department, Footscray Institute of Technology, P.O. Box 64, Footscray, Victoria 3011, Australia. 2CSIRO Division of Mineral Chemistry, P.O. Box 124, Port Melbourne, Victoria 3207, Australia.
ABSTRACT The products of surface oxidation of pyrite and pyrrhotite have been determined from analysis of linear potential sweep voltammograms. Pyrite oxidizes to both sulphur and sulphate.
The formation of sulphur is
restricted to the order of a monolayer at pH 9.2 and 13, but significant yield occurs at pH 4.6.
The proportion of sulphate formed increases rapidly with
increase in potential. Sulphur is the major product of pyrrhotite oxidation at pH 4.6, 9.2 and 13. Sulphate is also formed in significant quantities, particularly in the alkaline solutions.
Oxidation of pyrrhotite is strongly inhibited by the surface ferric
oxide produced.
INTRODUCTION Iron forms many compounds with sulphur which have a range of stoichiometries and crystal structures (ref. i).
The iron compounds are the most abundant of
the metal sulphides existing in nature, the predominant species being pyrite (FeS 2) and pyrrhotite (Fel_xS).
Iron sulphides constitute a binary boundary of
many multicomponent base metal systems of economic importance, e.g., copper, nickel and zinc, and they exist widely in sulphide ore bodies.
A knowledge of
their properties is important, therefore, in mining the ore, separating the valuable components from the mineral assemblages and in processing to win the metal values. Fine-grained iron sulphides oxidize rapidly in air; dust explosions (ref. 2) and "hot ground" problems due to spontaneous combustion (refL. 3-5) have occurred during mining.
On the other hand, large crystals of these sulphides, particular-
ly pyrite, retain their natural metallic sheen and, because of this, are used in
0022-0728/81/0000--0000/$02.50, © 1981, Elsevier Sequoia S.A.
328
jewellery.
This contrast in reactivity between the finely divided and massive
forms is similar to that observed for iron metal.
The resemblance between the
iron sulphides and the constituent metal suggests similar involvement of surface films in inhibiting sustained oxidation.
In fact, it has been demonstrated
(ref. 6) that a thin iron oxide layer forms spontaneously when pyrite is exposed to air. Pyrite is a good electrocatalyst for oxygen reduction (refs. 7,8) and hydrogen evolution (ref. 9), being much more active than pyrrhotite and most other metal sulphides (refS. 8,10).
Galvanic coupling between pyrite and other sulphide
minerals can result in an increased rate of oxidation of these minerals because oxygen reduction on the pyrite provides a relatively facile cathodic process (ref. Ii).
Indeed, the ability of pyrite to enhance the oxidation rate of other
sulphides has been known for many years (ref. 12).
Coupling with pyrite leads to
sulphur being formed on the surface of base metal sulphides (ref. ii) and, because sulphur is hydrophobic, this results in difficulties in the selective flotation of complex sulphide ores.
On the other hand, excessive oxidation, in which
sulphur is lost from the surface as soluble oxy-sulphur ions leaving a hydrated metal oxide layer, results in inhibition of flotation (ref. 13). Chalcopyrite (CuFeS2) has been shown to float in the absence of flotation collectors under oxidizing but not reducing conditions (ref~. 14,15). behaviour has been ~ttributed oxidation. 16).
This
(ref. 15) to the formation of sulphur by surface
Other sulphides also float naturally under certain conditions (ref.
Thus the products of surface oxidation can play an important role in
determining the flotation properties of sulphide minerals. The oxidative leaching reactions of pyrite and pyrrhotite have been studied in some detail (ref. 17), these minerals being widely present as impurities in sulphide concentrates.
The conclusions reached from hydrometallurgical studies
are supported by analysis of oxidation reactions at sulphide electrodes (refs. 18-20).
However, these investigations have identified bulk oxidation products
whereas it is the surface reaction products formed initially which are important in flotation. In this communication, we present an investigation of the surface oxidation of pyrite and pyrrhotite employing linear potential sweep voltammetry.
EXPERIMENTAL The pyrite and pyrrhotite electrodes were prepared from natural massive specimens of these minerals.
The pyrite was of Spanish origin and the pyrrhotite
came from Mt Isa, Queensland, Australia.
Pyrite has a stoichiometry which is
always very close to FeS2, while pyrrhotite has a range of stoichiometries with x in Fel_xS varying from 0 to 0.14 (ref. 21).
The ratio of sulphur to iron in
the pyrrhotite employed in the present investigations was found by analysis to
329 be 1.13.
Thus, we have assumed the pyrrhotite
Selected pieces were cut from the mineral epoxy resin as described
previously
to be FeSI.13 , that is, x = 0.115.
specimens and encapsulated
(refs. 7,8).
in an
The sulphide electrodes were
mounted on a Beckman rotated electrode assembly. In order to avoid any oxidation of the electrode
surface before the experiments
were run, all operations were carried out in a nitrogen atmosphere A fresh surface was produced on the mineral
in a glove bag.
surface before each electrochemical
run by wet grinding on 600 grade silicon carbide paper in the glove bag.
The
freshly ground surface was transferred
cell
which was also contained
immediately
to the electrochemical
in the glove bag.
The glove bag and the cell solutions were purged with nitrogen purified by passage over heated copper deposited on kieselguhr. Potentials were measured and converted
against a saturated calomel reference
to the standard hydrogen
electrode
has a potential of 0.245 V against the SHE (ref. 22). are on the SHE scale.
Current-potential
and charge-potential
on a Hewlett-Packard
programmed
and serviced by a current integrator.
and integrator were designed
and constructed
relationships
(SCE)
the SCE
All potentials reported
The potential was controlled by a potentiostat
with a Utah 0151 sweep generator potentiostat
electrode
(SHE) scale, assuming
The
in these laboratories.
were recorded
simultaneously
7046A XYIY 2 recorder.
Experiments were carried out in solutions of (i) pH 4.6; 0.5 M C H 3 C O O H +
0.5 M
CH3COONa , (ii) pH 9.2; 0.05 M Na2B407 and (iii) pH 13; 0.i M NaOH. In all the figures presented
in this paper, positive currents are anodic and
negative currents cathodic.
RESULTS AND DISCUSSION For the investigation
of pyrite and pyrrhotite
techniques used in previous
studies
surfaces,
we have employed
(refs. 15,23) of sulphide surfaces.
the
These
involve oxidation of the surface during a linear potential sweep and analysis of the products of anodic oxidation on a subsequent reverse potential behaviour
from the characteristics
scan.
Interpretation
of their reduction
of the voltammetric
requires a knowledge of the possible reactions which involve the
iron sulphides
and iron-and sulphur-containing
Thermodynamics
of the Fe-S-H20
Eh - pH diagrams
species.
system
for the S-H20 , Fe-H20 , FeS2-H20 and FeI_xS-H20
10 -3 M dissolved species are presented in Fig.
i.
systems
for
These diagrams were computer
generated employing the CSIRO-NPL Thermodata System (ref. 24). The diagram for the S-H20 system, Fig. above pH 4-5.
However,
la, shows that sulphur is not stable
electrode reactions
involving sulphate are irreversible
and sulphur can be formed from oxidation of sulphide species at high pH.
We have
330
1,0
i
i
i
]
i
1,0 :
i
I
I
I
I
i
i
(b)
(o)
0.5
0,5
Eh
H
=
S
SO,~"
~
0
0
-0.5
~
-I.0
0
~
'
~ Fez+
Fe
0.5
HS-
I
J
I
I
I
I
2
4
6
8
I0
12
i
I
I
I
I
14
1.0
J
I
I
I
i
I
2
4
6
8
I0
12
I
I
I
]
I
I
pH
pH 1.0 I ~
1,0
(c)
o L ,:A Eh
~
(
(OHa )
'
(d}
0.5 Fe (OH)s
i
FeZ+
H)3
0
~
__~
-~
(;
-05
• 0.5
.Fe(OH)2
~Fe(CH)= Fe -I.0 0
;
I
I
]
L
]
2
4
6
8
I0
12
pH
14
-I.0
I
I
I
J
I
I
2
4
6
8
I0
12
14
pH
Fig. i. Eh-pH diagrams for the (a) S-H20 , (b) Fe-H20 , (e) pyrite (FeS~)-HoO o m and (d) pyrrhotite (FeSI.143)-H20 systems at 25 C and 10-3M dissolved specles. Diagrams computer generated employing the CSIRO-NPL Thermodata system (ref. 24).
331
included in the S-H20 diagram the boundary between metastable sulphur and HS(dashed line) in addition to the thermodynamically more favourable equilibria. The diagram also indicates that sulphate in solution could be reduced to H2S or HS- at potentials within the range covered by the voltammograms.
However, this
process is highly irreversible and has a negligible rate at these potentials. Peters (ref. 25) points out that zinc electrowinning from sulphate solutions takes place at an overpotential of % 1 V
with respect to H2S formation, but no
trace of this compound is detected in practice. The Fe-H20 diagram, Fig. ib, includes the hydrated iron oxides, since they are more likely to be formed in the present environment than anhydrous oxides. The pyrite diagram, Fig. ic, shows the domain of stability of this mineral. The sulphur-containing products are dissolved sulphide species below, and sulphate ions abovejthe pyrite domain, i.e., in the stability regions of these species in Fig. la.
Oxidation of sulphides to sulphate generally requires high overpotentials
(ref. 9) and occurs at significant rates only at potentials at which sulphur can also be a product.
Hence, the boundary to which pyrite is metastable before
oxidation to Fe 2÷ or Fe(OH) 3 and S is included (dashed lines) in addition to the thermodynamically favoured equilibria. Figure id shows the Eh-pH diagram for pyrrhotite. not been included.
Oxidation to pyrite has
Pyrite can be formed in hydrothermal systems involving H2S ,
e.g., it forms through a series of iron sulphides in the corrosion of iron (refs. 26,27).
However, the oxidation step from pyrrhotite to pyrite is slow even at
100°C and occurs through a dissolution, reprecipitation mechanism rather than a transformation of the solid (ref. 28).
Furthermore, the reaction of pyrrhotite
with dissolved sulphide species is not relevant to the interpretation of the voltammograms presented here, since no sulphide species were present in solution at the commencement of the potential sweeps.
Conversion to pyrite by a reaction
such as,
2Fel_xS
÷
~l-2x)Fe 2+ + FeS 2 + 2(l-2x)e
(i)
has been neglected since there is no evidence of pyrite formation during the leaching of pyrrhotite (ref. 29).
In fact, reactions in nature which give rise
to pyrite are known to be particularly slow (ref. 30).
The anodic reactions
presented are for the formation of sulphate and metastable sulphur (dashed lines). The data in Fig. id correspond to pyrrhotite with composition FESI.143; the diagram for FeSI.13 will not be significantly different.
Oxidation of pyrite Voltammograms for a stationary pyrite electrode are shown in Fig. 2.
The
potential sweeps were commenced at the open circuit potential marked * and each
332 I
I
,
I
,
']
2.0
20 1.5 f
J
1
u
f
f
E
E
"~
1.0
==
g
I0
I
I
0.5
)
8 5
- J~
4
I
-0.2
- 0.4
m 2 I
I
I
I
0
0.2
0.4
0.6
Potential
I 0.8
V vs SHE
Fig. 2. Voltammograms for a stationary pyrite electrode at pH 4.6. Linear potential sweeps at 20 mV s -I reversed at different upper potential limits. Each v o l t a m m o g r a m relates to a freshly ground surface. Dashed curve is the recorded charge on the voltammogram, curve 4.
sweep was carried out on a freshly ground surface. positive
going scan,
It can be seen that,
for a
there is a steady rise in anodic current as the potential
is increased. It is apparent Fig.
FeS 2
2 to either
÷
from Fig.
ic that pyrite can be oxidized at the potentials
in
sulphur,
Fe 2+ + 2S + 2e
(2)
or sulphate,
FeS 2 + 8H20
÷
Fe 2+ + 2S042- + 16H + + 14e
There is a step on the curve at ~0.4 V. (see Fig.
(3)
This can be assigned
to the reaction
ib),
Fe 2+ + 3H20
÷
Fe(OH) 3 + 3H + + e
(4)
o
333 Investigation
of the effect of electrode rotation enables a distinction
to be
made between reactions which result in soluble products and those in which only surface species are produced
(refs.
15,23,31).
For a stationary electrode,
dis-
solved products remain in the vicinity of the surface and can take part in further electrochemical are dispersed
reactions.
When the electrode
is rotated,
the soluble products
as they are formed.
The step assigned This observation
to reaction
(4) is absent on rotating
is consistent with reaction
the pyrite electrode.
involving dissolved
species.
cathodic peak at %0.25 V, which is also absent when the electrode arises from the reverse of reaction
(4).
reduction of iron oxide on chalcopyrite It is to be expected state (Fig.
ib).
A similar peak is observed for the (ref.
(v.i.).
15) and pyrrhotite
that all iron formed above ~0.4 V will be in the ferric
However,
the charge under the cathodic peak at %0.25 V is insuf-
ficient to account for all the iron species arises from dissolution
The
is rotated,
formed anodically.
This probably
of Fe(OH) 3 at this pH to form species such as FeOH 2+ and
Fe(OH)2+ which can diffuse away from the electrode The cathodic peak at lower potentials
surface.
can be assigned
sulphur which had been formed by reaction
(2).
to the reduction of the
In acid solution,
the reduction
will produce H2S.
S + 2H + + 2e
÷
H2S
(5)
When the scan is reversed again at -0.35 V (Fig. 2), the H2S is oxidized back to sulphur by the reverse of reaction 0 V.
(5) and this gives rise to the anodic peak at
The effect of electrode rotation supports
this view.
unaffected but the anodic peak is eliminated because
The cathodic peak is
the H2S produced
cathodically
is dispersed. The charge associated with the lower cathodic peak is a measure of the quantity of sulphur formed anodically by reaction sum of reactions
(2),
(3) and
(4).
(2).
The anodic charge is due to the
If we assume that the iron product
ferrous state below 0.4 V and ferric above this potential, to determine
is in the
then it is possible
the number of moles of pyrite oxidized to sulphur and to sulphate
from integration
of the anodic and cathodic
charges.
An example of the recording of the charge is shown in Fig. 2. anodic charge is given by the maximum on the charge-potential
The total
curve.
The Fe(OH) 3
is given by the step at 0.25 V and the sulphur by the difference between the charge at 0.15 V and the minimum on the second positive going scan. Reduction (Fig.
of pyrite to a lower iron sulphide is expected from the Eh-pH diagram
Ic) and it was not possible
reduction
to set a lower scan limit which allowed complete
of sulphur on the negative-going
reduction of the pyrite itself.
sweep without causing considerable
However a lower limit can be chosen such that
334 reduction of sulphur is completed during the sulphur peak includes cathodic scan.
The sulphur charge was corrected
tion of the mineral by subtracting
the return scan.
Thus integration
charge passed at the beginning
of
of the reverse
for the small amount of background
reduc-
the charge passed during a negative-going
applied to a freshly ground pyrite surface
sweep
(curve i, Fig. 2).
The quantity of pyrite oxidized to sulphur and to sulphate on a potential sweep is presented as a function of the upper potential
limit of the sweep in
Fig. 5. In basic solution,
oxidation of pyrite takes place at potentials
stability domain of hydrated
in the
ferric oxide and hence the oxidation reactions
will be,
FeS 2 + 3H20
+
(6)
Fe(OH) 3 + 2S + 3H + + 3e
and
FeS 2 + IIH20
+
The reduction Figs.
Fe(OH) 3 + 2S042- + 19H + + 15e
(7)
of both Fe(OH) 3 and S occurs in a similar potential range
(see
la and ib) and hence a single cathodic peak appears on the voltammogram
(Figs. 3 and 4).
It is unlikely that HS- would be released
iron oxides are present on the mineral that an iron sulphide will be produced.
surface.
Rather,
to the solution since
it is to be expected
The cathodic peak occurs at potentials
at which pyrite itself is reduced to give a small but significant i, Fig. 3).
Thus, the likely product of sulphur reduction
current
iron oxides is FeS.
This conclusion
product precipitated
from aqueous solutions of iron and sulphide species.
There is negligible 4) when the electrode involve solid species. oxidation
substantiating
at pH 9.2 and 13 (Figs. 3 and
the conclusion
that all reactions
We have taken the reduction of the products of anodic
to be represented
Fe(OH) 3 + S + 3e
is supported by the fact that FeS is the
effect on the voltammograms is rotated,
(curve
in the presence of
by,
*
FeS + 3OH-
(8)
+
Fe(OH) 2 + H20
(9)
and
Fe(OH) 3 + H + + e
If we assume that x moles of pyrite are oxidized sulphate by reactions the potential
(6) and
(7) respectively,
to sulphur and y moles to
then the anodic charge passed on
sweep will be (3x + 15y)F coulombs.
The cathodic charge arising
335 I
I
I
I
I
1
1.5
i%
I.O -
o
6
E >,
•"o
0.5
C
8 6 0
6 I
I
-0.4
-0.2
[
I
I
I
0
0.2
0.4
0.6
Potential
V vs SHE
Fig., 3, Voltammograms for a stationary pyrite electrode at pH 9.2. Linear potential sweeps at 20 mV s -I reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
from reaction
(8) will be 6xF coulombs.
moles of which 2x are consumed
The total Fe(OH) 3 produced
in reaction
(8).
Thus the charge due to reaction
(9) will be (y - x)F and the total cathodic charge
(Sx + y)F.
for the oxidation of pyrite at pH 9.2 and 13, determined are also presented
is (x + y)
Values of x and y
from these equations
in Fig. 5.
It can be seen from Fig. 5 that, in all three solutions, product is sulphur with very little sulphate being formed. proportion of sulphate produced
increases as the potential
the initial oxidation In each case, the is taken to higher
values. At pH 4.6, the sulphur yield also increases with increase in potential,
but
sulphate becomes the dominant product above 0.8 V. At pH 9.2 and 13, the sulphur yield remains constant at ~0.45 x 10 -8 moles cm -2 , while the sulphate produced
increases rapidly.
This suggests that the
first anodic peak (0 V, pH 9.2; -0.2 V, pH 13) constitutes with the steady increase in current at high potentials sulphate.
being oxidation to
The charge passed in the production of the sulphur layer is similar
to that found assigned
sulphur formation
(ref. 23) for a prewave
to monolayer
in the oxidation of galena which was
formation of PbO + S.
It is reasonable
to assume that the
336 [
I
f
I
I
2.0
1.5
0
<~ E
1.0
"tD C L
8
0.5
0
-_iji;5
-0.6
;
I
-0.4
i
-0.2 Potenfiol
J
0
04
0.2
V vs S H E
Fig. 4. Voltammograms for a stationary pyrite electrode at pH 13. Linear potential sweeps at 20 mV s -I reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
sulphur layer on pyrite is also of the order of a monolayer There appears to be a decrease potentials
on the surface.
in the quantity of sulphur present at high
in pH 9.2 and 13 solutions.
However,
this could arise from under-
estimation of the cathodic charge due to either the background reduction
of the
mineral being inhibited by the presence of surface oxides, or failure to reduce all the oxidation products before reversal of the sweep slows the reduction
to a
negligible rate.
Oxidation of pyrrhotite It is well established
(ref. 17) that pyrite oxidizes entirely to sulphate and
sulphur during leaching processes formed. ions.
and that no other oxy-sulphur
On the other hand, pyrrhotite However,
the differential
species are
is known to give rise to lower oxy-sulphur
increase in anodic charge on potential
sweeps to
337 I
%
I
I
I
2.0
o
1.5
/7
1.0
,
f~
0.5
•
. el 2 ' o S '
0 0
0.2 0.4 Potential Limit V vs SHE
0.6
0.8
-i Fig. 5. Quantity of pyrite oxidized to sulphur and sulphate on a 20 mV s potential sweep as a function of the upper potential limit of the sweep: (i), (i') pH 4.6, (2), (2') pH 9.2, (3), (3') pH 13; (i) O , (2) ~, (3) h , sulphur, (i') O, (2') I, (3') A , sulphate.
high potentials in alkaline solution (>0.7 V) is close to 10 times the corresponding increase in cathodic charge, which is the value expected if pyrrhotite is oxidized to sulphate (see reaction 13).
It is possible that the other oxy-
sulphur anions, formed in the leaching of pyrrhotite in ammoniacal and caustic solutions, arise from air oxidation of sulphur which is a product of the anodic oxidation of the mineral at lower potentials
(v.{.).
Voltammograms for a stationary pyrrhotite electrode at pH 4.6 are presented in Fig. 6.
We have analyzed these curves in terms of the anodic reactions,
FeSl.13
÷
Fe2+ + 1.13S + 2e
(i0)
*
Fe(III) + 1.13S + 3e
(li)
or FeSI.13
and
F e S l . 1 3 + 4.52H20
÷
Fe 2+ + 1.13S042- + 9.04H+ + 8.7Se
(12)
338
2.0
--
•
r
I
i
I
i
I0-
?
oE E
o
8 -I0
-
-2.0
-
I
-0.4
-02
I
I
I
I
0
0.2
04
06
Potential
08
V vs SHE
Fig. 6. Voltammograms for alstationary pyrrhotite electrode at pH 4.6. Linear potential sweeps at 20 mV s reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
or
FeSl.13 + 4.52H20
÷
Fe(lll) + 1.13SO42- + 9.04H + + 9.78e
(13)
The peak at 0.43 V in Fig. 6, which is absent when the electrode is equivalent of Fe 2+.
to the step observed with pyrite which is assigned
is rotated,
to the oxidation
The cathodic peak at ~0.2 V arises from reduction of Fe(0H) 3 by the
reverse of reaction
(4).
The charge associated with this peak does not account
for all the iron species formed above 0.4 V so that some of the iron(III) must dissolve.
The peak is not eliminated
diminished
to some extent.
on rotating the electrode,
this peak is formed directly and is not produced As with pyrite,
from a dissolved
the sulphur formed anodically
it is
is reoxidized
second positive going scan unless it is dispersed by rotating to the voltammograms
same manner as for pyrite in order to determine
species.
is reduced at the lower end of
the reverse scan and the H2S produced by this reduction
The charges corresponding
although
This suggests that the oxide which gives rise to
in Fig.
on the
the electrode.
6 were analyzed
in the
the number of moles of pyrrhotite
339 oxidized to sulphur and to sulphate. were assumed to arise from reactions
That is, the anodie currents below 0.4 V (10) and (12) and above this potential,
to
reactions
(ii) and (13), with the cathodic peak at low potentials arising from
reduction
of the product
Voltammograms Fig. 7.
sulphur by reaction
for stationary pyrrhotite
(5).
electrodes
at pH 9.2 are presented
Rotated electrodes produce almost identical curves.
at this pH are due to reactions
(ii) and (13), the Fe(lll)
in
The anodic currents
species being Fe(OH) 3.
The processes which give rise to the cathodic peak on the reverse scan will be the same as those for pyrite,
i.e., reactions
(8) and (9).
The first anodic
wave on the second positive going scan arises from re-oxidation produced on the preeeeding reverse scan by reaction
(9).
of the Fe(OH) 2
The peak at ~0.05 V
we assign to oxidation of the FeS formed on the reverse scan by reaction It is more active than the original pyrrhotite, to be rapidly inhibited by the products,
cathodic peak.
Although the same processes
pyrite and pyrrhotite
I
surfaces,
I
but the oxidation process seems
ferric oxide and sulphur,
total anodic charge to 0.2 V is significantly
(8).
since the
less than that of the preeeeding
have been assumed to take place at
this anodic peak is not clearly discernable
I
I
1
on
I
2.0
4
1.0 E
o 8 -I.0
I -0.4
I -0.2
I 0 Potential
I 0.2
I 0.4
I 0.6
V vs SHE
Fig. 7. Voltammograms for alstationary pyrrhotite electrode at pH 9.2. Linear potential sweeps at 20 mV s- reversed at different upper potential limits. Each voltammogram relates to a freshly ground surface.
0.8
340
volEammograms
for the former mineral.
This is because the quantity of sulphur,
and hence of FeS, formed on pyrite is much less than on pyrrhotite. The quantity of pyrrhotite
oxidized
in the same manner as for pyrite.
to sulphur and to sulphate can be obtained
Thus, if x moles form sulphur and y sulphate,
the anodic charge will be (3x + 9.78y)F coulombs and the cathodic,
(3.26 + y)F.
The latter value arises because 3.39e are involved in the reduction to FeS of the 1.13x moles of sulphur and the remaining Fe(OH)3,
viz.
(y - 0.13x) moles,
are reduced to Fe(OH) 2. Voltammograms presented
for stationary and rotated pyrrhotite
in Fig. 8.
at t0 V, followed by a steeply increasing
current.
results in inhibition of the anodic process. from the thermodynamic
electrodes at pH 13 are
The anodic oxidation curves are characterized
instability
I
I
Rotation of the electrode
This phenomenon probably results
of pyrrhotite
I
by a peak
at high pH (Fig. id).
Thus,
I
I
I
I
I
- 0.4
- 0.2
I 0
I 0.2
3.0
2.0
?
5 E
==
|.0
--
C
8 0
f -I.0
3j
J
__J I -0.8
i -0.6
Potential
V vs SHE
Fig. 8. Voltammograms for a pyrrhotite electrode at pH 13. Linear potential sweeps at 20 mV s -I with electrode - - stationary and - - - rotated at i00 rps. Each voltammogram relates to a freshly ground surface.
341 part of the reaction could proceed via dissolved
sulphide
dispersed
on stirring and this would lead to a decrease
current.
The observation
rotated electrode
is diminished,
but the following anodic peak due to oxidation constant,
would support a loss of sulphur with-
in Fig, 8, the anodic currents are greater on the
reverse than on the initial positive-going for chalcopyrite
electrodes.
peak for the
loss of iron on stirring.
At the higher potentials
observed
in the overall anodic
(Fig. 8) that the cathodic reduction
of Fe(OH) 2 remains approximately out a corresponding
species which will be
(ref.
scan.
18), pyrite
It can arise from nucleation
(ref.
This behaviour
has also been
19) and galena
(ref. 23)
of active sites or pits which act as
growth centres for further reaction. The voltammograms
at pH 13 were analyzed
in the same manner as at pH 9.2 to
determine the quantity of sulphur and sulphate produced The quantity of pyrrhotite oxidized
to sulphur and to sulphate is presented
in Fig. 9 as a function of the potential
I
1
anodically.
limit of the sweep.
1
It can be seen
I
I
o?
6.0 L) (31.
% x
¢1
4.0
0 O
¢-
0.. (/)
2.0
-
O
/
O 0
0.2
0.4
0.6
I_ 0.8
Potential Limit V vs SHE -i Fig. 9. Quantity of pyrrhotite oxidized to sulphur and sulphate on a 20 mV s potential sweep as a function of the upper potential limit of the sweep: (i), (i') pH 4.6, ( 2 ) , ( 2 ' ) p H 9.2, ( 3 ) , ( 3 ' ) p H 13; (i) O, (2) D , (3) A , sulphur, (i') Q , (2') • , (3') • , sulphate.
342
that sulphur is the major oxidation product in each solution at all potentials investigated and initially accounts for all the pyrrhotite oxidized.
The amount
of sulphate produced increases as the sweep is taken to higher potentials. However, part of the current at the higher potentials in Figs. 7 and 8 still goes to produce sulphur, the sulphur yield continually increasing with potential. We have found that sulphate is the dominant product on sweeps taken to much higher potentials. Oxidation of pyrrhotite is inhibited significantly in alkaline solution, suggesting that iron oxide on the surface retards further reaction.
The anodic
oxidation curves for pyrite are also less steep in the basic media than at pH 4.6 (Figs. 4-6), but the inhibition for this mineral is not large since a regular shift in the anodic wave is observed of ~60 mV per pH unit.
On the other hand,
the oxidation of pyrrhotite at pH 9.2 is actually less than at pH 4.6 at the same potential on the SHE scale, even though the equilibrium potentials of the anodic reactions
(Ii) and (13) (with Fe(OH) 3 as the Fe(III) species) shift to
lower potentials with increase in pH.
The difference between the behaviour of
the two iron sulphides regarding the inhibiting effect of the product iron oxide is also apparent from a comparison of Figs. 5 and 9.
It could arise from different
structures of the hydrated ferric oxide species formed.
Pyrite is known (ref. 17)
to oxidize in alkali to form y-Fe203 which covers the surface but allows further oxidation thromgh a solid-state topotactlc reaction, whereas pyrrhotite becomes coated by a layer of ferric oxide which does not contain any well-defined oxide structure.
REFERENCES 1 P . H . Ribbe, (Ed.), Sulfide Mineralogy, Mineralogical Society of America, Washington, 1974. 2 M.J. Taylor, CIM Bulletin, 70 (1977) 78. 3 G.M. Lukaszewski, in M.J. Jones (Ed.), Mining and Petroleum, I.M.M., London, 1970 pp. 803-19. 4 R.L. Hewett, Proc. Aust. Instn. Min. Met., 226 (1968) 73. 5 D.J.M. Farnsworth, CIM Bulletin, 70 (1977) 65. 6 D. Michel_land R. Woods, Aust. J. Chem., 31 (1978) 27. 7 T. Biegler, D.A.J. Rand and R. Woods, J. Electroanal. Chem., 60 (1975) 151. 8 T. Biegler, D.A.J. Rand and R. Woods, in J.O'M. Bockris, D.A.J. Rand and B.J. Welch (Eds.), Trends in Electrochemistry, Plenum, N.Y., 1977, pp. 291-302. 9 E. Peters, in J.O'M. Bockris, D.A.J. Rand and B.J. Welch (Eds.), Trends in Electrochemistry, Plenum, N.Y., 1977, pp. 267-90. i0 D.A.J. Rand, J. Electroanal. Chem., 83 (1977) 19. ii H. Majima, Canad. Met. Q., 8 (1969) 269. 12 H.A. Buehler and V.H. Gottschalk, Econ. Geol., 5 (1910) 28. 13 N. Arbiter, The Physical Chemistry of mineral-reagent interactions in sulfide flotation, proceeding of symposium held at College Park, Md., 1978 U.S.B.M. ICC 8818, pp. 144-9. 14 G.W. Heyes and W.J. Trahar, Int. J. Miner. Process., 4 (1977) 317. 15 J.R. Gardner and R. Woods, Int. J. Miner. Process., 6 (1979) i. 16 G.W. Heyes and W.J. Trahar, Int. J. Miner. Process., 6 (1979) 229. 17 A.R. Burkin, The Chemistry of Hydrometallurgical Processes, E. & F.N. Spon, London, 1966.
343
18 T. Biegler and D.A. Swift, J. Appl. Electrochem., 9 (1979) 545. 19 T. Biegler and D.A. Swift, Electrochim. Acta, 24 (1979) 415. 20 L.K. Bailey and E. Peters, Can. metall. Q., 15 (1976) 333. 2 1 G . Kullerud, P.H. Abelson (Ed.), Researches in Geochemistry, J. Wiley, Chichester, Vol. 2, 1967, pp. 236-333. 22 R.G. Bates, Determination of pH, Wiley, N.Y., 1964, pp. 458-483. 23 J.R. Gardner and R. Woods, J. Electroanal. Chem., i00 (1979) 447. 24 A.G. Turnbull, Chem. in Aust., 44 (1977) 334. 25 E. Peters, Internal Report, U.B.C. 26 J.S. Smith and J.D.A. Miller, Br. Corros. J., I0 (1975) 136. 27 A.G. Wikjord, T.E. Rummery and F.E. Doern, Can. Mineral., 14 (1976) 571. 28 P. Taylor, T.E. Rummery and D.G. Owen, J. Inorg. Nucl. Chem., 4 (1979) 1683. 29 E. Peters, in D.J.I. Evans and R.S. Shoemaker (Eds.), Intern. Symp. Hydromet., AIME, 1973, pp. 205-28. 30 P.B. Barton and B.S. Skinner, in H.L. Barnes (Ed.), Geochemistry of Hydrothermal Ore Deposits, Holt, Rinehart and Winston, N.Y., 1967 pp. 236-333. 31P.E. Richardson and E. Maust Jr., in M.C. Fuerstenau (Ed.), Flotation, A.M. Gaudin Memorial Volume, AIME, N.Y., 1976, Vol. i, Ch. 12.