ish Columbia, 309-6350 Stores Road, Vancouver, BC, Canada, V6T 1Z4

species (e.g., thiosulfate) that form as intermediate products of pyrite oxidation. In order to account quantitatively for this

(200D 2003 Elsevier B.V. All rights reserved.

Keywords: Pyrite; Pressure oxidation; Leaching; Kinetics; Passivation; Autoclave; Refractory gold; Hydrometallurgy

1. Introduction often viewed as a gangue mineral. However, it often

influences the recovery of associated metal valuesthe entire range.passivation phenomenon, a new passivating shrinking sphere model is proposed which fits the conversion data precisely overReceived 25 March 2002; received in revised form 15 July 2003; accepted 22 July 2003

Abstract

The oxidation kinetics of a massive pyrite (FeS2) sample from Zacatecas, Mexico were investigated in sulfuric acid solution

under oxygen pressure. The effects of temperature (170230jC), particle size (49125 Am diameter), agitation speed (650950 rpm), oxygen partial pressure (3451035 kPa) and pulp density (120 g/L) were evaluated. The catalytic effect of Cu(II)

was also observed.

Fe(III) was found to be the initial product of pyrite oxidation, although the proportion of total dissolved iron as Fe(III) did

reach a minimum during each test, indicating the generation and subsequent oxidation of reduced sulfur species. Pyrite

oxidation kinetics are limited by the rate of reaction at the pyrite surface, with an activation energy of 33.2 kJ/mol (7.9 kcal/mol)

with respect to dissolved oxygen concentration, or 41.7 kJ/mol (10.0 kcal/mol) with respect to oxygen partial pressure, over the

temperature range 170230jC. The reaction order with respect to oxygen partial pressure was found to be 0.5 at 210jC,indicating the first charge transfers of both the anodic and cathodic half-cell reactions to be the most likely rate-controlling

steps.

Conversion data conform to the shrinking sphere model initially, but deviate at higher conversions, indicating passivation of

the mineral surface, most likely by elemental sulfur, which may precipitate via the disproportionation of other reduced sulfurDepartment of Metals and Materials Engineering, University of BritPressure oxidation of pyrite in sulfuric acid media:

a kinetic study

Hu Long, David G. Dixon*Hydrometallurgy 73Pyrite (FeS2) is the most common iron sulfide

mineral. In most pyrite ores and concentrates, the

pyrite itself is rarely of economic importance and is

0304-386X/$ - see front matter D 2003 Elsevier B.V. All rights reserved.

doi:10.1016/j.hydromet.2003.07.010

* Corresponding author. Tel.: +1-604-822-3679; fax: +1-604-

822-3619.

E-mail address: dixon@interchange.ubc.ca (D.G. Dixon).www.elsevier.com/locate/hydromet

4) 335349such as gold, copper and zinc. In many sulfidic

refractory gold ores and concentrates, gold is often

found finely disseminated in sulfide minerals, par-

ticularly in pyrite (Gasparrini, 1983), rendering the

adequate liberation of gold impossible by grinding

alone. One method for recovery of gold from these

refractory gold ores and concentrates is pressure

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349336oxidation in sulfuric acid solution at high temper-

atures (above 180jC) and oxygen partial pressuresas a pretreatment prior to cyanidation. This process

has been practiced commercially since the early

1980s (Argall, 1986), with the object to enhance

exposure of the gold particles to the cyanide solu-

tion by breaking down the sulfide lattice.

Pyrite oxidation has been studied extensively

because of its importance in sulfide mineral separa-

tions by flotation, in the generation of acid in mine

waters and in leaching. Aqueous oxidation of pyrite

has been reviewed thoroughly by Lowson (1982)

and Hiskey and Schlitt (1982). In most of the

pressure oxidation experiments reported (McKay

and Halpern, 1958; Gerlach et al., 1966; Bailey

and Peters, 1976; Papangelakis and Demopoulos,

1991), the oxidation of pyrite was found to yield

only the following products: ferrous sulfate, ferric

sulfate, sulfuric acid and elemental sulfur. No sulfur

products of intermediate oxidation state, such as

thiosulfate (S2O32) or thionates (SnO6

2, n= 26),were detectable under any conditions, although

significant evidence exists to suggest that sulfur

does indeed pass through such intermediate states

during pyrite oxidation at lower temperatures (Con-

way et al., 1980; Goldhaber, 1983; Morse et al.,

1987; Kelsall and Yin, 1996). The pressure oxida-

tion of pyrite is typically represented by the follow-

ing two competing reactions (Bailey and Peters,

1976):

FeS2 7=2O2 H2O ! FeSO4 H2SO4 1

FeS2 2O2 ! FeSO4 S 2

As temperature increases, reaction (1) becomes

predominant. The Fe(II) produced is subsequently

oxidized to Fe(III):

2FeSO4 1=2O2 H2SO4 ! Fe2SO43 H2O 3

Ferric sulfate has been reported as an importantoxidizing agent for pyrite (Lowson, 1982; Garrels andThompson, 1960). The oxidation reactions may be

represented as follows:

FeS2 7Fe2SO43 8H2O ! 15FeSO4 8H2SO44

or

FeS24Fe2SO434H2O! 9FeSO44H2SO4S5

At high temperatures, ferric sulfate tends to hydro-

lyze, resulting in the precipitation of ferric oxide

(Fe2O3) or basic ferric sulfate (Fe(OH)SO4) depend-

ing on acidity (Tozawa and Sasaki, 1986).

Table 1 lists the experimental conditions, the orders

of reaction for oxygen partial pressure, and the

reported activation energies from the most relevant

studies on pressure oxidation of pyrite in sulfuric acid

solution at temperatures above 120jC. A number ofconclusions may be drawn from Table 1:

1. It is accepted by all researchers that the

pressure oxidation of pyrite is controlled by the

surface reaction rate. Reaction orders with respect

to oxygen partial pressure depend on both tempera-

ture and pressure. First-order dependence is exhibited

predominantly at lower oxygen partial pressures

( < 20 atm) and at all temperatures employed, sug-

gesting a mass transfer limitation. One-half order

dependence is found at higher oxygen partial pres-

sures and temperatures.

2. McKay and Halpern (1958), Gerlach et al.

(1966) and Cornelius and Woodcock, (1958) inter-

preted their experimental results as evidence of an

oxygen chemisorption mechanism followed by a slow

chemical reaction. However, Bailey and Peters (1976)

convincingly demonstrated the mechanism of pressure

oxidation of pyrite to be electrochemical, involving

coupled anodic (pyrite oxidation) and cathodic (oxy-

gen reduction) reactions.

3. Activation energies between 46 and 55 kJ/mol

are reported at temperatures below 160jC. Interest-ingly, Papangelakis and Demopoulos (1991) report an

activation energy of 110.5 kJ/mol over the range of

160180jC, which is more than twice the value

reported by them and others at temperatures below

rNatural pyrite 0.5 140160 520 1

160180 510 1

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349 337160jC. The reason for this significant shift in activa-tion energy is unclear.

Although extensive investigations have been per-

formed by previous researchers, all of the reported

work has been conducted at temperatures lower than

180jC. No data is available on the kinetics of pyritedissolution during acid pressure oxidation over the

temperature range from 180 to 230jC, which is therange employed by most commercial plants. From a

process optimization standpoint, it is important to

know the behavior of pyrite during acidic pressure

oxidation at temperatures above 180jC; hence, thenecessity of the present study.

The main objectives of our investigation were to

gather data on the pressure oxidation of pyrite over the

typical temperature range employed in the commercial

160180 1020 0.5Table 1

Review of pyrite pressure oxidation studies

Material Experimental conditions PO2

[H2SO4], M T, jC PO2, atmorde

Upgraded pyrite

concentrate

0.075 100130 04 1

Natural pyrite 0.2 60130 015.5 1

Natural pyrite 1.0 85130 020 1

2066.4 0.5pressure oxidation of refractory gold ores, and to

develop a reliable rate equation which may be used

to model the pressure leaching of refractory gold ores.

2. Experimental

2.1. Materials

High-grade massive pyrite specimens originating

from Zacatecas, Mexico were obtained from Wards

Natural Science Establishment, of Ontario, Canada.

The samples were crushed, ground and dry-sieved to

different narrow size fractions, namely 149 + 105, 105 + 74, 74 + 53 and 53 + 44 Am. The sam-ples were then washed to remove any fines adheringto particle surfaces which might interfere with the

interpretation of the results, and air-dried at ambient

temperature.

The samples were analyzed by two analytical labs

in order to obtain reliable assays. The results indicated

that the purity of the pyrite sample was 97F 1%, thatthe mole ratio of sulfur to iron was 1.92, which is very

close to the stoichiometric value for pyrite, and that

iron and sulfide sulfur composition in each size

fraction showed very little variation, with maximum

differences of only 0.6% for iron and 1% for sulfide

sulfur. The pyrite crystal structure was confirmed by

X-ray diffraction.

All solutions were prepared with reagent grade

chemicals and deionized water. Medical grade oxygen

gas was used in all experiments.

Activation Assumed mechanism References

energy,

kJ/mol

55.7 Chemisorption +

chemical reaction

McKay and Halpern, 1958

54.8 Chemisorption +

chemical reaction

Gerlach et al., 1966

51.1 Electrochemical Bailey and Peters, 1976

reaction

46.2 Electrochemical Papangelakis and

110.5 reaction Demopoulos, 19912.2. Apparatus

All pressure oxidation experiments were conducted

in a 2-L Parr titanium autoclave. The temperature was

monitored using a thermocouple probe to an accuracy

of F 1jC and maintained at desired set points byexternal heating and by passing cooling water through

the internal cooling coils. Cooling water flow was

regulated by a solenoid valve. Agitation was provided

by dual four-pitched-blade impellers. Temperature and

agitation speed were maintained using a Parr control-

ler unit. Oxygen was introduced into the autoclave

near the bottom through a dip tube, which was also

connected to an autoclave sampler. Heat resistant

superalloy valves were selected to avoid any corrosion

in contact with oxidizing acid solution.

2.3. Experimental conditions

The slurry used for most experiments had a stan-

dard composition of 1 g/L pyrite and 0.5 M H2SO4.

Based on the results of Tozawa and Sasaki (1986) and

on our own preliminary tests, the acidity was fixed at

0.5 M H2SO4 for all the experiments in order to obtain

a high solubility of Fe2O3 and to avoid Fe(OH)SO4precipitation, since dissolved iron was taken as the

indicator of pyrite dissolution at low pulp density. The

proper amount of pyrite to charge into the autoclave

was determined by a series of tests at different pulp

densities at 230jC. The results showed that when 1 g/L FeS2 was charged, even if the pyrite was dissolved

completely, the amount of precipitated iron (owing

mostly to solution splashing onto the heated walls of

the reaction vessel above the solution line) corre-

sponded to less than 5% conversion of the pyrite. At

lower temperatures the precipitated amount was con-

siderably less, so 1 g of FeS2/L of solution was used

in all low pulp density experiments. At higher pulp

densities, total mass flow of oxygen gas to the

autoclave was taken as the indicator of pyrite oxida-

tion, and so the extent of iron precipitation was of no

concern.

2.4. Test procedure

It is well known that titanium is corroded in high

temperature sulfuric acid solutions in the absence of

an oxidizing agent (i.e., Cu(II), Fe(III) or O2) (Uhlig

and Revie, 1985). Therefore, for the experiments in

the absence of Cu(II) at low pulp density, concen-

trated sulfuric acid was sealed in ampoules during

O2,

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349338Fig. 1. Effect of temperature on pyrite oxidation (800 rpm, 690 kPa

vs. r2, (d) [Fe(III)]/[FeT] vs. time.

74 + 53 Am). (a) Conversion vs. time, (b) r vs. time, (c) s(dr/dt)

autoclave heating to avoid the corrosion of titanium,

while for the experiments in the presence of Cu(II),

pyrite was sealed in an ampoule to avoid premature

oxidation.

A slurry with several borosilicate glass ampoules

containing sulfuric acid, or a solution with one boro-

silicate glass ampoule containing pyrite, was placed in

the autoclave, which was then sealed and heated.

Once the desired temperature was reached, an initial

sample was taken, then the unit was pressurized with

oxygen for 1 min. Finally, the ampoules were broken

by engaging the impellers and the test was thus

initiated.

In the tests at low pulp density, slurry samples were

withdrawn periodically from the autoclave through

the dip tube. The withdrawn sample was cooled

immediately (in less than 5 s) to room temperature

and filtered. At the end of the run, the system was

cooled quickly (in less than 15 min) to room temper-

ature with the cooling coil. The slurry was removed

and filtered, the residue was washed and dried, and

both solids and solution were analyzed.

In order to avoid contamination from any iron

residue precipitated onto the walls of the bomb and

internal parts, the bomb was soaked after each exper-

iment with 1.5 L of 15% HCl solution for 30 min at a

temperature of 80jC, then cooled to room tempera-ture. The solution was analyzed for total iron to

determine the extent of precipitation. Then, after being

washed with water, the bomb was soaked overnight in

kPa

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349 339Fig. 2. Effect of particle size on pyrite oxidation (800 rpm, 210jC, 690

[Fe(III)]/[FeT] vs. time.O2). (a) Conversion vs. time, (b) r vs. time, (c) s(dr/dt) vs. r2, (d)

a solution containing 10% HNO3 and 5 g/L iron as

ferric sulfate in order to encourage the formation of a

TiO2 film. Blank tests run on 1 L of 0.1 M H2SO4solution under conditions of 230jC and 800 rpmwithin the cleaned bomb showed that the total iron

concentration was less than 0.14 mg/L after 1 h.

2.5. Analytical methods

Solution samples from the low pulp density tests

were analyzed for Fe(II) by redox titration with

Ce(SO4)2 solution using ferroin as an indicator (Jeff-

ery et al., 1989), and for total iron by atomic absorp-

tion spectrophotometry (AAS). Fe(III) was determined

as the difference between total iron and Fe(II). At

higher pulp densities, total oxygen flow was recorded

digitally from Omega in-line digital oxygen mass

flowmeters.

3. Results and discussion

Pressure oxidation tests were performed under the

following conditions: agitation speeds from 650 to

950 rpm, temperatures from 170 to 230jC, meanparticle sizes (diameters) from 49 to 125 Am, oxygenpartial pressures from 345 to 1035 kPa (50150 psi)

and pyrite pulp densities of 1 and 20 g/L. The mean

particle size was taken as the square root of the

product of two adjacent sieve sizes (i.e., the geometric

pm, 2

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349340Fig. 3. Effect of oxygen partial pressure on pyrite oxidation (800 r s(dr/dt) vs. r2, (d) [Fe(III)]/[FeT] vs. time.

10jC, 74 + 53 Am). (a) Conversion vs. time, (b) r vs. time, (c)

mean sieve opening). The standard conditions for the

leaching experiments were as follows: agitation speed

of 800 rpm, temperature of 210jC, O2 partial pressureof 690 kPa (100 psi), particle size of 74 + 53 Am(62.6 Am mean size) and pyrite pulp density of 1 g/Lin 0.5 M H2SO4 solution.

All of the experimental results are presented in

Figs. 14 and all experimental conditions and calcu-

lated parameters are presented in Table 2. The effects

of temperature (Tests 1 through 4) are shown in Fig. 1,

particle size (Tests 2 and 5 through 7) in Fig. 2,

oxygen partial pressure (Tests 2, 8 and 9) in Fig. 3,

and pyrite pulp density and the addition of copper

(Tests 2 and 10 through 12) in Fig. 4.

For the tests at low pulp density, pyrite conversion

was calculated from the concentration of iron in

solution divided by the mass of iron in the head

sample divided by the original solution volume. The

ratio of Fe(III) to total dissolved iron ([Fe(III)]/[FeT])

was calculated from the concentration of Fe(III)

divided by the total iron concentration in solution.

The time elapsed from sampling to Fe(II) titration was

never more than 1 h, during which time no significant

difference in titration results was found from the

preliminary tests.

For the tests at high pulp density, a digital oxygen

mass flowmeter was used to record the consumption

of oxygen during pressure oxidation. Thus, pyrite

(800 r

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349 341Fig. 4. Effect of pulp density and copper addition on pyrite oxidationr vs. time, (c) s(dr/dt) vs. r2, (d) [Fe(III)]/[FeT] vs. time (low pulp de

pm, 210jC, 690 kPa O2, 74 + 53 Am). (a) Conversion vs. time, (b)

nsity only).

074 + 5

0

5

.9

21

438

0

74 + 5

5

5

.9

73

466

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349342conversion was calculated from the oxygen consumed

Table 2

Summary of experimental conditions and calculated parameters

Test. no. 1 2a 3

T, jC 230 210 19d0, Am 74 + 53 74 + 53 PO2, kPa 690 690 69

[FeS2], g/L 1 1 1

[CuSO4], g/L 0 0 0

[H2SO4], M 0.5 0.5 0.

s, min 25.6 38.7 59m 1.09 2.83 3.

rp2 0.342 0.398 0.

Test. no. 7 8 9

T, jC 210 210 21d0, Am 53 + 44 74 + 53 PO2, kPa 690 1035 34

[FeS2], g/L 1 1 1

[CuSO4], g/L 0 0 0

[H2SO4], M 0.5 0.5 0.

s, min 28.7 31.5 52m 3.22 2.74 2.

rp2 0.272 0.336 0.

a Standard conditions.according to reactions (1) and (3). Given that the bulk

of dissolved iron reported as Fe(III) in every test, this

method should have incurred, at most, 12% relative

error in pyrite conversion.

3.1. Effect of agitation speed

Agitation speed is often important in liquidsolid

reactions, especially those under diffusion control.

The initial rate of pyrite extraction was recorded at

650, 800 and 950 rpm under the standard conditions.

It was observed that agitation speed had no significant

effect on the initial rate of pyrite oxidation when

maintained at 800 rpm or higher. Based on these

results, a stirring speed of 800 rpm was chosen for

all subsequent experiments to minimize iron precipi-

tation by solution splashing.

3.2. Effect of temperature

The effect of temperature (Tests 1 through 4) on

pyrite conversion is shown in Fig. 1(a). At 230jC,nearly all of the pyrite dissolved within 20 min. Data

for the first 15 min of leaching at each temperature arewell described by the shrinking sphere model denot-

4 5 6

170 210 210

3 74 + 53 149 + 105 105 + 74690 690 690

1 1 1

0 0 0

0.5 0.5 0.5

99.1 77.3 53.3

1.73 2.43 1.70

0.734 0.415 0.473

10 11 12

210 210 210

3 74 + 53 74 + 53 74 + 53690 690 690

20 1 20

0 5 5

0.5 0.5 0.5

25.4 32.0 16.9

2.34

0.353 ing surface chemical reaction control:

dX

dt 31 X

2=3

sor

drds

1s

6

or, in integrated form:

X 1 1 ts

3or r 1 t

s7

where r=(1X)1/3 = d/d0 (particle diameter/initial par-ticle diameter), s = time for complete oxidation inminutes.

This fit is demonstrated by the linear plots of r vs.time, shown in the inset of Fig. 1(b). (Similar plots to

test the shrinking core model denoting product layer

diffusion control failed to yield straight lines, and

examination under a scanning electron microscope

(SEM) found no product layers on partially leached

pyrite surfaces.) The inverse slopes (or the x-inter-

cepts) of these r vs. time lines represent s, the time forcomplete reaction predicted by the shrinking sphere

model. (All parameters calculated from experimental

results are tabulated in Table 2.)

Plotting the natural logarithm of 1/s vs. the inverseabsolute temperature, 1/s, gives an Arrhenius plot, theslope of which represents E/R, where E is theapparent Arrhenius activation energy of pyrite oxida-

tion and R is the gas constant. As shown in Fig. 5, the

apparent Arrhenius activation energy by this method

is 41.7 kJ/mol (10.0 kcal/mol). That this is not the true

activation energy will be demonstrated below, in the

section on the effect of oxygen partial pressure.

Returning to Fig. 1(a) and (b), it may be seen that

the shrinking sphere model (represented by dotted

curves) deviates from the conversion and j data, mostmarkedly at the lowest temperature. From Eq. (7), the

quantity s(dr/dt) should always have a value ofone. However, as shown in Fig. 1(c), plots of s(dr/

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349 343dt) vs. r2 deviate negatively and (more or less)linearly from one below a certain value of r2. Weattribute this deviation to passivation of the mineral

surface, most likely by elemental sulfur. Indeed, traces

of elemental sulfur were detected in the residue of the

test run at 170jC, which also displayed the mostsevere passivation behavior. (It bears noting that the

values of dr/dt were obtained directly from the datausing weighted central differences of the form:

drdt

icti1 titi1 ti1

ri ri1ti ti1

ti ti1

ti1 ti1

ri1 riti1 ti

8

which ensures second-order level accuracy.)

Fig. 5. Arrhenius plot for determining the effect of temperature onpyrite oxidation.3.3. The passivating shrinking sphere model

If we define the point on the x-axis of Fig. 1(c)

representing the onset of passivation (i.e., the point of

deviation from s(dr/dt) = 1) as rp2, then assuming alinear deviation, to the left of this point the rate is

given thus:

s drdt

1 mr2p r2 9

where m is the slope of the line and represents the rate

of passivation. The integrated form of this passivat-

ing shrinking sphere model is

r 1 tsat rzrp 10

r k rp k tankuk rp tanku

at rVrp and mr2p < 1 11

r rp1 rpu at rVrp and mr

2p 1 12

r k rp k tanhkuk rp tanhku

at rVrp and mr2p > 1 13

where k=(j1mrp2j/m)1/2 and u =m[rp (1 t/s)].When mrp

2V 1, then the line intersects the y-axisand the reaction goes to completion, since the oxida-

tion rate is still non-zero at 100% conversion; hence,

incomplete passivation. However, when mrp2>1, then

the line intersects the x-axis and the reaction fails to

go to completion, since the oxidation rate falls to zero

before complete conversion is reached; hence, com-

plete passivation. Instead, a limiting particle size is

approached asymptotically:

limt!lr k at mr

2pz1 14

One can see both situations in Fig. 1(c). Complete

passivation is achieved at the three lowest temper-

atures, and this observation is borne out in the

tests, suggesting that the rate of passivation is a strong

function of temperature only. Also, the relative parti-

cle surface area at the onset of passivation, rp2, appears

to be only a weak function of particle size.

The relationship between the ratio of [Fe(III)]/[FeT]

and particle size is shown in Fig. 2(d). It was again

found that the ratio of [Fe(III)]/[FeT] decreased with

time at the beginning, and then increased asymptoti-

cally to about 80%. However, in this case, roughly the

same minimum [Fe(III)]/[FeT] ratio was achieved at

each particle size, at times roughly proportional to s.

3.5. Effect of oxygen partial pressure

The effect of oxygen partial pressure (Tests 2, 8 and

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349344conversion data shown in Fig. 1(a), and most clearly

in the r vs. time plots shown in Fig. 1(b), where thebold curves represent the passivating shrinking sphere

model. (It will be noted that the lines shown in Fig.

1(c) were calculated based on the passivation param-

eters which resulted in the best fit of the integrated

rate laws, given by Eqs. (11)(13), to the r vs. timedata given in Fig. 1(b). These parameters are also

tabulated in Table 2.)

This model is similar in concept to a model put

forward by Crundwell and Godorr (1997) to explain

the passivation of gold during cyanidation. However,

whereas the mathematical form of our model has been

derived from empirical considerations, their model

was derived based on the assumption that the growth

of the passivating film obeys the same rate law as

Langmuir adsorption. Both models give similar

results, although only ours allows for the onset of

passivation well after the commencement of leaching,

and only ours is amenable to a closed-form integral

solution.

Fig. 1(d) shows the effect of temperature on the

ratio of Fe(III) to total dissolved iron. The ratios of

[Fe(III)]/[FeT] decreased with time at the beginning,

then increased, indicating that the iron in pyrite is

most likely oxidized to Fe(III) directly under these

conditions, and not Fe(II). Fe(III) is subsequently

reduced to Fe(II). While the rate of Fe(II) oxidation

increased with increasing temperature, more than 40%

Fe(II) was detected at the end of the test at 170jC andabout 20% Fe(II) remained even after complete oxi-

dation of the pyrite at 230jC.In order to clarify the role of pyrite in the oxidation

of Fe(II), a blank test was run under the standard

conditions with 0.6 g Fe(II) added to the initial

solution. The results of this test are shown in Fig. 6.

While Fe(II) was oxidized (slowly) in the absence of

pyrite, again, nearly 20% Fe(II) remained in solution

after pressure oxidation. Hence, the high percentage

of Fe(II) remaining in solution after pressure oxidation

of pyrite is not due to the disappearance of pyrite

(although pyrite definitely catalyzes Fe(II) oxidation)

but appears to be simply an equilibrium effect. (Of

course, high levels of Fe(II) in the autoclave discharge

following acidic pressure oxidation of refractory gold

ores are to be avoided, and may necessitate an alkaline

pre-aeration step prior to gold cyanidation, since ironforms strong complexes with cyanide.)3.4. Effect of particle size

The effect of particle size (Tests 2 and 5 through 7)

on pyrite conversion is shown in Fig. 2(a). Again,

reaction times were calculated from the slopes of the jvs. time lines shown in the inset of Fig. 2(b). A plot of

log 1/s vs. log d0, shown in Fig. 7, yields a straightline with a slope of 1.027, which represents thereaction order with respect to initial particle diameter.

This is very close to the value of 1.0 expected fromthe shrinking sphere model.

Plots of s(dr/dt) vs. r2 for the different particlesizes are shown in Fig. 2(c). Similar passivation

behavior may be observed, although the (somewhat

random) spread in the values of the passivation slope

m is much narrower at the single temperature of these

Fig. 6. Fe(II) oxidation kinetics in the absence of pyrite (800 rpm,

210jC, 690 kPa O2, 0.6 g/L Fe).9) on pyrite conversion is shown in Fig. 3(a) and on r

then one would expect a reaction order with respect

to oxygen partial pressure of 0.25. (For an enlighten-

ing discussion of the orders of multistep electrochem-

ical reactions, interested readers are referred to

Section 9.1 of Bockris and Reddy, 1970.)

Now that we have confirmed that the pyrite oxi-

dation rate is 1/2-order with respect to oxygen partial

pressure, it is possible to correct the Arrhenius acti-

vation energy for the temperature effect on the satu-

ration concentration of dissolved oxygen. According

to the recent model of Tromans (1998) (shown here in

slightly modified form), the (molal) saturation con-

centration of dissolved oxygen may be calculated

relative to the oxygen partial pressure (in atm) at

any temperature T (in K) thus:

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349 345in Fig. 3(b). A plot of log 1/s vs. log PO2, shown inFig. 8, yields a straight line with a slope of 0.47

(nearly 0.5), which represents the reaction order with

respect to oxygen partial pressure. This finding is

consistent with the anodic dissolution model of pyrite

in acidic media suggested by Mishra and Osseo-Asare

(1988), which assumes the first charge transfer step,

involving the deprotonation of adsorbed water, to be

the rate-controlling step of the anodic process:

FeS2 H2O ! FeS2 OH H e 15

It is also consistent with the cathodic reduction model

of oxygen on pyrite in acidic media put forward by

Biegler (1975), which assumes the first charge trans-

Fig. 7. Log log plot for determining reaction order with respect to

particle size.fer step, involving the reduction of adsorbed oxygen

molecules, to be the rate-controlling step of the

cathodic process:

O2 e ! O2 16

Assuming symmetry factors of approximately 0.5,

then the first charge transfer step must be the rate-

controlling step in both half-cell reactions, or else the

reaction order with respect to oxygen partial pressure

would report as some value other than 0.5. If, for

example, the second charge transfer step of the anodic

process was the rate-controlling step, as Biegler and

Swift (1979) surmised, then one would expect a

reaction order with respect to oxygen partial pressure

of 0.75. Alternatively, if the second charge transfer

step of the cathodic process were rate-controlling,O2PO2

wexp

0:046T2 203:35T lnT 1430:55T 68669

8:3143T

17

where w is a concentration-dependent parameterwhich, for sulfuric acid alone in molal concentration

units, is expressed as:

w 11 2:01628H2SO41:253475

( )0:16895418

As may be seen from the results of Tromans model

shown in Fig. 9, the saturation concentration of

Fig. 8. Log log plot for determining reaction order with respect tooxygen partial pressure.

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349346dissolved oxygen is a fairly strong increasing function

of temperature over the range relevant to this study.

Hence, one would expect the temperature dependence

of oxygen saturation to obscure significantly the

true temperature dependence of the pyrite oxidation

reaction.

Since s is inversely proportional to the square rootof [O2], then replacing ln 1/s with ln 1/s 1/2ln[O2]on the y-axis of the Arrhenius plot gives the true

activation energy of pyrite oxidation, which is 33.2

kJ/mol (7.9 kcal/mol) as shown in Fig. 5. This is an

important distinction to make, since in general the

dissolved oxygen concentration in a full-scale auto-

Fig. 9. Tromans model for oxygen solubility as a function of

temperature and [H2SO4].clave will not correspond to the saturation concentra-

tion due to the finite rate of gasliquid mixing. This is

especially true in the critical first compartment of any

continuous autoclave. Hence, in order to model such a

situation rigorously, the rate law of pyrite leaching

must be expressed in terms of dissolved oxygen

concentration, and not of partial pressure.

Plots of s(dr/dt) vs. r2 for the different oxygenpartial pressures are shown in Fig. 3(c). Here, the

passivation slope m is a virtual constant, suggesting

that the rate of passivation is not a function of oxygen

partial pressure. However, the point at which passiv-

ation begins, rp2, appears to be a linear function of

oxygen partial pressure, at least at lower pressures.

A partial explanation for the dependence of the

onset of passivation on oxygen partial pressure may

be found in the plots of the [Fe(III)]/[FeT] ratio shown

in Fig. 3(d). At the lowest pressure, the [Fe(III)]/[FeT]ratio falls to very low values shortly after oxidation

begins, corresponding to the earliest onset of passiv-

ation. Hence, it may be that a decrease in the electro-

chemical potential at the mineral surface, reflected by

the decrease in the [Fe(III)]/[FeT] ratio, initiates the

formation of a passivating film. However, the rate at

which this film then grows would seem to have little

to do with the surface potential.

3.6. Effect of pulp density

The effect of pulp density (Tests 2 and 10) on

pyrite conversion at 1 and 20 g/L, under the standard

conditions, is shown in Fig. 4(a) and on r in Fig.4(b). Increasing pulp density has a fairly dramatic

beneficial effect on the rate of pyrite oxidation.

However, perhaps the most significant effect of

increasing pulp density is that it seems to prevent

passivation, as shown most clearly in the plots of

s(dr/dt) vs. r2 in Fig. 4(c). The reason for thisbehavior is not completely clear. However, it may

depend on the fact that the solubilities of both Fe(II)

and Fe(III) at these temperatures are very low (on the

order of 1 g/L) (Zakharchenko and Tsitsorin, 1949).

Hence, given the increase in catalytic (pyrite) surface

area relative to the concentration of iron in solution, it

may be that the potential near the pyrite surface is

maintained at levels high enough to avoid the onset

of passivation.

3.7. Effect of copper addition

The effect of copper ions (Tests 2, 11 and 12),

added as 5 g/L anhydrous CuSO4 (thus giving 2 g/L

Cu), on pyrite conversion and r are also shown inFig. 4(a) and (b), respectively, at both low (1 g/L

FeS2) and high (20 g/L FeS2) pulp densities under

the standard conditions. The effect of copper addi-

tion at low pulp density is fairly modest, causing a

slight increase (roughly 14%) in the oxidation rate,

and no substantial change in the passivation behav-

ior, as shown in Fig. 4(c). However, the effect at

high pulp density is much more dramatic, increasing

the oxidation rate by over 50%.

It is doubtful whether copper is actually involved

in any galvanic interactions at the pyrite surface.

Instead, its catalytic effect probably stems from thefact that the reaction between Cu(I) and dissolved

(7) and (8). For the particular pyrite sample studied,

at a pulp density of 1 g/L and in the absence of

H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349 347oxygen under these conditions is virtually instanta-

neous (Kimweri and Dixon, 2001), while that between

Fe(II) and dissolved oxygen is very slow, as shown in

Fig. 6. Also, it is well known that Cu(II) acts as a

redox catalyst for dissolved oxygen in the oxidation of

Fe(II) (Dreisinger and Peters, 1989). Hence, in the

presence of copper, a larger proportion of Fe(II)

oxidation can occur homogeneously, in the bulk

solution rather than at the pyrite surface. This effec-

tively decreases the number of electron transfer steps

involved in the anodic process, which in turn serves to

increase the potential at the pyrite surface. This

increase in potential is also reflected by an increase

in the terminal [Fe(III)]/[FeT] ratio at low pulp density

from about 80% in the absence of copper to almost

95% in the presence of copper, as shown in Fig. 4(d).

At high pulp density, given the higher ratios of pyrite

surface area to dissolved iron, the catalytic effect of

copper is amplified.

3.8. Explanation of the passivation phenomenon

Several investigators have postulated the mecha-

nism of pyrite oxidation as involving the formation of

sulfoxy intermediates, most notably thiosulfate (Con-

way et al., 1980; Goldhaber, 1983; Morse et al., 1987;

Kelsall and Yin, 1996). Although the mechanism is

quite complex, involving multiple adsorption, proton-

ation, hydration, and desorption steps, the overall

stoichiometry of the strictly electrochemical compo-

nent of pyrite oxidation can probably be represented

thus:

4FeS2 7O2 4H ! 4Fe3 4S2O23 2H2O19

Thiosulfate is highly unstable in acid and dispropor-

tionates rapidly into elemental sulfur and sulfite:

S2O23 ! S0 SO23 20

The resulting sulfite is further oxidized by Fe(III) to

sulfate:

SO23 2Fe3H2O ! SO24 2Fe2 2H 21

The sulfur, most likely in the form of polysulfidesadsorbed onto the pyrite surface, is also oxidized,dissolved copper, the time constant s may be calcu-lated thus:

1

s 6:50 10

4O21=2d0

exp 3993s

22

where s is in minutes, [O2] is the molal concentrationof dissolved oxygen as given by Tromans modelpossibly through one or more thionate intermediates.

However, if the rate of this oxidation were to

become too slow, owing perhaps to a decrease in

potential near the pyrite surface due to the local

buildup of Fe(II) as indicated in Figs. 1(d), 2(d), 3(d)

and 4(d), then elemental sulfur would tend to accu-

mulate on the pyrite surface, ultimately forming a

passivating film.

4. Conclusions

The oxygen pressure oxidation kinetics of pyrite in

sulfuric acid media were studied in the temperature

range of 170230jC. A number of conclusions maybe drawn from the results:

1. Pyrite oxidation during acidic oxygen pressure

leaching at high temperature exhibits 1/2-order de-

pendency on dissolved oxygen concentration, sug-

gesting an electrochemical leaching mechanism.

2. Fe(III) appears to be produced directly from the

oxidation of pyrite by dissolved oxygen, with a portion

being subsequently reduced to Fe(II), most likely

during the oxidation of reduced sulfur intermediates.

3. At low pulp density, pyrite forms a passivating

film which impedes oxidation, and which may cause it

to cease altogether, depending on the extent of oxi-

dation at the onset of passivation and the rate of

passivation. The former would appear to be a function

of electrochemical potential at the pyrite surface (as

indicated by the ratio of Fe(III) to total iron in

solution), while the latter is mostly a function of

temperature. Increasing the pulp density appears to

prevent the onset of passivation.

4. At least until the onset of passivation, pyrite

oxidation conforms to the shrinking sphere model,

denoting surface reaction control, as given in Eqs.(Tromans, 1998), d0 is in microns and T is in Kelvin.

(9)(13).

R gas constant (8.314 J/gmol K)

u parameter of the passivation shrinking sphere

194122-96) for their financial support.

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H. Long, D.G. Dixon / Hydrometallurgy 73 (2004) 335349 349

Pressure oxidation of pyrite in sulfuric acid media: a kinetic studyIntroductionExperimentalMaterialsApparatusExperimental conditionsTest procedureAnalytical methods

Results and discussionEffect of agitation speedEffect of temperatureThe passivating shrinking sphere modelEffect of particle sizeEffect of oxygen partial pressureEffect of pulp densityEffect of copper additionExplanation of the passivation phenomenon

ConclusionsAcknowledgementsReferences