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A kinetic study on the acid pressure leaching of pyrrhotite
hydrometallurgy Hydrometallurgy
47 ( 1997) I- 18
A kinetic study on the acid pressure leaching of pyrrhotite Dimitrios Filippou, Rao Konduru, George P. Demopoulos Dept.
*
ofMining and Metallurgical Engineering, McGill University, 3610 llnimwity St., Montreal Que.. Canada H3A 2B2 Received 20 March 1997; accepted 5 May 1997
Abstract Natural monoclinic pyrrhotite particles (Fe, _,S> were subjected to pressure leaching by oxygen in sulphuric acid solutions at temperatures ranging between 353 and 453 K (80-180°C). For temperatures below the melting point of sulphur (392 K), the rate of pyrrhotite oxidation shows a moderate dependence on temperature, while it is totally independent of sulphuric acid concentration. Nonetheless, in the absence of oxygen, as much as 30% of the mineral can be dissolved in 0.5 mol/l H,SO,. The conversion data were found to fit well to a shrinking-core model with mixed control by half-order surface reaction and oxidant diffusion though a product layer. Despite the high initial reactivity of pyrrhotite, complete oxidation of the mineral was never achieved at temperatures below 393 K, apparently due to an impervious sulphur product layer covering the particles. Complete pyrrhotite oxidation was achieved at temperatures above the melting point of sulphur and only with the use of lignin sulphonate as dispersant of molten sulphur. By analysing the conversion data with the shrinking-core model, pyrrhotite oxidation in the high temperature range (403-453 K) was found to be surface-reaction controlled and of first order with respect to oxygen pressure. 0 1997 Elsevier Science B.V.
1. Introduction Pressure leaching technology has become the preferred choice for the treatment of non-ferrous metal ores in a number of metallurgical plants. Two good examples of successful application of this technology are the oxidative pre-treatment of refractory gold ores and concentrates [l-4] and the oxidative treatment of zinc sulphide concentrates [2,5-81. In both cases, sulphuric acid solutions, oxygen and temperatures above
Corresponding
author. E-mail: [email protected]
0304-386X/97/$17.00 0 1997 Elsevier Science B.V. All rights reserved. PII SO304-386X(97)00034-0
2
D. Filippou et al./ Hydrometallurgy 47 (1997) I-18
the melting point of sulphur (392 K or 119’C) are employed. The two processes differ in that in the former one all sulphur content of gold-bearing sulphide minerals is oxidised to sulphate ion (SO:-), while in the latter the sulphur content of zinc sulphide is oxidised to its elemental form (SO). In general, temperatures above 433 K (160°C) are applied in the case of refractory gold ores (usually 453-463 K or higher), while strict temperature control in the range of 413-428 K (140- 155°C) is exercised in the case of zinc sulphide pressure leaching. Pyrrhotite is a common sulphide mineral which is often associated with refractory gold and zinc sulphide concentrates which are subjected to pressure oxidative leaching. Several types of pyrrhotite are known to exist in nature having the general formula Fe, _,S with x I 0.13. The most common variety is monoclinic Fe,S,, while a more rare hexagonal stoichiometric pyrrhotite (FeS) is known as troilite [9]. In comparison to pyrite (FeS,) or arsenopyrite (FeAsS), the two other common sulphides present in refractory gold ores, pyrrhotite is known to be highly reactive [ 101, but its exact kinetics have not yet been fully understood and quantified. Thermodynamic considerations via E,-pH diagrams, suggest that at pH less than 2, pyrrhotite will dissolve with or without the presence of an oxidant [l I]. Nevertheless, from a process point of view, it is important to know the exact rate at which pyrrhotite dissolves under different conditions of temperature, solution composition and oxygen pressure. This necessity arises from the fact that, for optimum operation of pressure leaching processes like refractory gold pre-treatment and zinc sulphide treatment, it is highly desirable to know the chemistry and kinetics of the various mineral phases of the process feed. The experimental determination of the rate law of each mineral reaction is also a prerequisite for the development of kinetics-based mathematical models which are suitable for computer simulation of industrial pressure leaching processes. Apparently the first study on pyrrhotite pressure oxidative leaching was published in 1955 by Downes and Bruce [12], but it was Gerlach et al. [13] who first examined the kinetics of the process in 1965. In particular, Gerlach et al. observed that at temperatures 303-353 K, l-50 atm pressure of oxygen and 0.2- 1.4 mol H 2SO,, the dissolution of stoichiomettic pyrrhotite (FeS) proceeds mostly towards the formation of elemental sulphur and to less extent towards the formation of sulphate ion. Minimal amounts of hydrogen sulphide gas (H 2 S) were also detected and were attributed to direct acid attack on pyrrhotite. The authors concluded that, under the conditions they used, the pressure oxidative leaching of pyrrhotite is a process controlled by chemical reaction on the mineral surface and that it proceeds through the intermediate formation of H,S which is then oxidised by dissolved oxygen either to elemental sulphur or to sulphate ion. The activation energy was estimated equal to 71.5 kJ/mol, while the order of reaction with respect to oxygen pressure was found to be equal to l/2. Shneerson et al. [14] examined the role of pyrrhotite in the pressure oxidative leaching of mixed sulphide concentrates (pentlandite (Fe, N&S,, chalcopyrite CuFeS, and pyrrhotite). These investigators found as well that the oxidative dissolution of such mixed sulphides at about 383 K and 20 atm partial pressure of oxygen, is a process controlled by surface reaction with relatively high activation energy (59-84 kJ/mol) and of half-order dependency on oxygen partial pressure. Yan and Xianguang [15] investigated the kinetics of pressure leaching of pyrrhotite
D. Filippou et al./Hydrometallurgy
47 (1997) l-18
3
concentrates containing nickel, copper and large amounts of magnesium (more than 10% w/w>. Their study dealt with weak acidic solutions spiked with 0.005-0.054 mol/l ferrous sulphate (FeSO, . 7H,O), S-20 atm partial pressure of oxygen and temperatures between 363 and 388 K. Under these conditions pyrrhotite was oxidised to give a goethite precipitate ( a-FeOOH), dissolved ferrous sulphate (FeSO,) and elemental sulphur. The yield of elemental sulphur was found to vary between 50 and 80% and it was observed that, as the value of pH drops upon leaching, the yield of elemental sulphur increases. They reported that leaching is mainly controlled by surface reaction during its early stages, then switches to mixed surface-reaction and product-layer diffusion control and ends being controlled solely by diffusion through a surface product layer. The reaction order with respect to oxygen was estimated to be equal to 0.44, while the activation energy was reported to be 61.8 kJ/mol in the early stages of surface-reaction control and 12.1 kJ/mol in the later diffusion controlled stage. They explained that the formation of a product layer of elemental sulphur and goethite was the reason that caused the process to be diffusion controlled in its later stages. Nicholson and Scharer [16] examined the oxidation of hexagonal pyrrhotite (F,S,,) suspended in HNO,-NaOH solutions of pH 2-6 and saturated with oxygen under conditions of ambient temperature and pressure. They observed that under these conditions, the oxidation of pyrrhotite is almost independent of pH, but shows a strong dependency on temperature (activation energy 50-100 kJ/mol). Contrary to other investigators [ 12,13,15], Nicholson and Scharer found that most of sulphur from pyrrhotite was oxidised to sulphate ion and not to elemental sulphur. One possible explanation for this could be the presence of nitric acid (HNO,) which can act as oxidant, and the relatively high pH which seems to promote sulphate over sulphur formation [ 171. Apart from those studies on the kinetics of pyrrhotite pressure oxidation, a number of researchers have reported on the solid products that are formed by this process. Especially in the former USSR, where nickeliferous pyrrhotite has been treated on an industrial scale for sulphur and nickel recovery by pressure oxidation [IS], several studies have been carried out on the characterisation of solid products formed upon pyrrhotite pressure oxidation [19-211. According to these studies, a number of amorphous and/or crystalline iron oxide/hydroxide compounds are formed as intermediate products until the ultimate formation of hematite in a sequence: FeS(s)+:lSo(s)
+ Fe’+(aq)+~~Fe(OH),(amorph.)
+
.
-H,O
+ a-FeOOH(
s)
+
c-w-Fe,O,( s)
The above reaction mechanism seems rather simplistic when compared to the findings of Pratt et al. [22]. These authors have recently examined by various advanced microscopy and spectroscopy techniques the surfaces of natural pyrrhotite grains (Fe,S,) which had been subjected to reaction in dilute sulphuric acid solutions (pH 3.0) at 293 K. Thus, they found that acid-leached pyrrhotite surfaces are chemically stratified. Immediately adjacent to the unreacted pyrrhotite core, there is a zone of high S:Fe ratio which forms in response to outward diffusion of iron ions; next to this zone, there is a
D. Filippou et al./Hydrometallurgy
4
47 (1997) I-18
zone of antipathetic O-S concentrations; finally, the outermost zone on the pyrrhotite surface was identified to be a layer of ferric oxyhydroxides presumed to be formed by iron dissolution and oxidation. It is important to note that almost all past investigations on pyrrhotite pressure oxidation focused on conditions of relatively low temperatures (lower than the melting point of sulphur) and low acid concentrations (generally pH > 1.0). Such conditions are far from those employed in industrial pressure leaching operations. In addition, most of these studies do not provide a rate equation suitable for modelling of processes such as pressure oxidation of gold ores or zinc sulphide concentrates. In fact, even the two rate equations that are given in the bibliography cannot be used in modelling as they deal with very specific materials (i.e. stoichiometric pyrrhotite in Ref. [13] and high magnesium concentrate in Ref. [15]) and relatively low temperatures conditions (below 392 K). Over the past few years, a considerable amount of work has been performed by Demopoulos and co-workers at McGill University on the study of reaction kinetics of pressure leaching of sulphide minerals associated with refractory gold ores and concentrates [23,24]. This work led to the development of mathematical models for the computer simulation of industrial refractory gold-ore leaching operations [25-271. Demopoulos and co-workers have also developed a mathematical model for zinc sulphide pressure leaching [28-301. As an integral part of this ongoing work on pressure leaching kinetics and modelling, the kinetics of pyrrhotite oxidation under pressure leaching conditions had to be investigated once the data provided by other authors were considered insufficient. The results of this new experimental study are reported and discussed below.
2. Experimental The mineral specimens used in this work were nickeliferous pyrrhotite crystals originating from Falconbridge mines (Sudbury, Ontario) and supplied by Ward’s Natural Sciences (St. Catherines, Ontario). The specimens were analysed at the Centre de Recherches MinCrales (Ste-Foy, Quebec) and they were found to contain, beyond iron and sulphur, small amounts of nickel and copper (Table 1). Monoclinic pyrrhotite was identified by X-ray diffractometry to be the principal mineral, but pentlandite, chalcopy-
Table 1 Chemical
composition
of pyrrhotite
specimens % w/w
Fe s (total) Ni CU SiO,
40.0 25.8 3.47 0.91 28.4
D. Filippou et al./Hydrometallurgy
47 (1997) I-18
5
Table 2 Approximate mineralogical composition of pyrrhotite specimens. The following assumptions have been made: (i) all Ni is associated with Fe and S in pentlandite which has a composition typical of Sudbury ores 1311; (ii) all Cu is associated with Fe and S in chalcopyrite and (iii) the amount of Fe and S which is not accounted for pentlandite and chalcopyrite is present as pyrrhotite % w/w % of total Fe Pyrrhotite, Fe, _ ).S Pcntlandite, (Fe,,,, Ni, 7s)Ss C‘halcopyrite, CuFeS, Quartz, SiO,
51.9 9.6 2.6 28.4
90.6 7.4 2.0
Total
98.5
100.0
rite and quartz (SiO,) were also found to coexist in smaller quantities. In Table 2, an approximate mineralogical composition is given based on the results of chemical analysis and X-ray diffractometry. The major equipment used in the experimental work was a 2 1 titanium autoclave consisting of a reactor chamber and a head assembly, both of which could be combined and tightened with a leak-proof closure. Heat was provided to the autoclave by an electrical heating mantle that was connected to a thermocouple and controlled by an external temperature controller. The autoclave was further equipped with a solids feeding device that allowed for convenient injection of minerals or sulphuric acid into the reactor chamber once the desired temperature had been reached. The autoclave was directly connected to an oxygen cylinder in order to pressurise it to the required oxygen pressure. Prior to each experiment, the autoclave was tested for leaks by pressurising with nitrogen gas at ambient temperature. Following a successful leak test, the autoclave was depressurised, filled with 1 1 of sulphuric acid solution and sealed again. The chamber was then purged several times with nitrogen in order to flush out any oxygen that might have been present in the reactor chamber. A properly weighed sample of pyrrhotite was placed inside the solids discharge tube, which was then fitted firmly to the autoclave head and connected to a mini tank of oxygen. After a short heating-up period and once the reactor temperature had been stabilised at the desired set-point, the mineral solids (typically 2.0 g) were injected into the solution by being pushed through a pressurised oxygen stream. At the same time, oxygen was allowed to flow into the reactor from the main oxygen cylinder, which was preset to the required operating pressure, so as to start the reaction. Liquid samples were collected at regular time intervals in a 2 h total operation. The samples were diluted and analysed for iron by flame atomic absorption spectrophotometry (AAS). In the experimental procedure, two special precautions were taken so as to make it safer. First, as past studies have shown [13], pyrrhotite may be attacked by sulphuric acid to release H,S, a reducing gas which can be catastrophic for titanium at elevated temperatures. Hence, in order to prevent any titanium corrosion, the acid solutions had to be spiked with about 0.023 mol/l copper sulphate (5.7 g/l CuSO, . 5H,O) [32]. The effect of copper sulphate on the rate to pyrrhotite oxidation was experimentally
6
D. Filippou et al. / Hydrometallurgy 47 (1997) l-18
examined at low temperatures (below 373 K) and was found to be negligible (results not shown here). The second precaution addressed the problem of molten sulphur agglomeration at temperatures above the melting point of sulphur (392 K) [12]. An attempt was made to solve this problem by adding about 4 g of fine quartz per g of pyrrhotite in the feed solution [33]. However, it was found that the addition of quartz was ineffective in preventing particle agglomeration. Particle agglomeration was completely eliminated only by the addition of lignin sulphonate (about 0.2 g/l), a dispersant of molten sulphur which is currently used extensively in zinc pressure leach operations [5,34]. The conversion X, i.e. the fraction of pyrrhotite dissolved, at a given time was taken equal to the fraction of total iron dissolved and was calculated by taking into account the mass of all iron dissolved at that time, i.e. X = (mass of Fe in the present flush and sample collected + mass of Fe in the previous flushes and samples collected + mass of Fe in the solution remaining /(mass
of Fe in the mineral initially
in the reactor) fed to the reactor)
(1)
Under all conditions tested, all leached iron remained in solution and no appreciable precipitation of iron compounds was noticed. Small errors possibly introduced by the dissolution of other iron-containing minerals (Table 2) at different rates [12,35] than pyrrhotite, were ignored.
3. Results and discussion As mentioned, the principal aim of this work was to establish the rate equation which describes the kinetics of pyrrhotite oxidation by oxygen in sulphuric acid media under high temperature conditions (i.e. above the melting point of sulphur or 392 K). However, before the high temperature tests were conducted, several tests were also performed at temperatures below 392 K in an effort to become better acquainted with the system’s behaviour. 3.1. Low temperature
regime
Typical dissolution curves obtained at various oxygen pressures in 0.25 mol/l H,SO, at 373 K, are shown in Fig. 1. It is clear that after the initial fast dissolution stage, the reaction rate slows down reaching a plateau of maximum conversion of less than 100%. In the absence of oxygen, a maximum conversion of about 30% was recorded. Apparently, elemental sulphur is formed (this was clearly evident at the end of each experiment) which leads to a build up of impervious layers around each reacting particle, thereby arresting the reaction progress. In light of the results of Yan and Xianguang 1151, it is expected that the yield of elemental sulphur was more than 80%. Similar dissolution behaviour was noticed when the temperature was varied from 353 to 383 K (Fig. 2). On the other hand, the stirring speed and the concentration of acid
D. Filippou et al. / Hydrometallurgy 47 (1997) I-18
Partlal
pressure ,
Of oxygen: 0
atm
f 2.99
atm
d 5.51 atm ??
10atm 18.98
atm
L’
30
o*
0
m----
60
~ ----’ 90
120
Time (min) Fig. I. The effect of oxygen partial pressure on pyrrhotite conversion at temperatures below 392 K. Conditions: T = 373 K, pyrrhotite size fraction + 53-105 pm, 0.5 mol/l H,SO,. The lines drawn are fittings to Eq. (1 l), with the exception of that at P = 0 atm which is only indicative.
were found to have no effect on the rate of pyrrhotite oxidation in the presence of oxygen. The results shown in Fig. 3 seem to be in contradiction with the data shown in Fig. 1 for zero oxygen pressure. However, this may not be necessarily so. Acid may be thought to attack pyrrhotite initially as FeS(s)
+ 2H+(aq)
--f Fe’+(aq)
+ H,S(aq)
(2)
Temperature: ‘ 353
K
363
K
* 373
K
* 383 K
Fig. 2. The effect of temperature
atm, pyrrhotite
size fraction
on pyrrhotite conversion at temperatures below 392 k. Conditions: P = 10 + 53-105 pm, 0.5 mol/l H,SO,. The lines drawn are fittings to Eq. (1 I ).
D. Filippou et al. /Hydrometallurgy
47 (1997) 1-18
I
~~~ ~-
Sulphuric acid concentration: * 0.125 mol/L
1
0.25 m&L * 0.5 mol/L
0
30
60
90
120
Time (min) Fig. 3. The effect of acid concentration on pyrrhotite conversion T = 373 K, P = 10 atm, pyrrhotite size fraction +53-105 pm.
but direct oxidation tion process FeS(s)
by dissolved
+ 1/2O,(aq)
at temperatures
below 392 K. Conditions:
oxygen quickly takes over and dominates
+ 2H+(aq)
--) Fe*+(aq)
the dissolu-
+ So(s) + H,O
(3)
Since H,S gas has been detected in pressure oxidation experiments, some researchers [13] came to the conclusion that Eq. (2) is the initial stage of the overall leaching reaction, which undoubtedly involves a series of elementary steps. In the presence of oxygen, any H,S which forms initially is further oxidised to elemental sulphur H,S(aq)
+ 1/20,(aq)
*So(s)
+ H,O
(4)
In fact, very low oxidation potentials in the range of - 0. I5 to 0.0 V are known to suffice in oxidising H,S to elemental sulphur [37]. The severe retarding effect that sulphur brings about in the progress of the reaction, confirms that sulphur is formed and built up on the particle surface rather than in the bulk of the solution. This is to say that even Eq. (4) is most likely heterogeneous with all reactant species being in an adsorbed state on the mineral surface. Eventually, Fe” is oxidised to Fe3+ by dissolved oxygen 2Fe*‘(aq) Then, Fe3’(aq) FeS(s)
+ 1/20,(aq)
+ 2H+(aq)
* 2Fe3+(aq)
most likely attacks pyrrhotite + 2Fe3+(aq)
+ 3Fe*+(aq)
in a reaction
+ So(s)
+ H,O
(5)
parallel to Eq. (3) (6)
Volumetric analyses of samples at the end of a series of tests showed that 90% of the dissolved iron was present in the solution as Fe3+. As Fe*+ is the product of the reactions in Eqs. (3) and (61, the relative contribution of each of these reaction paths cannot be quantified. Knowing, however, that the reduction couple Fe3+/Fe2+ exhibits
D. Filippou et al./Hydrometallurgy
47 (1997) I-IX
9
a very low overpotential in comparison to the couple O,/H,O, it can be safely assumed that Eq. (6) plays a critical role as it does in the corresponding zinc sulphide pressure oxidation system [51. Therefore, it is concluded that Eq. (3) simply gives the overall stoichiometry of pyrrhotite pressure oxidation, while Eqs. (2), (4)-(6) are amongst the elementary steps of a complex reaction sequence. The conversion versus time data were fitted to the characteristic functions g(X) and l?(X) of the well-known shrinking-core model [36] for surface-reaction control and diffusion-through-product-layer control, respectively: g(X)
= I -(l
p(X)
= 1 -3(1
-X)1’3=k,t -x)2’3+2(1
(7) -X)
=k,t
(8)
where t is time and k, and k, are apparent rate constants. This analysis indicated that neither surface-reaction control, nor diffusion-through-product-layer control can adequately fit the complete kinetic data of pyrrhotite dissolution at temperatures below 392 K. This conclusion matches with that of Yan and Xianguang [15] who have observed that the pressure oxidation of pyrrhotite is controlled by mixed kinetics, i.e. by surface chemical reaction and diffusion through a product layer. An attempt was also made to fit the conversion data to a function which is characteristic for mixed control kinetics when the chemical reaction is first order with respect to the oxidant concentration (or partial pressure in the case of oxygen):
.f(X) =dX){l
+ PmPW~ =knt
(9)
where Y is a constant defined as a ratio of diffusion resistance over reaction resistance [36] and k, an apparent rate constant. However, even f(X) does not fit appropriately the kinetic data of pyrrhotite pressure oxidation at temperatures below 392 K. The reason for failing in all these efforts to fit the conversion data to simple shrinking-core model equations (Eqs. (7)-(9)) IS most likely due to the fact that the process is controlled equally by chemical reaction and oxidant diffusion through the product layer, but with the reaction order with respect to oxygen being one-half [14.15] and not one. To verify this hypothesis, the conversion data were fitted to a more complicated mixed-control shrinking-core mathematical model in which it is assumed that the reaction order with respect to oxygen is one-half. Assuming a pseudo-steady state for each reacting particle of initial radius rO, the basic equation of this model is rate of diffusion which becomes
through product layer = rate of surface reaction
a differential
equation
[36] as
where De is the effective diffusivity of the oxidant (dissolved oxygen or ferric ions). b a stoichiometric coefficient, k the intrinsic reaction rate constant, c the oxidant concentration at a radius Y inside the product layer shell and c, the oxidant concentration at the surface of the unreacted core which has radius r,.
D. Filippou et aE./Hydrometallurgy 47 (1997) I-18
10 Table 3 Estimated
values of D,* and k * at temperatures
Experimental
below 392 K
De* (mol/min
conditions
T (K)
P (atm)
rij ( Fm)
313 373 313 373
2.99 5.51 10.00 18.98
45 45 45 45
1.02x 0.93x 0.68X 0.43 x
353 363 383
10.00 10.00 10.00
45 45 45
0.38X lo-” 0.54x lo-“’ 0.89X lo-“’
lo- ‘” 10~“’ IO-” lo-‘0
cm atm)
k a (mol/min
1.17x 0.84X 0.79x 0.90x
1om4 1o-1 lo-’ 10-4
0.39x JOY” 0.61 X 10m4 l.llxIo-”
cm2 atm”‘)
R2
0.9973 0.997 1 0.9973 0.9974 0.9986 0.9990 0.996 1
Provided that the oxidant concentration, no matter if this is dissolved oxygen or ferric ions, is proportional to the partial pressure of oxygen, the above equation may be re-written as
(11) where De* is a modified effective diffusivity, k * a modified intrinsic reaction rate constant, P the partial pressure of oxygen which corresponds to a certain concentration of oxidant at a radius Y inside the product layer shell and P, the partial pressure of oxygen which corresponds to the concentration of oxidant at r = rs. Since the reaction rate is proportional to the square root of P, the differential equation (Eq. (11)) cannot be solved analytically. Therefore, fittings of experimental data to Eq. (11) to obtain values for k * and De* were done by a least squares method coupled with numerical integration (see Appendix A). As shown in Figs. 1 and 2, these fittings were very good, despite some small overestimation at very high oxygen pressures towards the end of the reaction. Values of estimated correlation coefficients (R2), k * and De* are given in Table 3. The capability of this model in predicting the conversion of pyrrhotite as a function of time has been checked against the conversion data obtained for three other size fractions. The experimental results show that the particle size has a negative effect on the rate of pyrrhotite oxidation. In Fig. 4, it is shown that this effect is predicted reasonably well (with some small overestimation) by integrating numerically the halforder-reaction mixed-control model. It is worth noting here that, for the estimation of the conversion of a size fraction ‘+ d in-d,,,ax’, the value of the initial particle radius r. has been taken equal to 1/2/a: ( more details are given in the Appendix A). From the slopes of Arrhenius plots of In De* and In k * against the reciprocal of temperature (1 /T) at P = 10 atm, activation energies were estimated. Therefore, the activation energy for De* was calculated equal to 31 .l + 1.3 kJ/mol (R* = 0.9962) while the activation energy for k * was calculated equal to 38.1 t 2.5 kJ/mol ( R2 = 0.9917).
D. Filippou et al. / Hydrometallurgy 47 (I9971 l-l8
1I
Size
fraction:
t.53.75
pm
+ 75-105
pm
* t105-150pm - +150-210
/m
Time (min) Fig. 4. The effect of particle size on pyrrhotite conversion at temperatures below 392 K. Conditions: 7 = 373 K. P = 10 atm, 0.5 mol/l HaSO,. The lines drawn are predictions obtained by numerical integration of Eq. ( I 1) assuming k * =0.79X 10d4 mol/(min cm atm), and 0,‘ = 0.68 X lo- ‘” mol/min cm2 atm”‘.
3.2. High temperature
regime
In contrast to the previously studied systems of arsenopyrite pressure oxidation [23] and pyrite pressure oxidation [24], complete oxidation of pyrrhotite was not possible at temperatures close to or above the melting point of sulphur (392 K). As shown in Fig. 5,
1
_-30
_~_1
60
90
120
Time (mln) Fig. 5. The effect of temperature on pyrrhotite conversion at temperatures sulphonate addition. Conditions: P = 10 atm, pyrrhotite size fraction +53-105 lines drawn are only indicative.
above 392 K without lignin pm, 0.5 mol/l H,SO,. The
12
D. Filippou et al./Hydrometallurgy
47 (1997) l-18
, Partlal
pressure
I
of oxygen:
* 5.51 atm k 10
Fig. 6. The effect of oxygen partial pressure on pyrrhotite conversion sulphonate addition. Conditions: T = 403 K, pyrrhotite size fraction lines drawn are predictions obtained by Eq. (12).
I I
aim
i 17.21
atm
at temperatures above 392 K with lignin + 53-105 pm, 0.5 mol/l HzSO,. The
the conversion in the high temperature range drops significantly and the reaction ceases prematurely, if no liquid sulphur dispersants are added. At the end of these experiments, most of the mineral particles were found fully agglomerated into a single cemented ball at the bottom of the reactor chamber. This was obviously due to the formation of liquid sulphur which acts as an agglomerating agent.
,
Temperature: * 403 K
0
15
30
45
60
?? 433
K
A 453
K
75
90
Time (min) Fig. 7. The effect of temperature on pyrrhotite conversion addition. Conditions: P = 10 atm, pyrrhotik size fraction are predictions obtained by E$. (12).
at temperatures above 392 K with lignin sulphonate + 53-105 Wm. 0.5 mol/l H,SO,. The lines drawn
D. Filippou et al. / Hydrometallurgy 47 (1997) I- I8 Table 4 Estimated
values of k, and k, at temperatures
above 392 K
k, (min-‘)
R2
k, (min- ‘1
R2
17.21
45 45 45
0.0073 0.0161 0.0284
0.7743 0.9532 0.8913
0.0074 0.0156 0.0260
0.9652 0.9829 0.9926
10.00 10.00
45 45
0.0501 0.1642
0.9959 1.0000
0.0489 0.1641
0.9960 0.9999
Experimental
conditions
T (K)
P (atm)
r, ( pm)
403 403 -io3
5.51
433 133
10.00
As already mentioned, the problem of sulphur agglomeration was solved only by adding to the leach solution some lignin sulphonate as liquid sulphur dispersant. The dissolution data obtained at high temperatures with the addition of lignin sulphonate are shown in Figs. 6 and 7. These data were easily fitted to Eqs. (7) and (8) giving in most cases better R2 for control by product layer diffusion (Table 4). However, from an Arrhenius plot of In k, and In k, against l/T, the activation energies were estimated to be equal to 68.5 _I 11.2 kJ/mol (R’ = 0.9739) and 69.2 + 1 I .7 kJ/mol ( R2 = 0.9721). respectively. Such high values of activation energy usually justify chemical reaction control. Considering also that Eq. (7) has a simpler form to work with in reactor modelling, it was decided to retain it as the most suitable equation to describe the kinetics of pyrrhotite pressure oxidation at temperatures above the sulphur melting point. The reaction order with respect to oxygen partial pressure at temperatures above 392 K was obtained by linear regression of In k, against In P. The slope of the resulting line gave an estimate of the reaction order 1.19 f 0.08 with R2 = 0.9953. All these findings have been summarised in a single equation which represents the kinetics of pyrrhotite pressure oxidation at temperatures between 403 and 453 K (the temperature range investigated): 1 - 3.3 X 103t( P1.19/ro)
exp
i
68.5 kJ/mol RT
’
(12)
where R is the universal gas constant. This formula can be used conveniently to predict the conversion of pyrrhotite at high temperatures (Figs. 6 and 7). A similar equation which was derived from a first rough analysis of the experimental data has already been used successfully in the modelling of the Sherritt zinc pressure leaching process [28,30].
4. Conclusions In the present study, the pressure oxidation of pyrrhotite was studied at temperatures below and above the melting point of sulphur (392 K). In the low temperature region (353-383 K), pyrrhotite reacts first very rapidly but, with the progress of the reaction, the rate of pyrrhotite oxidation slows down significantly and reaches a plateau considerably below 100% conversion in 2 h time. Apparently, the overall reaction involves a
D. Filippou et al./Hydrometallurgy
14
47 (1997) l-18
number of elementary steps which lead to the formation of a sulphur layer that covers the partially reacted pyrrhotite particles and arrests the progress of the reaction. As a result, a very complex kinetic behaviour is exhibited with mixed control by reaction at the surface of the unreacted particle core and diffusion through the product layer. The experimental data of conversion versus time fit well to a shrinking-core model differential equation in which it is assumed that the reaction order with respect to the oxidant is one-half. The validity of this equation has been confirmed by comparing its predictions against data from experiments with various pyrrhotite size fractions. In the high temperature region (403-453 K), the pressure oxidation was found to be totally inhibited due to molten sulphur formation which creates severe agglomeration problems. The liquid sulphur was effectively dispersed with the use of lignin sulphonate and, thus, complete conversion could be accomplished within a very short time (e.g. in about 6 min at 453 K). The high temperature experimental data could be equally fitted to the shrinking core model equation for control either by surface chemical reaction or for diffusion though a product layer. The equation for chemical reaction control was finally retained for its simplicity. The high activation energy (68.5 & 11.2 kJ/mol) also implies that the process is controlled by surface chemical reaction. In contrast to its half-order dependency at temperatures below, in the high temperature region, the reaction rate exhibits first-order dependency on oxygen partial pressure.
5. Notation
M n
P
P r
R
R2 S
stoichiometric coefficient of mol pyrrhotite per mol oxidant oxidant concentration (mol/l) diameter (cm) effective diffusivity (l/min cm) modified effective diffusivity (mol/min cm atm) conversion function defined by Eq. (9) conversion function defined by Eq. (7) intrinsic reaction rate constant (I” mol’ -“/min cm2) modified intrinsic reaction rate constant (mol/min cm2 atm’/*) apparent diffusion rate constant defined as (6 bMD, c>/( p r,‘> (min- ’> apparent mixed-control rate constant defined as (bMkc)/( p r,,) (min- ’) apparent reaction rate constant defined as (bMkc”)/( pr,) or (Mk * P”>/( pr,) (mini’) pyrrhotite molecular mass (87.91 g/mol) reaction order with respect to c or P conversion function defined by Eq. (8) oxygen partial pressure (atm) radius (cm) universal gas constant (8.31434 J/mol K) correlation coefficient sum of residuals
D. Filippou et al./Hydrometallurgy
1
T X Y
time (mm) temperature (K) conversion ratio of resistances
47 (19971 I-IX
15
defined as kr,/D,.
Greek letters A i’
difference pyrrhotite
mass density (4.74 g/cm31
Subscripts cal i J
max min s 0
calculated integer index integer index maximum minimum at the surface of the unreacted initial
core
Acknowledgements The authors are indebted to the Natural Sciences and Engineering Research Council of Canada and to Natural Resources Canada for funding this research. Dr. Gerry Bolton (Sherritt International Consultants, Fort Saskatchewan, Alberta, Canada) is thanked for providing a sample of lignin sulphonate. Finally, the laboratory assistance of Dr. Birendra Jena and Mr. Paulo Goncalves is greatly appreciated.
Appendix A The iterative procedure by which the values of k* and De* have been evaluated involves a series of steps for each set of experimental data. If the raw data of one experiment are a set of points {(t,, X,1,. . . ,(tj, X,1. . . } obtained at given conditions of then the estimation of kx and De* is done as T, P and particle size ‘+dmin-d,,,‘, follows: (1) First, guesses are made for the values of the k * and D * (2) The initial particle radius rO is taken equal to l/2 r-d,,,d,,, Then, starting from rg at t = 0, a new particle radius is estimated at a new time t, = At as dr r,=r,+-
At dt o
(A.1)
with At being a small time interval (e.g. 0.1 min) and dr -=
dt
n
-W’dk*JP,,.
(A.2)
D. Filippou et al. / Hydrometallurgy 47 (19971 l-18
16
Since it is assumed that there are no external mass-transfer limitations, F’s,0 can be taken equal to the partial pressure of oxygen P. Once the value r, is estimated, then the conversion X, at t, = At can be calculated as x, = I - (q/r,)’
(A.3)
(3) Eqs. (A.l), (A.2) and (A.3) can be used for the estimation of X at a new time t2 = t, + At, etc. However, at each time step tj = iAt, the value of P,,,_, must be evaluated by solving the equation:
(A.41 where yi_ I is the particle radius at t = (i - l>At. The left-hand side of Eq. (A.4) is obtained by double integration of the left-hand-side of the basic differential equation (I 1)) [36]. (4) The results of this integration at certain times t, can be compared experimental ones and an error estimate can be made as [38] (Eq.
p=
C(y - x,,J2.
A correlation
coefficient
R2 = 1 -P/cX;.
(measure
to the
(‘4.5) of fit) can be also estimated
as [38]: (A4
(5) The whole procedure from step 1 is repeated again with new guesses for k * and 0,” for the minimisation of 9’. This can be done by numerical optimisation methods
[381. The basic differential equation (Eq. (11)) can be also used for the prediction of X as a function of t for given conditions of T and P. In this case, numerical integration must be performed for given values of k * and D,* following the steps (1) to (3) of the above procedure.
References [l] G.P. Demopoulos, V.G. Papangelakis, Recent advances in refractory gold processing, CIM Bull. 82 (9231 (19891 85-91. [2] R.M.G.S. Berezowsky, M.J. Collins, D.G.E. Kerfoot, N. Torres, The commercial status of pressure leaching technology, JOM 43 (2) (1991) 9-15. [3] K.G. Thomas, Alkaline and acidic autoclaving of refractory gold ores, JOM 43 (2) (1991) 16-19. [4] J.A. King, D.A. Knight, Autoclave operations at Portega, Hydrometallurgy 29 (1992) 493-511. [5] M.T. Martin, W.A. Jankola, Cominco’s trail zinc pressure leach operation, CIM Bull. 78 (876) (1985) 77-81. [6] H. Veltman, G.L. Bolton, Direct pressure leaching of zinc blende with simultaneous production of elemental sulphur: A state-of-the-art review, Erzmetall 33 (2) (1980) 76-83. [7] D.W. Ashman, W.A. Jankola, Recent experience with zinc pressure leaching at Cominco, in: T.S. Mackey, R.D. Prengaman (Eds.), Lead-Zinc ‘90, TMS, Warrendale, PA, USA, 1990, pp. 253-275.
D. Filippou et al. / Hydrometallurgy 47 (I 997) I- IX
17
[8] M.J. Collins, E.J. McConaghy, R.F. Stauffer, G.J. Desroches, B.D. Krysa, Starting up the Sherritt zinc pressure leach process at Hudson Bay, JOM 46 (4) (1994) 5 I-58. [9] W.A. Deer, R.A. Howie, J. Zussman, Rock-forming Minerals, vol. 5, Non-Silicates, Longmans, London, 1962, pp. 145-157. [IO] H. Majima, E. Peters, Oxidation rates of sulfide minerals by aqueous oxidation at elevated temperatures, Trans. Metal]. Sot. AIME 236 (1966) 1409-1413. [I I] R.M. Garrels, CL. Christ, Solutions, Minerals and Equilibria, Freeman, Cooper and Company. San Francisco, CA, USA, 1965, pp. 178-229. 1121 K.W. Downes, R.W. Bruce, The recovery of elemental sulphur from pyrite and pyrrhotite, Trans. Can. Inst. Mining Metal]. 58 (1955) 77-82. [ 131J. Gerlach, H. Hlhne, F. Pawlek, Beitrag zur DruckIaugung von Eisensulfiden I: Zur Kinetik der DruckIaugung von Pyrrhotin, Erzmetall 18 (1965) 73-79, (in German). [ 141 Ya.M. Shneerson, I.Yu. Leshch, L.M. Frumina, Role of pyrrhotite in oxidative autoclave leaching of sulphides, Trudyi Proektnogo Nauchno-Izsledovatel’skogo Instituta ‘Gipronikel’, no. 29, 1966, pp. 24-38 (English abstract of Russian original in: Chemical Abstracts, Abstract No. 61809w., vol. 68, 1968). [ 151 S. Yan, L. Xianguang. Study on the kinetics of autoclave leaching of pyrrhotite containing nickel and large amount (sic) of magnesium, in: Y. Zheng, J. Xu (Eds.), Proceedings of the 1st International Conference on Hydrometallurgy (ICHM ‘88), International Academic Publishers, Beijing, 1989. pp, 135-139. [ 161 R.W. Nicholson, J.M. Scharer, Laboratory studies of pyrrhotite oxidation kinetics, in: C.N. Alpcrs. D.W. Blowes (Eds.), Environmental Geochemistry of Sulfide Oxidation, ACS Symposium Series 550. .4CS, Washington, DC, USA, 1994, pp. 14-30. [ 171 H.G. Linge, Anodic oxidation of pyrrhotite in simulated CIP liquors, Miner. Eng. 8 (I 995) 795-806. [ 181 A.B. Voronov, V.G. Popovich, V.D. Shakov, Experimental two-year operation of the Nadezhdinskii metallurgical plant, Tsvetn. Metall. 9 (1982) 18-22, (in Russian). [ 191 A.G. Kitai, V.I. Goryachkin, V.P. Korneev, A.V. Dolenko, V.A. Isaev, Composition of products obtained following autoclave oxidation leaching of pyrrhotite concentrates, Tsvetn. Metal]. I6 t I) (I 975) 1% 15, (English translation). [20] A.V. Medvedev, A.G. Kitai, V.A. Isaev, Mechanism of the transformation of iron oxides in the process of oxidising leaching of pyrrhotite concentrates, Tsvetn. Metal]. 26 (6) (1985) 30-32, (English translation). [21] A.V. Vanyukov, Yu.B. Voitkovskii, N.N. Razumovskaya, Oxidation of pyrrhotites in aqueous pulps, Tsvetn. Metal]. 19 (2) (1978) 9-10, (English translation). [22] A.R. Pratt, H.W. Nesbitt, I.J. Muir, Generation of acids from mine waste: Oxidative leaching ol pyrrhotite in dilute H,SO, solutions at pH 3.0, Geochim. Cosmochim. Acta 58 (1994) 5147-5159. [23] V.G. Papangelakis, G.P. Demopoulos, Acid pressure oxidation of arsenopyrite (Parts I-II), Can. Metall. Q. 29 (19901 l-20. [23] V.G. Papangelakis, G.P. Demopoulos, Acid pressure oxidation of pyrite: Reaction kinetics. Hydrometall Iurgy 26 (1991) 309-325. [25] V.G. Papangelakis, D. Berk, G.P. Demopoulos, Mathematical modeling of an exothermic leaching reaction system, Metall. Trans. B 21B (1990) 827-837. [26] V.G. Papangelakis, G.P. Demopoulos, Reactor models for a series of CSTRs with a gas-solid-liquid leaching system (Parts I-III), Metall. Trans. B 23B (1992) 847-877. [27] S.A. Baldwin, G.P. Demopoulos, V.G. Papangelakis, Computer simulation and analysis of prrssure oxidation autoclaves using REACTSIM, in: V.G. Papangelakis, G.P. Demopoulos (Eds.), Modelling, Simulation and Control of Hydrometallurgical Processes, The Metallurgical Society of CIM. CIM, Montreal, 1993, pp. 123-143. [2X] G.P. Demopoulos, S.W. Baldwin, D. Filippou, V.G. Papangelakis, Development of predictive models for the simulation of industrial zinc leaching processes, in: T. Azakami, N. Masuko, J.E. Dutrizac, E. Ozberk (Eds.), Zinc and Lead ‘95, MMPIJ, Sendai, Japan, 1995, pp. 103-I 17. 1291 S.A. Baldwin, G.P. Demopoulos, V.G. Papangelakis, Mathematical modeling of the zinc pressure leach process, Metal]. Mater. Trans. B 26B (1995) 1035-1047. [30] S.A. Baldwin, G.P. Demopoulos, Assessment of alternative iron sources in the pressure leaching of zinc concentrates using a reactor model, Hydrometallurgy 39 (1995) 147- 162. [3 I ] J.E. Hawley. The Sudbury ores: Their mineralogy and origin, Can. Miner. 7 (I) (1962) I-207.
18
D. Filippou et al. /Hydrometallurgy
47 (1997) l-18
[32] R.W. Schutz, L.C. Convington, Hydrometallurgical applications of titanium, Paper presented at the 3rd ASTM Conference on Titanium and Zirconium in Industrial Applications, New Orleans, Louisiana, USA, Sept. 21-23, 1982, 24 pp. [33] Ya.M. Shneerson, A.A. Onatskaya, L.V. Volkov, A.Yu. Lapin, M.V. Klement’ev, S.Ya. Vecher, Regularities of behaviour of molten sulphur during autoclave leaching of pyrrhotite concentrates, Tsvetn. Metall. 11 (1995) 14-17, (in Russian with English abstract). [34] G. Owusu, E. Peters, D.B. Dreisinger, Surface tensions and contact angles due to lignin sulphonates in the system: Liquid sulphur, aqueous zinc sulphate and zinc sulphide, Can. J. Chem. Eng. 70 (1992) 173-180. [35] T. Deng, Upgrading nickeliferous pyrrhotite concentrates by aqueous oxidation coupled with precipitation and flotation, Int. J. Miner. Proc. 43 (1995) 91-98. 1361 J.M. Smith, Chemical Engineering Kinetics, 3rd ed., McGraw-Hill Book Company, New York, 1981, pp. 636-663. 1371 P.D. Scott, M.H. Nicol, The kinetics and mechanisms of the non-oxidative dissolution of metal sulphides, in: .I. O’M. Bockris, D.A.J. Rand, B.J. Welch (Eds.), Trends in Electrochemistry. Plenum Press, New York, 1978, pp. 303-316. [38] A.K. Sen, MS. Srivastava, Regression Analysis: Theory, Methods, and Applications, Springer-Verlag, New York, 1990, pp. l-27.
47 ( 1997) I- 18
A kinetic study on the acid pressure leaching of pyrrhotite Dimitrios Filippou, Rao Konduru, George P. Demopoulos Dept.
*
ofMining and Metallurgical Engineering, McGill University, 3610 llnimwity St., Montreal Que.. Canada H3A 2B2 Received 20 March 1997; accepted 5 May 1997
Abstract Natural monoclinic pyrrhotite particles (Fe, _,S> were subjected to pressure leaching by oxygen in sulphuric acid solutions at temperatures ranging between 353 and 453 K (80-180°C). For temperatures below the melting point of sulphur (392 K), the rate of pyrrhotite oxidation shows a moderate dependence on temperature, while it is totally independent of sulphuric acid concentration. Nonetheless, in the absence of oxygen, as much as 30% of the mineral can be dissolved in 0.5 mol/l H,SO,. The conversion data were found to fit well to a shrinking-core model with mixed control by half-order surface reaction and oxidant diffusion though a product layer. Despite the high initial reactivity of pyrrhotite, complete oxidation of the mineral was never achieved at temperatures below 393 K, apparently due to an impervious sulphur product layer covering the particles. Complete pyrrhotite oxidation was achieved at temperatures above the melting point of sulphur and only with the use of lignin sulphonate as dispersant of molten sulphur. By analysing the conversion data with the shrinking-core model, pyrrhotite oxidation in the high temperature range (403-453 K) was found to be surface-reaction controlled and of first order with respect to oxygen pressure. 0 1997 Elsevier Science B.V.
1. Introduction Pressure leaching technology has become the preferred choice for the treatment of non-ferrous metal ores in a number of metallurgical plants. Two good examples of successful application of this technology are the oxidative pre-treatment of refractory gold ores and concentrates [l-4] and the oxidative treatment of zinc sulphide concentrates [2,5-81. In both cases, sulphuric acid solutions, oxygen and temperatures above
Corresponding
author. E-mail: [email protected]
0304-386X/97/$17.00 0 1997 Elsevier Science B.V. All rights reserved. PII SO304-386X(97)00034-0
2
D. Filippou et al./ Hydrometallurgy 47 (1997) I-18
the melting point of sulphur (392 K or 119’C) are employed. The two processes differ in that in the former one all sulphur content of gold-bearing sulphide minerals is oxidised to sulphate ion (SO:-), while in the latter the sulphur content of zinc sulphide is oxidised to its elemental form (SO). In general, temperatures above 433 K (160°C) are applied in the case of refractory gold ores (usually 453-463 K or higher), while strict temperature control in the range of 413-428 K (140- 155°C) is exercised in the case of zinc sulphide pressure leaching. Pyrrhotite is a common sulphide mineral which is often associated with refractory gold and zinc sulphide concentrates which are subjected to pressure oxidative leaching. Several types of pyrrhotite are known to exist in nature having the general formula Fe, _,S with x I 0.13. The most common variety is monoclinic Fe,S,, while a more rare hexagonal stoichiometric pyrrhotite (FeS) is known as troilite [9]. In comparison to pyrite (FeS,) or arsenopyrite (FeAsS), the two other common sulphides present in refractory gold ores, pyrrhotite is known to be highly reactive [ 101, but its exact kinetics have not yet been fully understood and quantified. Thermodynamic considerations via E,-pH diagrams, suggest that at pH less than 2, pyrrhotite will dissolve with or without the presence of an oxidant [l I]. Nevertheless, from a process point of view, it is important to know the exact rate at which pyrrhotite dissolves under different conditions of temperature, solution composition and oxygen pressure. This necessity arises from the fact that, for optimum operation of pressure leaching processes like refractory gold pre-treatment and zinc sulphide treatment, it is highly desirable to know the chemistry and kinetics of the various mineral phases of the process feed. The experimental determination of the rate law of each mineral reaction is also a prerequisite for the development of kinetics-based mathematical models which are suitable for computer simulation of industrial pressure leaching processes. Apparently the first study on pyrrhotite pressure oxidative leaching was published in 1955 by Downes and Bruce [12], but it was Gerlach et al. [13] who first examined the kinetics of the process in 1965. In particular, Gerlach et al. observed that at temperatures 303-353 K, l-50 atm pressure of oxygen and 0.2- 1.4 mol H 2SO,, the dissolution of stoichiomettic pyrrhotite (FeS) proceeds mostly towards the formation of elemental sulphur and to less extent towards the formation of sulphate ion. Minimal amounts of hydrogen sulphide gas (H 2 S) were also detected and were attributed to direct acid attack on pyrrhotite. The authors concluded that, under the conditions they used, the pressure oxidative leaching of pyrrhotite is a process controlled by chemical reaction on the mineral surface and that it proceeds through the intermediate formation of H,S which is then oxidised by dissolved oxygen either to elemental sulphur or to sulphate ion. The activation energy was estimated equal to 71.5 kJ/mol, while the order of reaction with respect to oxygen pressure was found to be equal to l/2. Shneerson et al. [14] examined the role of pyrrhotite in the pressure oxidative leaching of mixed sulphide concentrates (pentlandite (Fe, N&S,, chalcopyrite CuFeS, and pyrrhotite). These investigators found as well that the oxidative dissolution of such mixed sulphides at about 383 K and 20 atm partial pressure of oxygen, is a process controlled by surface reaction with relatively high activation energy (59-84 kJ/mol) and of half-order dependency on oxygen partial pressure. Yan and Xianguang [15] investigated the kinetics of pressure leaching of pyrrhotite
D. Filippou et al./Hydrometallurgy
47 (1997) l-18
3
concentrates containing nickel, copper and large amounts of magnesium (more than 10% w/w>. Their study dealt with weak acidic solutions spiked with 0.005-0.054 mol/l ferrous sulphate (FeSO, . 7H,O), S-20 atm partial pressure of oxygen and temperatures between 363 and 388 K. Under these conditions pyrrhotite was oxidised to give a goethite precipitate ( a-FeOOH), dissolved ferrous sulphate (FeSO,) and elemental sulphur. The yield of elemental sulphur was found to vary between 50 and 80% and it was observed that, as the value of pH drops upon leaching, the yield of elemental sulphur increases. They reported that leaching is mainly controlled by surface reaction during its early stages, then switches to mixed surface-reaction and product-layer diffusion control and ends being controlled solely by diffusion through a surface product layer. The reaction order with respect to oxygen was estimated to be equal to 0.44, while the activation energy was reported to be 61.8 kJ/mol in the early stages of surface-reaction control and 12.1 kJ/mol in the later diffusion controlled stage. They explained that the formation of a product layer of elemental sulphur and goethite was the reason that caused the process to be diffusion controlled in its later stages. Nicholson and Scharer [16] examined the oxidation of hexagonal pyrrhotite (F,S,,) suspended in HNO,-NaOH solutions of pH 2-6 and saturated with oxygen under conditions of ambient temperature and pressure. They observed that under these conditions, the oxidation of pyrrhotite is almost independent of pH, but shows a strong dependency on temperature (activation energy 50-100 kJ/mol). Contrary to other investigators [ 12,13,15], Nicholson and Scharer found that most of sulphur from pyrrhotite was oxidised to sulphate ion and not to elemental sulphur. One possible explanation for this could be the presence of nitric acid (HNO,) which can act as oxidant, and the relatively high pH which seems to promote sulphate over sulphur formation [ 171. Apart from those studies on the kinetics of pyrrhotite pressure oxidation, a number of researchers have reported on the solid products that are formed by this process. Especially in the former USSR, where nickeliferous pyrrhotite has been treated on an industrial scale for sulphur and nickel recovery by pressure oxidation [IS], several studies have been carried out on the characterisation of solid products formed upon pyrrhotite pressure oxidation [19-211. According to these studies, a number of amorphous and/or crystalline iron oxide/hydroxide compounds are formed as intermediate products until the ultimate formation of hematite in a sequence: FeS(s)+:lSo(s)
+ Fe’+(aq)+~~Fe(OH),(amorph.)
+
.
-H,O
+ a-FeOOH(
s)
+
c-w-Fe,O,( s)
The above reaction mechanism seems rather simplistic when compared to the findings of Pratt et al. [22]. These authors have recently examined by various advanced microscopy and spectroscopy techniques the surfaces of natural pyrrhotite grains (Fe,S,) which had been subjected to reaction in dilute sulphuric acid solutions (pH 3.0) at 293 K. Thus, they found that acid-leached pyrrhotite surfaces are chemically stratified. Immediately adjacent to the unreacted pyrrhotite core, there is a zone of high S:Fe ratio which forms in response to outward diffusion of iron ions; next to this zone, there is a
D. Filippou et al./Hydrometallurgy
4
47 (1997) I-18
zone of antipathetic O-S concentrations; finally, the outermost zone on the pyrrhotite surface was identified to be a layer of ferric oxyhydroxides presumed to be formed by iron dissolution and oxidation. It is important to note that almost all past investigations on pyrrhotite pressure oxidation focused on conditions of relatively low temperatures (lower than the melting point of sulphur) and low acid concentrations (generally pH > 1.0). Such conditions are far from those employed in industrial pressure leaching operations. In addition, most of these studies do not provide a rate equation suitable for modelling of processes such as pressure oxidation of gold ores or zinc sulphide concentrates. In fact, even the two rate equations that are given in the bibliography cannot be used in modelling as they deal with very specific materials (i.e. stoichiometric pyrrhotite in Ref. [13] and high magnesium concentrate in Ref. [15]) and relatively low temperatures conditions (below 392 K). Over the past few years, a considerable amount of work has been performed by Demopoulos and co-workers at McGill University on the study of reaction kinetics of pressure leaching of sulphide minerals associated with refractory gold ores and concentrates [23,24]. This work led to the development of mathematical models for the computer simulation of industrial refractory gold-ore leaching operations [25-271. Demopoulos and co-workers have also developed a mathematical model for zinc sulphide pressure leaching [28-301. As an integral part of this ongoing work on pressure leaching kinetics and modelling, the kinetics of pyrrhotite oxidation under pressure leaching conditions had to be investigated once the data provided by other authors were considered insufficient. The results of this new experimental study are reported and discussed below.
2. Experimental The mineral specimens used in this work were nickeliferous pyrrhotite crystals originating from Falconbridge mines (Sudbury, Ontario) and supplied by Ward’s Natural Sciences (St. Catherines, Ontario). The specimens were analysed at the Centre de Recherches MinCrales (Ste-Foy, Quebec) and they were found to contain, beyond iron and sulphur, small amounts of nickel and copper (Table 1). Monoclinic pyrrhotite was identified by X-ray diffractometry to be the principal mineral, but pentlandite, chalcopy-
Table 1 Chemical
composition
of pyrrhotite
specimens % w/w
Fe s (total) Ni CU SiO,
40.0 25.8 3.47 0.91 28.4
D. Filippou et al./Hydrometallurgy
47 (1997) I-18
5
Table 2 Approximate mineralogical composition of pyrrhotite specimens. The following assumptions have been made: (i) all Ni is associated with Fe and S in pentlandite which has a composition typical of Sudbury ores 1311; (ii) all Cu is associated with Fe and S in chalcopyrite and (iii) the amount of Fe and S which is not accounted for pentlandite and chalcopyrite is present as pyrrhotite % w/w % of total Fe Pyrrhotite, Fe, _ ).S Pcntlandite, (Fe,,,, Ni, 7s)Ss C‘halcopyrite, CuFeS, Quartz, SiO,
51.9 9.6 2.6 28.4
90.6 7.4 2.0
Total
98.5
100.0
rite and quartz (SiO,) were also found to coexist in smaller quantities. In Table 2, an approximate mineralogical composition is given based on the results of chemical analysis and X-ray diffractometry. The major equipment used in the experimental work was a 2 1 titanium autoclave consisting of a reactor chamber and a head assembly, both of which could be combined and tightened with a leak-proof closure. Heat was provided to the autoclave by an electrical heating mantle that was connected to a thermocouple and controlled by an external temperature controller. The autoclave was further equipped with a solids feeding device that allowed for convenient injection of minerals or sulphuric acid into the reactor chamber once the desired temperature had been reached. The autoclave was directly connected to an oxygen cylinder in order to pressurise it to the required oxygen pressure. Prior to each experiment, the autoclave was tested for leaks by pressurising with nitrogen gas at ambient temperature. Following a successful leak test, the autoclave was depressurised, filled with 1 1 of sulphuric acid solution and sealed again. The chamber was then purged several times with nitrogen in order to flush out any oxygen that might have been present in the reactor chamber. A properly weighed sample of pyrrhotite was placed inside the solids discharge tube, which was then fitted firmly to the autoclave head and connected to a mini tank of oxygen. After a short heating-up period and once the reactor temperature had been stabilised at the desired set-point, the mineral solids (typically 2.0 g) were injected into the solution by being pushed through a pressurised oxygen stream. At the same time, oxygen was allowed to flow into the reactor from the main oxygen cylinder, which was preset to the required operating pressure, so as to start the reaction. Liquid samples were collected at regular time intervals in a 2 h total operation. The samples were diluted and analysed for iron by flame atomic absorption spectrophotometry (AAS). In the experimental procedure, two special precautions were taken so as to make it safer. First, as past studies have shown [13], pyrrhotite may be attacked by sulphuric acid to release H,S, a reducing gas which can be catastrophic for titanium at elevated temperatures. Hence, in order to prevent any titanium corrosion, the acid solutions had to be spiked with about 0.023 mol/l copper sulphate (5.7 g/l CuSO, . 5H,O) [32]. The effect of copper sulphate on the rate to pyrrhotite oxidation was experimentally
6
D. Filippou et al. / Hydrometallurgy 47 (1997) l-18
examined at low temperatures (below 373 K) and was found to be negligible (results not shown here). The second precaution addressed the problem of molten sulphur agglomeration at temperatures above the melting point of sulphur (392 K) [12]. An attempt was made to solve this problem by adding about 4 g of fine quartz per g of pyrrhotite in the feed solution [33]. However, it was found that the addition of quartz was ineffective in preventing particle agglomeration. Particle agglomeration was completely eliminated only by the addition of lignin sulphonate (about 0.2 g/l), a dispersant of molten sulphur which is currently used extensively in zinc pressure leach operations [5,34]. The conversion X, i.e. the fraction of pyrrhotite dissolved, at a given time was taken equal to the fraction of total iron dissolved and was calculated by taking into account the mass of all iron dissolved at that time, i.e. X = (mass of Fe in the present flush and sample collected + mass of Fe in the previous flushes and samples collected + mass of Fe in the solution remaining /(mass
of Fe in the mineral initially
in the reactor) fed to the reactor)
(1)
Under all conditions tested, all leached iron remained in solution and no appreciable precipitation of iron compounds was noticed. Small errors possibly introduced by the dissolution of other iron-containing minerals (Table 2) at different rates [12,35] than pyrrhotite, were ignored.
3. Results and discussion As mentioned, the principal aim of this work was to establish the rate equation which describes the kinetics of pyrrhotite oxidation by oxygen in sulphuric acid media under high temperature conditions (i.e. above the melting point of sulphur or 392 K). However, before the high temperature tests were conducted, several tests were also performed at temperatures below 392 K in an effort to become better acquainted with the system’s behaviour. 3.1. Low temperature
regime
Typical dissolution curves obtained at various oxygen pressures in 0.25 mol/l H,SO, at 373 K, are shown in Fig. 1. It is clear that after the initial fast dissolution stage, the reaction rate slows down reaching a plateau of maximum conversion of less than 100%. In the absence of oxygen, a maximum conversion of about 30% was recorded. Apparently, elemental sulphur is formed (this was clearly evident at the end of each experiment) which leads to a build up of impervious layers around each reacting particle, thereby arresting the reaction progress. In light of the results of Yan and Xianguang 1151, it is expected that the yield of elemental sulphur was more than 80%. Similar dissolution behaviour was noticed when the temperature was varied from 353 to 383 K (Fig. 2). On the other hand, the stirring speed and the concentration of acid
D. Filippou et al. / Hydrometallurgy 47 (1997) I-18
Partlal
pressure ,
Of oxygen: 0
atm
f 2.99
atm
d 5.51 atm ??
10atm 18.98
atm
L’
30
o*
0
m----
60
~ ----’ 90
120
Time (min) Fig. I. The effect of oxygen partial pressure on pyrrhotite conversion at temperatures below 392 K. Conditions: T = 373 K, pyrrhotite size fraction + 53-105 pm, 0.5 mol/l H,SO,. The lines drawn are fittings to Eq. (1 l), with the exception of that at P = 0 atm which is only indicative.
were found to have no effect on the rate of pyrrhotite oxidation in the presence of oxygen. The results shown in Fig. 3 seem to be in contradiction with the data shown in Fig. 1 for zero oxygen pressure. However, this may not be necessarily so. Acid may be thought to attack pyrrhotite initially as FeS(s)
+ 2H+(aq)
--f Fe’+(aq)
+ H,S(aq)
(2)
Temperature: ‘ 353
K
363
K
* 373
K
* 383 K
Fig. 2. The effect of temperature
atm, pyrrhotite
size fraction
on pyrrhotite conversion at temperatures below 392 k. Conditions: P = 10 + 53-105 pm, 0.5 mol/l H,SO,. The lines drawn are fittings to Eq. (1 I ).
D. Filippou et al. /Hydrometallurgy
47 (1997) 1-18
I
~~~ ~-
Sulphuric acid concentration: * 0.125 mol/L
1
0.25 m&L * 0.5 mol/L
0
30
60
90
120
Time (min) Fig. 3. The effect of acid concentration on pyrrhotite conversion T = 373 K, P = 10 atm, pyrrhotite size fraction +53-105 pm.
but direct oxidation tion process FeS(s)
by dissolved
+ 1/2O,(aq)
at temperatures
below 392 K. Conditions:
oxygen quickly takes over and dominates
+ 2H+(aq)
--) Fe*+(aq)
the dissolu-
+ So(s) + H,O
(3)
Since H,S gas has been detected in pressure oxidation experiments, some researchers [13] came to the conclusion that Eq. (2) is the initial stage of the overall leaching reaction, which undoubtedly involves a series of elementary steps. In the presence of oxygen, any H,S which forms initially is further oxidised to elemental sulphur H,S(aq)
+ 1/20,(aq)
*So(s)
+ H,O
(4)
In fact, very low oxidation potentials in the range of - 0. I5 to 0.0 V are known to suffice in oxidising H,S to elemental sulphur [37]. The severe retarding effect that sulphur brings about in the progress of the reaction, confirms that sulphur is formed and built up on the particle surface rather than in the bulk of the solution. This is to say that even Eq. (4) is most likely heterogeneous with all reactant species being in an adsorbed state on the mineral surface. Eventually, Fe” is oxidised to Fe3+ by dissolved oxygen 2Fe*‘(aq) Then, Fe3’(aq) FeS(s)
+ 1/20,(aq)
+ 2H+(aq)
* 2Fe3+(aq)
most likely attacks pyrrhotite + 2Fe3+(aq)
+ 3Fe*+(aq)
in a reaction
+ So(s)
+ H,O
(5)
parallel to Eq. (3) (6)
Volumetric analyses of samples at the end of a series of tests showed that 90% of the dissolved iron was present in the solution as Fe3+. As Fe*+ is the product of the reactions in Eqs. (3) and (61, the relative contribution of each of these reaction paths cannot be quantified. Knowing, however, that the reduction couple Fe3+/Fe2+ exhibits
D. Filippou et al./Hydrometallurgy
47 (1997) I-IX
9
a very low overpotential in comparison to the couple O,/H,O, it can be safely assumed that Eq. (6) plays a critical role as it does in the corresponding zinc sulphide pressure oxidation system [51. Therefore, it is concluded that Eq. (3) simply gives the overall stoichiometry of pyrrhotite pressure oxidation, while Eqs. (2), (4)-(6) are amongst the elementary steps of a complex reaction sequence. The conversion versus time data were fitted to the characteristic functions g(X) and l?(X) of the well-known shrinking-core model [36] for surface-reaction control and diffusion-through-product-layer control, respectively: g(X)
= I -(l
p(X)
= 1 -3(1
-X)1’3=k,t -x)2’3+2(1
(7) -X)
=k,t
(8)
where t is time and k, and k, are apparent rate constants. This analysis indicated that neither surface-reaction control, nor diffusion-through-product-layer control can adequately fit the complete kinetic data of pyrrhotite dissolution at temperatures below 392 K. This conclusion matches with that of Yan and Xianguang [15] who have observed that the pressure oxidation of pyrrhotite is controlled by mixed kinetics, i.e. by surface chemical reaction and diffusion through a product layer. An attempt was also made to fit the conversion data to a function which is characteristic for mixed control kinetics when the chemical reaction is first order with respect to the oxidant concentration (or partial pressure in the case of oxygen):
.f(X) =dX){l
+ PmPW~ =knt
(9)
where Y is a constant defined as a ratio of diffusion resistance over reaction resistance [36] and k, an apparent rate constant. However, even f(X) does not fit appropriately the kinetic data of pyrrhotite pressure oxidation at temperatures below 392 K. The reason for failing in all these efforts to fit the conversion data to simple shrinking-core model equations (Eqs. (7)-(9)) IS most likely due to the fact that the process is controlled equally by chemical reaction and oxidant diffusion through the product layer, but with the reaction order with respect to oxygen being one-half [14.15] and not one. To verify this hypothesis, the conversion data were fitted to a more complicated mixed-control shrinking-core mathematical model in which it is assumed that the reaction order with respect to oxygen is one-half. Assuming a pseudo-steady state for each reacting particle of initial radius rO, the basic equation of this model is rate of diffusion which becomes
through product layer = rate of surface reaction
a differential
equation
[36] as
where De is the effective diffusivity of the oxidant (dissolved oxygen or ferric ions). b a stoichiometric coefficient, k the intrinsic reaction rate constant, c the oxidant concentration at a radius Y inside the product layer shell and c, the oxidant concentration at the surface of the unreacted core which has radius r,.
D. Filippou et aE./Hydrometallurgy 47 (1997) I-18
10 Table 3 Estimated
values of D,* and k * at temperatures
Experimental
below 392 K
De* (mol/min
conditions
T (K)
P (atm)
rij ( Fm)
313 373 313 373
2.99 5.51 10.00 18.98
45 45 45 45
1.02x 0.93x 0.68X 0.43 x
353 363 383
10.00 10.00 10.00
45 45 45
0.38X lo-” 0.54x lo-“’ 0.89X lo-“’
lo- ‘” 10~“’ IO-” lo-‘0
cm atm)
k a (mol/min
1.17x 0.84X 0.79x 0.90x
1om4 1o-1 lo-’ 10-4
0.39x JOY” 0.61 X 10m4 l.llxIo-”
cm2 atm”‘)
R2
0.9973 0.997 1 0.9973 0.9974 0.9986 0.9990 0.996 1
Provided that the oxidant concentration, no matter if this is dissolved oxygen or ferric ions, is proportional to the partial pressure of oxygen, the above equation may be re-written as
(11) where De* is a modified effective diffusivity, k * a modified intrinsic reaction rate constant, P the partial pressure of oxygen which corresponds to a certain concentration of oxidant at a radius Y inside the product layer shell and P, the partial pressure of oxygen which corresponds to the concentration of oxidant at r = rs. Since the reaction rate is proportional to the square root of P, the differential equation (Eq. (11)) cannot be solved analytically. Therefore, fittings of experimental data to Eq. (11) to obtain values for k * and De* were done by a least squares method coupled with numerical integration (see Appendix A). As shown in Figs. 1 and 2, these fittings were very good, despite some small overestimation at very high oxygen pressures towards the end of the reaction. Values of estimated correlation coefficients (R2), k * and De* are given in Table 3. The capability of this model in predicting the conversion of pyrrhotite as a function of time has been checked against the conversion data obtained for three other size fractions. The experimental results show that the particle size has a negative effect on the rate of pyrrhotite oxidation. In Fig. 4, it is shown that this effect is predicted reasonably well (with some small overestimation) by integrating numerically the halforder-reaction mixed-control model. It is worth noting here that, for the estimation of the conversion of a size fraction ‘+ d in-d,,,ax’, the value of the initial particle radius r. has been taken equal to 1/2/a: ( more details are given in the Appendix A). From the slopes of Arrhenius plots of In De* and In k * against the reciprocal of temperature (1 /T) at P = 10 atm, activation energies were estimated. Therefore, the activation energy for De* was calculated equal to 31 .l + 1.3 kJ/mol (R* = 0.9962) while the activation energy for k * was calculated equal to 38.1 t 2.5 kJ/mol ( R2 = 0.9917).
D. Filippou et al. / Hydrometallurgy 47 (I9971 l-l8
1I
Size
fraction:
t.53.75
pm
+ 75-105
pm
* t105-150pm - +150-210
/m
Time (min) Fig. 4. The effect of particle size on pyrrhotite conversion at temperatures below 392 K. Conditions: 7 = 373 K. P = 10 atm, 0.5 mol/l HaSO,. The lines drawn are predictions obtained by numerical integration of Eq. ( I 1) assuming k * =0.79X 10d4 mol/(min cm atm), and 0,‘ = 0.68 X lo- ‘” mol/min cm2 atm”‘.
3.2. High temperature
regime
In contrast to the previously studied systems of arsenopyrite pressure oxidation [23] and pyrite pressure oxidation [24], complete oxidation of pyrrhotite was not possible at temperatures close to or above the melting point of sulphur (392 K). As shown in Fig. 5,
1
_-30
_~_1
60
90
120
Time (mln) Fig. 5. The effect of temperature on pyrrhotite conversion at temperatures sulphonate addition. Conditions: P = 10 atm, pyrrhotite size fraction +53-105 lines drawn are only indicative.
above 392 K without lignin pm, 0.5 mol/l H,SO,. The
12
D. Filippou et al./Hydrometallurgy
47 (1997) l-18
, Partlal
pressure
I
of oxygen:
* 5.51 atm k 10
Fig. 6. The effect of oxygen partial pressure on pyrrhotite conversion sulphonate addition. Conditions: T = 403 K, pyrrhotite size fraction lines drawn are predictions obtained by Eq. (12).
I I
aim
i 17.21
atm
at temperatures above 392 K with lignin + 53-105 pm, 0.5 mol/l HzSO,. The
the conversion in the high temperature range drops significantly and the reaction ceases prematurely, if no liquid sulphur dispersants are added. At the end of these experiments, most of the mineral particles were found fully agglomerated into a single cemented ball at the bottom of the reactor chamber. This was obviously due to the formation of liquid sulphur which acts as an agglomerating agent.
,
Temperature: * 403 K
0
15
30
45
60
?? 433
K
A 453
K
75
90
Time (min) Fig. 7. The effect of temperature on pyrrhotite conversion addition. Conditions: P = 10 atm, pyrrhotik size fraction are predictions obtained by E$. (12).
at temperatures above 392 K with lignin sulphonate + 53-105 Wm. 0.5 mol/l H,SO,. The lines drawn
D. Filippou et al. / Hydrometallurgy 47 (1997) I- I8 Table 4 Estimated
values of k, and k, at temperatures
above 392 K
k, (min-‘)
R2
k, (min- ‘1
R2
17.21
45 45 45
0.0073 0.0161 0.0284
0.7743 0.9532 0.8913
0.0074 0.0156 0.0260
0.9652 0.9829 0.9926
10.00 10.00
45 45
0.0501 0.1642
0.9959 1.0000
0.0489 0.1641
0.9960 0.9999
Experimental
conditions
T (K)
P (atm)
r, ( pm)
403 403 -io3
5.51
433 133
10.00
As already mentioned, the problem of sulphur agglomeration was solved only by adding to the leach solution some lignin sulphonate as liquid sulphur dispersant. The dissolution data obtained at high temperatures with the addition of lignin sulphonate are shown in Figs. 6 and 7. These data were easily fitted to Eqs. (7) and (8) giving in most cases better R2 for control by product layer diffusion (Table 4). However, from an Arrhenius plot of In k, and In k, against l/T, the activation energies were estimated to be equal to 68.5 _I 11.2 kJ/mol (R’ = 0.9739) and 69.2 + 1 I .7 kJ/mol ( R2 = 0.9721). respectively. Such high values of activation energy usually justify chemical reaction control. Considering also that Eq. (7) has a simpler form to work with in reactor modelling, it was decided to retain it as the most suitable equation to describe the kinetics of pyrrhotite pressure oxidation at temperatures above the sulphur melting point. The reaction order with respect to oxygen partial pressure at temperatures above 392 K was obtained by linear regression of In k, against In P. The slope of the resulting line gave an estimate of the reaction order 1.19 f 0.08 with R2 = 0.9953. All these findings have been summarised in a single equation which represents the kinetics of pyrrhotite pressure oxidation at temperatures between 403 and 453 K (the temperature range investigated): 1 - 3.3 X 103t( P1.19/ro)
exp
i
68.5 kJ/mol RT
’
(12)
where R is the universal gas constant. This formula can be used conveniently to predict the conversion of pyrrhotite at high temperatures (Figs. 6 and 7). A similar equation which was derived from a first rough analysis of the experimental data has already been used successfully in the modelling of the Sherritt zinc pressure leaching process [28,30].
4. Conclusions In the present study, the pressure oxidation of pyrrhotite was studied at temperatures below and above the melting point of sulphur (392 K). In the low temperature region (353-383 K), pyrrhotite reacts first very rapidly but, with the progress of the reaction, the rate of pyrrhotite oxidation slows down significantly and reaches a plateau considerably below 100% conversion in 2 h time. Apparently, the overall reaction involves a
D. Filippou et al./Hydrometallurgy
14
47 (1997) l-18
number of elementary steps which lead to the formation of a sulphur layer that covers the partially reacted pyrrhotite particles and arrests the progress of the reaction. As a result, a very complex kinetic behaviour is exhibited with mixed control by reaction at the surface of the unreacted particle core and diffusion through the product layer. The experimental data of conversion versus time fit well to a shrinking-core model differential equation in which it is assumed that the reaction order with respect to the oxidant is one-half. The validity of this equation has been confirmed by comparing its predictions against data from experiments with various pyrrhotite size fractions. In the high temperature region (403-453 K), the pressure oxidation was found to be totally inhibited due to molten sulphur formation which creates severe agglomeration problems. The liquid sulphur was effectively dispersed with the use of lignin sulphonate and, thus, complete conversion could be accomplished within a very short time (e.g. in about 6 min at 453 K). The high temperature experimental data could be equally fitted to the shrinking core model equation for control either by surface chemical reaction or for diffusion though a product layer. The equation for chemical reaction control was finally retained for its simplicity. The high activation energy (68.5 & 11.2 kJ/mol) also implies that the process is controlled by surface chemical reaction. In contrast to its half-order dependency at temperatures below, in the high temperature region, the reaction rate exhibits first-order dependency on oxygen partial pressure.
5. Notation
M n
P
P r
R
R2 S
stoichiometric coefficient of mol pyrrhotite per mol oxidant oxidant concentration (mol/l) diameter (cm) effective diffusivity (l/min cm) modified effective diffusivity (mol/min cm atm) conversion function defined by Eq. (9) conversion function defined by Eq. (7) intrinsic reaction rate constant (I” mol’ -“/min cm2) modified intrinsic reaction rate constant (mol/min cm2 atm’/*) apparent diffusion rate constant defined as (6 bMD, c>/( p r,‘> (min- ’> apparent mixed-control rate constant defined as (bMkc)/( p r,,) (min- ’) apparent reaction rate constant defined as (bMkc”)/( pr,) or (Mk * P”>/( pr,) (mini’) pyrrhotite molecular mass (87.91 g/mol) reaction order with respect to c or P conversion function defined by Eq. (8) oxygen partial pressure (atm) radius (cm) universal gas constant (8.31434 J/mol K) correlation coefficient sum of residuals
D. Filippou et al./Hydrometallurgy
1
T X Y
time (mm) temperature (K) conversion ratio of resistances
47 (19971 I-IX
15
defined as kr,/D,.
Greek letters A i’
difference pyrrhotite
mass density (4.74 g/cm31
Subscripts cal i J
max min s 0
calculated integer index integer index maximum minimum at the surface of the unreacted initial
core
Acknowledgements The authors are indebted to the Natural Sciences and Engineering Research Council of Canada and to Natural Resources Canada for funding this research. Dr. Gerry Bolton (Sherritt International Consultants, Fort Saskatchewan, Alberta, Canada) is thanked for providing a sample of lignin sulphonate. Finally, the laboratory assistance of Dr. Birendra Jena and Mr. Paulo Goncalves is greatly appreciated.
Appendix A The iterative procedure by which the values of k* and De* have been evaluated involves a series of steps for each set of experimental data. If the raw data of one experiment are a set of points {(t,, X,1,. . . ,(tj, X,1. . . } obtained at given conditions of then the estimation of kx and De* is done as T, P and particle size ‘+dmin-d,,,‘, follows: (1) First, guesses are made for the values of the k * and D * (2) The initial particle radius rO is taken equal to l/2 r-d,,,d,,, Then, starting from rg at t = 0, a new particle radius is estimated at a new time t, = At as dr r,=r,+-
At dt o
(A.1)
with At being a small time interval (e.g. 0.1 min) and dr -=
dt
n
-W’dk*JP,,.
(A.2)
D. Filippou et al. / Hydrometallurgy 47 (19971 l-18
16
Since it is assumed that there are no external mass-transfer limitations, F’s,0 can be taken equal to the partial pressure of oxygen P. Once the value r, is estimated, then the conversion X, at t, = At can be calculated as x, = I - (q/r,)’
(A.3)
(3) Eqs. (A.l), (A.2) and (A.3) can be used for the estimation of X at a new time t2 = t, + At, etc. However, at each time step tj = iAt, the value of P,,,_, must be evaluated by solving the equation:
(A.41 where yi_ I is the particle radius at t = (i - l>At. The left-hand side of Eq. (A.4) is obtained by double integration of the left-hand-side of the basic differential equation (I 1)) [36]. (4) The results of this integration at certain times t, can be compared experimental ones and an error estimate can be made as [38] (Eq.
p=
C(y - x,,J2.
A correlation
coefficient
R2 = 1 -P/cX;.
(measure
to the
(‘4.5) of fit) can be also estimated
as [38]: (A4
(5) The whole procedure from step 1 is repeated again with new guesses for k * and 0,” for the minimisation of 9’. This can be done by numerical optimisation methods
[381. The basic differential equation (Eq. (11)) can be also used for the prediction of X as a function of t for given conditions of T and P. In this case, numerical integration must be performed for given values of k * and D,* following the steps (1) to (3) of the above procedure.
References [l] G.P. Demopoulos, V.G. Papangelakis, Recent advances in refractory gold processing, CIM Bull. 82 (9231 (19891 85-91. [2] R.M.G.S. Berezowsky, M.J. Collins, D.G.E. Kerfoot, N. Torres, The commercial status of pressure leaching technology, JOM 43 (2) (1991) 9-15. [3] K.G. Thomas, Alkaline and acidic autoclaving of refractory gold ores, JOM 43 (2) (1991) 16-19. [4] J.A. King, D.A. Knight, Autoclave operations at Portega, Hydrometallurgy 29 (1992) 493-511. [5] M.T. Martin, W.A. Jankola, Cominco’s trail zinc pressure leach operation, CIM Bull. 78 (876) (1985) 77-81. [6] H. Veltman, G.L. Bolton, Direct pressure leaching of zinc blende with simultaneous production of elemental sulphur: A state-of-the-art review, Erzmetall 33 (2) (1980) 76-83. [7] D.W. Ashman, W.A. Jankola, Recent experience with zinc pressure leaching at Cominco, in: T.S. Mackey, R.D. Prengaman (Eds.), Lead-Zinc ‘90, TMS, Warrendale, PA, USA, 1990, pp. 253-275.
D. Filippou et al. / Hydrometallurgy 47 (I 997) I- IX
17
[8] M.J. Collins, E.J. McConaghy, R.F. Stauffer, G.J. Desroches, B.D. Krysa, Starting up the Sherritt zinc pressure leach process at Hudson Bay, JOM 46 (4) (1994) 5 I-58. [9] W.A. Deer, R.A. Howie, J. Zussman, Rock-forming Minerals, vol. 5, Non-Silicates, Longmans, London, 1962, pp. 145-157. [IO] H. Majima, E. Peters, Oxidation rates of sulfide minerals by aqueous oxidation at elevated temperatures, Trans. Metal]. Sot. AIME 236 (1966) 1409-1413. [I I] R.M. Garrels, CL. Christ, Solutions, Minerals and Equilibria, Freeman, Cooper and Company. San Francisco, CA, USA, 1965, pp. 178-229. 1121 K.W. Downes, R.W. Bruce, The recovery of elemental sulphur from pyrite and pyrrhotite, Trans. Can. Inst. Mining Metal]. 58 (1955) 77-82. [ 131J. Gerlach, H. Hlhne, F. Pawlek, Beitrag zur DruckIaugung von Eisensulfiden I: Zur Kinetik der DruckIaugung von Pyrrhotin, Erzmetall 18 (1965) 73-79, (in German). [ 141 Ya.M. Shneerson, I.Yu. Leshch, L.M. Frumina, Role of pyrrhotite in oxidative autoclave leaching of sulphides, Trudyi Proektnogo Nauchno-Izsledovatel’skogo Instituta ‘Gipronikel’, no. 29, 1966, pp. 24-38 (English abstract of Russian original in: Chemical Abstracts, Abstract No. 61809w., vol. 68, 1968). [ 151 S. Yan, L. Xianguang. Study on the kinetics of autoclave leaching of pyrrhotite containing nickel and large amount (sic) of magnesium, in: Y. Zheng, J. Xu (Eds.), Proceedings of the 1st International Conference on Hydrometallurgy (ICHM ‘88), International Academic Publishers, Beijing, 1989. pp, 135-139. [ 161 R.W. Nicholson, J.M. Scharer, Laboratory studies of pyrrhotite oxidation kinetics, in: C.N. Alpcrs. D.W. Blowes (Eds.), Environmental Geochemistry of Sulfide Oxidation, ACS Symposium Series 550. .4CS, Washington, DC, USA, 1994, pp. 14-30. [ 171 H.G. Linge, Anodic oxidation of pyrrhotite in simulated CIP liquors, Miner. Eng. 8 (I 995) 795-806. [ 181 A.B. Voronov, V.G. Popovich, V.D. Shakov, Experimental two-year operation of the Nadezhdinskii metallurgical plant, Tsvetn. Metall. 9 (1982) 18-22, (in Russian). [ 191 A.G. Kitai, V.I. Goryachkin, V.P. Korneev, A.V. Dolenko, V.A. Isaev, Composition of products obtained following autoclave oxidation leaching of pyrrhotite concentrates, Tsvetn. Metal]. I6 t I) (I 975) 1% 15, (English translation). [20] A.V. Medvedev, A.G. Kitai, V.A. Isaev, Mechanism of the transformation of iron oxides in the process of oxidising leaching of pyrrhotite concentrates, Tsvetn. Metal]. 26 (6) (1985) 30-32, (English translation). [21] A.V. Vanyukov, Yu.B. Voitkovskii, N.N. Razumovskaya, Oxidation of pyrrhotites in aqueous pulps, Tsvetn. Metal]. 19 (2) (1978) 9-10, (English translation). [22] A.R. Pratt, H.W. Nesbitt, I.J. Muir, Generation of acids from mine waste: Oxidative leaching ol pyrrhotite in dilute H,SO, solutions at pH 3.0, Geochim. Cosmochim. Acta 58 (1994) 5147-5159. [23] V.G. Papangelakis, G.P. Demopoulos, Acid pressure oxidation of arsenopyrite (Parts I-II), Can. Metall. Q. 29 (19901 l-20. [23] V.G. Papangelakis, G.P. Demopoulos, Acid pressure oxidation of pyrite: Reaction kinetics. Hydrometall Iurgy 26 (1991) 309-325. [25] V.G. Papangelakis, D. Berk, G.P. Demopoulos, Mathematical modeling of an exothermic leaching reaction system, Metall. Trans. B 21B (1990) 827-837. [26] V.G. Papangelakis, G.P. Demopoulos, Reactor models for a series of CSTRs with a gas-solid-liquid leaching system (Parts I-III), Metall. Trans. B 23B (1992) 847-877. [27] S.A. Baldwin, G.P. Demopoulos, V.G. Papangelakis, Computer simulation and analysis of prrssure oxidation autoclaves using REACTSIM, in: V.G. Papangelakis, G.P. Demopoulos (Eds.), Modelling, Simulation and Control of Hydrometallurgical Processes, The Metallurgical Society of CIM. CIM, Montreal, 1993, pp. 123-143. [2X] G.P. Demopoulos, S.W. Baldwin, D. Filippou, V.G. Papangelakis, Development of predictive models for the simulation of industrial zinc leaching processes, in: T. Azakami, N. Masuko, J.E. Dutrizac, E. Ozberk (Eds.), Zinc and Lead ‘95, MMPIJ, Sendai, Japan, 1995, pp. 103-I 17. 1291 S.A. Baldwin, G.P. Demopoulos, V.G. Papangelakis, Mathematical modeling of the zinc pressure leach process, Metal]. Mater. Trans. B 26B (1995) 1035-1047. [30] S.A. Baldwin, G.P. Demopoulos, Assessment of alternative iron sources in the pressure leaching of zinc concentrates using a reactor model, Hydrometallurgy 39 (1995) 147- 162. [3 I ] J.E. Hawley. The Sudbury ores: Their mineralogy and origin, Can. Miner. 7 (I) (1962) I-207.
18
D. Filippou et al. /Hydrometallurgy
47 (1997) l-18
[32] R.W. Schutz, L.C. Convington, Hydrometallurgical applications of titanium, Paper presented at the 3rd ASTM Conference on Titanium and Zirconium in Industrial Applications, New Orleans, Louisiana, USA, Sept. 21-23, 1982, 24 pp. [33] Ya.M. Shneerson, A.A. Onatskaya, L.V. Volkov, A.Yu. Lapin, M.V. Klement’ev, S.Ya. Vecher, Regularities of behaviour of molten sulphur during autoclave leaching of pyrrhotite concentrates, Tsvetn. Metall. 11 (1995) 14-17, (in Russian with English abstract). [34] G. Owusu, E. Peters, D.B. Dreisinger, Surface tensions and contact angles due to lignin sulphonates in the system: Liquid sulphur, aqueous zinc sulphate and zinc sulphide, Can. J. Chem. Eng. 70 (1992) 173-180. [35] T. Deng, Upgrading nickeliferous pyrrhotite concentrates by aqueous oxidation coupled with precipitation and flotation, Int. J. Miner. Proc. 43 (1995) 91-98. 1361 J.M. Smith, Chemical Engineering Kinetics, 3rd ed., McGraw-Hill Book Company, New York, 1981, pp. 636-663. 1371 P.D. Scott, M.H. Nicol, The kinetics and mechanisms of the non-oxidative dissolution of metal sulphides, in: .I. O’M. Bockris, D.A.J. Rand, B.J. Welch (Eds.), Trends in Electrochemistry. Plenum Press, New York, 1978, pp. 303-316. [38] A.K. Sen, MS. Srivastava, Regression Analysis: Theory, Methods, and Applications, Springer-Verlag, New York, 1990, pp. l-27.