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Iron removal during oxidative, acid pressure leaching of a zinc sulphide concentrate
International Journal of Mineral Processing, 25 {1989) 241-260
241
Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands
Iron Removal During Oxidative, Acid Pressure Leaching of a Zinc Sulphide Concentrate FIONA M. DOYLE, HERNANDO ARAUCO*' and LUCIA M. FENG
University of Cali[ornia at Berkeley, Department of Materials Science and Mineral Engineering, Hearst Mining Building, Berkeley, CA 94720 (U.S.A.) (Received September 4, 1987; accepted after revision June 9, 1988)
ABSTRACT Doyle, F.M., Arauco, H. and Feng, L.M., 1989. Iron removal during oxidative, acid pressure leaching of a zinc sulphide concentrate. Int. J. Miner. Process., 25: 241-260. Zinc sulphide leaching is an example of an oxidative, acid pressure leaching process that must be followed by neutralization of excess acid, in this case to remove iron from solution. The leaching behaviour of a zinc sulphide concentrate in moderately acidic solutions (35 g/1 H2S04) has been studied at temperatures from 110 ° to 180 ° C to determine the extent of neutralization that can be achieved by the reaction: Z n S + 2 H + + ½02 = Z n ~+ + H 2 0 + S ° Extensive neutralization was observed at 150°C, even at pH values well outside of the thermodynamic stability region of sulphur. Neutralization was less effective at the other temperatures. Iron precipitated from solution as jarosite. The lead and silver content of these jarosites decreased markedly with increasing temperature. Substitution of other metals in jarosites was studied in hydrolysis tests, and also decreased with increasing temperature.
INTRODUCTION
Acid pressure leaching is being used increasingly in the commercial processing of sulphide materials. Extensive development work has been done on acid pressure leaching of zinc sulphide concentrates (Forward and Veltman, 1959; Stanczyk and Rampacek, 1961; Romanikaw and de Bruyn, 1963; Mackiw and Veltman, 1967; Scott, 1973; Doyle et al., 1978; Bolton et al., 1979; Parker and Romanchuk, 1980; Parker, 1981; Parker et al., 1983; Johnston, 1983; Nattrass et al., 1983; Martin and Jankola, 1985; Verbaan and Crundwell, 1986; Crundwell, 1987a,b; Tozawa and Piao, 1987) and the oxidative pressure leaching of arsenical and pyritic gold ores (Berezowsky and Weir, 1983; Weir and Bere*'Present address: International Rectifier, 233 Kansas Street, E1 Segundo, CA 90245 (U.S.A.)
0301-7516/89/$03.50
© 1989 Elsevier Science Publishers B.V.
2-12
zowsky, 1983; Argall, 1986; Monhemius, 1987 }. The advantages of' acid pressure leaching include the ability to produce elemental sulphur, rather than sulphate or SO~, and the potential use for mixed sulphide concentrates. However, neutralization is often needed after the leach stage; for example, zinc pressure leaching is followed by neutralization to precipitate iron prior to electrowinning, while gold ores must be neutralized before cyanide leaching. The use of an external neutralizing agent such as lime can add significantly to the operating costs of a process, and increase the volume of waste requiring disposal. It has been proposed that a two-stage, countercurrent leach could be used for partial neutralization in zinc sulphide leaching (Bolton et al., 1979; Parker, 1981 ). This relies on acid consumption in the overall leaching reaction: Z n S + H + + 1/20~ = Z n '-'+ + H 2 0 + S ~
(1)
Elemental sulphur, however, is only thermodynamically stable over a narrow range of potentials and pH values. Although this stability range is extended, in practice, because of the high overpotential needed for formation of sulphate (Peters, 1984, 1986), significant limits on the extent of neutralization that could actually be achieved might be expected, particularly at elevated temperatures where kinetic effects are less dominant. Neutralization would promote precipitation of iron as jarosite, a basic iron sulphate. Jarosites may contain appreciable levels of lead and silver, and smaller amounts of other metals that are likely to be present in mixed concentrates. This leads to a loss of values, which may seriously affect the economic viability of' pressure leaching as a means of treating these increasingly common materials. The objective of this work was to determine the extent of neutralization that can be achieved in zinc pressure leaching, and to examine the effects of simultaneous leaching and neutralization on the purity of the iron residues.
Leaching of zinc sulphide Two Canadian zinc producers, Cominco and Kidd Creek Mines, have now integrated direct leaching in their hydrometallurgical zinc circuits (Parker et al., 1983; Johnston, 1983; Martin and Jankola, 1985). Both use a high temperature, oxidative acid pressure leach; ground concentrate is fed into an autoclave with hot, spent electrolyte, typically containing 120 g/l H2SO4 and 50 g/1 Zn as zinc sulphate, to give a pulp density of 145 g/1. The oxygen partial pressure is maintained at 750 kPa (109 psi). Reaction 1, which predominates under these conditions, is strongly exothermic, and the temperature rises to about 145 °-150 ° C. Quebracho and lignin sulphonate are used to prevent the sulphur, which melts at 119 ° C, from occluding unleached sulphide particles, thereby preventing zinc extraction beyond about 50% (Doyle et al., 1978). The temperature is not allowed to exceed 155 °C because polymerization of the sul-
243
phur becomes problematic. The residence time is typically about 90 min. The leach residue is treated to separate elemental sulphur from unleached sulphides, lead sulphate and iron residues, while the pregnant leach solution, containing about 120 g/1 Zn, 35 g/1 H2SO4 and 7 g/1 Fe(III), goes to the main circuit for neutralization with calcine and purification prior to electrowinning (Parker et al., 1983; Martin and Jankola, 1985). The Eh-pH diagrams for the Zn-S-HeO system at 25 ° and 150 ° C are shown in Figs. 1 and 2. These have been calculated using the data in Table I. Although commercial leach solutions have a typical concentration of 2.0 M ZnSO4, the mean activity coefficients of concentrated ZnS04 solutions are very low (only 0.043 at 1.0 M) (Weast 1978), hence the activity of ZnS04 increases very slowly with increasing concentration. Activities of 0.1 M for all soluble sulphur and zinc species were, therefore, used for these diagrams. Under acidic conditions, there are three different leaching reactions for ZnS: (i) Z n S + 2 H + =Zn z+ "bH2S
(2)
t.5
1.0
Hs0~ Zn 2+
0.5
s042-
o3
'"s i....
d 0
rtlJ
ZnO
0.0
-0.5
ZnS -t
... HS-
.0 I
I
0
14
pH Fig. 1. E h - p H diagram for the Z n - S - H 2 0 system at 25 ° C. Activities of all soluble zinc and sulphur species = 0.1M.
244
1.5
i'O I 4C0~
0"5 t
Zn 2+
I-_J
0> tLU
ZnO
0.0 \
c
\
N \ \ \
\ \ \
--0.5
\ \
b~s
\
ZnS
\ \
\.
"x
--i.0 ,
2
%
4
i , i 6 8
,'0
,'2
,4
pH Fig. 2. Eh-pH diagram for the Zn-S-H20 system at 150°C. Activities of all soluble zinc and sulphur species= 0.1M. This occurs at pH < - 1.36 at 25 °C and pH ~< - 0 . 7 2 at 150 °C at 0.1 atm H2S, and thus does not appear on Figs. 1 and 2. (ii) Z n S = Z n 2+ + S ° + 2 e
(3)
(iii) Z n S + 4 H 2 0 = Z n 2+ + S O 2- + 8 H + +Be
(4)
which is also equivalent, in terms of neutralization behaviour, to: (iiia) Z n S + 4 H 2 0 = Z n 2+ +HSO~- + 7 H + + 8 e
(5)
In a chemical leaching system the anodic half-cell reactions 3 and 4 must be coupled with a cathodic reaction. The direct reduction of oxygen, for which the overall reaction is: 02 +4H + + 4 e = 2H20
(6)
occurs as a series of single-electron transfer steps, with high activation energies (Wadsworth, 1972 ), and is subject to strong activation polarization. Low over-
245 TABLE I Thermodynamic data for the Z n - S - H 2 0 system at 298 and 423 K z~G298 o .~
Species
8°98 *1
(kcal/mol) (cal/mol K } (unless specified) H÷ H=, O~ Zn Slrh,,mb) H~O ZnS ZnO H2Sta~ HS S ~HSO~ SO~Zn" ÷ HZnO~ ZnO~-
0.00 0.00 0.00 0.00 *'~ 0.00 - 56.69 - 48.11 *'~ - 76.08 *'~ -6.54 3.01 22.10 - 179.94 - 177.34 - 35.14 - 109.26 *'~ - 9 1 . 8 5 *'~
0.00 31.211 49.003 9.95 *5 7.62 16.716 13.80 .5 10.43 *'~ 29.20 14.60 - 6.40 30.32 4.10 - 26.80 10.00 *s - 23.00 *s
423 Cp0 129s -(cal/mol K)
33.00 *:~ 6.90 *4 8.30 *4 6.07 (298) *5 6.40 *4 18.14 *~ 11.00 ( 298 ) .5 9.62 (298)'5 61.00 *6 - 62.00 *:~ - 61.00 .7 - 18.00 *:~ - 105.00 *:~ 67.71"7 - 121.65 .7 - 162.51.7
0 ,2 zJG42:~
(kcal/mol) - 0.765 - 4.061 - 6.318 - 1.384 -- 1.101 - 59.200 - 50.090 - 77.607 - 11.60 2.622 24.313 - 183.313 - 175.420 - 33.359 - 107.691 -85.210
*'Latimer, 1952. *~Calculated from zlG~2.~= zIG°gs - (423 - 298)zlS°gs - ~C ° L~ ( 4 2 3 - 2 9 8 - 4231n423/298). *:~Criss and Cobble, 1964a. *4Perry, 1963. *'~Wagman et al., 1968. *(~Cobble, 1964. *TCalculated from Cpo k29s 42:~= t~ (T) + fl(T ) 8298(abs) o using constants from Criss and Cobble, 1964b. *SEstimated using the relationship S°gs( ..... ) = 4 3 . 5 - 46.5 ( z - 0.28n) given by Connick and Powell (1953} for oxyanions (z is the ionic charge, n is the number of oxygen atoms not included in hydroxide groups). a l l l e a c h i n g r a t e s a r e o b s e r v e d i n s y s t e m s w h e r e o x y g e n is t h e o n l y a v a i l a b l e oxidant. The reduction of iron (III): F e 3+ + e = F e
2+
(7)
shows minimal polarization until the current density becomes high enough for concentration polarization to be important (Peters, 1984, 1986). Much higher leaching rates are observed for concentrates containing pyrite, which releases i r o n ( P a r k e r a n d R o m a n c h u k , 1 9 8 0 ; P a r k e r e t al., 1 9 8 3 ) . H o w e v e r , t h e s e f a s t e r leaching rates may be due to effects other than simple depolarization of the cathodic reaction. Eh-pH d i a g r a m s f o r p y r i t e a t 25 ° a n d 1 5 0 ° C s h o w t h a t p y r i t e is o x i d i z e d t o F e 2+ a n d e l e m e n t a l s u l p h u r a t a b o u t 0 . 2 5 a n d 0 . 1 5 V, respectively, potentials that are fairly close to those for the same reaction of s p h a l e r i t e ( F e r r e i r a , 1 9 7 5 ) . T h e m e a s u r e d r e s t p o t e n t i a l s o f p y r i t e a t 25 ° C a n d
2,1 6
low pH, however, are about: 0.62 V (Sato, 1960; Peters, 1984 t. The oxidation of pyrite involves a volume increase, and the discrepancy between the predicted and measured potentials has been ascribed by Peters (1984) to passiration by the compact surf'ace layer of elemental sulphur. Sato (1960) was unable to measure the rest potential of sphalerite in acid solutions, because of' its low conductivity, but it is likely to be much closer to the predicted value, because oxidation involves a volume decrease, leading to a porous sulphur layer. Galvanic effects resulting from the potential difference between pyrite and sphalerite are also likely to be responsible for the faster leaching rates observed for concentrates containing pyrite (Peters, 1984). It has also been observed that the leaching rate of pure sphalerite increases linearly with increasing lattice iron (Frenay, 1984; Piao and Tozawa, 1985; Tozawa and Piao, 1987; Kammel et al., 1987 ), because the iron reduces the band-gap energy of the sphalerite and supplies holes in a d-orbital impurity band (Crundwell, 1987a,b). Iron (II) resulting from reaction 7 is oxidized in the bulk solution: 4Fe 2+ + 0 2 + 4 H + =4Fe :~+ +2H~O
(8)
Thus the F e ( I I I ) / F e ( I I ) couple is merely an intermediate, with the overall cathodic reaction being reduction of oxygen. Combining eqs. 3 and 4 with 8 gives: Z n S + 1/202 + 2 H + = Z n 2+ + S U + H 2 0
(9)
and: ZnS + 202 = Z n 2+ +SO~-
(10)
Clearly, reactions 2 and 9 will neutralize excess acid, whereas reaction 10 will not. Simple dissolution (reaction 2) occurs at pH's too low to be of any practical utility. Thus an additional leach stage will only achieve the desired goal of neutralizing acid if elemental sulphur is the main reaction product. It is clear from Fig. 1 that at 25 °C this is only thermodynamically favourable at pH values below 1. At 150°C (Fig. 2) the sulphur stability region contracts, and reaction 9 is only favourable at very low pH. Thus, it may be concluded that little neutralization is thermodynamically possible in an oxidative, acid pressure leach. Experimentally, however, it is observed that below its melting point, elemental sulphur ($8) is much more stable with respect to oxidation than is indicated in Fig. 1 (Biegler and Swift, 1979; Corriou and Kikindai, 1981; Lowson, 1982). Peters (1986) has accounted for the high oxidation overpotential by destabilizing sulphate by 300 k J / m o l (72 kcal/mol) when calculating E h pH diagrams for sulphide minerals at 25 ° C. This extends both the sulphur and the sulphide mineral stability regions; his diagram for the Z n - S - H 2 0 systems indicates that at 25 ° C, ZnS will leach according to reaction 9, and hence neutralize excess acid, up to about pH 5.5. However, from the Butler-Volmer equa-
247 tion, the oxidation overpotential will decrease dramatically with increasing temperature. Thus a much closer correlation is expected between the observed behaviour and that predicted by the Eh-pH diagram at 150 ° C than at 25 ° C.
Behaviour of iron during pressure leaching Although it is desirable to leach with appreciable concentrations of Fe (III) in solution, to increase the leaching kinetics, neutralization during leaching will lead to concurrent precipitation of iron. Phase diagrams for the Fe203-SO3-H20 system at temperatures from 25 ° to 200°C (Posnjak and Merwin, 1922; Walter-Levy and Quemeneur, 1964, 1966 ) indicate that hydronium jarosite, H3OFe3 (SO4) 2(OH) 6, is the stable solid phase in equilibrium with concentrated, acidic iron (III) sulphate solutions. Hydronium jarosite is one of the jarosites, which have the general composition MFe3(SO4)2(OH)6, where M may be K +, Na +, NH + , 1/2Pb 2+ and Ag +, as well as H30 +. Potassium jarosite is more stable than ammonium or sodium jarosite, and all are more stable than hydronium jarosite, although the stability of hydronium jarosite increases with increasing temperature (Dutrizac, 1983). In the jarosite process used in conventional electrolytic zinc processing, alkali salts are usually added to decrease the jarosite solubility, thereby promoting precipitation of iron. However, there is a continuous solid solution series between the jarosites (Dutrizac, 1980), and the jarosites produced in metallurgical processes are usually deficient in alkali, even when formed in the presence of excess alkali, due to incorporation of small amounts of H30 +. The extent of hydronium substitution has been reported in different studies to either increase or decrease with increasing precipitation temperature (Dutrizac, 1983 ). Increased substitution would be consistent with the increase in stability of hydronium jarosite with increasing temperature. The behaviour of impurities during jarosite precipitation has been studied at 90 ° C (Yaroslavtsev et al., 1975) and 97 ° C (Dutrizac, 1984). Losses of divalent metals to alkali jarosites increase with increasing divalent metal concentration, increasing pH or decreasing Fe (III) concentration. Because of their similar ionic radii, the divalent metals substitute for Fe 3+ in the jarosite lattice, the charge balance being maintained by loss of O H - . The extent of divalent metal incorporation increases as the concentration of alkali metal during precipitation increases, with losses being greatest for potassium, then ammonium and sodium jarosites. This is because substitution of divalent metals in the jarosite gives a thermodynamic weakness, thus only limited substitution occurs. Substitution in the alkali jarosites, and hence losses, are greatest when the jarosite is most stable, namely when there are high concentrations of alkali, especially potassium, in solution. Losses vary for different metals, but rarely exceed 3%. The descending order for losses is:
248
Fe :~+ >> Cu e+ > Zn '-'+ > Co-' + > Ni e+ > Mn 2* > Cd'-'~
{11)
which is the same as the order fbr decreasing hydrolysis constants. Lead and silver jarosites are clearly undesirable, because neither metal can be recovered without dissolution or recrystallization of the residue. Moreover, losses to lead jarosite are particularly severe, because there are vacant alkali metal sites; when divalent metals replace Fe :~+ the charge balance is maintained by including additional lead, hence the crystal lattice is not disrupted as severely as in alkali jarosites, where O H - is eliminated for neutrality. The order for losses is the same as that for alkali jarosites. Copper levels of 8% have been reported; at this concentration the composition of the precipitate approaches that of beaverite, PbCuFe~(SO4)e(OH)6 (Dutrizac and Dinardo, 1983). Lead sulphate is very insoluble, thus the formation of lead jarosite is kinetically unfavourable at 90 :-100 ° C, and has hitherto been unimportant in jarosite precipitation processes. Lead jarosite does form during pressure leaching at 145 ~-155 ~C, however. Lead jarosite formation could be suppressed by operating at temperatures of 180 °-200 ° C, above the jarosite stability region (Dutrizac et al., 1980), but this is unlikely to be feasible when leaching sulphide concentrates because of sulphur polymerization. Although the equilibrium behaviour of impurities at 90 ° and 97 °C has been thoroughly studied, there are likely to be significant quantitative differences for jarosites precipitated during leaching at higher temperatures, because of the different relative stabilities of the aqueous species and solid phases, and because of the changing solution composition. At temperatures above 150~C, iron may also precipitate as iron(II) sulphate; Bruhn et al. (1965) reported FeSO4 solubilities of 6.5, 2.0, 0.6 and 0.7 g/100 g H~O at 150 ° , 160 ° , 170:' and 180°C. The solubility at 180°C, the highest temperature used in this study, is equivalent to a solubility product of 2.12.10 -:~. Thus the equilibrium Fe z+ concentration in a 2.0 molal sulphate solution would be 1.06" 10- :~M, or about 0.06 g/1. Fe (II) produced by reaction 7 is thus likely to precipitate before it can be reoxidized to Fe (III). Crundwell ( 1987a,b ) has found that the Fe ( III )-Fe ( II ) couple is not reversible on sphalerite in sulphate solutions. Thus the redox potential at the sphalerite surface, and hence the leaching rate, is not directly determined by the Fe ( I I I ) / F e (II) ratio, although Fe(II) does affect the leaching rate through altering the concentrations of the electroactive Fe (III) species. Haung and Bernal (1984) found that most of the iron was present as ferrous during acid pressure leaching of sphalerite. Various workers have observed that the kinetics of reaction 8, oxidation of Fe(II) by oxygen, initially show a second-order (or close to this) dependence on the Fe(II) concentration (Huffman and Davidson, 1956; Mathews and Robins, 1972; Iwai et al., 1982; Chmielewski and Charewicz, 1984; Verbaan and Crundwell, 1986). If similar kinetics apply at 180 ° C, continual
249
precipitation of FeSO4 would therefore lead to very low concentrations of Fe (III), which in turn would decrease the leaching rate. Iron leaching and precipitation will affect the acid balance in a leaching system. FeS2 is generally oxidized directly to sulphate (Lotens and Wesker, 1987), which releases acid: FeS2+H20+7/202
Fe2++2SO42-+2H +
(12)
Precipitation of any hydrolyzed iron phase also releases acid, for example: 3Fe 3+ +2SO~- +6H20=H3OFe3(SO4)2(OH)6 + 5 H +
(13)
Thus continual neutralization is needed to maintain steady precipitation. Conversely, the precipitation of FeSO4 is a simple crystallization reaction, and does not release hydrogen ions. EXPERIMENTALPROCEDURE
Materials
Pressure leach tests used a Sullivan zinc concentrate provided by Cominco. The principal minerals present were sphalerite, marcasite and galena. The chemical composition of the concentrate is given in Table II. 2.16% of the zinc was water-soluble. After further grinding of the as-received concentrate, size analysis with a Microtrac particle size analyzer indicated the 10th, 50th and 90th percentiles of the cumulative size distribution to be 6.6 pm, 23.6 ttm and 59.2 ttm, respectively. The ground concentrate was stored under argon. Leaching tests
Pressure leaching tests were done in a 300-ml Autoclave Engineers stainless steel autoclave equipped with a 150 ml stainless steel, heated, pressurized reservoir. Acid and pulp were heated separately in the reservoir and autoclave, respectively, then mixed and pressurized with oxygen to initiate leaching. (A moderate partial pressure of oxygen was maintained in the reservoir during heating, to maintain passivity of the steel.) Tests were done at 110 °-180 ° C, nominally with 35 g/1 H2SO4, 7 g/1 Fe(III), pulp densities of 141 or 193 g/l, initial zinc concentrations of 0 or 120 g/1 (as sulphate), oxygen partial presTABLE II Chemical analysis of Sullivan concentrate (wt%) Zn
Pb
Cu
Fe
Ni
Co
Mn
Cd
Ag
Mg
A1
K
Ca
Na
S
SiO2
50.1
4.06
0.25
13.6
.004
.005
0.24
0.13
.007
0.20
0.02
0.03
0.08
0.09
29.6
1.48
25()
sures of 345 or 750 kPa, oxygen flow rates of 35-45 cm:~/min or 15-25 cm:~/ rain, using 0.1 g/1 magnesium lignin sulphonate and 0.2 g/1 quebracho as surihctants in some tests. The injection procedure made precise control over the final solution composition difficult; t = 0 data reported in the figures give the exact compositions.
Hydrolysis test~' To determine the effect of temperature alone on the deportment of impurities to the jarosite residues, sulphate solutions containing 147 g/1 Zn 2+, 18.9 g/1 Fe ~+, 17.5 g/1 Cu 2+, 0.5 g/1 Cd 2+, 0.4 g/1 Ag +, 0 g/l H2SO4 and PbS04 equivalent to 52 g/1 Pb 2+ were held at temperature for 90 min. The final solution and precipitates were analyzed.
Analytical procedures Major mineral phases were characterized by X-ray powder diffraction, using a diffractometer. Samples of concentrate were washed with water to determine the concentration of water soluble zinc, digested in HCI and HN03, then treated with ammonium acetate to dissolve any lead. Residues and precipitates were treated with carbon disulphide or by flotation, to remove elemental sulphur {where appropriate), followed by digestion in HNO3. The solutions generated by these procedures were analyzed by atomic absorption spectrophotometry. Solution samples withdrawn during leaching were centrifuged to remove all solids, the room temperature pH was measured, then metal concentrations were determined by atomic absorption spectrophotometry. Sulphate was determined by barium sulphate precipitation, using disodium rhodizonic acid as an indicator for barium. Total acid was determined by potentiometric titration with KOH. RESULTS AND DISCUSSION
Kinetic data for leaching Sullivan concentrate at 110 ° and 150°C under conditions representative of a first stage are shown in Fig. 3. There was a slight increase in the initial rate of zinc and iron leaching with increasing temperature. The pH after 90 min {measured at room temperature) was slightly higher at 150 ° C, but the pH was never high enough at either temperature to promote precipitation of iron. Similar results were obtained at 120 ° and 135 ° C. Fig. 4 shows kinetic data obtained using similar initial concentrations of acid, zinc and iron, but without surfactant. At l l 0 ° C the leaching rates were slightly faster than those shown in Fig. 3, and the final pH was somewhat higher. At 150 ° C, the absence of surfactant gave significantly faster leaching rates. After about 60 min iron started precipitating as jarosite. The pH in-
251
I
2.6
(a)
(b)
0.25 2.4
O. 2 0 v~E
:E . u c o (J
0.i5 o u 0~ LL
2.2
f
Z'f
c
O. iO
N
2.0
0.05 i.B
o.oo
2'o 4'o 8'o 8o TIME (mins)
2o
7o e'o s'o TIME (mins)
2.0
(c)
i.5
Icl
t.0
0.5 • •
o.o
liO ° 1500
go
go
TIME (mins) Fig. 3. Leaching kinetics for Sullivan concentratewith surfactant: [Zn]o= 120 g/l, [H2SO4]0=35 g/l, 141 g/1 concentrate, 750 kPa 02. creased rapidly to 1.44, then decreased slightly as jarosite precipitated (reaction 13). At 180°C the iron leaching rate was significantly faster, but the zinc concentrations, which are not shown, were inconclusive, because of pronounced scatter. The room temperature pH increased to 3.5 before apparent precipitation of iron, which is inconsistent with the solubility product of jarosite. The
252
2.6 F - - . . . . . . . . . . . . . . . . . . . . . . . .
0.25 2.4 2;"
I i
•
//
0.20 2.2
0.15 o {.9
\ b_
~=-i
/ii1 /
!
0.10
ou c N
./° 2.O
0.05 (a) 0.00
1.8
2~0
4=0 TIME
6'0
B'O
I/ 2=0
(mins)
4LO 6'0 8'0 TIME (mins)
3.5 3.0 2,5 2.0
1.0 • • •
0.5
0. C
2'0
4'0
610
tt0 ° i50 ° t80 °
8=0
TIME ( r a i n s ) Fig. 4. L e a c h i n g k i n e t i c s f o r S u l l i v a n c o n c e n t r a t e g/1, 141 g / l c o n c e n t r a t e , 750 k P a 02.
with no surfaet~t:
[ Zn ] o = 120
g/l,
[ H~SO4 ] o = 35
leach residue from the 180 ° C test was agglomerated, indicating that a substantial amount of elemental sulphur had formed and polymerized. The variation in the zinc and iron concentrations, and the unusually high pH are thought to be due to precipitation of zinc and ferrous sulphate. The solubility of ZnSO4 at 180°C is 34.8 g/100 g H20, or 2.16 molal (Bruhn et al., 1965). This solubility would be reduced by the high total sulphate concentration, and zinc sulphate
253 would .have precipitated as leaching proceeded. Any solid zinc sulphate withdrawn with aqueous samples would have redissolved on cooling. Thus the apparent concentration would have been determined largely by the amount of solid ZnS04 withdrawn in each sample. As discussed earlier, the low solubility of FeSO4 at this temperature would have caused continual depletion of both Fe (II) and Fe (III), leading to slower leaching. Higher pH values were reached before precipitation of Fe (III), because of its low concentration. Surfactants are used in commercial zinc sulphide pressure leaching to prevent occlusion of unleached particles by liquid sulphur (Doyle et al., 1978; Parker, 1981 ). Although faster leaching and neutralization rates were observed in this work in the absence of surfactant, the zinc extraction did not exceed 50%, the level at which liquid sulphur has been reported to prevent further reaction (Doyle et al., 1978). Crundwell (1987a) reported that addition of 0.05 wt% MAGNAFLOC 351, and industrial flocculating agent, guar gum or sodium isobutyl xanthate, a flotation reagent, all decreased the rate of oxidative leaching of sphalerite at 78 ° C, and attributed this effect to a decrease in surface area available for charge transfer resulting from adsorption of the organic material. It appears that the quebracho and lignin sulphonate used here acted similarly. These surfactants have a large number of polar groups, and are used as depressants in flotation (Pradip, 1981). They are thought to promote coalescence of sulphur by adsorbing on the sulphide minerals, preventing wetting by the hydrophobic sulphur. Even if the surfactants were oxidized during leach.ing, they would probably act similarly, since the carboxylic acid groups that would form from the primary alcohol groups are also strongly polar. The presence of the surfactant was least pronounced at l l 0 ° C . At this temperature much of the surfactant would be expected to desorb from the porous, solid sulphur layer forming on the mineral surfaces. It is therefore concluded that in a two-stage leach process, surfactant should be added to the residue from the first stage leach, not to the feed. Kinetic data for tests done with no zinc sulphate initially, but similar concentrations of acid and Fe2 (SO4)3 to those used in the other tests, are given in Fig. 5. Comparison with Figs. 3 and 4 indicates that iron leaching was somewhat faster in the absence of zinc sulphate, whereas the zinc leaching rate was relatively unaffected. The pH was significantly higher in these tests, and this gave earlier precipitation of iron. However, the leaching rates suggest that the higher pH is not due to faster neutralization; instead the pH is thought to be higher because of the decrease in H2SO4 activity with increasing Z n S Q concentration reported by Majima and Awakura (1986). The fact that iron precipitated at l l 0 ° C in the presence of surfactant indicates that the surfactant did not interfere with the nucleation of hydrolyzed iron (III) phases; the lack of precipitation in the tests shown in Fig. 3 was due only to the low pH. This was confirmed by adding 85 g/1 sodium jarosite seed to two leaching runs using
254 0.30 ............................ 0.8
0.25
J
0.20
~ o61
(3 0 . t 5 C o [9 O9 b_ 0 . i 0
u°
0.4
0.05 (a) ~ ~ _ 0.00
2`O
4`O
60
8`O
TIME (rains)
o.o~ 2'o 4'o 6'o 8'o TIME (rains)
2.0
(c)
1.5
i.0
0.5 $iiO
0.0
° no surfsctant
•
l i O ° 8urfactant
•
i50° no surfactant
2'0 2o ~`O ~`O TIME
(rains)
Fig. 5. Leaching kinetics for Sullivan concentrate with no surfactant: [Zn ] o = 0 g/l, [ H2SO 4 ] o -- 35 g/l, 141 g/1 concentrate, 750 kPa 02.
the same conditions as those shown in Fig. 3; again there was no precipitation of iron. Increasing the pulp density from 141 to 193 g/1 did not give significantly faster neutralization at 110 ° or 120 ° C. A reduction in the oxygen partial pressure to 345 kPa, and an increase in flow rate to 35-45 cm3/min also had little effect on the leaching and neutralization rates.
I
255
The molar ratio of zinc to iron in the concentrate was 3.15. The ratios of zinc to iron entering the solution during leaching (before any precipitation of iron) are given in Table III. It is clear that there was preferential leaching of sphalerite at l l 0 ° C , indicating a pronounced galvanic interaction. This was expected; at ambient temperatures, marcasite has a rest potential of 0.57 V, only slightly lower than that of pyrite (Sato, 1960). At 150 °C the zinc to iron ratios entering the solution were much closer to the ratio in the concentrate, indicating that galvanic interactions were non-existent or much weaker. This is entirely consistent with Peters' ( 1984 ) postulation that the potential differences leading to the galvanic effect are due to the difference in permeability of the sulphur layer forming on pyrite and sphalerite; these potential differences would disappear above the melting point of sulphur. Fig. 6 shows the effect of temperature on the composition of the jarosite precipitated during leaching tests with 120 g/1 Zn 2+, and no surfactant. It is evident that the lead and silver concentrations decreased markedly with increasing temperature, while the iron concentration increased, due to higher proportions of hydronium jarosite. These analyses were consistent with the Xray powder diffraction patterns. These indicated that the precipitates formed at 110 ° and 120 °C were solid solutions of hydronium, lead and silver jarosites, based on the diffraction patterns of Dutrizac and Kaiman (1976), whereas the material precipitated at 135 ° and 150°C only showed diffraction lines corresponding to solid solutions of hydronium and lead jarosites. The low levels of copper in the residues reflect the low copper level in the Sullivan concentrate, and do not necessarily reflect the maximum substitution that could occur. The increasing importance of hydronium jarosite at the higher temperatures is consistent with the increased stability of this phase in relation to other jarosites at higher temperatures, but the high concentration of lead in the jarosites formed at the lowest temperatures had not been expected, given the low solubility of lead sulphate, and the fact that lead jarosite prepared by Dutrizac and Kaiman (1976) at 95 ° C contained only 8.35 wt% Pb. Lead sulphate (anglesite ) is unlikely to account for these high lead concentrations, since the precipitate had been washed with ammonium acetate before analysis, and no lead TABLE III Ratio of zinc and iron entering solution during leaching
Surfactant present No surfactant [Zn]o=0
[Zn] (M)/[Fe] (M) .1 ll0°C
[Zn] (M)/[Fe] (M) 150°C
8.86 8.69 5.13
3.03 5.00 2.89
,1 [Zn] (M)/[Fe] (M) in concentrate is 3.15.
25(~ 4o!
5. (>
i
~ ~01 z 4.0
30
i
3.0 .~ a,e
• Fe
4~
• Pb
~" 20 0_
(.9
6 h
2.0
"I------"i~ 10
c~
1.0
110
120
l ~ , 130 140
~150
Temperature, °C
Fig. 6. Effect of temperature on compositionof jarosite precipitated after 90 min of leaching: no surfactant, [Zn](,= 120 g/l, [H~SO4]o=35g/l, 141 g/1 concentrate, 750 kPa 0.2. sulphate lines were evident in the diffraction pattern. Although Dutrizac and Kaiman noted lower silver concentrations in jarosites prepared below 140 °C than in material precipitated at higher temperatures, they were investigating pure silver-hydronium solid solutions with much higher silver concentrations than those developed in these leach solutions. The composition of iron precipitates shown in Fig. 6 reflects the effect of both temperature and changing leach solution chemistry. The composition of jarosites precipitated at different temperatures, from solutions of constant composition, are shown in Fig. 7. Although the lead and silver content of these precipitates also decreased with increasing hydrolysis temperature, the effect was not as marked as in Fig. 6. The concentrations of zinc and copper substituting for Fe (III) also decreased with increasing temperature. The final concentration of H2SO4 was measured after hydrolysis for 90 min at different temperatures. This concentration, which increases with increasing temperature, reflecting the increase in the equilibrium constants for hydrolysis, is plotted in Fig. 8, along with the final H2SO4 concentration from 90 min leaching-precipitation tests using an initial H2SO4 concentration of 35 g/l. The final concentration in the hydrolysis tests is indicative of the amount of acid released by iron hydrolysis and precipitation when leaching with no surfactant. Although the lowest H~SO4 concentration of 7.1 g/1 at 135 ° and 150 ° C is too high for complete iron precipitation, it does represent substantial neu-
257 40
5.0
[] Zn
• Fe
AAO
I~
~7 CU
•
30
~,. O + + O
4.0
S042-0H -
3.0
~=
if3
"T 2O 2.0
123 O3
x~ Q_ " U_
tO 1.0
I
I
I
I
I
li0
120
~3o
140
150
Temperature,
°C
Fig. 7. Effect of temperature on jarosite precipitated in hydrolysis tests, after 90 m L m [Zn 2÷ ]o = 147
g/l, [Pb2+]o=52 gll (as PbSO4), [Fe3+]o=18.9 g/l, [Cu2+]o=17.5 g/l, [Cd2+]o=0.5 g/l, [Ag + ] =0.4 g/l, [H2S04]o=0 g/1.
40
30
LJ
r- 2O 0 (J
J fu --r
to
I
I
I
I
t20
t40
t60
t80
Temperature,
°C
Fig. 8. Effect of temperature on final acid concentration:
Hydrolysis: [Zn 2+ ]o= 147 g/l, [Fe :~÷ ] 0 = 8 g/l, [H2S04]o=0 g/l; Leaching: [Zn 2+ ]o = 120 g/l, [Fe 3+ ] 0 = 8 g/l, [H2SO4]o=35 g/1.
tralization. Since n e u t r a l i z a t i o n results f r o m t h e t b r m a t i o n of e l e m e n t a l sulphur, which as seen in Fig. ,,,o is n o t t h e r m o d y n a m i c a l l y f a v o u r a b l e at such a high p H , the o v e r p o t e n t i a l for s u l p h a t e f o r m a t i o n m u s t clearly be high, even at 1 5 0 C . CONCLUSIONS S u b s t a n t i a l n e u t r a l i z a t i o n of acid, a n d p r e c i p i t a t i o n of iron c a n be a c h i e v e d in oxidative p r e s s u r e leaching of a zinc sulphide c o n c e n t r a t e . It t h u s a p p e a r s t h a t an a d d i t i o n a l leach could be used to a p p r e c i a b l y lower the n e e d for external n e u t r a l i z i n g a g e n t s in a f r e e - s t a n d i n g p r e s s u r e leach plant. U n d e r t h e conditions tested, t h e m i n i m u m final acid c o n c e n t r a t i o n was a c h i e v e d b y leaching at 150 ~C, with no s u r f a c t a n t . T h e toss of values to t h e j a r o s i t e residue f o r m e d d u r i n g leaching d e c r e a s e d w i t h i n c r e a s i n g t e m p e r a t u r e . T h e r e was no a d v a n tage in leaching at 180 ° C, however, b e c a u s e a g g l o m e r a t i o n of t h e u n l e a c h e d particles, a n d p r e c i p i t a t i o n of FeSO4, gave slower l e a c h i n g a n d n e u t r a l i z a t i o n kinetics. A galvanic i n t e r a c t i o n b e t w e e n s p h a l e r i t e a n d FeS2 p r o m o t e d zinc leaching at 110 ° C, b u t was less p r o n o u n c e d , or a b s e n t , at higher t e m p e r a t u r e s . ACKNOWLEDGEMENT S u p p o r t of this w o r k by t h e N a t i o n a l Science F o u n d a t i o n , u n d e r G r a n t C P E 8404653, is gratefully a c k n o w l e d g e d . REFERENCES Argall, G.O., Jr., 1986. Perseverance and winning ways at McLaughlin gold. Eng. Min. J., Oct. 1986, pp. 26-32. Berezowsky, R.M.G.S. and Weir, D.R., 1983. Pressure oxidation for treating refractory uranium and gold ores. 22nd Annu. Conf. Metallurgists of CIM, Edmonton, August, 1983. Biegler, T. and Swift, D.A., 1979. Electrochim. Acta, 24: 415. Bolton, G.L., Zubryckyj, N. and Veltman, H., 1979. Pressure leaching process for complex zinclead concentrates. Proc. 13th Int. Mineral Processing Congress, Warsaw, Vol. 1, pp. 581-607. Bruhn, G., Gerlach, J. and Pawlek, F., 1965. Untersuchungen iiber die LSslichkeiten von Salzen und Gasen in Wasser und wi~ssrigen LSsungen bei Temperaturen oberhalb 100 ° C. Z. Anorgan. Allgem. Chem., 337: 68-79. Chmielewski, T. and Charewicz, W.A., 1984. The oxidation of Fe (II) in aqueous sulphuric acid under oxygen pressure. Hydrometallurgy, 12: 21-30. Cobble, J.W., 1964. The thermodynamic properties of high temperature aqueous solutions, VI. J. Am. Chem. Soc., 86: 5394-5401. Connick, R.E. and Powell, R.E., 1953. The entropy of aqueous oxy-anions. J. Chem. Phys., 21: 2206-2207. Corriou, J.P. and Kikindai, T., 1981. J. Inorg. Nucl. Chem., 43: 9-15. Criss, C.M. and Cobble, J.W., 1964a. The thermodynamic properties of high temperature aqueous solutions, V. J. Am. Chem. Soc., 86: 5390-5393. Criss, C.M. and Cobble, J.W., 1964b. The thermodynamic properties of high temperature aqueous solutions, IV. J. Am. Chem. Soc., 86: 5385-5390. Crundwell, F.K., 1987a. The Effect of Flocculant Additions to the Leaching Solution on the Kinetics of the Oxidative Dissolution of Sphalerite. Mintek Report M298, Council for Mineral Technology, Randburg, South Africa.
259 Crundwell, F.K., 1987b. Kinetics and mechanism of the oxidative dissolution of a zinc sulphide concentrate in ferric sulphate solutions. Hydrometallurgy, 19: 227-242. Doyle, B.N., Masters, I.M., Webster, I.C. and Veltman, H., 1978. Acid pressure leaching of zinc concentrates with elemental sulphur as a by-product. Proc. 1 lth Commonwealth Mining and Metallurgical Congress, Hong Kong. IMM, London, pp. 645-653. Dutrizac, J.E., 1980. The physical chemistry of iron precipitation in the zinc industry. In: J.M. Cigan, T.S. Mackey and T.J. O'Keefe (Editors), Lead-Zinc-Tin '80. AIME, pp. 532-564. Dutrizac, J.E., 1983. Jarosite compounds and their application in the metallurgical industry. In: K. Osseo-Asare and J.D. Miller (Editors), Hydrometallurgy - - Research, Development and Plant Practice. AIME, pp. 531-551. Dutrizac, J.E., 1984. The behaviour of impurities during jarosite precipitation. In: R.G. Bautista (Editor), Hydrometallurgical Process Fundamentals. Plenum Press, New York, N.Y., pp. 125169. Dutrizac, J.E. and Dinardo, 0., 1983. The co-precipitation of copper and zinc with lead jarosite. Hydrometallurgy, 11: 61-78. Dutrizac, J.E. and Kaiman, S., 1976. Synthesis and properties of jarosite-type compounds. Can. Mineral., 14: 151-158. Dutrizac, J.E., Dinardo, O. and Kaiman, S., 1980. Factors affecting lead jarosite formation. Hydrometallurgy, 5: 305-324. Ferreira, R.C.H., 1975. High temperature E-pH diagrams for the systems S-H20, Cu-S-H20 and Fe-S-H20. In: A.R. Burkin (Editor), Leaching and Reduction in Hydrometallurgy. IMM, London, pp. 67-83. Forward, F.A. and Veltman, H., 1959. Direct leaching zinc sulphide concentrates by Sherritt Gordon. J. Met., 11" 836-840. Frenay, J.N., 1984. Leaching of sphalerite: influence of its iron content. In: W.C. Park, D.M. Hansen, R.D. Hagui (Editors), Applied Mineralogy. Proc. 2nd Int. Congr. Applied Mineralogy in the Minerals Industry, Los Angeles, 1984. TMS-AIME, pp. 531-545. Haung, H.H. and Bernal, J.E., 1984. Kinetic study on direct leaching of sphalerite in sulphuric acid solution using ferrous sulphate as the catalyst. In: P.E. Richardson, S. Srinivasan, R. Woods (Editors), Proc. Int. Symp. Electrochemistry in Mineral and Metal Processing. The Electrochemical Society, pp. 469-485. Huffman, R.E. and Davidson, N., 1956. Kinetics of the ferrous iron-oxygen reaction in sulphuric acid solution. J. Am. Chem. Soc., 78: 4836-4842. Iwai, M., Majima, H. and Awakura, Y., 1982. Oxidation of Fe (II) in sulphuric acid solutions with dissolved molecular oxygen. Metall. Trans. B, 13B: 311- 318. Johnston, B.H., 1983. The application of Sherritt zinc pressure leach technology at the Kidd Creek zinc plant. Zinc '83, 13th Annu. Hydrometallurgical Meeting. CIM Metallurgical Society, Edmonton, Alta., August 1983, Paper 11. Kammel, R., Pawlek, F. and Simon, M., 1987. Oxidizing leaching of sphalerite under atmospheric pressure. AIME Annual Meeting, Denver, TMS Paper A87-12. Latimer, W.M., 1952. Oxidation Potentials. Prentice-Hall, New York, N.Y., 392 pp. Lotens, J.P. and Wesker, E., 1987. The behaviour of sulphur in the oxidative leaching of sulphidic minerals. Hydrometallurgy, 18: 39-54. Lowson, R.T., 1982. Chem. Rev., 82 (5): 461-499. Mackiw, V.N. and Veltman, H., 1967. Recovery of zinc and lead from complex low-grade sulphide concentrates by acid pressure leaching. CIM Bull., 70 (Jan.): 16-21. Majima, H. and Awakura, Y., 1986. Water and solute activities of H2SO4-Fe2(SO4).~-H20 and HC1-FeCI:~-H20 solution systems, II. Activities of solutes. Metall. Trans., B, 17B: 621-627. Martin, M.T. and Jankola, W.A., 1985. Cominco's Trail zinc pressure leach operation. CIM Bull., (April), pp. 77-81. Mathews, C.T. and Robins, R.G., 1972. The oxidation of aqueous ferrous sulphate solutions by molecular oxygen. Proc. Aust. Inst. Min. Metall., 242.
Monhemius, A.J., 1987. Recent advances in the treatment of refractory gold ores. lI Meeting of the Southern Hemisphere on Mineral Technology, Rio de Janeiro, May 1987. Nattrass, W.A., Coetzee, C.F.B. and Verbaan, B.. 1983. Pressure leaching of sphalerite below the melting point of sulphur: an integrated process. Zinc '83, 13th Annu. Hydrometallurgical Meeting, CIM Metallurgical Society, Edmonton, Alta., August 1983, Paper 13. Parker, E.G., 1981. Oxidative pressure leaching of zinc concentrates. CIM Bull., 74 tMay): 145 150. Parker, E.G. and Romanchuk, S., 1980. Pilot plant demonstration of zinc sulphide leaching. In: J.M. Cigan, T.S. Mackey, T.J. O'Keefe (Editors), Lead-Zinc-Tin '80. AIME, pp. 407 425. Parker, E.G., McKay, D.R. and Salomon-De-Friedberg, H , 1983. Zinc pressure leaching at Cominco's Trail operation. In: K. Osseo-Asare and J.D. Miller (Editors), Hydrometallurgy, Research, Development and Plant Practice. AIME, pp. 927-940. Perry, J.H., 1963. Chemical Engineers Handbook. McGraw-Hill, New York, N.Y., 4th ed. Peters, E., 1984. Electrochemical mechanisms for decomposing sulphide minerals. In: P.E. Richardson, S. Srinivasan and R. Woods (Editors), Proc. Int. Symp. Electrochemistry in Mineral and Metal Processing. The Electrochemical Society, pp. 343-361. Peters, E., 1986. Leaching of sulfides. In: P. Somasundaran (Editor), Advances in Mineral Processing. SME, Colorado, pp. 445 462. Piao, S.Y. and Tozawa, K., 1985. Effect of iron content in zinc sulphide concentrates on zinc extraction in oxygen with elemental sulphur. J. Min. Metall. Inst. Jpn., 101: 795-800. Posnjak, E. and Merwin, H.E., 1922. The system, ferric oxide-sulphur trioxide-water. J. Am. Chem. Soc., 44: 1965-1994. Pradip, T., 1981. The Surface Properties and Flotation of Rare-Earth Minerals. Ph.D. Thesis, University of California, Berkeley. Romanikaw, L.T. and de Bruyn, P.L., 1963, Kinetics of dissolution of zinc sulphide in aqueous sulphuric acid. In: M.E. Wadsworth and F.T. Davis (Editors), Unit Processes in Hydrometallurgy. Gordon and Breach, London, pp. 45-66. Sato, M., 1960. Oxidation of sulphide ore bodies, II. Oxidation mechanisms of sulphide minerals at 25°C. Econ. Geol., 55: 1202-1231. Scott, T.R., 1973. Continuous, co-current, pressure leaching of zinc-lead concentrates under acid conditions. In: D.J.I. Evans and R.S. Shoemaker (Editors), Int. Symp. Hydrometallurgy, Chicago, AIME, pp. 718-750. Stanczyk, M.H. and Rampacek, C., 1961. Dissolution of zinc from sphalerite at elevated temperatures and pressures. U.S. Bureau of Mines, R.I., 5848. Tozawa, K. and Piao, S., 1987. Effect of iron content in zinc sulfide concentrates on zinc extraction in oxygen pressure leaching with elemental sulfur. Metatl. Rev. MMIJ, 4 (2): 89-105. Verbaan, B. and Crundwell, F.K., 1986. An electrochemical model for the leacing of a sphalerite concentrate. Hydrometallurgy, 16: 345-359. Wadsworth, M.E., 1972. Advances in the leaching of sulphide minerals. Miner. Sci. Eng., 4 (4): 36-47. Wagman, D.D., Evans, W.H., Parker, V.B., Halow, 1., Baily, S.M. and Schumm, R.H., 1968. Selected Values of Chemical Thermodynamic Properties. NBS Tech. Note, 270-3. Walter-Levy, L. and Quemeneur, E., 1964. Sur l'hydrolyse du sulfate ferrique ~ 100 °. C.R. Acad. Sci. Paris, 258: 3028-3029. Walter-Levy, L. and Quemeneur, E., 1966. Etude de l'hydrolyse du sulfate ferrique de 25 ~ 200 °. Bull. Soc. Chim. France, 1947-1954. Weast, R.C., 1978. CRC Handbook of Chemistry and Physics. CRC Press, Inc., Boca Raton, Fla., 59th ed. Weir, D.R. and Berezowsky, R.M.G.S., 1983. Gold extraction from refractory concentrates. 14th Annu. Hydrometallurical Meeting. CIM Metallurgical Society, Timmins, Oct. 1983. Yaroslavtsev, H.S., Getskin, L.S., Usenov, A.U. and Margulis, E.V., 1975. Behaviour of impurities when precipitating iron from sulphate zinc solutions. Tsvetn. Met., 16 (4): 40-41.
241
Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands
Iron Removal During Oxidative, Acid Pressure Leaching of a Zinc Sulphide Concentrate FIONA M. DOYLE, HERNANDO ARAUCO*' and LUCIA M. FENG
University of Cali[ornia at Berkeley, Department of Materials Science and Mineral Engineering, Hearst Mining Building, Berkeley, CA 94720 (U.S.A.) (Received September 4, 1987; accepted after revision June 9, 1988)
ABSTRACT Doyle, F.M., Arauco, H. and Feng, L.M., 1989. Iron removal during oxidative, acid pressure leaching of a zinc sulphide concentrate. Int. J. Miner. Process., 25: 241-260. Zinc sulphide leaching is an example of an oxidative, acid pressure leaching process that must be followed by neutralization of excess acid, in this case to remove iron from solution. The leaching behaviour of a zinc sulphide concentrate in moderately acidic solutions (35 g/1 H2S04) has been studied at temperatures from 110 ° to 180 ° C to determine the extent of neutralization that can be achieved by the reaction: Z n S + 2 H + + ½02 = Z n ~+ + H 2 0 + S ° Extensive neutralization was observed at 150°C, even at pH values well outside of the thermodynamic stability region of sulphur. Neutralization was less effective at the other temperatures. Iron precipitated from solution as jarosite. The lead and silver content of these jarosites decreased markedly with increasing temperature. Substitution of other metals in jarosites was studied in hydrolysis tests, and also decreased with increasing temperature.
INTRODUCTION
Acid pressure leaching is being used increasingly in the commercial processing of sulphide materials. Extensive development work has been done on acid pressure leaching of zinc sulphide concentrates (Forward and Veltman, 1959; Stanczyk and Rampacek, 1961; Romanikaw and de Bruyn, 1963; Mackiw and Veltman, 1967; Scott, 1973; Doyle et al., 1978; Bolton et al., 1979; Parker and Romanchuk, 1980; Parker, 1981; Parker et al., 1983; Johnston, 1983; Nattrass et al., 1983; Martin and Jankola, 1985; Verbaan and Crundwell, 1986; Crundwell, 1987a,b; Tozawa and Piao, 1987) and the oxidative pressure leaching of arsenical and pyritic gold ores (Berezowsky and Weir, 1983; Weir and Bere*'Present address: International Rectifier, 233 Kansas Street, E1 Segundo, CA 90245 (U.S.A.)
0301-7516/89/$03.50
© 1989 Elsevier Science Publishers B.V.
2-12
zowsky, 1983; Argall, 1986; Monhemius, 1987 }. The advantages of' acid pressure leaching include the ability to produce elemental sulphur, rather than sulphate or SO~, and the potential use for mixed sulphide concentrates. However, neutralization is often needed after the leach stage; for example, zinc pressure leaching is followed by neutralization to precipitate iron prior to electrowinning, while gold ores must be neutralized before cyanide leaching. The use of an external neutralizing agent such as lime can add significantly to the operating costs of a process, and increase the volume of waste requiring disposal. It has been proposed that a two-stage, countercurrent leach could be used for partial neutralization in zinc sulphide leaching (Bolton et al., 1979; Parker, 1981 ). This relies on acid consumption in the overall leaching reaction: Z n S + H + + 1/20~ = Z n '-'+ + H 2 0 + S ~
(1)
Elemental sulphur, however, is only thermodynamically stable over a narrow range of potentials and pH values. Although this stability range is extended, in practice, because of the high overpotential needed for formation of sulphate (Peters, 1984, 1986), significant limits on the extent of neutralization that could actually be achieved might be expected, particularly at elevated temperatures where kinetic effects are less dominant. Neutralization would promote precipitation of iron as jarosite, a basic iron sulphate. Jarosites may contain appreciable levels of lead and silver, and smaller amounts of other metals that are likely to be present in mixed concentrates. This leads to a loss of values, which may seriously affect the economic viability of' pressure leaching as a means of treating these increasingly common materials. The objective of this work was to determine the extent of neutralization that can be achieved in zinc pressure leaching, and to examine the effects of simultaneous leaching and neutralization on the purity of the iron residues.
Leaching of zinc sulphide Two Canadian zinc producers, Cominco and Kidd Creek Mines, have now integrated direct leaching in their hydrometallurgical zinc circuits (Parker et al., 1983; Johnston, 1983; Martin and Jankola, 1985). Both use a high temperature, oxidative acid pressure leach; ground concentrate is fed into an autoclave with hot, spent electrolyte, typically containing 120 g/l H2SO4 and 50 g/1 Zn as zinc sulphate, to give a pulp density of 145 g/1. The oxygen partial pressure is maintained at 750 kPa (109 psi). Reaction 1, which predominates under these conditions, is strongly exothermic, and the temperature rises to about 145 °-150 ° C. Quebracho and lignin sulphonate are used to prevent the sulphur, which melts at 119 ° C, from occluding unleached sulphide particles, thereby preventing zinc extraction beyond about 50% (Doyle et al., 1978). The temperature is not allowed to exceed 155 °C because polymerization of the sul-
243
phur becomes problematic. The residence time is typically about 90 min. The leach residue is treated to separate elemental sulphur from unleached sulphides, lead sulphate and iron residues, while the pregnant leach solution, containing about 120 g/1 Zn, 35 g/1 H2SO4 and 7 g/1 Fe(III), goes to the main circuit for neutralization with calcine and purification prior to electrowinning (Parker et al., 1983; Martin and Jankola, 1985). The Eh-pH diagrams for the Zn-S-HeO system at 25 ° and 150 ° C are shown in Figs. 1 and 2. These have been calculated using the data in Table I. Although commercial leach solutions have a typical concentration of 2.0 M ZnSO4, the mean activity coefficients of concentrated ZnS04 solutions are very low (only 0.043 at 1.0 M) (Weast 1978), hence the activity of ZnS04 increases very slowly with increasing concentration. Activities of 0.1 M for all soluble sulphur and zinc species were, therefore, used for these diagrams. Under acidic conditions, there are three different leaching reactions for ZnS: (i) Z n S + 2 H + =Zn z+ "bH2S
(2)
t.5
1.0
Hs0~ Zn 2+
0.5
s042-
o3
'"s i....
d 0
rtlJ
ZnO
0.0
-0.5
ZnS -t
... HS-
.0 I
I
0
14
pH Fig. 1. E h - p H diagram for the Z n - S - H 2 0 system at 25 ° C. Activities of all soluble zinc and sulphur species = 0.1M.
244
1.5
i'O I 4C0~
0"5 t
Zn 2+
I-_J
0> tLU
ZnO
0.0 \
c
\
N \ \ \
\ \ \
--0.5
\ \
b~s
\
ZnS
\ \
\.
"x
--i.0 ,
2
%
4
i , i 6 8
,'0
,'2
,4
pH Fig. 2. Eh-pH diagram for the Zn-S-H20 system at 150°C. Activities of all soluble zinc and sulphur species= 0.1M. This occurs at pH < - 1.36 at 25 °C and pH ~< - 0 . 7 2 at 150 °C at 0.1 atm H2S, and thus does not appear on Figs. 1 and 2. (ii) Z n S = Z n 2+ + S ° + 2 e
(3)
(iii) Z n S + 4 H 2 0 = Z n 2+ + S O 2- + 8 H + +Be
(4)
which is also equivalent, in terms of neutralization behaviour, to: (iiia) Z n S + 4 H 2 0 = Z n 2+ +HSO~- + 7 H + + 8 e
(5)
In a chemical leaching system the anodic half-cell reactions 3 and 4 must be coupled with a cathodic reaction. The direct reduction of oxygen, for which the overall reaction is: 02 +4H + + 4 e = 2H20
(6)
occurs as a series of single-electron transfer steps, with high activation energies (Wadsworth, 1972 ), and is subject to strong activation polarization. Low over-
245 TABLE I Thermodynamic data for the Z n - S - H 2 0 system at 298 and 423 K z~G298 o .~
Species
8°98 *1
(kcal/mol) (cal/mol K } (unless specified) H÷ H=, O~ Zn Slrh,,mb) H~O ZnS ZnO H2Sta~ HS S ~HSO~ SO~Zn" ÷ HZnO~ ZnO~-
0.00 0.00 0.00 0.00 *'~ 0.00 - 56.69 - 48.11 *'~ - 76.08 *'~ -6.54 3.01 22.10 - 179.94 - 177.34 - 35.14 - 109.26 *'~ - 9 1 . 8 5 *'~
0.00 31.211 49.003 9.95 *5 7.62 16.716 13.80 .5 10.43 *'~ 29.20 14.60 - 6.40 30.32 4.10 - 26.80 10.00 *s - 23.00 *s
423 Cp0 129s -(cal/mol K)
33.00 *:~ 6.90 *4 8.30 *4 6.07 (298) *5 6.40 *4 18.14 *~ 11.00 ( 298 ) .5 9.62 (298)'5 61.00 *6 - 62.00 *:~ - 61.00 .7 - 18.00 *:~ - 105.00 *:~ 67.71"7 - 121.65 .7 - 162.51.7
0 ,2 zJG42:~
(kcal/mol) - 0.765 - 4.061 - 6.318 - 1.384 -- 1.101 - 59.200 - 50.090 - 77.607 - 11.60 2.622 24.313 - 183.313 - 175.420 - 33.359 - 107.691 -85.210
*'Latimer, 1952. *~Calculated from zlG~2.~= zIG°gs - (423 - 298)zlS°gs - ~C ° L~ ( 4 2 3 - 2 9 8 - 4231n423/298). *:~Criss and Cobble, 1964a. *4Perry, 1963. *'~Wagman et al., 1968. *(~Cobble, 1964. *TCalculated from Cpo k29s 42:~= t~ (T) + fl(T ) 8298(abs) o using constants from Criss and Cobble, 1964b. *SEstimated using the relationship S°gs( ..... ) = 4 3 . 5 - 46.5 ( z - 0.28n) given by Connick and Powell (1953} for oxyanions (z is the ionic charge, n is the number of oxygen atoms not included in hydroxide groups). a l l l e a c h i n g r a t e s a r e o b s e r v e d i n s y s t e m s w h e r e o x y g e n is t h e o n l y a v a i l a b l e oxidant. The reduction of iron (III): F e 3+ + e = F e
2+
(7)
shows minimal polarization until the current density becomes high enough for concentration polarization to be important (Peters, 1984, 1986). Much higher leaching rates are observed for concentrates containing pyrite, which releases i r o n ( P a r k e r a n d R o m a n c h u k , 1 9 8 0 ; P a r k e r e t al., 1 9 8 3 ) . H o w e v e r , t h e s e f a s t e r leaching rates may be due to effects other than simple depolarization of the cathodic reaction. Eh-pH d i a g r a m s f o r p y r i t e a t 25 ° a n d 1 5 0 ° C s h o w t h a t p y r i t e is o x i d i z e d t o F e 2+ a n d e l e m e n t a l s u l p h u r a t a b o u t 0 . 2 5 a n d 0 . 1 5 V, respectively, potentials that are fairly close to those for the same reaction of s p h a l e r i t e ( F e r r e i r a , 1 9 7 5 ) . T h e m e a s u r e d r e s t p o t e n t i a l s o f p y r i t e a t 25 ° C a n d
2,1 6
low pH, however, are about: 0.62 V (Sato, 1960; Peters, 1984 t. The oxidation of pyrite involves a volume increase, and the discrepancy between the predicted and measured potentials has been ascribed by Peters (1984) to passiration by the compact surf'ace layer of elemental sulphur. Sato (1960) was unable to measure the rest potential of sphalerite in acid solutions, because of' its low conductivity, but it is likely to be much closer to the predicted value, because oxidation involves a volume decrease, leading to a porous sulphur layer. Galvanic effects resulting from the potential difference between pyrite and sphalerite are also likely to be responsible for the faster leaching rates observed for concentrates containing pyrite (Peters, 1984). It has also been observed that the leaching rate of pure sphalerite increases linearly with increasing lattice iron (Frenay, 1984; Piao and Tozawa, 1985; Tozawa and Piao, 1987; Kammel et al., 1987 ), because the iron reduces the band-gap energy of the sphalerite and supplies holes in a d-orbital impurity band (Crundwell, 1987a,b). Iron (II) resulting from reaction 7 is oxidized in the bulk solution: 4Fe 2+ + 0 2 + 4 H + =4Fe :~+ +2H~O
(8)
Thus the F e ( I I I ) / F e ( I I ) couple is merely an intermediate, with the overall cathodic reaction being reduction of oxygen. Combining eqs. 3 and 4 with 8 gives: Z n S + 1/202 + 2 H + = Z n 2+ + S U + H 2 0
(9)
and: ZnS + 202 = Z n 2+ +SO~-
(10)
Clearly, reactions 2 and 9 will neutralize excess acid, whereas reaction 10 will not. Simple dissolution (reaction 2) occurs at pH's too low to be of any practical utility. Thus an additional leach stage will only achieve the desired goal of neutralizing acid if elemental sulphur is the main reaction product. It is clear from Fig. 1 that at 25 °C this is only thermodynamically favourable at pH values below 1. At 150°C (Fig. 2) the sulphur stability region contracts, and reaction 9 is only favourable at very low pH. Thus, it may be concluded that little neutralization is thermodynamically possible in an oxidative, acid pressure leach. Experimentally, however, it is observed that below its melting point, elemental sulphur ($8) is much more stable with respect to oxidation than is indicated in Fig. 1 (Biegler and Swift, 1979; Corriou and Kikindai, 1981; Lowson, 1982). Peters (1986) has accounted for the high oxidation overpotential by destabilizing sulphate by 300 k J / m o l (72 kcal/mol) when calculating E h pH diagrams for sulphide minerals at 25 ° C. This extends both the sulphur and the sulphide mineral stability regions; his diagram for the Z n - S - H 2 0 systems indicates that at 25 ° C, ZnS will leach according to reaction 9, and hence neutralize excess acid, up to about pH 5.5. However, from the Butler-Volmer equa-
247 tion, the oxidation overpotential will decrease dramatically with increasing temperature. Thus a much closer correlation is expected between the observed behaviour and that predicted by the Eh-pH diagram at 150 ° C than at 25 ° C.
Behaviour of iron during pressure leaching Although it is desirable to leach with appreciable concentrations of Fe (III) in solution, to increase the leaching kinetics, neutralization during leaching will lead to concurrent precipitation of iron. Phase diagrams for the Fe203-SO3-H20 system at temperatures from 25 ° to 200°C (Posnjak and Merwin, 1922; Walter-Levy and Quemeneur, 1964, 1966 ) indicate that hydronium jarosite, H3OFe3 (SO4) 2(OH) 6, is the stable solid phase in equilibrium with concentrated, acidic iron (III) sulphate solutions. Hydronium jarosite is one of the jarosites, which have the general composition MFe3(SO4)2(OH)6, where M may be K +, Na +, NH + , 1/2Pb 2+ and Ag +, as well as H30 +. Potassium jarosite is more stable than ammonium or sodium jarosite, and all are more stable than hydronium jarosite, although the stability of hydronium jarosite increases with increasing temperature (Dutrizac, 1983). In the jarosite process used in conventional electrolytic zinc processing, alkali salts are usually added to decrease the jarosite solubility, thereby promoting precipitation of iron. However, there is a continuous solid solution series between the jarosites (Dutrizac, 1980), and the jarosites produced in metallurgical processes are usually deficient in alkali, even when formed in the presence of excess alkali, due to incorporation of small amounts of H30 +. The extent of hydronium substitution has been reported in different studies to either increase or decrease with increasing precipitation temperature (Dutrizac, 1983 ). Increased substitution would be consistent with the increase in stability of hydronium jarosite with increasing temperature. The behaviour of impurities during jarosite precipitation has been studied at 90 ° C (Yaroslavtsev et al., 1975) and 97 ° C (Dutrizac, 1984). Losses of divalent metals to alkali jarosites increase with increasing divalent metal concentration, increasing pH or decreasing Fe (III) concentration. Because of their similar ionic radii, the divalent metals substitute for Fe 3+ in the jarosite lattice, the charge balance being maintained by loss of O H - . The extent of divalent metal incorporation increases as the concentration of alkali metal during precipitation increases, with losses being greatest for potassium, then ammonium and sodium jarosites. This is because substitution of divalent metals in the jarosite gives a thermodynamic weakness, thus only limited substitution occurs. Substitution in the alkali jarosites, and hence losses, are greatest when the jarosite is most stable, namely when there are high concentrations of alkali, especially potassium, in solution. Losses vary for different metals, but rarely exceed 3%. The descending order for losses is:
248
Fe :~+ >> Cu e+ > Zn '-'+ > Co-' + > Ni e+ > Mn 2* > Cd'-'~
{11)
which is the same as the order fbr decreasing hydrolysis constants. Lead and silver jarosites are clearly undesirable, because neither metal can be recovered without dissolution or recrystallization of the residue. Moreover, losses to lead jarosite are particularly severe, because there are vacant alkali metal sites; when divalent metals replace Fe :~+ the charge balance is maintained by including additional lead, hence the crystal lattice is not disrupted as severely as in alkali jarosites, where O H - is eliminated for neutrality. The order for losses is the same as that for alkali jarosites. Copper levels of 8% have been reported; at this concentration the composition of the precipitate approaches that of beaverite, PbCuFe~(SO4)e(OH)6 (Dutrizac and Dinardo, 1983). Lead sulphate is very insoluble, thus the formation of lead jarosite is kinetically unfavourable at 90 :-100 ° C, and has hitherto been unimportant in jarosite precipitation processes. Lead jarosite does form during pressure leaching at 145 ~-155 ~C, however. Lead jarosite formation could be suppressed by operating at temperatures of 180 °-200 ° C, above the jarosite stability region (Dutrizac et al., 1980), but this is unlikely to be feasible when leaching sulphide concentrates because of sulphur polymerization. Although the equilibrium behaviour of impurities at 90 ° and 97 °C has been thoroughly studied, there are likely to be significant quantitative differences for jarosites precipitated during leaching at higher temperatures, because of the different relative stabilities of the aqueous species and solid phases, and because of the changing solution composition. At temperatures above 150~C, iron may also precipitate as iron(II) sulphate; Bruhn et al. (1965) reported FeSO4 solubilities of 6.5, 2.0, 0.6 and 0.7 g/100 g H~O at 150 ° , 160 ° , 170:' and 180°C. The solubility at 180°C, the highest temperature used in this study, is equivalent to a solubility product of 2.12.10 -:~. Thus the equilibrium Fe z+ concentration in a 2.0 molal sulphate solution would be 1.06" 10- :~M, or about 0.06 g/1. Fe (II) produced by reaction 7 is thus likely to precipitate before it can be reoxidized to Fe (III). Crundwell ( 1987a,b ) has found that the Fe ( III )-Fe ( II ) couple is not reversible on sphalerite in sulphate solutions. Thus the redox potential at the sphalerite surface, and hence the leaching rate, is not directly determined by the Fe ( I I I ) / F e (II) ratio, although Fe(II) does affect the leaching rate through altering the concentrations of the electroactive Fe (III) species. Haung and Bernal (1984) found that most of the iron was present as ferrous during acid pressure leaching of sphalerite. Various workers have observed that the kinetics of reaction 8, oxidation of Fe(II) by oxygen, initially show a second-order (or close to this) dependence on the Fe(II) concentration (Huffman and Davidson, 1956; Mathews and Robins, 1972; Iwai et al., 1982; Chmielewski and Charewicz, 1984; Verbaan and Crundwell, 1986). If similar kinetics apply at 180 ° C, continual
249
precipitation of FeSO4 would therefore lead to very low concentrations of Fe (III), which in turn would decrease the leaching rate. Iron leaching and precipitation will affect the acid balance in a leaching system. FeS2 is generally oxidized directly to sulphate (Lotens and Wesker, 1987), which releases acid: FeS2+H20+7/202
Fe2++2SO42-+2H +
(12)
Precipitation of any hydrolyzed iron phase also releases acid, for example: 3Fe 3+ +2SO~- +6H20=H3OFe3(SO4)2(OH)6 + 5 H +
(13)
Thus continual neutralization is needed to maintain steady precipitation. Conversely, the precipitation of FeSO4 is a simple crystallization reaction, and does not release hydrogen ions. EXPERIMENTALPROCEDURE
Materials
Pressure leach tests used a Sullivan zinc concentrate provided by Cominco. The principal minerals present were sphalerite, marcasite and galena. The chemical composition of the concentrate is given in Table II. 2.16% of the zinc was water-soluble. After further grinding of the as-received concentrate, size analysis with a Microtrac particle size analyzer indicated the 10th, 50th and 90th percentiles of the cumulative size distribution to be 6.6 pm, 23.6 ttm and 59.2 ttm, respectively. The ground concentrate was stored under argon. Leaching tests
Pressure leaching tests were done in a 300-ml Autoclave Engineers stainless steel autoclave equipped with a 150 ml stainless steel, heated, pressurized reservoir. Acid and pulp were heated separately in the reservoir and autoclave, respectively, then mixed and pressurized with oxygen to initiate leaching. (A moderate partial pressure of oxygen was maintained in the reservoir during heating, to maintain passivity of the steel.) Tests were done at 110 °-180 ° C, nominally with 35 g/1 H2SO4, 7 g/1 Fe(III), pulp densities of 141 or 193 g/l, initial zinc concentrations of 0 or 120 g/1 (as sulphate), oxygen partial presTABLE II Chemical analysis of Sullivan concentrate (wt%) Zn
Pb
Cu
Fe
Ni
Co
Mn
Cd
Ag
Mg
A1
K
Ca
Na
S
SiO2
50.1
4.06
0.25
13.6
.004
.005
0.24
0.13
.007
0.20
0.02
0.03
0.08
0.09
29.6
1.48
25()
sures of 345 or 750 kPa, oxygen flow rates of 35-45 cm:~/min or 15-25 cm:~/ rain, using 0.1 g/1 magnesium lignin sulphonate and 0.2 g/1 quebracho as surihctants in some tests. The injection procedure made precise control over the final solution composition difficult; t = 0 data reported in the figures give the exact compositions.
Hydrolysis test~' To determine the effect of temperature alone on the deportment of impurities to the jarosite residues, sulphate solutions containing 147 g/1 Zn 2+, 18.9 g/1 Fe ~+, 17.5 g/1 Cu 2+, 0.5 g/1 Cd 2+, 0.4 g/1 Ag +, 0 g/l H2SO4 and PbS04 equivalent to 52 g/1 Pb 2+ were held at temperature for 90 min. The final solution and precipitates were analyzed.
Analytical procedures Major mineral phases were characterized by X-ray powder diffraction, using a diffractometer. Samples of concentrate were washed with water to determine the concentration of water soluble zinc, digested in HCI and HN03, then treated with ammonium acetate to dissolve any lead. Residues and precipitates were treated with carbon disulphide or by flotation, to remove elemental sulphur {where appropriate), followed by digestion in HNO3. The solutions generated by these procedures were analyzed by atomic absorption spectrophotometry. Solution samples withdrawn during leaching were centrifuged to remove all solids, the room temperature pH was measured, then metal concentrations were determined by atomic absorption spectrophotometry. Sulphate was determined by barium sulphate precipitation, using disodium rhodizonic acid as an indicator for barium. Total acid was determined by potentiometric titration with KOH. RESULTS AND DISCUSSION
Kinetic data for leaching Sullivan concentrate at 110 ° and 150°C under conditions representative of a first stage are shown in Fig. 3. There was a slight increase in the initial rate of zinc and iron leaching with increasing temperature. The pH after 90 min {measured at room temperature) was slightly higher at 150 ° C, but the pH was never high enough at either temperature to promote precipitation of iron. Similar results were obtained at 120 ° and 135 ° C. Fig. 4 shows kinetic data obtained using similar initial concentrations of acid, zinc and iron, but without surfactant. At l l 0 ° C the leaching rates were slightly faster than those shown in Fig. 3, and the final pH was somewhat higher. At 150 ° C, the absence of surfactant gave significantly faster leaching rates. After about 60 min iron started precipitating as jarosite. The pH in-
251
I
2.6
(a)
(b)
0.25 2.4
O. 2 0 v~E
:E . u c o (J
0.i5 o u 0~ LL
2.2
f
Z'f
c
O. iO
N
2.0
0.05 i.B
o.oo
2'o 4'o 8'o 8o TIME (mins)
2o
7o e'o s'o TIME (mins)
2.0
(c)
i.5
Icl
t.0
0.5 • •
o.o
liO ° 1500
go
go
TIME (mins) Fig. 3. Leaching kinetics for Sullivan concentratewith surfactant: [Zn]o= 120 g/l, [H2SO4]0=35 g/l, 141 g/1 concentrate, 750 kPa 02. creased rapidly to 1.44, then decreased slightly as jarosite precipitated (reaction 13). At 180°C the iron leaching rate was significantly faster, but the zinc concentrations, which are not shown, were inconclusive, because of pronounced scatter. The room temperature pH increased to 3.5 before apparent precipitation of iron, which is inconsistent with the solubility product of jarosite. The
252
2.6 F - - . . . . . . . . . . . . . . . . . . . . . . . .
0.25 2.4 2;"
I i
•
//
0.20 2.2
0.15 o {.9
\ b_
~=-i
/ii1 /
!
0.10
ou c N
./° 2.O
0.05 (a) 0.00
1.8
2~0
4=0 TIME
6'0
B'O
I/ 2=0
(mins)
4LO 6'0 8'0 TIME (mins)
3.5 3.0 2,5 2.0
1.0 • • •
0.5
0. C
2'0
4'0
610
tt0 ° i50 ° t80 °
8=0
TIME ( r a i n s ) Fig. 4. L e a c h i n g k i n e t i c s f o r S u l l i v a n c o n c e n t r a t e g/1, 141 g / l c o n c e n t r a t e , 750 k P a 02.
with no surfaet~t:
[ Zn ] o = 120
g/l,
[ H~SO4 ] o = 35
leach residue from the 180 ° C test was agglomerated, indicating that a substantial amount of elemental sulphur had formed and polymerized. The variation in the zinc and iron concentrations, and the unusually high pH are thought to be due to precipitation of zinc and ferrous sulphate. The solubility of ZnSO4 at 180°C is 34.8 g/100 g H20, or 2.16 molal (Bruhn et al., 1965). This solubility would be reduced by the high total sulphate concentration, and zinc sulphate
253 would .have precipitated as leaching proceeded. Any solid zinc sulphate withdrawn with aqueous samples would have redissolved on cooling. Thus the apparent concentration would have been determined largely by the amount of solid ZnS04 withdrawn in each sample. As discussed earlier, the low solubility of FeSO4 at this temperature would have caused continual depletion of both Fe (II) and Fe (III), leading to slower leaching. Higher pH values were reached before precipitation of Fe (III), because of its low concentration. Surfactants are used in commercial zinc sulphide pressure leaching to prevent occlusion of unleached particles by liquid sulphur (Doyle et al., 1978; Parker, 1981 ). Although faster leaching and neutralization rates were observed in this work in the absence of surfactant, the zinc extraction did not exceed 50%, the level at which liquid sulphur has been reported to prevent further reaction (Doyle et al., 1978). Crundwell (1987a) reported that addition of 0.05 wt% MAGNAFLOC 351, and industrial flocculating agent, guar gum or sodium isobutyl xanthate, a flotation reagent, all decreased the rate of oxidative leaching of sphalerite at 78 ° C, and attributed this effect to a decrease in surface area available for charge transfer resulting from adsorption of the organic material. It appears that the quebracho and lignin sulphonate used here acted similarly. These surfactants have a large number of polar groups, and are used as depressants in flotation (Pradip, 1981). They are thought to promote coalescence of sulphur by adsorbing on the sulphide minerals, preventing wetting by the hydrophobic sulphur. Even if the surfactants were oxidized during leach.ing, they would probably act similarly, since the carboxylic acid groups that would form from the primary alcohol groups are also strongly polar. The presence of the surfactant was least pronounced at l l 0 ° C . At this temperature much of the surfactant would be expected to desorb from the porous, solid sulphur layer forming on the mineral surfaces. It is therefore concluded that in a two-stage leach process, surfactant should be added to the residue from the first stage leach, not to the feed. Kinetic data for tests done with no zinc sulphate initially, but similar concentrations of acid and Fe2 (SO4)3 to those used in the other tests, are given in Fig. 5. Comparison with Figs. 3 and 4 indicates that iron leaching was somewhat faster in the absence of zinc sulphate, whereas the zinc leaching rate was relatively unaffected. The pH was significantly higher in these tests, and this gave earlier precipitation of iron. However, the leaching rates suggest that the higher pH is not due to faster neutralization; instead the pH is thought to be higher because of the decrease in H2SO4 activity with increasing Z n S Q concentration reported by Majima and Awakura (1986). The fact that iron precipitated at l l 0 ° C in the presence of surfactant indicates that the surfactant did not interfere with the nucleation of hydrolyzed iron (III) phases; the lack of precipitation in the tests shown in Fig. 3 was due only to the low pH. This was confirmed by adding 85 g/1 sodium jarosite seed to two leaching runs using
254 0.30 ............................ 0.8
0.25
J
0.20
~ o61
(3 0 . t 5 C o [9 O9 b_ 0 . i 0
u°
0.4
0.05 (a) ~ ~ _ 0.00
2`O
4`O
60
8`O
TIME (rains)
o.o~ 2'o 4'o 6'o 8'o TIME (rains)
2.0
(c)
1.5
i.0
0.5 $iiO
0.0
° no surfsctant
•
l i O ° 8urfactant
•
i50° no surfactant
2'0 2o ~`O ~`O TIME
(rains)
Fig. 5. Leaching kinetics for Sullivan concentrate with no surfactant: [Zn ] o = 0 g/l, [ H2SO 4 ] o -- 35 g/l, 141 g/1 concentrate, 750 kPa 02.
the same conditions as those shown in Fig. 3; again there was no precipitation of iron. Increasing the pulp density from 141 to 193 g/1 did not give significantly faster neutralization at 110 ° or 120 ° C. A reduction in the oxygen partial pressure to 345 kPa, and an increase in flow rate to 35-45 cm3/min also had little effect on the leaching and neutralization rates.
I
255
The molar ratio of zinc to iron in the concentrate was 3.15. The ratios of zinc to iron entering the solution during leaching (before any precipitation of iron) are given in Table III. It is clear that there was preferential leaching of sphalerite at l l 0 ° C , indicating a pronounced galvanic interaction. This was expected; at ambient temperatures, marcasite has a rest potential of 0.57 V, only slightly lower than that of pyrite (Sato, 1960). At 150 °C the zinc to iron ratios entering the solution were much closer to the ratio in the concentrate, indicating that galvanic interactions were non-existent or much weaker. This is entirely consistent with Peters' ( 1984 ) postulation that the potential differences leading to the galvanic effect are due to the difference in permeability of the sulphur layer forming on pyrite and sphalerite; these potential differences would disappear above the melting point of sulphur. Fig. 6 shows the effect of temperature on the composition of the jarosite precipitated during leaching tests with 120 g/1 Zn 2+, and no surfactant. It is evident that the lead and silver concentrations decreased markedly with increasing temperature, while the iron concentration increased, due to higher proportions of hydronium jarosite. These analyses were consistent with the Xray powder diffraction patterns. These indicated that the precipitates formed at 110 ° and 120 °C were solid solutions of hydronium, lead and silver jarosites, based on the diffraction patterns of Dutrizac and Kaiman (1976), whereas the material precipitated at 135 ° and 150°C only showed diffraction lines corresponding to solid solutions of hydronium and lead jarosites. The low levels of copper in the residues reflect the low copper level in the Sullivan concentrate, and do not necessarily reflect the maximum substitution that could occur. The increasing importance of hydronium jarosite at the higher temperatures is consistent with the increased stability of this phase in relation to other jarosites at higher temperatures, but the high concentration of lead in the jarosites formed at the lowest temperatures had not been expected, given the low solubility of lead sulphate, and the fact that lead jarosite prepared by Dutrizac and Kaiman (1976) at 95 ° C contained only 8.35 wt% Pb. Lead sulphate (anglesite ) is unlikely to account for these high lead concentrations, since the precipitate had been washed with ammonium acetate before analysis, and no lead TABLE III Ratio of zinc and iron entering solution during leaching
Surfactant present No surfactant [Zn]o=0
[Zn] (M)/[Fe] (M) .1 ll0°C
[Zn] (M)/[Fe] (M) 150°C
8.86 8.69 5.13
3.03 5.00 2.89
,1 [Zn] (M)/[Fe] (M) in concentrate is 3.15.
25(~ 4o!
5. (>
i
~ ~01 z 4.0
30
i
3.0 .~ a,e
• Fe
4~
• Pb
~" 20 0_
(.9
6 h
2.0
"I------"i~ 10
c~
1.0
110
120
l ~ , 130 140
~150
Temperature, °C
Fig. 6. Effect of temperature on compositionof jarosite precipitated after 90 min of leaching: no surfactant, [Zn](,= 120 g/l, [H~SO4]o=35g/l, 141 g/1 concentrate, 750 kPa 0.2. sulphate lines were evident in the diffraction pattern. Although Dutrizac and Kaiman noted lower silver concentrations in jarosites prepared below 140 °C than in material precipitated at higher temperatures, they were investigating pure silver-hydronium solid solutions with much higher silver concentrations than those developed in these leach solutions. The composition of iron precipitates shown in Fig. 6 reflects the effect of both temperature and changing leach solution chemistry. The composition of jarosites precipitated at different temperatures, from solutions of constant composition, are shown in Fig. 7. Although the lead and silver content of these precipitates also decreased with increasing hydrolysis temperature, the effect was not as marked as in Fig. 6. The concentrations of zinc and copper substituting for Fe (III) also decreased with increasing temperature. The final concentration of H2SO4 was measured after hydrolysis for 90 min at different temperatures. This concentration, which increases with increasing temperature, reflecting the increase in the equilibrium constants for hydrolysis, is plotted in Fig. 8, along with the final H2SO4 concentration from 90 min leaching-precipitation tests using an initial H2SO4 concentration of 35 g/l. The final concentration in the hydrolysis tests is indicative of the amount of acid released by iron hydrolysis and precipitation when leaching with no surfactant. Although the lowest H~SO4 concentration of 7.1 g/1 at 135 ° and 150 ° C is too high for complete iron precipitation, it does represent substantial neu-
257 40
5.0
[] Zn
• Fe
AAO
I~
~7 CU
•
30
~,. O + + O
4.0
S042-0H -
3.0
~=
if3
"T 2O 2.0
123 O3
x~ Q_ " U_
tO 1.0
I
I
I
I
I
li0
120
~3o
140
150
Temperature,
°C
Fig. 7. Effect of temperature on jarosite precipitated in hydrolysis tests, after 90 m L m [Zn 2÷ ]o = 147
g/l, [Pb2+]o=52 gll (as PbSO4), [Fe3+]o=18.9 g/l, [Cu2+]o=17.5 g/l, [Cd2+]o=0.5 g/l, [Ag + ] =0.4 g/l, [H2S04]o=0 g/1.
40
30
LJ
r- 2O 0 (J
J fu --r
to
I
I
I
I
t20
t40
t60
t80
Temperature,
°C
Fig. 8. Effect of temperature on final acid concentration:
Hydrolysis: [Zn 2+ ]o= 147 g/l, [Fe :~÷ ] 0 = 8 g/l, [H2S04]o=0 g/l; Leaching: [Zn 2+ ]o = 120 g/l, [Fe 3+ ] 0 = 8 g/l, [H2SO4]o=35 g/1.
tralization. Since n e u t r a l i z a t i o n results f r o m t h e t b r m a t i o n of e l e m e n t a l sulphur, which as seen in Fig. ,,,o is n o t t h e r m o d y n a m i c a l l y f a v o u r a b l e at such a high p H , the o v e r p o t e n t i a l for s u l p h a t e f o r m a t i o n m u s t clearly be high, even at 1 5 0 C . CONCLUSIONS S u b s t a n t i a l n e u t r a l i z a t i o n of acid, a n d p r e c i p i t a t i o n of iron c a n be a c h i e v e d in oxidative p r e s s u r e leaching of a zinc sulphide c o n c e n t r a t e . It t h u s a p p e a r s t h a t an a d d i t i o n a l leach could be used to a p p r e c i a b l y lower the n e e d for external n e u t r a l i z i n g a g e n t s in a f r e e - s t a n d i n g p r e s s u r e leach plant. U n d e r t h e conditions tested, t h e m i n i m u m final acid c o n c e n t r a t i o n was a c h i e v e d b y leaching at 150 ~C, with no s u r f a c t a n t . T h e toss of values to t h e j a r o s i t e residue f o r m e d d u r i n g leaching d e c r e a s e d w i t h i n c r e a s i n g t e m p e r a t u r e . T h e r e was no a d v a n tage in leaching at 180 ° C, however, b e c a u s e a g g l o m e r a t i o n of t h e u n l e a c h e d particles, a n d p r e c i p i t a t i o n of FeSO4, gave slower l e a c h i n g a n d n e u t r a l i z a t i o n kinetics. A galvanic i n t e r a c t i o n b e t w e e n s p h a l e r i t e a n d FeS2 p r o m o t e d zinc leaching at 110 ° C, b u t was less p r o n o u n c e d , or a b s e n t , at higher t e m p e r a t u r e s . ACKNOWLEDGEMENT S u p p o r t of this w o r k by t h e N a t i o n a l Science F o u n d a t i o n , u n d e r G r a n t C P E 8404653, is gratefully a c k n o w l e d g e d . REFERENCES Argall, G.O., Jr., 1986. Perseverance and winning ways at McLaughlin gold. Eng. Min. J., Oct. 1986, pp. 26-32. Berezowsky, R.M.G.S. and Weir, D.R., 1983. Pressure oxidation for treating refractory uranium and gold ores. 22nd Annu. Conf. Metallurgists of CIM, Edmonton, August, 1983. Biegler, T. and Swift, D.A., 1979. Electrochim. Acta, 24: 415. Bolton, G.L., Zubryckyj, N. and Veltman, H., 1979. Pressure leaching process for complex zinclead concentrates. Proc. 13th Int. Mineral Processing Congress, Warsaw, Vol. 1, pp. 581-607. Bruhn, G., Gerlach, J. and Pawlek, F., 1965. Untersuchungen iiber die LSslichkeiten von Salzen und Gasen in Wasser und wi~ssrigen LSsungen bei Temperaturen oberhalb 100 ° C. Z. Anorgan. Allgem. Chem., 337: 68-79. Chmielewski, T. and Charewicz, W.A., 1984. The oxidation of Fe (II) in aqueous sulphuric acid under oxygen pressure. Hydrometallurgy, 12: 21-30. Cobble, J.W., 1964. The thermodynamic properties of high temperature aqueous solutions, VI. J. Am. Chem. Soc., 86: 5394-5401. Connick, R.E. and Powell, R.E., 1953. The entropy of aqueous oxy-anions. J. Chem. Phys., 21: 2206-2207. Corriou, J.P. and Kikindai, T., 1981. J. Inorg. Nucl. Chem., 43: 9-15. Criss, C.M. and Cobble, J.W., 1964a. The thermodynamic properties of high temperature aqueous solutions, V. J. Am. Chem. Soc., 86: 5390-5393. Criss, C.M. and Cobble, J.W., 1964b. The thermodynamic properties of high temperature aqueous solutions, IV. J. Am. Chem. Soc., 86: 5385-5390. Crundwell, F.K., 1987a. The Effect of Flocculant Additions to the Leaching Solution on the Kinetics of the Oxidative Dissolution of Sphalerite. Mintek Report M298, Council for Mineral Technology, Randburg, South Africa.
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