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O2 as an oxidant in hydrometallurgy
Minerals Engineering, Vol. 13, No. 13, pp. 1319-1328, 2000
Pergamon 0892,--6875(00)00115-1
© 2000 Elsevier Science Ltd All fights reserved 0892-6875/00/$ - see front matter
SO2/O2 AS AN OXIDANT IN HYDROMETALLURGY* W. ZHANG II, P. SINGH I and D.M. MUIR ~ Dept. of Chemistry and Mineral Science, Murdoch University, WA 6150, Australia Email: wzhang @central.murdoch.edu.au; § CSIRO Minerals, PO Box 90, Bentley, WA 6982, Australia (Received 4 May 2000 ; accepted 20 July 2000)
ABSTRACT
Oxygen mixed with S02 can function as a powerful oxidant in acid solution. This paper summarises basic studies on the mechanism of $02/02 oxidation, the kinetics of oxidation in acidic media, and potential applications in hydrometallurgy. The SOz/02 system readily oxidises Fe(lI) to Fe(IlI) and provides a simple method of regenerating Fe(III) for leaching minerals like Cu2S or U3Os. It also oxidises As(IIl), which in the presence of iron(Ill) enables a stable ferric arsenate compound to be precipitated.. A method for the leaching and removal of arsenic from smelter fumes is proposed. Oxidation of Mn(II) to Mn(III)/Mn(1V) oxides by S02 /02 allows manganese impurity to be removed from leach liquors and electrolytes. © 2000 Elsevier Science Ltd. All rights reserved.
Keywords Mineral processing; oxidation; leaching; reaction kinetics; hydrometallurgy
INTRODUCTION It is well known that pure sulphur dioxide is a reducing agent used in a wide range of industries [Zhdanov (1973), Well and Sandier (1997)]. In hydrometaUurgy, pure sulphur dioxide solutions are used for leaching of manganese ores [Pahlman and Khalafalla (1979), Abbruzzese (1987), Kanungo and Das (1988), Tindall (1998),] and certain metal sulphides under reducing conditions [Tiwari (1976), Thom and Waters (1978), Sohn and Wadsworth (1980), Brual et al. (1983),]. It also oxidises at an anode and acts as a depolariser in the electrowinning of copper from leach solutions containing iron [Pace and Stauter (1974), Subbaiah et al. (2000)]. However when combined with oxygen over a particular range of compositions, and in the presence of certain transition metals, sulphur dioxide acts as a stronger oxidising agent than oxygen [Tiwari et al. (1978), Devuyst et al. (1982), Sato et al. (1984), Krause (1988), Nami (1988), Kwateng (1990)]. Some reported applications of the SO2/O2 mixture in hydrometallurgy are summarised in Table 1. Whilst INCO [Devuyst et al. (1982), Ettel et al. (1979)] has successfully exploited the use of SO2/O2 as an oxidant for cyanide in alkaline media, there has been relatively little development in the use of SO2/O2 as an oxidant for metal ions in acidic media. Initial studies on the oxidation of Fe(II) were conducted around 1915 [Du Faur (1915)] and later between 1919 and 1927 [Ralston (1927)]. Although a number of other studies on the oxidation of Fe(II) with SO2/O2 have been carried out in recent years, there is inconsistency in the kinetic observations and the proposed mechanisms. This study therefore focuses on gaining a fundamental, understanding of the iron and manganese catalysed SO2/O2 oxidising systems in acidic media
* Presented at Hydromet 2000, Adelaide, Australia, April 2000
1319
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W. Zhang et al.
and examines three selected processes in hydrometallurgy where this system may be applied more efficiently than oxygen. These processes involve the leaching of chalcocite, the oxidation and precipitation of arsenic from arsenical fume leachates, and the oxidation and precipitation of manganese impurity from nickel/cobalt solutions. These represent widely different yet potentially significant applications.
TABLE 1 Applications of SO2/O2 systems in hydrometallurgy Systems
Working pH range
Oxidation Products
0-3
Fe 2+~ Fe 3+ S(IV) ~ SO42-/$2062-
Potential Applications [refs] Leaching of metal sulphides
Fe-SO2/O2
[*]
Oxidation of As(III)
[**]
Mn 2+--> Mn(III)/Mn(IV)
Mn-SO2/O2
3-6
Ni-SO2/O2
7-8.5
Oxidative Precipitation of Mn [this study] Production of Ni(OH)3 [Ettel et al. (1979)]
S(IV) ~ 5042-/52062-
Ni(II) ~ Ni(III) S(IV) ~ SO42CU+ ~ Cu 2+
Oxidation of CNS(IV) --~ SO42[Devuyst et al. (1982)] CN- ~ CNO*: Tiwari et al. (1978), Devuyst et al. (1982), Sato et al. (1984), Krause (1988), Nami (1988), Kwateng (1990), this study. **: Nishimura et al. (1996), Wang et al. (1998), Khoe et al. (1997), this study. Cu-SOz/O2
9-10
KINETICS AND MECHANISM OF Fe(II) OXIDATION WITH SO2/O2 The kinetics of the S 0 2 / 0 2 oxidising system were studied under a range of conditions: 0-0.5 M Fe(II), 033% SO2/O2, 25-80°C, pH 0-3. Figure 1 shows typical plots for Fe(II) oxidation versus SO2/O2 ratios.
0.1
4%so , I 2
0.08 ~-~ 0.06 Lr~
'i
.~ 0.04 ©
...........
0.02
'ii°i
Pure 0 2 0
30
60
90
120
Tmle, rrm Fig. 1 Typical plots showing rate of oxidation of Fe(II) with varying % SO2. (80°C and pH 0.5). As can be seen in Figure 1, the oxidation of Fe(II) is extremely slow with O2 alone (0% SO2). With SO2/O2 mixtures, the oxidation of Fe(II) is dramatically promoted, and the oxidation rate increases with SO2 up to an optimum SO2/O2 ratio above which the oxidation rate decreases due to significant reduction of Fe(III) by excess SO2. The optimum SO2/O2 ratio decreases with increasing temperature due to the significant decrease in 02 solubility. It also varies with pH, due to variation in the fraction of reactive sulphite and bisulphite species, and with the rate of 02 mass transfer. In this work it is found that both Fe(III) and S(IV)
SOyO2 as an oxidant in hydrometallurgy
1321
species are involved in the initial step of oxidation and that the rate is first order with respect to [Fe(III)] up to 0.02 M. A comparison of the kinetic features observed in the oxidation of Fe(II) with SO2/O2 and with 02 are summarised in Table 2.
TABLE 2 Kinetic observations in Fe(H) oxidation by SO2/Oz and comparison with 02 Kinetic Observations Fe-SO2/O2 Fe-O2 [refs#] r = kl [Fe(III)] [S(IV)]/f([H+]) Empirical rate r = k2 [Fe2+]2[Oz] up to 0.02M [Fe(III)] and optimum SO2/O2 ratio, expression f([I-I+]) is function of [I-I+] Decrease with increasing temperature, also depends Optimum SO2/O2 ratio on pH and 02 mass transfer 22 at 40-80 ° 51-74 at 25-100°C EAet (kJ/mol) Reaction Product orS(IV) < Optimum 502/02 ] SO4z- only, and tFe(IIX)]:tSO?-a = 2 I > Optimum SCh/O2 [ [$2062-]/[504 z-] ratio increase with SO2/O2 Refs~: George (1954), Huffman and Davidson (1956), Cornelius and Woodcock (1958), McKay and Halpern (1958), Keenan (1969), Mathews and Robins (1972), Chmielewski and Charewicz (1984), Vraca and Cerovic (1997).
Factor
I
As can be seen in Table 2, the kinetics of Fe(II) oxidation with SO2/O2 is totally different from that with 02, indicating an entirely different mechanism. It is observed that the rate of oxidation of Fe(II) with SO2/O2 is markedly inhibited by addition of small amounts of hydroquinone which has been reported to be an effective free radical schevenger [Brandt et al. (1994), Huie and Neta (1985)]. Furthermore, the maximum potential measured for the Fe--SO2/O2 system is about 1.4 volts at pH 2 which is well above the standard potential of O2/H20 couple, suggesting the formation of peroxy intermediate species. A radical chain lreaction mechanism is proposed in this study with the following steps [Zhang et al. (2000(a))]. SO2.H20 ,~ HSO3- + H + HSO3- ~' 5032- + n + Fe 3+ + 5032- ,0~FeSO3 + FeSO3 ÷ --~ Fe2+ + SO3"SO3"- + 02 -~ SOs"Fe 2+ + SO5"- + H + ~ Fe 3+ + HSO52Fe2÷+ HSOs- + H+ ~ 2Fe3*+ SO42-+ H20 2HSO5- -~ SO42- + 02 + 2H+
Slow Fast Fast Fast
(1) (2) (3) (4) (5) (6) (7) (8)
It involves the slow initial formation of a ferric sulphite complex and decomposition to produce sulphite radical SO3"-. This is followed by a fast reaction with 02 to form a peroxo-monosulphate species SOs'-, and subsequently HSOs- which are responsible for the oxidation of Fe(II) and sulphite species. The derived rate expression essentially agrees with the observed rate law and the kinetic observations can be rationalised based on the proposed mechanism.
EFFECT OF COPPER ON OXIDATION OF Fe(H) The effect of copper on the oxidation of Fe(II) by 502/02 has been investigated in the range 0.1--0.5 M Fe(II)/Cu(II). Copper catalyses the oxidation by 02 alone and the reaction exhibits a second order rate with respect to Fe(II) concentration. A mechanism that involves the initial formation of complex CuO2+ and
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W. Zhanget al.
generates H202 as the oxidant for Fe(II) has been proposed, as discussed by Zhang et al. (2000(b)). In contrast, the presence of > O.1M Cu(II) inhibits the oxidation by SO2/O2, presumably due to its effect on the decomposition of the peroxy intermediate species. Figure 2 compares the effect of Cu(II) on the oxidation of Fe(II) with SOjO2 and with 02 alone in the presence of 0.5 M CuSO4 at 80°C. The rate of oxidation by the Fe/Cu/SO2/O2 system is much higher than the corresponding Fe/Cu/O2 system. The presence of Cu(II) catalyses the Fe/O2 system quite significantly, however, its presence in large amount has the opposite effect on the Fe/SO2/O2 system. The effect of Cu(II) on the Fe/Cu/SO2/O2 system is characterised by two simultaneously occurring mechanisms. As can be seen in Figure 2, the oxidation of Fe(II) in the Fe/Cu/SO2/O2 system is initially dominated by the Cu/O2 mechanism, indicated by parabolic kinetics, but is overtaken by the Fe/SO2/O2 mechanism as the concentration of Fe(II) falls, indicated by linear oxidation kinetics in the later stages. Thus it may be concluded that the Cu/SO2/O2 system oxidises Fe(II) at a faster rate than the Cu/O2 system, particularly at low concentrations of Fe(II), which should maintain higher Fe(III) levels during continuous oxidative leaching of copper bearing minerals or concentrates.
0.I j" SO2/O2,~k/ / ~ l n i t i a l l y
0.1M Fe(ll)
0.08
0.04
~ ~ 0.06
/ /
Cu-SO2/O2
0.02
o2 Only
0
I
I
I
20
40
60
80
100
Time, min Fig.2
Comparison of oxygen systems on rate of Fe 2÷ oxidation. (initial 0.1 M Fe 2÷ , 0.5M CuSO4; 2% SO2/O2; pH 1.5; 80°C). LEACHING OF COPPER SULPHIDES WITH Fe/SO2/O2
The ferric leaching of naturally occurring chalcocite and chalcocite concentrate was compared at low Fe(II) concentrations (<0.1 M), with Fe(III) regenerated by SO2/O2 and 02. As can be seen in Figure 3, copper extraction after 7 hours with SO2/O2 is significantly higher than that with 02 alone. The results shown in Figure 4 indicate that under similar conditions for the Fe/SO2/O2 system, a higher concentration of Fe(III) is maintained in the solutions as compared to the Fe/O2 system. This confirms that the regeneration of Fe(III) by SO2/O2 during the leaching process is much faster than by 02. It has been reported earlier that at higher [Fe(II)], the rate of Fe(III) regeneration by both Cu/O2 and Fe/SO2/O2 is similar (Zhang e t al. (2000 b)). Furthermore the rate of second stage leaching of predominantly covellite (CuS) is slower and less dependent upon Fe(III) concentration. However, it is also observed that the presence of 0.5M CI- (NaC1) enhanced the recovery of copper from leaching of Cu2S by about 5% using either 502/02 or 02. This agrees with the observation by Cheng and Lawson (1991) with 02 as the oxidant.
SOYO2 as an oxidant in hydrometaUurgy
1323
100 80 d~
02 Only
60 40 7 Hours Leach, 80"C
20
0
I
I
I
I
0.02
0.04
0.0,6
0.08
0.1
Total IronConcentration,M Fig.3
Effect of Fe(II) concentration on Cu extraction after 7 h leaching with SO2]O 2 and with 0 2 alone. (4% Pulp density; 80°C; pH 1).
100
4 ° ° I, iFirst stage leach
/
o
80
::'7'
60
20
~
-
_
|
,
0
m
0
~
~'~P
60
I
I
I
I
120 180 240 300
I
360 420
Tine, rain Fig .4
Comparison of %Fe(III) present during leaching of Cu2S and continuous sparging by either SO2/O2 or 02. ( initial 0.1 M Fe(II)).
It is found that about 2 moles of H2SO4 is consumed using 02 alone when about 0.9 mole of Cu2S is leached in the above experiments. This measured ratio is very close to the stoichiometric requirement according to equettions (9) and (10) Cu2S + 4Fe 3+---}2Cu2++ 4Fe2+ + S° 4Fe 2+ + 02 + 2H2SO4 --~ 4Fe 3++ 2H20
(9) (10)
However, little H2SO 4 w a s required for the similar amount of Cu2S leached with SO2/O2, agreeing with the following reactions for generation of Fe(III) and H2SO4: 2Fe2+ + 502/02 --~ 2Fe3+ + 5042+ 502/O2 -F n20 "+ H2504
(11) (12)
1324
W. Zhanget al.
Thus the main benefit with SO2/O2 appears to be the simultaneous production of H2SO4 needed for leaching and oxidation, and leaching with lower iron in the circuit.
As(III) OXIDATION WITH SO2/O2 AND APPLICATION TO THE PROCESSING OF ARSENICAL SMELTER FUME The oxidation of As(III) is found to occur at the same time as the oxidation of Fe(II) using the SO2[O 2 system and has a similar activation energy [Zhang e t al. (2000(c))]. Khoe e t al. (1997) also found that UV light catalysed the oxidation of Fe(II) and As(III) using 02 alone. Compared with the SO2/O2 system, the rate of oxidation of As(III) by UV/O2 in the presence of trace Fe(III) is found to be similar at low initial concentrations of As(III) around 0.1 mM, but much slower at initial As(III) concentrations above 6.5 mM as shown in Figure 5.
100 with F
80
t,¢)
e
/
S
~
60
<
~D .o~
40
© 20 ,. 0
20
40
60
Tune, min Fig.5
Comparison of oxygen systems for oxidation of As(III). (6.5 mM As(III), 10 mM Fe(III); 25°C; pH=l.5).
Table 3 compares the kinetic features of As(III) oxidation by Fe/SO2/O2 and UV/Fe/O2.Fundamentally the Fe/SO2/O2 system is more suitable for oxidising more concentrated As(III) solutions found in hydrometallurgical processes. However, in the presence of low concentrations of Fe(III), the combination of both SO2/O2 and UV light, offers faster kinetics for the oxidation of dilute As(III) solutions than either of the individual systems. TABLE 3 Comparison of kinetic features of As(HI) oxidation Oxidising system [As(III)] (initial)
Fe/SO2/O2
UV/Fe/02
Rate increases over range: 0-100 mM
rate increases over range: 0-2.5 mM
[Fe(III)]
Rate increases over range: 0-10 mM
Temp.
Rate increases over range: 25-50°C
pH
Rate increases over range: pH 0-3
Rate law d[As(III)]oc [Fe(III)][As(llI)]°.5 (empirical) dt * From ANSTO Patent [Khoe e t al. (1997)]
rate increases over range: 0-1.5 mM excess [Fe(III)] reduces the rate* rate decreases with increasing temperature above 25 *C* rate increases with decrease in pH from 3 to 2* _ d[As(l]/)] oc[Fe(m)]°.5[As(lII)] °.z5 dt
so7]02 as an oxidantin hydrometallurgy
1325
In order to apply this observation, a method for removing arsenic from a typical smelter fume containing iron and valuable niickel has also been investigated in this study [Zhang et al. (1998)]. Smelter fume often contains arsenic as As203 as well as significant amounts of base metals and iron. Hence the general method consists of dilute ac,id leaching, oxidation of As(III) and Fe(II) using SO2/O2, and precipitation of iron(HI) arsenate as shown in Figure 6.
Fume STEP 1
Dilute
I
|
Recycle 4 .............................. .-.-..
LEACH
H2SO4
.... '
Species
i
llo o o m a l l l a l i i o e q l m g w i w m lie
Residue ~ ~FLiquor
Smelter
STEP 2
SO2/O2
i
_.[
t
~,T.2+
i
~
r~l+,t~o+
[..... : I
iii iiiiii
d
-
mloom m m mQ
2+
:
• -
Fe /Fe
i...A,s!.V).!..A:s..(!I!!. j
Oxidation
[
.....
.... i;;';:":;
"Fe . . . . . .3~" ..
COI ~ As(III)-~As(V) .PRECIPITATION u. ~.........................
i ..
~l~lemllllmm•wmmmmemoelIQliO
iii
Tailings "*"l F
i
T
~ ) e l e e Q m l l o l l m m w l l l ~
...I):
Precipitate~ •
Liquor ........
[ . . m e.
:
: FeAsO4 •
•
. m ,. m t .
Bleed and Rer~ove Nn, Co
".................. " •
io
.l•
.........
• n
"
Fig.6 Proposed flowsheet for processing arsenical smelter fume. For this particular nickel smelter fume with a high Fe:As ratio, dilute H2SO4 extracted 25-30% of the As(III) and sufficient Fe(II) to allow oxidation and precipitation of FeAsO4 to occur at pH 3.3 with minimal neutralisation cost. By using SO2/O2 as the oxidant, the reaction could be carded out rapidly at 30-38°C unlike oxidation with 02. This approach appears attractive to treat smelter fumes containing significant levels of arsenic and excess iron phases that can be leached with SO2 or dilute sulphuric acid.
1326
W. Zhang etaL
OXIDATIVE PRECIPITATION OF MANGANESE (H) WITH SO2/O2 The kinetics of oxidation of Mn(II) with SO2/O2 and subsequent precipitation of MnO2 have been investigated in the pH range 1-6 and temperature range 25-80°C. It is found that the mechanism of oxidation of Mn(II) is similar to the Fe(II) oxidation, with Mn(III) acting as the catalyst which disproportionates almost spontaneously to MnO2/Mn203 [Zhang (2000)]. The oxidation of Mn(II) is slow at pH < 3, but at pH > 4 Mn(II) is readily oxidised. When applied to nickel laterite leach solutions containing Co and Ni, the oxidative precipitation of Co as Co(OH)3 became significant at pH > 4 and the precipitation of Ni as Ni(OH)3 was significant at pH > 6.5. Therefore, the selective oxidative precipitation of Mn occurs at pH 3-4, and provides a useful means for the removal of manganese impurities from nickel/cobalt leach liquors. Figure 7 shows the precipitation of Mn from a synthetic Co sulphate leach liquor containing 0.1M Co(II) and 0.01 M Mn(II)/Ni(II)/Fe(II)/Cu(II)/Zn(II). As can be seen, 10 mM manganese was selectively precipitated from nickel and cobalt within ten minutes in the pH range 3--4 with less than 1% coprecipitation of Ni and Co. In the presence of 0.01M Fe(II), iron was oxidatively precipitated as Fe(OH)3 before manganese.
100
~" Mn
/I
*
1
80
/Cu
..~ ~o 60 40
I0 minutes for each run at 5 0 " C
/
/
20 Ni, Co m
3
"1=
I
I
I
I
3.5
4
4.5
5
5.5
6
pH Fig.7
Effect of pH on the precipitation of Mn, Cu, Zn and Ni from a Co leach liquor. (10 minute oxidation with 2% SO2/O2; 50°C).
SUMMARY
The results of this study demonstrate that SO2/O2 can function as a strong oxidant in acidic solution for the oxidation of Fe(II), Mn(II) and As(III). It can also function as a means of generating H2SO4 from the autoxidation of SO2 once Fe(II) is oxidised. The SO2/O2 system, catalysed by iron(III) or manganese(III) should be useful for a number of other hydrometallurgical applications. For example, for generating Fe(III) and leaching copper sulphide minerals or uranium(IV) minerals requiring both oxidant and H2SO4. It is also useful for the oxidative precipitation of As(III) and Fe(II) from leach solutions as stable ferric arsenate, and for the oxidative precipitation of Mn(II) as MnO2 from Co(II) and Ni(II) leach liquors at around pH 3.
REFERENCES
Abbruzzese C., Aqueous SO2 processing of manganese ores, in Separation Processes in Hydrometallurgy, Society of Chemical Industry, London, 1987, pp. 76-87.
soz/oz as an oxidantin hydrometallurgy
1327
Brandt C., F~ibi~inL, and van Eldik R., Kinetics and mechanism of the iron(III)-catalysed autoxidation of sulphur(IV) oxides in aqueous solution. Evidence for the redox cycling of iron in the presence of oxygen and modelling of the overall reaction mechanism, Inorg. Chem,, 1994, 33,687-701. Brual G. B., Byerley J. J. and Rempel G. L., Kinetic and mechanistic study of FeS dissolution in aqueous sulfur dioxide solution, Hydrometallurgy, 1983, 9, 307-331. Cheng C.Y. and Lawson F. The kinetics of leaching covellite in acidic oxygenated sulphate-chloride solutions, HydrometaUurgy, 1991, 27,269-684. Chmielewski T. and Charewicz W., The oxidation of Fe(II) in aqueous sulphuric acid under oxygen pressure, Hydrometallurgy, 1984, 12, 21-30. Cornelius R. J. and Woodcock J. T., Pressure leaching of a manganese ore. Part I: Kinetic aspects, Proc. Australas. Inst. Min. Met., 1958, 185, 65-78. Devuyst E. A., Mosoiu A. and Krause E., Oxidising properties and application of the SO2-O2 system, in HydrometaUurgy Research, Development and Plant Practice, (Eds. Osseo-Asare, K. and Miller J. D), Proceedings of the 3rd International Symposium on Hydrometallurgy, The Metallurgical Society of AIME, 1982, pp. 391--403. Du Faur J. B., Australian Patent 15862, 1915. Ettel V. A., Mosoiu M. A. and Devuyst E. A., A novel oxidant for nickel hydrometallurgy, Hydrometallur~:y, 1979, 4, 247-258. George P., The oxidation of ferrous perchlorate by molecular oxygen, J. Chem. Soc., 1954, 4, 4349--4359. Huffman R. E. and Davidson N. J., Kinetics of the ferrous iron-oxygen reaction in sulphuric acid solution, J. Amer. Chem. Soc., 1956, 78, p. 4836. Huie E. and Neta P., One-electron redox reactions in aqueous solution of sulphite with hydroquinone and other hydroxyphenols, J. Phys. Chem., 1985, 89, 3918-3921. Kanungo S.B. and Das R. P., Extraction of metals from manganese nodules of the Indian Ocean by leaching in aqueous solution of sulphur dioxide, Hydrometallurgy, 1988, 20, 135-146. Keenan E. A., A bacterial beneficiation of uranium materials, PhD thesis, University of New South Wales, Australia, 1969. Khoe G. H., Zaw M., Prasad P. S., and Emett M. T., "Photo-assisted oxidation of inorganic species in aqueous solutions", Intl Application Number: WO 99/05065, February 1999. Krause E., The oxidation of ferrous sulphate solutions by sulphur dioxide and oxygen, PhD thesis, University of Waterloo, Canada, 1988. Kwateng O., A kinetic study of the dissolution of nickel sulfide in acidified ferrous sulfate solution with a gas mixture of oxygen and sulfur dioxide, Ph.D Thesis, Columbia University, New York, 1990. Mathews C.T. and Robins R. G., Oxidation of aqueous ferrous sulphate solution by molecular oxygen, Proc. Australas. Inst. Min. Metall., 1972, No. 242, 47-56. McKayD. R. and Halpern, J., Kinetic study of the oxidation of pyrite in aqueous suspension, Trans. Metall. Soc. AIME, 1958, 212, 301-309. Nami F., The kinetics of zinc sulfide leaching by oxygen, sulfur dioxide and ferrous sulfate, PhD Thesis, Columbia University, New York, 1988. Nishimura T., Wang Q. and Umetsu Y., Removal of arsenic from process liquors by oxidation of iron(II), arsenic(III) and sulphur(IV) with oxygen, in Iron Control and Disposal, (Eds: Dutrizac J. E. and Harris G. B.), Proceedings of the Second International Symposium, CIM, Montreal, 1996, pp. 535-548. Pace G. F. and Stauter J. C., Direct electrowinning of copper from synthetic pregnant leach solutions utilising SO2 and graphite anodes-Pilot-plant results, CIM Trans., CIM Bull., 1974, LXXVII, 51-57. Pahlman J. E. and Khalafalla S. E., Selective recovery of nickel, cobalt, manganese from sea nodules with sulfurous acid, US Patent No. 4,138,456,1979. Ralston O. C., Theeferric sulphate-sulphuric acid process, U.S.B.M. Bulletin, 1927, 260, 1-61. Sato T., Goto T., Okabe T. and Lawson F., The oxidation of iron(II) sulphate with sulphur dioxide and oxygen mixtures, Bull. Chem. Soc. Japan, 1984, 57 (8), 2028-2086. Sohn H-J and Wadsworth M. E., Reduction of chalcopyrite with SO2 in the presence of cupric ions, J. of Metals, Noveraber, 1980, pp. 18-22. Subbaiah T., Singh P., Hefter G., Muir D. and Das R.P., Electrowinning of copper in the presence of anodic depolarisers - Part I - A review, Mineral Processing and Extractive Metallurgy Review, 2000, (in press). Thom G. C. and Waters P. F., A study of the metal sulfide-sulfur dioxide reaction in aqueous media by reaction pressure characterisation and ultraviolet spectrophotometry, HydrometaUurgy, 1978,3, 373396.
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Tindall G. P., High temperature of acid leaching of Western Australia laterites, PhD Thesis, Murdoch University, Perth, 1998. Tiwari B. L., The kinetics of oxidation of zinc sulfide and hydrogen sulfide by sulfur dioxide in aqueous sulphuric acid, PhD Thesis, Columbia University, New York, 1976. Tiwari, B. L., Kolbe J., and Hayden H. W., Oxidation of ferrous sulfate in acid solution by mixtures of sulfur dioxide and oxygen, MetalL Trans. B, 1978, 10B, 607 - 612. Vracar R. ~. and Cerovic K, P., Kinetics of oxidation of Fe(II) ions by gaseous oxygen at high temperatures in an autoclave, Hydrometallurgy, 1997, 44, 113-124. Wang Q., Demopoulos G. P., and Harris G. B., A novel process for arsenic oxidation and fixation, in Waste Processing and Recycling in Mineral and Metallurgical Industries III, (Eds. Rao S. R., Amaratunga L. M., Richards G. G. and Kondos P. D), Proceedings of the International. Symposium, Metallurgical Society of CIM, Montreal, 1998, pp. 375-387. Well E. D. and Sandier, S. R., Sulfur dioxide, In Kirk-Othmer Encyclopedia of Chemical Technology, 4th edn., John Wiley & Sons, 1997, 23,299-312. Zhang W., Muir D and Singh P., Leaching and recovery of arsenic from smelter fume, in Waste Processing and Recycling in Mineral and Metallurgical Industries III, (Eds. Rao S. R., Amaratunga L. M., Richards G. G. and Kondos P. D.), Proceedings of the International Symposium, Metallurgical Society of CIM, Montreal, 1998, pp 259-271. Zhang W., Singh P. and Muir D., Iron(II) oxidation by SO2/O2 in acidic media. Part I. Kinetics and mechanism, Hydrometallurgy, 2000(a), 55(3), 229-245. Zhang W., Muir D. and Singh P, Iron(II) oxidation by SO2/O2 in acidic media. Part II. Effect of copper.
Hydrometallurgy, 2000(b ), accepted for publication Zhang W., Singh P. and Muir D., Kinetics of oxidation of As(III) with SO2/O2 and UV light, in Minor Elements 2000: Processing and Environmental Aspects of As, Sb, Se, Te, and Bi, (Ed. Young C.), SME, Littleton CO.,2000(c), pp 333-344. Zhang W., SO2/O2 as an oxidant in hydrometallurgy, Ph.D Thesis, Murdoch University, Perth, Western Australia, 2000. Zhdanov S. I., Sulphur, Chapter IV-6, in Encyclopedia of Electrochemistry of the Elements, Marcel Dekker, 1973, p. 275.
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SO2/O2 AS AN OXIDANT IN HYDROMETALLURGY* W. ZHANG II, P. SINGH I and D.M. MUIR ~ Dept. of Chemistry and Mineral Science, Murdoch University, WA 6150, Australia Email: wzhang @central.murdoch.edu.au; § CSIRO Minerals, PO Box 90, Bentley, WA 6982, Australia (Received 4 May 2000 ; accepted 20 July 2000)
ABSTRACT
Oxygen mixed with S02 can function as a powerful oxidant in acid solution. This paper summarises basic studies on the mechanism of $02/02 oxidation, the kinetics of oxidation in acidic media, and potential applications in hydrometallurgy. The SOz/02 system readily oxidises Fe(lI) to Fe(IlI) and provides a simple method of regenerating Fe(III) for leaching minerals like Cu2S or U3Os. It also oxidises As(IIl), which in the presence of iron(Ill) enables a stable ferric arsenate compound to be precipitated.. A method for the leaching and removal of arsenic from smelter fumes is proposed. Oxidation of Mn(II) to Mn(III)/Mn(1V) oxides by S02 /02 allows manganese impurity to be removed from leach liquors and electrolytes. © 2000 Elsevier Science Ltd. All rights reserved.
Keywords Mineral processing; oxidation; leaching; reaction kinetics; hydrometallurgy
INTRODUCTION It is well known that pure sulphur dioxide is a reducing agent used in a wide range of industries [Zhdanov (1973), Well and Sandier (1997)]. In hydrometaUurgy, pure sulphur dioxide solutions are used for leaching of manganese ores [Pahlman and Khalafalla (1979), Abbruzzese (1987), Kanungo and Das (1988), Tindall (1998),] and certain metal sulphides under reducing conditions [Tiwari (1976), Thom and Waters (1978), Sohn and Wadsworth (1980), Brual et al. (1983),]. It also oxidises at an anode and acts as a depolariser in the electrowinning of copper from leach solutions containing iron [Pace and Stauter (1974), Subbaiah et al. (2000)]. However when combined with oxygen over a particular range of compositions, and in the presence of certain transition metals, sulphur dioxide acts as a stronger oxidising agent than oxygen [Tiwari et al. (1978), Devuyst et al. (1982), Sato et al. (1984), Krause (1988), Nami (1988), Kwateng (1990)]. Some reported applications of the SO2/O2 mixture in hydrometallurgy are summarised in Table 1. Whilst INCO [Devuyst et al. (1982), Ettel et al. (1979)] has successfully exploited the use of SO2/O2 as an oxidant for cyanide in alkaline media, there has been relatively little development in the use of SO2/O2 as an oxidant for metal ions in acidic media. Initial studies on the oxidation of Fe(II) were conducted around 1915 [Du Faur (1915)] and later between 1919 and 1927 [Ralston (1927)]. Although a number of other studies on the oxidation of Fe(II) with SO2/O2 have been carried out in recent years, there is inconsistency in the kinetic observations and the proposed mechanisms. This study therefore focuses on gaining a fundamental, understanding of the iron and manganese catalysed SO2/O2 oxidising systems in acidic media
* Presented at Hydromet 2000, Adelaide, Australia, April 2000
1319
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W. Zhang et al.
and examines three selected processes in hydrometallurgy where this system may be applied more efficiently than oxygen. These processes involve the leaching of chalcocite, the oxidation and precipitation of arsenic from arsenical fume leachates, and the oxidation and precipitation of manganese impurity from nickel/cobalt solutions. These represent widely different yet potentially significant applications.
TABLE 1 Applications of SO2/O2 systems in hydrometallurgy Systems
Working pH range
Oxidation Products
0-3
Fe 2+~ Fe 3+ S(IV) ~ SO42-/$2062-
Potential Applications [refs] Leaching of metal sulphides
Fe-SO2/O2
[*]
Oxidation of As(III)
[**]
Mn 2+--> Mn(III)/Mn(IV)
Mn-SO2/O2
3-6
Ni-SO2/O2
7-8.5
Oxidative Precipitation of Mn [this study] Production of Ni(OH)3 [Ettel et al. (1979)]
S(IV) ~ 5042-/52062-
Ni(II) ~ Ni(III) S(IV) ~ SO42CU+ ~ Cu 2+
Oxidation of CNS(IV) --~ SO42[Devuyst et al. (1982)] CN- ~ CNO*: Tiwari et al. (1978), Devuyst et al. (1982), Sato et al. (1984), Krause (1988), Nami (1988), Kwateng (1990), this study. **: Nishimura et al. (1996), Wang et al. (1998), Khoe et al. (1997), this study. Cu-SOz/O2
9-10
KINETICS AND MECHANISM OF Fe(II) OXIDATION WITH SO2/O2 The kinetics of the S 0 2 / 0 2 oxidising system were studied under a range of conditions: 0-0.5 M Fe(II), 033% SO2/O2, 25-80°C, pH 0-3. Figure 1 shows typical plots for Fe(II) oxidation versus SO2/O2 ratios.
0.1
4%so , I 2
0.08 ~-~ 0.06 Lr~
'i
.~ 0.04 ©
...........
0.02
'ii°i
Pure 0 2 0
30
60
90
120
Tmle, rrm Fig. 1 Typical plots showing rate of oxidation of Fe(II) with varying % SO2. (80°C and pH 0.5). As can be seen in Figure 1, the oxidation of Fe(II) is extremely slow with O2 alone (0% SO2). With SO2/O2 mixtures, the oxidation of Fe(II) is dramatically promoted, and the oxidation rate increases with SO2 up to an optimum SO2/O2 ratio above which the oxidation rate decreases due to significant reduction of Fe(III) by excess SO2. The optimum SO2/O2 ratio decreases with increasing temperature due to the significant decrease in 02 solubility. It also varies with pH, due to variation in the fraction of reactive sulphite and bisulphite species, and with the rate of 02 mass transfer. In this work it is found that both Fe(III) and S(IV)
SOyO2 as an oxidant in hydrometallurgy
1321
species are involved in the initial step of oxidation and that the rate is first order with respect to [Fe(III)] up to 0.02 M. A comparison of the kinetic features observed in the oxidation of Fe(II) with SO2/O2 and with 02 are summarised in Table 2.
TABLE 2 Kinetic observations in Fe(H) oxidation by SO2/Oz and comparison with 02 Kinetic Observations Fe-SO2/O2 Fe-O2 [refs#] r = kl [Fe(III)] [S(IV)]/f([H+]) Empirical rate r = k2 [Fe2+]2[Oz] up to 0.02M [Fe(III)] and optimum SO2/O2 ratio, expression f([I-I+]) is function of [I-I+] Decrease with increasing temperature, also depends Optimum SO2/O2 ratio on pH and 02 mass transfer 22 at 40-80 ° 51-74 at 25-100°C EAet (kJ/mol) Reaction Product orS(IV) < Optimum 502/02 ] SO4z- only, and tFe(IIX)]:tSO?-a = 2 I > Optimum SCh/O2 [ [$2062-]/[504 z-] ratio increase with SO2/O2 Refs~: George (1954), Huffman and Davidson (1956), Cornelius and Woodcock (1958), McKay and Halpern (1958), Keenan (1969), Mathews and Robins (1972), Chmielewski and Charewicz (1984), Vraca and Cerovic (1997).
Factor
I
As can be seen in Table 2, the kinetics of Fe(II) oxidation with SO2/O2 is totally different from that with 02, indicating an entirely different mechanism. It is observed that the rate of oxidation of Fe(II) with SO2/O2 is markedly inhibited by addition of small amounts of hydroquinone which has been reported to be an effective free radical schevenger [Brandt et al. (1994), Huie and Neta (1985)]. Furthermore, the maximum potential measured for the Fe--SO2/O2 system is about 1.4 volts at pH 2 which is well above the standard potential of O2/H20 couple, suggesting the formation of peroxy intermediate species. A radical chain lreaction mechanism is proposed in this study with the following steps [Zhang et al. (2000(a))]. SO2.H20 ,~ HSO3- + H + HSO3- ~' 5032- + n + Fe 3+ + 5032- ,0~FeSO3 + FeSO3 ÷ --~ Fe2+ + SO3"SO3"- + 02 -~ SOs"Fe 2+ + SO5"- + H + ~ Fe 3+ + HSO52Fe2÷+ HSOs- + H+ ~ 2Fe3*+ SO42-+ H20 2HSO5- -~ SO42- + 02 + 2H+
Slow Fast Fast Fast
(1) (2) (3) (4) (5) (6) (7) (8)
It involves the slow initial formation of a ferric sulphite complex and decomposition to produce sulphite radical SO3"-. This is followed by a fast reaction with 02 to form a peroxo-monosulphate species SOs'-, and subsequently HSOs- which are responsible for the oxidation of Fe(II) and sulphite species. The derived rate expression essentially agrees with the observed rate law and the kinetic observations can be rationalised based on the proposed mechanism.
EFFECT OF COPPER ON OXIDATION OF Fe(H) The effect of copper on the oxidation of Fe(II) by 502/02 has been investigated in the range 0.1--0.5 M Fe(II)/Cu(II). Copper catalyses the oxidation by 02 alone and the reaction exhibits a second order rate with respect to Fe(II) concentration. A mechanism that involves the initial formation of complex CuO2+ and
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W. Zhanget al.
generates H202 as the oxidant for Fe(II) has been proposed, as discussed by Zhang et al. (2000(b)). In contrast, the presence of > O.1M Cu(II) inhibits the oxidation by SO2/O2, presumably due to its effect on the decomposition of the peroxy intermediate species. Figure 2 compares the effect of Cu(II) on the oxidation of Fe(II) with SOjO2 and with 02 alone in the presence of 0.5 M CuSO4 at 80°C. The rate of oxidation by the Fe/Cu/SO2/O2 system is much higher than the corresponding Fe/Cu/O2 system. The presence of Cu(II) catalyses the Fe/O2 system quite significantly, however, its presence in large amount has the opposite effect on the Fe/SO2/O2 system. The effect of Cu(II) on the Fe/Cu/SO2/O2 system is characterised by two simultaneously occurring mechanisms. As can be seen in Figure 2, the oxidation of Fe(II) in the Fe/Cu/SO2/O2 system is initially dominated by the Cu/O2 mechanism, indicated by parabolic kinetics, but is overtaken by the Fe/SO2/O2 mechanism as the concentration of Fe(II) falls, indicated by linear oxidation kinetics in the later stages. Thus it may be concluded that the Cu/SO2/O2 system oxidises Fe(II) at a faster rate than the Cu/O2 system, particularly at low concentrations of Fe(II), which should maintain higher Fe(III) levels during continuous oxidative leaching of copper bearing minerals or concentrates.
0.I j" SO2/O2,~k/ / ~ l n i t i a l l y
0.1M Fe(ll)
0.08
0.04
~ ~ 0.06
/ /
Cu-SO2/O2
0.02
o2 Only
0
I
I
I
20
40
60
80
100
Time, min Fig.2
Comparison of oxygen systems on rate of Fe 2÷ oxidation. (initial 0.1 M Fe 2÷ , 0.5M CuSO4; 2% SO2/O2; pH 1.5; 80°C). LEACHING OF COPPER SULPHIDES WITH Fe/SO2/O2
The ferric leaching of naturally occurring chalcocite and chalcocite concentrate was compared at low Fe(II) concentrations (<0.1 M), with Fe(III) regenerated by SO2/O2 and 02. As can be seen in Figure 3, copper extraction after 7 hours with SO2/O2 is significantly higher than that with 02 alone. The results shown in Figure 4 indicate that under similar conditions for the Fe/SO2/O2 system, a higher concentration of Fe(III) is maintained in the solutions as compared to the Fe/O2 system. This confirms that the regeneration of Fe(III) by SO2/O2 during the leaching process is much faster than by 02. It has been reported earlier that at higher [Fe(II)], the rate of Fe(III) regeneration by both Cu/O2 and Fe/SO2/O2 is similar (Zhang e t al. (2000 b)). Furthermore the rate of second stage leaching of predominantly covellite (CuS) is slower and less dependent upon Fe(III) concentration. However, it is also observed that the presence of 0.5M CI- (NaC1) enhanced the recovery of copper from leaching of Cu2S by about 5% using either 502/02 or 02. This agrees with the observation by Cheng and Lawson (1991) with 02 as the oxidant.
SOYO2 as an oxidant in hydrometaUurgy
1323
100 80 d~
02 Only
60 40 7 Hours Leach, 80"C
20
0
I
I
I
I
0.02
0.04
0.0,6
0.08
0.1
Total IronConcentration,M Fig.3
Effect of Fe(II) concentration on Cu extraction after 7 h leaching with SO2]O 2 and with 0 2 alone. (4% Pulp density; 80°C; pH 1).
100
4 ° ° I, iFirst stage leach
/
o
80
::'7'
60
20
~
-
_
|
,
0
m
0
~
~'~P
60
I
I
I
I
120 180 240 300
I
360 420
Tine, rain Fig .4
Comparison of %Fe(III) present during leaching of Cu2S and continuous sparging by either SO2/O2 or 02. ( initial 0.1 M Fe(II)).
It is found that about 2 moles of H2SO4 is consumed using 02 alone when about 0.9 mole of Cu2S is leached in the above experiments. This measured ratio is very close to the stoichiometric requirement according to equettions (9) and (10) Cu2S + 4Fe 3+---}2Cu2++ 4Fe2+ + S° 4Fe 2+ + 02 + 2H2SO4 --~ 4Fe 3++ 2H20
(9) (10)
However, little H2SO 4 w a s required for the similar amount of Cu2S leached with SO2/O2, agreeing with the following reactions for generation of Fe(III) and H2SO4: 2Fe2+ + 502/02 --~ 2Fe3+ + 5042+ 502/O2 -F n20 "+ H2504
(11) (12)
1324
W. Zhanget al.
Thus the main benefit with SO2/O2 appears to be the simultaneous production of H2SO4 needed for leaching and oxidation, and leaching with lower iron in the circuit.
As(III) OXIDATION WITH SO2/O2 AND APPLICATION TO THE PROCESSING OF ARSENICAL SMELTER FUME The oxidation of As(III) is found to occur at the same time as the oxidation of Fe(II) using the SO2[O 2 system and has a similar activation energy [Zhang e t al. (2000(c))]. Khoe e t al. (1997) also found that UV light catalysed the oxidation of Fe(II) and As(III) using 02 alone. Compared with the SO2/O2 system, the rate of oxidation of As(III) by UV/O2 in the presence of trace Fe(III) is found to be similar at low initial concentrations of As(III) around 0.1 mM, but much slower at initial As(III) concentrations above 6.5 mM as shown in Figure 5.
100 with F
80
t,¢)
e
/
S
~
60
<
~D .o~
40
© 20 ,. 0
20
40
60
Tune, min Fig.5
Comparison of oxygen systems for oxidation of As(III). (6.5 mM As(III), 10 mM Fe(III); 25°C; pH=l.5).
Table 3 compares the kinetic features of As(III) oxidation by Fe/SO2/O2 and UV/Fe/O2.Fundamentally the Fe/SO2/O2 system is more suitable for oxidising more concentrated As(III) solutions found in hydrometallurgical processes. However, in the presence of low concentrations of Fe(III), the combination of both SO2/O2 and UV light, offers faster kinetics for the oxidation of dilute As(III) solutions than either of the individual systems. TABLE 3 Comparison of kinetic features of As(HI) oxidation Oxidising system [As(III)] (initial)
Fe/SO2/O2
UV/Fe/02
Rate increases over range: 0-100 mM
rate increases over range: 0-2.5 mM
[Fe(III)]
Rate increases over range: 0-10 mM
Temp.
Rate increases over range: 25-50°C
pH
Rate increases over range: pH 0-3
Rate law d[As(III)]oc [Fe(III)][As(llI)]°.5 (empirical) dt * From ANSTO Patent [Khoe e t al. (1997)]
rate increases over range: 0-1.5 mM excess [Fe(III)] reduces the rate* rate decreases with increasing temperature above 25 *C* rate increases with decrease in pH from 3 to 2* _ d[As(l]/)] oc[Fe(m)]°.5[As(lII)] °.z5 dt
so7]02 as an oxidantin hydrometallurgy
1325
In order to apply this observation, a method for removing arsenic from a typical smelter fume containing iron and valuable niickel has also been investigated in this study [Zhang et al. (1998)]. Smelter fume often contains arsenic as As203 as well as significant amounts of base metals and iron. Hence the general method consists of dilute ac,id leaching, oxidation of As(III) and Fe(II) using SO2/O2, and precipitation of iron(HI) arsenate as shown in Figure 6.
Fume STEP 1
Dilute
I
|
Recycle 4 .............................. .-.-..
LEACH
H2SO4
.... '
Species
i
llo o o m a l l l a l i i o e q l m g w i w m lie
Residue ~ ~FLiquor
Smelter
STEP 2
SO2/O2
i
_.[
t
~,T.2+
i
~
r~l+,t~o+
[..... : I
iii iiiiii
d
-
mloom m m mQ
2+
:
• -
Fe /Fe
i...A,s!.V).!..A:s..(!I!!. j
Oxidation
[
.....
.... i;;';:":;
"Fe . . . . . .3~" ..
COI ~ As(III)-~As(V) .PRECIPITATION u. ~.........................
i ..
~l~lemllllmm•wmmmmemoelIQliO
iii
Tailings "*"l F
i
T
~ ) e l e e Q m l l o l l m m w l l l ~
...I):
Precipitate~ •
Liquor ........
[ . . m e.
:
: FeAsO4 •
•
. m ,. m t .
Bleed and Rer~ove Nn, Co
".................. " •
io
.l•
.........
• n
"
Fig.6 Proposed flowsheet for processing arsenical smelter fume. For this particular nickel smelter fume with a high Fe:As ratio, dilute H2SO4 extracted 25-30% of the As(III) and sufficient Fe(II) to allow oxidation and precipitation of FeAsO4 to occur at pH 3.3 with minimal neutralisation cost. By using SO2/O2 as the oxidant, the reaction could be carded out rapidly at 30-38°C unlike oxidation with 02. This approach appears attractive to treat smelter fumes containing significant levels of arsenic and excess iron phases that can be leached with SO2 or dilute sulphuric acid.
1326
W. Zhang etaL
OXIDATIVE PRECIPITATION OF MANGANESE (H) WITH SO2/O2 The kinetics of oxidation of Mn(II) with SO2/O2 and subsequent precipitation of MnO2 have been investigated in the pH range 1-6 and temperature range 25-80°C. It is found that the mechanism of oxidation of Mn(II) is similar to the Fe(II) oxidation, with Mn(III) acting as the catalyst which disproportionates almost spontaneously to MnO2/Mn203 [Zhang (2000)]. The oxidation of Mn(II) is slow at pH < 3, but at pH > 4 Mn(II) is readily oxidised. When applied to nickel laterite leach solutions containing Co and Ni, the oxidative precipitation of Co as Co(OH)3 became significant at pH > 4 and the precipitation of Ni as Ni(OH)3 was significant at pH > 6.5. Therefore, the selective oxidative precipitation of Mn occurs at pH 3-4, and provides a useful means for the removal of manganese impurities from nickel/cobalt leach liquors. Figure 7 shows the precipitation of Mn from a synthetic Co sulphate leach liquor containing 0.1M Co(II) and 0.01 M Mn(II)/Ni(II)/Fe(II)/Cu(II)/Zn(II). As can be seen, 10 mM manganese was selectively precipitated from nickel and cobalt within ten minutes in the pH range 3--4 with less than 1% coprecipitation of Ni and Co. In the presence of 0.01M Fe(II), iron was oxidatively precipitated as Fe(OH)3 before manganese.
100
~" Mn
/I
*
1
80
/Cu
..~ ~o 60 40
I0 minutes for each run at 5 0 " C
/
/
20 Ni, Co m
3
"1=
I
I
I
I
3.5
4
4.5
5
5.5
6
pH Fig.7
Effect of pH on the precipitation of Mn, Cu, Zn and Ni from a Co leach liquor. (10 minute oxidation with 2% SO2/O2; 50°C).
SUMMARY
The results of this study demonstrate that SO2/O2 can function as a strong oxidant in acidic solution for the oxidation of Fe(II), Mn(II) and As(III). It can also function as a means of generating H2SO4 from the autoxidation of SO2 once Fe(II) is oxidised. The SO2/O2 system, catalysed by iron(III) or manganese(III) should be useful for a number of other hydrometallurgical applications. For example, for generating Fe(III) and leaching copper sulphide minerals or uranium(IV) minerals requiring both oxidant and H2SO4. It is also useful for the oxidative precipitation of As(III) and Fe(II) from leach solutions as stable ferric arsenate, and for the oxidative precipitation of Mn(II) as MnO2 from Co(II) and Ni(II) leach liquors at around pH 3.
REFERENCES
Abbruzzese C., Aqueous SO2 processing of manganese ores, in Separation Processes in Hydrometallurgy, Society of Chemical Industry, London, 1987, pp. 76-87.
soz/oz as an oxidantin hydrometallurgy
1327
Brandt C., F~ibi~inL, and van Eldik R., Kinetics and mechanism of the iron(III)-catalysed autoxidation of sulphur(IV) oxides in aqueous solution. Evidence for the redox cycling of iron in the presence of oxygen and modelling of the overall reaction mechanism, Inorg. Chem,, 1994, 33,687-701. Brual G. B., Byerley J. J. and Rempel G. L., Kinetic and mechanistic study of FeS dissolution in aqueous sulfur dioxide solution, Hydrometallurgy, 1983, 9, 307-331. Cheng C.Y. and Lawson F. The kinetics of leaching covellite in acidic oxygenated sulphate-chloride solutions, HydrometaUurgy, 1991, 27,269-684. Chmielewski T. and Charewicz W., The oxidation of Fe(II) in aqueous sulphuric acid under oxygen pressure, Hydrometallurgy, 1984, 12, 21-30. Cornelius R. J. and Woodcock J. T., Pressure leaching of a manganese ore. Part I: Kinetic aspects, Proc. Australas. Inst. Min. Met., 1958, 185, 65-78. Devuyst E. A., Mosoiu A. and Krause E., Oxidising properties and application of the SO2-O2 system, in HydrometaUurgy Research, Development and Plant Practice, (Eds. Osseo-Asare, K. and Miller J. D), Proceedings of the 3rd International Symposium on Hydrometallurgy, The Metallurgical Society of AIME, 1982, pp. 391--403. Du Faur J. B., Australian Patent 15862, 1915. Ettel V. A., Mosoiu M. A. and Devuyst E. A., A novel oxidant for nickel hydrometallurgy, Hydrometallur~:y, 1979, 4, 247-258. George P., The oxidation of ferrous perchlorate by molecular oxygen, J. Chem. Soc., 1954, 4, 4349--4359. Huffman R. E. and Davidson N. J., Kinetics of the ferrous iron-oxygen reaction in sulphuric acid solution, J. Amer. Chem. Soc., 1956, 78, p. 4836. Huie E. and Neta P., One-electron redox reactions in aqueous solution of sulphite with hydroquinone and other hydroxyphenols, J. Phys. Chem., 1985, 89, 3918-3921. Kanungo S.B. and Das R. P., Extraction of metals from manganese nodules of the Indian Ocean by leaching in aqueous solution of sulphur dioxide, Hydrometallurgy, 1988, 20, 135-146. Keenan E. A., A bacterial beneficiation of uranium materials, PhD thesis, University of New South Wales, Australia, 1969. Khoe G. H., Zaw M., Prasad P. S., and Emett M. T., "Photo-assisted oxidation of inorganic species in aqueous solutions", Intl Application Number: WO 99/05065, February 1999. Krause E., The oxidation of ferrous sulphate solutions by sulphur dioxide and oxygen, PhD thesis, University of Waterloo, Canada, 1988. Kwateng O., A kinetic study of the dissolution of nickel sulfide in acidified ferrous sulfate solution with a gas mixture of oxygen and sulfur dioxide, Ph.D Thesis, Columbia University, New York, 1990. Mathews C.T. and Robins R. G., Oxidation of aqueous ferrous sulphate solution by molecular oxygen, Proc. Australas. Inst. Min. Metall., 1972, No. 242, 47-56. McKayD. R. and Halpern, J., Kinetic study of the oxidation of pyrite in aqueous suspension, Trans. Metall. Soc. AIME, 1958, 212, 301-309. Nami F., The kinetics of zinc sulfide leaching by oxygen, sulfur dioxide and ferrous sulfate, PhD Thesis, Columbia University, New York, 1988. Nishimura T., Wang Q. and Umetsu Y., Removal of arsenic from process liquors by oxidation of iron(II), arsenic(III) and sulphur(IV) with oxygen, in Iron Control and Disposal, (Eds: Dutrizac J. E. and Harris G. B.), Proceedings of the Second International Symposium, CIM, Montreal, 1996, pp. 535-548. Pace G. F. and Stauter J. C., Direct electrowinning of copper from synthetic pregnant leach solutions utilising SO2 and graphite anodes-Pilot-plant results, CIM Trans., CIM Bull., 1974, LXXVII, 51-57. Pahlman J. E. and Khalafalla S. E., Selective recovery of nickel, cobalt, manganese from sea nodules with sulfurous acid, US Patent No. 4,138,456,1979. Ralston O. C., Theeferric sulphate-sulphuric acid process, U.S.B.M. Bulletin, 1927, 260, 1-61. Sato T., Goto T., Okabe T. and Lawson F., The oxidation of iron(II) sulphate with sulphur dioxide and oxygen mixtures, Bull. Chem. Soc. Japan, 1984, 57 (8), 2028-2086. Sohn H-J and Wadsworth M. E., Reduction of chalcopyrite with SO2 in the presence of cupric ions, J. of Metals, Noveraber, 1980, pp. 18-22. Subbaiah T., Singh P., Hefter G., Muir D. and Das R.P., Electrowinning of copper in the presence of anodic depolarisers - Part I - A review, Mineral Processing and Extractive Metallurgy Review, 2000, (in press). Thom G. C. and Waters P. F., A study of the metal sulfide-sulfur dioxide reaction in aqueous media by reaction pressure characterisation and ultraviolet spectrophotometry, HydrometaUurgy, 1978,3, 373396.
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Tindall G. P., High temperature of acid leaching of Western Australia laterites, PhD Thesis, Murdoch University, Perth, 1998. Tiwari B. L., The kinetics of oxidation of zinc sulfide and hydrogen sulfide by sulfur dioxide in aqueous sulphuric acid, PhD Thesis, Columbia University, New York, 1976. Tiwari, B. L., Kolbe J., and Hayden H. W., Oxidation of ferrous sulfate in acid solution by mixtures of sulfur dioxide and oxygen, MetalL Trans. B, 1978, 10B, 607 - 612. Vracar R. ~. and Cerovic K, P., Kinetics of oxidation of Fe(II) ions by gaseous oxygen at high temperatures in an autoclave, Hydrometallurgy, 1997, 44, 113-124. Wang Q., Demopoulos G. P., and Harris G. B., A novel process for arsenic oxidation and fixation, in Waste Processing and Recycling in Mineral and Metallurgical Industries III, (Eds. Rao S. R., Amaratunga L. M., Richards G. G. and Kondos P. D), Proceedings of the International. Symposium, Metallurgical Society of CIM, Montreal, 1998, pp. 375-387. Well E. D. and Sandier, S. R., Sulfur dioxide, In Kirk-Othmer Encyclopedia of Chemical Technology, 4th edn., John Wiley & Sons, 1997, 23,299-312. Zhang W., Muir D and Singh P., Leaching and recovery of arsenic from smelter fume, in Waste Processing and Recycling in Mineral and Metallurgical Industries III, (Eds. Rao S. R., Amaratunga L. M., Richards G. G. and Kondos P. D.), Proceedings of the International Symposium, Metallurgical Society of CIM, Montreal, 1998, pp 259-271. Zhang W., Singh P. and Muir D., Iron(II) oxidation by SO2/O2 in acidic media. Part I. Kinetics and mechanism, Hydrometallurgy, 2000(a), 55(3), 229-245. Zhang W., Muir D. and Singh P, Iron(II) oxidation by SO2/O2 in acidic media. Part II. Effect of copper.
Hydrometallurgy, 2000(b ), accepted for publication Zhang W., Singh P. and Muir D., Kinetics of oxidation of As(III) with SO2/O2 and UV light, in Minor Elements 2000: Processing and Environmental Aspects of As, Sb, Se, Te, and Bi, (Ed. Young C.), SME, Littleton CO.,2000(c), pp 333-344. Zhang W., SO2/O2 as an oxidant in hydrometallurgy, Ph.D Thesis, Murdoch University, Perth, Western Australia, 2000. Zhdanov S. I., Sulphur, Chapter IV-6, in Encyclopedia of Electrochemistry of the Elements, Marcel Dekker, 1973, p. 275.
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