THE REMOVAL AND STABILIZATION OF ARSENIC FROM AQUEOUS PROCESS SOLUTIONS: PAST, PRESENT AND FUTURE G.B. Harris Hatch et Associés 5, Place Ville Marie, Bureau 200, Montréal (Québec), H3B 2G2, Canada Tel: +1-514-861-0583; Fax: +1-514-397-1651; e-mail:
[email protected] ABSTRACT The removal of arsenic from metallurgical process liquors continues to be a subject that receives appreciable attention. In this paper, the situation regarding the disposal of arsenic as ferric arsenate is updated, and recent data offering potentially more options to the metallurgist are presented. The pre-eminent positions of both the high-iron (Fe/As molar ratio >4) ambient temperature amorphous arsenical ferrihydrite precipitates, as well as scorodite (Fe/As molar ratio of 1) or other crystalline minerals from high temperature processes, as vehicles for the environmentally safe disposal of arsenic, are once again highlighted. More recent work showing that arsenic can be precipitated as scorodite from both chloride and sulphate media at 80-95oC is discussed, since this offers the non-autoclave operator the option of producing scorodite, at the same time reducing the requirements for iron and lime, and hence the costs of disposal. Finally, a brief review of trivalent arsenic oxidation is presented. It is shown that the more exotic oxidants, such as ozone and elemental chlorine are effective in acidic media and moderate to low temperatures, and that hydrogen peroxide can achieve virtually quantitative oxidation at 95oC at pH 1-3, but its use is expensive. However, most recently, and somewhat interestingly, SO2-O2 gas mixtures have been shown to be effective at temperatures and pHs of interest. The implications of these observations are discussed in the context of future operating practice.
Minor Elements 2000: Processing & Enviroronmental Aspects of As, Sb, Se, Te, Bi Courtney Young, Editor, SME-AIME, 2000, p.3 2000 SME Annual Meeting & Exhibit February 28 - March 1, 2000 Salt Palace Convention Center, Salt Lake City, Utah, USA
INTRODUCTION Over the past decade, in particular, the environmental pressures on the minerals industry have gradually increased to the extent that, today, no plant can afford to emit any kind of even remotely toxic effluent, whether it be gaseous, liquid or solid. However, accidents such as the cyanide spill in August, 1995 at the Omai Gold Mine in Guyana(1,2) and others of recent memory are much more disastrous than merely the accident itself. These events provide fuel for those opponents who, in some cases, will go to extraordinary lengths to ensure that mining will not only not pollute, but also will not even spoil the local environment. Accidents and the threat of pollution also provide fuel for tougher regulations. Most jurisdictions, especially in the so-called developed countries, are presently formulating lists of “toxic” elements which are to be subject to the most stringent regulations. For example, in September, 1993, the Ontario Ministry of the Environment issued draft regulations which were claimed to “set the toughest enforceable limits in the world.”(3) These regulations were aimed specifically at Ontario mines and refineries, with the objective of reducing the emissions of a number of elements, including arsenic, by over 40%. However, they are not unique. The mining industry has responded to these challenges, usually positively, and considerable effort and resources have been expended to reduce toxic emissions. In an address to the Canadian Financial Post Business Outlook Conference in 1992, a senior Canadian mining executive stated, ... practices that were acceptable even ten years ago are simply not tolerated today. As a result, we are designing our new mines (and processing plants) in radically new ways.(4) At the same time, the industry finds itself having to manage with increasingly impure feedstocks. Despite the discovery of the Voisey’s Bay nickel deposit, the days of rich, relatively pure ore deposits belong to history. Many copper, nickel and gold ores, not to mention those of cobalt, lead, zinc and uranium, contain significant amounts of arsenic. Table I lists some of the many arsenic minerals which are known, all of which can find their way into process solutions and effluents. This list is far from exhaustive, but is presented as an indicator of the variety and relative abundance of arsenic minerals. Concurrently, a dramatic improvement in the technology used for extraction processes has resulted in much greater efficiency in recovering both desirable and non-desirable elements, and the extractive industry is faced with the necessity of disposing of its (often toxic) waste products in an environmentally safe manner. This is not a new problem, but in this "green age" it is one which is now gaining increasing public and political awareness. The "traditional" methods of ponding or dumping are no longer automatically acceptable, and the industry worldwide now must not only produce safe waste materials, but also clearly demonstrate that these products are safe. This is apparent in the Toxic Waste Leaching Protocols to be found in most countries of the western world, which are designed to "prove" that a given waste material will not permit the passage into the surrounding environment of any of the toxic elements contained in its mass. In order to survive in this climate, and at the same time remain profitable in a fluctuating market which can see the price of commodity metals vary by 50% within a year, the industry collectively has to address the problems of the toxic impurity elements which are present in every process.
Table I. Some Naturally-Occurring Arsenic Minerals(5)
Mineral Oxides Arsenolite Claudetite Arsenides Allemondite Algodonite Domeykite Niccolite Sperrylite Loellingite Safflorite Rammelsbergite Skutterudite Huntilite Leucopyrite Smaltite Whitneyite Sulphides Realgar Orpiment Dimorphite Cobaltite Arsenopyrite Proustite Xanthoconite Enargite Smithite Tennantite Seligmannite Yrbaite
Formula As2O3 As2O3 AsSb Cu6As Cu3As NiAs PtAs2 FeAs2 (Co,Fe)As2 NiAs2 CoNiAs3 Ag3As Fe3As4 CoAs2 Cu9As As4S4 As2S3 As4S3 CoAsS FeAsS Ag3AsS3 Ag3AsS3 Cu3AsS4 AgAsS2 (Cu,Fe)12As4S13 PbCuAsS3 TlAs2SbS5
Mineral Simple Arsenates Scorodite Kankite Karibibite Angelellite Pharmacosiderite Symplesite Pharmacolite Haidingerite Hoernesite Olivenite Euchroite Cornwallite Clinoclasite Lavendulite Trichalcite Tyrolite Leucochalcite Sarkinite Adamite Complex Arsenates Arseniosiderite Bukovskyite plus many others Chloroarsenates Mimetite plus many complex halo-arsenates
Formula FeAsO4.2H2O FeAsO4.3½H2O Fe2As4O9 Fe4As2O11 Fe3(AsO4)2(OH)3.5H2O Fe3(AsO4)2.8H2O CaHAsO4.2H2O CaHAsO4.H2O Mg3(AsO4)2.8H2O Cu2(AsO4)OH Cu2(AsO4)OH Cu5(AsO4)2(OH)4.H2O (CuOH)3.AsO4 Cu3(AsO4)2.2H2O Cu3(AsO4)2.5H2O Cu(CuOH)4(AsO4)2.3½H2O Cu2(OH)AsO4.H2O Mn2(AsO4)OH Zn2(AsO4)OH
Ca3Fe(AsO4)3.3Fe(OH)3 Fe2(AsO4)(SO4)OH.7H2O
Pb5(AsO4)3Cl
Whilst it may not be the most toxic of elements, particularly in its pentavalent state(6), arsenic always has and always will attract attention simply because of its notoriety as a poison both in history and in popular literature. In the famous British play, and later film, of the 1940s, Arsenic and Old Lace, it is much more likely, as noted by Hopkin,(7) that it was the strychnine content of the poison concoction (elderberry wine, strychnine, arsenic and cyanide) rather than the arsenic which was the lethal agent, yet it was arsenic which received prime billing. In dealing with arsenic, the emphasis to date has of necessity been on solid arsenical residues for long term safe disposal. As noted below, maximum stability and maximum removal from liquid effluents are achieved when both iron and arsenic are in their highest oxidation states. Since many metallurgical process streams have arsenic in a lower oxidation state, attention has recently been turned to its oxidation and removal from solution, whilst at the same time
maintaining the stability of the residue formed. This paper briefly examines the current state of knowledge in this respect; however, it is once again pertinent to first review of the situation regarding ferric arsenates. ARSENICAL SOLID WASTES A considerable research effort has been expended by the Minerals Industry over the last twenty years in order to understand and develop processes for the precipitation of arsenic from aqueous process liquors in a stable, environmentally-acceptable form for disposal. It is now generally accepted that the crystalline and high-iron amorphous ”ferric arsenates,” in their various forms, are the most appropriate compounds to form, and that these have a very low solubility under normal aqueous disposal conditions.(8) However, as will be discussed later, the disposal environment itself also has a significant role to play. The research effort has been one of considerable magnitude, with major studies in one form or another having been conducted at three centres in Canada (Noranda, Inco and McGill University), in Australia (Robins, firstly at the University of New South Wales and more latterly with the Aquamet Science Consortium), in Japan (Tohoku University), in Greece (the now-defunct Metba), and the MIRO (Mineral Industry Research Organization)-sponsored study at the Royal School of Mines (Imperial College) in the U.K. Additionally, there have been several minor studies, all of which have contributed to the knowledge base on this topic. The present author attempted to collate all of the available data regarding ferric arsenates in 1993.(8) Whilst most of the conclusions drawn at that time remain valid, research has continued and it is appropriate to update the situation. The study of ferric arsenate precipitation has primarily been undertaken in two distinct ways; its precipitation as a ”high iron” or ”basic” ferric arsenate (both of which are now more correctly referred to arsenical ferrihydrite(9)) from process liquors under ambient pressures at temperatures from 25 to 95oC, and its precipitation as a crystalline “mineral” at the autoclave temperatures generally encountered in refractory gold pressure leaching (180 to 225oC), although it should be pointed out that recent work has shown the possibility of forming crystalline ferric arsenate (scorodite) under non-autoclave conditions. Low-Temperature Ferric Arsenates Robins,(5,10,11,12,13,14,15,16,17,18,19,20,21,22) beginning in the mid-1970s, and following upon work conducted by Tozawa et al.,(23,24) led the study into the stability of metal arsenates, and in particular into the stability of ferric arsenates. Much of the early work, however, was theoretical, in that it was based on published thermodynamic data, and indicated that simple (molar Fe/As ratio of 1) amorphous ferric arsenates were not sufficiently stable for safe environmental disposal. Such conclusions have since been verified in laboratory studies.(25,26,27) Since the publication of the early thermodynamic diagrams, both the method of calculation and the accuracy of the data have significantly improved with new measurements and a more complete understanding of the systems involved. Figure 1 shows the stability-area diagram for six of the more commonly encountered base metals, based on up-to-date, computer-generated calculations conducted by Twidwell and his group at Montana Tech.(28) Solution concentrations of 1 g/L for both As(V) and the metal ion have been chosen, quite arbitrarily, but the program can develop stability regions for any concentration.
F e(III)-Cu-Zn-P b-C a-Cd-As-H2O A: Fe(III) B: Cu(II) C: D: E: F: G:
A
F
100 10
1
Conditions
D
C
Fe(III) 1g/L Cu(II) 1g/L Zn(II) 1g/L Pb(II) 1g/L Cd(II) 1g/L
B G
0.1
E
Ca(II) 1g/L As(V) 1g/L
A
0.01
Zn(II) Pb(II) Cd(II) Ca(II) As(V)
0.001
D 0
2
4
6
8
10
12
14
pH Figure 1. Calculated Stability Area of Some Simple Base Metal Arsenates at 25oC(28) It is worth noting here that the multi-client research program being carried out by MIRO is now (July, 1999) aiming to develop a set of tools with which to be able to accurately model both the high and low temperature systems where iron and arsenic are present. The objective is to be able to predict the process chemistry and operating conditions which will give the most stable arsenic precipitate, without the need to carry out an exhaustive testwork program. At about the same time that Robins began to examine the theoretical aspects of arsenate compounds, work was initiated at the research facilities of both Noranda(25,26) and Inco(27,29,30,31,32) to study the long term stability of arsenical solid wastes. These studies, along with all of the recent published literature, are essentially in agreement that non-iron-containing arsenates, and the whole family of arsenites, are not sufficiently stable for disposal purposes, although it should be noted that barium arsenate has been used to remove and stabilize arsenic.(33) The validity of this, however, has been questioned.(34) It is beyond the scope of this paper to reproduce the data for the non-iron and arsenite systems, mainly because they are only of passing interest, having no real practical use. The data to show that such compounds should be avoided can be found in the Noranda and Inco papers referenced above, as well as in some of the early Robins work. One point, however, should be noted - that atmospheric carbon dioxide can and does destabilize simple metal arsenates, such as those of cadmium, copper, zinc and especially calcium, all of which have carbonates with a significantly lower solubility product than their corresponding arsenate. Such is not the case with the ferric arsenates. The findings of the laboratory studies have been confirmed with available plant data,(8) clearly
demonstrating that co-precipitation of pentavalent arsenic with ferric iron is an effective vehicle for the environmental immobilization of arsenic, provided that the Fe/As molar ratio is at least three, and preferably >4. Additionally, relatively recent plant data published by Noranda(35) from work with the weak acid bleed at its Horne smelter, show that the incorporation of base metals into the arsenic precipitate has an additional stabilizing effect at ambient temperature. This demonstrates in practice the observations first reported by Harris and Monette.(26) Recently, and somewhat surprisingly given the body of literature in the public domain, a variation of the iron precipitation process for the removal of arsenic from wastewaters has been patented by McClintock.36 This process removes low levels of arsenic from wastewaters by the addition of iron and an oxidant (sodium hypochlorite), apparently much in the same way as described by Inco and Noranda. High Temperature Precipitates High temperature work was initiated at the Royal School of Mines (RSM), part of Imperial College (IC), in London, U.K., initially to look at the structures of the arsenate minerals formed. This prolonged study has examined the formation, structures and stability of ferric arsenates using a variety of instrumental techniques.(37,38,39,40,41,42,43) The work is being undertaken as a co-operative research program through MIRO (the UK-based Mineral Industry Research Organization), with sponsorship at various times by fourteen companies worldwide, and is presently into its fourth phase. Initial work concentrated on low temperature precipitates, but it became apparent that further work should focus on high temperature iron/arsenic precipitates, such as those generated during the treatment of refractory gold ores in autoclaves, due to the crystallinity of the compounds formed. The work has shown the formation of some iron-arsenic minerals, including scorodite, but also others distinctly different, the compositions of which depend upon factors such as temperature and the molar ratios and concentrations of iron, arsenic and sulphate.(39) Similar minerals, such as zykaite {Fe4(AsO4)3(SO4)(OH).15H2O} and bukovskyite {Fe2(AsO4)(SO4)(OH).7H2O}, both containing sulphate, have been found in old dumps as secondary minerals formed from the weathering of arsenopyrite, suggesting stability under normal environmental conditions. The current phase of the study aims at further investigation of these dumps, since identification of the same minerals in old dumps as produced during processing will, effectively, demonstrate their long-term stability, without the need for additional long-term testing. The MIRO studies have focused upon high temperature materials, since the iron/arsenic system initially appeared to be less complex at high temperatures. However, the data that have been generated suggest that there are a number of different crystalline and semi-crystalline minerals which can be formed, some of which are very stable and others which, although stable, are less so. The current phase of the study aims at determining the fields wherein these different minerals are formed, to correlate the stability fields with both feed composition and process operating conditions, and hence to develop a predictive model. Scorodite (FeASO4.2H2O) The studies at the Royal School of Mines suggest that although crystalline scorodite has the lowest solubility of the arsenic minerals produced, scorodite, the Type 2 mineral (as designated by the Royal School of Mines, and also referred to as Phase X) and the high iron ferric arsenates all have a comparable order-of-magnitude solubility. This is demonstrated in Figure 2, where, in this case, crystalline scorodite represents all of the above, and is some two orders of magnitude less soluble than the corresponding amorphous compound also with an Fe/As
molar ratio of 1.
Figure 2. Solubility of Amorphous and Crystalline 1:1 Ferric Arsenates(8) More recently, as discussed later, studies have shown that it is also possible to form scorodite under non-autoclave conditions. Practical adoption of such a process should have considerable economic benefits to those operations with an iron deficiency, such as the Noranda Horne smelter referred to above, where iron-containing solutions from mine run-off are pumped several kilometres. The formation of crystalline scorodite has been reported by various researchers from a number of different chemical environments. In 1988, Dutrizac and Jambor(44) synthesised crystalline scorodite from a nitrate solution at pH=7 and 160oC, equilibrated for 24 hours in an autoclave, where As(V) concentrations greater than 15 g/L were needed to ensure a pure product. Papangelakis and Demopoulos(45) reported its formation during acid pressure oxidation of arsenopyrite in a 0.5M H2SO4 medium. Conditions favouring the build-up of Fe(III) and As(V) in solution (>160oC and high slurry-densities) resulted in the precipitation of well-grown scorodite crystals with a diameter of >20m. Dutrizac et al.(46) also produced crystalline scorodite during jarosite precipitation (24 h) at 150oC from 0.6M SO42- solutions containing 5 g/L As(V) at an initial pH=1.3. However, at 97oC, an amorphous ferric arsenate was co-precipitated with the jarosite.(46) Precipitation of scorodite from a chloride (and also nitrate) medium has been found to be somewhat easier than from sulphate. Dutrizac et al.(47) precipitated crystalline scorodite from 4.5M Cl- solutions at 97oC, initial pH=1 and 24 h equilibration. Dove and Rimstidt(48), Kunter and
Bedal(49), and Demopoulos and Kondos(50) have each reported the formation of crystalline scorodite from chloride media. However, all of these were either under pressure or required very long equilibration times at 100oC. Robins(21) prepared crystalline scorodite from a sulphate medium at pH=1 by refluxing the solution at an undisclosed temperature for several days, presumably at ambient pressure. However, this resulted in very fine scorodite crystals of approximately 60nm. More recently, Demopoulos et al.(51,52) described a novel technique in which crystalline scorodite could be precipitated at ambient pressure and lower temperatures (80-95oC) from chloride media. The precipitation method was based on strict supersaturation control in combination with scorodite seeding. Attempts were then made to reproduce this from sulphate media under the sponsorship of a MITEC (Mining Industry Technology Council of Canada) multi-client project. Sulphate solutions were investigated, since such predominate in the metallurgical processing industry. Initially, it was found to be much more difficult to precipitate crystalline scorodite, the technique developed for chloride media proving not to be possible. This was attributed to the fact that the ferric ion has a much more complex chemistry in a sulphate solution, forming a number of sulphato- and hydroxy- complexes, the number and complexity increasing as the pH rises.(53,54) However, further work at McGill University showed that it was indeed possible to precipitate crystalline scorodite from sulphate media, at 95oC over a period of 1-2 hours, provided that (i) sufficient scorodite seed was added, and (ii) that there was a deficit of iron in the starting solution, a deficit being defined as Fe/As molar ratio 3, and hydrothermal production of scorodite or other arsenic minerals at elevated temperatures (>160oC) and pressures (above atmospheric).
DISCUSSION AND CONCLUSIONS An attempt has been made here to cover the vast ground of arsenic removal from metallurgical process liquors. Inevitably, any review such as this can only scratch the surface, but there are general trends which can serve as useful guidelines for both the researcher and the plant operator. Options Available This brief review has shown that whilst former conclusions remain valid, new process options are being introduced which should lead to more effective removal and disposal from both technical and economic viewpoints. Much has been written regarding the stability of ferric arsenates, whether they be the low temperature, high-volume, amorphous arsenical ferrihydrites generated by ambient precipitation methods, or the formation of scorodite and other crystalline arsenic minerals at autoclave temperatures. New work has shown that scorodite can now be produced from both chloride and sulphate media at ambient pressures (80-95oC) by controlled precipitation techniques. Further work is required in these areas to demonstrate their effectiveness under plant conditions. Stability testing, such the US EPA TCLP, suggest that both the low and high temperature materials are environmentally acceptable, as shown in Table 2 below.(43) It is worth pointing out the very high values obtained for the calcium arsenate phases. Although not shown in the data of Table 2, the values obtained for scorodite and Type 2 are generally «1 mg/L, as shown for scorodite by Krause and Ettel,(31) but do depend to some extent on the method of preparation. Figure 3, taken from the work of Monhemius and Swash,(9) shows a simplified, approximate speciation diagram for initial solution Fe/As ratio versus temperature. This diagram is useful, since when read in conjunction with Table 2, it shows that there are conditions in autoclave processing that should be avoided if a stable arsenical residue is desired. If the Fe/As ratio is too low, then the Type 1 mineral, which is relatively soluble, predominates. Similarly, conditions that favour the formation of basic iron sulphates should also be avoided. The precipitation work has also demonstrated the necessity of having iron(III) and arsenic(V) present together in solution at the time of precipitation. However, under many situations, this is not the case, and it is therefore necessary to effect in-situ oxidation of one or both iron and arsenic. The more exotic oxidants, such as ozone and elemental chlorine have been shown to be very effective, and more recent work has shown that hydrogen peroxide can be used almost quantitatively to effect oxidation at 80-95oC in acidic media. Even more surprisingly, SO2-O2 mixtures have been shown to be effective oxidants for arsenic at 50oC at pH values as low as 2, and for both arsenic(III) and iron(II) at 95oC at even lower pH values. The net result is that there are now a number of processing options available for the oxidation, removal and disposal of arsenic from metallurgical process liquors.
Table 2. Relative TCLP Solubilities of Precipitated Iron and Calcium Arsenates(43) TCLP solubility (As mg/L in filtrate)
Precipitated Phase Scorodite
3000
* precipitated calcium arsenate phases
Type-2
225
Type-1
Basic iron sulphate
Temp. of synthesis
Type-2
200
Basic Basic iron iron sulphate sulphate 175
Scorodite
Type-1
150
1:0 Fe
9:1
4:1
2.3:1
1.5:1
Decreasing Fe:As Increasing As levels in starting solutions
1:1 As
E.g. A solid precipitated from a 9:1 solution at 225°C yields ~90% basic iron sulphate and ~10% Type 2
Figure 3. Distribution of compounds formed from Fe-AsO4-SO4 solutions (pH4, to be stable for at least eight years in an aqueous environment. Conversely, data reported from Nishimura(7) showed a solubility of 375 mg/L As after three years for an arsenical ferrihydrite of Fe/As molar ratio of 5, which had an initial solubility of only 0.01 mg/L As. As pointed out by Dutrizac and Jambor, in an excellent review of ferrihydrite chemistry,(68) and others,(21) the iron component may ultimately crystallise over time to goethite or another iron mineral, with the (probable) desorption and release of arsenic. However, that this will happen is far from certain, since the presence of foreign ions such as arsenic have been shown to significantly retard this crystallisation process.(69,70) It is also conceivable, though perhaps less likely, that the iron and arsenic will react to form scorodite. Long term reflux tests conducted by Krause and Ettel(29) several years ago suggest that desorption is more likely,
whereas more recent work by Swash and Monhemius(9) has shown that with prolonged heating, the initial amorphous soluble material appeared to crystallise and yield solids with significantly lower solubilities Nishimura and Robins(71) observed the formation of three crystalline products, when initially amorphous Fe/As molar ratio of 1.01 solids were equilibrated for up to three months in arsenic acid solutions - FeAsO4.2H2O (scorodite), FeH3(AsO4)2.10H2O and Fe(H2AsO4)3.5½H2O. The second of these was considered only to be metastable, although solubility data showed all three to give
