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Chemistry of the drainage from a waste dump at BHP-Utah Mines Ltd, Island Copper mine Li, Michael Guoqing 1991

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CHEMISTRY OF THE DRAINAGE FROM A WASTE DUMPAT BHP-UTAH MINES LTD, ISLAND COPPER MINEByMICHAEL GUOQING LIB.Eng., Wuhan University of Iron and Steel Technology, China, 1982A THESIS SUBMITTED IN PARTIAL FULFILLMENT OFTHE REQUIREMENTS FOR THE DEGREE OFMASTER OF APPLIED SCIENCEinTHE FACULTY OF GRADUATE STUDIESTHE DEPARTMENT OF MINING AND MINERAL PROCESS ENGINEERINGWe accept this thesis as conformingto the required standardTHE UNIVERSITY OF BRITISH COLUMBIADecember, 1991© Michael Guoqing Li, 1991In presenting this thesis in partial fulfilment of the requirements for an advanceddegree at the University of British Columbia, I agree that the Library shall make itfreely available for reference and study. I further agree that permission for extensivecopying of this thesis for scholarly purposes may be granted by the head of mydepartment or by his or her representatives. It is understood that copying orpublication of this thesis for financial gain shall not be allowed without my writtenpermission. (Signature)Department of MINING AND MINERAL PROCESS ENGINEERINGThe University of British ColumbiaVancouver, CanadaDate nErEMRER 31, 1991DE-6 (2/88)ABSTRACTThis thesis examines the chemistry of the drainage around the North Dump at BHP-UtahMines Ltd., Island Copper Mine. It also presents the calculations of the pathway of waterfrom precipitation and the acid generation and consumption of some areas on the NorthDump.Information used in this study includes mine site fresh water quality monitoring records,acid-base accounting results of the North Dump drilling samples, flow rate measurementsof drainage streams around the North Dump and of pit dewatering, physical parametersof the North Dump, surface topographical maps, meteorological records, etc.It has been found that a typical contaminated drainage has a total ionic strength of 0.0426and that the activity coefficients of the dissolved species in consideration depart markedlyfrom unity. Most dissolved species are found to be under-saturated with respect to theirmost insoluble compounds, although over-saturation also exists. Levels of dissolved speciesare controlled by their rates of release and the frequency and amount of precipitation.Redox reactions and solubility control are investigated. Chemical reactions involvingminerals present in the waste rock are considered in view of their effects on acid generationor consumption and on release rates of dissolved species. Eh-pH diagrams and solubilitydiagrams are constructed at conditions specific to the drainage under study.iiTABLE OF CONTENTSPageABSTRACT ^  iiLIST OF TABLES  vLIST OF FIGURES ^  viiLIST OF ABBREVIATIONS ^  xSPECIAL NOTES ^  xiACKNOWLEDGEMENT  xii1.0 INTRODUCTION ^ 12.0 LITERATURE REVIEW 52.1 Acid Generation Mechanisms ^ 52.2 Thiobacillus ferrooxidans 192.3 Literature on Island Copper Mine 243.0 RESEARCH SITE^ 253.1 General Description ^ 253.2 Geology 283.3 The North/Old North Dump ^ 293.4 Hydrology^ 333.5 Land Dump Monitoring 354.0 RESEARCH METHODS ^ 384.1 Data Collection 384.2 Drainage Sample Collection ^ 394.3 Methods of Chemical Analysis and Physical Measurement ^ 395.0 RESULTS AND DISCUSSION ^ 405.1 Acid Generation Assessment of the North Dump ^ 405.1.1^Acid-Base Accounting 405.1.2^Flowrate Analysis ^ 405.1.3^Water Balance Calculation ^ 475.1.4^Calculation of Acid Generation Life of EMO and Caps ^ 48iii5.2 Land Dump Drainage Monitoring Records ^  495.3 Chemistry of North Dump Drainage Water  595.3.1 Mineral Constituent Mobilization and Dissolved SpeciesRemoval in General ^  615.3.2 Redox Reactions  715.3.2.1 Sulfur  835.3.2.2 Carbon ^  865.3.2.3 Calcium  915.3.2.4 Magnesium  955.3.2.5 Copper ^  975.3.2.6 Iron  1075.3.2.7 Lead  1195.3.2.8 Manganese ^  1245.3.2.9 Molybdenum  1315.3.2.10 Zinc  1335.3.2.11 Cadmium ^  1385.3.2.12 Arsenic  1405.3.3 Solubility Control  1425.3.4 Other Processes ^  1655.3.5 Bacterial Action  1735.3.6 Kinetic Considerations  1755.3.7 Underwater Disposal Considerations ^  1786.0 CONCLUSIONS ^  182BIBLIOGRAPHYAPPENDIX IAPPENDIX II 187Calculation of the Pathway of Rain Water in 1987 ^ 194Calculation of Acid Generation and Consumption of EMO andCaps ^  196ivLIST OF TABLESPageTable 1 Physical Parameters of Designated Regions on North Dump ^ 31Table 2 Island Copper Regular Fresh Water Sampling Stations  37Table 3 Acid-Base Accounting Results of North Dump Drilling Samples ^ 41Table 4 Volumetric Flow Rate Monitoring Data of Some Drainage Streams ^ 44Table 5 Summary of Flowrate Data Analyses ^  45Table 6 Pathway of Rain Water Falling on an Enclosed Area on North Dump ^ 47Table 7 Pathway of Rain Water Falling on EMO ^  48Table 8 North Dump Drainage (NDD) Water Quality 1980-1989 ^ 50Table 9 Twin Lake Discharge (TLD) Water Quality 1986-1988  51Table 10 North Dump Subterranean Flow (NDS) Water Quality 1988 ^ 52Table 11 Ten Eighty Ditch (TED) Water Quality 1986-1989 ^  53Table 12 Trey Creek (TCR) Water Quality 1987-1989  54Table 13 Eastern Most Outslope (EMO) Water Quality 1987-1989 ^ 55Table 14 East Drainage Lake (EDL) Water Quality 1988-1989  56Table 15 East Dump Drainage (EDD) Water Quality 1981-1988 ^ 57Table 16 East Dump Drainage (EDD) Water Quality 1988-1989  58Table 17 Range of Contaminant Concentrations January 1988 - January 1989. ^67Table 18 Composition of the Typical Drainage ^  69Table 19 Solubility Products of Salts and Bases of Interest at 25°C and 1 atm ..^145vTable 20 pH Ranges of Metal Ion Activity Control by Sulfates, Carbonates andHydroxides^  156Table 21 State of Saturation of Metal Ions in TED 1988 ^  159viLIST OF FIGURESPageFigure 1^Location Map ^  26Figure 2^Location of North Dump at Island Copper Mine ^  30Figure 3^Regions of North Dump and Locations of Drill Holes  32Figure 4^Pre-Mining Surface Topography ^  34Figure 5^North Dump Drainage Collection Ditch System and Fresh WaterSampling Stations ^  36Figure 6^Contribution of Fresh Water Streams to Pit Dewatering Pond ^ 46Figure 7^Distribution of Eh-pH Measurements of Natural AqueousEnvironments^  75Figure 8^Approximate Positions of Some Natural Environments asCharacterized by Eh and pH ^  76Figure 9^Change in pE (Eh) of a Fresh Water as a Function of the Amount ofOrganic Matter Decomposed^  78Figure 10 Approximate pE (Eh) Values at which Various Redox Reactions Occurin Water at pH 7 and 25°C ^  79Figure 11^Some Possible Redox Buffers in a Groundwater System ^ 79Figure 12^Eh-pH Diagram for Part of the System S-O-H at ES=10 -3 ^  84Figure 13^Eh-pH Diagram for Part of the System S-O-H at ES=10 -2 ^  85Figure 14 Eh-pH Diagram for Part of the System C-O-H ^  87viiFigure 15^Logarithmic Concentration Diagram of Species in Equilibrium for 10 -3M Carbonic Acid^  89Figure 16 Eh-pH Diagram for Part of the System Ca-C-O-H-S ^ 92Figure 17 Eh-pH Diagram for Part of the System Ca-C-O-H-S at a HigherActivity ^  93Figure 18 Eh-pH Diagram for Part of the System Mg-O-H ^  96Figure 19 Eh-pH Diagram for Part of the System Cu-O-H  98Figure 20 Eh-pH Diagram for Part of the System Cu-S-O-H^ 99Figure 21 Eh-pH Diagram for Part of the System Cu-S-C-O-H  100Figure 22 Eh-pH Diagram for Part of the System Fe-O-H Assuming Fe(OH) 3 asStable Fe(+III) Phase ^  108Figure 23 Eh-pH Diagram for Part of the System Fe-S-O-H Assuming Hematiteas Major Fe(+III) Phase ^  109Figure 24^Activities of Fe Species in Equilibrium with Fe(OH)3 ^  118Figure 25 Eh-pH Diagram for Part of the System Pb-S-C-O-H  120Figure 26 Eh-pH Diagram for Part of the System Mn-O-H ^ 125Figure 27 Eh-pH Diagram for Part of the System Mn-C-S-O-H ^ 126Figure 28 Eh-pH Diagram for Part of the System Mo-S-O-H  132Figure 29 Eh-pH Diagram for Part of the System Zn-C-S-O-H ^ 134Figure 30 Eh-pH Diagram for Part of the System Cd-C-S-O-H  139Figure 31^Eh-pH Diagram for Part of the System As-S-O-H ^ 141Figure 32^Logarithmic Concentration Diagram Showing the Activities of SomeMetal Ions in Equilibrium with Their Sulfates at 25°C. ^ 147Figure 33^Logarithmic Concentration-pH Diagram Showing the Activities ofviiiVarious Metal Ions in Equilibrium with Their Hydroxides at 25°C. . 149Figure 34^Logarithmic Concentration Diagram Showing the Activities of VariousMetal Ions in Equilibrium with Their Carbonates at 25°C. ^ 150Figure 35^^Logarithmic Concentration-pH Diagram Showing the Activities ofVarious Metal Ions in Equilibrium with Their Carbonates at 25°C. . 151Figure 36 Logarithmic Concentration-pH Diagram Showing the Activities ofVarious Metal Ions in Equilibrium with Both Their Carbonates andTheir Hydroxides at 25°C.   154Figure 37 Logarithmic Concentration-pH Diagram Showing the Activities ofVarious Metal Ions in Equilibrium with Both Their Carbonates andTheir Hydroxides at 25°C (Higher Carbon Activity)   155Figure 38 Activity-pH Diagram Showing the Maximum Activities of VariousMetal Ions as Controlled by Their Equilibrium with Sulfates,Carbonates and Hydroxides at 25°C  157Figure 39 Schematic Diagram Showing the Rate of Acid Production over Time 177ixLIST OF ABBREVIATIONSFRESH WATER SAMPLINGSTATIONSOTHER TERMSAAS^- Atomic Absorption SpectrometryBLA - Bay Lake (Francis Lake) (A.A.S.)EDD - East Drainage Ditch ACP - Acid Consuming PotentialEDL - East Drainage Lake AMD - Acid Mine DrainageEMO - Eastern Most Outslope APP - Acid Producing PotentialNDD - North Drainage Ditch ARD - Acid Rock DrainageND S - North Dump Subterranean MPN - Most Probable NumberPDP - Pit Dewatering Pond NNP - Net Neutralizing PotentialPDW - Pit Dewatering SHE - Standard Hydrogen ElectrodePTP - Pit Treatment PondSCR - Stephens CreekTCR - Old Trey CreekTLA - Twin Lake Discharge(TLD)SPECIAL NOTESThe data collected in this study are up-to-date to January, 1989. The first draft of thisthesis was completed in December, 1989. Therefore, the conclusions and statements madein this thesis are only up-to-date to January, 1989, although the thesis is dated October,1991.Both the Metric and the Imperial systems of measures are used in this thesis, since theoriginal data exist in both systems. Whenever a calculation involved figures in bothsystems, the figures in the Imperial system were converted to those in the Metric system,and the result was presented in the Metric system. Special attention is required todistinguish between the Metric ton, spelled as tonne (tonnes), and the Imperial ton, spelledas ton (tons).References are identified in the text with the author's name followed by the year ofpublication. References without named authors are identified with the title of thepublication.xiACKNOWLEDGEMENTI would like to express my sincere thanks to Dr. George Poling for the excellent academicguidance, insightful suggestions and inexhaustible patience he has given me during thecourse of this study. My thanks also go to Ron Hillis and Ian Horne of BHP-Utah MinesLtd., Island Copper Mine for their generous help in the data collection process.I am very grateful to BHP-Utah Mines Ltd., Island Copper Mine and the Department ofMining and Mineral Process Engineering at the University of British Columbia for theirfinancial support.3di1.0 INTRODUCTIONAcid rock drainage or ARD, sometimes referred to as acid mine drainage, or AMD, is aserious problem facing today's mining industry. ARD results from the exposure by miningof some reactive sulfide minerals, primarily pyrite, to air and water. Prior to mining,sulfide minerals are usually in a state of slow oxidation or no oxidation at all under burialconditions. Mining activities bring a large quantity of sulfides into direct contact withabundant oxygen in air and water from precipitation, resulting in a remarkable increasein the surface area of sulfides. Oxidation accelerates considerably and ARD results. ARDgeneration is inevitably catalyzed by bacteria except at extreme conditions such as freezing.Unfortunately, sometimes ARD is unavoidable because sulfides occur with many mineralsof mining interests and many metals exist and are mined as sulfides. Sources of ARD ina mine include open pit walls, waste dumps, access roads, stock piles, tailings facilities, andunderground workings.ARD gives rise to two kinds of problems: productional and environmental. ARD corrodesmining equipment and, when cleared by pumping from underground workings or from openpits during production, also corrodes pumps and pipes, thus increasing the operational cost.-1-The environmental damage that ARD causes is serious. The low pH in ARD helps dissolvemetals such as Ca, Mg, Hg, Zn, Cd, Pb and non-metal elements such as As, P, and S fromrock. Some of these elements are toxic or accumulatively toxic to animals and men. ARDis thus often an aqueous mixture characteristic of high total dissolved solids, high acidity,and low pH values. If ARD infiltrates into ground water, the quality of the ground waterwill deteriorate and become unsuitable for drinking and other purposes. If ARD is releaseduntreated to natural waters, it can kill aquatic lives, overthrow ecological balances, andmenace human beings through the food chain.While ARD adds a little cost to the production, its environmental implications concern notonly mining companies but also governments and the public. The treatment of ARD inorder to conform to environmental standards could easily turn a profitable mining operationinto a money-losing business. The probability of ARD production and remediation shouldbe well examined during the planning stage of a new mine. The aftermath of anirresponsibly (or improperly) abandoned mine that produces ARD could be even more costlyand often becomes a financial burden of local governments, as illustrated by the RumJungle Rehabilitation Project in Australia (Northern Territory Department of Mines andEnergy, 1986) and the Mount Washington Acid Mine Drainage Abatement Program inBritish Columbia (Steffen Robertson and Kirsten et. al., 1987). For these reasons and forthe purpose of environmental protection, both the Canadian and the U.S. governments haveissued stringent standards for mining and milling effluent to be released to biota-containingwaters. These standards usually specify the minimum allowable pH and maximumallowable levels of total suspended solids and a few dissolved metals. The British Columbia2New Mine Development Review Process prescribes the examination of ARD issue to be apre-requisite in the application of a new mine.In light of the seriousness of the ARD problem, studies of various aspects of ARD prosperedin the last three decades and are widely dispersed in the literature. These studies include,among others, prediction and modelling of ARD in acid-generating environments, laboratorysimulation of the acid generation process, the bacterium Thiobacillus ferrooxidans, variousremedial methods, and different preventive techniques. However, studies of the solutionchemistry specific to ARD are relatively scarce. This thesis examines from athermodynamic point of view the solution chemistry of the drainage emanating from theNorth Dump at BHP-Utah Mines Ltd., Island Copper Mine located on the northern end ofVancouver Island.The present thesis evolves from a research project entitled "Acid Mine Drainage Study ofthe North Dump" undertaken jointly by the Department of Mining and Mineral ProcessingEngineering at the University of British Columbia and BHP-Utah Mines Ltd., IslandCopper Mine (Acid Mine Drainage Study of the North Dump, 1991). The objectives of theIsland Copper Mine project were to:■^determine how widespread and significant acid generation is in the Old NorthDump, the North Dump and the Eastern Most Outslope regions of the North Dump;3■ help develop both short and long-term procedures to mitigate or abate acidgenerations from these waste dumps, and help develop plans for abandonment of theabove mentioned waste dumps.Data collected in the Island Copper Mine project are used in this thesis. The objectives ofthe present thesis are to:■ investigate the mechanism controlling the levels of dissolved species in the drainageby means of equilibrium chemistry and examine the interactions among differentdissolved species in the drainage;■ study the solubilization of metal species from waste rock and the removal of metalspecies from the drainage solution;■ infer the status of the dump with respect to oxidation and acid generation and helpdevelop conceptual mitigation measures;■ calculate the acid generation life and the acid consumption life of the EMO and theCaps and evaluate the water balance of two self-enclosed areas;■ explore the role played by bacteria.42.0 LITERATURE REVIEWAcid mine drainage probably has as long a history as the mining industry itself. However,it was not until 1919 that the involvement of bacteria in acid mine drainage formation wasrecognized when Parr and Powell reported that the coal inoculated with unsterilized ferroussulfate solution produced higher sulfate concentration than did the sterilized one.In 1947, the bacteria Thiobacillus ferrooxidans were first isolated by Colmer and Hinkle(Colmer and Hinkle, 1947). At this time, Thiobacillus ferrooxidans were merely consideredto oxidize iron by precipitating Fe(OH) 3. Since then, the study of acid mine drainage hasflourished. The following sections give an account of some aspects of acid mine drainagestudies related to this thesis.2.1 Acid Generation MechanismsIn 1970, Singer and Stumm showed that natural oxygenation was too slow to account forthe rapid acid generation in acidic mine environments, therefore they proposed that5Thiobacillus ferrooxidans catalyzed the pyrite oxidation by converting ferrous iron to ferriciron, which in turn acted as an oxidant to pyrite (Singer and Stumm, 1970).Kleinmann and Crerar demonstrated that Thiobacillus ferrooxidans can survive at pHvalues as high as 6.0 and at a pH slightly above 4.0 can reproduce at a relatively rapid rate,given the condition of minimum water flow (Kleinmann and Crerar, 1979). They inferredthat Thiobacillus ferrooxidans directly oxidized pyrite, created a semi-isolated micro-environment, and spread once the local pH was lowered to roughly 4.0. They concludedthat at such conditions the precursor Metallogenium, which was believed to mediate theacid formation process by bringing pH from 4.5 to 3.5 (Walsh and Mitchell, 1972), was notnecessary.A pH of 2.0 is not unusual in acid mine drainage. The lowest recorded pH is 0.5 (Rawat,1983). The initial acidification process in acid mine drainage formation may proceedabiotically or with the help of bacteria. Abiotic acidification is possible but very unlikelyto bring pH to below 4.0 because of its relatively slow reaction rate, its slowdown bydecreases in pH and the lowering of acidity by frequent washing effect of rainfalls. Onescenario where abiotic oxidation could bring pH to below 4.0 would be a case whereprecipitation were very little so that the acidity could accumulate, where the temperaturewere so low that it virtually prevented bacteria growth by freezing. The acid mine drainageformation at Faro in Yukon, Canada may have taken this route due to its geographicallocation.-6Bacterially enhanced initial acidification is by far the most common mechanism in acidmine drainage environments. This bacteria-catalyzed acidification process can be expressedby the following three-stage reaction mechanism (Kleinmann and Crerar, 1979; Kleinmann,Crerar and Pacelli, 1981):FeS2 + 7/2 02 + H2O , Fe++ + 2 SO4 + 2 H+^(1)(abiotic or by T.1)Fe++ + 5/2 H2O + 1/4 02 -4 Fe(OH)3(s) + 2 H+^(2)(abiotic)Stage one: Abiotic OxidationMechanism: Reaction (1) proceeds both abiotically and by direct bacterial oxidation.Reaction (2) proceeds abiotically, slows down when pH falls.Chemistry: pH above approximately 4.5, high sulfate, low iron, little or no acidity.FeS2 + 7/2 02 + H2O , Fe' + 2 SO4= + 2 11+^(1)(by T. f. or abiotic)Fe++ + 5/2 H2O + 1/4 02 -, Fe(OH)3(s) + 2 II+^(2)(by T. f..Stage two: Bacterially Catalyzed OxidationMechanism: Reaction (1) proceeds by direct bacterial oxidation as well as abiotically.Reaction (2) proceeds at a rate determined by the activity of T. ferrooxidans.Chemistry: Approximate pH 2.5-4.5, high sulfate, increasing acidity and total iron, lowferric/ferrous ratio.72 Fe++ + 1/2 02 + 2 11+ --. 2 Fe' + H 2O^ (3)(by T.f)FeS2 + 14 Fe' + 8 H2O , 15 Fe++ + 2 SO4 + 16 H+^(4)(abiotic)Stage three: Bacterially Catalyzed CycleMechanism: Reaction (3) proceeds at a rate totally determined by the activity of T.ferrooxidans. Reaction (4) proceeds at a rate primarily determined by therate of reaction (3).Chemistry: pH below approximately 2.5; high sulfate, acidity, total iron and ferric/ferrousratio.The three stages are defined by the role that bacteria, presumably Thiobacillusferrooxidans, play. In the first stage, the fine-grained pyrite is oxidized by atmosphericoxygen and/or by Thiobacillus ferrooxidans with the production of sulfate ions and acidity.If the pH is high enough the oxidation and hydrolysis of iron proceed. Thiobacillusferrooxidans are important in the continuation of the process, because they accelerate thedecline of pH. Each rainfall interrupts the previous acidity accumulation, increasing thepH value. Without the involvement of bacteria, the pH would be likely to stay at a valuenear neutrality if rainfalls are frequent enough. At this stage, it is possible to stabilize thepH at a neutral value by neutralization. As long as available alkalinity exceeds acidity, theonly downstream effect is an increase in SO 4 '.8If acidity exceeds alkalinity, the pH begins to fall continuously until it reaches 3.0. At thispoint, the abiotic oxidation of Fe' considerably slows down, since the reaction rate is asecond order function of OH - concentration. As the abiotic reaction (2) tends to cease,Thiobacillus ferrooxidans take over the role to oxidize Fe'. Although the overallstoichiometry remains the same, the turnover of reaction (2) from abiotic to biologicalsymbols the onset of stage two. Again, the pH of the system could be stabilized withphysical means such as sealing from oxygen, neutralization, etc., if conditions favour suchtreatments (low permeability, low surface pyrite exposure, etc.). It is difficult to bring thesystem back to stage one without appealing to bactericides.If no preventive measures are taken at this point, the pH of the pyritic system will furtherdecrease. As pH drops below 3.0, the increasing solubility of Fe results in a great increasein Fe/Fe" ratio, whereby reactions (3) and (4) are turned on, stage three begins.Once the pH reaches 2.5, which provides an optimal bacteria multiplication environment,an oxidation cycle of pyrite to SO4= and Fe' by Fe' and of Fe' to Fe by bacteriacommences and gives rise to a great amount of acid production.Kleinmann's three stage acidification model was based on laboratory studies of simulatedcoal mine environments. With the difference between coal mine environments and metalmine environments in mind, one can apply the model to a metal mine environment toachieve insight into the acid generation mechanism.9However, laboratory simulations are idealized and simplified versions of real fieldconditions. Many natural (random) variations can not be simulated in a laboratory setup.While the general framework of Kleinmann's acidification model may remain valid at realfield conditions, details will vary from one case to another. It is also possible that the realconditions are so far removed from the conditions Kleinmann simulated that the three stagemodel fails completely.Some researchers have studied or observed acid drainage generation in field conditions.Rawat and Singh reviewed some of the literature on acid mine drainage (Rawat and Singh,1983) and concluded that in coal mines: (1) organic sulfur is structurally bound and isinsignificant in contributing to acid mine drainage; (2) pyrite is the major source of acidmine drainage; (3) framboidal (raspberry-like) pyrite is the most reactive form to produceacid mine drainage; and (4) bacteria are important and Thiobacillus ferrooxidans enhancethe rate of pyrite oxidation.From a different perspective, a group of Australian researchers investigated themicroenvironment and the transport of oxygen to reaction spots within waste rock dumpsundergoing pyrite oxidation (Bennett, et. al. 1987, 1988; Harries, Hendy and Ritchie, 1987;Harries and Ritchie, 1981, 1983, 1985, 1986, 1987). They studied the relationships amongoxidation rates, temperature profiles, heat transfer, gas transport, and bacterial populationdensities in different regions and depths of two waste dumps. What they have achieved issummarized below:■^The rate of oxidation of pyrite materials depends on the rate of supply of oxygen.- 10 -■ Oxygen is transported into the dump in gas phase by diffusion, convection andadvection. The extent to which one process dominates reflects the proximity of theedge of the dump and the variability of the gas transport properties of different partsof the dump. The differences in gas transport properties reflect in turn thedifferences in the types of material excavated from the pit, differences in theweathering properties of the excavated material, and the way in which the dumpswere constructed.■ Convection is the major transport mechanism of oxygen into those regions deepwithin the dumps where there is a high rate of pyrite oxidation and therefore a hightemperature. The oxygen appears to be transported horizontally from the sides ofthe dumps, and then vertically through the hot oxidizing regions. This transportmechanism can penetrate into the dump up to 70 meters from the edge.■ Diffusion is the dominant process away from the edges of the dump and is very slow.■ Advection caused by atmospheric pressure variations occurs preferentially along highpermeability paths in some locations.■ High carbon dioxide and low oxygen levels were found within the dump which couldbe attributed to slow exchange rates between pore spaces and atmosphere, to oxygenconsumption by pyrite oxidation, and to the release of carbon dioxide by dolomiteweathering.■ The wet/dry seasonal variations of the measured physico-chemical parameters(temperature profiles and pore gas concentrations) were insignificant, indicating thatthere was little variation in pyritic oxidation rate.■ Temperature profiles can be used in a heat transfer mathematic model to find heatgeneration (i.e. intense oxidation) sources.■ There were no measurable changes in bacterial population to correspond changes inthe intensities of oxidation.Acidification process is only one of the many processes involved in the acidic drainagegeneration. For an acidic seepage to occur from a waste dump, the following processes arealso involved: the infiltration of precipitation into the dump, transport of free acidity by theinfiltrating water, on-the-route neutralization of the migrating acidic water by acid-consuming materials, leaching of mineral constituents by the migrating acidic water,chemical interactions between different species in the migrating solution (such asprecipitation of insoluble salts), and all the effects caused by the fluctuation of ground watertable and movement of ground water if these are applicable, i.e., if there is a ground watertable present within the dump. The quality of the drainage and ground water that leavethe dump is determined by all of the above processes.Comprehensive studies on the migration of acidity and dissolved species by flushing rainwater and ground water in a waste dump environment are lacking. The geochemistry ofthe weathering of various minerals and of the neutralization of acidity in water paths- 12 -within a dump is poorly understood. It is understood, however, the percolation of rainwater through waste dumps tends to follow some established channels. Probably only afraction of the acidity generated within a dump is flushed out by rain water due tochannelling. For a coal refuse dump, approximately 75% of the water intake during a rainstorm is flushed through within 24 hours; 25% is retained in the dump for a much longertime (Dugan, 1987). A general discussion of various processes involved in acid minedrainage formation can be found in Steffen Robertson and Kirsten (B.C.), et. al., 1989.The extent to which pyritic oxidation is accelerated by bacteria has been reported to varywidely, from several times to several thousands of times. This reflects the influence ofenvironmental conditions on the effectiveness with which bacteria catalyze acid generationprocesses, as well as differences in research methodology.While many researchers believe that the bacteria Thiobacillus ferrooxidans play a dominantrole in accelerating pyrite oxidation, others disagree (Muyzer et. al., 1987). Using a specialtechnique they developed, called combined immunofluorescence-DNA-fluorescence stainingtechnique which readily differentiates Thiobacillus ferrooxidans from other bacteria species,Muyzer and his fellow workers demonstrated that Thiobacillus ferrooxidans were not ableto multiply due to the competition of other bacteria in an unsterilized pyritic coal slurryinoculated with both a pure culture of Thiobacillus ferrooxidans and a mixed culturecontaining a variety of bacteria species. Thiobacillus ferrooxidans were only able to growsignificantly when competing bacteria were absent, i.e., on a sterilized coal slurry inoculatedwith pure Thiobacillus ferrooxidans only. In addition, the total iron leached from pyrite bypure Thiobacillus ferrooxidans only was much less than that by mixed cultures. Muyzer- 13 -suggested that many "pure cultures" used by previous investigators were probablycontaminated with other bacteria species, thus leaching effects caused by a mixedpopulation of a variety of bacteria may have been incorrectly attributed to Thiobacillusferrooxidans.Muyzer's argument was also supported by other experiments. Harrison examined 23strains of Thiobacillus ferrooxidans culture in 1982 and found that one quarter werecontaminated with heterotrophic bacteria (Harrison, 1982). Lobos et. al. isolatedAcidiphilium organovorum from a culture of Thiobacillus ferrooxidans which had beengrown autotrophically on FeSO4-basal salts medium for several years (Lobos et. al., 1986).As is understood now, real pure culture of Thiobacillus ferrooxidans is difficult to obtainand several species, both autotrophic and heterotrophic, or even mixotrophic are quiteresistant to traditional purification techniques.Muyzer's experiment probably represents one extreme while assuming complete dominationof Thiobacillus ferrooxidans the other. The real bacterial population composition probablylies somewhere between these extremes and varies from place to place according toenvironmental conditions. Evaluation of bacterial population composition in a typical metalmine acidic waste rock environment is a rather difficult task and yet to be carried out.In a more general sense, the B.C. AMD Task Force Draft Technical Guide Volume I (SteffenRobertson and Kirsten et. al., 1989) describes the major chemical factors which determinethe rate of acid generation in the absence of significant bacterial enhancement as follows:a) pH,- 14 -b) Temperature,c) Oxygen content of the pore gas phase if saturation is less than 100%,d) Oxygen concentration in the water phase,e) Degree of saturation of dumps (or tailings) with water,0 Chemical activity of Fe",g) Surface area of exposed metal sulfides,h) Chemical activation energy required to initiate acid generation.Strictly speaking, these factors determine the rate of sulfide oxidation rather than acidgeneration. Acid generation involves, in addition to oxidation, the neutralization of acidityby acid-consuming materials and hydrolysis of dissolved species in the solution phase.Therefore, the rate of acid generation from a dump is determined by the rate of oxidationas well as the rate at which acidity is consumed by neutralization or generated byhydrolysis. For example, if an originally acidic solution from sulfide oxidation contains highconcentration of Fe' and Al', upon neutralization in a different region of the dump throughwhich the water is percolating, Fe 3+ and Al' will be precipitated as Fe(OH)3 (ferrichydroxide) and Al(OH)3 (gibbsite) by hydrolysis, contributing acidity to the solution:Fe' + 3 H2O 4=2 Fe(OH)3 (s) + 3 H+Al3+ + 3 H2O4-----■ Al(OH)3(s) + 3 H+Among the listed factors that determine the rate of sulfide oxidation, factor d), oxygenconcentration in the water phase, only operates when some part of the dump or tailings is- 15 -immersed in a water table. If a dump is completely above water table, factor d) hasnegligible effect on the rate of oxidation.Factor 1), the chemical activity of Fe", or more straightly the concentration or activity ofdissolved Fe', rarely amounts to such an extent that it has a significant effect in the rateof oxidation in nature, if bacterial involvement is negligible.Factor g), the surface area of exposed metal sulfide can be broken down to a few sub-factors:the size distribution of sulfide materials, the interlocking of sulfides with inert minerals,the dissemination of sulfide minerals, and the forms of existence of sulfide minerals(euhedreal, amorphous, aggregative, etc.). Some of these may not have an immediate effecton the initial total surface area of sulfide minerals (which depends largely on the rock sizedistribution and exposure of sulfide on broken rock surfaces), they will determine the totalsurface area of exposed sulfides when rocks disintegrate through weathering.Activation energy of a chemical reaction refers to the energy barrier that the reactants mustovercome in order to be transformed into products. It measures the ease with which thechemical reaction can proceed. Since the initial acidification process is a quite complicatedone which involves many possible reaction mechanisms rather than a simple chemicalreaction, the activation energy (in a strict chemical sense) of a dump or tailings system isnot defined. If factor h), chemical activation energy required to initiate acid generation(oxidation), is used in a broader sense to represent the ease with which the initialacidification process proceeds, it should cover the following factors (see above discussion ofKleinmann's three-stage model of acidification): the total annual precipitation, the- 16 -frequency of precipitation, the total sulfide sulfur and its total exposed surface area, thekinds of metal sulfides present, the form of existence of sulfides, the availability of acidconsuming materials and their properties, etc. It can be seen that some of these factors arealready represented in the listed factors.There is a very important factor that affects the rate of oxidation in a waste dump ortailings but that is not listed: the time that has past since the initial deposition. Thisaffects the rate of oxidation in two ways: First, time is required for the acidity to build up.When acidity builds up, bacterial oxidation accelerates. This was demonstrated at IslandCopper Mine by the fact that the EMO, the Old North Dump and the North West Dump alltook a few years to acidify. Second, as time passes rocks are weathered, disintegrated toexpose more sulfide minerals.It is believed that bacteria are indigenous to sulfide waste environments. In most cases,bacteria will become involved sooner or later, even though their role in the early stage isinsignificant. If bacteria are involved significantly to accelerate the oxidation process, thefollowing factors, listed in the B.C. AMD Draft Technical Guide, also act to determine theoverall rate of oxidation in addition to the factors that determine the abiotic oxidation listedpreviously.a) Biological activation energy,b) Population density of bacteria,c) Rate of population growth,d) Nitrate concentration,- 17 -e) Ammonia concentration,f) Phosphorus concentration,g) Carbon dioxide content,h) Concentrations of any bacterial inhibitors.Among these factors, d), e) and f) reflect the nutrient requirements of the autotrophicbacteria. Although the growth of bacteria require other nutrients (such as some metalelements that are used in special enzymes), the amounts needed are so small that they veryrarely become limiting factors of growth. We can add the following factors to the above listto make it more complete:i) Interaction of autotrophic bacteria with heterotrophic bacteria, molds and fungi,for example, the competition for growth-limiting nutrients,j) The bacterial composition of the microflora.Finally, the total acidity and pollutants carried away by seepage and ground water dependon, in addition to the factors mentioned above which determine the chemical and biologicaloxidation rate, the following factors:a) The neutralization capacity of the waste materials,b) The tortuosity of the waste dump or tailings,c) The types and mineral composition of the waste rock,d) The solution chemistry of the percolating water.- 18 -2.2^Thiobacillus ferrooxidansDue to the involvement of Thiobacillus ferrooxidans in acid mine drainage formation andtheir utilization in bioleaching of valuable metals from low-grade ores, Thiobacillusferrooxidans have been extensively studied. Their morphology, metabolism and culturingtechniques are quite adequately established. Some of their genetic nucleotide sequenceshave been decoded. Many strains of different characteristics have been isolated. Thissection briefly summarizes what we currently know about this bacteria (Blancarte-Zuritaet. al., 1986; Blancarte-Zurita, 1988; Clum, et. al., 1985; Collis, 1969; Cooper et. al. 1987;Gloxhuber, 1980; Hendy, 1987; Kleinmann and Crerar, 1979; Kleinmann, Crerar andPacelli, 1981; Lobos et. al., 1986; Rawat et. al., 1983; Steffen Robertson and Kirsten et. al.,1989).Thiobacillus ferrooxidans are rod-shaped (bacillus means "rod"), gram-negative, non-sporeforming, aerobic (oxygen-dependent), motile bacteria. They measure 1.2-3.2 micrometersin length and 0.5-1.25 micrometers in diameter. They are acidophilic bacteria. They growoptimally in the pH range 1.5-4.0 and can survive in the pH range 0.5-7.0.Thiobacillus ferrooxidans are mesophilic (medium temperature-loving) bacteria. Theymultiply most readily in the temperature range 20-35°C and can survive 15-40°C. However,strains that are able to grow at a temperature as low as 6°C has been reported (Ferroni,1986). Harries et. al. reported a thermophilic bacteria species which was isolated from awaste dump undergoing intense oxidation (Harries, Hendy and Ritchie, 1987). The bacteria- 19 -had a optimum growth rate at pH 1.4 between 50-55°C and was able to oxidize iron.Whether the bacteria were Thiobacillus ferrooxidans was unclear.Thiobacillus ferrooxidans derive energy for metabolism from the oxidation of sulfide tosulfate and of ferrous iron to ferric iron (hence Thio- meaning "sulfur" and ferrooxidansmeaning "oxidizing iron"). They are also called chemoautotrophic (meaning "feed oninorganic carbon and derive energy from chemical reactions"), lithoautotrophic (meaning"derive energy from rock"), or chemolithotrophic. When the above mechanisms are notavailable, they are also able to oxidize other reduced forms of sulfur (such as thiosulfate)to sulfate to extract energy. They have been reported to use copper as an energy source(Nielsen and Beck, 1972; Tuovinen and Kelly, 1972; Ehrlich, 1978; Tuovinen et. al., 1978;all as cited by Blancarte-Zurita, 1988).Thiobacillus ferrooxidans obtain carbon by fixing atmospheric carbon via the Calvin cycle.So they are autotrophic (meaning "feed on inorganic carbon"). Thiobacillus ferrooxidanseither extract nitrogen from an inorganic source such as ammonium salts or nitrates orfrom fixation of atmospheric nitrogen (Blancarte-Zurita, 1988; Pretorius et. al., 1987).The cell composition of Thiobacillus ferrooxidans has been reported to be 44% protein, 26%lipid, 15% carbohydrate and 10% ash. The cell envelope consists of three zones: acytoplasmic membrane which constitutes the inner layer of the envelope bordering thecytoplasm, a middle zone that consists of a rigid layer of peptidoglycan and the periplasmicspace, and an outer layer which contains lipopolysaccharide and lipoprotein. Elemental- 20 -analysis of the dried microorganism was 47.6% carbon, 10.1% nitrogen and 7.3% hydrogen(Blancarte-Zurita, 1988).Thiobacillus ferrooxidans are usually cultivated in the 9K medium established by Silvermanand Lundgren in 1959, or a modification of it. Heterotrophic bacteria are supposedlyeliminated by serial culturing of Thiobacillus ferrooxidans in an organic carbon-free mediasuch as 9K media. Other autotrophic bacteria that oxidize reduced forms of sulfur but notferrous iron can be eliminated by growing Thiobacillus ferrooxidans culture on a mediumthat contains ferrous iron as the only energy source.In addition to numerous antibiotics, Thiobacillus ferrooxidans are inhibited by simpleorganic acids such as malonic acid, itaconic acid, citric acid (Tuttle et. al., 1976), benzoicacid, sorbic acid (Onysko et. al., 1984); by some surfactants such as sodium lauryl sulfate(SLS), akylbenzene sulfonate (ABS) and alpha olefin sulfonate (AOS) (Kleinmann et. al.,1982, 1983; Onysko et. al., 1984); by certain metals such as mercury (Balch et. al., 1987),molybdenum and tellurium (Kleinmann et. al., 1983), and by chlorides.Tuttle et. al. examined the inhibition of bacterial oxidation of iron and sulfur by a series oforganic acids (Tuttle et. al., 1976). They concluded that:■^Any organic substance which has an ability to interact abiotically with ferrous ironcan be an apparent inhibitor of iron oxidation;- 21 -■ The relative electronegativity of the inhibitor molecules appears to correlate with theinhibition effectiveness;■ A portion of inhibitive activities of organic acids can be attributed to the directinhibition of the iron-oxidizing system;■ At a concentration of about 2 - 4x10-5 M, malonic, itaconic, dihydrowyfumaric andhydroxypyruvic acids are inhibitory to the autotrophic growth. 10-fold greaterconcentrations of citric, transaconitic, isocitic, alph-keto-glictaric, malic and succinicacids have similar effect;■ Bacteria can accumulate organic substances which adversely affect the growth of thebacteria themselves.Kleinmann et. al. examined in laboratory and field-tested the bactericidal effect of SLS,ABS and AOS and found that the most effective biocidal agent was SLS.Microorganisms' tolerance to metals varies from strain to strain. It is possible to selectivelycultivate stains of bacteria that are highly tolerant to certain metals (Clum et. al., 1985)for use in bioleaching.Estimation of Thiobacillus ferrooxidans population density can be performed by directmethods or indirect methods. Direct methods include Most Probable Number test (MPN),direct plate count, microscopic count with staining technique, etc. Indirect methods include- 22 -measurements of cell component such as total cell carbon weight and total cell proteinweight, weight of metabolic products such as ATP (adenosine triphosphate), or consumptionof growth essentials such as carbon dioxide or oxygen.Bacteria can be killed in two ways: cell disintegration and/or irreversible damage to thegenome (Mason et. al., 1986). The mechanism by which surfactants inactivate or killThiobacillus ferrooxidans can be depicted as follows: Although Thiobacillus ferrooxidansare an acidophilic chemoautotroph, their internal pH is approximately neutral. This isattained through the cellular control of the semi-permeability of the cell wall as well as thecytoplasmic membrane. Anionic surfactants, which alter the semi-permeability of thecytoplasmic membrane, result in the seepage of W into the cell. The change in internal pHinactivates the pH-sensitive enzymes in the cell, which in turn causes a slowdown of theFe++ oxidation catalyzed by the enzymes, thus induces bacteriostatic effects. Highconcentrations of surfactants kill the bacteria by permanently damaging the enzymes andby disintegrating the membrane material and the cell walls.Thiobacillus ferrooxidans are very adhesive to solids such as rock particles (Bagdigian,1986). Most of the bacteria population are probably resident on the solid surfaces ratherthan in liquid phase in a pyritic coal slurry. It is very difficult to detach themicroorganisms from solid surfaces even with the help of such physical means as vigorousshaking and ultrasonic washing.2.3^Literature on Island Copper MineThere is abundant literature on the Island Copper Mine. Young described the explorationand geology of the Island Copper Mine property (Young and Rugg, 1971). Evans et. al.documented the establishment of the monitoring program for the underwater discharge oftailings into Rupert Inlet (Evans et. al., 1972, 1979). Drinkwater examined the patternsof tidal mixing in Rupert and Holberg Inlets and their environmental implications(Drinkwater, 1973). Hay carried out a comprehensive study of the underwater transportof tailings, the formation and evolution of submarine channels and the turbidity currentsin Rupert Inlet using acoustic sensing technique (Hay, 1981). Waldichuk et. al. gave acritical review of the environmental significance of the underwater tailings disposal atIsland Copper (Waldichuk and Buchanan, 1980). A comprehensive list of pre-1980literature entries can be found in Waldichuk et. al., 1980.However, the above literature is mostly concentrated on various aspects of the underwaterdisposal of tailings at Island Copper Mine, since the acid mine drainage problem was notrealized until 1985. The acid mine drainage problem was addressed in the report by BHP-Utah and Rescan, 1988, which presented the study of the South Wall Pushback Project andrelated environmental issues.3.0 RESEARCH SITE3.1^General DescriptionIsland Copper Mine, owned and operated by BHP-Utah Mines Ltd., is one of Canada'slargest open pit, low-grade copper mining operations. It is located on the north shore ofRupert Inlet, about 16 kilometres south of Port Hardy at the north end of Vancouver Island(Figure 1). The revenue-producing metals are copper, molybdenum, gold, silver andrhenium.The operation at Island Copper Mine consists of open pit mining by truck and shovelmethod in conjunction with the use of an in-pit crushing/conveying system, and milling inits concentrator by a conventional grinding/flotation circuit.At Island Copper Mine, waste rock is disposed of in waste dumps at various locations nearthe pit, both on land and into the sea. As of January, 1989, four waste dumps have beendeveloped on land since the commencement of operation: the North Dump (84.2 x 10 6 tons),the North West Dump (1.1 x 106 tons), the West Dump (6.6 x 10 6 tons) and the South Dump(6.6 x 106 tons). The majority of the waste rock (510 x 10 6 tons) has been placed on theBeach Dump along the north shoreline of, and extending into Rupert Inlet.- 25 -12 ° W 126°W I 24 ° We*040^Port H or d yAirportISLAND COPPERMINE— 5 I ° N NPOW ELLRIVER50 ° N":"VANCOUVER0CANADAU.S.A•— 49 °2FIGURE 1 LOCATION MAP0^10 20 40MILES0^20^40KILOMETERSAlthough assessment of acid mine drainage problems has today become an integral part ofthe environmental assessment required by government regulatory agencies, the AMDproblem was not widely recognized at the time of design and construction of the IslandCopper Mine. In fact, the waste rock at ICM was all expected to be non-acid generating.Little detailed consideration was given to the aspect of AMD prevention and handlingduring the early stage of waste deposition.The generation of AMD was not fully appreciated at Island Copper until 1982. Then thewest pit wall, after being exposed by mining activity and left undisturbed, showedacidification with pH values as low as 3.5; and visible yellow stains (Fe(OH) 3) startedstringing out on the wall. In December 1985, the seeps from the Old Marginal Dump at thenorth western corner of the North Dump - the oldest land dump - began to show asignificantly elevated level of zinc, indicating the oxidation of waste material in the dump.Concurrent with the increasing interest in AMD by industry and academia, the advance ofour understanding of acid generation process in nature and the development of aciddrainage at several metal mines in B.C., and prompted by the indications of potential acidproduction problems, Island Copper expanded its monitoring program of land dumpdrainage with a purpose of pin-pointing acid-generating spots. This resulted in the findingof a few active areas, the most serious of which was the Eastern Most Outslope, followedby the Old North Dump.3.2^GeologyThe Island Copper deposit is underlain by a thick section of Bonanza Volcanic Formation,which strikes about N60°-70°W and dips 30 0-40° to the south. The principal ore control isa dyke-like quartz-feldspar porphyry intrusion, Jurassic in age, which has invaded theTriassic Bonanza tuffs. The porphyry dyke is reversed-U shaped, strikes N70°W over 6,000feet, dips 50°I•TNE, and is usually 400-500 feet thick, although its maximum thicknessreaches 900 feet. Ore is concentrated along both the foot wall and the hanging wall of thedyke structure, but occasionally occurs in some places entirely across the dyke.The weakly mineralized porphyry dyke separates the ore body into two limbs, which areconnected on the east and open on the west. The ore zone averages about 1,100 feet inthickness from the hanging wall to the foot wall boundaries, and extends approximately3,000 feet along the dyke, with the north limb extending 1,000 feet further west than thesouth limb.Copper and molybdenum mineralization occurs in all rock types but is found primarily inthe fractured and silicified andesite as fillings and smears. About one half of the ore grademineralization is found in the biotite and transition zone andesite, and the remainder inthe crackle zone, the rotational breccia and the dyke.The mineralization in the deposit consists of pyrite, magnetite, chalcopyrite, some borniteand molybdenite in order of decreasing percentages. Pyrite is found in all types of rockspresent in finely disseminated forms and also as thin and randomly oriented seams. The- 28 -pyrite content of the ore body generally varies from two to five percent but there are localconcentrations, primarily fracture-controlled, of 15 to 20 percent.Alterations of several types have affected the volcanic host rocks. In order of decreasingintensity they are: 1) silicification 2) argillization 3) saussuritization and 4) biotitization.In addition, some potassium feldspar has been introduced and carbonate, talc, chlorite andpyrophyllite have been developed in varying amount. Silicification is generally pervasivethrough the mineralized zone but is most intense in the brecciated volcanics along the dykemargins. The clay alteration is quite closely associated with the silicification and is mostevident in the fault zones. Saussuritic alteration is quite pervasive and may be largelydeuteric. The potassium feldspar alteration is closely associated with silicification and isparticularly intense near the porphyry-dyke contacts. Some carbonate is present infractures. Talc and chlorite are the common alteration products of the mafic elements inthe andesite and are probably of both deuteric and hydrothermal origins.The Island Copper deposit is in general covered with a glacial till layer with a depthranging from a few feet to as deep as 250 feet.3.3 The North/Old North DumpThe North Dump is the largest of the four dumps developed on land at the Island CopperMine (Figure 2). The initial deposition of waste on the North dump began in April, 1971and the dump was completed in August, 1985. During the fifteen years, 84.2 million tonsof waste rock and till were disposed of in an area of 140 hectares.- 29 -Figure 2^Location of the North Dumpat Island Copper Mine•••••■•0r0•••••••••4Several regions of the North Dump possess their own common characteristics and weredesignated by different names. They are the Caps, the Eastern Most Outslope (EMO), theOld Marginal Dump and the Old North Dump (Figure 3). The physical parameters of theseregions are given in Table 1.Table 1^Physical Parameters of Designated Regions on North DumptNorth Dump The Caps EMO OldMarginalP.N.D.E.C.E.*Total Mass(milliontonnes)76.4 9.12 4.67 1.28 62.6% of Till 36.5 0.0 2.42 65.4 44.4Area (Ha.) 140.0 40.0 15.2 N/A 84.8Total Volume(x106 m3)38.9 4.7 2.5 0.65** 31.6Average Bulk 1965 1928 1840 >1981 1981Density(kg/m3)Average 21.0 27.3 30.3 <19.0 19.0Porosity(%)Figures in this table are taken from Acid Mine Drainage Study of the North Dump -Final Report, University of British Columbia, Department of Mining and MineralProcess Engineering, 1991.* The Part of the North Dump Excluding the Caps and EMO.** Calculated using a bulk density of 1981 kg/m3.NOTE: Coordinate scale follews that of ICMFigures are in feetFigure 3^Regions of North Dump and Lacations of Drill Holes3.4 HydrologyThe North Dump area at ICM is located mainly within the historic End Creek watershed.A small part of the North Dump intrudes Stephens Creek watershed on the northwest andanother small part lies within the historic Trey Creek watershed. Figure 4 is a pre-miningsurface topographical map, on which surface water courses and watershed boundaries aredrawn according to the topography. The current pit boundary, the Beach Dump and theNorth Dump are superimposed on the map to illustrate their locations relative to thehistoric watersheds. Before mining, rain water received by what is now the North Dumparea reported to three locations: 1) the East Twin Lake, which was connected to the WestTwin Lake, which fed to the Bay Lake, which was the headwaters of Stephens Creek; 2)End Creek, which collected the majority of the water from the area and reported to RupertInlet; and 3) Trey Creek and other small streams which also flowed into Rupert Inlet.The development of the North Dump has significantly altered the surface hydrology of thearea in two ways: The deposition of waste materials has redistributed the rain water andthe drainage management has completely changed the natural water courses. The rainwater redistribution by the existence of waste materials depends on several factors, suchas the topography of the waste dump, the porosity distribution of the dump, presence ofcompacted layers and the size of the rain storm.- 34 -The drainage management system has evolved over the past years. At the time of writing(1990), all drainage running off or flowing through the North Dump, except that lost toground water and a small subterranean flow which leaks some drainage from NDD to theTwin Lakes, is collected by a ditch system around the dump, treated by mixing with the pitdewatering water and released to Rupert Inlet through the Beach Dump. This ditch systemconsists of NDD, TED, TCR, EMO, EDL and EDD (Figure 5).NDD, which flows into TED, intercepts the flow emanating from the north western part ofthe North Dump and previously reporting to Twin Lakes. TED intercepts the Old EndCreek and other subterranean flows and surface runoffs within the End Creek watershed.All the old water courses within the North Dump area have been redirected to thetreatment pond on the Beach Dump. It is also noted that the part of the End Creekwatershed to the north of the North Dump contributes some water to the North Dump,mostly as subterranean flows through the North Dump.3.5 Land Dump MonitoringThe Island Copper Mine Environmental Department has operated a comprehensive freshwater monitoring program since the pre-operational phase of the mine. The monitoringscheme generally satisfied our research requirements. Some modifications to themonitoring scheme were made corresponding to the research needs. During the past years,ICM regularly sampled the streams listed in Table 2 for water quality analysis. Locationsof the sampling stations are shown in Figure 5.- 35 -a/eUP.ERTFigure 5 ICM North Dump Drainage Collection Ditch System andFresh Water Sampling Stations- 36 -NOD Pb1--5.••••••P7/NE L51 1-'147TZ,2 SAMPL.M6 STATIONSL.SENOa ,/4 -II, .WArriAr DArwv..40.14"X WE/k Pi-0)^4f ,T1CUNI(E-A/T WA'rc■T COL 0?.. gS_-„- -•^>11.47 Eif CoBHP - Utah Mines Ltd.ISLAND COPPER MINEp... IV.... LCTable 2^Island Copper Regular Fresh Water Sampling StationsStation Code^Station Name^Sampling FrequencyNDD^North Drainage Ditch^MonthlyTED Ten Eighty Ditch MonthlyEDD^East Drainage Ditch^MonthlyPDW Pit Dewatering MonthlyBLA^Bay (Francis) Lake^MonthlyTLD Twin Lakes Discharge MonthlySCR^Stephens Creek MonthlyJLD Joe Lake Drainage^MonthlyTCR^Old Trey Creek MonthlyEDL East Drainage Lake^MonthlyEMO^Eastern Most Outslope MonthlyPDP Pit Dewatering Pond^MonthlyV-notched weir flowmeters have been installed on six streams that flow through the NorthDump area. The six stations are: NDD, EDL, EMO, EDD, TED, TCR. Instantaneous flowrates are read on the flowmeters. Flow rate data are however sparse up to the time ofwriting because of the short time of data accumulation and irregularly-scheduledmeasurements. Two in situ measurements - flowrate and conductivity - were taken weeklyon these streams. Frequency was increased during the wet seasons, especially during majorrain storm events.4.0 RESEARCH METHODS4.1^Data CollectionData collected at the research site that are relevant to this study include:■ All available water quality analyses of streams around the mine site,■ Pre-operational baseline monitoring records of fresh water streams,■ All available flowrate measurements of drainage ditches and pit dewatering,■ Acid-base accounting of land dump drill hole samples,■ Mine site and Port Hardy meteorological records,■ Mining records, including the North Dump development records, month-to-monthbenchwise mining operation records, truck haulage records, etc.,■ Aerial photographs of the pit and dump areas,■ Land dump drainage management history,■ Settling pond water quality,■ Various maps, including location maps, topographical (contour) maps and surveymaps, all drawn to scale,■ South Wall Pushback project study report.- 38 -4.2^Drainage Sample CollectionDrainage sample collection, measurement of physical parameters and chemical analysiswere performed by qualified personnel from ICM environmental laboratory. Temperatureand pH were measured in situ at each fresh water monitoring station. A one-litre samplewas then taken directly from each stream into an acid-washed polyethylene bottle. Sampleswere returned to the laboratory for immediate filtration through 0.45 ilm membrane filters.The filtrates were stored prior to analysis in a second set of acid-washed polyethylenebottles and preserved by the addition of 7 ml of concentrated nitric acid. Separate sampleswere collected for the analyses of arsenic and mercury, although these were treated in thesame manner.4.3 Methods of Chemical Analysis and Physical MeasurementTotal dissolved solids was measured gravimetrically.^Sulfate was determinedturbidimetrically. Total alkalinity was determined using a potentiometric titration.Calcium and magnesium were measured by flame atomic adsorption spectroscopy (AAS).Dissolved cadmium, copper, iron, lead, manganese, molybdenum and zinc were analyzed byflame atomic absorption spectrometry (AAS) after digestion and concentration withHNO3 :HC1O 4 :HF. Total and dissolved arsenic were determined by reduction to As(III) usingKI:HC1 and analyzed using hydride generation AAS. All analytical methods were based onthose described in APHA, 1985 and were documented in detail in the ICM EnvironmentalDepartment Methods Manual of 1986.- 39 -5.0 RESULTS AND DISCUSSION5.1 Acid Generation Assessment of the North Dump5.1.1 Acid-Base AccountingA drilling program was conducted on the North Dump to sample the dump material. Sevenwells were percussion-drilled with a Becker drill. Drill cuttings were sampled in an eight-foot interval and submitted for acid-base accounting. The locations of the drill wells areillustrated in Figure 3 and the acid-base accounting results are tabulated in Table Flowrate AnalysisTable 4 presents all the available flowrate data on fresh water streams. These data weremeasured by weirs except those of PDW which were pit dewatering pumping records. Ananalysis was performed on the data and the results are presented in Table 5 and Figure 6.Table 3^Acid-Base Accounting Results of North Dump Drilling SamplesDepth(ft)Hole AMD #1 AMD #2APP^ACP NNP APP ACP NNP0 - 6 38.8^17.3 -21.5 14.7 8.1 -6.66^-14 32.3^16.3 -16.0 15.4 9.3 -6.114-22 46.4^21.7 -24.7 86.7 14.0 -72.722-30 51.3^14.5 -36.8 93.6 9.1 -84.530 - 38 36.7^14.7 -22.0 83.1 14.0 -69.138 -46 55.6^20.6 -35.0 89.8 6.0 -83.846-54 53.3^17.5 -35.854-62 53.4^22.7 -30.762-68 40.6^11.2 -29.4Average 45.4^17.4 -28.0 63.9 10.1 -53.8AMD #3 AMD #4Depth(ft) APP^ACP NNP APP ACP NNP0 - 6 62.8^17.3 -45.5 32.9* 28.3 * -4.6*6^-14 41.3^32.8 -8.5 30.2 26.4 -3.814-22 25.9^57.5 31.6 17.0 8.1 -8.922-30 50.1^30.0 -20.1 7.2 4.4 -2.830 - 38 4.8 3.4 -1.438-46 14.3 11.4 -2.946-54 21.5 9.5 -12.054-62 12.6 12.1 -0.562-70 18.6 23.6 5.070-78 16.1 28.6 12.578 - 86 12.9 24.6 11.786 - 94 15.1 28.1 13.094-102 36.2 28.1 -8.1Average 45.0^34.4 -10.6 18.4 18.2 -0.2* Sample was from 3 - 5 feet.(Continued fromDepth(ft)last page)AMD #5 AMD #6APP ACP NNP APP ACP NNP0 - 6 24.4 31.4 7.0 35.3 33.8 -1.56 -14 32.3 27.6 -4.7 43.0 48.2 5.214-22 57.0 20.6 -36.4 20.4 44.6 24.222 - 30 34.4 27.5 -6.9 28.6 41.2 12.630 - 38 17.0 22.9 5.9 62.7 41.9 -20.838-46 21.5 30.1 8.6 26.9 56.9 30.046 -54 21.9 54.0 32.154-62 16.9 49.3 32.462 - 70 20.1 37.7 17.670-78 - -78 - 86 8.2 19.4 11.286-94 3.9 22.0 18.194-102 3.9 34.6 30.7102-110 2.2 13.3 11.1Average 31.1 26.7 -4.4 22.6 38.2 15.6AMD #7 AMD #8Depth(ft) APP ACP NNP APP ACP NNP0 - 6 33.7 37.0 3.3 65.8 37.2 -28.66 -14 45.0 36.4 -8.6 - - -14-22 49.6 41.2 -8.4 43.4 42.2 -1.222 - 30 64.1 43.5 -20.6 22.0 83.4 61.430-38 72.6 51.3 -21.3 55.6 40.4 -15.238-46 35.8 62.7 26.9 61.3 31.2 -30.146 - 54 65.2 32.5 -32.7 46.4 44.6 -1.854-62 41.6 52.7 11.1 53.6 37.5 -16.162-70 35.5 67.1 31.6 58.2 38.8 -19.470-78 18.8 55.8 37.0 54.4 29.9 -24.578-86 26.9 45.9 19.0 55.8 34.6 -21.286 - 94 26.4 42.5 16.1 31.5 23.3 -8.294-102 - 43.8 27.0 -16.8102-110 33.5 53.7 20.2 76.3 18.6 -57.7110-118 23.8 45.1 21.3 26.2 36.7 10.5118-126 22.9 25.0 2.1 51.0 38.0 -13.0Average 39.7 46.2 6.5 49.7 37.6 -12.1(ContinuedDepth(ft)from last page)AMD #9 AMD #10APP ACP NNP APP ACP NNP0 - 6 48.6 21.4 -27.2 67.4 53.1 - 14.36 -14 64.8 28.5 -36.3 60.9 55.4 -5.514-22 63.2 18.7 -44.5 62.2 57.7 -4.522-30 35.9 - 77.3 57.7 -19.630-38 33.1 - 68.5 47.9 -20.638 - 46 45.6 8.0 - 37.6 101.9 26.7 -75.246-54 52.5 19.9 - 32.6 97.5 71.2 -27.354-62 65.8 15.9 -49.9 101.5 58.7 -42.862-70 93.4 71.4 -22.0 72.0 58.7 -13.370-78 87.0 68.3 -18.7 68.4 70.8 2.478-86 86.4 79.0 -7.4 43.5 42.4 -1.186-94 62.1 20.5 -41.6 71.2 50.7 -20.594-102 99.5 64.3 -35.2 70.8 59.1 -11.7102-110 74.7 38.6 -36.1 68.1 16.1 -52.0110-118 60.6 24.8 - 35.8 66.2 18.7 - 47.5118-126 85.6 45.9 -39.7 56.1 16.3 - 39.8126-134 90.4 88.5 -1.9 50.1 16.1 -34.0134-142 99.1 69.4 -29.7 77.3 17.5 -59.8142-150 86.3 71.2 -15.1 15.4 -Average 67.2 44.3 -22.9 71.2 42.6 -28.6Notes: APP - Acid Producing Potential, kg H 2SO 4/tonneof Waste Material.ACP - Acid Consuming Potential, kg CaCO 3 /tonne.Note that 1 kg H 2SO 4/tonne = 1.02 kg CaCO 3 /tonne.NNP - Net Neutralizing Potential, Unit Same as ACP.Table 4^Volumetric Flow Rate Monitoring Data of Some Drainage StreamsCOLUMN NUMBER 1I2I3I4 5 6 7i8YEAR^DATE PDW(PUMP NDD TED EST. EMO EDL EDD TCR PDPRECORD) WEIRS OR WEIRS WEIRS WEIRS WEIRS WEIRS CALC.1987^NOV.^2 N/A 1100 7200 654 926 1600 N/A N/A9 11600 760 6300 300 900 1200 N/A N/A13 21000 3800 25400 2534 5723 8300 N/A N/A20 22000 4100 N/A 2726 4088 6800 N/A N/A30 29000 2500 16300 397 899 1300 N/A N/ADEC.^11 N/A 3500 N/A 1963 3270 5200 N/A N/A14 35700 3500 23600 2180 3542 5700 N/A N/A21 N/A 1700 N/A 626 1470 2100 N/A N/A28 14400 1700 11300 354 926 1280 N/A N/A1988^JAN.^4 9000 900 6000 190 436 660 N/A N/A14 17600 1800 11500 1771 2998 4700 N/A N/A26 17800 1900 N/A 409 926 1300 N/A N/AFEB.^1 11500 1100 7265 305 626 930 N/A N/A8 6600 600 N/A 163 495 660 N/A N/A22 32500 1900 N/A 354 763 1100 N/A N/A23 N/A 1700 N/A 300 626 900 N/A N/A29 11400 1100 N/A 131 300 430 N/A N/AMAR.^4 8500 N/A N/A 223 626 850 N/A N/A7 N/A 2200 14700 1444 2235 3700 N/A N/A10 19700 1700 N/A 229 393 620 N/A N/A24 24700 1800 N/A 899 1689 2600 N/A N/AAPR.^4 19600 2500 16900 N/A N/A 5000 N/A N/A13 27800 1500 N/A N/A N/A 770 N/A N/A18 8300 900 6200 N/A N/A 360 N/A N/A29 6500 400 N/A N/A N/A 700 N/A N/AMAY^9 6100 400 2700 N/A N/A N/A N/A N/A25 4200 230 N/A N/A N/A 200 N/A N/AJUNE^6 N/A 400 2700 N/A N/A 690 N/A N/A13 N/A 230 1900 N/A N/A 260 N/A N/A20 N/A 170 1300 N/A N/A 120 N/A N/AJULY^11 N/A 40 550 N/A N/A 60 N/A N/A19 N/A N/A 230 N/A N/A 80 N/A N/A27 N/A 40 140 N/A N/A 50 N/A N/AAUG.^2 2100 20 130 N/A N/A 20 N/A N/A15 2200 N/A 280 N/A N/A 42 N/A N/A22 2500 N/A 280 N/A N/A 78 N/A N/ASEPT.^1 2600 14 200 7 30 37 N/A N/A6 2500 5 200 7 33 40 N/A N/A12 2300 5 160 5 33 38 N/A N/A19 2900 5 250 5 76 82 N/A N/A27 9400 5 280 7 109 116 N/A N/A28 9400 1254 10100 N/A N/A N/A N/A N/A29 9400 572 6200 N/A N/A N/A N/A N/AOCT.^4 7400 343 2700 N/A N/A N/A N/A N/A11 5100 131 1600 N/A N/A N/A N/A N/A17 9300 234 1900 N/A N/A N/A N/A N/A21 9300 627 4500 354 736 1090 570 1546023 22900 4252 16000 4360 8722 13100 3500 5550024 22900 2180 8300 1254 2507 3760 1600 3656025 22900 1962 10900 627 N/A N/A N/A N/A26 22900 1670 7600 463 N/A N/A N/A N/A31 23800 981 7600 207 927 1100 900 33400NOV.^3 23800 2889 15800 2235 4088 6300 3500 494007 N/A 2235 12800 981 2235 3200 3200 N/A14 N/A 2235 14700 899 1908 2800 3200 N/A20 N/A 3216 16000 2453 3816 6300 5500 N/A28 N/A 2180 15800 899 2998 3900 3200 N/A- 44 -Table 5^Summary of Flowrate Data AnalysesDescription PDWTEDPDPNDD OtherTCREMO EDLAnnual:Volume (m3) 4,560,000 2,598,000 365,000 534,000 8,056,000392,000 2,206,000 221,000 445,000% of TED - 100.0 - -14.4 85.6 - -%ofEDD - - 100.0 -- 33.2 67.0% of PDP 56.6 32.3 4.5 6.6 100.04.7 27.6 22 4.4Wet Season:Volume (ti 3) 3,876,000 2,445,000 345,000 508,000 7,185,000- 218,000 424,000% of TED 100.0% of EDD - 100.1 -33.1 67.0% of PDP 54.0 34.2 4.8 7.1 100.0- 2.4 4.7% of Wet in WholeYear, Same Stream85.0 94.5 94.5 95.3 89.2- - 98.7 95.3Dry Season:Volume (m3) 684,000 143,000 20,000 25,000 872,000- 3,000 21,000% of TED - 100.0 - -%ofEDD - 100.0 -- 12.3 87.7% of PDP 78.4 16.4 2.3 2 9 100.0- 0.4 2.5% of Dry in WholeYear, Same Stream15.0 5 5 5.5 4 7 10.81.3 4.7Wet Season = Nov. 20, 1987 - May 8, 1988; Sept. 28, 1988 - Nov. 19, 1988. Dry Season = May 9, 1988 - Sept. 27, 1988. Annual = Nov. 20, 1987 -Nov. 19, 1988.- 45 -^EMO^EDD^PDP33.00% 6.6% 100%EDL67.00%TCR4.50%NDD^TED14.40% 32.3% NDS9OthersPDW56.60%Figure 6 Contribution of Fresh Water Streams to TED5.1.3 Water Balance CalculationWe calculated the pathway of the total precipitation in the period September 27, 1987 -September 24, 1988 on the basis of an enclosed area on the North Dump. It is assumedthat all the rain water falling on this area either evaporated, or reported to TED, orpercolated into the ground water system. The result is presented in Table 6. A similarpathway calculation was also performed on the rain water falling on EMO. The resultsappear in Table 7. Details of calculation and assumptions for Tables 6 and 7 are includedin Appendix I.The percentage of the rain water recharging the ground water in EMO, 7.3%, is much lowerthan that of the first calculation. This difference may be due partly to variations inenvironmental conditions between the two areas, and partly to assumption errors.Table 6 Pathway of Rain Water Falling on an Enclosed Area on North DumpVolume of Water (m 3) Percentage (%)Precipitation 3,609,000 100.0Evaporation 783,000 21.7*To TED (asrunoff andseepage)2,210,000 61.2Percolationto Ground618,000 17.1Water* Calculated based on the pit area, precipitation and dewatering data for the period Sept.27, 1987 - Sept. 24, 1988. Details can be found in Acid Mine Drainage Study of theNorth Dump - Final Report. It is felt that this figure is probably too high.- 47 -Table 7^Pathway of Rain Water Falling on EMOBreakdown^Volume of Water (m3 )^Percentage (%)Precipitation 311,200^100.0Evaporation^ 67,500 21.7*To EMO Ditch 221,000^71.0(As Seeps orRunoff)To Ground^ 22,600^ 7.3* See the footnote for Table Calculation of Acid Generation Life of EMO and CapsThe duration of acid generation in both EMO and the Caps has been calculated using thedata presented in the above three sections. The details of calculation are documented inAppendix II. The calculation, however, is on an entirely static basis. It did not take intoconsideration the changes in the rate of acid production and consumption in the future. Toforecast the changes in acid production rate and acid consumption rate requires amathematical model, which was not attempted in this study. The results of the calculationcan nevertheless be used in guiding the abandonment planning if caution is exercised.It is found that the EMO contains enough acid-producing material to produce sulfuric acidat the current rate (approximately 370 tonnes H 2SO4 / year) for 650 years. Presently, 80%of the sulfuric acid produced by the EMO is neutralized by acid-consuming materials within- 48 -the dump. The acid-consuming materials will be depleted in 300 years. If the current rateof acid release continues after that, the pH of the EMO drainage will significantly drop andperhaps will approach the lower limit of 1.5.Based on the current rate, the Caps is estimated to be capable of producing acid for 650years at a rate of 730 tonnes per year, of which approximately 87% is consumed within thedump. The acid-consuming material will last almost as long as the acid-producing material.Therefore, the major downstream effect during the entire acid-generating life of the Capswill be elevated levels of sulfate and some easily soluble metals.5.2 Land Dump Drainage Monitoring RecordsTables 8 through 16 present the available water quality data of drainage streamsemanating from the North Dump. These data are used in the following sections to examinethe chemistry of the dump drainage at Island Copper Mine.Table 8^North Dump Drainage (NDD) Water Quality 1980-1989Date^pHDiss.Solids(mg/1)Sulph.(mg/1)Alkal.(mg/1)TotalCa(mg/1)TotalMg(mg/1) Cu Fe PbDissolvedMnMetalsMo(p.g/1)Zn Cd AsDec^80^7.7 303 110 153 37.0 5.4 3.0 630 2.0 400 2.5 11 0.20Jun 81^8.0 182 40 72 28.0 3.2 4.4 324 2.5 75 3.0 13 0.10Mar 82^7.5 306 140 56 61.0 10.0 3.4 217 1.0 281 3.0 30 0.20Jun 82^7.4 436 215 98 93.0 12.0 3.6 38 4.3 409 3.0 9.9 0.10Dec 82^7.1 400 158 94 93.0 13.0 2.4 111 1.0 1178 5.8 53 0.50Jun 83^7.7 330 101 135 85.0 13.0 1.5 78 2.2 130 3.0 2.2 0.10•Dec 83^6.8 440 190 91 58.0 12.0 3.0 210 1.0 120 3.8 77 0.70Mar 84^7.1 231 200 60 77.0 9.0 4.7 76 1.0 120 20.0 170 1.20Jun 84^7.4 460 230 90 92.0 11.0 2.7 120 1.0 64 6.0 160 0.80Sep 84^7.1 560 300 120 180.0 20.0 2.3 180 1.0 260 9.1 92 0.30Dec 84^8.0 401 160 62 79.0 3.3 4.3 360 1.0 130 802.0 13 0.10Mar 85^8.1 558 250 89 100.0 7.1 2.8 94 2.4 65 8.5 6.5 0.30Jun 85^8.2 570 310 140 120.0 904.0 4.3 35 6.3 24 9.0 10 0.10Sep 85^8 220 40 200 61.0 804.0 1.8 1.2 2.4 1.5 6.0 2.4 0.10Dec 85^6.8 770 460 100 160 22 9.6 190 1.2 240 9.2 500 2.10Jul 86^6.9 630 280 130 140 17 6.8 450 2.5 630 2.4 1500 7.10••Sep 86^6.9 1400 1100 230 280 44 6.9 6000 3.7 2900 10.7 680 3.10Dec 86^6.8 560 340 52 100 9.6 49.0 50 3.8 620 9.6 1500 9.60Mar 87^6.6 570 300 68 120 14 23.0 80 0.5 520 5.6 1500 8.90Jun 87^6.2 1100 300 24 130 14 7.6 59 2.5 780 7.5 1800 6.70Sep 87^6.8 1600 820 240 360 47 6.8 9200 3.8 100 13.0 650 2.30"*N0v87 4.8 630 -100 190 21 2700.0 320 4000 9500 58.00Dec 87^6.7 510 310 34 110 11 51.0 130 4.8 820 9.7 2100 13.00Jan 88^6.5 620 420 19 130 16 350.0 220 6.4 1800 4.0 4100 21.00Feb 88^6.1 690 480 21 140 18 160.0 140 7.3 1300 4.8 3600 23.00Mar 88^7.0 550 380 31 120 13 62.0 130 1.0 1200 4.9 2400 15.00 0.2Apr 88^7.2 700 340 66 140 18 11.0 70 1.0 770 3.9 1700 9.20May 88^7.0 600 410 78 130 15 6.9 4.7 0.5 17 4.4 23 0.05Jun 88^7.0 600 350 77 130 15 7.3 22 1.2 760 4.5 1700 9.60 0.3Jul^88^6.8 450 140 170 20 5.0 330 1.0 1200 4.7 1200 4.40Aug 88^7.4 760 450 170 190 22.0 5.4 90 1.2 400 3.6 910 3.80Sep 88^6.3 1000 400 41 180 15 110.0 460 1.0 2500 6.6 16000 51.00 1.5Oct^88^4.4 1000 140 210 27 2400.0 480 67.0 5600 1.2 18000 110.00Nov 88^4.4 980.0 180 23.0 4300 4.5 9580 51.00Jan 89 280.0 300 4.9 1890 3.0 4420 23.00tICMPermit 6.5- 50 50 500 10 100Level^11.5* Old North Dump reclaimed. 80% flow diverted into TED. ** 100% drainage diverted into TED. ** After the first major rainstorm 1987.t ICM's permit for the release of effluent to sea water. This is included here for reference only.Table 9^Twin Lake Discharge (TLD) Water Quality 1986-1988Date^pHDias.Solids(mg/1)Sulph.(mg/1)Alkal.(mg/1)TotalCa(mg/1)TotalMg(mg/1) Cu Fe PbDissolved MetalsMn^Mo(pg/l)Zn Cd AsMar 86^7.0 380 81 30 37.0 4.3 3.0 120 2.0 150 2.6 100.0 0.40 0.11Jun 86^7.0 200 110 43 38.0 4.3 4.1 270 2.0 410 2.4 61.0 0.20 0.05Sep 86^7.3 310 150 81 66.0 7.2 4.2 190 1.2 2100 2.5 28.0 0.10 0.14Dec 86^7.5 78 73 11 11.0 1.3 2.8 130 2.5 47 1.2 43.0 0.20 0.2Mar 87^7.6 87 40 22 19.0 2.0 5.5 210 1.0 140 1.0 57.0 0.10Jun 87^7.7 110 40 34 24.0 2.4 6.6 170 1.0 290 1.3 30.0 0.10 2.5Sep 87^7.1 150 46 42 28.0 3.0 6.8 180 2.5 1100 3.6 46.0 0.40 0.3Jan 88^7.8 72 56 140 15.0 1.9 2.6 130 0.5 170 0.5 34.0 0.05Mar 88^7.2 120 56 23 21.0 2.3 2.5 69 1.0 62 0.5 31 0.05 0.1Jun 88^6.9 200 73 37 33.0 3.2 3.3 110 1.2 210 0.5 70 0.10 0.1Sep 88^6.0 87 13 1 14.0 1.5 33.0 390 1.0 310 0.5 160 0.80 0.2Dec 88^6.2 90 32 17 14.0 1.6 0.1*FreshWater^7.1 69 4.53 7.83 1.83 5.4 248 3.6 57 4.1 32 0.85 1.3BaselinetICMPermit 6.5- 50 50 500 10 100Level 11.5These are the annual averages of the year 1971 of twelve fresh water streams around the minesite. Data were collected in the pre-operational monitoring program. See BHP-Utah, ICM andRescan, April, 1988.t^See the footnote for Table 8.Table 10^North Dump Subterranean Flow (NDS) Water Quality 1988Date^pHDiss.Solids Sulph.(mg/1)^(mg/1)Alkal.(mg/1)TotalCa(mg/1)TotalMg(mg/1) Cu Fe PbDissolved MetalsMn^Mo(pg/l)Zn Cd AsJan^6.7 440 280 190 98.0 11.0 37.0 89 0.5 59 6.0 550 2.90Feb^6.5 510 320 47 110.0 11.0 58.0 41 4.9 150 6.6 730 4.50 0.2Mar^7.4 442 260 58 93.0 10.0 31.0 82 1.0 29 6.7 500 3.00Apr^7.1 700 360 81 150.0 19.0 51.0 35 1.0 160 4.5 1100 3.70Jun^6.9 600 350 85 120.0 14.0 45.0 32 4.3 73 7.7 800 4.10 0.1Sep^6.8 1100 700 90 260.0 24.0 110.0 18 1.0 440 11.0 2400 14.00Oct^7.0 500 64 180.0 19.0 77.0 34 2.0 150 9.5 1300 7.50Nov^6.6 73.0 49 3.5 260 5.7 1360 6.90 0.1Jan^'89^6.9 1310 33.0 56 1.0 54 5.4 1120 4.80•^*ICMPermit 6.5- 50 50 500 1 0 100Level^11.5**FreshWater^7.1 69 4.53 7.83 1.83 5.4 248 3.6 57 4.1 32 0.85 1.3BaselineSee the fotenote for Table 8.See the footnote for Table 9.Table 11^Ten Eighty Ditch (TED) Water Quality 1986-1989Solids^Sulph.^Alkal.^Ca^MgDiss. Total^Total^ Dissolved^Metals^(pg/1)Date^pH^(mg/1)^(mg/1)^(mg/1)^(mg/1)^(mg/1)^Cu^Fe^Pb^Mn^Mo^Zn^Cd^AsMar^86^8.2^690^340^130^150.0^17.0^15.0^36^2.4^170^19.0^170^1.00^0.4Jun^86^8.0^1100^550^190^220.0^27.0 18.0^16^3.6 77^22.0^140^0.50^0.6Sep^86^8.0^1200^680^210^210.0^36.0^10.0^12^1.2^26^25.0 49^0.10^0.5Dec^66^7.9^660^440^79^120.0^14.0 19.0^50^3.8^250^16.0^400^3.50 1.3Mar^87^7.8^820^370^120^150.0^18.0^22.0^34^1.2^260^16.0^330^2.10Jun^87^8.0^720^610^190^240.0^29.0 15.0^18^2.5^300^19.0 93^0.60^0.1Sep^87^8.2^1300^690^220^290.0^38.0^18.0^16^3.8 130^18.0^99^0.40 1.4Dec^67^7.9^1200^690^97^250.0^27.0 10.0 47^2.4^170^12.0^950^4.40Jan^88^8.1^640^420^29^160.0^17.0^31.0^32^1.0^400^13.0^870^5.40Feb^88^7.1 570^400^69^150.0^16.0 23.0^26^2.0^320^14.0^790^5.90Mar^88^8.1^590^380^87^140.0^15.0^16.0^29^2.0^280^13.0^480^3.60^0.4Apr^88^8.4 800^380^130^180.0^22.0 12.0 14^3.1 350^11.0^400^2.40May^88^8.6^900^480^160^180.0^22.0^10.0^12^2.5^380^13.0^210^1.50Jun^88^8.5^700^420^150^170.0^21.0 9.0 15^3.0^340^14.0 190^1.80^0.7Jul^88^8.0 460^230^190.0^32.0^11.0^16^1.0^110^17.0^29^0.10Aug^88^8.1^910^460^220^200.0^32.0 13.0 14^1.2 14^17.0 20^0.10Sep^88^4.5^840^600^-59^190.0^15.0^280.0^400^1.0^560^4.0^900^5.80Oct^68^6.8 700^47^230.0^25.0 75.0^11^3.2^1500^14.0^4000^21.00Nov^88^6.9 22.0^11^2.4^870^13.0^3450^9.20Dec^88^7.5^710^489^87^160.0^17.0 0.2Jan^89 28.0^54^1.0^510^13.0^1100^5.80**FreshWater^7.1^69^4.53^7.83^1.83^5.4^248^3.6^57^4.1^32^0.85^1.3Baseline44e.1CMPermit^6.5- 50^50^500^10^100Level^11.5See the fotenote for Table 9.See the footnote for Table 8.Table 12^Trey Creek (TCR) Water Quality 1987-1989Date^pHDiss.Solids(mg/1)Sulph.(mg/1)Alkal.(mg/1)TotalCa(mg/1)TotalMg(mg/1) Cu Fe PbDissolvedMnMetalsMo(Ng/1)Zn Cd AsSep 87^8.0 1800 970 200 390.0 48.0 15.0 9.3 2.5 20 12.0 260 0.70 0.1Nov 87^7.5 570 74 210.0 23.0 11.0 45.0 2.3 88 6.2 590 2.50Dec 87^8.0 1200 740 100 270.0 29.0 9.9 110.0 7.3 170 9.9 950 4.60Jan 88^7.9 1600 1050 20 370.0 40.0 6.0 26.0 0.5 150 5.6 1200 3.20Feb 88^7.5 1560 1050 90 370.0 38.0 7.1 22.0 1.0 250 4.8 1500 4.50Mar 88^8.1 1700 1200 118 410.0 42.0 5.1 8.7 1.0 27 6.1 790 3.50 0.1Apr 88^8.4 1400 830 150 350.0 39.0 6.2 17.0 2.6 50 6.4 940 1.30May 88^8.5 1900 1200 190 420.0 47.0 6.8 9.7 1.2 78 7.6 570 1.20Jun 88^8.4 1900 1200 180 410.0 46.0 3.0Jul 88^8.0 1300 200 440.0 59.0 10.0 6.7 1.0 42 8.1 190 0.70 0.1Aug 88^8.0 2000 1250 190 440.0 55.0 14.0 20.0 1.9 15 7.9 240 0.70Sep 88^7.9 1600 1000 150 360.0 41.0 11.0 31.0 1.0 94 10.0 24 2.40Oct 88^7.6 1300 100 180.0 18.0 6.5 32.0 520 5.7 2400 4.20Nov 88 900 68 310.0 35.0 13.0 25.0 2.9 560 4.0 2330 6.20Dec 88 1199 130 450.0 52.0 0.1Jan 89it Fresh4.7 19.0 1.0 150 3.9 2540 3.70Water^7.1 69 4.53 7.83 1.83 5.4 248 3.6 57 4.1 32 0.85 1.3BaselineiNt ICMPermit 6.5- 50 50 500 10 100Level 11.5See the fotenote for Table 9.See the footnote for Table 8.Table 13^Eastern Most Outslope (EMO) Water Quality 1987-1989Date^pHDiss.Solids(mg/1)Sulph.(mg/1)Alkal.(mg/1)TotalCa(mg/1)TotalMg(mg/1) Cu Fe PbDissolvedMn- ^MetalsMo(pornZn Cd AsSep 87^4.8 2100 1400 390.0 67.0 85.0 170.0 3.3 9200 310 7.60Oct^87^4.7 2000 1000 390.0 65.0 78.0 470.0 8000 4300 6.90Nov^87^4.8 1400 -150 510.0 33.0 870.0 170.0 3700 8000 44.00Dec^87^4.5 1700 1000 -110 360.0 26.0 430.0 170.0 12.0 2900 4500 22.00Jan^88^4.8 1800 1300 -110 420.0 30.0 370.0 100.0 5.1 2900 5.4 4800 13.00Feb 88^4.7 1600 500 -84 370.0 24.0 360.0 100.0 5.5 1600 5.3 3700 13.00Mar^88^4.5 2000 1400 -173 450.0 40.0 340.0 200.0 2.5 2700 3.7 5800 14.00 0.1Apr^88^4.4 2400 1250 -380 390.0 53.0 520.0 290.0 1.0 6800 3.4 10000 10.00May 88^4.3 2600 2200 -406 430.0 67.0 520.0 830.0 4.5 8600 4.1 1000 23.00Jun^88^4.2 2800 1800 -404 440.0 68.0 510.0 780.0 6.1 9400 3.9 10000 31.00 0.1Jul^88^4.4 2000 -460 420.0 82.0 360.0 380.0 1.0 16500 4.7 3000 6.90Aug^88^4.3 2900 1900 -374 430.0 79.0 290.0 240.0 0.5 11000 3.1 11000 13.00Sep 88^4.4 1900 1400 -180 440.0 35.0 500.0 400.0 5.0 25000 11.0 36000 33.00 0.1Oct^88^4.3 1900 -250 540.0 45.0 1200.0 120.0 7.5 5400 4.6 12000 64.00Nov^88^4.3 980.0 120.0 3.8 4100 3.9 7710 14.00Jan 89 980.0 290.0 2.5 7000 4.8 9300 14.00*FreshWater^7.1 69 4.53 7.83 1.83 5.4 248 3.6 57 4.1 32 0.85 1.3Baselinefrit ICMPermit 6.5- 50 50 500 10 100Level^11.5See the fotenote for Table 9.See the footnote for Table 8.Table 14^East Drainage Lake (EDL) Water Quality 1988-1989Date^pHDiss.Solids(mg/1)Sulph.(mg/1)Alkal.(mg/1)TotalCa(mg/1)TotalMg(mg/1) Cu Fe PbDissolvedMnMetalsMo(Ng/1)Zn Cd AsJan^7.8 75 460 55 160.0 12.0 37.0 37 0.5 660 6.7 440 2.50Feb^7.0 760 520 46 160.0 13.0 65.0 33 1.0 790 6.0 560 3.30Mar^7.6 930 600 63 200.0 15.0 27.0 17 1.0 500 8.0 520 2.80 0.1Apr^8.2 1100 500 73 220.0 19.0 20.0 17 1.0 450 5.6 370 0.80May^8.3 1400 880 120 280.0 23.0 12.0 6.5 0.5 200 8.3 120 0.40Jun^8.3 1500 900 130 290.0 25.0 13.0 17 1.8 93 9.0 40 0.30 0.1Jul^7.5 800 130 260.0 24.0 9.7 8.3 1.0 19 8.7 33 0.05Aug^7.9 1500 950 150 310.0 26.0 8.1 12 0.5 15 12.0 22 0.05Sep^8.0 1500 1000 170 350.0 27.0 11.0 18 1.0 64 10.0 28 0.05 0.1Oct^7.9 800 130 270.0 23.0 9.8 24 2.9 140 9.8 92 0.20Nov 500 96 220.0 16.0 11.0 36 1.0 190 3.9 110 0.40Dec^7.1 640 489 58 120.0 11.0 0.1Jan^'89 44.0 20 1.0 1490 4.2 1190 5.20A FreshWater^7.1 69 4.53 7.83 1.83 5.4 248 3.6 57 4.1 32 0.85 1.3Baseline4* ICMPermit 6.5- 50 50 500 10 100Level 11.5See the fotenote for Table 9.See the footnote for Table 8.Table 15^East Dump Drainage (EDD) Water Quality 1981-1988Date^pHDiss.Solids(mg/1)Sulph.(mg/1)Alkal.(mg/ 1 )TotalCa(mg/ 1 )TotalMg(mg/ 1 ) Cu Fe PbDissolvedMnMetalsMo(wg/1)Zn Cd AsDec^81^8.1 363 80 135 48.0 7.1 9.0 300 3.0 339 4.0 18.0 0.10Mar 82^8.0 366 156 114 62.0 5.5 3.4 217 1.0 130 7.3 9.4 0.20Jun^82^8.2 513 203 148 82.0 6.7 4.6 107 1.0 36 7.0 3.5 0.20Sep 82^8.0 470 170 172 82.0 7.2 5.2 61 2.0 8 5.9 6.6 0.20Dec 82^7.8 500 178 153 84.0 7.8 3.0 367 2.0 365 15.0 6.2 0.10Mar^83^7.8 101 218 91 62.0 4.9 4.9 120 4.3 130 34.0 6.4 0.10Jun^83^8.1 480 206 153 38.0 10.0 3.4 34 4.4 15 19.0 3.0 0.20Sep^83^8.1 660 310 181 110.0 10.0 5.5 42 4.3 15 18.0 80.0 0.10Dec^83^8.3 480 190 120 76.0 8.5 3.0 220 2.0 87 20.0 3.6 0.10Mar^84^7.9 342 190 92 73.0 5.8 3.7 180 1.0 170 23.0 15.0 0.10Jun^84^8.0 570 290 142 110.0 7.7 3.0 41 2.0 38 16.0 10.0 0.20Sep 84^8.3 680 310 170 210.0 16.0 3.0 73 2.0 22 9.1 4.9 0.10Dec 84^8.0 401 160 95 79.0 3.3 4.3 360 2.0 130 9.9 13.0 0.10Mar 85^8.1 558 250 120 100.0 7.1 2.8 94 2.0 65 11.0 6.5 0.30Jun^85^8.2 570 310 150 120.0 9.4 4.3 35 5.0 24 15.0 10.0 0.10cn Sep 85^8.0 220 40 170 61.0 8.4 1.8 1.2 2.4 2 7.2 2.4 0.10-4 Dec^85^7.5 410 200 91 73.0 6.5 3.1 36 1.0 36 4.0 1.0 0.10Mar^86^8.2 580 310 140 130.0 9.7 4.4 43 2.0 210 10.0 16.0 0.10Jun^86^8.1 1000 640 150 190.0 14.0 6.6 29 4.8 15 18.0 13.0 0.10Sep 86^8.0 570 270 140 110.0 11.0 3.6 16 2.4 4 6.1 10.0 0.10Dec 86^7.3 640 370 77 140.0 12.0 17.0 79 2.5 78 5.9 33.0 0.10Mar^87^7.7 950 560 94 190.0 15.0 23.0 30 1.0 200 9.2 110.0 0.70Jun^87^8.1 1500 920 290.0 25.0 8.1 11 3.8 17 9.6 41.0 0.10Sep 87^8.2 1800 980 150 360.0 32.0 10.0 1.6 1.3 490 11.0 170.0 0.20Dec 87^6.9 1300 830 • 42 260.0 21.0 36.0 21 2.0 1100 7.1 1200.0 8.40Mar^88^5.6 1300 1000 5 280.0 23.0 110.0 110 2.5 2000 6.1 2100 9.50 0.1Jun^88^4.9 2500 1400 -94 360.0 34.0 170.0 170 2.0 333 5.1 3600 6.70 0.1Sep 88^7.8 1500 1000 150 340.0 28.0 7.8 14 2.0 330 9.4 1400 5.00 7.0Dec^88^4.8gFresh1000 647 2 210.0 19.0Water^7.1 69 4.53 7.83 1.83 5.4 248 3.6 57 4.1 32 0.85 1.3BaselineICMPermit^6.5- 50 50 500 10 100Level^11.5See the fotenote for Table 9.See the footnote for Table 8.Table 16^East Dump Drainage (EDD) Water Quality 1988-1989Date^pHDiss.Solids(mg/1)Sulph.(mg/1)Alkal.(mg/1)TotalCa(mg/1)TotalMg(mg/1) Cu Fe PbDissolvedMnMetalsMo(pg/l)Zn Cd AsJan 88^6.5 1100 740 11 260.0 20.0 98.0 52 2.5 1600 6.7 2000 7.70Feb 88^6.3 990 700 18 230.0 16.0 120.0 38 2.4 1200 6.6 1600 7.80Mar 88^5.6 1300 1000 5 280.0 23.0 110.0 110 2.5 2000 6.1 2100 9.50 0.1Apr^88^6.2 1200 700 13 240.0 25.0 54.0 2.0 1200 3.9 2000 3.20May 88^6.6 1700 1100 23 310.0 33.0 20.0 70 2.0 2300 6.3 2000 6.30Jun 88^4.9 2500 1400 -94 360.0 34.0 170.0 170 2.0 333 5.1 3600 6.70 0.1Jul^88^4.6 1300 -117 320.0 51.0 160.0 65 2.7 6700 4.7 5500 6.40Aug 88^7.1 1100 1100 68 320.0 35.0 11.0 11 1.0 2200 10.0 1100 3.20Sep 88^7.8 1500 1000 150 340.0 28.0 7.8 14 2.0 330 9.4 1400 5.00Oct 88^6.8 1000 49 350.0 28.0 220.0 12 2.0 1500 5.7 3800 9.30Nov 88^6.0 130.0 19 1.0 1240 4.5 2210 8.80Dec 88^4.8 1000 647 2 210.0 19.0 7.0Jan 89 300.0 86 1.0 2940 4.2 3160 13.00*FreshWater^7.1 69 4.53 7.83 1.83 5.4 248 3.6 57 4.1 32 0.85 1.3Baseline1CMPermit 6.5- 50 50 500 10 100Level^11.5See the fotenote for Table 8.See the footnote for Table 9.5.3 Chemistry of North Dump Drainage WaterAs is observed by researchers, many natural waters are not in thermodynamic equilibrium.The drainage streams at ICM may be no exception. Realizing this limitation, we cannevertheless employ equilibrium chemistry in our aid in understanding the processes inAMD generation. As has been demonstrated in many natural water studies,thermodynamics can deliver very good, sometimes excellent answers to many questions.No one can avoid the problem of activity coefficients while dealing with electrolyte solutionchemistry. What we measure with analytical tools such as AAS and ICP is theconcentration. What we use in the entire kingdom of equilibrium chemistry (that is,thermodynamics) is the activity. The bridge linking these two is the activity coefficient.In very dilute solutions, such as those with a total molar ionic strength of 10 or less, thedistinction between activity and concentration disappears. But in the case of ICM drainage,whose total ionic strength is around 0.0426 as demonstrated in the following section,activity coefficients deviate markedly from unity and therefore have to be taken intoaccount. This is accomplished by first estimating the total ionic strength of a typicaldrainage and then calculating the individual ion activities using the improved Debye-Hiikelformula. Today's science provides very exact but complicated empirical expressions for thecalculation of activity coefficients. They were not adopted because firstly any gain inprecision would be outweighed by the inaccuracy of estimated total ionic strength which allthe empirical expressions are based on and secondly the extended Debye-Hiikel formuladoes give an adequate accuracy up to a total ionic strength of 0.1.The activity of water can be taken to be unity in all cases without incurring noticeableerror. The water activity in sea water is 0.98 at an ionic strength of about 0.7. All thedrainage at ICM has an ionic strength less than 0.1, therefore a water activity greater than0.98 but less than unity.Although working with individual ion activities in solving electrolyte aqueous solutionproblems is convenient, individual ion activities are not experimentally measurable. Whatwe can measure experimentally is the mean activity coefficient. To partition the meanactivity coefficient of ionic salts into individual ion activity coefficients, we need to adopta more or less arbitrary reference point. We shall follow the convention that = TC1- =7x,-, where y, represents the mean activity coefficient and yi the individual ion activitycoefficients. Under this reference frame, the individual ion activities can be calculated bythe Debye-Hiikel formula, which was originally developed for calculating mean activitycoefficients. The constants A, B and a can be viewed as data fitting constants although theyalso have physical interpretations. The Debye-Hukel formula has the formAzT/I-log yi =1 + AAPwhere I is the total ionic strength and z i the charge on the ion i. The total ionic strengthof an aqueous electrolyte solution is defined asI = 1/2/(mizT)where mi is the molality of ion i and the summation is over all ionic species. In the caseof ICM drainage, we can neglect the deviation of drainage density from unity and thus usemolarity in lieu of molality. Values of A, B and A for various ionic species are tabulated inGarrels, 1965. The above formulae will be used in the next section to calculate the totalionic strength and activity coefficients of a typical ICM drainage.Colloids and suspended solids may play at times important roles in the metal balance ofdrainage water. They are discussed in a later section very briefly because we have onlydata on dissolved species.5.3.1 Mineral Constituent Mobilization and Dissolved Species Removal in GeneralThe earth surface has been undergoing weathering for billions of years. Weathering isresponsible for the formation of soil which supports the great majority of advanced lives andplants, for the earth's surface landscape (plate tectonics also contributes to this), and forthe saline nature of the ocean. Weathering can be broadly divided into physical weatheringand chemical weathering. Physical weathering is the disintegration of rocks by physicalmeans such as temperature fluctuation, the action of flowing water, wind action, frostaction, abrasion and so on. One of the important effects of physical weathering is that itaccelerates the rate of chemical weathering by increasing the total surface area of rocksupon which chemical weathering attacks. We will not discuss physical weathering furthersince this is rather irrelevant to our subject. Chemical weathering is the change of mineralcomposition of rocks through chemical reactions. The composition of water contained in theocean, rivers, and lakes is largely determined by chemical weathering. Chemical- 61 -weathering is a very broad field and has attracted tremendous interest of study. We willdiscuss briefly the chemical weathering processes that are important in our present study -the chemical weathering processes of silicates, sulfides, oxides and carbonates.Igneous rock makes up 95% of the outer 15 km of the earth's crust and 25% of the earth'ssurface. The vast majority of igneous rock-forming minerals are silicates. Silicate mineralsare thermodynamically stable at the conditions of their formation: high temperature, highpressure, and low oxidizing potential. They become thermodynamically unstable when theyare exposed at the earth surface, which is characterized by low temperature, low pressure,high oxidizing potential and presence of water. The stability of silicates is not muchaffected by the change of redox potential from the earth interior to the surface, since all theelements in silicates are already in their most stable oxidation states under earth's surfaceconditions. The weathering of silicates, therefore, is not a redox reaction. Thethermodynamic instability is brought about by the change in temperature, pressure andchemical environments, most importantly by the contact with water. The silicate mineralsweather by releasing alkaline or alkaline earth metal cations (IC -, Na+, Ca++, Mr), whichare subsequently carried to the ocean, leaving behind aluminum oxides (gibbsite, bauxite),aluminosilicates (clay minerals), silicon oxides (silica, crystalline or amorphous), and others,all of which contribute to the formation of soil. The alkaline and alkaline earth metalcations are displaced by W; thus the chemical weathering of silicate is usually an acidconsuming process. The weathering of silicates via the above mechanism can be regardedas hydrolysis. It is important to remember that the weathering of silicates is usually veryslow. Different silicates weather at different rates. It has been observed that the relativerates of weathering of silicates can be described by the reverse of Bowen's Reaction Series:the earlier a silicate mineral appears in Bowen's Reaction Series, the faster its weatheringrate.Sulfide minerals are generally of igneous origin. Their formation is characterized bystrongly reducing conditions and usually high temperature and pressure. But high pressureand high temperature are not essential. In contrast to silicates, the instability of sulfidesunder earth surface conditions is caused by the change in redox conditions. The oxidationstate of sulfur in primary sulfides is usually -I or -II while the only stable oxidation stateof sulfur at earth's surface is +VI. The weathering of sulfides is unexceptionally redoxreactions. These redox reactions convert reduced sulfur into S(+VI) in the form of sulfate.The occurrence of sulfides in igneous rock masses is usually localized. The rate ofweathering of sulfides is much greater than that of silicates. The weathering of sulfides isfrequently self-enhancing, because the release of acidity from previous weathering, if notremoved, tends to make the environment more oxidizing, and a more oxidizing environmentspurs a faster release of acidity.Soil is a product of rock weathering and biological activities. The type of soil formed in aparticular area is dependent of five broad variables. climate (temperature andprecipitation), parent rock, biological activity, topography, and the time during which thesoil has been forming (Carroll, 1970). Variations in the five variables have resulted innumerous soil types, which are classified virtually according to chemical compositions.Having said this much about weathering, we now turn our attention to see how chemicalweathering is related to the investigation of acid mine drainage and water pollution.Usually, natural weathering does not result in great concentrations of toxic species in- 63 -natural streams because of the low surface area to mass ratio of undisturbed rock masses.Water pollution refers to the release of environmentally toxic species in concentrations thatgreatly exceed the natural levels due to man's activities. Since our goal is to find out howand why the pollutants in ICM drainage come to the present levels and what the futuretrend is, only the aqueous species whose concentrations greatly exceed natural levels,namely pollutants concern us. We use the phrase "mineral constituent mobilization" todescribe the transfer of chemical elements, toxic or not, from minerals into solution.Mineral constituent mobilization differs from the phrase "chemical weathering" in that thelatter has a broader meaning than the former. In addition to the release of mineralconstituent elements to aqueous solution, chemical weathering also concerns about thetransition between different minerals, formation of soil, etc., which are out of our scope ofinterest.There are also some differences in the interest of study between chemical weathering andour present acid drainage study. Chemical weathering studies concern the effects ofmineral breakdown or mineral transition on the water composition of river systems, on thesoil formation and on new mineral formation. Our acid mine drainage study concerns aboutthe effect of mineral breakdown or mineral transition on the release of pollutants atconcentrations greatly exceeding the natural levels. Chemical weathering studies are oftencarried out on very large catchment basins while our study is on a very small and localizedarea.Water draining waste dumps is in constant motion and the waste dumps receiveintermittent renewal of fresh water from precipitation. The products of mineral breakdownare carried away from the reaction spot by flushing water. Therefore the concentrations- 64 -of pollutants in the acid mine drainage from waste dumps are affected by the rate ofmineral breakdown, the removal of pollutants by various mechanisms such as precipitation,coprecipitation, adsorption to solid surfaces, etc. and the pattern of precipitation (frequency,amount, interval, etc.) in a relatively complex manner. The mechanisms of mineralbreakdown that are relatively fast are more important than those that are relatively slow.We will examine the mechanisms of mineral breakdown next.Keller pointed out in 1955 the various types of chemical reactions involved in chemicalweathering. These include hydrolysis, ion exchange, oxidation, carbonation, hydration,chelation, dialysis, simple dissolution and chemical reconstruction. To this list we must addcomplexation. Of these reaction types, hydrolysis, oxidation, carbonation and simpledissolution are more important ones in environmental considerations of the study of acidmine drainage. They are considered in further detail in the following sections. Cationexchange and chelation are more profound near the surface zone of plant growth. Thesetwo processes contribute an insignificant amount to the total release of mineral constituentsinto acid mine drainage because chemical weathering takes place throughout the entiredepth of the waste dump, of which the surface zone is only a very small portion. Hydrationand mineral reconstruction do not release mineral constituents into solution. The contribu-tion of dialysis is also regarded insignificant in the present study.Natural waters draining an igneous formation contains in significant concentrations KF,Na+, Ca2+, Me, H4SiO4 (SiO2 .2H20), SO:, HCO3 and C1. The first five species in theabove list come from the chemical weathering of silicate rocks. Bicarbonate ions derivefrom the dissolution of atmospheric CO 2 and subsequent dissociation of carbonic acid, andof carbonate minerals. Cl - is primarily from precipitation. SO4 originates from the- 65 -dissolution of gypsum and the oxidation of primary sulfides when available. Because of theigneous nature of ICM waste rocks, all these species are expected to be present in thedrainage from the waste dumps at ICM.Due to the high acidity and low pH in some of the ICM waste dumps, some metal speciesare mobilized into solution to significant concentrations. These are Cu, Fe, Mn, Mo, Zn, andCd. Pb and As are not in concentrations much higher than natural levels, but areconsidered because firstly they are very environmentally toxic and secondly ICM hasmonitored them. In addition, SO 4-, Ca", Mg' are in concentrations much greater thannatural levels owing to the oxidation of sulfides and the accelerated silicate mineral weat-hering. The concentration of Ca' and Mg' is probably also attributable to the destructionof calcite and dolomite by acid generated in the dump.Table 17 presents the results of a survey of the concentrations of various species in the ICMdrainage streams for the period from January, 1988 to January, 1989. The factors toconvert original concentrations of various species in mg/1 or pg/1 to those in mo1/1 are listedin the table. The conversion was performed for EMO drainage. Since the concentrationsof various species vary in a wide range from stream to stream and within a stream fromtime to time, it is unrealistic to consider the chemistry of all streams and for all variations.We choose to take EMO drainage to carry on our study of solution chemistry. Thetreatment of other streams is similar except the range of contaminant concentrations. TheEMO drainage is chosen because firstly it is the most seriously polluted drainage and moreefforts are needed in alleviating its pollution and secondly the water quality of otherdrainage streams may approach that of the EMO drainage in the future when the sulfideStream^pH^SO4(mg/1)^Total^TotalCa^Mg(mg/1)^(mg/1)Dissolved Metals (Ng/1)Cu^Fe^Pb^Mn^Mo^Zn^Cd^AsTable 17^Range of Contaminant Concentrations January 1988 - January 1989TLDBLASCRNDDNDSTEDTCREMOEDLEDD6.0-7.8 13-110^14-46^1.5-12^1.6-33^69-390 0.50-3.1 62-21406.0-6.7 1.0-38^9.2-15^1.3-11^1.7-4.6^68-300 0.50-1.2 11-18006.4-7.6 0.70-16^5.8-9.9 0.80-1.4 0.60-2.0 92-310 0.50-2.6 3.0-994.4-7.2 380-1000 120-210 13-27^5.4-2400 4.7-480 0.50-67^17-56006.5-7.4 260-700 93-260 10-24^31-110^32-89^0.50-4.9 29-4404.5-8.6 380-700 140-230 15-32^9.0-280^11-400 1.0-3.2^14-15007.5-8.5 830-1300 180-450 18-59 5.1-14 6.7-32 0.50-3.0 15-5604.2-4.8 500-2200 370-540 24-827.0-8.3 460-1000 120-350 11-274.6-7.8 650-1400 210-360 16-510.50-1.9 31-160^0.05-0.80 0.1-0.20.50-1.9 12-130^0.05-0.40 0.1-0.10.50-1.0 7.4-60^0.05-0.10 0.1-0.11.2-6.6^23-18000^0.05-110 0.2-1.54.5-11^500-2400^2.9-14^0.1-0.24.0-17^20-4000^0.10-21^0.2-0.73.9-10^24-2500^0.70-6.2^0.1-0.11000-36000 6.9-64^0.1-0.122-1200^0.05-5.2^0.1-0.11100-5500^3.2-13^0.1-0.7290-1200 100-830 0.50-7.5 1600-25000 3.1-118.1-65^8.3-37^0.50-2.9 15-1500^4.2-127.8-300^11-170^1.0-2.7^330-6700^3.9-10Range 4.2-8.6 0.7-2200 5.8-540 0.80-82 0.60-2400 4.7-830 0.50-67^3.0-25000 0.50-17^7.4-36000 0.05-110 0.1-1.51CMPermit50 50 500 10 100(To convert the above figures into mo1/1, multiplythem by the following factors)1.04E-5^2.50E-5 4.12E-5^1.57E-8^1.79E-8 4.83E-9^1.82E-8^1.04E-8^1.53E-8^8.90E-9^1.34E-8EMO^(4.2-4.8)5.2E-3-^9.3E-J- 9.91-4-^4.6E-6-^1.8E-6- 2.4E-9-^2.9E-5-(m01/1)^2.3E-2^1.4E-2^3.4E-3^1.9E-5^1.5E-5^3.6E-8^4.6E-4Typical (4.5)^1.0E-2^1.0E-2^1.0E-3^1.0E-5^1.0E-5^1.0E-8^1.0E-4Drainage(mol/1)3.2E-8-^1.53E-5-^6.1E-8-^1.3E-9-2.0E-7^5.5E-4^5.7E-7^1.3E-91.0E-7^1.0E-4^1.0E-7^1.0E-9* See the footnote for Table 8.oxidation proceeds to more advanced stages. The ranges of concentration variations ofvarious aqueous species in the EMO drainage are shown in mo1/1 in Table 17. From theEMO drainage, we construct a "typical drainage" (hereforth referred as "the typicaldrainage") using the orders of magnitude of the concentrations of the aqueous species in theEMO drainage. The composition of this typical drainage is shown in Table 18.We have shown above that the ICM drainage contains in significant concentrations K+, Na+,Ca2+, Mg2+, H4SiO4 ( Si 02 .2H20), SO4-, H2CO3, HCO3 , C1, Cu, Fe, Pb, Mn, Mo, Zn, Cd, andAs. For the typical drainage, we already know the concentrations in mo1/1 of the eightdissolved metal species (including As), SO:, Ca 2+, Mg2+, and pH. In order to calculate thetotal ionic strength of the typical drainage, we need to estimate the concentrations of HCO3,and Cl -. H4SiO4 and H2CO3 are neutral species and do not contribute to the total ionicstrength. The concentrations of carbonate species in solution depend on the solution pH(see the Eh-pH diagram and concentration-pH diagram of C-O-H system presented inSection At the pH of the typical drainage (4.5), the prevailing species is carbonicacid H2CO3 ; the concentration of bicarbonate ion HCO3 is two orders of magnitude lowerthan that of H 2CO3 and the concentration of carbonate ion CO 3= is almost nil. If we assumethe total concentration of carbonate species in solution to be 10 -3 moll, which is a commonvalue for natural waters draining non-carbonate rock formations with some surfacevegetation (Dreyer, 1982), the concentration of aqueous H 2CO3 would be 10-3, that of HCO310-5 and that of CO3 nearly zero. For Cl - , we will adopt the order of magnitude in anaverage precipitation, 10 -5 mo1/1. The rain water usually also contains nitrogen species, pre-sumably NO3, in a concentration comparable to that of chloride ion (10 -5 mo1/1). Biologicalactivities can add much more nitrogen species, such as NO 3- and NH4+ to the solution. Thisis ignored since the biological activities are only appreciable near the surface while the-68-Table 18^Composition of the Typical DrainageSpecies pH^SO4 H4SiO4 H2CO3 HCO3 Cl NO3 Diss.N 2 Diss.02Concentration(mo1/1)4.5^10 2 10 3 10 3 10 5 10 5 10 5 9.5x10 4 3.1x10 -4Activity 0.86^0.467 1.10 1.10 0.83 0.82 0.82 1.10 1.10CoefficientActivity 2.7x^4.67x 1.10x 1.10x 8.3x 8.2x 8.2x 1.05x 3.4x10 5^10 3 10 -3 10 3 10 -6 10-6 10-6 10 -3 10 -4Species^+^+^K Na Ca ++ Mg ++ Cu++ Fe++ Pb++ Mn++ aMo04 Zn++ Cd ++ H2AsO4Concentration(mo1/1)10-4 10 -4 10-2 10 -3 10 5 10 5 10-8 10 -4 10-7 10 -4 10 -7 10-9Activity 0.82^0.83 0.50^0.53 0.50 0.50 0.48 0.50 0.83 0.50 0.49 0.83CoefficientActivity 8.2x^8.3x 5.Ox^5.3x 5.Ox 5.Ox 4.8x 5.Ox 8.3x 5.Ox 4.9x 8.3x10-5 10 5 10 3^10-4 10 6 10 -6 10 9 10 5 10 -8 10 -5 10 -8 10 -10Total Ionic Strength:^0.0426chemical weathering takes place in the entire dump. In cases where the rock containsphosphate minerals, phosphorous species may be present in significant concentrations. ICMwaste rocks do not contain significant amount of phosphate minerals therefore phosphorousspecies need not be considered.In Table 18, the total ionic strength is calculated from the molarities of all ionic species andthe activity coefficients are calculated from the total ionic strength and the concentrationsin molarity of various species using the extended Debye-Iliikel formula. Activity coefficientsof neutral species are calculated using the equation (Dreyer, 1982):7 . 100.11where I is the total ionic strength. Dissolved N2 concentration is calculated from thesolubility of N2 at 25°C (23.3 cm 3/1). Dissolved 02 concentration is based on measurementsmade at ICM.Dissolved species are removed from solutions by a number of chemical, physical, orbiochemical processes: over-saturation with respect to a highly insoluble compound, redoxreactions that yield a highly insoluble compound, complexation that results in a highlyinsoluble compound, adsorption onto solid or colloid surfaces, and biological uptake. In thestudy of ICM acid mine drainage, chemical removal of dissolved species is by far the mostimportant.Usually, the reactions that remove ionic species from solutions are much faster than thereactions that bring these species from mineral lattice to solutions. As a general rule,reactions between dissolved species in solutions are always much faster than solid-liquid- 70 -interfacial reactions, which is in turn faster than solid-solid reactions. In a solution, allionic or molecular particles of dissolved species take part in the chemical reaction while ina solid-liquid interfacial reaction only particles on the solid surface are involved. Forreactions in solutions, non-redox reactions are usually faster than redox reactions.5.3.2 Redox ReactionsReduction-oxidation, or redox reactions are very common in geochemical processes and areresponsible for the weathering of many minerals, including sulfides.There are two very basic processes in the realm of chemistry: electron transfer and protontransfer. Eh measures the potential of electron loss or gain and pH measures the potentialof proton loss or gain. We can speak of Eh of an aqueous solution, of a half-cell reaction,or of a redox reaction. The Eh of a half-cell reaction or a redox reaction is clearly definedin terms of the activities of the species involved in the reactions. As to pH, it always refersto aqueous solutions. The pH is clearly defined at any moment for any homogenous solution(at equilibrium or semi-equilibrium) as the negative common logarithm of the activity offree hydrogen ions.The Eh of aqueous solutions, however, deserves some clarification. It usually refers to thepotential of an inert electrode (such as a platinum electrode), which is dipped into thesolution being measured, relative to the Standard Hydrogen Electrode (SHE). Before themeasurement is taken, the inert electrode is usually allowed some time to reach"equilibrium" with the solution. If the solution being measured is at real equilibrium, all- 71 -possible redox pairs in the solution will define the same redox potential. Now if theplatinum electrode and the measuring process itself exert a negligible influence on theequilibrium of the solution and therefore do not upset the equilibrium of the solution (whichis equivalent to saying that the electric current that occurs during the measurement of thepotential between the inert electrode and the reference electrode is infinitesimal and thatthe insertion of the inert electrode does not bring any foreign substances), and if we furtherassume that the inert electrode does achieve equilibrium with the solution, the potentialso measured can be attributed to the solution without ambiguity. However, if the solutionbeing measured is not at equilibrium (which, as we have mentioned earlier, is a commonsituation in natural waters), what is the meaning of such a measurement? Since thesolution itself is not in thermodynamic equilibrium, such a measurement is called by somegeochemists a "formal potential". In the solution, the concentrations of many redox species(thus the potentials between redox pairs) are constantly changing in such a way as to bringthe system towards equilibrium. But some of these changes take place so slowly that theyare often at disequilibrium with the potential established commonly by those fast-reactingspecies, which we refer to as "semi-equilibrium potential". The semi-equilibrium potentialis usually determined in the early stage in the course from disequilibrium to equilibriumby a few fast-reacting redox pairs. If an inert electrode is immersed into such a system, andif the open circuit emf is measured across the inert electrode and the reference electrode,the potential obtained should be in the proximity of the semi-equilibrium potential;otherwise if the potential between the inert electrode and the reference electrode ismeasured using an electrometer with an external current (however small it may be) thatis great enough to upset the local semi-equilibrium around the inert electrode, the potentialobtained will be dominated by the fastest-reacting redox pair or pairs.In reality, most of the Eh measurements of natural waters are probably made by the secondmethod mentioned above; so these Eh values should reflect the potentials dominated by thefast-reacting redox pairs. The values of such measurements could be very different fromthose calculated from the concentrations of a particular redox pair if the solution is not inequilibrium, which is what is meant by disequilibrium after all. Nevertheless, the formalpotential Eh represents the mainstream redox potential of a solution. The usefulness of theformal (measured) Eh of natural water bodies is that, first, using the formal Eh as ayardstick, we can designate a water body as oxidizing water, or reducing water, ortransitional water; second, knowing the Eh, we can frequently predict the redox actions thatwill take place if a new species is introduced, as well as the stable species in a solution. Inthe following discussions, we will make extensive use of Eh-pH diagrams. One more pointto make about Eh-pH diagrams is that, being constructed from thermodynamic data, theyrepresent the equilibrium information only for the species considered, subject to theaccuracy of the thermodynamic data employed. They tell nothing about species notconsidered in their construction. Therefore, some caution has to be exercised to choose thespecies to be considered according to the application. For instance, when we construct anEh-pH diagram for iron (Fe), there are two species of Fe(III) oxides to choose from, Fe 203(hematite) and Fe(OH) 3 (ferric hydroxide). We know that hematite is thermodynamicallymore stable than amorphous Fe(OH) 3, which can be written alternatively as hydrated ferricoxide Fe203 .31-120:2Fe(OH)3 (amorphous) —> Fe203(hematite) + 3H20AG° = -63.7 kJ/molTherefore, given enough time, ferric hydroxide will dehydrate to convert to hematite. Wealso know from experience that the species to form if ferric ion Fe' is precipitated fromsolution by raising pH is ferric hydroxide. Given the above information, we will chooseFe(OH)3 in constructing the Eh-pH diagram if the removal of metal species from pollutedwater by chemical treatment is to be examined, although Fe(OH) 3 is only metastable andwill slowly revert to Fe203. On the contrary, if our interest is to investigate the geologicaloccurrence of iron oxides in an area, we will consider Fe 203 in the Eh-pH diagram, sinceFe(OH)3 would have had enough time to be converted to hematite in the geologic past. Theeffect on the Eh-pH diagram of choosing Fe(OH) 3 instead of Fe203 is, all other things beingequal, that ferric hydroxide occupies a smaller stability field on the Pourbaix diagram. Thisis expected because ferric hydroxide is thermodynamically less stable than hematite. Thestructure of the Eh-pH diagram will remain the same as if hematite were used.Another benefit in using Eh-pH diagrams is that when the total activities of some specieschange, there is usually only a slight parallel shift of the boundaries involving the specieswhose activities change, other boundaries are unaffected, and the general structureremains. Therefore, an Eh-pH diagram can be used effectively for solutions whose speciesactivities are in the neighbourhood of those assumed in constructing the Eh-pH diagram.They need not be exactly the same.Redox potentials of natural waters vary widely according to their environments. Figures7 and 8, both taken from Garrels, 1965, demonstrate the natural variations in Eh-pHmeasurements and his effort to zone the Eh-pH diagram according to the origins of thewaters in which the Eh and pH were measured.pHFigure 7 Distribution of Eh-pH Measurements of Natural Aquous Environments(After Garrels, 1965 with modifications)- 75 -- 1.06^P H^80^2^4110^12 14Figure 8 Approximate Positions of Some Natural Environments as Characterizedby Eh and pH.. (After Garrels, 1965 with modifications)- 76 -In addition to Eh, another important concept in electrochemistry is redox buffering. Asystem is buffered with respect to redox processes if oxidizable and reducible compounds(or species) are present that prevent a significant change in Eh in response to additions ofsmall amounts of strong oxidizing or reducing agents. Redox buffering is a close analogyto acid-base buffering. Just as acid-base buffering is realized by the presence in comparableconcentrations of a buffering pair, such as H 2CO3/HCO3 and CH 3COOH/CH3C00 -, redoxbuffering is facilitated by the presence of a redox buffering pair, such as 0 2/H20, SO42-/H2Sand H20/H2, although not necessarily at comparable concentrations. Figure 9 illustratesschematically how the redox pair 0 2/H20 buffers the change in Eh. It is assumed that thewater was initially in equilibrium with atmospheric oxygen, but no additional oxygen isadded as the organic matter decomposes. It can be seen that as long as free dissolvedoxygen is present, Eh remains high, but as soon as free oxygen is depleted, the Eh of thesystem drops dramatically to that buffered by the redox pair SO 42412S.Natural waters contain simultaneously many buffering pairs. The question arises in regardto which of these buffering pairs dictates the measured Eh, thus the overall redoxconditions. It appears justified to say, from the view point of the foregoing discussion ofsemi-equilibrium in natural waters, that the dictating redox pair (or pairs) would be that(or those) whose redox reaction(s) is the fastest in the environment concerned. Many suchpairs have been identified in different natural environments. Figure 10 shows some of theimportant natural redox buffering pairs along with their buffered Eh at pH 7.0 and 25°C.The buffered Eh of many of the natural buffering pairs changes with pH due to theinvolvement of II+ in redox reactions. This point is illustrated in Figure 11, which depictssome possible redox buffers in a ground water environment.20 —1.015 --,-1 0 —0.5a) 50 ^ 0_5 —_.-0.5-100L151 0a) 5a0—51FermentatiOn•••••■02 -*H 2 0SO4- H2 S0^1^2^3^4^5Amount of organic matter reacted(units are mmol C/2 H 2 0)Figure9 Change in pE (Eh) of a Fresh Water as a Function of.theAmount of Organic Matter Decomposed. (Dissolved oxygen= 10 mg/1, dissolved SO:= 96 mg/1, pH is assumed constantat 7.0. Reactions involving nitrogen compounds mayprovide a small amount of buffering between the 0 2/H20 andthe SO4/H2S levels, but they are ignored here. Also notethe pE-Eh convertion chart at 25 °C on the right.)(After Dreyer, 1982)15—— 02 /H2 0NO3 /N2— (N O3 IN H4 )5—a)0/ H 2 SCO2 /CH 4H 2 0/H 2Fermentationreactions—5—10 —10—5^6^7pHFigure 10 Approximate pE (or Eh)Values at which Various RedoxReactions Occur in Water at pH 7and 25°C. See Figure 34 for pE-Ehconversion. (After Dreyer, 1982.)Figure 11 Some Possible RedoxBuffers in a Groundwater System.Solid/solution boundaries are drawnfor activity of solute = 10'. SeeFigure 34 for pE-Eh conversion.(After Dreyer, 1982.)- 79 -Dreyer discussed the problem of dominant redox pairs of oxygenated (surface) water. Theoverall reduction of oxygen^02 + 4W + 4e - 4=* 2H20^ (1)Eh = Eh° 2 '303RT logaH2 2CAaH4 + TOO]nF1.229 - 0.05917pH^(2)= 0.815V (at pH7.0)generally does not occur as a single step, but as two separate reactions02 + 2H+ + 2e - -;----' H202^(3)RTEh = Eh° 2.303 ^log[aH202/(a4,1302)]nF— 0.682 - 0.05917 pH (4)= 0.27 (V) (at pH 7.0)andH202 + 2H+ + 2e- 4='-. 2H20^(5)Breck (as cited by Dreyer) pointed out that reaction (5) was much slower than reaction (3)and hence Eh was essentially controlled by reaction (3) (In accordance with the semi-equilibrium theory we discussed earlier). The effective Eh for oxygenated water would then- 80 -be 0.27 V (at pH 7.0) as predicted by reaction (1). Stumm (also as cited by Dreyer)disagreed and pointed out that some natural redox systems seemed to respond as if reaction(1) determined Eh, and some as if reaction (3) controlled Eh. Dreyer favoured Stumm'sopinion that the use of a single Eh to characterize oxygenated waters was meaninglessbecause the various redox couples in natural waters are not in equilibrium with each other.Stamm argued that since a single Eh for all redox systems could not be defined, it isprobably best to think of Eh as "high" without specifying an exact number. Since in realitymost of the Eh measurements are performed with a meter that employs an electrometer,we believe that the measured Eh of natural waters is determined by the most reactive redoxpair(s) (one that facilitates the largest exchange current) and is close in value to that of thesemi-equilibrium.We have plotted equations (2) and (4) on Figures 7 and 8. Equation (2) is the upper limitof water stability, and equation (4) is represented by the line labelled 0 2/11202 in Figure 7.The vast majority of the measured Eh-pH conditions of natural oxygenated (surface) waterfall between the two lines defined by equations (2) and (4). This is to say that reactions (1)and (3) actually define a boundary of occurrence of natural surface water (mine water, rain,streams, sea water, etc.) Eh-pH conditions. It seems that Breck's argument that the redoxpotentials of surface oxygenated water are controlled by reaction (3) is closer to reality,since the average line of the measured Eh-pH points falls closer to the line predicted byreaction (3) than to that predicted by reaction (1). Furthermore, almost no measurementsof natural waters fall within the band approximately 0.2V wide below the upper limit ofwater stability line.One more observation to make on Figure 7 and 8 is that the trend followed by themeasurements on natural surface aerated waters parallels the line of water upper stabilitylimit and the line defined by reaction (3). This implies that the more acidic a naturalsurface water, the more oxidizing it is. This explains from one particular angle why theoxidation of primary sulfides is favoured by low pH values.Based on the foregoing discussion of the measurements of natural aqueous environmentsand the considerations on the ICM environmental conditions, we can now "quantify" on anempirical basis a field on the Eh-pH diagram to encompass the possible Eh-pH conditionsin ICM drainage and within ICM waste dumps that are undergoing active oxidation. Thisarea is represented by the parallelogram ABCD, as shown on Figures 7 and 8. Theboundaries AB and CD are defined by equations pH=2.0 and pH=8.0, respectively. This pHrange is believed to cover most of the possible pH values in the drainage and the wastedumps at ICM.Line EF represents the line along which the Eh-pH measurements of natural waters appearmost frequently. It is the same line drawn on Figure 8 by Garrels and has an empiricalequationEh = 0.863 - 0.0592 pH^ (6)The lower boundary line AD is defined by the reduction of free oxygen to peroxide, reaction(3), in the form ofEh = 0.682 - 0.0592 pH^ (4)- 82 -The use of equation (4) as the lower boundary of measured natural Eh-pH conditions issubjective, thus somewhat arbitrary, although one can justify such a choosing on the basisof Figure 7 in that this line actually does represent a lower limit for the occurrence ofnatural Eh-pH measurements.The upper limit, line BC, is also chosen subjectively and can be justified on the basis ofFigure 7. Line BC is chosen such that it is parallel to line EF and is the same distanceabove line EF as line AD is below line EF. The equation for line BC is thereforeEh 1.044 - 0.0592 pH^ (7)In the sections to follow, we superimpose the defined parallelogram ABCD on some of theEh-pH diagrams of our interests to make observations and to draw thermodynamicconclusions.^SulfurThe Eh-pH diagrams for the system S-O-H at activities of dissolved sulfur of 10 -2 and 10-3are presented in Figures 12 and 13. The sulfate activity of the typical drainage is 4.67x10 -3 .Sulfite, thiosulfate, and other sulfur species rarely encountered in natural waters are notconsidered in the construction of the Eh-pH diagrams. Note that the only differencebetween the two diagrams is that the solid phase of elemental sulfur field is slightly biggerin the more concentrated solution.- 83 ->_c 0.2w0.0-0.2-0.4-0.6-0.8 0^2^4^6^8.^1 0^12^14pHFigure 13 Eh-pH Diagram for Part of the System S -O -H. Theassumed activity of dissolved ES=10 2 .- 84 -1 0I^I^I^1^I^12^4^6^8pH1.2SYSTEM^S-O-H25°C, 1 bar1. 0.8 0HS 1-12 14Fig-Limn Eh-pH Diagram for Part of the System S-0-H. Theassumed activity of dissolved ES=10 -3 .- 85 -From Figures 12 and 13, it can be seen that in the Eh-pH conditions we can possiblyencounter in ICM drainage (defined by the parallelogram ABCD), the predominant speciesis SO4-. Sulfuric acid is a very strong diprotonic acid, its first dissociation constant is verylarge, therefore in the pH range 0-14, aqueous H 2SO4 does not appear. In very reducingconditions close to the lower water stability limit, reduced sulfur species are stable andexist as aqueous H2S or HS, depending on pH. The sulfide ion (S=) field lies to the rightof that of HS when pH is greater than 14. In anoxic environment such as the bottom ofstratified lakes, sulfate can be reduced to H 2S or HS. This reduction can be greatlycatalyzed by sulfate-reducing bacteria. Sulfur species play a very important role ingeochemistry. The stability of various sulfides and sulfates will be discussed with differentmetal elements.^CarbonFigure 14 presents the Eh-pH diagram of C-O-H system with a total dissolved carbonactivity of 10-3. The total activity of carbon species in the typical drainage is 1.11x10 -3(Table 18).It can be seen from Figure 14 that in the Eh-pH conditions we are likely to encounter(parallelogram ABCD), the predominant species are H 2CO3 and HCO3. Only when pH isgreater than 10.3 is CO3= the predominant species. This configuration of the Eh-pHdiagram of carbonate species (i.e., both 11 2CO3 and HCO3 are available in the pH range ofnatural waters (pH 5-8) in comparable concentrations) renders H 2CO3 the most important- 86 -HCO33■^I2 -CO3C91.— 0.2_c0.0-0.2-0.4-0.6- 0.8012^4^6.^8^10^12^14pHSYSTEM^C-O-H2 5°C, 1 bar -Figure 14 Eh-pH Diagram for Part of the System C-0 -H. Theassumed activity of dissolved EC=10 -3 .- 87 -weathering agent of igneous silicate minerals in the nature. Carbon dioxide (CO 2) from theatmosphere or from soil horizons dissolves in water that subsequently penetrates igneousformations. Part of the CO 2 dissolved forms carbonic acid (H 2CO3). When igneous silicateminerals weather, they extract 11+ from H 2CO3, converting H 2CO3 to HCO3. The 11 1- ionsare used to displace alkaline or alkaline earth metal cations, which are released intosolutions; and to combine with the residue to form new minerals such as clay minerals. Inturning H2CO3 into HCO3 -, igneous silicate minerals have neutralized part of the aciditystored in carbonic acid. The weathering of silicate minerals by carbonation is an acid-consuming process. In cases where large concentrations of free 1-1 1- are present (low pH),such as the EMO drainage at ICM, silicate minerals need not extract W from carbonic acid,they instead utilize the readily available free II+ to displace 1C+, Na+, Me. In doingso, they have consumed the free W acidity in the water.Figure 15 is a logarithmic concentration-pH diagram for 10 -3 M (activity) carbonic acid. Itshows the variations of activities of various species in the system with pH. Comparing thisdiagram with Figure 14, we can better understand what the predominant fields of speciesmean. Note that the predominant field of a species A does not mean that it is the onlystable species in the solution. It only means that species A has the greatest activity of allstable species and is therefore the most stable species. In the predominant field of speciesA, away from the field boundaries, species A which we label the field with is in muchgreater activity (concentration) than all other species. Close to the boundaries, thispredominance decreases until the activity of species A is equal to that of another speciesB on the AB boundary.- 88 -I^I^1 6 8^10pH12 14Figure 15^Logarithmic Diagram for 10 -3 M Carbonic Acid- 89 -Across the boundary is the predominant field of species B, which has the greatest activityof all species in the solution, including A. Species A still exists but not in the greatestactivity. Another important point to remember about Eh-pH diagrams is that both Eh andpH are logarithms of activities (that is, pE=-log(ae_) and pH.-log(aH,)). Therefore, when wemove along either axis, the actual changes in activity of 11 -f ore- are an exponential function(with a base of 10) of the unit changes on the scale. The activities of dissolved specieschange in a similar manner, that is to say, when we move away from the A/B boundary intothe predominant field of A, the activity of species B will diminish exponentially. Except forareas very close to boundaries, the activity of A is usually several orders of magnitudegreater than that of all others - species A predominates its stability field.The above points can best be illustrated by comparing Figures 14 and 15. Within the waterstability field, from pH 0.0 to 6.4 is the predominant field of H 2CO3. At pH 2.0, which isfar away from the H2CO3/HCO3 boundary, the activity of H 2CO3 is 10 -3, that of HCO3 10 -73(4.3 orders of magnitude less) and that of CO3 104" (12.7 orders of magnitude less).Moving toward the H2CO3/HCO3 - boundary, at pH 5.0, the activity of H2CO3 is 10-302 , thatof HCO3  10'33 (1.36 orders of magnitude less) and that of CO 3= 109' (6.67 orders ofmagnitude less). At the H 2CO3/HCO3 - boundary, the activity of H2CO3 equals that of HCO3(both 0.5x10 -3, or 10 -3), the activity of CO3 is 10 72. From pH 6.4 to pH 10.3 is thepredominant ground of HCO3. At pH 8.3, approximately the middle of this field, theactivity of HCO3 is 10 -3, the activity of H2CO3 equals that of CO 3= (both 10-4.94 , 1.94 ordersof magnitude less). From pH 10.3 starts the stability field of CO 3=, in which the activity ofHCO3 is much less than that of CO 3=, and that of H 2CO3 is nearly zero.- 90 -^CalciumFigures 16 and 17 show the Eh-pH diagrams of the Ca-C-O-H-S system at two sets ofactivities: i) ICa=10 -2-5, ES=10-3, EC=10-3; ii) ECa=10 -2, ES=10 -2, EC=10-3. In the typical ICMdrainage, the activities are ECa=10 -2-3, ES=10 -2-33 , EC=10-2-".Calcium has two oxidation states: 0 and +II. The only species of Ca with an oxidation stateof 0 is metal Ca, which is not stable in the stability field of water. Only Ca(+II) appearsin the Eh-pH diagram of Ca.At the first set of activities, both Ca' and SO4 are under-saturated with respect to gypsumCaSO4 .2H20, therefore gypsum phase does not appear in the Eh-pH diagram. At thesecond set of activities, the solution becomes over-saturated with respect to gypsum whenpH is greater than 1.9, but at pH>7.4, gypsum is less stable than calcite, therefore gypsumoccupies the field between pH 1.9 and pH 7.4.Solubilization of Ca At ICM dissolved Ca2+ derives from two sources: weathering ofplagioclase and dissolution of calcite and/or dolomite. The weathering of igneous silicateminerals will be discussed in later sections. The fate of calcite depends essentially on threevariables of the solution with which it is in equilibrium: pH, concentration of Ca 2+, andconcentration of SO 4-. Larger concentrations of Ca', lower SO 4- levels and higher pHfavour the stability of calcite. At the conditions of the typical drainage, the solution isunder-saturated with respect to both calcite and gypsum, therefore calcite will dissolve and- 91 -SYSTEM^Ca-C-O-H-S25°C, 1 bar -El0.6^AL1. 0.4-0.6- 0.8 0N^ ., I'N. ..^-. .1. . IFCa2+^...^IIN^I.ACaCO34^6^8^10^12^14pHDFigure 16^Eh-pH Diagram for Part of the System Ca-C-O-H-SAssumed activities for dissolved species are:ECa=10 -2 ' 5 , ES=10 -3 , EC=10- 92 - 0.2- 0.4- 0.6- 0.8 0BEASYSTEM^Ca-C-O-H-S25°C, 1 bar _0•^N.0cf)0'ICD CaCO 32^4^6^8^10^12^14pHFigurel7 Eh-pH Diagram for Part of the System Ca-C-S-O-H.Assumed activities of dissolved species are: ECa=10 -2 ,ES=10 -2 , EC-10-3.- 93 -gypsum will not form. The CO; ions released from calcite add one W to form HCO3 ions,which add another II+ to form H 2CO3 , which is the most stable species and is in equilibriumwith atmospheric carbon dioxide (see Figure 14). Excessive CO3 will be eliminated fromthe solution through the above chain into the atmosphere as carbon dioxide. In the process,each CO; ion consumes two W ion to form one H2O. This is what "carbonate alkalinity"refers to. One interesting point to note is that the presence of sulfate ions causes calciteto act as an acid consumer at a higher pH. For example, at fixed total activities of C=10 -3 ,Ca=10-2 and S=0 in the solution, calcite becomes an acid consumer at pH<7.0, while in thepresence of total activity of S=10 -2, calcite becomes an acid consumer at pH<7.4. Thefundamental reason for this phenomenon is that the occurrence of gypsum field pushes thestability field of calcite to the right. It is obvious then that the solution must be over-saturated with respect to gypsum in order for the effect to appear.Removal of Dissolved Ca' The Eh-pH diagrams indicate that within the pH range 4-7,which covers the majority of ICM drainage, Ca2+ ions released either from plagioclaseweathering or calcite (or dolomite) dissolution only precipitate as gypsum when the sulfateactivity is greater than approximately 10 -2, which corresponds to a sulfate concentration of2056 mg/1 if the activity coefficient is assumed to be 0.467 (Table 18). This concentrationhas not been reached in ICM drainage streams except EMO drainage, which occasionallyhas sulfate concentrations greater than 2000 mg/l. Ca 2+ ions are not removed as calcitebelow pH^MagnesiumFigure 18 presents the Eh-pH diagram for part of the system Mg-O-H at activities ofdissolved EMg=10-2 ' -3. The EMg activity in the typical drainage is 10 -3.28 .The Eh-pH diagram of Mg is very simple. Like calcium, magnesium has only one stableoxidation state in the water stability field: +II. Due to the high solubility products ofmagnesium carbonate (1.0x10 -8 versus 3.84x10-8 for calcium carbonate) and magnesiumsulfate (3.09x10 -8 versus 2.5x10 -8 for calcium sulfate), at a total Mg activity of 10 -3, thepresence of S species in an activity less than 10-2 and/or C species in an activity less than10-2 will not produce new phases on the Eh-pH diagram within the pH range 0-14. Theysimply exist as aqueous species. The Eh-pH diagram we have presented here can actuallybe viewed as that of the system Mg-C-S-O-H as far as Mg species are concerned.Mg' ions are primarily released into solution by ferromagnesian silicate mineralweathering and perhaps by dolomite dissolution.For 10-3 total magnesium activity, Mg 2+ ions are removed from solution as hydroxide whenpH is greater than 9.8. 10-3 Mg activity corresponds to 45.85 mg/l. On the other hand, thehighest concentration recorded at ICM (that of EMO) is 82 mg/1, which gives an activity of1.79x103. For Mg2+ ions at this activity to be precipitated as hydroxide, the pH of thesolution at least has to be great than 9.4 (at which 10 -2 Mg starts to precipitate, see Figure18.). Since the pH of ICM drainage streams has never exceeded 8.6 (Table 17), we cansafely conclude that all the magnesium solubilized has been carried out in drainage streams- 95 - 0.4-0.6- 0.80SYSTEM^Mg-O-H25°C, 1 barB r.^ioo1^-'^,D^ 7i .E IN^..I .^. .-. .I^.^.A.L . c.^.N1. .dFMg 2+Dt lO OMg(OH) 2I I^I^►^1 2^4^6^8^10^12^14pHFigure 18 Eh-pH Diagram for Part of the System M9-0-H. Theassumed activity of dissolved EMT-10 2 ' .(excluding any adsorption onto solid and colloid surfaces). The removal of Mg' as sulfatesand carbonates is negligible.^CopperFigures 19 through 21 present the Eh-pH diagrams for the systems Cu-O-H, Cu-S-0-H, andCu-S-C-O-H. The assumed total activities are ECu=10-6 for the first system; ECu=10 -6,ES.10-3 for the second system; and ECu=10 -6 , ES=10 -3 , EC=10-1-3 for the third system. Inthe ICM typical drainage, ECu=5.0x10 -6, ES.4.67x10-3 and EC=1.1x10-3. Another elementthat greatly affects the Eh-pH diagram of Cu-S-O-H is iron (Fe). When Fe is introduced tothe Cu-S-O-H system, the Eh-pH diagram becomes extremely complicated, such as thosein Garrels, 1965, pages 231 and 231.For the system Cu-O-H, copper exhibits three stable oxidation states within the stabilityfield of water: +II, +I and 0. In the absence of S and C species, native (elemental) copperis a stable species. In the oxidizing-reducing transition zone, copper(+I) (cuprous) oxideCu2O (the mineral cuprite) is the most stable species in alkaline conditions. In the Eh-pHarea of our interest, parallelogram ABCD, the most stable oxidation state is +II. Below pH6.5, all 10-6 copper(+II) stays in solution as Cu' (cupric) ions. Above pH 6.5, cupric ions aretransformed to CuO (the mineral tenorite). However, the transformation from aqueous Cu 24-ions to CuO is not a redox reaction, Cu 2+ can not combine directly with free 02 ; it has to finda non-redox route to CuO. Actually Cu 2+ first combines with two OH - to form amorphous-97-1210I^I^I^1^I^12^4^6^8pHSYSTEM^Cu-O-H25°C, 1 bar -14Figwel9 Eh-pH Diagram for Part of the System Cu-O-H. Theassumed activity of dissolved ECu=10-6.- 99 - 0.2- 0.2- 0.4- 0.6-0.80^2^4^6^8^10^12^14pHFivre20 Eh-pH Diagram for Part of the System Cu-S-0-H. Theassumed activities of dissolved species are: ECu=10 -6 ,ES=10 -3 .2^4^6^8pHFigwe21 Eh-pH Diagram for Part of the System Cu-S -C -O-H. Theassumed activities of dissolved species are: ECu=10 -6 ,ES=10 -3 , EC=10-3.- 100 -cupric hydroxide, which dehydrates one water molecule to form CuO, as the followingreactions show:cu2+ + 20H - ---> Cu(OH)2 (amorphous) (very fast)^ (8)Cu(OH)2 (amorphous) —> CuO(crystal) + H 2O (slow)^ (9)The first reaction, precipitation of amorphous cupric hydroxide, does not occur until pH isgreater than 7.1, given a total Cu activity of 10 -6. Therefore, the conversion of Cu 2+ to CuO(namely the removal of Cu2+ from solution) can not occur until pH is greater than 7.1. Thefirst reaction is labelled "very fast" because it is an ion-ion reaction in solution. The secondone is slow because it is a conversion from one solid to another. Note that the conversionfrom Cu" to CuO is an alkalinity-consuming reaction and is favoured in higher pH.The reverse conversion, i.e., from CuO to aqueous cupric ions does occur at pH 6.5 as theEh-pH diagram shows through the reactionCuO(crystal) + 211+ —> Cu' + H2O (not very fast)^ (10)which is a hydrolysis reaction. It is also an acid-consuming reaction. It is not very fastbecause it is a surface reaction. It can be fast if the CuO crystal is finely powdered, whichusually does not happen in nature.- 101 -The introduction of S into Cu-O-H system completely alters the lower half of the Eh-pHdiagram, but the upper half remains identical (see Figure 20). The stability field of nativecopper becomes much smaller, and the predominant grounds of minerals chalcocite (Cu2S)and covellite (CuS) occur near the lower water stability limit. Since both of these stabilityfields are far removed from the parallelogram ABCD, which encompasses surface waterconditions, the sulfides are unstable at earth's surface. They oxidize to convert S(-II) toS(+VI) in SO4= and to convert Cu(+I) as in Cu2S to Cu(+II), either as aqueous Cu" or asCuO depending on pH values. The oxidation of copper sulfides obeys mechanisms similarto those of the oxidation of pyrite in general, that is to say, copper sulfides are oxidizedeither by free oxygen 0 2 or by in-solution oxidizing agents such as Fe', both are oftenpromoted by bacteria. The oxidation of pyrite is explained in detail in the literature review.Without going into details, we write the overall oxidation reactions as follows:Chalcocite:Cu2S + 2 1/202 + 2H+ --> 2Cu2+ + SO4= + H2OCu2S + 41120 + 10Fe' ---> 2Cu2+ + SO4= + 10Fe" + 8H+^(12)Covellite:CuS + 202 —> Cu' + SO4=^(13)CuS + 8Fe' + 4H 20 —> Cu2+ + SO4= + 8Fe' + 8W^(14)Note how the above oxidations differ in the view point of acid-base balance: the oxidationof covellite by oxygen (in air) has no effect on acid-base balance, the oxidation of chalcociteby oxygen consumes acid, and the oxidation of both by Fe 3+ generates acid. (This is a very- 102 -superficial statement and not exactly true. See the discussion on the oxidation ofchalcopyrite below for explanation). Other primary copper sulfides may also occur inigneous formations, they will oxidize in a manner similar to chalcocite or covellite.The oxidation of copper sulfides is always undergoing at surface conditions therefore therelease of copper is continuous. Like the oxidation of pyrite, the oxidation of copper sulfideis a surface reaction thus is generally slow. The oxidation rate is highly dependent of thesurface area to mass ratio and the total percentage of copper sulfide in the waste dump.The concentration of cupric ions in drainage is a result of the balance between oxidationrate, frequency and amount of precipitation, equilibrium control, etc.In Figure 21, carbon is added to the system Cu-S-H-O. In the new diagram, the lower part(reducing conditions) remains unchanged, the upper part is reshaped with the newmalachite (Cu2(OH)2CO3) phase replacing tenorite (CuO). Malachite is an oxidized coppermineral formed at or near earth surface conditions. At a total Cu activity of 10 -6, it is stableat pH>7.2 but dissolves at pil_7.2 to yield Cu 2+, hydroxide alkalinity and carbonatealkalinity. Malachite acts as an acid neutralizer at a higher pH (7.2) than tenorite does(6.5). At plI11.5, malachite dissolves to form CuO2 in a highly alkalinity consumingreaction.In the Eh-pH diagram of Cu-Fe-S-H-O system (not provided in this thesis but can be foundin Garrels, 1965, pages 231 and 232), at very reducing conditions very close to but withinthe lower water stability limit, one can find a variety of stable sulfides of copper and/or irondepending on pH. These include, approximately in the order from low to high pH,- 103 -chalcocite Cu2S, covellite CuS, pyrite FeS 2, Cu9S5, Cu5FeS4, chalcopyrite CuFeS 2, and theoxide magnetite Fe304. It is interesting that the oxide magnetite can form at such highlyreducing conditions. In the area defined by parallelogram ABCD, the predominant stablespecies are aqueous Cu', Fe' and Fe' at very acidic pH values; aqueous Cu' and solidFe2O3 (or Fe(OH) 3) at less acidic pH values; and solid CuO and Fe 2O3 (or Fe(OH)3) at pHvalues approximately ranging from neutrality to very high (between 12 and 13). At evenhigher pH values, CuO starts to dissolve to yield Cu0 2=, a soluble copper species. Thisstructure of Eh-pH diagram is expected from our knowledge of the Eh-pH diagrams of thesystems Cu-S-H-O and Fe-S-H-O.Upon exposure to earth's surface conditions, the various sulfides will oxidize to a commonproduct SO4-, as well as dissolved metal ions at lower pH or metal oxides (or other formssuch as malachite) at higher pH. Magnetite will slowly transform to hematite. Among theoxidation reactions is that of chalcopyrite as discussed below.The oxidation of chalcopyrite by oxygenCuFeS2 + 402 --> Fe' + Cu' + 2SO4= (slow)^ (15)does not produce acid itself, but the oxidation followed by hydrolysis of the ferrous ionsdoes:Fe' + 21H20 + 1402 ---) Fe(OH)3 + 2H+^ (16)- 104 -If microbiological oxidation is pronounced, Fe' will be oxidized by bacteria, mainlyThiobacillus ferrooxidans, to Fe:Fe' + 1/402 + H+ --> Fe' + 1/21120^ (17)(by bacteria)When the activity ratio of Fe 3+/Fe' in solution is great enough, the oxidation of chalcopyriteby Fe' takes place significantly:CuFeS2 + 8H20 + 16Fe3+ ---) 17Fe' + Cu' + 2SO4= + 16H+^ (18)Note that although this reaction shows 16 H+ ions being generated, they will be consumedwhen 16 Fe2+ ions of the 17 Fe2+ ions are oxidized by bacteria back to Fe' via reaction (17).Upon a complete circle consisting of reactions (17) and (18), the net result (the sum ofreactions (17) and (18)) is reaction (15): one CuFeS 2 is broken down to one Cu 2+, one Fe'and two SO4=. It is the Fe2+ generated in this process that gives up two II+ ions (acid)through reaction (16). The role of bacteria and the 16 Fe2+/Fe3+ is exactly that of a catalyst- speeding up the reaction but yielding no net change in themselves. From the sameperspective, the oxidation of chalcocite and covellite by Fe', when mediated by bacteria,will not generate any acidity, but rather the overall result would be that delineated by theequations (11) and (13).Solubilization of Cu At ICM, secondary oxidized copper minerals such as tenorite ormalachite are absent. So neither can copper be released from the hydrolysis of oxidized- 105 -Cu(+II) minerals nor can alkalinity or acid-consuming capability derive from them. Theaqueous Cu2+ ions come from the oxidation of primary sulfide minerals, the most abundantof which is chalcopyrite. The breakdown of copper sulfides or copper-iron sulfides yieldcupric ions, ferrous ions and sulfate ions. The oxidation of chalcocite consumes acid; thatof covellite neither consumes nor generates acid; and that of chalcopyrite generates acid.The acid is generated by oxidation followed by hydrolysis of aqueous Fe' ions to solidFe(OH)3, which is the short term, metastable compound of iron at earth surface conditionswhen pH is approximately greater than 3.0 and which will slowly dehydrate to transformto hematite in the long run (see next section for detailed discussion of iron species).Removal of Dissolved Cu' In an short term, at a Cu' activity of the order of magnitudeof 10", all cupric ions will stay in solution when pH is less than 7.0, even though they mightbe over-saturated with respect to tenorite or malachite. At greater pH values, part of thedissolved Cu2+ will be removed from solution as hydroxide. In a long term and at the sameactivity level, all Cu2+ will remain in solution at surface conditions if pH is approximatelyless than 6.5. At greater pH, part of the Cu' will be removed either as tenorite or asmalachite, depending on the total activity of carbon species. The short term removal of Cu 2+has importance in considering chemical treatment plant as a pollution abatement optionwhile the long term removal may be of significance in evaluating passive treatment ofdrainage.The copper activity of 10' translates to a concentration of 128 pg/l. This value has beenexceeded in NDD and EMO drainage by a factor of 10-20 for some of the high-concentrationdrainage flows. For example, the highest concentrations in NDD and EMO for the period- 106 -January, 1988 - January, 1989 are 2400 pg/1 and 1200 pg/l. For these concentrations ofdissolved copper to be stable in solution, the pH must be 0.3-0.5 unit less than pH 7.0. Thisproblem is examined again in the section on solubility control. IronFigures 22 and 23 provide the Eh-pH diagrams for the Fe-O-H and Fe-S-O-H systems ata total dissolved activity of EFe=10-6 and at total dissolved activities of EFe=10 -6, ES=10-3,respectively. In the typical drainage at ICM, EFe=5.0x10 6, ES=4.67x10 3 .The element iron has three oxidation states: +III, +II and 0. The oxidation state 0 isunstable in contact with water thus does not appear in the Eh-pH diagrams in the stabilityfield of water. Fe(+III) is stable at relatively oxidizing conditions while Fe(+II) is stable atrelatively reducing conditions.The introduction of S into the Fe-O-H system only changes the lower part of Eh-pH diagramnear the lower water stability limit, the upper part is unaffected. The upper parts aredifferent in the two diagrams because of the different choices of the stable solid iron species.In Figure 22, Fe(OH)3 and Fe(OH)2 are chosen as stable species while in Figure 23 hematite(Fe2O3) and magnetite (Fe304) are selected. Since hematite is thermodynamically morestable than Fe(OH)3 , it pushes the Fe3+/Fe203 boundary, the Fe2+/Fe203 boundary and the- 107 -—0.8 0^2^4^6^8pH1410 12SYSTEM^Fe —0—H2 5°C, 1 bar _Fe 022-Figure22^Eh-pH Diagram for Part of the System Fe-O-H AssumingFe(OH) 3 as Stable Fe(+III) Phase. The assumedactivity of dissolved EFe=10-6.-108-1.21.011•■•••••Fe 3+0.8 0.6 —- 0.4 —Fe304- 0.6 —F.S-0.80^2^4^6^8pH14Fe 2+0.4F• 2++-S°SYSTEM Fe-S-0-H25C4 1 barIFNNNA.C1 203NFeS 2-0.2 ---Figure23 Eh-pH Diagram for Part of the System Fe-S -O -H AssumingHematite as the Major Fe(+III) Phase. The assumedactivities of dissolved species are: EFe=10 -6 , ES=10-3.- 109 -D-c 0.2Fe304/Fe203 boundary outwards to occupy a larger territory. One can construct an Eh-pHdiagram of Fe-S-O-H system with Fe(OH)3 and Fe(OH)2 being the major Fe(+III) and Fe(+II)solid phase species by superimposing the lower part of Figure 23 onto Figure 22. Figure23 can be used to assess the dissolution of the minerals hematite and magnetite for variousEh-pH conditions. Figure 22 is suitable for evaluating the precipitation of dissolved ironfrom solutions, because observations have shown that when the pH of a solution containingdissolved iron is raised, Fe3+ and Fe' will be removed as Fe(+III) and Fe(+II) hydroxides(usually amorphous), not hematite or magnetite (both usually crystalline). This can beexplained as follows: The formation of Fe(OH)3 and Fe(OH)2 is simply the combination ofFe' and Fe' ions with OH - ions in solution due to electrostatic attraction; all three ions arehighly dispersed hence the reaction is very fast. However, Fe3+ and Fe2+ can not formhematite or magnetite directly, because firstly they are already in oxidized forms thus cannot combine with free dissolved oxygen for which an electron donor must be available inorder to form ion oxides, secondly a free oxide ion 0= is not present in naturalenvironments. They must extract 0= from some species that contains it: This is usually OH -(or H2O) in natural environments. Fe' and Fe 2+ combine with OH - to form hydroxides:Fe' + 30H - -+ Fe(OH)3 (solid)^ (19)Fe' + 20H- —) Fe(OH) 2 (solid)^ (20)Reaction (20) takes place when pH is approximately greater than 8.8 and the Eh is very lowat an activity of dissolved Fe 2+.10 -6 and is favoured by higher pH values. In contrast,- 110 -reaction (19) occurs at a pH as low as 3.1. In fact, reaction (19) is written for alkalinesolutions where OH- activity is reasonably high. In acidic solutions where OH - activity isextremely low, Fe 3+ hydrolyses to break down water into H+ and OH -, OH- combines withFe' while H+ is left behind in the solution, giving rise to acid mine drainage. Reaction (19)thus becomes:Fe' + 3H20 Fe(OH)3(solid) + 3H+^ (21)In the pH range from 3.1 to 8.8, Fe2+ is reduced to ferric hydroxide:Fe' + 21/2H20 + 1/402 ---> Fe(OH)3(solid) + 2H+^ (22)Eventually, however, given enough time, Fe(OH) 3 and Fe(OH)2 will transform to hematiteor magnetite depending on Eh-pH conditions. The transformation of ferric hydroxide issimply a dehydration process:2 Fe(OH)3(solid) ----> Fe203(hematite) + 3H 20^ (23)The direct dehydration of ferrous hydroxide yield ferrous oxide FeO (the mineral wuestite):Fe(OH)2(solid) —> FeO + H2O^ (24)Ferrous oxide is not a stable species (see Figures 22 and 23), thus in nature the above routeis not taken. Ferrous hydroxide either oxidizes completely to hematite at relative oxidizingconditions:2Fe(OH)2(solid) -4 Fe203(hematite) + H 2O + 2H+ + 2e^ (25)or oxidizes partially to magnetite at relative reducing conditions:3Fe(OH)2(solid) -3 Fe304(magnetite) + 21120 + 2H+ + 2e -^(26)Reactions (25) and (26) are only half reactions, they require electron acceptors to complete.When free oxygen (whether gaseous or dissolved) are available, oxygen acts as the electronacceptor:1/202 + 2H+ + 2e - --> H2O^ (27)The coupling of the half reaction (27) with half reactions (25) and (26) gives rise to:2Fe(OH)2(solid) + V202 —> Fe203(hematite) + 211 20^ (28)3Fe(OH)2(solid) + 1/202 -f Fe304(magnetite) + 311 20^ (29)However, in the natural environments where Fe(OH) 2 is formed, the condition is oftenanoxic, that is, without free oxygen. An example of anoxic conditions is that of the bottom- 112 -of a stratified lake where biological activities are abundant. At such conditions, some otherelectron acceptor other than free oxygen must be available to facilitate the transformationsrepresented by reactions (25) and (26). The oxidized species in the redox pairs that appearin Figure 10 may act as such electron acceptors at appropriate Eh-pH and activity settings.For example, if ammonium ion NH 4+ is in significant concentration (such as due to thebiological activity at the bottom of an anoxic lake), it can be reduced to nitrate to take upelectrons:NH4+ + 3H20 + 8e ---> NO3 - + 10H+^ (30)Coupling of this half reaction with half reaction (26) yields:12Fe(OH)2 + NH4+ —> 4Fe3O4 (magnetite) + NO3 - + 51120 + 1811+^(31)One more possibility is the direct oxidation of ferrous ions by oxygen to hematite ormagnetite:2Fe2+ + 21120 + 1/202 ---> Fe203(hematite) + 411+^ (32)3Fe2+ + 3H20 + 1/202 --> Fe304(magnetite) + 6H+^ (33)Reactions (32) and (33) are unlikely to take place in nature to any significant degree in ashort term, because the formation of crystalline hematite and magnetite is an extremely- 113 -slow process. Much faster precesses, such as that represented by reaction (20), (22) or theoxidation of Fe' to Fe' in solutionFe' ---) Fe' + e^ (34)predominate in nature. Note that when reaction (34) is coupled with reaction (27), reaction(17) results, but not necessarily with the involvement of bacteria all the time.From Figure 22, it can be seen that in the parallelogram ABCD, which represents thepossible conditions in surface waters, the most stable species of iron can be Fe", Fe 2÷, andFe(OH)3 . The Fe(OH)3 field occupies the majority of the area in ABCD, thus it is likely themost stable species to form in ICM drainage flows. As has been discussed above, althoughFe' should transform to hematite at pH values approximately greater than 1.3 accordingto Figure 23, this does not take place in a short term for kinetic considerations. Instead,Fe' precipitates as ferric hydroxide at pH values roughly greater than 3.1 at 10 -6 total Feactivity. The ferric ion activity in even the most acidic drainage (that of EMO with a pHof 4.2, see Table 18) should be three orders of magnitude less than 10 -6 (since [Fe'HOHT= constant, OH- activity one order of magnitude greater corresponds to Fe" activity threeorders of magnitude less), namely, 10 -9, which corresponds to 0.215 pg/1 assuming theactivity coefficient of Fe 3+ to be 0.26 (computed from Debye-Hiikel formula). Theconcentrations in other drainage streams with higher pH values are virtually zero. Theabove, however, are thermodynamic conclusions and should be applied with caution. Theabove rules are expected to be followed in ordinary conditions due to the rapid rate ofhydroxide precipitation except where the bacteria Thiobacillus ferrooxidans are extremely- 114 -active. The bacteria are believed to be able to maintain the Fe' concentration somewhathigher than that dictated by thermodynamic equilibrium.In the Eh-pH range defined by parallelogram ABCD, dissolved ferrous ions at an activityof 10-6 starts to transform to ferric hydroxide (reaction (22)) at pH values ranging from 3.1to 5.4 depending on Eh values of the solution. The appropriate pH of the precipitation ofFe' can be assumed to be 3.9, that along the line EF in Figure 22, due to lack of Ehmeasurements in the ICM drainage. Under this assumption, all 10 -6 activity of Fe" (112pg/I in concentration) stays in solution below pH 3.9, but at pH 5.0 the activity of Fe 2+ willdecrease to approximately 10-8, which corresponds to 1.12 pg/I. However, reaction (22)involves the oxidation of Fe(+II) to Fe(+III), therefore is much slower than the formationof ferric hydroxide from ferric ions, which involves only electrostatic attraction. Somedeviation in Fe' activity from the thermodynamic equilibrium is expected in reality. Thisis why in the ICM drainage concentrations of Fe 2+ as high as 830 pg/I have been detected(see Table 17).Figure 23 demonstrates that pyrite occupies most of the area near the lower water stabilitylimit. Pyrite is the most widespread primary sulfide in the world. Since the stability fieldof pyrite is too remote from the area defined by ABCD, which covers ordinary surfaceaerated water Eh-pH conditions, pyrite is not stable at earth's surface. Almost any oxidantsat earth surface conditions can play the role of oxidizing pyrite. The most common one isof course oxygen. According to Kleinmann, Fe" becomes the major oxidant in place ofoxygen when pli2.5 and Eh?_650 my (approximate) (Kleinmann, 1981). The oxidation ofpyrite is continuously undergoing at the earth's surface. It is clear that the rate of release- 115 -of iron and sulfur is not limited by thermodynamic equilibrium, but by kinetics of thereactions. The biochemistry of pyrite oxidation has been discussed in the literature reviewof this thesis. At ICM, the dissolved iron are determined by the reconciliation among therate of pyrite oxidation (which is determined by the total sulfur content, surface area tomass ratio of pyrite, bacteria activity, supply of water and oxygen, Eh-pH conditions, etc.),the frequency and amount of rain, and the rate of Fe 2+/Fe' removal by precipitation asferric hydroxide.Solubilization of Fe and S At ICM, the rate of release of Fe 2+ and SO4- depends on the rateof pyrite oxidation. The concentrations of Fe" and Fe' reflects the balance amongthermodynamic trends, bacteria activity, and kinetic considerations. Currently, the totalFe concentration in all the ICM drainage streams is basically that of Fe 2+, the concentrationof Fe+ is extremely low due to the Eh-pH conditions in the present drainage. Theconcentration of Fe2+ (hence the total Fe concentration) is often higher than that governedby the equilibrium because the removal of Fe2+ via reaction (22) is slower than thegeneration of Fe' by the oxidation of pyrite, which in practice has bee proven to berelatively fast, especially when bacteria is actively involved.The steady-state cycling of Fe2+/Fe" by bacteria oxidation of Fe 2+ and pyrite reduction ofFe' generate neither net Fe' or Fe' nor net acidity. It is the oxidation of one unit ofpyrite FeS2 , no matter abiotic or biological, that liberates two sulfate ions, two 1-1+ and oneFe2+ ion, which upon hydrolysis generates another two W ions (reaction (22)). The kineticsof pyrite oxidation will be further discussed in the section headed by "kinetic considera-tions".- 116 -Removal of Dissolved Fe The state of existence of iron mobilized from pyrite depends on theEh-pH conditions of the solution (see Figure 22). Assuming that line EF represents theaverage Eh-pH conditions that are likely to occur in the nature, we have calculate theequilibrium activities of dissolved iron species in the pH range 2-9. Note that thecomplexation of dissolved iron species with dissolved sulfur species is not considered. Theresult of the calculation is shown in Figure 24. It can be seen that at pH values greaterthan 2 the dissolved ferrous iron is so dominant that the total Fe activity is indiscerniblefrom the Fe2+ activity on the plot except near pH 2. The activities of dissolved iron speciesincrease exponentially with the lowering of pH. For example, the total dissolved iron willbe less than 10' in activity (112 pg/1 in concentration) when pH is greater than 3.9. At pH2, the drainage solution can hold 2.88x10 -3 Fe3+ (620,000 pg/1 in concentration) and 7.50x10 -3Fe2+ (840,000 pg/1), the ferric/ferrous ion concentration ratio is roughly one. At this pH,however, the formation of soluble complexes (actually ion pairs) such as FeSO4+, FeSO4°,FeHSO42+ and FeHSO4+ becomes significant and increases as pH is further lowered due tothe increasing iron activity. The concentrations of these ion pairs are not limited byequilibria with solid phase Fe(OH) 3 but by equilibria with the ions they consist of.Therefore, the total dissolved iron concentration can be much higher than the sum of theequilibrium concentrations of ferric and ferrous ions. At pH 4.5, the activity of Fe' inequilibrium with Fe(OH) 3 is 10-713 (8.4 pg/1) and that of Fe' is virtually zero (1040). In moresensible but less accurate words, we can state that the thermodynamics of the drainagesystem is such that under pH 2.5, all the iron mobilized by pyrite oxidation stays in thedrainage; between pH 2.5 and 4.5, the iron released by pyrite is partially precipitated; andabove pH 4.5, virtually all the iron is removed as ferric hydroxide. Two points should beemphasized, however. The first is that it is assumed that the occurrence of natural Eh-pH- 117 -9876543- 4- 6- 10-12- 142pHFigure 24^Activities of Fe Species in Equilibrium with Fe(OH) 3- 118 -conditions follows line EF; deviations from line EF (as natural waters always do) lead tosomewhat different answers, but the methodology still applies. The second is that these arethermodynamic (equilibrium) conclusions.Data collected at ICM suggest that the drainage are over-saturated with respect to ferrichydroxide. According to the thermodynamics, the total iron concentration should not bedetectable in all drainage streams but EMO which frequently had pH values below 4.5.This deviation from equilibrium, as has been mentioned above, is due to the kinetics: Theiron release is faster than its removal. We can use our understanding of the system inpractical applications. One example is that drainage samples for chemical analysis shouldbe acidified upon collection to maintain highly oxidizing and acidic conditions to preservethe dissolved iron in solution (although other considerations also apply). This has beenalmost a standard practice for environmental sampling. Another example is that almostall the dissolved iron (and some other dissolved metal species) can be removed from adrainage simply by aging in a pond, although deliberate aeration can accelerate the process.Once a drainage water is no longer in contact with pyrite, the supply of fresh Fe' iseliminated, thermodynamic equilibrium will be reached in time so that Fe' will be removedas Fe(OH)3 .^LeadFigure 25 displays the Eh-pH diagram for the system Pb-S-C-O-H at assumed activities ofdissolved species EPb=10 -", ES=10 -3, EC=10 - 3. The same activities in the ICM typical- 119 --120-0^2^4^6^8^10pH12^ 0.2- 0.2-0.4- 0.6- 0.8Figure25 Eh-pH Diagram for Part of the System Pb-S-C-0-H. Theassumed activities of dissolved species are:EPb=10 -8 ' 6 , ES=10 -3 , EC=10-3.Pb304PbOMC 03SYSTEM Pb-S-C-O-H25°C, 1 barPb0 2A■ IFIP bSAPb't-- ---._^Idrainage are EPb=4.8x10, ES=4.67x10 and EC=1.11x10 (Table 18). Pb level in thedrainage to be released to the environment is regulated at ICM to be less than 50 pg/l.With a few exceptions, the Pb levels in ICM land drainage are within the permit value.Pb is a very heavy (molar mass 207 g/mol), environmentally toxic, group W metal element.Its toxicity to man is cumulative. It has four oxidation states: 0, +II, +III, +IV. In thepresence of S, Pb(0) is not stable in contact with water thus does not appear in the stabilityfield of water. At EPb= 10-8, Pb2+ occupies a large field up to pH 8.1. Between pH 8.1 and9.4, PbCO3 (the mineral cerussite) is the stable phase provided the condition is notextremely oxidizing. When pH is greater than 9.4, PbCO 3 transforms to PbO (massicot).At very oxidizing, near-neutral to alkaline conditions, Pb(+II) is oxidized to Pb(+III) asPb304 (minium) and Pb(+IV) as Pb02 (plattnerite). Another Pb(+II) species PbSO4(anglesite) which would be stable at acidic, oxidizing conditions does not occur in the Eh-pHdiagram because EPb is too low. At higher EPb and/or ES activities, the P13 2+/PbCO3boundary will shift to the left by a great deal, and anglesite will occur in the diagram whenEPb is high enough (approximately 10). At reducing conditions, PbS (galena) is the stablespecies in the pH range 0-14. Within the parallelogram ABCD, the most stable species isPb2+.Lead can occur in nature as galena, anglesite, cerussite, massicot, minium and plattnerite,depending on the conditions of formation. ICM rocks are igneous thus we expect theprimary sulfide galena to be the dominant Pb mineral. Galena probably occurs inassociation with other sulfides such as pyrite and chalcopyrite, which occupy a similar fieldin the Eh-pH diagram. Galena is unstable at surface oxygenated conditions as designated- 121 -by parallelogram ABCD, therefore it oxidizes readily. Like the situation with pyrite, therelease of Pb" and SO4= is not equilibrium-controlled, but kinetics-controlled. It isnecessary to differentiate two kinetics measured on different scales. First of all, the word"kinetics" means in chemistry the study of the rate of a chemical reaction, the ratecontrolling factors, the rate's mathematical expression (the reaction orders), the reactions'selementary reactions, the reaction's activation energy and the rate's temperature response.In the context of acid mine drainage study, the kinetics either can be measured as the rateof the release of reaction products (such as Fe' or Pb") per unit weight of pure solidreactants (such as pyrite or galena) (the first one); or can be measured as the rate of therelease of reaction products to each unit volume of drainage (the second one). Thedifference between the two is that the second is affected by all the factors that affect thefirst one, as well as the over-all content of solid reactants (such as pyrite), the frequencyand amount of flushing water. We use the second interpretation of the word kinetics. Theconcentration of a reaction product (such as Fe" or Pb") is usually not directly related tothe kinetics of the reaction, since once a reaction product is released into the drainage, itcan be removed by other mechanisms that are not part of the kinetics of the reaction thatrelease the product.According to Figure 25, the concentration of Pb" in equilibrium with cerussite is 10' inactivity (4.3 lag/1 in concentration) at pH 8.1, and 10' (431 lig/1) at pH 6.5. Similarcalculations result in the following: 10' (1.6 pg/1) at pH 8.6, 10 -6' (50 pg/1) at pH 7.2.Solubilization of Pb The solubilization of Pb from galena is kinetics-controlled. Anexamination of the drainage pH values and Pb" concentrations in the ICM drainage- 122 -streams that derive directly from the North Dump reveals that the Pb 2+ in all drainageflows except NDD and TED are under-saturated with respect to cerussite. Thus the EPbin all drainage streams except NDD and TED represents 100% of the Pb mobilized fromgalena. The slow release rate (hence the low EPb concentration) in drainage is due to thescarcity of lead minerals in the ICM waste rock, because if the slow rate of Pb release weredue to other limiting factors other than the low galena content, we would expect Pb releaserate to increase in strongly acidified areas such as EMO where the release rates of othermetals (such as Fe, Mn, Zn) have greatly increased. But in fact the Pb 2+ level in EMO isapproximately the same as other drainage streams (except NDD). The only drainage thatcontains EPb in concentrations greater than the permit level (50 pg/1) is NDD, whichusually contains about ten times greater EPb concentration than other streams. The originof the NDD dissolved lead can almost certainly be traced to the Old Marginal Dump, whichis very likely to contain much higher percentage of galena than average waste rocks dueto its higher degree of sulfide mineralization. The PIP - in NDD is sometimes over-saturatedwith respect to cerussite. So is the Pb2+ in TED to which NDD reports along with the Pb 2+it carries. The over-saturation is possible for kinetic reasons, as has been discussedpreviously.It seems safe to concluded therefore that EPb does not constitute a concern at ICM (exceptNDD), since the limiting factor responsible for low Pb concentration is the availability, thusthe low Pb release rate can not be significantly accelerated by further acidification. TheOld Marginal Dump deserves some attention since it is a warehouse of Pb for NDD.- 123 -Removal of Dissolved Pb At ICM, the concentration of EPb has been very low, mostlywithin the range 0.5-7.5 pg/l. At such low concentrations, PIP will neither be removed ascarbonate (EC= 10-3) below pH 8.0 nor be removed as sulfate (ES=4.7x10 -3). To reduce thedissolved EPb below 50 pg/1 (current ICM permit level) by cerussite precipitation, a pHgreater than 7.2 and time to allow equilibrium to be reached is required. The removal ofPb' as hydroxide is discussed in the solubility control section.^ManganeseFigures 26 and 27 present the Eh-pH diagrams of the systems Mn-O-H and Mn-C-S-O-Hat an assumed activity of EMn=10 -" and activities EMn =10 -6, EC=10-3 and ES= 10 -3 ,respectively. The activities in the ICM typical drainage are EMn=5.0x10 -5, ES=4.67x10 -3 andEC=1.1 lx10 -3 (Table 18).Mn is a very active transition metal. Its oxidation states are 0, +II, +III, +IV, +VII. Theoxidation states 0 and +VII are not stable in the water stability field thus do not occur inthe Eh-pH diagrams. The Mn2+ field occupies much of Eh-pH diagram in both figures fromacidic to relatively alkaline conditions except at high Eh. In parallelogram ABCD, Me isthe most stable species except the most oxidizing and alkaline corner. Under oxidizing andvery alkaline conditions, Mn forms oxide minerals of oxidation states +IV (Mn0 2, pyrolusite,todorokite, etc.) and +III (Mn203). At transitional Eh and very high pH, the stable phaseis Mn304 (hausmannite), an oxide of mixed oxidation states (+III and +II). At reducing andalkaline conditions hydroxide and carbonate are formed. The sulfide of Mn (MnS,- 124 - 0.20.0-0.2-0.4- 0.6- 0.80 2^4^6^8^10^12^14pHF4we26 Eh-pH Diagram for Part of the System Mn-O-H. Theassumed activity of dissolved EMn=10 -6 .- 125 ->0.2w0.0- 0.2-0.4-06- 0.80 2^4^6^8^10^12^14pHFigure27 Eh-pH Diagram for Part of the System Mn-C-S-0-H. Theassumed activities of dissolved species are: EMn=10 -6 ,ES=10 -3 , EC=10 -3 .-126-alabandite) occupies only a very small field. This explains the rare occurrence of alabanditein the nature.The nature of Mn-bearing minerals at ICM can not be determined from availableinformation with certainty. The oxidation of MnS neither generates nor consumes acid:MnS + 202 —> Mn2+ + SO4^ (35)The reactions that release Mn2+ from all other minerals that occur in Figure 27 are acid-consuming (or alkalinity-generating):Mn02 + 4H+ + 2e - —> Me + 21120 (36)Mn203 + 6H+ + 2e —> 2Mn2+ + 311 20 (37)Mn304 + 8H+ + 2e —> 3Mn2+ + 4H 20 (38)MnCO3 + H+ -4 Mn2+ + HCO3 (39)HCO3 + H+ —> H 2CO3 (40)>^+I^CO2(g)^H2OMn(OH)2 + 2H+ -f Me + H 2O (41)- 127 -Of these reactions, (36), (37) and (38) are reduction half reactions. Since the EMnconcentration in the ICM drainage is increasing very rapidly recently, it is worthwhile toexamine the role of above acid-consuming reactions in the overall acid-base balance of thedump-drainage system if the major Mn mineral at ICM were oxides or carbonate orhydroxide. For convenience, let us assume the major Mn mineral is rhodochrosite (MnCO 3)and take EMO drainage as an example. Other minerals are similar. The average EMnconcentration in 1988 in EMO drainage is 4600 pg/1 (8.37x10' molar concentration), thehighest EMn is 25,000 pg/1 (4.55x10-4). The pH range of EMO in the same period is 4.2-4.8,within which H2CO3 (aqueous) is the predominant carbonate species. Each MnCO 3 releasesone Mn2+ and consumes two II+ via Reactions (39) and (40). Thus each mole of MnCO 3 isequivalent to one mole of CaCO 3 with respect to acid consumption. Using the total annualdrainage volume 243,600 m 3 and the total tonnage 4,670,000 tonnes for EMO, the alkalinityreleased by MnCO 3 is 2,034 kg CaCO 3 (using average EMn) and 11,084 kg CaCO 3 (usinghighest EMn). These are equivalent to 4.36x10 -4 kg CaCO3/tonne and 2.37x10-3 kgCaCO3/tonne. The total annual release of alkalinity (used in consuming acidity) is 300tonnes CaCO3 in 1988 (Appendix II). The percentages of the alkalinity released by MnCO 3are 0.68% and 3.7% respectively. These percentages are just marginally significant.It is also possible that Mn appears as silicates (such as spessartite, Mn3AL2Si3O 12) or asreplacement constituents of more complicated silicate minerals (such as metamorphicmineral tourmaline), or as constituents of other minerals. We will not discuss theseminerals in further detail except noting that the release of Mn from silicates is usually non-redox, slow, incongruent dissolution processes that consume acidity.- 128 -An interesting feature of the Mn-C-S-O-H system is that the stable phases at alkaline, fromvery oxidizing to transitional Eh conditions are all at a higher oxidation state than Mn'.These oxides therefore will act as oxidizing agents from near-neutral to acidic conditionsin the same Eh range, as illustrated by Reactions (36), (37) and (38). The consequence ofthis is that, if the Mn oxides exist in appreciable quantities in the waste rocks, theoxidation of sulfides continue even oxygen is not present, since the Mn oxides replaceoxygen as the major oxidants. This may take two routes: First, if water is available, watermay be oxidized to release oxygen:2H20 4W + 02 + 4e^ (42)The coupling of reactions (36), for example, with (42) yields:2MnO2 + 4W 2Mn2+ + 02 + 2H20^ (43)The oxygen so produced will be transported to other locations to be used to oxidize sulfide,for instance, pyrite:FeS2 + 33402 + 31H 20 ---> Fe(OH)3 + 2SO4 + 4H+^ (44)The overall reaction, assuming all the oxygen produced in Reaction (43) is used in Reaction(44), isFeS2 + 71/2Mn02 + 11H+ —> Fe(OH) 3 + 71/2Me + 2SO4= + 41120^(45)- 129 -It can be seen from Reactions (42)-(45) that water is only needed to initiate the reaction.The overall reaction generates water so the water is self-sufficient. The problem is 11+ ions,for each FeS2 oxidized consumes 11 II+ ions. If there is no net acidity input into the system,the acidity is soon exhausted, raising the pH until the Mn oxides are stable thus no longerreduced to Mn2+.In the second route, the Mn oxides may oxidize sulfide directly (also take the reactionbetween pyrite and Mn0 2 as an example) in one single step. The reaction is the same asReaction (45), but without water and oxygen being electron transferring media. Althoughthe limitation imposed by acidity availability still applies, the more stringent requirementis imposed by the reaction itself: It is a solid-solid-solute electron transfer reaction thattakes geologic time to occur substantially, thus is not considered in our study.The above discussion has potential practical significance. For example, if the waste rockcontains Mn02 (which could even be intentionally added), after the dump is sealed off fromair, the pyrite oxidation (which does not produce but rather consumes acid in the absenceof oxygen, Reaction (45)) may continue for a short time, during which the pH quickly picksup, then the oxidation will stop completely. The Mn0 2 function as a "negative feedback":the faster the pyrite is oxidized by Mn02, the faster the pH rises, and thus the sooner theoxidation stops. Mn oxides are as efficient acid consumers as CaCO 3 in that each Me orCa' released to a solution will eliminate two II+ from it. The difference between limestoneand Mn oxides is that the breakdown of Mn oxides (except MnO) is a redox process whilethat of CaCO3 a simple dissolution process. The idea of using Mn oxides as acid productionabatement agents has not been seen elsewhere. Since the above discussion is largely based- 130 -on thermodynamics, the feasibility of this idea needs to be evaluated based on kineticstudies.Solubilization of Mn In the parallelogram ABCD, Me is the most stable species. At ICM,the EMn concentrations in all the drainage streams are under-saturated with respect to Mnoxides. The concentration of Me therefore is kinetics-controlled rather than equilibriumcontrolled. The identities of the major Mn-bearing minerals have not been determined.Removal of Dissolved Mn Me is a quite soluble species and is not removed from the ICMdrainage as hydroxide or carbonate (see Section 5.3.3).^MolybdenumFigure 28 presents the Eh-pH diagram for the system Mo-S-O-H at assumed dissolvedspecies activities of EMo=10 -8, ES=10 -3. The activities in the ICM typical drainage areEMo=8.3x10 -8, ES=4.67x10 -3 .Molybdenum is one of ICM's revenue-generating metals. The mine effluent to be releasedto the environment is required to contain less than 500 i e by the environment permit.So far, the EMo levels in the drainage streams emanating from the North Dump have beenfar below the permit level.- 131 -1210 . 14SYSTEM Mo-S-O-H2^4^6^8pH0Figure28 Eh-pH Diagram for Part of the System Mo-S-O-H. Theassumed activities of dissolved species are: EM0=10 8 ,ES=10-3.- 132 -The metal Mo (oxidation state 0) is not stable in contact with water. The oxidation statesthat occur in the water stability field are +IV, +V and +VI. The Mo(+VI) occupies themajority of the Eh-pH diagram as two oxo-anions. The sulfide MoS2 (molybdenite) is thestable species at reducing conditions. At transitional Eh and low pH conditions, Mo(+V) asMo02+ is the stable species. The mineral ilsemannite (Mo308, a Mo(+V)-Mo(+VI) mixture,or Mo203 •Mo03) occupies a triangle area near the MoS2 field, between Mo(+V) and Mo(+VI).The major Mo-bearing mineral at ICM is certainly molybdenite. In the parallelogramABCD, the stable species is the soluble oxo-anion HMo04. In the activity range of ourinterest, EMo is obviously under-saturated, therefore the concentration of EMo is controlledby the kinetics of MoS2 oxidation. The low concentrations of 1Mo in ICM drainage areobviously limited by the low molybdenite content in the waste rock, just like the situationwith lead (see Section The EMo in the drainage will not be elevated by furtheracidification, as proved by the low EMo level in EMO drainage. Therefore, EMo does notconstitute an environmental concern at ICM. ZincFigure 29 presents the Eh-pH diagram of the system Zn-C-S-O-H at the assumed activitiesEZn=10 -", EC=10 -3, ES=10-3. In the typical ICM drainage, these activities are 5.0x10',1.1x10-3, 4.67x10 -3, respectively.- 133 -zn022-lott0co01.— 0.20.0-0.2-0.4-0.6= 0.8 0Figure 29SYSTEM^Zn-O-H-S-C25°C, 1 bar -B 0EIs'■AZn2+fI^1^ 1^1 2^4^6^8^10^12^14pHEh-pH Diagram for Part of the System Zn-C-S-O-H.assumed activities of dissolved species are:EZn=10 -4 ' -6 , ES=10 -3 , EC=10 -3 .0 FZnCO3----44DZnSlcThe- 134 -Although zinc only exhibits one oxidation state Zn(+II) in nature, the system is affected byEh because of the stability of sphalerite ZnS. Sphalerite is the most stable species betweenthe lower water stability limit and the sulfate-sulfide boundary at pH>2.1 for 10 -6 EZnactivity or at pH>1.14 for 10 -4 EZn activity. Zn2+ is very soluble and occupies a largestability field from very acidic to near-neutral conditions above the sulfate/sulfide boundary.From near-neutral to alkaline conditions and above the sphalerite field, the stable Zn(+II)species are ZnCO 3 (smithsonite), ZnO (zincite) and the soluble complex ZnO 22-. TheZnCO3/ZnO boundary is located at pH 8.1 and is not dependent of EZn activity. At 10 -4 EZnactivity, ZnCO3 occupies a narrow field between pH 7.5 and 8.1 while Zn022- predominatesat pH>13. At an 10-6 EZn activity, the Zn2+/ZnCO3 border shifts to approximately pH 8.6(to the right of the ZnCO 3/ZnO boundary), therefore ZnCO 3 is unstable. The Zn 2+/ZnOborder appears at pH 8.8. The ZnO/Zn0 22- boundary shifts to the left to pH 12. In theparallelogram ABCD, assuming EZn activity to be 10 -4, Zn2+ is stable except at pH valuesgreater than 7.5.Zinc commonly occurs in nature either as the primary sulfide sphalerite or as secondaryminerals such as smithsonite and zincite. Other rarer zinc minerals also exist in nature.At ICM, the major zinc-bearing mineral is probably sphalerite due to the igneous natureof the formation. Like calcite, smithsonite and zincite behave as acid consumers belowapproximately neutral pH:ZnCO3 + 11+ ---> Zn2+ + HCO3^(46)ZnCO3 + 2H+ -3 Zn2+ + H2CO3^(47)1I----> CO2 + H2O- 135 -ZnO + 2 H+ --) Zn 2+ + H2O^ (48)The oxidation of sphalerite neither produces nor consumes acid:ZnS + 202 -4 Zn2+ + SO42-^(49)The oxidation of sphalerite may also be facilitated by ferric ions with the involvement ofbacteria in a fashion similar to the bacteria-enhanced oxidation of pyrite. An examinationof the pH values and EZn concentrations recorded for the drainage streams around theNorth Dump reveals that all the drainage streams are under-saturated with respect toZnCO3, ZnO and Zn(OH) 2 , even for the EMO drainage whose dissolved EZn once peaked to36,000 pg/1 (2.75x10 -4 M activity). This fact implies that the concentration of EZn in theICM drainage is not controlled by equilibria between dissolved EZn and secondaryprecipitates, neither is it controlled by the equilibrium between dissolved EZn and theprimary sulfide sphalerite, which requires the total destruction of sphalerite. The EZnlevels are limited by the kinetics of the oxidation of sphalerite.The concentration of EZn has rapidly increased lately. The levels of EZn are in the sameorders of magnitude as those of EMn. It is estimated that if the major zinc-bearingminerals were secondary minerals (smithsonite or zincite), the alkalinity released by thedissolution of such minerals would have accounted for 0.5-5% of the total alkalinity releasedin the EMO drainage in 1988. But the major zinc-bearing minerals are more likely to beprimary minerals, especially sphalerite. If this is true, the solubilization of zinc will haveno influence on the acid-base balance of the drainage, but will contribute a small percentage- 136 -of sulfate to the drainage. For example, in 1988 the EMO drainage contained 500-2,200mg/1 (5.2x10' - 2.3x10-2 mo1/1) sulfate (Table 17) and 1,000-36,000 pg/1 (1.53x10 -5 - 5.5x104mol/1) dissolved EZn. That is to say, approximately 0.3-2.4% of the total sulfate derivedfrom sphalerite. The largest contribution to total sulfate by the oxidation of Mn-bearingminerals would occur when we assume all the dissolved EMn derived from MnS (which isvery unlikely). This percentage is 2.0% for EMO drainage in 1988. The input of sulfatefrom sources other than pyrite, sphalerite and alabandite is insignificant. Therefore, thecontribution to total sulfate by sulfides other than pyrite is at most 5%. This provides thejustification for an assumption we use elsewhere in this thesis for calculation purposes. Theassumption is that all the sulfate in the drainage derives from the oxidation of pyrite.Solubilization of Zn The concentration of EZn is controlled by the kinetics of sphaleriteoxidation. More advanced stages of pyrite oxidation will be accompanied by more advancedstages of sphalerite oxidation because of their common rate-limiting factors.Removal of Dissolved Ze So far in ICM drainage streams, all the zinc released byoxidation has been carried out, removal of dissolved zinc by precipitation is negligible. Thisis further discussed in the section on solubility control.^CadmiumFigure 30 presents the Eh-pH diagram of the system Cd-C-S-O-H for assumed activities ofECd=10-8, EC=10-3, ES=10-3. In the ICM typical drainage, these activities are: ECd=4.9x10 -8,EC=1.1x103, and ES=4.67x10-3 .The permitted level of ECd in the mine drainage to be released to the environment is 10lig/1 because of the environmental toxicity of cadmium. This level has been marginallyexceeded in NDD, TED, NDS, EMO, and EDD. The Eh-pH diagram of cadmium is verymuch like that of zinc (Figure 29). Cadmium has only two oxidation states in nature, Cd(0)and Cd(+II). CdS (greenockite) occupies all the area below the sulfate/sulfide boundaryexcept at very acidic conditions. Experience has shown that in nature Cd rarely forms itsown sulfide but rather occurs in other sulfide minerals such as sphalerite as impurities,perhaps because the element is too sparse in the crust. In oxidizing conditions, Cd 2+ reignsa large field from near-neutral to acidic conditions. In mildly alkaline to alkaline conditionsare the predominant fields of CdCO 3 and Cd(OH)2, which dissolves to yield the complex oxo-anion Cd022- at very alkaline conditions. In the parallelogram ABCD, Cd 2+ is thepredominant species.Solubilization of Cd At ICM, Cd released to the solution is not removed either as carbonateor as hydroxide. High concentrations of ECd are observed to almost always correspond tohigh concentrations of EZn, which can indicate an association of Cd with Zn minerals. Theconcentration of ECd is controlled by the kinetics of oxidation and at the same time issubject to the availability of Cd-bearing minerals.- 138 -1.2SYSTEM Cd-C-S-C-H25•C, 1 bar1.0Br`...,.0.8 — ^N N.^*El .N„^.I^N ,.^.I N^. .0.6 — I^2...Ai, Cd + .^IC'N.N.-,,^.^I^. .. .^I0.4^ ..^. 1Cd0 22 -7^ -P.,(SO42- )CdC O 3.c 0.2(42S)Cd (OH )2-0.2CdS-0.4-0.6I^I^1^I^I^12 4 6^8 10^12^14pHFigure 30 Eh-pH Diagram for Part of the System Cd-C - S -O-H. Theassumed activities of dissolved species are: ECd=10 -8 ,EC=10 -3 , ES=10-3.-0.8- 139 -Removal of Dissolved Cd Cd2+ is under-saturated with respect to either carbonate orhydroxide at ICM. To reduce the Cd 2+ concentration in the solution to under 10 pg/1 (permitlevel) as carbonate, a pH approximately greater than 8 is required (assuming dissolved ECactivity to be 10 -3). To achieve the same purpose by hydroxide precipitation, pH should begreater than^ArsenicFigure 31 presents the Eh-pH diagram for the system As-S-O-H at assumed activitiesEAs=10 -6, ES=10 -3. The corresponding activities in the ICM typical drainage areEAs=8.3x10-1°, ES=4.67x10 -3.Arsenic is very poisonous and is restricted at ICM by the permit to below 100 mil in thedrainage to be released to the environment. EAs levels in current ICM drainage flows arecommonly two orders of magnitude lower than the permit level. Arsenic is a metalloidelement and exhibits six oxidation states in nature: -II, -I, 0, +II, +III, and +VI. Theoxidation states -II and -I do not appear in the Eh-pH diagram. In reducing conditions Asoccurs as AsS (realgar), As2S3 (orpiment) and natural As. In approximately transitionalzone it exists as As203 (arsenolite), arsenous acid and its various dissociation species. Inoxidizing conditions, As occurs as arsenic acid and its dissociation products, all of which aresoluble. When iron is introduced into the system, new stable phases include FeAsS(arsenopyrite), FeAs 2 (loellingite) and Fe(As04 )-2H20 (scorodite). At ICM, arsenic probablyexists in various sulfide minerals that contains iron.- 140 -Figure 31 Eh-pH Diagram for Part of the System As-S -0-H. Theassumed activities of dissolved species are: EAs=10 -6 ,ES=10 -3 .- 141 -Solubilization of As At a EAs activity of 10 arsenic is totally soluble as arsenic acid or itsionization products. The EAs in the ICM drainage flows is well under-saturated. Thereason for this is obviously the scarcity of As-bearing minerals in ICM waste rocks.Therefore the advancement of acidification will not cause arsenic to be mobilized faster toelevate the EAs concentration in the ICM drainage, as demonstrated by the fact that EMOdrainage contains minimum arsenic despite its very acidic pH.Removal of Dissolved As At ICM, dissolved EAs is not removed by precipitation and neednot be removed to satisfy the permit since the EAs in the drainage is already so low andwill not rise significantly in the future. Dissolve /As does not constitute a concern at ICM.5.3.3 Solubility ControlSolubility product control, or more exactly equilibrium ion activity product control refersto the limitations set on the activities of ions in a solution by relatively insoluble compoundswhich are composed of these ions. For example, the equilibrium chemistry requires thatthe activities of Fe3+ and OH - in a solution must be such that their ion activity product[Fe31-] •[0111 3 should be less than or equal to 3x10 -39, the solubility product (or equilibriumion activity product) of Fe(OH)3 . One particular ion may be a constituent of many insolublecompounds and must obey the activity restrictions set by all such insoluble compounds. Themaximum allowable activity of an ion in a solution at thermodynamic equilibrium isdetermined by the compound that requires lowest equilibrium activity of the ion. When theactivity of an ion is smaller than the maximum allowable activity, the solution is under-- 142 -saturated with respect to all solid phases containing the ion. A solution under-saturatedin a certain ion can be in equilibrium when all solid phases containing the ion as a constit-uent are absent, otherwise the solid will dissolve to raise the activity of the ion, thus thesystem is not at equilibrium. A solution-solid system can be over-saturated in some ionsdue to kinetic considerations (such as extremely slow reaction rate, lack of crystallizationseeds, etc.). Such systems are never at true thermodynamic equilibrium.We will only consider the solubility control of ion activities by ionic or highly polar covalentcompounds whose elements do not change oxidation states upon dissolution. Thesecompounds will dissolve to yield the ions they are composed of in solution and are referredto as strong electrolytes. Highly insoluble electrolytes are expect to exercise major ionactivity control in many solutions where the total content of organic substances is low,partly because the precipitation of such solids involves only electrostatic attractions betweenions thus is rapid relative to redox reactions which involve electron transfer, and otherreactions. Oxides are not considered because the metal ions in basic oxides (such as CaO,Fe2O3 etc.) are most frequently removed from solutions as hydroxides (which are in facthydrated oxides and are considered) while acidic oxides are often weak electrolytes and donot include the metal elements in which we are interested. We have also excluded thecompounds whose dissolutions are redox reactions, since these have already been consideredin the previous sections. Complexation of ionic metal species with organic ligands alsoexerts an important, even primary control over ionic activities of metal ions as well as totalactivities of metals under certain conditions. This is treated in the next section.- 143 -The previous discussion of the Eh-pH diagrams has in fact dealt implicitly with thesolubility controls of many metal ions by hydroxides, oxides, carbonates (when carbonspecies are considered in the Eh-pH diagrams), and sulfates (when sulfur species are con-sidered). This section will view the solubility control systematically and find out the majoractivity controllers for some of the important metal species in the ICM drainage.Table 19 is a compilation of the solubility products of some strong electrolytes formed bythe metals of our interest and anions present in the ICM drainage. In Section 5.3.1, wementioned that the major anions in the ICM drainage include HSO4, SO:, 011 - , HCO3,CCV, Cl -, and perhaps NO 3- and the ionization products of phosphoric acid H 21304, HPO42-and PO43-. Of these, nitrates, chlorides, hydrogen sulfates and bicarbonates of all the metalswe are interested in are too soluble to commit any significant ion activity control, so theywill be dropped from the list. Although phosphate could impose control over some ion activ-ities when present in significant activities (concentrations), especially in alkaline conditions,we believe that the total phosphate concentration is too low in the ICM drainage to assumethis responsibility. In addition, at the pH range of our concerns, the predominantphosphate species are hydrogen phosphate ions and di-hydrogen phosphate ions whose saltsare usually soluble, and the activity of phosphate ion is at least four orders of magnitudeless than the total inorganic phosphate activity below pH 8.5. For example, let us make anapproximate calculation using Pb 3(PO4)2, the phosphate salt that has the lowest solubilityproduct (7.9x10 13) in Table 19. Assuming a total inorganic phosphorus activity of 10' M,at pH 8.5, the activity of PO,' would be 10 41. From [Pb2+ ]3 {PO4312 = 7.9x10-43, [Pb2+]--.2x10 -7. This activity is greater than the activity of Pb 2+ in equilibrium with lead carbonateat pH 8.5 (assumed activity of EC=10 -4, see Figure 35) so the activity of Pb2+ will be- 144 -Table 19^Solubility Products of Salts and Bases of Interest at 25°C and 1 atmAnionsCations Carbonates^Chlorides^Hydroxides^Sulfates^Sulfides^Phosphates*Ca 2. 3.84x10 9^3.63x10 ♦11 4.0 x10 -5 2.5 x10 5 d. 1^x10-26Mg^1^x10 -5^1.76x10432*^7.1 x10-12^3.09x10 -5*^d.^1^x10 -32Al 3+ N/A s.^3.5 x10 -34 v.s.^d.^5.8 x10-19Fe2+ 2.1 x10 -11^9.14x10 .7 8^x10-16 sl.s.(FeS0 4 'II20) 4.2 x1017 (FeS) N/C, ins.v.s.(FeS0 4 7H 20) (Fe3 (PO4 ) 2 '8H 20)8.3x10Fe3.^ +11*N/A^ 3^x10-39 v.s.^sl.d.^2.08x10-41*(Fe2 (50 4 ) 3 . 9H20)2.3 x10 -10^ -20Cu 2+ 6.82x10 +3* 1.3 x10^4.43x10+2* 6^x10-36 (CuS) N/C,ins.Cu '^N/C,ins.^1.2 x10-6 N/A d.^3^x10 48 (Cu 2 S) N/AZn 2+^1.7 x10 -11 1.14x10+7* 3.3 x10-13 1.23x10i3. 2^x10 -24 (ZnS) N/C,ins.Mn 2+^5^x10 -10^v. s.^2^x10-13^2.45x10 +2*^2.3 x10 13 (MnS) 1^x10-22Mo54 N/A d. 2.3 x10 *(Mo205 . 3H 20)N/A N/C,ins.(Mo2S 5 . 311 20)N/AMo 6 N/A^s. d.^N/C,ins.^N/A^N/C,s1.s.(MoS3) N/A(Mo03 2H 20)_8Pb2+ 7.4 x10 -14 2^x10-5 1.2 x10-15^ *1.82x10^1^x10 -28 (PbS) 7.9 x10 -43Cd2.^1.0 x10 -12^2.75 8.1 x10 -15^1.13 *^2^x10-28 (CdS) N/C,ins.Notes: 1. In constructing this table, the following references are consulted: Bailar et. al. 1984; CRC Handbookof Chemistry and Physics, 47th edn.; Dreyer, 1982; and Brookins, 1988. See Bibliography at the backof this report.2. Entries indicated with an "*" are computed from Standard Free Energies of Formation at 25 °C foundin the References listed above.3. Abbreviations used: v.=very, s.=soluble, ins.=insoluble, d.=decompose, sl.=slightly, N/A notavailable, usually because the substance does not form or is not stable, N/C=non-computable due to lackof thermodynamic data.4. Scope of phrasal expressions: very soluble = greater than 10 grams per 100 ml of water, soluble =between 1 and 10 grams per 100 ml, slightly soluble = approximately around 0.1 gram per 100 ml.- 145 -controlled by carbonate instead of phosphate. Therefore, the role of phosphates in solubilitycontrol is also insignificant at ICM. What we have left to consider are carbonates,hydroxides and sulfates. Some words about sulfides before we proceed. Sulfides of metalsare generally very insoluble in water (in the sense that they dissolve to give M" and S=, seethe solubility products in Table 19) except those of alkaline and alkaline earth metals,which are unstable in contact with water. At reducing conditions in the water stability fieldbelow the sulfate/sulfide boundary, the predominant species are aqueous H 2S at OR? andHS - at 7<pH<14 (Figures 12 and 13). The S= ions become the predominant species only atpH>14. Assuming a ES activity of 10' M, we find the activity of S= to be 5x10 -11 M at pH7. Even at such a small activity, sulfide is able to control the activities of some metals dueto the extremely small solubility products of these metals. For example, with a solubilityproduct of 2x10-28, sphalerite (ZnS) can have at most an activity of 4x10 -18. However, theabove picture changes completely in oxidizing conditions, i.e., above the sulfate/sulfideboundary. At the same ES activity (10' M) and pH (7.0), the predominant sulfur speciesis SO4= at pil>1.8. The activity of S= ions in equilibrium with the 10' M SO 4= at pH 7 iscalculated to be 1.47x10 5° M. At such an activity, S= will definitely not exercise any controlover any metal ion activities via the formation of sulfides.The equation of solubility product for sulfates [M 21 {SO421.1<si, can be written aslog[M21,-logK8  + p[SO4], thus the plot of log[M 2+] versus p[SO41(which is defined as-log[SO4]) is a straight line. Using the Ksi, values given in Table 19, we have constructedFigure 32. It turned out that only Pb 21-, Ca' and Mg2+ appear in the M2145041 range ofour interest. The straight lines represent the metal ion activities in equilibrium with theirsulfates.- 146 -01- 2- 37-40)0- 5- 6- 7- 80^1^2^3^4^5pSO4-- (= p(rotal S) at pH>2)6Figure 32 Logarithmic Concentration Diagram Showing the Activities of Some MetalIons in Equilibrium with Their Sulfates at 25°C.- 147 -For metal hydroxides, [M"].[OH1' 1 = Ksp, thus log[M°1 = logKsp - nlog[OH1, or log[Mn - ]logKsp + 14.n - n.pH. Using the K, 1, in Table 19, log-Nn-1 is plotted against pH in Figure33. A point on a particular straight line indicates the activity of the metal ion concernedin equilibrium with its hydroxide at the corresponding pH. Note the difference in slopebetween +2 ions and +3 ions.Figure 34 is constructed in the same way as Figure 32. Much more metal ions appear inthis figure than in Figure 32 because generally carbonates are less soluble than sulfates.Figure 34 may be difficult to use because the CO 3= is a function of both IC activity and pHdue to the equilibrium between CO3= and HCO3 and between HCO3 and H2CO3 (Figure 15).A diagram showing the activities of various metal ions in equilibrium with theircorresponding carbonates for a specific IC activity is more convenient. Figure 35 isconstructed for this purpose.To construct Figure 35, we need to choose a IC activity. For this reason, let us firstconsider the IC activity in a water in equilibrium with atmosphere at pH 7. The partialpressure of CO2 in the standard atmosphere at sea level is quite low, 0.0003 (or 10 -3.52) atm(which is also equal to the mole fraction of carbon dioxide in the atmosphere). Accordingto Henry's law, at equilibrium, the activity of carbonic acid (dissolved carbon dioxide) isaH2CO3 = PCO2.1<CO2, where Kco2=10-1.47 at 25°C. Therefore aH2c03 = 104' 10-5. The totalcarbon activity is calculated as:0-90^2^4^6^8pH10^12^14Figure 33 Logarithmic Concentration-pH Diagram Showing the Activities of VariousMetal Ions in Equilibrium with Their Hydroxides at 25°C.- 149 --10 ^-8^-6^-4log [003- -1Figure 34^Logarithmic Concentration Diagram Showing the Activities of Various MetalIons in Equilibrium with Their Carbonates at 25°C.- 150 -5^6^7^8^9^10^11^12pHFigure 35^Logarithmic Concentration-pH Diagram Showing the Activities of VariousMetal Ions in Equilibrium with Their Carbonates at 25°C.- 151 -amco3( 1+^+ K1 .1(2/aH+2)= 5.3x10-5= 10-4.28where Ki and K2 are the first and the second dissociation constants of carbonic acid andequal to 4.3x10-7 and 4.7x10 -11 at 25°C. The same calculation at pH 8 leads to a EC activityof 10-3.35 . Considering the flushing water in the dumps to be approximately neutral, thedumps to be well aerated and the flushing water is approximate in equilibrium with thepore air in the dumps, the EC activity in the flushing water will be in the order ofmagnitude of 10 -4. This activity value is assumed in Figure 35. In reality, the flushingwater could be much more acidic on one hand, which will reduce the EC activity in waterin equilibrium with atmosphere, on the other hand, the partial pressure of carbon dioxidein the pore air in the dumps could be much higher than that in the atmosphere due tobiological activities and/or release of carbon dioxide by the acid consumption reactions ofcarbonate (such as calcite), which tends to increase the EC activity. Therefore, the ECactivity in real situations probably fluctuates between 10 -2 and 10. The construction of adiagram similar to Figure 35 with a different EC activity involves only a verticaldisplacement of all the curves but that of CO3= by a distance determined by the differencebetween the new EC activity and 10 -4. For example, if a similar diagram is to beconstructed for 10 -3 EC activity, all the curves in Figure 35 simply shift down by one unit,except the curve of CO 3= which shifts up by one unit.Using Figures 32 through 35, we can determine the saturation state (i.e. over-saturation,under-saturation, or equilibrium) of a particular metal ion with respect to sulfate if we- 152 -know ES activity (calculated form concentration) and pH, or SO4 activity; or with respectto hydroxide if we know the pH; or with respect to carbonate if we know either CO 3= activityor the EC activity along with pH. We can also determine the conditions for precipitationof various ions as hydroxides, carbonates and sulfates.An comparison of Figure 32 with Figures 33 and 35 reveals that at the conditions typicalof the ICM drainage, i.e., ES=10 -2, EC=10 -3, the maximum activity of Ca' is controlled byits sulfate at pH<7.5 and by its carbonate at pH>7.5; that of Mg2+ by its sulfate at pH<9.7and by its hydroxide at pH>9.7; and that of Pb' by its sulfate below pH 6.3 and by itscarbonate above pH 6.3. If the activity of EC decreases to 10 -4 but other conditions remainunchanged, Me is not affected, but Ca 2+ and Pb2+ are. The pH values of the transitionfrom sulfate control to carbonate control become pH 8.5 for Ca' and pH 7.1 for Pip'. Theactivities of all other ions will be controlled either by carbonates or by hydroxides, but notby sulfates.In Figures 36 and 37, we have combined Figure 33 and 35 into one plot at two different ECactivities. The intersection point of the hydroxide line and the carbonate curve of aparticular metal ion defines the transition pH, above which its hydroxide controls itsactivity while below which its carbonate does. The information obtained in Figures 36 and37, along with that from the analysis of activity control by sulfates, are presented in Table20.Figure 38 summarizes the above discussions graphically. It shows in bold lines themaximum activities of various ions as allowed by their equilibrium with sulfates, carbonates- 153 -pHFigure36 . Logarithmic Concentration-pH Diagram Showing the Activitiesof Various Metal Ions in Equilibrium with Both Their Carbonates andTheir Hydroxides at 25 °C. Assumed activity of EC=10 - 4.2 1 0 128 1A0— 1-2-3C -4cp -50- 6-7- 8- 90PHFiglze37^Logarithmic Concentration-pH Diagram Showing the Activitiesof Various Metal Ions in Equilibrium with Both Their Carbonates andTheir Hydroxides at 25 °C. Assumed activity of EC=10-3.CaCO, MnCO,CuCO,Table 20^The pH Ranges of Metal Ion Activity Control by Sulfates, Carbonates andHydroxides (ES activity is assumed to be 10')IonsSulfate Ctrl. Carbonate Ctrl. Hydroxide CtrlEC=10-3 EC=10 -4 EC=103 EC=104 EC-10-3 EC=10 -4Ca 2+ 2-7.5 2-8.5 7.5-14.5 8.5-14.0 >14.5 >14.0Me 2-9.7 2-9.7 Nil Nil >9.7 >9.7Al 3+ Nil Nil Nil Nil All pH All pHFe Nil Nil <10.1 <9.2 >10.1 >9.2Fe Nil Nil Nil Nil All pH All pHCe Nil Nil Nil Nil All pH All pHZe Nil Nil <11.6 <11.1 >11.6 >11.1Mn2+ Nil Nil <10.7 <10.1 >10.7 >10.1Pb 24- 2-6.3 2-7.1 6.3-11.6 7.1-11.1 >11.6 >11.1Cd 2+ Nil Nil <11.4 <10.9 >11.4 >10.90)0-1-2-3-4- 5- 6- 7-82 64 8 10 12^14gCO 3CaCO,MnCO, CuCO, Fe CO,ZnCO,M SOCaSPbSO,-n••-g ^0pHFigure 38^Activity-pH Diagram Showing the Maximum Activities of VariousMetal Ions as Controlled by Their Equilibrium with Sulfates, Carbonatesand Hydroxidesat 25 °C. Assumed activities are: EC=10 -3 , ES=10-2.and hydroxides for a selected set of conditions. As a matter of fact, we can modify Figure38 for any combinations of EC and ES activities corresponding to those of particulardrainage steams , then plot the pH and the activities of various metal ions in the streamsonto the diagram to determine their saturation states.Applying the composition of the typical ICM drainage (Table 18) to Figure 38, we find thatall the metal ions are under-saturated except Ca', which is approximately in equilibriumwith gypsum (CaSO4).An examination of the analytical results of the EMO drainage samples in 1988 using thesame principles as above shows that the drainage is under-saturated in all the metalcations except Ca2+, which is over-saturated with respect to gypsum by a factor from 1 to3 in most samples and is slightly under-saturated in the rest of the samples. This mayimply that the sulfate level and the calcium level in the EMO drainage is controlled bygypsum precipitation. The major reason for the under-saturation of the metal cations is thelow pH in EMO (4.2-4.8), which raises the activities of cations in equilibrium with theircarbonates and hydroxides but has no effect on the activities of cations in equilibrium withsulfates approximately above pH 2.In a similar exercise performed on the water quality data of NDD in the period January -October, 1988, we have found the following: In cases where pH is near-neutral (6.5-7.5),the activities of all metal ions except Zn2+, Ca' and Mg' in the drainage are 1-3 (mostly2) orders of magnitude lower than the equilibrium activities (namely under-saturated). Zn 2+is also under-saturated but its activity is in the same order as the equilibrium activity.- 158 -Ca2+ and Mg' are under-saturated with respect to their sulfates by one order. In the onlycase where pH is 4.4, however, all the metals except Ca 2+, Mg2+, Pb2+ become under-saturated by 3-6 orders. Pb' and Mg2+ are under-saturated with respect to their sulfatesby one order of magnitude while Ca' is under-saturated by a factor of 2.The importance of pH on the saturation states of various metal ions in the ICM drainagecan best be demonstrated in Table 21, which is an analysis similar to the ones above carriedout on the TED drainage data collected in 1988. TED is the stream that has exhibited thehighest pH values in all the drainage streams originating directly from the North Dump.The following have been observed from Table 21:Table 21 State of Saturation of Metal Ions in TED 1988w.r.t. Sulfate w.r.t. Activity Controlling CompoundDate pH Su1ph. Ca Mg Pb Cu Fe Pb Mn Zn CdJan 88 8.1 U U-4 U-1 U-2 0+2 U-1 U-1 U-1 0+0 U-1Feb 88 7.1 U U-4 U-1 U-2 U-1 U-2 U-1 U-2 U-1 U-2Mar 88 8.1 U U-4 U-1 U-2 0+1 U-1 U-1 U-2 0+0 U-1Apr 88 8.4 U U-3 U-1 U-2 0+2 U-1 E U-1 0+0 U-1May 88 8.6 U U-3 U-1 U-1 0+2 U-1 E U-1 0+0 U-1Jun 88 8.5 U U-3 U-1 U-1 0+2 U-1 E U-1 E U-1Jul 88 8.0 U U-4 U-1 U-1 0+1 U-1 U-1 U-2 U-1 U-2Aug 88 8.1 U U-4 U-1 U-1 0+1 U-1 U-0 U-3 U-1 U-2Sep 88 4.5 U U-4 U-1 U-1 U-5 U-5 U-3 U-6 U-5 U-6Oct 88 6.8 U U-3 U-0 U-1 U-0 U-3 U-2 U-2 E U-2Note: U=Under-saturated, 0=Over-saturated, E=Approximate Equilibrium+2=Over-saturated by 2 orders, -3=Under-saturated by 3 orders, +0 or-0 = Over- or Under-saturated, but in the same order.w.r.t.= With respect to.- 159 -■ All the over-saturation and equilibrium states occurred above pH 8.0 (except onecase). In other words, below pH 8.0, all the metal ions are under-saturated, thushave not been precipitated prior to their appearance in the seeps.■ All the metal ions except Ca' and Mg2+ are dramatically much more under-saturatedin the only drainage with a pH of 4.4 than in all others whose pH values range from6.8 to 8.6. The equilibrium activities of Ca 2+ and Mg2+ ions are controlled by theirsulfates and do not change in the pH range 2-7.■ The only metal ions that achieved either equilibria or over-saturation are Cu', Pb 2+,and Zn2+, and they did so at pH values greater than 8. It appears that the activitiesof Pb2+ and Ze are controlled by their equilibria with their activity-limitingcompounds, since the two were either under-saturated, or at equilibria, or slightlyover-saturated, but never over-saturated by one order of magnitude or more. Cu 2+is different, however, in that it was over-saturated for many times by 1-2 orders ofmagnitude with respect to its equilibrium activity defined by its activity-limitingcompound. This is not surprising because it has been well documented that innatural waters some solutes are often over-saturated by a factor of 1-1000, but rarelyby a factor greater than 1000 (Dreyer, 1982). This phenomenon is again attributableto slow reaction kinetics.The conclusions are clear from the discussion of solubility control and the examination ofthe three representative streams at ICM (EMO - lowest pH and heaviest contamination,- 160 -NDD - heavy contamination in metals, pH near-neutral but less than 8.0, TED -intermediate metal contamination and wide pH variation 4.4-8.6):■ Below pH 8.0, no metal ions are removed from any of the ICM drainage streams.This is perhaps contrary to our chemical intuition but is theoretically and practicallytrue.■ The ability of the drainage in holding dissolved metal species (excluding Ca", Mg'and Pb' due to sulfate control) is greatly enhanced by decreases in pH. Between theapproximate pH range 6-8, the drainage is able to carry metal ions in activities 1-2orders of magnitude higher than the present levels. Below pH 6.0, the drainage'scapability of holding dissolved metal ions is virtually unlimited as far as themobilization of metal constituents from minerals is concerned.■ Above pH 8.0, there is probably some precipitation of Cu(OH) 2, PbCO 3 and ZnCO3in the dump or drainage courses to reduce the levels of Cu', Zn2+, and Pb2+ in theICM drainage.■ In acidic to approximately neutral pH values, the activities of Ca', Mg 2+ and PIPare controlled by the formation of their sulfates, as the sulfate activity is generallyquite high in the ICM drainage, although the activity of Pb2+ has never reached sucha magnitude as to be precipitated as lead sulfate despite the high sulfate activities(see Section In drainage streams such as EMO, when sulfate, calcium andmagnesium are high, the sulfate control mechanism will operate to adjust their- 161 -activities (concentrations). Gypsum (or jarosite if iron concentration is high and pHis low) can precipitate within the dump or in the drainage ditches.■ Below pH 10, the activity controlling compounds are carbonates rather thanhydroxides for most metals (excluding Cu"). To reduce effectively the activities(concentrations) of all metal ions in concern at ICM to satisfactory levels by theprecipitation of hydroxides, the pH has to be raised above 11.Having presented the principles of the mechanisms that affect the concentrations of variousmetals in a solution, we now use the element iron to exemplify how these principles operateinteractively to determine its actual concentration in drainage streams. All the concernedmetals should be examined in such a comprehensive way in order for their actualconcentrations in the drainage to be explained by chemical principles. Otherwise,conclusions based on incomplete considerations may be erroneous and misleading.Now assume redox reactions and simple ionic compound solubility control are the majormechanisms of controlling the activities of iron species in solution. Neglect the ion pairsFe(OH)+, Fe(OH) 2+, Fe(OH)2+, and others (they can be shown to be in negligible concen-trations), the iron species in solution are Fe' and Fe'. The processes in determining theactivities of ferrous and ferric ions in solution include:(1) The solubilization of iron from sulfides (essentially pyrite) by oxidation. Thisis a relatively slow, surface redox reaction. The reaction rate is determinedby a variety of factors as discussed previously. The reaction is verythermodynamically favourable. There will be no equilibrium between the- 162 -sulfides and surface oxygenated water before all sulfides are completely oxidized, sincesulfides are unstable at such conditions.(2) The control of Fe' activity by ferric hydroxide (Figure 33). The precipitationof ferric hydroxide is fast.(3) The control of Fe' by iron(II) carbonate below pH 10.2 (approximate) andferrous hydroxide above pH 10.2 (Figure 38). These reactions are fast.(4) The control of Fe' by Fe(OH) 3 (Figure 24). The reaction is slow because ofits redox nature.(5) The equilibrium between Fe" and Fe', which is dependent of the Eh of thesolution. The reaction rate is intermediate when abiotic and fast whenbiological.(6)^The control of both ferric and ferrous ions by hematite. This is ignored dueto the slowness of the reactions.The compounds formed in (3) (ferrous carbonate or hydroxide) are always only meta-stablewith respect to ferric hydroxide, but they do form due to their fast reaction kinetics whenconditions are appropriate. Now the activities of Fe 2+, Fe' and total Fe are determined asfollows. If the release of Fe (process (1)) is so slow that both Fe 2+ and Fe3+ are under-saturated with respect to processes (2), (3) and (4) above (note the ferric/ferrous ratio isdetermined by the Eh of the solution), no iron will be removed from solution. The activityof total iron is solely determined by the rate of iron solubilization while the activities offerrous and ferric ions are determined by the solution Eh. At pH 4.5 and Eh 0.6 V(approximate) which are typical for the EMO drainage, the above conditions are satisfiedwhen the Fe' activity is less than 10 -7 and Fe' activity is less than 10' (Figure 24). These- 163 -activities translate to concentrations 11 pg/1 Fe' and 0.02 pg/1 Fe'. At pH 6 and Eh 0.5,the same calculation leads to 10-101 Fe' in activity (0.009 pg/1 in concentration) and 10 -'3Fe' in activity (0.000001 pg/1). These conditions are certainly unrealistic to fulfil, even inthe EMO drainage (see Table 13), which is the most acidic of all the ICM drainage streams.The pH values of drainage have to be below 4.0 for the total iron concentration to be solelycontrolled by the oxidation of pyrite.Next, if the release of iron is fast enough so that the iron activity in the solution is over-saturated with respect to processes (2) and (4) but not (3) above, we would expect theactivity of Fe" to be roughly in equilibrium in solution with Fe(OH) 3 while Fe' is over-saturated to a certain degree with respect to Fe(OH) 3. At this time, Fe" will oxidize bothto Fe(OH)3 directly by oxygen and to Fe', which will be precipitated as Fe(OH)3. At pH 4.5,the concentrations of iron species in solution should be: Fe' approximately 0.02 pg/1, Fe 2+1.1 pg/1 - 11 g/1. It is clear at this pH Fe" - will not be over-saturated with respect to process(3). At pH 8.0, the concentration of Fe" should be virtually zero, and the concentration ofFe' should be below 450 pg/1.Last, if the iron release by pyrite oxidation is fast and the pH of the drainage is high, theiron in solution may be over-saturated with respect to processes (2), (3) and (4). Note thisover-saturation is impossible at low pH values. At pH 8.0, the over-saturation of Fe' withrespect to carbonate occurs when the release of iron exceeds 450 pg/l. The ferrous ions willbe removed as ferrous carbonate (total C activity is assumed at 10 -3).- 164 -Examination of the total Fe concentrations shows that all the ICM drainage streams areunder-saturated relative to carbonate and ferrous hydroxide, but most are over-saturatedwith respect to Fe(OH)3, even the most acidic EMO drainage. The aging of the ICMdrainage, therefore, should bring the total iron activity towards that defined in Figure Other ProcessesIn accounting the acid-base balance and the contaminant balance, processes other thanredox reactions and salt precipitation (both discussed above) are also involved. Theseinclude complexation, chelation, hydration, simple dissolution, hydrolysis, carbonation, ionexchange, adsorption-desorption, and co-precipitation. They are briefly discussed below.Complexation, Chelation and HydrationComplexation is the chemical association of a metal atom or ion with a ligand, both of whichare capable of independent existence, through coordination covalent bonds to form newspecies (complexes). A coordination covalent bond is formed between a donor atom or ion,which has an unshared pair of electrons, and an acceptor atom or ion, which has unoccupiedorbitals that can accommodate the electron pair. All metal ions have the ability to formcomplex ions. Those with small radius and high charge, especially those that also havevacant d orbitals - the transition metal ions - do so most readily. But even the sodium ion,with its comparatively large radius and its small charge, forms complexes. The differencelies in the stability.- 165 -Complexation is important in considering the contaminant levels of a natural water,because it can alter the total concentration of metals (which we measure in the laboratorychemical analysis) remarkably. If the complex formed is very soluble, the totalconcentration of the metal involved in the complex formation will be greatly increased. Thecomplex is a new substance thus does not have to satisfy the solubility and redoxrequirements that the metal has to before complexation. For example, the formation of thecomplex [Au(CN)2]2- brings about the total dissolution of otherwise insoluble elemental gold.The complexation of Zn2+, Cu2+ and Al' with OH - tremendously increases the totalconcentration of these metals in alkaline conditions relative to near-neutral conditions.However, if the complex formed is highly insoluble, the total concentration of the metalinvolved in the complex formation may be controlled by the dissociation constant of thecomplex rather than by the solubility products of the relatively insoluble compounds of themetal.Common inorganic ligands in nature are: F, C1, Br- , r, CN -, NCS-, NO2, oil-, CO;, SO4- ,SO3=, NH3, H2O, CO, NO. Many kinds of organic ligands exist which usually bond to metalatoms or ions through N atoms or 0 atoms.If an organic ligand has swveral electron donor atoms which are simultaneously bound toa central metal species to form a complex, the ligand is then called a chelating agent, or abi-, tri-, quadri-dentate according to the number of coordinating atoms in the ligand. Thecomplex so formed contains organic rings around the central metal species. The chemicalagent EDTA is an example of quadridentate. In nature, some plants and microorganismsare capable of excreting chelating agents, therefore the chelation process may be important- 166 -in the destruction of some minerals by bringing the metal constituents into solution bychelation. At ICM, we believe chelation play an insignificant role in the overall metallicspecies (contaminants) balance.Although water is a common ligand capable of complex formation, we rarely view theresulting species as complexes. The word "hydration" is used to refer to the association ofchemical species (such as ions, molecules, or salts) with water. Hydration can take placethrough two mechanisms: that caused by the electrostatic attraction between chargedspecies (such as Ca2+, Cu' and SO4=) and the water dipoles, and that caused by thecoordination covalent bonding between a metal species (such as Cu 2+, Fe2+, etc.) and watermolecules (as a ligand). The latter belongs to the category of complexation and is relativelystronger than the electrostatic hydration. In this regard, hydration and complexation areoverlapping concepts. When metal ions are precipitated from solution, all or part of thecoordinated water molecules will be retained and incorporated in the crystalline saltsformed, depending on the abundance of free, unbounded water molecules in the solutionfrom which the salts are precipitated. Such is the mechanism of the formation ofZnSO4 .7H20, ZnSO4 .6H20, CuSO4 6H20 etc. The crystalline water molecules in thesecompounds can be gradually driven off by intensive heating, first the relatively weaklybound ones at a lower temperature, then the more strongly bound ones at a highertemperature. The great availability of free water molecules in dilute solutions (such as theICM drainage) makes water the most competitive ligands to prevent the formation ofcomplexes of metal ions with other ligands which are usually at much lower activities. Thehydration of metal ions, either by electrostatic association or by complexing, usually doesnot alter the thermodynamic properties of the metal ions in solution, since it has been dealt- 167 -with in the formulation of activity coefficients and the equilibrium constants. Therefore,hydration does not comprise a concern in considering the metal species (contaminants)balance at the ICM drainage. There is another kind of hydration which involves thecombination of water into the crystalline structure by anhydrous minerals formed in theabsence of water. This process modifies the crystalline structure, but does not affect thesolution chemistry.Of the common ligands listed above, only Cl -, Off, CO3=, SO4 , and H2O are present in theICM drainage in detectable concentrations. Of these, C1, CO 3- and SO4 are weak ligandsand act as significant complexing agents only when they are very concentrated, which is notthe case in the ICM drainage. OH - does not form complexes significantly with the metalions of our interest below pH 9. H 2O has been discussed above. Therefore, complexation,hydration and chelation will not affect the consideration of metal species balance in theICM drainage. However, these processes do play a role in chemical treatment plants ofpolluted water. The pH of the precipitation tanks may be high enough to solubilize Zn andCu as complexes. Chelation may also increase the total concentrations of some metals.These factors should be taken into consideration in the design of the treatment plant if sucha option is selected.Simple Dissolution, Hydrolysis and CarbonationAgain, these are not mutually-exclusive, but rather overlapping concepts defined accordingto different criteria. Simple dissolution refers to the transfer of mineral constituents fromthe solid mineral phases to the solution as dissolved species without the involvement of- 168 -redox reactions. In congruent dissolution, all mineral constituents are transferred intodissolved species, such as the dissolution of rock salt (NaC1) and other brine minerals, ofgypsum, calcite and dolomite, and of some magnesian minerals. Incongruent dissolutionconverts some of the mineral constituents into dissolved species while leaving others in asolid residual mineral, such as the dissolution of most aluminosilicates in nature.Hydrolysis refers to the process in which amphoteric species (usually an ion) from mineraldissolution combine with either 11+ or 011 - to form a weakly dissociated species or a highlyinsoluble compound, resulting a shift in pH of the solution. The formation of Fe(OH) 3 fromFe' and gibbsite Al(OH)3 from the dissolution of aluminosilicates are examples ofhydrolysis. Carbonation is actually a hydrolysis process involving carbonate species. It isused to refer to the weathering of calcium-, magnesium- and iron-bearing rocks by theformation of carbonate or bicarbonate salts of these metal elements, which are relativelysoluble. Its use has largely been discontinued due to the ambiguity in its definition.The importance of simple dissolution of minerals lies in the fact that most igneous silicateminerals weather by this mechanism. Our primary interests in silicate mineral weatheringare how it affects the acid-base balance and the metal contaminant balance of the drainage.Most of the silicate dissolution processes are acid-consuming. In nature, silicate mineralsare commonly weathered by the attack of dissolved carbon dioxide (or carbonic acid), whichis universally available from the atmosphere and biological activities. The carbonic acid isconverted to bicarbonate ions (at near-neutral conditions) or carbonate ions (at alkalineconditions) with the release of II+ ions, which is necessary for the dissolution (oftenincongruent) of silicate minerals. It has been observed that the weathering of silicateminerals by dissolved carbon dioxide is a slow process and can be considered negligible in- 169 -contaminant release (if it were not, the majority of rivers and streams would have beenpolluted), and most probably also negligible in acid consumption, insomuch as acid minedrainage generation is concerned. However, theoretically the silicate weathering processshould be accelerated by the lowering of pH as much more Fr and acidity becomesavailable. 11+ penetrates the outer residual shell of silicates being weathered (as manyresearchers believe to be present) readily and diffuses to the fresh silicate surface rapidlydue to its tiny size. At ICM where some streams (such as EMO) have low pH values,whether the silicate weathering has been accelerated to such a degree as to deserveconsiderations in the acid-base balance and the metal contaminant balance still remainsto be investigated.Ion ExchangeAs the name states, this refers to the exchange of an ion in solution for an ion with thesame type of charge in a solid (often crystalline) structure. Ion exchange is often driven bya concentration gradient across the solution-solid interface. In the weathering of rocks, weare primarily concerned about cation exchanges. Within the soil layer penetrated by plantroots, cation exchange is often an important mechanism to deliver mineral nutrients toplants for their growth. In the process, the W ions present in plant roots exchange for ionssuch as K+, Na, Ca', etc. with the interstitial water, which may in turn exchange withsurrounding clay minerals. If the exchange chain travels far enough to reach some clayminerals in contact with unweathered silicate minerals, exchange may take place betweenthe two, thus accelerating the breakdown of the silicate minerals into clay minerals. Thenet result of ion-exchange between plant roots and silicate minerals is neither acid-- 170 -consuming or acid-generating, nor contaminant-releasing (although it does result in newmineral formation), thus is ignored in this thesis. The ion-exchange reactions betweensilicate minerals and fresh water is acid-consuming as the silicates absorbs H+ from waterto make the water alkaline. But this exchange only takes place in the initial stage whendry, fresh silicate minerals are brought into contact with water; it reaches saturation veryquickly as silicate minerals are in continued contact with water through the weatheringprocess. The net results brought about by this reaction is negligible in a long term.Adsorption, Desorption and CoprecipitationAdsorption is the attachment of dissolved species to the surfaces of solids, often colloids;desorption is the reverse process. They can be physical or chemical. Chemisorptions areactually surface chemical reactions while physical sorptions are mostly of electrostaticnature. In reality, these two are often indistinguishable. Coprecipitation occurs when adissolved species is incorporated as a minor component in a solid phase as that phase itselfis precipitated. Coprecipitation is often indistinguishable from adsorption thus sometimesis included by the word adsorption. In natural waters transition metals are stronglyadsorbed by oxides and hydroxides (particularly of aluminum, iron and manganese). Thisadsorption is most pronounced when the oxides or hydroxides are finely divided, e.g., ascolloids. It has been found that in some natural waters the adsorption of many transitionmetals by hydrous manganese and iron oxides exerts primary controls over theconcentrations of these metals (Dreyer, 1982). As a general rule, the adsorption oftransition metal ions on to oxides and hydroxides is most remarkable in alkaline conditions,decreases as pH drops, and disappears totally when pH is low enough. The transition from- 171 -high adsorption to non-adsorption occurs in a narrow pH range. This can be explained (atleast partially) by two arguments: first, at higher pH (greater than the point of zerocharge), oxides and hydroxides bear negative surface charges, which favour the adsorptionof positively charged metal ions or metal complexes which are positively charged; second,transition metals usually do not adsorb as metal ions due to their large hydrated shells, butadsorb as metal-hydroxy complexes (such as Zn(OH)+, Pb(OH)+, etc.) which only form whenpH is greater than a particular value for a particular metal and which do not possess alarge hydrated shell. The pH characteristic of the transition from non-adsorption to highadsorption is often the pH at which the metal-hydroxy complex becomes predominant.At ICM, the influence of the adsorption of transition metals by oxides and hydroxides isignored, primarily because we do not have any information on the composition of thesuspended solids and the colloidal portion of the drainage. We believe this will notintroduce significant errors into our conclusions regarding the solution chemistry due to thefollowing considerations:■ Some of the ICM drainage streams, such as EMO drainage which is our primaryconcern have low pH values (4.2-4.5 for EMO), which prohibits the adsorption.■ In the ICM drainage, the total suspended solids and perhaps colloids are relativelylow while the total dissolved metal species relatively high (such as Mn, Zn, Cu, andFe in EMO), so that even the solid surfaces are saturated with adsorption of thesetransition metals, the amount adsorbed would still account only a small portion ofthe total dissolved species. Therefore, if the total concentration of a metal is greatlyin excess of the natural (baseline) level, the concentration of the metal can not have- 172 -been controlled by its adsorption on to any solids, despite that it may be heavilyadsorbed.■^For those drainage streams whose pH values are high enough for the adsorption tooccur, we would expect that only the one or two most competitive metals areadsorbed significantly, while others are relatively unaffected.However, we can not rule out the possibility that the total concentrations of some metalsare controlled by their adsorption on to solid surfaces, particularly those whose totalconcentrations are very small.5.3.5 Bacterial ActionThe involvement of bacteria in the generation of acid mine drainage has been reviewed inSection 2.1 of this thesis. The goal of the present section is to evaluation the degree ofbacteria involvement in ICM drainage, with the emphasis being given to EMO drainage.Let us first discuss the relationship of drainage pH and interstitial pH within the porespaces in an acid-generating dump. The moisture retained in a dump between rain stormsacidifies quickly and experiences a great dilution when the next rain storm comes. Theinterstitial pH within the dump should be lower than the drainage pH before the acidityis diluted. But the interstitial pH can not be predicted from the drainage pH by simplyconsidering the dilution effect. For example, one can not conclude that the interstitial pH- 173 -is 2 units lower than that of the drainage purely based on a 100-fold dilution. Other factorshave to be considered, such as the buffering capacity of various minerals. For example,Fe(OH)3 starts to consume acid when pH reaches roughly 3.0, the interstitial pH at placeswhere Fe(OH)3 is present can not be much lower than 2.5 before Fe(OH) 3 is completelydissolved. Kleinmann found in laboratory experiments that interstitial pH values weretypically 0.2-0.5 below the corresponding drainage pH (Kleinmann, Crerar and Pacelli,1981). As to pH control, he also demonstrated that the drainage pH of a dump is controlledprimarily by the role played by bacteria, presumably Thiobacillus ferrooxidans. Heproposed the three-stage acidification model to characterize the acid generation from pyrite-containing materials and divided the three stages at pH 4.5 and 2.5, which were consideredto correspond to major changes in the role of bacteria (see Section 2.1). Others havedemonstrated the step-like control of drainage pH in the course of acidification of pyriticdumps which contain readily available neutralizing materials (SRK and Norecol, 1989). Inaddition, each rain storm disrupts the lowering of pH in pore spaces such that at the endof each rain event, the pH of the pore spaces which have been flushed will approach thatof the drainage. It seems fair to say that the interstitial pH values are normally less thanthat of the drainage but the difference is small. The value of this statement is that we canuse the pH of the drainage to represent that of the pore spaces without incurring a largeerror.In Section 5.3.2, we assumed that the redox potential of natural waters is controlledprimarily by the redox pairs H 20/02 and H202/02. However, Kleinmann presented evidencesto show that the Eh seems to be dominated by the Fe 2+/Fe(OH)3 redox pair in many AMDenvironments (Kleinmann, Crerar and Pacelli, 1981). He also showed that increased- 174 -bacteria involvement in the acid production will cause the Eh of the system to rise, the pHto decrease and the total Fe concentration to increase.The total dissolved iron and pH values in the EMO drainage all indicate that EMO iscurrently in stage II of the acidification process in view of Kleinmann's three stage model.Bacteria are involved in both the direct oxidation of pyrite and the oxidation of Fe' to Fe".The contributions of chemical oxidation and biochemical oxidation of pyrite are bothsignificant in this stage. The steady state cycling of ferrous and ferric ions by bacterialoxidation of ferrous ions to ferric ions and pyrite reduction of ferric ions to ferrous ions hasnot been reached. In other words, the bacterial involvement in pyrite oxidation has notreached its maximum. If there is enough pyritic material available, the bacterial oxidationwill increase and the pH will decrease, provided the acid consuming potential will begradually exhausted. The onset of stage III is signalled by an enormous increase in totaldissolved Fe (to hundreds to thousands of mg/1), by a high ferrous/ferric ratio (greater than1), and by a pH lower than 2.5. In stage III, the abiotic oxidation of Fe' to Fe+ ceases; theabiotic, direct oxidation of pyrite becomes insignificant; all the acidity is generated by theoxidation of pyrite (and other sulfides) by Fe', which itself is produced by bacteria fromFe2+.5.3.6 Kinetic ConsiderationsIt is unfortunate that general quantitative descriptions of the rate of pyrite oxidation, acid-neutralization and acid generation are so scarce, especially for real-field situations, given- 175 -the vital importance of such information in the planning of pollution control and abatement.The limited existing quantitative expressions are almost unexceptionally established incontrolled laboratory conditions. Qualitatively, the rate of acid generation is considered tofollow the curve outlined in Figure 39 (after SRK and Norecol, 1989). The controllingfactors of acid generation rate include sulfide content, presence of neutralizing materials,bacterial activity, availability of oxygen and water, pH, passage of time and other lessimportant ones such as climate and surface vegetation. Some of the above factors areinteractive with the rate of oxidation, such as pH and bacterial activity. The maximumoxidation rate is achieved when the steady-state cycling of ferrous and ferric ions bybacterial oxidation and pyrite reduction is reached.Prediction of future acidity and metal release requires a set of kinetic tests to start with.Such information is not available in this study. The calculation in Appendix II is based onthe assumption that the current rate of acidity release will continue in the future. Inreality, the release rate is likely to increase in both EMO and the Caps due to theavailability of acid producing materials and all the necessary ingredients of pyrite oxidation,if these dumps are not treated one way or another. Therefore, the duration of acidgeneration is probably shorter than that calculated in Appendix II, while the acidity andmetal load in the drainage will be greater than the present level.It took from five to ten years at ICM for various acid-generating dump areas to go acid orshow indications of acid generation in the drainage (metal level elevation). EMO was themost acidic dump and was built in 1981. It was found generating acid in 1986. The OldNorth Dump was deposited in early 70s and showed no indication of acid generation until- 176 -DRAINAGE pH]7^-----,/ •/ \/////^/^ X^/ 1^/ ^pH OF UNIMPACTEDENVIRONMENTRATE OF ACIDGENERATIONTIME(days,years,centuries)Figure 39 Schematice Diagram Showing the Rate of Acid Production over Time(After SRK and Norecol, 1989)- 177 -mid 80s. The initial period is necessary because oxidation is slowest at the very beginningbecause of low pH and little bacteria involvement; because the pH decrease and the aciditybuildup is often disrupted by rainstorms; and the most reactive acid-consuming materialsare available at the initial stage. The length of initial acidification stage depends on manyvariables such as S content, the content and nature of acid-consuming materials, dumpporosity, total surface area of reactive sulfides, etc. After the initial acidification, acidgeneration will be accelerated due to lowering of pH and increased bacteria involvement.5.3.7 Underwater Disposal ConsiderationsIf the finished pit is to be flooded either with fresh water or with sea water, underwaterdisposal of reactive wastes or polluted drainage is among the potential options of acid minedrainage abatement at ICM.Acid mine drainage abatement measures must be based on the fact that in the North Dumpproblematic areas are isolated and account for only a small portion of the entire NorthDump. The majority of the North Dump only gives rise to mild AMD problems.Underwater disposal options may include the removal of the reactive acid-generating wastematerials from the North Dump and their placement in the pit bottom before flooding thepit, or the collection of polluted drainage followed by its introduction into the pit bottomafter the pit is flooded.- 178 -The first option above may be enhanced by the placement of a layer of crushed limestoneon the top of the reactive wastes. The purpose of this limestone layer is to trap the metalpollutants in the bottom before they are stabilized by reduction to sulfides, or carbonates,or oxides in the initial stage of flooding. This option is most effective in preventing furtheracid and pollutant production but is costly.The purpose of the second option above is to retain the pollutants and the acidity in thedrainage in the flooded pit bottom. The pit bottom is generally expected to be anaerobicand reducing. The acidity input from the drainage will not make the pit bottom acidicbecause of its relatively small quantity compared to the pit and will be gradually consumed,for example, by the reduction of sulfate by decomposing organic debris below thesulfate/sulfide boundary but above the carbon dioxide/methane boundary, or by thereduction of some dissolved metal ions (such as Fe' and Cu24-) to their sulfides (pyrite andchalcocite) which are acid-consuming reactions.There are a few methods of flooding the pit after the mine's closure. They are:1) Turn the pit into a fresh water lake closed from the sea by natural drainagewater accumulation, or more quickly by pumping water from fresh watersources such as Marble river or Francis Lake.2) Turn the pit into a salt lake closed from the sea by first introducing sea waterfrom Rupert Inlet by a channel then disconnecting the channel.3) Turn the pit into a salt lake open to the sea by a connecting channel.4) Turn the pit into a salt lake semi-closed to the sea by a controlled connectingchannel.- 179 -Water balance has to be studied to prevent the lake from overflowing or drying up in 1) and2) above. The surface fresh water drainage courses around the pit area should then bealtered accordingly. The conditions at the flooded pit bottom depends primarily on thecirculation patterns of water and biological activity (supply of organic matters to the lakebottom). Generally, the lake bottom will be anoxic and reducing in 1) and 2) above, perhapsin 3) and 4).The underwater disposal of reactive wastes is not greatly influenced by the choice offlooding methods since its purpose is to seal the acid generating materials from oxygen. Forthe polluted drainage to be disposed of in the flooded pit bottom, the pit bottom should beanaerobic and reducing, preferably with a relatively high pH (>9) so that the acidity inputfrom the drainage will be neutralized instantly and the dissolved metals will be removedimmediately as carbonates or hydroxides (Figure 38), which later will be reduced to varioussulfides or carbonates or oxides. It may be beneficial to reduce the total amount of polluteddrainage to be put in the pit bottom by diverting the unpolluted water such as that to thenorth east of the North Dump within the End Creek watershed away from the North Dumparea. Another consideration in this option is the difference in the densities of the pitbottom water and the incoming drainage water. This difference should be reduced to suchan extent that the incoming drainage water will not trigger an upward density current tospread the polluted water throughout the entire water body. This problem may be solvedby mixing the drainage water with the pit bottom water (or other water of similar density)prior to the placement of the drainage water in the pit bottom. If the pit is going to befilled with fresh water, the density problem is less marked.- 180 -The introduction of acidic, polluted drainage into the pit bottom can be realized by usinga corrosion-resistant pipeline, or by constructing a tunnel connecting the pit bottom and asurface drainage discharge point. There is still much to be investigated if this option ischosen.6.0 CONCLUSIONS1. A typical, contaminated ICM drainage in resemblance to EMO has a pH of 4.5, atotal ionic strength of 0.0426 and a few species in concentrations much higher thannatural levels. The activity coefficients depart markedly from unity for all thespecies in consideration, thus have to be taken into consideration when examiningthe solution chemistry of the drainage.2. Metal solubilization is caused by the oxidation of sulfides and less importantly thedissolution of silicate minerals. The rate of metal release depends on many factorsincluding sulfur content, surface area to mass ratio of reactive sulfides, bacterialactivity, frequency and amount of rain, Eh-pH conditions, surface vegetation,temperature, etc.3.^The important processes in determining the levels of dissolved species in ICMdrainage include rates of mineral constituent release and removal of dissolvedspecies by precipitation of highly insoluble compounds, as well as the amount ofprecipitation. The removal of dissolved species by adsorption is not considered. Themost important insoluble compounds in the ICM drainage environment arecarbonates, hydroxides, and sulfates.- 182 -4. The typical drainage is under-saturated with respect to most of the metal speciesconsidered, usually by more than one order of magnitude. This implies that forthese metal species, chemical equilibrium with their most insoluble compounds is nota limiting factor of their levels. They can be virtually as high as their release ratecan afford. It is over-saturated, however, with respect to Cu in some streams, andis in approximate equilibrium with Pb and Zn in other streams. This means theequilibrium chemistry will set an upper limit for the levels of these metals.5. Levels of dissolved species can often be up to 2 orders of magnitude higher than isdictated by the equilibrium with their most insoluble species, as has beendemonstrated in this study.6.^The removal of metal contaminants by formation of carbonates and hydroxidesusually does not happen below pH 8. In acidic conditions, Pb and Ca may beremoved from solution by the precipitation of their sulfates when sulfateconcentration is high; all other metals are highly soluble. Above pH 8 and below pH10, most metals are removed as carbonates when they are concentrate enough,except Cu2+ and Fe', which are removed as hydroxides in this range. Increased totalcarbon activity will assist this carbonate removal process to function more efficientlyand at a lower pH. Therefore, growth of surface vegetation helps the removal ofmetal contaminants by increasing the total carbon activity. Above pH 10, mostmetals will be removed as hydroxides, but such a pH is not likely to occur in thedump or drainage.- 183 -7.^There is no barrier to sulfide oxidation at present. Oxygen and water, the two majorreactants for sulfide oxidation can enter the dump almost freely.8. Thiobacillus ferrooxidans are believed to have been involved significantly in thesulfide oxidation process in EMO. The steady-state Fe 2+/Fe3+ cycling, which is themost advanced stage of bacterial involvement in pyrite oxidation, has not beenreached in EMO.9. The Caps are weakly acid-generating, very porous, and contain no glacial till. Theyare however underlain by a layer of glacial till-rich layer of older North Dump, whichhelps neutralize the acidity in the drainage that leaves the Caps mostly fromunderneath. The Caps have a thermodynamic potential to generate acid for 650years. The current acid release rate is 730 tonnes H2SO4/year, of which 87% isneutralized within the dump by 640 tonnes CaCO 3 equivalent/year acid-consumingmaterials. The APP and the ACP are consumed at a proportional rate so that theacid-consuming material will last almost the entire acid-generating life of the dump.10.^EMO is very acid-generating and contains virtually no glacial till. It is notunderlain by a cushion of till-rich, acid-consuming layer, as is other areas in theNorth Dump. Therefore it is the most problematic area in the North Dump. Itsoxidation will continue for 600 years at the current rate. The current rate of acidrelease is 370 tonnes H 2SO4/year, of which 80% is being neutralized within the dumpby 300 tonnes CaCO 3 equivalent/ year acid consuming materials. The ACP is being- 184 -used up faster than the APP so that after 300 years the ACP within the dump willbe exhausted. The pH could drop dramatically after that time.11. The Old North Dump area, especially the Old Marginal Dump is related to the highloads of Pb and Mo in NDD. This indicates that active oxidation is undergoing insome locations, which mobilizes the metal contaminants; but the acidity isneutralized before the seepage reaches NDD.12. Other areas of the North Dump are less problematic than the ones discussed abovebecause the dump materials are a mixture of rock and glacial till, which is underlainby a till-rich layer.13. Without some kind of abatement measures, the North Dump will produceapproximately 3,500,000 cubic meters of contaminated drainage per year which willnot meet the provincial objectives with respect to some heavy metals. By divertingEDL drainage and TCR drainage which are very slightly polluted, the contaminateddrainage can be reduced to 2,600,000 cubic meters per year.14. As acid generation proceeds and pH drops in the future, the concentrations of Zn,Mn, Cu, and perhaps Cd (all of which are much more soluble in acidic conditions) inthe ICM drainage could increase dramatically, since the only control over theseconcentrations when pH is below 4.5 are the oxidation rate, which will acceleratesin the future. The concentration of Fe will only increase remarkably when pH is lessthan 3.0. The levels of As, Mo, and Pb will not be elevated significantly relative to- 185 -the present levels since their concentrations are limited by their low content in theICM waste rocks. Therefore they do not constitute concerns in environmentalpollution in the ICM drainage. The only exception may be Mo and Pb in the OldMarginal Dump, which probably contains much more of these metals than the wasterocks. The concentrations of Ca, Mg, Pb, and SO4 are mutually-limiting by thecorresponding solubility products of their sulfates, which do not change with pH.15. At ICM, the initial acidification of potential acid producing materials took from fiveto ten years. After the initial acidification, the acid generation process willaccelerate. The Caps were deposited on the North Dump in 1985. According to thelength of acidification demonstrated at ICM (5-10 years) and taking intoconsideration of the fact that the Caps is weakly acid-generating, highly porous, andcontains no till, we expect that the Caps will start to generate acid around 1992.The direct indication of this happening will be the elevation of dissolved metals inNDS and TCR.16. Effective abatement measures must be applied to the EMO area, the Caps area andthe Old North Dump area to avoid further drainage quality deterioration and tosatisfy the environmental standards for drainage quality. The rest of the NorthDump can be treated with less stringent abatement measures. 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Anionic Surfactants - Biochemistry, Toxicology, Dermatology,Marcel Dekker, Inc., New York.Goodman, A.E., A.M. Khalid, and B.J. Ralph, 1981. Microbial Ecology of Rum Jungle I -Environmental Study of Sulfidic Overburden Dumps, Experimental Heap Leach Pile andTailing Dam Area, Australian Atomic Energy Commission Report AAEC/E4aX.Goodman, A.E., A.M. Khalid, and B.J. Ralph., December, 1981. Microbial Ecology of RumJungle II - Environmental Study of Two Flooded Opencuts and Smaller, Associated WaterBodies, Australian Atomic Energy Commission Report AAEC/E-527.Harries, J.R. and A.I.M. Ritchie, 1981. "The Use of Temperature Profiles to Estimate thePyrite Oxidation Rate in a Waste Dump from an Opencut Mine", Water, Air and SoilPollution, vol. 15, pp.405-423.Harries, J.R. and A.I.M. Ritchie, May, 1983. "The Microenvironment within Waste RockDumps Undergoing Pyritic Oxidation", Progress in Biohydrometallurgy.Harries, J.R. and A.I.M. Ritchie, August, 1985. 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Inc. in association with Norecol EnvironmentalConsultants Ltd. and Robinson Dames and Moore, July 1987. Acid Mine Drainage Study,Mount Washington, Vancouver Island, B.C., Report 62701/1 Prepared for B.C. Ministry ofEnvironment and Parks.Steffen Robertson and Kirsten (B.C.) Inc. (SRK) in association with Norecol EnvironmentalConsultants Ltd. and Gormely Process Engineering, August, 1989. Draft Acid RockDrainage Technical Guide, Volume I, B.C. AMD Task Force Report.Stumm, W. and J.J. Morgan, 1981. Aquatic Chemistry, 2nd edn., Wiley Inc., New York.Tuttle, J.H. et. al., 1976. "Inhibition of Growth, Iron, and Sulfur Oxidation in Thiobacillusferrooxidans by Simple Organic Compounds", Can. J. Microbiol., Vol.26.- 192 -1984 Annual Environmental Assessment1984 Annual Environmental Assessment1985 Annual Environmental Assessment1985 Annual Environmental Assessment1986 Annual Environmental Assessment1986 Annual Environmental AssessmentUtah Mines Ltd., Island Copper Mine, June, 1985.Report, Volume I.Utah Mines Ltd., Island Copper Mine, June, 1985.Report, Volume II.Utah Mines Ltd., Island Copper Mine, June, 1986.Report, Volume I.Utah Mines Ltd., Island Copper Mine, June, 1986.Report, Volume II.Utah Mines Ltd., Island Copper Mine, June, 1987.Report, Volume I.Utah Mines Ltd., Island Copper Mine, June, 1987.Report, Volume II.Waldichuk, M. and R.J. Buchanan, November, 1980. Significance of EnvironmentalChanges due to Mine Waste Disposal into Rupert Inlet, Fisheries and Oceans Canada,British Columbia Ministry of Environment.Walsh, F. and R. Mitchell, 1972. "A pH-dependent Succession of Iron Bacteria", Environ.Science and Technology, Vol.6, pp809-812.Walsh, F. and R. Mitchell, August 1975. "Mine Drainage Pollution Reduction by Inhibitionof Iron Bacteria", Water Research, Volume 9.Walsh, F., 1978. "Biological Control of Mine Drainage", Chapter 14 in R. Mitchell (ed.)Water Pollution and Microbiology.Watzlaf, G.R, March, 1986. "Control of Acid Drainage from Mine Waste Using BacterialInhibitors", Paper presented at the 1986 National Meeting of the American Society forSurface Mining and Reclamation, Jackson, Mississippi, U.S.A.Weast, R.C. (editor-in-chief), et. al., 1966. CRC Handbook of Chemistry and Physics, 47thedn., The Chemical Rubber Co., Cleveland, Ohio.Young, M.J. and E.S. Rugg, February, 1971. "Geology and Mineralization of Island CopperDeposit", Western Miner.APPENDIX I^Calculation of the Pathway of Rain Water in 1987I. Calculation of the Pathway of an Enclosed AreaIn this appendix, significant digits in all figures are underlined. Insignificant digits areincluded here to reduce intermediate round-off errors. When actually reporting the figures,one should only quote the underlined digits, after rounding off the first non-underlineddigit, as the significant figures.An area which is believed to be a self-contained watershed is chosen on the map. This areais located to the north east of the pit, and the boundary of this area is defined by End Creekwater shed boundary on the north, the west and the east, by the End Creek/Trey Creekwater shed boundary on the south, and by the Ten Eighty Ditch (TED) on the south west.The area of this region is 1,767,424 m2 (19,024,400 square feet). In the production year1988, i.e., from September 27, 1987 to September 24, 1988,Total Mine Site Precipitation^= 2042.2 (mm)Total Volume of Rain and Snow Water Falling on the Area= 2.0422 x 1,767,424= 3,609,433 (m3)Assume 21.7% evaporation,Total Evaporation^= 3,609,433 m3 x 21.7%= 783,247 m3Total Volume of Water Running off or Percolating into the Dump within theEnclosed Area = 3,609,433 - 783,247= 2 ,826,186 m3Total TED Flow in this Period = 2,598,090 m3Since TED flow data is not available for September, 1987 - September, 1988, the data forNovember, 1987 - November, 1988 was used above. This was the only available data of thiskind. Of the total TED flow, 15% comes from NDD,Flow from NDD^= 2,598,090 x 0.15= 389,714 in3Therefore,Total Volume of Runoff and Seepage from the Enclosed Area= 2 , 598,090 x 0.85= 2,208,377 m3- 194 -The Volume of Water that Goes into the Ground Water System= 2,826,186 - 2,208,377= 617,548 m3II. Calculation of the Pathway of the Rain Water Falling on EMOThe technique used in this calculation is more or less the same as that employed in theabove calculation. The basic assumption in this calculation is: All the rain water fallingdirectly on the EMO area either evaporates or infiltrates or runs off. The runoff, if any,reports to the EMO ditch. Furthermore, the part that infiltrates either goes to recharge theaquifer or seeps into the EMO ditch.From September 27, 1987 to September 24, 1988,Mine Site Precipitation = 2042.2 mmTotal Area of EMO = 152,320 m2Annual Total Volume of RainWater Falling on EMO = 152,320 x 2.0422= 311,150 m3Adopt an evaporation rate of 21.7% (from previous calculation),Rain Water Evaporated = 311,150 x 21.7%= 67,520 m3Rain Water Infiltration intoand Runoff from EMO = 311,150 - 67,520= 243,630 m3Total Volume of Drainage fromEMO into EMO ditch = 221,050 m3The last figure above is for the period November 20, 1987 - November 19, 1988, thus a timeoffset exists between this figure and other figures we use in this calculation. We had to usethis figure in the calculation because this was the only available complete year drainagevolume. An error is introduced in doing so. However, if the total precipitation and thepattern of precipitation are similar in 1987 and 1988 for the months October and November,this error will be small. The precipitation record in 1988 was not yet available at the timeof the calculation. Therefore,Rain Water that Goes intoGround Water Aquifer^= 234,630 - 221,050= 22,580 m3-195-APPENDIX II^Calculation of Acid Generation and Consumption of EMO and CapsIn this appendix, significant digits in all figures are underlined. Insignificant digits areincluded here to reduce intermediate round-off errors. When actually reporting the figures,one should only quote the underlined digits, after rounding off the first non-underlineddigit, as the significant figures.I. Eastern Most OutslopeArea of Upper EMO^= 340,000 ft2Area of Lower EMO^= 1,300,000 ft2Total Area of EMO = 1,640,000 ft2= 152,361 m2= 15.24 Ha.Average Depth of Upper EMO = 40.0 ftAverage Depth of Lower EMO = 58.3 ftVolume of Upper EMO^= 340,000 ft2 x 40.0 ft= 13 600,000 ft3Volume of Lower EMO^= 1,300,000 ft2 x 58.3 ft= 75 790,000 ft3Total Volume of EMO^= 13 600,000 ft3 + M 790,000 ft3= 89 390,000 ft3= 2,531,239 m3If we use the average bulk density of the entire North Dump, 1,965 kg/m 3, we obtainMass of Waste Material in EMO, Calculated from Volume and Average Bulk Density= 2,531,239 m3 x 1.965 tonne/m3= 4,973,846 tonnesMass of Material in EMO from Haulage Records= 5,147,577 tons x 0.9071847 tonne/ton= 4,669,803 tonnesRelative Error of the Two Tonnages Obtained Above= ( 5,482,727 - 5,147,577 ) /[(5,482,727 + 5,147,577 ) / 2 ]= 6.306%- 196 -This error is presumably due to the error in the assumption of the average bulk density.Of course, we will adopt the tonnage from the Haulage Records for the followingcalculations.Average Acid Production Potential (APP)= 51.15 kg H2SO4 /tonneAverage Add-Consuming Potential (ACP)= 18.67 kg H2SO4 /tonnePrecipitation at the Mine Site from September 27, 1987- September 24, 1988= 2,042.2 mmAnnual Volume of Rain Falling on EMO= 152,361 m2 x 2.0422 m= 311 152 m3At 21.7% evaporation rate, 78.3% will percolate (and slightly run off) EMO.Total Percolation and Runoff^= 311,152 m3 x 78.3%= 243,632 m3Average Sulfate Concentration in EMO Drainage in 1988= 1,500 mg SO4= /1= 1.50 kg SO4 /m3Sulfate Released in 1988^= 1.50 kg SO4 / m 3 x 243,632 m 3= 365,448 kg SO4== 365.4 tonnes SO4Assume that all sulfate comes from the oxidation of sulfides and that sulfate removal fromdrainage by precipitation and other mechanisms is negligible. The above sulfate release canbe converted intoSulfuric Acid Leached out from Sulfides in 1988= 365,448 kgSO4= x ( 98.0 kgH2SO4 / 96.0 kgSO4)= 373,062 kg H 2SO4= 373.1 tonnes H2SO4Total Potential Tonnage of Sulfuric Acid Contained in EMO= 4,669,803 tonnes x 51.15 kg H 2SO4/tonne= 238 860,424 kg H2SO4= 238,860 tonnes H 2SO4If EMO continues to release sulfuric acid at the present rate ( 373.1 tonnes of H 2SO4 peryear),- 197 -Number of Years Needed for EMO to Release All the Sulfuric Acid Contained inSulfides^ = 238,860 tonnes H 2SO4 / (373  tonnes H2SO4/year)= 640.4 yearsIn 1988, the Average Acidity of the Drainage from EMO= 300 mg CaCO 3 equiv./1= 0.300 kg CaCO3 equiv./m3Total Annual Acidity Released in 1988= 243,632 m3 x 0.300 kg CaCO3 equiv./m3= 73,090 kg CaCO 3 equiv.= 73.09 tonnes CaCO3 equiv.If we assume the acidity in EMO is all from sulfuric acid,Acidity Consumed by Acid-Consuming Material in EMO in 1988= 373.1 tonnes - 73.09 tonnes= 300.0 tonnes(Note: 1.00 tonne 112SO4 — 1.00 tonne CaCO 3 equiv.)Percentage of Acidity Consumed by EMO Material in Total Acidity Released bySulfides= 300.0 tonnes / 373.1 tonnes= 80.41%Average Acid Consuming Potential of EMO= 18.67 kg CaCO 3 equiv./tonneTotal Acid Consuming Capability Contained in EMO= 18.67 kg CaCO3 equiv./tonne x 4,669,803 tonnes= 87 185,222 kg CaCO3 equiv.= 87,185 tonnes CaCO3 equiv.If the acid consuming capability of EMO continues to be consumed at the current rate of300.0 tonnes CaCO3 equiv. per year,Number of Years Needed before the Acid Consuming Capability of EMO Is Depleted= 87,185 tonnes CaCO3 equiv. /(300.0 tonnes CaCO 3 equiv./year)= 290.6 yearsAfter 291 years, because the acid consuming capability of EMO will have been depleted, all373 tonnes of H 2SO4 will appear in the EMO drainage for the remaining 349 years. Thisis equivalent toConcentration of H2SO4 in the Drainage Emanating from EMO- 198 -= 373.1 tonnes H2SO4 / 243,632 m3 drainage= 373 000,000 g / 243 632,000 1=1.531 g/1Since the molar mass of H 2SO4 is 98.0 g/mol,Molar Concentration of H2SO4 in the Drainage= (.1.531 g/1) / (98.0 g/mol)= 0.01562 mollMalar Concentration of II+ in the Drainage= (0.01562 mol H2SO4 /1) x ( 2 mol /mol H 2SO4)= 0.03124 mol 11-71Therefore,pH of the Drainage^=^ 0.3124)= 1.505Of course this is an idealized calculation and the pH 1.51 should be regarded as a lowerlimit.II. The CapsTotal Area of CapsTotal Tonnage of CapsAverage APP of the CapsAverage ACP of the Caps= 4,318,943 ft2= 401,243 m2= 40.124 Ha.= 10 051,791 tons= 9,118,831 tonnes= 52.29 kg H2SO4/tonne= 43.51 kg CaCO 3 /tonneAnnual Precipitation from September 27, 1987 -September 24,1988^= 2042.2 mmAnnual Total Volume of Rain Water Falling on Caps= 401,243 x 2.0422= 819,418 m3At an evaporation rate of 21.7%, 78.3% of total rain water will percolate or run off the Caps.Rain Water Infiltration into and Runoff from the Caps= 819 418 m3 x 78.3%= 641,604 m3- 199 -Assume that the sulfate load in the drainage along the Old Trey Creek (TCR) representsthat in the seeps form the Caps. (The Caps are underlain by a thick layer of porous olderdump material thus there is no stream emanating exclusively from the Caps ). The annualmean of sulfate in TCR is 1120 mg SO4 /1.Sulfate Concentration in Drainage Seeping from the Caps= 1120 mg SO4= /1This translates toSulfuric Acid Released by the Caps= 1120 x (98.0/96.0)= 1143 mg H2SO4 /1= 1.143 kg H2SO4 / m3Total Sulfuric Acid Contained in the Caps= 52.29 kg H2SO4 /tonne x 9,118,831 tonnes= 476,824 tonnes H2SO4Annual Release of H2SO4 from the Caps= 1.143 kg H2SO4 /m3 x 641,604 m3/year= 733.4 tonnes H2SO4 /yearNumber of Years for the Acid Producing Capability to be Exhausted= (476,824 tonnes H2SO4 )/(733.4 tonnes H2SO4/year)= 650.2 yearsBefore appearing as seeps along the Old Trey Creek water course, the drainage from theCaps has already been neutralized by the underlying layer composed largely of till whichhas a relatively large acid consuming potential. The tills are effective in neutralizing acidsbecause of their fine particle size. We could use the Old Trey Creek sulfate load torepresent that of the Caps because the sulfate released by the Caps is not removed by itsflowing through the till-rich layer (or the removal is negligible). This is not the case withacidity. All the acidity contained in the drainage immediately beneath the Caps is lost enroute before its reporting to TCR ditch. We have no information on the acidity of the waterdraining the Caps, therefore can not make a justifiable assessment as to how many yearsare needed before the acid consuming capability of the Caps is depleted.Nevertheless, if we were to make such an assessment regardless of what has beenmentioned above, we can use the acidity in EMO, which is 300 mg CaCO 3 equiv./1, as aguide. The APP of EMO is 51.15 kg/tonne and that of the Caps 52.29 kg/tonne. Theconcentration of sulfate in EMO ditch is Imo mg/I and that from the Caps 1120 mg/l. Weobserve from these that roughly equal APP resulted in roughly equal sulfate release. Nowthat the ACP of EMO is 18.67 kg/tonne and that of the Caps 43.51 kg/tonne, it is almostcertain that the acidity in the Caps drainage will be less than that in EMO drainage. For- 200 -calculation purpose, we will assume arbitrarily that the acidity in the Caps drainage is halfof that in EMO, namely 150 mg/l.Acidity in the Caps Drainage^= 150 mg CaCO 3 equiv./l= 0.150 kg/m3Total Acidity Released to the Caps Drainage= 0.150 kg/m3 x 641,604 m3 /year= 96.24 tonnes/yearBy the same reasoning as in EMO calculation,Acid Consumed by Acid Consuming Material in the Caps= 733.4 tonnes/year - 96.24 tonnes/ year= 637.2 tonnes/yearwhich is numerically equal to the consumption of acid-consuming material.Percentage Acidity Consumed = 637.2/733.4= 86.88 % Total Acid-Consuming Capability of the Caps= 43.51 kg/tonne x 9,118,831 tonnes= 396,760 tonnes CaCO 3 equiv.Number of Years Needed before Acid Consuming Capability in the Caps isExhausted= 396,760/637.2= 622.7 yearsThis is to say, the acid-consuming material will last almost as long as the acid-producingmaterial does. If this were true, the drainage from the Caps would carry only a smallacidity through its acid-generating life, and its pH will not lower considerably. Meanwhile,after the journey within the till-rich cushion layer, the drainage will hardly contain anyacidity or perhaps contain some alkalinity, the only effect down stream would be elevatedlevels of sulfate and relatively soluble metals, such as Mn and Zn, much like the waterquality exemplified by today's TCR drainage.However, this is only a very rough estimate and one has to bear in mind the assumptionswe have made in the calculation when applying these figures. It will be very beneficial tothis calculation if an analysis of sulfate and total acidity could be performed on the seepsdirectly draining the Caps, so that this calculation can be updated with real parameters ofthe Caps drainage.- 201 -


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