Patent Application: US-22403905-A

Abstract:
acidic metal - bearing wastewaters are treated to produce a finished water of sufficient purity to meet discharge standards while recovering metals removed in forms which are commercially valuable . the metals are selectively precipitated , either in a batch or in a continuous system , for removal of individual metal products in a specific sequence of steps from the wastewater . in each step , the ph is adjusted to the specific ph range and sulfide ion is introduced to precipitate the metals , excepting the removal of ferric iron and aluminum which is achieved using hydroxide precipitation . bioconversion process using unique equipment converts sulfate in the wastewater to the hydrogen sulfide gas required for the precipitation process . this bioconversion process reduces the sulfate in the wastewater so that the water can be directly discharged or used for agricultural applications .

Description:
the following experimental studies illustrate the present invention , but are not intended to be limiting of the scope of the invention . all of the experimental studies were conducted with acid mine drainage from the berkeley pit , located in butte , mont ., unless otherwise noted . the berkeley pit contains over 30 billion gallons of acid mine water with some 3 to 5 million gallons being added each day , depending upon operating conditions . this water contains relatively dilute levels of heavy metals . the ph ranges from 2 . 2 to 2 . 7 . table 3 gives the concentrations of the dissolved metals present in the berkeley pit water . this source was chosen because it represents the largest single source of acid mine drainage in the u . s ., and has been classified as the largest superfund site in the nation . however , the experimental findings in the present application can be used to recover metals from any bearing - bearing waste stream and / or acid mine drainage from sources other than the berkeley pit . i . metal precipitation studies ii . application of membranes for biological sulfate reduction ; the hydrogen sulfide produced by sulfate reduction can be used for metal precipitation iii . precipitate settling studies iv . conversion of iron sulfide to other products this process and subsequent experimental study was conducted to establish the precipitation conditions required to obtain high purity precipitates from acid mine drainage . fig1 shows a schematic of the batch precipitation process . feed water is pumped into a reservoir , then pumped to a reactor 10 . the left - hand side of the reactor is where the primary reaction takes place . upon addition of hydrogen sulfide to berkeley pit water at the proper conditions , metal sulfides were precipitated from the water . a mixture of hydrogen sulfide and carbon dioxide gas ( to simulate the product of srbs ) was bubbled into the solution through a gas sparger 12 ( which simultaneously provides mixing for the reaction ). the eh probe 14 and ph probe 15 are used to control the eh and ph of the solution during the reaction . a ph controller and pump maintain the solution at the desired ph for precipitation by addition of sulfuric acid ( 3 . 70 m ). the right - hand side of the reactor is designed to allow settling time to minimize carry - over of solids from one reactor to another . the sulfide concentration was maintained by controlling the gas flowrate bubbling through the reactor liquid . the batch process was operated sequentially to obtain the various metal sulfide / hydroxide precipitates and the water was analyzed to assure mass balance . table 4 summarizes the operating conditions of the various stages , operated sequentially using the same reactor apparatus . samples were analyzed for thirteen metals using icp ( method sw - 846 - 6010b ). sulfate ion was measured using method sw - 846 - 903b . chloride ion was measured using epa method 325 . 3 , and the ph was measured by epa method 150 . 1 . table 5 shows the metal recoveries obtained at the end of each stage ( step ). it can be seen that as long as the proper operating conditions were maintained in the batch reactor , almost complete precipitation of each metal was obtained . further , precipitate analysis showed that the metal precipitates were very pure . special conditions were obtained in the case of ferric and ferrous sulfide precipitation . further , fig2 shows a eh - ph diagram to illustrate the very low oxidation state required for the stability of iron monosulfide , fes , and the general redox dependence of fes and pyrite , fes2 . this is a simulated result obtained using the experimental data obtained from the batch reactor system as input to the geochemist &# 39 ; s workbench ( bethke 1992a , b ; 1996 ). in this experiment , the tests were conducted batch - wise sequentially rather than simultaneously . the amd was blanketed under argon gas and stirred continuously . the precipitation was conducted in batch - wise stages , using the filtrate liquid from the previous stage . the ph at each stage was adjusted before precipitation using a caustic solution , either sodium or potassium hydroxide . the ph values selected for each batch stage are given in table 6 . once the correct ph value had been obtained , the amd solution was introduced into a hermetically sealed apparatus that provided capabilities for measuring the pressure of the head space , the removal of gas from the head space , and the introduction of hydrogen sulfide gas from an external source . caustic was used as the precipitating agent in the aluminum removal stage . the head space within the reactor was then evacuated by the vacuum pump . once evacuated , the flow of hydrogen sulfide / carbon dioxide was started . this flow continued until a predetermined head space pressure was obtained . the hydrogen sulfide gas was introduced at 110 % of the stoichiometric amount needed for the target metal was used . the effluent solution was filtered through 0 . 2 micron filter paper . the precipitate and filtrate samples were analyzed for the metals present by icap ( inductively coupled argon plasma ) spectrometer . the gas used to precipitate the other metals in the preliminary experiment was a 50 / 50 percent mixture of carbon dioxide and hydrogen sulfide , a composition comparable to that produced biologically by sulfate reducing bacteria using an organic compound as a feedstock . it was found that the carbon dioxide in the gas used for this experiment formed carbonates in the presence of ferric ions and interfered with effective metal precipitation . further experimentation of metal precipitation in the presence of ferric ion was performed using 100 % pure hydrogen sulfide gas as the precipitating agent was conducted to circumvent this interference . this approach corresponds with the composition of the product gas from an srb reactor utilizing hydrogen consuming bacteria fed with hydrogen and carbon dioxide as feedstock . each precipitation was conducted step - wise in a single vessel , using the filtrate liquid from the previous step . the ph at each step was adjusted before the precipitation , using hydroxide . the precipitation was achieved using either sodium hydroxide or hydrogen sulfide , as appropriate . the optimal ph values selected for each batch stage are given in table 6 . table 7 gives the resultsof the initial and final concentrations of each metal , and the results of mass balance calculations , using 1000 liters of amd as the basis . fairly pure precipitates are obtained , as indicated by the summary shown in table 7 . table 7b summary of the amounts of each metal precipitated or removed in each stage and the corresponding percentage removals . al zn cu ni co fe mn as cd metal components inlet amount ( g ) 241 . 920 265 . 700 173 . 410 1 . 231 1 . 930 237 . 800 85 . 800 0 . 007 1 . 543 stage 1 2 . 210 0 . 800 167 . 820 0 . 000 0 . 030 0 . 600 0 . 000 0 . 000 0 . 051 stage 2 4 . 11 32 4 . 39 0 0 122 0 0 . 007 0 . 002 stage 3 0 . 000 231 . 430 1 . 200 0 . 000 0 . 00 52 . 630 0 . 000 0 . 000 0 . 950 stage 4 212 . 810 0 . 710 0 . 000 0 . 000 0 . 000 0 . 000 0 . 800 0 . 000 0 . 520 stage 5 0 0 0 0 . 131 1 . 01 0 59 . 87 0 0 stage 6 20 . 490 0 . 760 0 . 000 1 . 000 0 . 570 1 . 600 81 . 400 0 . 000 0 . 020 total removed ( g ) 239 . 620 265 . 700 173 . 00 1 . 131 1 . 610 236 . 700 82 . 200 0 . 007 1 . 543 amt . in effluent 2 . 3 0 0 0 . 1 0 . 32 1 . 1 3 . 6 0 0 % removal stage # 1 0 . 9 0 . 3 96 . 8 0 . 0 1 . 6 0 . 3 0 . 0 0 . 0 3 . 3 stage # 2 1 . 7 12 . 0 2 . 5 0 . 0 0 . 0 51 . 3 0 . 0 100 . 0 0 . 1 stage # 3 0 . 0 87 . 1 0 . 7 0 . 0 0 . 0 22 . 1 0 . 0 0 . 0 61 . 6 stage # 4 88 . 0 0 . 3 0 . 0 0 . 0 0 . 0 0 . 0 0 . 9 0 . 0 33 . 7 stage # 5 0 . 0 0 . 0 0 . 0 10 . 6 52 . 3 25 . 2 0 . 0 0 . 0 0 . 0 stage # 6 8 . 5 0 . 3 0 . 0 81 . 2 29 . 5 0 . 0 0 . 0 0 . 0 0 . 0 total 99 . 0 100 . 0 100 . 0 91 . 9 83 . 4 99 . 5 95 . 8 100 . 0 100 . 0 this experiment demonstrates that a batch process is able to produce fairly high purity precipitates , and the final effluent meets epa &# 39 ; s gold book standard . it should be noted that the above numbers are based on mass balance calculations , especially for the intermediate and final effluents . actual experimental analysis of the final effluent water is shown in table 8 . although metals can be precipitated as sulfides using hydrogen sulfide gas , either alone or in a mixture , there are several problems associated with using sparged gas . these problems include the following issues . 1 . the unused hydrogen sulfide gas must be recycled into the precipitator , which requires a recycle compressor . 2 . the use of a recycle compressor introduces high investment and operating costs for the mechanical equipment because hydrogen sulfide is a corrosive gas . 3 . it is critical to control the rate of hydrogen gas dissolution in water and subsequent reaction with the metal sulfates to form insoluble sulfides that are of sufficient size to settle rapidly . in a sparged system , it is difficult to control the addition of hydrogen sulfide gas in stoichiometric or controlled amounts to the acid mine drainage liquid , since the usual method is simply to bubble the gas through the liquid . 4 . the formation of colloidal metal sulfide particles , which are difficult to settle and which require special additives to agglomerate , is common in sparging and results in both increased capital and operating costs . fig3 shows the distribution of particle sizes that was obtained in a previous study ( govind , et al . 1999 ) using a 50 - 50 mixture of hydrogen sulfide and carbon dioxide gases which was bubbled through acid mine drainage . as can be seen from fig3 , the particle size distribution ( weight fraction ) obtained . by bubbling a 50 - 50 mixture of hydrogen sulfide and carbon dioxide gas mixture through acid mine drainage is bimodal , with the bulk of the precipitate shifted into the sub - micron range . it can be seen that colloidal metal sulfide particles are produced , which are difficult to settle without the use of additives . fig4 shows a membrane system used to precipitate , metal sulfides using hydrogen . sulfide gas , either alone or in mixtures of gases . the system ( 160 ) comprises a vessel ( 161 ) for retaining the membranes ( 162 ) with an inlet ( 163 ) for acid mine drainage and an outlet ( 162 ) for treated acid mine drainage . precipitated metal sulfides ( 166 ) exit the vessel at the bottom thereof , ( 165 ). 1 . because it eliminates bubble formation , no hydrogen sulfide recycle is required , since there is no gas bubbling through the acid mine drainage . 2 . only a low gas pressure is needed to introduce the hydrogen sulfide gas into the acid mine drainage with dissolution at the membrane interface and subsequent reaction . 3 . it provides a very high contact surface area between the gas and the liquid because of the small pore sizes in the membrane hollow fiber . 4 . it results in the formation of particles from the metal sulfide precipitation having desirable settling characteristics . preliminary experimental studies were conducted using an apparatus as detected in fig5 . the apparatus includes a single hollow fiber ( 2 mm internal diameter , 0 . 2 microns average pore size , polypropylene material ) which was 22 . 4 cm long inserted in a flask ( 190 ). the hollow fiber was affixed to a header ( 194 ). a cylinder containing pure hydrogen sulfide gas was connected to the hollow fiber ( 193 ), which allowed the hydrogen sulfide gas to flow through the fiber and diffuse through the micro - pores along the length of the fiber . this system was placed onto a magnetic mixer ( 192 ) for stirring the reaction mixture . liquid samples were periodically withdrawn from the side port of the flask , filtered and analyzed for metal concentration using icp analysis . samples were withdrawn at different time intervals , while hydrogen sulfide gas was allowed to diffuse and react with the metal sulfate in acid mine water . experimental measurements of copper concentrations in acid mine drainage as a function of time were made using the above - described apparatus . the experimental conditions were ph = 2 . 4 , temperature 25 ° c . fig6 shows the experimental data as a function of time . it can be seen from fig6 that the initial rate of precipitation , fitted by a straight line , is much higher than the rate achieved after 25 minutes , as shown by the second line . this decrease in rate of precipitation , which results in decreasing copper concentration in the liquid phase , is caused mainly by pore plugging of the membrane pores by the deposited copper sulfide particles . as the membrane pores become plugged , the contact surface area between the hydrogen sulfide gas and the acid mine drainage decreases , thereby causing the rate of precipitation to decline . furthermore , since the rate of precipitation is constant with time , the precipitation process is mass transfer controlled , and the reaction kinetics forming copper sulfide from copper sulfate is much faster than the rate of mass transfer . after the experiment , when the membrane fiber was withdrawn from the liquid , it was discolored by a black copper sulfide precipitate , again indicating that surface and pore precipitation of copper sulfide had occurred in the membrane fiber . the particle size distribution of the metal sulfide precipitate was determined and is shown in fig7 . the particle size distribution obtained with bubbling of hydrogen sulfide gas is also shown on the same plot for comparison . it can be seen that the particle size obtained using a membrane precipitator is much larger than the size obtained by bubbling hydrogen sulfide gas . this larger size of this precipitate facilitates its settling and removal from the precipitation system . experiments were also conducted with membrane precipitator and encapsulated srb bacteria in gel beads . active sulfate reducing bacteria ( srbs ) were encapsulated in two kinds of gels : ( 1 ) silica gel ; and ( 2 ) polyvinyl alcohol gel . silica gel beads were made as follows : 3 % alginate solution and distilled water were added into the colloidal silica solution so that the final concentration of alginate , biomass and colloidal silica would become 1 . 5 %, 5 % and 5 - 20 %, respectively . the ph of the colloidal silica solution was maintained between about 6 - 7 . the solution thus prepared was dropped into a 5 % cacl 2 solution to form beads of about 0 . 4 cm in diameter . then the beads were cured for 5 hours in the solution containing equal concentration of biomass in gel in order to prevent bacteria from diffusing out into the liquid solution . polyvinyl alcohol beads cross - linked with sodium nitrate were synthesized as follows : polyvinyl alcohol . ( pva ), 80 g , with 99 - 100 % saponification and about 1 , 500 - 2 , 500 degree of polymerization was mixed with 6 g of sodium alginate and diluted with deionized water to 500 ml . the mixture was heated until all the material dissolved to form a homogeneous mixture . the solution was then cooled down and mixed with 500 ml of sre biomass suspension containing about 300 mg / l dry weight of cells . the final mixture contained about 6 - 10 % ( w / v ) pva , 0 . 3 - 0 . 6 % ( w / v ) sodium alginate , and 150 - 250 mg / l of active microorganisms . the mixture was then added drop - wise into a solution of . sodium nitrate ( 50 % w / v ) and calcium chloride ( cacl 2 ) ( 1 % w / v ) and immersed for 1 hour to form pva - sodium nitrate beads . the main advantages of the above two synthesis methods for making silica gel and pva beads are as follows : 1 . beads are made at ambient temperature and neutral ph condition , which does not harm the active cells during the synthesis of the gel beads ; 2 . the two gels offer good mechanical strength and durability for use in membrane precipitators and reactors ; 3 . the gels are non - toxic to microorganisms ; 4 . the beads do not agglomerate after synthesis and hence can be easily mixed in liquid phase systems ; and 5 . the solution used for gelation of the gel , such as sodium nitrate and calcium chloride are also non - toxic to microorganisms . experiments were conducted with these gel beads , about 0 . 5 - 2 mm in diameter , for reducing metal sulfates in acid mine drainage into insoluble metal sulfides . a membrane precipitator was assembled using a 2 liter volume glass reactor , as shown in fig4 . the reactor consisted of 300 ml total volume of 0 . 5 mm gel beads , and acid mine drainage was added resulting in a total volume of 1 . 5 liters . membrane fibers 162 were immersed in the membrane precipitator , through which a mixture of hydrogen and carbon dioxide was passed . the gel beads enabled sulfate to diffuse into the gel and is then converted by the active srbs , encapsulated inside the gel beads , to form sulfide . this sulfide then diffuses out of the gel bead and reacts with the metal ions in acid mine drainage to form insoluble metal sulfides . the hydrogen and carbon dioxide , provided by the membrane fibers , dissolved in the water and diffused into the gel beads . the main advantages of using the gel beads were as follows : 1 . the gel beads made it possible to maintain a high concentration of active srbs inside the reactor - precipitator system , thereby accelerating the reducing reaction ; 2 . the gel beads protected the active srbs from the low ph acid mine drainage ; 3 . the metal precipitation of metal sulfides occurred outside the beads and outside the membrane hollow fibers , thereby preventing fouling of the membranes that occurred when using the membrane precipitator described in this application ; and 4 . the metal precipitates were easily settled from the liquid , indicating that they were larger in size than the precipitates created by bubbling hydrogen sulfide through acid mine drainage . 1 . experiments with precipitating pure metal sulfides from acid mine drainage by conducting the experiments at selected phs ;. and 2 . experiments with precipitating all the metal sulfides in one single step , where the purity of metal sulfide precipitates was not important . if the objective is to produce pure metal precipitates with the intention of recycling the metals , and the metal sulfate concentrations of the desired metals in the acid mine drainage are significant , the former strategy can be used to obtain pure metal sulfides . however , if the objective is to produce treated water for discharge and the metal sulfate concentrations in acid mine drainage are low , then the latter strategy may be adopted . sequential batch experiments were conducted with 2 - l reactors , as shown in fig8 . 1 . 5 l of acid mine drainage , obtained from the berkeley pit , was added with 300 ml of total volume gel beads . this acid mine water was drawn from a depth of 200 ft in the berkeley pit and all of the iron present was in ferrous ( fe + 2 ) form . a mixture of silica gel beads and pva gel beads were used , to check the stability of the bead materials in the presence of acid mine drainage . no attempts were made to compare the performance of the two different gel beads in this study . the phs used in each step of this batch experiment were as follows : 1 . initial ph of acid mine drainage , which was 2 . 3 . 2 . after the precipitation of the first metal sulfide , the ph was increased to 4 . 0 , by adding sodium hydroxide ; 3 . after the precipitation of the second metal sulfide , the ph was increased to 6 . 0 by adding more sodium hydroxide . the system was operated at each selected ph for about 1 hour and the metal sulfide precipitates were filtered out after each step . the gel beads were not changed at any intermediate step . a gas mixture of 50 % hydrogen and 50 % carbon dioxide was passed through the hollow fiber immersed inside the batch reactor . step # 1 : ph of 2 . 3 , which was the initial ph of the acid mine drainage . the metal sulfide precipitate obtained was mainly copper sulfide with a purity of 95 % with 5 % of zinc sulfide . all of the copper sulfate in the acid water was precipitated . the metal sulfide precipitate was 99 % zinc sulfide and all of remaining zinc sulfate was precipitated from the water in this step . there was no other metal sulfide produced . the metal sulfide precipitate was 85 % ferrous sulfide with 15 % aluminum hydroxide . all of the iron was precipitated in this step . this experiment showed the feasibility of using gel beads with membranes to precipitate metal sulfides from acid mine drainage . finally , an experiment was conducted at ph of 8 . 0 , using a new charge of acid mine water . after 3 hours of mixing the beads with the acid mine drainage , and passing the hydrogen - carbon dioxide gas mixture through the hollow fiber , 99 . 9 % of all of the metals present were precipitated as a mixture , indicating the feasibility of producing treated water using this kind of system . the only metal remaining in the water was manganese , which would have precipitated completely at a slightly elevated ph . in the precipitation schemes described in section 0088 , a standard laboratory hydrogen sulfide gas was used . this experiment demonstrates that this hydrogen sulfide gas mixture can be generated using a membrane reactor utilizing sulfate reducing bacteria to remove sulfate from amd ( or other waste waters containing sulfate ) to produce hydrogen sulfide gas . the water that results after removal of the metals by precipitation will contain excess sulfate and is suitable for processing by the membrane reactor system . sulfate reduction to hydrogen sulfide gas can be achieved in a conventional stirred tank or packed reactor , using sulfate reducing bacteria ( sre ) species , or in a membrane reactor . bioreactors can be operated either with an organic source , such as acetate , or a gaseous mixture of hydrogen and carbon dioxide . the species of srb that use organic nutrient sources , such as acetate , are different from those that use hydrogen and carbon dioxide gas mixture . studies have shown that , for large systems , it is more expensive to employ sulfate reducing bacteria that utilize acetate or other organic sources than those that utilize a gaseous mixture of hydrogen and carbon dioxide , which can be obtained by steam reforming of natural gas . since acetates and most organics useful in this process are liquids , they can be simply added to the sulfate contaminated feed water , and membrane systems are not necessary . however , when a gaseous mixture of hydrogen and carbon dioxide is used in an srb reactor , the use of a membrane system can offer distinct advantages . in this study , a novel membrane bioreactor system was used to biologically reduce sulfate to hydrogen sulfide gas , which can then be used to precipitate the metals from acid mine drainage . membrane reactors have been used in a variety of applications , including waste water treatment , chemical processing , and air pollution control ( govind and itho , 1989 ). existing technologies for using a hydrogen and carbon dioxide gas mixture for sulfate reduction using hydrogen utilizing srb is based upon the use of gas sparged reactors ( dupreez et al ., 1991 ). the gas mixture is bubbled through the reactor liquid , with the liquid bubbles rising through the liquid containing active sulfate reducing bacteria . the gases dissolve and diffuse to the active cells , resulting in the formation of sulfides . since hydrogen is rather insoluble in water , the unreacted gases exiting the reactor are re - pressurized and recycled . the main disadvantages of the sparged gas reactor system are as follows : 1 . because hydrogen gas has a very low solubility in water , a tall sorption tower is required to provide the mass transfer area required for a minimal hydrogen sorption because of the low solubility of hydrogen in water . 2 . the sorption is so limited even with the use of a tall sorption tower that a large hydrogen gas mixture recycle is required to maximize hydrogen utilization . 3 . these factors result in a substantial gas - phase pressure drop , which in turn requires the use of large recycle gas compressors to recover and return the hydrogen to the sorption tower . 4 . managing hydrogen gas compression for recycle to the sorption tower introduces safety issues . 5 . because of mass transfer limitations , sparged gas reactors have significantly higher volume than membrane reactors , and the operating costs of sparged reactors is higher compared to membrane systems mainly due to gas recompression and recycle costs . membrane reactors have been used in a variety of applications , including wastewater treatment , chemical processing and air pollution control ( itoh , 1989 ). ( membrane reactor technology , rakesh govind and naotsugu itoh , editors aiche symposium series , american institute of chemical engineers , 1989 ). recent literature provides information on the use of membrane bioreactors in wastewater treatment , and in biological sulfate removal as alternative systems of conventional bioreactors ( govind et al ., report to epa on “ studies on metal recovery from acid mine drainage and production of useful products , part 3 : membrane reactor studies , 2003 ). in these studies , experimental data are reported on the use of biofilms in the membrane bioreactors where the biofilms are attached to the membranes and actually grow in the pores of the membranes . in a recent publication ( tabak et al ., 2004 ), a membrane reactor was used to achieve biological sulfate reduction . this membrane reactor had the following major disadvantage : the attached biofilms exhibited washout at moderate reynolds number . at higher liquid flow rates , liquid shear resulted in removing the attached biofilm from the membrane surface , and a decrease in performance . at lower flow rates , the biofilm remained attached to the membrane , and superior performance was observed . a larger scale membrane module ( 40 ) was purchased and assembled as shown in fig8 . this system included a liquid reservoir ( 41 ) for the metal - free wastewater , which liquid reservoir was placed onto a magnetic stirrer ( 42 ). argon gas was introduced at the top of the liquid reservoir to prevent oxidation of the metals by air . wastewater was removed from the liquid reservoir by a pump ( 47 ) through a membrane filter ( 48 ) to a membrane reactor ( 44 ). hydrogen and carbon dioxide were introduced to the membrane reactor from a gas cylinder ( 45 ). further characteristics of this module are shown in table 8 . the hollow fiber module was operated at various liquid reynolds numbers by varying the liquid flow rate . fig1 shows the effect of reynolds number on the efficiency of sulfate reduction . it can readily be seen that at a reynolds number less than 500 , the membrane bioreactor behaves as a biofilm system , wherein the biofilms are retained on the outside surface of the hollow fibers within the module . however , as the reynolds number increases , which occurs when the liquid flow rate through the shell side of the membrane module is increased beyond the critical velocity , the biofilm begins to slough off , resulting in a combination of a biofilm and mixed reactor , wherein the active biomass is to some extent present as a biofilm outside the hollow fibers as well as suspended in the shell side liquid . as the reynolds number is increased further , the system behaves as a mixed reactor . at yet higher flow rates , significant amount of biomass is washed out of the membrane reactor and is removed by the external filter . the removal efficiency of sulfate declines as the membrane reactor changes from a biofilm system to a mixed reactor system . the performance of the membrane reactor at a reynolds number of 300 is shown in fig1 . as time increases , the efficiency of sulfate removal by conversion to hydrogen sulfide increases , until all the sulfate in the reservoir is converted to sulfide . the use of a membrane reactor system , shown in fig9 , overcomes the problems involved in using gas - sparged reactors and previously studied membrane reactors . the main advantages of this new membrane reactor system are as follows : 1 . the microporous membrane surface ( 30 ) presents a very large surface area to the liquid phase , resulting in high mass fluxes , compared to the surface area of the much larger rising gas bubbles in the sparged reactor system . 2 . hydrogen sulfide gas is formed outside the membrane and hence does not mix with the pressurized gas inside the hollow fibers , as shown in fig9 , so that there is no contamination of the hydrogen sulfide gas with carbon dioxide gas present on the membrane side , while the hydrogen sulfide is produced on the shell or liquid side , which is outside the membrane ( 32 ). 3 . there is no requirement for a gas recycle compressor , which is a major advantage in particular because of the safety issues concerned with hydrogen gas compression . 4 . the gel beads provide a suitable support for immobilization of active srb , preventing the problem of clogging biofilms on the membrane surface . the concentration of active srb present as biofilms is substantially greater than the concentration that can be achieved in suspended culture gas - sparged reactors , resulting in substantially higher sulfate reduction rates . these type of method to prevent membrane fouling can also be used in conventional wastewater treatment systems for reducing soluble bod in wastewater . 5 . use of encapsulated bacteria prevents washout problems associated with suspended culture reactors and previously operated membrane reactors . 6 . the investment and operating cost projected for the reactor are significantly lower than for a tall liquid - phase sparged reactor system . as described in previous experiments with metal precipitation using gel beads and membranes , the same apparatus can also be used in a bioreactor configuration to reduce sulfate solution to hydrogen sulfide gas , which can be then used to precipitate the metal sulfides from acid mine drainage . experiments were conducted using the silica and pva gel beads to quantitate the reaction rates in such a membrane reactor system . sulfate reduction can be performed by using either soluble organic substrates , such as acetate , alcohols , etc . or by using hydrogen - consuming srbs . the main issue with hydrogen consuming srbs is the low aqueous solubility of hydrogen in water . experimental data obtained in the lab shows the following characteristics of gel bead membrane reactors : 1 . the sulfate reduction rates are at least 3 - 10 times higher , mainly due to higher concentration of srbs in the encapsulated gel bead systems when compared with suspended cultures of srbs ; 2 . the problem of washout of the active cultures of srbs from the reactor system is eliminated using gel bead encapsulated srbs ; and 3 . the srbs are protected from the outside harsh environment , such as low ph , as in the case of acid mine waters . encapsulated bacteria using the gel beads can be used for many other applications : 1 . improve the performance of existing compost / soil biofilters for treating emission of odors and volatile organics ; 2 . nitrification of waters containing ammonia ; in nitrification the ammonia is converted to nitrate in water ; 3 . denitrification of nitrate in water to nitrogen gas ; 4 . treatment of trichloroethylene in groundwater ; 5 . improvement of activated sludge wastewater treatment system by putting the beads into the aeration basins ; and 6 . enhancing the operation of any biological treatment system , whether it is for air , water , soil or sediments . back - pulsing has been found effective in preventing excess accumulation of biomass outside the membrane hollow fibers when the module is operated as a biofilm system with a liquid reynolds number less than 500 . back - pulsing can be achieved by using a cylinder and piston arrangement , which is attached to the inlet gas flow line of the membrane module . by moving the piston , the gas pressure inside the hollow fibers be increased or decreased . when the pressure is increased inside the hollow fibers , the liquid present in the membrane pores is pushed out , which causes the excess biofilm to slough off of the fibers . however , since it is desirable to maintain a biofilm outside the hollow fibers , a low frequency and low amplitude pressure pulse is used to remove only the excess biomass and leave a thin active biofilm on the membrane surface . experimental studies were conducted on achieving sulfate reduction using a polypropylene hollow fiber membrane reactor system using hydrogen - consuming srb . master culture reactor studies showed that hydrogen - consuming srb could be cultured from anaerobic digested sludges . the nutrient medium used was adequate for growing hydrogen consuming srb , and biokinetic studies showed that the yield of the bacterial culture was very low . membrane reactor studies conducted using the hydrogen - consuming srb showed that the reactor is capable of reducing sulfate efficiently in a short residence time . the present invention provides a chemical treatment process for environmental clean up of acidic concentrations of ferrous ions from the open berkeley pit mine waters . this process involves the quantitative conversion of ferrous sulfate to a filtererable and non - colloidal ferrous monosulfide . precipitation followed by clean separation of ferrous sulfide during environmental separation procedures of the berkeley pit waters has been difficult to achieve because of the undesirable formation of non - filterable , colloidal ferrous polysulfides . these particles , are often produced by the reaction of sulfide ions upon the acidic waters of the berkeley pit with reagents such as hydrogen sulfide , sodium sulfide , or sodium hydrosulfide that are added to precipitate acidic ferrous sulfate at a certain ph using the sequential separation procedures used for metal ions . these colloidal particles are problematical even when other heavy metals such as copper and zinc sulfides are removed quantitatively at different phs . colloidal particles such as those formed as fes x ( where x is greater than one ) are difficult to isolate for purification if they are gelatinous and minute . after washing and drying , analysis revealed the presence of fes x , which formulation suggests the presence of 12 . 8 sulfur atoms to one ferrous atom . the desired formulation is an iron to sulfur ratio of one . the present invention solves this problem . the acid mine water treated was an acidic ( ph about 2 . 2 ), metal - bearing wastewater generated by the aqueous oxidation of metallic sulfides ( fes ) by the action of certain bacteria in active and abandoned mining operations . sodium sulfide and sodium hydrosulfide as well as the positive “ counter ions ” of quaternary ammonium - halides , such as cetyl trimethyl ammonium chloride , cetyl pyridinium bromide , benzal alkonium chloride , and mixed alkyl trimethyl ammonium chloride were purchased from aldrich . ferrous salts were purchased from fisher scientific . the colloidal material was determined to be negative at the interfacial surfaces . its negativity was established by addition of positive “ counter ions ”, such as those produced by benzyl alkonium chloride , dodecyl trimethyl ammonium chloride , cetyl pyridinium bromide , and mixed alkyl trimethyl ammonium bromide . when tested individually , the counter ions effected the rapid precipitation of the suspended colloidal particles . experimentally , 25 mg of each of the above positive counter ion quaternary ammonium halides was added to each vial containing 20 grams of berkeley pit mine waste waters having suspended ferrous polysulfide molecules . addition of the above first - named counter ion induced precipitation within 10 second . however , the resulting precipitated colloids with positive ions were not crystalline . it should be noted that an essential characteristic of these particles is that they must be crystalline in order to provide facile separation and purification of the ferrous sulfide compounds . a volume of 500 ml of berkeley pit mine waste water was used , containing 471 mg / l of ferrous ions , equivalent to 8 , 4 millimoles of ferrous ions / l . to this amount was added 2 . 016 grams ( 10 millimoles ) of sodium sulfide monohydrate along with 71 . 4 mg of potassium nitrite . to 500 ml ( 4 . 2 millimoles of ferrous ions ) of acid mine water ( from which zinc , copper and aluminum ions were removed ) were added 71 . 4 mg ( 0 . 84 millimoles ) of potassium nitrite and 2 . 08 g ( 10 millimoles of nas . 9h 2 o ), as well as 3 . 8 ml t - butyl cresol . the kno 2 was added in one portion along with the aqueous solution of di - t - butyl cresol . the reaction was run for 40 minutes at 45 ° c ., during which time the nas 9h 2 o was added in small increments . heating and stirring were discontinued . no phase separation nor precipitation was noted for one hour . the potassium nitrite ( 0 . 8 millimole ) showed no beneficial effect under the operating conditions used , as the filtration was slow and settling time took several hours . this reaction required several hours for two distinct phases to be noticeable in the flask filled with nitrogen and an anti - oxidant . the iron to sulfur ratio was 1 : 1 . 9 , suggesting the formation of polysulfide in the mixture . ten millimolar percent equivalent of sodium sulfite proved to be highly effective in remediating iron pyrite . the remediation treatment by sodium sulfite was conducted as follows : to the above reaction berkeley pit pyrites were added 10 mole percent of 8 . 4 millimolar of sulfite ions in 15 ml of water , and the sulfite was added in ten portions over the one hour period at 55 ° c . stirring and heating were terminated . this reaction displayed two distinct phases within about fifteen minutes . from this reaction , 1 . 078 grams of black precipitate was obtained . thus potassium nitrite ( kno 2 ) by itself , does not seem effective under conditions used in preparing a rapidly - settling precipitate , ferrous sulfide ; while the addition of 10 millimolar percent of sodium sulfite appears to be highly effective as shown below : the analysis of iron to sulfur ratio can be reduced as follows : 223 / 55 . 85 53 . 15 / 32 . 06 fe / 3 . 9928 s / 1 . 6578 when the above numbers are divided by their respective atomic weights to find milliequivalents of iron and sulfur and simplified to their lowest common denominator in the ratio , the iron to sulfur ratio was 2 . 41 : 1 , indicating that polysulfides were not formed . as stated by secor ( chem . rev . ), some degree of selective crystallization of one desirable form of crystals may be induced to form by the introduction of selected crystals . into a one liter round bottom flask with three necks was introduced 450 ml of doubly distilled water . then , 13 . 8997 grams ( 0 . 05 mole ) of ferrous heptahydrate was introduced into the flask followed by the addition of 211 mg of black ferrous sulfide . heat was applied to raise the temperature of the contents of the flask to 38 ° c . the reaction was discontinued after 45 minutes , at which time stirring and heating were discontinued . after 12 minutes , two clean phases were noted , indicating conversion of a polysulfide to the monosulfide . precipitation of iron as ferrous monosulfide and its conversion to iron products the ratio of iron to sulfur is calculated to be approximately 1 . 0 showing the absence of ferrous polysulfide , and the presence of ferrous monosulfide , as shown in table 9 . since the berkeley pit &# 39 ; s acid mine drainage contains large amounts of iron in the form of ferrous sulfate , experiments were conducted on converting this ferrous sulfate , once precipitated as ferrous sulfide , into iron products , such as alpha - goethite , alpha - magnetite , etc . synthesis procedures that had been earlier followed for converting pure ferrous sulfate to alpha - goethite and magnetite were applied to the ferrous precipitation obtained from the acid mine drainage . a major problem encountered was the slow dissolution rate of ferrous sulfide obtained from the precipitation strategies discussed above . even in the presence of strong acids , only small amounts of ferrous sulfides can be reacted and dissolved as ferrous ions . the reaction of hydrogen sulfide with ferrous sulfate occurs according to the following equation : this reaction is reversible , and when a high partial pressure of hydrogen sulfide gas is present , the dissolution of ferrous sulfide to from ferrous ions does not occur , even with strong acids . when hydrogen sulfide is bubbled through acid mine drainage , since both iron and sulfur can exist in multiple oxidation states , iron polysulfides ( fes x ) are formed during precipitation . the presence of excess sulfide results in further reaction between ferrous monosulfide and sulfide to form iron polysulfides . iron polysulfides are insoluble in most acids , nitric acid being the exception , and the rate of dissolution of iron polysulfides to form ferrous ions is very slow . this poses a major problem in forming iron products from the iron polysulfide precipitates formed in the precipitation processes disclosed above . before using ferrous sulfide produced from berkeley pit water , granular , reagent - grade ferrous sulfide was used to determine feasibility . initially , the desired concentration was a 0 . 40 mol / l ferrous solution ( t . wang et al ., 1998 ) and a volume of 700 ml . the ferrous sulfide was added to the water - filled reactor maintained at 40 ° c . the solution turned a cloudy gray color because some particles became suspended in solution . a stoichiometric amount of sulfuric acid was added while argon was bubbled through the reactor . reaction occurred , as evidenced by the odor of hydrogen sulfide gas . however , the reaction did not proceed to a high conversion rate because most particles never dissolved . thus , more acid was added . each successive addition of acid slightly increased conversion , but complete dissolution was never achieved . in this experiment , a total of 160 ml of concentrated sulfuric acid ( 12n ) was added to 6 . 1 grams of ferrous sulfide and reacted for 24 hours . in addition to using concentrated sulfuric acid , both concentrated hydrochloride and nitric acids were tried as reactants . small amounts of ferrous sulfide were placed into a 40 ml vial . excess acid was added and the mixture was brought to a boil . even under these extreme conditions , the ferrous sulfide did not completely react , and it left a porous black solid . experiments were conducted on dissolving the precipitated iron sulfide using oxidizing agents . initially , 15 % by weight hydrogen peroxide solutions were added to a mixture of iron polysulfide and 1m nitric acid . however , the dissolution reaction was slow and resulted in the formation of some iron oxides as follows : however , when a commercial oxidizing agent , paratene shp ™ ( woodrising resources , ltd ., calgary , albert ), which is a mixture of hydrogen peroxide and a stabilizer , was used , rapid dissolution of the precipitated ferrous polysulfide was obtained . the effect of the stabilizer in hydrogen peroxide prevents the hydrogen peroxide from decomposing in the presence of metal ions , which allows more hydrogen peroxide to react with the iron sulfides . further , paratene shp ™ immobilizes the iron as an acid soluble salt , and prevents further oxidation of the iron with hydrogen peroxide to form iron oxides , as was the case with hydrogen peroxide alone . two hundred mg of iron polysulfide precipitated from berkeley pit acid mine drainage using hydrogen sulfide gas in the precipitation experiment was mixed with 5 g of paratene shp ™ diluted with water in a ratio of 1 : 1 . the solution was heated to 40 ° c . and continuously stirred using a magnetic mixer . complete dissolution of iron polysulfides occurred in about one hour of mixing , indicating that all of the polysulfide had reacted to form a stable solution of ferrous iron . one n sulfuric acid was then added to obtain a clear solution of ferrous sulfate . similar results were also obtained with commercially produced iron sulfide or pyrites . the advantages of using iron sulfide precipitated with hydrogen sulfide gas , rather than the ferrous sulfate solution obtained from acid mine drainage are : 1 . the volumetric flow rate of acid mine drainage through the precipitation process does not affect the conversion process of ferrous polysulfide , since this step is conducted after the iron polysulfide precipitate is removed from the precipitation system ; 2 . the time taken for the iron polysulfide to react with the oxidizing agent solution does not affect the precipitation process ; and 3 . no oxidizing chemicals need be added to acid mine drainage , which can result in not only increasing chemical costs significantly , but also results in oxidizing manganese ions present in solution , forming a manganese oxide precipitate during precipitation of iron product . since the flow rate of acid mine drainage can be quite large ( 3 to 5 million gallons per day in the case of the berkeley pit ), the cost of adding any oxidizing agents to this flow can be prohibitively expensive , and result in impure precipitate . conversion of ferrous sulfate solution from acid mine drainage to iron products despite the disadvantages of converting ferrous sulfate solution obtained from acid mine drainage directly into iron products , this option was investigated . this approach may be particularly desirable when the flow rate of the acidic metal - bearing waste stream is not large and the concentration of manganese ions is small . goethite and magnetite are iron products that have been used commercially as pigments as well as for other products . when ferrous solutions are slowly oxidized by air bubbling , one or several of the following products may form : goethite ( alpha - feooh ), lepidocrocite ( gamma - feooh ), magnetite ( fe 3 o 4 ) and hematite ( alpha - fe 2 o 3 ). rapid oxidation using hydrogen peroxide leads to the precipitation of feroxyhyte ( delta - feooh ) [ frini et al ., 1997 ]. goethite is of particular interest , primarily because of its use as a precursor for synthesizing acicular iron pigments needed in magnetic recording media [ pozas et al ., 2002 ]. goethite is the alpha phase of iron oxyhydroxide and is produced both synthetically and naturally . goethite varies in color from yellow to dark brown , but the color by transmitted light is often blood red . it crystallizes in the orthorhombic system , with a mohs hardness of about 5 - 5 . 5 and a specific gravity of about 4 - 4 . 4 [ tottle , 1984 ]. goethite is chemically identical to lepidocrocite and pyrosiderite , differing only in crystalline structure . goethite has been successfully synthesized in the laboratory from both ferrous and ferric solutions . schwertmann described a method by which pure goethite is synthesized from ferrous iron as follows : 9 . 9 g of unoxidized crystals of fecl 2 . 9h 2 o was dissolved in one liter of deionized water . the ferrous solution was held in a wide - mouth two liter bottle . to the ferrous iron solution was added 100 ml of 1 . 0m nahco 3 . then the solution was aerated at a flow rate of between thirty and forty cubic centimeters per minute . oxidation of the ferrous iron was complete after 48 hours . the ph of the solution was maintained around seven by buffering with nahco 3 . both the ferrous iron solution and the sodium bicarbonate solution should be sparged with nitrogen gas to remove any dissolved oxygen prior to reaction ( schwertmann et al ., 2000 ). in addition to preparation from a ferrous iron solution , goethite may be produced directly from a basic ferric iron solution . boehm described a method by which goethite is produced from fe ( no 3 ) 3 . one hundred ml of 1 . 0m ferric nitrate solution was poured into a two liter polyethylene flask . to this solution , 180 ml of 5m koh solution was added rapidly with stirring . the resulting solution was immediately diluted to two liters with twice distilled water . the polyethylene flask was closed and held at 70 ° c . for 60 hours , during which time the reddish - brown solution was converted to a compact , yellow - brown precipitate of goethite . after the 60 hours , the solution was filtered and the resulting goethite filter cake was washed with twice distilled water to remove excess oh — and no 3 — ions ( boehm , 1925 ). additionally goethite may be synthesized from an acidic ferric solution . 283 grams of fe ( no 3 ) 3 . 9h 2 o was dissolved in 350 ml of 2m hno 3 . this solution was diluted with 1 . 4 liters of distilled water to which was added 1 . 4 liters of 1 . 0 m naoh with vigorous stirring . this yielded a final solution with hydroxide to iron ratio of approximately 2 . 0 . the ph of the solution was between about 1 . 7 and 1 . 8 . yellow goethite began to precipitate from solution after 50 days . the solution was then filtered , and the filter cake washed ( morup et al ., 1983 ; schwertmann et al ., 2000 ). x - ray powder diffraction is a useful method for determining the crystalline composition of various iron oxides . in the powder method , the substance to be examined is reduced to a very fine powder . the sample is then placed in a holder and inserted into a beam of monochromatic x - rays . the holder is then rotated under the monochromatic x - rays , and the diffracted waves are intercepted by the detector and measured . different crystal structure material show different peaks when collected by x - ray powder diffraction , and it is these distinct peaks at different angles that allow for the differentiation of materials . tables 10 - 13 below show the peak intensity and relative peak intensity versus 2θ for different types of iron oxyhydroxides and iron oxides . the xrd spectra described above were obtained using a siemens diffractometer . the two 2θ values ranged from 10 to 80 at a step size of 0 . 5μ with a scan time of one second . the spectra obtained are compared with the values given in the tables below to determine the mineral derived . goethite was produced from a fe 2 + system based on the aforementioned procedure by schwertmann ( schwertmann et al ., 2000 ). approximately 13 . 9 g of feso 4 . h 2 o was dissolved in 100 ml of deionized water through which nitrogen had been sparged for one hour to remove any dissolved oxygen . the solution was placed into a 2 . 0 liter erlenmeyer flask . a solution of sodium bicarbonate was made by dissolving approximately 9 . 2 g of nahco 3 in 110 ml of deionized water which had previously been sparged with nitrogen gas for one hour . the sodium bicarbonate solution was then added to the ferrous sulfate solution with rapid magnetic stirring . after the sodium bicarbonate solution had been added , air was sparged through the solution at a flow rate of between 30 - 40 cc / minute . the flow rate of the air was monitored by a rotameter on the air line . the solution was sparged for 48 hours . a precipitate formed , and , after the aeration was complete , the solution was filtered through 1 . 2 micron glass fiber filter paper . the filter cake was dried , weighed , and a small portion was taken for x - ray diffraction analysis to determine if goethite was produced . in addition , a new value called the r value is given . the r value is the ratio of the moles of bicarbonate to the moles of iron present in solution . r = nahco 3 n n fe = nahco 3 m ⁢ nahco 3 mw ( x fe ⁢ v ) ( 1000 * mw fe ) where : n nahco3 is the number of moles of bicarbonate in solution ( mol bicarbonate ) n fe is the number of moles of iron in solution ( mol iron ) m nahco3 is the mass of sodium bicarbonate added ( g ) mw nahco3 is the molecular weight of sodium bicarbonate ( g / mol ) x fe is the concentration of iron at the time of bicarbonate addition ( ppm ) v is the volume of solution ( l ) mw fe , is the molecular weight of iron ( g / mol ) goethite was produced from solutions of feso 4 . 7h 2 o with the addition of sodium bicarbonate and aeration for 48 hours . x - ray diffraction confirmed that the product was goethite . both precipitates were a tan color . two approaches were investigated in attempts to make goethite from berkeley pit amd . the first approach ( experiments 1 - 3 ) were conducted using raw berkeley pit amd as the starting material in each and treating with the bicarbonate method of schwertmann which involves bicarbonate addition followed by aeration ( schwertmann , et al . 2000 ). in the second approach , the synthesis of goethite from green rusts formed from berkeley pit amd was attempted in experiments 4 - 6 . in experiment 1 , the raw ( as received ) berkeley pit amd water was directly treated prior to the removal of any of the metals . in order to investigate the potential of improved segregation of metals , further experiments ( experiments 2 and 3 ) were conducted after removal of copper and zinc as their corresponding metal sulfides and removal of aluminum as aluminum hydroxide . the difference between the two experiments is their r value . experiment 1 : in this procedure one liter of berkeley pit amd water was placed into a two liter erlenmeyer flask and was aerated with nitrogen gas for one half hour to remove dissolved oxygen . the ph was approximately 2 . 6 . then , 1 . 97 g . of sodium bicarbonate was dissolved in 110 ml of deionized water through which nitrogen gas had been sparged for one half hour . the sodium bicarbonate solution was added , and aeration was begun at a flow rate of 45 cc / minute . the solution was aerated for 48 hours . once aeration was complete , the solution was filtered through 1 . 2 micron glass fiber filter paper , dried , and a small sample was taken for x - ray diffraction analysis . a sample of the filtrate was collected and taken for icap analysis . for experiments 2 and 3 , the raw amd was pretreated to remove the copper , zinc , and aluminum as follows . one liter of berkeley pit amd was treated to adjust the ph to 4 . 0 with koh in a two liter erlenmeyer flask . at this point , the solution was sparged with a pre - made mixture of 50 % hydrogen sulfide / 50 % carbon dioxide for one half hour in an attempt to remove the copper and zinc from the berkeley pit amd as copper sulfide and zinc sulfide . after one half hour of sparging , during which the solution was stirred , the solution was filtered through 1 . 2 micron glass fiber filter paper into a receiving flask . a sample of the filtrate was collected and analyzed by icap spectrometry . the filtrate from the copper / zinc removal precipitation stage was then treated for aluminum removal . the solution was sparged with nitrogen gas to remove any excess hydrogen sulfide and any dissolved oxygen . the ph of the solution was adjusted with koh to approximately 5 . 2 . the ph of the solution began to drop as the aluminum in the berkeley pit amd formed aluminum hydroxide . more koh was added to raise the ph back to 5 . 2 . this process was continued until the ph of the solution did not drop any further . the solution was then filtered through 1 . 2 micron glass fiber filter paper . a sample of the filtrate was collected and taken for icap analysis . the filtrate from the aluminum removal stage was then treated by the bicarbonate method of schwertmann . an amount of sodium bicarbonate was added to the solution , and the solution was aerated at a flow rate of between 30 and 40 cc per minute ( the schwertmann method ). the solution was aerated for 48 hours . once aeration was complete , the solution was filtered through a 1 . 2 micron glass fiber filter paper . the filter cake was dried , and a portion was taken for x - ray diffraction analysis . a sample of the filtrate was collected for icap analysis . the experimental conditions for experiment 1 are given in table 14 . this experiment was simply the addition of sodium bicarbonate to raw berkeley pit amd followed by aeration as previously described . the ph of the berkeley pit amd at the start of the experiment was 2 . 66 . table 15 filtrate metal concentrations : sodium bicarbonate addition and aeration to raw berkeley pit amd . sample sample sample concentration , ppm location number cu + 2 zn + 2 fe + 2 al + 3 mn + 2 aeration 1 0 . 1192 15 . 556 0 . 2848 4 . 460 153 . 0 outlet 2 0 . 1284 16 . 292 0 . 3308 5 . 024 161 . 2 3 0 . 1472 17 . 028 0 . 4080 5 . 180 162 . 2 average 0 . 13 16 . 29 0 . 34 4 . 89 158 . 8 error 0 . 01 0 . 64 0 . 05 0 . 33 4 . 4 as was expected , all of the metals were precipitated from solution as metal oxides . in this case , no goethite was formed , as evidenced by x - ray diffraction analysis . however , the iron product produced may have some commercial value . experiments 2 and 3 were then conducted using a feed to the aeration stage that was prepared using a sequential batch - wise treatment for removal of copper , zinc , and aluminum as previously described . the conditions are described in table 16 . table 17 shows the filtrate metal concentrations at the various stages of metal removal from the berkeley pit acid mine water for experiment 2 . the most notable aspect of this table is the finding that although all of the metals were not removed prior to aeration , the zinc and iron appeared to be co - precipitating in the aeration stage . as in experiment 1 , the aeration stage precipitate obtained from experiment 2 showed no sign of goethite formation . 7 . filtrate metal concentrations for experiment 2 : copper , zinc , aluminum removed ; sodium bicarbonate addition and aeration . sample sample sample concentration , ppm location number cu + 2 zn + 2 fe + 2 al + 3 mn + 2 cu / zn 1 bdl 343 . 88 325 . 08 180 . 80 139 . 80 precipitator 2 0 . 015 365 . 28 337 . 60 190 . 08 145 . 28 filtrate 3 0 . 034 372 . 68 340 . 28 192 . 76 146 . 04 average 0 . 02 360 . 61 334 . 32 187 . 88 143 . 71 error 0 . 02 12 . 95 7 . 03 5 . 44 2 . 95 al 1 bdl 403 . 2 375 . 24 23 . 54 167 . 4 precipitator 2 bdl — — — — filtrate 3 bdl 435 . 2 389 . 36 25 . 14 173 . 4 average 0 . 00 419 . 2 382 . 30 24 . 34 170 . 4 error 0 . 00 — — — — aeration 1 bdl 1 . 382 0 . 2896 4 . 756 10 . 144 filtrate 2 bdl — — — — 3 bdl 1 . 566 0 . 3000 5 . 288 10 . 480 average 0 . 00 1 . 474 0 . 2948 5 . 022 10 . 312 error 0 . 00 1 . 211 0 . 2408 4 . 118 8 . 423 experiment 3 — this experiment repeats the procedure of experiment 2 . table 18 shows the experimental conditions for the experiment . table 19 shows the filtrate metal concentrations at the various stages of removal from the berkeley pit amd for experiment 3 . the r value is determined from the concentration of iron in the aluminum filtrate . the most notable aspect of the above table is the fact that although all of the metals are not removed prior to aeration , the zinc and the iron appeared to co - precipitate upon aeration . this behavior was the same as in experiment 2 . batch experiments : goethite from berkeley pit amd green rusts ( experiments 4 , 5 , and 6 ). green rust is a term used to described fe 2 + = fe 3 + hydroxide salts appearance in the equilibrium state . when these salts are in solution , they dissociate into their cation salt and fe 2 + and fe 3 + hydroxides . there was an attempt to produce green rust from the berkeley pit amd as a precursor to goethite . one liter of berkeley pit amd as added to a two liter erlenmeyer flask , and the ph was adjusted to 4 . 0 with koh . at this point , the solution was sparged with a pre - made mixture of 50 % hydrogen sulfide / 50 % carbon dioxide gas for one half hour in an attempt to remove the copper and zinc from the berkeley pit amd was copper sulfide and zinc sulfide . after one half hour of sparging , during which the solution was magnetically stirred , the solution was filtered through 1 . 2 micron glass fiber filter paper . a sample of the filtrate was collected and taken for icap analysis . the filtrate from the copper / zinc removal precipitate stage was then re - administered into a clean and dry two liter erlenmeyer flask . the solution was sparged with nitrogen gas to remove any excess hydrogen sulfide and any dissolved oxygen . the ph of the solution was adjusted with koh to approximately 5 . 2 . as aluminum hydroxide was formed , the ph of the solution began to drop , therefore continual additions of koh were required to maintain the ph in the ph range of about 5 . 2 . at the conclusion of the reaction , the solution was then filtered through glass fiber filter paper . a sample of the filtrate was collected and taken for icap analysis . the concentrations of metal species in solution were analyzed using an icap spectrometer . liquid samples were filtered through a 0 . 22 micrometer membrane filter to remove solids and diluted by 25 percent with concentrated nitric acid to avoid precipitation of metals from changes in ph . table 20 shows the experimental conditions for experiment 4 . the r value is determined from the iron concentration in the aluminum precipitation filtrate . table 21 shows the filtrate metal concentrations at the various stages or removal from the berkeley pit amd for exp4amd . the most notable aspect of table 21 is that though all of the metals are not removed prior to aeration , the zinc and iron appeared to be co - precipitating . the was the same phenomenon observed with experiments 1 , 2 , and 3 . the aeration stage precipitate obtained from exp4amd showed no sign of goethite formation . table 22 experimental conditions for experiment 5 : copper , zinc , aluminum removed , green rust formed , sodium bicarbonate addition and aeration . moles of con - moles bicarbonate centration of iron in solution , r of iron in volume of in solution , n nahco3 ( mol solution solution n fe ( mol ( mol nahco 3 / mg / l ) ( l ) nahco 3 ) nah co 3 ) molfe ) error r 328 . 48 1 . 187 0 . 01 0 . 02 2 . 48 0 . 02 table 23 shows the filtrate metal concentrations at the various stages of removal from the berkeley pit amd for experiment 5 . the most notable aspect of table 23 is the fact that , although all of the metals are not removed prior to aeration , the zinc and iron appear to co - precipitate . this is the same behavior as was observed in experiments 1 through 4 . table 24 experimental conditions for experiment 6 : copper , zinc , aluminum removed , green rust formed , sodium bicarbonate addition and aeration . moles of con - moles bicarbonate centration of iron in solution , r of iron in volume of in solution , n nahco3 ( mol solution solution n fe ( mol ( mol nahco 3 / ( mg / l ) ( l ) nahco 3 ) nah co 3 ) molfe ) error r 267 . 96 1 . 1984 0 . 01 0 . 03 5 . 88 0 . 03 table 25 shows the filtrate metal concentration at the various stages of removal from the berkeley pit amd for experiment 6 . even though all of the metals are not removed prior to aeration , the zinc and iron appeared to co - precipitate , as was noted in all of the experiments conducted previously ( experiments 1 - 5 , above ). the precipitate obtained from the aeration stage in experiment 6 showed no sign of goethite formation . experiments were also conducted on ferrous sulfate solutions , obtained from amd , beginning with a berkeley pit amd solution obtained from a deeper section of water within the berkeley pit . this amd has a significantly higher concentration of ferrous sulfate and less than 10 ppm of ferric sulfate . two experiments were conducted at 39 ° c . the first experiment was conducted using an aluminum settler effluent containing 677 ppm fe . the second experiment was conducted using an aluminum settler effluent containing 620 ppm fe . because of the low level of ferric relative to ferrous iron , it was assumed that the iron present in solution was entirely ferrous ions . the initial rector volume of 700 ml was degassed at reaction temperature for one hour with nitrogen . at time t = 0 , 17 ml of nitrogen - degassed , 1n sodium hydroxide solution was added , producing an fe2 / oh — ratio ranging from 0 . 50 and 0 . 41 for the 677 ppm and 620 ppm solutions , respectively . air was then bubbled into the reactor at approximately 30 ml per minute for one hour . after one hour , the solutions were drained from the reactor and filtered . the filtrate from both reactions was orange in color and non - magnetic . for these two experiments , samples were taken from the reactor to measure the ferrous conversion . a 5 ml sample was removed from the sample port using a syringe . this sample was then added to 20 ml of dilute sulfuric acid and mixed . because the solubility product of goethite is approximately 10 - 44 ( t . wang et al ., 1998 ), converted ferrous ion can be filtered using a 0 . 2 micron filter . during the reaction , a yellow solid was retained on the filter . the filtrate was then analyzed to determine the concentration of un - reacted iron in the system , from which the reaction conversion is calculated . table 26 shows the iron concentration at various time increments for the two experiments . the low conversion and amorphous nature of the product obtained indicated that acicular goethite was not produced . several experiments were conducted at 60 ° c . the initial iron concentration of the aluminum settler effluent for these experiments was 625 ppm . table 27 gives the conditions used for each experiment . first , the reactor was filled with aluminum settler effluent ( the overflow ) and heated to reaction temperature while degassing with argon . this process took one hour to complete . after this first hour , sodium hydroxide solution was added , and the solution was mixed under argon at a flow rate of approximately 100 ml / minute for one more hour . then , air was bubbled through the mixture for two hours at a rate that varied between 100 and 200 ml / minute . the solutions were drained from the reactor , allowed to cool , and filtered using a 1 . 2 micron filter . for all experiments , two distinctly colored solids were visible . the greatest portion of the solids produced was gelatinous and colored orange . however , a few milligrams of a yellow solid were also visible . domingo et al . discuss the effect of chloride ions in solution . according to their research , goethite is formed in solutions containing both sulfate and chloride ions ; however , the particles formed are spherical . according to the crc handbook of chemistry and physics ( lide , 1990 ), ferric oxide ( fe 2 o 3 ) is reddish - brown , amorphous and gelatinous . this description matches the solids obtained above . it is to be understood that the phraseology or terminology employed herein is for the purpose of description and not of limitation . the means and materials for carrying out disclosed functions may take a variety of alternative forms without departing from the invention . thus , the expressions “ means to . . . ” and “ means for . . . ” as may be found the specification above , and / or in the claims below , followed by a functional statement , are intended to define and cover whatever structural , physical , chemical , or electrical element or structures which may now or in the future exist for carrying out the recited function , whether or not precisely equivalent to the embodiment or embodiments disclosed in the specification above , and it is intended that such expressions be given their broadest interpretation . adams , d . “ did toxic stew cook the goose ?” high country news , 27 ( 23 ), dec . 11 , 1995 . allen , j ., r . govind , r . scharp , h . tabak and f . bishop . “ metal recovery and reuse from acid mine drainage .” presented at the 1999 aiche annual meeting , technical program , november 1999 . unpublished . baltpurvins , k . a ., r . c . burns , and g . a . lawrance . “ heavy metals in wastewater : modelling the hydroxide precipitation of copper ( ii ) from wastewater using lime as the precipitant .” waste management , 16 ( 8 ), 717 - 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