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Hydrolysis of aluminum sulphate solutions at high temperatures Nikolic, Cvetko 1971

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HYDROLYSIS OF ALUMINUM SULPHATE SOLUTIONS AT HIGH TEMPERATURES BY CVETKO NIKOLIC Dipl. Ing. (Chemical Engineering) University of Belgrade, 1963 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF APPLIED SCIENCE In the Department of METALLURGY We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA February 1971 In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. Department of Metallurgy  The University of British Columbia Vancouver 8, Canada Date April 28, 1971 - ii -ABSTRACT Normal and acid aluminum sulphate solutions containing about 6.0 gr/1 of aluminum and up to 50 gr/1 SO^ were hydrolyzed until equilibrium was reached in the temperature region 125-250°C. Under the equilibrium conditions the only stable solid phase observed in equilibrium with a liquid phase of various compositions was basic aluminum sulphate with nominal formula 3A1203.4S03.9E^0. A small portion of the ternary diagrams for the system A^O^-SO^-K^Q at 225°C and 250°C was constructed. A mixture of aluminum sulphate and other metal sulphates, K„S0., NaoS0,, Li.SO., FeSO. and CuSO. 24 24 24 4 4 i.e. was hydrolyzed at 225°C in order to find the effect of these salts on hydrolysis. The overall hydrolysis reaction was found to occur according to the chemical equation; ekl^ + 4HS0. + 14H„0 • 3Alo0o.4S0„.9Ho0 + 14H+ 4 2 2 3 3 2 The equilibrium constants at 125, 150, 175, 200, 225 and 250°C were determined. Finally a mechanism for the hydrolysis of aluminum sulphate was proposed. - iii -ACKNOWLEDGEMENT The author wishes to express sincere thanks to Dr. E. Peters for his continuing guidance and interest in this project. His thanks is extended to fellow graduate students, particularly Mr. D. Jones and the technical staff of the Department of Metallurgy for their helpful discussions and assistance. Financial support from the National Research Council of Canada in the form of a Research Assistantship is gratefully acknowledged. - iv -TABLE OF CONTENTS Page 1. INTRODUCTION 1 2. A REVIEW OF THE LITERATURE .. . 3 2.1 The Fe203-S03-H20 System 3 2.2 The A1203~S03-H20 System 6 2.3 The M20-A1203-S03-H20 System 10 3. EXPERIMENTAL 12 3.1 Experimental Technique 13.2 Materials Used 18 3.3 Preparation of Solutions for Hydrolysis 21 3.4 Chemical Analysis 22 4. RESULTS AND DISCUSSION 4 4.1 The Effect of Temperature on Hydrolysis of Aluminum Sulphate Solutions 25 4.1.1 Hydrolysis of Aluminum Sulphate Solutions at 125°C 25 4.1.2 Hydrolysis of Aluminum Sulphate Solutions at 150°C 28 4.1.3 Hydrolysis of Aluminum Sulphate Solutions at 175°C 30 4.1.4 Hydrolysis of Aluminum Sulphate Solutions at 200°C 32 4.1.5 Hydrolysis of Aluminum Sulphate Solutions at 225°C 34 - V -Page 4.1.6 Hydrolysis of Aluminum Sulphate Solutions at 250°C 36 4.2 The Effect of Sulphuric Acid Concentration on Hydrolysis of Aluminum Sulphate Solutions 38 4.2.1 Hydrolysis of Aluminum Sulphate Solutions with Initial S03:A1203 Ratio > 3.0 at 250°C. 38 4.2.2 Hydrolysis of Aluminum Sulphate Solutions with Initial SO^Al^ Ratio > 3.0 at 225°C... 40 4.3 The Equilibrium Constants 42 4.4 Ternary Diagrams for the System A1203-S03-H20 . 51 4.5 Mechanism of the Hydrolysis of Aluminum Sulphate ... 55 4.6 The Effect of Alkali Metal Sulphates on Hydrolysis of Aluminum Sulphate at 225°C 56 4.6.1 The Effect of Lithium Sulphate on Hydrolysis of Aluminum Sulphate at 225°C 58 4.6.2 The Effect of Sodium Sulphate on Hydrolysis of Aluminum Sulphate at 225°C 60 4.6.3 The Effect of Potassium Sulphate on Hydrolysis of Aluminum Sulphate at 225°C 61 4.7 The Effect of Divalent Metal Sulphates on Hydrolysis of Aluminum Sulphate Solutions at 225°C 65 4.7.1 The Effect of Copper Sulphate on Hydrolysis of Aluminum Sulphate Solutions at 225°C 67 4.7.2 The Effect of Ferrous Sulphate on Hydrolysis of Aluminum Sulphate Solutions at 225°C 69 - vi -Page 4.8 Application of the Hydrolysis Process 71 5. CONCLUSIONS . 74 APPENDIX I SOLUBILITY OF ALUMINUM SULPHATE 75 APPENDIX II ANALYSIS OF THE SOLID PHASE 6 APPENDIX III X-RAY DIFFRACTION PATTERNS 77 LITERATURE 81 - vii -LIST OF TABLES Table Page 1 Hydrolysis of aluminum sulphate solutions at 125°C... 26 2 Hydrolysis of aluminum sulphate solutions at 150°C... 28 3 Hydrolysis of aluminum sulphate solutions at 175°C... 30 4 Hydrolysis of aluminum sulphate solutions at 200°C... 32 5 Hydrolysis of aluminum sulphate solutions at 225°C... 34 6 Hydrolysis of aluminum sulphate solutions at 250°C... 36 7 The effect of excess sulphuric acid on hydrolysis of aluminum sulphate at 250°C 40 8 The effect of excess sulphuric acid on hydrolysis of aluminum sulphate at 225°C 42 9 Equilibrium constant for the reaction HSO^ J H+ + Su^T • 45 10 Hydrolysis constant of water 46 11 Equilibrium hydrolysis of normal aluminum sulphate in the temperature range 125-250°C ^7 12 Equilibrium constants for hydrolysis of aluminum sulphate solutions with excess sulphuric acid at 225°C and 250°C 47 13 Equilibrium constants for the hydrolysis reaction of aluminum sulphate 49 14 Solution.composition of 250°C 51 15 Solution composition at 225°C 2 I | | 16 Equilibrium constants for reaction Al + ^0 £ AKOH)^4" •+ H+ 56 l - viii -Table Page 17 Hydrolysis of aluminum sulphate-lithium sulphate solution at 225°C 58 18 Hydrolysis of aluminum sulphate-sodium sulphate solution at 225°C 61 19 Hydrolysis of aluminum sulphate-potassium sulphate solution at 225°C 63 20 Ionic radii of the ions involved in the system 65 21 Hydrolysis of aluminum sulphate-copper sulphate solution at 225°C 67 22 Hydrolysis of aluminum sulphate-ferrous sulphate solution at 225°C 69 I-A Solubility of Alo(S0.)o in water 75 2 4 3 I-B Solubility of aluminum sulphate in aqueous solutions of sulphuric acid at 25°C 75 I-C Solubility of aluminum sulphate in aqueous 10% ^SO^. 75 II Analysis of the solid phases 76 III-A Diffraction pattern of the precipitate obtained at 150°C 77 III-B X-Ray diffraction pattern of the precipitate obtained at 175°CIII-C X-Ray diffraction pattern of the precipitate obtained at 200°C 78 III-D X-Ray diffraction pattern of the precipitate obtained at 225°CIII-E X-Ray diffraction pattern of the precipitate obtained at 250°C 79 - ix -Table Page III-F X-Ray diffraction pattern of the precipitate obtained from FeSO.-Al-CSO.K solution at 225°C 79 4 2 4 3 III-G X-Ray diffraction pattern of potassium alunite obtained at 225°C 80 II1-H X-Ray diffraction pattern of sodium alunite obtained at 225°C 8III-I X-Ray diffraction pattern of lithium alunite obtained at 225°C 0 - X -LIST OF FIGURES Figure Page 1 Fe203-S03-H20 system at 200°C 4 2 Fe203-S03-H20 system at 140°C ' 5 .3 A1203-S03-H20 system at 220°C 7 4 Experimental Apparatus 13 5 YSI-602 Termistor Calibration 14 6 Transducer Calibration 1.5 7 Pressure Dependence on Temperature 16 8 Heating Time to Reach the Working Temperature 17 9 Hydrolysis of Aluminum Sulphate at 125°C 27 10 Hydrolysis of Aluminum Sulphate at 150°C 29 11 Hydrolysis of Aluminum Sulphate at 175°C 31 12 Hydrolysis of Aluminum Sulphate at 200°C 33 13 Hydrolysis of Aluminum Sulphate at 225°C 35 14 Hydrolysis of Aluminum Sulphate at 250°C 37 15 Hydrolysis of Acid Aluminum Sulphate at 250°C 39 16 Hydrolysis of Acid Aluminum Sulphate at 225°C 41 17 Phase Diagram of the Al^-SO^H 0 System 43 18 Equilibrium Hydrolysis 48 19 The Change of Equilibrium Constant with Temperature .. 50 20 A1203-S03~H20 System at 250°C 53 21 A1203-S03-H20 System at 225°C 4 22 Hydrolysis of Lithium-Aluminum Sulphate Solution at 225°C 59 23 Hydrolysis of Sodium-Aluminum Sulphate Solution at 225°C 62 - xi -Figure Page 24 Hydrolysis of Potassium-Aluminum Sulphate Solution at 225°C 64 25 Hydrolysis Yields Dependence on Ionic Radii 66 26 Hydrolysis of Copper-Aluminum Sulphate Solutions at 225°C 68 27 Hydrolysis of Iron-Aluminum Sulphate Solutions at 225°C 70 28 D.T.A. of 3Al203.4SO3.9H2O 72 1. INTRODUCTION The aluminum industry was the first to use a pressure hydro-metallurgical process (The Bayer process) on a large scale. Caustic soda solutions are used to dissolve aluminum from bauxites and alumina is separated. Even though caustic solutions are not very selective and the Bayer process is limited to low silica ores, alumina is produced commerically only by t-his process. This has imposed on the aluminum industry the task of transportation of bauxties over a long distance because the producing countries do not have bauxites of their own. The production of aluminum in countries such as Canada and the United States is today very sensitive to the availability of bauxites from foreign countries. With a desire to recover aluminum from clay minerals, processes have been considered,from time to time, that are applicable to common North American clays for the domestic production of alumina. Such processes might also be useful in the treatment of certain mine waters from copper dump leaching, which contain dissolved aluminum sulphate, among other sulphate salts. Any new process producing alumina must compete in price with the Bayer process. This limits new leaching processes to cheap reagents. Sulphuric acid is the favourite reagent considered for alternate - 2 -* alumina production processes and an acid process developed by C.S.I.R.O. has been found to produce aluminum sulphate solutions which can be hydrolysed to yield basic aluminum sulphate. Most of the work in hydrolysis has been done with basic aluminum sulphate solutions, or solutions containing alkali metal sulphates and under "process conditions". Because of this there is practically no data on equilibrium hydrolysis available in the literature. Some equilibrium hydrolysis work was carried out on the Fe^O^-SO^-H^O system. With some exceptions A^O^-SO^-H^O and Fe20.j-S03-H20 systems precipitate similar compounds during hydrolysis. The separation of basic aluminum sulphate salts from aluminum sulphate solutions is not well understood either from a kinetic or thermodynamic point of view. This work is concerned with hydrolysis reactions of aluminum sulphate solutions that occur in the temperature range 125-250°C, and attempts were made to study both precipitation rates and the final equilibria. Commonwealth Scientific and Industrial Research Organization -Melbourne - 3 -2. A REVIEW OF THE LITERATURE It is generally known that above a temperature which is a characteristic of the metal ion involved, a sulphate salt becomes less soluble as the temperature is further increased. E.T. Carlson and C.S. Simons"^ have reported the necessity of temperature increase from 200 to 275°C to achieve selective extraction of nickel and cobalt from laterite ore. At 200°C,even though nickel and cobalt sulphate solubility was high enough to give good recoveries, the selectivity was rather poor because of the high solubility of aluminum and iron sulphates. They noticed that the SO^tA^O^ ratio in the residue was about 1:1 and suggested that Al^O^ and SO^ exist in it as a compound. 2.1 The Fe203-S03-H20 System At elevated temperatures ferric sulphate solutions hydrolyze to 2 precipitate ferric oxide or basic salts. Posnjak and Merwin have studied the Fe^^-SO^-H^O system at temperatures up to 200°C. At 200°C only three solid phases exist in equilibrium with various compositions of the liquid phase. ¥e^0^ is in equilibrium with dilute solutions of ferric sulphate. The stable phase in equilibrium with a somewhat higher ^concentration of sulphuric acid in solution is a basic salt of the composition Fe„0„.2S0_.Ho0. Finally, at high concentrations Fe2°3~S03" H2° SYSTEM AT 2 0 0 C° - 5 -Fe203 Fiq. 2 • Fe203- S03- H20 SYSTEM AT !40C° - 6 -of sulphuric acid in solution the stable phase is the normal salt, Fe203.3S03 (Fig. 1). With decreasing temperature the isotherms become more complicated and at 140°C, there are five solid phases in equilibrium with different compositions of the liquid phase (Fig. 2). With further decrease in temperature the equilibrium isotherms become still more complicated. At 130°C, ferric oxide monohydrate replaces the anhydrous ferric oxide, and at 110°C anhydrous ferric sulphate disappears, and two acid salts are stable (Fe^. 4S03> 3H20 and Fe^. 4S03.9H20) . 2.2 The A1203-S03-H20 System 3 P.T. Davey and T.R. Scott have studied the A^C^-SC^-H^O system under process conditions in order to determine the optimum conditions for production of alumina. Their experiments approached equilibrium to varying degrees. Solutions with initial S03:A1203 ratio less than 3.0 and high in aluminum content were hydrolyzed at temperatures of 100-220°C. In this temperature range the only stable phase observed, in equilibrium with various liquid phases, was basic aluminum sulphate of the nominal formula 3A1203«4S03>9H20. At these temperatures, the precipitated phase was in equilibrium with acid solutions of aluminum' sulphate. Only a limited number of experiments were performed for periods of time long enough to reach the equilibrium. The equilibrium saturation curve for this system at 220°C is given in Fig. 3. This saturation curve was obtained by hydrolyzing aluminum sulphate solutions with initial S0„:A1„0„ ratio less than 3.0. - 8 -4 Basset and Goodwin have studied the system Al^-SC^-t^O at 25°C at this temperature concentrated solutions of aluminum sulphate can be crystallized to yield either hydrated aluminum sulphate A^ (SO^) ^. 16^0 or the basic salt Al2O2.2SO2.HH2O depending on SO^iA^O^ ratio of the initial solution. These crystalline materials tend to occlude mother-liquor and are readily soluble in water, which makes it difficult to remove impurities by washing without redissolving a substantial amount of the solid phase. 5-9 T.R. Scott et al. have proposed an acid process for recovery of aluminum from ores low in bauxite content and high in silica. The stages of the new process are described as follows: a. Digestion Recycled liquor containing aluminum sulphate is used to recover extra aluminum from an ore at 180°C. b. Modification A liquor from the digestion stage is treated with fresh bauxite to give a solution of basic aluminum sulphate with SO^A^O^ ratio less than 3.0. c. Reduction Modified liquor is treated with sulphur dioxide to reduce soluble iron to the ferrous state at 100°C. d. Hydrolysis Reduced liquor is hydrolyzed at 220°C in the absence of oxygen to produce basic aluminum.sulphate. e. Calcination Basic aluminum sulphate is heated, preferrably in stages at 1200 °C-1300°C, to yield the required a-alumina and a mixture of sulphurous gases for recycling. J.L. Henry, and G.B. King"^ have studied the system A^O^-SO^-E^O - 9 -at 60°C. At this temperature two basic salts were found to be in equilibrium with solutions of aluminum sulphate having an SO^A^O^ ratio very close to 3.0. The salt Al^O^.2S0.J.HH^O is in equilibrium with solutions of pH about 2.5 and A^O^. SO^. 61^0 is in equilibrium with solutions of pH 2.5-3.20. At very low concentrations of free sulphuric acid in solution normal aluminum sulphate is the stable phase at 60°C. R.A. Chaves, V.V. Kavelin and B.P. Sobolev"'"^ have studied the side reactions occurring in the sulphuric acid process for extracting nickel and cobalt from Cuban laterites and have revealed the presence of the following compounds: hematite a-Fe^O^, boehmite y-A^O^.H^O, aluminochromite FeCCrAl^O^, basic iron sulphate Fe^O^. 2SO.J.H2O, basic aluminum sulphate 3AI2O.J. ASO^. 9^0 and hydronioj arosite 3Fe203.4S0.j. 9H20. The largest amount of basic aluminum sulphate is formed in the autoclave where the preheated pulp at 250°C mixes with concentrated sulphuric acid,indicating very high reaction rate at this temperature. Beside aluminum and ferric sulphate salts, other trivalent and 12 quadrivalent metal sulphates can be hydrolyzed. At room temperature violet-coloured chromium sulphate solutions are stable, but at high temperatures these are converted into green complexes in solution. When SO^iC^O^ ratio in the initial solution is less than 3.0 hydrolyzed chromic oxide is precipitated but only at temperatures higher than 180°C. Even at high temperatures solutions of lanthanum sulphate do not hydrolyze readily below pH 7. - 10 -Vanadium is almost completely precipitated even in the presence of free sulphuric acid to yield ^O,.. In the absence of oxygen at 100°C, the solutions are stable, but precipitate V^O^ at temperatures above 200°C. Scandium could be separated from rare earths provided that the solubility product of the scandium compound was exceeded at the temperature and in the solutions used. Indium can be separated from zinc electrolysis liquors under the same conditions as for scandium. Titanium and zirconium sulphates hydrolyze readily without recourse to autoclave conditions. 2.3 The M^A^O -SOy-H^O System* Under otherwise the same conditions, the presence of alkali metal sulphates will increase the hydrolysis yield of a basic salt from aluminum sulphate solutions. V.S. Sazhin, A.K. Zapolskii and N.N. 13 Zaklarova have studied the influence of ammonium, sodium, and potassium sulphates on hydrolysis of aluminum sulphate solutions. The degree of hydrolysis was increased in the following series: (NH^^SO^ < Na.SO. < KoS0.. 2 4 2 4 14 S. Bretsznajder, J. Boczar, J. Piskorski and J. Porowski have studied the hydrolysis of aluminum sulphate solutions with addition of sodium hydroxide at 179-285°C for 0-240 min. The precipitated solid phase was found to be Na20.3A1203.4S03>7H20. Excess sodium hydroxide * M20 = alkali oxide; M = Li, Na, K, NH^ - 11 -greatly increased the yield. Higher temperature and more dilute solutions gave higher yields. There was little change in yield for periods longer than 30 min. 15—18 A.K. Zapolskii et al. studied the effect of potassium sulphate from 0-167.5 gr/£ K^SO^ on hydrolysis of aluminum sulphate solution containing 308 gr/£ Al^SO^ at 170-250°C. The degree of hydrolysis sharply increased at all temperatures for K^SO^ concentration from 0-1 mole/1. Further increase in K^SO^ concentration had little influence on hydrolysis. In the absence of K^SO^ basic aluminum sulphates of constant composition (3kl^0^. ^SO^. 7^0) were formed. Potassium alunite K2[A16(S04)4(0H)12]'> was preferentially formed at 230-250°C from solutions containing 1 mole of K^SO^, or more, per mole of A^CSO^)^-At temperatures 170-190°C a less basic salt S^SO^.3A1203.4S039H20 was observed. The system Na20-A12Q3-S03-H20 was studied at 200°C under process conditions in order to determine the conditions for recovery 19 of alumina from ores. The precipitated compound was found to have a formula very close to Na20.3A1203>4S03.6H20. K2[A16(S04)4(0H)12] = K20.3A1203.4S03.6H20. - 12 -3. EXPERIMENTAL 3.1 Experimental Technique A shaking autoclave (Fig. 4) made of zirconium was used for all hydrolysis experiments. The autoclave was placed in an aluminum block containing five heaters (100 W. each) to provide good heat distribution. The temperature was controlled by a 71 model "Thermis-temp" temperature controller to within ± 6.1 °C. The temperature controller was used with a YS1-602 sensing probe with a range of 80 to 250°C, and was independently calibrated against a precision mercury thermometer in oil for each temperature used. The calibration curve is shown in Fig. 5. In the low pressure region the pressure in the system was monitored by a transducer and recorded by a "Heath" servo-recorder. For the high pressure region, a pressure gauge was used. A calibration curve for the transducer is shown in Fig. 6. The variation of pressure with temperature is given in Fig. 7. The heating time varied with temperature and this variation is shown in Fig. 8. For each run, the autoclave was filled with 75 ml of solution leaving about 30 cc of gas volume in the autoclave. re corder tronsduc er aotoclov e J? temperature controler I c ooIe r sampling tube Fig. 4- EXPERIMENTAL APAR AT US - 16 -•00 150 200 250 C° Fjg. 8. HEATING TIME TO .REACH THE TEMPERATURE - 18 -Samples of hydrolyzed solution were taken from the cooling system and the sampling time x^as between 2 and 3 minutes depending on temperature. To make sure that the samples represented the solution from the autoclave, the first 3 ml of solution from the sampling tube were discarded. Immediately after the sample was taken it was filtered and its pH measured. Most of the hydrolysis product was precipitated on the autoclave walls and remained inside after the liquid was removed. The autoclave was then cooled with running water to room temperature and opened. The precipitate was removed from the autoclave walls mechanically by shaking the autoclave with glass balls and a few milliliters of distilled water. After being filtered and washed, the precipitate was dried for 24 hrs at 105°C. After the solid phase was removed the autoclave was washed with sodium hydroxide solution and then successively with dilute sulphuric acid and water. To evaporate the residual water, it was then heated and again allowed to cool to room temperature. 3.2 Materials Used Aluminum sulphate, crystal, reagent, B&A, Al (S04) .18H20; M.W. 666.45 Maximum Limit of Impurities Insoluble 0.005% Free acid (H SO.) 0.20% Chloride (Cl) 0.002% Arsenic (As) 0.00005% - 19 -Heavy metals (as Pb) 0.001% Iron (Fe) 0.002% Substances not precipitated by NH^OH as sulphates 0.20% Lithium sulphate, granular reagent B & A LioS0..H.O M.W. 127.96 2 4 2 Maximum Limit of Impurities Insoluble 0.010% Chloride (Cl) 0.002% Nitrate (N03) 0.001% Heavy metal (as Pb) 0.0005% Iron (Fe) 0.0002% Potassium (K) 0.05% Sodium (Na) 0.10% Potassium sulphate, crystal, reagent B & A K.S0.. M.W. 174.27 2 4 Maximum Limit of Impurities Insoluble 0.005% Chloride (Cl) 0.001% Nitrogen Compounds (as N) 0.005% Arsenic (As) 0.00005% Calcium, Magnesium and R 03 precipitate 0.020% Heavy metals (as Pb) 0.'0005% Iron (Fe) 0.0002% Sodium (Na) 0.005% - 20 -Sodium sulphate, anhydrous, granular reagent B & A Na„S0. M.W. 142.05 z 4 Maximum Limit of Impurities Insoluble 0.010% Loss on ignition 0.50% Chloride (Cl) 0.002% Nitrogen compounds as (N) 0.0005% Arsenic (As) 0.0001% Calcium, magnesium and R 03 precipitate 0.010% Heavy metals (as Pb) 0.0005% Iron (Fe) 0.0005% Ferrous sulphate, crystal, reagent B & A FeSO..7Ho0 M.W. 278.03 4 2 Maximum Limit of Impurities Insoluble 0.005% Chloride (Cl) 0.001% Phosphate (PO^) 0.001% Copper (Cu) 0.005% Ferric- Iron (Fe'' ') 0.01% Manganese (Mn) 0.05% Substances not precipitated by NH.0H 4 0.050% Zinc (Zn) 0.005% - 21 -Cupric sulphate, granular crystals, B.D.H. reagent CuSO..5H„0 M.W. 249.69 4 2 Alkalies (suphated) not more than 0.5% Chloride (Cl) not more than 0.005% Iron (Fe) not more than 0.08% Aluminum "AnalaR" B.D.H. Al M.W. 26.98 Acid insoluble matter passes test Iron (Fe) 0.004% Copper (Cu) 0.005% Silicon (Si) 0.01% Total nitrogen (N) 0.002% 3.3 Preparation of Solutions for Hydrolysis The aluminum sulphate solution was prepared by dissolving hydrated aluminum sulphate Al^(SO^)^.I8H2O in distilled water. A solution of approximately 6.0 gr/£ of aluminum was made having 302^1^02 ratio of about 3.0, and a pH of 3.10. The solution was filtered before it was used for hydrolysis. Solutions of the lowest pH (pH = 0.70) used in this work were prepared by dissolving A^C^SO^) ^. I8H2O in dilute sulphuric acid. Any solution between these two limits (pH = 0.70 and pH = 3.10) was prepared by mixing selected ratios of the two. Solutions containing alkaline metal sulphates and divalent metal sulphate salts were prepared by mixing the solution of the corresponding salt with aluminum sulphate solution. Solutions of normal aluminum sulphate were discarded after 7 days of - 22 -storage. Aluminum sulphate solutions containing excess sulphuric acid were kept for a few weeks. Solutions containing alkaline metal sulphates were prepared just before using them. Initial and final solutions were analyzed for aluminum, sulphate, and added metal ions. The pH of the solutions was measured before and after hydrolysis. The sulphate content was determined by precipitation as BaSO^. 20 * Aluminum was determined by complexing with EDTA and then back titrating the excess EDTA with zinc sulphate using dithizone as an indicator and 1 M CHoC00H-l M CHoC00NH. solution as a buffer. 3 3 4 Standard aluminum sulphate solution was made from "Analar" aluminum by dissolving it in sulphuric acid. 3.4 Chemical Analysis * disodium salt of ethylene diamine Na-OOCH„C CHoC00H H00CHoC CH^COONa diphenylthiocarbazone .H H II S II H H Divalent iron in the aluminum sulphate solution was analyzed by I |  | i titrating Fe with CeCHSO^)^ to oxidize it to Fe using 1,10-orthophenanthroline ferrous sulphate (Ferroin) as an indicator. In this case the aluminum content of the solution was determined from the difference in EDTA used to titrate both iron and aluminum. Copper and aluminum sulphate solution was analyzed by titrating liberated 1^ with ^2820^ to determine the copper content. The aluminum content was determined from the difference in EDTA used to titrate both copper and aluminum. The precipitated phase was also analyzed for aluminum and sulphate. The solid phase is soluble in strong sulphuric acid and strong sodium hydroxide, and therefore the same methods as for the solutions were used for analysis. Water content was taken as the difference between the total weight and the aluminum and sulphate content. Li, Na and K content was determined by the "EEL" flame photometer. ,c—c c— C.> // \ // \ c c—C \ / \ ' C=N N=C C [Fe(C12HgN2)3] ++ cation; red colour 41 + 2Cu ++ 2CuI + I2 I 2 21 - 24 -4. RESULTS AND DISCUSSION The intention of the work presented in this thesis was to study the change in equilibrium hydrolysis of aluminum sulphate solutions with temperature, sulphuric acid concentration and other cations. It is generally known that the amount of precipitated phase is determined by the concentration of aluminum and sulphate in the initial solution 3 for a given temperature. In the present work the concentration of aluminum in the initial solution was not varied in order to make it possible to establish the effect of other factors. Results obtained are discussed under the following classifications: 4.1 The effect of temperature on hydrolysis of aluminum sulphate solutions. 4.2 The effect of sulphuric acid concentration on hydrolysis of aluminum sulphate solutions. 4.3 The equilibrium constants 4.4 Ternary diagrams for the A^O^-SO^-f^O system. 4.5 Mechanism of the hydrolysis of aluminum sulphate. 4.6 The effect of alkali metal sulphates on hydrolysis of aluminum sulphate. 4.7 The effect of divalent metal sulphates on hydrolysis of aluminum sulphate solutions. 4.8 Application of the hydrolysis process. - 25 -4.1 The Effect of Temperature on Hydrolysis of Aluminum Sulphate Solutions At ordinary temperatures, concentrated solutions of aluminum sulphate will only crystallize a basic salt if the SO^iA^O^ ratio in the initial solution is less than 3.0. 14 S. Bretsznajder has obtained normal aluminum sulphate from clay (calcined at 770-820°C for 1 hr)by leaching it with dilute sulphuric acid for 10-12 hrs at 80°C and cooling the filtrate to crystallize 3 the normal aluminum sulphate hydrate. It was reported by T.R. Scott that at temperatures higher than 110°C hydrolysis of aluminum sulphate solutions with initial SO^A^O^ ratio less than 3.0 precipitates crystalline material "basic aluminum sulphate" (B.A.S.). This compound is related to the mineral alunite, having a formula expressed as 3AI2O2.4SO2.9H2O with a theoretical composition as follows: ^L2°3 ~ 38.8%; S03 = 40.6%; HyO = 20.6%. The temperature effect on hydrolysis was determined by hydrolyzing dilute solutions of aluminum sulphate with initial S02:Al202 ratio of about 3.0 in the temperature range 125-250°C. Retention time in these experiments was as long as 16 hrs. The equilibrium was determined from the hydrolysis curves when there was no further change in the aluminum concentration with prolonged time. 4.1.1 Hydrolysis of Aluminum Sulphate Solutions at 125°C Hydrolysis experiments at this temperature were done with unfiltered solutions of pH = 3.10. Results obtained were not reproducible. When the same solutions were filtered, and then hydrolyzed, results were more uniform. Hydrolysis yields at equal time intervals were less for - 26 -filtered solutions, and, in both cases, not sufficient for chemical analysis. For times which were long enough to establish equilibrium, hydrolysis yields were almost the same for both the filtered and the unfiltered solutions, thus suggesting that a nucleation process plays an important role in -the kinetics of the reaction at this temperature. Results of the analysis of hydrolyzed solutions are shown in Table 1 and Fig. 9. '» Table 1. Hydrolysis of aluminum sulphate solutions at 125 C. Time at 125°C hrs Al in sol. gr/Ji SO. in sol. 4 gr/i> Al prec. gr/£ SO^ prec. pH S03/A1203 in the pre-1.00 6.09 - 0.01 . - 2.98 -3.00 5.89 - 0.21 - 2.53 -5.00 5.80 31.70 0.30 0.70 2.24 1.31 6.00 5.69 31.33 0.41 1.07 1.80 1.47 7.00 5.68 31.34 0.42 1.06 1.80 1.41 12.50 5.76 31.30 0.34 1.10 1.78 1.82 13.75 5.65 31.38 0.45 1.02 1.79 1.27 16.50 5.68 31.38 0.42 1.02 1.78 1.36 Starting solution : 32.40 6.10 + + 0.3 gr/£ 0.05 gr/£ S0.= A pH = 3.10 - 28 -4.1.2 Hydrolysis of Aluminum Sulphate Solutions at 150°C Filtered solutions of aluminum sulphate with initial SO^rA^O^ ratio of about 3.0 were hydrolyzed at 150°C for different times. Results obtained are presented in Table 2 and Fig. 10. From Fig. 10 equilibrium seems to be reached in 6.0 hrs and the SO^tA^O^ ratio of the precipitated compound remains unchanged with prolonged times. The amount of precipitated compound was not sufficient for chemical analysis but the X-ray diffraction pattern indicated that basic aluminum sulphate existed at this temperature (see Appendix III-A). Table 2. Hydrolysis of aluminum sulphate solutions at 150°C* Time at 150°C Al in sol. SO, in sol. 4 Al prec. SO. prec. 4 PH so3 :A1203 hrs gr/£ gr/£ gr/£ in the prec. 1.00 5.28 30.44 0.82 1.96 1.67 1.34 2.00 5.07 29.96 1.03 2.44 1.55 1.33 3.00 4.84 - 1.26 - 1.45 -4.00 4.76 - 1.34 - 1.45 -6.00 4.64 28.65 1.46 3.75 1.40 1.44 10.00 4.64 28.60 1.46 3.80 1.38 1.46 16.00 4.65 28.64 1.45 3.46 1.44 1.34 Starting solution: 32.40 gr/2, SO 6.10 gr/£ Al pH = 3.10 - 30 -4.1.3 Hydrolysis of Aluminum Sulphate Solutions at 175°C Filtered solutions of aluminum sulphate with initial SO^iA^O^ ratio of about 3.0 were hydrolyzed at 175°C and results obtained are presented in Table 3 and Fig. 11. The precipitated phase has a ratio of SOyA^O^ of about 1.32 (calculated from Fig. 11) corresponding to the ratio of basic aluminum sulphate. Analysis of the solid phase confirmed the presence of the basic aluminum sulphate (see Appendix II, for Chem. Anal, and Appendix III-B for X-ray diff. patterns.). Table 3. Hydrolysis of aluminum sulphate solutions at 175 C Time at Al in sol. SO in sol. Al prec. SO prec. pH SO :A190» 175°C 4 * hrs. gr/Jl gr/>! gr/£ gr/£ in the prec. 1.00 4.45 28.53 1.54 3.87 1.30 1.32 2.00 3.96 27.26 2.14 5.14 1.27 1.35 4.00 3.68 26.76 2.42 5.64 1.18 1.31 6.00 3.56 26.39 2.54 6.01 1.15 1:33 10.00 3.55 26.40 2.55 6.00 1.15 1.32 16.30 3.50 26.40 2.60 6.00 1.15 1.30 * Starting solution: 32.40 gr/£ S04 6.10 gr/£ Al pH = 3.10 J ' " 1 -J 1 I > I ! * ' ' ' ' I L_ 0 • 2 3 4 5 6 7 8 9 10 II 12 13 14 15 16 Fig II. HYDROLYSIS OFALUMINIUM SULPHATE AT 175 C° timeinhrs. - 32 -4.1.4 Hydrolysis of Aluminum Sulphate Solutions at 200°C Solutions with intial SO^ikl^O^ ratio of about 3.0 were hydrolyzed at 200°C for as long as 18 hrs. Results obtained are shown in Table 4 and Fig. 12. The equilibrium hydrolysis yield is higher than at 175°C. The SO^iAl^O^ ratio of the precipitated compound calculated from Fig. 12 is about 1.31. Analysis of solid phase shows that the precipitated compound is the basic aluminum sulphate (see Appendix II for Chem. anal, and Appendix III-C for X-ray diff. pattern). Table 4. Hydrolysis of aluminum sulphate solutions at 200°C * Time at 200°C hrs. Al in sol. gr/2, SO, in sol. 4 gr/£ Al prec. gr/£ SO, prec 4 gr/£ pH S03:A1203 in the prec. 1.00 3.30 25.78 2.80 6.62 1.07 1.33 2.00 3.01 24.98 3.09 7.42 1.03 1.35 4.00 2.71 24.32 3.39 8.08 0.99 1.34 6.00 2.51 23.91 3.59 8.49 0.98 1.33 6.45 2.51 23.89 3.59 8.51 0.98 1.33 18.00 2.40 23.65 3.70 8.75 0.98 1.33 Starting solution: 32.40 gr/SL S0~ 6.10 gr/£ Al pH = 3.10 0 I 23456789 10 II Fig.12. HYDROLYSIS OF ALUMINIUM SULPHATE AT 200C 12 13 14 15 ° timsinhrs. 16 - 34 -4.1.5 Hydrolysis of Aluminum Sulphate Solutions at 225°C Filtered solutions of aluminum sulphate with initial SO^iA^O^ ratio of about 3.0 were hydrolyzed at 225°C and results obtained are shown in Table 5 and Fig. 13. Hydrolysis yields are very high. The precipitated phase has an SO^Al,^ ratio 1.33 (calculated from Fig. 13). Analysis of the solid phase showed the composition of basic aluminum sulphate (see Appendix II for Chem. anal, and Appendix III-D for X-ray diff. pattern). Table 5. Hydrolysis of aluminum sulphate solutions at 225°C Time at 225°C hrs. Al in sol. gr/l SO. in sol. 4 gr/H Al prec. gr/n SO, prec. 4 gr/X, pH SO in 3:A12°3 the prec. 1.00 1.89 22.45 4.21 9.95 0.89 1.33 1.00 1.88 22.35 4.22 10.05 0.92 1.34 2.00 1.80 22.31 4.30 10.09 0.92 1.32 2.00 1.73 21.55 4.37 10.85 0.89 1.39 3.00 1.65 21.96 4.45 10.44 0.88 1.32 5.00 1.58 21.40 4.52 11.00 0.89 1.37 8.40 1.44 21.36 4.66 11.04 0.85 1.33 16.00 1.48 21.63 4.54 10.77 0.85 1.33 * Starting solution: 32.40 gr/j> S04 6.10 gr/Z Al pH = 3.10 Fig. 13- HYDROLYSIS OF ALUMINIUM SULPHATE AT 225 C* timemhrs - 36 -4.1.6 Hydrolysis of Aluminum Sulphate Solutions at 250°C Filtered solutions of aluminum sulphate were hydrolyzed at 250°C and it was found that the amount of precipitated basic aluminum sulphate on heating the autoclave to this temperature was comparable to the amount of basic salt precipitated in 16 hrs at 225°C. This suggests that the reaction rate above 225°C must be very high. Results obtained at this temperature are shown in Table 6 and Fig. 14. The SO^A^O^ ratio of the precipitated compound is 1.33. Analysis of the solid corresponds to the basic aluminum sulphate (see Appendix II for Chem. anal, and Appendix III-E for X-ray diff. pattern). Table 6. Hydrolysis of aluminum sulphate at 250°C* Time at 250°C hrs. Al in sol. gr/£ SO, in sol. 4 gr/£ Al prec. gr/£ SO^ prec. gr/£ PH S03:A1203 in the prec. 0.00 1.52 21.57 4.58 10.83 0.85 1.33 1.00 0.97 20.27 5.13 12.13 0.80 1.33 3.30 0.88 20.27 5.22 12.13 0.80 1.31 6.30 0.91 20.13 5.19 12.27 0.79 1.33 * Starting solution: 32.40 gr/2, SO 6.10 gr/£ Al pH = 3.10 o ca < 6.0 J 0 I v o \ ll 5.0 1 4.0 3.0 H 2.0 H I.OJ \ \ O Al Q S0| V PH PH 3.0 o 1-30 2.0 4 -28 - 26 24 1.0 -I o Co 22 h20 18 16 1 1 ' 1 • 1 1 L 1 1 I I I • • 3 1 2 3 4 5 6 7 8 9 10 II 12 13 14 15 16 Fig. 14. HYDROLYSIS OF ALUMINIUM SULPHATE AT 250C° timeinhrs. - 38 -4.2 The Effect of Sulphuric Acid Concentration on Hydrolysis of Aluminum Sulphate Solutions Most of the work on hydrolysis of aluminum sulphate was done on solutions having initial SO^Al^O^ ratio less than 3.0. The reason for this was to obtain high yields in a short time and at low tempera tures. If hydrolysis is to be used as a purification process for leach solutions higher concentrations of sulphuric acid would be expected. In the work presented in this thesis a series of experiments with solutions containing about 6.0 gr/£ aluminum and different concentrations of sulphuric acid was done. Hydrolysis curves in the previous part were used to determine the retention time for the two temperatures 225°C and 250°C. These two temperatures were chosen because of the high yields at short times with solutions having an SO^rA^O^ ratio of about 3.0. 4.2.1 Hydrolysis of Aluminum Sulphate Solutions with Initial S03:A1203 Ratio > 3.0 at 250°C Filtered solutions with excess sulphuric acid were hydrolyzed at 250°C for 4 hrs. The highest S03:A1203 ratio used in this set of experiments was about 4.63. Results obtained are shown in Table 7 and Fig. 15. As the concentration of excess acid in the initial solution was increased the amount of precipitated basic aluminum sulphate decreased. i ! - 39 -~~i r~ 22 24 —T" 28 n 1 1 1 1 r 30 32 34 36 38 26 '0 o o0 v SO^jn sol. gr/1. O Al 0 S°4 VPH IO 2.0 —L— —r-1.0 3.0 4.0 AI in sol. gr/1. + Fig.15. HYDROLYSIS OF ACID ALUMINIUM SULPHATE AT2 50C* 2.0 finalpH - 40 Table.7. The effect of excess sulphuric acid on hydrolysis of aluminum s ulphate at 250°C. pH SOT init. 4 Al in sol. SO, in sol. 4 Al prec. SO^ prec. pH initial gr/£ gr/Jl gr/£ gr/£ gr/A final 0.70 50.08 3.31 43.43 2.79 6.65 0.56 1.00 44.26 2.17 34.97 3.93 9.29 0.60 1.50 37.05 1.25 25.45 4.85 11.60 0.70 2.00 34.08 1.00 22.05 5.10 12.03 0.76 2.50 33.21 0.98 21.13 5.12 12.08 0.78 3.10 32.40 0.91 20.13 5.19 12.27 0.80 4.2.2 Hydrolysis of Aluminum Sulphate Solutions with Initial S03:A1 03 Ratio > 3.0 at 225°C Filtered solutions of aluminum sulphate were hydrolyzed at 225°C for 6 hrs. Initial solutions used in this set 0f experiments were the same as solutions used at 250°C. Hydrolysis yields increased with i^preasing pH of the initial solution. Results obtained from this set of experiments are presented in Table 8 and Fig. 16. - 41 -X o. 3.0 H 22 ~i— 26 30 I i p 34 —r— 38 i i 42 SO4 in soLgr/l. 1.0 -4 20 3.0 • 4.0 Al in sol.gr/1. f-,-. . ~ IX) 20 pH final Fig. 16. HYDROLYSIS OF ACID ALUMINIUM SULPHATE AT 2 25C° - 42 -Table 8. The effect of excess sulphuric acid on hydrolysis of aluminum sulphate at 225°C. pH SO. init. 4 Al in sol. SO. in sol. 4 Al prec. SO^ prec. pH initial gr/£ gr/£ gr/a gr/A gr/£ final 0.70 50.08 4.94 47.04 1.16 3.04 0.65 1.00 44.26 3.62 38.35 2.48 5.91 0.70 1.50 37.05 2.33 28.37 3.77 8.68 0.78 2.00 34.08 1.87 24.10 4.23 9.98 0.83 2.50 33.21 1.63 22.60 4.47 10.61 0.84 3.10 32.40 1.50 21.52 4.60 10.88 0.85 4.3 The Equilibrium Constants Application of the phase rule to the system Al^O^-SO^-H^O requires that at constant temperature, a solution composition exists which is invariant when in equilibrium with two solid phases, e.g. A^O^nR^O and 3A1203.4S03.9H 0 (B.A.S. - the most basic salt). If the starting solution has a composition below this point, the first phase to precipitate is always the oxide, but the solution composition will change in the direction of the invariant point until the basic salt also precipitates. The location of the invariant point P (Fig. 17) will change with the temperature. Fig.I 7. PHASE DIAGRAM OF THE AI2O3-SO3-H2O SYSTEM - 44 -If the solution composition is below the invariant point P at the temperature T and above the invariant point at any temperature higher than T, then the oxide will precipitate in the early stages of heating and at later stages the basic salt will precipitate while the oxide will tend to redissolve. Eventually the oxide may completely disappear if the composition of the solution is such that it is above the invatiant point. In the work of T.R. Scott, a second phase (A^O^m^CO was not observed because the hudrolysis was carried out at high concentrations of aluminum in solutions and it can only be expected to precipitate from dilute solutions of basic aluminum sulphate. The X-ray diffraction patterns did not show the presence of A^O^ in the solid phase and therefore it can be concluded that the solution composition is above the invariant point where only one phase exists in equilibrium with the liquid phase. The X-ray diffraction pattern of the precipitate obtained at 125°C, was very poor so that no conclusion about the presence of aluminum oxide could be made. If there is only one solid phase in equilibrium with liquid phase at all temperatures from 125°C up to 250°C, then the same hydrolysis reaction should occur. For such a reaction, the slope of a plot of log [equilibrium constant] vs. 1/T°K should be linear. 21 R.G. Robins has reported values for the second dissociation constant of sulphuric acid up to about 200°C which are in good agreement with the values calculated from the correspondence principle. Values of these constants, are given in Table 9. The presence of aluminum sulphate will shift the equilibrium to the left in the reaction 4SC>4" J H+ + SO^". Therefore only HSO^ ions will exist in solution. Assuming that aluminum sulphate is completely dissociated, and that only HSO, ion can exist - 45 -- -»• + Table 9. Equilibrium constants for the reaction HSO, H + SO. 4 4 T°C 125 150 175 200 225 250 log K -3.53 -3.89 -4.40 -4.72 -5.95 -6.05 in solution in the temperature range .125-250°C the hydrolysis reaction can be written as: 6A1444" + 4HS0. + 14KL0 —— 3Alo0„.4S0o.9Ho0 + 14H+ (1) 4 2 — 2 3 3 2 The equilibrium constant for this reaction can be expressed as: K = (2) [A1W]D[HS04"]^ An expression for the variation of the equilibrium constant K with temperature is derived by combining -AG0 = RT in K (3) and 9T ^ T ;JP T2 K ' Therefore ,9 in K. = d&n K = AH°_ ,,-\ 9T ;P dT 2 UJ Kl assuming thatAH° is virtually independent of temperature. Equation (5) can be also written as - 46 -din K AH° . . dl/T = R {b) Thus, if £n K is plotted against 1/T the slope of the curve at any point is equal to -AH°/R. From the expression for the equilibrium constant (2) it follows that -log K = 14pH + 6 log [Al+++] + 4 log [HS04~] (7) The pH of the hydrolyzed solutions was measured at room temperature and as such it is different from the pH at which the reaction occurs. The change in the dissociation constant of water (Table 10) is such that the pH of the solution should decrease with increasing temperature. Table 10. Hydrolysis constant of water T C° 25 60 100 150 200 250 300 350 -log K 14.0 13.05 12.21 11.65 11.30 11.18 11.19 11.33 True pH value at any temperature can be defined thermodynamically but it cannot be easily measured. The pH change which occurs in heating or cooling an electrolyte solution is due to the hydrolysis and change in activity of all ions in solution. Assuming that the pH of hydrolyzed solutions remains unchanged with temperature the equilibrium constant for the hydrolysis reaction can be calculated at each temperature. - 47 -Results of the equilibrium hydrolysis of normal aluminum sulphate solutions in the temperature range 125-250°C are presented in Table 11 and Fig. 18. Results in Table 7 and Table 8 were also used to calculate the equilibrium constants for hydrolysis of aluminum sulphate solutions with excess sulphuric acid. Calculated values of log K are presented in Table 12. Table 11. Equilibrium hydrolysis of normal aluminum sulphate in the temperature range 125-250°C. Temp. Aluminum in sol. Sulphate in sol. pH °C gr/a gr/Ji 125 5.68 31.35 1.78 150 4.65 28.65 1.42 175 3.50 26.40 1.15 200 2.45 23.80 0.98 225 1.46 21.60 0.85 250 0.90 20.20 0.79 Table 12. Values of log K for the hydrolysis of aluminum sulphate solutions with excess sulphuric acid at 225°C and 250°C Init. pH 0.70 1.00 1.50 2.00 2.50 3.10 log K225 -3.43 -2.96 -2.42 -2.16 -1.93 -1.77 log K250 -1.00 0.00 0.513 0.503 0.374 0.35 - 48 -- 49 -Mean equilibrium constants for 225°C and 250°C were calculated from Table 12 and were found to be: log K225<>c = -2.44 K = 3.62 x 10 log K250oc = 0.302 K = 2.0 Table 13 presents the equilibrium constants for hydrolysis reaction of aluminum sulphate in the temperature range 125-250°C. Table 13. Equilibrium constants for the hydrolysis reaction of aluminum sulphate 1/T x 104 log K 1/T°K T°C 25.10 -18.92 1.2 x 10"19 125 23.60 -13.20 6.3 x IO-14 150 22.30 - 8.63 2.3 x 10"9 175 21.10 - 5.05 8.9 x IO-6 200 20.05 - 2.44 3.6 x 10~3 225 19.10 +• 0.302 2.0 . 250 Since the plot of log K against 1/T gives a straight line (Fig. 19) it is most probable that there is only one solid phase in equilibrium with the liquid phase. Therefore the solution composition is above the invariant point P (Fig. 17) at all temperatures. ,9 ~^~0 3 Zi ?i 24 .^5 Fig. 19.THE CHANGE OF EQUILIBRIUM CONSTANT WITH TEMPERATURE - 51 -4.4 Ternary Diagrams for the System Al^O^-SOy-H^O Some data is available on the hydrolysis of basic aluminum sulphate 3 at temperatures up to 220°C, but no useful data is available on hydrolysis of acid aluminum sulphate solutions. Experimental data from Section 4.2 was used to calculate the equilibrium liquid composi tion in acid aluminum sulphate solutions at 225 and 250°C. Ternary diagrams were constructed for the water corner of the A1203"S03-H20 system at 225 and 250°C. Tables 14 and 15 show the composition of the initial and final solutions at 250 and 225°C. Table 14. Solution composition at 250°C Starting liq. comp. wt % Equil. liq. comp. wt % A1203 so3 H20 A1203 so3 H20 . 1 1.12 4.10 94.78 0.51 3.52 95.98 2 1.12 3.59 95.29 0.40 2.84 96.76 3 1.12 3.00 95.88 0.23 2.07 97.70 4 1.12 2.77 96.11 0.184 1.79 98.03 5 1.12 2.69 96.20 0.18 1.72 98.10 6 1.12 2.64 96.24 0.168 1.63 98.20 - 52 -Table 15. Solution composition at 225°C Starting liq. comp. wt % Equil. liq. comp. wt % A1203 so3 H20 A1203 so3 HgO 1 1.12 4.10 94.78 0.91 3.81 95.28 2 1.12 3.59 95.29 0.66 3.11 96.23 3 1.12 3.00 95.88 0.43 2.31 97.26 4 1.12 2.77 96.11 0.345 1.96 97.69 5 1.12 2.69 96.20 0.30 1.84 97.86 6 1.12 2.64 96.24 0.276 1.75 97.97 The data from Tables 14 and 15 are presented graphically in Fig. 20 and 21. The intersection of the initial and final liquid composition lines suggests the possibility of the presence of a second phase in the system at 250 and 225°C. The second phase is probably a normal aluminum sulphate but it could not be detected in any of the experiments from the X-ray diffraction patterns of the solid. Although a number of experiments with solutions of the composition corresponding to the points A^ and A2 in Fig. 20 and 21 were done, the presence of a second phase could not be unambiguously proven. Acid aluminum sulphate solutions (compositions A^ and A2 - Figs. 20 and 21) were hydrolysed for predetermined times, and solution samples were taken. The autoclave and the remaining material was cooled to room temperature, and the resulting solution was analyzed. From the difference in solution compositions, the S03:A1203 ratio of the re-dissolved precipitate was calculated. It was - 53 -- 54 -- 55 -found to be slightly increased but it was far from the SO^.-AJ^O^ ratio of a normal salt. 4.5 Mechanism of the Hydrolysis of Aluminum Sulphate In aqueous solutions aluminum sulphate is at least partially dissociated into its constituents. Al2(S04)3(aq) ^ 2A1 + 3S04 (8) The Al ion in solution will react with ^0 to give: Al"1"1"1" + H20 * AKOH)"1^ + H+ (9) The concentration of A1(0H) in solution will depend on the concentra--H-tion of Al and H in the intial solution. A1(0H) cation in solution reacts with HS04 ion to give the basic aluminum sulphate: eAKOH)"1"1" + 4HS04" + 8H20 —»- 3A1203.4S03.9H20 + 8H+ (10) 22 The equilibrium.constant for reaction (9) was reported by Helgeson -4. 75 and it is 10 * for dilute solutions. Because of the change in the dissociation constant of water with temperature this constant will change, and become more and more positive as the temperature is increased. Helgeson gives the equilibrium constant values for the reaction: i A1(0H) >- Al + OH (11) - 56 -Reaction (9) is simply a combination of reaction (11) and H20 -—v H+ + 0H~ (12) Knowing the equilibrium constants at temperatures up to 300°C for reactions (11) and (12) the equilibrium constant for reaction (9) can be calculated from the relation: log K9 = log K12 - log Kn Calculated values for are given in Table 16. I | | Table 16. Equilibrium constants for reaction Al + R^O «-A1(0H)++ + H+ Temp °C 25 50 60 100 150 200 . 250 300 log K -4.75 -3.87 -3.53 -2.26 -0.84 +0.63 +1.97 +3.31 Precipitation reaction (10) is a fast reaction as can be seen from the precipitation curves at high temperatures where it reaches equilibrium during the heating period. 4.6 The Effect of Alkali Metal Sulphates on Hydrolysis of Aluminum Sulphate at 225°C When sulphates of Li, Na or K are present in aluminum sulphate solutions during hydrolysis, corresponding alunites of the general form - 57 -M2O.3Al2O2.4sO3.6H2O are precipitated instead of basic aluminum sulphate. Since the solubility of these compounds is less than the solubility of basic aluminum sulphate ("Hydrogen Alunite") the hydrolysis yields are higher and very close to 100%. 3 P.T. Davey and T.R. Scott have reported complete precipitation of aluminum in 30 min.at 220°C from solutions with initial SO^iA^O^ ratio of. 2.87. The ratio of I^SO^:Al2(SO^)3 calculated on the basis of aluminum present in starting solution was 2.16. The precipitated compound was of a composition corresponding to K2O.3Al2O3.4SO3.6H2O. 13 V.S. Sazhin, A.K. Zapolskii and N.N. Zakharova have studied hydrolysis of aluminum sulphate solutions having an initial molar ratio of I^SO^iA^CSO^^ of 0.33. Concentration of aluminum sulphate in solution was 308 gr/& (0.9 M), having S03:A1203 ratio of 3.0. Such solutions were hydrolyzed for 1 hr in the temperature range 175-250°C. Maximum hydrolysis yield at 250°C was about 82%. Their results show the presence of two solid phases. In the temperature range 175-190°C, basic potassium-aluminum sulphates of the form K^SO^.3AI2O3.4SO3.9H2O are precipitated. At temperatures higher than 230°C, alunites are formed with the formula K^O .3AI2O3.4SO3.6H2O. In the temperature range 190-230°C they have reported a mixture of the two. In order to see the effect of alkali metal sulphates on hydrolysis of dilute solutions of aluminum sulphate, solutions with M2S04:Al2(30^)3 ratio of about 0.5 containing 6.10 gr/£ aluminum with S03:A1203 ratio of about 3.0 were hydrolyzed at 225°C until equilibrium was established. ^SO^ corresponds to I^SO^, Na2S0^ and K2S04. - 58 -4.6.1 The Effect of Lithium Sulphate on Hydrolysis of Aluminum  Sulphate at 225°C Aluminum sulphate-lithium sulphate solution containing 6.12 gr/£ aluminum, 38.30 gr/£ SO^ and 0.815 gr/£ lithium was hydrolyzed at 225°C until equilibrium was reached. The results obtained are presented in Table 17 and Fig. 22. Table 17. Hydrolysis of aluminum sulphate-lithium sulphate solution at 225°C* I | | = — Time at Al in sol SO, in sol Li in sol Al prec SO, prec Li prec pH 225°C hrs gr/£ gr/£ gr/£ gr/£ gr/£ ,gr/£ 0 1.74 27.87 0.44 4.38 10.43 0.375 0.82 1 1.29 26.71 0.40 4.83 11.59 0.415 0.78 2 1.16 26.48 0.39 4.96 11.82 0.425 0.75 3 1.07 26.36 0.39 5.05 11.94 0.425 0.71 3 1.11 26.36 0.39 5.01 11.94 0.425 0.73 5 1.05 26.37 0.39 5.07 11.93 0.425 0.72 8 1.05 26.36 0.39 5.07 11.94 0.425 0.72 Starting solution: 38.30 gr/£ SO 6.12 gr/£ Al 0.815 gr/£ Li pH = 3.05 o c < 6.0 -I ?\ DLi -1.0 36 S.O'-3.0-1.0 -O Al 0 SO4 40_| ;i\ V v V — _ 30 ^| -0.6 2 8 o o c (0 » . x — c: o ^ a. •_"] o to 34 V PH _08 3 2 26 20-j V • n n n n -0.4 24 -O O - 22 --0.2 2 0 • VO 0 12 3 4 5 6 7 8 9 10 II 12 13 time in hrs. Fig. 2 2. HYDROLYSIS OF LITHIUM-ALUMINIUM SULPHATE SOLUTION AT 2 25C° - 60 -The precipitated compound was analyzed (see App. II) and its composition was found to be Li2O.3Al2O2.4SO2.6H2O. In a four component system such as this, it is normal to have at least two solid phases in equilibrium with the liquid phase, but since the concentration of Li+ ions in solution is high enough to precipitate all the aluminum present in the starting solution as lithium alunite, only one solid phase is precipitated. The X-ray diffraction pattern for both basic aluminum sulphate and basic lithium-aluminum sulphate is almost the same, so it is difficult to distinguish between the two (see App. III-K and III-E). 4.6.2 The Effect of Sodium Sulphate on Hydrolysis of Aluminum  Sulphate at 225°C. Aluminum sulphate-sodium sulphate solution containing 6.12 gr/£ Al4-1"4", 2.625 gr/£ Na+ and 38.14 gr/£ S04= was hydrolyzed at 225°C. About 5.0 gr/£ of aluminum was precipitated while heating the system to 225°C. When the equilibrium was reached about 98% of the aluminum present in the initial solution was precipitated as a basic salt. The results obtained in this set of experiments are presented in Table 18 and Fig. 23. Analysis of precipitated solid phase shows the composition corresponding to Na2O.3Al2O2.4SO2.6H2O (see App. II). The X-ray diffraction pattern of this compound is the same as for basic aluminum sulphate and lithium alunite (see App. III-H). - 61 -Table 18. Hydrolysis of aluminum sulphate-sodium sulphate solution at 225°C* Time at 225°C hrs Al in sol gr/a SO. in sol 4 gr/il Na+ in sol gr/il Al prec gr/il SO^ prec gr/il Na prec gr/il pH 0 1.05 26.06 1.44 5.07 12.08 1.185 0.77 1 0.41 24.58 1.00 5.71 13.56 1.625 0.70 2 0.24 24.19 0.925 5.88 13.95 1.70 0.69 2 0.32 24.06 0.925 5.80 14.08 1.70 0.69 4 0.17 24.04 0.925 5.95 14.10 1.70 0.63 8 0.13 23.84 0.925 5.99 14.20 ' 1.70 0.63 * = Starting solution: 38.14 gr/Jl SO^ 6.12 gr/Jl Al 2.625 gr/il Na pH = 3.09 4.6.3. The Effect of Potassium Sulphate on Hydrolysis of Aluminum  Sulphate at 225°C Potassium sulphate-aluminum sulphate solution containing 6.12 gr/il aluminum, 4.465 gr/£ potassium and 38.14 gr/£ SO^ was hydrolyzed at 225°C. About 97% precipitation occurs in heating the system to 225°C. 100% precipitation occurs in only a few minutes at 225°C. Results obtained are presented in Table 19 and Fig. 24. Thev.precipitated compound was analyzed and it was found to correspond to the formula K20.3A1203.4S03.6H20 (see App. II). 0 I 2 3 4 5 6 7 8 9 10 II 12 13 timeinhrs. Fig. 2 3. HYDROLYSIS OF SODIUM-ALUMINIUM SULPHATE SOLUTION AT225C° - 63 -Table 19. Hydrolysis of aluminum sulphate-potassium sulphate solution at 225°C* Time at 225°C hrs Al in sol gr/£ SO, in sol 4 gr/£ K in sol gr/Jt Al prec gr/£ SO, prec 4 gr/£ K prec gr/£ pH 0 0.21 .24.06 1.61 5.91 14.08 2.855 0.69 1 o.oo • 23.54 1.51 6.12 14.60 2.955 0.67 2 0.00 23.56 1.51 6.12 14.58 2.955 0.66 3 0.00 23.54 1.51 6.12 14.60 2.955 0.66 Starting solution: 38.14 gr/£ SO^" 6.12 gr/il Al 4.465 gr/£ K pH = 3.09 When alkali metal sulphate-aluminum sulphate solutions are hydrolyzed at 225°C only one solid phase is found to be in equilibrium with the liquid phase. The solid phase corresponds to the formula 13 M20.3A1203.4S03<6H20. This is contrary to the results of V.S. Sazhin. If there is any temperature limit in formation of M20. 3A1203« 4S0.J. 6H20 it should be below 225°C and not 230°C. The most probable.second phase in such mixed solutions would be the basic aluminum sulphate under the condition that there is Insufficient M+ present in the initial solution to precipitate all aluminum in the form of alunite. The reaction for • the hydrbthermal precipitation of alunites can be written as ekl^ + 4HS0 " + 2M+ + 12H.0 *• Mo0.3A1.0o.4S0„.6Ho0 + 16H+ 4 2 -< 2 2 3 3 2 (13) o « -< 5.0 A 4.0 -3.0 -2.0 -1.0 -0.0-0 6 k V \ ^o—o—o \ \ Ov ^•—o—o—o-O Al _: o v •» o c - n"t~: o O ^ « pH (A o c x - 36 D K 3.0 -0.8 34 0 S04 - 32 -2.0 -0.6 30 -V PH *• • • — 28 • 1.0 -0.4 26 "• - 24 OJO -0.2 2 2 • 20 ON 4>-0 I 2 3 4 56 7 8 9 10 11 12 time in hrs. Fig. 24. HYDROLYSIS OF POTASSIUM—ALUMINIUM SULPHATE SOLUTION AT 225 C° - 65 -The effect of alkali, metal sulphate salts is such that the hydrolysis yields are increased in the following series K > Na > Li. This series is the same if the ionic radii of the Li, Na, and K are compared. Table 20 shows the ionic radii of the ions involved in this system. Fig. 25 shows the straight line relation between the ionic radii and precipitated amount of aluminum in heating the system to 225°C. Table 20. Ionic radii of the ions involved in the system Element Type of Ionic radius Al prec. Radius A° % H 1+ 0.00 48.00 Al 3+ 0.57 Fe 3 + 0.67 Li 1+ 0.68 71.80 Fe 2+ 0.80 60.50 Cu 2+ 0.80 61.40 Na 1+ 0.98 83.00 K 1+ 1.33 96.80 4.7 The Effect of Divalent Metal Sulphates on Hydrolysis of Aluminum Sulphate Solutions at 225°C As a general rule, ions of divalent metals do not precipitate with basic aluminum sulphate. Many trivalent or quadrivalent metals precipitate as oxides, hydroxides or basic salts under the same - 66 -- 67 -conditions as for hydrolysis of aluminum sulphate. The effect of copper and iron was investigated at 225°C mostly because of their presence in many leach solutions. 4.7.1 The Effect of Copper Sulphate on Hydrolysis of Aluminum  Sulphate Solutions at 225°C A copper- sulphate-aluminum sulphate solution containing 6.10 gr/£ aluminum, 3.58 gr/£ copper and 37.92 gr/£ SO^ was hydrolyzed at 225°C. There was practically no copper precipitation detected in these experiments. Results obtained are presented in Table 21 and Fig. 26. Table 21. Hydrolysis of aluminum sulphate-copper sulphate solution at 225°C* Time at 225°C Al in sol Cu in sol SO. in sol 4 Al prec Cu prec SO^ prec pH hrs. gr/il gr/Jl gr/£ gr/jl gr/£ gr/£ 0 2.35 3.56 29.02 3.75 0.02 8.90 0.94 1 1.48 3.58 26.92 4.62 0.00 11.00 0.85 2 1.23 3.57 26.32 4.87 0.01 11.60 0.82 4 1.12 3.56 26.06 4.98 0.02 11.86 0.82 6 1.10 3.58 26.06 5.00 0.00 11.86 0.82 Starting solution: 6.10 gr/£ Al 3.58 gr/£ Cu 37.92 gr/£ S04 pH = 3.0 ON 00 I 0 I 2 34 56 7 8 9 10 I I 12 time in hrs. Fig.26. HYDROLYSIS OF COPPER-ALUMINIUM SULPHATE SOLUTION AT 225 C° - 69 -4.7.2 The Effect of Ferrous Sulphate on Hydrolysis of Aluminum  Sulphate Solutions at 225°C A ferrous sulphate-aluminum sulphate solution containing 6.12 gr/£ aluminum, 3-26 gr/SL iron and about 38.29 gr/£ SO^- was hydrolyzed at 225°C. Some iron was precipitated in the form of ¥e^0^ due to atmospheric oxidation of ferrous iron. Results obtained in this set of experiments are presented in Table 22 and Fig. 27. The increased precipitation of aluminum can be attributed to the common ion effect. The bivalent ions are not compatible with the alunite lattice and do not tend to replace the hydrogen ions in basic 3 aluminum sulphate.. T.R. Scott has reported that no contamination of basic aluminum sulphate occurs in the presence of Mg, Cd, Zn and Ni even at high concentrations. Table 22. Hydrolysis of aluminum sulphate-ferrous sulphate solution at 225°C* Time at Al in sol Fe in sol SO. in sol Al prec Fe prec SO, prec pH hrs gr/£ gr/£ gr/Jl . gr/£ gr/Jl gr/Jl 0 2.42 2.69 29.44 3.70 0.57 8.85 0.97 1 1.57 2.68 27.48 4.55 0.58 10.81 0.82 2 1.40 2.69 27.04 4.72 0.57 11.25 0.80 4 1.15 2.66 26.42 4.97 0.60 11.87 0.82 6 1.12 2.69 26.40 5.00 0.57 11.89 0.80 Starting solution: 6.12 gr/£ Al; 3.26 gr/l Fe; 38.29 gr/£ SO pH = 3.0 I 8 --i 1 i —J 1 i '«»'•'•''' 0 I 23 4 56 7 89 10 II 12 time in hrs. Fig. 27 HYDROLYSIS OF IRON-ALUMINIUM SULPHATE SOLUTION AT 22 5C° - 71 -4.8 Application of the Hydrolysis Process Basic aluminum sulphate of the form 3A1203.4S03.9H20 is the product of high temperature hydrolysis of aluminum sulphate solutions with a wide range of initial SO^Al^ ratios. A basic salt of the form M2O.3Al2O2.4sO2.6H2O is a high temperature hydrolysis product of aluminum sulphate solutions containing alkali metals (Li, Na or K). Thermal decomposition of such basic salts yields alumina as a final product which can be used for aluminum production. Thermal decomposi tion of a few samples of basic aluminum sulphate was followed by D.T.A. A typical D.T.A. curve is shown in Fig. 28. The first endothermic peak in the temperature range 200-400°C corresponds to the loss of interstitial lattice water.9 The peak at about 450°C corresponds to the dehydroxylation process. This process is completed at temperatures around 600°C where anhydrous basic aluminum sulphate 9 exists. At temperatures above 600°C, SO^, SO2 and O2 are evolved yielding amorphous alumina which subsequently crystallizes to y-alumina at higher temperatures. Y-Alumina is transformed to a-alumina at temperatures of 1000°C and higher. A few samples of K2O.3Al2O2.4SO2-6H2O were calcined at about 1000°C for 2 hrs. The decomposition of these samples was followed by the weight loss. Results of one of the decomposition experiments is shown below. 1.0811 gr K20.3A1203.4S03.6H20 > 0.5857 gr K2S04 + Al 0 + , sulphurous t gases 0.5857 gr K2S04 4- A1203 iffg^ 0.365 gr A1203 o 200 400 600 Fig. 2 8.D.T.A. OF 3AI2O34SO39H2O - 73 -In leaching processes at high temperatures where aluminum is present in solution it precipitates as a basic salt in autoclaves and pipelines"'"''" which necessitates "forced" shutdowns of the process in order to clean autoclaves and/or pipelines. With increasing sulphuric acid concentration and decreasing temperature the solubility of basic aluminum sulphate was found to increase. Therefore running concentrated solutions of sulphuric acid through the system at low temperatures should dissolve the precipitate. In dump leaching solutions aluminum is present as aluminum sulphate. If aluminum is not recovered the viscosity of the solution will increase. To remove the, aluminum from solution, a part of the liquor after cementation can be treated in an autoclave to precipitate basic aluminum sulphate. Liberated acid from the hydrolysis can be recycled to the heap. - 74 -5. CONCLUSIONS 1. Dilute solutions of normal and acid aluminum sulphate hydrolyze when heated above 125°C to yield a basic salt of the nominal composition 3A1203,4S03.9H20. 2. Hydrolysis yield is a function of the temperature and initial concentration of sulphuric acid in solution. 3. Only one solid phase was found to be in equilibrium with the liquid phase in the temperature region 125-250°C, but there are some indications that a less basic salt may exist in the system. 4. When Li, Na or K ions are present in solutions of aluminum sulphate a basic salt of the form M20.3A1203«4S03.6H20 is precipitated. 5. The hydrolysis yields are increased in the presence of alkali metal sulphates in the following series: K > Na > Li. I | 6. Aluminum can be selectively precipitated in the presence of Cu and Fe"1"1" as 3A1203> 4S03-9H20. 7. Both basic aluminum sulphate and alunites can be calcined at temperatures above 1000°C to yield alumina. - 75 -APPENDIX I SOLUBILITY OF ALUMINUM SULPHATE I-A Solubility of Al (SO.K in water Temp. °C 20 30 40 50 60 70 80 90 100 gr A12(S04)3 26.7 28.8 31.4 34.3 37.2 39.8 42.2 44.7 47.1 100 gr sat. sol. I-B Solubility of aluminum sulphate in aqueous solutions of sulpuric acid at 25°C. gr A12(S04)3 27.82 29. 21 26.2 19.5 11.6 4. 8 1. 5 1.0 2.3 4.0 100 gr sat. sol. gr H2S04 0.0 5. 73 10.0 20.0 30.0 40. 0 50. 0 60.0 70.0 75.0 100 gr sat. sol. I-C Solubility of aluminum sulphate in aqueous 10% H SO Temp. ° C 30 42 50 gr A12(S04)3 14.52 16.45 18.77 100 gr sat. sol. - 76 -APPENDIX II ANALYSIS OF THE SOLID PHASES Time at Temp. % Al 0 % SO % M00 % HO Starting hrs. C pH 6.00 175 38.7 37.2 - 24.1 3.10 10.00 175 38.5 39.6 - 21.9 3.10 16.30 175 38.7 39.7 - 21.6 3.10 4.0 200 38.6 39.5 - 21.9 3.10 6.0 200 38.8 39.9 - 21.3 3.10 18.00 200 38.7 40.1 - 21.2 3.10 1.00 225 38.8 40.2 - 21.0 3.10 5.00 225 38.9 40.8 - 20.3 3.10 16.00 225 39.1 41.6 - 19.3 3.10 0.00 250 38.8 40.3 - 20.9 3.10 1.00 250 38.7 40.2 - 21.1 3.10 3.30 250 39.0 40.5 - 20.05 3.10 6.30 250 39.2 40.9 - 19.9 3.10 6.00 225 39.0 . 40.6 - 20.4 2.5 6.00 225 38.7 40.6 - 20.7 2.0 6.00 225 38.9 40.8 - 20.3 1.5 6.00 225 . 38.9 40.2 - 20.9 1.0 6.00 225 39.2 41.7 - 19.1 0.70 4.00 250 38.9 40.6 - 20.5 2.5 4.00 250 38.8 40.7 - 20.5 2.0 4.00 250 38.9 41.2 - 19.9 1.5 4.00 250 39.2 41.6 - 19.20 1.0 4.00 250 38.9 40.8 - 20.3 0.70 Li20 3.00 225 39.90 42.00 4.00 14.10 3.05 Na20 2.00 225 38.40 40.10 7.70 13.80 3.09 8.00 225 38.50 40.20 7.80 13.50 3.09 K20 0.00 225 37.05 38.39 11.36 13.20 3.09 0.00 225 36.91 38.91 11.38 12.80 3.09 1.00 225 37.12 38.48 11.30 13.10 3.09 - 77 -APPENDIX III X-RAY DIFFRACTION PATTERNS III-A Diffraction pattern of the precipitate obtained at 150°C Prec. obtained in 10 hrs. Prec. obtained in 16 hrs d A° I/xo d A° I/To 5.71 7 5.71 7 5.03 2 5.03 2 3.55 5 3.57 5 3.01 1 3.01 1 2.27 4 2.27 4 1.909 3 1.908 3 1.76 5 1.755 5 1.654 7 1.659 7 1.566 7 1.570 7 1.496 5 1.500 5 1.325 7 1.329 7 1.290 6 1.299 6 1.220 7 1.217 7 1.171 7 1.171 7 1.150 7 1.147 7 III-B X-ray diffraction patterns of the precipitate obtained at 175 Prec. obtained in 2 hrs. Prec. obtained in 16.5 hrs d A° I/I0 d A° 1/To 5.71 7 5.75 1 5.034 2 5.12 2 3.56 5 3.56 5 3.01 : 1 3.04 1 2.87 7 2.486 8 2.27 4 2.30 4 2.22 8 1.908 3 1.920 3 1.757 5 1.770 5 1.654 7 1.654 7 1.566 7 1.570 7 1.495 5 1.500 5 1.380 8 1.325 7 1.330 7 1.293 6 .1.290 6 1.219 7 1.217 7 1.174 7 1.172 7 1.150 7 1.150 7 - 78 -III-C X-ray diffraction pattern of the precipitate obtained at 200°C. Prec. obtained in 18 hrs. d A° . I/I 5.72 7 5.03 2 3.56 5 3.04 1 2.87 7 2.27 4 2.22 8 1.910 3 1.756 •5 1.654 7 1.565 7 1.490 5 1.387 8 1.290 6 1.217 7 1.171 7 1.151 7 III-D X-ray diffraction pattern of the precipitate obtained at 225°C Prec. obtained in 3 hrs at 225°G d A° I7I0 d A° I/I 5.61 8 1.736 6 5.02 2 1.666 8 4.80 2 1.641 8 3.43 6 1.550 8 3.03 1 1.546 8 2.932 1 1.479 4 2.806 4 1.4227 8 2.440 8 1.362 8 2.304 4 1.316 8 2.241 3 1.281 4 2.191 8 1.205 8 1.925 8 1.195 8 1.880 3 1.161 8 1.767 8 1.132 8 1.758 6 - 79 -III-E X-ray diffraction pattern of the precipitate obtained at 250°C Prec. obtained in 0.0 hrs. Prec. obtained in 3.5 hrs. d A° ^0 d A° I/To d A° I/xo d A° l/1o 5.61 8 1.765 8 5.61 8 4.02 2 1.736 6 5.01 2 1.742 6 4.80 2 1.665 8 4.85 2 1.646 8 4.27 8 1.557 8 - 1.555 3.40 6 1.479 4 3.46 6 1.483 4 3.01 1 1.424 8 3.01 1 -2.91 1 1.367 8 2.95 1 1.369 8 2.79 4 1.317 8 2.81 4 1.317 8 2.45 8 1.286 4 2.46 8 1.286 4 2.296 4 1.208 8 - 1.208 8 2.24 3 1.200 8 2.24 3 -2.187 8 1.167 8 1.895 8 1.167 8 1.929 8 . 1.138 8 1.139 8 1.879 3 III-F X-ray diffraction pattern of the precipitate obtained from from FeSO -Al„(SO 4 2 solution at 225°C Prec. obtained in4.0hrs at 250°C d A° 1/To d A° I/I 5.08 2 1.558 4 4.80 2 1.501 5 3.03 1 1.481 2 2.915 1 1.423 5 2.294 5 1.371 3 2.234 1 1.314 4 1.918 4 1.209 3 1.888 2 1.1969 5 1.768 - 5 1.1867 5 1.737 2 1.1364 5 1.637 4 1.1054 5 - 80 -III-G X-ray diffraction pattern of potassium alunite obtained at 225°C Prec. obtained inO.Ohrs. at 225°C d A° I/I d A° I/Ir 5.71 5 1.560 8 5.01 2 1.546 8 3.49 3 1.489 3 3.01 1 1.422 8 2.90 8 1.380 8 2.47 8 1.316 8 2.29 3 1.281 5 2.20 5 1.205 5 1.888 2 1.161 6 1.736 2 1.132 6 1.641 8 III-H. X-ray diffraction pattern of sodium alunite obtained at 225°C Prec. obtained in 2.0 hrs. at 225°C d A° X/I0 d A° 5.70 5 1.746 2 5.01 2 1.644 8 3.496 3 1.554 8 3.01 1 1.538 8 2.91 8 1.505 3 2.45 8 1.281 5 2.212 3 1.205 5 1.896 2 1.161 6 III-K X-ray diffraction pattern of lithium alunite obtained at 225°C Prec. obtained in 3.0 hrs. at 225°C d A° : Wo d A° w„ 5.68 : 5 1.736 2 5.01 2 1.641 8 3.491 3 1.554 8 3.01 1 1.535 8 2.90 8 1.489 3 2.44 8 1.281 5 2.20 3 1.205 5 1.887 2 1.161 6 - 81 -LITERATURE 1. E.T. Carlson, C.S. Simons, Met. Soc. of A.I.M.E., Extractive Metallurgy of Copper, Nickel and Cobalt. International Symposium, New York, February 15-18, 1960. 2. Posnjak, E., Mervin, H.E., J. Am. Chem. Soc. 44, 1965-1994 (1929). 3. P.T. Davey and T.R. Scott, Aust. J. Appl. Sci. 13, 229-241 (1962) 4. Basset, H. and Goodwin, T.H., J. Chem. Soc., 2239-2279 (1949). 5. T.R. Scott, Paper presented at the Annual Conference of the Institute, August 14-24, 19|>3, Melbourne. 6. T.R. Scott, Extractive Metallurgy of Aluminum, Vol. 1, Alumina, International Symposium on the Extractive Metallury of Aluminum, February 18-22 (1962), New York. 7. T.R. Scott, Research Appl. Ind. 14, 50-54 (1961). 8. P.T. Davey and T.R. Scott, Nature 195, 376 (1962). 9. P.T. Davey, G.M. Lukaszewski and T.R. Scott, Aust. J. Appl. Sci. 14, 137-154 (1963). 10. Henry, J.L. and G.B. King, J. Am. Chem. Soc. 72_, 1282-1286 (1950). 11. R. Acosta Chaves, V.V. Karelin, and B.P. Sobolev, Tsvet. Met. 1968, 41 (4), 50.53. 12. T.R. Scott, Internal Report, Division of Mineral Chemistry C.S.I.R.O., Melbourne. 13. V.S. Sazhin, A.K. Zapolskii and N.N. Zakharova, Zh. Prikl. Khim., 41 (7), 1420-1423 (1968). 14. S. Bretsznajder, J. Boczar, J. Piskorski, and J. Porowski, Prezmysl. Chem. 11, 89-93 (1955). 15. A.K. Zapolskii and G.I. Tsarenko, Ukr. Khim. Zh. 1969, 35 (8), 866-868. - 82 -16. V.S. Sazhin, A.K. Zapolskii, N.N. Zakharova and A.I. Volkovska, Ukr. Khim. Zh. 1966, 32 (1), 95-100. 17. A.K. Zapolskii, Ukr. Khim. Zh. 33 (8), 805-809 (1967). 18. A.K. Zapolskii, V.S. Sazhin, N.N. Zakharova and A.I. Volkovska, Ukr. Khim. Zh. 32 (11), 1222-7 (1966). 19. H.G. Iverson and H. Leitch. Report of Investigations 7162 (1968). 20. Erkki Wanninen and Anders Ringbom, Analytica Chimica Acta Vol. 12 (1955), 308-318. 21. R.G. Robins, LR 80 (NST) Report of the Mineral Sci. and Technology Div. of the Warren Spring Laboratory. 22. H.C. Helgeson, Am. J. of Sci. Vol. 267, 1969, 729-804. 

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