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The role of surfactants in the leaching of zinc sulphide minerals at temperatures above the melting point… Owusu, George 1993

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We accept this thesis as conformingto the required standardTHE UNIVERSITY OF BRITISH COLUMBIAJune 1993© George Owusu, 1993THE ROLE OF SURFACTANTS IN THE LEACHING OF ZINC SULPHIDEMINERALS AT TEMPERATURES ABOVE THE MELTING POINT OF SULPHURByGeorge OwusuB.Sc.(Hons) Metallurgical Eng., University of Science & Tech., Kumasi, Ghana, 1985M. A. Sc. Metallurgical Eng. University of British Columbia, Vancouver, Canada, 1989A THESIS SUBMITTED IN PARTIAL FULFILLMENT OFTHE REQUIREMENTS FOR THE DEGREE OFDOCTOR OF PHILOSOPHYinTHE FACULTY OF GRADUATE STUDIESMETALS AND MATERIALS ENGINEERINGIn presenting this thesis in partial fulfilment of the requirements for an advanced degree atthe University of British Columbia, I agree that the Library shall make it freely availablefor reference and study. I further agree that permission for extensive copying of thisthesis for scholarly purposes may be granted by the head of my department or by hisor her representatives. It is understood that copying or publication of this thesis forfinancial gain shall not be allowed without my written permission.Metals and Materials EngineeringThe University of British Columbia2075 Wesbrook PlaceVancouver, CanadaV6T 1W5Date: (1(4,/7(3ABSTRACTThe role of the surfactants lignin sulfonic acid, cocoamido hydroxyl sulfobetaine (CAHSB),tallowamido hydroxyl sulfobetaine (TAHSB), cocoamido betaine (CAB), naphthalene sul-fonic acid-formaldehyde condensates and orthophenylene diamine (OPD) in the liquidsulfur-aqueous zinc sulphate-zinc sulfide mineral system has been studied. The effectof the various surfactants on the liquid sulfur-aqueous system interfacial tensions andthe liquid sulfur-mineral contact angle were measured in a high temperature-high pres-sure apparatus. In the absence of any surfactant, the liquid-liquid interfacial tensionmeasured 54±1 mN/m, and the liquid sulfur-mineral contact angle was 80±5°. In thepresence of 0.3 g/L lignin sulfonic acid, the interfacial tension decreased to about 29±1mN/m; the interfacial tension was not sensitive to any further increase in the surfactantconcentration. The contact angle measured 145±5° in the presence of 0.07-0.3 g/L ofthe surfactant. The presence of OPD in the solution did not have any effect on the liq-uid sulfur-aqueous solution interfacial tension but increased the contact angle to about130±5°. At a concentration of 0.1 g/L, naphthalene sulfonic acid-formaldehyde con-densates decreased the interfacial tension to 21±1 mN/m. Beyond this concentration,the liquid-liquid interface was not sensitive to any further increase. The contact angleincreased to 155±5° even at concentrations as low as 0.05 g/L. CAHSB exhibited justabout the same efficiency as lignin sulfonic acid. TAHSB and CAB required higher con-centrations in order to effect any substantial changes in the liquid sulfur-aqueous solutioninterfacial tension. They also effected a contact angle increase to 135±5°. With the ex-ception of OPD, all the surfactants adsorbed at both the liquid-liquid and aqueous-solidinterfaces. OPD, on the other hand adsorbed at only the aqueous-solid interface.iiWhen introduced into leaching systems, the surfactants influenced the metal extrac-tion to different extents. OPD and lignin sulfonic acid were the most efficient surfactants(or dispersants for liquid sulfur). The extent of zinc extraction was enhanced to about94 % for lignin sulfonic acid and over 99 % for OPD under both low or high pulp densityconditions. The other surfactants were unable to enhance zinc extraction even underlow pulp density conditions. Application of metaphenylene diamine as dispersant forliquid sulfur under both low and high pulp density leaching conditions resulted in zincextractions in excess of 98 %. A combination of the interfacial studies and the leachingexperiments indicates that the most important characteristic for any effective surfac-tant under pressure leaching conditions is for the surfactant to adsorb at the aqueousphase-solid interface and increase the liquid sulfur-mineral contact angle.Infrared spectral analysis of both sphalerite leach residue and pulped sphalerite in-dicates that the lignin sulfonate ions are adsorbed both physically and chemically bythe mineral. The chemically adsorbed species form a Zn(II)-based complex that has anorganic portion similar to the bulk surfactant structure. OPD, on the other hand, isadsorbed chemically through the interaction of the C-N functional group with the metalions forming a metal-amine complex.iiiTable of ContentsABSTRACT^ iiList of Tables viiList of Figures^ xiList of Symbols xviACKNOWLEDGEMENT^ xx1 INTRODUCTION 12 LITERATURE SURVEY 42.1 ELECTROCHEMICAL PROPERTIES OF METAL SULFIDES ^ 52.2 LEACHING OF SULFIDE MINERALS ^ 62.2.1^Leaching of Zinc Sulfide Minerals 92.2.2^The Behaviour of Sulfur in Oxidative Leaching of Sulfide Ores 222.2.3^Effect of Sulfur in Blocking Leaching Reactions ^ 242.2.4^Interfacial Phenomena ^ 252.3 PROPERTIES OF LIQUID SULFUR 262.4 SURFACTANTS AND INTERFACIAL PROPERTIES ^ 272.4.1^Characteristic Features of Surfactants ^ 272.4.2^Surfactants Identified for Study 302.4.3^Properties of Individual Surfactants ^ 31iv2.5 INFRARED SPECTROSCOPY ^ 362.5.1^Theory of Infrared Spectroscopy 372.5.2^Molecular Force Constants ^ 382.5.3^Absorption of Infrared Radiation 402.5.4^Infrared Techniques ^ 423 OBJECTIVES AND EXPERIMENTAL APPROACH 513.1 OBJECTIVES ^ 513.1.1^Scope of Study and Approach ^ 524 EXPERIMENTAL TECHNIQUES 534.1 INTERFACIAL PHENOMENA ^ 534.1.1^Interfacial Tension 564.1.2^Contact Angle Measurements ^ 584.1.3^Work of Adhesion ^ 644.2 PRESSURE LEACHING OF SPHALERITE, (ZnS) ^ 654.2.1^Low Pulp Density Leaching ^ 664.2.2^High Pulp Density Leaching 674.3 INFRARED STUDIES ^ 704.3.1^KBr Technique 714.3.2^Attenuated Total Reflection Technique ^ 715 RESULTS AND DISCUSSION 805.1 INTERFACIAL PHENOMENA ^ 805.1.1^Results ^ 805.1.2^Discussion 845.2 PRESSURE LEACHING OF SPHALERITE CONCENTRATE ^ 1015.2.1 Results ^  1015.2.2 Discussion  1095.3 INFRARED SPECTROSCOPY ^  1325.3.1 Results ^  1325.3.2 Adsorption Studies in the Presence of Orthophenylene Diamine^1566 CONCLUSIONS AND RECOMMENDATIONS^ 1756.1 CONCLUSIONS ^  1756.2 RECOMMENDATIONS  178Bibliography^ 180viList of Tables2.1 Electronic properties of selected sulfide minerals  ^55.2 Surface activity of OPD at a concentration of 0.3g/L in the presence of[Fe2+] = 0.05 M,^= 0.12 M, [H 2 SO4] = 0.2 M, and [ZnSO 4 ] = 1.2M, Temp.= 135±5°C  815.3 Effect of silicate ions (0.15g/L) on the surface activity of OPD (0.3 g/L)in the presence of [H 2 SO4 ] = 0.2 M and [ZnSO 4 ] = 1.2 M, Temp.= 135±5°C 815.4 Contact angle measurements in the presence of the various surfactants:conditions - [ZnSO 4] = 1.2 M, Temp.= 135±5°C ^  825.5 Contact angle measurements in the presence of different surfactants andionic species; conditions: [ZnSO4 ] = 1.2 M, [Fe 2+] = 0.05 M, [Fe 3+] = 0.12M and [H2 SO4] = 0.2 M, Temperature = 135°C   835.6 Work of adhesion, Wa , in the presence of 0.3g/L of different surfactants;conditions: [ZnSO 4 ] = 1.2 M, Temperature = 135°C; no other ionic speciespresent in solution   845.7 Work of adhesion, Wa , in the presence of 0.3g/L surfactants; conditions:[ZnSO 4 ] = 1.2 M, [H 2 SO 4] = 0.2 M, [Fe 2+] = 0.05 M and [Fe'l] = 0.12 M,Temperature = 135°C.   845.8 Resultant work, W, for complete sulfur wetting (a hypothetical situation)in the presence of 0.3 g/L of surfactant and [ZnSO 4] = 1.2 M ^ 905.9 Variation of 7MA - 7MS with respect to aqueous condition; [ZnSO 4] = 1.2M; other ionic species absent ^  91vi i5.10 Extent of zinc extraction using different surfactants under low pulp densityconditions and nitrogen atmosphere; initial conditions: [ZnSO 4] = 2 M;[H2 SO4] = 0.5 M; [Fe3+]=[Fe2+]=0.15 M; Temp.= 140±1°C; PN2 = 1100kPa; B-batch addition; C-continuous addition   1025.11 Effect of naphthalene sulfonic acid on the extent of zinc extraction underlow pulp density conditions; initial conditions: [ZnSO 4] = 2 M; [H 2 SO 4 ]= 0.5 M; [Fe3+1=4Fe21=0.15 M; Temp.= 140±1°C; PN2 = 1100 kPa; B-batch addition; C-continuous addition   1025.12 Extent of zinc leaching using OPD as surfactant under low pulp den-sity conditions; initial conditions: [ZnSO 4] = 2 M; [H 2SO 4] = 0.5 M;[Fe31.[Fe21.0.15 M; Temp.= 140±1°C; PN2 = 1100 kPa; B-batch ad-dition; C-continuous addition   1035.13 Extent of zinc extraction using MPD and PPD as surfactants (batch ad-ditions) under low pulp density conditions; initial conditions: [ZnSO4] =2 M; [1-1 2 SO4] = 0.5 M; [Fe3+]=[Fe21=0.15 M; Temp.= 140±1°C; PN2 =1100 kPa.   1035.14 Extent of zinc extraction after 30 minutes of low pulp density leachingunder oxygen pressure; concentrate sample size: 8-10 g; initial condi-tions: [ZnSO 4] = 2 M; [H2 SO4] = 0.5 M; [Fe3+]=1Fe21=0.15 M; Temp.=140±1°C; P02 = 1100 kPa; Batch addition of surfactants   1055.15 Weight loss due to leaching in the absence and presence of different sur-factants under high pulp density conditions; initial conditions: [ZnSO 4] =1.27 M; [H 2 SO4] = 1.12 M; [Fe 2+]=0.14 M; Temp.= 140±1°C; P 02 = 1100kPa; Batch addition of surfactants   1205.16 Extent of zinc extractions (computed from residue analysis) and the changesin Zn2+ (aq) concentration of the filtrate in the presence of different surfac-tants and under high pulp density conditions; initial conditions: [ZnSO 4 ]= 1.27 M; [H 2 SO4] = 1.12 M; [Fe 2+]=0.14 M; Temp.= 140+1°C; P02 =1100 kPa; batch addition of surfactants. * Residue analysis   1205.17 Iron extractions (based on residue analysis) in the presence of differentsurfactants and under high pulp leaching conditions; initial conditions:[ZnSO 4] = 1.27 M; [H 2 SO4 ] = L12 M; [Fe2+]=0.14 M; Temp.= 140+1°C;P02 = 1100 kPa.   1215.18 Oxygen and acid (as H+ ions) requirements per 100 gram sphalerite con-centrate and 1.5 L leach solution assuming complete oxidation of concen-trate after test period. "x" represents the fractional conversion of sulfurto sulfate   1215.19 Experimental data for oxygen and acid (measured as H+) consumptionsin the presence of different surfactants; initial conditions: [ZnSO 4] = 1.27M; [H 2 SO4] = 1.12 M; [Fe2+]=0.14 M; Temp.= 140+1°C; P 02 = 1100 kPa. 1225.20 Elemental sulfur analysis of leach residue and filtrate and the extent oftotal sulfide sulfur oxidation to sulfate (as %); initial conditions: [ZnSO 4 ]= 1.27 M; [H 2 SO4] = 1.12 M; [F0+]=0.14 M; Temp.= 140+1°C; P02 =1100 kPa.   1225.21 Calculated oxygen and acid (as H+) requirements for the different experi-mental conditions, assuming NO elemental sulfur oxidation   1235.22 Comparison between experimental and calculated oxygen requirements forthe different experimental conditions; Temp.= 140+1°C; P02 = 1100kPa 123ix5.23 Total leaching efficiencies and iron:zinc ratios; initial conditions: [ZnSO 4 ]= 1.27 M; [H 2 SO4] = 1.12 M; [Fe 2+]=0.14 M; Temp.= 140±1°C; P02 =1100 kPa.   124List of Figures2.1 Current density- potential diagram; after Wadsworth [1]  ^472.2 Viscosity of liquid sulfur; after Tuller [78].  ^472.3 Possible structure of lignin sulfonate; after Pearl [86].  ^482.4 Structure of OPD; after reference [92].  ^482.5 Structure of naphthalene sulfonic acid—formaldehyde condensates; afterRosen [83] ^  492.6 Atomic energy levels; after Hair [121] ^492.7 Internal reflection effect; after references [121] ^504.8 Profile of pendant drop.  ^754.9 Contact angle between a liquid and solid surface: a) in the absence of anysurfactant; b) in the presence of a surfactant.  ^754.10 Cross—section of high temperature and pressure stainless-steel bomb.^764.11 Experimental set—up for monitoring oxygen flow rate^ 774.12 Injection Unit.  ^784.13 KBr Transmission. ^  784.14 ATR Accessory  795.15 Liquid sulfur—aqueous solution interfacial tensions in the presence of dif-ferent surfactants; [ZnSO 4 ] = 1.2 M; Temperature = 135+5°C; P NZ = 620kPa.   96xi5.16 Liquid sulfur—aqueous solution interfacial tensions in the presence [Fe 3+]= 0.12 M, [Fe 2+] = 0.05 M and 0.3 g/L naphthalene sulfonic acid; [H2SO4 ]= 0.2 M, Temperature = 135±5°C; PN2 = 620 kPa  965.17 Liquid sulfur-zinc sulfide mineral contact angle in surfactant-free system;[ZnSO 4 ] = 1.2 M, Temperature = 135°C; PN 2 = 620 kPa  975.18 Liquid sulfur-zinc sulfide mineral contact angle in the presence of 0.3g/Lnaphthalene sulfonic acid—formaldehyde condensates, [ZnSO 4 ] = 1.2 M,Temperature = 135°C; PN2 = 620 kPa  975.19 Liquid sulfur-zinc sulfide mineral contact angle in the presence of 0.15g/LOPD; conditions: [Fe 21 = 0.05 M, [Fe3+] = 0.12 M, [H 2 SO4] = 0.2 M,Temperature = 135°C; P N2 = 620 kPa  985.20 Photograph of liquid sulfur suspended in surfactant-free zinc sulfate solu-tion: conditions; [ZnSO 4] = 1.2 M, Temperature = 135±5°C; P N2 = 620kPa.   985.21 Photograph of liquid sulfur suspended in aqueous solution in the presenceof 0.3g/L naphthalene sulfonic acid—formaldehyde condensates; conditions:[ZnSO 4 ] = 1.2 M, Temperature = 135°C; PN 2 = 620 kPa  995.22 Different stages of wetting   995.23 Interfacial excess of naphthalene sulfonic acid—formaldehyde condensatesat the liquid sulfur—aqueous solution interface; conditions: [ZnSO 4] = 1.2M, Temperature = 135±5°C; PN 2 = 620 kPa.   1005.24 Interfacial excess of lignin sulfonic acid at the liquid sulfur—aqueous so-lution interface; conditions: [ZnSO 4 ] = 1.2 M, Temperature = 135±5°C;PN2 = 620 kPa  100xii5.25 Extent of zinc extraction in the presence of OPD, MPD, PPD and naph-thalene sulfonic acid-formaldehyde condensates; Temp.= 140±1°C; P02 =1100 kPa; sample weight --, 12 grams^  1255.26 Extent of zinc extraction under high pulp density leaching conditions inthe presence of OPD; Temp.= 140±1°C; Po, = 1100 kPa. ^ 1255.27 Instantaneous oxygen flow rate in the absence of any surfactant; Temp.=140±1°C; P02 = 1100 kPa. ^  1265.28 Instantaneous oxygen flow rate in the presence of lignin sulfonic acid;Temp.= 140±1°C; P02 = 1100 kPa. ^  1265.29 Instantaneous oxygen flow rate in the presence of OPD; Temp.= 140±1°C;P02 = 1100 kPa. ^  1275.30 Instantaneous oxygen flow rate in the presence of 0.1 g/L MPD; Temp.=140±1°C; Po, = 1100 kPa. ^  1275.31 Cumulative oxygen consumed in the absence and presence of lignin sulfonicacid; Temp.= 140±1°C; Po, = 1100 kPa. ^  1285.32 Cumulative oxygen consumed in the presence of OPD; Temp.= 140±1°C;P 02 = 1100 kPa. ^  1285.33 Cumulative oxygen consumed in the presence of MPD; Temp.= 140±1°C;Po, = 1100 kPa. ^  1295.34 Relative consumptions of oxygen in the presence of the different surfac-tants; Temp.= 140±1°C; Po, = 1100 kPa. ^  1295.35 A photograph of zinc sulfide leach residue in the absence of a surfactantafter 60 minutes; Temp.= 140±1°C; Poe02^1100 kPa^ 1305.36 A photograph of zinc sulfide leach residue in the presence of 0.1 g/L ligninsulfonic acid after 60 minutes;Temp.= 140±1°C; Po, = 1100 kPa^ 1305.37 A photograph of zinc sulfide leach residue in the presence of 0.1 g/L OPDafter 60 minutes; Temp.= 140±1°C; Po, = 1100 kPa^ 1315.38 A photograph of zinc sulfide leach residue in the presence of 0.1 g/L MPDafter 60 minutes; Temp.= 140±1°C; P 02 = 1100 kPa^ 1315.39 Infrared spectrum of sphalerite using KBr pellet technique^ 1465.40 ATR spectrum of wet sphalerite. ^  1465.41 KBr disc spectrum of lignin sulfonic acid. ^  1475.42 KBr disc spectrum of Zn—based lignin sulfonate complex^ 1485.43 KBr pellet spectrum of Na—based lignin sulfonate complex  1495.44 ATR spectrum of lignin sulfonic acid solution; solution condition: [Fe 3+]= 0.025 M, [H2 SO4 ] = 0.036 M, Temperature = 25°C; reference spectrum:solvent.   1505.45 ATR spectrum of sphalerite conditioned in surfactant solution for 20 min-utes; solution condition: [Fe31 = 0.025 M, [H2 SO4 ] = 0.036 M, Temp. =25°C: (a) reference: solvent spectrum (b) less solution spectrum  1515.46 ATR spectrum of washed sphalerite after conditioning in surfactant/ferricsolution, reference spectrum: solvent^  1525.47 ATR spectrum of sphalerite conditioned in ferric free—surfactant solutionfor 20 minutes; solution condition: [H 2 SO4] = 0.036 M, Temp.= 25°C. (a)reference: solvent spectrum (b) less the solution spectrum.   1535.48 ATR spectrum of washed sphalerite after contact with ferric free—surfactantsolution, reference spectrum: solvent^  1545.49 KBr pellet spectrum of leached sphalerite (in the presence of surfactant)residue; leaching performed at 140±1°C for 15 minutes under PN 2 =1100kPa; lignin sulfonate solution concentration: 5 g/L at flow rate =2 mL/min  155xiv5.50 KBr pellet spectrum of OPD. ^  1685.51 ATR spectrum of 5 g/L solution of OPD; solution condition: [H 2 SO4 ] =0.036 M, Temperature = 25°C; reference spectrum: solvent. ^ 1695.52 ATR spectrum of sphalerite conditioned in OPD solution (no ferric pre-treatment) for 20 minutes; solution condition: [H2 SO4 ] = 0.036 M, Temp ^= 25°C (a) reference spectrum: solvent (b) less solution spectrum^ 1705.53 ATR spectrum of sphalerite washed after conditioning in 5 g/L OPD so-lution for 20 minutes, reference spectrum: solvent^  1715.54 ATR spectrum of sphalerite conditioned in 5 g/L OPD solution (afterferric pre-treatment) for 20 minutes; solution condition: [H 2 SO4] = 0.036M, Temp.= 25°C (a) Reference spectrum: solvent (b) less the solutionspectrum.   1725.55 ATR spectrum of washed sphalerite after contact with ferric/OPD solu-tion, reference spectrum: solvent. ^  1735.56 Metal—OPD complex on interaction.  174xvList of Symbols^ai^activity of species ia- surface area of solid substrate in contact with liquid, cm 2 or m2A^integrated absorbance; cm'^CRXZ , c2^concentration of solute in bulk solution, mol/Le — concentration of absorbing species; molesc — speed of light; 3x10 1 ° cm/sec^de^equatorial diameter of drop, cm^ s^diameter of drop at a distance D e from apex, cm^d'en^magnified diameter of pendant drop, cm^Eh^electrode potential; Ve — electron^E°^standard electrode potential; VE^energy; ergs, joules or eVEtotal resultant molecular energy; ergs, joules or eV^Eelec^electronic energy contribution; ergs, joules or eVEvib vibrational energy contribution; ergs, joules or eVErotn. rotational energy contribution; ergs, joules or eVEtransl. translational energy contribution; ergs, joules, or eV^fi^activity coefficient of component iF — Faraday's constantg^acceleration due to gravity, 9.81 m/s 2 or 981 cm/s 2xvi^-AG^interfacial free energy change; joules or ergsh — Planck's constantH — dimensional parameter as a function of drop shape^Ia^anodic current density; A/m 2cathodic current density; A/m2^I^intensity of transmitted radiationintensity of incident radiation^K„^equilibrium constantk^absorption coefficient of sample; mol" liter cm'^1^path length of radiation through absorbing sample, cmm — ionic speciation parameter of surfactantm l , m 2 — atomic masses^MS^metal sulfiden — number of electronsn2 , n 1 — refractive indexes of sample material and reflector respectivelyn21 — ratio of sample to IRE refractive indexes^p^differential pressure, atmospheres or kPaR1 , R2^radii of curvature, cm^R^universal gas constant^RX^ionic surfactant^S°^elemental sulfur^ ^shape ratio of pendant dropT - transmittanceT — temperature; °Cxviiwo - original diameter of drop forming tip, cmmagnified diameter of drop forming tip, cmWa - reversible work of adhesion per unit area; ergs/cm2 or joules/cm 2^WD^dispersion force component of work of adhesionWa - polar force component of work of adhesion^Wa^hydrogen bonding force component of I/V:^wfD^dipole-dipole interactions component of W:Wa^electrostatic interactions component of W:^Wa^r bonding component of 14/:- molar concentration ratio of inorganic to organic saltsY, X - dimensionless variablesy - vertical height of liquid measured from apex of drop, cm or mz - valency of organic surface active ionxviii0 - mineral - liquid sulfur contact angle, degrees7MA - mineral - aqueous solution interfacial tension, mN/m or dynes/cm7sA - liquid sulfur - aqueous solution interfacial tension, mN/m or dynes/cm-yms - mineral - liquid sulfur interfacial tension, mN/m or dynes/cmr i^relative surface/interface excess of solute at the interface per unit area, mol/m 2II 2 - chemical potential of solute- activity coefficient parameter for component iCo.^surface/interfacial potentialBoltzman's constant- anodic overpotential; V• - cathodic overpotential; V• frequency of vibration; sec -1force constant or strengthreduced mass of atomswavenumber; cm -1A^wavelength of incoming (incident) radiation; cmE - extinction coefficient; cm -10, - critical angle of IRE; degrees- interfacial tension, mN/m or dynes/cmp - difference in density between liquids under test, kg/m 3 or g/cm3# - dimensionless parameterxixACKNOWLEDGEMENTI wish to express my sincere thanks to my research supervisors Drs. David Dreisingerand Ernest Peters for their patience and guidance during the course of my research.I will like to express my gratitude to NSER.0 and the Department of Metals andMaterials Engineering of this university for financial support during my professionalstudies. I also wish to extend my sincere thanks to any person who in any way contributedto making this work a succcess.Lastly, I owe much gratitude to my uncle and mother for their moral and materialsupport throughout my education.xxChapter 1INTRODUCTIONSulfide minerals are the main source of most non-ferrous base metals, and can be animportant source for some precious metals as well. The most important source of zinc isthe mineral sphalerite (ZnS). Chalcopyrite (CuFeS 2 ) is the predominant copper mineral,although copper may be extracted from other sulfide minerals such as bornite (Cu 5 FeS 4 ),chalcocite (Cu2S), and covellite (CuS). Until the last forty years or so, sulfide mineralswere treated mainly by roast-leach-electrowin (for Zn) and roast smelt-convert or flash-smelt-convert (for Cu) technologies with the resultant production of large volumes ofsulfur dioxide that must be collected to minimize environmental damage. In most cases,the collected sulfur dioxide is converted to sulfuric acid in acid plants located on site.The cost of recovering SO 2 and converting it into acid, the SO 2 pollution problemsassociated with incomplete SO 2 recovery, and the increasing demand for metals, alongwith the progressive depletion of high grade ores require the development of new, cleanand efficient metal extraction techniques for the treatment of sulfide minerals and low—grade ores. Much time and money have been spent to understand hydrometallurgicalsulfide decomposition reactions in the last forty years and this work has had a greatimpact on the research and development of hydrometallurgical processes for sulfide oresand concentrates. The direct acid pressure leaching of sulfide minerals, notably ZnS andcopper sulfides - Cu 2 S, CuFeS 2 and FeS 2 has certain potential advantages over the roast-leach and roast-smelt technologies. For example, in the Sherritt Gordon acid pressureleach for zinc the process achieves high metal extraction (about 96-98 %) and goodChapter 1. INTRODUCTION^ 9recovery of elemental sulfur (more than 85 %) without any sulfur dioxide production.Pyrite, when present as an impurity can vary in reactivity under leaching conditions.The oxidation of the impurity pyrite can influence the leaching kinetics as a source ofiron catalyst. Furthermore, pressure leaching of pyrite during zinc pressure leaching yieldsonly 10-50 % elemental sulfur (depending on leaching conditions) with the remainder ofthe mineral sulfur being oxidized to the sulfate form.The physical state of the elemental sulfur depends on the operating conditions. Attemperatures higher than 119°C, elemental sulfur exists as a liquid which tends to wetand occlude any unoxidized sulfide mineral particles, thus slowing down the reactionrate at higher extractions. At temperatures below 119°C, sulfur forms a porous solidproduct layer on the reacted mineral surface. Complete reaction can be achieved but thepresence of the sulfur layer impedes the reaction and therefore lengthens the retentiontime for acceptable metal extractions. This sulfur issue was one of the main obstacles'to commercializing the zinc pressure leach process during the early days of its invention.It continued to hamper the development of the process for at least 20 years until finally,it was found that ligninsulphonic acid or its salts are effective surface active agents(surfactant) for dispersing liquid sulfur in the process, thus preventing it from coatingthe unreacted sulfide particles, and thereby enhancing the metal dissolution rate.Surfactants are important reagents in the metal extraction industry. They are usedas wetting agents, dispersants, flotation reagents, etc. in mineral beneficiation processes.However, the role of surfactants in leaching processes has not been widely studied. Thegoal of increasing metal sulfide recovery and the desire to minimize cost and pollution,and increase extraction efficiency require a thorough understanding of mineral-surfactantinteractions during pressure leaching. It is therefore imperative that surfactants with'The problem of liquid sulfur prevented the early development of the commercial zinc pressure leachon a commercial scale until about 1977; the process as currently practiced takes place at temperaturesbetween 135 and 155°C.Chapter 1. INTRODUCTION^ 3properties that can affect mineral decomposition in aqueous solutions should be studied,and a data base needs to be accumulated on such surfactants that improve (or impede)extractions, affect leaching kinetics, or play a role in separations between minerals andtheir residual or precipitated leach products.Unlike flotation systems which have been well documented, there has been no exten-sive investigation of the nature of molecular interactions that take place between surfac-tants and sulfide minerals in leaching reactors. It is the objeCtive of this research to studythe effect of surfactants on mineral-solution-sulfur systems with a view to understandingtheir behaviour during zinc pressure leaching.Chapter 2LITERATURE SURVEYLeaching is basically the act of selectively dissolving metals from an ore or concentrate.Generally, sulfide minerals are semi—conductors and thus, the oxidative or reductive leach-ing of metal sulfides may be an electrochemical process which involves the transfer of elec-trons across phase boundaries. Since most sulfide minerals occur as impure materials, theimpurities may affect the electrochemical leaching kinetics of the process. The leachingof metal sulfides such as ZnS, FeS, CuFeS 2 , FeS 2 under particular conditions results inthe generation of elemental sulfur which, depending on the operating temperature, canbe either liquid or solid.The presence of liquid sulfur tends to slow or stop the metal dissolution process ifthere is complete wetting of reacting mineral particles by the sulfur. This slow downcan be prevented by an interfacially active reagent that will create an unfavourable ther-modynamic environment for wetting. The nature of the surfactant—mineral interactionsunder leaching conditions has not been extensively studied (unlike flotation systems),so this study seeks to focus on the basic mechanisms of the interactions in the leachingsystems. One of the direct methods of surface analysis of solids (eg. for adsorption ofsurface active species) is by employing infrared spectral analysis and this approach isadopted in this study.4chapter 2. LITERATURE SURVEY^ 5Table 2.1: Electronic properties of selected sulfide mineralsMineral Resistivity, Ohm-M Conductor typeCu 5 FeS 4 10-3 - 10 -6 pCu2S 4(10-2) - 8(10 - 5 ) pCuFeS 2 2(10') - 9(10 -3 ) nCuS 8(10-5) - 7(10 -7 ) metallicFeS 2 3(10-2) - 1(10 -3 ) n SL pZnS 1 - 1(104 ) n2.1 ELECTROCHEMICAL PROPERTIES OF METAL SULFIDESMetal sulfide minerals are good electronic conductors. Table 2.1 summarizes the elec-tronic properties of some sulfide minerals [1, 2]. Generally, they are semi-conductingl,and hence can be used as electrode materials in laboratory studies. The kinetics oftheir dissolution are often controlled by electrochemical processes. Oxidative leachingproceeds by the anodic dissolution of the sulfide mineral. The electronic conductivityof metal sulfide minerals can be attributed to the covalent character which results innon-localization 2 of charge. Semi-conducting minerals possess "rest" potentials'. Theequilibrium potential of a mineral electrode depends on the solution composition as wellas on the exact composition of the mineral. This observation holds true for the dissolu-tion of metals too as has been indicated by Jun Li et al. [3]. The electronic propertiesof most metal sulfides and certain oxides readily lead to corrosion and galvanic couples'There is no reason for semi-conductors to behave differently from conductors unless the conductivityis so low that cathodic and anodic reactions are polarized by electron or electron—hole transport throughthe mineral grain.'Motion of free—charge carriers which may be either electrons in the normally empty conductionband or holes in the normally full valence band; conductivity may be referred to as a n-type or p-typedepending on whether electrons or holes dominate the conduction.3 "Rest" potential is the potential at which there is no net anodic or cathodic current and can be amixed potential or an equilibrium potential.Chapter 2. LITERATURE SURVEY^ 6in aqueous systems. As a result of their electronic character, the kinetics of the leach-ing of sulfides can be related to the potential of the solid in contact with an aqueouselectrolyte. Thermodynamically, dissolution reactions are strongly favoured so that theeffects of reverse reaction kinetics are usually ignored 4 . However, the build-up of productions in the solution may markedly affect the kinetics by influencing the potential at themineral-solution interface [1]. The reversible electrode potential, Eh, of a sulfide mineralin equilibrium with elemental sulfur can be written in a Nernst equation form, viz:Mn+(aq) S° ne MSn RT , amn+.a s oEh E-^ InnF^amsa i = activity of species i.2.2 LEACHING OF SULFIDE MINERALSThe leaching of metal sulfides is often considered electrochemical in nature and theextensive work done by different investigators [1]-[19] has led to the advancement andunderstanding of the theory. Under oxidative leaching conditions, a sulfide mineral ispresumed to give up electrons to an oxidant e. g. Fe3+, Cu 2 +, dissolved oxygen, etc. Thereactions occurring during sulfide mineral decomposition in an aqueous environment canbe classified as follows [9]:• molecular mechanism in which oxygen and sulfur transfer and where H 2 S is con-sidered as an intermediate product in the leaching process. H 2 S is considered to beimmediately oxidized to elemental sulfur by ferric ions and dissolved oxygen duringthe oxidation of sulfide minerals:MS + 2H+ <^  /1/2+ + H2 4,04 This is not a general statement. Depending on the system under consideration, the reverse reactionmay or may not be ignored.Chapter 2. LITERATURE SURVEYThe hydrogen sulfide is distributed between the aqueous and gaseous phases asH2 S(g) -,-=H2S ( a q )thenH2 S(aq) 2Fe a+q) -- 2H(+aq) + 2Fefa+q) + S°and/orH2 S ( aq ) + 1/20 2^H2O + S°• electrochemical mechanisms promoted by high overvoltages as a result of high redoxpotentials (in acidic ferric ion solution); the decomposition reaction would be:MS(0-4 Mn + + S° + neandnFe 3+ +^nFe2+withMS + Fe3+ -4 Mn+ + S° + nFe 2 +as the overall reaction• and/or a mechanism that generates sulfateMS(0+ 4112 0 —+ AIn+ + S0,21 - +8H+ + (n 6)eand^6)Fe3+^6)e^(n 6)Fe2+with^MS + 4H20 (n 6)Fe3+^Mn+ .504- + 8H + (n 6)Fe 2+^(2.1)Chapter 2. LITERATURE SURVEY• a combination of the aboveIf the electrochemical process prevails during sulfide leaching, the leaching rate isdetermined entirely by the physical and structural nature of the mineral surface and bythe electrochemical potential of the surface. The mineral surface may be clean (i.e. afreshly fractured surface) or it may be coated with elemental sulfur and/or metal oxideor hydroxide solids. These coatings may have varying degrees of porosity and electricalconductivity. The electrode (mineral) potential would be determined in most cases by asteady-state process involving the anodic and cathodic reactions.The cathodic reaction of an oxidant operating on a sulfide mineral surface is governedby the mass transfer rate of the species 5 and/or the activation polarization of the electrodereaction (which can vary with the identity of the cathode surface). The most importantoxidants in hydrometallurgy are 0 2 , Fe3+ ions, Cl 2 (or C10 - ) and HNO 3 acid and thesehave very different reduction kinetics. Oxygen reacts cathodically as:02 + 4H+ + 4e —÷ 2112 0but this process is very irreversible and slow unless overvoltages are large. Ferric ions arevery reactive on a conducting surface and so react almost reversibly. The consequenceis that in ferric ion leaching the limiting rate is either mass transfer of ferric ion to themineral surface (concentration polarization) or slow anodic dissolution of the mineral.Aqueous chlorine is also a powerful oxidant which exhibits high Eh values, negligibleactivation polarization, and high leaching rates. The disadvantage of using chlorine forthe leaching of metal sulfides is that chlorine will oxidize sulfur to sulfate s .Thus, the potential of the mineral may be governed by the electrochemical kineticsof both anodic and cathodic processes. Mixed potentials' are determined by balancing5 This can lead to concentration polarization.6 Not if Cl 2 oxidizes another species that in turn exerts a lower Eh.'At the mixed potential the sum of the anodic current is counterbalanced by the sum of cathodiccurrent i. e. E/a^—E./a ; different minerals have different mixed potentials unless in galvanic contactChapter 2. LITERATURE SURVEY^ 9the current between separated anodic and cathodic processes. Potentials more positivethan the mineral rest potential drive the half-cell reaction in the net anodic direction andpotentials more negative than the rest potential drive the reaction in the net cathodicdirection', Figure 2.1.The anodic overvoltage is 1/a E - Eel and the cathodic overvoltage, Tie E - Ee Leaching of Zinc Sulfide MineralsOxidative dissolution is an important part of the hydrometallurgical extraction of metalsfrom sulfide ores/concentrates and a lot of energy and time has been spent in order todevelop a process for the leaching of ZnS in and SO 42- media using an oxidizing agentsuch as ferric ions or oxygen.Ferric sulfate leaching of sphaleriteFerric sulfate - sulfuric acid (0.3 NI each) leaching of sphalerite was carried out by Perezand Dutrizac [34]. Their data display an increase in the metal dissolution rate with tem-perature (50-90°C). The kinetic data show that the rate determining factor in sphaleriteleaching is the surface chemical reaction (i.e. the data fit the shrinking core model). Therate of leaching increases with increasing iron content of the sulfide mineral.It has been suggested by Scott and Nicol [35] that during ferric sulfate leaching ofsphalerite, the first step in the series of reactions is the non-oxidative reaction whichproduces H 2 S as an intermediate product which is then oxidized to elemental sulfur byeither dissolved oxygen or ferric ions in solution. The molecular mechanism may berepresented by [35, 9]:ZnS + 2H+<=>- Zn 2+ + 112 S( a , )^(2.2)with each other.8 Figure 2.1 depicts a hypothetical polarization curve for the dissolution of a metal sulfide in thepresence of a cationic oxidant Nn+, eg. Fe.3+.^Chapter 2. LITERATURE SURVEY^ 10H2S(a q) 2Fe3+ ^ 9H+ + 9 Fe2+ S °^(2.3)The H 2 S has to be dissolved before it can be oxidized. From reaction 2.2, the extentof metal dissolution is sensitive to the acidity of the process but Peters and Doyle [9]have reported that in spite of the fact that this mechanism is easy to understand, it isinconsistent with practical oxidative acid leaching. In practice, there is no dependenceof the rate on acid concentration. Secondly, according to reaction 2.2, an increase in thezinc ion concentration of the leach solution should lower the activity of H 2 S at the mineralsurface. This will put a constraint on reaction 2.3 particularly if it is the rate-determiningstep. Thus Peters and Doyle [9] have indicated that the oxidative acid leaching of sulfidesis best described in terms of an electrochemical mechanism and this is consistent withobservations made by others, [1, 6, 7, 11, 19]. The presence of Fe 3+ and Fe 2+ in the leachsolution contributes to the kinetics through the ferric/ferrous redox couple [9, 19] andthe electrochemical dissolution of the metal is thus written as:ZnS^Zn2+ S° 2e^ (2.4)2Fe3+ 2e--> 9Fe 2+ (2.5)The overall reaction is:^ZnS 2Fe3+^Zn2+ 9Fe2+ S°^(2.6)These reactions are equivalent to the molecular version, i.e. reactions 2.2, and 2.3 and theoverall reaction are the same in both mechanisms. The presence of the ferric ions in theleach solution enhances the metal dissolution rate by oxidizing H 2 5 to elemental sulfurand ferrous ions'. Verbaan and Crundwell [19] observed a linear relationship between theinitial metal dissolution rate and the redox couple's potential and an activation energyof 79 kJ/mol was obtained for their studies in the temperature range 25-85°C.9 This oxidation reaction is very fast even at room temperature.Chapter 2. LITERATURE SURVEY^ 11In a report on the oxidative dissolution of sphalerite concentrate in ferric sulfate me-dia, Crundwell [36] indicated that in the early part of the process, the rate-limiting stepwas the surface chemical reaction. The extent of leaching increased with increasing ferricion concentration and decreased with increasing ferrous concentration and an activationenergy of 46 kJ/mol was obtained for this rate control regime. As the sulfur product layergrows in thickness, diffusion through the product layer becomes the rate-controlling step.The initial metal dissolution (reaction) rate was insensitive to the presence of 0.05 and0.01 M zinc sulfate in the leach solution. This is consistent with an earlier observationmade by Dutrizac and MacDonald [37] and the comment made by Peters and Doyle [9]viz, that the electrochemical mechanism for the ferric ion leaching of sphalerite (involv-ing an electron pair as the intermediate product) is more favoured over the molecularpathway which involves H 2 S as the intermediate product'.The refractory nature of two sphalerite concentrates leached in 0.5 M ferric sulfate-0.1M sulfuric acid media resulted in not more than 45 % zinc extraction [38] after 5 hoursof leaching at 78°C. This was assigned to the formation of insoluble lead sulfate and/orplumbojarosite on the particles' surface.In an independent study into the influence of flocculants on the extent of metal extrac-tion, Crundwell [39] reported that the addition of organic surfactants such as flocculantslowers the rate of sphalerite dissolution. The decline was attributed to a decrease in thesurface area available for charge transfer processes (otherwise the chemical reaction) asa result of the surfactant adsorption.' °The reason is that with an increase in zinc ion concentration in solution, equation 2.2 should shiftthe equilibrium state in such a way as to re-establish equilibrium according to:Keg = [ Zn21.[H2Sei9][H + ] 2this will mean decreasing the rate of either metal dissolution or H 2 S generation either of which will resultin a decrease in the extent of metal extraction.^Chapter 2. LITERATURE SURVEY^ 12Ferric chloride leaching of sphaleriteDespite the success and the apparent advantages of sphalerite leaching in either sulfateor chloride media, most of the work has been in sulfate media, probably because of theestablished commercial technology for electrowining zinc from purified sulfate solutions.Numerous similarities exist in the reactions occurring in the ferric chloride and ferricsulfate leaching of sphalerite [40] - identical free energies, enthalpies and entropies ofleaching. The reaction is electrochemical in nature and controlled by charge transferprocesses at the sphalerite surface (as has been reported by others [1, 7, 9, 11, 19]) withthe potential of the sphalerite being dependent on the ferric/ferrous couple at the surface.The overall reaction for ferric chloride leaching of zinc sulfide isZnS 9Fe 3+ Zn2+ + 2Fe 2+ + S°In ferric chloride leaching of complex sulfide ores containing pyrite, sphalerite, mi-nor amounts of chalcopyrite, galena, etc., the abundant pyrite is relatively inert to thelixiviant resulting in the preferential leaching of the sphalerite component. The mostimportant reactions taking place during ferric ion (chloride) leaching of complex sulfideconcentrates are the oxidation of CuFeS 2 , PbS, ZnS, FeS 2 , etc. generating elementalsulfur and/or sulfate and ferrous ions. The kinetic data obtained by Ngoc et al. [41]on the leaching of a complex sulfide ore suggested that the simultaneous dissolution ofchalcopyrite, sphalerite and galena in ferric chloride occurred as if each individual min-eral was leached separately. The mechanism of Cu and Zn dissolution was reported tobe controlled by surface chemical reaction.Majima et al. [15] have studied the non-oxidative leaching of sphalerite in aqueoushydrochloric acid solutions at 75°C and reported activation energy values of about 83.4kJ/mol, indicative of surface chemical reaction control kinetics. De Bruyn and his col-leagues [42, 43] reached similar conclusions in their investigations into the non-oxidativeChapter 2. LITERATURE SURVEY^ 1:3sulfuric acid dissolution of zinc sulfide minerals. A similar study by Crundwell and Ver-baan [44] in sulfuric acid media at 25-65°C had a similar conclusion. On the additionof ferric ions to the leach solution, Crundwell and Verbaan observed an enhanced metaldissolution rate as the ferric ions oxidized 11 2 5 to elemental sulfur and ferrous ions.The addition of NaC1 to the chloride leaching media has been observed to enhance theleaching rate of sphalerite [15, 41] probably due to the formation of chlorocomplex specieson the sulfide mineral surface as a result of the increased chloride ion concentration oractivity. The non-oxidative acidic leaching of sphalerite is represented by reaction 2.2.Evidence presented by Dutrizac and MacDonald [45] who carried out test work onthe leaching of sphalerite in 0.1 M FeC13 - 0.3 M HC1 solution at 85°C for a period of0-500 hours indicate that most of the sulfide sulfur reported as porous elemental sulfurin the leach residue. Both elemental sulfur and sulfate were produced. The elementalsulfur was the predominant product - (85-95 % conversion of sulfide sulfur to elementalsulfur). A similar study by Jin et al.[46] resulted in less than 5 % sulfide sulfur oxidationto the sulfate. The rate-determining step for the process was found to be the chemicalreaction at the solid surface; this observation is consistent with that made by Ngoc[41] and others [15, 34, 42, 43, 46]. Dutrizac and MacDonald indicated that: (a) theleaching rate is enhanced by the presence of CuCl 2 in the system, the rate increasingwith increasing Cu concentration; (b) the build-up of FeC1 2 tends to slow the reactionrate - Su [47] had made a similar observation"; (c) higher HC1 concentrations increasedthe rate of reaction because of increased direct acid attack of the sphalerite 12 ; at HC1concentrations in excess of 1 M, H 2 S was generated, part of which was, however, oxidizedto elemental sulfur by the ferric ion.11 This implies that in commercial practice, the continuous build—up of ferrous chloride in the sys-tem will have an adverse effect on the process efficiency unless a way is sought to continuously elimi-nate/reduce the ferrous chloride build—up.12 This observation should not be confused with comments by Peters and Doyle [9] who were referringto industrial practice.Chapter 2. LITERATURE SURVEY^ 14Perez and Dutrizac [34] observed a shrinking core leaching behaviour (i.e. surfacechemical reaction being the rate-determining factor) and linear leaching kinetics in theferric chloride leaching of sphalerite. This observation is in agreement with that of others.The rate of sphalerite dissolution was found to be related to the iron content of the sulfidesolid solution. They indicated that the magnitude of the increase in the metal extractionrate was approximately the same as in the analogous ferric sulfate system.The ferric chloride leaching of zinc sulfide (size range -100+150 mesh size) has beeninvestigated by Bobeck and Su [47]. These workers observed a change in the kineticbehaviour of the system with time - in the early stages of the process the rate-determiningregime was surface chemical reaction whilst in the latter stages when substantial surfaceelemental sulfur had been formed, there was a shift in the rate-controlling mechanism toproduct layer (elemental sulfur) diffusion control.Venkataswamy and Khangaonkar [48] studied the kinetic behaviour of sphaleriteleaching in ferric chloride media in the presence of tetrachloroethylene as a solvent forelemental sulfur. In the presence of about 100 mL solvent, approximately 95 % Zn ex-traction was achieved in 3 M FeCl 3 at 104-106°C after six hours of leaching as comparedto 68 % achieved under similar conditions but in the absence of the organic solvent. Theuse of the solvent also permitted almost complete recovery of sulfide sulfur. The solventdissolves the sulfur from the mineral surface. A mixed rate-control regime (i.e. com-bination of diffusion and chemical reaction) is proposed, i.e. their data fit both surfacechemical reaction and diffusion control models over the entire test time. This observationis in contrast to that made by others.In contrast to the surface chemical reaction control regime proposed by most investi-gators, Rath et al. [49] proposed a diffusion control kinetic regime (of either Fe 3+, Fe2 +or Zn 2+ through the product layer). However, the activation energy value of 90 kJ/molobtained by them is inconsistent with diffusion control kinetics.Chapter 2. LITERATURE SURVEY^ 15In their studies on the kinetics of ferric chloride leaching of iron-activated bulk sul-fide (Cu-Zn) concentrate (size range -104 + 53 microns) at 95°C, Neou-Singouna andFourlaris [50] made some observations which are in contrast to those reported by others.These authors reported an increase in the Cu dissolution rate with increasing ferric ionconcentration between 0.5 and 2.5 NI in a 0.6 M HC1 solution. With regards to sphaleritedissolution, there are some contradictions. They mentioned an increase in percent Zndissolution in 0.5 to 1 M ferric chloride, beyond which the rate decreases. This is insharp contrast to the observations made by Dutrizac [45] and others [41, 46, 48] whoobserved increasing zinc dissolution for ferric ion concentrations' as high as 4 M. Themetal extractions versus HC1 concentration plots presented by the authors also indicatean increasing rate with increasing HC1 up to about 1 M concentration beyond whichthe extent of extraction begins to decrease. This effect is more pronounced in the zincdissolution curve. Again this observation does not agree with the work of Dutrizac etal. [45] and others [15] who indicated that a HC1 concentration of unity or more in-creases the sphalerite dissolution rate due to increased direct acid attack of the mineral.On the rate-determining mechanism for the sphalerite dissolution, Neou-Singouna et al.again disagreed with others with the exception of Rath et al. [49] - that the rate isdependent on product layer diffusion. An activation energy of 32 kJ/mol computed forthe sphalerite dissolution over a period of one hour is too high to be consistent with adiffusion-controlled mechanism"."Moreover, Majima et al. [15] and Ngoc et al. [41] have pointed out that increasing the Cl - ionconcentration or activity increases the rate of metal dissolution probably through the formation ofchlorocomplex species on the mineral surface. Since an increase in ferric ion concentration of the system isaccompanied by a simultaneous increase in the chloride ion concentration, a decrease in metal dissolutionrate is not expected.14 0ne of the effects of temperature may be on the sulfur morphology. If at higher temperatures thesulfur layer has larger pores, the diffusion rate will increase through this layer more than the effectof temperature alone, yielding an apparently high activation energy when porosity is not adjusted fortemperature. 32 kJ is too high for diffusion, but only a bit too high (20 kJ is usually an upper limit) .When the diffusion is through a porous layer, the effect of temperature on porosity needs to be added tothe effect of temperature on the diffusion constant in aqueous media. Due to these effects, the reactionChapter 2. LITERATURE SURVEY^ 16One consistent observation made by all authors is the formation of a surface laver ofelemental sulfur on leaching zinc sulfide particles. At temperatures below the meltingpoint of sulfur, the sulfur product layer is reported to be porous. At these temperatures,Jin et al. [46] report that the elemental sulfur layer is porous enough that it provides noresistance to the diffusion of reacting ionic species across the layer.Oxygen pressure leaching of sphaleriteThe pressure leaching of zinc has been reported as far back as 1954 by Bjorling [51].Further work by later investigators [52] led to more understanding of the chemistryinvolved in the zinc pressure leaching. Forward and Veltman [52] found that at a reactiontemperature of 140°C and above, the leaching process ceased almost completely in aboutten minutes giving 65-72 % zinc dissolution. Further leaching of concentrate for upto six hours at 140°C resulted in only a minor increase in dissolved zinc. Jan et al.[53] also observed an increase in the zinc extraction rate up to 120°C beyond whichthe rate decreased and the leach residue appeared as several lumps. These were theresult of molten sulfur occluding unoxidized zinc sulfide particles. However, with theintroduction of lignin sulfonic acid or its salt' into the system (after 20 years of sporadicresearch) this problem no longer exists in the zinc pressure leach. There have beenmany investigations and publications [54]-[65] on the successful use of this surfactant asa dispersant for liquid sulfur in the oxidative pressure leaching of zinc sulfide minerals.The agent dispersed the liquid sulfur that previously inhibited high zinc extraction andso made it possible to commercialize the process'. The process as currently practicedtakes place at temperatures between 135 and 155°C at a total pressure of about 14mechanism is obscured by the activation energy values computed."This reagent is marketed with the trade name lignosol.16 Quebracho can also be added as a supplement to boost the effectiveness of the lignin sulfonate.Chapter 2. LITERATURE SURVEY^ 17atmospheres, with a chemical stoichiometry close toZnS + H2 SO4 + 1/20 2 —+ ZnSO4 + Si + H2 O^(2.7)as the overall reaction. The overall reaction as given by equation 2.7 is impractically slowin the absence of dissolved iron, which acts as a catalyst. Bernal [65] tried to use othermetals such as Ni 2+, Co2+, and Cu 2+ as catalysts only to find out that dissolved iron(ferric/ferrous) was the most effective catalyst of all the dissolved metals investigated.Jan et al. also observed that the presence of ferric ions in the leach solution increased theleaching rate dramatically and that the rate-determining step was the chemical reactionoccurring on the mineral surface.The sequence of reactions including iron catalysis can be represented as [64]:a) gas-liquid mass transfer of oxygen0 2(9 as ) —+ ° 2(aq)^ (2.8)b) homogeneous oxidation of ferrous ions2Fe 2+ + 1/20 2 (aq) + 2H+--+ 2Fe3+ +112 0^(2.9)c) ferric ion leaching of zinc sulfide2Fe3+ ZnS--* Zn 2+ + Si +2Fe 2+^(2.10)with cyclic regeneration of Fe3+ by reaction 2.9. Normally, there is sufficient acid-solubleiron in the concentrate (in the form of marmatite(ZnFe)S, pyrite(FeS 2 ), or pyrrhotite(FeS))to satisfy the iron requirements for the catalysis. If copper and lead are present in theconcentrate (often in the form of chalcopyrite and galena), they may also dissolve ac-cording to:CuF eS2 + 2H2SO4 + 0 2^CuSO4 + FeSO 4 + 2112 0 + 2S°^(2.11)Chapter 2. LITERATURE SURVEY^ 18PbS H2 SO4 +1120 2 —4 PbSO 4 (s) H20 S °^(2.12)Hydrolysis of iron may occur as the leaching reaction progresses (at low acidity), throughthe formation of hydronium jarosite and plumbojarosite:3Fe2(SO4)3 + 141120^21130Fe3(SO4)2 (011)6 5112SO4^(2.13)3Fe2(SO4)3 PbSO4 + 12112 0 —4 PbFe6 (SO4 ) 4 (OH)12 + 61/25'04^(2.14)It has been reported by Dreisinger et al. [54] that under zinc pressure leachingconditions, lignin sulfonic acid or its salt is degraded by ferric ions according to thereaction:Fe3+ lignosol --4 Fe 2 + + degraded lignosol^(2.15)Bernal [65] has also reported that in the presence of cupric ions as catalyst (instead ofdissolved iron), zinc extraction in their study almost ceased after 40 % dissolution (in thepresence of lignin sulfonate). However, they did not assign any reason to this behaviour.It is possible that the decreased metal dissolution rate was due to enhanced degradationof lignin sulfonate by the cupric ions in solution. It is unclear at what point in thedegradation process the sulfur again occludes the mineral surface.Scott and Dyson [66] investigated the catalytic role of different metal ions on thepressure leaching of zinc sulfide in 0.1 N sulfuric acid, at 250 psi oxygen partial pressureand 113°C (in the absence of a surfactant). After one hour of leaching in the presence of0.6 mg Cu (as the catalyst) per gram of ZnS, 92 % zinc extraction was achieved whilst 14mg Fe per gram of ZnS yielded only 51 % extraction under the same leach conditions ascompared to only 6 % (this value seems very low) in the absence of any catalyst. Jan et al.[53] proposed a chemical reaction which involves H 2 S as an intermediate product whichis further oxidized indirectly by oxygen through the help of the metal catalyst such asferric ions to elemental sulfur with H 2 S oxidation by the catalyst as the rate—controllingChapter 2. LITERATURE SURVEY^ 19step. The proposed mechanism is:ZnS + H2 SO4^ZnSO 4 + H2 S<^> (2.16)Fe2 (SO4 ) 3 +112 S--> Fe2 +^H2 SO 4^S° (2.17)(2.18)FeSO 4^20 2^H2 SO4^Fe 2 (SO 4 ) 3^H2 O+1/^+ —4^+But, as has been pointed out by Peters and Doyle [9], the electrochemical mechanism ismore favoured as the above molecular mechanism is inconsistent with practical experi-ence.Previous work by Owusu [55, 56] has shown that lignosol acts on both the solution-mineral and solution-liquid sulfur interfaces, to lower the work of adhesion between liquidsulfur and zinc sulfide minerals.De Nys and Terwinghe [67] have also reported that the oxidative pressure leachingof zinc sulfide concentrate was successfully done at a temperature of about 150°C in thepresence of orthophenylene diamine (OPD) as surfactant. The authors report that thereagent was used to enhance the transfer of reactants to the mineral surface. However,they did not elaborate on the mechanisms involved.Effect of solid solution iron content on sphalerite leaching behaviourInvestigations carried out by different workers [15, 18, 34, 68, 69], on different spha-lerite concentrates and ores indicate that the leaching rate of sphalerite increases withan increase in the solid-solution content of iron. The metallic impurity is believed toincrease the ionic character or the electrical conductivity of the mineral. This results inan increase in charge transfer across the mineral-solution interface (i.e. electron transferfrom mineral to solution) and hence an increasing rate of dissolution. Iron increases themineral conductivity through the creation of holes in the crystal lattice [18].Chapter 2. LITERATURE SURVEY^ 20Majima et al. [15] reported the leaching of four different zinc sulfide concentratesin 2 M HC1 media at 55°C in the absence of ferric ions and observed a linear leachingrate increase with increasing iron content of the solid; there was approximately a sixfold increase in the rate with the iron content varying from 0.28-2.15 %. There was nonoticeable effect of the iron content on the system's activation energy.A similar linear relationship between the metal (Zn) extraction rate and the ironcontent of the mineral in 0.2 M FeC1 3 - 0.3 M HC1 - 2 M NaC1 at 45, 65, and 85°C hasbeen reported by Crundwell [68] in a study of five sphalerite minerals of different ironcontents ranging from 0.55 to 8.6 %. There was approximately a twenty-fold increasein the rate over this iron content range. Activation energies obtained indicated a slightdecline with increasing iron content of the mineral.Piao and Tozawa [69] studied the leaching of sphalerite samples of different ironcontents in oxygen-sulfuric acid media at a temperature of 150°C. They reported aseven-fold linear increase in the leaching rate as the solid solution iron content increasedfrom 1.87-8.9 %.A linear correlation between the sphalerite leaching rate and the solid solution ironcontent had been reported by Kametani et al. [70] from the work done on the leaching ofsix sphalerite concentrates of different iron contents. The study was conducted in 10 g/Lferric sulfate-100 g/L sulfuric acid media at 90°C. There was a seventeen-fold increasein the dissolution rate in moving from 0.79 to 13.4 % iron content. The rate of leachingdepended on the ferric ion concentration for [Fe3+] < 10 g/L. No dependence on the ferricion concentration was found at [Fe 3+] > 10 g/L.Perez and Dutrizac [34] studied the leaching of fifteen sphalerite mineral samples attemperatures of 20-90°C in ferric chloride and sulfate media. The increasing leaching rate- increasing Fe content showed a linear correlation. A decrease in the apparent activationenergy for leaching with increasing iron content of the sulfide was noted. Bobeck and SuChapter 2. LITERATURE SURVEY^ 21[47] also reported a decreasing activation energy with increasing iron content of sphaleritemineral. Perez and Dutrizac observed similarities in the kinetics in both media, therate-determining factor being the transfer of charge at the sphalerite surface. A similarmechanism has been mentioned by Crundwell [68] who postulated that "the sphaleritedissolution rate is directly dependent on the number of occupied sites in the d orbitalconduction band i.e. to the iron content in the absence of other impurities".Xia Guang-xiang et al. [62] have also reported that under the conditions of theirstudy on oxidative zinc pressure leaching, the rate of oxygen consumption is linearlyrelated to the iron content (percentage Fe) of the concentrate as:R„y. = 30 * [Fe] mol/m 2 minuteThe increased rate of oxygen consumption (with increasing iron content) is a directreflection of the increase in the rate of zinc dissolution.In contrast to the linear correlations reported by most authors (as mentioned above),Kammel et al. [71] reported that in the absence of Cu 2+, there is a logarithmic relation-ship between the extent of zinc extraction (from sphalerite) and the iron content in therange tested (0.11 to 10.3 %). The test was done in 5 g/L ferric sulfate - 50 g/L sulfuricacid at 70°C. The logarithmic correlation reported is%Zn extraction = A + B * log[%Fe]. where A and B are constants.Chapter 2. LITERATURE SURVEY^ 222.2.2 The Behaviour of Sulfur in Oxidative Leaching of Sulfide OresThe behaviour of sulfur in the oxidative leaching of sulphidic minerals is not clearlyunderstood. The overall oxidation' reaction is usually given by:^MS^M2+ + 8 0 + 2e^ (2.19)The formation of sulfate occurs to a lesser extent:^MS + 4H2 0^+ 50 24 - + 8H+ + 8e^(2.20)Peters [7] has reported that under acidic conditions e.g. at 1 M sulfuric, perchloric orhydrochloric acid, under oxygen pressure, minerals such as PbS, ZnS, CuS and CuFeS 2can yield nearly 100 % elemental sulfur. Under the same conditions, FeS 2 yields 30-50 %elemental sulfur and MoS 2 leaching results in the complete conversion of sulfide sulfur tothe sulfate. The leaching of pyrite and molybdenite by ferric ions at temperatures below120°C is extremely slow.In their study on the oxidative pressure leaching of pyrite, McKay and Halpern [12]reported that in the absence of initial sulfuric acid in the reactor, all sulfide sulfur wasoxidized to sulfate (with half appearing as sulfuric acid and the other half as ferric andferrous sulfates) with no elemental sulfur being present. However, when 0.15 M sulfuricacid was initially added to the system, 50 % elemental sulfur was realized from theoxidation of the pyrite, the rest occurring as ferric and ferrous sulfates. Thus, the yieldof elemental sulfur in pyrite leaching is influenced by the acidity of the leach liquor andincreases as the pH drops 18 to below unity.In a series of investigations carried out by Lotens and Wesker [20] on ZnS, PbS, andFeS 2 in chloride media, at pH=2 (i.e. roughly 0.01 M H+ concentration) and at 30-60°C17The mechanisms involved in the sulfide sulfur oxidation process during the oxidative leaching processare detailed in ref. [20].18 1n virtually all reports of the oxidative acidic pressure leaching of ZnS and copper sulfides, theacidity is very high i. e. pH < 1.0Chapter 2. LITERATURE SURVEY^ 23temperature, using chlorine as the primary oxidant, they reported that pyrite was theonly exception where virtually all sulfide sulfur was converted to sulfate. The sulfuryield from ZnS leaching under similar conditions was 50 %. Lotens and Wesker [20]have reported that in addition to elemental sulfur and sulfate, other intermediate sulfurcompounds are present in the leach liquors, eg. thiosulphates and sulphites during theleaching of sulfide minerals. This has been confirmed by other investigators [22]-[25].In a study on the ferric ion leaching of chalcopyrite (in 0.5 M sulfuric acid) Dutrizac[26] reported that at temperatures below its melting point, elemental sulfur is hardlyoxidized, if at all. Similar sulfur behaviour was observed when the leaching was repeatedin acidic ferric chloride (0-2 M) media [27] where at temperatures less than 100°C,elemental sulfur conversion in excess of 95 % (5 % sulfate) was consistently realized. Thisis consistent with an earlier study done by Paynter [28] on the ferric chloride leaching ofchalcopyrite. Corriou and Kikindai [31] arrived at similar conclusions (for temperaturesbelow the melting point) when they tried to oxidize elemental sulfur in ferric ion-sulfuricacid medium over a wide temperature range up to 230°C.At temperatures greater than 159°C, polymeric S8 rings break down to polymericchains which are believed to exhibit higher oxidation rates than the S8 rings (in a givenaqueous environment). This statement is consistent with the activation energy numbersreported by Corriou and Kikindai [31]: between 125 and 160°C, the activation energyreported is 119.2 kJ/mol and between 175 and 230°C the number reported is 64.8 kJ/mol.The authors also reported that an increase in the sulfuric acid concentration of theaqueous phase slows down the rate of sulfur oxidation. These observations imply thatonce the elemental sulfur is formed in the leaching process' (below 160°C) it hardlyoxidizes. On the basis of this argument, the oxidation of sulfide sulfur to sulfates inthe oxidative leaching of pyrite and MoS 2 (at below 159°C) does not necessarily involve'This is true for pyrite, sphalerite, or any other sulfide mineral.Chapter 2. LITERATURE SURVEY^ 24elemental sulfur as an intermediate product.In ammoniacal systems (i.e. at high pH's), mineral or elemental sulfur is oxidized tooxy-sulfur products such as S20 32- , SO42- , Sn062- or NH2 S0; [9]. However, as has beenobserved by Reilly [32], an elemental sulfur by-product in oxidative ammonia leaching canbe recovered as such in the presence of C2 C/6 , a non-aqueous solvent for sulfur. Valensi[33] describes the most probable dissolution path for sulfur in this aqueous medium as:4S° + 6011 - (aq) —4 S2 03- (aq) 2S2-4S° + S2- (aq)equilibrium being attained at a pH of about 9. These species are oxidized slowly bymolecular oxygen in the absence of copper ions or at low temperatures [9].2.2.3 Effect of Sulfur in Blocking Leaching ReactionsAt temperatures less than 119°C, the morphology of the sulfur product formed duringacid leaching of sulfide minerals can have a marked influence on the extent of reaction.Peters [7] reports that if the formation of sulfur is accompanied by a positive molarvolume change (eg. as with pyrite), the sulfur layer is protective at extremely smallthickness whereas those with negative molar volume changes (e. g. sphalerite, galena,chalcopyrite, etc.) develop cracks and pores and the sulfur is not protective but onlyacts to slow the transport processes through the sulfur layer. At temperatures greaterthan the melting point of sulfur, liquid sulfur spreads over the unreacted mineral surfaceand inhibits further reaction. Since sulfur is a non-conductor, both anodic and cathodicprocesses in sulfide leaching should take place underneath any sulfur layer that is formed,and its presence is bound to slow down both anodic and cathodic processes by limitingmass transfer and possibly by decreasing the available solution-mineral interface area.Chapter 2. LITERATURE SURVEY^ 252.2.4 Interfacial PhenomenaA recent study by Owusu et al. [54, 55, 56] shows that when lignin sulfonic acid or its saltis introduced into a leaching reactor as a surfactant (or dispersant) during zinc pressureleaching, it tends to adsorb at the liquid sulfur - aqueous solution and on the solid mineral- aqueous solution interfaces. By so doing, the liquid sulfur - aqueous solution interfacialtension decreases from 54-55 mN/m in the absence of the surfactant to 28-30 mN/m inthe presence of 0.3 g/L lignin sulfonate concentration, beyond which no further increasein the surfactant dosage effects any change. The liquid sulfur - zinc sulfide mineralcontact angle also increases from about 80±5° in the absence of the surfactant to 145±5°in the presence of the surfactant. These phenomena resulted in the reduction of the workof adhesion2° from about 64 mJ/m 2 in the absence of any surfactant to about 5 mJ/m 2in the presence of 0.3 g/L lignin sulfonic acid. When 0.3 g/L quebracho was added as asupplement to the lignosol, the work of adhesion as computed was about 2 mJ/m2 . Thequebracho by itself effected only minor changes in the interfacial tension-it dropped from54-55 mN/m to about 45 mN/m in the presence of 0.3 g/L quebracho; but the changein the contact angle was quite significant under similar aqueous conditions. There wasan increase in the contact angle from about 80° in the absence of a surfactant to about126° in the presence of surfactant. In the presence of equal amounts (0.3 g/L each) ofquebracho and lignosol, the contact angle increased to about 156° whereas the liquidsulfur-aqueous solution interfacial tension obtained was no different from that obtainedin the presence of 0.3 g/L lignosol. The presence of ionic species such as ferric, ferrousand hydrogen ions in solution did not influence either the contact angle nor the computedwork of adhesion in the presence of the surfactants.20 The work of adhesion is defined as the reversible work required to separate a unit area of liquid fromthe solid substrate.Chapter 2. LITERATURE SURVEY^ 26The effect of tergito1 21 was also reported. The presence of 0.5-1.0 g/L tergitol low-ered the interfacial tensions to about 42-43 mN/m and increased the contact angle toabout 105°. However, the presence of 0.5 g/L each of tergitol and lignosol decreased theinterfacial tensions to about 27-28 mN/m but could not effect any further change in thecontact angle measured, the measured value being about 107°. This observation impliesa higher affinity of lignosol for the liquid sulfur-aqueous phase interface and a higheraffinity of tergitol for the aqueous phase-solid mineral interface when these two surfac-tants compete for adsorption sites in a system. In the presence of these two surfactants,the reported work of adhesion was about 19 mJ/m2 .2.3 PROPERTIES OF LIQUID SULFURThe melting point of sulfur could be either 113 or 119°C depending on the crystallineallotrope. Below 94°C, sulfur exhibits a rhombic structure and above 94°C, a monoclinicstructure is stable. However, the transition from rhombic to monoclinic is slow and if therhombic form is heated rapidly, it can easily reach its melting point of 113°C and spreadover any surface before recrystallizing to the monoclinic form (melting point 119°C).Thus, it is easy for the sulfur produced in the pressure leaching process to be molten,wetting the mineral and impeding the leaching reaction.At the melting point, sulfur forms a light yellow, low viscosity liquid. On heating, theliquid darkens, and becomes deep orange. Nickless [77] reports that viscosity increaseswith temperature and passes through a maximum at 159°C, beyond which there is dis-continuity in all physical properties 22 . This observation is in contradiction to that madeby Tuller and others [78]-[81] who have reported that the viscosity drops with increasing21 Tergitol is a non-ionic surfactant which has been reported to slow down the Zn extraction rate inthe zinc pressure leach [57].22 The viscosity behaviour reported by Nickless is undoubtedly a mistake.Chapter 2. LITERATURE SURVEY^ 27temperature up to about 159°C above which any further increase in temperature makesthe liquid more viscous up to about 200°C beyond which it begins to fall again withincreasing temperature up to its boiling point, (Figure 2.2). The increase in viscosityat temperatures above 159°C is the result of the rupturing of the puckered sulfur rings(S8 ) to form chains of varying lengths. Above 200°C, the long chains begin to break upresulting in decrease in viscosity with temperature.The measurement of surface tension on highly purified sulfur (i.e. sulfur-air interface)was carried out by Fanelli [82]. He concluded that the surface tension decreases linearlywith temperature, but that there is a sharp discontinuity at about 159°C, the rate of fallbeing smaller above this temperature. Based on Fanelli's results, Tuller [78] derived arelationship between the surface tension of sulfur and temperature as: below 159°C^= 73.4 — 0.105t^ (2.21)above 159°C^7 = 65.7 — 0.0566t^ (2.22)where t is the temperature in degrees Celsius, and y is the surface tension in mN/m.The density of liquid sulfur is reported [78] as decreasing with increasing temperature.2.4 SURFACTANTS AND INTERFACIAL PROPERTIES2.4.1 Characteristic Features of SurfactantsA surface-active agent (or a surfactant) is a substance that, when present at low con-centration in a system, has the property of adsorbing onto the surfaces or interfaces ofthe system resulting in a lowering of one or more of the surface or interfacial free ener-gies. Surfactants always act to reduce interfacial free energy rather than to increase itChapter 2. LITERATURE SURVEY^ 28although there are occasions when they are apparently used to increase it [831. 23 Sur-factants become significant (or applicable) when the interfacial phenomena occurring atphase boundaries within a system are so unusual relative to the expected bulk phase in-teractions that the entire system behaviour becomes dependent on the interfacial process(eg. corrosion, detergency or flotation).Surfactants have a characteristic molecular structure consisting of a structural groupthat has very little attraction for the solvent, known as a lyophobic group, together with agroup that has strong attraction for the solvent, called a lyophilic group. This is known asan amphipatic structure. In solution, the presence of the lyophobic group in the interior ofthe solvent causes a distortion of the solvent—liquid structure, increasing the free energyof the system. In an aqueous solution of a surfactant, this distortion of the water bythe lyophobic (hydrophobic) group of the surfactant and the resulting increase in thefree energy of the system when it is dissolved means that less work is needed to bringa surfactant molecule than a water molecule to the surface. The surfactant thereforeconcentrates at the surface. This decreases the surface/interfacial tension. On the otherhand, the presence of the lyophilic (hydrophilic) group prevents the surfactant from beingexpelled completely from the aqueous phase. Thus, the amphipatic structure causes thereduction of the surface or interfacial tension of the solvent as well as orientation of themolecules at the surface or interface with its hydrophilic group in the aqueous phase andits hydrophobic group oriented away from it.The chemical structure of groupings suitable as the lyophobic and lyophilic portionsof the surfactant molecule vary with the nature of the solvent and the conditions of use.As the temperature and use conditions (eg. presence of electrolyte or organic additives)23 Because surface tension is a free energy, a 'surfactant' that raises the surface/interfacial tensionmust necessarily interact with, and reduce the adsorption of a previously present (though not necessarilydetectable) surfactant. A surfactant that can displace another (by substitution) will always lower th,surface tension. Only when a surfactant is destroyed (precipitated or oxidized) by a second reagent canthe surface tension rise.Chapter 2. LITERATURE SURVEY^ 29vary, modifications in the structure of the surface-active agent become necessary. Thus,for surface activity in a particular system the surfactant molecule must have a chemicalstructure that is amphipatic in that solvent under conditions of use.The hydrophobic group is usually a long chain hydrocarbon residue, less often ahalogenated. The hydrophilic group is an ionic or highly polar group. Depending on thenature of the hydrophilic group, surfactants are grouped as:1. anionic - the surface-active portion of the molecule has a negative charge e. g.ligninsulphonates, xanthates, and carboxylates.2. cationic - the surface-active portion bears a positive charge e. g. RN(CH3 )3C/-(quaternary ammonium chloride).3. zwitterionic - both positive and negative charges may be present in the activeportion e. g. sulfobetaines.4. nonionic - the active portion bears no apparent ionic charge e. g. tergitol ( a deriva-tive of ethylene oxide ) but is polar enough to bind to aqueous solution.The choice of a particular surfactant depends on availability, cost and efficiency. Thisrequires knowledge of : a) characteristic features of available surfactants (physical andchemical properties and uses), b) interfacial phenomena involved in the job to be doneand the role of the surfactant and c) the surface chemical properties of various structuraltypes of surfactants and the relationship of the structure of a surfactant to its behaviourin various interfacial phenomena [83, 84, 85].In comparing the performance of different surfactants, it is necessary to distinguishbetween the amount of surfactant required to produce a given amount of change in thephenomenon under investigation and the maximum change in the phenomenon that theChapter 2. LITERATURE SURVEY^ :30surfactant can produce, regardless of the amount used. The former parameter is knownas the efficiency of the surfactant and the latter its effectiveness.2.4.2 Surfactants Identified for StudyIn this study the chemicals were chosen based on their chemical and physical propertiescoupled with the conditions prevailing during the pressure leaching of zinc sulfide min-erals and concentrates. In the zinc pressure leach process, any surfactant introduced inthe leaching stage should be neutralized by the time the slurry reaches the solution pu-rification, sulfur recovery, and zinc electrowining stages. Thus, the degradation propertyof the reagents selected is an advantage provided these reagents are effective during theleaching. Their ability to degrade under oxidizing conditions implies that there will beno residual surfactant to interfere with the subsequent extraction steps. This is one ofthe factors considered in the selection of the reagents. The chemicals identified include:1. Lignin sulfonate/lignin sulfonic acid2. Orthophenylene diamine (OPD)3. Water-soluble product of naphthalene sulfonic acid-formaldehyde condensates4. Cocoamido hydroxyl sulfobetaine (CAHSB)5. Tallowamido hydroxyl sulfobetaine (TAHSB)6. Cocoamido betaine (CAB)The only commercially important surfactant among these is lignin sulfonic acid, whichis currently being used in the pressure leaching of zinc sulfide minerals.Chapter 2. LITERATURE SURVEY^ 312.4.3 Properties of Individual SurfactantsLignin Sulfonates/Lignin Sulfonic AcidsOrganic ChemistryLignin sulfonates are derivatives of lignin whose exact definition and structure still re-mains a matter of debate. They are complicated condensation polymers, fig. 2.3, whoseexact structure is not completely known.The most common lignin sulfonic acids have a molecular weight of about 4000. How-ever, molecular weights can vary from 200 to 100,000. Sulfonate groups are attachedto the carbon chains but not to the benzene ring. The sulphonatable carbon atoms arein alpha-positions to the benzene ring. Secondary and tertiary OH- groups are presentin the side chains and these augment the hydrophilic properties of the lignin sulfonatesprovided by the sulfonic group.Lignin sulfonic acids are only soluble in water or in a solvent where hydrogen bondingis very strong. They undergo hydrolysis or oxidation. Lignin sulfonates are anionicpolyelectrolytes [88].Adsorption CharacteristicsThe surface activity of lignin sulfonates is affected by the number of hydrogen bondinggroups present in the molecule, such as side chain hydroxyl, phenolic and carboxyl groups.The adsorption of lignin sulfonates is known to be determined by the nature of thehydrogen bonding. Lignin sulfonates do not form micelles. Lignin sulfonates are notcapable of forming any oriented layer when adsorbed upon a liquid or solid surface [88].Chapter 2. LITERATURE SURVEY^ 39Uses and Applications of Lignin SulfonatesThe major quantity of purified lignin sulfonates is used in oil-well drilling mud formula-tions. The lignin sulfonates are effective in controlling the fluidity of the drilling mud.keeping mineral salt contaminants encountered in drilling suspended and preventing themfrom flocculating the mud while, at the same time, performing a sequestering function[88].Lignin sulfonates play very important roles in the mining industry, functioning as ore-flotation agents for a number of minerals. They are used as depressants. They are alsovery good dispersing agents for slimes and are used for this purpose in both the flotationand tabling of slimy ores. In tabling operations, lignin sulfonates disperse the slime andprevent occlusion of ore particles. In flotation operations the dispersing action on slimeincreases the degree of separation, which results in an increased recovery and bettergrade of concentrate. Lignin sulfonates are used in the zinc industry as dispersants fordispersing liquid sulfur from the surface of the unreacted mineral particles, thus enhancingthe extent of zinc extraction at temperatures above the melting point of elemental sulfur.The surfactant degrades in the presence of oxidizing agents. It has been reported [54]that at zinc pressure leaching temperatures (and aqueous conditions), lignin sulfonicacid has a half life of ten minutes. A study of the flocculating properties of ligninsulfonates shows that low-molecular weight lignin sulfonates act as dispersants at allconcentrations. High-molecular weight lignin sulfonates, (as any other polyelectrolyte),in contrast, act as flocculating agents at low concentrations and as dispersants at highconcentrations. The flocculating effect increases with the molecular weight of the ligninsulfonate [88]. Reducing the pH enhances the dispersing action of the low molecularweight lignin sulfonates. Other uses of lignin sulfonates as effective dispersants are in theblending of carbon black into rubber, the distribution of vat dyes in the textile industry.Chapter 2. LITERATURE SURVEY^ 33the preparation of wettable agricultural powders and sprays, and the manufacture ofportland cement.Lignin sulfonates are good sequestrants for heavy metal ions such as ferric, cupric, andstannous, and (less effectively) for cobalt, manganese, nickel, silver, stannic, uranium,and zinc.Orthophenylene DiamineOrthophenylene diamine is sometimes called orthamine or 1,2 benzenediamine. It has thechemical formula C6 H4 (NH2 ) 2 and the structure shown in Figure 2.4 and belongs to thearomatic amine family. It is colourless, forms monoclinic crystals, and darkens in air. Ithas a molecular weight of 108.1. The melting point is 102-104°C and the boiling point is252-258°C. It is soluble in alcohol, ether, chloroform and water. The solubility in water isreported to be about 3g/100g of water [91] but Dupont Chemicals [92] reports a solubilityof 17 % at 60°C. It is stable at normal temperatures and conditions of storage. Moist airand excess heat cause product quality to degrade. It is incompatible with oxidizing agentsand decomposes at high temperatures. It can be used as an oxidation inhibitor. Morrisonand Boyd [93] report that any functional group attached to the aniline structure' in theortho position imparts more acidic characteristics to the molecule irrespective of whetherit is electron donating e.g NH 2 or withdrawing, e.g NO 2 ; however, the effect is moreintense if it is an electron withdrawing group.Cocoamido Betaine (CAB)CAB is a member of the surface-active betaines and is derived from coco-fatty acids.Betaines are pH-sensitive zwitterionics (or amphoterics) which show the properties ofanionics at high pHs and those of cationics at low pHs. CAB has a chemical formuld24 The aniline structure consists of an aromatic ring with one amino group.Chapter 2. LITERATURE SURVEY^ 3 1R— CONH — (CH2 ) 3 — N+(CH3 ) 2 C H2 — COO - . It is a clear yellow liquid with a faintodour. The specific gravity is 1.05 and the boiling point > 100°C. It is highly solublein water. Slightly better wetting and foaming properties are reported at acidic pH's. Itis stable under normal conditions but incompatible with strong oxidizing and reducingagents [83, 94, 95] and degrades if overheated. Betaines are used in the textile industryas assistants in vat dyeing and printing. They are also used as detergents, foaming andwetting agents.Cocoamido- & Tallowamido—Hydroxyl Sulfobetaines (CAHSB & TAHSB)CAHSB and TAHSB belong to the amphoteric group of surface active agents calledsulfobetaines, with the general chemical formula' RN+(CH3 ) 2 (CH)(OH)(CH2) x ,503where R is either a coco-fatty or tallow-fatty derived alkyl group. They are also referredto as pH-insensitive zwitterionics, i.e. at no pH do they possess only cationic or anioniccharacters. They are therefore "ionic neutral", stable and substantially non-reactive inacid, neutral, and alkaline solutions and have strong surface activity. They adsorb ontocharged surfaces at all pH's without forming a hydrophobic film [96]. Sulphobetainesare very soluble in water and strong electrolyte solutions and are not sensitive to hardwater. They are compatible with all other types of surfactants, but are insoluble inmost organic solvents. The effect of pH and added electrolytes on the surface activeproperties of sulfobetaines is generally minimal. Sulphobetaines are partially degradedin the presence of oxygen producing carbon dioxide. This leads to a loss of surface activity[97].Sulphobetaines form good wetting, foaming, flotation, dispersing, dyeing agents or asemulsifiers, detergents and as a component of laundry and cleaning agents. They exhibit25 Structurally, a sulfobetaine molecule has two different hydrophilic parts - one positive and the othernegative situated at different parts within the molecule.Chapter 2. LITERATURE SURVEY^ 35good complexing properties with respect to divalent ions. Sulphobetaines are graduallyreplacing petroleum suifonate' in the petroleum industry for enhanced oil recovery 27 .Sulphobetaines, present in only very small concentrations, lower the interfacial tensionbetween mineral oil and water, especially salt-containing water. They are compatiblewith anionics, cationics or nonionic surfactants as well as with organic thickening agents.Hydroxylated sulfobetaines have improved solubility [96]-[102].Naphthalene Sulfonic Acid - Formaldehyde CondensatesNaphthalene sulfonic acid-formaldehyde condensates have the chemical structure shownin Figure 2.5 and the formula [C10 11-6 ,503NaCH2],i . The average molecular weight is ap-proximately 2000. These products are used as wetting agents for powders (agriculturalwettables, powdered pesticides) and paint formulation. Their uses are similar to those ofligninsulphonates [96]. They are completely soluble in water. They are suited for applica-tions in highly concentrated aqueous solutions and dispersions and retain their dispersingability over a wide pH and temperature range. They can be used in combination withanionic or nonionic surfactants when lower surface or interfacial tensions are required.They degrade in the presence of oxidizing agents and undergo thermal degradation underextremely high temperatures greater than 230°C. These reagents when used will stabilizeand reduce the viscosity of thick pastes and produce smooth, uniformly wet paste with-out separation of the solid particles. They are used in the leather industry as dyeing andlevelling assistants. These surfactants can be used to produce very high strength concreteby reducing the water-to-cement ratio and reducing the concrete's water requirement byabout 20-30 %.26 This replacement is due to the high sensitivity of the petroleum sulfonates to divalent cations pro-ducing precipitates during the progress of the surfactant solution in the reservoir.27 Lignin sulfonate is also used in the petroleum industry for the same purpose.Chapter 2. LITERATURE SURVEY^ 362.5 INFRARED SPECTROSCOPYOne of the most direct approaches to the study of adsorbed molecules on solid surfaces isby infrared spectral analysis of the surface. The spectrum provides information about theadsorbed molecules and the nature of surface bonding or interactions. As far as leachingsystems are concerned, no infrared investigations have been reported that characterizethe possible mechanisms of interactions between any commercially important surfactant'and the mineral (or adsorbent) on a molecular level. However, there have been manyinfrared studies [103]-[120] to investigate reagent-mineral molecular interactions in flota-tion systems. Several in-situ mineral-aqueous solution studies have been reported in thelast 10-15 years in the literature [107]-[120] and they have contributed to explainingthe mechanisms and the nature of interactions on a molecular scale between flotationreagents and their mineral substrates. Some ex-situ work has been performed by others[103]-[106]. Sulfide minerals are among the few which have been thoroughly investigatedunder flotation conditions. With the exception of sphalerite (or ZnS), all other sulfideminerals reported in the literature were readily rendered floatable on reagent adsorption;on the other hand, sphalerite must be activated first with heavy metal ions eg. Cu 2 +,and Fe3+, before the mineral adsorbs the flotation reagent and is rendered hydrophobicand floatable. The metal ions, eg. Cu 2+, transform the surface of the ZnS mineral intoa chalcocite-like surface before it can adsorb the reagent [105, 113, 118, 119, 120].Termes and Richardson [113] have reported that natural sphalerite absorbs at 4800-2500 cm' and below 800 cm' in the infrared region. This observation makes it easier tocharacterise the mineral-surfactant interface since most of the characteristic absorptionsdue to the surfactant occurs in the transparent region. Any infrared absorption band28 So far the only commercially important surfactants in the zinc pressure leaching systems are lignin-sulphonic acid or its salt and quebracho.Chapter 2. LITERATURE SURVEY^ 37in the 2500-800 cm -1 region would be either due to impurities in the mineral' or tomolecules adsorbed on the surface of the mineral.A significant portion of the present research was concerned with the study of theadsorption mechanism of leaching reagents (surfactants, notably ligninsulphonic acid,and OPD) on the surface of sphalerite. This objective was achieved by employing in-situattenuated total reflection' (ATR) - Fourier transform infrared (FTIR) spectroscopytechnique.The conventional transmission technique' was used in this work primarily to recordthe reference spectra of sphalerite and the surfactants studied.2.5.1 Theory of Infrared SpectroscopyAll spectra are the result of absorption or emission of electromagnetic radiation (byatoms or molecules) that occurs between definitely quantized energy levels. The transferof energy between the molecule and the electromagnetic field is given by [121]:DE = 1wwhere h is known as Planck's constant and v is the frequency of incident radiation. Apositive AE implies absorption of radiation by a molecule resulting in an absorptionspectrum. A negative zIE implies that a molecule emits radiation and hence representsan emission spectrum. This results in a spectrum being unique to the molecule underinvestigation, and the spectrum is represented on a light intensity - frequency plot, withpeaks occurring wherever the above conditions are satisfied.Both atoms and molecules give rise to spectra. The atomic spectrum consists of aseries of sharp lines, whereas a molecular spectrum consists of a series of hands. The29 Naturally occurring minerals will always contain some amount of impurities.30 This is sometimes referred to as internal reflection spectroscopy (IRS)."This technique is the KBr pellet/wafer method.Chapter 2. LITERATURE SURVEY^ 38difference between atomic and molecular spectra depends on the different energy levelsinvolved in the transitions. Atomic spectra are the results of transitions between thedifferent allowed energy levels for the orbital electrons, whereas for molecular spectra,the atoms within the molecule vibrate and the molecule as a whole rotates, and thus theresultant energy contributions are given by [121]:Etotal = Ee l ec Evib. Erotn. Etransl.The translational energy, Etransi . is usually extremely small, and hence can be ignored.The isolation of the electronic, vibrational and rotational energies is possible becausethe separation between electronic energy bands is much greater than separation betweenvibrational levels which in turn is larger than the separation between rotational levels,see Figure 2.6.Thus, a series of vibrational levels is associated with each electronic level and a seriesof rotational levels is associated with each vibrational level. Electronic level transitionsresult in electronic spectra which are observable in the visible or UV regions of thespectrum; vibrational level transitions in the same electronic state give absorption in theinfrared region of the spectrum. These vibrational transitions are usually accompaniedby rotational transitions. At low frequencies, in the far infrared (below 500 cm -1 ), purerotational spectra are observed.2.5.2 Molecular Force ConstantsThe vibration of atoms within a molecule can be equated to a simple harmonic oscillator.whose frequency of vibration is given by classical mechanics as [121, 122, 123]:1v = —27rtift mChapter 2. LITERATURE SURVEY^ 39Or1^1A^27rcwhereK = force constant or strength= reduced mass of atoms= wavenumberv = frequency of vibrationA = wavelengthc = speed of lightThe reduced mass is given by:1^1^1m 1^77/2where m 1 and m 2 are the masses of the two atoms involved in the particular bonding inthe molecule. The force constant of the harmonic oscillator is considered to be analogousto the strength of the chemical bond, and hence n is a measure of the strength of achemical bond. It is easy to conclude that larger force constants are indications of strongerchemical bonds. Changes in the positions of the vibrational frequency upon adsorptionof a molecule on a surface can be directly attributed to changes in the strength of thebond that gives rise to the particular vibration. Shifts to lower wavenumbers indicateweakening of the bond (a lower force constant) and shifts to higher wavenumbers indicatestrengthening of the bond under consideration. Frequency is also dependent upon themass of atoms that constitute the molecule. Chemical bonds involving atoms of low masswill have higher vibrational frequencies than those involving atoms of high mass for equalforce constants.Chapter 2. LITERATURE SURVEY^ 40Spectral Unitsv = frequency, sec-1A = wavelength (cm)Av = c, the velocity of light = 3*10 10 cm/sec= 10 -4 cm =--  io4 Awavenumber, cm-1d(cm-i) =  1  _ 104A(cm)^A(p)Hence,hcE = hv =^= hcdE = ergs, joules or eVh = Planck's constant = 6.63*(10') J s.2.5.3 Absorption of Infrared RadiationElectrodynamically, a molecule can absorb radiation if: i) it possesses a vibrational orrotational frequency the same as that of the incoming electromagnetic radiation; and ii)a change in the magnitude and/or direction of the dipole moment occurs. [121]-[124].When the electrical field of the molecule rotates or vibrates at the same frequencyas the incident electromagnetic radiation, transfer of energy takes place'. The dipolemoment acts as a coupler, coupling the energy from the incident radiation to the molecule.Absorption band intensities are related to the concentration of the absorbing species.The intensity of light (I) transmitted through an absorbing medium at a wavenumberis given by Beer's law:I = lac —katwhere32 This transfer of energy is usually absorption of radiation by the molecule.Chapter 2. LITERATURE SURVEY^ 41Io = intensity of incident radiationconcentration of absorbing species in the sample1 = path length of radiation in the samplek = absorption coefficient, which is characteristic of the absorbing speciesThe definitions of transmittance and absorbance are given as:- transmittance, T 'IL or percent transmittance, % T^1001//ati- absorbance, A = In^= clwhere c = kc , an "extinction coefficient" which has units of cm' [121].Adsorption ProcessesThere are two distinct types of adsorption processes: a) physical adsorption (or physisorp-tion) - a process in which the bond between adsorbent and adsorbate is considered tobe due to Van der Waals or electrical forces and permanent dipoles; b) chemisorptionin which the adsorbate goes through chemical interaction with the adsorbent; it may bedissociative, non-dissociative or reactive in nature [121].A physically adsorbed species is usually considered to be an adsorbed material thatcan be completely removed from the surface without decomposition.The infrared spectrum of a molecule is the result of vibrations of the atoms within themolecule; the number and frequency of vibrations being determined by the symmetry andbond strength of the molecule as a whole. The infrared spectrum of physically adsorbedspecies is only slightly altered and small frequency shifts are observed. In chemisorption,a completely new infrared spectrum may be observed and band shifts and intensitiesare far removed from those of the independent bulk species [121]. Both physical andchemical adsorption processes can take place simultaneously on a surface, especially atlower temperatures.Chapter 2. LITERATURE SURVEY^ 42The Fingerprint Region of an IR Spectra.At frequencies greater than 1500 cni -1 , assignment of absorption bands to functionalgroups is easy. Below 1500 cm -1 , band assignment is possible (hut not easy) and thisregion is usually known as the "fingerprint region" of the spectrum [121]-[124 Thevibrations of the molecule as a whole give rise to a series of absorption bands at lowenergy, below 1500 cm', the positions of which are characteristic of the molecule understudy.Vibrational Changes on AdsorptionThe spectrum of any given molecule is unique and depends on the vibrations of the atomswithin the molecule and the rotation of the molecule as a whole. On adsorption of themolecule on a surface, the rotational movement is severely restricted and if the processof adsorption is chemisorption, the vibrational modes will be severely affected as well.If the process of adsorption is physical adsorption, much of its rotational movement willbe lost but the vibrational movements will be little affected by the process. Thus, theadsorption bands before and after adsorption should be similar.2.5.4 Infrared TechniquesKBr Pellet/Wafer Transmission TechniqueBefore the advent of the attenuated total reflection - FTIR technique'', the KBr windowtransmission method was the standard infrared practice. In the transmission measure-ment or spectral recording, the radiation of wavelength A passes through a sample (e.g.KBr pellet) and the absorption of the radiation by the sample causes its intensity to dropfrom I, to I. The KBr window per se is transparent to infrared radiation. This method33This technique is described in the next section.Chapter 2. LITERATURE SURVEY^ 1:3is applicable to solids and involves the use of KBr salt 34 to prepare sample pellets ordiscs'. The sample disc can then be scanned for infrared absorbance or transmittance.The technique is based on the assumptions that the KBr does not interact with the sam-ple. The KBr technique is a useful means of studying adsorbed reagents on minerals butremains an "ex-situ" technique, so that any details about the adsorbate in an in-situsituation are lost [116].Attenuated Total Reflection (ATR) TechniqueThe ATR technique otherwise referred to as the Internal Reflection Spectroscopy (IRS),presents an opportunity for an in-situ study of the mineral-aqueous solution interface 36 .It is a technique for recording the optical spectrum of a sample material that is in contactwith an optically denser but transparent medium [116, 127]-[130]. The ATR technique,has been successfully employed in in-situ studies of flotation systems. The principle ofthis method is that an infrared beam should be internally reflected through a trapezoidof an internal reflecting element (IRE) 37 . For total internal reflection to take place, theincident beam should have an incident angle greater than the critical angle, 0,, of thecrystal, so that (coupled with the high refractive index of the reflecting crystal), it canbe reflected by the flat faces of the IRE. The critical angle is given by:0, = sin -1 —n2n lwhere n 1 and n 2 are the refractive indexes of reflector and sample respectively.34 The salt should be of infrared grade.35 Adequate precautions should be taken to avoid contamination of the pellet.36 1t is important to note that all techniques (eg. KBr) which deal with the adsorbent being transferredfrom an aqueous (or liquid) medium to a gaseous environment are liable either to pick up species adsorbedat the liquid-gas interface, or to shed some of the adsorbate onto the liquid-gas interface. This mean ,that the interpretation of adsorption at the solid-liquid interface cannot be ascertained unless done b^an in—situ method and only the ATR method is capable of handling that situation [116].3 'This can be any crystal that is transparent to infrared rays and has a higher refractive index dial,the sample under investigation.Chapter 2. LITERATURE SURVEY^ 44Theoretically, during the course of the beam being reflected by the IRE, it appearsto penetrate slightly beyond the reflecting surface and into the rarer medium's, Figure2.7.A material (or sample) placed in contact with the IRE will selectively absorb radiationcorresponding to its vibrational frequencies owing to interactions with the beam. Theoutgoing beam will have an absorption spectrum characteristic of whatever is next tothe crystal face e.g. surface coating. As a result, only the layer situated very close tothe surface may be examined spectroscopically. The thickness of the layer may be afraction of wavelength or several wavelengths. Increasing the number of reflections (i.e.multiple reflections) is equivalent to increasing the path length in transmission studies39 .This increases the amount of the beam absorbed by the sample. The depth of beampenetration is a function of a) beam wavelength, b) refractive index of both sample andreflector (IRE), and c) angle of incidence.The advent of Fourier Transform Infrared (FTIR) spectrometers coupled with com-puters has made the application of the ATR technique in surface chemistry easier. FTIRspectroscopy can give an average infrared spectrum of the compounds present on the sur-face of minerals, even in very low concentrations. FTIR spectrometers have high energythroughput, high signal-to-noise ratio, high sensitivity, high data aquisition speeds andflexibility to manipulate the data.Internal reflection spectroscopy can take three different approaches: in one method themineral (or sample) is deposited by vacuum evaporation as a thin layer on the reflectingelement and then used as the adsorbent. The second approach uses finely ground sam-ple/mineral particles conditioned externally and introduced into the ATR cell. The thirdThe rarer medium mentioned here represents the working sample material or any other surfacecoating on the IRE.39 The ATR technique as practiced nowadays makes use of multiple internal reflections and this resultsin better sensitivity.Chapter 2. LITERATURE SURVEY^ 45approach involves direct adsorption of surfactants from solution onto reactive mineral—reflecting element substrate. The use of vacuum evaporation/deposition of mineral sam-ple has its own problems. The vapour deposited mineral will certainly be different froma natural mineral. The second approach permits spectral measurement of a natural min-eral surface in the presence of an aqueous solution. Direct adsorption onto a reactivemineral IRE provides more quantitative information regarding the nature, structure andorientation of adsorbed species than the other two methods.SUMMARYIt can be seen from the literature that much work has been done in the area of sphaleriteleaching in ferric sulfate and ferric chloride media and in the presence of oxygen. Allof these studies are related to the kinetics of the leaching process and the role of liquidsulfur in the process both at temperatures lower and greater than the melting pointof elemental sulfur (i.e. 119°C). The leaching mechanism is usually electrochemical innature. It is noted that the presence of ferric ions in the leach solution and the solidsolution iron content tend to speed up the metal dissolution rate. The leaching of ZnSin an acidic medium at temperatures less than 159°C converts almost all of the sulfidesulfur to the elemental state with minor oxidation to the sulfate form. Depending on thetemperature, the elemental sulfur can slow clown the reaction rate or have a negligibleeffect. At temperatures greater than 119°C, liquid sulfur slows the rate down by completewetting of the particles and the only way to avoid this wetting is by the introduction ofa surface active agent (e.g. lignin sulfonic acid) which plays a very important interfacialrole by creating unfavourable thermodynamic conditions for the wetting by liquid sulfurto continue.The mechanisms of the surfactant adsorption process on a molecular scale in theleaching system are not well understood. Unlike flotation systems which have been wellChapter 2. LITERATURE SURVEY^ 46studied spectroscopically, infrared studies of surfactant adsorption in leaching systemsare not reported in the literature.Thus, it is the objective of this study to further our knowledge of mineral-surfactantinteractions in zinc pressure leaching systems. High temperature surface tension and con-tact angle measurements in the liquid sulfur-mineral-aqueous system will be employedto determine the macro effects of a range of surfactants. Laboratory scale zinc pressureleach experiments will be performed to measure surfactant performance in relation tomeasured interfacial activity. On a molecular scale, ATR-FTIR methods will be em-ployed to further our understanding of mineral-surfactant interactions in an aqueousenvironment.POTENTIALE = Mixed PotentialE el= Equilibrium PotentialMS re++ S °+me0.20 -2.0100 159 200 300 400Temperature 0C3.02.00.351.00aa(1)-1.0Chapter 2. LITERATURE SURVEYEel= Equilibrium PotentialN (n-1)+= N n++ eFigure 2.1: Current density- potential diagram; after Wadsworth [1].47Figure 2.2: Viscosity of liquid sulfur; after Tuller [78].Chapter 2. LITERATURE SURVEYFigure 2.3: Possible structure of lignin sulfonate; after Pearl [86].48N H 2Figure 2.4: Structure of OPD; after reference [92].ElectronicTransitionV=1Vibrational^groundTransition ystateE'v'—nv'=211= 0Chapter 2. LITERATURE SURVEY^ 49M 4- -03Figure 2.5: Structure of naphthalene sulfonic acid-formaldehyde condensates; after Rosen[83].v"-1^v"=0 ^ E   excitedstateEJnJ" etc^J =4, etc^J"-2^J"=0 ^^J'=4, etc^J'-2^ J'=0^E o zero pointenergyRotational #Transition 4Figure 2.6: Atomic energy levels; after Hair [121].Chapter 2. LITERATURE SURVEY^ 50INCIDENT BEAM^ REFLECTED BEAM(CONTAINS INFORMATION)Figure 2.7: Internal reflection effect; after references [121].Chapter 3OBJECTIVES AND EXPERIMENTAL APPROACH3.1 OBJECTIVESIn—situ investigations of solid—aqueous solution interactions are becoming very importantin the development of novel technological and scientific applications. Until recently, notvery much was known about the nature and structure of adsorbed species formed onthe surfaces of solids fully immersed in liquids. Recent studies on in—situ solid-solutioninteractions have contributed tremendously to the understanding of the mechanisms andthe nature of interactions occurring between flotation reagents and their substrates. Itis evident from the literature that, unlike flotation systems, not much is known aboutmineral—surfactant solution interactions in leaching systems on a molecular level. A crit-ical review of the literature has shown that very little work has been done on the leachingbehaviour of sulfide minerals (particularly ZnS) with commercially important surfactantsor dispersants, in spite of the fact that many studies have been carried out in terms ofprocess kinetics and chemistry in the different leaching media. It is time that the fun-damental mechanisms governing surfactant—mineral interactions in leaching systems belooked into as has been done in the past decade or more for flotation systems. Practicalleaching processes are complex phenomena in which numerous variables come into play,and where possible, these variables need to be studied. The growing importance of pres-sure leach processes at temperatures greater than the melting point of elemental sulfurcreates opportunities for research and development of other surfactants as substitutes for51Chapter 3. OBJECTIVES AND EXPERIMENTAL APPROACH^52lignin sulfonic acid. This objective can be achieved only when the fundamental chemistryinvolving surfactant-mineral interactions is well understood.The present study builds on an earlier investigation of mineral-surfactant interactions[54, 56].It is the objective of this study to investigate the molecular interactions that occurbetween zinc sulfide (or sphalerite) minerals and surface active agents during oxidativepressure leaching and to understand the basic principles governing the adsorption andeffectiveness of surfactants or dispersants in the zinc pressure leach process.3.1.1 Scope of Study and ApproachThe scope of the study involved identifying some selected surfactants' and applying themin the following areas of study:• interfacial tension and contact angle measurements: the measurements made in-clude liquid sulfur-aqueous solution interfacial tensions and liquid sulfur-solid min-eral contact angles both in the absence and presence of surfactants employing thependant drop and sessile drop techniques respectively.• leaching studies: these involve the use of the selected surface active agents asdispersants for the oxidative pressure leaching of the ores or concentrates.• infrared studies: these involve the study of the nature of bonding or interactionsbetween the surface of the mineral and any added surfactant in the solution. Thein-situ ATR-FTIR technique was employed in this study.'These were selected based on their physical and chemical properties.Chapter 4EXPERIMENTAL TECHNIQUESThe experimental study was divided into three distinct phases:1. Interfacial phenomena investigations in the aqueous - mineral - liquid sulfur - sur-factant system.2. Pressure leaching experiments designed to evaluate surfactant performance.3. Infrared spectroscopy of mineral - surfactant systems.4.1 INTERFACIAL PHENOMENAThe purpose of this study was to investigate the interfacial behaviour exhibited by theliquid sulfur - aqueous zinc sulfate - zinc sulfide mineral system under simulated zincpressure leaching conditions in the absence and the presence of different surfactants. Thisrequires the measurement of the liquid sulfur - aqueous solution interfacial tensions andliquid sulfur - zinc sulfide - aqueous solution contact angles in the absence and presenceof the different surfactants mentioned. There are numerous techniques for determiningthese quantities; however, each of these has its own limitations. A thorough review of theliterature [132]-[151] shows that for the purpose of this study, the pendant drop techniquehas a number of outstanding advantages over other methods for measuring the interfacialtension. Advancing contact angles were measured using the sessile drop approach. Thependant drop technique involves taking a photograph of liquid sulfur suspended in anaqueous solution of zinc sulfate while in the sessile drop method, a photograph is taken53Chapter 4. EXPERIMENTAL TECHNIQUES^ 54of a drop of liquid sulfur droplet resting on a flat mineral surface fully immersed in azinc sulfate solution. The interfacial tension and the advancing contact angle can bedetermined by measuring the shape and position of the sulfur droplet. The aqueoussolution may contain other solutes which are typical of zinc pressure leach solutions.Such measurements require the use of a pressurized apparatus capable of raising theboiling point of the solution to temperatures above the melting point of elemental sulfur.These methods are static: once the surface of the drop is formed, equilibrated, andphotographed, the data for surface tension measurement are complete and not affectedby any outside influence prior to or during the measurement.The pendant drop technique is suitable for measuring the interfacial tensions betweentwo liquids. In brief, the pendant drop method consists of suspending a small drop of oneliquid (the denser liquid) from the end of a vertical cylindrical tube in a second liquid(the lighter liquid). The shape and size of the drop can be determined by measurementsmade on a photograph. The drop formed is influenced by two opposing forces, namelysurface/interfacial tension and gravitational forces. The effect of the presence of theforce of gravity is to elongate the drop so that its length exceeds its width. Without thepresence of a gravitational force, the drop assumes a shape of minimum surface area orenergy (i.e. a sphere). Figure 4.8 shows the profile of a pendant drop.Advancing contact angles were measured using the same experimental setup as thatfor the interfacial tension determinations. However, instead of a suspended drop as inthe interfacial tension determinations. the liquid sulfur drop was made to rest on the flatsurface of a zinc sulfide mineral specimen that was fully immersed in an aqueous solutionand a picture taken of the drop. Once in contact with the sulfide mineral, the sulfurdroplet begins to spread and the pictures were taken when this lateral movement ceased.Figure 4.9 shows the profile of the drop shape under different aqueous environmentalconditions.Chapter 4. EXPERIMENTAL TECHNIQUES^ 55ApparatusFigure 4.10 shows the details of the equipment used for the interfacial tension and contactangle measurements. The bomb was heated externally by wrapping it with a heatingtape followed by an insulating blanket. The inner chamber of the bomb was totallyenclosed to prevent loss of liquid by evaporation. The tip was made from tubing having acircular cross-section. The end was cut off perpendicular to the vertical axis. Microscopicperfection was not essential since the liquid tends to bridge over any minor irregularities.ChemicalsThe chemicals used were: reagent grade zinc sulfate, sulfuric acid, sulfur flowers, fer-ric sulfate, ferrous sulfate, sodium silicate, zinc sulfide mineral (marmatite) specimens(obtained from Cominco Ltd). The surfactants and the sources of supply were:• Calcium—based lignin sulfonic acid obtained through the courtesy of Cominco Met-als Ltd, Trail, British Columbia (supplied by Reed Inc., Quebec).• Orthophenylene diamine (often referred to as OPD in this work) obtained throughthe courtesy of Dupont Chemicals, Wilmington, Delaware, U.S.• Sodium salt of naphthalene sulfonic acid - formaldehyde condensates supplied byHandy Chemicals Ltd., Laprairie, Quebec• Cocoamido hydroxyl sulfobetaine (CAHSB)• Tallowamido hydroxyl sulfobetaine (TAHSB)• Cocoamido betaine (CAB)The last three surfactants were obtained through the courtesy of Alkaril Chemicals Ltd.,Mississauga, Ontario.Chapter 4. EXPERIMENTAL TECHNIQUES^ 56EXPERIMENTAL METHODSThe bomb and its accompanying accessories were carefully cleaned' and dried to getrid of any unwanted material or dirt. A thermocouple was constructed from type E(chromel-constantan) wire and calibrated. After the cleaning process, the equipmentwas assembled and calibrated by measuring the interfacial tension of water and benzeneat room temperature2 . After this test, the equipment was thoroughly cleaned again anddried to make sure there was no residual benzene in the system to affect the subsequentdeterminations. The focus of the study was on the interfacial role of different surfactantsin the liquid sulfur - aqueous zinc sulfate - zinc sulfide system under conditions similar tothose prevailing in the zinc pressure leaching process. This involves the measurement ofliquid sulfur - aqueous solution interfacial tensions and liquid sulfur - zinc sulfide mineral- aqueous solution contact angles.4.1.1 Interfacial TensionAll experiments were carried out at a pressure of 620 kPa and a. temperature of 130+5°Cusing 1.2 mol/L zinc sulfate solution'. The conditions of the aqueous solution environ-ment (due to the presence of added solutes) were varied in accordance with the experi-mental objectives. These included:• the effect of different surfactants and dosages• the presence of ferric and ferrous ions'The cleaning was performed using acetone and distilled water.2 The purpose of this test was to determine how the measurement deviates from the theoreticallylisted interfacial tension for water—benzene systems.'The choice of these conditions was based on results obtained from earlier work [55, 56] in whichit was observed that above a zinc sulfate solution concentration of 1.0 mol/L, there was no change inthe liquid sulfur-aqueous solution interfacial tensions (in the absence of any surfactant) and within apressure range of 276-620 kPa the effect of pressure on the interfacial tensions was negligible.Chapter 4. EXPERIMENTAL TECHNIQUES^ 57• the presence of acidIn each test, the syringe was partially filled with sulfur. The rectangular optical cellwas carefully mounted on the adjustable seat, K, in the bomb such that the sapphirewindows of the pressure vessel were parallel to the optical sides of the cell. The cellwas then carefully filled with the aqueous solution. The syringe was then mounted andthe thermocouple inserted into the thermowell to record the temperatures as the bombwas being heated. The heating rate was controlled using a powerstat. The chamberwas closed and sealed and then subjected to nitrogen pressure at 620 kPa. Since theoperating temperature was higher than the melting point of elemental sulfur 119°C),the sulfur melted in the syringe. At the operating temperature, the drive-screw of thesyringe was moved in a downward direction until the sulfur formed pendant drops at thetip immersed in the aqueous solution. The design and construction of the syringe wassuch that it allowed the position (i. e. height) of the tip to be adjusted. This enabledthe drop to be positioned appropriately to be photographed. Several drops were formedand photographed for each experimental condition after which one of the controllingvariables was changed and the sequence of operations was repeated. All experimentswere done in duplicate and some in triplicate to ensure that any data obtained werereproducible. The results of repeated experiments were within +2 %. The measurementswere made from the photographs of drop profiles and interfacial tensions were calculatedusing these profile measurements. The measurements were made at points on the dropwhich were remote from the end of the tip and where the influence of imperfections inthe shape of the tip was normally negligible. Adequate precautions were taken to ensureattainment of interfacial equilibrium. The required dimensions were measured with a rulegraduated to 0.5mm. These dimensions were estimated to the nearest 0.1mm. Repeatedmeasurements of the same drop were reproducible to within ±0.1mm. Some of theChapter 4. EXPERIMENTAL TECHNIQUES^ 58dimensions were remeasured using a computer scan to determine the accuracy of themanual measurements. The variation was within ±2 %.The magnification of the pendant drop was determined from a knowledge of theenlarged diameter of the capillary tip (actual diameter was known to within ±0.002mm).4.1.2 Contact Angle MeasurementsAfter the surface tensions experiments, a series of experiments was carried out to deter-mine the zinc sulfide mineral - liquid sulfur - aqueous solution advancing contact anglesin the absence and presence of the different surfactants. Determinations were carried outin the following environments:• zinc sulfate (1.2 M) solution• 1.2 M zinc sulfate, 0.05 M Fe 2+, and 0.12 M Fe 3+ in 0.2 M sulfuric acid solutionThe technique used was similar to that employed in the previous work. The differencewas that a picture was taken of liquid sulfur resting on a solid mineral surface (insteadof being suspended) in an aqueous environment.The marmatite mineral specimen was carefully cut into a rectangular shape. Thesurface was carefully polished using a 600 grit sand paper. The specimen was then rinsedseveral times with distilled water. The mineral was carefully mounted horizontally in theoptical cell. The cell was filled with the aqueous solution and the set up completed asdescribed above. The drops were carefully controlled in order to determine the advancingcontact. angles by gradually increasing the droplet size. Once in contact with the sulfidemineral, the sulfur droplet begins to spread and the pictures were taken when this lateralmovement ceased. As has been pointed out by Good and Oss [152], the system does notneed to attain thermodynamic equilibrium before taking the pictures for the necessarymeasurements. After each run, the mineral was carefully cleaned (as described above)Chapter 4. EXPERIMENTAL TECHNIQUES^ 59to remove any residue that might have been deposited on the surface during the courseof the experiment and the series of measurements repeated. The contact angles weremeasured directly from the photographs.Quantifying Interfacial ParametersInterfacial TensionsUnlike contact angles which can be measured directly from the photographs, the in-terfacial tensions have to be computed [133]—[150]. There are three techniques for themathematical treatment of the pendant drop profile: a) the method of the plane of inflec-tion b) the equation of state approach [153] c) the method of a selected plane. However,the method of the plane of inflection has two severe limitations which make its applica-tion. very tedious and inaccurate. These are i) the true plane of inflection of the dropmust be accurately located (which is graphically difficult) since a slight deviation fromthe true plane might introduce a large error and ii) the volume of the drop must becomputed from its profile. These difficulties limit the precision of the method [134, 137].The second method involves rigorous mathematical formulation but this is simplified bybeing incorporated into a computer software. It involves direct digitizing of the profile ofthe liquid droplet resting on a solid surface. This method measures the interfacial ten-sion and the contact angle simultaneously. The drop is considered to be at equilibriumwith its surroundings. It results in a precise solution. The third treatment is an indirectapproach and also leads to a more rapid and precise solution. In this method, the dropprofile, (Fig. 4.8), is considered to be at equilibrium with its surroundings, symmetricaland to possess two radii of curvature, H 1 and R2 (as with any curved surface). At theapex (or bottom of the drop). 0, the two are the same. The mathematical treatment isbased on the assumption that the pressure difference developed at any curved surface dueto the surface or interfacial tensions is proportional to the sum of the principal curvaturesChapter 4. EXPERIMENTAL TECHNIQUESof the surface, i. e.or60(4.23)( 1P = UT + R21 ) (4.24)where 7 is the interfacial tension in mN/m. This equation is known as the Laplaceequation.If y is the vertical height of the liquid measured from the apex of the drop profilewhere the two radii of curvatures are equal to each other i. e. R 1 = R2 = h, then thehydrostatic pressure due to the curvature of the hanging drop at any point is given by:27P =^- PgYCombining equations (4.24) and (4.25)(4.25)( 1^1^27U= i + il2 } = h PgYwhere p is the density difference between the liquids in question. From geometry' [133,160, 161](1 ^0)213 1 2Rl^ dx =^d2y dx2(4.27)(4.26)andR2 =x^r(1 + (m2)1/2)sin dydx(4.28)Therefore'The geometrical derivation is described in appendix AChapter 4. EXPERIMENTAL TECHNIQUES^ 61hence,d2 y dx2 dx r21 3 2^21112{1 + (2- ) j^X Li^(r)d2 y  2(dy) ^11^2dx 2^x (^ day^th2 pgyhdY)2]3/2dx(4.29)(4.30)This second-order, second-degree differential equation cannot be easily integrated; how-ever it can be simplified by transforming it into a non-dimensional form such that:Y" + — [1 + (Y') 2] = [2 — 9Y][1 + y2]3/2where/(3 = gPh 2andX = x/h, Y = y/h, ' dY/dX, and Y" = d 2 Y/dX 2From equation (4.32),(4.31)(4.32)gh2p= ,^L)(4.33)The dimensionless parameter ,3 and the radius of curvature h cannot be precisely deter-mined for a pendant drop from a picture of the drop.In order to solve for the interfacial tension, values must be assigned to /9 and h inequation 4.32): consider the drop profile in figure 4.8 - the size is dependent on itsequatorial diameter, de , and the shape is also dependent on the ratio of the diametersmeasured at two different horizontal planes [134]. The line CD is at a distance de fromChapter 4. EXPERIMENTAL TECHNIQUES^ 62the apex of the drop, where de is the equatorial (or maximum) diameter of the drops.From the measured values of de and ds , the shape ratio,S = ds /de^(4.34)is computed for each drop profile. The value of the shape ratio is dependent on the drop.Another parameter, H, can also be defined as a function of the drop shape such that:(4.35)(4.36)H is considered to be a function of S. Tables of values of H as a function of S areavailable [132, 133] from which y can be quantified from the photograph. Rewritingequation (4.36)=pg(de)2(4.37) wherep = density difference between the two liquids in question;g = acceleration due to gravity;de =equitorial diameter of the drop and is obtained from:de =^den (4.38)wo = original diameter of the tiptr, is the magnified diameter of the tip as measured from the photographdm is the magnified diameter of drop as measured from the picture.The density of the liquid sulfur with temperature was obtained from the literature[78].'The diameter de is first measured, then by constructing the base line AOB, a parallel line CD, (d,),is drawn at a distance d e from AOB.Chapter 4. EXPERIMENTAL TECHNIQUES^ 63The density of the aqueous solutions were determined at room temperature and cor-rected for high temperatures (i. e. at the operating temperature) on the assumption thatthe solutions have about the same coefficient of expansion as water (though this mightintroduce a slight error).Contact AngleContact angles were measured directly from the photographs with a fair degree of ac-curacy. The contact angle between a liquid and any other phase(s) in contact with itat equilibrium' is related to the interfacial tension or free energy per unit area of thosephases. In the case of sulfur in contact with a marmatite mineral surface and the aqueousphase these can be related as:7Ms 'Ysitcos 0 (4.39)andynrA — 7A1s = -ysAcos 0 (4.40)where -ymA is the mineral - aqueous solution interfacial tension, -yms is the mineral -liquid sulfur interfacial tension, 7SA is the liquid sulfur - aqueous solution interfacialtension, and 0 is the contact. angle measured in the liquid. The above equation is knownas Young's equation [134]-[150], [162].7/1 f A — ^ym s can be computed for the different systems. This will help in identifying theinterface(s) at which surfactant adsorption takes place in the system. For no adsorptionat the mineral surface, -)A/A — 1A1S has to be constant for all systems irrespective of theaqueous environment.'The contact angles measured in this study are not really thermodynamic equilibrium values. Theyare pseudo equilibrium or steady state values. However, for the purpose of computation, the contactangles will be assumed to be equilibrium valuesChapter 4. EXPERIMENTAL TECHNIQUES^ 644.1.3 Work of AdhesionIn the initial stages of a batch zinc pressure leach, liquid sulfur is non-existent. However,it immediately begins to form in the system due to the reaction. In a continuous process,there is a steady state of liquid sulfur present at all times. The sulfur tends to displacethe aqueous phase from the solid surface by adhesional wetting (i. e. liquid sulfur formson the surface, or makes contact and adheres to the surface). The surface free energychange associated with this wetting by liquid sulfur is:AG-tv^a(71V1,4 + 7SA — 7111.9)^(4.41)where a is the surface area of the solid substrate (i. e. mineral) in contact with an equalsurface of the liquid (sulfur) after adhesion. The driving force of this phenomenon is7mA + 7SAThe work of adhesion, Wa , is the reversible work required to separate a unit area ofliquid from the solid substrate:Wa = 7mA + 7SA^=^ (4.42)This is known as Dupre's equation [83, 133, 149, 151]. In this process, any reductionof the mineral - sulfur interfacial tension results in an increased tendency for adhesionto occur, but reduction of either the liquid-liquid or mineral-aqueous interfacial tensionsdecreases the adhesive tendencies.Substituting for IMA —^from equation 4.40 into equation 4.42,Wa^'"),9A COS 0 + 7SA^ (4.43)= "1' sA(1^cos 0)^ (4.44)Chapter 4. EXPERIMENTAL TECHNIQUES^ 65Wa can be computed from a knowledge of -ysA and cos 0. The driving force in adhe-sional wetting can never be negative and is equal to zero only when the contact angle is180°, which is never attainable in practice.4.2 PRESSURE LEACHING OF SPHALERITE, (ZnS)A series of leaching experiments was performed in the absence and presence of the dif-ferent surfactants as dispersants for liquid sulfur. The objective of this study was toinvestigate the performance of each of these surfactants in leaching systems in relation totheir interfacial activities and compare the findings to those obtained with lignin sulfonicacid.Chemical Reagents and Surfactants usedThe following chemicals were used in the preparation of the leaching solution: reagentgrade zinc oxide, sulfuric acid, ferrous sulfate and hydrogen peroxide. The concentrateused for these studies was obtained through the courtesy of Snow Lake Mines - a sub-sidiary of Hudson Bay Mining and Smelting Co., Manitoba, Canada. The surfactantsused were the same as mentioned earlier.Chemical analysis of the concentrate gave the following elemental compostion:• 49.2 % zinc• 11.4 Vc iron• 0.6 % copper• 0.6 (7/ lead and• 31.6 (7( sulfurChapter 4. EXPERIMENTAL TECHNIQUES^ 66X-ray diffraction analysis of the concentrate indicated that the bulk of the concentratewas present as ZnS, the rest being mainly pyrite and copper sulfide minerals.Two sets of leaching experiments were performed - low pulp density tests in which theamount of concentrate used was between 1-12 grams and high pulp density leaching testsin which the amount of concentrate used was 100+5 grams for each test. The leachingtests were performed in a 2-L Parr titanium autoclave. The operating temperature foreach test was 140+1°C.4.2.1 Low Pulp Density LeachingThe size range of the concentrate used for these tests was -270+325 mesh. The variableswhich were investigated were: type and concentration of surfactant in question, time ofleaching, and the mode of introduction of surfactant (either batch or continuous) into thesystem. For those surfactants available in only the liquid state, the tests were done withcontinuous addition of the surfactant. All batch additions were injected into the autoclavetogether with the concentrate sample. Prior to the introduction of the concentrate intothe autoclave, about 1 or 1.5 L of the pre-prepared leaching solution (or lixiviant) washeated to the operating temperature of about 140+1°C in the autoclave. The leachingsolution was made up of:• ti 2 mol/L zinc ions (or zinc sulfate)• ti 0.15 mol/L ferric ions• 0.15 mol/L ferrous ions• ti 0.5 mol/L sulfuric acidOnce the operating temperature was reached, the concentrate was injected into thesolution under oxygen or nitrogen pressure of about 1100 kPa using a specially designedChapter 4. EXPERIMENTAL TECHNIQUES^ 67injection unit attached to the autoclave. The experiments lasted for a predeterminedtime of up to 30 minutes after which the autoclave was quickly quenched in an ice waterbath until cooled to 75-80°C after which it was de-pressurised. The pulp was quicklyfiltered using a pressure filtration unit. Samples of the filtrate were collected for chemicalanalysis. After each run, care was taken to ensure that all of the leach residue wasrecovered for subsequent chemical analysis. The residue was dried in a vacuum drier forabout 3-4 days, chemically digested', and analysed for metal content. The extent of zincextraction was determined from the residue analysis. The change in zinc concentrationin solution was too small to determine the zinc extraction values accurately enough. Theother forms of phenylene diamine: meta- and para- phenylene diamines were also testedunder low pulp density conditions.4.2.2 High Pulp Density LeachingHigh pulp density leaching was carried out in the presence and absence of a surfactant.The principal surfactants used as dispersants were lignin sulfonic acid and orthopheny-lene diamine, although some experiments were performed using metaphenylene diamine(MPD) as the surfactant. The decision to use these surfactants was based on the resultsof the low pulp density tests'. The leaching time was 35-60 minutes. The particle sizedistribution of the concentrate sample used was:• 3.4 % +170 mesh• 43 % -170+230 mesh• 23.8 % -230+325 mesh and'The digestion was performed with concentrated sulfuric acid and dropwise nitric acid additions toa beaker on a hot plate.In the low pulp density leaching experiments only lignin sulfonic acid, OPD and MPD gave favourableresults.Chapter 4. EXPERIMENTAL TECHNIQUES^68• 29.8 % -325 meshEquipmentThe details of the equipment and the setup used for the high pulp density pressureleaching are shown in Figure 4.11.Experimental ApproachChemical analysis of the stock (or leach) solution for the oxidative pressure leaching was:• 0.14 mole/L Fe 2+• 1.27 moles/L Zn 2+ (in the form of sulfate)• 1.12 moles/L (or 109 g/l) sulfuric acid which dissociated incompletely to give• 1.95 moles/L total acidity' (H+).The recipe was designed so that enough sulfuric acid would be left in the system at theend of the process even if all of the soluble zinc and iron in the concentrate were dissolvedin order to prevent or minimize the oxidation of elemental sulfur to sulfate, and also toprevent the hydrolysis of ferric ions.A standard leaching procedure was followed throughout the course of the study.About 100±5 grams of concentrate was weighed into an autoclave and 1.5 L of leachsolution was added. The autoclave was sealed air-tight and flushed with nitrogen gas w .The autoclave was connected to the computer and the gas mass flow meter (as shown'The method of determination of total acidity is described in the appendix B.'This was to make sure there was no residual oxygen present in the space above the solution in theautoclave before the introduction of commercial grade oxygen. This was to ensure that at least withinthe limits of the experimental technique there was no reaction taking place during the heating period.Chapter 4. EXPERIMENTAL TECHNIQUES^ 69in Fig. 4.11) via the gas tubing to measure the oxygen flow rate at 1 second time inter-vals. Since any leakages in the system were registered on the computer, adequate stepswere taken to eliminate leaks. The autoclave was heated to the working temperatureof 140+1°C. In the runs with surfactants, a predetermined amount of the surfactant"under study was weighed into the injection unit. A brief sketch of the injection unitand its operation is shown in Fig. 4.12. At the working temperature, the agitator wasstarted and oxygen was admitted into the autoclave via valve A and through the injectionunit to pressurize it initially (valve C should be closed when performing this operation,B was opened all the time). Immediately after this operation, valve A was closed andC opened to allow oxygen to flow through the flow meter and into the autoclave as itis being consumed in the leaching reaction. The computer then began to register theoxygen flow rate. All these operations were completed within a time window of less than8 seconds and it was assumed that the errors introduced into the data as a result ofsome reactions occurring in this time frame were very minimal as compared to the totalextent of reaction in the predetermined time of 60 minutes or 35 minutes for the process(from the data obtained, it was in fact, estimated to be within +0.4 %). The oxygenwas delivered at a pressure of 1100 kPa. The use of the injection device ensured that thesurfactant was introduced into the autoclave at the same time as the oxygen (when thereaction was assumed to begin).At the end of the pre-determined time (for leaching), the oxygen flow was stopped andthe autoclave was quickly quenched in an ice water bath to about 75-80°C and filtered.The filtrate was analysed chemically for metal, sulfur and total acid content. The leachresidue was washed and dried in a vacuum oven for about 4 days and samples were takenfor chemical analysis.Surfactants were added in batch form for all the tests carried out under high pulp density conditions.Chapter 4. EXPERIMENTAL TECHNIQUES^ 704.3 INFRARED STUDIESInfrared studies were performed using the attenuated total reflection-Fourier transforminfrared technique (ATR-FTIR). The application of this technique permitted the studyof the surfactant adsorption "in-situ". The transmission KBr disc/wafer technique wasalso employed to record the infrared spectra of the concentrate and the surfactants usedbefore they were introduced into the solution.ChemicalsThe chemicals used were potassium bromide of infrared quality, sphalerite concentrate'',lignin sulfonic acid and OPD. All other chemicals were of reagent grade.EquipmentThe ATR accessory had a germanium crystal internal reflection element (of 45° angle ofincidence). The FTIR equipment used was a Bomem MB 100 series spectrophotometerwith a frequency range of 6000-200 cm -1 .The accessory was a model 11086 Specaclamp ATR. It had a micrometer controlledclamp-down facility. The sample material in the form of a paste/film or fine powder wasevenly spread on the reflecting crystal and then sandwiched between the crystal and atop pressure plate to which an even pressure was applied via a micrometer thumb screwand a leaf spring. The applied pressure could be set with the micrometer.'The concentrate was subjected to 35-40 minutes of fine (dry) grinding to expose fresh mineral surface.Chapter 4. EXPERIMENTAL TECHNIQUES^ 714.3.1 KBr TechniqueThe FTIR equipment was first calibrated using polystyrene disci' prior to any sampleanalysis. Regular calibration of the equipment was adopted since any change in atmo-spheric conditions could influence the spectra.The absorption spectra of the organic surfactant and the sphalerite concentrate wererecorded by both KBr transmission and ATR-FTIR technique.About 3 mg of the reference material' and 300 mg of KBr were thoroughly mixedand ground in an agate mortar and pressed into a disc at a pressure of about 28000-34000kPa. Two pellets/discs were made for each material to ensure reproducibility.Infrared spectra were recorded from the discs with the spectrophotometer, the samplesbeing mounted in a holder in the path of the radiation beam, (figure 4.13). The inabilityof the KBr to absorb any infrared radiation between 4000-400 cm' wavenumber meansthat any absorption band within this region was due to the sample. Prior to the recordingof any sample spectrum, the reference spectrum (of the sample chamber) was taken.The sample spectrum obtained was then plotted using a plotter connected to the FTIRcomputer.4.3.2 Attenuated Total Reflection TechniqueThe ATR-FTIR technique was employed to investigate the nature of the surfactant,adsorption at the solid-liquid interface in-situ. It permitted the recording of the spectraof the sphalerite conditioned in different aqueous phases. This technique was used forthe spectral recording of both surfactant solutions and the conditioned (wet) sphaleritemineral. Figure 4.14 shows the set-up of the AIR accessory used.13 This disc was supplied, together with its spectrum by the manufacturers of the equipment"The reference materials referred to here are the lignin sulfonic acid, OPD and the dry sphaleritecon centrate.Chapter 4. EXPERIMENTAL TECHNIQUES^ 72A spectral record was first taken of the solvent (distilled water containing 0.036 mol/Lacid) by spreading a thin film of the solvent on the germanium reflecting crystal andmounting the accessory in the infrared beam path.A 3-4 gram sample of the concentrate was conditioned' in the solvent (acidifieddistilled water) for about 20 minutes. This time window was selected based on the effectof time on the extent of zinc extraction from the leaching studies. A paste of the wetsample was then uniformly spread on the germanium crystal' and mounted in the beampath to record the spectrum of the wet sample. These spectra served as the blank orreference spectra of the solvent and natural sphalerite when wet.The scanning was done at a resolution of 4 cm' and 100 scans were done for eachsample. After each test, the reflecting crystal was thoroughly cleaned using warm chromicacid followed by cold dilute nitric acid and distilled water in order to remove any residualchromate ions.Lignin sulfonic acid testsTwo solutions, each containing 5.0 g/L of lignin sulfonic acid, and 0.036 mol/L sulfuricacid were prepared. One solution contained 0.025 mol/L ferric ions (added in the formof sulfate). The other solution did not contain ferric ions. The spectra for the surfac-tant solutions were recorded' by spreading a thin film of the solutions on the internalreflection element. Spectra of 3-4 grams of concentrate conditioned in the appropriatesurfactant solution for about 20 minutes were then recorded. After each conditioning inthe surfactant solution, the sample was cold washed 3-4 times with the solvent (distilledwater) and then subjected to a more drastic washing with boiling distilled water followedby cold washing again with the solvent. The infrared spectra of the washed samples were'The concentrate was suspended for about 15 minutes and then allowed to settle.'Sample is sandwiched between the reflecting crystal and the pressure plate.''These were used for the appropriate spectral subtraction.Chapter 4. EXPERIMENTAL TECHNIQUES^ 73also recorded. Prior to each scanning, the background spectrum of the solvent had to berecorded.Below is the list of the different conditions studied:• KBr analysis of lignin sulfonic acid• ATR analysis of 5.0 g/L surfactant and ferric solution• ATR analysis of sphalerite conditioned in surfactant/ferric solution• ATR analysis of washed sphalerite after conditioning in surfactant/ferric solution• ATR analysis of sphalerite conditioned in surfactant solution in the absence of ferricions• ATR analysis of washed sphalerite after contact with surfactant solution in theabsence of ferric ions• KBr analysis of leached sphalerite residueOrthophenylene diamine testsA surfactant solution containing 5.0 g/L of OPD, and 0.036 mol/L sulfuric acid wasprepared. A 0.025 mol/L ferric solution (in the form of sulfate) was also prepared.The spectrum for the surfactant solution was recorded' by spreading a thin film ofthe solution on the internal reflection element as described earlier. About 3-4 gramsof concentrate was conditioned in the surfactant solution (with and without ferric ionpre-treatment") for about 20 minutes. The spectra for the conditioned samples were'These were used for the appropriate spectral substraction.'In some of the tests carried out, the mineral concentrate was first preconditioned in a solutioncontaining 0.025 mol/L ferric ions and 0.036 mol/L acid for 15-20 minutes, washed 3 times with thesolvent before being introduced into the surfactant solution.Chapter 4. EXPERIMENTAL TECHNIQUES^ 74then recorded. After contact with the surfactant solution, the sample was subjected tothe washing procedure described above and scanned as well. The background spectrum(of the solvent) had to be taken prior to each test.Listed below are the different conditions studied:• KBr analysis of orthophenylene diamine. This was compared with available spectrain the literature.• ATR analysis of 5.0 g/L surfactant.• ATR analysis of sphalerite conditioned in surfactant solution (no ferric pre-treatment).• ATR analysis of washed sphalerite after conditioning in surfactant.• ATR analysis of sphalerite conditioned in surfactant solution after pre-treatmentwith ferric ions.• ATR analysis of washed sphalerite after contact with ferric/surfactant solution.Chapter 4. EXPERIMENTAL TECHNIQUES^ 75deFigure 4.8: Profile of pendant drop.Figure 4.9: Contact angle between a liquid and solid surface: a) in the absence of anysurfactant; b) in the presence of a surfactant." D'^ OHLS = liquid sulphurZM = marmatiteCV0Chapter 4. EXPERIMENTAL TECHNIQUES^ 76P: converging lens ;^I: gas inletL: light source for backlightingB: stainless-steel bomb, (7.6 x 20.3 cm)T: thermowell (for thermocouple)C: camera for taking pictures0: rectangular optical glass cellA: drive-screw syringe with plunger DV: replaceable pyrex glass tubing tip(for creating liquid sulphur drops)S: sapphire windows (1.3 cm diameter)K: adjustable height seatH: pressure gaugeFigure 4.10: Cross-section of high temperature and pressure stainless-steel bomb.Chapter 4. EXPERIMENTAL TECHNIQUES^ 77A,B,C - VALVESL - PORTABLE COMPUTERD - DATA ACQUISITION BOARDM - GAS MASS FLOW METERT - THERMOCOUPLEI - INJECTION UNITP1,P2 - PRESSURE GAUGESE - AUTOCLAVEG - OXYGEN TANKFigure 4.11: Experimental set-up for monitoring oxygen flow rate.Flush water held hereTo openChapter 4. EXPERIMENTAL TECHNIQUES^ 78AUTOCLAVEFigure 4.12: Injection Unit.SampleIncident radiation DetectorFigure 4.13: KBr Transmission. ■■ ■LEAF SPRING ON PRESSURE PLATEChapter 4. EXPERIMENTAL TECHNIQUES^ 79MICROMETER SCREWTHUMB SCREW TO SECURE PRESSURE PLATPRESSURE PLATEGERMANIUMINTERNAL REFLECTION ELEMENTINCIDENT RADIATIONM1 - INPUT MIRRORM2 - OUTPUT MIRRORM2OUTGOING BEAMSAMPLE FILMra7,1‘ . --74111,VATIIMILTAWNII=MIFigure 4.14: ATR Accessory.Chapter 5RESULTS AND DISCUSSIONThe results and complete discussion of each area of study are presented here and attemptsare made to correlate the results.5.1 INTERFACIAL PHENOMENA5.1.1 ResultsInterfacial TensionFigure 5.15 shows the results obtained using the different additives as surfactants in theinterfacial tension measurements. Included in the figure are the data obtained (from anearlier work) using lignin sulfonic acid [54, 55].In the absence of any surfactant, the liquid sulfur-aqueous solution interfacial tensionmeasured 54-55 mN/m. It can be observed that irrespective of how much OPD is addedto the solution, there was no significant change in the interfacial tension. However, theother surfactants influenced the interfacial tension to different degrees, depending on thesurfactant. The most efficient surfactant in terms of interfacial tension reduction wasnaphthalene sulfonic acid-formaldehyde condensate. It reduced the interfacial tensionsto values lower than those achieved with equal concentrations of lignin sulfonic acid(about 20-22 mN/m as compared to 28-30 mN/m for lignosol). It required only aminimum dosage of about 0.1 g/L to effect the maximum influence whereas lignin sulfonicacid required a concentration of 0.3 g/L to effect its maximum influence. Cocoamido80Chapter 5. RESULTS AND DISCUSSION^ 81Table 5.2: Surface activity of OPD at a concentration of 0.3g/L in the presence of [Fe 2+]= 0.05 M, [Fe3-1 = 0.12 M, [H 2 SO 4 ] = 0.2 M, and [ZnSO 4] = 1.2 M, Temp.= 135+5°CExpt. 1 2 3 4 5 6 Mean value Std. Deviationy, mN/m 52.7 54.4 54.2 54.9 55.5 54.7 54.4 0.86Table 5.3: Effect of silicate ions (0.15g/L) on the surface activity of OPD (0.3 g/L) inthe presence of [H 2 SO4 ] = 0.2 M and [ZnSO 4 ] = 1.2 M, Temp.= 135+5°CExpt. 1 2 3 Mean value Std. Deviationy, mN/m 53.6 52.4 52.5 52.8 0.54hydroxyl sulfobetaine (CAHSB) has about the same effect as lignin sulfonic acid up to aconcentration of about 0.3 g/L beyond which it further reduces' the interfacial tensionuntil it attains a maximum effect at about 0.5 g/L.Table 5.2 shows the surface activity of OPD in the presence of acid, ferrous andferric ions. The data show that the presence of these ionic species did not change thebehaviour of OPD at the liquid sulfur - aqueous solution interface. Table 5.3 shows theeffect of silicate ions on the surface activity of OPD 2 . Again, it was observed that underthe conditions of the experiment, these species do not have any effect on the liquid-liquidinterfacial tensions. In the absence of a surfactant but in the presence of ferric, ferrousand hydrogen ions (i.e. an acid), the liquid-liquid interfacial tension measured 53-54mN/m.The effect of ferric/ferrous ions on the efficiency of naphthalene sulfonic acid-formaldehydecondensates as a surfactant was significant. Figure 5.16 shows how the surfactant de-grades with time (in terms of an increase in the liquid-liquid interfacial tensions) in the'At surfactant concentrations greater than 0.3 g/L, lignin sulfonic acid does not exert any noticeableinfluence on the liquid sulfur-aqueous solution interfacial tension.These tests were performed in order to find out if the enhanced leaching rate observed with OPDwas due to the presence of these species (or not) which are typical of zinc pressure leaching conditions.Chapter 5. RESULTS AND DISCUSSION^ 82Table 5.4: Contact angle measurements in the presence of the various surfactants: con-ditions - [ZnSO 4 ] = 1.2 M, Temp.= 135+5°CSurfactanttypeSurfactant concentration, g/L0.05 0.07 0.15 0.3 0.33 0.6Naphth. 154+5 148+5 150+4 155+5 - -CAHSB - 143+4 148+5 152+4 148+5OPD 124+5 125+6 127+5 125+4 -TAHSB - 133+5 138+5 135+4 -CAB - 135+5 136+6 134+6Lignin. - 144+5 146+5 143+5 -presence of ferric and ferrous ions in the solution in one of the measurements made, wherethe time in minutes includes the time taken to heat the system from room temperature tothe working temperature' of 130+5°C, and the period over which the photographs weretaken of the pendant drops. The data indicate that this surfactant has a fairly high rateof degradation. The degradation could be due to the surfactant reacting with either theferrous or the ferric ions in solution'. As one of its chemical properties, this surfactant isknown to degrade rapidly in the presence of an oxidizing agent [83, 163]. This propertyprobably indicates a ferric reaction degradation mechanism.Contact Angle MeasurementAdvancing contact angles have been measured in the liquid sulfur - aqueous zinc sulfate -marmatite system both in the absence and presence of the various additives or surfactants.Tables 5.4 and 5.5 show the negligible changes effected in the liquid sulfur - aqueouszinc sulfate - marmatite mineral contact angle in the presence of a given surfactant(at varying concentrations) and ionic species. All measurements made had a spread of'It takes roughly 13-15 minutes of heating to reach this temperature.'No attempt was made to determine which of the iron species was involved in the degradation process.Chapter 5. RESULTS AND DISCUSSION^ 83Table 5.5: Contact angle measurements in the presence of different surfactants and ionicspecies; conditions: [ZnSO 4] = 1.2 M, [Fe2+] = 0.05 M, [Fe 3+] = 0.12 M and [H 2 SO 4 ] =0.2 M, Temperature = 135°CSurfactanttypeSurfactant concentration, g/L0.15 0.3Naphth. 149+5 148±5CAHSB 145+4 147±5OPD 136±5 135+3TAHSB 116±5 113+6CAB 119±5 122+5Lignin. 146±5 148+5within ±6°. In the absence of any of the surfactants, the contact angles measured wereabout 80±5° and 98+5° respectively in the absence and presence of ionic species suchas H+, ferric and ferrous, Figure 5.17. Again, it was observed that the largest increasein contact angle was achieved with naphthalene sulfonic acid-formaldehyde condensatesand CAHSB with orthophenylene diamine effecting the least increase (Table 5.4). Inthe presence of other ionic species (Table 5.5), the contact angles remained almost thesame for the naphthalene sulfonic acid-formaldehyde condensates, lignin sulfonic acidand CAHSB. However, there was a small increase with respect to OPD while therewas a decrease in the case of TAHSB. Considering these increases in the contact anglecoupled with the interfacial tension reductions it is very easy to conclude that naphthalenesulfonic acid and CAHSB should be more efficient as surfactants in the zinc pressureleach process than lignin sulfonic acid and OPD should be the least efficient surfactant.Tables 5.6 and 5.7 summarize the work of adhesion values for the liquid sulfur-aqueoussolution-marmatite system in the presence of the different surfactants and ionic species'.These tables show that under pressure leaching conditions', it is much easier to displace'The mineral is considered as a two-dimensional plane. The figures shown in the tables will bedifferent if the mineral is treated as a three-dimensional object.Chapter 5. RESULTS AND DISCUSSION^ 84Table 5.6: Work of adhesion, IV„, in the presence of 0.3g/L of different surfactants;conditions: [ZnSO 4 ] = 1.2 M, Temperature = 135°C; no other ionic species present insolutionSurfactant None Lignin. Naphth. OPD CAHSB TAHSB CAB147„(mJ/m2 ) 64.0 5.3 2.3 25.4 4.0 9.0 10.7Table 5.7: Work of adhesion, 147,, in the presence of 0.3g/L surfactants; conditions:[ZnSO 4] = 1.2 M, [H 2 SO4 ] = 0.2 M, [Fe 2+] = 0.05 M and [Fe31 = 0.12 M, Temperature= 135°C.Surfactant None Lignin. Naphth. OPD CAHSB TAHSB CABWa (mJ/m2 ) 46.2 4.6 3.6 16.0 4.5 25.0 18.8liquid sulfur from the particle surface in the presence of naphthalene sulfonic acid, ligninsulfonic acid, CAHSB, TAHSB than it is with OPD and CAB and thus effect high zincdissolution rates. Figures 5.17-5.19 show photographs of liquid sulfur resting on themarmatite mineral surface in different aqueous environments.5.1.2 DiscussionSurface tension is the boundary tension between a liquid and a gas or vapour, and in-terfacial tension is the boundary tension at a phase boundary between two immisciblecondensed phases. Surface/interfacial tension is a measure of the free energy of thefluid interface and at equilibrium it is the same at every point in all directions alongthe surface/interface. It is defined as the work required to increase the area of a surfaceisothermally and reversibly by a 1111 it amount. For solid systems, it is defined as the restor-ing force necessary to bring freshly exposed surface molecules to equilibrium position. Ithas units of dynes/cm (cgs units) or milli-Newton/in (SI units). Any substance which6 i.e. if these data are projected into the zinc pressure leaching conditions.Chapter 5. RESULTS AND DISCUSSION^ 85lowers the surface/interfacial tension of a system tends to accumulate at the interfacesuch that its concentration at the interface is higher than the bulk concentration.The measurements reported here indicate the roles of the individual surfactants atthe interfaces present in the liquid sulfur-aqueous solution-sulfide mineral system (whichcan be projected into the industrial zinc pressure leach process). With the exceptionof orthophenylene diamine, all surfactants used here reduced the liquid sulfur-aqueoussolution interfacial tension from a maximum' of 54-55 mN/m in the absence of any sur-factant and at a zinc sulfate concentration of 1.2 mol/L to their respective minimumvalues: lignin sulfonic acid effects a maximum decrease to 28-30 mN/m in the presenceof 0.3 g/L of surfactant; naphthalene sulfonic acid-formaldehyde condensates effected amaximum drop to 20-22 mN/m at a concentration of 0.1 g/L. Beyond these surfactantdosages, the interfacial tensions were insensitive to any further increase. Among the am-photeric surfactants, cocoamido hydroxyl sulfobetaine (CAHSB) was the most effective.It has about the same effect as lignin sulfonic acid up to a concentration of about 0.3g/L beyond which it further decreases the interfacial tension to about 19-21 mN/m atabout 0.5 g/L; additional increase of concentration beyond this point did not introduceany further change in the liquid sulfur-aqueous solution interfacial tension. In spite ofthe fact that cocoamido betaine (CAB) and tallowamido hydroxyl sulfobetaine (TAHSB)were effective in reducing the interfacial tension, their working efficiencies are lower thaneither lignin sulfonic acid or CAHSB or naphthalene sulfonic acid, Figure 5.15.With an increase in the surfactant concentration, the amount of the additive adsorbedat the liquid-liquid interface increases (accompanied by a corresponding decrease in theinterfacial tension) until a state is reached (i.e. the limiting state) when the interfaceis virtually completely covered with a layer (or layers) of the surfactant molecules. It7 Fanelli [82] has reported that liquid sulfur has a surface tension of 57-61 mN/m in the temperaturerange 120-150°C, decreasing with increasing temperature. Interfacial tension values usually lie betweenthe individual surface tension values for the individual condensed phases.Chapter 5. RESULTS AND DISCUSSION^ 86is at this point that increasing the surfactant concentration further does not effect anyadditional influence on the interface. This limiting point corresponds to about 0.3 g/Lfor lignosol, 0.5 g/L for CAHSB, 0.1 g/L for naphthalene sulfonic acid-formaldehydecondensates, etc.The interfacial activity of OPD (at the liquid-liquid interface) is insensitive to thepresence and absence of solutes such as ferric, ferrous, hydrogen and silicate ions.The adsorption of a surfactant at the solid-liquid interface is probably the mostversatile phenomenon involved in surfactant applications. The adsorption of a surfactantby a solid can transform the solid surface from a hydrophobic to a hydrophilic one andvice versa. Adsorption can cause finely divided solids to be readily dispersible in a liquidor make the solid particles flocculate. Surfactant adsorption is important in the flotationseparation of mineral particles.. It can also cause the solid to be easily wetted by a givenliquid or very difficult to wet by sulfur. It is the last of these properties that allows forthe development of an efficient leaching process for metal sulphides in the presence ofliquid elemental sulfur.Ideally, the ability of these surfactants to reduce the liquid sulfur-aqueous solutioninterfacial tension (with the exception of OPD) to the lowest value possible coupled withthe ease of increasing the liquid sulfur-mineral contact angle (all surfactants increased thecontact angle from about 80+5° in the absence of any additive), Figures 5.17-5.19, createsa favourable environment for the liquid sulfur to be dispersed from the mineral surface.This phenomenon 'rolls' the liquid sulfur film (on the surface of the sulfide mineral)into droplets which are easily removed from the solid surface by vigorous agitation anddispersed in the liquid. This is the basic theoretical principle underlying the removal of aliquid from the surface of a solid substrate. The process of wetting of the mineral particlesby the liquid sulfur is the result of the relatively high degree of attraction between thesolid and the liquid molecules. A high contact angle value indicates that the liquid sulfurChapter 5. RESULTS AND DISCUSSION^ 87has a low affinity for the mineral surface whereas a low value for the contact angle is anindication of the liquid sulfur's ability to spread over the mineral surface and hence makethe mineral less accessible to the leaching solution. Complete wetting of the solid by theliquid sulfur implies a contact angle of 0° and absolutely no wetting corresponds to anangle of 180°.When this phenomenon is projected into the zinc pressure leach process, a surfac-tant that increases the contact angles tends to lower the work of adhesion' which is thereversible work required to separate a unit area of the liquid from the solid substrate.Tables 5.6 and 5.7 show the work of adhesion calculated for the different aqueous envi-ronmental conditions. The work of adhesion, Wa , between the two dissimilar phases canalso be expressed as the sum of the different intermolecular forces acting between theliquid and the solid mineral [149, 152] such thatWa = WD+ Wa^ (5.45)where WD is due to the dispersion force component of the intermolecular forces whichis always present in all systems and Wr is due to the polar component of the interact-ing forces depending on the polar nature of the phases and occurs to different degreesdepending on the system and is given by=,^TIT)^-F^+  ^(5.46)-= hydrogen bonding force component= bonding force component due to dipole-dipole interactionsWaE = bonding force component due to electrostatic interactions14": = 7- bonding component of the adhesional work, etc.For non-polar systems, only dispersive forces are present and hence such a phase interacts'The contact angles give us an idea about how much energy is necessary to remove a drop of liquidsulfur of unit area from the surface of the mineral substrate.Chapter 5. RESULTS AND DISCUSSION^ 88with polar systems only through WD, e. g. zinc sulfide mineral (polar) and liquid sulfur(non-polar).If the zinc sulfide particles are treated as hypothetical three-dimensional objects, cubes of unit volume or dimensions [154], a similar pattern in the work of adhesionvalues are obtained. Consider Figures 4.9 and 5.22. From equation 4.40^7mA - 7MS = -rsA cos 0^ (5.47)neglecting the vertical force component' of lisA. Assume the particle in Figure 5.22 iscompletely wetted by the aqueous solution (stage I) and that stages II, III, and IV arethe different stages that each particle goes through in order to completely spread liquidsulfur all over the surface. By moving from I to II, adhesional wetting of the particle byliquid sulfur takes place such that the work of adhesion W1 is given by:^Wr = (57mA + 7Ms) — (67mA^+ 7 sA)^(5.48)= "OviS^— tSA (5.49)which coupled with equation 5.47 gives^= -7sA (1 + cos 0)^ (5.50)This implies that in order to prevent wetting of the surface (stage II) by the liquidsulfur, an amount of work (positive work) equal to or greater than -ysA (1 + cos 0) mustbe done10 . The expression in equation 5.48 indicates that there is a change in the netsurface energy when the system consisting of mineral-aqueous solution and liquid sulfur-aqueous solution interface changes to one consisting of a liquid sulfur-mineral interface(stage II). If")'MA^-)SA > 7MS^ (5.51)'Chattoraj and Birdi [149] have reported that on a rigid solid surface, the vertical component isnegligible but on soft surfaces is not negligible.10 This work had been compiled in Table 5.6.Chapter 5. RESULTS AND DISCUSSION^ 89adhesional wetting of the solid will occur. Thus, in order to prevent adhesional wettingfrom taking place (in physical terms), an applied surfactant must be able to increase -ymsand/or decrease either -01A or -ysA or both.During stage II, the process is said to be immersion wetting where the particle iswetted on all sides but one by liquid sulfur. The accompanying work of immersion is:which is equivalent toorIVII = 5 7MS + 1A1A) — (57.111.4 yMS)= 4( -YAts — 7ArA)Wil = — 21-ysA cos 0(5.52)(5.53)(5.54)from equation 5.47 above. This implies that if -o f s is less than -ymA , the second stagesulfur wetting will occur. To prevent this immersional wetting from taking place, anequivalent amount of work must be done on the system.During the third and final stage of the sulfur wetting process, the particle is completelycovered by the sulfur and the process is referred to as spreading wetting. Under zincpressure leaching conditions, this is the stage when the leaching of the sulfide particleby the leaching solution is almost stopped because the aqueous solution has negligibleaccess to the mineral surface. The accompanying work of spreading is given by:" TM^(6111S^"Y SA )^( 5-)A1S^ (5.55)From equation 5.47, this simplifies to^TV111 = —^;sA (COS 0 — 1)^ (5.56)Thus. for complete wetting of the particle by liquid sulfur, the resultant work ishh = TV/ +^+ 117111^ (5.57)Chapter 5. RESULTS AND DISCUSSION^ 90Table 5.8: Resultant work, W, for complete sulfur wetting (a hypothetical situation) inthe presence of 0.3 g/L of surfactant and [ZnSO 4 ] = 1.2 MWork(mJ/m2 )Surfactant typeNone Lignin. Naphth. OPD CAHSB TAHSB CABW / -64 -5.3 -2.3 -23.4 -4.0 -9.0 -10.7W1 1 -37.5 95 74.5 125.3 91.6 104 103.2Wm 44.9 52.8 40.0 85.9 49.9 61.0 61.5W -56.6 142.5 112.2 187.8 137.5 156 154This implies that in order to prevent complete wetting of the zinc sulfide particles inthe zinc pressure leaching process (assuming the particles are hypothetically of cubicshape and unit dimensions), an amount of work equivalent to W must be performedon the system. The presence of stable and efficient surface active agents in the systemdrastically reduces the required work to a lower value, Table 5.8, due to the ability ofthe surfactant to reduce the liquid—liquid interfacial tension and/or increase the contactangle. A negative W value implies spontaneous wetting of particles by the sulfur. Thisis most likely the condition present in the pressure leaching system in the absence ofa suitable surfactant where the particles become completely occluded with liquid sulfur(with time) after enough sulfur has been formed in the system. This leads to the formationof agglomerates and large lumps of unreacted sulfide particles, Figure 5.35.When W is positive (Table 5.8), sulfur wetting can only take place after performingan equivalent amount of work on the system i.e. there is a spontaneous and voluntarydispersion of liquid sulfur from the surface of the particles. This is what actually takesplace in the presence of a suitable (and stable) surfactant under leaching conditions.Considering the results (interfacial tensions and contact angles) obtained in thesestudies, it is very easy to conclude from the adhesional work data that surfactants withhigh interfacial tension reductions and high contact angle increases, e.g. naphthaleneChapter 5. RESULTS AND DISCUSSION^91Table 5.9: Variation of 7mA — 'yms with respect to aqueous condition; [ZnSO 4 ] = 1.2 M;other ionic species absentAqueousConditionNoSurfactantSurfactant (0.3 g/L)Lignin. Naphth.^OPD CAHSB TAHSB CAB7mA - yms, mN/m 16.8 -18.5 -16.0^-20.8 -19.2 -19.7 -19.3sulfonic acid-formaldehyde condensates, will be the most effective dispersant for liquidsulfur. In contrast, a reagent such as OPD which increases the contact angle but has noeffect on the interfacial tension should be the least effective dispersant in the pressureleach process. The reason is that (from the tables), less effort and energy should berequired to displace liquid sulfur of unit surface area from a solid substrate of an equalarea in the presence of surfactants such as lignin sulfonic acid, naphthalene sulfonic acid-formaldehyde condensates, CAHSB, CAB, and TAHSB than in the presence of OPD. Thisshould translate to higher zinc extractions in the presence of these surfactants than inthe presence of OPD in the zinc pressure leach process.It is concluded from the interfacial tensions and contact angle measurements that withthe exception of OPD, all other surfactants applied are adsorbed at both mineral-aqueousand aqueous-sulfur interfaces. OPD on the other hand adsorbs at only the solid-liquidinterface thus increasing the contact angle from 80+5° in the absence of the surfactantto 125-135+5°. From equation 5.47 above, it is realised that if the respective surfactants(with the exception of OPD) are adsorbed only at the liquid-liquid interface, then -0 1Aand yms must remain unchanged and 1 A/A - '1Ais should also be constant' irrespectiveof the aqueous conditions (for a given surfactant). It so happens that this is not thecase, (Table 5.9). The interfacial tension is dependent on the solution composition,confirming that the respective surfactants are adsorbed at both solid-aqueous and liquidthe value of this parameter in the absence and presence of a surfactant should be the same.Chapter 5. RESULTS AND DISCUSSION^ 92sulfur-aqueous interfaces.Generally, the adsorption of surfactants onto solids proceed through two steps:• the adsorption of the surfactant species through interactions such as electrostaticand/or van der Waals etc., between surfactant species and solid surface• the adsorption of a surfactant through interactions (polar, hydrophobic, hydro-gen bonding etc.) between the already adsorbed surfactant and the approachingmolecules.In the first step, the surface active species are individually adsorbed (as ions or molecules)in the first adsorbed layer on the solid surface through electrostatic ( present in only ionicsurfactants where the surface active ions are adsorbed on the oppositely charged solidsurface) and/or specific (i.e. van der Waals) attraction. In the second step, the adsorptionincreases dramatically as hemimicelles 12 form on the adsorbent through association orhydrophobic interaction between the hydrocarbon chains of the surface active species,and each of the ions or molecules adsorbed in the first steps provide a possible activecenter for aggregation.Surface Excess at the Liquid Sulphur-Aqueous solution InterfaceThe Gibb's general thermodynamic expression for adsorption of a solute from dilutebinary solutions at liquid-liquid interfaces at constant temperature and pressure is givenby [149, 150]:- d-) = 1' 12 .4 2^(5.58)orc2 . d-y^1^dy^ i^.^-2 ^RT dc2^RT • dlnc2(5.59)12 Hemimicelles refer to the formation of surfactant aggregates on the surface of the adsorbent.Chapter 5. RESULTS AND DISCUSSION^ 93wherec2 = concentration of solute in the bulk solution= relative surface/interface excess of solute at the interface per unit area of interface;it is an estimate of the concentration of the adsorbed species at the interface in excess ofthe bulk concentration of the solute; component 1 is considered to be the solvent whoseinterfacial concentration is assumed to be the same as in the bulk.T = operating temperature and R = universal gas constant, 8.32 Joules/(mole-K).For a multicomponent solution, e.g. solution containing an organic electrolyte (sur-factant) and inorganic electrolyte (such as zinc sulfate)— dy = RTE ir^ (5.60)where is a parameter for the activity coefficient of component i and is given bydln f,4.i =1 + ^dine, (5.61)fi = the activity coefficient of component i of concentration c,. For the adsorption of anelectrolyte (i.e. an ionic surfactant) at a liquid-liquid interface, equation 5.60 is modifiedto take into account the ionic speciation of the surfactant such that:— dy = niRTFR.y z .dlncRx.^ (5.62)where RX, represents the ionic surfactant. The details of the derivation of equation5.62 are available in appendix C. For dilute solutions of surfactants, the value of m isestimated from either Helmholtz or Gouy's model of the electrical double layer in relationto Gibb's adsorption equations [149]. Gouy estimates thatZ 2772 = 1 +  ^ (5.63)z^eewhere50012oT)is the ratio of molar concentrations of the inorganic to organic salts (surfactant) presentChapter 5. RESULTS AND DISCUSSION^ 94in the mediumz is the valency of the organic surface active anion= surface/interfacial potentiala = Boltzman's constant.In the absence of any inorganic electrolyte e.g. zinc sulfate in our situation, ± = 0 and m= 1 z. When the concentration of the organic surfactant ( electrolyte) is very small asopposed to the inorganic component e.g. 0.3 g/L surfactant versus 345 g/L zinc sulfatesalt, ± tends to infinity and m is equal to unity, resulting inOrCRXz dyFRX z = -^•RT deRx z(5.64)1^d-yFRXZ = ^ (5.65)2.303RT . d/ogcRx zat constant temperature. The valency of the organic electrolyte disappears since in isequal to unity under the conditions of this study. Thus, the interfacial surfactant excess,FRxx , can be computed for each concentration of surfactant provided the molecular weightof the surfactant is known.To compute interfacial excess for lignin sulfonic acid used in this study, a molecularweight of 4,000 was assigned'. The interfacial excess was computed for the naphthalenesulfonic acid-formaldehyde condensates adsorption as well; it has a molecular weight ofabout 2000. The interfacial excess plots are shown in Figures 5.23 and 5.24. These datashow the concentration of the surfactant at the liquid sulfur-aqueous solution interface inexcess of the bulk concentrations. Similar computations could be made for the amphotericsurfactants if information on their molecular weights was available'.13 Lignin sulfonic acid has no definite molecular weight: it varies from 4000 for the lower series to ashigh as 100,000 for the high molecular weight groups; however, the most common species are reportedto have molecular weights of about 4,000.14 The R component of the structure can have between 12 and 24 or more carbon atoms.Chapter 5. RESULTS AND DISCUSSION^ 95The critical micelle concentrations for the lignosol and the naphthalene sulfonic acid—formaldehyde condensates are shown on the plots'. The critical micelle concentration(cmc) is the limiting concentration point at which the interfacial concentration of thesurfactant reaches saturation. The cmc is about: 0.5-0.6 g/L for CAHSB; 0.1 g/L fornaphthalene sulfonic acid—formaldehyde condensates, and 0.3 g/L for lignin sulfonic acid.SummaryThe results of the interfacial studies have shown that of all the surfactants used, OPDwas the least effective in terms of interfacial tension reduction and contact angle effects,both in the presence and absence of other ionic species which are typical of zinc pressureleaching conditions. Both naphthalene sulfonic acid formaldehyde condensates and co-coamido hydroxyl sulfobetaine were quite effective in terms of interfacial activity. Theirinterfacial activities compete very well with that of lignin sulfonic acid. These observa-tions can easily lead one to conclude that they can be effective as dispersants for liquidsulfur in the zinc pressure leach process. Again, OPD could be expected to be leasteffective of all the reagents studied. However, as has been observed with naphthalenesulfonic acid formaldehyde condensates, the efficiency of a given surfactant is related tothe stability of the surfactant under the operating conditions, (Figure 5.16).'These points occur in Figure 5.15 at the points where maximum effect is attained irrespective ofsurfact ant concentration.60• 50• 400c020-c 10„ Naphth., CAHSBL• ignin.•^14^15^16^17^18^19^20 22216050EE 402 20a)10Chapter 5. RESULTS AND DISCUSSION 96605040302010o o0.2^0.4^0.6^0.8^1• TAHSBCAB• OPD0.2^0.4^0.6^0.8Surfactant conc. g/L Surfactant conc. g/LFigure 5.15: Liquid sulfur-aqueous solution interfacial tensions in the presence of differentsurfactants; [ZnSO4 ] = 1.2 M; Temperature = 135+5°C; PN2 = 620 kPa.Time of heating system (min)Figure 5.16: Liquid sulfur-aqueous solution interfacial tensions in the presence [Fe3+]= 0.12 M, [Fe 2+} = 0.05 M and 0.3 g/L naphthalene sulfonic acid; [H 2 SO 4 ]^0.2 M.Temperature = 135+5°C;^= 620 kPa.Chapter 5. RESULTS AND DISCUSSION^ 97Figure 5.17: Liquid sulfur-zinc sulfide mineral contact angle in surfactant-free system;[ZnSO4] = 1.2 M, Temperature = 135°C; PN2 = 620 kPa.Figure 5.18: Liquid sulfur-zinc sulfide mineral contact angle in the presence of 0.3g/Lnaphthalene sulfonic acid—formaldehyde condensates, [ZnSO 4 ] = 1.2 M, Temperature =135°C; PN2 = 620 kPa.Chapter 5. RESULTS AND DISCUSSION^ 98Figure 5.19: Liquid sulfur-zinc sulfide mineral contact angle in the presence of 0.15g/LOPD; conditions: [Fe2+] = 0.05 M, [Fe3+] = 0.12 M, [H 2SO4] = 0.2 M, Temperature =135°C; PN2 = 620 kPa.Figure 5.20: Photograph of liquid sulfur suspended in surfactant-free zinc sulfate solution:conditions; [ZnSO 4 ] = 1.2 M, Temperature = 135±5°C; P N2 620 kPa.Chapter 5. RESULTS AND DISCUSSION^ 99Figure 5.21: Photograph of liquid sulfur suspended in aqueous solution in the presenceof 0.3g/L naphthalene sulfonic acid—formaldehyde condensates; conditions: [ZnSO 4] =1.2 M, Temperature = 135°C; PN2 = 620 kPa.Figure 5.22: Different stages of wetting1.41.2 -MC1.0 -o^0.8 -E0.6La 0.40.21^2^3Surfactant concentration, mol/L (x 10 4 )1.0esl0.80.60Eq?x0.40.200^0.5^ 1.0^1.5Concentration of lignosol, mol/L (x 10 -4)Chapter 5. RESULTS AND DISCUSSION^ 100Figure 5.23: Interfacial excess of naphthalene sulfonic acid—formaldehyde condensates atthe liquid sulfur—aqueous solution interface; conditions: [ZnSO 4 ] = 1.2 M, Temperature= 135±5°C; PN2 = 620 kPa.Figure 5.24: Interfacial excess of lignin sulfonic acid at the liquid sulfur—aqueous solutioninterface; conditions: [ZnSO 4 ] = L2 M, Temperature = 135±5°C; P N2 = 620 kPa.Chapter 5. RESULTS .A ND DISCUSSION^ 1015.2 PRESSURE LEACHING OF SPHALERITE CONCENTRATE5.2.1 ResultsThree sets of leaching tests were performed.• Low pulp density leaching under a nitrogen atmosphere• Low pulp density leaching using oxygen overpressure• High pulp density leaching in the presence of oxygen overpressureLow Pulp Density Leaching Under Nitrogen PressureThe primary aim of the low pulp density tests was to investigate the extent of zinc ex-traction in the presence of the different surfactants, relate the extraction data to theinterfacial behaviour of the surfactants and then compare the results with the data ob-tained using lignin sulfonic acid under similar conditions. The amount of concentrateused for these tests was 1-3 grams in 1 L of leach solution. The variables investigatedwere: time of leaching, type and concentration of surfactant and the mode of introduc-tion of the surfactant into the system, either bat ch (B) or continuous (C). The averagesurfactant solution flow rate was 1.5 ml/minute.Results for the experiments using the different additives are shown in Tables 5.10-5.12.The reason for the high leaching rates in the absence of any surfactant is that theliquid sulfur generated as a result of the zinc dissolution reaction is being dispersed fromthe mineral surface by the continuous agitation as a result of the low solid percentagein the system. The low zinc extraction values obtained in the presence of the othersurfactants (apart from lignin sulfonic acid and the phenvlene diamine) suggest thatthese surfactants may be chemically unstable under the leaching conditions, as shownChapter 5. RESULTS AND DISCUSSION 102Table 5.10: Extent of zinc extraction using different surfactants under low pulp den-sity conditions and nitrogen atmosphere; initial conditions: [ZnSO 4 ] = 2 M; [H2SO4] =0.5 M; [Fe3+]=[Fe2+]=0.15 M; Temp.= 140±1°C;P ]^*2 = 1100 kPa; B-batch addition;C-continuous additionSurf. Surf.conc. (g/L) Addition method Time (min) Zn extraction (%)None 20 70±0.4None 30 72±1.0CAHSB 10.0 C 30 75±0.5Lignin. 5.0 C 30 94±1.0Lignin. 0.05 B 20 86±0.8Lignin. 0.1 B 30 90±0.5Queb. 0.05 13 20 90±0.4Queb. 0.05 B 30 96±0.5CAB 10.0 C 30 51±0.6CAB 6.0 C 20 45±0.3TAHSB 10.0 C 20 40±0.3Table 5.11: Effect of naphthalene sulfonic acid on the extent of zinc extraction un-der low pulp density conditions; initial conditions: [ZnSO 4 ] = 2 M; [H 2 SO 4 ] = 0.5M; [Fe3+]=[Fe 2+]=0.15 M; Temp.= 140±1°C; = 1100 kPa; B-batch addition;C-continuous additionSurfactant conc.(g/L) Addition method Time min.) Zn extraction (%)1.0 C 30 70±0.32.0 C 30 72±0.55.0 C 30 67±0.80.05 B 30 64±0.30.1 B 20 53±0.30.1 B 30  64±0.40.3 B 20 51±1.00.3 B 30 65±0.50.5 B 20 48±0.40.5 B^_30 63±0.30.7 B 20 48±0.30.7 B 30 61±0.2Chapter 5. RESULTS AND DISCUSSION^ 103Table 5.12: Extent of zinc leaching using OPD as surfactant under low pulp densityconditions; initial conditions: [ZnSO 4 ] = 2 M; [H 2 SO 4] = 0.5 M; [Fe3+]=[Fe2+]=0.15 M;Temp.= 140±1°C; PAT, = 1100 kPa; B—batch addition; C—continuous additionSurfactant conc.(g/L) Addition method Time (min) Zn extraction (%)0.05 B 30 95+0.20.1 B 20 75+0.40.1 B 30 94+0.35.0 C 20 78+0.35.0 C 30 87+0.4Table 5.13: Extent of zinc extraction using MPD and PPD as surfactants (batch addi-tions) under low pulp density conditions; initial conditions: [ZnSO 4 ] = 2 M; [H 2 SO 4] =0.5 M; [Fe3+].[Fe 2+]=0.15 M; Temp.= 140+1°C; PAT, = 1100 kPa.Surfactant Conc. (g/L) Time (min.) Zn extraction (%)MPD 0.1 30 95+0.5MPD 0.3 30 97+0.2PPD 0.1 30 96+0.2PPD 0.3 30 97+0.6in Figure 5.16. This could be due to these surfactants reacting with or sequestering theferric ion oxidizing agent thus depriving the mineral of the oxidant.Based on the results obtained from leaching in the presence of OPD, the other formsof phenylene diamine i.e. the meta— and para-phenylene diamines (MPD PPD) wereapplied as surfactants in some low pulp density leaching tests, (they were not part ofthe primary list). The purpose of these experiments was to determine if the position ofthe amine functional group on the aromatic ring influences the extent of zinc extraction.Under the low pulp density conditions, the data obtained in the presence of these addi-tives were quite comparable to those of lignin sulfonic acid and orthophenylene diamine.Table 5.13 shows the results obtained in the tests where these reagents were used as theChapter 5. RESULTS AND DISCUSSION^ 104surfactants.Low Pulp Density Leaching under Oxygen PressureBased on the results obtained from leaching under a nitrogen pressure, some of thesurfactants were not used under oxygen pressure. The reason was that they did not giverise to zinc extractions greater than those obtained in the absence of any surfactant.These surfactants were CAHSB, CAB, and TAHSB. It was presumed that these degradein the presence of oxidizing agents such as ferric ions and that the presence of oxygenwill also catalyse the degradation process rendering the surfactants ineffective. Thesesurfactants are reported to decompose in the presence of oxidizing agents. All surfactantsused in oxygen pressure leaching were batch additions. The amount of concentrate usedfor each of these tests was 6-12 grams in 1.5 L leach solution. Table 5.14 shows the resultsof some of the experiments. The leaching data presented in this table were obtained after30 minutes of leaching. As Table 5.14 shows, pressure leaching with oxygen resulted inhigher zinc extractions than leaching with ferric ions under a nitrogen atmosphere, thoughthe differences are not large. The differences may be due to either one or both of thefollowing reasons:• the continuous oxidation of ferrous to ferric ions in the system (due to the presenceof oxygen) ensures that ferric ions are available throughout the leaching period andare never depleted• the presence of oxygen in the system enhances the surface/interfacial activity ofthe surfactants in the systern 16 .The extent of zinc extraction achieved in the presence of the phenvlene diamines comparesvery well with that of lignin sulfonic acid which have been well documented. Figure 5.25] 'This is only a speculation which has not been confirmed in any of our studies.Chapter 5. RESULTS AND DISCUSSION^ 105Table 5.14: Extent of zinc extraction after 30 minutes of low pulp density leaching underoxygen pressure; concentrate sample size: 8-10 g; initial conditions: [ZnSO 4 ] = 2 M;[11 2 SO 4] = 0.5 M; [Fe3+]=[Fe 2+] ,---0.15 M; Temp.= 140+1°C; Po, = 1100 kPa; Batchaddition of surfactantsSurfactant Conc. (g/L) Zn extraction (%)None 83+0.7Lignin. 0.1 98+0.4Naphth. 0.1 80+0.3OPD0.05 97+0.50.1 97+0.30.3 98+0.2MPD0.05 99+0.20.1 99+0.20.3 99+0.4PPD0.05 91+0.30.1 98+0.30.3 99+0.2shows that 99 % zinc extraction can be achieved within 20 minutes of leaching in thepresence of the amine additives (under low pulp density conditions 17 ). The plot showsalso the extent of zinc extraction at different leaching times in the presence of naphthalenesulfonic acid—formaldehyde condensates.It should be noted that, at low pulp densities, the oxygen mass transfer rate in theprocess (and subsequent oxidation of ferrous ions to ferric ions by oxygen) is almost neverrate determining, while at high enough pulp density, oxygen transfer may be sufficientlylimited to result in a substantial lowering of the ferric:ferrous ratio.MPD and PPD have been extensively studied, much cannot be said about their performanceunder high pulp density conditions.Chapter 5. RESULTS AND DISCUSSION^ 106High Pulp Density Leaching under Oxygen PressureThe results obtained from the chemical analysis of both the filtrate and the leach residues' 8for the different conditions of leaching are shown in Tables 5.15-5.20. The extent of zincextraction in the presence of OPD was not sensitive to the OPD concentrations usedor the time of leaching (i.e. 35 and 60 minutes). However, the same cannot be saidof lignin sulfonic acid. MPD was also quite effective as a surfactant under the leachingconditions adopted in this work. In contrast, the extent of iron extraction was dependenton surfactant concentration and time of leaching. The extent of iron extraction from theconcentrate during leaching is shown in Table 5.17. It can be seen from the tables thatleaching in the presence of OPD effects higher metal extractions than in the presence ofthe other surfactants used in these tests.Shown in Figure 5.26 is the extent of zinc leaching with respect to time in the presenceof OPD. The figure shows that within about 35 minutes, over 99 % zinc extraction wasachieved in presence of either 0.06 or 0.1 Of, surfactant.During the high pulp density tests, the instantaneous oxygen consumptions rateswere monitored during the course of the leaching period. Figures 5.27-5.30 show theinstantaneous oxygen flow rate for the leaching of sphalerite in the absence and presenceof lignin sulfonic acid, OPD and MPD'. These figures show that in the absence of anysurfactant, the flow of oxygen into the system comes to almost a complete stop afterabout 30-35 minutes of leaching whereas in the presence of a surfactant oxygen is stillflowing at a low rate after the 35-minute mark. This indicates that in the absence of anysurfactant, all reactions requiring oxygen cease almost completely after about 35 minutesof leaching.Figures 5.31-5.31 show the cumulative oxygen consumed (in moles) under the differentThe details of the chemical analysis procedure are given in appendix B.'These were the only surfactants used in the high pulp density leaching testsChapter 5. RESULTS AND DISCUSSION^ 107test conditions. These figures show that the gas mass flow rate is independent of OPDconcentration (between 0.06 and 0.1 g/L) used while the rate is influenced to some extentby the lignin sulfonic acid concentration (also between 0.06 and 0.1 g/L).It can be assumed that the reactions taking place during the leaching period are thefollowing:1) the homogeneous oxidation of ferrous ions present in the stock solution' to the ferricstate:Fe2+ + H+ + 1/40 2^Fe3+ + 1/21120^(5.66)2) the oxidative dissolution of the metal sulphides such as:Zn(S) +2H+ + 1/20 2 —+ Zn 2+ (S) + H2 O^(5.67)Fe(Sx) + 3H+ + 3/40 2^Fe3+ (Sx) 3/2H20^(5.68)Cu(S) + 2H+ + 1/20 2 —+ Cu 2+ (S) + H2 O^(5.69)Pb(S) + 211+ + 1/202^Pb2+ (S) + H2 O^(5.70)and(3) the oxidation of elemental sulfur to sulfateX(S) + x112 0 + 3x/202^2xH+ + xSO4-^(5.71)where x is the number of moles of elemental sulfur oxidized to sulfate. Based on thesereactions, equations can be written for the theoretical acid and oxygen consumptions fora sample size of 100 grams, Table 5.18. Ideally, the theoretical acid to oxygen (consumed)ratio should he equal to the experimentally measured values. Thus, from Table 5.18II+ 2.351 — 2x (5.72)02(,-pt.)^0.5876^1.5x2c111 the 1.5 L solution used for each leaching test, the total ferrous concentration was 0.21 mole.Chapter 5. RESULTS AND DISCUSSION^ 108Thus, using this relationship, the extent of elemental sulfur oxidation through reaction5.71 can be computed for each experiment.Table 5.19 shows the experimental results for oxygen and acid consumption. Table5.20 shows the experimental results for the elemental sulfur analysis of the leach residue,and a comparison between the experimental (using the barium sulfate method for theanalysis of the filtrated ) and calculated extent of elemental sulfur oxidation to the sulfatethrough reaction 5.71 for the measured extent of metal extractions. The calculated datawere computed using the following approach: for each experiment, assuming completemetal dissolution (for a 100—gram sample), equation 5.72 can be used to calculate theextent of sulfur oxidation, x, under the prevailing conditions (given the H+/0 2 ratio).Using this value of x and the elemental sulfur content from the residue analysis, theextent of elemental sulfur oxidation, y, can be obtained for each experiment using thefollowing equation':(5.73)0.986^B ywhere B (moles) is the quantity of elemental sulfur in the residue for a given experimentand is shown in table 5.20 as S°. The experimental data are almost equal to or dif-fer slightly from the computed values. The differences between the calculated (throughreaction 5.71 and equation 5.73) and the experimental values can be attributed to inde-terminate analytical errors.Also shown in the same table is the extent of total sulfide sulfur (in the concentrate)oxidation to sulfate (% S tjiajation) based on the barium sulfate figures. The extent of sulfuroxidation in the presence of OPD is about one and half to two times that observed withlignin sulfonic acid. The extent of elemental sulfur formation due to the metal extractions'The purpose of this filtrate analysis was to determine the total change in the sulfur content of thesolution after the experiment.22 The total 'X sulfur oxidation in terms of x and y should be equal for each set of conditions.0.986 — xChapter 5. RESULTS AND DISCUSSION^ 109is also shown'. It can be seen that the presence of OPD speeds up the formation ofsulfate due to the formation of sulfur particles of very fine sizes', (Figures 5.35-5.38).These finely—sized sulfur particles have a high specific surface area.5.2.2 DiscussionLow Pulp Density LeachingIt has been observed that of all of the surfactants studied, the phenylene diamines andlignin sulfonic acid seem to be the most efficient surfactants (or dispersants) under theconditions of leaching. The amphoterics and naphthalene sulfonic acid—formaldehydecondensates were very poor dispersants.A close look at the interfacial activities of the different surfactants suggests thatnaphthalene sulfonic acid and CAHSB should be very efficient in the leach and OPDthe least efficient. However, this was not the case. The data collected show that thediamines and lignin sulfonic acid were the most efficient. This observation could lead tothe conclusion that there may be other mechanisms operating during the zinc pressureleaching process (involving surfactants) that influence the interfacial phenomena eitherdirectly or indirectly. Thus, the interfacial property of the surfactant may deviate fromthe ideal situation, e.g., a surfactant may be very effective in reducing the liquid sulfur—aqueous solution interfacial tension and increasing the contact angle (a process which,theoretically favours high leaching rates), but the chemical may not be very efficientin enhancing the leaching process. One explanation is that the surfactant in questionmay not be stable under the oxidative pressure leaching conditions. From Table 5.11, it"Ideally, the extent of total sulfide sulfur oxidation to sulfate and the extent of elemental sulfurformation should add up to 100 %; however, due to analytical descrepancies, they differ by 1-2 %.24 This is quite unexpected since it is believed that the inability of OPD to reduce the interfacialtension should result in the coalescence of sulfur and hence large particle sizes.Chapter 5. RESULTS AND DISCUSSION^ 110appears that even with an increase in naphthalene sulfonic acid formaldehyde conden-sates concentration, relatively low zinc extractions were still achieved. This suggests thatthe surfactant probably depletes ferric ions that directly attack the zinc sulfide mineral.This phenomenon is displayed graphically in Figure 5.16 for naphthalene sulfonic acid-formaldehyde condensates under simulated zinc pressure leach conditions. It can be seenfrom the plot that as the residence time increases, the interfacial tension also increases.This suggests that the naphthalene sulfonic acid is going through some form of degra-dation process in the presence of ferrous and ferric ions. Within a residence time of 15minutes', in the presence of 0.05 M Fe 2+, 0.12 M Fe3+, 0.2 M sulfuric acid and 0.3 g/Lnaphthalene sulfonic acid-formaldehyde condensates, the liquid sulfur-aqueous solutioninterfacial tension rises to about 27-2S mN/m from the 20-22 mN/m value recorded inthe absence of the ferric and ferrous ionic ions. This implies that the surface activityof 0.3 g/L surfactant had been reduced to the equivalent of 0.05 g/L in a matter of 15minutes by heating the solution from 25°C to 130°C in the presence of ferrous and ferricions. A series of similar experiments was performed to confirm the above observation. Inthese tests, once the operating temperatures had been reached (after 13-15 minutes ofheating), the system was heated for an additional 2-5 minutes at the same temperaturebefore the pendant drop pictures were taken for the interfacial tension measurements.Additional increments of 4 and 5 minutes of maintaining the temperature resulted ininterfacial tensions of 36-38 and 40-42 mN/m respectively. In the absence of ferric andferrous ions, a surfactant concentration of 0.01 g/L gave an interfacial tension of 42.5mN/m. This implies that the surfactant (of starting concentration 0.3 g/L) had beenreduced to the equivalent of 0.05 g/L in about 15 minutes and shortly after, (about 4-525 11 takes roughly 13-15 minutes to heat the pressure vessel from room temperature to the operatingtemperature of about 130°C.Chapter 5. RESULTS AND DISCUSSION^ 111minutes later), further degraded to the equivalent of 0.01 g/L. At the leaching tempera-ture of 140°C, the degradation rate in the presence of Fe3+ and/or Fe 2+ would probablybe faster and hence render the naphthalene sulfonic acid ineffective as a sulfur dispersantwithin a short period of time. The tendency for a surfactant to degrade under leachingconditions could explain why this surfactant is such a poor dispersant (for liquid sulfur)for the leaching process. Similar observations have been made on lignin sulfonic aciddegradation [55].Dreisinger et al. [54] have reported that lignin sulfonic acid has a half-life of tenminutes under oxidative pressure leaching conditions. This means that after every tenminutes, the capability of the surfactant is degraded to a level equivalent to half its pre-vious strength. Unlike naphthalene sulfonic acid—formaldehyde condensate which seemsto degrade within a short time at high enough temperatures, it seems to take a relativelylonger time for lignin sulfonic acid to be entirely degraded, and so it can still enhancethe leaching rate.All the amphoteric or zwitterionic surfactants performed very poorly in the leachingtests. Though they are all used as dispersants, detergents and wetting agents (like ligninsulphonates) in the laundry detergent industry, these reagents did not fare well in theleaching studies. It is believed that these surfactants undergo similar degradation processunder oxidative leaching conditions, hence their inability to enhance the leaching rates.These three examples indicate that amphoteric surfactants might not do well in leachingunless the chemical stability of these materials can be enhanced.A brief study was made to investigate the stability of OPD during use. In this study,the leach (or pregnant) solution from a previous experiment (12 grams concentrate inlow pulp density conditions) was recycled. The solution already contained 0.1 g/L OPDfrom the previous experiment, and no surfactant was injected with the concentrate. After30 minutes of leaching (amount of concentrate used was about 12 grams), about 98.4 7cChapter 5. RESULTS AND DISCUSSION^112zinc extraction was achieved. This indicates that OPD seems to survive a batch leachingexperiment and so might survive in the full zinc circuit that includes electrowinning.Normally, it is desirable for the surfactants to be degraded before the leach solutionsare passed to the very sensitive zinc electrowinning process. If OPD is to be introducedcommercially, it is important to study further the degradation kinetics of OPD as wellas the impact (if any) of OPD on zinc electrowinning.Orthophenylene diamine was the only surfactant which had no effect on the liquid-liquid interfacial tension, yet managed to enhance the leaching rate to levels comparableto those obtained with lignin sulfonic acid. However, it increased the contact anglefrom 80 to 125-135°. These observations suggest that the most important criteria for asurfactant to be able to disperse liquid sulfur from the mineral surface are:1. its ability to increase the sulfur-mineral contact angle and2. remain kinetically stable under the conditions of zinc leachingwith the liquid sulfur-aqueous interfacial tension reductions being of secondary impor-tance. The relative stability of the surfactant should be such that it does not interferewith the subsequent metal extraction stage. The interfacial tension reduction dispersesthe liquid sulfur into small droplets and the increase in contact angle 'rolls' the liquidinto a spherical shape, which when combined with the continuous agitation of the pulpdetaches the liquid from the solid surface. There remains another explanation for thehigh rates of zinc leaching in the presence of OPD: that the surfactant may form anintermediate product during the leaching process and this intermediate product has theability to decrease the liquid sulfur-aqueous solution interfacial tension and disperse theliquid sulfur as well as to increase the contact angle resulting in an enhanced leachingrat e26.26 This is only a speculation since there is no experimental proof.Chapter 5. RESULTS AND DISCUSSION^ 113High Pulp Density LeachingFrom Table 5.16, OPD was the most efficient surfactant in the tests carried out underhigh pulp density conditions. About 99 % zinc extraction was achieved at a surfactantconcentration of 0.06 g/L within 35 minutes, whereas, with lignin sulfonic acid (which isthe commercially accepted surfactant for the process), a concentration of 0.1 g/L gaveabout 91 and 94 % zinc extractions after 35 and 60 minutes respectively. With MPD as asurfactant, 95 and 98 % zinc extractions were achieved at 35 and 60 minutes respectivelyat a surfactant concentration of 0.1 g/L.Figure 5.26 shows that with OPD as a dispersant for liquid sulfur, it is possible toachieve about 95-96 % zinc extraction in about 25 minutes in the high pulp densityleaching, which is higher than the extent of extraction achieved with lignin sulfonic acideven after 60 minutes of leaching. This is comparable to the 20 minutes the process takesto achieve 99 % zinc dissolution under the low pulp density conditions, Figures 5.25.From Table 5.17 OPD may be expected to be more effective as a surfactant in theleaching of iron sulfide minerals, though the effectiveness may not he the same as thatfor zinc sulfide minerals. Tests conducted with OPD resulted in about 79-80 % ironextraction after 60 minutes as compared to 70 % iron extraction with 0.1 g/1_, ligninsulfonic acid for the same time period. The results obtained as to the extent of ironleaching are not strictly comparable with those that may be obtained with regard to thedissolution of individual iron sulfide minerals'. This argument may he supported bythe observation made by Ngoc et al. [41} in the leaching of complex sulfide concentratecontaining CuFeS 2 , ZnS. PbS, and FeS 2 in which they observed that the leaching ofCuFeS 2 , ZnS and PbS out of the complex ore was as if each individual mineral was beingleached separately. FeS 2 , however, did not follow that pattern.27 Galva.nic effects between minerals may be important.Chapter 5. RESULTS AND DISCUSSION^ 114With OPD the extent of metal extraction was not sensitive to the surfactant concen-tration used. However, the same cannot be said for lignin sulfonic acid. The literatureindicates that lignin sulfonic acid concentrations in excess of 0.3 g/L have no furtherinfluence on the zinc dissolution rate.The extent of zinc and iron extraction (and hence the rate of reaction) for the differentconditions are accurately correlated to the cumulative oxygen plots, Figure 5.34 andTable 5.19; the higher the extent of metal extraction, the higher the cumulative amountof oxygen consumed.The instantaneous oxygen consumption plots show that in the absence of any sur-factant (Figure 5.27), the oxygen flow into the reactor comes to almost a complete haltafter about 35 minutes of leaching. This is believed to be due to liquid sulfur stabilityand to the encapsulation of the sulfide mineral particles. This implies that no furtherreaction (that requires oxygen) occurs beyond 35 minutes. On this basis the total sulfuroxidized to sulfate (Table 5.20) is partly due to pyrite oxidation and/or sulfur oxida-tion by molecular oxygen, (reaction 5.71). It is well accepted that under equal oxidativeacidic conditions of leaching, pyrite is more likely to generate sulfate than sphalerite orpyrrhotite.Figure 5.28 shows the typical behaviour of the system (in terms of instantaneousoxygen consumption) for the tests conducted with lignin sulfonic acid as the surfactantfor dispersing liquid sulfur. The continuing flow of oxygen after 35 minutes shows thatthe zinc and iron dissolution reactions (and any other side reaction) were still in progressand hence the need for the continuous supply of oxygen. Data shown in Tables 5.16,5.17, 5.19 and 5.20 are self—explanatory. The extent of zinc and iron extraction increasesfrom 86 to 93 and 59 to 70 Vc respectively between 35 and 60 minutes of leaching. Thereis an increase in the acid consumption as well as the sulfur oxidation.The continuing flow of oxygen for OPD and MPD leaching (Figures 5.29 and 5.30)Chapter 5. RESULTS AND DISCUSSION^ 115even after achieving 99 % zinc extraction (after 35 minutes) is due to the oxidation ofimpurity minerals such as pyrite taking over from the main zinc dissolution reaction.Table 5.17 shows that the percent iron dissolution increased from about 71 (after 35minutes of leaching) to 81 % (after 60 minutes of leaching with OPD) and 63 to 70% (for MPD leaching). It is also possible that there is some elemental sulfur oxidationoccurring which usually takes place in a low acid medium and at high temperature. Thus,besides the main zinc dissolution reaction taking place in the system, other minor sidereactions due to the presence of other minerals or sulphides contribute to the consumptionof the reactants.The cumulative oxygen consumptions for the different leaching conditions are shownin Figures 5.31-5.34. The plots shown are in agreement with the data, in Tables 5.16and 5.17. As mentioned earlier, approximately 0.0525 mole of these cumulative oxygenconsumptions is used up in the oxidation of 0.21 mole ferrous ions present in the 1.5 Lstarting solution to the ferric state. From Figure 5.34. it can be observed that for leachtimes under ten minutes, the cumulative oxygen consumption is about the same for thedifferent conditions. After the ten minute mark, the curves begin to diverge. It is properto assume that at this point, enough liquid sulfur has been formed in the system suchthat, in the absence of surfactant in the system, the liquid sulfur has taken over thecontrol of the reacting particles surfaces resulting in a slowing down of the reaction.The acid:oxygen ratio shows that the lower the ratio, generally, the higher the sulfuroxidation to sulfuric acid to supplement the acid concentration in the starting solution.If it is assumed that there is no elemental sulfur oxidation taking place during theleaching of the concentrate, the acid:oxygen ratio for the extent of metal (iron and zinc)dissolutions is equivalent to 4. Table 5.21 shows the theoretical oxygen and acid require-ments as well as the acid to oxygen ratios for the extent of metal dissolution observed inChapter 5. RESULTS AND DISCUSSION^ 116the absence of any sulfur oxidation. A comparison of the ratios in Table 5.21 to those ex-perimentally measured, (Table 5.19), suggests about 10-20 % elemental sulfur oxidation.However, due to factors such as analytical discrepancies and the effect of side reactions(which have been ignored) the extent of the determined sulfur oxidation is less.Physical examination of the leach residues (after the filtration) showed that the el-emental sulfur by—product from the leaching experiment with OPD addition formed afilter cake with very fine particle sizes, whereas the residue produced with lignin sulfonicacid addition was of a coarse texture, (Figures 5.35-5.38); the residue produced in theMPD leaching experiment had an intermediate particle size. Fine particles sizes havelarge specific surface areas and high specific rates of reaction, hence the elemental sulfurproduced in the OPD experiment will be more prone to oxidation. This explains thehigher sulfur oxidation observed in the OPD leachings, Table 5.20.Ideally, the inability of OPD to reduce the interfacial tension should enhance coa-lescence of sulfur and produce larger final drop sizes. On the other hand, if the OPDremoves the sulfur from the solid surface faster than does the lignin sulfonic acid, thisshould result in the finer particles. However, then the sulfur particles should coalesce toform larger globules since they cannot he dispersed in the aqueous phase by the surfac-tant. The production of fine particles indicates the presence of other interfacial effects.There is a possibility that because OPD fails to degrade substantially during leaching.this could cause the decrease in the final drop sizes. For example, the lignin sulfonatedegradation products may affect the coalescence of sulfur drops favourably (though wedo not have any physical evidence of this). It is also possible that the interaction betweensurfactants and precipitated solids (or materials) such as jarosite, silica, lead sulfate, etc.during leaching could influence the coalescence of liquid sulfur in the system'.2ftnfortunately. this effect had not been considered in our interfacial studies and this is one area thatneeds to be studied.Chapter 5. RESULTS AND DISCUSSION^ 117Table 5.22 shows the comparison between the theoretical (calculated based on the ex-tent of metal extraction and sulfur oxidation, y„ ic and experimental oxygen requirements.The theoretically calculated oxygen consumptions and the experimentally measured dataare within +2 % of each other. This is a very good agreement. The theoretical oxygenconsumptions were calculated based on reactions 5.66, 5.67, and 5.68 and 5.71, neglectingreactions 5.69 and 5.70 which might occur to a very minor extent. Without consideringany sulfur oxidation in the process (in the runs without any surfactant) the calculatedoxygen requirements for the extent of metal extractions is about 0.288 mole, Table 5.21.When the calculated oxidized sulfur is taken into consideration through reaction 5.71,the total theoretical oxygen requirements for the reactions is about 0.308 mole. Thisvalue compares quite well with 0.306 mole, the experimental value, 5.22. Again, a closelook at Tables 5.19 and 5.21 shows that the theoretical acid (as H+ ions) requirementsfor the extent of reaction (based on reactions 5.66, 5.67, 5.68) are slightly higher thanthe experimentally measured data. The shortfall is caused by the oxidation of sulfur.Shown in Table 5.23 is the total leaching efficiency as well as the Fe:Zn ratios in thepresence of the different surfactants. The total leaching efficiency is given by:Total leaching efficiency,  % = (Zn Fe)leached  x100(712 F el concentrate(5.74)As expected, the highest efficiencies were obtained in the presence of orthophenylenedia.mine and the lowest occurred in the absence of any surfactant. The Fe:Zn ratiosshow the extent of iron extraction for each mole of zinc dissolved in the process. In theconcentrate, the Fe:Zn ratio is 0.272. If these two metals dissolve at the same rate, theratio should be constant throughout.. The lower numbers shown in Table 5.23 suggestthe zinc dissolves faster than the iron. This is not unexpected since pyrite (the impuritymineral and the major source of iron in the concentrate) is known to dissolve at a slowerrate than sphalerite under the same conditions. It is also possible that dissolved ironChapter 5. RESULTS AND DISCUSSION^ 118may be precipitated in the form of jarosite and this will also result in low Fe:Zn ratios.The information gathered from the above study makes OPD a potential supplementor substitute for lignin sulfonic and quebracho in the zinc pressure leaching systems.SummaryThe information gathered so far indicates that in spite of the fact that OPD was notvery effective in terms of interfacial tension reduction as compared to the other surfac-tants, (though it increased the liquid sulfur-mineral contact angle from about 80+5° to135+5° in the presence of ionic species), OPD proved to be a very effective surface activeagent (dispersant for liquid sulfur) in the pressure leaching experiments. Naphthalenesulfonic acid-formaldehyde condensates and CAHSB which were very efficient in termsof interfacial tension reductions and contact angle increases, did not perform well in theleaching experiments, even in the low pulp density runs, due to high degradation ratesunder the leaching conditions. The metal dissolution rates obtained in the presence ofOPD are very comparable to those obtained with lignin sulfonic acid. Thus, it is properto conclude that the most important criterion for a surfactant to be a good dispersant isits ability to increase the solid-liquid sulfur contact angle, provided the said surfactantis stable under the operating conditions. This is achieved by adsorption of the surfactantonto the mineral. The increase in contact angle does not necessarily need to be veryhigh provided the increase is enough to exceed a critical contact angle which is the angleabove which sulfur dispersion can be achieved and below which the liquid sulfur is notdispersed from the mineral surface. Beyond that critical contact angle, the tendency forliquid sulfur to adhere to the solid surface will be very low and thus a high zinc extractioncan be achieved. Met aphenylene diamine was also quite effective in the high pulp densityChapter 5. RESULTS AND DISCUSSION^ 119experiments and needs to be studied further 29 . Despite the fact that both OPD andMPD proved to be very efficient in the pressure leaching experiments, nothing is knownabout their effects on the electrowinning of zinc from the leach solution. This is onearea that needs to be investigated before any commercial application of these surfactantscan be made. Lignin sulfonic acid is known to go through a degradation process by thetime the zinc sulfate leach solution gets to the tank house and hence does not interferewith the electrowinning process. Thus, there is a need to investigate the effect of resid-ual OPD on the electrowinning process. Nothing can be said about the effectiveness ofparaphenylene diamine under the high pulp density conditions since this aspect was notinvestigated. However, it did prove to be very efficient on a low pulp density scale, withvery reproducible results.29 No studies were performed to determine NIPD's interfacial activity at the liquid sulfur-aqueousphase-mineral interfaces, but it is believed it might fall into the same category as OPD since the -1\11 2groups attached are activating groups and hence tend to activate the benzene ring, though the activationeffect is more pronounced in the ortho positions.Chapter 5. RESULTS AND DISCUSSION^ 120Table 5.15: Weight loss due to leaching in the absence and presence of different surfactantsunder high pulp density conditions; initial conditions: [ZnSO 4] = 1.27 M; [11 2 SO 4 ] = 1.12M; [Fe21.0.14 M; Temp.= 140+1°C; P02 = 1100 kPa; Batch addition of surfactantsSurfactant Conc. wt.(g/L) Time(min) Weight loss (%)None 60 31+0.3Lignin.0.0635 53+0.260 58+0.40.1035 57+0.460 59+0.3OPD0.0635 61+0.560 63+0.60.1035 61+0.560 63+0.3MPD0.1035 58+0.360 61+0.5Table 5.16: Extent of zinc extractions (computed from residue analysis) and the changesin Zn 2+ (aq) concentration of the filtrate in the presence of different surfactants andunder high pulp density conditions; initial conditions: [ZnSO 4 ] = 1.27 M; [H 2 SO 4 ] = 1.12M; [Fe2+]=0.14 M; Temp.= 140+1°C; P 02 = 1100 kPa; batch addition of surfactants. *Residue analysisSurfactant Conc.(g/L)Time(min)Zn leached*(moles)AZn filtrate(moles)% ZnextractionNone 60 0.353 0.335 49±0.5Lignin.0.0635 0.616 0.599 86+0.560 0.675 0.654 92+0.30.1035 0.653 0.637 91+0.660 0.688 0.672 93+0.2OPD0.0635 0.733 0.730 99±0.560 0.741 0.735 99+0.20.1035 0.740 0.703 99+0.260 0.743 0.740 99+0.4MPD 0.1035 0.707 0.679 95±0.360 0.731 0.732 98±0.4Chapter 5. RESULTS AND DISCUSSION^ 121Table 5.17: Iron extractions (based on residue analysis) in the presence of differentsurfactants and under high pulp leaching conditions; initial conditions: [ZnSO 4] = 1.27M; [H 2 SO4] = 1.12 M; [Fe 2+]=0.14 M; Temp.= 140±1°C; P0 2 = 1100 kPa.Surf. Conc.(g/L) Time(min) Fe extraction (%)None 60 40+0.2Lignin.0.0635 59+0.360 64+0.30.1035 59+0.260 69+0.3OPD0.0635 70+0.760 78+0.40.1035 72+0.560 79+0.5MPD 0.1035 62±0.560 69+0.5Table 5.18: Oxygen and acid (as H+ ions) requirements per 100 gram sphalerite con-centrate and 1.5 L leach solution assuming complete oxidation of concentrate after testperiod. "x" represents the fractional conversion of sulfur to sulfateElement Source Conc'n. (moles) 0 2 (moles) fl+ (moles)Fe (as Fe2 +) 1.5 L sol'n.  0.21 0.0525 0.21Zn concentrate 0.751^_0.3755 1.503Fe ,, 0.204 0.153 0.612Cu 0.01 0.005 0.020Pb , 0.003 0.0015 0.006S 0.986 1.5x -2xTotal 2.164 0.5876 + 1.5x 2.351 -2xChapter 5. RESULTS AND DISCUSSION^ 122Table 5.19: Experimental data for oxygen and acid (measured as H+) consumptions inthe presence of different surfactants; initial conditions: [ZnSO 4 ] = 1.27 M; [H 2 SO4] =1.12 M; [Fe 2+]=0.14 M; 140±1°C; Po, = 1100 kPa.Surf. Conc.(g/L) Time(min) 0 2 (Moles) H+ (Moles) H + /02None 60 0.306±0.003 1.10±0.01 3.61Lignin.0.0635 0.480±0.005 1.708±0.03 3.5660 0.517±0.003 1.867±0.01 3.620.1035 0.503±0.003 1.832±0.02 3.6460 0.557+0.005 1.898±0.03 3.41OPD0.0635 0.571+0.004 2.000±0.03 3.5060 0.620±0.007 2.020±0.02 3.260.1035 0.574+0.006 2.030±0.03 3.5460 0.619±0.003 2.040±0.02 3.30MPD0.1035 0.538±0.005 1.960±0.02 3.6460 0.577±0.005 2.030±0.01 3.52Table 5.20: Elemental sulfur analysis of leach residue and filtrate and the extent of totalsulfide sulfur oxidation to sulfate (as %); initial conditions: [ZnSO 4] = 1.27 M; [H 2 SO 4 ]1.12 M; [Fe 2+]=0.14 M; Temp.= 140±1°C; Po, = 1100 kPa.Surf.Conc.(g/L)Time(min)S!7 esiclue(Moles)Y calc(Moles)AS f iltrate(Moles)% S tjiadiationSre' ,du,S, .^-1-,A;T e ° due^ftltrate%None 60 0.381±0.008 0.013 0.013±0.001 1.4 97.0Lignin.0.0635 0.563+0.01 0.021 0.022±0.001 2.3 96.060 0.722+0.008 0.024 0.034±0.001 3.6 96.0Lignin0.1035 0.658±0.01 0.020 0.02.5+0.002 2.3 97.060 0.755±0.01 0.041 0.037±0.001 3.8 95.0OPD0.0635 0.736±0.01 0.032 0.056±0.002 5.7 93.060 0.786±0.02 0.054 0.064±0.002 6.6 93.0OPD0.1035 0.769±0.01 0.030 0.051±0.001 5.2 94.060 0.807±0.01 0.052 0.064±0.002 6.5 93.0MPDI^0.1035 0.690+0.008 0.020 0.034±0.001 3.5 95.060 0.735+0.006 0.030 0.047+0.001 4.8 94.0Chapter 5. RESULTS AND DISCUSSION^ 123Table 5.21: Calculated oxygen and acid (as H+) requirements for the different experi-mental conditions, assuming NO elemental sulfur oxidationSurfactant Conc.(g/L) Time(min) O2(Moles) H+ (Moles) H + /02None 60 0.288 1.153 4.00Lignin.0.06 35 0.447 1.788 4.0060 0.486 1.944 4.000.10 35 0.475 1.900 4.0060 0.501 2.002 4.00OPD0.06 35 0.526 2.102 4.0060 0.541 2.164 4.000.10 35 0.532 2.130 4.0060 0.544 2.176 4.00MPD0.10 35 0.501 2.004 4.0060 0.523 2.092 4.00Table 5.22: Comparison between experimental and calculated oxygen requirements forthe different experimental conditions; Temp.= 140+1°C; P o, = 1100kPaSurfactant Conc.(g/L) Time(min) 02a l c . (Moles) Or t • (Moles)None 60 0.308 0.306Lignin.0.06 35 0.478 0.48060 0.522 0.5170.1035 0.498 0.50360 0.562 0.557OPD0.0635 0.574 0.57160 0.622 0.6200.10 35 0.572 0.57460 0.622 0.619MPD0.10 35 0.5:31 0.53860 0.568 0.577Chapter 5. RESULTS AND DISCUSSION^ 124Table 5.23: Total leaching efficiencies and iron:zinc ratios; initial conditions: [ZnSO 4 ] =1.27 M; [H 2 SO 4 ] = 1.12 M; [F0+]=0.14 M; Temp.= 140+1°C; Po, = 1100 kPa.Surfactant Conc.(g/L)Time(min)Fe(Moles)Zn(Moles)Fe:Zn % Total leachingefficiencyNone 60 0.079 0.353 0.224 47.7Lignin.0.06 35 0.115 0.616 0.187 80.460 0.128 0.675 0.189 86.50.10 35 0.122 0.667 0.182 85.460 0.139 0.688 0.202 88.5OPD0.06 35 0.142 0.733 0.194 92.860 0.157 0.741 0.212 95.30.10 35 0.147 0.740 0.198 93.460 0.160 0.743 0.215 95.6MPD0.10 35 0.127 0.707 0.179 88.360 0.140 0.731 0.191 92.210080604020Oo^10^15^20^25Leaching time (minutes)30 30 4010^20Leaching time (minutes)No surfactant0.05 g/L OPD0.1 g/L OPD,0.1 g/L Naphth.0.06 g/L OPDg/L OPD10080604020Chapter 5. RESULTS AND DISCUSSION^ 125Figure 5.25: Extent of zinc extraction in the presence of OPD, MPD, PPD and naphtha-lene sulfonic acid-formaldehyde condensates; Temp.= 140±1°C; P0 2 = 1100 kPa; sampleweight --, 12 grams.10^20^30^40^50^60^70Leaching time (minutes)Figure 5.26: Extent of zinc extraction under high pulp density leaching conditions in thepresence of OPD; Temp.= 140±1°C; P 02 = 1100 kPa.0.06 g/LChapter 5. RESULTS AND DISCUSSION^ 12620E 15Ea)cvCC 100a)x0cr) 5"PI 0 0^10^20 30^40^50 60Time (min)Figure 5.27: Instantaneous oxygen flow rate in the absence of any surfactant; Temp.=140±1°C; Po, = 1100 kPa.20U)E ' 5a)CC100a)Oria604/.1 g/L10 20^30^40Time (min)Figure 5.28: Instantaneous oxygen flow rate in the presence of lignin sulfonic acid:Temp.= 140±1°C: Po, = 1100 kPa.50^60Chapter 5. RESULTS AND DISCUSSION^ 12715a)CC100a)5O10^20^30^40^50^60Time (min)Figure 5.29: Instantaneous oxygen flow rate in the presence of OPD; Temp.= 140+1°C;Po, = 1100 kPa.2015CC1000) 5O10^20^30^40^50^60Time (min)Figure 5.30: Instantaneous oxygen flow rate in the presence of 0.1 g/L MPD; Temp.=110-±1°C; P0 2 = 1100 kPa.No Surfactant—0.1 g/L lignin0.06 g/L lignin.g/L lignin.g/L lignin.70_It_ 0.06 g/L0.06 g/L0.1 g/L0.1 g/L_A_ 0.1 g/L10^20^30^40^50^60Time (minutes)Chapter 5. RESULTS AND DISCUSSION^ 1280.70-(7); 0.6000.50ca)0.400• 0.30•4=co 0.20E0.10010^20^30^40^50Time (minutes)60^70Figure 5.31: Cumulative oxygen consumed in the absence and presence of lignin sulfonicacid; Temp.= 140±1°C; Po, = 1100 kPa.0.800.70E• 0.60a9 0.50rn>ft.x• 0.40a)• 0.30E 0.20Eo.i oFigure 5.32: Cumulative oxygen consumed in the presence of OPD; Temp.= 140±1°C;Po, = 1100 kPa._,_No surfactant_A_0.1 g/L OPD_•_0.1 g/L Lignin.g/L MPDChapter 5. RESULTS AND DISCUSSION^ 1290.70'Z- 0.60 -a)0E 0.500)0.40 -a) 0.30 ->ca-5 0.20 -E0.1000 0^10^20^30^40Time (minutes)- ,,,_ 0.1 g/L- 0.1 g/L50^60^70Figure 5.33: Cumulative oxygen consumed in the presence of MPD; Temp.= 140+1°C;Po, = 1100 kPa.0.80— 0.70a)E° 0.60as 0.50a)0 0.40a)*icTi 0.30E 0.200 0.100^10^20^30^40Time (minutes)50^60^70Figure 5.34: Relative consumptions of oxygen in the presence of the different surfactants;Temp.= 140±1°C; P02 = 1100 kPa.Chapter 5. RESULTS AND DISCUSSION^ 130Figure 5.35: A photograph of zinc sulfide leach residue in the absence of a surfactantafter 60 minutes; Temp.= 140±1°C; Po, = 1100 kPa.Figure 5.36: A photograph of zinc sulfide leach residue in the presence of 0.1 g/L ligninsulfonic acid after 60 minutes;Temp.= 140±1°C; Po, = 1100 kPa.Chapter 5. RESULTS AND DISCUSSION^ 131Figure 5.37: A photograph of zinc sulfide leach residue in the presence of 0.1 g/L OPDafter 60 minutes; Temp.= 140±1°C; Po, = 1100 kPa.Figure 5.38: A photograph of zinc sulfide leach residue in the presence of 0.1 g/L MPDafter 60 minutes; Temp.= 140±1°C; Po, = 1100 kPa.Chapter 5. RESULTS AND DISCUSSION^ 1325.3 INFRARED SPECTROSCOPY5.3.1 ResultsCharacterization of SphaleritePrior to any adsorption studies, the infrared spectral data of the sphalerite concentratewas recorded using:• KBr pellet transmission technique (on dry sphalerite)• ATR analysis (of wet concentrate)Theoretically, natural sphalerite is transparent in the region 2500-800 cm -1 of aninfrared spectrum [113]. This implies that any absorption in this region must be dueto either impurities in the concentrate or to surface oxidation products such as Zn(II)oxy-sulfur products.All of the accompanying spectra records are reported in absorbance'.Figure 5.39 shows the FTIR spectrum of natural sphalerite using the conventionalKBr pellet transmission technique. The band at 1619-1617 cm -1 is most likely due tothe bending vibrations of water absorbed by the KBr and/or water of crystallizationof the concentrate. The absorption bands at 1381 to 1024 cm' may be due to oxysulfur oxidation products formed on the surface of the mineral particles due to contactwith the atmosphere. The broad peak frequency at 3500-3300 cm -1 is associated withthe symmetric and asymmetric stretching vibrations of physically adsorbed water by theconcentrate and possibly going through intermolecular hydrogen bonding.Figure 5.40 shows the ATR-FTIR spectrum of natural sphalerite conditioned in dis-tilled water for 20 minutes with the solvent serving as the background spectrum. This'It should also be noted that the absorption of infrared radiation by the solvent is not quantitativelycompensated so spectral subtraction may be over compensated in some cases.Chapter 5. RESULTS AND DISCUSSION^ 133means that any absorbance due to the solvent in the suspension is cancelled out. The pHof the water was adjusted to that of the working solution (pH = 1.37) to ensure similaraqueous pH conditions throughout the study. The sharp absorption band at 1103-1102cm' and the weak shoulder at 1024 cm -1 are due to the surface oxidation productswhich appeared at 1118 cm' (Figure 5.39). The shift from 1118 to 1103 cm' may beattributed to the solvent-concentrate interactions. The broad and weak intensity bandat 1658-1620 cm-1 is due to the bending vibrations of water of hydration. The broadand weak absorption bands at 3490-2860 cm' may be designated to the asymmetric andsymmetric stretching vibrations of the OH group of the water of crystallization and/orhydration.Adsorption Studies in the Presence of Lignin Sulfonic AcidListed below are the different conditions employed in the infrared spectroscopic studies• KBr wafer spectral analysis of lignin sulfonic acid• ATR spectral recording of 5 g/L surfactant (lignin sulfonic acid) and ferric solution• ATR spectral recording of sphalerite concentrate conditioned in surfactant/ferricsolution• ATR spectral analysis of washed sphalerite concentrate after conditioning in sur-factant/ferric solution• ATR infrared analysis of sphalerite concentrate conditioned in surfactant solutionin the absence of ferric ions• ATR recording of washed sphalerite concentrate after contact with surfactant so-lution in the absence of ferric ionsChapter 5. RESULTS AND DISCUSSION^ 134• KBr disc infrared recording of leached sphalerite concentrate residueKBr Pellet Transmission Spectrum of Lignin Sulfonic AcidThe structure of lignin and hence lignin sulfonate/lignin sulfonic acid has not been welldefined by wood chemists. Different chemists have put forward different structures forlignin sulfonate/lignin sulfonic acid. Figure 2.3 is one such structure [86].Figure 5.41 shows the KBr pellet transmission spectrum of lignin sulfonic acid. Thebroad and strong peak frequency at 3402-3388 cm -1 is designated to the asymmetric andsymmetric stretching vibrations of the OH functional group which are going through in-termolecular hydrogen bonding. Water absorbed by the KBr window may also contributeto this band. The asymmetric stretching vibrations of the -CH 2- group are representedby the absorption band at 2937 cm' whilst the shoulder band at 2842 cm" indicates thestretching vibrations of the -OCH3 group in the molecular structure [123, 124, 170, 171].The most important absorption bands or diagnostic peak frequencies occur in the region1773-1039 cm'. The bands at 1773-1716 cm -1 are representative of the stretching vibra-tions of the carbonyl (C=0) group. The band progression between 1653 and 1423 cm'are assigned to the skeletal vibrations of the molecule - the stretching C=C vibrationsof the aromatic ring being responsible for the 1653-1463 cm -1 whilst C-H deformationsaccount for the 1424-1423 cm -1 vibrational band. The absorption bands at 1370-1369and 1266-1265 cm -1 are assignable to the phenolic OH deformation (or bending) vibra-tions (of the structure). The peak frequencies at 1211-1160 cm' are the result of theasymmetric stretching vibrations of the sulfonate (-S0 3 H/S0) group; the same groupis responsible for the symmetric stretching vibrations at 1040-1039 cm", and the peakfrequency at 650-649 cm" [123, 124], [165]-[171]. The absorption band at 527-526 cm -1is characteristic of the type of lignin sulfonate (165]-[169] - in this case a hardwood type,and it is of no diagnostic importance. The absorption bands at 864-724 cm' are dueChapter 5. RESULTS AND DISCUSSION^ 135to the aromatic ring substitution pattern; though they are not of diagnostic importancetheir presence in a spectrum will help to confirm the presence of the reagent.The KBr disc transmission spectrum for a zinc-chelated (or complexed) lignin sulfonicacid obtained through the courtesy of Georgia Pacific, Bellingham, WA. is shown inFigure 5.42. Basically, this spectrum has a lot in common with that for the lignin sulfonicacid. In spite of the similarities, there are some differences. There is the disappearanceof the absorption bands due to the carbonyl group, and there are shifts in some of thebands compared to the spectrum for the lignin sulfonic acid.The KBr infrared spectrum of sodium-based lignin sulfonate complex is shown inFigure 5.43.ATR Technique for Surfactant SolutionsFigure 5.44 shows the ATR-FTIR absorption spectrum of 5 g/L lignin sulfonic acidsolution 31 . The aqueous solution contained 0.036 mol/L sulfuric acid (pH=1.37), and0.025 mol/L ferric ions (added as ferric sulfate). Compared to the KBr wafer spectrum,it is observed that some of the absorption bands have shifted in positions to either lower orhigher frequencies. Others have disappeared from the solution spectrum as the result ofsolvent-surfactant molecular interactions'. The absorption bands at 1644, 1543, 1514-1513 (not shown on plot due to crowding) and 1464-1463 cm -1 are the result of stretchingvibrations of the aromatic rings. These bands are well-defined with the 1514-1463 cm -1bands appearing in the same wavenumber region as in the KBr spectrum. The band shiftfrom 1653/1605 to 1644 is most likely due to the solvent-surfactant interactions; however,the bending vibrations of water molecules could contribute to this band. The weakintensity shoulder band at 1739-1738 cm -1 may represent the stretching vibrations of the31 Preliminary tests proved that lower surfactant concentrations of 1-1.5 g/L were not sufficient toachieve sensitivity in the recorded range.'The background spectrum is taken to be that of the solvent which is distilled water and acid.Chapter 5. RESULTS AND DISCUSSION^ 136carbonyl, ez---0, group which is capable of hydrogen bonding. The shoulder at 1372-1371cm' represents the OH bending vibrations. The asymmetric stretching vibrations of thesulfonate group are designated by the absorption band at 1194-1193 cm" (a significantshift from 1211 cm' in the solid) while the symmetric stretching vibrational modes arerepresented by the weak shoulder at 1050-10,19 cm'. The broad absorption band atwavenumbers 3500-3100 cm' is the result of the asymmetric and symmetric stretchingvibrations of the OH groups involved in hydrogen bonding. The strong absorption peakat 1106 cm' is due to the presence of a sulfate radical or ions in the solution whichoriginate from the ferric sulfate.Characterization of Surfactant AdsorptionDue to strong infrared absorption by natural sphalerite at wavenumbers lower than 800cm -1 , the spectral analysis of surfactants or organic reagents in this region (in the pres-ence of the concentrate) is less useful. Shown in Figure 5.45(a) are the spectral dataof natural sphalerite conditioned in the ferric/lignin sulfonate solution for 20 minuteswhere the background spectrum (sometimes referred to as reference) is the solvent.The spectrum shown in Figure 5.45(b) is the result of the spectral subtraction of theferric/surfactant solution spectrum from the sphalerite-surfactant (solution) spectrum.Both spectra display most of the characteristic peak frequencies of the surfactant solutionspectrum. The spectrum of the adsorbate on the adsorbent agrees very well with thatof the adsorbate solution from the high frequency end of the spectrum to about 1371cm' (in both Figures (a) and (b)), beyond which there are some new developments, in-dicating the presence of new surface products due to the solid-liquid interactions. Thesespectra are very useful especially in the fingerprint region, where most of the diagnosticabsorptions bands are located.Chapter 5. RESULTS AND DISCUSSION^ 1:37r figure 5.45(b): Region 3600 - 1265 cm- 1The absorption bands at 3500-3100 cm -1 correspond to the asymmetric and symmetricstretching vibrations of the OH groups and the band progression from 1372-1265 cm - 'are also designated to the OH bending vibrations. There are changes in the relativeintensities of the peaks as well. The vibrational bands at 1647-1265 cm -1 are enough toconfirm the presence of the surfactant (an aromatic compound) on the mineral surface[123, 170, 171].Region 1265 - 650 cm - 1There is the appearance of a sharp absorption band at 1098-1097 cm-1 . This may beclue to the formation of a soluble zinc sulfate product on the mineral surface. This wasnot confirmed experimentally. The asymmetric stretching vibrations of the sulfonategroup at 1193-1192 cm' in the bulk solution spectrum has disappeared, and there isa significant shift in peak frequency of the symmetric vibrational band at 1050 (in thebulk solution) to 1013 cm'. These observations are indications of a chemical adsorptionmechanism and the formation of a new surface compound. The bands between 870  -750 cm -1 (in the bulk solution spectrum, Fig.5.44) which are due to the substitutionpattern on the aromatic rings (though diagnostically are less useful) have all changedtheir frequency positions following adsorption of the surfactant. These changes indicatethat the lignin sulfonic acid is chemisorbed to some extent. Physical adsorption mayco-exist with chemisorption.Figure 5.46 is the ATR-FTIR absorption spectrum of the natural sphalerite con-ditioned in the ferric/surfactant solution and then washed 3-4 times with the solventfollowed by more drastic washing in boiling distilled water for 25-30 minutes and coldwashed again 3-4 times. The absorption bands from the high frequency end to 1266cm' still persisted after washing, though there were some minor shifts in some of theChapter 5. RESULTS AND DISCUSSION^ 138nand positions. The bands at 835 to 734 cm- 1 disappeared with new ones appearing at826 and 774 cm -1 . There was the formation of new bands at 1148- 1147.1035-1031, and988 cm' as well. These new developments are indicative of the fact that the layer ofadsorbate on the surface is a chemisorbed product and that it could not be removed bythe washing procedure. The shoulder at 1148-1147 cm" is indicative of the presence ofZn-complexed compound on the surface of the mineral particles ( as can be seen in thespectrum for the zinc-complexed lignin sulfonate). The weak bands at 2921 cm -1 whichrepresent the stretching vibrations of -CH 2- also indicate the presence of the surfactant.The band at 1640-1639 may be a contribution from both co-adsorbed water and thesurfactant molecule.Adsorption Tests in Ferric-free Surfactant SolutionFigure 5.47(a) shows the spectral record of a natural sphalerite conditioned in the ferric-free 5 g/L surfactant solution for 20 minutes (background spectrum being the solvent).The spectrum in Figure 5.47(b) is the result of the appropriate spectral subtraction (i.e.adsorbent-adsorbate spectrum minus adsorbate spectra). Like the ferric solution, thesulfonate absorption bands have disappeared from both spectra.In Figure 5.47(a) the vibrational modes indicative of the presence of the aromaticcompound [170, 171] form the series of weak bands at 3056-2930 cm' and 1550-1460cm-1 in the fingerprint region. In Figure 5.47(b) the diagnostic bands present in theregion 1741-1100 cm' coupled with the band at 2928 cm -1 ( which is due to asymmetric-CH2- stretching vibrations) are indicative of the presence of the adsorbate moleculeson the surface of the mineral particles. The band positions in the fingerprint region ofthe spectrum are very much the same as those obtained in the ferric-surfactant system,though there are some differences between the two in other regions of the spectra.Chapter 5. RESULTS AND DISCUSSION^ 139Figure 5.48 shows the spectral profile of a natural sphalerite conditioned in the ferric-free surfactant solution, washed several times with acidified distilled water followed byfurther washing in boiling water for 25-30 minutes and cold washed again. With theexception of the vibrations at 1608-1607 cm -1 , the absorption hands in the fingerprintregion 1608-1032 cm -1 and 776-775 cm -1 are in just about the same positions as thoseshown in Figure 5.46 (within the limits of instrument sensitivity) with minor frequencyshifts. The vibrations at 1608-1509 cm -1 are due to the skeletal C=C stretching vibra-tions of the aromatic rings. The weak shoulder band at 1267-1266 cm' is due to thephenolic OH group. The absorption hands at 1143 and 1031 cm -1 may be attributed tothe formation of a zinc-complexed surface compound. These observations together withthose in the ferric system imply that the lignin sulfonate molecule is adsorbed both inthe presence and absence of ferric ions in the leaching systems.The decreases in the adsorption intensities or heights after the washing process stagesindicate the removal of a physically adsorbed surfactant layer'. Once this layer is re-moved by the washing, the chemisorbed layer is exposed and this is hard to remove.This observation implies that both physical and chemical adsorption processes takeplace during the adsorption, with physical adsorption probably occurring after the firstchemisorbed monolayer is formed.DiscussionIdeally, shifts in frequency or changes in relative intensities of absorption indicate chemisorp-tion. Physically adsorbed species and pure reagent species are not expected to displayany spectral differences. However, there are occasions when chemically adsorbed speciesdo not show much spectral difference. The intensity of an absorption band is very muchdependent on the magnitude of t he dipole change during the vibrations, the larger the33The physically adsorbed layer can he Inuit ilayers.Chapter 5. RESULTS AND DISCUSSION^ 110change, the stronger the band.Certain functional groups may not always give rise to absorption bands. even thoughthey may be present in the sample.Washing of the mineral particles/adsorbed species did not bring about significantchanges in the positions of the absorption frequencies but showed a decrease in theintensities or height of the peaks (Figures 5.45(b) and 5.46). However, in the ferric-freesystem (Figures 5.47(b) and 5.48) some bands in the fingerprint region disappeared afterthe washing. This may be due to:• the signals at those vibrational frequencies are so weak that the instrument couldnot respond to them34• and/or the absence of ferric ions in the system affected the adsorption pattern tosome extent.This behaviour persisted for different samples and repeated washings of some of them.Because of the drastic washing process adopted, it appears therefore that the species thatremain on the surface after the washing/boiling can be attributed to chemical adsorption.It is reasonable to argue that the surface adsorbed species/layer after the washing processis due to chemical interaction, since any physically adsorbed layer is supposed to bedissolved or removed during the washing/boiling. The absorption bands at 1148-1147and 1143-1142, 1034-1031, and 776-773 cm' can therefore be assigned to the presenceof a zinc complexed lignin sulfonate, (Figures 5.42, 5.46 and 5.48).Some of the peak frequencies obtained in the unwashed and washed samples occur inalmost the same positions. However, the bands at 1148-1143, 1034-1031, 988, and 776-773 cm -1 are new (in the washed spectra). The appearance of these new bands indicatesthat underneath the top physically adsorbed layer lies a surface product(s) which give(s34 Some signals will not show up in the spectrum if they are very weak.Chapter 5. RESULTS AND DISCUSSION^ 111rise to these new vibrations. The unwashed spectra consist of chemically and physicallyadsorbed species but the fraction due to physical adsorption is removed by the washingprocess. It can thus be concluded that the peaks in the region 1741-1265 cm-1 (Figures5.45(a) and 5.47(a)) represent both lignin sulfonate adsorbed in bulk form (i.e. physicaladsorption) and chemisorbed zinc-complexed form which may differ in structure from thebulk compound. The disappearance of the strong sharp vibrational band at 1101-1097cm -1 indicates the removal of soluble surface compounds such as oxysulphur productsfrom the solid surface by the washing process.A final confirmation of this spectral interpretation is obtained from the KBr disctransmission spectrum of sphalerite leached for 15 minutes in the presence of lignin sul-fonic acid. The lignosol concentration was 5 g/L in solution and was added continuouslyat a rate of 2.0 mL/min. The amount of concentrate used was 2 grams. The leachingsolution contained approximately 2 mol/L zinc ions, 0.4 mol/L sulfuric acid, 0.1 mol/Leach of ferric and ferrous ions. The leaching was done under a nitrogen atmosphere.The leached residue was thoroughly washed several times with about 8 litres of distilledwater, vacuum dried and a KBr disc made from the dried sample. No organic solventcould be used for further washing of the residue because lignin sulfonic acid is insolublein organic liquids 35 . The residue disc was then mounted in the path of the IR beamto record the spectral profile (Figure 5.49). Compared to Figure 5.46, it exhibits mostof the characteristic peak frequencies of the ATR-FTIR spectrum. The appearance ofbands at 1460-1459, 1509-1508, 1265-1264, 1149-1148 and 1032-1031 cm -1 agrees verywell with the ATR spectra, indicating the presence of (a) similar compound(s) in bothspectra (as well as Figure 5.48). The absorption band at 1149-1148 cm -1 present in bothspectra is characteristic of a zinc-complexed compound. The differences which are shown"This was confirmed in the lab as attempts to dissolve 0.25 grams lignin sulfonate in 100 ml acetonefailed.Chapter 5. RESULTS AND DISCUSSION^ 112ay be attributed to the different sample preparation procedures and the environmen-tal conditions (i.e KBr - dry and ATR - wet); otherwise the two are very similar. Thevibrational modes due to the skeletal rings of the surfactant molecule are well-definedand occur in the same frequency range in the two spectra, the only exceptions being thebands at 1603-1602 cm -1 (Figure 5.49), and 1641-1640 and 988 cm -1 (in Figure 5.46).There exist, however, significant shifts from 1716 to 1687 cm'. The appearance of anabsorption band at 1418-1417 cm -1 (Figure 5.49) compared to 1424-1423 cm' ( figure5.41) further confirms the presence of the adsorbate molecules on the solid surface. Thisobservation further confirms the chemisorption argument.The large size of the lignin sulfonate structure could create the condition for eachreactive site on the solid surface to be occupied by only one lignin sulfonate ion. Byinteracting with the aqueous phase, electrical neutrality is maintained on the surface.The adsorbent-solvent interactions will result in an adsorbate structure which will bedifferent from the bulk material and hence can be expected to show some differences inthe spectral profile. This explains the minor differences between Figures 5.46 and 5.49on one hand and Figure 5.42 on the other hand, e.g. the bands at 1033 and 1043 cm -1respectively. The physically adsorbed lignin sulfonate will not be affected by the sulfidemineral surface and hence will exhibit the same spectrum as the bulk.It is likely that the adsorbate is adsorbed on the ZnS mineral as a bulk-like Zn(II)-surfactant compound. The surface compound chemically formed has an organic portionwhich is very similar to the bulk adsorbate.The shifts in the absorption bands of the OH stretching vibrations (in the upper fre-quency region) and the asymmetric and symmetric stretching vibrations of the sulfonategroups from 1161-1160 to 1148 and 1143, and 1010-1039 to 1034 and 1031 cm" (figures5.41, 5.46, 5.48 and 5.49) indicate chemical interactions of the OH and SO3 groups withthe solid surface.Chapter 5. RESULTS AND DISCUSSION^ 143It is believed that under the washing conditions adopted, no lignin sulfonate (or phys-ically adsorbed species) could survive and hence the spectrum observed must thereforebe due to chemisorbed species.Although the absorption bands at 1603-1602, 1220-19 and 1509-1508 cm -1 (Figure5.49) occur in the same frequency range as those of the bulk lignin sulfonic acid, therelative intensities of these bands have changed from more intense 1211 cm -1 in the bulkspectrum to almost equally intense bands in the leach residue.The assignment of the peak frequencies between 1640-1264 cm -1 (in Figures 5.45(b)and 5.47(b)) to both forms of adsorbed species is due to:• after the washing/boiling process, which is presumed to remove any physicallyadsorbed species, most pre-washing peak frequencies remained in almost the samepositions (with the exception of the 1103 cm -1 which disappeared) although theirrelative intensities changed• Zn-, Na-based lignin sulfonate have peaks in this region (just like the bulk ligninsulfonic acid) suggesting that metal-lignin sulphonates complexes also have ab-sorption bands in that region and therefore that a surface metal-lignin sulfonatecomplex could have an absorption band in that region• under the conditions of the washing process, physically adsorbed species could notbe present and hence the layer is due to chemisorption.In the KBr spectrum of the leach residue (Figure 5.49), there are some absorptionbands present in the 3500-3100 cm -1 range. However, these bands have intensities similarto the noise background in that spectral region. Therefore, they cannot be regarded astrue peak frequencies in spite of the fact that they occur in the appropriate wavenumberregion.Chapter 5. RESULTS AND DISCUSSION^ 1_1_tIt is possible that mineral-surfactant reaction products can co-exist with mineralsurface oxidation product on the mineral surface where the mineral-surfactant productmay exist as small islands on the oxidized mineral surface. However, hydrophilicity isenhanced on dissolution of any soluble surface oxidation product of sphalerite followedby adsorption of the surfactant.The elemental sulfur present in the zinc leaching system is hydrophobic. This hy-drophobic character makes it possible to interact with the hydrocarbon portion of thesurfactant molecule, with the hydrophilic groups projecting into the aqueous phase. Thus,much as lignin sulfonic acid interacts with the mineral surface, it also interacts with theliquid sulfur through the hydrocarbon end of the molecular structure. This behaviourtends to lower the liquid sulfur - aqueous phase interfacial tensions or free energies ashad been reported earlier [55].SummaryThe presence of multiple functional groups, OH - , SiOt, and C=O, (all hydrophilic groupsand capable of hydrogen bonding) and the three-dimensional nature of the lignin sul-fonate structure suggests that during the process of surfactant-mineral interactions, oneor more hydrophilic groups interact with the mineral surface whilst the other group(s)interact(s) with the aqueous phase through hydrogen bonding. Thus, the surfactant actsas a bridge between the solid and liquid phases. The disappearance of the sulfonate ab-sorption bands and the shifts in the hydroxyl group frequencies (in the upper frequencyregion) implies that the surfactant molecule interacts with the solid surface most likelythrough them. This role played by the surfactant continuously renders the sphaleritesurface hydrophilic and prevents it from becoming sulphophilic. Since lignin sulfonicacid or its salt does not form any oriented layer on adsorption. it is most likely that themolecule sits on the mineral surface as a net" or a sponge - when adsorbed. FormationChapter 5. RESULTS AND DISCUSSION^ 115_ islands of metal-surfactant species on the oxidized mineral surface is probably enoughto make the mineral particles sufficiently hydrophilic.The complexities involved in the satisfactory analysis of more than one componentor functional group from a given absorption band remain a problem even with the bestresolution. In view of this difficulty it is not easy for one to predict with certainty thenature of the adsorbed species.In conclusion, the above observations indicate the presence of physically and chemi-cally adsorbed species on the mineral surface in an acidic medium. However, it is verydifficult to discriminate between physically and chemically adsorbed molecules from thespectrum both qualitatively and quantitatively.Chapter 5. RESULTS AND DISCUSSION4000^3100^2200^1300^400Wavenumbers (cm-1)Figure 5.39: Infrared spectrum of sphalerite using KBr pellet technique.0^rn3000^ 2000 ^ 1000Wavenumbers (cm--1)Figure 5. , 10: .DTI{ specs ruin of wet sphalorite.U0116Chapter 5. RESULTS AND DISCUSSION^ 1470 NI4000^3100^ 2200^1300^ 400Wavenum.bers (cm-1)Figure 5.41: KBr disc spectrum of lignin sulfonic acid.Chapter 5. RESULTS AND DISCUSSION^ 1-1ScoCO3000^ 2000^ 1000Wavenumbers (cm-1)Figure 5.42: KBr disc spectrum of Zn—based lignin sulfonate complex.Chapter 5. RESULTS AND DISCUSSION^ 1494000^3100^2200^1300^ 400Wavenurnbers (cm-1)Figure 5.43: IKBr pact spectrum of Na- based^sulphonate complex.Chapter 5. RESULTS AND DISCUSSION^ 1503800^3012.5^2225^1437.5^65(1Wavenumbers (cm-1)Figure 5.44: ATR spectrum of lignin sulfonic acid solution; solution condition: [Fe 3+] =0.025 M, [H 2 SO 4 ] = 0.036 M, Temperature = 25°C; reference spectrum: solvent.3800 3012.5^2225^1437.5Wavenumbers (cm-1)(a)650OChapter 5. RESULTS AND DISCUSSION 15101 (SDCNI•-•-)^tn Oaoc,,17^CO3800 3000^2200Wavenumbers (cm-1(b)1400^600Figure 5.45: ATR spectrum of sphalerite conditioned in surfactant solution for 20 min-utes; solution condition: [Fe 3+] = 0.025 M. [H 2 50,4 ] = 0.036 M, Temp. = 25°C: (a)reference: solvent spectrum (b) less solution spectrum.1"-)LC)^.--1"--4.6 ..7- szt _^CV ,MCCCO^CD'CV —•Chapter 5. RESULTS AND DISCUSSION^152C.)CZJ=21-4OCJD3800^3012.5^2225^1437.5^650Wavenumbers (cnn-1)Figure 5.46: ATR spectrum of \vashed splialerite after conditioned in surfactant/ferricsolution, reference spectrum: solvent.OlQ6CnCnCC1Chapter 5. RESULTS AND DISCUSSION3500^2787.5^2075^1362.5^650Wavenumbers (cm-1)(a)O1533500^2787.5^2075Wavenumbers cm-1)1362.5^650Figure 5.47: ATR spectrum of sphalerite conditioned in ferric free—surfactant solution for20 minutes; solution condition: [H 2SO 4 ] = 0.036 M, Temp.= 25°C. (a) reference: solventspectrum (b) less the solution spectrum.Chapter 5. RESULTS AND DISCUSSION^ 154t17CNIOCt) 3500^2787.5^2075^1362.5^650Wavenumbers (cm-1)Figure 5.4S: ATR spectrum of washed sphalerite after contact with ferric free—surfactantsolution, reference spectrum: solvent.Chapter 5. RESULTS AND DISCUSSION^ 1552500^1975^1450^925^400Waver umbers (cm-1)Figure 5.49: KBr pellet spectrum of leached sphalerite (in the presence of surfactant)residue; leaching performed at 140+1°C for 15 minutes under P .hr2 1100kPa; ligninsulfonate solution concentration: 5 g/L at flow rate 2 mL/min.Chapter 5. RESULTS AND DISCUSSION^ 1565.3.2 Adsorption Studies in the Presence of Orthophenylene DiamineListed below are the different studies:• KBr analysis of orthophenylene diamine• ATR recording of 5 g/L surfactant solution• ATR recording of sphalerite concentrate conditioned in surfactant solution (no ferricpre-treatment)• ATR analysis of sphalerite concentrate washed after conditioning in surfactant• ATR analysis of sphalerite concentrate conditioned in surfactant solution after pre-treatment with ferric ions• ATR recording of sphalerite concentrate washed after contact with ferric/surfactantsolutionKBr Pellet Absorption Spectrum of OPDThe KBr-FTIR spectrum of OPD powder is shown in Figure 5.50. The absorptionbands above 3236 cm -1 represent the asymmetric and symmetric stretching vibrationsof N41 bonds due to the -NH 2 functional group. The absorption band at 3036-3035cm' represents the stretching vibrational modes of =C-H bonds of the aromatic ring.The shoulder band at 2922-2921 cm' is also attributed to -C-H asymmetric stretchingvibrations [123, 124, 170, 171]. The characteristic (or diagnostic) bands of the moleculeas a whole are found at 1632 cm -1 and below. The peak frequencies in the fingerprintregion are assigned to the following functional groups [124, 170, 171]:Chapter 5. RESULTS AND DISCUSSION^ 1.57• 1632- 1589 cm', due to the internal deformations (of -N-H bonds) of the -NII 2functional group; stretching vibrations by the C=C bonds in the aromatic skeletonoverlaps with the -NH 2 vibrations at 1590-1.589 cm -1• 1497 - 1461 cm -1 , the result of stretching vibrations of the aromatic ring (or C=Cbonds)• 1333 - 1268 cm -1 are attributed to the C-N stretching vibrations• 1152 - 1029 cm" are assigned to the C-H in-plane deformations of the ring• the 926 -925 cm -1 band is representative of the 1,2 substitution pattern on thering• bands between 828 - 711 cm" are assigned to either out-of-plane C-H ring defor-mations and/or -NH 2 out-of-plane bending vibrations i.e. there is an overlap ofvibrational bands from these two functional groups in this region• 544 - 439 cm -1 bands are designated to C-H out of plane deformationsThe absorption bands at the wavenumbers 2780-1781 cm' are due to overtones andcombinations [123, 124, 126].Characterization of OPD solutionFigure 5.51 shows the ATR-FTIR spectrum of 5 g/L OPD in a solution containing 2mL/L (about 0.036 mol/L) sulfuric acid. Compared to the spectrum obtained fromthe KBr disc, (Figure 5.50) some absorption bands observed in the solid spectrum havedisappeared. Others have also shifted slightly in position. All these changes are clueto surfactant-solvent interactions. However, a few bands occurred at almost the samepositions in both spectra e.g. the stretching vibrational modes at 3036 - 3031 cm -1 . TheChapter 5. RESULTS AND DISCUSSION^ 158stretching absorption bands at 3084 -3035 cm -1 are designated to the =C-H functionalgroup of the aromatic ring. These bands, together with the stretching vibrations at 1593- 1483 cm -1 are representative of the aromatic molecular structure'. The asymmetricand symmetric stretching peak frequencies at 3130 - 3247 cm" are assigned to the -N-Hbonds (of the -NH 2 group) which normally go through hydrogen bonding in the aqueousphase. The absorption frequency at 1294 cm" and the shoulder band at 1194 cm" areindications of the C-N bonds (or groups) while the absorption band at 885 - 827 cm'may be assigned to the -NH 2 deformations. The -C-H bonds in-plane deformations maybe responsible for the vibrations at 975-974 cm -1 [123, 124, 126, 170, 171). The reference(or background) spectrum is that of the solvent'.Results of Adsorption testsFigures 5.52(a) and 5.52(b) show the absorption spectra of sphalerite concentrate sus-pended in the 5 g/L OPD solution for 20 minutes in the absence of any ferric ions orferric solution pre-treatment. In Figure 5.52(a), the reference/background spectrum isthat of the solvent. The appropriate spectral subtraction' is shown in Figure 5.52(b).The band at 3231 - 3230 cm -1 (Figure 5.52(14) is a shift from 3247-46 cm" (in Figure5.51), whilst the stretching band at 3420 - 3419 cm -1 (in Figure 5.52) occurs as a weakintensity band at 3429-3428 cm -1 in the bulk solution spectrum. A detailed comparisonof Figure 5.51 to 5.52(b) shows that there are shifts in some of the absorption frequencies.others have disappeared and there are some new developments as well. The appearance ofthe asymmetric stretching vibrational mode at 2928 - 2927 cm -1 (Fig. 5.52(b)) indicatesthe presence of a -C-H bond related compound on the solid surface. The appearance ofstretching vibration absorption bands at 1601 - 1490 cm', the shoulder band at 1444'The N—H group also contributes to the absorption frequency at 1593-1592 cm-1 .37 The solvent as often used in this text refers to distilled water containing 0.036 mol/L sulfuric acid.38 The bulk solution spectrum is subtracted from that of the mineral conditioned in the solution.Chapter 5. RESULTS AND DISCUSSIONcm -1 together with the weak absorption hand near :3059-3058 cm' are attributed tothe presence of an aromatic compound on the solid surface while the vibrational modesat 1274-1273 cm -1 represent the stretching vibrations of the C-N functional group ofthe compound. The -NH 2 in-plane deformations also contribute to the absorption bandat 1601-1599 cm -1 . The 3059-3058 cm' absorption hand could be a shift from 3036-3035 cm' in the bulk solution spectrum due to perturbations upon solid - surfactantinteractions.Figure 5.53 shows the ATR - FTIR spectral record of the natural sphalerite condi-tioned in OPD solution (without ferric ions), cold-washed several times with the solvent,followed by further washing in boiling water for 20-25 minutes and then cold-washedagain. The background spectrum is that of the solvent, implying that any absorptiondue to the solvent content of the sample paste is cancelled out by the background (orreference) spectrum.The absorption bands at 1388-1386 cm -1 indicate the presence of -C-H bonds. Thepresence of the skeletal aromatic ring i.e. C=C bonding, on the solid surface even afterthe drastic washing process is indicated by the series of stretching vibrational modes at1568-1566 (a weak shoulder), 1539-1538, and 1457-1456 cm -1 and the weak stretchingvibrations near 3030-3029 cm -1 which is designated to the =C-H groups of the molecularstructure. The skeletal ring vibrations at 1568-1456 cm' are absent in Figure 5.52(b)but surfaced after the washing.The -NH2 group deformation vibrational bands at 835-754 cm -1 in the pre-washspectrum disappeared after the washing process, with the appearance of new ones at870-869, 816-815 and 731 cm 1 .The development of the stretching shoulder band at 1300 - 1299 cm -1 representsthe formation of a C-N surface product. These observations point to the fact that theadsorbed layer is firmly held onto t he solid surface such that even the boiling water couldChapter 5. RESULTS AND DISCUSSION^ 100not remove the species.In the higher frequency range, the -N-I-I (of -NH 2 ) stretching vibrations occur inthe same frequency range (within the limits of instrumental sensitivity) in Figures 5.30,5.52, and 5.53, further suggesting the presence of an -NH 2 related compound on the solidsurface. The same can be said for the band at 3238 cm -I .Figures 5.54(a) and (b) show the absorption spectra of ZnS concentrate conditioned in0.025 mol/L ferric solution (containing 0.036 mol/L sulfuric acid) for 15 minutes prior tocontact with the 5 g/L OPD solution. The surfactant solution conditioning time was 20minutes. The conditioning was done at room temperature The sample was cold-washed3 times with the solvent in between the conditioning solutions. The spectral record ofthe washed concentrate after the surfactant conditioning period is shown in Figure 5.55.With the exception of the vibrations near 1098-1095 cm -1 , the rest of the absorptionband positions (in the fingerprint regions) in Figures 5.54(a) and (b) are different fromthose of the bulk solution spectra, Figure 5.51. Between Figures 5.54 and 5.55, there aresome new developments in the after-wash spectrum.DiscussionNo ferric ion pre-treatmentFigure 5.52(a) displays absorption bands at wavenumbers 3268-3267, 3189-3188, and1503-1502 cm -1 which are absent from Figure 5.52(b). On the other hand, Figure 5.52(b)shows vibrations at such frequencies as 1491-1490, and 1444-1443 cm' (in the fingerprintregion) which are also absent from Figure 5.52(a). Between Figures 5.51 and 5.52, it canbe observed that most of the absorption bands in the bulk solution spectrum shift tohigher wavenumbers on surfactant adsorption e.g. 975 shifts to 996-991 cm', 159:icm-1 shifts to 1601-1600 , 1483 cm' shifts to 1502-1491 cm', 1442 cm' shift -Chapter 5. RESULTS AND DISCUSSION^ 161to 1445 cm', 3036 cm' increases to 3060-3059 cm', and 3084 cm' shifts to 3096-3095(not shown on plot) cm -1 . These shifts in absorption frequencies to higher wavenumberssuggest the formation of strong chemical bonds between surfactant molecules and themineral and hence chemically adsorbed species on the mineral surface.The disappearance of the absorption bands at wavenumbers 1294 and 1194 cm -1 (inthe bulk solution spectrum, Figure 5.51) due to C-N vibrations and the development ofa new peak frequency at 1275-1273 cm' (on adsorbate adsorption) further support thechemisorption concept. This development indicates that the C-N group is involved inchemical interactions with the mineral. This is most likely due to the overlapping of thenitrogen lone pair of electron with the metal ion orbitals in the sulfide lattice. Figure5.56 shows the possible molecular structure of the amine related complex that can beformed between the metal sulfide lattice and the surfactant molecule.The appearance of an absorption band at 3364-3363 cm' (due to N-H asymmet-ric and symmetric stretching vibrations) further confirms the presence of the surfactantmolecule on the mineral since this band is one of the prominent vibrations in the solidOPD spectrum (at 3368-3367 cm' in the KBr spectral record, Figure 5.50). The de-crease in absorption frequency is expected (or normal) since a C-N group involved ina chemical interactions with a metal ion will tend to stretch the N-H bond and hencelower the wavenumber or frequency of vibration. Further evidence in support of thisis obtained by making reference to the vibrational bands at 3429-3428 and 3246-3245cm -1 (Figure 5.51) and 3419-3416 and 3240-3230 cm' (Figure 5.52). These bands areall due to the -N-H stretching vibrations. This development can be given a mathematicalinterpretation as well. Consider the equation [121, 122, 123]:1v = 5-7i= (5.75)Chapter 5. RESULTS AND DISCUSSION^ 162or(5.76)K = force constant or bond strengthitm = reduced mass of atomsd = wavenumberv = frequency of vibrationand1^1^1— = —rn^m1^M 2(5.77)m 1 and m 2 are the masses of the two atoms involved in the particular bonding in themolecule. As mentioned earlier, a strong chemical bonding between the nitrogen andmetal ions, Figure 5.56, results in the increase in the -N-H bond length and a lower svalue for that bond. This will cause a decrease in the the frequency of vibration for thatbond.The C-N bond is also stretched or lengthened on bonding resulting in a low # valuefor the C-N bonding leading to a lower v value after the contact period. This explainsthe shift in frequency from 1294 cm' (in the bulk solution spectrum) to 1275-1273 cm'on adsorption.Ideally, Zn 2+ (or other metal) ions will be more reactive than hydrogen towards thenitrogen (with a lone pair of electrons). The formation of hydrogen bonding between the-N-H groups and the aqueous phase (i.e. water molecules) can also contribute to thelowering of the absorption band frequencies.The appearance of the 3144-3143 cm" (Figure 5.52) absorption band may be at-tributed to the formation of a new amine related chemisorbed species.The shift in the =C-H group wavenumbers from 3036 and 3084 (in Figures 5.50and 5.51) to 3060-58 and 3096-3095 cm -1 (not shown on plot) respectively (in FiguresChapter 5. RESULTS AND DISCUSSION^ 1635.52 (a) and (b) ) can also be explained on the basis of chemical interaction concepts.The weakening and hence the lengthening of the C-N bond results in this functionalgroup having less influence on the adjacent group and hence results in higher C-Hvibrations. Mathematically, the factor ' for the C-H bond increases leading to highervibrational frequencies.In the fingerprint region (Figures 5.52(a) and (b)), the absorption bands at 1667-1275 and 996-740 cm" are all new developments. In figure 5.51, the bands at 885-827cm -1 were assigned to the bending vibrations of the -NH 2 group. These bands havedisappeared with the formation of new ones at lower frequencies, 835 and 795 cm',which is in agreement with the mathematical expression above and the chemisorptionconcept. These observations indicate that whatever surface layer or product is formedon the pre-wash mineral surface is a chemically adsorbed species and has molecularvibrational characteristics quite different from the bulk surfactant solution or material.On washing the mineral, a whole new set of absorption bands is revealed, especiallyin the diagnostic region of the spectrum, (Figure 5.53). This suggests that two or moredifferent layers of surface chemical species are formed, and that the top layer(s) is notas firmly held to the adsorbent surface as the bottom layer(s). The differences in theabsorption band positions, particularly in the diagnostic region, indicate that the bottomlayer may have a different molecular arrangement and/or vibrations from the top layerthough both are chemisorbed species. The presence of stretching vibrational modes at1540-1384 cm' (in the diagnostic region and assignable to molecular vibrations) and theweak shoulders near 3030 and 3095 cm" (assignable to bonds) are fingerprints ofthe adsorbate aromatic skeleton on the adsorbent surface. This is further supported bythe presence of the stretching -N-H bond vibrations at 3367 and 3238 cm -1 in the upperfrequency region. The drastic shifts in absorption bands due to the aromatic structureto 1539, 1457, and 1385 cm -1 after the washing process also supports the chemisorptionChapter 5. RESULTS AND DISCUSSION^ 164mechanism.There is the development of a new vibrational mode ( weak shoulder) at 1300-1299ern' which is different from the C-N vibrations at 1333 and 1270 cm' (Figure 5.50),1294 cm -1 (Figure 5.51) and 1275-1274 cm -1 (Figures 5.52 (a) and (13)) and may beassigned to one of the characteristic molecular vibrations of the new surface product. Thisdevelopment further confirms the involvement of the C-N group in chemical interactions.The appearance of new vibrational bands at 1149-1000 cm' is attributed to theperturbation of the C-H in-plane deformations as a result of the adsorption process.The new absorption bands in the lower wavenumber region reinforce the proposedchemisorption route. Thus, in the absence of ferric ion conditioning, the OPD (surfactant)molecules are chemisorbed, resulting in the formation of two or more layers of chemisorbedspecies on the adsorbent with the top layer(s) being relatively loosely held on the surfacecompared to the lower layer(s).Ferric ion pre-treatmentThe spectral record shown in Figure 5.54(a) has the solvent as the background spectrumwhile that in Figure 5.54(b) is the spectral difference between the adsorbent-adsorbate(with prior conditioning of the concentrate in ferric solution) and bulk surfactant solutionspectra. The two are very similar in shape and absorption band positions. However,there are a few differences e.g. the bands at 855 and 741 cm -1 (in Figure 5.54(a)).The absorption bands at 999-995, 1097-1095 and 3419-3363 cm -1 are present in all fourspectra'. However most of the absorption bands are different from those shown in Figure5.52.In the fingerprint region, the absorption frequencies are quite different compared tothe same region in the bulk solution spectrum (Figure 5.51). This suggests that the39 1n comparison with those obtained in the absence of ferric conditioning.Chapter 5. RESULTS AND DISCUSSION^ 165region 1644-711 cm" (Figures 5.54(a) and (h)) represent chemically adsorbed species.The disappearance of the absorption band at 1294 and the shoulder at 1194 cm"(present in Figure 5.51) or 1333-1268 cm-1 (in Figure 5.50) assigned to the C-N func-tional group from Figures 5.54 (a) and (b) supports the existence of chemically adsorbedspecies on the adsorbent surface.The vibrations occurring at 1643-1420 cm -1 are fingerprints from the aromatic skele-ton of whatever surface product is formed. The -NH 2 group also has a strong absorptionat 1643-1639 cm -1 .Figure 5.55 shows the absorption spectrum of the sample washed after the condi-tioning process. Between Figures 5.51 and 5.55, there are upward shifts in some of theabsorption bands in the fingerprint region. These upward shifts in frequency positionsfurther support the process of chemisorption.Just as there is reappearance of some of the absorption bands (present in Fig. 5.54),e.g. 1644-1640 cm -1 , others have disappeared with the emergence of new ones. Thisdevelopment of new vibrational modes suggests the formation of two (or more) layersof chemically adsorbed species with different vibrational characteristics (just like theprevious condition) and that the top layer(s) is relatively loosely held to the mineralsurface. The lower layer is, however, strongly bonded to the mineral surface such thateven the drastic washing procedures adopted could not dislodge it from the adsorbent.The band at 1775-1773 may be due to overtones and/or combinations. The absorptionband at 1641 cm' (Fig. 5.55) may be a contribution from the bending vibrations of-NH 2 (or -N-H) and co-adsorbed water molecules. The peak frequencies at 1641-139 Icm' are assigned to vibrational characteristics of the aromatic structural component ofwhatever surface product is being held on the surface after the washing process, thoughthe characteristic aromatic absorption wavenumber positions are different from those (,'the KBr wafer and the bulk solution spectra.Chapter 5. RESULTS AND DISCUSSION^ 166The appearance of weak shoulders near 3040 and 3008 cm -1 coupled with the bandsat 1641-1240 cm -1 confirms the presence of the adsorbate(aromatic)-related species onthe washed mineral surface.The disappearance of the stretching C-N. vibrations (1294 in Figure 5.51 or 1333and 1269 cm -1 Figure 5.50) and the appearance of new bands at the wavenumbers 1291and 1240 cm -1 further confirm the involvement of this functional group in chemicalinteractions, most probably with the nitrogen trying to share its free electron pair withthe metal ions of the sulfide lattice. In acidic media (e.g. under the zinc pressure leachingconditions), metal sulphides carry positive surface charges and the surfactant moleculewill act as a base. The basic character of the amine group implies that it is willing to sharethe valence electrons with any electrophile, thus, it either picks up protons in solutionand becomes protonated or interacts with the metal ions in the crystal lattice. With zincions being more reactive than hydrogen ions, the surfactant will be more willing to shareits free electrons with the zinc than the hydrogen ions.The disappearance of the band at 1098-1096 is due to the removal of soluble surfaceoxidation products from the sample material during the washing process.A comparison between Figures 5.53 and 5.55 shows that the two have a lot of char-acteristic absorption bands in common (within the limits of instrument sensitivity) suchas the bands at 1541-1538, 1300-1295, 1074-1073, and 998-994 cm -1 in the diagnos-tic region. This display of agreement in the fingerprint region indicates that whateveradsorbed species is formed on the adsorbent surface is common to both conditioningmethods inspite of the differences in the upper frequency- regions of the spectra. On theother hand, it may be speculated that this agreement is probably due to the chemicalcomposition of the concentrate. The concentrate contains about 11 % iron mostly in theform of pyrite and probably this minor amount of iron in the sulfide matrix is just enoughto affect any surface characteristics and hence the adsorption mechanism such that theChapter 5. RESULTS AND DISCUSSION^ 167presence of externally introduced ions on the surface is not reflected'.SummaryThe mechanism of adsorption of the surfactant orthophenylene diamine onto zinc sul-fide mineral (i.e. sphalerite) is through chemical interactions of the adsorbate with theadsorbent and this is independent of the pre-adsorption surface characteristics of theconcentrate. The study indicates the formation of two (or more) different layers ofchemisorbed species with the upper layer being vulnerable to removal by washing whilstthe lower component is firmly bonded to the adsorbent surface. The absorption spectrastrongly indicate that the OPD molecule interacts with the mineral through the C-Nfunctional group where the nitrogen free-electron pair overlaps with the metal ions ofthe sulfide mineral lattice leading to the formation of an amine-related surface species.At the same time, the hydrogen atoms of the -NH 2 functional group interact with theaqueous phase by forming hydrogen bonds, thus ensuring that the solid surface is contin-uously hydrophilic. This process tends to render the solid surface more hydrophilic thansulphophilic and causes an increase in the contact angle of the three-phase system [55].40 This is only a speculation.Chapter 5. RESULTS AND DISCUSSION^168 033500 2725 1950Wavenumbers (cm-1)O1175^400Figure 5.50: I<F3r pellet spectrum of OPD.CJILc;^•C--.4"--,12 co c., co roMI<I-(Oa) 'CD_,.,-, LC), 0 ,,,,, 0, ,,,i.0 „ei,o-,cO ff-, c--4^ r---0h'^ — ,,,c0odC.s.1at")Chapter 5. RESULTS AND DISCUSSION^ 169'3000^ 2000^ 1000Wavenumbers (cm-1)Figure 5.51: ATR spectrum of 5 g/L solution of OPD; solution condition: [1 -12SO4] =0.036 M, Temperature = 25'('; reference spectrum: solvent.-.41- 6M t044441-— 4444"]C ^0•140 .-7444,-41h44-4,04-10cv CO LaO4444,CrsO,d O CP -44-eh3500^27.67.5^2075Wavenumbers (cm-1)ti3-O1362.5^650Chapter 5. RESULTS AND DISCUSSION3500^2775^2050^1325^600Wavenumbers (cm-1)(a)170(b)Figure 5.52: ATR spectrum of sphalerite conditioned in OPD solution (no fenpre-treatment) for 20 minutes; solution condition: [H 2 SO 4 ] = 0.036 M, Temp. = 25(a) reference spectrum: solvent (b) less solution spectrum.Chapter 5. RESULTS AND DISCUSSION 1713500 2787.5 20'75Wavenumbers (cm-1)1362.5 650Figure 5.53: ATR spectrum of sphalerite washed after conditioning in 5 g/L OPD solutionfor 20 minutes, reference spectrum: solvent.Chapter 5. RESULTS AND DISCUSSIONC)C-)CS0CC).0 1793500 2775^2050^1325Waveriumbers (cm-1)600(a)-5 ,CCS000LC,<7,3500 2775^2050IN'avenurnbers (cm-1)1325^600Figure 5.54: AIR spectrum of sphalerite conditioned in 5 g/L OPD solution (after ferricpre-treatment) for 20 minutes; solution condition: [1-1 2 SO 4 ] = 0.036 M, Temp.= 25°C (a 1Reference spectrum: solvent (6) less the solution spectrum.OCVOCVOOciO 010^r•-).--, 01 hr)tr) ,= •nr 01 CD..— CO Cr/ CV -o--,r ro — c.,2 ---Chapter 5. RESULTS AND DISCUSSION^ 173r7 m'^ '3000 2000^ 1000Wavenurnbers (cm-1)OFigure 5.5.5: A'I'R spectrum of washed sphalerite after contact with ferric/OPD solution,reference spectrum: solvent.N HHH^ N^-M2+HHChapter 5. RESULTS AND DISCUSSION IT tFigure 5.56: Metal-OPD complex on interaction.Chapter 6CONCLUSIONS AND RECOMMENDATIONS6.1 CONCLUSIONSThe studies conducted have measured the interfacial activity of various surface activeagents and how these interfacial characteristics are related to the usefulness of thesereagents as dispersants for liquid sulphur in the oxidative pressure leaching of sphalerite.The fact that a surfactant can be very efficient in terms of interfacial tension reductionsand contact angle increases is no guarantee that the surfactant can be efficient in theleaching process. Besides the interfacial properties exhibited by the different surfaceactive agents other variables such as reagent stability come into play if a surfactant is tobe effective under the zinc pressure leaching conditions. The following are the conclusionsfrom the present studies:• In terms of interfacial tension reductions, OPD was the least effective. It effectedvirtually no change in the liquid sulphur—aqueous phase interfacial tension but actedto increase the sulphur—mineral contact angle. This implies the surfactant is not ad-sorbed at the liquid sulphur—aqueous interface but adsorbed at the liquid—mineralinterface. All other surfactants adsorbed at both interfaces, naphthalene sulphonicacid—formaldehyde condensates being the most reactive reagent at the interfaces.It effected the highest change in the interfacial tension and the contact angle at aconcentration of about 0.1 g/L. CAHSB exhibited about the same effect as lignin175Chapter 6. CONCLUSIONS AND RECOMMENDATIONS^ 176sulphonic acid up to about 0.3 g/L. However, unlike the behaviour of lignin sul-phonic acid, the interfacial tension decreased further when CAHSB concentrationwas increased to about 0.5 g/L beyond which the interface was insensitive to anyfurther increase in surfactant concentration. TAHSB and CAB effected an interfa-cial tension decrease to a lesser extent than the other reagents with the exceptionof OPD.• Just like the interfacial tension effects, lignin sulphonic acid, naphthalene sulphonicacid-formaldehyde condensates, and CAHSB effected the most contact angle in-crease from 80 to 145-155°; OPD effected an increase to 125-135°.• Without adsorption at the mineral-liquid interface, ^01A and -y ms will not be af-fected by the presence of any surfactant and hence ysA cos 0 (or y?IA - -yms ) shouldremain constant; however, -y SA cos 0 varies with surfactant addition. From Table5.9, -y,vis far exceeds the value of MA (compared to that obtained in the absence ofany surfactant), since a surfactant can only act (or accumulate) at an interface tolower the interfacial free energy, and not raise it. Thus -yms cannot be raised in thepresence of the surfactants. Hence the surfactants act to lower "rmA and the resultis an increase in the contact angle. The effect of lowerin g u!IA and lisit (exceptfor OPD), and hence increasing the contact angle, increases the tendency for liquidsulphur to detach (or roll) from the mineral surface, making the mineral accessibleto the aqueous phase.• Under leaching conditions, the effectiveness of the surfactants (used as dispersantsfor liquid sulphur) does not necessarily exhibit the same behavioural pattern as theirinterfacial activities. Thus, naphthalene sulphonic acid-formaldehyde condensates.CAHSB, CAB, and TAHSB were poor dispersants even under low pulp densityconditions. On the other hand, OPD was a very good dispersant for liquid sulphur.Chapter 6. CONCLUSIONS AND RECOMMENDATIONS^ 177the results being comparable to those obtained with lignin sulphonic acid underboth low and high pulp density leaching conditions. This indicates that the mostimportant criterion for a surfactant to be effective under leaching conditions isfor the surfactant to be able to increase the liquid sulphur-mineral contact angle(by being adsorbed on the mineral-liquid interface). The angle does not necessarilyneed to be very high provided a certain critical angle (greater than 90 0 ) is exceeded.• The presence of oxidizing ionic species such as ferric ions can have an adverseeffect on how the surfactants function. Such species can enhance the kinetics ofdegradation of the surfactants e.g. the effect of ferric/ferrous ions on the interfacialactivity of naphthalene sulphonic acid-formaldehyde condensates and the low %zinc extraction (even under low pulp density leaching conditions) in the presence ofCAB, CAHSB, TAHSB and naphthalene sulphonic acid-formaldehyde condensates.• The adsorption of lignin sulphonic acid proceeds through both physical and chemi-cal mechanisms. The surfactant interacts with the mineral through the hydrophilicgroups and acts as a bridge between the mineral and the aqueous phase. The in-teraction results in the formation of a bulk-like Zn(II)-surfactant complex havingan organic portion similar to the bulk surfactant structure.• OPD is adsorbed onto the solid mineral adsorbent through chemical interactionand this adsorption is independent of the pre-adsorption surface characteristics ofthe mineral. The infrared spectra suggest the formation of amine-related surfacespecies through the C-N functional group of the surfactant. The nitrogen freeelectron pair possibly overlaps with the metal ions of the sulphide lattice, with thehydrogen atoms of the -NH 2 group interacting with the aqueous phase through hy-drogen bonding, rendering the mineral surface hydrophilic. NIPD possibly interactsthrough a similar mechanism.Chapter 6. CONCLUSIONS AND RECOMMENDATIONS^ 178• It is believed that the formation of pockets of metal—surfactant species/complexeson the mineral surface is probably enough to make the mineral particle sufficientlyhydrophilic.6.2 RECOMMENDATIONSThese studies have shown that OPD (and NIPD) can be potential surfactants (or dis-persants) for the zinc pressure leach process. This finding could lead to the commercialapplication of OPD for zinc pressure leaching. However, more investigations need to hecarried out on the use of such surfactants. The impact of residual OPD and MPD onthe subsequent purification and electrowinning stages of zinc recovery are two areas ofconcern that need to be addressed before any full scale commercial application is initi-ated. The environmental effects of these reagents are also very important and requireextensive investigations. It is believed that lignin sulphonic acid degrades before the zincsulphate leach solution reaches the tank house, thus, not interferring with the process.The current study should be expanded to include other surfactants not mentioned in thisstudy.This thesis was designed to address the role of surfactants in the liquid sulphur—aqueous zinc sulphate interfacial system. The effect of interactions between the surfac-tants and precipitated solids such as jarosite, lead sulphate, silica, etc. on the coalescence(or dispersion) of liquid sulphur was not studied (or considered) and hence nothing canbe said about the role of such precipitated solid materials on the coalescence of sulphurin the presence of the surfactants under leaching conditions, e.g. though OPD was in-effective on the liquid sulphur—aqueous solution interface, it produced sulphur particlesof very fine sizes while lignin sulphonic acid which effected a reduction in the interfacialtension (i.e. sulphur dispersion) gave rise unexpectedly to coarse size sulphur particlesChapter 6. CONCLUSIONS AND RECOMMENDATIONS^ 179in the leaching. Thus a thorough study needs to be conducted on the effect of precipi-tated solids (on interaction with surfactants) on the coalescence (or dispersion) of liquidsulphur droplets.The pressure leaching of copper sulphide and pyrrhotite minerals above the meltingpoint of sulphur also experiences the sulphur occlusion problem which tends to reducethe extent of metal extraction. This has prevented the development of a high temper-ature process (above the melting point of elemental sulphur) for copper recovery fromconcentrates. It is strongly recommended that future research be directed towards iden-tifying surfactants that disperse liquid sulphur from the surface of industrially importantcopper minerals. Surface tension and contact angle measurements, leaching studies andATR—FTIR investigations will be required to develop a complete picture for the coppersystem.Bibliography[1] M.E. 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Parfitt, Dispersion of Powders in Liquids, John Wiley & Sons, N.Y., 1969,pp.1-17.[155] J. Basset, R.C. Denney, G.H. Jeffrey, J. Mendham, Vogel's Textbook of QuantitativeInorganic Analysis, 4th edition, Longman, London, 1978.[156] J.S. Fritz and G.H. Schenk, Quantitative Analytical Chemistry, Allyn & Bacon,Inc., Boston,1979.[157] L.F. Hamilton and S.G. Simpson, Quantitative Chemical Analysis, MacMillan Co.,N.Y., 1958.Bibliography^ 197[158] R.A. Day, Jr. and A.L. Underwood, Quantitative Analysis, Prentice-Hall, NewJersey, 1986.[159] A. Doren and P.C. Ronxhet, Infrared Study of the Adsorption on Silica of a Col-lector with an Oxyethylene chain, Int. Journal of Min. Proc., 9(1982), pp.343-351.[160] E.J. Purcell, Calculus with Analytic Geometry, Appleton-Century Crofts, N.Y.,1965, pp.584-589.[161] G.B. Thomas, Jr. Calculus and Analytic Geometry, Addison-Wesley Company, Inc.,1960, pp.585-590.[162] D.J. Shaw, Introduction to Colloid and Surface Chemistry, Butterworths, London,3rd. ed.[163] Manual on Naphthalene Sulfonic Acid-Formaldehyde Condensates supplied byHandy Chemicals Ltd., Laprairie, Quebec, Canada.[164] K. Osseo-Asare, Interfacial Phenomena in Leaching Systems, in Hydrometallurgi-cal Process Fundamentals, ed. R.G. Bautista, Plenum Press, 1982, pp.227-269.[165] H.I. Bolker and N.G. Sommerville, Infrared Spectroscopy of Lignins. Part II:Lignins in Unbleached Pulp, Pulp and Paper Magazine, vol. 64, 1963, pp. T-187-193.[166] A.N. James and P.A. Tice, The Presence of Carboxylic Groups in Lignosulfonatepreparations, Tappi, vol.48, 1965, pp.239-244.[167] A.J. Michel, A.J. Watson and H.G. Higgins, An Infrared Spectroscopic Study ofDelignification of Eucalyptus Regnans, Tappi, vol 48, 1965, pp.520-532.Bibliography^ 198[168] H.L. Hergert, Infrared Spectra of Lignin and Related Compounds II. Conifer Ligninand Model Compounds, Journal of Organic Chemistry, vol. 25, 1960, pp. 405-413.[169] K.V. Sarkanen and C.H. Ludwig, Lignins, Occurrence, Formation, Structure andReactions, Wiley-Interscience, N.Y.[170] L.J. Bellamy, Infrared Spectra of Complex Molecules, John Wiley & Sons, N.Y.,1958.[171] G. Socrates, Infrared Characteristic Group Frequencies, John Wiley^Sons, N.Y.,1980.[172] T.W. Healy and M.S. Moignard, A Review of Electrokinetic Studies of Metal Sul-phides, in Flotation Volume I - AIME New York, 1976, ed. M.C. Fuerstenau,pp.275-297.[173] M.J. Moignard, D.R. Dixon and T.W. Healy, Electrokinetic Properties of the ZincSulphide-Water and Nickel Sulphide-Water Interfaces, in Proceedings of AIMM.,Sept. 1977, pp.31-38.[174] G. Owusu, D.B. Dreisinger, and E. Peters, Oxidative Acid Pressure Leaching ofZinc Sulfide Minerals in the Presence of Different Surfactants, presented at theCanadian Chemical Engineering Conference, Toronto, October 1992.Appendix ARADII OF CURVATUREThe derivation of radii of curvature is based purely on analytical geometry and appliedto the pendant drop. The curvature, K, of a curve at a point P(x,y) is the rate of changeof 0 per unit change in arc length along the curve at that point i. e.dOdswhere 0 is the angle in Figure A.57 and s is the arc length measured along the curve.Radius of curvature is the reciprocal of curvature:= 11K^ (A.79)Thus, if one of these quantities is known, the other can be easily obtained. Consider theequationp= „y ( 1 1 - + 1 2 )^(A.80)where R 1 and R2 are the radii of curvature. Consider the curve C (i. e. the profile ofthe pendant drop) to be represented by a twice differentiatable function y = f(x). LetP(x,y) be any arbitrary point on the curve, C, and 0 is the inclination of the tangentto the curve at P (i. e. the angle between a tangent and the horizontal). At P, R 1 canbe represented as being in the plane of the paper and R2 perpendicular to the plane ofthe paper. R2 is given by NP (on the diagram) which rotates about 00. At the apex(origin, 0) R 1 = R2. In two-dimensional space, x, y, the direction of a line is given byits inclination. The tangent to the curve at any arbitrary point P is given by:tan 0 = f'(x) = y'^ (A.81)K = (A.78)199Appendix A. RADII OF CURVATURE^ 200Figure A.57: Section of pendant drop profileTherefore, sinceK,=dq=dstan -1dqdx .f(x)dxdsdx(A.82)(A.83)dx tan -1 r(x)ds(A.84)However (from calculus),d^_ 1^,-1 y"(x)c7x- tan (x) = 1 + (y/(x))2^(A.85)Meanwhile, -(!p; can be derived from the curve considering the movement of P(x,y)such thatdsdx ^1^(Yi(x))2(A.86)Appendix A. RADII OF CURVATURE^ 201=-- y" (x)^1 1 + (y'(x))2 ^ + (y'(x)) 2y" (x)'V1^(y'(x)) 2The radius of curvature R1 , can be obtained as the reciprocal of K.From Figure A.57(A.87)(A.88)R2 =x^x[1^(yi (x)) 2 } 1 /2(A.89)sin c/5 y' (x)Appendix BCHEMICAL ANALYSISB.1 Analysis of Fe3+ and Fe 2 + ionsAbout 55-60 mL of deionized water is added to 5 mL of sample solution (i.e. filtrate).The diluted solution is titrated against 0.1 or 0.05 M ceric sulphate solution (i.e. Ce 4+)using and an Autotitrator with Pt. and calomel electrodes. The Fe 2+ ion concentrationis then computed once the end point of the titration is reached.The ferric ion concentration is calculated from the total iron concentration of thesample solution. About 15-25 mL of 6 M HC1 solution is added to 2 mL of the samplesolution and heated to about 80°C. Any ferric ion is reduced to the ferrous state byadding 0.5 N SnCl2 solution until the solution turns colourless. Excess SnC1 2 solution isadded. The solution is cooled and then diluted with 55-60 mL of deionized water; 10 mLof 5 % HgC1 2 solution is also added and allowed to stand for about 5 minutes and thentitrated with the standard ceric sulphate solution to an end point of about 740 mV. Thetotal iron concentration is then computed. The difference between this and the ferrousion concentration gives the total ferric ion concentration.B.2 Total Acidity AnalysisTo a 2 mL sample solution, 10 mL potassium hydrogen oxalate solution, KHC 2 O 4 , (atpH = 5-6) is added to complex any metals in solution and then diluted with about55-60 mL of deionized water. The solution is then titrated against 0.5 M NaOH using202Appendix B. CHEMICAL ANALYSIS^ 203the Autotitrator (with a combination pH electrode - glass+calomel). The volume of thetitrant is then read from the chart. A blank titration is done by titrating 10 mL potassiumhydrogen oxalate solution against the standard NaOH solution. The difference betweenthe two volumes give the NaOH requirement of the solution acid (H+) concentration.The total acid concentration of the sample solution is then calculated. The reactionsoccurring are:K H C2 04 -4 K+^(B.90)For the blank titrationHC2 04 + 0H^C20?_+ H2O^(B.91)Sulphuric acid dissociates according toH2 SO4 = H+ HSO:i^ (B.92)ThenH+^HC204^ (B.93)andHS0:1 d-C2 0?1— SO?, - HC2 04-^(B.94)Reaction B.91 then proceeds from reactions B.93 and B.94. This gives the base require-ments by the acid.The metal ions are complexed according to:Zn 2+^Zn(C204 42n-2)-^(B.95)andFez+ + nC2 W1 - ,== Fe(C204) (2n z)^(B.96)These two reactions do not make any H+.Appendix B. CHEMICAL ANALYSIS^ 204B.3 Sulphur AnalysisDetermination of elemental sulphur in leach residueAdd 50-60 mL of ti 2.1 mol/L Na 2 S to 0.5-0.8 grams of sample (i.e. leach residue) andheat to 70-80°C. After about 15 minutes, filter and wash residue 2-3 times with about 0.5mol/L NaOH solution followed by distilled water. Dry residue and weigh. The differencebetween residue and sample weights gives the weight of elemental sulphur.Determination of sulphate sulphur in leach solutionTo 5-10 mL of sample solution, add 50 mL of water and about 1 gm. of hydroxylaminehydrochloride and gently heat to dissolution of salt. Filter into 500 mL beaker and washthoroughly. Adjust the pH of the filtrate to about 7 using ammonium hydroxide andthen add 5 mL of HC1. Dilute the solution to about 300-400 mL, and heat to boiling.Add 25 mL of barium chloride solution while boiling or stirring and allow to stand forseveral hours or overnight on a low warming plate. The barium chloride solution willprecipitate the sulphate sulphur as barium sulphate. Test if solid precipitation is completethrough drop-wise addition of barium chloride to the supernatant liquid. If precipitationis complete filter and thoroughly wash residue with 5 to 6 times with hot water. If theprecipitation is not complete, add more chloride solution to complete it. Transfer theresidue to a ceramic crucible and ignite in muffle furnace at 850-900°C for one hour.Cool the solid, weigh, and calculate sulphur content as follows:Sulphur, % = A * 0.13735 * 100Or1000g/L Sulphur = A* 0.13735 * C * DBwhereAppendix B. CHEMICAL ANALYSIS^ 205• A = weight of barium sulphate found (g)• B = sample weight (g) (if starting material is solid and has to be digested beforethe sulphur determination)• C = volume of sample solution used (mL)• D = dilution (as ratio of diluted volume to initial volume)Determination of sulphide sulphur in concentrateAbout 10 mL each of bottle strength HNO 3 and HCl are added to 0.3-0.5 gram ofconcentrate sample in a beaker and then heated until contents boil, shaking the beakerfrom time to time. The beaker is cooled and about 1 mL of liquid bromine is added tooxidize the sulphide sulphur to the sulphate form. The beaker is allowed to stand coldfor about 10 minutes with gentle shaking from time to time. The beaker is gently heatedto evaporate any residual bromine. The solution is boiled down to 10-15 mL, allowed tocool and then carefully diluted to 50-60 mL and filtered into a 250 mL volumetric flask.The filter is washed several times and the solution made up to volume. The sulphurdetermination is carried out as described in the previous section.Appendix CSURFACE EXCESSThermodynamically, the Gibbs adsorption equation for an interface in an i componentsystem is given by [149]:— d-y = SLdT^ (C.97)where Siint = interfacial excess entropy per unit areaT = temperature= chemical potential of ith component in the systemAt constant temperature— dry =^d^ (C.98)The chemical potential, p i , of the ith solute in the solution is given by:^= + RT1na i^ (C.99)or^tc i = /.4 + RT1n fi ci^ (C.100)where ,a`i = standard chemical potential of component i.Differentiating the above equationdit i = RT(dln f^dlnc i )^ (C.101)RT (1dln fi dlnci (C.102)dlnc i) = RT4-i (dlnci)^ (C.103)206Appendix C. SURFACE EXCESS^ 207For i components system, equations C.98 and C.103 give— d-y = RT(116d1nc i^116clinc2 + • • • +^(C.104)Component 1 is considered to be the solvent whose interfacial excess is assumed benegligible. This means thatIi = 0and rZ can be written as r i . Equation C.104 can thus be simplified to^F36d/nc3^r 4 4-4 dInc4— d-y^RTF22d1nc2 (1 + • +  ^(C.105)T,^^2C2at c2^1 2C2ainc2^12c2u,nc2This simplifies to^— d-y = inRTF 2 .dlnc2^(C.106)where=• ClF36d/nc3 F46^+ • +dInc4^F4dlnciF2 C2ClinC2^r,2C2Utne2 2C2aInc2For a solution containing an organic electrolyte (or surfactant) RX, and inorganicelectrolyte ZnSO 4 , it can be assumed that they both dissociate completely according to:zX +ZnSO4^Zn2+ S0,21wherez = valency of organic surface active anionR = is the surface active anion.For a multicomponent solution like this one—^+ rznso4ditznso4and(C.107)itRXz = PR+ Zilx+Appendix C. SURFACE EXCESS^ 208itznsat = ttzn2-4- + Itso -Mathematically, it is shown that [149]:rZnSO4 = r SO42-Similarly,rRxz = F RThus, the relative surface excess of the surfactant can be computed using equationC.106 if the following speciation assignments are made:el assigned to the solvent (water)c2 assigned to R (or RX,)c3 can be assigned to X+c4 represents Zn 2+ andc5 represents SO4-.


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