Heavy Metals Removal from Effluents by Adsorption on
Activated Silica Sols
by
Elias El-Ammouri
A thesis submitted io the FacuIîy of Craduate Studies
and Raearch of McGN Universi@ in partial fuI/dlment of the
requirements for the degree of Doctot of Phüosuphy.
Department of Mining and Metallurgical Engineering
McGill University, Montreal
January, 2ûûû
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Abstract
Effluents are neutralized with lime forming a voluminous jaste sludge which may
slowly redissolve in groundwater. T'lis sludge may also be a significant secndary metd
source. The present research aims to develop a concept for metals removal and recovery
using stabilized colloidal suspensions of activated silica sol, which acts as a neutralizing
agent and an adsorption medium.
Silica sol was prepared by fintly acidifjhg a sodium silicate solution to initiate
polyrnerization, which would ultimately yield a gel. Polyrnerization was then mested
before gelation by water dilution giving a stabilized, negatively charged sol. Adsorption
and precipitation were distinguished by comparing the coloured products from adding
silica sol, lime or sodium hydroxide to cobaltous sulphate solution.
Studies of dissolved coppedsilica sol interactions showed copper
adsorption/desorption to be a reversible pH-controlled process, which was monitored 'in-
situ' using a cupric specific ion electrode. Dissolved copper values determined by the
electrode were identical to those measured by atomic absorption spectrometry ( A M ) for
solutions at below pH 7. At above pH 7, the electrode showed zero dissolved cupric, but
AAS reported some dissolved copper, attributed to desorbed colloidal cupnc hydroxide.
A copper/sol underflow sharply divided fiom a clear ovefflow were the adsorption
products. Addition of concentrated sulphuric acid to centrifbged underflow produced a
concentrated copper solution and regenerated sol. The separation of two metals (cupric
and ferric) using pH controi was also demonstrated.
ii
Adsorptionfdesorption cycles were used to remove iron from synthetic effluents
producing concentrated iron products. Three consecutive cycles were required to produce
30 g/L dissolved iron, suitable for sewage treatment, fiom 0.5 g/L iron effluent. Final pH
control with lime precipitated desorbed colloidal hydroxides. Good settling
characteristics following adsorptionlprecipitation from synthetic and actual mine effluents
were observed. Problems included gypsum build-up, while excess silica addition gave
dispersion, high silica losses and no phase separation.
Nickel was selectively recovered over iron and magnesium from a tailings pond
sludge by acid leaching then solution treatment with activated silicdlime. One
adsorptioddesorption cycle produced up to a 17 g/L Ni product from 2 g/L Ni leach
solution.
iii
Résumé
Lonque des effluents sont neutralisés par traitement à la chaux. des résidus
volumineux de boue sont produits, pouvant se dissoudre lentement dans les eaux des sols.
Cene boue peut être une source secondaire et importante de métaux. La présente
recherche tend à développer un nouveau concept pour le retirement et la récupération des
métaux, en utilisant des suspensions colloïdales stabilisées de sol de silice actifs. agissant
comme un moyen pour la neutralisation et l'adsorption.
On prépare les sols de silice en acidifiant une solution de silicate de sodium pour
initier la polymérisation qui pourrait éventuellement produire un gel. Cependant. la
polymérisation est arrêtée par dilution. cn ajoutant de l'eau. et ceci avant la gélation. Cet
ajout résulte en un sol stabilisé et chargé négativement. L'adsorption et la précipitation
sont différenciées en comparant la couleur des produits obtenues. après l'addition de sois
de silice, de chaux ou d'hydroxyde de sodium. à une solution de sulfate de cobalt.
Les études des interactions du cuivre dissout avec les sols de silice indiquent que
l'adsorption/désorption du cuivre est un procédé réversible contrôlé par le pH, qui est
suivie in situ par une électrode spécifique d'ions cupriques. À un pH inférieur à 7. la
concentration du cuivre dissout déterminée par l'électrode était identique à celle mesurée
par spectrométrie d'adsorption atomique (SAA). À des pH supérieurs à 7. l'électrode
mesurait une concentration nulle du cuivre dissout. Par contre la SAA indique une faible
quantité de cuivre attribuée à le désorption de l'hydroxyde de cuivre colloïdal. Le produit
de l'adsorption résultait en une phase inférieure cuivre/sols de silice, nettement séparée du
iv
surnageant. L'addition de i'acide sulfûrique à la phase inférieure centrifugée a produit une
solution concentrée de cuivre et a régénéré les sols de silice. On a pu séparer les metaux
(cupriques et ferriques) en ajustant le pH.
Les cycles d'adsorptioddésorption ont été employés pour retirer le fer d'effluents
synthétiques, produisant ainsi des solutions concentrées de fer. À partir &un effluent de
fer de 0.5 g/L, trois cycles consécutifs ont été requis pour générer 30 g/L de fer dissous.
convenables au traitement des eaux usées. Le contrôle final du pH à la chaux précipitait
les hydroxydes colloïdaux désorbés. On observe une bonne sédimentation suite à
I'adsorption/précipitation des effluents synthétiques et des mines. Le problème se résume
à I'accumulation du gypse alors que l'addition excessive de silice produit une dispersion.
une importante perte de silice et une non- séparation des phases.
D'un étang de boue. le nickel a été sélectivement récupéré et séparé du fer et du
magnésium par lessivage acide. suivi d'un traitement de la solution à la silice active et la
chaux . Un cycle d'adsorption/désorption générait jusqu'à 17 g/L de nickel d'une solution
lessivée de 2 g/L de nickel.
Acknowledgements
I would like to express my profound appreciation and gratitude to my supervisor
Prof. P.A. Distin while working with him for the past seven years (Master's and Ph.D),
especially for his kindness and gentleness in answering my questions. Also, for his
guidance, advice and moral support throughout this research.
I would like to thank National Silicates Ltd. for their financial support, and the
guidance and efforts of Dr. Graham Hagens for proposing new ideas. The help of Barbara
Lernpka and Vicky Sidorkiewicz is also appreciated.
At McGill University, hydrometallurgy and mineral processing groups, department
of mining and metallurgical engineering, special thanks to my colleagues and to Dr. S. R.
Rao for introducing new ideas to part of this project, and of come. Prof. J.A. Finch for his
guidance and support during the past two years.
Finally, thanks to my wife. Marlene, for her love and moral support during this
entire period.
Table of Contents
. . .............................................................. Abstract n ................................... ................. Rdsurnh ............. iv
................................................ Acknowledgements vi ............................................... List of abbreviations x
Nomenclature ............ ..... .................................. xi List of Figures ...................................................... xii List of Tables ....................................................... xvii
CHAPTER 1 ............................................................... 1 Introduction
CHAPTER 2 ............................................................... 6 Literature Review
....................................................................... 2.1 Introduction 6 ............ 2.2 Problems Associated with Metal Precipitation as Hydroxides 6
........ 2.3 ~mprovements/Altematives to Hydroxide Precipitation Processes 9 2.4 Metals Removal from Effluents using Silica (quartz, silica gels, silica
....................................................................... resins or filters) 12 ................................. Production and Usage of Silicate Materials 13 ................................... Aqueous Chemistry of Silica and Silicates 15
................................................. Structure of Monosilicic Acid 17 ......................................................... Water Soluble Silicates 18
................................................ Polymerization of Silicic Acid 20 ......................................................................... 2.10 Gelation 23
.................................... 2.1 1 Dissolved MetaYSilica Sol Interactions 28 ......................... 2.12 Adsorption in Dissolved Metal/' Silica' Systems 31
CHAPTER 3 ............................................................. 36 Experimental Procedure
........................................................ 3.1 Reagents and Materials 36 .................................................................. 3.2 Sol Preparation 37
vii
3.3 Metals Recovery Stage ......................................................... 39 3.4 Metals Redissolution Stage .................................................... 40 3.5 Settling Stage .................................................................. 42 3.6 Sludge Treatment ............................................................. 43 3.7 Analytical Techniques ......................................................... 44
3.7.1 Dissolved Metai Concentrations ................................. 44 3 .7.2 Dissolved Cupric Concentrations .................................. 45
........................................................ 3.7.3 Zeta potentials 46 .............................................................. 3.7.4 Viscosity 46
3.7.5 Surface Tension ...................................................... 47
CHAPTER 4 ............................................................. 48 Characteristics of Dissolved CopperIActivated Silica In teractions
4.1 Introduction ...................................................................... 48 4.2 Gel Time Measurements ...................................................... 48 4.3 Calibration of Cupric Ion Electrode .......................................... 55 4.4 Copper Adsorption on Silica Surfaces (quartz. silica gel. activated
.................................................................................. silica) 59 4.5 Activated Silica as Copper Adsorber and pH Controller .................. 65 4.6 Copper Recovery using Adsorption/Desorption Cycle ..................... 74 4.7 CopperlIron Separation using Activated Silica .............................. 77
CHAPTER 5 ............................................................. 79 Iron Recovery from Acid Mine Drainage using Activated Silica
5.1 Introduction ...................................................................... 79 5.2 Iron Recovery Step (Fig . 3.1). ................................................. 79 5.3 Thickening (Fig 3.1) ............................................................ 89
5.3.1 Effect of Container Size ............................................. 90 5.3.2 Effect of Degree of Gelation ........................................ 92
.......................................................... 5.3.3 Effect of pH 92 5.3.4 Effect of Silica Concentration ...................................... 95 5.3.5 Effect of Recycling Regenerated Silica (after redissolution
......................................................................... step) 98 5 -3.6 Effect of Recycling Iron-'loaded' Silica (afier recovery step) . . 99 5.3.7 Settling Tests with 'Balsam Stream' Effluent .................. 101
5 -4 Iron Redissolution (Fig . 3.1). ................................................ 103 ......................................................... 5 -5 E ffect of Temperature 109
.............................................................. 5.6 Gypsum Formation 112
viii
CHAPTER 6 ........................................................... 121 Nickel Recovery from Tailings Pond Sludge by Acid Leaching and Treatment with Activated Silica
.................................................................. 6.1 Sludge Source 121 ................................................ 6.2 Sludge Treatment Flowsheet 122
............................................. 6.3 Process Conditions and Results 124 ............................................................ 6.3.1 Acid leach 124
6.3.2 Nickel Recovery onto Activated Silica (Fig . 6.2). .............. 124 ............................................... 6.3.3 Thickening (Fig . 6.2). 128
6.3.4 Nickel Redissolution (Fig 6.2) ..................................... 133 .................................. 6.4 Effluent Treatment Using Activated Silica 136
CHAPTER 7 .......................................................... 138 Summary
...................................................................... 7.1 Conclusions 138 ......................................................... . 7.1 1 Introduction 138
............................................... 7.1.2 Fundamental Studies 139 ............................................. 7.1.3 Potential Applications 141
...................... Iron Removal from Acid Mine Drainage 141 Selective Nickel Recovery from Tailings Pond Sludge ...... 143
.......................................................... 7.2 Claims to Originality 144 ...................................... 7.3 Suggestions for Further Investigations 145
............................................................... References 147
List of Abbreviations
AAS atomic absorption spectrometry
LSM liquid surfactant membranes
ISE ion specific electrode
PZC point of zero charge
XRD X-ray di fiaction
EDXRA energy dispersive X-ray analysis
Nomenclature
- m
- 'N' silicate
- 'D' silicate
- wlw %
- v/v %
- 1 % sol
- 1
weight ratio SiOdNazO
Silicate with m = 3.22 contains (w/w %) 8.90 % Na20,28.7 % Si02
Silicate with m = 2.00 contains (wlw %) 14.7 % Na2Q29.4 % Si02
weight percent
volume percent
1 .O w/w % Si02 sol
ionic strength
activity coefficient
ionic activity
molar concentration of species i
charge of species i
Nernst potential
Stem potential
zeta potential
moles of metal adsorbed
moles of hydrogen ion released
List of Figures
Figure 2.1 111: Solubility of hydroxides versus pH ..................................... 8
Figure 2.2 1121: Lime neutralization process alternatives .............................. I O
Figure 2.3 1411: Process for production of sodium silicate products .................. 14
Figure 2.4 [5. 51). Proposed mechanism of dissolution of silica in water in the presence of hydroxyl ions . Dotted lines represent the interface between silica on the left and water on the right ............................................................. 16
............. Figure 2.5 15.63. 641. Solubility diagram for amorphous silica at 25 OC 18
Figure 2.6 (801: Structure of a charged silicate polymer ............................... 20
Figure 2.7 [SI: Sol vs . gel vs . precipitate ................................................. 21
Figure 2.8 1811: Change in pH with time after addition of mineral acids to sodium silicate (Si02Ma20. m = 3.22 by weight. 20g SiO2/L) .................... 23
Figure 2.9 [SI]: Gel times of sodium silicate neutralized with different acids . (SiOdNazO. m = 3.22 by weight. 20g Si021L) ........................ 24
Fipre 2.10 1821: Gel times at 25 O C for various concentrations of sodium silicate (SiOflazO. m = 3.22 by weight) when mixed with sulphuric acid to give initial pH values as shown ................................................... 24
Figure 2.11 186-871: The electrical double layer (Stem model) and potential at a particle surface in solution ................................................ 27
Figure 2.12 1801: Production of "stabilized silica sol for water clarification ...... 28
Figure 2.13 1921: Copper ion activity vs . pH after the addition of sodium silicate or NaOH solution to Cu(ClO& solution at pH 4.0 ........................... 30
Figure 2.14 [IOSJ: Proposed mechanisrn for copper adsorption on silica (quartz) surface ........................................................................ 35
Figure 2.15 198.10 0.1021. Metals adsorption ont0 silica particles vs pH for various metais ........................................................................... 35
x i i
Figure 3.1 : Conceptual flowsheet for metals recovery/redissolution stage ................ using activated silica sol (one cycle), - could be almost any metal.. 38
Figure 3.2: Metals recovery/redissolution experimental set-up (with the cupric ion electrode). ............................................................................. 40
Figure 3.3: Size descriptions for Pyrex glas containers used for settling experiments.. ................................................................................. -43
Fipre 3.4: Typical absorbance vs. concentration calibration curves obtained when using the atomic absorption spectrometer.. ...................................... 45
.......... Figure 4.1 [5,63,641: Solubility diagram for arnorphous silica at 25 OC.. 50
Figure 4.2: Effect of 'N' silicate aging on gel time.. ................................... 5 1
Figure 4.3: Viscosity of 'N' silicate sols stabilized at different gel time fractions. Conditions as in Table 4.1 ....................................................... 54
Figure 4.4: Surface tension of 1 .O w/w % Si02 sol stabilized at different .......................................................................... gel time fractions.. 55
Fipre 4.5: Calibration curves for the copper electrode. Sodium sulphate (2 M) added as ionic strength adjuster.. ................................................. 57
Figure 4.6: Effect of water addition on the sensitivity of the copper ion selective electrode at pH 5.5.. ............................................................ 58
Fipre 4.7: Zeta potential of 'N' silicate sols and copper concentration as a function of pH. 1 .O w/w Sioz sol stabilized at half total gel time of 18 minutes. 2 I mL sol added to 100 mL 63.5 ppm cupric solution cupric at pH 5.5 initially. pH is adjusted downward with 10 wlw % sulphuric acid after sols addition. Copper analyses by A. A.S ................................................................. 6 1
Figure 4.8 11121: Zeta potential of quartz as a function of pH in the presence -4 o f P l 0 MPbC12 ......................................................................... 64
Figure 4.9: Zeta potential of 'N' silicate sols stabilized at different fractions of total gel time. 1 .O wlw % Si02 sols with total gel time of 18 minutes.. ............ 64
x i i i
Figure 4.10: Dissolved cupric concentration vs. pH. Effect of reagent addition on copper concentration. 1 .O wlw % Si02 sol ('N' silicate) stabilized at half total gel time of 18 minutes. 1 .O wlw % Si02 sol ('D' silicate) stabilized at half total gel time of 2 1 minutes. Lime added as 28 g CaOL suspension. Analyses by A.A.S .......................................................................... 67
Figure 4.1 1 a): Products obtained at pH 8.05 after the addition of 1 .O wlw % Si02 sol (left) or sodium hydroxide (right) to a 2 g/L cobalt solution. AAer fiee settling, a pink overflow is shown for both samples while the underfiow is pink after sol addition and blue after sodium hydroxide addition.. .................. 67
Figure 4.1 1 b) Products obtained at pH 9.5 after the addition of 1 .O wlw % Si02 sol (right) or lime (left) to a 2 g/L cobalt solution. AAer fiee settling, a clear overflow is shown for both samples while the undertlow is pink afier sol addition and green after lime addition.. ............................................... 68
Figure 4.1 1 c) Centrifuged products obtained at pH 9.5 after the addition of (lefi to right) 1 .O W/W % Si02 sol, sodium hydroxide, lime to a 2 g/L Cobalt solution.. ........................................................................... 68
Figure 4.12: pH vs. volume of activated silica sol ('N' or 'D') used for results obtained in Figure 4.10.. ................................................................ 69
Figure 4.13: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1 .O wlw % Sior sols ('N' silicate) with total gel time of 18 minutes. Analyses using cupric ion electrode.. .................................... 72
Figure 4.14: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1 .O wlw % Si02 sols ('N' silicate) with total gel time of 40 minutes. Analyses using cupric ion electrode. .............................. 72
Figure 4.15: Dissolved cupnc concentration vs. pH. Effect of polymer size on copper adsorption. 1 .O wlw % Si02 sols ('N' silicate) with total gel time of 105 minutes. Analyses using cupric ion electrode.. ................................. 73
Figure 4.16: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1 .O w/w % Si02 sols ('N' silicate) with total gel time of 18 minutes. Analyses by A.A.S.. ................................................ .73
Figure 4.17: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1 .O w/w % Si02 sols ('D' silicate) with total gel time of 2 1 minutes. Analyses by A.A.S.. ................................................ -74
xiv
Fipre 4.18: Products corresponding to Table 4.6 for the 3 15 ppm copper feed after centrifugation. The lefk side is the blue copper4oaded' sol (bottom phase) and clear effluent (top phase) after adsorption at pH 7.0. The right side is the sol for recycle (bottom phase) and concentrated solution (top
....................................................... phase) after desorption at pH 4. 14. 77
Fipre 5.1: Effect of container size on settling rate. Conditions: 558 pprn Fe feed, 1 .O w/w % Si02 'N' sol stabilized at 50 % of gel tirne, 200 mL sol11 L feed solution, pH adjusted to 4.0 with lime.. ................... 91
... Figure 5.2: Volume % as underflow vs. time for results shown in Figure 5.1 . . . 92
Fipre 5.3: Effect of degree of gelation on settling rate. Conditions: 558 pprn Fe feed, 1.0 w/w % Si02 'N' sol, 200 mL soVl L feed
................................................. solution, pH adjusted to 4.0 with lime.. 93
Figure 5.4: Effect of final solution pH on settling rate ('Medium' container). Conditions: 558 ppm Fe feed. 1 .O w/w % Si02 'N' sol stabilized at 50 % of gel time, 200 mL soVl L feed solution, pH adjusted with lime ............................. 94
Figure 5.5: Effect of final solution pH on settling rate ('Large' container). Conditions: Same as in Figure 5.4.. ....................................................... .94
F i p n 5.6: Effect of sol concentration on settling rate. Conditions: 558 pprn Fe feed, 'N' sol stabilized at 50 % of gel time, pH
................................................................. adjusted to 4.0 with lime. 97
Figure 5.7: Effect of sol concentration on settling rate, for results in Figure 5.7. R= Molar ratio (Si02 / Fe). ................................................................. 97
Figure 5.8: Effect of recycling regenerated sol on settling rate. Conditions: 558 pprn Fe feed, 1.0 w/w % Si02 'N' soi stabilized at 50 % of gel time, 200 mL soVl L feed solution, pH adjusted to 4.0 with lime.. ................... 99
Figure 5.9: Effect of recycling iron-loaded sol on settling rate. Fresh sol is 1 .O wlw % Si02 'N' sol stabilized at 50 % of gel time, 1 L 558 pprn Fe feed at pH 2.4. Final pH adjustment to 4.0 with lime for each cycle. Stirred slurry in excess of 1 L rejected before settling.. ....................... 10 1
Fipre 5.10: Effect of sol concentration on settling rate for 'Balsam Stream' (Inco) effluent (60 pprn Fe, 27 ppm Ni, 1.3 ppm Cu). 1 .O w/w % Si02 'N' sol
.......................... stabilized at 50 % of gel time. pH adjusted to 5.5 with lime.. 102
Fi y n 5.1 1: Typical sarnples obtained from the tests in Tables 5.7 - 5.9 afler centrifugation. a) and b) are samples fiom the 'recovery step' for 5 5 8 ppm Fe and 1 1.2 g/L Fe feed respectively , and show iron-loaded sols beneath clear overfiow. In b) a separate gypsurn layer was produced. Figure c) shows a concentrated iron product (about 35 g/L Fe) above centri fuged underflow (white silica/gypsum mixture) fiom the
...................... redissolution step following recovery fiom 1 1.2 g/L Fe feed..
Fipre 5.12: Femc removal fiom solution as a function of temperature. Conditions: 558 ppm Fe feed, 1 .O w/w % Si02 'NT sol stabilized at 50 % of gel
............. time, 200 rnL solll L feed solution, final pH adjusted to 4.0 with lime.. 1 1 I
Figure 5.13: Redissolution of femc as a function of temperature. Iron loaded sol generated as in Figure 5.12 at 25 OC then stored for one day before redissolution.. ,112
Figure 5.14: Photornicrographs of centri fùged underflow afier iron redissolution as in Table 5.9, column a). ............................................ .117- 1 19
....................... Figure 5.15: Recovery/Redissolution circuit with silica wash.. 120
. Figure 6.1 11221: Waste treatment operations of inco Ltd., Sudbury, Ontario.. 1 23
Figure 6.2: Proposed process for sludge treatrnent.. .................................. 123
Fipre 6.3: Settling of metal 'loaded' activated silica. Effect of feed composition. 200 mL 1 .O w/w % Si02 sol additions to 1 L feed. Final pH adjustment with lime ............................................................ 129
Figure 6.4: Settling of metal 'loaded' activated silica. Effect of silicdfeed . ratio. 1 .O w/w % Si02 sol additions to 1 L feed. Final pH adjustrnent with lime.. 132
.............. Figure 6.5: Existing and proposed processes for emuent treatment.. 137
List of Tables
Tabk 2.1 1111: Provincial effluent quality limits.. ..................................... 9
Table 4.1: Wavelength and bandpass used for metal concentration analyses by atomic absorption.. ...................................................................... ..44
Table 4.1: Total gel times of silicates (m = 3.22 and 2.0) after the addition of 10 WIW % (1.02 M) to 10 mL 50 v/v % silicate with 79.5 mL deionized water giving unstabilized activated sol containing 2 wlw % SiO2. A further 50 v/v % dilution with water would have given stabilized 1.0 wlw % Si02 sol.. ........................................................................... 5 1
Table 4.2: Volumes used for the preparation of stabilized 'activated' silica containing 1.0 wlw %, 2.0 w/w % and 4.0 w/w % SiOz . Sols were stabilized
............... at 50 % of total gel time and derived from silicate with m = 3.2 2... 53
Table 4.3: Times required for complete gelation of sols prepared as in Table 4.2. Effect of pH for 1.0 wlw % SiO2 is also shown.. ........................... 54
Table 4.4: Comparison of analysis methods. Electrode measurements vs. atomic absorption. Sol added as 1.0 wiw % Si02 with m = 3.22
............. ('N' silicate) stabilized at 0.5 gel time (total gel time 18 minutes).
Table 4.5: Unadsorbed copper measured after 20 minutes and one day following addition of 1 .O wfw % SiOz sol ('N' silicate) stabilized at 0.5*gel
..................................................... time (total gel time 18 minutes).
Table 1.6: Copper concentrations obtained after a single 'adsorption/ desorption' cycle. 100 mL initial volumes at pH = 5.5.20 rnL sol (1 .O WIW % Si02, stabilized at 50 % of gel time) + NaOH for pH adjustment. Concentrated acid was used for desorption.. ............................................. 76
Table 4.7: Copperliron (femc) separation by selective iron adsorption on stabilized sol (1.0 wlw % Si02, stabilized at 50 % of gel time). Initial CulFe solution at pH 3.09 preneutralized with lime to pH 3.20.. ............................... 78
Table 5.1 : Femc recovery from 100 mL of 5 5.8 to 1 1 1 7 ppm Fe feeds using lime or acid for pH control ............................................................ 8 1
xv i i
Table 5.2: Ferric recovery from 100 mL 55.8 to 1,117 ppm Fe feeds using NaOH or acid for pH control.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 84
Table 53: Femc recovery fiom 100 mL 1 1.17 g/L Fe feed at pH 1.59 using lime (28 gR. Ca0 suspension) for pH adjustment to 2.3.. . . . . . . . . .. . . . . . . . . . . . . . . . . . .. 86
Table 5.4 : Femc recovery tiom 100 mL 1 1.17 g/Ji Fe feed at pH 1.59 using lime (28 g/L Ca0 suspension) for pH adjustment to 2.8.. . .. . . .. . . . . .. . . . . . . .. . 89
Table 5.5: Underflow properties corresponding to settling tests ofFigure5.6 ................................................................................. 96
Table 5.6: Undefflow properties corresponding to settling tests of Figure 5.9.. . ... 100
Table 5.7: Iron recovery/mdissolution cycles applied to 558 ppm Fe feed at pH 2 .45 . . ... ...... . .. .. .... ...... . .. . ... .. . ... .. ......... . ... .. ...... .. ... .... .-. ...
Table 5.8: Iron recovery/redissolution cycle applied to 1.12 g/L Fe feed at pH 2.27.. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Table 5.9: Iron recovery/redissolution cycles applied to 1 1.17 g/L Fe feed at pH 1.59 .......................................................................... ....
Table 6.1: Dependence of nickel recovery and nickellrnagnesium separation on pH afier 24 h reaction time. 100 mL feeds: 3.0 g/l Ni, 10 g/L Mg, 0.0 g/L Fe, pH 3.0 .........................................................................
Table 6.2: Dependence of nickel recovery and nickelfmagnesium separation on pH after 24 h reaction time. 100 mL feeds: 2.0 g/l Ni, 4.0 g/L Mg, 0.10 g/L Fe, pH 3.0 ................................ .......................................
Table 6.3: Dependence of nickel recovery on pH after 1 h and 24 h reaction times. 100 mL feeds: 2.0 g/l Ni. 4.0 g/L Mg, 0.0 g/L Fe. pH 3.0 .... .. .. .... .. .. .. ..
Table 6.4: Dependence of nickel recovery and nickellrnagnesium separation on reaction time at pH 8.5. 100 mL feeds: 3.0 g/l Ni, 4.0 g/L Mg, 0.1 g/L Fe, pH 3.0 ..........................................................................
Table 6.5: Percent feed volume rejected to the effluent after 30 minutes of free settling and before underfiow centrifugation for tests of Figures 6.3 and 6.4. Feeds contained 2.0 g/L Ni, 4.0 g/L Mg, 0.0 or 0.1 gR. Fe ............................. 133
Table 6.6: Cornparison of a nickel recovery/redissolution cycle for synthetic and sludge leach solutions. Synthetic solution feed: 2.0 g/L Ni, 4.0 g/L Mg, 100 mg/L Fe, pH 3.0. Sludge leach solution feed: 3.1 g L Ni, 5.2 g/L, Mg, 7.1 mgL Fe, pH 3.0.. . . . .. .... . 134
CHAPTER 1
Introduction
Environmental contamination by aqueous solutions of heavy metals (Le. copper,
iron, nickel, cadmium, lead ...) from industrial waste strearns such as mine effluents and
acid mine drainage remains a serious problem. Despite significant advances in the
treatment of such strearns in recent years, there is no satisfactory method for removal of
heavy metals and other toxins such that environmental regulations are completely
satisfied.
Acid mine drainage and mine effluents are universally treated by mixing with lime
and flocculants to a pH between 9 and Il . A voluminous sludge is formed containing
metal hydroxides, prirnarily ferric hydroxide. This procedure is becoming undesirable
since the sludge is not ultimately stable due to a steady decline in pH by acid generation
in the sludge which leads to the dissolution of heavy metals. These metals affect the
quality of drinking water and are very deleterious to aquatic life. Also, in some cases, a
sludge carries a significant quantity of metal which was not recovered during the ore
processing steps. This makes the sludge a good secondary source for metal recovery.
The focus of this work is to develop a new effluent treatment concept that would
supplement the lime treatment method. The new process should be technically feasible,
economically achievable and environmentally safe. An alternative to conventional lime
treatment is to precipitate metals as silicates which are thought to be highly insoluble with
more stable structures than hydroxides (11. Thus, in 1990- 199 1, National Silicates Ltd., a
1
manufacturer of soluble silicates, initiated an ' in-house' project where solubilities of
heavy metals were determined as a hnction of pH upon addition of lime and/or soluble
silicates with various Siofla20 weight ratios. In general, silicate addition permitted
more efficient metal removal at lower pH than achievable using lime only. However, this
process was determined to be expensive compared to the traditional lime treatment
process since soluble silicate solutions are higher in price than lime on a dry basis (almost
10 times).
The present study is a development of the previous work by National Silicates
Ltd., an important objective k ing to Find ways to overcome the cost barrier. A key
feature of this development is the use of silicate in a recyclable form. Direct soluble
silicate addition must operate as a single pass system, since the silicate component cannot
be recovered as a reusable reagent. In the present work, silicate is added as a recyclable,
activated silica sol.
For purposes of clarity, the following bief summary is presented describing the
preparation and properties of activated silica sols ('stabilized silica'), and their interaction
with metal-bearing solutions. A detailed discussion is included in Chapters 2 and 4.
Addition of sulphuric acid to an alkaline soluble silicate solution initiates
polymerization ('gelation') of silicate species, ultimately leading to formation of a gel
which no longer has regular fluid flow properties. This polymerization process can be
arrested (for a long period) by water dilution to make a colloidal suspension containing
stabilized silicate polymers of a given size distribution. This distribution is dictated by the
fraction of the polymerization time at which the process is halted.
When activated sol is added to a heavy-metal solution to reach a controlled pH, a
metal-bearing silica layer will form and readily settle to give a sharp interface and a clear
upper solution. For pH control, lime cm also be added. Metals adsorption ont0 silica sol
is similar to precipitation of hydroxides in the sense that each metal reacts at a specified
pH, thus metals separation into groups is possible. In the present work, it is proposed that
metal is adsorbed by silica sol as M(oH)"' in a manner similar to that generally accepted
for M"+ adsorption ont0 silica particles (i.e gel), which occurs under more acidic
conditions than with sol.
Since metals adsorption ont0 silica sol is pH dependent and reversible (as with
silica gel), acidification of the metal-bearing silica layer (after discarding the upper phase)
with concentrated sulphuric acid solution leads to metals desorption. After centrifugation,
a concentrated metal solution is produced and silica sol is regenerated as a separate layer
for recycle back to the metals adsorption step. When lime is added for pH control during
the adsorption step, the term 'recovery' (or 'removal') is used instead of adsorption, and
redissolution instead of desorption. This is because lime usage produces hydroxide
precipitate in addition to metal ion adsorption ont0 sol.
This thesis focuses on using the metals recovery/centrifugation/redissolution cycle
such that recyclable sol is generated and concentrated metal solutions (much higher than
the feed) are produced. Also, when using this system, waste sludge production is much
reduced when compared with the conventional lime neutralization system.
Chapters 2 and 3 consist of a literature review and description of experimental
procedures respectively. Initial studies were canied out with copper (Chapter 4)
primarily because of the availability of a specific ion electrode for copper (cupric ions).
niis device permits 'in-situ' measurements of the interaction of copper and activated
silica, and can distinguish 'free' dissolved copper (CU*) from copper adsorbed by the
silica sol. Thus Chapter 4 deals mainly with f'undamental issues, and is not related to a
particular industrial application. The dependence of copper adsorption ont0 activated
silica polymer size is studied along with the response of silicates with different
Si02/Naz0 weight ratios. Chapter 4 also includes viscosity and surface tension
determinations for activated silica, along with results from zeta potential measurements
(colloid surface charge). Finally in Chapter 4, copper adsorption/centrifugation/desorption
cycles are demonstrated along with an example of metals separation (cupriclfemc) by pH
control.
Chapters 5 and 6 deal with potential applications to specific industrial problems.
Iron is the major contaminant in many effluents and must be removed. Chapter 5
describes femc removal from 50 to 1 1,000 ppm iron feeds using activated sol (1-4 weight
% SiOz). Free settling characteristics were determined, since in an industrial process,
thickener/clarifiers would be required. In addition to synthetic solutions. results were
obtained for an industrial effluent taken from the 'Balsam Street' effluent treatment plant
of Inco Ltd.. Rates of ferric removai and redissolution were rneasured as a function of
temperature to simulate seasonal changes. Chapter 5 finishes with results for a system of
3 consecutive cycles which progressively upgrade a 50 ppm iron feed to a 40 gR iron
solution, a product that has potential use in sewage treatment.
Chapter 6 studies a possible role for activated silica in treating a tailings pond
sludge. The particular sludge treated in the present work originated fiom an Inco Ltd.
tailings pond, and was high in nickel content. This material was first leached with
sulphuric acid at about pH 3.5, which generated an iron-fiee, dilute (2-3 g/L) nickel
solution with magnesium as the major impurity. Treatrnent with activated silica showed
that a relatively concentrated nickel solution (20 g/L) could be produced in one
recovery/redissolution cycle, which also could separate nickel from magnesium by pH
control. Results were compared with synthetic solutions.
Finally Chapter 7 outlines the most important conclusions and claims to
originality, while recomrnendations for future work are summarized.
CHAPTER 2
Literature Review
2.1 Introduction
In addition to lime neutralization treatment. silicon. the second most abundant
element in the earth's cmst [2] is also used in its various forms (silica gels. soluble
silicates, etc..) for the removal of metals from effluents. Siiica cycles in nature have been
described by Siever [3] and extensive literature reviews for silica including data on
properties such as, phase diagrams. pH. density. viscosity and solubility were assernbled
by Iler [4,5], Vaii [6 ] , Falcone [7], and Weldes and Lang [8].
In this Chapter, a brief review of problems and alternative solutions to metals
rernoval from effluents using lime is given. General usage of siiica and silicates.
especially for effluent treatment. are summarized. then the chemistry of silica and the
production of soluble silicates are reviewed. Since the system studied in the present work
involves interactions of dissolved metal species ( ~ e - . CU". ~ i - and M~*) with
activated silica sol polymers in solution. the structure and preparation of activated sol are
described. Also. descriptions of charge developrnent on surfaces (oxides, hydroxides and
silica) and adsorption mechanisms for dissolved metal-silicate systems are given.
2.2 Problems Associated with Meta1 Precipitation as Hydroxides
Presently. the most widely used method for the removal of heavy metals (e.g.
CU*, ~ i * , Znw, ~b*) fiom effluent streams is by neutralization of acid to a pH at which
metal hydroxides precipitate. Extensive reviews for the treatments of acidic minera1
emuents cm be found in the literature such as in Skelly and Loy. 1973 [9]. and Ritcey.
1989 [IO] In the case of acidic solutions, hydrated lime (Ca(0H)z) is used since it is
considered to be the best and most economical technology available. Since most effluent
streams contain several metals and each metal has a pH at which it is l e s t soluble. as
shown in Figure 2.1 [ I l , an optimum pH is selected (from pH 9 to 1 1 ) when treating an
effluent where a compromise minimum solubility of al1 metal hydroxides is achieved. A
mixed metal hydroxide sludge is then precipitated and the supernatant liquid is
discharged (or reused).
However. due to the amphoteric nature of heavy metals (hydroxides can act as
acids and bases). heaw metals present in sludges (i. e. landfills and lagoons) redissolve
and enter ground water when encountering a pH range other than 9 to 1 1. Afier several
years in a landfill. in most cases a disposed sludge will have a low pH ranging from 5 to
6.5 due to the acidity generated inside the sludge as a result of many factors (Le.
oxidation of sulphides, presence of acid foming bacteria found in a landfil1 during the
normal decomposition of organic materials). Therefore, when solid wastes are disposed
of. they are not ultimately safe since safe disposa1 means stable sludge for a very long
penod of time. A~so, when treating plant wastes, additional chernicals such as floccuiants
must be added to ensure complete settling of solids. These additives certainly add to the
already existing probiem.
Recently, environmental regulations require that most heavy metals in
process effluents be reduced to < 1 ppm before discharge as seen in Table 2.1 [ I l ] . These
new low levels are not easily achievable in many cases with the traditional methods.
Figure 2.1 11 1: Solubility of hydroxides versus pH.
Table 2.1 (1 II: Provincial effluent quality limits.
Parameter (ml&)
As (total)
1
As (trivalent)
CU
Cd
CN
Fe
Pb
Hg
Ni
Zn
2.3 Im provements/Alternatives to Hydroxide Precipitation Processes
During lime neutralization processes, separation of the arnorphous hydroxide
sludges from water is the most difficult task. Over the years, many techniques were
developed to improve solid/liquid separation to produce underflows with higher percent
solids' content (fiom 2 % to 30 %). These techniques are generally sumrnarized as
shown in Figure 2.2 [12] where alternatives: a) Direct lime addition. and b) Lime addition
with air (for oxidation) followed by a settler/clarifier are self-explanatoiy. Option c), the
high density sludge process, requires lime and air additions to pH 7.7 (with agitation),
New- foundland
0.5
0.3
0.05
0.2
0.005
0.5
Quebec
0.5
0.3
0.001
1.5
3 .O
0.2
0.00 1
0.5
0.5
Ontario
0.5
0.3
0.001
1 .O
1 .O
0.2
0.5
0.5 L
Saskatchewan
0.5
0.3
1 .O
0.2
0.5
0.5
British Columbia
0.10-1.0
0.05-0.25
0.05-0.3
0.0 1-0.1
0.1-0.5
0.3- 1 .O
0.05-0.2
N IL-0.005
0.2- 1 .O
0.2- 1 .O
New Brunswick, Nova Scotia, Manitoba
0.5
0.3
I
1
0.2
4
0.5
0.5
followed by flocculant addition in a settler/clarifier to get good settling, then the
underflow (sludge) is recycled up to 25 times to get a good final solids' content. The
third route, even though expensive, is more efficient in terms of lime utilization, has good
filtering and settling propties, and produces sludges with much higher densities (20-30
%). Finally, in some cases the use of vacuum filters for dewatering lime sludges can
generate a cake containing up to 45-50 % filter cake.
c) Type 3: Lime addition to neutralization reactors. Solid'liquid separaiion in chrifiers.
-
Figure 2.2 (121: Lime neutralization process alternatives.
OVlntLOW
SLUDQI CON00
a) Type 1: Lime addition to effluent or tailings lines.
in an attempt to solve the problem of stability of existing hydroxide sludges
(tailing sites), improvements have k e n studied such as the use of artificial covers. These
covers prevent oxygen penetration thus reducing acid generation which results in metal
dissolution. CANMET developed a program that predicts acid generation in sulphide
tailings with or without the use of coven [13]. For hydroxide sludges being currently
discharged, special lined Iagoons and basins are used as dumping sites where the acid is
collected at the bottom and treated (101. However, these solutions are not optimal since
maintenance is aiways required.
Many alternative techniques for the removal of heavy metals from effluents can be
found in the literature. For example, since metal hydroxide precipitation is pH
dependent, metals can be recovered by selective precipitation as proposed by Jenke and
Diebold [14]. The use of sodium hydroxide can generate a cake with about 20 % solids
which does not contain gypsurn [15,16]. Finally, silica and silicates have been used as
discussed in Section 2.1. However. these methods suffer from several disadvantages
(i.e. expensive, difficult to control on a large scale).
Other aitemative technologies are the use of liquid surfactant membranes (LSM)
[17-201 and ion exchange resins [21,22] for the r.xtraction and separation of metal ions
such as CU", ~ i " and ~ n * fiom solution which have been suggested as alternatives to
lime treatment [23]. However. these processes are not realistic alternatives simply
because of the high tonnage of mine effluents, low capacity of such processes, and their
high cost of investrnent and maintenance. Finally. these methods treat only specific
selected ions or complexes. Such processes are only viable when treating value-
containing and/or radioactive waste.
2.4 Metals Removal from Effluents using Silica (quartz, silica gels,
silica resins or filters)
Silica, silica gel resins [24], silica gel filters [25], and morphous silicate gels [26-
281 have been used for the removal of heavy metals from certain effluents. Also, for the
treatment of nuclear wastes, silica gels are used for the removal of uranyl and neptunyl
cations [29]. The production of silica gel particles of a controlled size is described by
various authors [5,7]. However. many problems are associated with the use of "pure"
silica products, where it is known that these processes are relatively slow and silica gel is
not only fragile but is also expensive. Also, since silica gel is a weak acid, the solution
pH declines as metals are adsorbed therefore desorption becomes favoured unless pH is
controlled. Finally, these methods do not solve the problem of waste production.
The structural stability of silica gels and resins may be enhanced by adding
various chernicals which become bound to the gel or resin structure during preparation.
The production of finely divided silica gel particles bound to quatemary ammonium
compounds (e.g. triaikylbenzylammonium, polydimethyldiallylarnmonium) is described
by Cotta et al. [30.31]. A summary of many complex forming reagents supported on
silica gel was published by Terada [32]. Examples of the application of silica-based gels
and resins for heavy metals adsorption include the use of dipicolylarnine bound to a
silicate support [33]. calcium phosphate or pyridylazo naphthol adsorbed on silica gel
[34.35]. tin and titanium amorphous silicates [36]. magnesium silicate [37] (to prevent
beer contamination with metal ions) and amino acid (glycine) addition to silica gels [38].
Finally for gaseous wastes, odorous gases (H2S and NH3) are removed from gaseous
effluents by a combination of silica gel-alkaliheavy metal salt [39].
2.5 Production and Usage of Silicate Materials
Sodium or potassium silicates have been produced since the seventeenth century.
The production of soluble silicates was reviewed by Williams [40]. Sodium silicate
products are manufactured by reacting sand and soda ash in different proportions at high
temperature via the process shown in Figure 1.3 [4 1 1. The reaction is:
m Sioz + NarCO3 -, CO: + (SiOz)m.Na20 (3.1)
Different proportions of sand and soda ash will result in different end-product ratios of
SiOzMazO. Physicai and chernical properties depend on the composition.
Liquid sodium silicates are solutions of water soluble glasses which have a weight
ratio (m) of SiOzNazO ranging from 0.5 to 3.25. Generally soluble silicate products do
not exceed a SiOtMa2CO3 molar ratio of 4 because of the very low solubility of fused
silicates above this ratio. The price for soluble silicates increases as the ratio m decreases
from 3.22 to 2.00 which makes the silicate with m = 3.22 the most suitable for many
industrial applications, and especiaily for effluent treatment.
b
Dissolver b
Gnnd and b
classi& J - - Soiutions -p
v Anhydrous gIass 4 NaOH
powdcr v Hydrous
CrystaIiize. *. Pau* I . J h b t
blend ~iquid silicate *O treatment
1 1 1 liquid poducy
(hthosilicrit Metasilicak
Figure 2.3 1411: Process for production of sodium silicate products.
The properties and general usage of soluble silicates are sumrnarized by Weldes
and Lange [8]. Applications for heavy metals removal from waste watcr solutions
include the addition of silicates with no other additives [42] or with further pH
adjustment [43], and the use of silicates and flocculant as a second stage after lime
treatment [44]. Another method of treatrnent is the use of silicates with a setting agent
(Le. Chemifix process) to produce a highly insoluble solid matrix [1.45-491. This process
is presentiy in use. but again. none of these techniques eliminate the problem of waste
production.
Silicates are versatile products with a wide range of applications outside metals
removal from waste waster. Sodium silicates are used in detergents to enhance the
performance of the detergent system. Silicates, with a ratio of SiOz/NazO ranging fiom
2.8 to 3.2, are used as adhesives or binders due to their conversion frorn liquid to solid
with only a small rernoval of water. Also, when these silicates are combined with cernent
they react chemically forming products having strong bonding properties. and are
sometimes used to treat fresh concrete surfaces for longer durability.
In coatings, dried films of silicates are unaffected by fats, greases and oils and are
fire and water resistant. Silicates, when converted to activated sifica sols. are used in
conjunction with alum. femc salts or other coagulants as floccuiants to promote removal
of fine suspended solids from industrial and municipal water supplies. Sodium silicates
mixed with portland cernent make good self hardened foundry molds. In ore
beneficiation. sodium silicates are used both in notation [50] to disperse undesired
siliceous particles, and in ball milling to reduce ball Wear due to corrosion.
2.6 Aqueous Chemisty of Silica and Silicates
Silica is the origin of al1 silicate solutions. The dissolution of silica can most
simply be represented as follows [5 ] :
Si02 + 2 H20 -, Si(0H)r (3.2)
Reaction (3.2) occurs most readily in alkaline solution in which the hydroxyl ion acts as a
catalyst. The hydroxyl ion chemisorbs on the silica surface and increases the
coordination nurnber of a surface silicon atom to more than four. thus weakening its O-Si
bonds (Figure 2.4 [5.5 1 j) and creating a negatively charged surface. The silicon atom
then initially dissolves as the species Si(0H)s- as shown in Figure 2.4. Si(0H)r- is in
equilibrium with soluble Si(0H)d and OH - ions as follows:
Si(0H)s F Si(0H)r + OH -
Si(0H)r predominates at pH below about 1 1.
Many authors [52-601 have show that various factors influence the rate of
dissolution such as pH, temperature, type of silica used (i.e. degree of crystallinity). prior
mechanicaiheat treatment. previous exposure to water, alkali, or acid, mass of silica in
solution, particle size and nature of the electrolyte. In amorphous silica. the structure is
an open arrangement where large spaces between oxygen atoms exist on the surface to
accommodate hydroxyl ions. which promotes more rapid dissolution.
Figure 2.4 [5,511: Proposed mechanism of dissolution of silica in water in the presence
of hydroxyl ions. Doned lines represent the interface between silica on the left and water
on the right.
2.7 Structure of Monosilicic Acid
The soluble form of silica at low (< 0.002 M) concentrations is monomeric and is
usually fotmulated as Si(OH)4. This is referred to as monosilicic acid or orthosilicic acid.
The structure involves one silicon atom coordinated with four oxygen atoms as in
crystalline quartz. The hydration state is not accurately defined although at high pressure
therc is an indication that one water molecule is linked to each OH group probably by
hydrogen bonding (Si(OH:H20)4) [6 1 1. However, silicon is coordinated with six oxygen
atoms in rare minerals such as stishovite [62].
Figure 2.5 [5.63,64] is an equilibrium diagrarn showing the dissolution behaviour
of arnorphous silica at 25'C. It is a good approximate description of the system at
different concentrations and pH's. At below 0.002 M SiOz, nonionic Si(0Hh
predorninates in neutral and weakly acidic or weakly alkaline solution. Si(0H)s- becomes
important at pH > 1 1. At above 0.002 M, the insolubility domain of Figure 2.5 is entered.
and Si02 precipitates. In alkaline solutions (pH > 8). a stable multimeric domain exists in
which monosilicic acid polymerizes initially fonning polysilicic acids of low molecular
weight, as described in Section 2.9. Under these conditions (i.e. within stable multimeric
domain of Figure 2.5) the term soluble silica usually includes monosilicic acid and low
polymers such as tetramer and decarner. At higher concentration above the upper lirnit of
the multimeric domain, larger polymenc species are found consisting of colloidal
particles. Colloidal silica comprises high molecular weight polymers in the form of
particles larger than 20-50 A (molecular weight as Si02 of 40,000-100.000). Polysilicic
acid and/or coIloidal silica are loosely referred to as silica sol. The solubility of
monosilicic acid and its polymerization rate both increase rapidly with increasing
temperature. klthough Si(OH)4 i s non volatile at room temperature. the presence of
monosilicic acid in the vapour phase at elevated temperatures has been reponed [6?.66].
Stable multimeric domain
r / * e O 5 10 - lnsolubility domain 2 u e 4, U
7" C
10-2- L a œ rn rn - - - /y-" k c l r w a l l
8 Mononuclear domain
Figure 2.5 [5,63,61]: Solubility diagram for amorphous silica at 25 O C .
2.8 Water Soluble Silicates
A water soluble silicate such as sodium metasilicate. NazSi03 ((SiOzNazO) = 1)
dissolves in water giving an alkaline solution in which the following equilibcîa are
O bserved:
~ i 0 3 ' - + ~ 2 0 F HS~OJ- + OH-
Equilibriurn constants for these reactions were measured by Roller and Enin [67] at
25 OC. Soluble silicates are generally those of potassium or sodium. The major ions
present in solution obviously depend on concentration of added silicate and pH. For
example. a dilute solution of sodium silicate (0.002-0.02 M SiO?) with a pH ranging fiom
8 to 1 1, consists rnainly of rnonorneric s~o,' '. HSi03-. siz0s2 -and Si(OHk generated by
the above equilibria. As the concentration of added silicate increases (or a 'higher'
silicate, such as Na2Siz05 with SiO?/Na?O of 2/1, is dissolved), the additional
silica/silicate ions tend to fom extremely small three-dimensional. intemally condensed
polymer-ions which. with fiuther concentration increases. become negatively charged
colloid particles.
The size of polyrneric colloidal species in sodium silicate solutions with rn = 3.22
(6.67 M SiOz. pH 11.3) was estimated by Bacon and Wills [68] to be 11 4 (molecular
weight of 930 with about 15 Si atomdparticle). The bulk of the present expenmental
work was carried out using silicate solutions with m = 3.22. Debye and Nauman [69]
used light scattering measurements to determine that the average molecular weights for
silicates with m = 2.03 (7.52 M SiOl, pH 12.4) and 3.32 were 1 50 and 325 respectively.
Iler [ 5 ] noted for example that in a silicate solution with m = 3.22, 39 % of silica is
polyrneric while 61 % is monomenc at 6.67 M Sioz and pH 11.3 which gives a position
on the edge of the stable multimeric domain of Figure 2.5.
2.9 Polymerization of Silicic Acid
Study of silica sols and gels began in 1865 by Graham [70,71], but Carmen [72]
was the fint to clearly state that "silicic acid polyrnerizes to discrete particles which then
aggregate into chains and networks". Baylis [73-761 showed that addition of sulphuric
acid to soluble silicate solution initiates polymerization and c m produce sols which could
be added in conjunction with alum to improve the quality of water. The use of this
addition to drinking water has been reviewed by several authors [77-791 and very recently
by Gibson [80].
Silicate polyrners, within the stable multimeric domain of Figure 2.5 grow in size
upon contact with a minera1 acid. This is the process of polymerization (gelling).
Eventually, the solution becomes one single polyrner at which point a gel has fomed. A
silica gel has been described [81] as " a heavily hydrated intrrlaced fibrillar or brush heap
structure of very large polysilicic acid molecules with the spaces filled with water or
dilute silicate solution". During polymerization. the electrostatic repulsion of negatively
charged polysilicate sol species is outweighed by sorption effects that permit growth and
flocculation. Figure 2.6 [80] shows the structure of a charged silicate polymer. Figure 2.7
[ 5 ] shows the differences between particle configuration when in a sol (separate
particles), gel (linked chains), or precipitate (discrete aggregates).
Figure 2.6 [sol: Structure of a charged silicate polymer.
Sol Gel Precipitate
Figure 2.7 15): Sol vs. gel vs. precipitate.
The self condensation of monomer is catalyzed by the OH - ion and is commonly
witten as [j]:
2 Si(OH)r (HO)> SiOSi(OH)3 + &O (3-8)
In a structural sense, this cm be shown as:
Rates for reaction (3.8) are given by Iler [SI in the pH range from 2 to 10. The order of
reactiûn 3.8 is second order in Si(OH)4 at pH > 2. and third order at pH < 2 [j].
The polymerization reactions were also described by Lange et al. [8 11 as:
Lange et al. [8 11 also proposcd that rcaction (3.10) contebutes to polymerization:
Here Si0 is a functional group presurned to exist in alkaiine solution at the surface of
colloid particles.
Reactions (3.8) and (3.9) are the condensation of silanol groups to siloxanes with
the production of water molecules. Reaction (3.10) is the silanol group and a silanolate
anion condensing to siloxane with the production of a hydroxyl ion. The relative
importance of nactions (3.8) or (3.9) and (3.1 0) depend on the pH at the start of the
polymerization process. Figure 2.8 [8 11 shows that if gelation is initiated at about pH 8.3,
the pH rises becaw reaction (3.10) is predominant. At a lower initial pH of 6, the pH
nmains essentially constant, suggesting reaction (3.8) or (3.9) is dominant. In general,
sois an prepared under alkaiine conditions where reaction (3.10) predorninates.
5.00 ---- 1 l I 2 5 10 20 50 '00
Contact time (minutes)
Figure 2.8 (811: Change in pH with time afier addition of minera1 acids to sodium
silicate (Si02MazO. m = 3.22 by weight, 20g SiOzL).
2.10 Gelation
The continued progression of polymerization eventually leads to gelation. Total
gel time is lengthened as one goes fiom hydrochloric to sulphuric and then to phosphoric
acid at constant pH and silicate concentration as s h o w in Figure 2.9 [81]. The primary
difference between these systems is the neutral salt produced. Here, sodium chloride is
much more effective than sodium sulphate or sodium phosphate in reducing the surface
charges that inhibit polymerization [8 11, although the mechanism is unclear.
Gel time depends not only on acid type and pH but also on silica concentration.
Figure 2.10 [82] is an example of how gel time varies with both pH and silica level. As
Figure 2.9 1811: Gel times of sodium silicate neutralited with different acids.
(SiOuNazO, rn = 3.22 by weight, 20g SiOfi).
Figure 2.10 1821: Gel times at 25 OC for various concentrations of sodium silicate
(SiOz/NazO, m = 3.22 by weight) when mixed with sulphuric acid to give initial pH
vaiues as shown.
expected. an increase in silica concentration at a given pH gives decreasing gel times.
The fastest polymerization. or gelling, rate should occur when the repulsion forces
between colloidal particles are minimal. Surface charge development models have been
proposed for insoluble oxide and hydroxide suspensions in an aqueous medium whereby
particle surfaces tend to become negatively charged under alkaline conditions and
positively charged in acidic solutions [83-851:
M (OH)surfa~.il + Hz0 F MO - SUT^ + ~ 3 0 ' (3.1 1)
iM (0H)surrarr + &0+iq F M(OH2)+ + Hz0 (3.12)
In reaction (3.1 l), surface hydration is followed by dissociation of surface
hydroxyl groups to yield a negatively charged surface in alkaline solution. while in
reaction (3.12). acidic conditions lead to protonation of the surface. The point of zero
charge (PZC) is defined as the pH at which the surface charge is zero. The PZC is not
necessarily at pH 7. For exarnple. the PZC has been reported as being at pH 4.8 for
n a t d a-Fez03 (hematite) and at pH 8.8 for amorphous femc hydroxide [83]. By
analogy with oxidehydroxide suspensions. it would be anticipated that surface charges on
silicate polymer particles are determined by the pH which also influences the progress of
the polymerization reactions themselves. Figure 2.10 suggests that the various reactions
leading to surface charges produce a net minimum charge at pH between 7 and 8, where
gelation is fastest.
Whenever the surface acquires a finite charge. a potential with respect to the
solution is established. The surface charge is compensated by an equal and opposite
charge distribution in the aqueous phase. The charge in solution facing the surface
charge is called the electrical double layer. A schematic representation of the double layer
and potential &op across it are presented in Figure 2.1 1 [86,87]. The potential
determining ions are at the surîàce, a layer of counter ions are anchored at the outer side
of the surface (Stem layer), then counter ions are arranged diffisely away from the
surface in solution. The counter ions are attracted to the surface by electrostatic
attractions. The potential difference between the surface and bulk solution is called the
total surface potential, (Nemst potential). The elecvical potential drops linearly from the
Nemst Potential (\Yo), to a value called Stem potential (Y*). then beyond this point, it
decreases exponentially with distance. The potential difference between the closest
distance of approach of hydrated counter ions to the surface (plane of shear on Fig. 2.1 1 )
and bulk solution is called the zeta potentiai, Zp, which can be measured.
Following acid addition to an alkaline silicate solution. polymerization. or
gelation, can be effectively halted by dilution, which transforms an activated sol to a
"stabilized sol. Dilution greatly reduces the polymerization rate because of the
decreased silica concentration. A gel may or may not eventually form depending on the
particular silica concentration/pH combination obtained (Fig 2.10). Stabilized sols
derived fiom dilution following sulphuric acid addition are apparently more stable than
those generated by other mineral acids [8 1,821. Dilution is often practised at half the gel
time when using the sol as a flocculant for drinking water clarification, where finely
suspended solids are removed. Figure 2.12 [80] shows the simple arrangement used to
produce stabilized sol for this latter application [80]. As part of an experimental study to
Q Stem layer y- + - 1 + + (Goy Diffise laper) layer
++l , ; -,>:me ofshear
+ 9
a) Disuibution of charges in the vicinity of a colloida1 particle.
Distance from particle surface
b) Distribution of potential in the clectncal double layer.
Figure 2.11 186-871: The electrical double hyer (Stem modei) and potential at a
particle surface in solution.
optimize stabilized sol production for water clarification. Gibson [80] showed ihat a 1.5
W/W % Si02 sol (SiOz/Na20 weight ratio 3 22) stabilized at 20 % to 50 % of its total gel
time has polymers with sizes ranging fiom 1 xl O" to 1 xl O" meters. This is of particular
relevance to the present work which is mainly based on the sarne silicate (SiOdNa20
weight ratio 3.22) also diluted at 50 % of gel time giving silica contents in the 1 :O to 4.0
w/w % range. The preparation of stabilized sols under a variety of conditions is also
described in several patents [88-PO].
Sulfuric Sulfuric acid / acid
silicate silicate Sodium Sodium
Stabilized "activated" silica
Figure 2.12 1801: Production of "stabilized silica sol for water clarification.
2.1 1 Dissolved MetaUSilica Sol Interactions
When silica sol is added to an acidic metal-sait solution. the pH increases and
adsorption sites are created for dissolved metal species. With a continuously increasing
20
pH, the reaction path followed by dissolved metal ions will pass through several
intennediate hydroxy complexes before reaching the point where precipitation would be
expected. ller [91] and Falcone [92] observed that sodium silicate solutions begin to
adsorb dissolved multivalent metal ions (in chloride or perchlorate solutions) at pH
values up to two units below the equilibrium pH at which metal hydroxide should .
precipitate. In these latter studies [91,92], the silicate solutions had not been deliberately
'activated' as in the present work, but some degree of natural polymenzation would have
taken place reflecting the particular conditions in the sodium silicateimetal-salt solutions.
Direct evidence of adsorption ont0 silicate polymers is given by Falcone [92],
who compared dissolved sodium hydroxide and silicate with various weight ratio
SiO:/Na?O when added to perchlorate solutions. When silicate (or NaOH) was added to
dissolved copper perchlorate at pH 4. there was a rapid initial reaction followed by an
apparent steady state after about 10 minutes. Figure 2.13 [92] shows the pH values
reached in this metastable condition following different volume additions of silicate (m =
3.8) or sodium hydroxide. Dissolved cupric ion levels were rneasured using a specitic
ion electrode for copper. Figure 2.13 also gives equilibrium cupric ion activities for the
following reaction:
CU" + Hz0 ZCuO, + 2 H ' (2.13)
Reaction (2.13) was taken by Falcone [92] to represent cupric precipitation as a function
of pH.
There are several important conclusions associated with Figure 2.13. Firstly,
while a precipitate fomed upon neutralization with sodium hydroxide. there was no
visible precipitate when using silicate. Secondly, a horizontal line drawn on Figure 2.13
readily demonstrates that silicate would remove dissolved copper at a lower pH than
would sodium hydroxide. Figure 2.13 also shows that residual dissolved copper levels
are much higher than the equilibrium values corresponding to precipitation. In fact
cupric ion activities following sodium hydroxide addition are even hirther displaced fiom
the CU"/CUO equilibrium than obtained using silicate, which shows the high degree of
supersaturation needed for precipitation.
After the silicate sarnples represented by the data of Figure 2.13 had been aged for
one rnonth, residual cupric activities decreased to levels that approached those for the
CU"/CUO, equilibrium. and a precipitate appeared. 1t should also be noted that the
nature of the precipitate was not examined by Falcone. although the precipitate colour
was observed to change with time and the SiO?/Na20 ratio of the silicate used.
Figure 2.13 1921: Copper ion activity vs. pH afier the addition of sodium silicate or
NaOH solution to Cu(C104)2 solution at pH 4.0.
2.12 Adsorption in Dissolved MetaU'Silica' Systems
Similar to most worken [93-961, Dugger et al. [97] and Schindler et al. [98] have
represented metals adsorption ont0 a 'silica' surface as:
I I A (n-m) + M n + + m ( - S i O H ) r M ( 0 S i - )
I + mH' (2.14)
I m
The tem 'silica' is used here as a convenience. because in reality the reacting specics is a
surface silanol grooup. Funhermore, reaction (2.14) has b e n proposed regardless of the
starting siliceous material (i.e. quartz, sodium silicate solution, silica gel powder).
because surface silanol grooups always form upon contact with an aqueous medium.
Reaction (2.14) is the sum of the following wo rerictions:
Vydra and Galba [99] have made an obvious refinement to reaction 2.14 and presented
the adsorption process as:
Here. L is a complex forming ligand.
Reaction (2.14) (or (2.17)) is a revenible pH-controlled ion exchange process.
32
Dugger et al (971, working with silica gel powder, have carried out extensive studies of
the stoichiornetry of reaction (2.14) when applied to a wide range of metals dissolved in
buffered nitrate or perchlorate solutions. Adsorption was camied out at low pH where n
could be taken as the charge of the simple unhydrolysed metal ion. Under these latter
conditions, the quantity of H' ions released per mole rnetal adsorbed was always such that
n = m in reaction (2.14) when at equiiibrium. There are no known exceptions to ihe n =
m rule.
Vydra and Galba [99] studied the stoichiometry of reaction 2.17 at pH levels
where hydrolysed metal ions would be expected to form. Metais were adsorbed ont0
silica gel powder from chloride or nitrate media buffered with solutions prepared from
chloroacetic, formic or acetic acid. By assuming the n = m rule applies to reaction 2.17,
Vydra and Galba were able to deduce the charge on the adsorbed complex through
measurements of CMJCH where C M ~ and CH are the moles of rnetal adsorbed and
hydrogen ions released respectively. For example. when dissolved aluminum is adsorbed
at pH between 3.8 and 5.8, CMJCH is 0.50 suggesting adsorption of a divalent cation.
Sometimes, CMJCH is a fractional number other than 1.0. 0.50, 0.33 etc. so indicating
adsorption of at least two cations of different valencies.
In principle, adsorption could also result from electrostatic interaction between I
dissolved cations and negatively charged silanolate. - S~O-, surface sites. However, the I
experimental evidence of Dugger et al. [97] and Vydra and Galba [99] conceming the
stoichiometries of reactions (2.14) and 2.17 strongly suggests ion exchange, not
electrostatic interaction, is the primary adsorption mechanism.
In a thermodynamic study of the ligand properties of surface silanol groups on
silica, Schindler et al. [98] showed a close correlation between the stability constants of I
the surface complexes fonned with metal ions (i.e. M( OSi-)m ("-w in reaction (2.14)) and 1
the corresponding values for simple metal hydroxy complexes. Schindler observed that
'this correlation indicates the ligand properties of the surface OH groups on silica are not
basically changed by king attached to silica'. This laner observation was considered
consistent with the occurrence of adsorption and precipitation at similar pH values. With
increasing pH, the silica surface effectively captures dissolved metal hydroxy complexes
just prier to nucleation of hydroxide precipitate.
James and Healy [ 100-1 OZ] studied dissolved metaVsilica interactions and
concluded that metals are adsorbed as cationic metal hydroxy complexes, and these
adsorbed species are separated fiom the surface by at least one layer of water molecules
thus preventing direct chernical bonding that would lead to silicate precipitation. It was
suggested (1 00- 1021 that these adsorbed metal ions do not lose their 'primary hydration
sheaths' and are in the same fom as in the aqueous solution. This latter observation has
been noted by several authors [103- 1061. In a project king carried out concunently with
the present work [107], activated silica, lime and sodium hydroxide were compared as
neutralizing agents for cobaltous sulphate solutions. When using lime or sodium
hydroxide, the original pink colour of the solution disappears leaving a bright blue
hydroxide precipitate in a colourless aqueous phase. In contrast, the pink colour of the
hydrated cobaltous ion is retained when cobalt is adsorbed ont0 activated silica, and the
system separates into a pi& cobalt-'loaded' silica underflow and a colourless overflow.
Palmer et al. [108] studied dissolved copper adsorption ont0 silica, along with
corresponding zeta potential measurements. It was suggested that CU(OH)' was the
adsorbed species, which attached to a silanol group by hydrogen bonding. The proposed
reaction mechanism, shown in Figure 2.14 [108], is inconsistent with the conclusions of
other authors [i.e. 97-99] who have demonstrated the ion-exchange nature of adsorption
on silica whereby hydrogen ions are generated.
The pH dependence of metal ion adsorption ont0 silica may be used to perfonn
crude separations of dissolved metals. Schindler et al. [98] and James and Heaiy [100-
1021 presented adsorption curves as a function of pH for several metais upon removal
from the solution ont0 silica particles (Fig. 2.1 5 [98,100- 1021). Femc (and aluminum
and chromium) are preferentially adsorbed at pH 1.0 to 3.5, 'heavy metal' divalent
cations adsorb at pH 4.0 to 8.5. while dissolved calcium (and magnesium) are removed at
pH 9.0 to 12.0
The most common practical application of activated silica is as a floccuiant. added
at 1-2 ppm Sioz levels, for removal of fine suspended solids from public water supplies.
in the present work. sufficient activated silica is added to raise the pH such that metal
hydroxy complexes are produced. These complexes are adsorbed ont0 a silica sol
substrate of sufficient volume that a discrete layer is formed. This silica layer creates an
underfiow that is readily separable fiom the feed solution. Chapter 4 deals with
fundamental measurements of dissolved copper (as sulphate) interactions with activated
silica when added in g/L amounts. Chapters 5 and 6 examine potential applications for,
respectively, treatment of an iron effluent and a nickel-bearing tailings pond sludge.
- O\ /
OH- - -0 -Cu +
Si H
-O' 'OH
a) H bonding occurs between adsorbed H ions and the hydroxy complex.
/ 0' 'OH
b) Adsorption of hydroxy complex by the formation and rejection of water.
Figure 2.14 (1081: Proposed mechanism for copper adsorption on silica (quartz) surface.
Figure 2.15 198,100-1021: Metais adsorption ont0 silica particles vs pH for various
metals.
CHAPTER 3
Experimental Procedure
3.1 Reagents and Materials
The starting material for the preparation of activated silica sols was a sodium
silicate solution provided by National Silicates Ltd.. Toronto. The solutions used had a
Si0fia20 weight ratio (m) of 3.22 or 2.00. Most work was camed out with silicate
solution of m = 3.22. due primarily to its lower cost, which would be a significant factor in
a commercial application involving effluentlwaste treatrnent. As-supplied silicate with m =
3.22 contained (w/w %) 8.90 % NazO, 28.7 % Si02 at pH 1 1 3. and was designated 'N'
silicate by National Silicates Ltd. Some work was camed out using silicate solution with m
= 2.00 ('Dg silicate), the composition of as-received solution being (w/w */O) 14.7 % NazO.
29.4 % Si02 at pH 12.4. The specific gravities of 'N' and 'D' silicates were 1.39 and 1 5 3
respective..
*.4queous solutions of metals were prepared by dissolving appropriate arnounts of
sulphate saits in deionized water fiom a two-colurnn Cole Parmer ion exchange system.
The required solutions were made up fiom F~?(SOJ)~.~H?O. CuSOa.SH20, NiS04.6HzO or
MgSO4.7H20 either as single metal solutions or as CulFe or NiMg sulphate mixtures for
metals separation tests. Sulphuric acid was used either in as-supplied, 93 w/w % H2S04.
form or after dilution (10 w/w %). Lime was used in the form of a 28 g CaO/L aqueous
suspension. Al1 chernicals were reagent grade (Fisher Sientific Ltd.).
Treatment of two industrial wastes was investigated. One of these was a sludge
fiom the 'upper pond' of the tailings treatment operations of Inco Ltd, Sudbury. As-
received sludge contained 86 wfw % water and the composition of the dried materiai was
(W/W %) 6.2 % Ni, 16.8 % Fe, 20.3 % Mg and 0.2 % Ca. This sludge was leached with
sulphuric acid solution to generate a feed for processing using activated silica. The
second waste studied was an effluent solution (acid mine drainage) obtained from the
' Balsam strearn' that feeds the water treatment plant of Inco Ltd.. This effluent contained
60 ppm Fe, 1.3 ppm Cu, 27 ppm Ni.
3.2 Sol Preparation
A 1.0 W/W % Si02 activated sol ('stabilized sol') derived from 'N' silicate was
chosen as the standard reagent for the bulk of the tests. Activated silica was prepared by,
fintly, diluting 'N' silicate by 50 v/v % with water to lower the viscosity and promote ease
of handling. The pH of the diluted 'N' silicate was 1 1.15. While stimng with a magnetic
stirrer, gelation was initiated by adding 10 w/w % sulphuric acid solution such that a 2.04
W/W % silica solution at pH 8.2 was produced. Agitation was stopped 45 seconds afier acid
addition. This gave a gel time of 18 minutes. The criteria for complete gelation were Ioss
of uniform fluid flow with simultaneous appearance of surface cracks and intemal breakage
planes when the mixture was tilted. Other characteristics of gelation are the adherence of
gel to the glas wall of a tilted beaker, and depression of the surface when a giass rod was
placed on the gel.
Gelation was arrested at various fractions of the gel time by 50 v/v % dilution with
water, so producing a stabilized sol. in most cases, the sol was stabiiized at half the gel tirne
(Le. 9 minutes), to give about 200 mL of product containing 1.0 wlw % silica. Activated
sols containing 2.0 wlw % and 4.0 wlw % silica were also used, and some additional tests
were carried out with activated silica prepared fiom ' D' silicate. The terni 'stabilized' sol
refen to 'activated' sol where polymerization was slowed down (arrested) for a very long
period.
Metal solution
lime Metal recovery onto
activated silica
Activated silica under flow Regenerated
activated silica 1
1 Stabilized silica I j Concentrated
Metal redissolution acid I
1 Concentrated metal Centrifige '
b solution
Figure 3.1: Conceptual flowsheet for metals recovery/redissolution stage using activated
siiica sol (one cycle), - could be almost any metal.
3.3 Metals Recovery Stage
Figure 3.1 shows the conceptual flow sheet used as the basis for the following
explanation. In the recovery stage, activated silica was added incrementally. using an
autotitrator, to about 100 mL of metalîontaining solution with agitation provided by a
magnetic stiner. The pH was continuousIy monitored, and in the case of cupnc solutions,
a cupnc ion-specific electrode was also used (Figure 3.2). Activated silica addition
increased the pH which resulted in metal adsorption ont0 the sol. Metal removal. as
M(oH)"', begins at pH levels up to 1 unit below values for hydroxide precipitation.
Apparent equilibnum was reached within about 20 minutes following each incremental
silica addition. The volume of added sol ranged from 5 mL to 80 mL for 100 mL of
metal sulphate solution, depending on the required final conditions of rnetal/silica ratio
and pH. When necessary, lime addition for pH control was preferred when the alternative
would have been extra silica, since excess silica promotes dispersion and a Mayer system
will not fom. At the end of the recovery stage, agitation was stopped, and a metal-
bearing silica layer (underflow) was fomed and readily settled to give a sharp interface
and a clear upper solution (overflow). The underflow colour depended on the metal
involved (brown for ferric, blue for cupric, green for nickel). The underflow was
centrihiged for 10 minutes at 2,000 rpm using a IEC International centrifuge, mode1 CL.
Sampies of overflow were retained for analysis. The reproducibility of metal
concentration measurements was * 5 % of the rnean value from replicated tests.
Although most experiments were performed at room temperature, a few tests of
femc sulphate/activated silica interaction were carried out at 5 OC and 45 OC to simulate
seasonal changes in outdoor water temperatures. For non-room temperature tests, the
reactor was held in a water bath placed on a magnetic stirrer. Temperature control was
maintained using a circulating water flow which passed through a Cole-Parmer polystat
heaterkooler. Femc sulphate solution and activated silica were stored at bath
temperature prior to an experiment.
' Activated* sol
4 Autotitrator
reference elect rode
cupric electrode - , Activated sol dropkts
7'
l n h n I w V u
stirrerhot plate
Figure 3.2: Metals recovery/redissolution experimental set-up (with the cupric ion
elec trode).
3.4 Metals Redissolution Stage
Up to 90 v/v % of the feed to metal recovery was rejected as 'effluent' after the
settling and centrifuge steps shown in Figure 3.1. Thus in a continuous process. 'metal
redissolution' is a much smaller volume operation than 'metal recovery'. Batch tests
using 100 mL of sulphate solution were satisfactory when studies were confined to metal
recovery and the following settlingfcentrifuge steps. However the amount of centrifuged
silica underflow made available for the redissolution stage was too small for meaningful
study. It was necessary to process up to 1 L batches of sulphate solution through 'metal
recovery' to generate suficient feed for 'metal redissolution'. This increased scale of
operation required a larger centrifuge (IEC international Mode1 K, size 2) than used for
the 100 mL batches.
The experimentd set-up for 'metal redissolution' was the same as for
'metal recovery' except that the autotitrator delivered concentrated (93 w/w %) sulphuric
acid to centrifuged, silica underflow (Figure 3.1). The volume of added acid was up to 5
mL depending on the desired final pH. The reaction rate increased with decreasing pH,
and conditions were usualiy adjusted such that the required retention time was between
20 minutes and 1 hour. By using concentrated acid, the dilution effect inherent in using
an acid solution is avoided, one goal of the overall process being to produce a
concentrated metal solution relative to the feed. Centrifugation of the 'metal
redissolution' products gives a concentrated metal solution overflow and regenerated
activated silica. In some cases, activated silica was recycled (after pH adjustment with
lime) to the metal recovery stage. The reproducibility of metal concentration
measurements was 10 % of the mean value fiom replicated tests.
3.5 Settling Stage
The settling characteristics of metai 'loaded' activated silica are important
because this material has to be separaied from a relatively large overflow volume with
minimal silica losses. Settling data were obtained after activated silica treatment of 1 L
samples of various synthetic solutions as well as the 'Balsam strearn' effluent supplied by
Inco Ltd.. Final pH adjustment prior to settling rate measurements was made with a lime
suspension (28 g CaOR). Three different Pyrex glass container sizes were used for
settling tests as show in Figure 3.3. It is important to define the dimensions of the
containers since settling rates are affected by the tendency for activated silica to cling to a
glass surface.
Afier the required reaction time following activated silica addition. the stirred
products were immediately transferred to the selected container. and the downward
progress of the silica/overflow interface was followed. The interface was always clear
and well-defined giving accurate results. The influence of several variables was studied,
including container dimensions, initial dissolved metal concentration, pH, sol polymer
site (i.e. fraction of gel time at which sol was stabilized), and the effect of using sol that
had been recycled several times.
container I I
'Large'
Toial volume = I L Total volume = 2 L Total volume = I f L Volume used = 1 L Volume uscd = 1 L Volume iued = 11.1 L
4 container
Figure 3.3: Size descriptions for Pyrex glass containers used for settling experiments.
35.5 Cm
3.6 Sludge Treatment
'Medium' container
Tailings pond sludge, supplied by Inco Ltd., was leached at room temperature
with sulphuric acid solution to provide feed for treatment with activated silica. This
preliminary step was carried out by the Mineral Processing group of the Department of
Mining and Metallurgical Engineering, McGill University. Wet sludge (80 g) was
slurried in tap water (400 mL). Concentrated sulphuric acid solution was added. with
mechanical stimng, until the pH reached 3.0 to 3.2. This pH Ievel was found to dissolve
nickel hydroxide, while femc hydroxide was retained as a solid. However significant
amounts of magnesium hydroxide also dissolved, so generating a feed which could be
tested using activated silica both to separate nickel and magnesium, and to produce a
v
concentrated nickel solution. Following the leaching step, samples of leach solution and
solid msidue were retained for c hemical analysis.
3.7 Analytical Techniques
3.7.1 Dissolved Metal Concentrations
A Perkin Elmer 31 10 atomic absorption spectrometer was used for routine
measurements of dissolved metal concentrations (CU", ~ i * , ~e"+ , ~ a " , ~ g * , si"").
Analytical samples and standards (Fisher Scientific Ltd.) were diluted to the Iinear region
of the absorbance vs. concentration plot using de-ionized water. Typical calibration lines
are s h o w in Figure 3.4 . Foc silicon analyses, a nitrous oxide/acetylene flame was used
instead of the aidacetylene mixture used for the other determinations. For calcium and
magnesium analyses, the conventional procedure of lanthanum chloride addition was
followed to suppress possible silicon interference. Table 4.1 shows wavelength and
bandpass used for most metal analyses.
Meta1
Fe
Table 4.1: Wavelength and bandpass used for metal concentration analyses by atomic
Si
absorption.
Wavelength (nm)
248.3
25 1.6
Bandpass (am)
0.30
Flame
AirIAcetyiene
0.30 Nitrous oxide/ Acety lene
O 1 2 3 4 5 6 7
Concentration (ppm)
Figure 3.4: Typical absorbance vs. concentration calibration curves obtained when
using the atomic absorption spectrometer.
3.7.2 Dissolved Cupric Concentrations
In the case of copper only. results obtained by atomic absorption spectrometry
were supplemenied by measurements with a cupric ion selective electrode. This device
provides instantaneous in-situ measurements of the approach to equilibrium when
activated silica is added to a cupric solution, or when acid is added to copper-'loaded'
silica. The electrode distinguishes 'fiee' dissolved cupric fiom other forms of copper,
such as cu(0I-Q' or copper adsorbed by activated silica. For experiments involving the
electrode (Figure 3.2), the initial cupric solution (100 mL) contained 0.001 M Cu (63.5
ppm) at about pH 5.5 and room temperature. Activated silica (or acid), stabilized at
various fractions of the gel time, was added to stirred cupric solution (or copper-loaded
45
silica). The addition rate was precisely controlled using the autotitrator. Both cupric
potentials and pH were continuously monitored, potentials king converted to activities
using calibration curves of the form: E (mV) = a [log (activity)] + b. where a and b are
constants. as discussed in Chapter 4.
3.7.3 Zeta potentials
Zeta potentials of sols were detemined using a Nikon 501 zeta potentiometer.
Each sample was held in a glass-covered ceil at room temperature and placed on the
instruments' microscope stage. Afier adjusting illumination. the microscope was focused
on randomly selected sol particles. which were rnoving in a certain direction. A voltage
was then applied across the ce11 to stabilize the observed particles in a fixed position.
This applied voltage corresponds to the zeta potential of the particles which was directly
registered in mV by the machine. Zeta potentials were measured for sols stabilized at
different percentages of the total gel time. Data were also obtained for sols before and
after copper adsorption. and at various pH values. The claimed accuracy of the
instrument is 5 % or * 1 mV.
3.7.4 Viscosity
A 5 0 4 7 0 5 Canon-Fenske viscorneter was used to determine viscosities of sols
stabilized at various fiaction of the gel time. Viscosity is denved by measuring the t h e
required for the sol to discharge by gravity from a fixed point on a giass bulb to a fixed
point on a comecting capillary. then comparing the result to that for water.
3.7.5 Surface Tension
Surface tension measurements for stabilized sols were obtained at room
temperature using a 'surface tensiomat 2 1 ' (Fisher Scientific Ltd.). This device consists
essentially of a platinurn ring attached to a torsion balance. The ring is initially immersed
in the sol sample, then raised to the surface. The force required to remove the ring from
the surface is then measured and recorded directly in dynelcm.
CHAPTER 4
Characteristics of Dissolved CoppedActivated Silica
Interactions
4.1 Introduction
Before attempting practical applications of silica sols for solution treatment, some
basic knowledge of a typical dissolved metallactivated silica system was obtained. The
metal chosen was copper because of the availability of a specific ion electrode for the
cupric ion. Initial studies focussed on some properties of activated silica mainly in a
metal-fiee system (gel times, viscosity, surface tension, zeta potential).
Subsequent studies of dissolved copper/activated silica interactions included the
pH dependence of dissolved copper transfer to and from the activated silica layer when
stabilized at different fractions of the gel time. The potential operation of copper
'adsorption/desorption' cycles was then investigated, these being carried out in a manner
that produced a concentrated solution (relative to the feed) and regenerated activated
silica (Fig. 3.1). Finally, the possibility of using pH control to separate two metals
(copper and iron) was demonstrated.
4.2 Gel Time Measurements
The composition and pH of as-supplied 'N' silicate (28.7 wiw % silica at pH
1 1.3) corresponds to a position in the 'insolubility domain' in the silica solubility diagram
shown in Figure 4.1 (same as Figure 2.5), but close to the 'insolubility dornain'!'stable
muhimeris domain' border. This location on Figure 1.1 irnplies that the as-supplied
reagent is colloidal. However, 'NT silicate is an impure commercial product - a small
increase in pH from 11.3 to 11.6 would place 'No silicate in the 'stable rnultimeric'
domain indicating total silica solubility. The equivalent position for 'Dg silicate. 29.1
w/w % Si02, pH 12.4, lies within the 'stable multimeric domain' representing a non-
colloidal solution. As described in Section 3.2. the production of stabilized sols involves
a dilution followed by acidification then another dilution, a sequence that would give
equilibrium products in the 'insolubility domain'. Figure 4.1 shows where the location of
the ' stabilized' activated silica used in the bulk of the enperimental work ( 1 .O w/w %
Si02 at pH 9.0) would be under equilibrium conditions. However this 1 wlw % Si02
reagent, when activated and stabilized. is a metastable, not equilibrium product. The
slow gelation process eventually yields insolubility in the form of a gel.
Gel times following acidification, but without the final 'stabilizing' dilution, are
shown in Table 4.1. For both silicates, the expected decrease in gel time with increasing
acid addition is seen. Table 4.1 gives the controlling parameter as volume of acid added
rather than pH. since the pH electrode does not give stable readings as gelation is
approached. The narrow volume ranges of acid addition were such that the silica content
for both 'NT and 'D' silicates was essentially constant (2.0 k 0.1 w/w % SiO2) over the
range of gel times shown (1 minute to > 120 minutes). 'N' silicate with the 18 minute
gel time (8.50 rnL acid addition) was the most frequently used. Figure 4.2 gives an
indication of gel time reproducibility, and also shows that the storage time (up to 1 2 days)
of 'N' silicate &er 50 v/v % dilution with water had negligible effect on gel time. The
significance of the aging of 50 v/v % 'N' silicate is that batches of the viscous as-received
reagent were diluted 50 vlv % with water and used as needed in the preparation of sols.
Stable multimeric *
lnsolubility domain
domain L
/
3
Mononuclear domain 10 - J -
Fipre 4.1 [5,63,641: Solubility diagram for amorphous silica at 25 OC.
1. As received 'N' silicate. 2. As received .De silicate.
3. 1 .O W/W % Sioz stabilized .activated' silica sol.
1 m = 3.22 ('N' silicate) 1 m = 2.0 ('D' silicate) 1
Table 4.1: Total gel times of silicates (m = 3.22 and 2.0) after the addition of 10 w/w % H2S04 (1.02 M) to 10 mL 50 vlv % silicate with 79.5 rnL deionized water giving unstabilized activated sol containing 2.0 wlw % Sioz. A further 50 v/v % dilution with water would have given stabilized 1 .O w/w % Si02 sol.
Volume of acid
(mL)
8.50
Conditions 8.5 m L 10 W/W ./a HISOI 10.0 ml, 50 V/V O h 'N' silicate 79.5 mL witer
O 2 4 6 8 10 12 t 4 16
Aging of 50 viv Oh 'N' silicate (days)
Total gel time
(minutes)
18.0
Figure 4.2: Effect of 'N' silicate aging on gel time.
Volume of acid
(mL) 16.24
Total gel time
(minutes)
1 .O
The preparation of 'stabilized' sols (i.e. those with additional water dilution
following acidification as in Table 4.1) is summarized in Table 4.2, which shows the
relative volumes of the various additives involved in preparing the 200 rnL batches
produced for the rnetals recovery tests in Chapters 5 and 6. With reference to Table 4.2,
samples of as-received 'N' silicate, diluted to 50 v/v % (column 2), were added to water
while stirring (column 3) followed by IO w/w % acid addition (colurnn 4). The arnounts
of acid were chosen to give gel times in the order of minutes (column 5). Gelation was
arrested ai 50 percent of total gel time (column 7) by about 50 vlv % dilution with water
(column 6) to give the final product at the pH values in column 8. A 4.0 w/w % Si02
stabilized sol is at the maximum practicable concentration when using 'N' silicate.
because at more than 4.0 wiw % Si02 , discrete particles are visible.
The term 'stabilized sol' does not mean gelation bas been halted. Stabilization by
water dilution greatly deiays gelation to the point that a metastable structure is retained
for a useful tirne period, so allowing practical application of the sois' properties. Thus the
stabilized sols also have a finite, but lengthy, gel time as seen in Table 4.3. The 1 .O %
sol (1 W/W 46 SiO2) at pH 8.3. refened to in Table 4.3, is the material used in most of the
experimental work in Chapters 5 and 6. lt is seen that storage for 42 days did not
produce gelation. However. acidification of the 1.0 % sol to pH 0.5 caused gelation in
21 days. A shorter sol life time at decreased pH means that in a pH controlled
adsorption/desorption cycle, the time spent at low pH for desorption (or in storage)
should be as short as possible consistent with good desorption results. If silica
concentration is increased to 2.0 % or 4.0 %. the gelation time decreases, as expected
(Table 4.3).
Figure 4.3 shows how viscosity increases for 1.0 w/w % Si02 sols stabilized at
increasing proportions of the gel tirne. This enhanced viscosity has been attributed to
increased degree of p l ymetization [SI. Viscosity measurernents could there fore be used
to characterize the percent gel time at which sols are stabilized. However, such
measurernents are unnecessary since the metal adsorption capability seems to be
insensitive to the gel time fraction at which the sol is stabilized (as discussed below in
Section 4.5). Measurements of surface tension as a fmction of 'percent gel time' (Figure
4.4) show that this property remains essentially constant at the value for water (75
dynekm) which is in agreement with literature pertaining to silicate solutions [SI.
Weight VO SiQI in stabilized sol after dilution
Table 4.2: Volumes used for the preparation of stabilized 'activated' silica containing
HzO (mL)
SO vlv % 'N' silicate
( m u
1.0 W/W %, 2.0 W/W % and 4.0 w/w % Sioz . Sols were stabilized at 50 % of total gel
time and derived fiom silicate with m = 3.22.
10 w/w % H2SOd ( m u
Total gel time (min)
H20 dilution
(mL)
Time before
dilution
PH after
dilution
Weight % Si02 in stabilùed soi after
dilution
1.0 %
l 2.0 % 1 9.9 1 29 days I
pH
1.0 %
5 days
Time for complete gelation
0.5
Table 4.3: Times required for cornplete gelation of sols prepared as in Table 4.2. Effect
of pH for 1 .O w/w % Si02 is also shown.
2 1 days
9.0
mL aci& gel time
no gelation afier 42 days
A Conditions 8.25 or 8.50 mL 10 wlw ?A H,SO, 10.0 mL 50 vlv % "N" silicate
Percent gel time
Figure 4.3: Viscosity of 'No silicate sols stabilized ai different gel time fractions.
Conditions as in Table 4.1 .
Conditions 8.3 mL 10 W/W % H,SO, 10.0 mL 50 vlv % 'N' silicate 79.5 mL watcr Gel time = 40 min
Percent gel time
Figure 4.4: Surface tension of 1.0 w/w % Si02 sol stabilized at different gel tirne
fractions.
4.3 Calibration of Cupric Ion Electrode
Before using the cupric ion electrode for rneasurements in the copper
sulphate/activated silica system, it was necessary to establish calibration curves. The
- electrode gives the potential corresponding to ionic activitv, a, where a = y C. y -
activity coefficient, C = molar concentration. Activity coefficients were determined using
Pitzers' model [109], which is valid for ionic strength, 1 < 0.50 molar, where:
2, is the charge on species I at concentration C,. The equation for Pitzers' model is:
For this system, A. in equation 4.2 is 0.392 and p is 14 [109].
For each copper concentration, the ionic strength was calculated fint using
equation 4.1, then for the measured potential, the cupric ion concentration was
determined using the calibration curves s h o w in Figure 4.5 where data points for a given
ionic strength produce equations of the form: E (mV) = a log a + b . Each line corresponds to a certain ionic strength and activity coefficient, as controlled by
adding sodium sulphate solution. Although in principle the activity coefficient correction
is necessary, values for y at up to only 0.001 M Cu, as in the present application, are
between 0.9 and 1.0, so that the enor in assuming activity equals molarity would in
reality be minor.
Accuracy was checked by comparing concentrations derived from electrode
readings with values obtained by 'AA' spectrometry. Table 4.4 shows good agreement
between the two analytical methods. Measurements by ' AA' spectrometry of less than 1
ppm Cu were not possible with the equipment available. The data in Table 4.4 were
obtained using solutions to which increasing amounts of sol had been added. M i l e the
'AA' spectrometry measurements were canied out on filtered solutions, the electrode
readings were made 'in situ' in the unfiltered system during sol addition. Thus the
electrode measures only fiee dissolved copper as distinct fiom adsorbed copper. If
agitation is stopped, adsorbed copper appears as a faintly blue, slightly opaque phase that
settles beneath clear solution. The lower, copper-containing activated silica layer is
visibly slightly more viscous than the upper solution.
Electrode sensitivity was checked by adding water to a solution of known copper
concentration, then, knowing the volumes involved, the electrode-derived values could be
compared to the me, calculated values. Figure 4.6 demonstrates good electrode
sensitivity.
Na,SO, addtd (mL)
Cupric concentration (ppm)
Figure 4.5: Calibration curves for the copper electrode. Sodium sulphate (2 M) added as
ionic strength adjuster.
1 Sol addition
.Table 4.1: Cornparison of analysis methods. Elecuode rneasurements vs. atomic
absorption. Sol added as 1 .O w/w % SiOr with m = 3.22 ('N' silicate) stabilized at 0.5 gel
time (total gel time 18 minutes).
Volume of water added (m L)
Figure 4.6: Effect of water addition on the sensitivity of the Coppet ion selective
electrode at pH 5.5.
4.4 Copper Adsorption on Silica Surfaces (quartz, silica gel, activated
silica)
Previous work with silica (quartz) and silica gel particles [97-99) has shown that,
when starting in an acidic solution, adsorption increases with increasing pH reaching a
maximum when concentration of dissolved rnetal hydroxy complexes also reaches a
maximum. For copper adsorbing ont0 a silica (quartz) surface, or ont0 silica pl
particles, maximum adsorption is at pH 7.0 [97-991.
In general, the adsorption reaction has been presented [97-991 as a reversible
cation exchange between M~ (or ML"+ where L is a ligand such as OH- ) and the silanol. I
-SiOH, group: f
M W + m ( - S i O H ) $ M(0Si-) (n-m) + + m H ' I 1 m (4.3)
Dugger et al. [97] have shown that, for silica gel, m = n regardless of the rnetal
cation involved. Thus copper adsorption can be represrnted as:
The sorption capacity (moles metal I g Si@) depends not only on pH but also on
the specific adsorbent used. For example, sorption capacity of a batch of 0.2 mm
diameter silica gel particles with pore size of 85 A and specific surface area of 568 m'/g
was shown to be between 10~' and 1 o4 mole metdg silica, depending on pH [99]. In the
work cited above [97-991, the arnount of quartz (or silica gel) present is constant, and the
pH is increased by adding buffer solutions prepared by neutralizing (with NaOH)
chloroacetic acid, formic acid, acetic acid or perchloric acid.
In the present work. addition of activated silica (1 w/w % Sioz stabilized at 50 %
of gel time) to 0.0010 M cupnc solution gave 30 ppm dissolved copper at pH 8.8 (top
half of Figure 4.7). In this latter Figure, the silica addition is fixed, and pH was adjusted
downwards using 10 w/w % sulphuric acid solution. Maximum copper adsorption was
at pH 7.0, as for silica (quartz) and silica gel [98j. The amount of copper adsorbed at pH
7.0 in Figure 4.7 is equivalent to 5*104 mole Cu/g SiOz, which is about an order of
magnitude higher than the sorption capacity reponed for silica gel powder [99]. A higher
sorption capacity would be expected for liquid activated silica than for solid silica gel
powder because of the improved contact in the former case. Further work at different
copper/activated silica ratios would establish how closely the adsorption data of Figure
4.7 represent true sorption capacities.
Earlier studies [98] do not include data for copper that extends to more alkaline
conditions than pH 7.0. In the present case, desorption was observed at pH above 7.0
(Fig. 4.7). The Pourbaix diagram for the coppedwater system [ 1 1 O] shows that dissolved
CUOZ- - can exist under alkaline conditions, but the amount that could form at pH 8-9
would be undetectable. The most probable explanation is that a colloidal product
(Cu(OH)2 or CuO.Hz0) forms at above pH 7 and desorbs from the silica surface.
The copper analyses in Fig.4.7 were determined by ' A.A' spectrornetry rather than
by the cupric ion electrode. At pH above 7.0, there was no visible precipitate and no
Figure 4.7: Zeta potential of 'N' silicate sols and copper concentration as a function of
pH. 1 .O wlw Si02 sol stabilized at half total gel time of 18 minutes. 21 mL sol added to
100 mL 63.5 ppm cupric solution cupric at pH 5.5 initially. pH is adjusted downward
with 10 wiw % sulphuric acid after sols addition. Copper analyses by A.A.S.
special filtration procedure was followed, so that a colloidal copper compound would be
reported as dissolved copper. While the cupric ion electrode gave almost identical results
to those obtained by 'A.A.' spectrometry in the pH 5 to pH 7 range, there was no
electrode response at above pH 7, indicating an absence of dissolved cupric ion, as
expected.
It is doubtful that cupric analyses in Figure 4.7 are true equilibrium values, but
rather those corresponding to a seemingly constant pH reached afier about 20 minutes. In
a fwidamental snidy of cupric adsorption by both dissolved and colloidal silicates added
to copper perchlorate. Falcone [92] found that dissolved cupric activity kvas still
decreasing afier one month. Falcones' studies [92] also indicated that colloidal silica
begins to adsorb dissolved metal ions at 1 to 2 pH units below the equilibrium value for
hydroxide precipitation. In the present case, at Say pH 6 (Fig. 4.7). the measured value of
44 ppm Cu compares with a calculated 55 ppm Cu that would be in equilibrium with
copper hydroxide [ I l l ] . aithough the usefulness of this kind of cornparison is
questionable.
The bottom half of Figure 4.7 shows zeta potentials corresponding to the
adsorption data on the top half of Figure 1.7. In general. the zeta potential is the potential
difference between the closest distance of approach of hydrated counter ions to a charged
surface and the bulk solution. These counter ions are of opposite charge to that of the
surface itself. Zeta potential measurements have been used to demonstrate the pH
dependence of hydroxy complex adsorption on silica (quartz) surfaces. For example
(Figure 4.8), the zeta potential of quartz when in 1 * 104 lead chloride at pH 5 has been
measured at -30 mV [117]. As the pH is increased, the zeta potential becomes positive
and reaches a maximum of + 30 mV at pH 8. A M e r pH increase to 1 1 causes the zeta
potential to reven to its original negative value at pH S. These potential vs. pH changes
are considered to represent increasing formation and adsorption of P~(OH)+ complexes
between pH 5 and pH 8. At above pH 8, desorption occurs. At pH 8, lead is known to
fomi a maximum concentration of P~(oH)* complexes [113]. Therefore. Fuerstenau et al.
concluded that PC(OH)+ is the adsorbed species. Similar observations were made for the
nickel-quartz system [114], the aiuminumquartz system [ 1 12,115], and the calciurn-
quartz system [ 1 1 61.
In the present case (Fig 4.7, bottom), the zeta potential also increases from a
negative value as copper hydroxy complexes are adsorbed with increasing pH from 5 to 7.
However at above pH 7.0, the potential does not reflect copper desorption by decreasing
back towards the original negative value. An important difference between quartz and
activated silica is that, surface effects aside, the quartz structure does not change with pH
within the range of interest (pH 5 to 9). In contrast. activated silica tends to
depolymerize with increasing pH (the opposite of gelation when acidified). Thus the
change in zeta potential in Figure 4.7 represents the net effect of both copper
adsorption/desorption and change in polymer size.
The effect of polymer size on zeta potential is seen in Figure 4.9. where
measurements have been made in the copper-free system ai essentially constant pH (8.9 * 0.1). Although the experimental scatter on Figure 4.9 is high. there are sufficient data to
show a clear trend towards higher zeta potential with decreasing 'percent gel time' at
whic h the sol has been stabilized (i.e. decreasing pol ymer size). This qualitatively
supports the findings in Figure 4.7, where zeta potential fails to return to negative values
at above pH 7.0 due both to decreased polyrner size and incomplete copper desorption.
An improved undentanding of the complex dependence of surface charge on polymer
size, amount of adsorbed metal and pH requires a more detailed fundamental study than
undertaken in the present work.
Figure 4.8 11 121: Zeta potential of quartz as a funftion of pH in the presence
of 1 + 1 O-' M P ~ C I ? .
Percent gel time
Figure 4.9: Zeta potential of 'N' silicate sols stabilized at different fractions of total gel
time. I .O w/w % Sioz sols with total gel time of i 8 minutes.
4.5 Activated Silica as Coppet Adsorber and pH Controller
The copper adsorption data of Figure 4.7 were obtained following a fixed silica
addition to reach pH 8.8 with subsequent pH adjustment using sulphuric acid solution. In
this situation, the activated silica acts solely as a medium for copper adsorption. It is also
possible to use silica both as a copper adsorber and pH controller, in which case the silica
addition is determined by the pH required regardless of the coppedsilica ratio.
Expenments carried out under these latter conditions gave the results shown in Figure
4.10, where use of activated silica prepared from 'N' and 'D' silicates (SiO1Na20 weight
ratio, m = 3.22 and 2.00 respectively) is compared. the copper analyses being performed
by 'AA' spectrornetry. In both cases, copper adsorption reaches a maximum at pH 7.0,
above which some desorption occurs. as also seen in Figure 4.1 O.
Also included on Figure 4.10 are results obtained by Lempka [ I l 71 using lime. It
is seen that the three reagents give comparable results in the pH range 5.5 to 7.0.
However, above pH 7 using lime, the colloidal copper compounds present when silica is
used disappear because a copper hydroxide precipitate f o n s instead. This important
behavioural diflerence is the reason lime is added for final pH control in Chapters 5 and 6
where practical applications involving iron and nickel recovery from an effluent and a
tailings pond sludge respectively are described. Other work [II71 carried out at National
Silicates Ltd., has also demonstrated the importance of lime addition to prevent high
effluent copper levels at pH above 7. In these latter studies (with copper nitrate
solutions), soluble silicates were added without deliberate 'silica activation'. However.
some uncontrolled activation would have occuned because the initial copper nitrate
solutions were acidic.
The difference between adsorption using silica and precipitation using lime (or
sodium hydroxide) is visually obvious when cobaltous sulphate is the solution being
treated. Figure 4.1 1 a) compares use of sol and sodium hydroxide when the pH is
deliberately adjusted for partial cobaltous removal, leaving a pink overfiow in each case.
The opaque underflow in the sample with sol retains its pink colouration because pink
hydrated cobaltous ion has been adsorbed fiom solution. The sample with sodium
hydroxide addition shows a blue undefflow due to cobaltous hydroxide precipitation.
Figure 4.1 1 b) compares sol and lime additions. The underflow using lime is now green,
representing a combination of crearny white gypsurn plus blue cobalt hydroxide. In
Figure 4.1 1 b). the overflows are colourless because the pH has been adjusted for total
dissolved cobaltous removal. M e n sarnples are centrifuged, the results are as seen in
Figure 4.1 1 c), which shows (fiorn left to right) centrifuged cobalt-ioaded sol. cobaltous
hydroxide and lime/gypsum plus cobaltous hydroxide. In general. the colour differences
in Figure 4.1 1 (a-c) are much more dramatic than with the cupric ion.
The relative amounts of 'N' and 'Dg activated silica needed to reach a specified
pH are seen in Figure 4.12. where the volumes of 1 .O w/w % SiO2 corresponding to the
data points of Figure 1.10 are shown. At a fixed silica level, 'Dg silicate is a more
powerful reagent for raising pH than 'N' silicate mainly because of the higher Na20
content in the former case.
Figure 4.10: Dissolved cupric concentration vs. pH. Effect of reagent addition on
copper concentration. 1 .O w/w ?/o Sior sol ('W' silicate) stabilized at half total gel time of
18 minutes. 1.0 w/w % Si02 sol ('D' silicate) stabilized at half total gel time of 21
minutes. Lime added as 28 g CaOIL suspension. .\nalyses by A.A.S.
Figure 4.11 a): Products obtained at pH 8.05 after the addition of 1 .O wlw % Sioz sol
(lefi) or sodium hydroxide (right) to a 2 gi2 cobalt solution. AAer free settling, a pink
ovefflow is shown for both samples while the underflow is pink after sol addition and
blue after sodium hydroxide addition.
1 .-. - - ? - - Figure 4.1 1 b) Products obtained at pH 9.5 after the addition of 1 .O w/w O/O Si02 sol
(right) or lime ( M t ) to a 2 g/L cobalt solution. Aftcr free settling, a clear overfiow is
s h o w for both samples while the underflow is pink afier sol addition and green aher
lime addition.
Figure 1.1 1 c) Centrifbged products obtained at pH 9.5 after the addition of (lefi to
righr) 1.0 W/W % Sioz sol. sodium hydroxide. lime to a 2 g/L Cobalt solution.
Activated silica sol added (mL)
Figure 4.12: pH vs. volume of activated silica sol ('N' or 'D') used for results obtained
in Figure 4.10.
Copper adsorption kinetics were measured for a typical test in which activated
silica was used both to adsorb copper and control pH. Table 4.5 shows the data obtained
with the cupric ion electrode following incremental sol additions. Afier 1 day, dissolved
cupnc levels are slightly below the 20 minute values, while after 3 days, the results are
essentially the same as after I day retention. In general, metal adsorption ont0 colloidal
silicates requires at least a month to reach equilibriurn [92]. For practical purposes,
adsorption reactions in the present work are considered complete after 20 minutes.
The bulk of the experimental work was carried out using sol stabilized at 50 % of
the gel time. This arbitrary choice was based on the unexpected finding that the
performance of activated silica as a copper adsorber is effectively independent of gel tirne
Table 1.5: Unadsorbed copper measured afier 20 minutes and one day following
addition of 1 .O wiw % SiOz sol ('N' silicate) stabilized ai O.j*gel time (total gel time 18
minutes).
Sol addition ( m u
3
6
9
12
15
18
fraction at which gelation is arrested (at between 10 % to 80 % of gel time). An
extensive study was performed to establish dissolved copper concentration vs pH curves
using sols prepared fiom 'N' silicate. and stabilized at different fractions of various gel
time (18 min.. 40 min.. 105 min.. gel rimes). The results. shown in Figures 4.1 3 to 4.15
were obtained using the cupric ion electrode. so that measurements extend only to pH
7.0-pH 8.0, above which there is no emf reading.
The insensitivity of copper adsorption to gel time fraction at which sol has been
stabilized can be readily seen in Figures 1.13 to 4.15. The reason the curves are
effectively the sarne for a given total gel time is that the silica addition is determined by
the pH required. and the mole Si02/mole Cu ratio is not controlled. The result is a very
After 20 minutes After one day
pH
6.14
6.17
6.23
6.45
6.76
7.48
pH
5.90
6.08
6.27
6.47
6.55
7.26
rnV
192.0
185.9
179.7
170.4
154.2
115.7
mV
190.5
185.1
178.5
168.5
148.8
116.1
Conc ( P P ~ )
50.24
36.34
22.70
11.27
3.37
0.13
Conc* (PW)
43.1 3
33.91
30.46
9.56
2.13
0.14
large molar excess of silica over copper. For example at pH 7-pH 8, where dissolved
copper removal is aimost complete, the mole Si02/mole Cu is about 30. This suggests
the capability of the added silica to adsorb copper greatly exceeds the amount actually
needed regardless of the gel tirne fraction at which stabilization has occurred.
It can be seen that Figures 4.13 to 4.15 refer to sols stabilized at 10 % to 80 % of
gel tirne. If non-activated W' silicate is added directly to a 'heavy metal' solution. there
will always be some natural, uncontrolled gelation because the latter solution will always
be acidic. However, the tendency to form a 'heaw metal' silicate precipitate will be
avoidad in favour of adsorption only if gelation has produced sufficient adsorption sites
within the polymeric network. Sols stabilized ai more than 80 % of gel time are too
viscous to be readily handled (Fig. 4.3).
Sets of experirnents similar to Figures 4.13 to 4.15 were also c k e d out with 'N'
and 'D' silicates (Si02lNazO weight ratio 3.22 and 2.0 respectively). with the difference
that copper analyses were perfomed by 'AA' spectmmetry instead of the copper
electrode (Figures 4.16, 4.17). Since the former method reports finite dissolved copper
levels under aikaline conditions, the data extend to about pH 8.8. At up to pH 7, the
dissolved copper vs pH curves follow almost the same path as established using the
cupric ion electrode (Fig. 4.13-4.19, and copper adsorption is, again, independent of
degree of gelation at which the sol is stabilized. If pH is increased to 8.8, 'AA'
spectrometry shows apparent dissolved copper at levels similar to those seen in Figure
4.7. As discussed previously in comection with Figure 4.7. this apparent increase in
dissolved copper under alkaline conditions is attributed to desorbed colloidal copper
hydroxide.
Figure 4.13: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1 .O w/w % Si02 sols ('N' silicate) with total gel tirne of 18 minutes. Analyses using cupric ion electrode.
Figure 4.14: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1.0 w/w % Sioz sols ('NT silicate) with total gel time of 40 minutes. Analyses using cupric ion elecuode.
Percent gd tirne nt which sol stibiliztd
-10
m20
40
Figure 4.15: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1.0 wlw % Si02 sols ('N' silicate) with total gel time of 105 minutes. Analyses using cupric ion electrode.
%
\\ whkh sol slabilircd 9% Percent gel timc a i
Figure 4.16: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1 .O w/w % Sioz sols ('NI silicate) with total gel time of 18 minutes. Analyses by A.A.S.
Percent gel timt at which sol strbilized
*20 810 5 0 x 6 0 r 80
Figure 4.17: Dissolved cupric concentration vs. pH. Effect of polymer size on copper adsorption. 1.0 wlw % Si02 sols (3' silicate) with total gel time of 21 minutes. Analyses by A. A. S.
4.6 Copper Recovery using Adsorption/Desorption Cycle
The global objective of the present research is to study the potential for
developing a solution treatrnent process based on a reversible pH-controlled
adsorption/desorption cycle using silica sol, as illustrated by Figure 3.1. Important
features are production of a concentrated product solution (relative to the feed), sol
regeneration in recyclable form and creation of an acceptable effluent. In Figure 3.1, the
term 'Recovery' is preferred over 'Adsorption' because lime (or NaOH) is needed for
final pH control to precipitate any colloidal copper cornpounds that desorb at pH 7 and
above.
74
After sol addition, separation into a purified emuent and concentrated copper-
loaded sol is greatly enhanced by centrifugation. Addition of a few drops (when on a lab
scale) of concentrated acid to reach about pH 4 will redissolve most of the copper. Use
of concentrated acid minimizes the dilution effect that would result from using a weak
acid solution, so promoting the concentration of copper. A second centrifugation step
separates a product solution from regenerated sol. The polymer size distribution in
regenerated sol would not be the sarne as that in fresh reagent because both acidification
for copper redissolution. and the concentration obtained b y centrifugation encourages
gelation. However, as seen in Figures 4.13 to 4.17, copper adsorption is insensitive to
degree of gelation. providing a large molar excess of activated silica over copper is
present.
Table 4.6 gives results for two single 'adsorption~desorption' cycles carried out
with feeds at 63.5 pprn Cu and 315 pprn Cu. The 'pprn copper in effluent after
centrifugation' is undesirably high at 1.60 ppm and 5.3 ppm. but these copper levels
could be reduced by raising the 'Adsorption' pH to about 7.5. The product solutions
contain 355 pprn Cu and 1 .O4 g/L Cu, which represent concentration factors ((pprn Cu in
product)/(ppm Cu in feed)) of 5.6 and 3.3 respectively. In ternis of mass balance. copper
losses to the effluent are nrgligible, while recovenes into the product solutions were 84 %
(63.5 pprn Cu feed) and 69 % (3 15 ppm Cu feed). The unrecovered copper would be
recycled with the sol, but could possibly be recovered by raising the acidity of the
'desorption' step. Figure 4.18 shows products corresponding to Table 4.6 for the 3 15
pprn copper feed. The left side of Figure 4.18 is the blue copper-'loaded' sol (bottom
phase) and clear effluent afler adsorption/centrifugation while the right side is the sol for
recycle (bottom phase) and concentrated solution afier desorptiodcentrifugation.
The two columns of Table 4.6, when taken together, suggest that a 63.5 ppm Cu
leed could be concentrated into a 1 .O4 g/'L Cu product using two cycles. It should also be
noted that up to six sol recycles have been successfully carried out without a noticeable
change in adsorption characteristics.
Feed with 63.5 ppm CU"
Feeà with 315 pprn CU*
1 'Adsorption' pH 1 ppm CU" in effluent
Volume of concentrated solution (after centrifugation)
(mL)
after centrifugation
pH for 'desorption'
ppm CU" in concentrated solution (after centrifugation)
1 .O0
Table 4.6: Copper concentrations obtained afier a single gadsorption/desorption' cycle.
5.3
4.36
100 mL initiai volumes at pH = 5.5. 20 rnL sol (1.0 w/w % SiOz, stabilized at 50 % of gel
4.14
time) + NaOH for pH adjustment. Concentrated acid was used for desorption.
Figure 118: Products corresponding to Table 4.6 for the 3 15 ppm copper feed after
centrifugation. The left side is the blue copper--1oaded' sol (bottom phase) and clear
effluent (top phase) afirr adsorption at pH 7.0. The right side is the sol for recycle
(bottom phase) and concentrated solution (top phase) after desorption at pH 4.14.
4.7 CopperlIron Separation using Activated Silica
Previous worken [98, 10 1 - 1031 have s h o w that certain metal group separations
are possible by pH-controlled selective adsorption onto silica gel powder. For example,
if the pH of an acidic, multi-metal solution is increased. iron (femc) and aluminum will
be adsorbed first onto solid silica gel. fdlowed by divalent k a i y metal' cations and
ultimately calcium and magnesium (Fig 2.15). .Adsorption data for silica gel powder
[98.10 1-1 031 also show that separations within a group (i.e. Cu* fiom ~ i - ) would not
be feasible.
Table 4.7 shows data for coppediron (ferric) separation by selective adsorption
ont0 activated silica from a solution initially containing 0.0010 M of each metal at pH
77
3.09. M e r pre-neutraiization with lime to pH 3.20, activated silica was added to reach
the pH values show in Table 4.7. The colour of the silica underfiows progressively
changed fiom brown (iron) to green (iron and copper) as the pH increased. Optimum
separation is at about pH 5.2, at which 95 % of the iron has been removed, leaving 90 %
of the copper dissolved. As expected, almost complete copper adsorption occurs at pH
7.0.
The data in Tables 4.6 (copper concentration) and 1.7 (coppediron separation) are
intended to illustrate the principles involved in the use of activated silica. rather than as a
practical alternative to conventional cementation and solvent extraction. Chapters 5 and
6 develop the concepts discussed in Chapter 4 in the context of processing waste
materials (iron effluent and a tailings pond sludge) for which no satisfactory treatment
presently exists.
Table 4.7: Copperliron (ferric) separation by selective iron adsorption on stabilized sol
(1.0 WIW % Si@, stabilized ai 50 % of gel time). Initial Cu/Fe solution at pH 3.09
preneutralized with lime to pH 3.20.
PH ppm ferric in solution
% femc removed
ppm cupric in solution
% cupric removed
CHAPTER 5
Iron Recovery from Acid Mine Drainage using Activated Silica
5.1 Introduction
The recovery of iron from synthetic and actual acid mine drainage was the fint
practical application studied for use of activated silica in generating a concentrated
dissolved metal product from a dilute effluent feed. National Silicates Ltd, the project
sponsor, had identified a potential market for impure concentrated femc solution that
would be suitable as a flocculant in sewage treatment. Also both sewage and activated
silica physically resemble a mud rather than a granular solid. Since centrifiges are used
on a large scale in sewage treatment, this approach seemed technically feasible for
dewatering activated silica after 'adsorption' and 'desorption', as indicated in Figure 3.1.
It is not realistic to expect concentration fiom about 100 ppm Fe (effluent) to 30
g L Fe (target minimum level in product) in a single 'adsorption/desorption' cycle. and it
was found that three consecutive cycles were needed. Thus the present Chapter
describes use of activated silica to concentrate iron from a wide range of feed levels (55
ppm Fr to 1 1,000 ppm Fe). Other aspects of Figure 3.1 that have been addressed include
settling characteristics of iron-loaded sol, seasonal temperature changes and the gypsum
formation problem that results from using lime for pH control in a closed
' adsorptioddesorption' loop.
5.2 Iron Recovery Step (Fig. 3.1)
Activated silica was prepared from 'NT silicate as described in Section 3.2 for the
production of the 'standard' 1 .O wlw % Si02 reagent used in the bulk of the experimental
79
work. However feed solutions containing fiom 55.8 ppm Fe to 11.0 g/'L Fe were
investigated. With this wide range of iron concentrations, it was convenient to have
available activated silica reagent containing 2 % (2.0 wlw %) and 4 % Sioz. The latter is
the maximum practicable concentration, since at more than 4 % Si02 , srna11 discrete
aggregates are visible. Presumably gelation of the 8 % Si02 starting sol was already well
advanced at half the overall gel time, at which point 50 vlv % dilution with water creates
the 4 % Si02 sol actually used. Details for the preparation of 1 %, 2 %, and 4 % activated
and stabilized sols are given in Table 4.2 and Section 4.2.
In the first series of iron recovery tests. 20 mL activated silica were added to 100
mL s h e d femc soiution containing 55.8 ppm to 1.1 17 ppm femc. Results were
compared at a constant final pH of 3.95 (t 0.10). In each case, pH adjustment to pH 3.95
with lime or sulphuric acid was needed aHer silica addition. The femdsilica mixtures
were stirred for 24 hours. with further pH adjustment if needed, then left for a final 2
houn. Samples of the stirred mixture (45 mL) were then centrifuged (2,000 rpm for 10
minutes) without prior settling. Quaiitatively, special note was made of the sharpness of
the interface and ovefflow clarity in the centrifuged material. It is important that an iron-
containing silica layer readily forms with a well-defined interface separating silica from a
clear overtlow. The iron concentration in the overflow was rneasured directly, while that
in the silica underflow was estimated by mass balance. Table 5.1 presents the most
important results.
Weight % Si02
1.0 O h
Initial conc., pH 55.85 pprn ~ e * (0.001 M), pH= 2.98 pH(+ sol) = 4.50 Final pH = 3.88 Excellent V (sol) = 8.0 % (light brown) Solution: 0.50 pprn Sol: 361 pprn (Silicdfemc) = 35.7 pH(+ sol)>>9.0 Final pH = 3.90 Acceptable
v (sol) = 9.3 (light brown) Solution : 19.50 pprn Sol: 438 pprn (Silicdfemc) = 7 1.4 pH(+ sol)>> 10.5 Final pH = 4.05' Acceptable
V (sol) = 10.7 (light brown) Solution: 30.50 ppm Sol: 291 pprn (Silicdferric) = 1 42.9
Initial conc., pH 5JS.S pprn ~ e * (0.01 M), pH= 2.45 pH(+ sol)=2.69 Final pH = 3.93 Excellent V (sol) = 8.7 (dark brown) Solution : 1.95 pprn Sol: 3477 ppm (Silica/ferric) = 3.6 pH(+ sol)=3 .O5 Final pH = 3.94 Excellent V (sol) = 13.3 (dark brown) Solution : 1.85 pprn Sol: 23 18 pprn (Silica/ferric) = 7.1 pH(+ so1)»7.5 Final pH = 3.95 Excellent v (sol) = 20.0 (light brown) Solution: 1 15.0 pprn Sol: 18 15 pprn (Silica/femc) = 14.3
Initial conc., pH 1117 pprn ~ e + " (0.02 M), pH-2.27 pH(+ sol)=2.50 Final pH = 3.97 Excellent V (sol) = 10.7 (light brown) Solution : 1.90 pprn Sol: 6968 pprn (Silicdfemc) = 1.8 pH(+ sol)=2.70 Final pH = 3 .9O Excellent V (sol) = 13.3 (dark brown) Solution : 4.20 pprn Sol: 6953 pprn (Silicafferric) = 3.6 pH(+ sol)=3 .O9 Final pH = 4.00 Excef lent V (sol) = 26.7 (light brown) Solution: 85.0 pprn Sol: 3207 pprn (Silicdfemc) = 7.1
Table 5.1: Femc recovery fiom 100 mL of 55.8 to 11 17 pprn Fe feeds using lime or acid for pH control.
Key: pH (+ sol) = After 20 mL sol addition. Final pH = After acid (10 wlw % H2SOj)flime addition. Excellent = Phase separation characteristics. V (sol) = VIV % of metal '1oaded'-sol phase d e r centrifugation in 45 mL flask (light brown) = colour of sol phase. Solution: x pprn = Ferric in top solution (ovefflow). Sol: y pprn = Femc in sol fiom mass balance. (Silicajfemc) = (SilicalFemc) weight ratio in solution after sol addition
The following explanation of Table 5.1 takes the top lefi-hand 'box' as an
example (55.8 ppm Fe feed at pH 2.98, 1 % Si02 added). Here the pH after silica
addition rose to 4.50, acid then being needed to reach the desired final pH of 3.88. Lime
was added in cases where final pH was greater than pH (+ sol). 'Excellent' means a
sharp interface was obtained with a hi& clarity, colourless overflow. The 'V (sol)' refen
to the volume percent occupied by the light brown bottom phase (underflow). Iron
contents measured in top phase (overflow) and undefflow (calculated) are also shown.
If the effect of siiica addition at constant initial femc is considered (i.e. vertical
columns of Table 5.1). it can be seen that the pH following silica addition increases with
increasing silica concentration. as expected. However, the silica/solution interface
becomes less sharply defined with a white cloudy overtlow when a high silica addition is
combined with a low initial iron level (4 % Sioz , 55.8 pprn Fe). In contrast. with zero
silica added (i.e. lime only), the interface is alrnost undetectable, and settles slowly
leaving a dirty brown overflow. Even centrifugation does not produce good solid/liquid
separation.
In general, as the pH rises following silica addition, femc hydroxy complexes
(F~(oH)". ~e(0I-l);) start to be adsorbed on the negatively charged silicate polymers,
then femc hydroxide foms if the final pH reaches about 4, as in Table 5.1. Al1 particles
and polymer species in the system then coagulate thus leading to colloid instability and
second phase formation. In the total absence of iron (or other heavy metal). the activated
silica is uniformily distributed as colloid. and there is no natural tendency to phase
separation. At any given femc level, it is readily possibie to add too much silica, which
is detrimental to adsorption and second phase formation. A high silica concentration in
solution creates the white opaque overflow mostly seen with 4 % Si02 addition, and iron
contents in the effluent increase. In this situation, the excess negatively charged silicate
is effectively acting as a dispersant thus reporting to the overflow with entrapped iron.
Concentration factors for iron ((a in silica)/(g/L in feed)) tend to be highest (about x 7)
at low (1 %) silica addition, at which the silica layer is more compact.
A parallel set of experiments was carried out in exactly the same manner as for
Table 5.1 except that sodium hydroxide was used instead of lime (Table 5.2) whenever an
upward adjustment in pH was required. The results are essentially the same when lime is
replaced with NaOH. Also, phase separation characteristics again deteriorate at high
Si02 I Fe ratios (i.e. results for 55 ppm Fe), but are excellent when this ratio decreases.
Thus the critical factor controlling separation behaviour is Si02 I Fe ratio regardlrss of
whether lime, sodium hydroxide or acid is used for adjustment of final pH.
Weight % Si02
1.0 %
2.0 O h
Initial conc., DH 55.85 pprn ~ e " (0.001 M), pH= 2.98 pH(+ sol) = 4.50 Final pH = 3.84 Poor separation V (sol) = 6.7 % (light brown) Solution: 23.0 pprn Sol: 391 pprn (Silicdfemc) = 35.7 pH(+ so1)>>9.0 Final pH = 3.90 No separation
Filtrate: 43 ppm
pH(+ sol)» 10.5 Final pH = 4.02 No separation
Filtrate: 42.0 pprn Sol: ---a-- PPm (Silicdfemc) = 142.9
(0.01 M), pH= 2.45 pH(+ sol)=2.69 Final pH = 3.90 Excellent V (sol) = 13.3 % (dark brown) Solution : 0.35 pprn Sol: 3488 pprn (Silicafferric) = 3.6
Initial conc., pH 558.5 ppm ~ e -
(0.02 M), pH=2.27 pH(+ sol)=2.50 Final pH = 3.82 Excellent V (sol) = 13.3 % (light brown) Solution : 0.90 ppm Sol: 6975 pprn (Silicdfemc) = 1.8
,, Initial conc., pH 1117 ppm ~ e *
pH(+ sol)=3 .O5 Final pH = 3.83 Excellent V (sol) = 18.7 % (dark brown) Solution : 0.85 pprn Sol: 2489 pprn (Silica/femc) = 7.1
pH(+ sol)=2.70 Final pH = 3.88 Excellent v (sol) = 2 1.3 % (dark brown) Solution : 0.9 pprn Sol: 4359 pprn (Silicdferric) = 3.6
pH(+ so1)»7.5 Final pH = 3.80 Excellent v (sol) = 20.0 % (light brown) Solution: 130.0 pprn Sol: 1749 pprn (Silicdfemc) = 14.3
pH(+ sol)=3 .O9 Final pH = 3.85 Excellent V (sol) = 26.7 % (light brown) Solution: 29.0 pprn Sol: 3393 pprn (Silicdferric) = 7. I
Table 5.2: Femc recovery from 100 mL 55.8 to 1,117 pprn Fe feeds using NaOH or acid
for pH control.
The ma..imum iron concentration in the feed solutions shown in Tables 5.1 and
5.2 is 1.12 g& Fe. Funher work focussed on 11 .O g/L Fe solutions, this being judged to
be an appropriate feed for the final stage of a 3 cycle process (Fig. 3.1) to produce about a
30 g/L Fe product from a 100 ppm-500 ppm Fe emuent.
Tests with silica addition were carried out in a similar way to that described for
Tables 5.1 and 5.2. Lime was prefened over sodium hydroxide for pH control, because
lime would be cheaper in an industrial application. Procedural differences were that
volume of SOI added, as well as its silica concentration were variables. and the amount of
lime needed for final pH adjustment was measured. Also. two different final pH levels
were studied.
In the first set of experiments with 11.0 g/L Fe feed (Table 5.3). final pH
adjustment with lime was to about pH 2.3. Aside from obvious trends (i.e. pH increases
with increasing sol addition), the most important conclusion is that pH 2.3 is too acidic
for recovering iron ont0 silica. At 2 % and 4 % Sioz additions. essentially al1 the iron
remains in the effluent. and the centrihiged underflows consist mainly of iron-free silica
with small amounts of gypsum. In al1 cases, a sharp interface and clear brown effluent
were produced.
Sol ( m u
PH (+ Sol)
Lime (mu
Final V(so1) Emuent PH ,
fi Fe % of Fe
2.28 8.0 3.44 43.9
1 Weight %
a = 0.58 g, 10 w/w % &SOJ b = 3.81 g, 10 wlw % H~SOJ
V (sol) = V/V % of metal '1oaded'-sol phase afier centrifugation in 45 mL flask.
Table 5.3: Ferric recovery from 100 mL 1 1.17 g/L Fe feed at pH 1-59 using lime (28 g/L
Ca0 suspension) for pH adjutment to 2.3.
In a subsequent series of tests. the final lime-adjusted pH was increased to 2.8
(Table 5.4) at which the maximum iron content of the effluent was only I l % of the
input. Since the conditions of Table 5.4 are reasonably suitable for iron recovery fiom
the effluent, the data have been analysed in more detail than for Table 5.3. As already
seen in Tables 5.1 and 5.2, iron levels in the emuent tend to increase with increasing
silica addition, which also gives more voluminous centrihged underfiows (v/v % sol
phase). A bulky undedow rneans production of a concentrated iron solution in
subsequent redissolution becomes more dificuit.
The best result was obtained with the smallest silica addition (20 mL 1 .O % SiOr)
at which 99 % of the iron was recovered to give a sol containing 37 g/L Fe (calculated by
mass balance), the concentration factor being 3.4 ((g/L Fe in sol) / (g/L Fe in feed)). It is
interesting that at the highest silica addition (80 mL, 4 % Sioz ), the 'concentrated'
product at 8.54 gR. Fe is actually more dilute than the 11.1 g/L Fe feed due to the
voluminous underflow produced and the dilution effect of silica and lime additions. For
al1 tests in Table 5.4, the overtlows were totally clear and free of suspended solids.
Another important parameter is the silica loss to the effluent. As expected,
increased silica addition, either through higher reagent concentration or volume added.
gives higher dissolved silicon contents. These losses arnount to between 13 % and 26 %
of the input silica over the range of experimental conditions. Fonunately the best
conditions for iron recovery are also those giving the lowest silica loss of 1 3 %.
Also included in Table 5.4 are calcium analyses of the effluents. Calcium
contents range fiom 365 ppm to 570 ppm, while saturation with respect to gypsurn is at
about 600 ppm Ca [118]. Since dissolved calcium represents between 5 W and 10 % of
the added lime, it is obvious that a high proportion of lime is converted to gypsum, which
will be recycled with regenerated silica. Potential methods for dealing with gypsum
build-up are discussed in Section 5.6.
a) 1.0 % Si02 sol I 1
Experimental data
b) 2.0 O/O Si02 sol
Calculated pararneters
Exnerimental data 1
Sol pH Lime Final V (sol) Effluent (mL) (+Sol) (mL) pH ppm Fe ppm Si ppm Ca
20 1.72 48 2.83 18.7 80 180 564 40 1.86 49 2.80 29.3 216 450 510 60 2.00 47 2.75 30.7 480 1,000 485
1 80 2.11 1 4 1 2.80 28.0 780 1,700 485 !
Lime (mL)
50 52 50 47
Sol (mL)
20 40 60 80
Final pH 2.80 2.85 2.78 2.84
% of Si02 in effluent
13.5 13.6 17.5 15.4
Sol (mL)
20 40 60 80
Calculated Daramefers
pH (+ Sol)
1.66 1.76 1.81 1.85
Sol Siofle gL Fe (g/L Fe in sol) / % of Fe in O/O of Si01 (mL) (molar) in sol (g/L Fe in feed) emuent in effluent
20 , 0.33 3 5.4 3.16 1 .O 13.2 40 O .67 19.7 1 1.76 2.6 16.1
V (sol)
17.3 17.3 17.3 20.0
SiOJFe (molar)
Effluent
g/L Fe ; in sol
ppm Fe 80 152 322 214 :
(g/L Fe in sol)/ (g/L Fe in feed)
O. 17 [ 37.6 0.33 1 32.7
ppm Si 90 160 3 15 317
% of Fe in emuent
3.37 2.93 2.6 1 2.13
0.50 0.67
ppm Ca 570 540 540 530
1 .O 2.2 5 .O 3.5
29.1 23.8
c) 4.0 % SiOs sol
Sol (mL)
20 40 60 80
V (sol) = V/V % of metal '1oaded'-sol phase aber centrifugation.
Sol (mL)
20 40 60 80
Tabk 5.4 : Femc recovery fiom 100 mL 1 1.17 g/L Fe feed at pH 1-59 using lime (28
g/L Ca0 suspension) for pH adjustment to 2.8.
pH (+Sol)
1.84
5.3 Thickening (Fig 3.1)
At many mine sites, large expensive settlers/clarifien (typically 150 fi. diameter)
are needed to thicken the waste sludge fomed by adding lime to effluents. Flocculant
addition (e.g. Percol 338) increases free settling rates which reduces the size and cost of
the thickeners required. In the present work, activated silica is effectiveiy replacing
Percol. while providing an adsorption medium when added in g/L quantities. A study of
settling characteristics fol lowing activated si lica addition is thus needed.
Settling tests were carried out with a synthetic feed containing about 500 ppm Fe.
and with effluent samples fiom the 'Balsam stream' (Inco Ltd.. Sudbury, Ont.). Unless
Siofle (molar)
0.66 1.33 2.00 2.66
Lime (mL)
45
g/L Fe in sol 29.7 1 1.3
Final pH 2.87
2.12 2.35
2.82 2.81 2.78
47 33
(g/L Fe in sol) I (g/L Fe in feed)
2.66 1 .O1
V (sol)
22.7
, 2.49 23
9.21 1 0.82 8.54 1 0.77
52.0 61.3
% of Fe in efiïuent
1 .O 1 . 1
, Effluent
% of Si02 in effluent
26.3 20.1
2.1 7.3
ppm Fe 74 136 310
19.7 25.3
60.0 1 1.000
ppm Si 770
1,670 2,650
ppm Ca 365 385 377
4.660 , 375
othenvise specified, the silica addition was the value found suitable for iron recovery
fiom 558 ppm Fe solution, as described in Table 5.1 (1 w/w % Si02 stabilized at 50 % of
gel time, 20 mL/lOO mL iron feed, final pH adjustment with lime to 4.0). This procedure
gave a slurry containing 2 g/L Si02 (as sol) and pH increases from about 2.2 (initial) to
2.7 (after sol addition) to 4.0 (after lime addition). Thus the underflows are a mixture of
adsorbed femc species. femc hydroxide, sol. gypsum and lime. In some cases,
underflow % solids were also measured either before or afier centrifugation. Settling data
are presented as 'Distance travelled' (downwards by interface) vs. time.
5.3.1 Effect of Container Size
For the synthetic feeds, three different containers were used (Fig. 3.3), the slurry
quantities being I L for the 'small' and 'medium' containers and 14.4 L for the large
container. The limited arnount of ' Balsarn Street' effluent available restricted sarnple
size to 500 mL. and settling tests were performed in a standard 500 mL laboratory beaker.
Figures 5.1 and 5.2 give the settling curves obtained for synthetic feeds, along with the
volume % appearing as underflow at the end of the test. The fiee settling rate (about first
10 minutes of settling) increases with increasing container size. Unfonunately, the
tendency for activated silica to adhere to a glas surface enhances the 'wall effect'
associated with settling tests. Thus, although the absolute values measured are not the
same as would be obtained in a commercial application. the observed settling rates can be
used in a relative sense for comparing effects of pH, silica concentration etc.. providing
the measurements are made in the sanie container.
While fiee settling rates depend on container size, the degree of compaction at the
end of settling is the same in al1 three cases (Fig. 5.2). Although the actual settling rate
measurements extended only to 150 minutes, the volume percentages occupied by the
underflows were detemined to be 36 vlv % to 37 v/v % after 24 houn. In the following
sections, the large container was rarely used, because the production of 14.4 L samples
was time consuming and cumbersome.
Time (minutes)
Figure 5.1: Effect of container size on settiing rate.
Conditions: 558 ppm Fe feed, 1 .O w/w % SiO: 'N' sol stabilized at 50 % of gel time. 200
mL sol11 L feed solution. pH adjusted to 4.0 with lime.
80 t medium (1 L)
70 t lrrga (14.4 L)
Time (minutes)
Figure 5.2: Volume % as underflow vs. tirne for results s h o w in Figure 5.1.
5.3.2 Effect of Degree of Celation
Settling tests were carried out. using the 'medium' container, with activated silica
stabilized at 20 %. 50 % and 80 % of the gel time (18 minutes). The data seen in Figure
5.3 show that the 50 % of-gel-time additive gives the best results. The reason for this
optimum performance with the 'half gelled' sol is unclear, but it should be noted that
activated silica used to remove suspended solids from dnnking water also performs best
when stabilized at 50 % of the gel time [80].
5.3.3 Effect of pH
Figures 5.4 and 5.5 show the influence of pH on settling behaviour. Both the
-medium' and 'large' containers were used. Final lime-adjusted p h levels were 4, 7 and
92
10. When using the 'medium' container, settling characteristics at pH 7 and 10 are
effectively the same and slightly better than at pH 4. With the large container. results are
also insensitive to pH. Thus settling behaviour is not a major factor in the choice of pH
for the preceding iron recovery step. The underflows (24 hou retention) occupied 38 %
to 44 % of the total volume.
Percent of total gel time + 20%
-m- 50%
-&- 80%
Time (minutes)
Figure 5.3: Effect of degree of gelation on settling rate.
Conditions: 558 ppm Fe feed, 1 .O w/w % Sioz 'N' sol, 200 mL solIl L feed solution, pH
adjusted to 4.0 with lime.
Final pH
Time (minutes)
Figure 5.4: Effect of final solution pH on settling rate ('Medium' container).
Conditions: 558 ppm Fe feed, 1 .O w/w % Si01 'N' sol stabilized at 50 % of gel time, 200
mL sol/ l L feed solution, pH adjusted with lime.
Final pH
Tim e (minutes)
Figure 5.5: Effect of final solution pH on settling rate ('Large' container).
Conditions: Same as in Figure 5.4.
5.3.4 Effect of Silica Concentration
Since activated silica acts as a flocculant, it is expected that the arnount added
wouid have a major effect on settling behaviour. Results were obtained for various silica
additions, as indicated in Figure 5.6, the arnounts added being similar to those studied in
the iron recovery tests of Table 5.1. It is seen that free settling rate ('medium' container)
decreases with increasing silica addition while volume % appearing as underflow
increases. These effects may be due to a progressively more negative net charge on the
activated silica, because the amount added is increasing relative to a constant iron
concentration. The repelling effect of negatively charged polymers would discourage
flocculation and lead to more voluminous underflows.
The results of Figure 5.6 are replotted as 'Distance travelled' (by interface) vs.
'Silica concentration', as seen in Figure 5.7. Each vertical line in Figure 5.7 represents
one of the settling tests in Figure 5.6, 'R' being the molar ratio Si02 / Fe . The "x" axis
of Figure 5.7 represents the calculated overall silica concentration in the slurry. This is
not a measurable quantity since immediately upon addition to the iron feed. the silica
segregates into a hi& silica underflow and low silica overflow.
The decreasing settling rates with increasing molar excess of silica over iron is
clearly seen in Figure 5.7. It is unlikely that settling behaviour is king influenced by
silicate formation. While ferrous iron readily foms simple silicates such as fayalite.
FeSi04 , femc iron is generally present in silicates only in combination with another
cation such as potassium, magnesiurn or calcium. The deterioration in settling properties
with increasing silica level is considered to be due mainly to the repulsive effect of excess
negative charge on silica surfaces.
With total absence of silica (lime addition only), a faint. ill-defined interface
foms which still persists d e r the standard 30 minute settling tests used in the present
study. If the sole purpose of adding silica is to act as a flocculant, Figure 5.7 suggests
that a few ppm would suffice, but the neutralizing effect would be negligible. Macro
amounts of activated silica are required to raise the pH and achieve the concepts of a
cecyclable neutralizing agent with production of a concentrated product solution. There
is thus a compromise between adding enough silica to form a metal-bearing silica
underflow while retaining reasonable settling rates.
The degree to which silica and iron have become concentrated in the underflow
influences the iron concentration achievable when subsequently redissolved. Underflow
properties before and after centrifugation are compared in Table 5.5. Centrifugation
produces concentration factors of about 4, yielding underfiows at 2.4 io 2.8 w/w % solids,
after oven drying. This approximates the range of solids content to be expected if,
during underflow drying, al1 the activated silica appears as a solid. and al1 the iron is
oxidized to hematite or goethite. Thus the measured underflow wlw % solids does not
tnily represent the actual underflow w/w % solids at the end of the settling test.
factor from solids after
Table 5.5: Underflow properties corresponding to settling tests of Figure 5.6.
U'flow vlv % after
centrifugation
Weight % S i 0 1 (znL sol added to 1 L
feed)
U'flow v/v % after settling for 30 min.
Sol (wlw VfSiO,) mL soVt L Fe fccd
Q-1 5 0
*l 100
Time (minutes)
Figure 5.6: Effect of sol concentration on settling rate.
Conditions: 558 ppm Fe feed, 'N' sol stabilized at 50 % of gel time. pH adjusted to 4.0
with lime.
Silica (SiO,) concentration (ppm)
Time (min)
Figure 5.7: Effect of sol concentration on settling rate. for results in Figure 5.7. R= Molar ratio (Si02 1 Fe).
5.3.5 Effect of Recycling Regenerated Silica (after redissolution step)
An important concept seen in the flowsheet of Figure 3.1 is the ability to
regenerate and recycle activated silica after metal redissolution. Figure 5.8 compares
settling curves for fresh and recycled silica additive. In the first cycle (fiesh silica), the
'standard' procedure was followed (drtails on Fig. 5.8). Afier settling and underflow
centrifugation, iron was redissolved from centnfuged underflow using concentrated
sulphuric acid undrr conditions established as suitable (pH 1.30, 30 min. contact), as
discussed below in Section 5.4. After a second centrifugation to separate relatively
concentrated iron solution (about 5 g/L Fe) from regenerated silica, the latter was recycled
and added to fresh iron feed for a second settling test carried out in exactly the same
rnanner as for the fint cycle.
Figure 5.8 shows that, although the settling rate with recycled silica is slightly less
than for fresh reagent. the underflow compaction afier 30 minutes of settling is essentially
the same in both cases (44 vlv % - 45 v/v % as underflow). It would be surprising if
identical settling rates had been obtained since one cycle produces a combined silica loss
of about I l % into the effluent and iron solution product, these streams being readily
identifiable in Figure 3.1. In addition, extra lime has to be added upon recycle to
neutralize the acid needed for iron redissolution. The resulting gypsurn build-up is
reflected in a higher w/w % solids in centrifbged underflow when using recycled silica
(3.97 wlw % solids) compared with the value fiom the initial cycle with fresh reagent
(2.39 w/w % soiids). The issue of gypsurn formation is discussed in Section 5.5.
Time (minutes)
Figure 5.8: Effect of recycling regenerated sol on settling rate.
Conditions: 558 ppm Fe feed, 1 .O w/w % Si02 'N' sol stabilized at 50 % of gel time, 200
mL solll L feed solution. pH adjusted to 4.0 with lime.
5.3.6 Effect of Recycling Iron-'loaded' Silica (after recovery step)
In conventional effluent treatment using lime, part of the underflow is recycled to
increase the solids content to a level suitable for tailings discharge. In the present case. it
is also beneficial to increasç undefflow wlw % solids so that a concentrated solution
product is more readily generated. Three consecutive settling tests ('medium' container)
were carried out in which the centrifuged underfiow fiom the first experiment was added,
dong with activated silica and lime? to fiesh iron feed. AAer a second settling test. the
centrifuged undeflow recycling procedure was repeated, and a third settling expenment
performed. In each case, settling data were obtained using the same slurry volume ( 1 L)
99
to make the results directly comparable. A constant volume could be maintained only by
compensating for the various additives required upon recycle (recycled undefflow,
activated silica, lime). Before the settling tests in cycles 2 and 3, the sluny was stirred
and simultaneously the amount in excess of 1 L was rejected.
Figure 5.9 shows that the increased solids loading in cycles 2 and 3 has a
negligible effect on settling rates, although the underflows are slightly more voluminous
than in cycle 1, as expected (Table 5.6). Table 5.6 also shows the build-up in wlw %
solids for uncentrifbged underflow when recycled. The value of 3.74 wlw % solids for
centrihiged underflow fiom cycle 3 compares with 2.4 to 2.8 w!w % for a single cycle
(Table 5.5).
[ Cycle 2 1 30.0 1 1.20 1 .
Cvcle i
1 Cycle 3 1 29.0 1 1.74 (3.74 after centrihigation) 1
Tabk 5.6: Underflow properties corresponding to settling tests of Figure 5.9.
U'flow v/v % after settling for 30 min.
22.0
U'flow w/w O h solids after settling for 30 min.
0.91
Sol additions to 1 L Fe feed
A
E c a I Q)
E w Q) 3 0 Ca Ci m E
Figure 5.9:
O\ Cyc le1 Cycle 2 cycle 3
1 \ - 50 nL 'fresh' sol - 50 nL 'frah' sol - 50 mL 'fresh' sol
- 150 mL Fe-loaded - 220 mL Fe-loaded Sol from cycle 1 Sol from cycle 2
-e Cycle 1
+Cycle 2
7 : O 5 10 15 20 25 30
Time (minutes)
Effect of recycling iron-loaded sol on settling rate.
Fresh sol is 1 .O wlw % Sioz 'N' sol stabilized at 50 % of gel time, 1 L 558 pprn Fe feed
at pH 2.4. Final pH adjustment to 4.0 with lime for each cycle. Stirred slurry in excess of
1 L rejected before settling.
5.3.7 Settling Tests with 'Balsam Stream' Effluent
Samples were obtained from the 'Balsam stream' of Inco Ltd, Sudbury. Ont. All
settling tests in Figures 5.1 to 5.9 were carried out with synthetic 558 pprn Fe feed to be
comparable with 'Baisam Stream' effluent. which was expected to contain about 500
pprn Fe. In reality, the latter actually analysed 60 pprn Fe (27 pprn Ni, 1.3 ppm Cu. pH
3 . 9 , aithough these concentrations change every month. Settling data (Figure 5.10) were
obtained for various silica additions in a similar manner to that applicable to Figures 5.6
and 5.7 for synthetic 558 pprn Fe solution. The silica additions for 60 pprn Fe 'Balsam'
101
effluent were adjusted to cover a range of R (molar ratio Si02 / Fe) similar to that for
synthetic 558 pprn Fe feed. Although Figures 5.6 and 5.10 are not quantitatively
comparable, the best results (fastest settling and most compact underflow) were again
obtained with the smallest silica addition. For the 'Balsarn' effluent, the fastest settling
rate was about 2.4 mih, which is about an order of magnitude higher than found for this
effluent in the regular IimefPercol system [119]. The uncentrifuged underflow
corresponding to this 2.4 m/h settling rate was the most compact of those measured for
'Balsarn Stream' effluent, at 8 v/v % of the total slurry.
Molar ratio (R) Sol addition1 l L Fe leed SiO, / Fe ppm SiO, i
+100 mL 14.2 909 f
-+ 80 mL 11.5 740 1 -t 40 mL 6.00 385 1
I
* 20 mL 3.06 1% 1
Time (minutes)
Figure 5.10: Effect of sol concentration on settling rate for 'Balsam Stream' (Inco)
effluent (60 ppm Fe, 27 ppm Ni, 1.3 ppm Cu). 1 .O wlw % Sior 'N' sol stabiiized at 50 %
of gel time. pH adjusted to 5.5 with lime.
5.4 Iron Redissolution (Fig. 3.1)
Afier thickening and rejection of low-iron overflow, the iron content of the
silica/gypsurn underflow can be recovzred, with concentration, by redissolution in
concentrated acid. Ii is unrealistic to expect production of a sufficiently concentrated
solution for sewage treatment (- 30 g/L, Fe) fiom, Say, a 500 ppm Fe effluent using only
one recovery/redissolution cycle. Thus a three cycle system was studied in which the
dissolved iron content was progressively built up fiom about 500 ppm Fe to at least 30
g/L Fe.
The first cycle was performed using a 558 ppm Fe feed to the initial recovery step,
from which the centrifuged underflow was treated for iron redissolution. In Table 5.7,
column a). iron recovery was carried out as described for the 558 ppm Fe, 1.0 % Si02
'box' of Table 5.1. Afier centrifugation, the residual solution containing 1.9 ppm Fe was
discarded. The pH of the undertlow was adjusted to 0.77 with a few drops of
concentrated sulphuric acid. After a contact time of 20 minutes, the system was
centrifuged again, giving a clear brown aqueous phase containing 8.0 g/L Fe, which
corresponds to a concentration factor of 14.3 ((a Fe in product)/(g/L Fe in feed)). An
opaque, white lower Iayer of activated silica (with gypsurn) was formed, this being
essentially iron-free. A contact time of 20 minutes was chosen based on preliminary
kinetic experiments in which analyticai samples were withdrawn at various time intervals
up to 1 hou. The iron distribution was 99.6 O/O into the product solution, 0.4 % into the
effluent and 0.0 % retained by the activated silica.
Column b) shows a similar cycle to that of column a), the important differences
k i n g scale-up (100 mL to 900 mL feed), and less acidic iron redissolution (pH 1.25
instead of 0.77). The significance of scale-up is that, afler the recovery step, a larger
centrifuge was required than used for 100 mL batches (Section 3.4). The 'big' centrifige
did not compact silica underflows as readily as the 'srnall' machine, so that more effluent
was carried through to iron redissolution. The net result was a 4.0 g/L Fe solution
(concentration factor 7.2) which is half that generated when using the small centrifuge.
The 4.0 g/L Fe product codd probably have been further concentrated by increasing the
r.p.m. and/or centrifugation time, but a concentration factor of 7.2 was considered
sufficient for demonstraiion purposes.
1 % of feed Fe in effluent 1 0.4 1 0.3
Recovery step a) b)
i
1 Final pH 1 0.77 [ 1.25
I , ppm Si in effluent , % of feed Si in effluent
Feed vol. (mL) 1.0 w/w % Si02 addition (mL) Final (lime adjusted) pH m m Fe in effluent
II % of feed Fe redissolved 1 99.6 1 99.7
35 4.1
Contact time (min) e/L Fe redissohd
1 O0 20
4.00 1.90
35 4.1
Table 5.7: Iron recovery/redissolution cycles applied to 558 ppm Fe feed at pH 2.45.
900 180
4.05 1.50
20 8 .O
1
20 4.0
Si dissolved COL feed Si dissolvrd 1.6 12
1.3 1 1
An increase in redissolution pH fiom 0.77 (column a)
gave an opaque, white silica layer. its colour being a good
redissolution eficiency. Even if redissolution is faster at a
to 1.25 (column b) again
qualitative guide to iron
lower pH level. the main
advantage of being able to raise redissolution pH is a decreased tendency for the silica to
gel. At pH 0.77, an unpumpable gel forms in a few houn. while at pH 1.25. gelation
occurs afier about one day. An increased gelation tendency at low pH means that in an
actual process (Fig. 3.1). the time required to recycle acidified silica back to the more
alkaline conditions of the iron recovery step is critical. Also. it would not be possible to
insert a surge tank between iron dissolution and iron recovery. unless the recycle Stream is
partially neutralized.
Silicon analyses in columns a) and b) for both the effluent after iron recovery and
product solution afier iron redissolution should be interpreted with caution. Although
there is a clear trend to increased silica loss with decreasing pH (i.e. 4 % loss at pH 4.0
and 12 % loss at pH 0.77 - Table 5.7), the dissolved silicon level also depends on
exposure time at the specified pH. especially under strongly acidic conditions. For
example. the 20 minute contact times shown in Table 5.7 for iron redissolution are
optimum for effectively total iron transfer into a concentrated solution.. Further retention
beyond 20 minutes serves only to increase g /L Si in the product solution.
A recoverylredissolution cycle was also carried out in a similar manner to those in
Table 5.7 but with 1.12 g/L Fe feed. The results are shown in Table 5.8. Iron recovery
was performed as described for the 1.1 17 ppm Fe, 1 .O % Si02 box of Table 5.1, giving
essentidl y iron- free effluent. The iron-bearing centri fùged underfiow was acidi fied to pH
0.95 for 30 minutes. AAer centrifugation, a well-defined interface was obtained dividing
a clear brown ovefflow fiom a silica underflow. This combination of pH and contact
tirne gave complete iron redissolution into a 12.0 g/L Fe product (concentration factor =
10.7).
1 1 Feed vol. (mL) 1 100 II 1.0 W/W O h Si02 addition (mL
4.05 1 ppm Fe in effluent % of feed Fe in emuent 0.05
. OS I / Redissolution s t e ~ 1
Table 5.8: Iron recovery/redissolution cycle applied to 1 .12 g/L Fe feed at pH 2.27.
, ina al p~ Contact time (min) pC/L Fe redissolved % of feed Fe redissolved
The final set of recovery/redissolution tests treated a 'high-iron' feed of 1 1.2 glL.
which represents the last stage of a three cycle system which would treat progressively
higher iron concentrations in each successive cycle. Table 5.9 gives the results obtained.
For al1 cycles in Table 5.9, iron recovery followed the procedure applicable to 1 1.2 g/L Fe
feed in Table 5.4, and almost total recoveries into the silica were again obtained.
Subsequent acidification of the centrifùged underflows was carried out at constant pH
(1.3) and contact time (30 minutes) to allow study of the effect of silica addition on iron
redissolution.
0.95 30
12.0 1 O0
Although the maximum '% of feed iron redissolved' (93.6 %) is obtained at high
siiica addition (40 mL 4.0 % Si02 in recovery step), the corresponding glL Fe (26.5 g/L)
is not the highest value in the series of tests in Table 5.9. This is because centrihigation
does not compress a bulky silica underflow as readily as that generated from a relatively
low silica addition. The most concentrated product, at 39.3 g/L Fe, yields a
concentration factor of 3.5. It should also be noted that the 39.3 g/L Fe product
contained 3.70 g/L Si, which represents a 46 % silica loss. Although dissolved silicon is
unlikely to create an unacceptable discharge, the economics of close to 50 % 'make-up'
upon recycle may be unattractive.
Figure 5.1 1 shows typical products obtained from the tests in Tables 5.7 - 5.9.
Centrifuged underflows fiom the recovery step are seen in Figures 5.1 1 a) and b) for 558
ppm Fe and 1 1.2 g/L Fe feed respectively. and show iron-loaded sols beneath clear
overflow. In the latter case, the lime addition was sufficiently high that a separate
gypsum layer was produced. Figure 5.1 1 c) shows a concentrated iron product (about 35
g/L Fe) above centrifuged undedow fiom the redissolution step following recovery From
11.2 g/L Fe feed. lron redissolution was essentially complete leaving a white
silicdgypsum mixture.
In surnmary, the results of section 5.4 suggest it would be possible to upgrade a
500 ppm Fe feed to about 40 g/L Fe using activated silica, with lime for pH control, for
iron recovery, foliowed by redissolution in concentrated acid. A three cycle system
would be required in which progressively more concentrated iron solutions are treated.
The most significant problems are especially evident in the final 'high iron' cycle.
Gypsum build-up is a major concem, and potential methods for handling this difficulty
are discussed in section 5.6. Although mainly an economic issue, the operation of a
'high iron' cycle is also likely to give major silica losses.
Figure 5.11: Typical samples obtained from the tests in Tables 5.7 - 5.9 afier
centrifugation. a) and b) are sarnples from the 'recovery step' for 558 ppm Fe and 1 1.2
d L Fe feed respectively. and show iron-loaded sols beneath clear overflow. In b) a - separate gypsurn layer was produced. Figure c) shows a concentrated iron product (about
55 g/L Fe) above centrifuged underflow (white silicalgeypsum mixture) from the
redissolution step following recovery from 1 1.2 g/L Fe feed.
5.5 Effect of Temperature
Effluent treatment is perfomed under outdoor arnbient conditions with seasonal
temperature variations. Many flocculants do not perform well over the temperature
range covering winter and summer so giving dirty overflows. lron recovery1redissolution
tests have k e n carried out at 5 OC, 25 OC and 45 O C , using the water bath arrangement
described in Section 3.3. A 558 ppm Fe feed was studied with silica additions and pH
values previously found to be suitable at 25 OC (Table 5.1 for iron recovery, Table 5.7 for
iron redissolution).
For the iron recovery step, the required solutions were placed in the water bath for
1 hour prior to the experiment. The data in Figure 5.12 were generated by adding silica
to stirred 558 pprn Fe solution followed by incremental lime addition over a 12 minute
period to reach pH 40. Time zero is the point at which al1 additions had been made. At
various time intervals. samples were withdrawn and centrifuged for five minutes to give
overflows for analysis. Initially, dissolved iron recovery follows the expected trend of
slowest rate at coldest temperature. Thermodynamically, the equilibrium femc
concentration should decrease with increasing temperature [IZO]. However the final,
120 minute, situation shows the room temperature result to be the best, with I pprn Fe
effluent, while the hot (45 OC) and cold (5 O C ) systems had 6 ppm Fe and 8 ppm Fe
respectively .
The shape of the 5 O C cuve shows that a steady state had not been reached after
120 minutes. At 45 OC. some fine particles were seen in the effluent. This suggests that
the final, measured dissolved iron level has been enhanced by colloidal ironîontaining
particles which have detached fiom the silica and have escaped filtration in preparation of
the analytical sample. The final dissolved femc concentration was also reported to
increase with increasing temperature during formation of hydroxide precipitates resulting
fkom lime addition [120].
Studies of iron redissolution as a hinction of temperature were carried out using
ironcontaining silica underflows conesponding to the 25 O C , test of Figure 5.12. Stirred
undefflows (uncentrifuged) were held in the water bath for one hour before concentrated
acid addition such that pH values reached about 1.2 (Fig. 5.13). Time zero on Figure
5.13 is at the instant of acid addition. Samples were withdrawn at timed intervals then
centnfuged to give product solutions for analysis. Redissolution was complete after
about 5 minutes within the temperature range 4 O C to 45 O C . The concentration factor of
5 is less than found previously (Table 5.7) because uncentrifuged underflow From iron
recovery was used for iron redissolution.
O 30 60 90 1 20 150
Time (minutes)
Figure 5.12: Femc removal fiom solution as a function of temperature.
Conditions: 5 58 ppm Fe feed, 1.0 w/w % Si02 'N' sol stabilized at 50 % of gel time, 200
rnL sol/ l L feed solution. final pH adjusted to 4.0 with lime.
Time (minutes)
Figure 5.13: Redissolution of femc as a function of temperature. Iron loaded sol
generated as in Figure 5.12 ai 25 OC then stored for one day before redissolution.
5.6 Gypsum Formation
Gypsum would rapidly build-up during operation of a closed cycle for iron
recovery/redissolution. because most of the acid needed for redissolution must be
neutralized with lime before reuse in the iron recovery step. This problem is especially
pronounced for iron, because the required redissolution pH is relatively low when
compared with values needed for most other metals, as demonstrated in the present work
for copper. nickel and magnesium.
A detailed evaluation of gypsurn/silica mixtures generated during iron
recove~redissolution cycles has been carried out at the Technical Center of National
Silicates Ltd. [121]. Techniques used included optical microscopy, X-ray diffraction
O(RD) and energy dispersive X-ray analysis (EDXRA). Examination by rnicroscopy
mainly involved use of polarized light to enhance the visual difference between gypsurn
and activated silica, both phases being whiteish. A few photographs were also taken
using a stereoscopic microscope which generaies 3-dimensional images. XRD identified
the crystalline compounds, while EDXRA qualitatively showed which elements were
present in individual phases.
The two samples examined were centrifuged underfiows from the redissolution
step prepared as described in Table 5.9, columns a) and e). Thus the sarnples would be
expected to contain gypsum (since lime was added in preceding recovery step), activated
silica and any iron compounds that had not redissolved. The column a) sample (sample
1) had three distinct layers (top-yellowish, middle-white, bottom-brown) and was chosen
to determine if this layering effect was due to silica/gypsum segregation. In sample 1.
about 70 % of the iron had redissolved. The column e) material (sample 2) was an
unlayered uniformly white product. At least 90 % of the iron had redissolved in sarnple
2. The sarnples chosen represent worst case scenarios in that the limelsilica ratio was
large and equivalent to a gypsumlsilica mixture of only about 5 w/w % Si02 (sample 1)
and 17 wlw % Si02 (Sample 2).
Figure 5.14 contains photomicrographs of the three layers of sample 1. Figure
5.14 a, b (top layer) are taken under polarized light (light background) and show,
respectively, activated silica containing trapped hydrated iron oxides, and acicular crystals
of gypsum. However, activated silica plus iron oxides often appear as a skin or c m t on
gypsum crystals, as seen in the stereoscopic image of the top layer in Figure 5.14 c. The
white rniddle layer consists mainly of aggregated, fine crystalline needles of gypsum, as
s h m in Figure 5.14 d, where crossed polarized light (dark background) was used.
Figures 5.14 e,f represent the brown bottom layer photographed with and without crossed
polars respectively. Use of crossed polarized light mainly shows aggregated gypsum
needles some of which sparkle due to a thin surface layer of activated silica. Without
crossed polars, brown iron oxides are seen attached to activated silica and gypsum. As
expected, XRD analysis confirmed the presence of gypsum. EDXRA showed elements
consistent with the acicularlneedle crystals being gypsum (e-g. Ca, S) with the
yellow/brown areas being silica and iron compounds.
Semi-quantitative microchemical analysis gave the following results for the three
layers of sample 1 [12 11:
Weigbt O h
Gypsum
Given the accuracy of the semi-quantitative analysis (* 5 w/w %), it can be seen that the
white middle layer is essentially silica-fiee gypsurn, while some concentration of silica
has occurred in the brown bottom layer. In sample 1, about 30 % of the iron did not
Silica + iron oxides
Top layer
90
10
Middle Iayer
95
Bottom Iayer ,
80
5 20
redissolve. It seems that most of this residual, undissolved iron is attached to silica, and
the extra weight is sufficient to carry a high proportion of the silica into the bonom layer.
In contrast, when iron dissolution is essentially complete (sample 2), the layering effect
is absent, and a visually uniform silicdgypsum mixture is obtained. Thus natural
segregation of silica and gypsum does not occur to a degree that would permit a physical
separation, while there is also some attachent of the solid phases present.
Photomicrographs of sample 2 did not show any features that were not observed in Figure
5.14.
The simplest way to avoid gypsum formation is to neutralize with sodium
hydroxide instead of lime. but this is uneconomic, especiaily with such a low value
product as ferric sulphate solution. Gypsum notation might be considered, given the
wide difference in surface properties between gypsum crystals and arnorphous silica.
However, many gypsum particles have a silica coating. Altematively. providing metal
redissolution is complete. the gypsum/activated silica mixture could be seen as a
marketable byproduct, such mixtures being used in the manufacture of cernent. The
most feasible option appean to be the introduction of an acid wash step such that the
excess acid is neutralized outsidr the recovery/redissolution cycle, as shown in Figure
5.1 5.
With reference to Figure 5.15, gypsum formation within the
bRecovery~/bRedissolution' circuit originates mainly From neutraiization of acid added in
'Redissolution' that does not leave with the *product' solution and remains trapped in the
centrifugeci underflow. In principle. this residual acid could be washed out of the silica
and neutralized outside the main process loop in a manner that is self-evident from Figure
S. 15. The volume of solution to be treated in a silica wash would be relatively small,
since about 90 vlv % of the 'Feed' has k e n rejected as 'Effluent' while most of the iron
has been removed as 'Product'. A cntical requirement is that the silica must retain about
1 00 ppm Fe after ' Redissolution' and 'Si02 wash', because totally iron- fiee activated
silica will not easily form a separate phase when centrifuged.
Acid removal by washing the silica would not be 100 % efficient. and lime would
still be needed for residual acid removal in 'Recovery', but the amount would be small
relative to that needed for the silica wash solution. There would always be a circulating
load of gypsum. but the arnount could be controlled and held at a roughly constant level
in contrast to the continuous build-up that would result from the absence of silica
washing. The effectiveness of a silica wash has been demonstrated previously in
co~ec t ion with cobalt recovery when following the steps of Fig 5.1 5 ( 1 071. About 85 %
of the acid added for 'Redissolution' could be washed fiom the silica and neutrdized
outside the Recovery/Redissolution circuit.
Figure 5.14 a) Top layer. polarized light, (x 150).
Figure 5.14 b) Top layer, polarized light, (x 600).
117
Figure 5.14 c) Top layer, stereoscopic image. (s 3).
Figure 5.11 d) Middle layer. crossed polarized light, (x 60).
Figure 5.14 e) Bottom layer, crossed polarized light. (x 60).
Figure 5.14 f) Bonom layer' polarized light. jn 60).
Figure 5.14: Photomicrographs of centrifuged underfiow after iron redissolution as in Table 5.9, column a).
'Small' lime addition Feed
Activated silica
Recovery 1
Centrifuge * Product
Centrifuge r-l' 'rJ -Largev lime addition I
Gypsum
SiO, wash 1 A u
Figure 5.15: Recovery/Redissolution circuit with silica wash.
Neutralization
CHAPTER 6
Nickel Recovery from Tailings Pond Sludge by Acid Leaching
and Treatment with Activated Silica
6.1 Sludge Source
An example of how direct addition of activated silica to an effluent can remove
and concentrate a dissolved metal has been described in Chapter 5, where the target metal
was iron. Another type of application relates to the treatment of sludges formed by
conventional lime addition to an effluent Stream. When acid mine drainage generated in
the tailings ponds of nickel sulphide processing operations is treated with lime, a
volurninous waste sludge is produced containing femc hydroxide and other metal
hydroxides, mainly of nickel and magnesium. The sludge may cany a significant quantity
of nickel (up to 6 % on a dry weight basis) which was not recovered during the ore
processing steps. This nickel content may be higher than in most natural ores, which
makes the sludge a good secondary nickel source.
The particular sludge involved in the present work came from the 'Upper pond' of
the waste treatment operations of Inco Ltd., Sudbury, Ontario. Figure 6.1 is a simplified
schematic [122] showing how the treatment of various waste materials are inter-related.
It c m be seen that 'Upper pond' sludge results fiom lime addition to a mixture of
effluents from the 'Lower pond' and the 'Tailings area'. Smelter slags are the main
component of the former, while the latter comprises flotation tailings, refinery wastes,
sludge from the waste water treatment plant, and mine water to which lime is added. The
'as received' sludge from the 'Upper pond' contained 86 w/w % water, and the
composition of the dried material was (by weight) 6.2 % Ni, 16.8 % Fe. 20.3 % Mg, 0.2
% Ca.
6.2 Sludge Treatment Flowsheet
The experimental work centered around the flowsheet show in Figure 6.2.
Although there are obvious similarities in general concepts to those seen in Figure 3.1
(Cu and Fe recoveries), some additional feanires on Figure 6.2 are specific to sludge
treatment. The solution to be treated using activated silica must first be generated by a
preliminary leaching step. In the present case, the solution was produced by sulphuric
acid leaching at a carefully controlled pH (3-4) such that nickel and magnesium dissolve
but iron (femc) remains in the leach tesidue. Thus iron is removed before activated silica
treatment. With appropriate pH control using lime. nickel can be recovered ont0
activated silica with partial selectivity ovet magnesium. The magnesium-containing
effluent resulting fiom thickening and centrifugation of activated silica underflow can be
generated as an essentially nickel-free solution suitable for direct discharge. Nickel can
be redissolved fiom activated silica by concentrated acid addition, and f i e r
centrifugation, the products are a concentrated nickel solution and regenerated activated
si 1 ica.
Mine watcr , Concentrate Lime
'Upper' pond Smelter slag
wastes
I Lime I
Sludge , Waste water treatment plant
Treated discharge
Figure 6.1 Il 221: Waste treatment operations of Inco Ltd., Sudbury, Ontario.
Sludgc
Residuc (Fe)
1 Solution (Ni. Mg)
Lime - 1-1 Ni rtcovtry oato
Regenerated activatcd
silica Thickening
t Emuent (Mg)
Activated silica under flow
Centrifuge
Conc.
1 ctntrilup -Concentrated Ni solution
Figure 6.2: Proposed process for siudge treatment.
6.3 Process Conditions and Results
6.3.1 Acid leach
The acid leach step was carried out by the Mineral Processing group. Department
of Mining and Metallurgical Engineering, McGill University. In each test. 80 g wet
sludge was slumed in 400 mL tap water. Sulphuric acid was added with mechanical
stimng, and at room temperature. until a pH between 3.0 and 3.2 was rnaintained.
Sulphuric acid requirement for essentially 80 % nickel recovery into an effectively iron-
fiee solution was 8 rnL per 80 g of sludge. The residual sludge was allowed to seule for
about 4 hours, and the overflow became the feed for activated silica treatment. A typical
sludge leach solution contained 3.1 g/L Ni. 5.2 g/L Mg and 7 mg/L Fe (ferric) at pH 3.2.
6.3.2 Nickel Recovery onto Activated Silica (Fig. 6.2)
The conditions for this step were fint optimized using synthetic solutions to
simulate a sludge leach product. These synthetic solutions contained 2.0 g/L to 3.0 g/L
Ni, 4.0 g/L to 10.0 g/L Mg, 0.0 g/L to 0.1 giL Fe (ferric) at pH 3.0 to 3.4. The
preparation of activated silica is descrîbed in detail in Chapters 3 and 4. For the tests
involving synthetic or actual leach solutions. the activated silica used was a I .O wlw %
Sioz sol, prepared from 'NT silicate (SiOzNa20 weight ratio = 3.22), stabilized afier half
the gel time of 18 minutes (Section 4.2). The scale of the tests was from 20 mL to 200
mL activated silica added to 100 mL to 1.00 L Ni/Mg/Fe solution. The addition always
satisfied the ratio 20 mL soi to 100 mL solution. which corresponds to a silicdnickel
molar ratio between 0.65 and 1.00.
M e r sol addition, the pH increased to about 6.0. Final pH adjustment was with a
suspension of 28 g/L Ca0 added in an amount that varied widely (0.3 rnL to 90 mL)
depending on the scale of the test and the desired pH. Reaction times studied were
between 20 minutes and 24 hours. The overflows (effluent) generated by thickening and
centrifugation accounted for about 90 % of the total volume.
Table 6.1 shows, for a synthetic feed, how nickel recovery (into silica) and
magnesium rejection (into effluent) changes with pH afier a lengthy (21 hour) retention
time. Nickel recovery increases with increasing pH. The optimum pH for NiMg
separation is about 8.5, at which nickel recovery is essentially complete, while 88 % of
magnesium has been rejected. A M e r pH increase beyond 8.5 results in decreased
selectivity for nickel recovery due to unwanted removal of magnesium 'hydroxide' into
the silica. It should be noted that. due to the dilution effect of silica and lime additions.
values for ' g / L Mg in effluent' in Table 6.1 cannot be linked directly to the 10 g/L Mg
feed for calculation of '% Mg rejected into effluent'.
1 Final pH 1 ppm Ni 1 g/L Mg 1 % Ni recovend 1 % M g rejected 1
Table 6.1 : Dependence of nickel recovery and nickeilmagnesium separation on pH after
24 h reaction tirne. 100 mL feeds: 3.0 g/l Ni, 10 g/L Mg, 0.0 g/L Fe. pH 3 .O.
1
7.17 7.59
in effluent 2,775 1.942
in elfluent 8.33 8.32
by silica 2.4
into effiuent 87.9
29.9 1 90.1
A second set of experiments, similar to those of Table 6.1, were cmied out as shown in
Table 6.2. Here the feed solution was different from that of Table 6.1, but still within the
range to be expected for a sludge leach solution. Optimum pH for nickel recovery and
nickeYmagnesiurn separation is again at pH 8.5. However, it is difficult to make
meaningful cornparisons between Tables 6.1 and 6.2 regarding magnesium rejection into
the effluent, especially at pH 8.5 and above. Aside from different magnesium feed levels
(10 g/L vs. 4 g/L), some magnesium adsorption by ferric hydroxide (Table 6.2 only) may
have reduced the dissolved magnesium content of the effluent below the value seen in the
absence of iron (Table 6.1). Another effect is the slow precipitation of magnesium
hydroxide during storage of the effluent (after separation from silica) prior to analysis.
Although acidification of the analytical sample could have prevented this tendency, the
main conclusions of Table 6.1 and 6.2 have not been affected (Le. the optimum pH values
for high and selective nickel recovery. and for production of < 1 ppm Ni effluent).
Table 6.2: Dependence of nickel recovery and nickellmagnesium separation on pH after
Final pH
7.50 8 .O0 8.25 8.50
24 h reaction time. 100 mL feeds: 2.0 g/l Ni. 4.0 g/L Mg, O. 1 O g/L Fe. pH 3.0.
ppm Ni in emuent
850.0 630.0 155.0 18.5
g/L Mg in eflîuent
3.18 3.12 3.36 2.59
9.00 1.66 1 .50
% Ni recovered by silica
55.1
99.9
O h Mg rejected
46.8
into emuent 84.0
66.3 91.6 99.0
'
83.5 91 .O 71.6
Further tests were carried out to compare 1 hour and 24 hour reaction times for
nickel recovery at several pH values (Table 6.3). At the optimum pH of 8.5. as deduced
from Tables 6.1 and 6.2. reaction for 1 hour is sufficient for effectively total nickel
recovery. However if the main objective is to produce a non-toxic, nickel-free (< lpprn
Ni) efliuent, the pH m u t be at least 9.5, in which case substantial arnounts of magnesium
will also be recovered into the silica.
1 Final 1 ppm Ni in 1 % Ni 1 ppm Ni in 1 % Ni
Table 6.3: Dependence of nickel recovery on pH after 1 h and 24 h reaction times. 100
mL feeds: 2.0 g/l Ni, 4.0 g/L Mg, 0.0 g/L, Fe. pH 3.0.
pH 7.65
The final experimental senes with synthetic solutions was perfomed with the
objective of obtaining detailed kinetic data for both nickel recovery and magnesium
rejection at the optimum pH of 8.5. Table 6.4 shows that, while 20 minute retention
allows good (92 %) rnagnesiurn rejection, some nickel loss is experienced (91 %
recovery) giving a 152 ppm Ni effluent. Afier reaction for an optimal 1 hour. nickel
recovery has improved to 98 % with magnesiurn rejection retained at 92 %. The major
r deterioration in degree
magnesium hydroxide
effluent (1 h) 1,548
effect of an increase in retention much beyond 1 hour is a steady
of nickeVmagnesium separation. A fine white precipitate of
recovered (1 b) 18.2
emuent (24 h) 1.355
recovered (24 h) 28.4
slowly foms giving a distinctive deposit above the greenish silica layer formed in the
first 20 minutes of reaction. The precipitation of magnesium hydroxide, shown in Table
6.4, occurs during reaction with silica as distinct fiom during storage of analytical
sarnples of effluent, as described above in connection with Tables 6.1 and 6.2.
Table 6.4: Dependence of nickel recovery and nickePmagnesium separation on reaction
timeat pH 8.5. 100 mL feeds: 2.0 dl Ni.4.0 g /L Mg.O.1 gL Fe. pH 3.0.
Time (h)
0.33 1 .O0
The data for synthetic solutions. as detailed in Tables 6.1 to 6.4. were considered
to be an adequate basis for tests with actual sludge leach solutions. The latter are
included in the discussion of nickel redissolution step where a nickel recoverylnickel
redissoluiion cycle (Fig. 6.2) is demonstrated for both synthetic and sludge leach
solutions.
ppm Ni in , elfluent
152 41.2
g/L M g in effluent
3.33 3.33
63.3 Thickening (Fig. 6.2)
After nickel recovery ont0 activated silica, the system must be separated into an
effluent and a siiica undertlow for further treatment. The free settling rates and
underflow water content obtained in p n m q thickening are critical parameters that
% Ni recovered by silica
91.6 97.7
% Mg rejected into emuent
92.0 92.0
determine the size of thickener (or 'clarifier') required, the capital cost of which would be
a major component of the overall equipment requirement.
Data were obtained for settling characteristics as outlined in Section 3.5. using a
'medium' size container (Fig. 3.3). Initial tests were camed out on 'metal-loaded' silica
prepared under optimum conditions as established in Tables 6.1 to 6.4. The scale of
operation was 1 L synthetic solution to which 200 mL activated silica had been added
(Le. sarne silica / (Ni/Mg/Fe) solution ratio as for Tables 6.1 to 6.4) with final pH
adjustment to 8.5 using lime (pH 7.0 for one test where the solution contained only iron).
Figure 6.3 gives the results in the fom of distance travelled downwards by the interface
vs. time.
Feed to recovcry step (g/L)
E Ni Mg Fe pH
O 1 + - - 0.06 7.0
7 .--- -
O 2 4 6 8 10 12 14 16 18 20 22 24
Time (minutes)
Figure 6.3: Settling of metal 'loaded' activated silica. Effect of feed composition. 200
mL 1 .O w/w % Si02 sol additions to 1 L feed. Final pH adjustrnent with lime.
The best fiee settlhg rate was with the nickeVmagnesiurn system without iron.
Here, fiee settling was at 0.6 mh. The presence of iron decreases the settling rate to
about 0.5 mh, which is applicable to both the NitMgFe and Fe only systems in Figure
6.3. The overflow (effluent) was exceptionally clear in al1 cases. Settling rates in the
range 0.5 m/h to 0.6 m/h are comparable to those found for the 'Balsam Street' water
treatrnent plant of Inco Ltd., Sudbury, Ont.. This large facility handles water run-off
from a variety of tailings ponds and uses lime for the entire neutralizatiodmetal removal
process. However, at the 'Balsam Street' plant, it is necessary to add a flocculant
('Percol') before settling rates of 0.5 m/h to 0.6 mih can be achieved.
It was also noted that the underflow generated after free settling in Ni/Mg/Fe
systems was mostly separated into two distinct sections. A bottom phase was pale green
containing nickel adsorbed on sol, nickel hydroxide and gypsum. A bulkier top layer was
yellow-green containing mostly femc adsorbed on sol or in hydroxide fom, with nickel
species in small arnounts. Thus it appears that the slightly decreased settling rate
associated with the presence of iron is due to ferric hydroxide which settles independently
and more slowly than the nickeVsilica component. At pH 8.5, it would be expected that
femc ion would tend to desorb h m activated silica and form a fiee hydroxide phase. In
contrast, an underflow generated as a result of adding only lime (no silica addition) does
not form a layered structure, but eventuaily appears as a roughly uniform mixture of
solids with a colour representing an average of the components. When adding only lime,
settling is extremely slow and 'duty' ovemows are generated.
The w/w % solids (or water content) of the underflow is important, because this
parameter fixes the feed to the centrifbge (Fig. 6.2), the product of which should have as
low a water content as possible to promote ultimate production of a concentrated nickel
solution. Figure 6.3 shows that settling essentially stops after about 30 minutes. At this
point the undertlow wlw % solids was found to be between 1 % and 3% , with the lower
end of this range corresponding to iron-containing systems. This is slightly better than
the 'Balsam Street' plant, which generates underfîows in the 0.5 to 1 w/w % solids range
(afler one cycle).
Further senling tests were carried out (Fig. 6.4) to study the effect of reducing the
silica addition below the 'standard' 20 mL 1 .O wiw % Si02 per 100 rnL metal-containing
solution applicable to Figure 6.3. The optimum pH of 8.5 was maintained which implies
an increased lime requirement to compensate for the reduced silica addition. Settling
rates are again in the 0.5 m/h to 0.6 m/h range (based on the 6 minute data points of Fig.
6.4). However, undefflow wiw % solids is slightly improved when silica addition is
decreased, as cm be seen qualitatively fiom the positions of the plateaux in Figure 6.4.
This latter e ffect reflects the qualitative observation that uncentri hged undefflow
becomes more voluminous and traps more water as silicdlime ratio increases.
O 2 Feed to recovery step (fi)
Tirne (minutes)
Figure 6.4: Settling of metal 'loaded' activated silica. Effect of silicaffeed ratio. 1 .O
W/W % Si02 sol additions to 1 L feed. Final pH adjustrnent with lime.
In addition to the settling data of Figures 6.3 and 6.4, measurements were made of
the % of the feed volumes rejected as effluent following the thickening step. This is
important because the throughput to be treated in the last stages of the process (Fig. 6.2)
is much smaller than the requirements for the initial stages. Table 6.5 gives results for %
of feed volume discarded as efnuent for a variety of situations following free settling for
30 minutes and before underflow centrifugation.
(Fig 6.3) (Fig. 6.3) (Fig 6.4) (Fig. 6.4)
Table 6.5: Percent feed volume rejected to the emuent after 30 minutes of free settling
and before underflow centrifugation for tests of Figures 6.3 and 6.4. Feeds contained 2.0
giLNi .4 .0gLMg ,0.0or0.1 @Fe.
% feed volume rejected
It can again be seen that an increased silica addition gives a more voluminous underflow
and less water rejected. Afier the 'standard' addition of 200 mL sol/ 1 L feed, % of water
rejected after settling approached 80 %, while almost 90 % was discarded if sol addition
was reduced to 50 mLI 1 L feed. In every case, the overall proportion of input volume
discarded rises to 92 % to 95 % following centrifugation.
In summary, the settling characteristics of 'Ni/Mg/Fe-loaded' activated silica are
comparable to or slightly better than experienced with conventional lime treatment using
a flocculant. The silica undefflow physically resembles a sewage sludge, for which
centrifuges are cornmonly used on a large scale.
6.3.4 Nickel Redissolution (Fig 6.2)
Feed for 'Ni redissolution can be generated only by passing through the
8 1 76 89 88
preceding steps of Figure 6.2. It is thus convenient io discuss 'Ni redissolution' as part
of a cornparison of synthetic and sludge leach solution behaviour in an overall process
cycle, as summarized in Table 6.6.
pH Retention Time (h) Vol. % Discarded as Effluent * % Ni Recovered ** into silica % Mg Reiected ** into effluent
g/L Ni in "Concentrated" Solution 1 17.6 1 11.8
Synthetic solution Recovery step
8.5 1 1.5 92
* After settling + u'flow centrihiging. ** Based on feed Ni (Mg) as 1ûO %.
Sludge leach solution Recovery step
8.67 1.5 84
98.9 75.1
Sludge leach solution Redissolution stop
1.68 1 .O
58.4 41.1
Synthetic solution Redissolution stop
Table 6.6: Cornparison of a nickel recovery/redissolution cycle for synthetic and sludge leach solutions. Synthetic solution feed: 2.0 g/L Ni, 4.0 g/L Mg, 100 mg/L Fe, pH 3.0. Sludge leach solution feed: 3.1 g/L Ni, 5.2 gR. Mg, 7.1 mg/L Fe, pH 3.0.
99.5 87.2
pH Retention Tirne (h) % of Ni in "Concentrated" Solution ** % of Ni in Recvcled Silica * *
'Ni recovery' results are discussed first since these influence the outcome of 'Ni
redissolution'. Only minor adjustments are needed for the recovery stage. which was
carried out following the same procedure as for Tables 6.1 to 6.4. At pH 8.5 1 (syntheiic
solution), 99 % of the nickel was recovered, with 75 O/O magnesium rejection into the
1.37 1 .O
61.4 37.5
effluent. A small pH increase to 8.67 (sludge leach solution) also gave 99 % nickel
recovery, with an improved magnesium rejection (to 87 %). It is unlikely that the
behaviour of the sludge leach solution would be exactly the sarne as for the synthetic
solution, because the former contains unknown levels of additional dissolved m e a s other
than nickel, magnesium and iron. The difference in 'Vol. % discarded as effluent' (92 %
vs. 84 %) may reflect a small variation in underflow w/w % solids fed to the centrifuge,
again bearing in mind the probable complexity of the sludge leach solution.
'Ni redissolution' was achieved using concentrated (93 %) sulphuric acid
addition, camied out as described in Section 3.4. The most important conclusion fiom
Table 6.6 is the unfavourable distribution of nickel between 'concentrated solution' and
'recycled silica' (i.e. for synthetic solution, 98.9 % nickel recovery has split into 61.4 %
in product solution and 37.5 % in recycled silica). Although nickel in recycled silica is
not a loss, a recycle of 37.5 % of the feed nickel is unacceptably high. Appropriate
methods for reducing the amount of recycled nickel would depend on the fom of the
nickel in the recycled silica (i.e. dissolved nickel, adsorbed nickel or precipitated nickel
hydroxide). If the recycled nickel is dissolved, improved centrifugation should increase
nickel distribution in favour of the concentrated solution product, but the g/L Ni in
'concentrated solution' would not change. If the recycled silica contains phar i ly
adsorbed nickel or precipitated nickel hydroxide, an increased acidity and/or longer
retention time coupled with irnproved centrifugation would both increase g/L Ni in the
concentrated solution and reduce the circulating load of nickel in recycled silica.
The redissolution step has not k e n explored as thoroughly as the recovery stage,
treatment of sludp leach solutions with activated silica being an ongoing project.
Nevertheless the results seem ~ ~ c i e n t l y encouraging to justiQ further work in this area.
6.4 Ef'fluent Treatment Using Activafed Silica
Chapter 6 is primady concemed with treatment of tailings pond sludge
generated through lime addition to various effluents (Fig. 6.1). By combining results
from Chapters 5 and 6. it cm be seen that in principle sludge formation can be greatly
reduced by directly treating an effluent with activated silica. As with silica gels. metals
separation into groups may be feasible through pH-controlled selective rernoval (Section
2.12). For exarnple. essentially complete ~ e - (~l*- . ~ r - ) removal occurs at pH 1.0.
~ i * (Cu-, Co-. ~ e * . ~ b - . cd") at pH 8.5 and ~ g * ( ~ a * ) at pH above 10. The fint
two groups of metals could be recovered then redissolved to generate more concentrated
solutions than the feed. leaving a ~ g - ( C a ) effluent. Figure 6.5 is a simple schematic
comparing current practice with a conceptual flowsheet using activated silica. Effluents
are presently treated with lime and flocculants (Le. 'Percol') to fonn a waste sludge at pH
9- 1 1 and a metal-free discharge.
The proposed process, using activated silica. involves initial femc removal at pH
4.0 using sol and lime, the sol phase then being treated for iron redissolution and
concentration. The effluent fiom this first step would contain nickel plus other divaient
'heavy' metals along with magnesium and cakium. and would pass to a second stage
where silica and lime are added to pH 8.5. The final products would be nickel-bearing
silica available for redissolution/concentration steps and a magnesium (calcium) effluent.
This activated silica process offers potential advantages such as production of separate
concentraed solutions and a low volume waste sludge.
Lime
Existing process
4 Waste sludge/recycied 3 timcs to incrtasc % solids
solution) ,T l! iT ,[ 1 ,
Proposed - process pH - 4.0 NiIMg rich pH-8J Mgrich solution emucn t
4 + Fe rich sols for production Ni rich sols for production
of conccntratcd solution of conctntrrtcd solution
Figure 6.5: Existing and proposed processes for effluent treatment.
CHAPTER
Summary
7.1 Conclusions
7.1.1 Introduction
The goal of the present work was to develop a conceptual process for 'heavy
metals' removal from dilute aqueous streams, such as mine effluents, using an alkaline
solution of ' stabilized' activated silica sol. Subsequent metal redissolution by
acidification of the sol would allow metal recovery and production of a lower volume
sludge than obtained by conventional practice using lime. Dissolved metal is removed
largely by adsorption giving a metal-loaded sol that settles to produce a sharp interface
dividing a metal-sol undefflow (bottom phase) and clear overflow (top phase). Final pH
adjustment using lime is beneficial becaw the metal content of the ovefflow is lower
than with sol alone, and settling rates are increased. Thus a proportion of metal
precipitates as hydroxide, depending on the amount of lime added.
Additional water rejection by underflow centrifugation facilitates production of a
concentrated solution (relative to the feed) which is created by acidification with
concentrated acid. Simultaneously, 'unloaded' activated silica is regenerated, and after
phase separation and neutralization with lime, the sol can be reused. Metal adsorbed on
silica sol is more readily redissolved than is hydroxide precipitate intimately mixed with
lime and gypsurn, as would be produced by neutralizing with lime only. The present
work comprised some initial fundamental studies, followed by investigation of the
practical application of silica sols in the treatment of iron effluents and nickel-containing
tailings pond sludge.
7.1.2 Fundamental Studies
Copper was the metal chosen for investigating key features of
adsorption/desorption with silica sols because of the availability of the cupric ion
selective electrode. This device permits instantaneous 'in-situ' measurement of
unadsorbed cupric ions, which can be distinguished from total dissolved copper content
as measured by atomic absorption spectrometry. Adsorption/desorption of copper is a
rapid, pH-dependent and reversible process. Copper adsorption increases with increasing
pH up to pH 7. Above pH 7, atornic adsorption spectrometry indicates some desorption
has occurred, while the electrode registers zero dissolved cupric ion. This desorbed
copper is attributed to colloidal copper hydroxide which would appear in the effluent in a
practical application. Final pH adjusment with lime removes colloidal hydroxide
species.
The distinction between adsorption and precipitation was demonstrated using
cobaltous sulphate solutions. When removed fiom solution using silica sol, the pink
colour of the dissolved cobaltous ion is transferred to the silica sol undefflow. indicating
an adsorption process, while the overflow becomes colourless. If dissolved cobaltous
ion is removed using sodium hydroxide, a blue cobaltous hydroxide precipitate forms.
Use of lime generates a greenish mixture of solid products.
The production of stabilized silica sol requires the initiation of pol ynerization
(gelation) by acidifying a silicate solution then dilution to anest pal-mer growth.
Polymerization is characterized by an increase in sol viscosity. For a given pH. copper
adsorption is independent of polymer site when silica addition is determined by the pH
required, which results in high SiOdCu molar ratios. This independence of copper
adsorption on polymer size is attributed to provision of excess adsorption sites when pH
is the controlling parameter for silica addition. Other parameters measured as a hinction
of polymer size were surface tension and zeta potential. 50th remain essentially constant
at 75 dynekm and -25 mV respectively. After sol addition to a cupric solution. the zeta
potential increases from -25 mV to zero as pH tises to 7, and remains at zero up to pH 8.8
due to only partial copper desorption. This is different to that reported for silica gel
particles. where zeta potential revens back to its original value afier total copper
desorption at pH 8.8.
At a fixed silica level. 'D' silicate (SiOz/NazO = 2.00) is a more powerful reagent
for raising pH than 'Na silicate (SiOzMa?O = 3.22) because of the higher Na20 content in
the former case. Use of 'D' silicate instead of 'N' silicate under otherwise the same
conditions gives the same arnount of copper adsorbed for a given pH. Although
maximum sorption capacity was not determined. sorption values (mole Cu/moie SiO?) up
to 10 times that reported for silica gel powden were measured in the present work. High
sorption values for liquid silica sol are attributed to better access to adsorption sites than
available using a solid powder.
The feasibility of concentrating a dilute feed using adsorption/desorption cycles
was demonstrated. Using single cycles for 63.5 pprn Cu and 315 pprn Cu feeds, 355
pprn Cu and 1,040 pprn Cu solutions respectively were produced suggesting that, with
two consecutive cycles, a 1 .O4 g/L Cu product could be generated fiom a 63.5 pprn Cu
feed. Optimum adsorption and desorption pH values are about 7.0 and 4.3 respectively.
The possibility of separating two dissolved metals with activated silica was shown
using iron (ferric) and copper as an exarnple. Selective femc removal fiom a
femckupric mixture was optimal at pH 5.2, and using one adsorption step, 95 % of femc
was rernoved with 90 % of copper remaining dissolved. The final pH was adjusted to
5.2 using lime. Addition of lime precipitates desorbed colloidal femc hydroxide that
forms as pH 5.2 is approached.
7.1.3 Potential Applications
Iron Removal from Acid Mine Drainage
At the request of the project sponsors, National Silicates Ltd., the first potential
application studied was iron removal, with concentration of dissolved iron, from acid
mine drainage. An impure concentrated ferric sulphate solution was considered to be
marketable as a flocculant for the sewage industry.
Experiments with femc feeds containing fiom 55 pprn Fe to 1,117 pprn Fe
showed that use of activated silica only for pH control gave excessive additions that
produced undesirable side effects related to sol dispersion. These effects included partial
iron desorption that gave entrapment of colloidal iron species in the 'adsorption overflow'
into which high silica losses were also observed (up to 26 % of added silica). In
addition, a second 'loaded' sol phase did not form at excessively high silica levels.
These problems were resolved by using lime for final pH control which resulted in some
femc hydroxide precipitation dong with femc adsorption. Use of sodium hydroxide
instead of lime was not completely satisfactory since iron-loaded sol settled slowly giving
a poorly defined overflowlundefflow interface.
Settling tests were performed for sol loaded from 558 ppm Fe feeds at pH > 4.0.
After metal removal under optimum conditions (1 wiw % Si02 sol, pH 4.0), the free
settling rate (initial 20 minutes) was 0.32 m/h which is similar to the value for
conventional lime neutralization of 0.5 g/L Fe solution using 'Percol' as a flocculant.
Free settling rates decrease and volume % occupied by settled underflow increases with
increasing silica addition. There is thus a compromise between adding enough silica for
iron removal to form a metal-bearing underflow while retaining reasonable settling
characteristics and avoiding sol dispersion. Settling rates using recycled sol were
essentially the same as with fiesh sol.
Underflow wfw % solids after iron rernoval under optimum conditions, and free
settling at 0.3 m/h, was 2.6 % solids afler centrihigation and oven drying. However the
measured wlw % solids includes the siliceous residue created by oven drying and does
not represent the true w/w % solids generated in the settling test itself. Apparent
undertlow wlw % solids increased fiom 2.60 % to 3.75 % after two recycles of loaded
sol. Settling characteristics for a real iron effluent when treated with activated silica/lime
(60 ppm Fe, Balsam stream, lnco Ltd., Sudbury, Ont.) were similar to those for synthetic
solutions.
It is possible to upgrade a 500 ppm Fe feed to about a 30 gR. Fe product using a
three-cycle system where progressively more concentrated solutions are treated. In the
ihird cycle, gypsurn build-up is a major concem because of the need to neutralize acid
used in redissolution of a large quantity of iron. High silica losses (up to 46 % of input)
are also observed under the strongly acidic conditions of third cycle redissolution.
O Changes in temperature between 5 C and 45 OC had only minor eflects on
removal/redissolution.
Selective Nickel Recoverv from Tailines Pond Sludee
Sulphuric acid leaching of a tailings pond sludge containing (dry basis) 6.2 % Ni,
16.8 % Fe, 20.3 % Mg produced solutions at about 3.0 g/L Ni, 5.0 gL Mg, 0.01 g/L Fe,
the final pH k ing controlled ai 3-4 to minirnize iron dissolution. Both sludge leach
solution and synthetic solution analogues were treated for selective nickel recovery over
magnesiurn. Addition of 1 w/w % Si02 sol, with lime for pH adjustment to 8.5, gave
optimal results of 98 % to 99 % selective nickel recovery ont0 silica, with 75 % to about
87 % magnesiurn rejection into an overflow. Good fiee settling rates of 0.5 m/h to 0.6
m/h were obtained (initial 10 minutes) giving about 75 % volume rejection to the
effluent. Percent volume discarded in overflow increased to 90 % upon underflow
centrifugation.
Preliminary testing of one recoverylredissolution cycle for synthetic and sludge
leach solutions gave 97 % and 87 % magnesiurn rejection respectively to the effluent, and
concentrated nickel solutions contained 17.6 g/L Ni (synthetic solution) and 11.8 g/L Ni
(sludge leach solution). The resdts obtained for both iron removal from synthetic
effluent and selective nickel recovery From nickellmagnesium sludge have been used as a
guide for proposing a conceptual effluent treatment process. The proposed procedure
involves two stages of selective metal recovery and concentration (iron at about pH 4,
nickel at pH 8.5) with sol recycle, instead of conventionai one step lime treatment
producing a mixed metal hydroxide/gypsum waste.
7.2 Claims to Origiaality
The work descnbed in this thesis aimed to develop the concept of using a
recyclable silica sol ('activated silica') for metals recovery from dilute streams such as
effluents, to generate a concentrated product. Most of the results obtained in this thesis
were the subject of a patent application by El-Ammouri et al. [123]. The following
elements of the conceptual process are considered to be original contributions to
knowledge:
Use of a stabilized liquid silica sol to adsorb metals from a sulphate solution
giving a metal loaded sol product that seules by gravity thickening to form a
discrete layer.
Use of lime for pH adjustment following silica sol addition to precipitate desorbed
colloidal hydroxides and to improve settling characteristics.
Production of a low volume concentrated dissolved metal product by desorption
using concentrated acid added to centrifuged undertlow produced in claim 1.
Regeneration of recyclable silica sol concurrently with metal desorption as in
daim 3.
Use of a cupric ion selective electiode to distinguish between fiee dissolved
cupric ion in a sol-copper mixture and other forms of copper (adsorbed copper,
desorbed colloidal or precipitated copper species). Demonstration of the
reversibility of the adsorptioddesorption cycle in-situ by following emf changes
as a Function of pH.
Use of sol and lime to separate metals in a mixed metals solution, using an iron
(ferric)/copper (cupric) solution as an exarnple.
Selective nickel recovery over magnesium from a tailings pond sludge by acid
leaching and sol/lime treatment to concentrate nickel into a product solution.
Use of silica sol (a non-toxic reagent) both as a replacement for Percol (organic
flocculant) and as an adsorbent when used with lime in effluent treatrnent for iron
removal, with the possibility of concentrating and recovering iron.
7.3 Suggestions for Furtber Investigations
The practical application of claims 1-8, listed in Section 7.2, would require further
work as suggested below:
1. If removaVredissolution cycles are operated continuously in a closed loop,
gypsum build-up would be a major difficulty. This problem becomes especially
acute when acid requirement for redissolution is high (e.g. production of about 30
g/L Fe solution - Section 5.4). Suggestions for handling gypsum build-up have
been proposed in Section 5.6. The easiest way of removing excess acid before
sol recycle to 'removal step' is probably a silica wash with neutralization of the
acidic wash water outside the main process loop. This should be checked
experimentally.
2. Even though silica losses do not create additional environmental problems,
optimization of silica addition for a given feed concentration and pH is necessat-y
fiom an economic standpoint to minimize losses to outgoing streams (effluent,
concentrated metal solution).
3. Recycled sol would contain locked-in metal due to aqueous phase entrapment,
which represents a recirculating load rather than a loss. Although some tests
were carried out using recycled sol. the circulating metal load was not moni tored.
while in other experiments, single cycles were used using fiesh sol. Tests should
be done in which sol fiom the first cycle should be used for the second cycle etc.
with metal entrapment in the sol being measured for each cycle.
4. Experiments at fixed metal/silica ratios with pH controlled independently of silica
addition such that adsorption capacity and reaction stoichiometry can be studied.
5. Further development of the process tested for nickel recovery from tailings pond
sludge should be pursued in two areas:
a) Tests for nickel redissolution at higher pH levels than actually used (pH 1.5) to
lower the lime requirement and gypsum production upon neutralization of
recycled sol.
b) Expenments in which the first cycle 15 g/L Ni product actually obtained is fed to
a second cycle with the objective of generating a sufficiently concentrated solution
for electrowinning (about 60 g/L Ni).
References
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