The Extraction of Zinc from Secondary Zinc Minerals

THE EXTRACTION OF ZINC

FROM SECONDARY ZINC MINERALS WITH

AQUEOUS SODIUM CYANIDE SOLUTIONS

Steve Bogdan Kesler, B. Sc(Eng), Ä. R. S. M.

April, 1976

A thesis submitted for the degree of Doctor of Philosophy of the University of London and for the Diploma of Imperial College.

Department of Mining and Mineral Technology, Imperial College, London S. W. 7.

Dvý`- -1-

ACKNOWLEDGEMENTS

I should like to thank Professor E. Cohen, all the staff and research students of the Department for their helpful discussions and the technical staff for their patience in building and repairing apparatus. In particular, I thank my colleagues Mr. R. F. Dougill for all his valuable time spent on my computing problems and Mr. J. R. J. Burley for his advice on mineralogy. The willingness of Mr. A. Read of Warren Spring and the Institute of Geological Sciences to spare me some of their time was also much appreciated.

Above all I thank Dr. H. L. Shergold, my supervisor, for his constant advice, tact and understanding and without whom this work could not have been accomplished.

I should also like to thank the Science Research Council for providing the financial assistance to enable me to carry out this research and last, but not least, Miss Christine Ball, for volunteering to type this thesis even after she had'seen my appalling handwriting. -2-

ABSTRACT

The possibility of recovering zinc from zinc oxide material with a sodium cyanide leaching stage has been investigated. Zinc oxide and the secondary zinc minerals all dissolve stoichiometrically in cyanide solutions. The predominant species in cyanide solutions saturated with

smithsonite or hemimorphite is the Zn(CN)42 complex, and a cyanide to zinc molar ratio close to 4/1 is obtained. A

lower ratio results from cyanide solutions saturated with

zinc oxide or hydrozincite because of the formation of zinc

hydroxy complexes.

The dissolution of smithsonite and hemimorphite is

mass transfer controlled, the former by transfer either of

cyanide ion to or, more probably, Zn(CN)42 away from the

reaction interface whilst the latter is controlled by the

diffusion of zinc cyanide complexes through a thin silica

surface layer. The dissolution reactions follow heterogeneous

reaction theory and empirical reaction models have been

derived.

Dissolution of smithsonite and hemimorphite is very

anisotropic the reactions being initiated at high energy surface

sites. Preferential dissolution at sub-grain boundaries

results in the disintegration of the smithsonite particles. -3-

The addition of sodium hydroxide to the cyanide solution increases the rate of dissolution of hemimorphite by thinning or removing the surface silica film.

Roasting the hemimorphite removed compositional

water from channels in the crystal structure and produced

an increased reaction surface area.

Zinc and free cyanide can be recovered from cyanide

solutions containing dissolved zinc oxide by electrolysis but

part of the cyanide is oxidised to cyanate at the anode.

The recovery of zinc from cyanide solutions containing

dissolved smithsonite, hemimorphite or hydrozincite is

possible by precipitation as the cyanide. The precipitate

can be dissolved in spent electrolyte from zinc sulphate

electrolysis allowing cyanide recovery and electrodeposition

of zinc from the electrolyte.

Reagent losses as cyanide or sodium hydroxide are

substantial but a cyanide leaching process might be

economically viable if cyanide oxidation during electrolysis

can be minimised and if cheap supplies of sodium hydroxide

are available. -4-

CONTENTS Page Acknowledgements. 1

Abstract. 2 Contents. 4 Chapter

INTRODUCTION 6

1.1. Occurrence of zinc oxide minerals 6 1.2. Conventional zinc mineral processing. 8 1.3. Alkaline leaching of zinc oxide material. 11 1.4. Reactions between cyanide and the zinc minerals. 18 2 EXPERIMENTAL 32 2.1. Materials. 32 2.2. Analytical techniques. 44 2.2.1. Zinc analysis. 44 2.2.2. Cyanide analysis. 45 2.2.3. Cyanide ion-selective electrode. 47 2.2.4. Potentiometric titrations. 55 2.2.5. Ultra-violet absorption spectrophotometry. 70 2.2.6. Conclusion. 75 2.2.7. Determination of the free cyanide concentration of leach liquors. 76 3 THE SOLUBILITY OF THE SECONDARY ZINC MINERALS IN AQUEOUS CYANIDE SOLUTIONS 82

3.1. Experimental procedure. 82 3.2. Results. 86 3.2.1. Influence of cyanide concentration. 86 3.2.2. Influence of temperature. 92 3.2.3. Influence of sodium hydroxide addition. 93 3.3. Discussion. 98 4 KINETIC STUDIES 112 4.1. Introduction. 112 4.2. Agitation system. 118 4.3. Experimental technique and data analysis. 120 4.4. The influence of agitation rate on the rate of dissolution of the secondary zinc minerals. 122 4.5. The influence of particle size on the rate of dissolution of smithsonite and hemimorphite in cyanide solutions. 132 -5-

Chapter Page

4.6. The influence of cyanide concentration on the rate of dissolution of the secondary zinc minerals. 139 4.7. The influence of temperature on the rate of dissolution of smithsonite and hemimorphite. 148 4.8. General rate equations for the dissolution of smithsonite and hemimorphite in cyanide. 154 4.9. Scanning electron microscope examination of smithsonite and hemimorphite after leaching in sodium cyanide. 156 4.10. Surface area changes during smithsonite dissolution. 179 4.11. Further dissolution studies on hemimorphite. 182 4.12. Addition of sodium hydroxide to the cyanide leach solvent. 185 4.13. Roasting of hemimorphite. 194 5 DISCUSSION OF KINETIC RESULTS. 204

5.1. Comparison of rate equations with the experimental results. 224 6 METAL AND SOLVENT RECOVERY 228 6.1. Introduction 228 6.2. Solution purification. 235 6.2.1. Experimental. 240 6.2.2. Results. 241 6.2.3. Conclusions. 242 6.3. Electrodeposition of zinc from cyanide solutions. 244 6.3.1. Experimental. 251 6.3.2. Results. 254 6.3.3. Discussion. 275 6.4. Zinc cyanide precipitation. 283 6.4.1. Experimental. 283 6.4.2. Results. 285 6.4.3. Discussion. 292 6.4.4. Conclusions. 295 7 PROCESS EVALUATION. 297 7.1. Direct electrolysis of zinc cyanide solution. 297 7.2. Precipitation of zinc as zinc cyanide. 302 8 FINAL CONCLUSIONS. 307

REFERENCES 312 -6-

1. INTRODUCTION

1.1. Occurrence of zinc oxide minerals

The major source of zinc is the sulphide minerals

sphalerite and wurtzite, however, significant amounts

have and are being recovered from zinc carbonate and

zinc silicate minerals. These latter minerals, known

as calamine ores, originated through the weathering

and oxidation of the primary sulphide deposits resulting

in the formation of a zinc sulphate solution which precipitated

on encountering carbonate rocks or available silica.

Many primary zinc deposits are, therefore, capped by

oxidised zinc ores containing smithsonite (zinc carbonate),

hydrozincite (basic zinc carbonate) and hemimorphite

(zinc silicate). In some cases, oxidised. zinc deposits

occur alone either as a result of the complete oxidation

of the sulphide minerals or the transport of the zinc

sulphate solution some distance from the original

primary ore body before precipitating.

The zinc minerals of commercial interest are

shown in Table 1.1. -? -

Table 1.1. Zinc minerals of commercial interest

I--- - T Composition Mineral Formula rTn(%)_ I Si(%)

Sphalerite ZnS 67.1 ý - Wurtzite ZnS 67.1 - Smithsonite ZnCO3 52.1 - H d i i te I 2Z 3 Z 59.6 y roz nc nC(ý . n(OH) 2 - H em i morp hi te Zn Si Cý (OH) H 54.3 11.7 62722 .O 180.3 ý Zincite* Zn - Willemite Zn SiO 58.7 12.6 (Fe, Franklinite* Mn Zn)(Fe, Mn)204 15-20 -

m Zincite and franklinite only occur in quantity at Franklin, New Jersey(, ).

The zinc sulphide minerals are generally

associated with lead minerals such as galena and

cerussite, other sulphides like pyrite and chalcopyrite,

silver minerals and less commonly gold, the latter

two greatly adding to the value of the ore. The most

common gangue minerals associated with zinc deposits

are calcite and dolomite with fluorite, quartz and

limonite. A number of metals are commonly found

substituted for zinc in the oxide zinc minerals and

these include copper, cobalt, cadmium, manganese,

lead, magnesium and iron, although usually only to a

small extent (less than 1%). The most important of

these substituted metals is cadmium which rarely -8-

occurs as a natural cadmium mineral and, indeed,

most of the world's production of this metal comes as a

by-product from the processing of zinc ores.

1,2, Conventional zinc mineral processin

The zinc sulphide minerals are conveniently

concentrated by normal sulphide flotation methods and

the zinc oxide minerals are often recovered by flotation

with a dodecylamine/fuel oil emulsion(2) a fatty acid(3)

Flotation the or after sulphidisation(4) . of zinc oxide

minerals, however, generally results in a poor

concentrate grade and recovery because the reagents

are not selective enough. More recently work has

been conducted whereby the zinc oxide minerals are

floated by the use ofchelating agents(5) which enable

more selective flotation. However, flotation is made

difficult by the similarity in the surface properties

of the zinc oxide minerals and the gangue and by the

presence of hydrated iron oxide slimes which coat

both gangue and zinc minerals alike.

Zinc has a low boiling point (9030 C) and is

often recovered from oxidised zinc deposits of

sufficiently high grade by fuming in retorts after

roasting or calcining. This process, however, has

many disadvantages not least being the requirement -9-

of a cheap source of fuel. The reduction can be summarised by the following two gas-solid reactions.

Zn0(s) + CO(g) -s Zn(g) + C02(g) (1, la)

C02(g) C(s) 2CO(g) (1. lb) + r-

Zinc oxide is relatively stable, however, and the reduction reaction is only thermospontaneous at 950°C.

On condensing the zinc, therefore, the reaction given by Equation 1.1 tends to reverse and large amounts of the zinc revert to zinc oxide. Other volatile metals such as Pb, Sn, Sb, Cd, Hg, As, Mo. and Ti are also recovered in the fume. The introduction of the Imperial

Smelting zinc blast furnace has partially solved the problem of zinc oxide formation by the shock chilling of the blast furnace gases. with a spray of unsaturated zinc in liquid lead at 550 to 6000C thus condensing and collecting the zinc before it has time to react with carbon dioxide to form zinc oxide. Even so appreciable zinc oxide (up to 11%) is still formed.

Although there is a current slump in zinc sales the declining mine output in existing mines and the shortage of new mines means that, in the future, an increasing amount of zinc must be supplied by recycling

Numerous secondary zinc (6)' sources of secondary zinc are available particularly from lead and copper -10-

smelters which often have slag fuming adjuncts to permit recovery of zinc as zinc oxide from the slag(7).

Complex zinc sulphide/oxide ores are often upgraded by the use of `Vaelz kilns and the zinc is again recovered as zinc oxide in a fume containing typically 55.6% zinc

13,6% lead(, Impure is, therefore, and ). zinc oxide

formed as an inevitable by-product in metal fuming

stages, blast furnaces and the Waelz kiln.

With the rapid increase in transportation costs

there is a tendency to produce concentrates locally and,

unlike blast furnaces which need large capacity to be

economic, hydrometallurgical operations can

economically treat small tonnages. Leaching, therefore,

offers a possible method of treating oxide zinc ores

and secondary zinc such as fume that cannot be bene-

ficiated by conventional mineral dressing processes.

The oxidised zinc ores can be leached with sulphuric

acid but as the gangue is usually calcareous or dolamitic

in character, large quantities of acid would be consumed

by the dissolution of valueless material. A sulphuric

acid leach of siliceous zinc ores or of oxide zinc ores

in the presence of a siliceous gangue results in the

formation of gelatinous silica which is difficult to

separate from the leach liquor(9) and consequently -11-

metal losses in the gel are high. Zinc oxide minerals,

zinc calcine and fume all contain minor amounts of

other metals which are also soluble in sulphuric acid

e. g. As, Sb, Cu, Fe, Cd, Co, Ni, Sn, Cre, Se, Te,

Al and if the zinc is to be recovered by conventional

electrolysis the solution purification procedure becomes

extremely important. Often the leaching methods are

influenced more by the necessity of minimising the

amount of impurities in solution than by any other

factor(7) Many deposits . oxidised contain a substantial

amount of iron oxide material and sulphuric acid

leaching in large iron in results quantities of solution (l0)'

1.3. Alkaline leaching of zinc oxide material

Under alkaline conditions a basic gangue and

the iron--oxide minerals would be essentially insoluble

and, therefore, by the choice of a suitable solvent it

should be possible to selectively dissolve the zinc oxide

minerals and fume. Many alkaline reagents are

available but the most suitable ones will be those which

form stable soluble complexes with zinc.

A transition metal can form coordinate bonds

where the metal ion accepts a share in a pair of electrons

donated by a non-metallic Lewis base. Metals such as

copper, nickel and zinc can coordinate with oxygen, -12-

carbon or nitrogen donors and complexes can, therefore, be expected with hydroxyl, amines, carbonyls and cyanides. The donor properties of nitrogen are stronger than oxygen(11) and strong complexes are

generally formed with amines and aliphatic amines although in the latter case the coordination ability of

the ligand decreases in the order primary, secondary

and tertiary amine, probably as a result of steric

factors(12). Nitrogen in organic nitriles also has

fairly strong donor properties. Cyanide has unshared

electrons on both carbon and nitrogen but coordination,

in simple mononuclear complexes, seems to be only

through carbon and isomeric series of complexes with

cyanides, corresponding to nitriles and isonitriles,

are not observed.

In complexes of transition metals with ligands

possessing free acceptor orbitals, back coordination

can be expected in addition to the usual 0- bond and,

hence, the stability of the complex will be enhanced

and the selectivity of the reaction increased. For

individual metals the order of increasing stability for

common ligands is given by the spectrochemical series(13)

I< Br < Cl

the high field strength of cyanide being a consequence -13-

of its high acceptor capabilities. Cyanide coordination

with transition metals is powerful and frequently

displaces all other groups in the coordination sphere.

Possible solvents for a selective alkaline

leaching process, therefore, include aqueous solutions

of sodium hydroxide, ammonia-ammonium mixtures,

aliphatic amines, cyanides and nitriles. The stability

constants for the various zinc-ligand complexes are

compared in Table 1.2.

Table 1.2. Stability of zinc-li and complexes

Log cumulative stability Ligand constant Ref. _ P1 13 R3 134 2

NH3 2.66 4.81 7.48 9.55 14 OH 6.31 11.19 14.31 17.70 15 CN (5.34) 1L'"07 16.05 19.62 16 Ethylenediamine 5.77 10.83 14.11 18,19 Diethyltriamine 8.9 14.5 20

Tetraethylene pentamine 154 I 21

The data presented is representative of a wide

range of reported values for the complexes (22,23) and,

generally the order of increasing stability for the

tetrahedral complex can be written as

NH3 < 0H ( CN

The leaching of copper and nickel ores with ammonia -14-

is well documented and is used on a commercial scale in the Sherritt-Gordon ammonia pressure leach the Arbiter process (24 25,26) and process(27). and ammonia has also been used to leach a dolomitic

The treatment by copper ore(28) . of zinc ores leaching has long been ammonia suggested (29,; 30) and more recently the pressure leaching of sphalerite was investigated(31). Wendt(32) has reported that zinc can be recovered from oxidised zinc ores by leaching with ammonia-ammonium carbonate but no large scale operations compared to those for nickel and copper have, however, been initiated.

Attention has also been given to the leaching of oxidised zinc ores with caustic solutions(33) more particularly by Merrill and Lang(34) who showed that the dissolution of zinc oxide in sodium hydroxide was not stoichiometric and an excess of sodium hydroxide

was necessary to prevent hydrolysis of the zincate ion.

They also found that the oxide zinc minerals were slow

to dissolve at room temperatures and boiling temperatures

were needed to achieve satisfactory dissolution. High

sodium hydroxide concentrations were used and typical

results gave a maximum solubility of hemimorphite

of 0.77M dissolved zinc in a 6. iM sodium hydroxide solution -15-

and a similar solubility of hydrozincite in a 4.25M sodium hydroxide solution. All lead and copper oxide minerals associated with the zinc oxide minerals were readily dissolved and hence the principle solution impurities were lead, copper, carbonate and silica.

Cyanide has long been used to recover gold

from deposits because and silver suitable (35,36) of the very stable complexes that cyanide forms with these metals (Ag(CN)2+, log B2 20; Au(CN)2+, lg B2 40) being and new plants are still commissioned (37)*

Many workers have noted the presence of zinc and

copper in the cyanide liquors from the leaching of precious metals and have reported that zinc entered

the mill solutions in two ways. Firstly through the

dissolution of the zinc dust commonly used to precipitate

the gold from solution(38) and secondly from the

dissolution of part of the zinc and copper minerals

associated the Hedley(39) with gold/silver ore(39 40)'

reported that the dissolution of copper minerals was

rapid and that a high degree of solubility was obtained

(1.7M dissolved copper in 6.41M sodium cyanide).

The presence of the copper cyanogen complexes was

considered to lead to difficulties in gold dissolution.

Leaver and Woolf(40) have shown that all zinc minerals -16-

are partially soluble in cyanide solutions but their work was conducted in dilute solutions applicable to precious metal cyanidation and the solubility of the zinc minerals in cyanide was not determined. Much more interest has been shown in cyanide leaching for copper

have been recovery and a number of patents granted (41,42) and furthermore some pilot plant scale studies have also been carried out(43). The use of nitriles for gold recovery has been suggested(44) but the lack of stability data for complexes with copper and zinc precludes any discussion of their usefulness, however, their cost would probably render the extraction of cheap metals like zinc uneconomic.

Iron oxide minerals such as hematite, goethite/ limonite, siderite and iron silicates are often found associated with zinc and copper minerals but cyanide has little dissolution action on them or on metallic iron and steel. Complex iron carbonates such as ankerite may, however, decompose to some extent and iron sulphides decompose appreciably(38). The main losses of cyanide in copper sulphide cyanidation are in the formation of ferrocyanides, thiocyanates and cyanates due to the reduction of Cu2+ to Cu+ in an alkaline environment. In the absence of sulphidic -17-

material the leaching of zinc oxide minerals would not result in any of these cyanide losses and cyanate would not be formed because zinc is not reduced from its 2+ oxidation state. As zinc only occurs in this oxidation state in zinc minerals and in solution the use of oxidising conditions during leaching are unnecessary.

The dissolution of zinc oxide minerals in cyanide has been represented by the following equations(39)

Zinc oxide:

ZnO + 4CN + H2O Zn(CN)42 + 20H (1.2a)

Smithsonite:

ZnCO3 + 4CN Zn(CN)42 + CO32 (1.2b)

Zinc silicate-

Zn Sio 8CN + H2O " 2Zn(CN)42 + Si032 + 20H (1.2c) 2 4+

The reactions produce, on dissolution, the zinc cyanide complex and an alkali which can react with more oxide mineral to form the soluble zincate and higher hydroxy complexes of zinc.

ZnO + 20H + H2O ZnO22 + 21120 1 Zn(OH)42 (1.3)

The solubility of zinc oxide minerals in cyanide solutions will, therefore, depend on both the cyanide and hydroxyl

concentrations. By suitably combining these two reagents,

leaching might be accomplished under less rigorous -18-

conditions than those employed in the caustic leaching

and ammonia leaching of zinc ores.

1.4. Reactions between cyanide and the zinc minerals

The actual equilibria pertaining to the dissolution

of the zinc minerals are not as simple as given in

Equations 1.2a to 1.3 because zinc forms a series of

stepwise complexes with both cyanide and hydroxide ions.

The dissolution of sodium cyanide in water results in

some hydrolysis of the cyanide ion to form hydrogen

cyanide(49)

CN J= + H2O ! HCN(aq)+ OH and logrC_N -9.32 + pH (1.4a) FC Nag

The boiling point of hydrogen cyanide is very low

(25-60C) and an appreciable partial pressure of hydrogen

cyanide can arise above the cyanide solution giving the

equilibrium

[HCN HCN(aq) - HCN(g) for log 1.31 + log fHCN which (aq) =

(1.4b)

The extent of these reactions depend on the pH of the

solution and this is illustrated in Fig. 1.1. which shows

that for a 1. OM sodium cyanide solution, at pH 11.32,

99% of the cyanide is present as the free ion but at pH 9.32,

50% of the cyanide is free and at pH 7.32 only 1% of the

cyanide is free. The equilibrium partial pressure of -19-

hydrogen cyanide is also given in Fig. 1.1. as a

function of pH and is shown to be 10-3.5 atm. pH 11.5 34 but as high as 10-1' atm. at pH values less than 7.

The extent of the hydrolysis of cyanide solutions

prepared in water at pH 7 is indicated in Fig. 1.2.

which shows that, for a 1. OM sodium cyanide solution,

greater than 99.5% of the cyanide is present as the

free ion and the pH of the solution is about 11.55. More

dilute solutions result in a smaller proportion of free

cyanide and a lower pH.

The equilibrium concentrations of the various

soluble species present in the dissolution of zinc

compounds in cyanide solutions can be calculated

provided that all the possible reactions together with

the relevant stability data are known. Manual solution

of the complex polynomial simultaneous equations

involved is tedious and in some cases extremely

difficult without the use of approximations. A computer

program written by Ingri(46) called HALTAFALL,

however, permits solution of even the most complex

solution equilibria problems. This program which

is written in ALGOL has been utilised throughout

this work.

Equilibrium constants usually relate to -20-

activities of the reactants and product rather than concentrations. Conversion of one to the other is then possible with activity coefficients but, unfortunately, information on the latter coefficients is not available for the various complex zinc cyanide and zinc hydroxy

species encountered in this work. In view of the fact

that the equilibrium constants quoted in the literature(22)

for a given reaction show wide variations, by orders of

magnitude in some cases, it was considered that any

deviation of the coefficient from unity was insignificant

compared to the variation in the equilibrium constants.

Concentrations instead of activities were, therefore,

used in the calculations.

The identity of the reaction species in solution

during the dissolution of zinc oxide and smithsonite

are known but in the case of the zinc silicate minerals

the exact nature of the silicate species in solution is

uncertain. It is also doubtful that equilibrium would

ever be attained in this case during the leaching time

envisaged for this study (a few days) and solution

equilibria calculations were, therefore, not made for

the zinc silicate minerals.

Using reactions Equation 1.5 to 1.24 and the

stability data summarised in Table 1.3. the solubility -21-

of zinc cyanide, zinc oxide and smithsonite (zinc

carbonate) as a function of cyanide concentration was calculated. The values obtained together with

the concentrations of the various zinc species present

are shown in Figs. 1.3. to 1.5. inclusive.

The solubility of zinc cyanide in sodium cyanide

solutions increases by 10 times for a tenfold increase

in cyanide concentration (Fig. 1.3). At cyanide

concentrations greater than 10-3M the tetrahedral

Zn(CN)42 complex predominates and in cyanide

concentrations greater than 10-1M the solubility of

zinc cyanide can be equated to the concentration of the

Zn(CN)42 complex. Zinc hydroxy complexation is

negligible. Fig. 1.4. shows the distribution of species

in cyanide solutions saturated with zinc oxide. At

total cyanide concentrations less than 10-2M the

concentration of the Zn(CN)42 complex corresponds

to the solubility of zinc oxide. At this concentration

the formation of zinc hydroxy complexes is small but

in more concentrated cyanide solutions appreciable 2- hydroxy Zn(OH)4 zinc complexation occurs, mainly as ,

and in cyanide solutions in excess of 4M the concentration

of Zn(OH)42 is approximately equivalent to that of

the Zn(CN)42 complex. At equilibrium the solution -22-

pH is very high i. e. a 1. OM sodium cyanide solution has an equilibrium pH of 12.3, which increases in more concentrated cyanide solutions. The formation of hydrogen cyanide is sinäll and in cyanide solutions -1 greater than 10 M the concentration of hydrogen cyanide approaches 10-6 M which is negligible compared to the total. The distribution of species in cyanide solutions saturated with smithsonite is shown in Fig. 1.5. At all cyanide concentrations the formation of zinc complexes other than Zn(CN)42 was very small and hence the concentration of this species determines the solubility of smithsonite in cyanide. The pH of the solutions are not less than 11.5 due to appreciable hydrolysis of the carbonate ion and, therefore, the formation of hydrogen -cyanide can be considered negligible.

The solution equilibria calculations show that

smithsonite and zinc oxide are soluble in cyanide

solutions and that at equilibrium, in the former case,

the cyanide to zinc molar ratio approaches 4/1 whereas,

in the latter case, a lower ratio is obtained becai se

of the formation of zinc hydroxy complexes. The

solubility of zinc oxide can, therefore, be expected

to be much greater than that of smithsonite. The pH -23-

of the leach solutions at equilibrium will always be

high enough so that the hydrolysis of the cyanide to

hydrogen cyanide will be very small and losses of

cyanide to the atmosphere neglible.

The use of cyanide as a solvent for zinc oxide

material seems to be the most promising for a practical

alkaline hydrometallurgical process for the following

reasons.

(1) The most stable zinc complexes are formed with

cyanide.

(2) Sodium cyanide is very soluble (58.3g NaCN/100g

water(45)) and high concentrations of zinc in solution

can be expected.

(3) An excess of cyanide over the stoichiometric

requirements is not necessary.

(4) The dissolution reactions are selective and only

metals that form stable cyanide complexes will be

present as impurities.

(5) Cyanide losses during leaching are expected to

be very small in the absence of sulphidic material

and due only to the formation of cyanate if copper

is present in the material to be treated.

(6) Plentiful supplies of cyanide are available. -24-

The dissolution of zinc oxide involves the following equilibria.

Zn(OH) Zn2+ + 20H KS (1.5) 2(s) 0 Zn(OH)2 Zn(OH) Zn(OH)2(aq) in KS1 (1.6) 2(s) saturated solution, Zn(OH)2 2+ K1 (1.7) Zn + OH Zn(OH)+ ++ Zn(OH) OH ± Zn(OH)2 K2 (1.8)

Zn(OH)2 + OH '- Zn(OH)3 K3 (1.9) 2- Zn(OH)3 + OH- Zn(OH) K4 (1.10) 2+ Zn + CN Zn(CN)+ K5 (1.11) ++ Zn(CN) CN K6 (1.12) v-"Zn(CN) 2

Zn(CN)2 + CN --"Zn(CN)3 K7 (1.13)

Zn(CN)3 + CN Zn(CN)4 K8 (1.14)

Zn(CN) = Zn2+ 2CN K (1-15) 2(s) + SO Zn(CN)2

Zn(CN) Zn(CN)2(aq) in KS1 (1.16) 2 saturated solution, Zn(CN) 2 H2O H+ + OH K (1.17) W NaCN * Na+ + CN (1.18)

H+ + CN HCN(aq) Kb/KW (1.19)

HCN(aq) - HCN(g) Kp,,,,,, (1.20) ý11 V1YI

For the dissolution of smithsonite, the following additional equilibria must also be considered 2+ ZnCO3(s) ' Zn + CO3 Kg0 (1.21) ZnCO3) H+ CO32 " HCO3 K9 (1.22)

H+ + HCO3 H2CO3 K10 (1.23)

11 -1 H2O + CO2 Kp(CO (1.24) 2C03 2) -25-

Table 1.3. Equilibrium data used in the solution equilibria calculations

Equilibrium log cumulative reference constant constant

Ks0(Zn(OH)2) 17 -16.76 I Ksl(Zn(OH)2) -5.57 IK 17 sO(Zn(CN)2) -15.48 Ksl(Zn(CN)2) -4.41

KsO(ZnCO3) -'FO. 85 K1 5.04 23 K2 11.19 15 K3 13.9 23 K4 15.. 1 23 K5 5.34 ( 17 K6 11.07 16 K7 15.06 16 1 19.62 16 1 K8 . K9 I10.3 3 14 4 K10 16.68 144 K 14 w - .0 Kb/Kw 9.32 145

Kp(J 140 cN. ) -1.31 Kp12(CO2) - 1.4 6 144

The values of the equilibrium constants at 250C were selected in accordance with the following principles:

(1) More recent data preferred.

(2) Zinc cyanide complex data from work that did not use a zinc electrode as zinc reacts with cyanide even in the absence of oxygen. -26-

C) b Zv c O ü -az th0 a, a, 3'

0- 0

"4 60 10 12 14 pH

Fig. 1.1. Influence of pH on the free cyanide concentration and equilibrium hydrogen cyanide partial pressure for a 1. OM NaCN solution. 100 114

CN-

co Ia3

60 2

^W. PH

40 7

lO 20

9 16-4 10-3 10-2 16-1 100 101 Total cyanide conch (M) Fig. 1.2. Natural pII and composition of cyanide solutions. -27-

v -; 2 solubility a U 2- (C N) Zn (C N) N n 43 r-- U 0 O. N 3 CN- A 0

4 HCN ''- OH-

Zn (CN)2 5

10-3 10-2 10-1 100 701

Total cyanide concn. (M )

Fig. 1.3. The distribution of species and maximum solubility of zinc cyanide in cyanide solutions. -28-

0 solubility n(CNy

-N 0H- 1

uC CN' O u 2

4J 4 Zn (OH)3 3 Zn(OH), -

Zn (CN) 4

ý-- HCN Zn (0H)2

i64 3 7ö 10'2 16-1 io0 poi Total cyanide concn. (M )

Fig. 1.4. The distribution of species and maximum solubility of zinc oxide in cyanide solutions. -29-

solubility Zn (CN -"--C03

HCO'

2 CN- -ý OH Zn (CN)3 ... Zn (ßH4 C U -3 Zn(OH) O U H ý_ Zn (OH) 'V .4 Cl. N Zn (CN)2

O HCN s

- [:

I H2 CO3 -7j' Zn(CN)+ Zn(9H)+ 0l234 Total cyanide concn. (M Fig. 1.5. The distribution of species and maximum solubility of smithsonite in cyanide solutions. -30-

The solution equilibria calculations indicated that zinc oxide and smithsonite were soluble in cyanide

but the calculations were made assuming activity

coefficients of unity and it is likely that deviations

from the calculated solubility of the minerals will

occur. Other factors such as temperature and pH

should also affect the solubility. Although solution

equilibria calculations indicate the extent to which

the reactions should occur they do not provide any

information as to whether equilibrium would, in fact,

be attained because they do not consider the kinetics

of the reactions.

The aims of the project were, therefore, to: -

(1) Study the dissolution of the zinc oxide minerals

in various cyanide solutions so that the solubility and

rate of dissolution of the minerals can be found.

(2) Quantitatively study the variables that influence

the solubility and rate of dissolution of the zinc oxide

minerals.

(3) Obtain some understanding of the reaction

mechanism involved and formulate rate equations for

the dissolution reactions.

(4) Study methods of recovering both zinc and cyanide -31-

from the leach liquors.

(5) Determine whether a cyanide leaching process is technically and economically feasible or not. -32-

2. EXPERIMENTAL

2.1. Materials

Hand picked samples of smithsonite, hydro-

zincite and hemimorphite were obtained from

Parkinson and Co. Ltd., Somerset, and smithsonite,

hemimorphite and willemite from David New, Utah.

Samples of zincite were not available and 'Analar'

zinc oxide was used instead in the dissolution studies.

1 The mineral lumps were broken to -. cm* in plastic

bags to avoid contamination by iron, and then

ground in an agate vibratory mill. To prevent the

production of excessive fine material the grinding

was interrupted at intervals and the -300

material removed by screening. This material was

screened into the various size fractions required for

the dissolution tests. All reagents were supplied by

Hopkin and Williams and were of 'Analar' grade

except for the chemically precipitated zinc carbonate.

Analysis of the -300 + 75 and -75

fractions of each of the minerals by classical and

X-ray fluorescence techniques gave the results

summarised in Table 2.1. -33-

Table 2.1. Composition of the secondary zinc minerals

-I - Composition (%) Mineral Element -300 + 75/1m -75n

+ 1 Smithsonite I Zn 49.82 0.1 49.43 0.1 (from Cd, Cu, Co, Fe 0.1 - 0.5 0.1-0.5 Mexico) Si 0.1 0.5 (XRF) 0.1-0.5 + 1.29-0.02 (Class) Theoretical Pb, As, Cr 0.05 0.05 content Ca 0.5 -2 (XRF) 0.5-2 52.1%Zn 0.98-0.02 (Class) I Al, K 0.1 0.1 I } Ag 0.1 - La - 0.05 Mn 0.05-0.1 0.08-0.3 Na 1 1

+ + Smithsonite 2 Zn I 47.90-0.1 46.94-0.1 (from Eire) Cd 0.3-0.8 0.3-0.8 Pb 0.05 0.03-0.08 Theoretical Fe 0.1-0.5 I 0.5-1 content Mn 0.05-0.2 0.05-0.2 52. l%Zn Cr 0.05 0.05 Ca 1-5 ( RF) 1-5 2.42-0.05* (Class) K 0.2 0.2 Si 0. (XRF) 0.5-2 0.77-0.02 (Class) Al 0.1 0.1-0.5 S 0.1 0.1 I- P_ 0.1 I Hydrozincite Zn 57.52±0.1 57.40±0.1 (from Cd - 0.01 Mexico) Pb 0.05-2 0.05-0.2 Theoretical Fe 0.5-2 1-3 content Cu Observed (a) 0.1-0.5 59.6%Zn Mn 0.05-0.2 0.05-0.2 Cr 0.05 0.05 Ca 0.3-0.8 0.3-0.8 K 0.1 0.1 S 0.2 0.2 0.1 ý- 0.1 - w... -34-

Table 2.1. Continued ...

Composition (%) Mineral Element -300 + -75pm

Hemimorphite Zn 43.29± 0.1 40.54+0.1 (from Mexico) Cd, Cr 0.05 0.05 As, Pb 0.1-0.5 0.1-0.5 Theoretical Fe 0.5-2 1-4 content Mn 0.05-0.3 0.5-1 54.3%Zn Ca 10-20 10-20 K 0.2 0.2 Cl' Al 0.5+ 0.5 ' Si 9.41 (Class 5-10 (XRF) 5-10 Mg, Na Up to 10 Up to 10(b)

Zinc Carbonate (chemically precipitated 50.68±0.1%Zn (Theoretical content 52. l%Zn)

Note: (a) = Obscured by zinc lines (b) = Obscured by calcium lines -35-

Smithsonite 2 contained appreciably more calcium than the purer smithsonite 1. There was no significant variation in the chemical composition between the two size fractions of each smithsonite sample, apart from the presence of a little more silica in the finer fraction of smithsonite 2. Both fractions of hydrozincite were quite pure and of similar composition, iron being the major contaminant. The sample of hemimorphite was of lower purity than the other minerals, the -75 pm fraction containing less zinc and more iron and manganese than the coarser fraction.

The mineral phases present in each sample were identified by means of microscopic examination by standard determinative techniques(47,48) and

X-ray powder diffraction analysis. A 'Geoscan' electron probe microanalyser was also utilised in the aquisition of further chemical and textural information.

Smithsonite 1

The mineral pieces were botryoidal with a fibrous internal structure, the fibres- being at right

angles to the surface. A white to pink colouration -36-

was indicative of zinc substitution in the crystal lattice by cobalt or manganese. Examination of thin and polished sections showed the major phase to be anhedral crystals of smithsonite (96%) with minor amounts of hemimorphite (2%), siderite (1%), a carbonaceous material (0.5%), wollastonite (0.5%) and possibly calcite.

The smithsonite appeared to exist in four distinctive crystal habits,, as single crystals with or without dark inclusions (Plate 2.1), compact crystalline aggregates (Plate 2.2) and microcrystalline aggregates (Plate 2.3). Examination with an electron probe micronanalyser, however, showed that the zinc content of the different grains was uniform across the grains and the same for each grain. Furthermore the apparent phase changes observed under polarised light were shown to be attributable to differences in the orientation of the smithsonite crystals.

Detailed examination of polished sections - with the 'Geoscan' proved useful in identifying the minor mineral phases intimately associated with the smithsonite. The black areas (A) in the smithsonite grain (Plate 2.4) were found to contain iron (45%Fe) -37-

,ex cM

Plate 2.1.1'i, oLomicroý_raph of snzithsouite I showing uniform crystal phase with some dark inclusions. (Ultraphot, plane polarised light, incident illumination, x 160)

Plate 2.2. Photomicrograph of smithsotlite 1 showing compact crystalline aggregates. (Ultraphot, cross-nicols, incident illumination, x 130) -38-

Plate 2.3. Photomicrograph of smithsonite 1 showing the presence of a microcrystalline aggregate. (Vickers, cross-nicols, transmitted illumination, x 100) -39-

in addition to zinc and were probably siderite inclusions. Scanning across X-X', the concentration of zinc decreased as the electron beam traversed the dark grey areas (B) whereas the concentration of calcium and silicon increased (Plate 2.5). Counts taken in this area showed that calcium and silicon were present in a molar ratio of 1: 1, thus providing further evidence of association of wollastonite with smithsonite. The edges of some of the grains proved to be zinc rich (60%Zn) indicating the possible presence of a little hydrozincite. Finely disseminated areas containing iron, zinc and manganese (up to 9%), probably as mixed carbonates, were also observed.

The. zinc rich areas at the edges of some grains might have been produced by the decomposition of part of the smithsonite during sample preparation.

Burton(49) has reported that mere grinding at room temperature can yield carbon dioxide because zinc carbonate has a relatively low decomposition temperature of 300°C.

The X-ray diffraction studies on smithsonite 1

indicated the possible presence of calcite. No discrete

grains of calcite were, however, found by 'Geoscan' -40-

Plate 2.4. Photomicrograph of smithsonite 1 showing the intimate association of mixed mineral phases. (Ultraphot, cross-nicols, incident illumination, x 160)

Plate 2.5. Polaroid photograph of 'Geoscan' backscattered electron image of smithsonite 1: Scanning across X-X' on Plate 2.4. for Zn, Ca and Si. -41-

Plate 2.6. Polaroid photographs of smithsonite 1 (a) 'Geoscan' backscattered electron image. (b) Zinc distribution. (c) Calcium distribution. -42-

examination although calcium was detected within the smithsonite grains apparently evenly distributed

(Plate 2.6 a, b, c). Exchange of zinc for calcium in the smithsonite lattice is unlikely because the ionic radius of calcium is much larger than that of zinc.

Calcite is, however, isostructural with smithsonite and it is, therefore, possible that the calcium peaks in the calcium trace shown in Plate 2.6 represent calcite that has grown isostructurally with the smithsonite. These calcite or calcium compound inclusions were present to a small extent in almost every smithsonite grain.

The 'Geoscan' examination also showed that

some substitution of zinc by cobalt and manganese had occurred to a small extent.

Smithsonite 2

The form of the mineral pieces was either as

smithsonite 1 or as surface encrustations on a matrix.

The predominant colour was yellow, rather than pink,

suggesting zinc substitution in the lattice by cadmium or iron rather than cobalt or manganese. Euhedral

crystals of galena and fluorite were present in the

sample but as much as possible was removed by -43-

handpicking prior to analysis and experimentation.

Microscopic examination showed that smithsonite,

as anhedral and micro crystalline grains was the

major phase (97%) with minor amounts of siderite

(1%), quartz (1%), calcite, fluorite (0.5%) and

galena (0.5%).

Examination by ' Geoscan' and microscope

showed that smithsonite 2 was very similar to

smithsonite 1.

Hydrozincite

The sample was a massive, friable, white

material with surface iron staining. The hydro-

zincite appeared to be extremely porous. The X-ray

diffraction pattern showed that hydrozincite, calcite

and iron oxides were the main mineral phases present.

Hemimorphite

The sample comprised orthorhombic prisms

of hemimorphite on a heavily weathered iron oxide

matrix. Microscopic examination and. X-ray powder

diffraction analysis showed that the matrix contained

specular hematite and calcite finely interspersed

with goethite/limonite. The hemimorphite crystals

were handpicked from the matrix prior to analysis -44-

and experimentation. The zinc concentration across

the hemimorphite crystals was found, by 'Geoscan'

examination, to be uniform and the only phases

intimately associated with the hemimorphite were

small inclusions of hematite and surface iron oxides.

Point counting resulted in the following mineral

0.5% balance: - hemimorphite 82%, hematite and the

calcite - goethite -limonite matrix 17.5%.

2.2. Analytical Techniques_

2.2.1. Zinc Analysis

The soluble zinc was determined with a

Perkin-Elmer model 290B Atomic Absorption

Spectrophotometer and a Hilger-Watts 'Atomspek'.

A (0 to 3.5 range p. p. m -) of standard zinc solutions

was prepared from a master solution of zinc acetate

(BDH standard solution for atomic absorption

spectrophotometry, lml = 1mg Zn). All standards

and solutions for analysis were made up with 1,000

p. p. m. excess cyanide. This procedure was found

to be necessary, because, in the absence of excess

cyanide the absorbance readings fluctuated and

drifted with time.

A concentration of 1,000 p. p. m. sodium -45-

cyanide was sufficient to swamp any cyanide

accompanying the zinc and also to effectively buffer

all the solutions at approximately pH 11. This latter

point is important because Dong(50) showed that in

the 6-10 pH range the absorbance of zinc varied with

the pH. Outside this range the absorbance was

independent of pH. The relative error for the use

of atomic absorption spectrophotometry for the

analysis of zinc in cyanide solutions was found,

after summation of errors, to be 1.5%.

2.2.2. Cyanide Analysis

To obtain a more comprehensive understanding

of the dissolution reactions, it was thought necessary

to follow the change in concentration of the reacting

free ions, CN the cyanide , as well as concentration

of soluble zinc.

Numerous ways of quantitatively determining

cyanide have been reported and these include

gravimetric(J 1) , conductometric (51) ,. amperometric

colorimetric(53) (52)' , polarographic (54,55)' and

titrimetric(56) methods. The latter technique has

achieved widespread practice in the form of the

Liebig(51) argentometric titration or the Deniges -46-

modification(51) as its precision and reliability are well known The use (57,58 and 59) " of potassium iodide indicator, however, has been reported(60) to result in higher estimations of cyanide concentration due to the occurrence of the end point at constant Ag+ concentration and not constant CN concentration. The potentiometric titration of cyanide solutions with silver nitrate has proved highly accurate and reliable(61).

More recently extensive use of specific ion- selective electrodes has been made for the analytical determination ions in of solution (62) and a cyanide

ion-selective electrode is now commercially

available. The application of ultra-violet absorption

spectrophotometry has also been suggested(63) as a

possible method for the determination of the free

cyanide concentration of cyanide solutions.

The following methods were, therefore,

investigated because they appeared to be the most

promising techniques for the analysis of CN

concentration in solutions containing both zinc and

cyanide.

(1) Specific cyanide electrode because it has the -47-

potential capability for the continuous

monitoring of the CN concentration in the

solvent.

(2) Potentiometric titrations because they have

proved highly accurate and reliable.

(3) Absorption spectrophotometry because it

provides a non-destructive means of analysis.

2.2.3. Cyanide ion-selective electrode

The membrane layer of the electrode contains

a slightly soluble silver halide precipitate. The

potential of the electrode surface has been stated,

by Toth to from the and Pungor(164) , result

continuous dissolution of the silver halide,

AgX(s) + 2CN Ag(CN)2 +X (2.1)

and the dissolution constant KXCNof the electrode

can be written as: - aAg(CN)2 ax 'Sc, (2.21 c- 2 a CN- . aAgx The liberated halide( ions determine the

interfacial potential of the menbrane electrode and,

therefore, in a solution containing cyanide the

potential can be written as: - -48-

RT In (aX + KX, a2CN- (2.3) E =E + CN ' o nF t

where Eo is the standard electröde potential,

ax is the total halide activity on the t- surface of the electrode,

aCN- is the cyanide ion activity,

KX, is the dissolution the and CN- constant of

electrode.

The total halide activity results from the halide ion released from the membrane by the cyanide (ax-) in addition to the halide ion activity in the solution

1aCN- Thus ax -= ax- + ax ax + (2.4) t Equation (2.4) is simplified when the solution contains only cyanide ions, and the potential can then be written as: -

E= Eo + RT In (iaCN + KX, (2.5) CN aCN-)-)

This equation has been proved to be valid in the range 10-4Mto 2M cyanide ion concentration, and 1M the electrode response to be linear up to 10 cyanide ion concentration..

The application of a cyanide ion-selective electrode to the analysis of free cyanide in zinc -49-

cyanide solutions was studied using an ORION

IONALYSER cyanide activity electrode model 94-06,

an ORION single junction reference electrode model

90-01, and a Pye Unicam model 290 pH meter. A

range of standard solutions (10-6M to 10-1M) was

prepared from ORION Ionalyser cyanide standard

solution (specification 0.01M KCN±O. 0001M in 0.1M KOH)

by dilution with 0.01M KOH. This concentration of

alkali prevented hydrolysis of the cyanide and also

maintained a constant ionic strength. The electrode

response was linear at cyanide concentrations

greater than 10-5M, with a slope of -56mV which

compares well with the Nernstian slope of -59mV.

To determine whether or not free cyanide

could be determined with the cyanide selective

electrode in the presence of zinc, the electrode

response was measured in solutions of different

cyanide to zinc molar ratios. This response was

converted to a free cyanide concentration using the

callibration curve obtained in the absence of zinc.

In all cases the solutions were made up in 0.1M KOH.

The results obtained are presented in Table

2.2 and they show that, in all cases, the determined -50-

free cyanide concentration was similar to the total cyanide. Calculations by HALTAFALL showed that in the presence of 0.1M KOH the zinc cyanide complexes were almost completely dissociated.

Table 2.2 Estimated free cyanide concentrations _

of zinc cyanide solutions

(ZnSO [CN ]* LNaCN KOH E `CN1 4ý est (mM) mM (mm (mm) (mV) (mm)

9.89 10 2.40 100 -258 9.8 I l 9.91 10 2.00 100 -258 9.8 1 9.93 10 1.67 100 -259 9.9

* calculated by HALTAFALL

An alternative method of providing a constant

ionic strength is to use an inert eletrolyte such as

potassium nitrate. However, in diluting the various

zinc cyanide solutions to fall within the recommended 5M operating range (10- to 10-3M) the pH would be

reduced and the zinc cyanide complexes would

dissociate. This point is emphasised in Fig. 2.1.

The cyanide selective electrode obeys equation

2.4 at cyanide concentrations up to 2M and therefore

the response of the electrodes was determined at -51-

cyanide concentrations above the normal operating range. Standard solutions were prepared with

0.1M KOH but the zinc cyanide solutions were not, although they were of a comparable ionic strength.

Calculations of the ionic strength in the latter case were made by HALTAFALL and the values obtained were used in conjunction with activity coefficient data supplied by Orion to convert concentrations into activities. The results obtained are summarised in Table 2.3 where it can be seen that in all cases the selective electrode gave free cyanide concentrations higher than those predicted from solution equilibria considerations. Clearly part of the zinc cyanide complexes must in some way be interacting directly, with the electrode. -52-

100 Zn(CN)2

Zn (CN)3 cýoF 1.0 N Zn (CN)4 4 ; OF V) 'pH C ro"5 10 IC PH WF O v HCN 10.0 0

CN " Ar L 104 103 102 101 10a Dilution

2.1. Equilibria changes on dilution of a solution containing 1. OAT NaCN and 0.2M ZnSO4 ((, a i-, mated by HHAL.TAFALL). -53-

C) o 0

U " O C) N ß) o LO w M

0 U' v C)1O N d+ N N ' 0 ul 1 a) Ü 0 M 0 O LO N Cd Lo M N r+

CM C) 'CIM 00 N N M C+9 CO U '> M ri 00 N M N r-i O C) M N N M M M co N NI - 1 1 1 1 1 1 1 1 1

cri tz 10 U L4 r, co N. O Cd N CM cu 10 ý] t . -1 N dý ci ý}I M M M cd L- t- C- c- N LN c- N LN W 4-1 O o O 0 0 o O o 0 0 H

0 .. a m ",1 .. ý f~ O O O r-1 O U) IR14 O Q) O O M O 0 $4 CV N M ij+ cr N I: - ' V-4 T-4 . -I r-1 r-1 V-4 r-i . -1 b0 O La j 0 F o O O O 0 O M 0 O 0 1 1 1 1 1 ct3 U ºv+ r-1 r-4 r-i r-1 Q) "-i _A N CD r7l 11 It Q ^ LL) N O O ý-r Cd Fr' 1 1 1 ! p O N to O Lc) a) CV N r-1 r-- r1 O 0 z o O o o 0 o O O cd O L1') -4 O C O O O LF-!. 1 v-I v-I e--1 ri r4 -54-

Toth and Pungor (64) also studied the effect of zinc ions on the response of a silver halide ion selective electrode to cyanide ion concentration and concluded that the zinc cyanide complexes participated in the electrode membrane reactions in a manner analogous to cyanide. In their work, however, more dilute -solutions at a pH of 11 was used throughout and under these conditions the formation of zinc hydroxy complexes is substantial. The high electrode response obtained by these authors need not, therefore, necessarily be due to zinc cyanide complexes but to the cyanide released as a result of the complex dissociating.

It was also stated by Toth and Pungor that the cyanide electrode was not responsive to HCN, but it has been suggested (65) that undissociated HCN can also generate halide ions by the reaction.

AgX(s) + 2HCN Ag(CN)2 + 2H+ +x (2.6)

An analogous reaction can also be postulated for the zinc cyanide complex. 2+ 2AgX(S) + Zn(CN)42 - Ag(CN)2- + 2X + Zn (2.7) -55-

KX, = aZn2+ a2 (2.8) and Zn(CN) a2Ag(CN) -. X- 42 aZn(CN)42-

2 Ks0 KZn(CN)42 (2.9)

K2 " Ag(CN)2

= 2.62 x 10-10

KX, is the dissolution constant of the where Zn(CN)4

AgX electrode.

Kso is the solubility product of AgX,

and KZn(CN) and KAg(CN - are the dissociation 42- for Zn(CN) constants 42 and

Ag(CN)2 respectively.

In conclusion, the results have shown that

the cyanide ion selective electrode cannot be used

to give a direct and accurate determination of the

free cyanide concentration in the presence of zinc

because of the interaction of the zinc cyanide

complexes with the electrode membrane.

2.2.4. Potentiometric Titrations

The potentiometric titrations of cyanide

solutions were conducted with a silver electrode, a

saturated calomel reference electrode (with a salt

bridge) and a Pye-Unicam model 290 pH meter. All -56-

solutions were freshly prepared and agitated with a magnetic stirrer.

A typical titration curve (Fig. 2.2) shows the presence of two well defined end points, the first corresponding to the completion of the reaction

2CN + Ag ? Ag(CN)2 (2.10) and the second to the complete precipitation of silver cyanide [Ag+ Ag(CN)2 + Ag+ Ag(CN)2-3 (2.11)

The first end point is. the more sensitive although both are satisfactory for the determination of free

cyanide concentration

The application of this method to the analysis

of free cyanide concentration of zinc cyanide solutions

has been investigated. Previous work by Willis and

Woodcock(60) led to the conclusion that it was not

possible to determine CN in excess of Zn(CN)4-

as there was no pronounced end point for the removal

of CN Their work was, however, only conducted in

dilute solutions where the total cyanide concentration

was approximately 10-2M.

To assess the feasibility of this method, a

detailed study of the titration equilibria was made by -57-

computer calculation with HALTAFALL and complete potentio: netric titration curves plotted by combining the results with the Nernst expression. The effect of the total cyanide concentration on the shape of the titration curves can be seen from Fig. 2.3. All solutions contained a cyanide to zinc molar ratio of 1M, 5: 1. At cyanide concentrations less than 10 dissociation of the zinc cyanide complexes was indicated by a series of weak, ill-defined inflexions, until the end point corresponding to the first total cyanide end point was reached. The remainder of the curve was as in Fig. 2.2. When the total cyanide concentration was greater than 10-1M, a definite end point corresponding to the precipitation of Zn(CN)2 was observed. Thus when all of the CN in excess of that complexed with zinc was reacted with Ag+ and the cyanide to zinc molar ratio had decreased to 4: 1, zinc cyanide precipitated. The precipitation reaction

can be written

Zn(CN)42 + Ag+ Ag(CN)2 + Zn(CN)2 (2.12)

Addition of more silver ions produced a

second end point that corresponded to an estimation

of the total cyanide minus that precipitated as zinc -58-

cyanide. With further addition the zinc cyanide precipitate dissolved and finally all the cyanide was [Ag+Ag(CN)2. ] precipitated as This final end point gave a good estimate of the total cyanide concentration.

The computer study, therefore, shows that the first end point can be used to determine the free cyanide concentration of a zinc cyanide solution provided that the total concentration of cyanide is greater than 10-1M.

To verify the titration model, a number of

10-1M total cyanide solutions of varying cyanide to zinc molar ratio were prepared (Table 2.4) and titrated with 10-1M silver nitrate. The titration curves obtained (Fig 2.4) exhibited the same shape as the calculated curves (Fig 2.3). However, the dull white zinc cyanide precipitate was formed at a higher potential than that predicted by HALTAFALL as can

be seen from Table 2.4. -59-

-6

-A

IL v

E W

0 10 20 30 40 0.5 M Ag NO3 (cm3 )

ii2.2. Poten, iometric titration curve for the titration of 30cm 0.5'MT NaCN with 0.5M AgNO3. -60-

. /ý /ý /ý

-M: v

05 10 15 20 25 Ag NO3 (cm3)

Fig. 2.3. Influence of total cyanide c9ncentration on the titration of cyanide solutions (25cm ) with AgNO (equal concentration to cyanide) at a cyanide tö zinc molar ratio of 5/1 (calculated by 1ALTAFALL). -61-

C.*) O,

N N

d N

(31

.. Ö

U'l

T E (m V) ref. S.C. E.

Fig. 2.4. Potenýiometric titration curves for the titration of 25cm zinc cyanide solutions with 0.117 AgNO3. -62-

r. O cd

10U . -4 aý Cd

> r-1 cN ce) G) M N Cl N C

41 cd 1) 1-4 w cd 14 0 . -1 4 co G) Cd ºýý+ Cl di CO cd 7-1 -4 Fi 4 4 U ýl w 1 1 1 1 1 .x

0 ^',, ^ O "O O O 9 co cq .I. Cl M M M O co O w I 1 1 111

O O ^ U] C O L& b r O CO Co to 1 0 14+ y r04 ý--1 Fý ý

N O O 4-4 LÖ N O 0 O N GO O L) 0 r-1 r-1 N Cl 0 N, r-1

U) O O O O O O O O O O O

0 U

N GQ U Q 0 w 1-4 ý .Qcd E-+ -63-

The theoretical potentials were always 125 to 135mV more negative than those found experimentally. Such a systematic discrepancy is not, however, surprising because of the uncertainty of the equilibrium constants for the zinc cyanide complexes and the solubility product of Zn(CN)2. An increase in the value decrease in Ks0 by of .84 or a an order of magnitude would give quite close agreement between the two sets of potentials and would fall well within the

84 KS0 in the literature(22) range of and . values quoted

The end points corresponding to free cyanide, total cyanide-cyanide as Zn(CN)2 precipitate and total cyanide were found from the experimental titration curves and are compared in Table 2.5 with the calculated end points. The second end point was of no analytical significance because part of the Zn(CN)2 precipitate redissolved before all the zinc had been precipitated.

The final end point corresponded extremely well to the total cyanide concentration. -64-

'",- o o O OO 0 o 0o 02 ci w 0 0 +- 0 o -4 ý. +; 'i o o c O O 0 Cd .0 O O O» O O a z4 V-4 v--4 r-4 r-4 U) 14-4 ý+ M a) O Z CO co c- . --1 O M Co CM 0 0 0 ri) l N N N N NU) U cd ".-4 Ü b Uo ..-ii rl fj -1 GO d4 cm N e Cda z CO Co C- N CD ab Cd u U It

O 0 'Z 4. . U) ý-i U) CD N o U, Co Co c- m cm ax , o ý U M ca s. . -4 cd M M Co LO .., N hA Co C) N N a) 4-4 O M M O O o o0 N a O a)

O O a) Ü CO O 0 Co LO M N o U] an d

0 ^ U] a Zý N N Co ,01 14 CO 0 CO 1f3 M 0 T 1 ^i M ca 1

0 0 z CO (1) Ln c43 Uf)_-J CO a) ;3 (n b M N CD N M w M N Ü M GD d1 CV N ri Cd U Qv .ýN E ni Cd -4 E-{ N b bß r. a) ö ý W U A W ý O -65-

The free cyanide concentrations determined from the first end point were in fairly good agreement with the calculated values although slightly higher.

The end point was not very well defined and it was noticed that there was a slight decrease in potential after the first precipitation of Zn(CN)2. This phenomenon was also noted by Willis and Woodcock who suggested that there was a supersaturation of zinc in the solutions which would account for the slightly high estimation of the free cyanide concentration.

Titrations of 0.5M total cyanide zinc cyanide solutions substantiated the results obtained with the

10-1M total cyanide solutions. The first and third end points were very distinct and the titres corresponded to that predicted from theory.

In conclusion, the potentiometric titration of zinc cyanide solutions with silver nitrate provides an analytical technique for the approximate determination of free cyanide concentration and an accurate estimation of the total cyanide concentration. The first end point corresponds to the free cyanide concentration of the solutions, but as the total cyanide concentration is reduced the detection of the end point -66-

becomes more difficult. At total cyanide concentrations less than 10-1M, the dissociation of the zinc cyanide complexes produces erroneously high estimates of the free cyanide concentration. The major source of error in this method is in the location of the end point and not in the experimental procedure.

Potentiametric titrations with cadmium nitrate

Although equilibrium data for cadmium are

it has been that uncertain, reported (17,22) cadmium forms cyanide complexes that are slightly less stable than the zinc cyanide complexes. Titrations with Cd2+ instead of Ag+ might, therefore, allow a more accurate determination of the free cyanide concentration because the zinc cyanide complexes would be less likely to dissociate. Cadmium cyanide is fairly soluble and would not precipitate during the course of the titration. The free cyanide end point would, therefore, correspond to the completion of the stepwise equilibria

Cd2+ + CN Cd(CN)+ Cd(CN)+ + CN = Cd(CN)2 Cd(CN) CN Cd(CN) 2+ 3 Cd(CN) 3+ CN ? Cd(CN)42 (2.13) -67-

A number of 10-1M total cyanide concentration solutions with various cyanide to zinc molar ratio were titrated with cadmium nitrate and the potential followed with a cadmium electrode and a saturated calomel reference electrode with a salt bridge. The cadmium electrode was observed to blacken during the titration, however the surface coating was removed by scrubbing before each titration. The titration curves (Fig. 2.5) show a series of weak inflexions until a precipitation occurred at a potential of about -550mV, even in the absence of zinc. A study of the titration equilibria by

HALTAFALL was made and the calculated potentiametric titration curves plotted. These curves showed only one inflexion point corresponding to the complete formation 2+ of Cd(CN)42-. A further addition of Cd resulted in the dissociation of the cyanide complex until there was a precipitation of Zn(CN)2 and in the absence of zinc

Cd(OH)2.

An estimation of the free cyanide concentration was made from the final titration end points and the values obtained compared to the theoretical values.

In all cases the estimated concentration and the electrode potential at the end point was higher than -68-

a, L

E LU

05 10 0.1M Cd(N03)2 (cm3)

Fig. 2.5. Potentiometric titration of 30cm zinc cyanide solutions 0. lM Cd(1703), with ). -69-

that predicted. These results indicate that the original assumption was incorrect and that cadmium cyanide complexes are more stable than those of zinc. Titrations with Cd(NO3)2, therefore, do not provide a useful technique for analysis of CN concentration in zinc cyanide solutions.

Otherotentiometric techniques

Back titrations of zinc cyanide solutions with cyanide after precipitation with silver nitrate proved unsuitable due to the lack of a distinct end point on the redissolution of the Zn(CN)2 precipitate.

Solution equilibria calculations suggested that titrating the zinc cyanide solutions with zinc ions

should provide an analytical method for the determination of the free cyanide concentrations. The end point should be marked by the precipitation of Zn(CN)2 which would only occur after all the free cyanide had 2 complexed with the added zinc to form Zn(CN)4ý.

Irreproducible results were, however, obtained with this method because of the blackening of the zinc

foil electrode. Zinc and zinc amalgam electrodes have been reported(66) to be attacked by aqueous

cyanide solutions and their use for potentiometric -70-

titrations in the presence of cyanide is, therefore,

not to be recommended.

2.2.5. Ultra-Violet Absorption Spectrophotome_

The chemical methods of analysis that were I investigated all led to changes in the equilibria of

the zinc cyanide solutions. Absorption spectrophoto-

metry provides a possible non-destructive method

for the determination of the free cyanide concentration.

Energy is absorbed only by an unfilled electron

energy level and as Zn2+ has a complete outer M shell,

it will not absorb radiation in solution. However, the

divalent carbon ion has an incomplete L shell and

cyanide and cyanide complexes will, therefore, be

absorbing species. The absorption spectrum of

K2 Zn(CN)42 in aqueous solution was studied by

Brigando (67) who found that the absorption by

K2 Zn(CN)42 did not follow Beer's Law and attributed

this variation of the coefficient of absorption with

to the dissociation the Zn(CN) concentration of 42 .

The absorption spectrum approached that of KCN

and was said to confirm the ionic character of the

bonding between the metal and cyanide groups. The

very low coefficients of extinction increased slowly -71-

on going from the visible to the ultra-violet with a peak at about 240nm. The absorbance of cyanide,

cyanate and formate, which are possible hydrolysis

products in cyanide solutions, were measured by

Simpson and `Vaind(68) in the wavelength range

320-220nm, and they concluded that the molar

extinction coefficients of cyanate and formate in the

ultra-violet region were too small to influence their

results.

Noblitt(63) has described the use of absorption

spectophotometry for the monitoring of cyanide

solutions in flotation pulps containing copper ions.

The complex absorbance curves of the copper cyanide

complexes were resolved into individual Gaussian

peaks by the use of a curve resolver after the calcu-

lation of the wavelength at which each species absorbed.

He concluded that unless excessive interference was

present, the free cyanide concentration could be

monitored by the use of a monochromator set at

208.6nm.

In view of the limited amount of information

available a study was made of the application of

ultra-violet spectrophotometry to the analysis of 'the -72-

free cyanide content of zinc cyanide solutions. The absorption spectra of zinc cyanide solutions were found with a Perkin Elmer model 214 double beam grating spectrophotometer using ultra-violet light and quartz cells. The machine performance was checked with a holmium oxide doped slide. At wavelengths shorter than about 200nm, oxygen within the instrument monochomator and cell compartment causes pronounced absorbance peaks, thus weakening the intensity of the ultra-violet light before it reaches the detector.

However, it was not necessary to purge with nitrogen at wavelengths longer than 196nm.

The absorption spectra obtained with different

concentrations of sodium cyanide is shown in Fig. 2.6.

As the concentration of cyanide increased, the absorbance peak shifted to longer wavelengths. These

results show, therefore, that even in the simplest

case with no zinc present, cyanide solutions do not

obey the Beer-Lambert law.

A number of zinc cyanide solutions were

prepared by the dissolution of zinc cyanide in sodium

The cyanide solution . absorption spectra obtained

(Fig. 2.7) all show three peaks at the wavelength . -73-

(a) 0.05M NQCN 2.0 (b) 0-5M NaCN

1.5 (c)

a, " 1.C c 0 I 0

Q i

0.:5 (b)

(a)

200 220 240 260 280 JULI Wavelength (nm)

Fig. 2.6. Absorption spectra for cyanide solutions at different concentrations. ' -74-

S I u

.Q O

-0 ";x

200 220 240 260 280 300 . Wavelength (nm)

Fig. 2.7. Absorption spectra of zinc cyanide solutions (a) 1. OM CN +0.1MZn (b) 1.0M'I CN + 0.25M Zn (c) 1.5M CN + 0.25NI Zn -75-

summarised in Table 2.6.

Table 2.6. Wavelength of absorption peaks of zinc

cyanide solutions

NaCN Zn Absorption peaks (M) (M) (nm)

(a) 1.0 0.10 212,265,284,

(b) 1.0 0.25 215,265,285,

Curve (b) was obtained under conditions where

the free cyanide concentration was small and the

complex Zn(CN)4 predominated. The peak obtained

at 215nm would, therefore, appear to correspond to

Zn(CN)42 However, this is . an oversimplification

because in both solutions (a) and (c) the free cyanide

concentration would be significant and considerable

absox'Pcion due to the presence of CN would be

expected. It, therefore, can be concluded that the

absorption peaks for Zn(CN)42 and C.N are so close

together that spectrophotometry cannot be readily

used to determine the free cyanide concentration.

2,2.6. Conclusion

The only practical method for the determination -76-

of free cyanide in zinc cyanide solutions appears to

be potentiometric titration with silver nitrate.

Although the method is unreliable in dilute solutions,

the errors. incurred at total cyanide concentrations

greater than 10-1M should not exceed 5% relative.

2.2.7. Determination of the free cyanide concentration of

leach liquors

The application of a potentiometric titration

with silver nitrate as the titrant to the determination

of the free cyanide concentration of leach liquors was

studied. Various amounts of zinc oxide, hemimorphite

and smithsonite were dissolved in a 1.00Pß sodium

cyanide solution and, after removal of any residue,

the clear solutions were titrated with 1. OOM silver

nitrate. The potential was followed with a silver

electrode and measured relative to a saturated

calomel electrode with a salt bridge.

The titration of cyanide solutions containing

dissolved hemimorphite (Fig. 2.8) resulted in a

gradual increase in silver potential until a gelatinous

brown precipitate was obtained at-the end point

corresponding to the precipitation of silver cyanide

which denoted the total cyanide concentration of the

solution. The precipitate was probably a mixture of -77-

silica gel and silver cyanide. There was no end-point corresponding to the free cyanide concentration of the solution.

Similarly, only the total cyanide end-points were observed when cyanide solutions containing zinc oxide were dissolved. Solution equilibria calculations

showed that this was due to the high pH of the solutions

and the formation of zinc hydroxy complexes which

allowed the zinc cyanide species to completely

dissociate, on addition of silver nitrate, without

precipitating zinc cyanide. During the titrations there

was a small amount of precipitation of zinc hydroxide.

Titrations of solutions containing dissolved

smithsonite resulted in a free cyanide end-point that

was not well defined and a sharp end-point for the

total cyanide concentration. The free cyanide end-

point was marked by a precipitation of zinc carbonate

probably according to the reaction.

Zn(CN)4w + CO3 + 2Ag+hAg(CN)2+ + ZnCO3(S) (2.14)

The titration curves showed that this precipitation

occurred at a potential about lOOmV more negative than

that calculated from solution equilibria considerations

(Fig. 2.9). This difference can, however, be accounted -78-

for by the spread in reported values for the equilibrium constants involved. The experimentally determined end-points corresponded to a free cyanide concentration rather higher than that predicted from theory. The error in estimating the end-point increased with an increase in the zinc concentration of the solution. This is shown in Table 2.7. where the experimentally determined free cyanide cmcentration is reported as a probable maximum and minimum limit.

Table 2.7. Free cyanide concentration of cyanide solutions containing dissolved smithsonite

Cam] [ Jptl [C Ilcd () (M) (M)

0 1.00 - 1.02 1.00 0.047 I 0.85 - 0.87 0.81 0.099 0.74 - 0.76 0.60 0.132 0.53 - 0.63 0.47 0.155 0.31 - 0.39 0.28 0.202 0.10 - 0.20 0.07 0.246 0.08 - 0.02 0.01

The high experimental results were probably

caused by the difficulty in judging the onset of precipitation.

Potentiometric titration with silver nitrate

provides, therefore, only an approximate method

for the determination of the free cyanide concentration -79-

of cyanide solutions containing dissolved smithsonite

and is unsuitable for solutions of high pH or those

containing dissolved silica.

The use of the computer program HALTAFALL

for solution equilibria calculations has proved

extremely useful as an aid to interpreting the

analytical results and can be used to calculate the

free cyanide concentration of zinc cyanide solutions

provided that knowledge of the species in solution is

available. For solutions containing dissolved

smithsonite or zinc oxide this method should give

fairly close approximations to the actual free cyanide

concentrations considering the limitations of the

equilibrium data. In the case of solutions containing

dissolved hemimorphite the free cyanide concentration

can be calculated if it is assumed that the pH change

on dissolution of the mineral in cyanide solutions is

small. -80-

_6 [Zn] [CN+1free (M) (M)

-500 1 0.02 0.91 2 0.04 0.05 3 0.06 0.77 4 0.07 0.71 -400 5 0.09 0.64 16 0.15 0.40 1 4.1 V In -300 precipitation point L 5 -200 E 6 W

0 2.0 1.OM 1.0AgNO3 (cm's )

Fig. 2.8. Titration curves of hemimorphite leach solutions with 1. OM AgNO3. -81-

eooi

600

400

eil 6200

0 :.Ih. E W ZOG

40C

012 1.OM AgNO3 (cm3)

Fig. 2.9. Comparison between theoretical and experimental titration curves for smithsonite. -82-

3. THE SOLUBILITY OF THE SECONDARY ZINC

MINERALS IN AQUEOUS CYANIDE SOLUTIONS

3.1. Experimental procedure

The solubility of each zinc mineral in a

cyanide solution was determined in terms of the

maximum amount of the mineral that could be

dissolved in a solution at a given total cyanide

concentration i. e. when the solution was saturated

with respect to the solid minerals. Values of the

solubility are quoted in terms of the zinc concentration

in solution corresponding to the maximum amount of

mineral dissolved.

The solubility determinations were conducted

in a 250cm3 'Quickfit' reaction vessel. The vessel

was flanged, 7.5cm in diameter and had a rounded

bottom. Multiple openings in the flanged lid of the

reactor vessel accommodated additional apparatus and

allowed the extraction of samples for analysis.

Agitation was supplied by an impellor that entered

the vessel through a mercury seal situated at the

centre of the lid. The impellor, which was connected

to a variable, constant speed motor, was a curved,

plastic blade that fitted almost flush to the bottom of -83-

the. reactor. This impellor design proved the most efficient me ans of ensuring total suspension of the mineral particles. The temperature of the solutions ± was controlled to within 0.1°C. of the desired temperature by immersing the reaction vessel in a constant temperature water-bath. Excessive evaporation of the leach solution was prevented by fitting a water- cooled condensor to the top of the reactor.

The solubility determinations were conducted by the stepwise addition of small quantities of the mineral to the reaction flask until no more w uld dissolve. The point of maximum dissolution was found by taking samples of the suspension at regular intervals. Each sample was centrifuged and then

0.100 to 0.250- 0.001 cm3 of the supernatant solution was removed by an 'Agla' micrometer syringe for zinc analysis. A small volume was taken to minimise errors caused by. pulp volume changes and also, by dilution of the solution to that required for zinc analysis.

The remaining solution and all solids were returned to the reaction vessel.

The cyanide leach solutions were prepared

immediately before the solubility determinations by -84-

dilution of a freshly made, standardised, stock solution of 'Analar' sodium cyanide. This procedure was found to be necessary because of the apparent instability of the cyanide solutions. It was not possible to keep concentrated stock solutions for any

length of time because, after a few days, they became yellow and a gradual decrease in the free cyanide

concentration was noted accompanied by the appearance

of a brown, flocculated precipitate. Similar behaviour has been reported by Williams(45) who showed that,

in cyanide solutions, hydrogen cyanide decomposed to

cyanate, formate and ammonia and eventually a brown

flocculated precipitate of azulmic acid was formed.

Freshly prepared solutions were made with a free ± cyanide concentration within 1% of that required.

It is well known that, according to Ostwald(69)

the solubility of colloidal sized particles is greater

than that of coarser particles. From thermodynamic

and Gibbs free energy considerations an expression

for the solubility of a crystalline solid can be derived

and for the case of spherical particles has been given

by Orr(70) as -85-

2O"mM a2 1_ log _1 rl (3.1) ý1 - 2.30SRT r2 where a1 and a2 are the activities of the solute in solution of the larger and small particles respectively, r1 and r2 are the radii of the larger and smaller particles respectively,

a-m is the mean interfacial tension,

M, the molecular weight of the crystalline material

P the density of the crystalline material,

R, the gas constant, and T, the absolute temperature.

Empirical equations have often been described

Schindler has that the in the and reported (23) change solubility product of zinc oxide with particle size is given by the equation

(3.2) log Ks0 = -16.82 + 50d-1 where d is the diameter of the particles

Furthermore, the solubility of a crystalline material has been shown(71) to vary with the crystal face exposed, surface roughness and irregularities.

The solubility of silica has been noted(72) as changing with particle size because of the presence of a -86-

disturbed or disordered layer on the surface of the

quartz particles.

Solubility tests on -300+75/um and -53pm

smithsonite, however, showed that under the conditions

used particle size did not have a significant effect on

solubility. In all subsequent mineral solubility tests,

-53pm material was used.

The leach liquors saturated with mineral were

analysed by X-ray fluorescence and classical methods

and the insoluble residue was examined by X=ray

diffraction.

3.2. Results

3.2.1. Influence of cyanide concentration

The solubilities of zinc oxide, chemically

prepared zinc carbonate, smithsonite 1 and 2,

hydrozincite and hemimorphite were determined in

solutions of varying cyanide concentration at a

temperature of 250 C. and the results are shown in

Fig. 3.1.

In a large excess of cyanide, all the zinc

minerals were completely dissolved. Zinc oxide was

the most soluble material but the solubility curves for

hydrozincite and the chemically prepared zinc carbonate -87-

/c/1

Q1 N

ýý '.:.'1^

V J 0 N

12345 cyanide concn. (M )

Fig. 3.1. The solubility of the secondary zinc minerals in sodium cyanide solutions. -88-

were very close and only slightly lower than that of zinc oxide. The solubility curves for these materials were not quite linear and a slight increase in the slope and hence decrease in the cyanide to zinc molar ratio was obtained as the cyanide concentration increased.

This point is demonstrated in Table 3.1.

Table 3.1_ The change in cyanide to zinc molar ratio with cyanide concentration

Cyanide Cyanide to zinc molar-ratio Concentration (M) Zinc oxide Hydrozincite

1.0 3.17 3. '34

2.0 3.01 3.08

4.0 ' 2.81 2.94

The smithsonite and hemimorphite samples

were less soluble than the other secondary minerals,

hemimorphite being the least soluble. The solubility

of both samples of smithsonite was linearly dependent

on the cyanide concentration up to a cyanide concentration

of 4. OM. At higher cyanide concentrations, the

solubility curves departed from a straight line and a

greater zinc concentration was found than expected -89-

from a linear dependency. When these leach liquors were centrifuged, filtered and the clear solutions allowed to stand, a fine white precipitate was observed after a few hours. The zinc concentration of the solutions was found to have decreased thus indicating that there was an apparent supersaturation of zinc in the leach solutions at cyanide concentrations greater than 4. OM.

Smithsonite 2 appeared to be slightly more soluble than smithsonite 1, but this difference was shown not to be significant after consideration of the cyanide to zinc molar ratio and standard deviation in each case i. e. for smithsonite 1 the cyanide to zinc molar ratio was 4.15 with a standard deviation of 0.14, and for smithsonite 2 the values were 4.06 and 0.15 respectively.

The hemimorphite was only slightly less soluble than the smithsonite samples but in this case equilibrium was only obtained, if at all, very slowly.

Maximum dissolution was assumed to have occurred when, on addition of fresh mineral, there was no measurable

change in the zinc concentration after 48 hours. The

solubility curve was linear with a slope corresponding -90-

to a cyanide to zinc molar ratio of 4.28 (standard

deviation 0.18).

The examination of the dissolution products of

smithsonite and hemimorphite gave the results

summarised in Tables 3.2 and 3.3 respectively. -91-

W F---T

N 0 Cd A V fý NQ "ý1 N . Q'i O 3 'a )? 41 cnx ö° "", cd d) 'c7 O r. 0 -U2U 1 4" o -4 cn J cd N av a) a)

U U

Co O a O 0 d O +I, 1 .4 033 N v C) vO aý +J NU N Cl) O U] -- 0

O O C; cn +I v1 U 0 p., U U b O 10 ä in Ei O

0 tv°-ý z 0 ö Cd N 0 `_ý, ýÜ +1 U) v 0 UÜÜ 0 Co m cd M N0N ý w b a) C: CD Ei 4 LO o E-{ o L, LO O cri M "ý 0O äý to M nnnU-ýC; C; M Ann ö Lo I Lo o + O Lt) O cd Lo0 cd c; C; vvv ýý Lo E-4 E-+ V vv -92-

The leach solutions obtained from the

dissolution of zinc oxide and chemically prepared

zinc carbonate contained no measurable impurities

whilst the hydrozincite solution had only a trace

amount of iron.

The only impurities detected in solution during

the dissolution of smithsonite were trace amounts of,

copper and sulphur. X-ray diffraction analysis of the

residue showed that during the dissolution of smithsonite

no significant conversion of the smithsonite to hydro-

zincite occurred.

The only impurities present in solution on

dissolution of the hemimorphite, apart from silica,

were trace quantities of copper, iron and phosphorus.

Zinc and silicon in solution were present in a molar

ratio very close to 2 to 1.

3.2.2. Influence of temperature

The influence of temperature on the solubility

of the secondary zinc minerals was investigated at a

cyanide concentration of 1. OM and at temperatures up

to 80°C (Fig. 3.2. ). In all cases an increase in

temperature produced an increase in solubility and

the most marked increase was obtained with -93-

hemimorphite. There was a tendency for the

temperature effect to be smaller at temperatures above

60 0C than below.

3.2.3. Influence of sodium hydroxide addition

In aqueous sodium cyanide solutions the

cyanide hydrolyses to form hydrogen cyanide and

hydroxyl ions. Zinc complexes with both cyanide and

hydroxyl and so the solubility of zinc oxide minerals

should be dependent on both the hydroxyl and cyanide

concentrations. Tests were, therefore, conducted to

determine the influence of hydroxide concentration on

the solubility of zinc oxide minerals in the presence and

absence of sodium cyanide. The results obtained are

shown in Fig. 3.3.

The addition of sodium hydroxide to the leach

solvent markedly increased the solubility of zinc oxide

and hemimorphite in 1. OM sodium cyanide but the

solubility of smithsonite was only slightly increased.

The slope of the solubility curve obtained with zinc

oxide decreased with an increase in hydroxyl concen-

tration whereas that obtained with hemimorphite

increased. In the presence of excess smithsonite, a

white precipitate was observed on the mineral surface, -94-

X-ray diffraction analysis indicated that the precipitate

was hydrozincite.

In the absence of cyanide the solubility of the

minerals only increased slowly with sodium hydroxide

concentration. Smithsonite was more soluble than

zinc oxide which in turn was more soluble than

hemimorphite. In the latter case it is possible that

equilibrium was not obtained because the dissolution

was very slow. With smithsonite the zinc concentration

in solution was dependent on time (Fig. 3.4. ) and in 1. OM

NaOH the concentration increased to a maximum of

2.9g14 and then decreased to 1.4g14. This latter

value is approximately the solubility of zinc oxide in

1. OM sodium hydroxide. A white precipitate was

observed in the smithsonite residue and X-ray diffraction

analysis indicated that significant amounts of hydrozincite

were present. -95-

ýý +. a c

V 0 h

20 30 40 50 60 70 80 temperature (° C)

Fig. 3.2. Effect of temperature on the solubility of the secondary zinc minerals. -96-

A -s

tr) c N

1ý 0 h

0 0.5 1.0 1.5 2.0 sodium hydroxide concn. (M )

Fig. 3.3. Solubility of the secondary zinc minerals in a solvent of sodium cyanide and sodium hydroxide. -97-

C -{1 U C O V v C ýN

0 10 20 30 40 50 Time (h)

Fig. 3.4. The dissolution of smithsonite in 1. OM sodium hydroxide in the presence of excess mineral. -98-

3.3 Discussion

For solution equilibria considerations zinc

oxide can be treated like zinc hydroxide because, on

immersion in aqueous solution, the surface will become

hydrated and resemble it.

OH ZnO ...... H\> Zný (3.3) OH OH

On dissolution in cyanide solution, therefore,

hydroxyl ions will be produced and soluble zinc

hydroxy complexes -will be formed in addition to zinc

cyanide complexes. The validity of this reasoning

is shown in Fig. 3.5. where the solubility curve of

zinc oxide is compared to that calculated for zinc

hydroxide. The latter is represented by a band

indicating the probable maximum and minimum

limits of solubility of zinc hydroxide based on the

spread of the equilibrium constants quoted in the

literature. The solubility curve for zinc oxide falls

well within the theoretical solubility limits for

zinc hydroxide.

The composition of the solutions at equilibrium

was presented in Fig. 1.4. where it was seen that at

cyanide concentrations greater than 10-1M there was

appreciable zinc hydroxy complex formation. The high -99-

hydroxyl concentration of the solutions ensures that loss of cyanide by hydrolysis is negligible.

Both samples of smithsonite exhibited solubility curves that corresponded very closely to that predicted from solution equilibria theory. The theoretical curve

(Fig. 3.5. ) is again shown as a band representing the maximum and minimum solubility limits. Theory shows (Fig. 1.5) that dissolution of the smithsonite in cyanide solutions alone does not lead to appreciable zinc hydroxy complex formation as hydroxyl ions are only produced by hydrolysis of the carbonate and cyanide ions. At saturation the predominant species is the

Zn(CN)4 complex at all cyanide concentrations studied, and the experimentally determined cyanide to zinc molar

ratio of approximately 4 to 1 supports this.

One of the reasons for the lower solubility of

smithsonite than that of zinc oxide is probably that in

the former case very little zinc hydroxy complexation

took place. The production of hydroxyl by hydrolysis

of the carbonate ion is sufficient to maintain the

solution at pH 11.5 or greater where cyanide loss can

be considered negligible, but insufficient to contribute

to the dissolution reactions. -100-

-ý

al c N

012345 cyanide conch. (M )

Fib. 3.5. Comparison between experimental and theoretical solubility of zinc oxide and smithsonite in sodium cyanide. -101-

The unreacted smithsonite was not altered to any detectable extent by hydrolysis of the mineral surface. These observations support solution equilibria calculations which showed that in cyanide concentrations greater than 0.5M, in the absence of added sodium hydroxide, precipitation of zinc hydroxide at the saturation point would not occur.

The composition of the chemically prepared zinc carbonate is given as ZnCO3.2ZnO. 3H20 which is similar to the accepted formula for hydrozincite,

2ZnCO3.3Zn(OH)2. The almost coincident solubility curves obtained are, therefore, not surprising. During the dissolution the hydrated zinc groups would lead to the presence of additional hydroxyl in solution with the consequential formation of zinc hydroxy complexes.

A solubility greater than that of pure zinc carbonate

(smithsonite) but less than that of pure zinc hydroxide

(or zinc oxide) would, therefore, be expected. The results show that this behaviour was in fact obtained.

A theoretical solubility curve for hemimorphite in cyanide solutions has not been presented owing to inadequate knowledge of the silica species in solution, the absence of equilibrium data and doubts as to whether equilibrium would ever be attained. The structure of -102-

hemimorphite Zn4Si2O7(OH)2H20, includes hydroxyl groups which would be expected to form zinc hydroxy complexes and also to be utilised in the formation of monomeric silicic acid, silicate ions and possibly polymeric silica in solution. The 2 to 1 zinc to silicon molar ratio obtained in solution indicates that hemimorphite dissolved stoichiometrically and the absence of willemite or crystalline silica forms in the residue suggest that a silica layer was not formed on the hemimorphite surface. It is, however, possible that an extremely thin silica rich film was present below the level of detection when the solubility limit of silica was approached.

The solubility of amorphous silica has been

reported by Alexander 0.012% SiO (73) as 2 at pH

values below 8 with the silica species existing in

solution as the momomeric silicic acid.

Si02 2H20 Si(OH)4 (3.4) + .-'r..

Alexander has shown, however, that at higher pH

values the solubility of silica increases due to the

formation of silicate ion in addition to silicic acid. -103-

Si(OH)4 + OH (HO)3SiO + H2O (3.5) {H]_1HO)3SiOj 8 where = 10-9' (3.6) {si(oHJ

and if St = total solubility of silica (SiO2)

(3.7) pH- 9.8 = log St - 0.012 0.012

A 1. OM cyanide solution has a pH 11.5 which,

assuming that the p11 does not change markedly during

the dissolution of the hemimorphite, would permit a

total silica solubility of 0.287% (Si). Analysis of the

solution obtained gave a silica concentration of t 0.293 0.003% w/v (Si). It is very likely, therefore,

that when the solution became saturated with silica

there was precipitation of silica from solution, probably

as a gel on the mineral surface.

Changes in solubility of minerals with variation

in temperature are to be expected because of the

temperature dependence of the equilibrium constants.

This dependence is given by the Van't Hoff isochore

which in its integrated form is

log I- log K298 = LI_H_ 1_1 (3.8) 2.30M 298 T

If the change in heat capacity is small, as has

been reported (74) for most complex reactions, and can -104-

be regarded as zero then, &H in Equation (3.8) can be replaced byQH298.. Thermochemical data is unavailable for many of the equilibria of concern but the enthalpy changes for some of the more important equilibria can be from data calculated published (75,22) 2+ Zn(OH)2 Zn + 20H (3.9) (s)

D 110 29.6 kJ 298 = mole-1 2+ ZnCO3 Zn + CO3 (3.10) (S) V_

H298 = 15.9 kJ mole-1 2+ Zn +4 CN-'-- Zn(CN)ý (3.11) L1H298 = -116.3 kJ mole-1

An increase in temperature results in a small increase in the solubility product of the minerals, and a decrease in the zinc cyanide stability constants. At temperatures up to 80°C. these changes are not large compared to the discrepancy in the published values for the equilibrium constants at 25°C.

The small increase in solubility of zinc oxide, hydrozincite and smithsonite with increase in temperature indicates that the overall effect of changes in the magnitude of the ejailibrium constants was slight. The solubility of hemimorphite increased, with an increase in temperature, more than the solubility of the other

zinc oxide compounds. This can be explained either, by -105-

a, greater change in the magnitude of the equilibrium constants or, more probably, by the hemimorphite dissolution 250 C. At not attaining equilibrium at . higher temperatures the dissolution would be faster and it is probable that a closer estimate of the solubility would be obtained.

The solubility of zinc oxide in sodium hydroxide solutions was found to be much lower than the solubility calculated from the solution equilibria considering the dissolution

(Fig. 3.6. ) Owing to the wide range in reported(22) values for the stability constants of the zinc hydroxy complexes, the theoretical solubility is shown as a band representing the maximum and minimum solubility limits. Although solubility determinations were only

conducted in low caustic concentrations, close agreement

by Merrill Lang was obtained with results reported and (34)

The solubility of zinc oxide in cyanide solutions

was in good agreement with those values calculated from

solution equilibria considerations. If the same reaction

sites are involved in the dissolution in sodium hydroxide

then one would not expect such large discrepancies

between the experimental and theoretical solubilities.

Zinc oxide and zinc hydroxides exist in a number -106-

-ý

Z-Vt c

0 0.5 1.0 1.5 2.0 2.5 sodium hydroxide conch. (M )

Fig. 3.6. Comparison of the solubility of zinc oxide in sodium hydroxide with published and theoretical values. -107-

of polymorphous modifications where the solubility product has been found(23) to be different for the

various modifications. Values for the solubility

products of these different forms are reproduced in

Table 3.4 where the value for zinc oxide can be seen

to be rather lower than that for amorphous zinc

hydroxide.

Table 3.4. Solubility products of zinc oxide and

zinc hydroxide modifications

(Modification 1 log K (25° C) sO Amorphous Zn(OH) 52 2 -15. Zn(OH) 24 2 -16. Zn(OH) -16. 26 Zn(OFI) -16. 15 Zn(OH)2 -16. 47 L ZnO -1 6. 83

Copper hydroxides have been found and nickel (76)

to exist as active or inactive forms, the active form

being considered as having a disordered lattice. This

form shows the maximum solubility due to a metastable

equilibrium existing between the active solid and the

solution, but with time it reverts to the inactive form.

Similar behaviour might be expected from the zinc

hydroxides and oxides.

These factors would explain the great difference -108-

in solubility between zinc oxide and zinc hydroxide but would not account for the solubility of zinc oxide being much lower than theory predicts. The rate of dissolution of zinc oxide in sodium hydroxide was slow and it is possible that either, only an apparent solubility, or the solubility of an 'inactive' zinc oxide was determined.

X-ray diffraction analysis on the smithsonite residue after dissolution in sodium hydroxide indicated the presence of hydrozincite. This can be considered as resulting from the following reaction on the surface of the mineral

ZnCO3(s) + 20H + C032 (3.12) --Zn(OH) 2(s) where an equilibrium constant can be written

K= Kso(ZnCO3)/KS 105 (3.13) 0(Zn(OH)2) =8x

This reaction leads to a decrease in hydroxyl concentration in solution and hence the zinc concentration in solution would be reduced after the maximum solubility had been obtained. This would account for the apparent time dependence of the smithsonite solubility. A similar decrease in zinc concentration after the maximum solubility of smithsonite in sodium hydroxide was also noted by Merrill and Lang(34). They showed that -109-

conversion of sodium hydroxide to sodium carbonate in the presence of excess mineral, resulted in a decrease in zinc concentration until eventually the solubility corresponded to a stable concentration represented by the zinc oxide solubility.

When the cyanide leach solutions were prepared with up to 2. OM sodium hydroxide a much greater increase in the solubility of zinc oxide and hemimorphite was achieved than expected from the limited solubility of the minerals in low caustic concentrations. The solubility of zinc oxide and smithsonite in solutions containing both cyanide and sodium hydroxide is compared in Fig. 3.7 with the theoretical solubility from solution equilibria calculations. The theoretical solubility is again presented as a band indicating the maximum and minimum solubility limits. The increase in solubility of zinc oxide can be attributed to the formation of additional zinc hydroxy complexes. The addition of sodium hydroxide alters the equilibrium of the solution and results in an increase in the free cyanide concentration which is then available to dissolve further zinc oxide. Solution equilibria calculations have shown this to be the case but again, the influence of sodium hydroxide addition is less than -110-

predicted.

Solution equilibria calculations suggest that the solubility of smithsonite should also be increased by the addition of sodium hydroxide to the cyanide solution. However, as can be seen in Fig. 3.7, this did not happen. The calculations also indicated that zinc hydroxide would not be precipitated from the solution but the presence of hydrozincite on the surface of the residual smithsonite was detected. That an increase in solubility was not achieved is not surprising because, in the presence of excess smithsonite, the reaction

given by Equation (3.12) can be expected to occur. The formation of zinc hydroxy complexes would take place

only to a very limited extent and the predominant zinc

species in solution would still be the Zn(CN)42-

complex.

The increased solubility of hemimorphite that

resulted from the addition of sodium hydroxide to the

cyanide solutions can be explained by the increased

solubility of silica at high pH. If the solubility of

hemimorphite is controlled by the solubility of silica

then the addition of sodium hydroxide v,o uld allow a

higher silica concentration in solution and consequently

a higher zinc concentration due to the formation of

additional zinc hydroxy complexes. -111-

i -S / / 0 / 01 / / i H / / / / ..ý / / -C /

0 / i h / /

o ex tl. 101 zinc oxide - theoretical

smithsonite o exptl. ---theoretical

0 0.5 1.0 1.5 2.0 2.5 sodium hydroxide addition (M)

3.7. Comparison of the solubility of zi: ic oxide and sniitlzsonite in a 1. OM cyanide solution containing added sodium hydroxide, with theoretical values. -112-

4. KINETIC STUDIES

4.1. Introduction

To obtain a proper understanding of a chemical

process it is not only necessary to find the extent of

the reaction but also the rate at which the reaction

proceeds. In the heterogeneous system, where a

solid dissolves after reaction with a solute in solution,

the reaction can be considered(77) as progressing in

a number of steps.

(i) Mass transfer of reactants to the solid-liquid

interface.

(ii) Adsorption of reactants on the surface.

(iii) Reaction at the surface.

(iv) Desorption of soluble reaction products.

(v) Mass transfer of soluble products away from

the solid-liquid interface.

Steps (ii), (iii) and (iv) are chemically controlled

processes whereas (i) and (v) are controlled by mass

transfer. Any of these might be the slowest step and

the rate of reaction might, therefore, be controlled

by chemical, mass transfer, or some intermediate

processes.

A chemically controlled reaction between a

solid and a reactant in solution can be regarded as a -113-

bimolecular collision process where the 'concentration' of the solid is represented by the surface area available

for reaction. The rate of the reaction is given by

- dCi kA Cn (4.1) dt -ci V

where c. is the concentration of the reactant at the 1 solid-liquid interface, A is the surface area of the

solid. V is the volume of solution, kc is the observed

reaction velocity constant per unit area at unit volume

and n is the order of the reaction with respect to the

reactant concentration.

In the case of amass transfer, however,

the reaction can be represented by the Nernst expression,

- dC = DA (C Ci) (4.2) dt - vb -

which describes the rate of the reaction if it were

dependent on the diffusion of reactant across a

stationary liquid layer of thickness (a) at the solid-

liquid interface when the concentration of reactant at

the surface is C. and that in bulk C. 1 solution

The Nernst expression indicates that the rate

of reaction should increase with a decrease in boundary

layer thickness. However, there is considerable

evidence(78) to suggest that a stationary layer does -114-

not exist at the solid-liquid interface and that fluid

flow continues right down to the solid surface. This

flow can be considered laminar with no fluid motion

perpendicular to the interface and mass transfer across

this region is supposed to take place by molecular

diffusion. Resistance to diffusion has been shown(79)

to be largely due to an eddy zone in the turbulent

stream. An effective film thickness can, therefore,

still be considered as providing a resistance to mass

transfer. The thickness of this 'boundary layer' will

be influenced by the efficiency of agitation which depends

on the rate of stirring and the dimensions and geometry

of the stirred system (80).

The hydrodynamics of an agitated system is

complex but a complete solution has been derived by

Levich(81) for convective diffusion to the surface of

a rotating disc under non turbulent conditions and when

the reaction is entirely controlled by mass transfer.

Application of the rotating disc to studies of dissolution

rates has received widespread attention for metal

dissolution in many solvents including cyanide (82,83)

A few investigations into the dissolution of minerals

in cyanide solutions have also been carried out(84) -115-

and practical details described(85)"

For Levich's hydrodynamic solution to be valid, the disc must have a diameter large enough for edge effects to be neglected so that the boundary layer can be assumed to be of uniform thickness over the whole area of the disc. The volume of solution must also be very large so that wall effects on the fluid do not influence the rate of mass transfer to the disc. Application of these conditions to the study of mineral dissolutions does, however, produce a number of difficulties. It is practically impossible to prepare large homogeneous discs of most minerals unless materials such as massive sulphides are being studied. Many minerals exhibit anisotropic leaching and hence a change in the surface area of the disc can be expected. Moreover, the active surface area of the mineral available for reaction is not necessarily equivalent to the area of the disc because of the presence of cracks and pores. Some of these difficulties have been overcome(85) by fashioning discs from pressed mineral particles but the rotating disc method still appears unsuitable to a study of the dissolution rates of the secondary zinc minerals because of the following reasons- -116-

(1) The hydrozincite was extremely friable.

(2) The hemimorphite crystals were platy prisms too small to fashion into a disc. The use of discs of pressed particles would not be convenient because of

difficulties inherent in ensuring random orientation

of the crystal faces.

(3) The smithsonite consisted of polycrystalline

aggregates which would probably lead to preferential

leaching along grain boundaries and material falling

from the disc into the solution.

(4) The anisotropic character of hemimorphite and

smithsonite could lead to a changing surface area during

leaching,

(5) Most importantly the surface area of the minerals

would not be equivalent to the area of the disc because

of the surface morphology and the presence of cracks

and pores. This can be seen from Table 4.1 where

the specific surface areas for the minerals are

compared to the calculated surface areas assuming

flat sided cubes. -117-

Table 4.1. A comparison between actual and theoretical specific surface areas of the minerals

Mineral Mean Size(1) Specific Surface area fraction (Nm) ------(ac uaIl) (calculgted) mg mg

I Smithsonite 1 75pm 120 0.166 0.01 14 I -300 + 1 1 -53 + 37pm 44 0.245 0.031

Hemimorphite 0.1 05 0.014 -300 + 75pn 120 0.031 -53 + 371pm 44 0.228

Hydrozincite 120 8.78 0.0114 L-300+75pm

(1) Calculated from the mean of inverse size as the specific surface area is inversely proportional to particle size.

(2) From BET krypton adsorption measurements.

The covering area of a krypton molecule has

been reported($O)to be from 0.17 to 0.22nm2 whereas

the cyanide ion has a diameter of 0.40nm and hence a

cross-sectional area of 0.13nm2. It is to be expected,

therefore, that the surface area available for reaction

would be not much different from the gas adsorption

area.

For the reasons outlined above the dissolution

rate of the secondary zinc minerals in cyanide solutions

was studied as a system of suspended particles in an -118-

agitated vessel. In this system, provided that agitation

is sufficient to ensure adequate suspension of the

material, the total surface area of the particles is

available for reaction and there is no preferred

orientation of the crystal faces. The variables studied

in the dissolution of the secondary zinc minerals were

stirrer speed, mineral particle surface area, reactant

concentration and temperature. No attempt was made

to vary the agitator size and geometry and these were

maintained constant because of their importance in mass

transfer controlled processes.

4.2. A itation system

The experimental apparatus was as described in

the determination of the solubility of the minerals. The

stirrer speed was controlled to within 8 r. p. m. with

a 'Berco' controller and the temperature to within ± 0.1°K by immersing the reaction vessel in a constant

temperature water bath. The reaction was started by

quickly inserting the mineral powder into the agitated

solution through an opening in the lid of the reaction

vessel.

A preliminary study was undertaken to ensure

that agitation was sufficient for complete suspension -119-

of the mineral particles. Impellors of varying sizes and shapes were investigated and it was found that the use of propellor type stirrers resulted in a small cone of material residing directly beneath the impellor even at a stirrer speed as high as 1,200 r. p. m. Although material was continually being swept up into the solution

it was thought unlikely that all the mineral surface was available for reaction. To prevent the formation of a

cone of material it was necessary to use a curved impellor

rotating almost flush to the bottom of the vessel. Surface

area determinations on -300 + 75pm fractions of smithsonite

and hemimorphite that had been agitated for 24 hours at

1,000 r. p. m. with this system showed that it did not

result in appreciable attrition of the minerals.

Baffles were not used in the reaction flasks

because it has been reported(70) that at agitation speeds

necessary for complete particle suspension they alter the

dissolution coefficient to only a small degree. They are

however, used in many cases because they tend to promote

particle fracture and permit a more direct input of

energy from an agitator to the pulp. In this study of

the dissolution rates of the secondary zinc minerals in

cyanide, it was decided that artificial particle fracture

was not required. -120-

4.3. Experimental technique and data analysis

It was the objective of the kinetic studies to not

only determine those variables that influenced the rate

of dissolution of the secondary zinc minerals but also

to quantify the various effects in the form of a reaction

velocity constant. From these observations and

calculations it was hoped to provide information concern-

ing the overall mechanism of dissolution and to derive

an overall mechanism of dissolution and to derive an

overall rate equation.

The course of the reactions was followed by taking

samples at intervals, in the manner previously described,

and analysing the solutions for product concentration (zinc

in solution). Although it was considered desirable to

follow the changes in the free cyanide concentration

directly this was not possible because no analytical

method was available. All free cyanide concentrations

were, therefore, calculated from the determinations of

the zinc concentration in solution and knowledge of the

solution equilibria of the system.

The half life method is frequently used to find

the order of a reaction. However, heterogeneous

reactions occur at an interface and are usually complex -121-

proceeding by stages, each step being an elementary

reaction. For complex or reversible reactions or

when more than one reactant is involved which are

present in unequal concentrations, then the concept of

a half life is less useful(87). It is possible to establish

the reaction order and mechanism by the initial rate

method i. e. by measuring the rate at t=0 for various

concentrations of each reactant. This method, although

powerful, is often limited by the experimental difficulties

in determining the initial reaction rates. From the

experimental points of a concentration-time curve the

initial reaction rate can be found by taking the tangent

to the curve at t=0 but this is not good practice as errors

caused by bias may be large. The fitting of a non-linear model(88) provides a useful means of fitting an equation to the data points. The general equation can be expressed as a polynomial of the form

y=A+Bx+Cx2+... +Dxn (4.3) and the power of the polynomial is taken as the lowest that will provide a good fit to the data, goodness of fit being judged by analysis of variance. Polynomial curve fitting has, therefore, been used to provide an unbiased method that also helps in eliminating random experimental errors by smoothing out the curve. A polynomial was -122-

fitted to the data provided that a minimum of eight

points were present on the initial smooth part of the

curve and the differential of the polynomial at t=0

used to give the initial reaction rate. However, when

the rate of reaction was so rapid as to prevent the

taking of sufficient samples for curve fitting, an

approximation to the initial reaction rate was made by

assuming the first part of the curve to be linear.

The polynomials were fitted to the experimental

data by multiple linear regression using a computer

program and it found that the data (89) was was usually

fitted adequately by a second or third order polynomial.

A typical regression curve has been fitted to the

experimental data points in Fig. 4.1. Throughout the

kinetic studies the initial reaction rates are expressed

in terms of zinc extracted into solution per unit volume

per unit time (mole Zn 1-1 s-1).

4.4. The influence of agitation rate on the rate of dissolution

of the secondary zinc minerals

All the rate tests were conducted under the

conditions summarised in Table 4.2. -123-

polynomial by regression 3 io y =0.08+ 3.421 -0.5612+0.049t 0

)ý O

v 31c

1' Q)

c 25 U c 0 U U C ýN -- curve fitting

o exptl. points

02345 Time(h)

1'iß. 4.1. Polynomial fitting by multiple linear regression to data points on the initial portion of the curve of smithsonite dissolution in 1. UM cyanide. -124-

Table 4.2.

Cyanide concentration 1. OOM Cyanide/zinc molar ratio 4/1 Particle size -300 + 75pm Temperature 298°K

The rate curves obtained for the dissolution of zinc oxide powder (Fig. 4.2) show that the dissolution was so rapid as to make the influence of agitation rate on the initial reaction rate impossible to determine.

The reaction was complete with 100% zinc extraction within 10 minutes even at the slowest stirrer speed.

An increase in the agitation rate increased the rate of dissolution of hydrozincite (Fig. 4.2) but the dissolution was too rapid for the effect to be determined quantitatively. The mineral was not completely suspended at stirrer speeds less than 200 r. p. m. and the much lower rate of dissolution at slower speeds can be attributed to inadequate presentation of the hydrozincite surface to the solvent. Complete dissolution of the mineral was obtained within 15 minutes at an agitation rate of 400 r. p. m.

The observed rate of dissolution of smithsonite increased with an increase in the agitation rate and this effect is shown in Fig. 4.3. Complete suspension of the -125-

mineral was obtained at stirrer speeds greater than

240 r. p. m. A sharp increase in zinc concentration was noticed after about 6 hours leaching, but after about 10 hours the rate of dissolution decreased so that even at the

faster agitation rates (greater than 510 r. p. m. ) the

smithsonite was only 92.5% dissolved after 24 hours.

Microscopic examination of the residue after this time

showed the presence of some large particles of hemi-

morphite that had been only partially leached but no

large particles of smithsonite remained.

The rate of dissolution of hemimorphite increased

with an increase in agitation rate (Fig. 4.4) but after

about 10 hours the rate of zinc extraction decreased

markedly. The shape of the dissolution curves can be

best described as almost parabolic. The rate of

dissolution of hemimorphite was very much slower than

that of the other zinc minerals and only 20% zinc extraction

was achieved after 24 hours.

The initial reaction rates per unit surface area

for all the minerals are plotted against the agitation rate in

Fig. 4.5. The specific surface areas of the minerals

were found by B. E. T. krypton adsorption and the values

obtained are presented in Table 4.5. In all cases the

initial dissolution rate increased markedly with an -126-

increase in the stirrer speed until the material was completely suspended at speeds between 200 and 250 r. p. m. Further increases in the agitation rate led to only slightly increased dissolution rates and eventually a maximum was reached. This would indicate that when the minerals were not completely suspended in the solution, all the active surface area of the particles was not available for reaction. Nienow(90) has reported that the rate at which themasstransfer coefficient increases with

stirrer speed falls markedly after the minerals become

fully suspended and that no further increase occurs when

aeration of the solution begins. The fluid flow at high

agitation rates appeared very turbulent but it was difficult

to judge the onset of aeration. The maximum dissolution

rate, however, evidently corresponded to a limiting mass

transfer coefficient for the agitation system. The

maximum dissolution rates and the stirrer speed at which

the limitingmass transfer coefficient was reached are

presented in Table 4.3. -127-

ezl% 01

u 0 O U C hydrozincite (Z 641 v U g) C ýN v o 70 RP.M. C 200 - N 4 300 of V 400 is 75

zinc oxide (2-035g)

o 80 R.PM. 2!

012345 Time(h)

Fig. 4.2. The influence of agitation speed on the dissolution of zinc oxide and hvdrozincite in cyanide. -123-

i Oº

C U C O U U C N

05 10 15 20 Time (h)

Fig,,. 4.3. The influence of agitation rate on the dissolution of smithsonite 1 in 1. OM1 cyanide. -129-

V 3- A th

V C *ZZ0 O V c v k a, N 2- U P4c

125 R PM 450 )z 510 0 600 v 720 o 850 Io 1000 ".

0 10 20 30 40 Time (h ) Fig. 4.4. The influence of agitation rate on he dissolul "on of hemimorphite in 1. OM cyanide. -130-

smithsoni te (3-260g)

N E.

c N 2 (U 0 E

0 4- U hydrozincite (2.641 t) 9) - -it= -- -

"ti / hemimorphite (3"777g ) a 4

op OF

0 500 1000 Agitation rate (R P.M. )

Fig. 4.5. The influence of agitation rate on the initial reaction rate of the secondary zinc minerals in 1. OM cyanide. (Mineral weights in brackets). -131-

Table 4.3. Maximum dissolution rates of -300 + 75 m mineral fractions in 1. OOM sodium cyanide

I Limiting Maximum initial I Mineral stirrer dissolution 'atf speed (mole Zn 1sm) -2 krp

Smithsonite 510 2.70 x 10 5 Hydrozi ncite(1) (400) (1.12 x 10- ) i 5 Hemimorphite 600 0.4 0x 10-

(1)approximate figures only

The dissolution rate per unit area of smithsonite

was much greater than that of the other minerals,

hemimophite being the slowest to dissolve. Although

the observed dissolution rate of hydrozincite was very

much faster than that of the other minerals this can be

attributed to the high porosity of the mineral particles

giving rise to a high surface area for reaction with the

solvent. Complete extraction of the zinc from

smithsonite took a long time and this is probably due to

the slow rate of dissolution of the small amounts of

hemimorphite present.

It is well known that the Reynolds number for

an agitated system is defined as: - -132-

Re =N d2 (4.4)

where N is the stirrer speed (rev sec-1), is the

liquid density, d is the diameter of the stirrer and

is the liquid viscosity.

The fluid flow is turbulent for Reynolds numbers

than 1,000 500 greater and at an agitation rate of r. p. m. ,

Re is 24,250 and therefore conditions within the reactor

were turbulent. All other parameters that influence the

dissolution rate of the minerals were hereafter studied

in the flow regime where increased agitation did not

increase the reaction rate.

Further investigations into the dissolution

kinetics were confined to smithsonite and hemimorphite

because the rate of dissolution of zinc oxide and

hydrozincite were too fast for accurate measurement.

4.5. The influence of particle size on the rate of dissolution

of smithsonite and hemimorphite in cyanide solutions

The dissolutions were carried out using the

conditions summarised in Table 4.4. -133-

Table 4.4.

I Smithsonite Hemimorphite

Cyanide concentratio n 1. OOM 1. OOM Cyanide to zinc molar ratio 4/1 4/1 Temperature 298°K 298°K Mineral weight 3.280g 3.777g- 4.035g Agitation rate 540 rpm 620 rpm

Various screened and sized fractions of the minerals were prepared. The rate curves for smithsonite (Fig. 4.6) show that the rate of dissolution increased with a decrease in particle size and that there was a sharp increase in zinc concentration for the size fractions coarser than

53 m after the reaction had proceeded to some extent.

The smaller the particle size, the shorter was the time

before this sharp increase occurred. The dissolution

rate of the fractions finer than 53 m was very rapid and no sudden increase in zinc concentration could be observed.

It was also noticed that the finer the particle size

the closer was the approach to maximum extraction.

This was undoubtedly due to the reduced particle size

of the hemimorphite in the smithsonite sample reacting

at a faster rate. It is also possible that some smithsonite

was present as inclusions in the hemimorphite phase

and that size reduction led to liberation of part *of this -134-

smithsonite.

The hemimorphite rate curves (Fig. 4.7) show that the rate of dissolution was increased by a reduction in particle size. After about 10 hours leaching the dissolution rate markedly decreased and the rate of zinc extraction from the hemimorphite showed almost parabolic leaching kinetics.

The specific surface areas of the mineral fractions were determined by B. E. T. krypton adsorption and are presented in Table 4.5.

Table 4.5 Specific surface area measurements of the mineral size fractions

2a- Si ze Specifi c surface area m fraction Smithsonite T Hemimorphite Hydrozincite pm i II

+ 75 0.166 0.115 0.105 8.78 -300 1 10.074

-300 + 21O 0 130 0.0675 ` -210 + 150 0.152 0.0809 0.102

-150 + 105 0.212 0.108 0.111 10.133 -105 + 75 0.250 0.124 1 -75 + 53 0.45210.284 0.154 I -53 + 37 0.8801 0.228 I

-53 1.16 0.848 0.893 -135-

v ^'ýi C ý,. Oý N

c Si 0 u U C ýN 0 -300 + 75 Nm -150 +105 pm o- -105 . 75 Nm 75 + 53 pm 53 + 37 pm 53 pm

05 10 15 20 Time (h

Fig. 4.6. The influence of particle size on the rate of dissolution of sinithsonite t in 1. OTIJ cyanide. -136-

V 9 O V V

D l0 20 30 40 Time(h)

Fig. 4.7. The influence of particle size on the rate of dissolution of hemimorphite in 1. OM cyanide. I! -137-

From heterogeneous reaction theory, the rate of mineral dissolution should be proportional to the surface area available for reaction. A plot of log rate against log surface area should, therefore, have a slope of unity if the theory is valid. The initial reaction rates for the various size fractions were found by regression analysis to be linearly correlated with the surface area of the minerals (Fig. 4.8).

For smithsonite the correlation coefficient was

0.98 which is highly significant at the 95% confidence level and the regression equation was found to be: -

d[Zn] = 2.97 x 10-5 A1.06 mole 1-1s-1 (4.5) ät~

and the 95% confidence interval for the exponent was

0.94

The correlation coefficient for hemimorphite was

1.00 which is highly significant at the 95% confidence

level and the regression equation was found to be: -

d[in] = 3.70 x 10-6 A0.91 mole 1-1s-1 (4.6) dt

and the 95% confidence interval for the exponent was

0.76

As an exponent of unity lies within the confidence

intervals for both minerals, the data can be forced to

fit linear a equation by regression and is given by: - III -138-

_y o smithsonite 0 5A1"06 R=2.97x10 0 Q,

0 O 0 V O L "S 0

o hemimor phit e

R^3.70x10-640.91

-0.4 -0.2 0 0.2 0.4 log (surface area A)

Fig. 4.8. Correlation between initial reaction rate and surface area for the dissolution of smithsonite 1 and heniimorphite in cyanide. -139-

for smithsonite: -

d[Zn] = 9.7 x 10-8 + 2.85 x 10-5 A mole 1-1s- (4.7) dt

and neglecting the intercept- [Zn] dd = 2.85 x 10-5 A mole 1-1s-1 (4.8) ät

and the 95% confidence interval for the constant term was

2.77 x 10-5< k<2.93 x 10-5

and for hemimorphite: -

d[ n]_ = 2.60 x 10-8 + 3.33 x 10-6 A mole 1-1s-1 (4.9) dt

and again neglecting the intercept: - [Zn] 6A l d = 3.33 x 10- mole 1- s-1 (4.10) dt

and the 95% confidence interval for the constant term was

2.94 x 10-6< k <3.72 x 10-6 A mole 1-1s-1

The dissolution of smithsonite and hemimorphite

in cyanide solutions follows what would be expected from

heterogeneous reaction theory.

4. G. The influence of cyanide concentration on the rate of

dissolution of the secondary zinc minerals

The dissolutions were conducted using the

conditions summarised in Table 4.6. -140-

Table 4.6.

FI Smithsonitel Hemimorphite Hydrozincite

Cyanide concn. 1. OOM 1. OOM 1. OOM Cyanide to zinc 4/1 4/1 4/10 molar ratio 0 Temperature 2980K 298 K 298 K Particle size -300 + 75pnn -53pn -300 + 75/.m 3.280g 4.035g 2.841g2 Mineral weight 2 2 Surface area 0.5444m 3.373m 24.9m Agitation rate 540 rpm 620 rpm 400 rpm

The rate of dissolution of hydrozincite in 1. OOM cyanide

was too fast to follow and the reactions were therefore

carried out at reduced cyanide concentrations. However,

at all cyanide concentrations down to 0.25M, the dissolution

reaction was completed within 15 minutes (Fig. 4.9). The

final concentration of zinc in solution corresponded to the

solubility of hydrozincite in the particular cyanide solution.

The smithsonite rate curves (Fig. 4.10) show that

the rate of dissolution increased with an increase in cyanide

concentration. The sudden rise in zinc concentration

previously described was again evident at cyanide

concentrations between 0.75M and 1.50M after about

5 to 6 hours of leaching. This sharp increase was not

observed at higher or lower cyanide concentrations. The

dissolution of smithsonite in 0.50M and 0.75M cyanide

solutions continued until a cyanide to zinc molar ratio -141-

in solution of 4.0 to 1 was obtained after about 24 hours.

This corresponded to the maximum solubility of smithsonite in cyanide solutions of that strength. At higher cyanide concentrations dissolution of the mineral continued to completion although a long time was required to extract the last 5% of the zinc from the sample. The presence of slowly dissolving hemimorphite would again account for this behaviour.

The rate curves showed that the sample of smithsonite 2 dissolved at a much faster rate than the smithsonite 1. A sharp increase in zinc concentration was also observed after smithsonite 2 had reacted to some extent and the general shape of the curves was similar for both samples. Previous microscopic examination had shown both samples to have a similar polycrystalline crystal habit. The specific surface area of the smithsonite 2 was rather less than that of the smithsonite 1 (Table 4.5) and hence the greater rate of reaction of smithsonite 2 can not be due to a larger surface area (Table 4.7). -142-

C U C O U U C "1 N

012 Time(h)

Fig. 4.9. The influence of cyanide to zinc molar ratio on the rate of dissolution of hvdrozincite. -143-

ti) 11-- Uc c 0 U u

05 10 15 20 Time (h )

4.10. The influence of cyanide to zinc molar ratio on the rate of dissolution of smithsonite 1 and smitlisonite 2. -144-

Table 4.7 A comparison of the reaction velocities for two sources of smithsonite in 1. OM cyanide

Surface Initial reaction rate per ar unit arg a12 (m ) (mole Zn 1sm) 5 I Smithsonite 1 0.5444 2.66 x 10_ Smithsonite 2 0.3975 6.21 x 10

An increase in the cyanide concentration resulted

in an increased rate of dissolution of hemimorphite

(Fig. 4.11). After a period of time the dissolution rate

decreased markedly and the dissolution curve followed

almost parabolic kinetics. The point at which the

dissolution rate was greatly decreased occurred after

a longer leaching time at increasing cyanide concentrations.

Very long contact times were required before the reaction

halted i. e. about 80 hours in 1. OOM cyanide when the

cyanide to zinc molar ratio in solution was 4.31 to 1 but

this was reduced to about 30 hours in 3.00M cyanide.

Even at the higher cyanide concentration the hemimorphite

was not completely dissolved.

The initial reaction rate can be expected to follow

the equation: -

1. Cn (4.11) A dt _=k -145-

00 0 i 0

^ý 0

c 1 v c T N r

C 50 V C

V C N 0 5ý

cyanide/zinc molar ratio

0 2/1 It 4/1 A ß/1 g 1211

0 10 20 30 40 Time(h)

4.11. The influence of cyanide to zinc molar ratio on the dissolution rate of hemimorphiLe. -146-

and a logarithmic plot of initial reaction rate against cyanide concentration should yield a straight line of slope n, the order of the reaction with respect to the cyanide concentrations. Fig 4.12 shows that a straight line was fitted to data by linear regression and analysis of variance for regression(91) indicated that a good fit was obtained.

The regression equations were found to be for

smithsonite: - 05 1 d[Zn ]=2.77 x 10-5 CO' (4.12) Ä dt

and the 95% confidence interval for the exponent was

0.93

1 d[Zn] = 3.90x10-6C0.91 (4.13) A dt

and the 95% confidence interval for the exponent was

0.61

The exponent of the cyanide concentration term

was very close to unity for the dissolution of both

smithsonite and hemimorphite showing that the reactions

can be considered first order with respect to the cyanide

concentration. The experimental data can be forced to

fit a first order rate expression and the following

equations were, therefore, found by linear regression: - -147-

o smithsonite

R=2.77 x 10 -SCO-95

-ý

a, ... I-

O 4. -5-

O a

o hemimorphite Q1 O R=3,90x10-600.91

-6

-0,5 0 0.5 log (cyanide concn. C)

Fig. 4.12. The relationship between the initial reaction rate and the cyanide concentration (by linear regression). -148-

for smithsonite: -

1_ d[Zn] = 2.58 x 10-5 C (4.14) A dt

and. the 95% confidence interval for the constant term

was 2.49 x 10-5

and for hemimorphite: -

1d [Zn] = 3.63 x 10-6 C (4.15) A dt

and the 95% confidence interval for the constant term was

3.11 x106

4.7. The influence of temperature on the rate of dissolution

of smithsonite and hemimorphite

The experiments were carried out using the

conditions summarised in Table 4.8.

Table 4.8.

' Smithsonite Hemimorphite

Cyanide concn. 1. OOM 1. OOM Cyanide to zinc molar ratio 4/1 4/1 Particle size -300 + 75pm -53/im Mineral 3.2 809 4.0358 weight 2 2 Surface area 0.5444m 3.373m Temperature 298°K 298°K Agitation rate 540 rpm 620 rpm -ýý

The results presented in Figs. 4.13 and 4.14 show

that a rise in temperature resulted in an increase in the

rate of dissolution of smithsonite and hemimorphite -149-

C w 10- U C -. N

5- 50

C-) C °C O v 25 V A 35 °C V 50°C "1C o N 0 70°C 5 1( 90 °C

05 10 15 20 Time (h )

Fig. 4.13. The influence of temperature on the rate of dissolution of srnithsonite 1 in 1. OiNI cyanide. -150-

15-

10- C I

O A 5 0 U C O V U C N 5- v 25°C n 35°C 0 50°C a 70°C 90 °C

0 10 20 30 40 Time (h )

Fig. 4.14. The influence temperature the dig tion of on rate of .>c: ý.: of hendmorphite in 1. OM cyanide. -151-

respectively. At temperatures greater than 308°K the sharp increase in zinc concentration observed previously in the smithsonite dissolution did not occur.

The zinc was eventually completely extracted from the smithsonite sample but, again, a lengthy period was needed to recover it from the hemimorphite phase within the smithsonite sample. At higher temperatures, however, this extraction was much faster and a closer approach to the maximum solubility was made.

Although an increase in temperature led to an increased dissolution rate of the hemimorphite a marked decrease in the rate of reaction was still observed after about 10 to 15 hours leaching, and the dissolution curves can be described as almost parabolic. This phenomenon was less noticeable, however, at higher temperatures.

The dissolution reactions for both these minerals can be considered as a second order elementary reaction, and the rate of the reactions can be expressed as: -

d[Zn] =kAC (4.16) dt where the velocity constant, k, can be written in the

Arrhenius form: -

k= Ale-E/RT (4.17)

where E is the experimentally determined critical -152-

increment of energy to the reaction and A' is the pre-exponential term.

From the theory of the transition state the reaction velocity constant can be written as :-

k=KkT eA5t/R e-&H'RT (4.18) h

X where is a transmission coefficient, k is Boltzmann's constant, h is Planck's constant, j S+ and L&H* are the entropy and enthalpy of activation respectively.

Equation (4.18) can be rewritten to include the experimental activation energy because

E =ýfi + RT (4.19) and, therefore,

E/RT k= e)ýICTedSt/R e- (4.20) h

The first part of Equation (4.20) includes a reaction frequency and an entropy term that results from steric influences and can be regarded in a similar way to the preexponential term of the Arrhenius expression.

A' and E from the Arrhenius equation can be considered independent of temperature to a first approximation and, hence, if the equation is followed a plot of log k against T-1 should result in a straight -153-

-. -x

4- O 4- N O

"IN V O

O u qp

riý

2.8 2.9 3.0 3.1 3.2 33 103/T ( °K"7 )

Fig. 4.15. Arrheniu $ plots for smithsonite and hemimorphite. -154-

line.

The Arrhenius plots for smithsonite and

hemimorphite are shown in Fig. 4.15 at temperatures

between 298 and 363°K. Analysis of variance showed

that a straight line provided an adequate fit to the data

and the following equations were found by linear

regression.

For smithsonite: -

k=0.146 e -21100/RT s-1m-2 (4.21)

and the 95% confidence interval for the critical increment 1. of energy was 19850

For hemimorphite: -

k=0.335e '28400/ßT s-1m-2 (4.22)

and the 95% confidence interval for the critical increment

of energy was 24100

The critical increment of energy for these

heterogeneous reactions is not the same as the

activation energy for homogeneous reactions. It is

only an apparent activation energy as the activation

energy in the adsorbed layer and the heat of adsorption

of the reactants and products should also be considered.

4.8. General rate equations for the dissolution of smithsonite

and hemimorphite in cyanide -155-

By combinding the influence of surface area, cyanide concentration and temperature, the following rate equations were derived.

For smithsonite: -

d IZrjl = 0.144 AC e-21100/RT (4.23) dt and the 95% confidence interval for the constant term was 0.112 < k< 0.176. and for hemimorphite: - [Z d n] = 0.352 AC e-28400/RT (4.24) and the 95% confidence limit for the constant term was

0.324

The smithsonite and hemimorphite samples were not pure, however, and the rate equations di ould, therefore, be adjusted to account for the true surface area available for reaction. In the absence of detailed information the specific surface areas of the various non- reacting mineral phases in the samples were assumed to be similar to that of the reacting phase. The rate constant k must be replaced by k/9 where 9 is the proportion of reacting mineral surface in the sample.

Equations (4.23) and (4.24) can, therefore, be rewritten as, for smithsonite: - -156-

[Zn d = 0.152 AR C e-21100/RT (4.25) dt

and the 95% confidence interval for the constant term

0.120

and for hemimorphite: -

jznj = 0.471 Ih C e-28400/RT (4.26) dt

and the 95% confidence interval for the constant term

0.443

area of the reacting mineral phase only.

Deviations of the experimental results from the

models are due to (apart from necessary experimental

errors): -

(1) Forcing the data to fit a first order heterogeneous

rate equation.

(2) Assuming the Arrhenius equation to be valid.

(3) Assuming that the specific surface area of the

reacting phase was similar to the sample as a whole.

4.9. Scanning electron microscope examination of smithsonite

and hemimorphite after leaching in sodium cyanide

Polished sections of smithsonite and hemi-

morphite grains (-300 + 75pm) set in Araldite were

prepared and then leached in 1. OM sodium cyanide at

room temperature under gentle agitation. Sections -157-

were removed at intervals, washed, dried and examined with a stereoscan scanning electron microscope,

so that a study could be made of the change in physical

characteristics of the minerals during dissolution.

Smithsonite

The fresh smithso nite surface (Plate 4.1) appeared to be relatively smooth with only a few small

cracks and polishing scratches present. There were no

observable phase or crystal boundaries within the grains.

After one hour leaching, however, a large number of

etch pits could be seen on the mineral surface (Plate 4.2).

The pits merged to form cracks and at these the rate of

dissolution was more rapid than at other parts of the

mineral. Although the dissolution rate was much faster

at the location of polishing scratches than on the general

surface, etch pits were not prevalent on the scratches

themselves. The production of dissolution cracks

seemed to occur in two ways, either by the merging of

etch pits which had been formed along a line or by a

number of etch pits nucleating at a point which then

developed into a line. These dissolution cracks

gradually extended over the entire surface of the

smithsonite (Plate 4.3) dividing the particles into -158-

smaller crystals.

Certain grains were noticed to dissolve at a much faster rate than others. The more rapidly dissolved grains developed a fibrous texture (Plate 4.4) with the fibres orientated parallel or slightly inclined

to the mineral surface whereas the grains which were

slower to dissolve either did not show this fibrous

texture or were orientated with the fibres nearly

perpendicular to the smithsonite surface. The

smithsonite appeared to be made up of polycrystalline

material (Plate 4.4), with the grains within a particle

often at differing orientations and the dissolution was

noted to be rapid at these sub-grain boundaries.

Plate 4.4 shows the fibrous structure of the smithsonite

where the dissolution rate was fast. Etch pits were not

greatly in evidence and a few perfect rhombs of material

could be seen. Considerably more etch pits were

observed when the smithsonite was in the 'ends-up'

position with fibres directed perpendicular to the

mineral surface which was littered with small perfectly

formed rhombohedral crystals. (Plate 4.5)

The effect of dissolution on the shape of the

smithsonite particles can be attributed to the anisotropic

leaching characteristics of this mineral. The crystal -159-

1 w 71-

Plate 4.1. Scanning electron photomicrograph of the fresh smithsonite surface. (x 100)

Plate 4.2. Formation of etch pits on the smithsonite surface (SEM) ]'referential leaching along scratch marks and cracks (x 1,000). -160-

ý,: ýi. -ý' ;-'

r `? _'º w- ". 1

Plate 4.3. Formation of dissolution cracks over the entire smithsonite surface (SEM) (x 250)

Plate 4.4. Development of a fibrous texture on the (SEM) The smithsonite grain . polycrystalline nature of smithsonite becomes obvious (x 250). -161-

Plate 4.5. The polycrystalline nature of smithsonite can be seen with rapid dissolution at grain boundaries (SEM). Note the 'end-up' orientation of the mineral and appearance of rhombs. (x 500).

Plate 4. G. Parts of the smithsonite have been leached at a faster rate than those areas orientated the fibrous 'end-up' Particle with structure . disintegration has occurred. (SEM x 250). -162-

Plate 4.7. All the smithsonite has dissolved and only perfect rhombohedral crystals remain. (SEAT x 2,500) -163-

faces parallel to the fibres were much more readily dissolved than those perpendicular to them as is evident from Plate 4.5 where dissolution of the particles was more rapid from the edge of the particle, where the fibres were available for reaction, than from the top surface. This would also account for the proliferation of etch pits on the surface of this grain.

Anisotropic dissolution of smithsonite is also illustrated in Plate 4.6 which shows that the smithsonite particle disintegrated after some time of leaching.

Eventually, all the smithsonite was dissolved leaving perfect rhombohedral crystals scattered about the hole remaining in the Araldite. These rhombs (Plate 4.7) were analysed with an energy dispersive unit attached to the scanning electron microscope and the characteristic wavelengths obtained showed that the rhombs consisted of a calcium compound with a very

small amount of magnesium. This material was undoubtedly calcite which has a perfect rhombohedral

crystal structure and grows isostructurally with

s-mithsonite.

Hemimorphite

Only a few small cracks were present on the

fresh hemimorphite surface and these were probably -164-

caused by the platiness of the crystal structure. After two hours leaching, well-defined etch pits could be seen located either individually or in clusters, all orientated in the same direction. The dissolution

(Plate 4 8) rate was faster at the edges and cracks in the grain and also at polishing scratches (Plate 4.9) than on the unmarked surface. Dissolution occurred, however, over the entire hemimorphite surface as indicated by the roughening of the grain in Plate 4.10. The line of large etch pits on this grain probably denoted the existence of a grain boundary or dislocation. A step- like the (Plate 4.11 ), extension across mineral surface . orientated in two directions, was apparently nucleated by the etch pits. Plate 4.12 shows that leaching was more rapid in certain directions and this can be clearly

seen from Plate 4.13 where dissolution was more rapid

on the crystal faces orientated approximately top to

bottom on the plate rather than those in the perpendicular

direction. The dissolution was very anisotropic.

Stalactitic forms were observed after 12 hours

leaching (Plate 4.14) on a crystal face apparently

advancing across the surface of the grain. The

formation of slot like pits should be noted on the upper -165-

right ofthe plate to be orientated in a direction parallel

to the advancing face. The elongation of the etch pits

also show that dissolution of the hemimorphite was

more rapid in this direction. After 12 hours leaching,

the entire surface of the hemimorphite was covered

by large pits (Plate 4.15) which gradually developed a

slot-like appearance. Dissolution was seen to be

faster on the crystal faces parallel to the orientation

of the slots indicated by the line A-A' on Plate 4.15.

Either the mineral surface became greatly

roughened (Plate 4.16) and any slot-like developments

were directed almost parallel to the surface, or a

layered structure was formed (Plate 4.17) presumably

due to the slots being inclined to the mineral surface.

Eventually, continued dissolution for 26 hours resulted

in the lattice-like structure of Plate 4.18 either as a

'honeycomb' (area A) or as a very fine platiness

(area B) until finally all the hemimorphite was

dissolved.

At no stage of the dissolution was there evidence

of a surface product being formed whether as a residue

or by precipitation from solution.

Just as the anisotropy of crystal growth is -1 GG-

1% 1001*'*-

o; 4, -,

, rw-

401

Plate 4.8. The formation of etch pits can be seen on the hemirnorphite surface. Faster dissolution was noted along cracks on the mineral surface (SEM x 1,000)

Plate 4.9. Preferential leaching at polishing scratches (S.IýNT x 2,500) -1.67-

Plate 4.10. Rou

-dM66

Plate 4.11. Dissolution initiated at etch pits which then form a series of steps eventually producing; cracks on the mineral surface (SEAI x 2,500) r- 0 ..;., I AR

_1 ,ý

Plate 4.12. Cracks form on lie-mimorphito surface (SEM x 500)

ppl-

Amok,

+r 1`-

40µj

Plate 4.13. Dissolution cracks ;; raciually extend across the mineral. grain (5L; 1-T x 250) Plate 4.14. Anisotropic dissolution of heminrnrphite. Stalaci. iti(, forms on an advancing crystal face. More rapid dissolution on crystal faces orientated to bottom right. (SEM x 2,500).

Plate 4.15. Entire surface covered by deep etch pits orientated in the same direction (S1, M x 2,500) -170--

Yýo "} _ rý , yam. ý" a

Plate 4.1(3 I-temimorpliite surface has become vet'y rouYli as dissolution takes place over entire grain (SEM x 1,000)

ýýý

ple '

Plate 4.17, Layered structure o(' hemimorphite (SJ! \-i x5 00) -171-

Plate 4.18 Lattice-like structure develops either as a honeycomb (area A) or a platiness (area I3). -172-

Plate 4.19. Polaroid photograph of back$cacrered electron irrvage of a sectioned lii , )ixnorphite grain after 65 hours leaching. Silicon and zinc traces are superimposed. -173-

detailed(92) the anisotropic well , many examples of

have been dissolution of crystals presented (93,94,95) and Gatos(96) has reported that many crystalline materials exhibit differences in chemical reactivity on different crystal planes. These differences were said to be due to differences in atomic spacing or packing between the various crystal planes affecting

Prosser however, the adsorption of reactants. X97), has also pointed out that dislocations may take preferred directions in the crystal and terminate at external

surfaces and grain boundaries and the possibility arises, therefore, of certain crystal planes having a

greater density of dislocations than others.

A good correspondence between centres of

preferred attack and regions of emergence of dis-

locations at a surface has been reported (98,99,100,101)

and consequently the etch pit technique has often been

utilized to observe the position of dislocations

Considerable strain energy is stored in the

elastically disturbed region around a dislocation line(101)

has been to lead to and the extra energy said (102)

enhanced reactivity because of the availability of this

energy to contribute to activation of the reactant Species.

Grain boundaries can also be treated as dislocations and -174-

Read has that the boundary (103) shown energy of a grain is proportional to the angular difference in orientation between the crystals. Regions of relatively high energy, therefore, have enhanced reactivity.

The mechanical working of solids has also been

to the reported (104) affect chemical reactivity of the solid and the presence of a disordered lattice

has been to structure on a quartz surface said (72,105) increase the rate of dissolution as well as the solubility of quartz in water.

It was not surprising, therefore, that the rate of dissolution of both the smithsonite and hemimorphite was initially greatest along polishing scratches where the crystal lattice would have been appreciably disordered. The lack of large etch pits on the scratches indicates that the locally disordered lattice was the contributing factor to the rapid rate of dissolution rather than the presence of dislocations emerging at the surface on the scratches.

Increased reactivity wi s very much in evidence at cracks and edges, where the surface free energy would be rather higher than at the rest of the mineral

surface. Although sub-grain boundaries were not.

observed by the scanning electron microscope on the -175-

smithsonite sections prior to leaching, the smithsonite grains were shown to be polycrystalline because of

cyanide attack along the grain boundaries. The presence of sub-grain boundaries were not generally

detected in the grains of hemimorphite.

The different reactivities of different crystal

faces was pronounced in the case of smithsonite and

was easily observed because of the gradual formation

of a fibrous texture on the mineral grains. The

smithsonite was more readily dissolved from the

crystal planes parallel to the fibres rather than from

the planes perpendicular to them, but as no well defined

crystallographic features could be distinguished the

crystal planes were not identified.

The formation of etch pits on the hemimorphite

surface was most probably located at the point of

emergence of dislocations at the mineral surface.

Development of step-like layers seemed to indicate

that dissolution of the layers was nucleated at the

dislocations in a similar manner to that noticed by

Warren(93) in the dissolution of hematite in acids.

The facets produced during the dissolution of the

hemimorphite most likely corresponded to faces at

which movement of the steps was most rapid. The -176-

in crystal face marked by the long edge of the slots the memimorphite surface was parallel to the face that was most rapidly dissolved, and it is most probable that preferential development of these faces resulted in the layered structure of the partially leached mineral.

Kostov(106)has described the lattice structure of hemimorphite which is given in Fig. 4.16. The

(Si207) group exists as two tetrahedrally coordinated

(SiO4) groups with a common corner. The tetrahedra are oriented with one corner pointing along the c axis and the zinc atoms are tetrahedrally co-ordinated with three oxygens of the silica tetrahedra and one hydroxyl.

The water molecules are in channels of this strcture and are parallel to the (110) planes which explains the perfect cleavage of hemimorphite on this plane. After initiation of dissolution at the mineral surface it would seem possible that preferred attack might occur along the direction of the channels in the lattice leaving the planes available for reaction. The formation of slots

in the hemimorphite after some time of leaching

together with the development of a platy structure

indicates that dissolution of this type might have 11101 occurred and that the layers denoted the crystal

plane. -177-

E C( M1

E c 0 n 0

O©

4.16. Lattice structure of hemimorphite showing channels on (110) plane. -178-

Continued leaching of the smithsonite grains resulted in eventual particle disintegration because of the rapid leaching along the sub-grain boundaries whereas in the case of hemimorphite this phenomenon did not occur due to the absence of any polycrystalline nature.

The complete dissolution of both smithsonite and hemimorphite supports the view that no solid reaction products were formed during leaching. The possibility remains, however, that a thin film of surface product forming during the dissolutions may have been removed by the washing of the sections in water prior to examination with the scanning electron microscope.

Both smithsonite and hemimorphite exhibit anisotropic leaching characteristics and the surface area available for reaction will, therefore, change markedly during the course of the dissolutions. The relationship between surface area and time of leaching is complex as the following factors illustrate: -

(1) The rapid dissolution of small particles and the

rounding of sharp corners and edges by preferential

leaching tend to decrease the specific surface' area. -179-

(2) Gradual dissolution of large particles generally

results in an increased specific surface area due

to size reduction.

(3) Anisotropic dissolution leads to roughening of

the mineral surface and consequently a greater

surface area.

(4) An increase in the specific surface area would

also be caused by particulate disintegration.

The interplay of these factors makes the

prediction of surface area changes during leaching

very difficult.

4.10. Surface area changes during smithsonite dissolution

Examination of smithsonite grains, that had

undergone partial dissolution in cyanide, with a scanning

electron microscope showed that after some time the

particles disintegrated. It was thought that this would

have a significant effect on the surface area available

for reaction.

Repeated dissolution tests were, therefore,

carried out under the conditions summarised in Table 4.9. -180-

Table 4_9.

Cyanide concentration 1. OOM Cyanide to zinc molar ratio 4/1 Particle size -300 + 75pm Mineral 3.280g weight 2 Surface area 0.5444m Temperature 298°K Agitation rate 540 rpm

Agitation was stopped after a certain time and the remaining solids were removed from the ?each solution by filtration. After washing and drying, the specific surface area of the residues were determined by B. E. T. krypton adsorption and hence the total surface area.

The results are presented in Fig. 4.17 and show that the shape of the dissolution curves closely followed those previously obtained. The specific surface area measurements given on the figure represent the mean of two determinations. As the reaction proceeded the specific surface area decreased from 0.166m2g-1 at the start to 0.114m2g-1 after 6 hours. Shortly afterwards the specific surface area increased to 0.140m2g-1 and the total surface area available for reaction also

increased. This corresponded to the region in which

the dissolution rate also increased. -1ß1- 00

C äi -(0.156m2g" 0.199 m2 )0 U C (0.140m2j'0.203 0 . M2) N ýý Ql (0.114m2g-1/0.172m2) O 50

C U (0.133m2g110.206m2) C 0 O V U C ýN

(0 166 m2g-/0 544m2)

05 10 15 20 Time (h) Fig. 4.17. Variation in the specific surface area of smithsonite during leaching. -182-

4.11. Further dissolution studies on hemimorphite

Samples of hemimorphite were leached in

in Table 4.10. cyanide under the conditions summarised

Table 4.10,

Cyanide concentration ý 1. OOM Cyanide to zinc molar ratio 4/1 75pm Particle size -300 + 3.777g Mineral weight 2 Surface area 0.397m 9K Temperature 298 Agitation rate 640 rpm

After periods of 48 and 65 hours the hemimorphite

was removed from solution by filtration, washed, dried

and made into polished sections such that the leached

grains were intersected. Leaching times of this order

resulted in about 30% and 40% respectively of the zinc

being dissolved from the sample.

The polished sections were examined with a

'Geoscan' microprobe analyser which showed that there

was no variation in either zinc or silicon concentrations

across the hemimorphite grains. A photograph of the

'Geoscan' trace, following the zinc and silicon

concentration superimposed on the backscattered

electron image of a hemimorphite grain leached for

65 hours, is presented in Plate 4.19. The silicon -183-

and zinc traces were well correlated and show that there was no differential dissolution of silicon or zinc.

No obvious surface product formation could be discerned but a very thin film would be beyond the resolution of the instrument.

Spot counts for zinc and silicon were taken about the grain and the results are summarised in Table 4.11.

Table 4.11 'Geoscan' anaysis of leached hemimorphite

Leachtime Zn Si (h nj (%) (%)

0 52+3 10+2 48 524 9-2 65 54-4 9-1 Theoretical 54 0 11.7

There was no significant difference in zinc and

silican concentration between unleached and partly

leached hemimorphite. All silicon counts gave slightly

lower analyses than determinations by classical methods.

Direct: evidence of a surface formation on

hemimorphite was not found by examination of the

partially leached material with either a scanning

electron microscope or a microprobe analyser. The

shape of the hemimorphite dissolution curves indicated,

however, that a solid surface product might be inhibiting -184-

the reaction and the apparent activation energy of

28.4kJ mole-1 suggested that this reaction was mass transfer controlled.

Henderson(72) has reported that washing quartz with water partially removes a disturbed surface layer thus increasing both the solubility and rate of reaction of quartz. All samples of hemimorphite were washed well, before the instrumental examinations, and the possibility of removal of a thin surface film can not be discounted. Consequently dissolution tests were carried out on hemimorphite using the conditions

summarised in Table 4.12.

Table 4.12

Cyanide concentration 1. OOM Cyanide to zinc molar ratio 4/1 Particle size -53 + 37/Im Mineral 3.777g weight 2 Surface area 0.861m Temperature 250C Agitation rate 640 rpm

After some time of leaching agitation was stopped

and the remaining solids were removed from the solution

by filtration, the solution being returned to the reaction

vessel. The solids were, first, washed in 0. iM sodium

hydroxide under gentle agitation for 5 minutes. The -185-

solid material was allowed to settle and the

supernatent solution was decanted. The material was

finally washed twice more in water and after decantation

of the water the hemimorphite was dried at 70°C and,

after cooling, returned to the reaction vessel.

Fig. 4.18 shows that, after a few hours leaching,

the reaction rate decreased markedly and also that an

almost parabolic dissolution curve was obtained. The

removal and subsequent washing of the hemimorphite

after 6 hours leaching resulted in an increased rate of

reaction on re-introduction to the solution. This

increased rate of dissolution was not, however, quite

as great as the initial reaction rate and eventually the

rate of zinc removal from the hemimorphite was

markedly decreased. The shape of the dissolution

curve after the surface treatment was similar to that

obtained before it.

4.12. Addition of sodium hydroxide to the cyanide leach

solvent

Dissolution tests were conducted on hemimorphite

and smithsonite in the presence of both sodium cyanide

and added sodium hydroxide. The addition of sodium

hydroxide to the hemimorphite system should result -186-

-S

O V V C \solids N washed

o 10 20 Time (h)

Fig. 4.18. The increase in the rate of dissolution of hemimorphite after the remainin? solids had been removed from the solution, washed in water, 0.1-NI 1aOll, and returned to the same solution. -187-

in the removal of any surface film of silica because the solubility of the silica increases with an increase in p1I

An increase in the rate of hemimorphite dissolution can, therefore, be reasonably expected.

The solubility studies with smithsonite (cf) indicated that in the presence of sodium cyanide - sodium hydroxide mixtures zinc hydroxide was precipitated on the surface probably by the reaction given by Equation

3.12. If such a reaction does occur to an appreciable extent then it is reasonable to assume that the addition of sodium hydroxide to the cyanide solution should influence the dissolution rate in some way.

The conditions used in these dissolution studies are summarised in Table 4.13.

Table 4.13.

Smithsonite Hemimorphite

Cyanide concn. 1. OOM 1. OOM Cyanide to zinc molar ratio 4/1 4/1 Particle size -300 + 75/im -53pm Mineral 3.280g 4.035g weight 2 2 Surface area 0.5444m 3.373m Temperature 298°K 298°K Agitation rate 540 rpm 620 rpm

In the absence of cyanide the rate of dissolution of hemimorphite in sodium hydroxide was slow (Fig. 4.19). -188-

On adding cyanide, however, the rate increased but an appreciable increase was only obtained at high hydroxide concentrations. This effect is illustrated in F ig. 4.19.

The addition of sodium hydroxide to the solvent resulted in the extraction of more zinc from the hemimorphite before the rate of the dissolution markedly decreased i. e. after about 10 hours. This point is emphasised in Table 4.14 where the time required to obtain a 50% extraction in the different solvents is shown. Addition of 0.5M sodium hydroxide to a 1. OOM cyanide solution resulted in a reduction in the t 50 by about 50%. Although the initial rate of zinc extraction from hemimorphite was greater in a

2. OM cyanide solution than in a 1. OM cyanide + 1. OM

sodium hydroxide the 't for the latter solution, 50 was appreciably shorter. -189-

lCýo

1 v

v . 5 0

C C.) [CN] [OH] C O (M) (M) U V C - 2.0 N 5 O 1.0 - m 1.0 0.5 Q 2.0 - p 1.0 l"0

0 10 20 30 40 Tirrrn(h)

Fib;. 4.19. The influence of sodium hydroxide addition on the rate of dissolution of hemimorplhite in 1. OM cyanide. -190-

Table 4_14 Influence of sodium hydroxide addition on the initial dissolution rate of hemimorphite

Sol Initial dissolution ýNaCNT ution [NaOH] rate (M) (M) (mole C1 s-lm-2) x 106 (hrs)

0 2.0 0.04 1.00 1 0 4.06 22.8 1.00 0.5 4.42 11.6 1.00 1.0 5.79 6.4 I 2.00 ý 0- 7.09 ?. 9

These results, therefore, further suggest that a

silica layer builds up at the hemimorphite interface and

prevents the diffusion of reactants and pro ducts to and

from the reaction interface. It should, however, be

stated that in the presence of hydroxide more free

cyanide will be available because of the formation of

zinc hydroxy complexes. This will also lead to an

increased rate of dissolution.

The influence of sodium hydroxide addition on

the rate of dissolution of smithsonite is shown in Fig. 4.20.

In sodium hydroxide solutions the rate of dissolution

was much slower than at equivalent cyanide concentrations.

After a time the zinc concentration reached a maximum

and then decreased. Similar behaviour had previously

been observed during the solubility studies. The addition -191-

of less than 0.5M sodium hydroxide to the cyanide solution had a negligible effect on the dissolution rate but at higher hydroxide concentrations the reaction rate increased. The initial reaction rates are presented in

Table 4.15.

Table 4.15 Influence of sodium hydroxide addition on the initial rate of dissolution of smithsonite

J Solution Initial dipsollutig rate5 _ [NaCN] [NaOH ] (mole 1 s m)x 10 (M) (M)

0 0.5 0. 31 0 1.0 0. 47 1.00 0 2. 63 I 1.00 0.1 2. 66 1.00 0.5 2. 59 1.00 1.0 2. 76 11.00 2.5 3. 70

After 24 hours leaching in a 1. OOM cyanide

solution only 92.5% of the zinc was extracted from the

smithsonite but the extraction neared completion with

a sodium hydroxide addition greater than 1. OM. This

was probably because of the increased rate of dissolution

of the minor hemimorphite phase at increased hydroxide

concentrations. The addition of sodium hydroxide to

the leach solutions results in an increased free cyanide

concentration due to zinc hydroxy complex formation. -192-

loc I 5

I 0 [C N] [0H] (M) (M) N 0 1.0 - x 1.0 0.1 50 A 1.0 0.5 m 1.0 1.0 1.0 2.5 O a-0.5 U e-1.0 U o

05 10 15 20 Time (h)

F_g;. 4.20. The influence of sodium hydroxide addition on the rate of dissolution of smithsonite 1 in 1.0': cyanide. -193-

2.5M [oHl

C U C 0 COM U 1.OM e

OC Ü 4 ýQ

O.SM [0

0&0. JMM[o

0 0.1 0.2 zinc concn. (M )

4.21. The influence of sodium hydroxide addition on the free cyanide concentration of a 1. OM cyanide solution containing zinc. -194-

Solution equilibria calculations (Fig. 4.21) have shown

that at sodium hydroxide concentrations greater than

1. OM, even at high zinc concentrations, the major

part of the cyanide would be present as the free cyanide

ion.

4.13. Roasting of hemimorphite

Hemimorphite loses its compositional water on

heating and is converted to willemite. The temperature

at which this reaction takes place is uncertain but it has

been reported by several authors to be in the range 5230

to 7730K(106,107) Dissolution tests " were carried out

to determine whether or not heating the hemimorphite

increased its rate of dissolution. Prior to conducting

the roasting and dissolution tests the hemimorphite was

characterized by thermogravimetric and differential

thermal analysis. The thermogravimetric curve (Fig. 4.22)

showed that there was a continuous weight loss from

573°K to about 673°K after which there was a sudden

loss in weight in a series of steps. Differential thermal

analysis showed that there was little deviation between

the sample and reference temperature until an endothermic

peak was obtained at about 10030K.

A differential thermal analysis curve for -195-

hemimorphite from Franklin, New Jersey has been presented by Zussman(l08) which exhibits an endothermic peak at about 1013°K and an exothermic peak at about 1173°K. The lower value is close to that obtained in this work. Zussman concluded that the hemimorphite dehydrates in several stages and that the exothermic peak was due to the formation of willemite.

Bragg(109)reported that Zambonini showed that when hemimorphite was heated to 773°K water was lost continuously without loss of crystal transparency. Only half of the water content was lost and the remainder was only removed by heating to much higher temperatures where destruction of the crystal occurred.

The weight loss of 3.73% after the first step in the thermogravimetric curve (Fig. 4.22) was equivalent to the loss (theoretical loss 3.75%) of one water molecule .

Samples of hemimorphite were examined microscopically,

by scanning electron microscope and by X-ray diffraction

after roasting for 17 hours at 673°K. No marked increase

in surface cracking or crystal opacity was observed.

The hemimorphite did, however, take a slight yellow

colouration. The X-ray diffraction pattern showed that

the mineral had retained its hemimorphite structure

and that no willemite was present. These factors -196-

indicated that the water loss was from the water loosely bound in the hemimorphite crystal structure.

Zussman(108) reported that goethite is dehydrated at a temperature of 673°K and the sudden weight losses that occurred in a series of steps was, therefore, probably due to the loss of water from the hydrated iron oxides present in the sample.

The endothermic peak obtained at 10030K can be attributed to the removal of the more strongly bound hydroxyl groups from the hemimorphite lattice.

The thermogravimetric curve was not continued beyond

800°K because of the difficulties in quantitative

interpretation of the decomposition of hemimorphite

caused by interference from the decomposition of the

iron oxides and calcite of the matrix.

Samples of -300 + 75pm hemimorphite were

roasted at 573°K and 673°K for various times and

dissolution studies were conducted after the hemimorphite

had been cooled to room temperature. The leaching

conditions that were used are summarised in Table 4.16. -197-

(3 te) Z

a ý. Eö

4 CO

0 0tl%

c q E b E V)

Clv

1_^

q `.vi

O E...

v ý o

F ig. 4.22. The loss in weight of hemimorphite with increase in temperature (thermogravimetric balance). -198-

Table 4.16

Cyanide concentration ' 1. OOM Cyanide to zinc molar ratio 4/1 Particle size -300 + 75um Mineral 3.777g weight 2 Surface area 0.397m Temperature 298°K Agitation rate 640 rpm

Fig. 4.23 shows that roasting at 573°K had very little effect on increasing the rate of dissolution of hemimorphite. Roasting at 673°K for increasing times, however, produced a. large increase in the rate of zinc extraction but roasting for more than 4 hours did not , increase the rate further (Fig. 4.24). In all cases, the rate of dissolution markedly decreased after about 10 hours leaching. The specific surface area of the samples after roasting were determined by BET krypton adsorption and the results are presented in Table 4.17 which shows that roasting produced an increase in the surface area available for reaction. -199-

A unroasted 20 15 300°C for 1h

400°C " 0 25 h

13 400°C 1h

4 400°C 4h 0 C 0 400°C " 17h 13 10 L J

V 0

u O V U C N5

0 10 20 30 40 rime (hj

Fig. 4.23. The influence of roasting temperature and time on the rate of dissolution of hemimorphite. -200-

h0- (3.14%) OZZ.

n, (0-97%) (3.6601 (3.39 "A (324%) a, O c 0

4 5 (1.22%) b .

C 573°K (0.23%).

lit, o5 17 45 Roasting time (h

1 :s. 4.24. The initial rate of dissolution of hemimorphite in 1. OAT cyanide after roastin, ` at (-i93 K. (The weight losses are shoe n in brackets). -201-

Table 4.17 Surface area measurements of roasted

hemimorphite

Roasting time I Specif ic surface area (hrs) (m g)

0 0. 106 0. 25 0. 337 1. 0 0. 494 4. 0 0. 462 17. 0 0. 482 45. 0 0. 508

The kinetic studies have shown previously that

the rate of dissolution of hemimorphite was directly

proportional to the surface area available for reaction

and hence the initial reaction rates were plotted against

the sample surface area in Fig. 4.25. A linear

relationship was found by linear regression to be

d Zn 4.66 x 10-6 A (4.27) dt

where the 95% confidence limit for the constant term

was 3.65 x 10-6 4k45.67 x 10-6.

The value of the constant term should be compared

with the value of 3.33 x 10-6 found from those tests

designed to determine the effect of surface area on the

dissolution rate (cf). Statistical analysis of the two

test series showed that the difference was significant. -202-

N 0 0 0 6+4.66x 6A E 0R= -0.28x 1Ö 10- ý ° Q) O oe c 0 - J 05 'expected from mode! ö i i c

o, 2 surface area A (m2)

Fig. 4.25. Regression line for the influence of surface area on the initial dissolution rate of roasted heinimorphite. -203-

This difference can be attributed to either an incorrect

determination of the specific surface area because of

rehydration on cooling or because heating the material

disturbed the crystal lattice so that more stressed areas,

dislocations or point defects were formed. Both these

factors are likely to affect the dissolution rate. The

former might give rise to irreproducibility unless the

cooling conditions are carefully controlled whereas the

latter will produce an increased rate of dissolution.

To summarise, the rate of dissolution of

hemimorphite in cyanide solutions is increased by

heating the hemimorphite prior to dissolution. Heating

produces an increased surface because of the area .

removal of water from channels within the crystal

and also a slightly more reactive surface. -204-

5. DISCUSSION OF KINETIC RESULTS

The results obtained in this kinetic study have

allowed an empirical model for the rate of dissolution

of smithsonite and hemimorphite in cyanide solutions

to be derived. These models are only valid, however,

for the reaction conditions utilised in the dissolution

tests i. e. where an increase in stirrer speed had no

further effect on the dissolution rate. Under such

conditions the mass transfer rate through the liquid

boundary layer at the mineral surface can be

considered to be either constant or else no longer the

rate controlling step of the dissolution.

The increase in the rate of reaction with an

increase in agitation rate was due, initially, to a

greater number of particles becoming suspended and

allowing free access of the solvent to the mineral surface.

When all the particles were suspended in the leach

solution, increased agitation led to only a slight- increase

in the dissolution rate which eventually became constant

as aeration and vortex formation developed. Such

features of solid-liquid mixing in agitated tanks are

well known

The rate of dissolution of zinc oxide and

hydrozincite was very rapid owing to their very great -205-

specific surface areas but for the latter mineral the reaction velocity constant was apparently less than that for smithsonite. It is possible that the presence of zinc hydroxide groups in the hydrozincite structure leads to a smaller reaction velocity constant, but it

is more probable that the surface area measured by krypton adsorption was not all immediately available

for reaction with the cyanide, a certain time being

necessary before the cyanide had diffused through the

pores to all parts of the mineral. Steric hindrance by

the tetrahedral zinc cyanide complex might conceivably

have played a part in preventing access of cyanide to

the internal surface of the hydrozincite if the pore size

was very small.

The dissolution reactions of smithsonite and

hemimorphite in cyanide followed heterogeneous

reaction theory closely, that is, the rate of dissolution

was directly dependent on the surface area of the

material availalbe for reaction.

The sudden rise in zinc concentration after a

certain time of leaching of smithsonite can be explained

by reference to the physical changes occurring during

the dissolution. Examination of the smithsonite with

a scanning electron microscope after partial leaching -206-

showed that dissolution was rapid along sub-grain boundaries within the polycrystalline smithsonite grains and eventually the particles disintegrated.

The sudden increase in surface area would result in an increased rate of reaction. Surface area determinations made on the smithsonite material at various stages in the disintegration gave results that showed an increase in specific surface area and a slight increase in the total surface area. The magnitude of the change, however, did not seem sufficient to cause such a marked effect on'the observed reaction rate. It is likely that, when the mineral grains disintegrated, the very small particles produced dissolved rapidly.

Consequently, the surface area of the material determined by gas adsorption was almost certainly very much less than that produced immediately upon particulate disintegration.

The sharp increase in zinc concentration occurred after a shorter leaching time with finer sized starting material. This was not surprising as the smaller particles can be expected to disintegrate much sooner than the coarser material. At very fine sizes the absence of any sharp increase in zinc concentration was not unexpected because most of the material would -207-

have already been broken into the sub-grain sized crystals.

The sharp increase in zinc concentration -mas not observed when the cyanide to zinc molar ratio was 2/1 or less owing to the dissolution reaction approaching equilibrium before the break-up of the smithsonite grains. At cyanide to zinc molar ratios greater than 8/1 the dissolution was so rapid as to preclude the observation of a sudden increase in zinc concentration.

The different rates of dissolution of the two sources of smithsonite was unlikely to be caused by differences in the reaction mechanism. Although the crystal structure of the two samples was similar, examination of the material with a scanning electron microscope showed that dissolution of smithsonite 1 was initiated at fewer surface sites than was observed in the case of smithsonite. 2. This difference must presumably be due to a greater defect density in the latter sample.

The dissolution rate of smithsonite was increased when the cyanide solutions contained greater than 1. OM sodium hydroxide. This was undoubtedly due to increased free cyanide concentration in the leach solution. The surface precipitation of zinc hydroxide, probably; has -208-

little effect on the reaction rate because the cyanide ion most likely reacts with surface zinc sites whether they are bonded with carbonate or hydroxide. If the rate controlling process is diffusion of cyanide or zinc cyanide complexes to or from the reaction interface then differences in the surface reaction rate would be unlikely to affect the overall rate of dissolution.

The hemimorphite dissolution curves did not exhibit any sudden increases in zinc concentration because the mineral grains did not totally disintegrate.

Examination of the grains by scanning electron microscopy after leaching for some time showed the absence of any

appreciable sub-grain crystal texture.

The effect of temperature and cyanide concentration

and particle size on the reaction rate of smithsonite and

hemimorphite was determined under conditions where

an increase in agitation rate had no further effect on

the dissolution. Independence of the dissolution rate on

the agitation rate implies that either the reaction is not

mass transfer controlled or else the hydrodynamic

efficiency of the agitation system has reached a maximum.

Bircumshaw Riddiford have and (77) reported

that diffusion controlled reactionq generally have -an -209-

activation energy between 10.5 and 27.2kJ mole-1 with many in aqueous solution having an activation

energy of about l6kJ mole-1. The experimental

activation energies obtained for the dissolution of

smithsonite (21.1±1.3kJ mole-1) and hemimorphite

(28.4± 4.3kJ mole-l) therefore point to diffusion

controlled reactions although the value obtained

in the latter case is rather higher than might be

expected.

The temperature dependence of the reaction

velocity constants have been expressed in terms of

the Arrhenius activation energy. However, this is only

an apparent activation energy and is not the same as

for homogeneous reactions because the following should

also be considered,

1) Activation energy in the adsorbed layer.

2) Heat of adsorption of the reactants.

3) Heat of adsorption of the products.

The measured activation energy does not,

however, distinguish between which mass transfer

stage involved in the reaction is rate determining or

whether the rate is controlled by an intermediate process

involving both a mass transfer and chemical process.

The dissolution curves showed that the rate of -210-

hemimorphite dissolution was markedly decreased

after a few hours leaching. The general shape of the

curves suggested an almost parabolic relationship

between the amount of -zinc extracted and the time of

leaching. Similar features have been noted in other

dissolution(, Peters studies of mineral 10) , (111)

has said that parabolic leach kinetics are usually due to

the formation of a thickening film of surface products

and that this film provides an ever increasing barrier

to diffusion of either the metal ion (M) across to the

interface II or else of reactant (R) to the mineral

interface I illustrated in Fig. 5.1.

Fig. 5.1. Schematic diagram of diffusion through a

growing porous film

M11 '[R R reactant concn.

solid porous film bulk solution

product concn. [p]11 ß[p11 [p]0 -211-

The results that have been presented have shown that the hemimorphite totally dissolved and that the zinc and silicon were in solution in the correct stoichiometric ratio of 2/1. Moreover, no evidence for a growing solid film on the hemimorphite surface could be obtained by microprobe analysis. The shape of the dissolution curve and the increase in reaction rate after washing the mineral surface does indicate, however, that some kind of surface product was formed during leaching which inhibited the dissolution reaction.

In addition to the formation of some insoluble layer on hemimorphite the possibility of the formation of insoluble zinc cyanide on smithsonite must be considered.

Burkin has insoluble (74) suggested that an metal cyanide layer influences the rate of dissolution of metals in cyanide solutions. He proposed that the metal would be unlikely to react in one step to give the soluble cyanide

M+ 2CN M(CN) (5.1) 2+e but would probably proceed via the stepwise reaction

M+ CN ý-= MCN +e (5.2)

If analagous reactions occur with smithsonite and hemimorphite, zinc cyanide should be present on the f mineral surfaces. Whether or not the species Zn(CN) -212-

exists in solution has been disputed by Izatt(16) who reported that zinc cyanide complexation from Zn2+ and

CN in solution takes place non-stepwise to form

Zn(CN) then to form the higher 2(aq) and stepwise zinc Collier(112) Persson cyanide complexes. and (17) however, have suggested that Zn(CN)+ exists in solution but that it has a rather narrow range of existence. It would seem likely, therefore, that dissolution of the minerals would take place through the stepwise reaction,

Zn CN Zn(CN)+ (5.3) surface+ without formation cf surface zinc cyanide. Detailed solution equilibria calculations showed that the conditions required for precipitation of the zinc cyanide were unlikely to be encountered at the mineral surface.

These calculations were supported by the lack of X-ray diffraction evidence for the presence of zinc cyanide on partially leached smithsonite grains.

If the dissolution of the smithsonite and

hemimorphite were controlled by an intermediate

process the rate of the appropriate chemical and mass

transfer processes would be of the same order of

magnitude. -213-

For a first order chemical reaction CCN d [Zn] =kA and for the diffusion controlled reaction

[CN J0 ICN1. (5.5) df rj =k,, - jA

[CN [CN li where and are the free cyanide 0 concentrations in bulk solution and at the reaction interface respectively.

At the steady state the overall reaction would be represented by

d[ZnL = k1A[CN]o (5.6) _ dt where kl = kc kT (5.7) kc +k T and the process is first order with respect to the

sample surface area and the cyanide concentration in accordance with the experimental results.

To determine whether or not the independence

of the dissolution rates on the agitation rate was

attributable to the attainment of maximum hydrodynamic

efficiency baffles were inserted in the reaction vessels

to increase the degree of turbulence and also to prevent

vortex formation in the leach solution. Two baffles,

1.0cm. wide, were attached opposite to each other to -214-

the walls of the reaction vessel. The baffles extended

from the top of the reactor almost to the bottom

leaving enough clearance for the impellor to rotate.

The dissolution tests were carried out using

the conditions summarised in Table 5.1.

Table 5.1.

Smithsonite Hemimorphite

Cyanide concn. 1. OOM 1. OOM Cyanide to zinc molar ratio 4/1 4/1 + 75prn Particle size -300 + 75pm -300 3.28 Og 3.777 92 Mineral weight 2 Surface area 0.5444m 3.373m °K ° Temperature 298 298 K

The rate of dissolution of hemimorphite was not

significantly increased in the presence of baffles

(Fig. 5.2) whereas that of the smithsonite was. These

results indicate, therefore, that in the case of

smithsonite the dissolution was controlled by mass

transfer of either the cyanide to the reaction interface

or of the zinc cyanide complexes away from it. The

size of the zinc cyanide complexes is greater than that

of the cyanide ion and it is possible, therefore, that

the transport of the zinc cyanide complexes from the

surface is the rate determining step. -215-

500 R.P M. 1 00

0 smithsonite (baffled) N-, % 6O hemimorphite ( ý" )

ýý p smithsonite (unbaffled) c N 5O hemimorphile ( "" ) v, 0 E

4- 4 I- c 0

. - C

05 10 15 20 agitation rote (R"P S" )

Fig. 5.2. Influence of inserting baffles into the reaction vessel on the initial rate of dissolution of hennirnorphite and smithsonite. -216-

In the case of hemimorphite the results indicate that either the dissolution was chemically controlled or that the rate determining step was diffusion of reactants or products through an insoluble surface layer. The

parabolic shape of the dissolution curve and the

behaviour of the hemimorphite on reinsertion into the

cyanide leach solution after thorough washing of the

mineral surface favours the latter process.

The dependence of the reaction rate of a

heterogeneous reaction on the stirrer speed is often

expressed in the power form

dC Na (5.8) Cýd dt

where N is the stirrer speed.

The influence of agitation rate on the dissolution

rate of smithsonite in a baffled reactor can, therefore,

be expressed, for stirrer speeds greater than 2 r. p. s., 47 as, drzn]c4 N0' (5.9) dt