Heavy metal ion activation of zinc sulphide: (1978)

HEAVY METAL ION ACTIVATION

OF ZINC SULPHIDE

A THESIS SUBMITTED BY

JOHN RALSTON

FOR THE DEGREE OF DOCTOR OF PHILOSOPHY

IN THE

UNIVERSITY OF MELBOURNE

JUNE 1978

(i)

ACKNOWLEDGEMENTS

This project has received financial support from the Australian Mineral

Industries Research Association Limited. Sincere thanks are extended to the

sponsors.

To my supervisor, Professor T. W. Healy, I owe a large debt of gratitude

for the encouragement, advice and stimulation that he has provided throughout

the duration of this work.

I should like to warmly thank both the academic and technical staff of

the Department of Physical Chemistry, Melbourne University and of the Depart-

ment of Applied Chemistry, Swinburne College of Technology for their kind

assistance and advice. Peter Alabaster of the Department of Physics, Swin-

burne College, introduced me to the inner workings of a mass spectrometer and

his help is much appreciated.

The members of the research group in Surface and Colloid Chemistry at

Melbourne University displayed good humour, critical scientific attitudes and

friendship during the years in which this work was carried out. I thank them

sincerely.

Mr. W. J. Trahar, CSIRO Division of Chemical Engineering, provided a

copy of his paper on the Natural Flotability of Chalcopyrite prior to its pub-

lication. I am grateful for this and for his thoughtful advice.

Thanks are also due to Pam Bain for her skilful typing of this thesis.

The fact that this work is completed is in no small part due to the

patience and understanding of my wife and small son, who rescued me from the

pits of despair into which an author sometimes sinks in the attempt to create

something worthwhile.

John Ralston.

JUNE 1978

ABSTRACT

This thesis deals with the activation of zinc sulphide by the heavy

metal ions Cu", Cd11 and PbII over a wide range of pH, metal ion concentration,

zinc sulphide concentration and incident light conditions.

At acid to neutral pH values the uptake of a heavy metal ion has been

observed to take place in two stages: a fast, initial step followed by a second,

slower step. Kinetic data have shown that heavy metal ion uptake generally

follows a logarithmic dependence on time.

At alkaline pH values heavy metal ion uptake is complicated by hydrolysis

effects.

A sensitive mass spectrometric technique has been developed and is cap-

able of quantitatively detecting down to 1% of a nominal monolayer of elemental

sulphur on a mineral of surface area 0.7 m2 g-1. Elemental sulphur was detected

on sphalerite surfaces activated by Cu11 and PbII

up to pH 6.6, both in the dark

and under UV irradiation. Elemental sulphur was only detected on CdII activated

and unactivated zinc sulphide surfaces under UV irradiation.

A mechanism for the activation of zinc sulphide is advanced at acid,

neutral and alkaline pH values.At acid to neutral pH values, metal ion uptake

occurs by an exchange reaction

i.e. ZnSsolid + Maq (Zn,M)Ssurface + Zn

aq2+

This is coupled with a surface redox process, which is linked with the semi-

conductor properties of zinc sulphide. As a result of this surface redox pro-

cess, elemental sulphur may form on the zinc sulphide surface, depending on the

type of metal ion involved and on the incident light conditions. It is proposed

that the rate determining step in the exchange reaction is the transfer of an

adsorbed metal ion from a surface, adsorbed site to a lattice site. This leads

to a logarithmic dependence of metal ion uptake on time, as observed experiment-

ally.

At alkaline pH values, the mechanism at acid to neutral pH values is

complicated by metal ion hydrolysis, adsorption and precipitation effects. An

overall mechanism is proposed which accounts for the observed behaviour.

(iv)

CHAPTER 1.

INDEX

OUTLINE OF THE THESIS Page 1

2

2

3

3

4

9

10

12

13

13

15

15

22

22

25

27

28

28

28

31

33

33

33

36

36

38

44

CHAPTER 2. A REVIEW OF THE ACTIVATION, SOLUBILITY, OXIDATION AND

2.1

2.2

2.3

2.4

2.5

CHAPTER 3.

SEMICONDUCTOR PROPERTIES OF ZINC SULPHIDE

Introduction

Heavy metal ion activation of zinc sulphide

(a) General features

(b) Kinetics

(c) Mechanism

(d) The nature of the activation products

(e) "Collectorless" flotation

The solubility of ZnS

(a) The solubility products of heavy metal sulphides

(b) Solubility and hydrolysis equations for zinc sulphide

Oxidation of sulphide minerals

The semiconductor properties of sulphides

(a) The structure of semiconductors

(b) Charge transfer at the semiconductor-water interface

(c) The photovoltaic effect in sulphide semiconductors

EXPERIMENTAL METHODS

3.1

3.2

3.3

Kinetics experiments

(a) Apparatus

(b) Procedure

Materials

(a) Water quality

(b) Chemicals

Analysis

(a) The type of ion selective electrodes used

(b) M2+ electrode performance

(c) AAS analysis

(v)

3.4 Subsidiary tests 48

(a) Pretreatment experiments 48

(b) Adsorption on filters 52

(c) Presence of H2Sgas and SO42aqueous 52

(d) Eh measurements 52

3.5 The determination of elemental sulphur 53

(a) Introduction 53

(b) Mass spectrometry 55

(c) Experimental procedure 60

CHAPTER 4. THE ACTIVATION OF ZnS WITH CuII 65

4.1 Introduction 65

4.2 Results 65

4.2.1 CuII uptake 65

(a) Preliminary work 65

(b) Results 69

4.2.2 Exchange ratio, ZnII release and pH change 102

(a) Method of data presentation 102

(b) Summary of major results 111

(c) Subsidiary effects 113

4.3 Rate equations 113

4.3.1 Acid pH (4.0 to 6.5) 113

4.3.2 Neutral and alkaline pH 120

4.3.3 Activation energy 120

CHAPTER 5. THE ACTIVATION OF ZnS WITH CdII 122

5.1 Introduction 122

5.2 Results 122

5.2.1 CdII uptake 122


(b) Results 124


(a) Method of data presentation

138

(b) Results

139

(vi)


5.3.1 Initial, rapid activation step 146

5.3.2 Second, slow activation step 150


CHAPTER 6. THE ACTIVATION OF ZnS WITH PbII 152

6.1 Introduction 152

6.2 Results 152

6.2.1 PbII uptake 152


(b) Results 154


(a) Method of data presentation 168

(b) Results 173


6.3.1 Initial, rapid activation step 175

6.3.2 Second, slow activation step 178


CHAPTER 7. MASS SPECTROMETRIC DETERMINATION OF ELEMENTAL

SULPHUR ON ZnS SURFACES

180

7.1 Identification of elemental sulphur 180

(a) Identification of major ions 180

(b) Identification of elemental sulphur through isotopic 182 abundance

7.2 Quantitative determination of elemental sulphur 184

(a) Analysis technique 184

(b) Detection limit 187

(c) Blank experiments 189

(d) CdII activation experiments 189

(e) Cuii activation experiments 189

(f) PbII activation experiments 193

7.3 Sample mass spectra 194

(vii)

CHAPTER 8. INTERPRETATION 198

8.1 On the stability of elemental sulphur (S°) 198

8.2 The significance of the logarithmic law 206

(a) Comparison with other activation studies 206

(b) The link with mechanism 208

8.3 Free energy calculations 210

(a) Oxidation 210

(b) Exchange reactions 211

(c) Selected redox reactions producing elemental 212

sulphur

8.4 A comparison between Cu", CdII and PbII 214

(a) Relative rate of uptake of CuII, CdII and PbII 214

(b) Correlation between MII uptake, quantity of S° 218

detected, R values and pH changes

8.5 Solid state properties and surface redox behaviour 220

(a) General remarks 220

(b) CuII 221

(c) The band gap at the surface of ZnS - the effect 222

of CuII, CdII and PbII

(d) The formation of S° by a surface state mechanism 228

8.6 pH changes

239

(a) Blank experiments 239

(b) Activation experiments 241

8.7 Overall mechanism at acid pH values 242

(a) Exchange reaction 242

(b) Surface redox behaviour 243

(c) Model for the logarithmic law 243

(d) Magnitude of heavy metal ion uptake in other studies 250

8.8 Overall mechanism at near neutral to alkaline pH values 251

(a) Blank experiments 251

(b) Electrokinetic studies 253

(c) Mechanism 256

8.9 Subsidiary effects

(a) Influence of So on M

II uptake - further comments

(b) Effect of pretreatment on subsequent Cu II

activation

(c) Fundamental solubility considerations

8.10 Summary

261

261

261

263

263

REFERENCES 269

CHAPTER 1. OUTLINE OF THE THESIS.

This thesis is divided into seven main chapters, the sequence and

organization of which require some preliminary comments.

Chapter 2 is a review chapter, concentrating mainly on existing know-

ledge relating to the heavy metal ion activation of zinc sulphide. The sol-

ubility characteristics of zinc sulphide are also considered here. Key points

regarding the oxidation and semiconductor properties of zinc sulphide are then

briefly summarized, for adequate reviews and textbooks already exist in these

areas.

Chapter 3 deals solely with experimental methods relating to kinetics

experiments, analytical techniques and mass spectrometry.

Chapters 4, 5 and 6 present, in sequence, the kinetics data and experi-

mental rate laws relating to the uptake of CuII, CdII and PbII by ZnS.

Chapter 7 deals specifically with the mass spectrometric determination

of elemental sulphur on ZnS surfaces, both from qualitative and quantitative

points of view.

Chapter 8 interprets the experimental results. In this chapter a

mechanism for the heavy metal ion activation of zinc sulphide is proposed, re-

lated to relevant literature studies and the main features are then summarised

in three tables at the end of the chapter.

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CHAPTER 2.

A REVIEW OF THE ACTIVATION, SOLUBILITY,

OXIDATION AND SEMICONDUCTOR PROPERTIES

OF ZINC SULPHIDE.

2.1. INTRODUCTION

The origins of the discovery of copper sulphate as a flotation activator

for sphalerite (cubic zinc sulphide) are rather obscure. According to Ralston

et al (117) its favourable action was discovered accidentally at a mill in

Mascot, Tennessee (U.S.A.) early in 1914. Laboratory test work on a sphalerite

ore yielded good flotation results and, on the basis of these results, a mill

was constructed. The subsequent flotation concentration of the ore was unfor-

tunately very poor. In searching for a reason for this discrepancy between the

laboratory and mill tests it was noticed that the laboratory test machine was

manufactured from bronze whereas the commercial mill machine allowed the flot-

ation pulp to contact iron and wood only. A plate of copper sheet hung in the

mill machine immediately improved the flotation results. Replacing the copper

plate with a small amount of copper sulphate, introduced into the flotation pulp,

resulted in an even more favourable flotation recovery of sphalerite. By the

end of 1914 copper sulphate was used in the acid flotation circuit of the Butte

and Superior Copper Company (U.S.A.). The person or persons responsible for

its discovery remain unknown and the discovery was not patented, despite its

great importance.

Copper sulphate, the cheapest available form of copper, is today common-

ly used as an activator in sulphide flotation circuits, activating sulphide min-

erals other than sphalerite e.g. chalcopyrite (1,2,4). Copper sulphate additions

of between 0.5 to 2.0 kg per tonne of ore treated are frequently used and

difficulties are still encountered in optimizing activator addition (4).

Apart from the vast and conflicting literature dealing with the role of

xanthates in sulphide flotation (e.g. 1, 119, 120, 121) there is also a

significant body of evidence concerning the mechanism and kinetics of heavy

metal, particularly copper, ion activation of zinc sulphide (e.g. 36, 44).

Yet comments such as ".... in spite of the lack of understanding of the

mechanisms underlying activation" (118) are all too common in recent liter-

ature. There is therefore a definite need to understand the kinetics and

mechanism of heavy metal ion activation of zinc sulphide, in particular, both

for practical and fundamental reasons. It is therefore the purpose of this

thesis to try and clarify the problems concerning the heavy metal ion activat-

ion of zinc sulphide.

2.2. HEAVY METAL ION ACTIVATION OF ZnS.

2.2(a) General features.

Sutherland and Wark tested some 26 metals in order to assess their activ-

ation behaviour with respect to zinc sulphide (128). Only those metals whose

sulphides were less soluble than zinc sulphide were found to be effective activ-

ators, with the exception of TZ+ and Sn2+. The latter hydrolyses readily at

the pH values used in the tests by Sutherland and Wark (129) whilst the anoma-

lous behaviour of TQ+ still remains obscure (55).

As the solubility of heavy metal sulphides is progressively reduced below

that of zinc sulphide, the metal ion becomes an increasingly better activator.

This solubility behaviour has generally been widely used as a criterion of a

metal ion's activation ability by many authors (36,97,122,124,131). Studies

dealing with the kinetics and mechanism of zinc sulphide activation have, in the

main, concentrated on copper, lead and silver as metal ion activators (36,97,98).

A recent review of the flotation of zinc sulphide has dealt in part with metal

ion activation (55).

2.2(b) Kinetics.

(I) CuII (note comment in the first paragraph of Chapter 4).

Most workers (25,26,36 to 39) who have studied the activation of ZnS

II by Cu identified two stages in the activation reaction: a rapid initial

step, completed within 1 to 5 minutes,forming 2 or more monolayers of CuII

followed by a second, slower step. Gaudin, Fuerstenau and Mao's results (36)

demonstrated that the system was not at equilibrium after 63 hours of contact

time whereas Pomianowski et al(125) and Wada and Okada (26) suggest that the

slower second step continues until equilibrium is reached. There is no con-

clusive experimental evidence demonstrating that CuII uptake stops while CuII

ions are present in solution.

Initial rapid step. Considerable confusion exists regarding the kinetics of

activation of ZnS by CuII during the initial stage of activation at acid to

neutral pH values. CuII uptake has generally been found to be dependent on the

initial CuII concentration (25, 36 to 39). First order kinetics involving the

concentration of copper together with the surface area of the solid is most

commonly proposed (25, 37 to 39):

dt = k A [CuII]

although Bazanova and Mitrofanov (38,39) also advanced a logarithmic relation-

ship:

r a + k l log t

where r is the CuII uptake per gram of solid in time t, A is the surface area

of the solid and a, k, kl are constants.

Generally these investigators have worked with natural, crushed sphaler-

ite samples of varying degrees of purity and of low surface area (frequently less

than 0.05 m2 g-1), the latter being determined by the BET method. Apart from

the errors inherent in determining such a low surface area (40, 41), the actual

-4-

-5-

surface area available for reaction in their experiments has often been low

(e.g. 450 cm2 in the study by Bazanova and Mitrofanov(38)). Loss of copper

by adsorption on to container walls etc. can then become an important problem,

particularly when CuII hydrolysis products are present. It is rare, indeed,

to find any mention of precision in these papers, except for some passing

comment in the work by Gaudin et al (Table 1, reference 36), leading the read-

er to treat the interpretations advanced with caution. For example 1.n Figure 2

of Mitrofanov and Bazanova's study of the adsorption of cupric ion by sphaler-

ite (37), there are two very divergent points below the 1 minute mark in their

Cu2+ adsorption curve at pH 5.7, yet 65% of the measured adsorption has occurred

in one minute. The general shape of their Cut+ uptake curve is certainly

logarithmic (e.g. 29) but the scatter in Mitrofanov and Bazanova's data meant

that there would be, of course, considerable confusion in deciding what rate

law was obeyed during this initial, rapid step. This point is borne out by the

two rate laws (first order and logarithmic) which were proposed in their sub-

sequent papers (38, 39).

Dixon's limited study (42, 43) of the activation of synthetic zinc

sulphide with CuII at pH 3 merits some comments here. His results indicated

that there was a rapid, single stage uptake of Cu". The measured uptake values

were far greater than reported here, indicating that more than 20 monolayers of

CuII were abstracted in about 3 minutes. These results were based on only 2

separate CuII uptake values, however (42) Dixon makes no comment as to why his

measured uptake values differed from those of other workers. The present author

feels that no useful conclusions can be drawn from this study, in view of its

very limited scope.

Gaudin et al (36) studied the abstraction of CuII by sphalerite at pH 6,

using a radiotracer technique to follow CuII uptake. Their first reading was

taken after a reaction time of 5 minutes which allowed these workers to suggest

-6-

the existence of a fast initial surface reaction, but not to determine its

rate law.

Mitrofanov and Bazanova (37-9) and Girczys et al (44) have studied CuII

activation of ZnS at alkaline pH values, particularly at pH 10. Radiotracer

techniques were used to detect the amount of Cuis adsorbed on the ZnS surface.

The adsorption data obtained by Girczys et al (44) is not determined as a funct-

ion of time and therefore does not allow a rate law to be formulated. Mitro-

fanov and Bazanova (37) worked with CuII concentrations in the range

1.9 x 10 5 to 6.3 x 10 -4 M with a solids:liquids ratio of 1:100. They main-

tain that the formation of visible precipitates was precluded at these low

concentrations. The measured CuII uptake above pH 8 ranges from 1 to 3 mg CuII

per g ZnS x 102. Calculation shows that Cull residual concentration ranges

from 1.4 x 10-5 M to 6.3 x 10-4 M.Figixre 4.1(a) of this present study clearly

shows that CuII concentrations in this region will result in precipitation of

Cu(OH)2solid above pH 6.5 to 7. Mitrofanov and Bazanova state that visible

precipitates were removed by filtration. Light scattering studies by the pres-

ent author (129) and by James (130) have demonstrated that precipitates which

may not be visible to the naked, unaided eye do, however, scatter the incident

light of a laser beam (detected with a sensitive photodiode). At best Mitro-

fanov and Bazanova may have removed Cu(OH)2solid. Other CuII hydrolysis species

which are dominant above pH 8 will still be present however. Therefore the CuII

uptake data obtained by these workers above pH 8 do not represent the exclusive

interaction of Cu2+ with sphalerite, as their paper suggests, nor does it en-

able a rate law describing this interaction to be formulated.

Second, slower step.

The few workers (25, 26, 36) who have studied the kinetics of the second

stage of activation concluded that CuII

uptake follows a parabolic rate law:

r2 = kt + c (k, c are constants).

-7

These workers maintain that the rate controlling step is diffusion through the

solid copper sulphide film, but make no comment on the fact that rate constant

k in the expression above is dependent on both the diffusion coefficient and

the concentration of the diffusing species(22). This suggested rate law is,

therefore, somewhat tentative.

Activation energy.

Mukai and Nakahiro (25) determined an activation energy of 23 kJ mol-1

for the initial activation step and associated it with diffusion control. This

result is in marked contrast to Wada and Okada's (26) observations that temp-

erature has virtually no influence on reaction rate.

Normally an activation energy of 25 kJ mol -1 for processes occurring in

aqueous solution is taken to be the demarcation between chemical and diffusion

control (31-34) i.e. a value in excess of 25 kJ mol -1 corresponds to chemical

control whilst a value less than 25 kJ mol -1 can be taken to mean that a re-

action is diffusion controlled. This simple criterion fails rather badly, how-

ever, in reactions which are subject to mixed control i.e. where the slow

chemical step and the slow diffusion step are comparable in magnitude. Moreover

in studies of metal displacement reactions Miller and Beckstead (35) have shown

that a high (42 kJ mol 1) activation energy may be obtained for a diffusion-

controlled reaction if the form of the deposit is temperature sensitive. Cer-

tainly for reactions which are subject to diffusion control, the expected rate

law shows a first order dependence on the diffusing species (34), however diffus-

ion controlled reactions are also sensitive to changes in stirring rate and

ionic strength (31, 34, 35), which were apparently not tested by Mukai and

Nakahiro (25).

Mukai and Nakahiro determined an activation energy of 36 kJ mole-1 for

the second stage of activation whereas Wada and Okada (26) obtained a value of

58 kJ mole-1 and associated the value with "structure sensitive diffusion". No

further interpretation was given.

-8-

No firm conclusions may readily be drawn from these measured activation

energies.

(II) PbII.

Fuerstenau and Metzger (97) have studied the kinetics of activation of

ZnS by PbII at an unspecified pH (probably around 5.5 to 6) using a radiotracer

technique. PbII uptake (I') was found to obey the following equation:

r = tn k

where t is the time in minutes; k and n are constants for a given initial

PbII concentration and n is generally around 0.28. No further interpretation

of this rate law was given (97).

II In contrast to Cu activation, Fuerstenau and Metzger identified only

a single stage reaction over the time span 2 to 10,000 seconds, with PbII up-

take reaching monolayer coverage between 15 to about 1000 seconds, depending

on the [PbII]/[ZnS] ratio. No temperature variation was reported. The surface

area available for activation in this study was about 0.75 m2, which should be

sufficient to give reliable abstraction data.

(III) AgI

Gaudin, Spedden and Corriveau (127) studied the uptake of AgI by ZnS

and identified a two stage reaction: a rapid, initial step followed by a slower

step. No rate law was proposed. In a subsequent paper, Gaudin, Fuerstenau and

I Turkanis (98) investigated Ag activation of ZnS at an unspecified pH (probably

5.5 to 6) with some experiments performed in the presence of cyanide at pH 9.

They identified only a single stage reaction and made no comparison with the

previous study (127). At high AgI concentrations, AgI uptake in mole g-1 (i)

was found to follow a logarithmic dependence oto time:

r = l log10t + I'1 k

where F1 is the mole of AgI

abstracted per gram of ZnS in 1 minute, t is the

time in minutes and k is a constant for a given set of conditions. Up to 20

monolayers of AgI were abstracted. Differences in initial rates were assigned

to boundary layer diffusion whilst the origin of the logarithmic law was ten-

tatively attributed to the existence of microcracks in the silver sulphide lay-

er (98). No temperature variation was reported. The surface area available for

activation in this study was only 0.063 m2 which may account for some of the

scatter apparent in the experimental AgI uptake data.

2.2(c) Mechanism.

Few workers have considered the question of mechanism (36,44,55,117) in

any depth at all. In acid to near neutral pH values it is generally accepted

that the overall reaction is an exchange reaction of the type:

M2+ + ZnS —+ MS + Zn2+

aq surface f surface aq

occurring at the ZnS surface. The evidence advanced for this proposed reaction

rests on two key points: Firstly the ratio of M2+abstracted

to Zn2+ released

is unity although the precision of much of the experimental data is questionable.

Secondly this exchange reaction has been shown to be thermodynamically favour-

able (e.g. 36, 55, 122). Kraus and Phillips (49) support this exchange react-

ion in their ion exchange study of the uptake of metal ions by amorphous,

freshly precipitated heavy metal sulphides at pH values near 3. These authors

found very high degrees of conversion of the amorphous sulphides to more

insoluble sulphides but were unable to satisfactorily explain their results.

Pomianowski et al's (125) results are interpreted in terms of an exchange

reaction, however they also noticed a substantial pH rise during the Cu" act-

ivation of sphalerite, but did not identify the cause.

Girczys, Laskowski and Lekki (44) propose that on the basis of the

different flotation behaviour of ZnS,activated by copper sulphate:at acid,

neutral and alkaline pH values, the following mechanism is applicable:

acid solution

ZnS +

Cuâg f Cussurface

+ Znäq

neutral pH

Zn-OH surface

Cu(OH) -O-Cu + H20 surface aq Fsurface

0

alkaline pH

ZnS + 20H [Zn(OH)2 + sl surface state

[Zn(OH)2 + S) surface state + Cuaq F Zn(OH)2 surface +Cus

surface " •(ii)

Their proposed mechanism at neutral and alkaline pH values is not supported by

any hard evidence and is purely speculative. At alkaline pH values the known

instability of elemental sulphur (14) suggests that (i) and (ii) are not viable.

Finkelstein and Allison (55) envisage two alternative mechanisms for the

exchange reaction at acid to neutral pH values,during CuII activation of ZnS,

but are unable to distinguish between them:

(a) adsorption, ion exchange, desorption

InSsurface + Cuaq <- - Ssurface Cu2+ CuSsurface + aq

(b) dissolution, surface-nucleated precipitation

qZnS +Cu aq

aqsurface +S +Cu f ZnS•CuSsurface+ surface â âq aq

The mechanism of metal ion activation of ZnS is obviously in need of

clarification.

2.2(d) The nature of the activation products.

The identification of the reaction products formed at the surface of

activated ZnS has received increasing attention (69,76,77,78,100,124). Cooke

(124) exposed sphalerite to 0.1 M solutions of Pb(NO3)2,AgNO3, CuC12 and HgCl2

for 50 days at 100 C. X-ray diffraction of the treated sphalerite revealed a

-10-

thick coating of covelline (CuS) while, under the same conditions the

sphalerite was totally replaced by Ag2S. For PbII and HgII there was very

little, if any, surface alteration. The experimental conditions used by Cooke

are far removed from those found in a flotation pulp so that assuming the same

product to be present in both cases is questionable.

Sato (100) used an electron diffraction technique to study the surfaces

of ZnS activated by CuII and AgI. While he was unable to draw any firm con-

clusions in the case of CuII activation Satodid identify some elemental silver

on the surface of ZnS activated by AgI.

Finkelstein and coworkers (76,77) have developed a gas chromatographic

technique for determining elemental sulphur on sulphide mineral surfaces.In none

of their systems, including those in which ZnS was activated with Cu", AgI

or Pb", was elemental sulphur detected.

Clifford and coworkers (69,78) have recently obtained some interesting

results in their attempts to characterize sulphide mineral surfaces in froth

flotation systems. These workers have used Electron Spectroscopy for Chemical

Analysis (ESCA) to identify the reaction products on activated ZnS and other

surfaces. At present the technique is only semiquantitative and there is some

concern that the species being detected may actually be formed by X-ray photo-

reduction in the spectrometer, although the authors maintain that this latter

process is apparently slower than the time which is necessary to actually per-

form the analysis (78). Nevertheless Clifford et al(69,78) have detected

both Cu' and elemental sulphur on the ZnS surface. Some recent electron spin re-

sonance spectroscopy (ESR) work by Storey and Platt (54) has shown that CuII

apparently exists on CuII activated sphalerite surfaces with as little as 10%

nominal monolayer coverage. ESR spectroscopy is insensitive to Cu', however,

so that it cannot be said that CuII is exclusively present. The results of

Clifford et al (69,78) together with those of Storey and Platt (54) suggest

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that Cu', CuII and elemental sulphur may well be present on CuII

activated

ZnS surfaces.

2.2(e) "Collectorless" flotation.

The flotation of sphalerite without the use of collectors has recently

experienced an upsurge of interest (86,134,135). Sphalerite has, of course,

been floated by the Potter process for many years (157). In this case flot-

ation is attributed to the presence of elemental sulphur, formed on the

sphalerite surface following reaction with a hot concentrated solution of nit-

ric and sulphuric acids.

Finkelstein and coworkers (85,135) activated sphalerite with copper sul-

phate under oxygen-free conditions at pH 8 and then carried out a series of

flotation tests in the pH range 5 to 11. Enhanced flotation recoveries of

sphalerite (50 to 70%) were observed over unactivated sphalerite. KMnO4 and

K2Cr2O7 were found to be effective depressants. The authors carried out exten-

sive tests in an attempt to detect elemental sulphur on the activated sphalerite

surface (85,105), but did not find any elemental sulphur present. No firm con-

clusions were reached regarding either the cause of this enhanced flotability

or the reasons why KMnO4 and K2Cr2O7 functioned as depressants.

Lepetic (134) floated chalcopyrite at pH 6 after dryjautogeneous grind-

ing using a small amount of frother as the only flotation reagent. Lepetic

suggested that molecular oxygen, adsorbed during dry grinding, accounted for

the observed flotability. Heyes and Trahar (86), in their recent paper on the

natural flotability of chalcopyrite, disputed Lepetic's interpretation and

conclusively demonstrated that the presence of oxygen was not a prerequisite

for the observed flotability of chalcopyrite. Heyes and Trahar worked with

carefully cleaned chalcopyrite and carried out extensive flotation tests at

pH 11 in the presence of polypropylene glycol as a frother (86). These authors

showed that chalcopyrite displayed flotability in an oxidizing environment

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(e.g. copper sulphate) and non-flotability in a reducing environment (e.g.

sodium dithionite). Furthermore this flotability was found to be rapidly

reversible as either an oxidant or reductant was added to the flotation cell.

Heyes and Trahar did not regard elemental sulphur as being responsible for

their observed flotability and stated that (86):

"While it is not possible at present to explain this behaviour in

terms of the fundamental principles involved, it appears that the sur-

face is rendered hydrophobic in an oxidizing environment and hydro-

philic in a reducing environment."

Whilst this investigation by Heyes and Trahar (86) contains some inter-

esting implications, much fundamental work remains to be done concerning

collectorless flotation and natural flotability.

2.3. THE SOLUBILITY OF ZnS.

2.3(a) The solubility products of heavy metal sulphides.

In principle, metal sulphides react with water to yield the hydrated

metal cations and sulphide ions. These sulphide ions are strong bases and are

hydrolyzed in aqueous solution, forming HS and H2S. The equilibrium concen-

tration of S2 ions in solution and, consequently, the solubility of the metal

sulphide is strongly dependent on the pH of the solution. The solubility pro-

ducts of the metal sulphides can be directly correlated with the electronic

configuration and charge of the metal cation (16,20,99). Solubility products

for various heavy metal sulphides, which are of primary concern in this present

investigation, are listed in Table 2.1.

Sulphides of heavy metal cations with the dl° electronic configuration

have smaller solubility products than those with s2p6 cations. This enhanced

stability is conventionally attributed (16,20,99) to back-donation of electrons

from the filled d orbitals of the cation to the empty d orbitals of the sul-

phide ion. Back-donation is particularly important if the cation has a low

-14-

charge density e.g. Cu+, Hg2+ and Ag+ have a very large affinity for sulphur.

The solubility products of transition metal sulphides where the metal

cation has the ndm (1 < m < 8) electronic configuration decrease with increas-

ing m of d electrons and with increasing n, where n is the principal quantum

number (e.g. Mn2+, Ni2+). Again this is due to back-donation of the d elect-

rons, which also accounts for the low solubility products of sulphides of

cations with the s2 electronic configuration such as Pb2+ (16,20,99).

TABLE 2.1. Solubility products for heavy metal sulphides

at 298 K.

The solubility product is defined for the reaction

2+ 2- MSSsolid M + S aq aq

KSO = aM2.4. • aS2_ (ai is the activity of species,i) aq aq

Sulphide pKSO

MnS 14.0

FeS (pyrite) 17.8

NiS 19.6

ZnS (cubic) 24.1

CdS 27.1

PbS 28.5

CuS 36.1

HgS (cinnabar) 52.1

N.B. reported pKSO values vary considerably, according

to their source (6,42,136,137). The values quoted

in Table 2.1 represent an average of a wide range

of values and are taken from ref. 42, 137. The

solid forms are, unfortunately,frequently not

specified.

2.3(b) Solubility and hydrolysis equations for zinc sulphide.

The basis for solubility calculations in systems containing complex

equilibria is well documented (5,7,91,137) and only brief, pertinent comments

are made here.

In calculating the solubility of ZnS, hydrolysis of both the metal and

sulphide ions must be considered, as well as soluble complexes of ZnS (138).

The relevant equilibria, together with their equilibrium constants are shown

in Table 2.2. •••

Thiocomplexes of the form M(SH)2 x are reported to exist for a number

of sulphides (5,138), however no information is apparently available for ZnS.

Furthermore polysulphides (S2 ) are known to exist in alkaline, aqueous sol-

utions (15,16,91) forming a complex dynamic equilibrium involving sulphides of

various chain lengths. Their concentrations are generally small (15) until

the aqueous solution becomes strongly alkaline (16) and may be neglected in

ZnS solubility calculations, at least in the pH range from 2 to 12.

The equilibrium solubility of ZnS is a closed system as a function of

pH is shown in Figure 2.1, calculated using the thermodynamic data and methods

given in Table 2.2. The diagram shows that the solubility of ZnS is essentially

constant over the pH range from 5 to 12, increasing with decreasing pH below

pH 5.

A diagram such as Figure 2.1 does not take into account oxidation effects

nor the marked increase in solubility of ZnS which is observed under UV irrad-

iation (43,63). In view of this it should be used as a guide to, rather than

as an absolute measure of, ZnS solubility.

2.4. OXIDATION OF SULPHIDE MINERALS.

The kinetics and mechanism of the oxidation of sulphide minerals have

been the subjects of extensive studies and reviews (141 to 155). Only the

6 10 12 8

o

2

6

pZn3r

8 —

10 —

12

14 ■

16

4

ZnS (solid)

Zn2+

/I 11I1111 11111111111I1I11 11

-16-

pH

FIGURE 2.1 The equilibrium solubility of ZnS in a closed system as a function of pH.

+

} +

� +

} +

Ka2,

aS2- _ 1 + [H+]

where + [H+]2 1-1

Kai Ka1Ka2

-17-

TABLE 2.2.

Notation solution

Solubility and hydrolysis data for ZnS at 298K

in a closed system.

according to Sillen and Martell (6). All species are in aqueous (aq) unless otherwise noted. S denotes solid phase.

(Data from refs. 6, 7, 136).

Equilibria log10 (constant)

Zn2+ + S2 ZnS s

ZnOs + H20

Zn2+

+ H20

Zn(OH) + H20

Zh(OH)2aq+ H20

Zn(OH) 3 + H20

ZnS + H2O

H2S

HS +

Zn2+ + 20H

Zn(OH) + + H+

f Zn(OH)2a q -> Zn(OH)3 + H+ +

+ Zn(OH)42 + H+

ZnSHOHaq

+ H + HS

S2 + H+

KSO, -26.9 (note 2)

KSO, -16.6

*Ki, - 9.5

*K2, - 6.5

*K3, -12.3

*K4, -12.8

- 5.87 (note 1)

- 7.0

-17.1 (note 2)

+H+

Method of calculation.

For a closed system (no loss of H2Sgas), total soluble zinc equals

total soluble sulphur. i.e. [Zn]T = [S]T

[Zn]T [Zn2+lazn2+

where aZn2+ = 1 + *K1 + *K*K2 + *K*K*K3 + *K*K*K*K4

1

[H+) + [H+] 3 [H] 4

Similarly [S]T = [S2 -)a-1_

thus log10 [Zn2+] = [-26.9 + log10aZnp+ -logaS21.I

Similarly for other species.

Note 1: K from reference 138. Sil

Note 2: This value is taken from the careful study in Ellis and Giggenbach (140) and results in the for sulphides increasing by about 3 (136,140) pK soof 24.1 becomes 26.9 for ZnS.

oxygen-free solutions by "accepted" pK values e.g. in Tableso2.1, the

-18-

major findings are summarised here.

It is generally accepted (e.g. 143,144,146) that the first step in the

oxidation of sulphides corresponds to

MS --> M2+ + S° + 2e 1--

which is coupled with reduction reactions such as

3102 + 2H+ + 2e H2O

or 2H20 + 2e H2 + 20H

so that in acid, aqueous solutions the overall reaction for the first step

corresponds to

MSsolid

+ i02aq + 2Haq Maq + Ssolid + H2Oliquid

Elemental sulphur (So) is a product of this first step. S0 is known to become

increasingly unstable with increasing pH, ultimately decomposing to form sul-

phate (14,143, Chapter 8). Plante and Sutherland (14) studied the oxidation

products of galena, pyrite, chalcopyrite and sphalerite in neutral and alkaline

solutions. Their results are summarised in Tables 2.3 and 2.4. For sphaler-

ite, thio-salts were found to be the principal oxidation products at pH 6 with

sulphate as a minor product. Both thio-salts and sulphate were found in

increased quantities at pH 10. Eadington and Prosser (146) found that the

principal oxidation product of galena depended on solution pH and the time of

exposure to oxygen. After many hours of oxidation they found that the princi-

pal oxidation products were sulphur in acid solutions, sulphate in neutral

solution and thiosulphate in alkaline. Clearly the type of oxidation product

depends on the particular mineral being studied, as is obvious from this latter

study, the work of Plante and Sutherland (141) and numerous other investigations

(142,145,149).

The kinetics of oxidation of sulphides has received a considerable

-19-

TABLE 2.3.

Reaction products formed by the oxidation of

aqueous suspensions of sulphide minerals.

(ref 141).

Mineral Products from originally neutral (pH 6) suspension

Products from originally alkaline (pH 10 to 11) suspension.

galena present: Cu2+,Fe2+,Fe3+, present: H+,Ag+,Pb2+,SO4 , 2-

SO4

absent : Pb2+,polythionate

thio-salts

absent : Fe2+,Fe3+,Cu2+

pyrite present: Fe2+, H+, SO4 -

present: H+ ,SO4

2 ,S3062-,

S4062-,S2032-,S032- O 2 .SO 2 32 ,SO 32 4 6

absent : Fe3+,Fe2+,C1.12+,S2 ,

absent : Fe3+,polythionate

55062-,52062-

chalcopyrite present: H+,Cu2+,Fe2+,Fe3+

S042- O42

absent : polythionate

present: H+,SO42 ,S2O32 ,

S4O62

absent : Cu2+,Fe2+,Fe3+,5032

sphalerite •resent: S042-,thio-salts present: H+,Zn2+,SO42 ,5032

absent : Zn2+, Fe2+, Fe3+, absent : S4062 -,S2032-

Cu2+

Oxidation product Concentration (mg per litre) of product from

pyrite chalcopyrite sphalerite galena

5042

thio-salts

Cu2+

Fe 3+

Feg+

Pb2+,Zn2+

final pH

30

nil

0.01

15

3.4

nil

4.2

1

24

2.6

2

24

nil

nil

nil

nil

6.0

(H) in alkaline solution, initial pH = 10.1

SO42

thio-salts

5032

82032

2+ 2+ 2+ 2+ ,Fe ,Cu ,Zn

Ag+,Pb2+

final pH

40

60

nil

nil

nil

5.3

20

90

nil

5

nil

7.2

30

70

nil

9.6

30

30

nil

3.6,1.1

5.5

1

nil

6.0

-20-

TABLE 2.4.

Quantity of soluble oxidation products formed

by oxidation of aqueous suspensions of sulphide

minerals. (ref. 141)

(A) in distilled water, initial pH = 6

Conditions: temperature: 288 to 291 K

-36 mesh mineral, deslimed.

pulp density: 1 part mineral, 5 parts water

CO2-free air bubbled through suspension for 5 days.

-21-

amount of attention (147,148,149). The rate of oxidation has been found to

depend on the surface area available for reaction and the partial pressure of

oxygen (147), the type and composition of the sulphide mineral (142), solution

pH (154) and temperature (155). The oxidation rate of complex sulphide ores

has been found to be enhanced if pyrite is present (145). Thus in order to

study the oxidation rates of sulphide minerals, it is of critical importance

to control the mineral composition, solution pH and partial pressure of oxygen

if any meaningful rate law is to be developed.

While they have not developed a rate law, Eadington and Prosser (146)

have carried out a very careful study of the rate of oxidation of high purity,

synthetic lead sulphide under controlled oxygen atmospheres of between 20 and

76 cm Hg. These authors found significant differences in the behaviour of

freshly precipitated lead sulphide and lead sulphide which had been stored for

several weeks in water under "oxygen free" nitrogen (static). This latter

material was called an "aged precipitate" and proceeded to oxidize at a lin-

ear rate immediately it was placed in the reaction vessel. In contrast, when

freshly precipitated lead sulphide was placed in the reaction vessel it showed

a long induction time of up to 40 hours (and not less than 5 hours), during

which no oxygen uptake could be detected (i.e.<1 x 10-6 g mole -1 PbS). Ead-

ington and Prosser attributed this induction period, where no oxidation pro-

ducts on the PbS powder or in solution could be detected, as a stage where the

outer layers of the PbS become supersaturated with sulphur due to the inter-

action of chemisorbed oxygen atoms and PbS according to:

PbS + 0(surface)

+ H20 -4-P b2+ +

S(PbS) + 20Haq

After this induction period has been completed, measurable oxidation begins

with the formation of an elemental sulphur phase at points on the surface and

the reduction of the sulphur content of the adjacent PbS. Eadington and

Prosser concluded that the reason for the absence of an induction time for

aged precipitates was that prolonged storage under "oxygen-free" nitrogen was

not, in fact, in the absence of oxygen, the induction period essentially hav-

ing taken place during storage.

Moignard, Dixon and Healy (9,150) have studied the electrokinetic

properties of high purity synthetic zinc sulphide (99.99% pure) in the pres-

ence and absence of cupric nitrate, natural marmatite containing 5.4% Fe and

11% Fe and synthetic nickel sulphide (> 96% pure). All of the experimental

work was performed under "oxygen-free" conditions. Pre-equilibration of ZnS

marmatite and nickel sulphide at acid pH values of 3 to 4 or less for 17 hours

yielded an isoelectric point of about pH 2.5 for ZnS, the natural marmatite

and NiS, as well as for ZnS subsequently activated with cupric ion at pH 4.

The authors attributed this common, low pH, of 2.5 to a partial surface iep

coating of elemental sulphur, formed according to (9):

MS 4--- M2+ + Sc + 2e

with 402 + 2H+ + 2e H2O

or 2H2O + 2e H2 + 20H -

or Cut+ + e —i- Cu+ 4---

Their conclusion, that exposure of a sulphide mineral to an aqueous system

containing "some" oxygen for long periods to time, results in the formation

of elemental sulphur, is consistent with the findings of Eadington and Prosser

(146) .

In summary, then, published evidence indicates that the rate of oxid-

ation of high purity, freshly formed synthetic sulphides is likely to be slow,

particularly in the presence of trace oxygen concentrations.

2.5. THE SEMICONDUCTOR PROPERTIES OF SULPHIDES.

2.5(a) The structure of semiconductors.

The quantum-mechanical description of solids is firmly established

-23-

(22,70,96,102,103,108). The salient features are presented briefly here. In

a solid, due to the overlap of atomic orbitals, the single electronic energy

levels from each constituent atom are combined and broadened into bands of

very closely spaced energy levels. For a semiconductor, the bands which are

of key importance are the band occupied by the valence electrons in the solid

and the next excited state. The band occupied by the valence electrons is

termed the valence band, and its uppermost energy level is called the valence

band edge (Ev). The next higher band, the excited statel is called the con-

duction band (since electrons introduced into this band are capable of con-

ducting electricity) and the lowest energy level is termed the conduction band

edge (E0). The forbidden region between the valence band and the conduction

band is referred to as the band gap.

Two types of semiconductors may readily be distinguished. In an intrin-

sic semiconductor, conduction electrons may be produced by direct excitation

(e.g. thermal) of electrons across the band gap from the valence band to the

conduction band. In this case the concentration of free electrons in the

conduction band equals the concentration of electron vacancies or holes in

the valence band. In practice crystals are not perfect (103) and usually con-

tain some foreign atoms which may be present in substitutional or interstitial

solid solution. These foreign atoms possess valence electrons which are bound

to their nucleus by forces different from those binding such electrons in the

other atoms i.e. additional quantum states are introduced into the crystal

which differ in energy from those already present. If the electrons occupying

these quantum states contribute to the conductivity in the crystal, it is

called an extrinsic semiconductor. For the case where there are donor impur-

ities in the crystal denoting excess electrons to the conduction band, the

crystal is an n-type semiconductor (i.e. conduction occurs by electrons). If

there are acceptor impurities capturing electrons from the valence band present,

-24-

the crystal is a p-type semiconductor. The unoccupied levels or holes left

in the valence band behave as unit positive charges in the conduction of

electricity through the crystal. Zinc sulphide is classified as an n-type

semiconductor (63) with a band gap of between 3.54 to 3.88 eV at room temper-

ature (51,108) and a resistivity reportedly ranging from 2.7 to 10-3 to

1.2 x 104 ohm.metre at 273 K (126), depending on composition, with the lower

value referring to "pure" zinc sulphide.

The occupancy of the various energy states in a semiconductor is des-

cribed by the Fermi-Dirac distribution function, which gives the probability,

f, that a given state will be occupied (22,102,103):

f = 1

exp E-EF + 1

kT

where E is the energy of the state, EF is the Fermi level, k is the

Boltzmann constant and T is the absolute temperature. The Fermi level is in

fact the electrochemical potential of the electrons or their partial molar

free energy. In a bulk semiconductor the position of the Fermi level depends

on the type of semiconductor (i.e. intrinsic or extrinsic), the concentration

and type of donors and/or acceptors, temperature etc. (102,102), and lies be-

tween the valence and conduction bands.

When the notion of "band structure" in semiconductors (and other solids)

was first introduced, the effect of the surface on the band structure was not

considered. A semiconductor, of course, is not infinite but is bounded by

surfaces. The surface atoms have fewer nearest neighbours than bulk atoms and

are in an asymmetric environment. The introduction of such a discontinuity at

the surface perturbs the periodic potential and results in solutions of the

Schradinger Wave Equation which do not occur for an infinite crystal (156).

These solutions predict the presence of electronic energy states, localized at

-25-

the surface, which can trap electrons or release them into the conduction

band.

These energy levels of the surface electronic states lie in the for-

bidden regions of the bulk band structure. Two types of surface states are

normally distinguished - "Tamm states" arising from the asymmetrical term-

ination of the crystal lattice at the surface and "Shockley states" arising

from the symmetrical termination of the lattice at the surface. An important

consequence of the presence of electronic surface states is that the electron

bands of the bulk semiconductor will be modified ("band-bending", up or down)

at the surface, even in the absence of a space charge, electron acceptor or

donor species (101.156). In the presence of an electron donor, for example,

the conduction and valence bands of an intrinsic semiconductor are "bent"

downwards at the surface, and upwards in the presence of an electron acceptor

(101,156). The occupancy and energy distribution of these electronic surface

states may be studied, for carefully specified semiconductors, by surface con-

ductance, field-effect and capacitance measurements etc.

2.5(b) Charge transfer at the semiconductor-water interface.

Charge transfer processes at the solid-liquid interface are of funda-

mental importance in understanding electrochemical corrosion, reduction,

oxidation and catalytic processes etc. (24,56 to 59). For a semiconductor

in contact with an aqueous solution containing a redox system (i.e. a reduct-

ant/oxidant couple), electron transfer may occur between the conduction band

or the valence band of the semiconductor and the ions of the redox system

according to the following mechanisms (24):

I M+ 1+= M2+ + e- conduction band mechanism i.e.

electron (e-) extraction or

injection

II M+ + e+ 1+, M2+

i_ valence band mechanism

i.e. hole (e+) extraction or injection.

•

A

D

hv

• 0 e

e+

e

(M is the ion in question; i+ is the anodic current ; i_ is the cathodic curr-

ent)

Much data has been accumulated for reactions occurring at germanium, zinc

oxide, silicon etc. electrodes. (24, 56 to 59). Morrison (56) has considered

the transfer mechanism between a surface state and the fluctuating energy levels

of an ion in solution in great detail and has particularly commended the current

doubling technique as a tool for studying hole or electron transfer steps.

The catalytic effect of irradiation upon oxidation/reduction reactions

at the semiconductor solution interface is well documented (42,43,57,58,59(a,c),

133). A simple example serves to illustrate the principles involved. If rad-

iation of energy (hv) equal to or less than the band gap excites an electron

from the valence band to the conduction band in a semiconductor, an electron-

hole pair is generated. If an acceptor species in solution (A) has an energy

level less than the conduction band (CB), electron transfer from CB to A is

possible. Similarly if a donor species (D) has an energy level above that of

the valence band (VB), electron transfer to the valence band may occur. The

process may be represented schematically as:

Energy Semiconductor Solution

•

CB

-27-

In such a case the semiconductor is acting as a catalyst and may be considered

ad a "photocatalytic pump". Examples like this are frequently encountered (42,

43,57,133) in semiconductor-aqueous systems.

2.5(c) The photovoltaic effect in sulphide semiconductors.

Williams has made a detailed study of the photovoltaic effect in binary

semiconductors (63) i.e. the electrical effect produced by light incident on

a semiconductor electrode which is in contact with an electrolyte solution.

Williams took crystals of CdS in the form of a thin plate and mounted it with

the bottom half in an electrolyte solution, its upper half remaining dry. The

thin plate of CdS was then connected through an electrometer to a calomel ref-

erence electrode and the broad face of the crystal was illuminated with light

from a tungsten lamp. A current of about 100 uA was produced under bright

illumination and a yellow solid formed on the face of the CdS plate upon which

the light was incident, but not on the face farthest from the lamp (63). The

yellow solid, through its solubility behaviour and UV spectrum, was shown to be

elemental sulphur. Williams concluded that the sulphur was formed by the

chemical decomposition of the CdS crystal under the influence of light. On

the basis of thermodynamic arguments Williams clearly demonstrated that for an

n-type semiconductor such as CdS, light raises electrons from the valence band

to the conduction band, leaving a deposit of elemental sulphur on the surface

and allowing Cd2+ ions to go into solution i.e.

CdS ;=--- Cdâq+ + S° + 2e

with an electron-consuming reaction at the calomel electrode. A similar mech-

anism was proposed for the n-type semiconductors ZnO and ZnS i.e. for ZnS

ZnS ± Zn2++ + S° + 2e aq

William's findings have, of course, been corroborated and elaborated on by

others (e.g. 56, 58, 59a) although the basic mechanism remains the same.

-28-

CHAPTER 3. EXPERIMENTAL METHODS.

3.1 KINETICS EXPERIMENTS.

3.1(a) Apparatus.

Kinetic studies were carried out in the thermostatted pyrex reaction

vessel shown in the photograph in Figure 3.1. Experiments were performed at

298 K, with the exception of some temperature dependence work at 308 and 318

K. The reaction vessel was designed to minimize the intrusion of atmospheric

02 and CO2 into the experimental system. It was fitted with a tightly fitting

perspex lid, which was clamped to the top of the reaction vessel by way of wing

nuts and brass rods, the latter being attached to a perspex base. The lid was

separated from the lightly polished, flat surface of the reaction vessel by a

rubber gasket, ensuring an efficient seal. Ports were drilled in the perspex

lid for the following purposes:-

nitrogen inlet and outlet

combined glass pH electrode

silver-silver chloride double junction reference

electrode (inner chamber 1 M KC1, outer chamber 1 M KNO3)

ion selective electrode

thermometer (°C)

A brass plug containing a thick rubber septum was screwed into the lid. This

was used when taking samples, with a needle and syringe, for Atomic Absorption

Spectroscopy (abbreviated to AAS) analysis. With the exception of the ion selective

electrode, which was sealed in its port by two tightly fitting rubber "0-rings",

all of the other ports were standard tapered joints. Before commencing a kin-

etics experiment all of these ports were protected further with adhesive tape,

attached at the junction between the upper surface of the lid and the electrode

body,for example.

An injection apparatus was designed to introduce ZnS powder into the system.

FIGURE 3.1 Photograph of the thermostatted reaction vessel, showing the essential features (N2 bubbler, ZnS powder, adhesive tape and uv lamp etc. not shown).

-30-

It consisted of a block of perspex with a hole drilled through it.Halfway

down this hole a channel was cut which allowed a closely fitting brass plate,

manipulated via an external handle (see Figure 3.1) to alternately open or

close the hole. The shaft of the handle passed through 2 rubber "0-rings" in

the perspex block before it reached the atmosphere. Above the brass plate

were 2 "O-rings", through which a standard glass syringe with a flat polished

end could be placed, hard and square against the brass plate when the latter

was in the closed position. A weighed quantity of ZnS powder was placed in the

barrel of the syringe and lightly compressed against the brass plate with the

syringe plunger, which moved through the Teflon ring at the top of the barrel.

The brass plate was moved to the open position and the ZnS "plug" remained in

place at the end of the barrel, held by adhesive forces. The syringe plus

contents could then be moved further into the reaction vessel. Fast, effic-

ient injection of the ZnS powder into the underlying liquid could then be

performed. Careful checking showed that, for a typical 2.00 g sample of ZnS,

the weight loss during this injection procedure (due to small quantities of

powder adhering to the walls of the syringe barrel) amounted to not more than

0.2%. The perspex block, syringe and external handle for manipulating the

brass plate are clearly visible in Figure 3.1.

pH and metal ion activator concentration measurements were carried out in

situ, with the relevant electrodes connected to Radiometer Model 28 and 26 pH

meters respectively. The output from these meters was fed into a two channel

Rikadenki Model B261 recorder.so that a continuoûs record of metal ion activ-

ator concentration and pH as a function of time was obtained.

Experiments were carried out in normal light, in the dark and under UV

irradiation. "In normal light" refers to experiments carried out in the re-

action vessel when it was unprotected from daylight and artificial laboratory

light. For experiments performed "in the dark" the reaction vessel was masked

with black tape and then shrouded with heavy, thick black cloth. "Under UV

irradiation" refers to experiments carried out under direct irradiation from

a 100 watt , "Ultraflux" Oliphant UV lamp. For these experiments the lamp

was mounted 25 cm from the external wall of the reaction vessel. The lamp and

vessel were then surrounded by reflecting aluminium foil. The spectrum of the

light emitted from this lamp, after passing through a pyrex beaker of water,

displayed strong intense peaks at 437, 406 and 367 nm.

3.1(b) Procedure.

400.0 cm3 of 10-1 M KNO3 was placed into the steamed reaction vessel, a

magnetic follower was introduced and the reaction vessel was placed on a mag-

netic stirrer. The perspex lid was fitted into place and clamped. A combined

glass pH electrode, previously calibrated at 298 K with standard Merck buffers

at pH 4, 7 and 9 was inserted (pH measurements were generally reproducible to

± 0.04 pH). The double junction reference electrode (with both chambers freshly

washed and replenished) and the ion selective electrode were introduced. The

temperature was set at 298 K and controlled to within + 0.2 K.

High purity N2, previously scrubbed through aqueous suspensions of TiO2

and SiO2 and then passed through conductivity water was bubbled through the

system until the pH reached 7 and the dissolved oxygen concentration was 0.05

ppm (i.e. 1.5 x 10-6 M) or less. The dissolved oxygen concentration was mon-

itored with a Yellow Springs Instrument Company Model 54 dissolved oxygen meter

- the probe was inserted through the hole in the injection system for this

measurement.

The dissolved oxygen probe was then removed and a highly accurate (0.1%

or better) Agla micrometer syringe, containing standard Cu(NO3)2, Cd(NO3)2 or

Pb(NO3)2 was used to carefully calibrate the relevant Cut+, Cd2+ or Pb2+ ion

selective electrode in the reaction vessel (see Section 3.3). The pH was then

-32-

carefully adjusted to the required value with small quantities of KOH or

HNO3 using 1 cm3 volume, all glass syringes. The total volume change follow-

ing calibration and pH adjustment was generally not greater than 0.5 cm3.

The Agla syringe was removed and the brass plate in the injection sys-

tem was moved into the "closed" position. A weighed quantity of ZnS (most

commonly 2.00 or 4.00 g) was placed in the syringe, compacted as previously

described and the syringe was lowered until the powder was some 3 cm above the

surface of the liquid. The exposed ZnS "plug" was then subjected to the con-

stant flow of high purity N2 flowing through the reaction vessel for 5 minutes,

in an effort to remove any residual "trapped" 02 or CO2 held in the "plug".

The system was then completely sealed and the N2 supply was shut off. The ZnS

was then injected into the liquid - the rapidly moving magnetic follower ensured

rapid, efficient dispersion in 4 seconds or less. Simultaneously with the ZnS

injection the recorder' was started and the metal ion concentration (i.e. Cu2 ,

Cd2+ or Pb2+) and pH was recorded as a function of time.

There was no suitable continuous method for analyzing for zinc as a

function of time, hence five 10 cm3 samples, containing both ZnS powder and

liquid, were removed with a needle and syringe through the rubber septum in the

lid of the reaction vessel at specified intervals of time. These samples were

then filtered through a 0.22i Millipore filter and analyzed for zinc and the

relevant metal ion by AAS. The latter analysis was carried out as a check on

the ion selective electrode results.

Thus a record of pH, residual metal ion activator concentration and con-

centration of zinc released was obtained as a function of time. Knowing the

initial concentration of the metal ion activator, the metal ion uptake per

gram of ZnS was readily computed.

3.2 MATERIALS.

3.2(a) Water quality.

Conductivity water was prepared by oxidizing once distilled water in

alkaline permanganate followed by a second distillation step. Distillation

and storage was carried out in all-glass apparatus. This conductivity water

had a specific conductivity of 1 x 10-6 0-1 cm -1 or less throughout the entire

experimental program, as shown by regular determinations. Conductivity water

was used in making up solutions for the entire experimental work. All apparatus

was thoroughly cleaned with nitric acid and steamed before use.

3.2(b) Chemicals.

All chemicals used throughout the experimental work were of AR grade or

better, unless otherwise noted.

(I) electrode calibration

Cu2+ electrode

A stock solution of Cu(NO3)2 was standardized against standard

sodium thiosulphate which, in turn had been previously standardized against

potassium iodate. It was found to have a molarity of 0.1025 ± 0.0002 M.

Cd2+ electrode

A stock solution of Cd(NO3)2 was standardized by gravimetric anal-

ysis of cadmium as the molybdate and was found to have a molarity of

0.0974 ± 0.0002 M.

Pb2+ electrode

A stock solution of Pb(N0i3)2 was standardized by gravimetric analy-

sis of lead as the chromate and was found to have a molarity of 0.0958 ± 0.0002 M.

N.B. These stock solutions were individually stored in steamed polypropylene

containers at pH 5 to prevent hydrolysis and were used in all subsequent kin-

etic and mass spectrometric experiments.

(II) AAs standards.

AAS standards of 1000 ppm (1000 pg/cm3) were prepared by acid

dissolution of 1.000 g of the relevant high purity metal followed by dilution

to 1 litre, according to standard procedures (111) and then stored in steamed

polypropylene containers. Prior to analysis, standards in the desired concen-

tration ranges were freshly prepared from these 1000 ppm solutions by dilution

and were stored in small polypropylene bottles.

(III) ZnS.

High purity (> 99.95%), electronic grade synthetic zinc sulphide,

manufactured by Koch-Light Laboratories (U.K.) was used in this study. It was

prepared by a precipitation technique, followed by firing to remove volatiles.

This synthetic zinc sulphide existed as the cubic crystalline form and a typi-

cal spectrographic batch analysis revealed trace quantities of Cd, Al, Si, K,

Na and Sn (approximately 20 ppm or less). The particles were regular in appear-

ance, as is shown by the electronmicrographs in Figure 3.2.

Five separate batches of ZnS were used in this experimental work.

They were kept in bottles with tightly fitting lids and were stored in a vacuum

desiccator over silica gel, under vacuum, to minimize the effects of air oxidat-

ion. The surface areas of these five samples were kindly determined by Tioxide

Australia Pty. Ltd. (Burnie, Tasmania) using the standard BET (41) technique

with N2 as the adsorbate. The respective surface areas, representing the mean

of 3 separate determinations, were found to be

-35—

FIGURE 3.2 Electronmicrographs of the high purity, synthetic zinc sulphide used in this investigation (top 5700 X; bottom, edges, 34000 X).

Sample 1 (S1) 0.72 ± 0.05 m2 g-1

Sample 2 (S2) 0.68 ± 0.05 m2 g-1

Sample 3 (S3) 0.74 ± 0.05 m2 g-1

Sample 4 (S4) 0.58 ± 0.05 m2 g-1

Sample 5 (S5) 0.57 ± 0.05 m2 g-1

Noting that either 2.00 or 4.00 g of ZnS was commonly used in a volume of

400 cm3 during the kinetic studies, the surface area available for activation

therefore commonly ranged from 1.14 to 2.96 m2.

3.3 ANALYSIS.

3.3(a) The type of ion selective electrodes used.

The following solid state ion selective electrodes, manufactured by Orion

Research Incorporated (U.S.A.), were used in this investigation:

Cu2+ Model 94-29A

Cd2+ Model 94-48A

Pb2+ Model 94-82A

The sensing membranes of these electrodes are composed of a Ag2S/MS mixture,

where M2+ represents the metal ion under study. The principles of construction

and behaviour of these solid state electrodes have been fully documented by

Durst (112) and others (113,114).

After heavy use the surface of these electrodes becomes pitted - a layer

of A42S is left after the MS is leached out (3). This necessitates light pol-

ishing of the membrane surface followed by several minutes immersion in 0.02 M

H2SO4, a method which Blaedel and Dinwiddie (3) have found to be satisfactory

with Orion Cu2+ electrodes used to measure submicromolar concentrations of Cu2+.

This proven technique was adhered to during this experimental program. In

fact, prior to each kinetics experiment, the Cu2+ electrode was cleaned with

0.02 M H2SO4 for two to three minutes, and then rinsed with conductivity water.

-37-

Rapid response, together with excellent reproducability and stability were

noted. The very low, controlled dissolved oxygen levels which applied in this

experimental program, as well as the much higher Cut+ concentrations used,

eliminated the slight air oxidation effects noticed by Blaedel and Dinwiddie

in their submicromolar study of Cut+ determination (3).

When using divalent metal ion electrodes it is necessary to measure the

electrode potential to ± 0.1 mV in order to obtain high precision in metal ion

concentration (or activity) measurements. This was achieved with the Radio-

meter Model 26 meter mentioned above, using the expanded scale facility.

Furthermore all of the kinetic experiments were performed in a background

electrolyte concentration of 10-1 M KNO3, with the exception of some documented

examples (Chapter 4) demonstrating that changing the ionic strength had no

detectable effect on the kinetics. This ionic strength of 10-1 M was similar

to that encountered in flotation pulps (1,2,4) and possessed two very important

characteristics:

(i) metal ion activity coefficients were maintained constant,

resulting in linear calibration curves

(ii) electrode potential fluctuations, caused by fluid flow past

the metal ion and reference electrodes due to the uneven stirr-

ing action of the magnetic follower were minimized. In fact

such potential fluctuations were less than ± 0.1 mV at 10-1 M

KNO3 (increasing to about ± 0.8 mV at 10-3 M KNO3) over the

range of metal ion concentrations studied.

KNO3 was chosen as the background electrolyte because it has been shown to be

indifferent (no specific adsorption) in electrokinetic studies of ZnS (9). No

other background species (e.g. pyrogallol to "scavenge" trace quantities of

dissolved 02) were introduced into the system for, as shown by cementation

studies (31) they can affect product morphology and reaction rate and are an

unnecessary complication.

3.3(b) M2+ electrode performance.

(I) pH dependence.

For the M2+ concentration ranges specified in Table 3.1 and with

a background ionic strength of 10-1 M KNO3 the Cu2+, Cd2+ and Pb2+ electrodes

proved to be insensitive to pH in the ranges also shown in Table 3.1.

TABLE 3.1.

pH dependence of the M2+ electrodes.

M2+ (M2+] range, mole 2,-1 pH range over which electrode reading was independent of pH

Cu2+ 5 x.10-6 to 1 x 10-4 3.8 to 6.5

Cd2+ 5 x 10-6 to 1 x 10-4 4.5 to 8.8

Pb2+ 5 x 10-6 to 1 x 10-4 4.0 to 7.8

Below the lower pH limits the electrode potential readings become progressively

more unstable as the pH was reduced, presumably due to leaching of the relevant

metal sulphide from the surface of the sensing membrane. Above the upper pH

limits hydrolysis of M2+ occurs, leading to a consequent apparent non-Nernstian

response. Metal ion concentrations generally fell within the specified concen-

tration boundaries so that pH changes within the ranges shown in Table 3.1 did

not affect the determination of Cu2+, Cd2+ or Pb2+. Outside these pH limits,

metal ion concentrations were determined by AAS.

(II) Effect of UV irradiation.

The M2+ electrodes proved to be very unstable under irradiation

from the 100 watt UV lamp used in the experimental program. For example when the

Cut+ electrode, immersed in a 5 x 10-5 M Cu2+ solution in the reaction vessel,

was exposed to UV radiation from the lamp, the potential decreased by 30

millivolt over several minutes and continued to decrease.

When the lamp was switched off the electrode immediately returned to

within 5 millivolt of its initial reading and then gradually stabilized at

a potential within 1 millivolt of this initial value. This effect persisted

over a range of pH and Cut+ concentration values. It meant that any attempts

to obtain kinetic data using a Cu2+ electrode in the presence of intense UV

radiation were fruitless, due to the marked change in the electrode's potent-

ial characteristics. Similar behaviour was shown by the Cd2+ and Pb2+ elect-

rodes. Hence for experiments performed in the presence of UV irradiation

metal ion analysis was carried out by AAS.

(III) Effect of inert solid on electrode response.

In order to determine whether or not the presence of suspended, in-

ert* solid particles affected the M2+ electrode response the following ex-

periments were performed:

The Cut+ electrode was calibrated in the reaction vessel, follow-

ing the normal procedure.

The initial Cue-1- concentration was 5.12 x 10-5 M at pH 4.0 in

10-1 M KNOB. The gradual addition of a high purity, large surface area 100 m2

g-1 SiO2 powder up to a solid concentration of 12 gil did not cause any

detectable alteration to the Cut+ electrode potential. Repeating this experi-

ment at pH 5.0 yielded the same result. Similarly the Cd2+ and Pb2+ electrode

potentials did not alter in the presence of the SiO2 powder.

It was therefore concluded that the presence of inert solid par-

ticles in suspension did not affect the M2+ electrode potential (i.e. there

was no "solid suspension effect" under the experimental conditions).

(IV) Electrode response time, stability and calibration behaviour.

Over the pH ranges studied the M2+ electrode response was rapid

provided that

* "Inert" = no reaction with M2+

(i) the mémbrane had been cleaned in the manner already

described

and (ii) the M2+ and reference electrodes were immersed in the

0.1 M KNO3 solution, contained in the reaction vessel

for at least half an hour prior to testing so that

thermal equilibrium was achieved.

Under these conditions repeated checking showed that the M2+ electrodes

reached stable potential readings (i.e. stable meaning not more than a 0.2 mV

drift over 2 hours) in 3 seconds or less for M2+ concentrations above

4 x 10-6 M.

Since the M2+ concentrations were normally within the concentration range

from 7 x 10-6 to 10-4 M,the ion selective electrodes could be used with

confidence to follow the reaction rate, as preliminary experiments had demon-

strated that the metal ion activation reactions occurred much more slowly than

the electrode response time (see also part (V) below).

Typical calibration curves for the Cut+, Cd2+ and Pb2+ electrodes are

shown in Figures 3.3, 3.4 and 3.5. They are linear for concentrations in the

range 5 x 10-7 to 10-4 M (and above) and have gradients of 30.0 ± 0.2 mV. A

practical detection limit of 10-7 M was set for Cut+, Cd2+ and Pb2+ — the

calibration curves are generally non-linear below about 5 x 10-7 M. Selected

linear portions of calibration curves like those illustrated in Figures 3.3 to

3.5 were plotted on expanded scales in order the achieve greater accuracy in

metal ion determinations. A precision of ± 2% or better was achieved.

(V) Comparison between metal ion concentrations determined by

the ion selective electrodes and by AAS.

Sample values taken from the kinetic experiments are given in Tables

3.2, 3.3 and 3.4. Due to the time involved in sampling and filtration before

AAS analysis could be performed, the metal ion concentrations apply at an

interval of 5 minutes and beyond after the commencement of a kinetics

120

110

100

90

80

70

60

50

-41-

104 10.6 10.6 10.4

Concentration of Cue+ (mole 2-1)

FIGURE 3.3 Typical Cue+ electrode calibration curve.

Cd2

' E

lec t

rode

pot

enti

al

(mill

ivo l

t)

-42—

10'6 10'5

Concentration of Cd2+ (mole 10 )

FIGURE 3.4 Typical Cd2+ electrode calibration curve

Pb2

+ E

lec t

rod

e pote

nt i

al (

mil

livo

lt)

-43-

Concentration of Pb2+ (mole Q" I )

FIGURE 3.5 Typical Pb2+ calibration curve.

experiment. Agreement between the AAS and electrode values was excellent

(± 3% or better) over the following metal ion concentration and pH ranges:

(i) [Cu?+] ti 5 x 10-6 to 1 x 10-4 M, pH 4.0 to 6.5

(ii) [Cd2+] ti 5 x 10-6 to 5 x 10-5 M, pH 5.0 to 8.6

(iii) (Pb]?+ ti 5 x 10-6 to 5 x 10-5 M, pH 4.0 to 7.8

There was no detectable difference between values determined in normal light

and those determined in the dark. Bearing in mind the excellent response

characteristics of the ion selective electrodes and combining this with the

close correlation between metal ion concentrations determined by the two

techniques, reliable and precise metal ion concentration values could be ob-

tained during the kinetics experiments. The results reported in Chapters 4

to 6 after a reaction time of 5 minutes are an average of the individual metal

ion concentrations obtained from the two techniques. Electrode instability at

low pH and metal ion hydrolysis at high pH meant that analysis by AAS was

performed in these pH regions, as discussed in Section I above.

3.3(o) AAS analysis.

Copper, cadmium, lead and zinc were determined using a Varian Techtron

Model 6 Atomic Absorption Spectrometer coupled to a chart recorder. The

following operating conditions were used (111):

Wavelength run

Spectral Band Pass nm

Typical Working range ppm

Flame

Cu 324.7 0.2 0.1 - 4.0 Lean air - acetylene

Cd 228.8 0.5 0.1 - 2.5 air-acetylene

Pb 217.0 1.0 2.0 -10.0 air-acetylene

Zn 213.9 0.2 0.1 - 2.5 air-acetylene

TABLE 3.2. DETECTION OF COPPER.

Comparison between Orion Cu2+ electrode (Model 94-29A) and Atomic Absorption Spectroscopy.

(N.B. Electrode values refer to Cu2+ in situ, AAS values apply after filtration).

(Cu2+],mole k-1 x 10+5

electrode

[Cu]TOTAL,mole k-1 x 10+5

AAS

pH

2.14 2.12 nk 4.2

1.47 1.57 nk 5.0

0.39 0.37 n k 5.1

0.25 0.28 n k 5.2

6.48 6.37 n 5.3

0.94 0.87 nk 5.6

6.12 6.12 n2. 5.8

2.56 2.52 nk 6.0

0.78 0.71 nk 6.3

0.77 0.79 nk 6.3

0.77 0.79 nk 6.5

1.28 1.23 nk 6.5

2.61 2.68 d 4.0

1.66 1.65 d 5.4

1.46 1.29 d 5.5

1.04 1.01 d 6.3

1.00 1.04 d 6.3

N.B. "ni" refers to experiments carried out in normal light,

"d" to those in dark.

TABLE 3.3. DETECTION OF CADMIUM

Comparison between Orion Cd2+ electrode (Model 94-48A) and Atomic Absorption.

(N.B. Electrode values refer to Cd2+ in situ ,1 S values apply after filtration).

(Cd2+],sole ß-1 x 10+5

electrode

[Cd]TOTAL,mole t-1 x 10+5

AAS

pH

0.97 1.04 nt 5.1

1.06 1.01 nt 5.3

1.94 2.02 nt 5.3

1.50 1.51 nt 6.2

0.75 0.82 nR. 6.3

0.90 0.96 nt 6.5

1.47 1.54 nt 6.5

1.76 1.63 nt 6.5

1.96 1.98 nR. 6.5

0.97 0.98 nR, 6.6

1.04 1.06 nR. 6.7

1.09 0.97 nt 7.4

1.52 1.60 nt 7.4

1.60 1.64 nR, 7.6

1.67 1.71 nt 7.7

1.72 1.72 nR. 8.4

1.80 1.83 nt 8.6

1.48 1.51 d 6.2

1.90 1.96 d 6.5

1.75 1.85 d 6.6

2.05 2.10 d 6.8

N.B. "nt" refers to experiments carried out in normal light, "d"

to those in the dark.

TABLE 3.4. DETECTION OF LEAD

Comparison between Orion Pb2+ electrode (Model 94-82A) and Atomic Absorption Spectroscopy.

(N.B. Electrode values refer to Pb2 ,in.situ,AAS values after filtration).

[Pb2+],mole L-1 x 10+5

electrode [Pb]TOTAL

,mole 2.-1 x 10+5

AAS

pH

3.90 4.03 ni 4.0

0.44 0.46 ni 5.2

3.87 4.01 ni 5.2

0.36 0.43 ni 5.3

4.02 3.93 ni 5.6

3.24 3.08 ni 6.7

3.00 3.16 ni 6.7

3.82 3.83 ni 6.8

3.78 3.86 ni 6.8

3.84 3.93 ni 6.8

3.92 3.78 ni 7.7

3.80 3.76 ni 7.8

4.20 4.22 d 5.2

4.30 4.23 d 5.2

4.18 4.20 d 5.3

4.10 4.21 d 5.3

N.B. "ni" refers to experiments carried out in normal light, "d"

to those in dark.

No chemical interferences were detected under the operating conditions,

in agreement with established behaviour (111). Non-atomic absorption was

absent (checked with hydrogen continuum lamp).

Representative Absorbance versus Concentration curves are shown in

Figures 3.6 to 3.9. Unknown solutions, when required, were diluted so that

their metal ion concentrations fell within the linear working range.

Detection limits of 0.01 ppm (10-7 M), representing a reading twice that

of the "background", were set for both zinc and copper. Similarly,detection

limits of 0.01 ppm (10-7 M) and 0.2 ppm (10-6 M) were set for cadmium and lead

respectively.

3.4 SUBSIDIARY TESTS.

3.4(a) Pretreatment experiments.

The type and degree of pretreatment can frequently have a dramatic effect

on the properties of mineral surfaces (9t 13, 42). In a practical flotation

circuit sphalerite is normally in aqueous suspension before the activating

CuII solution is encountered (1, 2, 4). Depending on the particular plant

operating conditions, the sphalerite may experience large changes in pH, come

into contact with a variety of metal and other ions together with organic

species and foreign mineral particles, experience changes in dissolved oxygen

content etc. before it is "activated" with CuII (1, 2, 4). From the vast num-

ber of experiments possible it was decided, in this current investigation, to

restrict the pretreatment experiments to those involving changes in either pH

or incident light.

Pretreatment at low pH should result in the formation of elemental sulphur

at the mineral surface (114 and Chapter 8) together with soluble products which

may be retained in the Stern Layer. On the other hand conditioning the mineral

Concentration of Copper (ppm)

100

90

80

70

60

50

40

30

-49-

1.0 2.0

3.0

Concentration of Zinc (ppm)

FIGURE 3.6 Typical AAS calibration curve for zinc

J20

10

a

o

a

•

$ ac

100 0 ô 90

G 80

70

60

50

40

30

20

10

0

FIGURE 3.7 Typical AAS calibration curve for copper

-50-

100

90

80

70

60

50

40

30

20 é a $ 10

0

10 , .

8 m ¢ r 100 r U

90

â 80

70

60

50

40

30

20

10

1.0 2.0

3.0

Concentration of Cadmium (ppm)

FIGURE 3.8 Typical AAS calibration curve for cadmium

1 2 3 4 5 6 7 8 9 10

Concentration of Lead (ppm)

FIGURE 3.9 Typical AAS calibration curve for lead

-51-

at high pH (pH 12 and above) can lead to the formation of polysulphides (15,

16), Sn2-, or polynuclear ZnII species (17) which may alter the activation

rate. The effects of irradiation on the surface properties of semiconductors

are well established (Chapter 2).

It is important to note that it was not possible to use the Cu2+ elec-

trode to follow the extent of reaction. Nor, for that matter, could Cd2+ or

Pb2+ electrodes be used. The Cu2+ electrode example is used here. In the

previous experiments fresh, dry ZnS had been injected into a system containing

a known concentration of Cu2+ and a calibrated Cu2+ electrode. If the ZnS

was pretreated as a suspension in the reaction vessel before the cupric nitrate

solution was added a complication arose. The Cu2+ electrode had to be placed

in the suspension prior to adding the cupric nitrate solution. Of course

immediately it contacted the ZnS suspension its potential altered rapidly as

the small number of Cu2+ ions in solution, which arose from membrane solubility,

were removed by the ZnS. Progressive leaching of Cu2+ from the membrane sur-

face layers then occurred. This depletion of the surface layers of the mem-

brane and the concomitant alteration in potential meant that subsequent con-

tinuous, quantitative determination of Cu2+ removal using the Cu2+ electrode

was impossible.

Pretreatment of the ZnS followed by filtration, drying and redispersion

was an undesirable procedure as it would probably have removed the effects of

pretreatment (except perhaps for insoluble products formed on, and attached to,

the ZnS surface).

The effect of pretreatment of the ZnS on the uptake of CuII was there-

fore determined by AAS analysis for copper and the influence of pretreatment

after a reaction time of 5 minutes was investigated. Pretreatment involved "con-

ditioning" the ZnS for 1 hour at the desired pH etc., followed by activation

with CuII under specified conditions for a further 2 hours. The pretreatment

time of 1 hour was simply a convenient choice - the effects of pretreatment

-52-

were certainly apparent at this time and are reported in Chapter 4.

3.4(b) Adsorption on filters.

Generally samples for AAS analysis were filtered through 0.22p Milli-

pore filters before storage and subsequent analysis. In pH regions where

metal ion hydrolysis was expected, samples were centrifuged and the super-

natant was extracted for AAS analysis. Values obtained after centrifugation

and after filtration could then be compared. This was carried out for copper

.(> pH 6.5), _ cadmium (> pH 7.8) and zinc (> pH. 8.5).A.greeMent between

the 2 separation methods was good, indicating that the metal ion hydrolysis

products formed, if any, were on the ZnS surface. At pH values less than

those specified above, the close correlation between the electrode and AAS

values showed that no detectable loss of copper, cadmium or lead occurred on

the filters. For zinc, careful checking with known zinc concentrations showed

that no detectable loss of zinc on the filters occurred below pH 8.5.

3.4(c) Presence of H2Sgas SO4

gas aqueous.

A trap containing 2 M AgNO3 inserted in one of the ports of the reaction

vessel during preliminary activation tests showed no detectable precipitate

of Ag2S (from H2Sgas) for pH values of 4 and above, nor was the odour of H2S

detectable. Similarly a standard nephelometric technique (46,141) indicated that

sulphate was not detectable (i.e < 10-5 M) in solution during these tests.

3.4(d) Eh measurements.

Eh measurements were performed using a bright Pt electrode and a double

junction calomel reference electrode. The electrodes were constantly checked

against Zobell's solution (14) and yielded the "expected" Eh of +0.430 volt

at 298 K.

3.5. THE DETERMINATION OF ELEMENTAL SULPHUR.

3.5(a) Introduction. The presence or absence of elemental sulphur on sul-

phide mineral surfaces is a contentious issue, as has already been discussed

in Chapter 2. The development of a sensitive analytical technique, which is

capable of the quantitative detection of very low levels of elemental sulphur

on sulphide mineral surfaces, is therefore of paramount importance. Only then

will the problems concerning elemental sulphur be at least partly resolved.

Broadly speaking, current methods for detecting elemental sulphur on

mineral surfaces fall into two classes:

(a) indirect methods which involve the extraction of sulphur from

the surface with a suitable solvent. Acetone, petroleum ether, pyridine

and carbon disulphide have been favoured to date. This extraction pro-

cess is then followed by quantitative detection of sulphur using colori-

metric (74,75) or gas chromatographic (76,77) techniques. Colorimetric

techniques suffer from interference from sulphur-containing compounds

and are of low sensitivity e.g. sulphur can only be determined above

5 x 10-4 g, which, for the purposes of this present study is too insen-

sitive. Allison and Finkelstein (76,77) have used gas chromatography to

detect elemental sulphur on mineral surfaces following extraction with

carbon disulphide. The extracted sulphur is concentrated and injected

into the gas chromatograph. Sulphur is quantitatively determined from

a calibration curve, after accounting for background corrections (in

part due to the presence of small amounts of sulphur which are likely to

be present in carbon disulphide). This method is reproducible and is

reported to have a detection limit of 5 x 10-6 gram of sulphur (20% of a

monolayer on an available area of 0.4 m2, assuming each sulphur atom

occupies an area of 8.7 Â2 per atom (76,77)). All of these

-54-

indirect techniques suffer from the disadvantage that there is always

some uncertainty whether or not sulphur extraction is 100% complete.

(b) direct methods involving solid-liquid separation, drying and

analysis of the solid surface. Clifford and Miller (69) used electron

spectroscopy for chemical analysis (ESCA) to determine sulphur on sphal-

erite and Clifford, Purdy and Miller (78) have extended the technique

to other sulphide minerals. The method at present is sensitive but is

only semiquantitative and, as recognized by the authors themselves, may

possibly form the species that are actually being analyzed for (by X-ray

photoreduction).

Direct methods have considerable advantages over indirect methods, par-

ticularly regarding sensitivity and completeness of extraction. Mass spectro-

metry, in particular, appears to be an ideal technique for the quantitative

detection of elemental sulphur, since

(i) the constituents of sulphur vapour have been studied by

mass spectrometry (79,81)

(ii) interchalcogen compounds (of sulphur and selenium) have been

qualitatively analyzed by mass spectrometry, using a direct

insert probe (82)

(iii) the technique is extremely sensitive - it is capable of

detecting down to 1 x 10-9 gram (83)

(iv) there is no direct electron bombardment of the solid surface

(reduction is therefore extremely unlikely).

Preliminary tests showed that the technique was in fact, very promising, des-

pite the fact that

(a) the sample must be dried before analysis

and (b) the technique has apparently not been used for the quantitative

analysis of solid mixtures to date (83). Mass spectrometry was therefore

used in this investigation to detect elemental sulphur on the surface of

ZnS under a variety of conditions. A description of the apparatus

-55-

and experimental technique is given below.

3.5(b) Mass Spectrometry.

(I) Apparatus

A Finnigan 1015 electrostatic mass spectrometer was used in this

investigation. It operates by passing a beam of ionized particles through a

hyperbolic electrostatic field. The direction of the field varies in such

a way that only ions of particular mass to charge ratio (M/e) can pass through

it. The intensity of the field is swept so that the spectrometer can scan

across a wide M/e range. The mass spectrometer performs four major functions:

(i) inlet and vapourization of the test sample

(ii) ionizing the sample vapour

(iii) filtering the sample ions in order to analyze them according

to their M/e ratio

(iv) detecting the filtered ions and generating an output signal

that can be suitably displayed.

A vacuum system is also incorporated into the mass spectrometer, together with

a high voltage source which drives the quadrupole mass filter. The major

components are shown in the block diagram (Figure 3.10).

The DC/RF generator has two functions - it generates a sweep voltage

to drive the readout system oscilloscope and it uses the same sweep voltage

to drive the quadrupole mass filter. The various major functions are described

in more detail below.

(i) Inlet systems.

In this study, the solid probe inlet was used to introduce solid

samples into the vacuum system. The probe is a stainless steel rod with

a holder for a small capillary tube in the end. The dried sample, which

was in the form of a powder in this investigation, is placed in the

-56—

DE

TE

CT

OR

FIG

UR

E 3

. 10 B

lock

dia

gram

of

the

Fin

niga

n 1

01

5 M

ass

Spe

ctro

met

er.

DC/R

F V

OL

TA

GE

S

J

w U

capillary tube, which in turn is placed in the solid probe. Then the

probe is inserted through a vacuum lock directly into the ionization

chamber. The sample is heated until it produces sufficient vapour for

analysis - the temperature is monitored and controlled from the front

panel. The probe design prevents local gas pockets in the capillary

tube from "blowing out" during heating - a small spring encloses the

capillary tube and holds it in place. The tip of the solid probe is

adjacent to the ionization chamber, so that as the molecules are vapour-

ized, they are directed into the ion source to achieve maximum sensitivity.

The capillary tube allows relatively large sample sizes to be used (see

below) .

(ii) Ion source.

The ion source is located on the quadrupole mass filter assembly,

as shown in Figure 3.11. The vapourized sample is directed into the ion

source. Electrons emitted from a ductile rhenium filament are focused

into a beam. The electron beam crosses the ionization chamber where it

bombards the sample molecules to form ions. The ions are then removed

from the ionization chamber, pass through a series of focusing apertures

and are then injected directly into the quadrupole mass filter.

(iii) Quadrupole mass filter.

The mass filter consists of two pairs of metal rods that are pre-

cisely aligned and held by two high-purity ceramic insulators. The rods

are electrically connected (in opposite pairs) to RF and DC potentials,

which form a hyperbolic electrostatic field (see Figure 3.11). The

amplitude of the DC and RF potentials increases with time for a "mass

scan". The RF and DC potentials increase uniformly in time to a maximum

value, drop suddenly to zero and then the sweep is repeated. Ions are

extracted from the ion source and focused by electrostatic lenses into

a beam that enters the hyperbolic field. Most of the ions that enter the

EL

EC

TR

ON

MU

LT

IPL

IER

Physical unit Corresponding signal flow

Corresponding function

0 1 2 w

4 o¢ 4 o H

Lv

wa

Q â ao ââ Cd

â ° ¢c

w Q â c

_ r

Wa

W

d

i-- !— ....1 -

.

H

0 . ...

8â —1 F— z

Îï z —11 w w --1 . o g IC7I __ _ _

w

Î\1 '''

w

I a ô

. •.l./ o a

Wa W • •'. ......'. ' ..•

ô w 2 ,N

. â ââ1 W

aw ooM 4

'p âôa H H c cn

w 0

F tai a1-1âw2

a 2

AM

PL

IFIC

AT

ION

F

ILT

ER

ING

FIG

URE

3.1

1

-59-

field are deflected excessively and assume an "unstable" trajectory.

They are attracted to one of the rods, where they are neutralized.

At any given time, ions of one specific M/e ratio are deflected

as much towards one rod as another. Ions of this particular M/e ratio

pass through the filter, but ions with other M/e ratios are momentarily

excluded. As the voltages sweep from zero to their maximum values the

entire mass range is scanned, starting at low mass numbers and increas-

ing to high mass numbers.

Ion achieving a stable trajectory (stable meaning that the amplitude

of the oscillations of the ion remain within the boundaries of the mass

filter rod surfaces) are able to traverse the length of the mass filter.

These stable ions are then focused into the detector.

(iv) Detector.

The detector is an electron multiplier and is part of the mass fil-

ter assembly, as shown in Figure 3.11. The electron multiplier, together

with the quadrupole assembly and the ion source, fit into the vacuum mani-

fold. Once the ions have passed through the mass filter, they are acceler-

ated from about 5 eV to 3 KeV and enter the electron multiplier. The

incident ions strike a surface in the multiplier called a "dynode", which

emits secondary electrons upon impact. An average of 2.6 electrons leave

the surface for each incident ion. These secondary electrons then cas-

cade through a series of dynodes, each dynode focus ing the electrons on

to the next dynode. The output of the electron multiplier is a current

that flows through a resistance selected by the sensitivity control on

the display unit. The voltage across this resistance is amplified by the

preamplifier in the readout system. The manner in which the quadrupole

assembly operates is shown in Figure 3.11.

(v) Display section.

The display section of the Finnigan 1015 comprises four major com-

ponents - a preamplifier, an oscilloscope, a light-beam oscillograph

and the total ion monitor. The signal from the preamplifier drives the

oscilloscope, oscillograph, total ion current monitor and an external

display. In this investigation an additional 100 X amplifier was insert-

ed between the signal from the preamplifier and the external display. The

latter consisted of either a Gould Instrument Company Model 220 dual

channel recorder or a Rikadenki Model TO2N1-H recorder. The selection of a

particular recorder depended on the requirements of the experiment being

performed e.g. whether a "fast" or "slow" scan was needed.

3.5(c) Experimental procedure.

(I) Preliminary studies.

Elemental sulphur (S°) exists in numerous complex forms (80). Fusion

normally occurs in the temperature range 383 to 393 K (P = 1.01 x 105 Pa),

depending on the crystalline form and vapourization occurs at 717.8 K

(P = 1.01 x 105 Pa) (80,84). At temperatures between 279 and 465 K, the sat-

urated vapour pressure of S° ranges from 1.33 x 10-5 Pa (10-7mm Hg) to .

1.33 x 10+2 Pa (1 mm Hg) (84) .

Brandt, Mohler and Dibeler (79) have shown that when S° is evaporated

from a heated tube directly into the ionization chamber of a mass spectrometer,

ions of the form Sx+ are produced, with x ranging from 1 to 8. S2+ is the

most abundant. This was confirmed in the current study - a few microgram of

Analar grade S° yielded a similar spectrum to that described above, with S2+

being by far the most intense peak. A similar result was obtained from activ-

ated ZnS samples where the presence of S° was suspected. As a result of these

preliminary studies, two lines of approach emerged:

-61-

(a) qualitative identification of So through observation of the ions

S+ to S8+ combined with a comparison of measured and calculated

peak intensities

(b) quantitative determination of So by comparing the intense S2+ peaks

of known standards with those of unknown samples.

The results of (a) and (b) for a variety of systems are reported in

detail in Chapter 7.

(II) Mass spectrometer.

The Finnigan 1015 mass spectrometer contains 3 mass ranges:

0-100, 10-220 and 50-750 mass numbers. Most of the work was confined to the

0-100 and 10-220 ranges. The mass meter was therefore calibrated with acetone

(parent peak at M/e = 58) and bromobenzene (parent peak at M/e = 156) respect-

ively.

For quantitative work, the mass range from approximately 62 to 67 was

scanned (the mass meter was carefully calibrated at mass number 58 and checked

repeatedly - no significant drift was observed). This corresponded to a very

"clean" part of the spectrum i.e. the background level was very low. Moreover

to ensure a low background level the vacuum manifold was baked for periods up

to 4 days.

The ionization chamber normally operates in the vicinity of 573K (lower

temperatures result in rapid contamination and a marked decrease in sensitivity).

Since the tip of the solid probe is adjacent to the ionization chamber, heat

transfer naturally occurs and the probe generally heated up from approximately

308 K (350 C) to 373 K (100°C) during the course of an experiment. Without

substantial modifications to the instrument, this was unavoidable. The lower

limit of about 308 K corresponds to withdrawing the !lot probe through the

vacuum lock, allowing it to cool, inserting a fresh sample and re-inserting the

probe. Higher temperatures could be attained by heating the probe.

-62-

Normal instrument operating conditions were as follows:

Ionization current : 250 UA Background current : 2 to 3 pA

Ionizing current 15 pA

Pressure : 1 x 10-6 mm Hg (1.33 x 10-4 Pa)

Electron energy 70 Volt

(III) Sample preparation.

All samples were placed in clean glass capillary tubes. The cap-

illary tubes were of similar wall thicknesses, sealed at one end and about

1 cm in length. The tubes were packed with ZnS powder to the same level,

using the same packing technique. All handling was performed with clean

tweezers and special care was taken throughout to avoid the risk of cross con-

tamination. Repeated (more than 20) weighings showed that each sample corres-

ponded to 8.0 ± 0.6 mg of ZnS powder.

(i) Standards.

The preparation of standard mixtures of elemental sulphur and zinc

sulphide by weighing and mixing proved ineffective, hence the following

procedure was devised. Analar grade sulphur was dissolved in Analar grade

benzene to produce sulphur-in-benzene standards at 1000, 100, 10 and 1 ppm.

Aliquots of these standards were taken and a paste of 0.500 g of ZnS

(S° or S°) was made in a small Petri dish. The benzene was then evaporated

under a light vacuum (water pump) at room temperature in a clean vacuum

desiccator over silica gel. The dry mixture of S° and ZnS was then light-

ly mixed, placed in a tightly sealed sample bottle and stored in a second

vacuum desiccator until required. There was no detectable difference in

S° content between "freshly prepared" and "stored" (several weeks) samples.

The absence of peaks arising from benzene in the mass spectrum indicated

that evaporation of the benzene was complete. In passing it is worthwhile

noting that the author observed that acetone does not appear to be a

good solvent for sulphur, in discord with some published reports (85,

86) e.g. for contact times of up ,to X10 minutes between S and acetone,

no detectable sulphur peaks appeared in the mass spectrum. In con-

trast S° was rapidly and completely soluble in benzene under the same

conditions.

Standards were prepared at 0.4, 0.1, 0.04, 0.02, 0.015, 0.01, 0.005,

0.002 and 0.001 weight % S°. The sensitivity of the mass spectrometer

was such that standards from 0.02 to 0.001 wt% were normally used. Sam-

ples for analysis were then placed in capillary tubes, as described above.

Reproducible results were obtained using this method.

(ii) Unknowns.

Samples were prepared in the sealed activation vessel previously

described at 298 K and at a background ionic strength of 10-1 M KNO3, un-

less otherwise noted. Activation was carried out for 30 minutes (no

detectable differences were noted during preliminary studies on Cu"

activation of ZnS at 60 to 120 minutes). At the end of this specified

time the particles were allowed to settle, then the supernatant was re-

moved by suction until only a small amount remained. The solid and resid-

ual supernatant were then carefully poured on to three thicknesses of

filter paper contained in a small Petri dish, washed with several cm3 of

the decanted supernatant and then several cm3 of Analytical grade methanol.

The influence of various washing techniques is discussed in Chapter 7. The

dish and its contents were then placed in a clean vacuum desiccator and

dried under a light water pump vacuum over silica gel. The time taken

for the transfer procedure (i.e. activation vessel to vacuum desiccator)

was less than 1 minute, during which time the system was exposed to air.

Carrying out the transfer process under high purity nitrogen in a glove-box

-64-

did not show any detectable differences in comparison with the first

technique. This second process was therefore abandoned in favour of the

first.

After drying, the unknown samples were lightly mixed and stored in

tightly sealed sample bottles in a vacuum desiccator over silica gel.

Again "freshly prepared" and "stored" (several weeks) samples showed no

detectable differences in S0 content. Analysis samples were subsequently

placed in capillary tubes, as described above.

A series of tests were carried out on the stability of S° at high pH.

Activation was performed at the specified pH for 30 minutes, the pH was

then increased to the required value and samples were withdrawn, with a

20 cm3 syringe, at specified times. The suspension was then filtered through

a 0.22p Millipore filter, dried and stored as before. A small but sufficient

quantity of solid was obtained in this fashion.

CHAPTER 4 THE ACTIVATION OF ZnS

WITH CuII

4.1 INTRODUCTION

N.B. In this and succeeding Chapters, MII should be understood to represent

the oxidation state +2 of the heavy metal M and may appear as free

aqueous M2+ or as a hydrolysis product such as M(OH)+ etc., depending

on pH. All species are considered to be in aqueous solution, unless

otherwise specified.

Copper in its +2 oxidation state is by far the most commonly used

activating agent in sphalerite flotation (1,2). The initial and major thrust

of this investigation has, therefore, been directed towards clarifying the

mechanism by which CuII reacts with sphalerite. This Chapter focuses on

the CuII-ZnS interaction and deals with kinetic information obtained under

a wide range of pH, solid concentration, CuII concentration and incident

light conditions.

4.2 RESULTS

4.2.1 CuII Uptake

4.2.1(a) Preliminary work.

(I) Planning In following the rate of reaction of CuII with ZnS,

several objectives must be considered.

Firstly, a measurable change in Cu2+ concentration must occur, yet be

consistent with the noted stability and response characteristics of the Cu2+

electrode. Secondly the "conditioning" period for CuII activation of sphal-

erite in a flotation circuit rarely, if ever, exceeds thirty minutes and is

generally shorter than this (1,2,4 ). Sphalerite flotation in practice is

therefore concerned with the properties of a "thin film" of "copper sulphide"

-66-

at the sphalerite-aqueous solution interface. Hence in this study an

arbitrary upper limit of two hours was set for the kinetic studies. Satis-

factory data were collected over the important 0 to 1 minute period. It

was felt that this time limit of 2 hours was sufficient to allow a study

of the formation of the "thin film", albeit in a model system and to relate

the results obtained to both existing studies and flotation practice. The

results reported below are consistent with these two objectives.

The influence of initial CuII concentration, ZnS concentration, init-

ial pH, ionic strength, incident light and pretreatment on CuII uptake is

reported in Section 4.2.1(b). The effects of these variables on exchange

ratios, ZnII release and pH changes are described in Section 4.2.2.

(II) Hydrolysis of CuII and ZnII

The hydrolysis behaviour of metal ions is well documented and a great

deal of information is available on the relevant stability constants (5,

6,7 ). The emphasis to date has been placed on the existence of mononuclear

hydroxo complexes (1,5,7 ). The importance of polynuclear hydroxo complexes,

whose formation may be slow (5 ) and which are frequently precursors to pre-

cipitate formation (5,7), must not be overlooked. It may be readily demon-

strated, however, that the concentration of such polymeric species is very

low in dilute systems (5,7). Furthermore the presence of these polymers

does not markedly alter the solubility characteristics of most oxides and

hydroxides (5). For the purposes of the present investigation, where only

moderately dilute systems have been considered, it was only necessary to

consider the existence of mononuclear hydroxo complexes when examining the

general hydrolysis behaviour of CuII and ZnII. Self-consistent stability

constants, developed by James (7) from Sillen and Martell's published data

(6), were used to plot the equilibrium diagrams for CuII and ZnII shown in

Figures 4.1(a) & 4.1(h). The data used to obtain these diagrams is given in

-67 -

4 8 10 12

pH

FIGURE 4.1 (a) Equilibrium diagram for Curas a function of pH.

2 4 6 8 10 12

-68-

10

12

14

16

pH

FIGURE 4.1 (b) Equilibrium diagram for ZnII as a function of pH.

-69-

Table 4.1 and, while subject to some uncertainty, is about the best avail-

able at the time of writing.

Table 4.1 Thermodynamic stability constants for

CuII and ZnII hydrolysis at 298K (7).

Symbols are those given by Sillen (6).

log10K

CuII ZnII

K so -19.0 -15.5

K1 6.8 4.9

K2 6.9 6.6

K3 1.3 2.5

K4 1.0 1.3

The boundary of the hatched region represents total solution CuII and

ZnII and is generated from an equation of the form:

I [ M

I ] TOTAL Kso

w

[1.+]2 + KsoK1 [H+] + Ksô 1K2

Kw

+ K K1K2K3 Kw + K K1K2K3K4 Kw2 SO [Hf+T so [H

+TE

These diagrams will be referred to below.

4.2.1(b) Results

(I) Method of data presentation

The data are presented in Figures 4.2 to 4.12 and Tables 4.2 to 4.13

inclusive. Zinc sulphide samples S1 (BET area 0.72 m2 g-1) and S2 (BET area

0.68 m2 g-1) were used. The CuII uptake data represent the mean of two or

more completely separate experiments.

Cua

UP

TA

KE

(MO

LE

G'1

ZnS

x 1

0+8

)

0.5

1.5

1.0

FR

AC

TIO

N O

F M

ON

OLA

YE

R

0.5

1.0

1.5

FR

AC

TI O

N O

F M

ON

OL

AY

ER

10 h(0.120 min 1

,f

-70-

1

2 3

4

5

TIME (MINUTES)

60 120 TIME (MINUTES)

FIGURE 4.2 Cuauptake by ZnS as a function of time — varying ZnS concentration.

Conditions: Initial pH - 5.00 [Cur] initial - 10.24 x 10'6 M (Refer Table 4.2)

0 ZnS -5.00gri (Si) A ZnS-10.00g11 (S1)

Culr

UP

TA

KE

(MO

LE

G"1

Z

nS

x10 +

6)

0

0.5

1.0

FR

AC

TI O

N O

F M

ON

OL

AY

ER

Cuir undetectable in solution after this point.

1 2 TIME (MINUTES)

FIGURE 4.2 (C) Cumuptake as a function of time, where all Cuis consumed in one minute.

Conditions:

Initial pH = 5.00, ZnS = 3.00 g1-1 (S11; [Cum]initial = 1.50 x 10'5 M

-/l-

10

a10-5 min )

Cull U

PT

AK

E (

MO

LE

G"1

Z

nS

x 1

0+

6)

1.5

1.0

FR

AC

TIO

N O

F M

ON

OL

AY

ER

0.5

1

2 3 4 5

TIME (MINUTES)

10

1.5

1.0

0.5

FR

AC

TI O

N O

F M

ON

OL

AY

ER

60

120

TIME (MINUTES)

FIGURE 4.3 Cum uptake by ZnS as a function of time — varying initial Cu-rr concentration.

Conditions: ZnS = 5.00 g1"1 (S1 ) Initial pH = 5.00 (Refer Tables 4.2(A), 4.3(B))

1 [COI initial _ 5.12 x 10"5 M A [CUT) initial = 10.24 x 10'5 M

a10. 5 min ) 10

Cull U

PTA

KE

(M

OLE

0 Zn S

x 1

0+6

)

1

2 3

4

5 TIME (MINUTES)

10

b(0.120 min )

- 1.5

60 TIME (MINUTES)

CC W }

1.0 o z o 2 u. o z o

0.5 cc u.

120

—73—

FIGURE 4.4 Cu uptakeby ZnS as a function of time — constant [CuK] initial/[ZnS] ratio

Conditions: Initial pH = 5.00 (Refer Tables 4.2(B), 4.3(B))

0 ZnS = 5.000-1 (S1) A ZnS = 10.00 gl'/ (S1)

[Cull] initial = 5.12 x 10"5 M [Cull] initial = 10.24 x 10.5 M

1.5

1.0

0.5

FR

AC

TIO

N O

F M

ON

OL

AY

ER

10

5

CuF

UP

TAK

E (M

OLE

0 Zn S

x 1

0+

6)

0

1.5

W

1.0 g 0 z 0 u. O z 0

0.5 cc u.

2 3 4 5

TIME (MINUTES) 1

a (0-5 min)

b (0-120 min:)

10

1.5

COI-

UP

TA

KE

(MO

LE

G'1

ZnS

x 10

+ e


FIGURE 4.5 Cull uptake by ZnS as a function of time — varying initial pH

Conditions: ZnS = 5.00 gr1(S1) [Cull initial = 5.12 x 10'6 M (Refer Table 4.3 (A to C))

o n41 = d nn o initial nH = Ç nn 0 Initial off = 7.2n

4 1 2 3

TIME (MINUTES)

b (0-120 min.)

120 60

TIME (MINUTES)

a (0-5 min)

5

10

1.5

FIGURE 4.6 Cu uptake by ZnS as a function of time — varying initial pH.

Conditions: ZnS = 10.0001 (S1) [Cult]initial = 10.24 x 10'5 M (Refer Tables 4.2(6), 4.4)

0 Initial pH = 5.00 A Initial pH = 4.00

Cul

t U

PT

AK

E (

MO

LE

0 Zn

S x

4 )

10

5

1.5

1.0

0.5

FR

AC

T!O

N O

F M

ON

OL

AY

ER

—75—

8.5

8.0

7.5._...

7.0

6.5

Cue

+ E

LE

CT

RO

DE

PO

TEN

TIA

L (M

I LL

IVO

LT)

40

30

20

10

6.0

40 60 80 100 120

TIME (MINUTES)

FIGURE 4.7 Cull activation of ZnS. Variation of Cue+ electrode potential with time at initial pH = 8.10. pH readjusted to near initial pH at two stages during experiment. (Refer Table 4.3(D))

Conditions:[Cu J initial ` 5.12 x 10- M

ZnS = 5.00 gl't (S2 )

A pH 0 Cue+ electrode potential

g10-5 min)

4 5

120 60

TIME (MINUTES)

FR

AC

TIO

NN

OF

MO

NO

LA

YE

R

1.0

0.5

2 3

TIME (MINUTES)

610-120 min.)

1

O

CulE

UP

TA

KE

(MO

LE

0 Zn S

x 1

0+8

)

1.5

10.

10

—77—

FIGURE 4.8 Clair uptake by ZnS as a function of time — varying ionic strength (I)

Conditions: Initial pH = 5.00 (Cur] initial = 5.12 x 10'6 M (Refer Table 4.6)

ZnS = 5.00 g11 (S2)

OI= 10'1 M KNO3 r= 10-3 M KNO3

1.5

1.0

0.5

FR

AC

TIO

N O

F M

ON

OL

AY

ER

5

1.5

¢ W }

g o- z

1.0 0 2

FR

AC

TIO

N OF

0.5

-78-

10

Cuu

UP

TAK

E (

MO

LE 0

ZnS

x 1

0+

4 o:

u. 0.5

1

2 3 4 TIME (MINUTES)

10


FIGURE 4.9 CuTuptake by ZnS as a function of time — varying initial pH

Conditions: [CA initial = 5.12 x 10"5 M ZnS = 5.00 el (S2) (Refer Tables 4.6(A), 4.7)

o Initial pH = 5.00 o Initial pH = 4.00

Culr

UP

TA

KE

(M

OL

E G

" 1 Z

nS

x 1

0 +6)

10 -

5

FR

AC

TIO

N OF

MO

NO

LA

YE

R

1.0

0.5

Cul

l U

PT

AK

E (M

OLE

G''

Zn

S x

10+

6)

1.5

0.5

1.0

FR

AC

TIO

N OF

MO

NO

LA

YE

R

-79-

CuT

IUP

TA

KE

(MO

LE

G'1

ZnS

x 1

0+6

)

1

2 3

4

TIME (MINUTES)

1.5

60

120

TIME (MINUTES)

FIGURE 4.10 CUM uptake by ZnS as a function of time — varying light conditions.

Conditions: Initial pH = 5.00 [CuTC1 initial = 5.12 x 10'6 M

ZnS = 5.00 g11 (S2 )

Refer Tables 4.6 (A), 4.8 (A), 4.9 (A) )

°> normal light A in dark

I I under UV

10

5

0

A • V

i

1.5

1.0

0.5

FR

ACT

ION

OF

MO

NO

LA

YE

R

CJ

EU

PTA

KE

(MO

LE

0 Z

nS

x 1

0+8

)

10

A

0

1.5

—80—

a(0-5 min )

1

2 3

4

5

TIME (MINUTES)

b(0. 120 min 1


FIGURE 4.11 Cumuptake by ZnS as a function of time — varying light conditions

Conditions: Initial pH = 4.00 [Cum) initial = 5.12 x 10-6 M

ZnS = 5.00 gl".' (S2 ) Refer Tables 4.7, 4.8(B), 4.9(B)

O normal light 0 in dark ❑ under UV

60 120

TIME (MINUTES)

FIGURE 4.12 (A) Cumuptake by ZnS as a function of time — varying pretreatment experiments.

Conditions: initial pH = 5.00 ZnS = 5.00 g1"1 (S2 )

[Cur] initial = 5.12 x 10"5 M

(Refer Tables 4.6 (A), 4.10)

0 fresh ZnS (no pretreatment) A ZnS pretreated at pH 5 D ZnS pretreated at pH 2

10

1.5

1.0

0.5

Cul

rU

PT

AK

E ( M

OL

E G'1

ZnS

x 10

+6

)

FR

AC

TIO

N O

F M

ON

OL

AY

ER

60

120 TIME (MINUTES)

FIGURE 4.12 (B) Cum uptake by ZnS as a function of time — varying pretreatment experiments.

Conditions: Initial pH = 5.00 ZnS = 5.00 gI"t (S2 )

[Curt] initial = 5.12 x 10"5 M

(Refer Tables 4.6 (A), 4.11)

0 fresh ZnS (no pretreatment) A ZnS pretreated at pH 2

Cut

rUP

TA

KE

(MO

LE

0 Z

nS

x 10+

6)

10

5

0

1.5

1.0

0.5 FR

AC

TIO

N O

F M

ON

OL

AY

ER

CIF

UP

TA

KE

(M

OL

E G

'1 Z

nS

x 1

0+

10

1.5

1.0

FR

AC

TIO

N O

F M

ON

OL

AY

ER

0.5

60

120 TIME (MINUTES

FIGURE 4.12 (C) CuII uptake by ZnS as a function of time — varying pretreatment experiments.

Conditions: initial pH = 5.00 ZnS = 5.00 gl'l (S2)

(Refer Tables 4.10 (A), 4.12, 4.13)

o ZnS pretreated at pH 5, normal light.

A ZnS pretreated at pH 5, in dark.

ZnS conditioned at pH 5, under UV.

0 ZnS pretreated at pH 2, no adjustment to pH 5

[CO II initial = 5.12 x 10'5

-83- TABLE 4.2

(A)

CuII ACTIVATION OF ZINC SULPHIDE

pHinitial = 5-00

Time minutes

II ( ]residual mole k-1 x 10 +5

II (Cu ]_exchang d mole Z-1 x 10 5

II CuII

mole g-1 x 10+6

pH Fraction of Monolayer

0.0 10.24 - - 5.00 -

0.25 8.96 1.28 2.56 0.45 0.50 8.45 1.79 3.58 0.62 0.75 8.10 2.14 4.28 0.75 1.0 7.65 2.59 5.18 0.90 1.5 7.20 3.04 6.08 1.1 2.0 7.04 3.20 6.40 1.1

3.0 6.97 3.27 6.54 1.1 4.0 6.92 3.32 6.64 1.2

5.0 6.86 3.38 6.76 5.78 1.2 15.0 6.65 3.59 7.18 5.80 1.3 30.0 6.38 3.86 7.72, 6.30 1.3 60.0 6.20 4.04 8.08 6.35 1.4

120.0 6.11 4.13 8.26 6.40 1.4

(ZnS] = 5.00 gQ 1(S1). Normal light. 298K

(B) CuII ACTIVATION OF ZINC SULPHIDE

pHinitial = 5.00

i

!1 Time ; [CuII]residual Minutes -1 10+5 mole R x 10+ [Cu'1]excha q d

-1 5 mole Q x 10 CuII Uptake

-1 6 mole g x 10+


0.0 10.24 - - 5.00 -

0.25 7.55 2.69 2.69 0.52 0.50 6.70 3.54 3.54 0.62 0.75 5.46 4.78 4.78 I 0.83 1.0 4.86 5.39 5.39 16.13 0.94 1.5 4.49 5.76 5.76 1.0 2.0 4.11 6.14 6.14 1.1 3.0 3.88 6.37 6.37 1.1 4.0 3.66 6.59 6.59 Î 1.2 5.0 3.50 6.75 6.75 6.21 1.2 15.0 2.80 7.44 7.44 16.21 1.3 30.0 2,28 7,96 7.96 i 6.30 1.4 60.0 2.16 8.08 8.08 6.38 1.4 120.0 2.10 8.14 8.14 6.40 1.4

[ZnS] = 10.00 gß-1 (S1) Normal light 298K

-84- TABLE 4.3

(A) CuII ACTIVATION OF ZINC SULPHIDE

pH initi = 4.00 al

Time minutes

[CuII] residua

mole Z-1 x 10+5

[CuII] exchang d mole 9.-1 x 10 5

CuII Uptake

mole g-1 x 10+6


0.0 5.12 - - 4.00 -

0.25 3.77 1.36 2.71 0.47

0.50 3.21 1.91 3.82 0.67

0.75 2.90 2.22 4.44 0.77

1.0 2.64 2.48 4.96 4.04 0.86

1.5 2.45 2.67 5.34 0.93

2.0 2.36 2.76 5.52 0.96

3.0 2.18 2.94 5.88 1.0

4.0 2.16 2.96 5.92 1.0

5.0 2.12 3.00 6.00 1.1

15.0 1.88 3.24 6.48 4.10 1.1

30.0 1.88 3.24 6.48 4.15 1.1

60.0 1.88 3.24 6.48 4.18 1.1

120.0 1.88 3.24 6.48 4.22 1.1

[ZnS] = 5.00 g!C 1 (S1) Normal light 298K

(B) CuII ACTIVATION OF ZINC SULPHIDE

pH. nitial - 5.00

Time [CuII] residua' mole Z-1 x 10 5 minutes

[CuII] e_xchang d

mole Z-1 x 10 5

CuII Uptake _ mole g-1 x 10

+6 pH Fraction of

Monolayer

0.0 5.12 - - 5.00 -

0.25 3.92 1.20 2.40 0.36

0.50 3.26 1.86 3.72 0.65

0.75 2.81 2.31 4.62 0.80

1.0 2.59 2.53 5.05 5.50 0.88

1.5 2.25 2.87 5.73 0.94

2.0 2.15 2.97 5.94 1.0

3.0 2.08 3.04 6.08 1.1

4.0 1.97 3.16 6.31 1.1

5.0 1.88 3.25 6.49 5.84 1.1

15.0 1.62 3.50 7.00 5.91 1.2

30.0 1.20 3.92 7.84 5.98 1.4

60.0 1.05 4.07 8.14 6.30 1.5

120.0 1.01 4.11 8.21 6.30 1.5

[ZnS] = 5.00 g2.-1 (S,) Normal light 298K

-85-

TABLE 4.3

(C)

Cu" ACTIVATION OF ZINC SULPHIDE

PH initial 7.2 al

Time minutes

[Cu k-1 x 10+5

II]exchanged [Cu

mole Z-1 x 10+5

Cu" Uptake

mole g-1 x 10+6


0.0 5.12 - _ 7.20 _

0.25 2.80 2.32 4.64 0.81

0.50 2.23 2.89 5.78 1.0

0.75 1.66 3.46 6.92 1.2

1.0 1.33 3.79 7.53 7.35 1.3

1.5 0.94 4.20 8.40 1.5

2.0 0.76 4.36 8.72 1.5

3.0 0.62 4.50 9.00 1.6

4.0 0.57 4.55 9.10 1.6

5.0 0.51 4.61 9.22 7.42 1.6

15.0 0.30 4.82 9.64 7.65 1.7

30.0 7.63

60.0 7.62

120.0 0.20 4.92 9.84 7.62 1.7

[ZnS] = 5.00 g2 1 (S1) Normal light 298K

-86-

TABLE 4.3

CuII ACTIVATION OF ZnS

(D)

Cu2+ electrode pH

0 50.3 8.10

0.5 34.8 8.36

1.0 26.7 8.44

2.0 19.0 8.60

3.0 15.0 8.67

4.0 12.5 8.72

5.0 10.5 8.75

6.0 8.9 8.78

7.0 7.3 8.79

8.0 6.2 8.80

9.0 5.8 8.80

10.0 5.0 8.80

11.0 4.3 8.85

26.0 0 8.85

pH reduced to 8.08

27.0 37.0 8.08

35.0 14.0 8.33

43.0 8.5 8.45

51.0 7.0 8.48

59.0 5.5 8.49

pH reduced to 8.03

60.0 18.0 8.03

89.0 14.5 8.11

120.0 13.2 8.11

Time minutes

[ZnS] = 5.00 0,-1 (S2) Normal light 298K

TABLE 4.4

Cu" ACTIVATION OF ZINC SULPHIDE

pHiniti = 4.00 al

Time minutes

[Cu II]residual

Q,-1 x 10+5 (CuII]exchanged mole R-1 x 10+5

CuII Uptake

mole g-1 x 10+6


0.0 10.24 - - 4.00 -

0.25 7.69 2.55 2.55 0.44

0.50 6.59 3.65 3.65 0.64

0.75 6.06 4.18 4.18 0.73

1.0 5.66 4.58 4.58 0.80

1.5 5.43 4.81 4.81 0.84

2.0 5.20 5.04 5.04 0.88

3.0 5.00 5.24 5.24 0.91

4.0 4.80 5.44 5.44 0.95

5.0 4.60 5.64 5.64 0.98

15.0 4.07 6.17 6.17 4.10 1.1

30.0 3.65 6.59 6.59 1.2

60.0 3.65 6.59 6.59 1.2

120.0 3.65 6.59 6.59 4.30 1.2

[ZnS] = 10.00 gk-1 (Si) Normal light 298K

-88-

TABLE 4.5


pHiniti = 5.00

al

Time minutes

[CuII]residual [Cull]

exchanged

CuII Uptake pH Fraction of Monolayer

0.0

0.25

0.50

0.75

1.0

1.5

2.0

3.0

4.0

5.0

15.0

30.0

60.0

120.0

1.03

0.86

0.71

0.65

0.58

0.49

0.46

0.40

0.37

0.37

0.32

0.29

0.26

0.22

-

0.17

0.32

0.38

0.45

0.54

0.57

0.63

0.66

0.66

0.71

0.74

0.77

0.81

-

1.7

3.2

3.8

4.5

5.4

5.7

6.3

6.6

6.6

7.1

7.4

7.7

8.1

5.00

5.15

-

0.30

0.56

0.66

0.78

0.94

0.99

1.1

1.2

1.2

1.2

1.3

1.3

1.4

[ZnS] = 1.00 g2-1 (S1) Normal light 298K

-89-

(A)

TABLE 4.6.


pHiniti = 5.00 al

Time minutes

[CuII] residual

mole Z-1 x 10+5

II [Cu J exchanged mole Z-1 x 10+5

Cu" Uptake

mole g-1 x 10+6


0.0 5.12 - - 5.00 - 0.25 4.04 1.08 2.16 0.40 0.50 3.27 1.85 3.70 0.68

0.75 2.79 2.33 4.66 0.86 1.0 2.52 2.60 5.20 5.32 0.96 1.5 2.31 2.81 ,5.62 1.0

2.0 2.29 2.83 5.66 1.0

3.0 2.02 3.10 6.20 1.1

4.0 1.95 3.17 6.34 1.2

5.0 1.88 3.24 6.48 5.33 1.2

15.0 1.58 3.54 7.08 5.40 1.3

30.0 1.50 3.62 7.24 5.90 1.3

60.0 1.28 3.84 7.68 6.36 1.4 120.0 1.18 3.94 7.88 6.40 1.5

[ZnS] = 5.00 gZ-1 (S2) Normal light 298K

(B) Cu" ACTIVATION OF ZINC SULPHIDE

pH .=- 5.00

Time minutes

[CuII] residual

mole R-1 x 10+5

[CuII] exchanged

mole R-1 x 10+5

CuII Uptake

mole g-1 x 10+6

pH Fraction of Monola er Y

0.0 5.12 - - 5.00 -

0.25 4.20 0.92 1.8 0.33

0.50 3.54 1.58 3.16 0.58 0.75 3.15 1.97 3.94 0.73 1.0 2.67 2.45 4.90 5.41 0.78 1.5 2.57 2.55 5.10 0.94

2.0 2.37 2.75 5.50 1.0 3.0 2.26 2.86 5.72 1.1 4.0 2.17 2.95 5.90 1.1 5.0 1.88 3.24 6.47 5.45 1.2 15.0 1.81 3.31 6.62 5.47 1.2 30.0 1.57 3.55 7.10 5.65 1.3 60.0 1.51 3.61 7.22 5.99 1.3 120.0 1.32 3.80 7.60 6.20 1.4

[ZnS] = 5.00 0-1 (S2) I = 10-3M KNO3 Normal light 298K

TABLE 4.7


pH initi = 4.00 al

Time [CuII] residual mole ß- x 10+5 minutes

[Cula] exchanged mole Z- x 10+5

CuII Uptake -1 +g mole g x 10


0.0 5.12 - - 4.00 -

0.25 3.91 1.21 2.42 0.45

0.50 3.20 1.92 3.84 0.71

0.75 2.98 2.14 4.28 0.79

1.0 2.74 2.38 4.76 0.88

1.5 2.62 2.50 5.00 0.92

2.0 2.48 2.64 5.28 0.97

3.0 2.30 2.82 5.64 1.0

4.0 2.22 2.90 5.80 1.1

5.0 2.19 2.93 5.86 4.03 1.1

15.0 1.95 3.17 6.34 4.04 1.2 30.0 1.91 3.21 6.42 4.06 1.2 60.0 1.91 3.21 6.42 4.10 1.2 120.0 1.91 3.21 6.42 4.19 1.2

[ZnS] = 5.00 gQ-1 (S2) Normal light 298K

-91- TABLE 4.8

CuII ACTIVATION OF ZINC SULPHIDE (A)

pHinitial - 5.00

Time [CuII] residual

mole k- x 10+5 minutes

[CuII] exchanged mole x 10+5

CuII Uptake + mole g_

1 x 10 6


0.0 5.12 - - 5.00 -

0.25 4.20 0.92 1.8 0.33

0.50 3.50 1.62 3.24 0.60

0.75 3.00 2.12 4.24 0.78

1.0 2.60 2.52 5.04 5.40 0.93

1.5 2.18 2.94 5.88 1.1

2.0 1.99 3.13 6.26 1.2

3.0 1.77 3.35 6.70 1.2

4.0 1.63 3.49 6.98 1.3

5.0 1.65 3.47 6.94 5.45 1.3

15.0 1.29 3.83 7.66 5.50 1.4

30.0 1.14 3.98 7.96 5.69 1.5

60.0 1.00 4.12 8.24 6.39 1.5

120.0 0.87 4.25 8.50 6.40 1.6

[ZnS] = 5.00 gk-1 (S2) In dark

298K

CuII ACTIVATION OF ZINC SULPHIDE (B)

pHiniti 4.00 al

Time minutes

[CuII] risidua+5

mole k- x 10+5

[CuII] exchanged mole k- x 10+5

CuII Uptake _1 +g mole g x 10+6


5.12 - - 4.00 -

0.25 4.10 1.02 2.04 0.38

0.50 3.50 1.62 3.24 0.60

0.75 2.88 2.24 4.48 0.83

1.0 2.61 2.51 5.02 4.04 0.92

1.5 2.39 2.73 5.46 1.0

2.0 2.22 2.90 , 5.80 1.1

3.0 2.00 3.12 6.24 1.2

4.0 1.92 3.20 6.4,0 1.2

5.0 1.87 3.25 6.50 4.04 1.2

15.0 1.64 3.48 6.96 4.05 1.3

30.0 1.59 3.53 7.06 4.07 1.3

60.0 1.58 3.54 7.08 4.10 1.3

120.0 1.58 3.54 7.08 4.15 1.3

[ZnS] = 5.00 gk-1 (S2) In dark 298K

TABLE 4.9

(A)


pH initi = 5.00 al

Time minutes

II [Cu ]

residual mole k-1 x 10+5

[CuII] exchanged

mole 2,-1 x 10+5

CuII Uptake

mole g-1 x 10+6

pH ..

Fraction of Monolayer Y

0.0 5.12 - - 5.00 -

5.0 0.98 4.14 8.28 5.76 1.5

15.0 0.84 4.28 8.56 5.98 1.6

30.0 0.80 4.32 8.64 6.30 1.6

60.0 0.77 4.35 8.70 6.60 1.6

120.0 0.77 4.35 8.70 6.70 1.6

[ZnS] = 5.00 gk-1 (S2) UV irradiation 298K

(B)


pH initi = 4.00 al

Time minutes

[CuII] residual mole k-1 x 10+5

[CuII] exchanged

mole k-1 x 10+5

CuII Uptake

mole g-1 x 10+6

pH Fraction of Monolayer Y

0.0 5.12 - - 4.00 -

5.0 1.83 3.29 6.58 4.01 1.2 15.0 1.83 3.29 6.58 4.02 1.2 30.0 1.85 3.27 6.54 4.10 1.2 60.0 2.00 3.12 6.24 4.20 1.2 120.0 2.00 3.12 6.24 4.30 1.2

[ZnS] = 5.00 gk-1 (S2) UV irradiation 298K

TABLE 4.10

(A) PRETREATMENT EXPERIMENT


pH initi = 5.00 al

Time minutes [CuII]residual

mole Z.-1 x 10+5

[CuII]exchanged mole k-1 x 10+5

CuII Uptake

mole 1 +6 g! x 10


0.0 5.12 - - 5.30 -

5.0 1.66 3.46 6.92 6.24 1.3

15.0 1.46 3.66 7.32 6.30 1.4

30.0 1.32 3.80 7.60 6.30 1.4

60.0 1.14 3.98 7.96 6.32 1.5

120.0 1.04 4.08 8.16 6.32 1.5

[ZnS] = 5.00 gZ-1 (S2) Normal light 298K

[ZnS pretreated at initial pH 5.0, then CuII added]

(B) PRETREATMENT EXPERIMENT


pH initial 5.00 al

Time minutes

[CuII]residual

5 mole Z.' x 10+5

[CuII] eichange+d5 2 x .10

II Uptake Cu _1

mole g x 104-6


0.0 5.12. - - 5.00 -

5.0 2.71 2.4 4.82 5.24 0.89

15.0 2.71 2.41 4.82 5.33 0.89

30.0 2.63 2.49 4.98 5.41 0.92

60.0 2.52 2.60 5.20 5.50 0.96

120.0 2.51 2.60 5.20 5.54 0.96

[ZnS] = 5.00 gQ 1 (S2) Normal light 298K II

(ZnS pretreated at pH 2.0 then pH increased to 5.0 and Cu added]

TABLE 4.11

PRETREATMENT EXPERIMENT


pH initi = 5.00 al

Time minutes

[CuII] residual

[Cu"] exchanged

CuII Uptake pH Fraction of Monolayer

0.0 5.12 - - 5.00 -

5.0 1.53 3.59 7.18 5.85 1.3

15.0 1.35 3.77 7.54 6.02 1.4

30.0 1.26 3.86 7.72 6.11 1.4

60.0 1.10 4.02 8.04 6.19 1.5

120.0 0.96 4.16 8.32 6.21 1.5

Normal light 298 K

(ZnS pretreated at pH 12.0, pH reduced to 5.0 and CuII added]

(ZnS] = 5.00 gk 1 (S2)

TABLE 4.12

(A)



PH initi = 5.00 al

Time minutes

[CuII] residual mole Z-1 x 10+5

(GolT]exchanqed mole 271 x 10+5

CuII Uptake

mole g-1 x 10+6


0.0 5.12 - - 5.19 -

5.0 1.61 3.51 7.02 6.30 1.3

15.0 1.42 3.70 7.40 6.30 1.4

30.0 1.23 3.89 7.78 6.30 1.4

60.0 1.04 4.08 8.16 6.32 1.5

120.0 1.01 4.11 8.22 6.30 1.5

[ZnS] = 5.00 g9.-1 (S2) in dark 298 K

[ZnS pretreated at initial pH 5.0, CuII added]

(B)



pHinitial = 5.00

f Time

minutes

-- [Cull]

residual mole Q-1 x 10+5

[CuII] exchanged

mole R-1 x 10+5

CuII Uptake

mole g-1 x 10+6


0.0 5.12 - - 7.15 - i 5.0 1.0 4.11 8.22 7.59 1.5

15.0 0.63 4.49 8.98 7.59 1.7 30.0 0.41 4.71 9.42 7.40 1.7 60.0 <0.01 5.11 10.2 7.20 1.9 120.0 <0.01 5.11 10.2 7.20 1.9

I

[ZnS] = 5.00 g2 1 (S2) UV irradiation 298 K

[ZnS pretreated at initial pH 5.0 under UV, then CuII added]

TABLE 4.13



pHinitial = 2.0

Time minutes

[CuII] residual

mole R-1 x 10+5

[CuII] exchanged

mole Q-1 x 10+5

CuII Uptake

mole g-1 x 10+6

pH Fraction of Nbnola yer

0.0 5.12 - - 2.00 -

5.0 1.18 3.94 7.88 1.98 1.5

15.0 1.44 3.68 7.36 1.98 1.4

30.0 2.42 2.70 5.40 1.98 0.99

60.0 2.83 2.29 4.58 1.98 0.84

120.0 3.43 1.69 3.38 1.98 0.62

[ZnS] = 5.00 gt-1 (S2) Normal light 298K

[ZnS pretreated at initial pH 2.0, then CuII added]

-97-

Figures 4.2 to 4.12 show CuII uptake as a function of time for the

time spans 0-5 and 0-120 minutes. CuII uptake is expressed as mole per

gram of ZnS and as the fraction of monolayer covered. Monolayer coverage is

based on an available area of 20.82 for each CuII species and is calculated

on the assumption that the ZnS surface is composed predominantly of 110

cleavage faces, each of which has two cationic sites per unit of surface

7.67 x 5.42 Â2 (8,36). For the zinc sulphide samples used, monolayer cover-

age occurs at the following CuII uptake values:

S1(0.72 m2 g-1) : 5.8 x 10-6 mole g-1

S2(0.68 m2 g-1) : 5.4 x 10-6 mole g-1

It should be noted that this calculation of "monolayer coverage" presupposes

precise knowledge of the surface geometry, a desirable but as yet unattainable

situation for systems such as those encountered in the present investigation.

The fraction of monolayer covered, therefore, merely serves as an estimate

of the depth of penetration of the zinc sulphide crystal lattice.

CuII uptake values, when expressed as mole per gram, are subject to

an error of ± 5%, a feature which must be borne in mind when interpreting the

kinetic data.

Tables 4.2 to 4.13 show CuII uptake (mole g-1 and fraction of mono-

layer) as a function of time together with residual and exchanged CuII

concentrations. The relevant pH values for the CuII-ZnS system are also tab-

ulated as an aid in commenting on CuII hydrolysis (pH changes during activat-

ion and for the blank are dealt with in Section 4.2.2).

(II) Summary of the major results

(i) During the CuII-ZnS activation reaction, CuII uptake occurs

in 2 stages - a fast step, generally completed within.1.5 minutes and

leading to monolayer coverage followed by a slow step which, up to 2

hours never exceeds 2 monolayers and is generally about 1.5. A pH

increase accompanies this Cu" uptake. The initially white ZnS turns

pale brown during activation; the intensity of the colour is more pro-

nounced at acid pH values, other conditions being held constant. X-ray

diffraction studies performed on dry, activated zinc sulphide revealed

no detectable changes in either lattice spacings or line intensity from

the untreated,cubic ZnS lattice.

(ii) Provided that the initial rapid step is complete (e.g. see

Figure 4.2(c)) Cu" uptake is independent of initial Cu11

concentration

and of ZnS concentration at the same initial acid pH values.

(iii) At acid pH values, Cu" uptake during the slow, second stage

of activation is greater at an initial pH of 5.0 than at 4.0, although

it is the same during the first step.

(iv) At acid pH values Cu" uptake is independent of ionic

strength and of stirring speed (provided that the ZnS remained in sus-

pension).

(v) At acid pH values Cu11 uptake is increased slightly in the

dark in comparison with uptake values registered in normal light during

the slow stage of activation; the initial step is unaffected. W

irradiation enhances Cu" uptake during the second stage of the reaction

at an initial pH of 5.0 , but gives similar Cu" uptake values to those

recorded in normal light and in the dark at an initial pH of 4.0.

(vi) Pretreatment of ZnS at initial pH values of 5.0 and 12.0

has essentially no effect on CuII uptake, nor does pretreatment in the

dark at an initial pH of 5.0. Pretreatment of ZnS under W irradiation

at an initial pH of 5.0 leads to enhanced CuII uptake. Pretreatment at

pH 2.0 followed by activation at pH 5.0 or at 2.0 leads to reduced

CuII uptake.

III Subsidiary effects

(i) If, as shown in Figure 4.2(c), the ZnS concentration was

fixed at 3.00 g2,-1 and the initial CuII concentration was reduced to,

say, 1.50 x 10-5 M, all of the available CuII was consumed in the init-

ial rapid phase of the reaction. (i.e. residual CuII was less than 10-7

M or 0.01 ppm). The measured CuII uptake values were, within experi-

mental error, the same as those measured for higher initial CuII

concentrations, up to the point where the CuII was consumed.

On the other hand, if the initial CuII concentration was increased

to 10-3 or 10-2 M, there was such a slight change in the CuII concentrat-

ion over the 2 hour period that any calculation of CuII uptake became

impracticable and the initial, rapid phase of the reaction was obscured.

At no stage can the systems be regarded as having reached equilibrium -

rather was reaction continuing at a slow, but detectable rate. For

example examination of a system originally having an initial CuII con-

centration of 5.12 x 10-5 M, initial pH of 5.0 and ZnS concentration of

5.00 gC 1 revealed that, after continuous agitation for four days, there

was no detectable(< 10-7 M Culi)left in solution (the final pH was 6.3).

A similar effect was noticed at an initial pH of 4.0.

(ii) For the systems shown in Figures 4.2 to 4.3, if, at the con-

clusion of the 2 hour reaction period the CuII concentration was returned

to its original value (by titrating the Cut+ electrode back to its

initial potential reading), it was discovered that there was no detect-

able alteration in reaction rate. A similar result was obtained at

shorter time intervals'(30 and 60 minutes) and at an initial pH of 4.0.

(iii) Opening several of the ports in the top of the reaction

vessel just prior to injecting the ZnS, and allowing these to remain

open for up to 10 minutes of the reaction period, showed that the ab-

sorption of some 02 and CO2 by the previously purged system had no

detectable effect on CuII uptake. This applied to ZnS concentrations of

5.00 anid 10.00 50.-1 at an initial pH of 5.0.. pH .changes above 5.5 were

reduced.

(iv) In a separate experiment, ZnS (5.00 g9 1) was activated at

an initial pH of 5.0 and initiai CuII concentration of 5.12 x 10-5 M

for 2 hours. The activated ZnS was then filtered through a 0.22µ Milli-

pore filter, washed with 15 cm3 of degassed, 0.1M KNO3 (adjusted to the

same pH as the activating system from which the ZnS was removed) and

partially dried in a vacuum desiccator. The damp powder was then in-

jected into a fresh activating solution of the same initial pH and CuII

concentration as that above. There was no detectable removal of CuII

from solution over a subsequent 2 hour period. In fact the Cut+ elec-

trode potential drifted only by 0.1 millivolt during this period; AAS

analysis revealed no change in CuII from the initial value. These

results suggest that the reaction product(s) formed when CuII interacts

with ZnS is, apparently, not readily removable and that further reaction

of CuII with the previously activated ZnS lattice is strongly retarded.

(v) alkaline initial pH (8.10) Figure 4.7, Table 4.3(D)

At a pH of 8.1 (or above) and a CuII concentration of

5.12 x 10-5 M, only a very small amount of free Cu2+ exists in aqueous

solution (refer Figure 4.16). The bulk of the Cu" is present as

Cu(OH)2 solid, in equilibrium predominantly with Cu(OH)2aq and Cu(OH)+.

For this reason it is impossible to follow quantitatively the rate of

reaction of Cu" with ZnS using a Cu2+ electrode.

Under the experimental conditions used, when ZnS was added to a CuII

solution at a pH of 8 or above, copper was not detectable in solution by

AAS (i.e. 10-7 M),(although it could be detected prior to ZnS addition)

This applied to filtered samples taken as soon as possible after ZnS

addition and to the analysis of samples of supernatant (obtained by allow-

ing the ZnS to settle out in the reaction vessel and sampling the super-

natant with a syringe).

The Cu2+ electrode is useful as a qualitative probe because it will

still respond to changes in Cu2+ concentration as the reaction proceeds.

Hence Figure 4.7 shows the variation in Cu2+ electrode potential and pH

as a function of time at an initial pH of 8.1. Initially there was a

rapid decrease in Cut+ electrode potential (i.e.a decrease. in [Cu2 )aq),

accompanied by a rapid increase in pH. Both rates decreased over 26

minutes. At this juncture the pH was decreased to about its original

value, whereupon the Cu2+ electrode potential rose rapidly to within

13 millivolt of its original value. The electrode potential then decreas-

ed (pH increased) over the next 33 minutes, more slowly than in the init-

ial 26 minute period. Decreasing the pH once again to near its initial

value at 59 minutes resulted in similar behaviour to that which followed

the first pH adjustment after the reaction commenced i.e. the electrode

potential very slowly decreased, whilst the pH increased, over the period

of 60 to 120 minutes.

-102-

4.2.2 Exchange ratio, ZnII release and pH change

4.2.2(a) Method of data presentation

Tables 4.14 to 4.18 inclusive contain additional information which is

complementary to that already shown in Figures 4.2 to 4.12 and Tables 4.2

to 4.13. The information may be readily correlated thus:

Table 4.14 (except D) relates to Figures 4.2(a & b) to 4.6 and/or

Tables 4.2 to 4.3(A to C), 4.4 and

4.5.

Tables 4.15 to 4.19 (except 4.15(E), 4.17(C)) relate to Figures 4.8 to

4.12 and/or Tables 4.3(D) and 4.6

to 4.13.

Tables 4.14(D), 4.15(E) and 4.17(C) contain relevant information whose

prior display has been inappropriate up to this point.

The data in Tables 4.14 to 4.18 show the pH during activation, pH of

the blank, total zinc, the zinc contribution from the blank, total zinc, the

zinc exchanged, residual copper, copper exchanged and the "exchange ratio" (R)

as a function of time. Generally the data have been determined at 5, 15, 30,

60 and 120 minutes with the exception of some early results shown in Table

4.14. Five minutes was, of course, the earliest convenient time at which

samples could be taken for AAS analysis. The "exchange ratio", is defined as

II] —

II R = [Cu i

nitial [Cu ]residual II — II

[Zn ]total [Zn ]blank

i.e. R = [Culll exchanged II

[Zn ]exchanged

Thus any contribution to [ZnII]total

from the blank is subtracted out in this

definition of R. For each system the data of course represent the mean of two

TABLE 4. 14

Cu"

ACTIVATION OF ZIN C SULPHIDE

-103—

-. C H EO r-, U I C th C d N "-" 0 01

U

O O

in

O O

ul

O O

O ri

O O

!")

1

z

0. 78 ± 0. 08 I

1

O O

O

+ 1 d' Co O

CO O

O

+ 1 r r

O

CO O

O

+ I 01 n

O 0.87 ± 0.09

I

CO 0 O

+I

M Co

O

7:1 O Mr) C+ rt0 .z r-i U X X r-, CI) I

2 H H (D

Z r-I U O — E

I

O Ln

• M

l O

• V'

r-I r-I

• d'

I

01 ill

M

V O

V

M ri

d' I

c)' V'

r

CO O

CO I

W In

• N

r-I Id in g+

R7 O -ri ri N X N •-+ $.I I

d H H (U

r♦ U Q

N ri

in

N 0

ri

Ln 0

ri

ri 0

ri

V N

O r-I

tf) W

W

O N

W

ri r-I

l0

d' N

O ri

O CO

N

W ri

N

N r-I

In

W In N

II

[Zn ]exchanged

mole 2,-1 x10+

5

I

r-I r) V'

O N

In

r-I N

In I

In N

d'

l0 N

tn

l0 N

Ln I

d' In

co

I CO

CO I

O ,--1

• m

u) -i-

.x O

td X ri-y .p 1

d H H u)

C r-i N

S

I

l0 co

O

d' O

rI

co rI

• ri

I

ko CO

• O

d' O

• ri

co r1

ri I

r- O

• ri

CO M

• ri

I

O CO

• O

II

[

] total

mole 2.-

1x10+

5

I

I r-I

In

d' N

l0

W fr) W

i ri r-1

• In

O M

W

d' d'

W I

ri l0

01

In N

O ri

I

O O1

• M

.X C b

,-i A

O O

in

N N

In

W N

tn

CO N

tn

O O

In

N N

Ln

W N

in

CO N

In

O O

In

O M

In

In d'

ln

O O

Li)

to N

In

O O

ul

r♦ 01

In

O M MO

O M

LD

O 0

N

O CO

In

In M

\o

O d lo

O O

tn

r-I N

LO

CO M

\O

O 0

tn

N 01

In

N Iq E C

-ri-I

• O In

H O

O

0 In ri

O l0

O

rN-I

O In r--I

0 l0

O O

r-I

U

A

TABLE 4.14

ACTIVATION OF ZINC SULPHIDE

H H

U

—104—

-- c --. tfl '-. U I c to c eq N ---- O 17 1

U

o O in

o O

O ri

o O tf1

0; I

O

O

i-1

01 CO

O

O

O

-I-I

CO CO

O

I I +1

0

O

O CO

O

I 1 I 1 1

H

O

-I-1

,-1 r-1

,-1

r0 tU tri in

+ N O x ri

X r-4 tU I

02 H H W

A ri U

a

I N

• co)

N •

M I

in •

0 I I I I I

01 •

Is

[Cu ]

residual

mole C lx10+5

(NI ri

tn

CO o

ri

00 CO

ri

V N

O r-i

in 0,

M

N ri

in I I I I

O

N O

ro W Min

g+ 0

,C ri

X O I

oz H H W

M r-1

p

I t 0

M

o

M I

k0 N

• CO

MI!

m

Cl.

•:11

[ZnII]blank

mole

9.71x10+5

1

O O

N

t" O

N I

•:1, CO

• N

I 1 I I 1

tf1

CO •

O

in +

FA O td ri O .1.)

rr H H N

O rQi N E

I t0

• in

CO

u 1 I

HO

• H r-I I I 1 I

CO

1

?C

194

Â

O O

•

01 O •

Cr

01 O •

d'

O 0

• d'

tn ,-1 • [p

in N .

-0 01

N 01

•

t0

O 01 .

t0

O n

tD

1f1 tn

t0

x a

O CO O ri

O N

.Cr

O O

O M

er

0 N

t`

N [N

r

tf1

t0

r

tr1 t0

h

N t0

n

N 10

r

N N

E+ -.1

O to N

ri

O

O in ri ei t0 N

ri

W

f

-105-

.1 Cf) N UI ) >~ VD f; 02

? N `-' Û CI,

J

O •

In

O

tn

z

IICi

H HO •

0

O

tfl

r

4 {

{

f cx

+ I

I

01

O

O

+1

o Ol

0

ma')

O O

O O

+I +I V .-+

fs 01 0 0

co

O

O

+1 N

cil 0

co

0

O

+1 V•

W

0

I

co

O

O

+I to

co

0

co

O

O

+1

W

0

CO

0

O

+I O co

0

co

O

O

+I O co

0

co

O

O

+I tn co

0

I

Ol

O

O

+1 H 01

0

ma)

O O

0 O

+I +I

r-I 01

01 W

0 0

W

O

0

+I tn CO

0

co

0

0

+I N 0O

0

[CuII exchanged

mole

9,-1x10+5

I

V N

cot

d'

in

K1

N

tfl el

V

CO

Cr)

V

W

re)

I

d'

N

f"1

r-I

re)

f r̀!

Lo

in

el

ri

to

M

O W

frl I

M Ol

•

N

N H

M

r-i N

fr)

r-I N

el

.-1 N

M

II

[ Cu

] residual

mole R,-lx10+5

NW

HI co

tn H

W

to

r--I

O

in

ri

W

N

HI

W

r-1

Hi

N

HI

tn

W

CO

H

ri

CO

HI

h

Ln

H

ri

tn

H

N

M

H

N

.--1

to

01

r-I

N

tn

O1

H

ri

O1

H

rl

O1

H

H

fT

H

[II

Zn exchanged

mole R-l x10+5

_

Î

I

O W .

M

CO h .

(r)

W IT .

el

01 4 -1

. V'

ri h

•

V'

t0 in

h W

. .

mm

I

N

V' .

V'

01

V' .

V'

01

V' •

V'

I

N

N .

fn

01

V . M

N

tO .

M

W

h .

M

O

h .

M

tn +

X o C ri V k

A I °2

H

H N

Ç'.. H

N g

I O

O • r-I

tn

r-I

r-i

h

r-I

r--I

t0

N

H

H

en

H

I V'

W 0

h

W

0

en

N 0

m

W

O

t0

W

O

I co

tY1

HI

O

cM

Hi

h

V'

H

el

Ln

H

el

01

H

II

[Zn

] total

mole C

1x10+5

1

O W

V

frl

01

V

Ln

r-I

in

Ln

V' to

N

O

tO I

O

d'

V'

N

W

V'

tn

r-1

Ln

N

en

tn

In

M

tn

I

O

to

V'

01

W

V'

01

O

to

ri

f`1

to

01

tO

in

x

>~

x ns

ara

A

o

O

to

V'

ri

Ln

M N

in

V'

N

in

t0

N

Lo

h

N

to

ON

OH

tn to

O fY1

NN

in to

tn Ln

NN

in in

0,-4

00

V' V'

N M

00

V' Ti.

fY1 cr

00

V' V'

X

Cl,

orna

O fn V'

to tn in

OW

W en

tn W

0

V

to

0

O

to

to

V

to

h

V'

in

Ln

t0

tn

01

01

Ln

O

N

to

O

O

V'

M

O

V'

V'

O

V'

t0

O

d'

O

rI

cr

01

HI

cr

al UI

É É

O Ln tn r-1

O en

O t0

O N H

0 in in

r O M

O to

O N H

O tn in

r--1 O el

O t0

O N H

W A H

d U

z H

W

O

z O

H

E >

H

U

L>7

TABLE 4.15


-106-

-

•- ---

ZnS

(S2 )

concn

gR-1 0

0 Lri

4-$ -I I 1 ! ; a

Btection lim

[ call]exchanged

mole

9.71 x10+5

CuII below d

e

[CuIIlresidual

mole R-1x10+5

1 -1-) U tll H

W N +

r-1 N • a

Ln U

N $ 1 Â in •r1 W N › R 4.) rti 3 O ...i

N S1 r-1 N

b dJ U 0E0 O oQ N O U N >`I ß4 b4.) — t0 r-I

b b N I`1 N tn 0 •r•1 to O › c".,

4.) (Ei 3 O •r•1

N + U tll 0+ â

f•1 r1 N › r-1 N N4-) U g O • g.- N O U N

ul U N ßi 13 -I-)•--' d' ri

II

[Zn ]exchanged

mole t-1 x10+5

O ri

1 • 1-1

h r-1 01 M

d r N

' CO -1 • • • •

.--1 1--1 N N

1

L En +

xo

g ie r-1 rr A 1

r-+ cbt H H N

C: r-i N

R

M d'

I 0

M d' 1/40 N d' d' d' 111

0 0 O O

tr) +

.--1 O (13 r•1 yd K O r. •P 1

OZ

H H N

C. r-1 N

g

r

M tf1

4 .-i

o t(1 Ln tn 01 cV 1,O r's

• ri N N N Z

nII beloy

.

Z b

Â

0 01 ri M

W

co in M .-1

N O co ul

. •

h r 10

O 01 01 r u1 N o 01 01 01 01 01

. • • • . .

O 01 01 01 01 01

R.

co N r-I lf1

CO OD

CO O Ln tn l0 r` r` t` CO CO CO CO

O d' d' d' N .-1 O 01 01 O1 01 01

. . • O 01 01 01 01 01 r1

N W eg

H O r-I M l0 N

O lll tfl O O O ri M 10 N

r-I

6

rzl

—107—

C W .-. U rr 0 N C: I N 01 O ot

— U tT

o 0

Ln

lY• I

O O H r--I

00

+ 1 +1

0h.1

O O

O O r-I r-1

00

+1 +1

O Ô

r- I r-1

O rl

0

+ 1

CO

O [C

u ] ex

chan

ged

mol

e

k-l x

10+5

I

h V rr)

f•1 W

M

CO ON

M

N ri

d'

ut N

.Zr 3.5

3

ri

5+ R7 O N X N r+ Iy 1

02 H H N

O r-I U g

N r-1

u1

Ill W

r-I

0'1 N

r-1

d' r--1 •

H

O O

• H

h CO

O

II

[Zn ]

exch

ange

d m

ole 2 lx

10+5

1

0 111

• in

co al

• fr1

o O

• d'

d' O

• d'

N M

• d'

I

0 N

• 01

01 r-1

• el

W N

• M

r el

• el

N N

• frl

u1 +

-lC 0

10 )S r-1 •--' A I

ot H H W

C ri

I

O O

r-I

rl O

r-1

CO 0

r-1

t0 r-I

H

ul N

rl I

Ill r1

• r1

O d

r-I

N d

r-I

el in

H

el co

H

N +

r-1 O

X 0 r•r 4-1I

C.

H H W Cr♦ N

8

lo til

d I

01 01

C

W O

O N

u1

r 1.17

I tn

I

111 u1

d'

01 1.11

[N

0 h

d'

O 01

•

d

1.11 O • u1

x x w Â

O O

;

ul r

u i N N

u i M N

v i 1/40 N

i

W N

O O

d.,

rI O

V. v i N O

V

N O

d'

N O

d,

N O

d

W °'

O O

ui

ul d' ui

O u1

v;

01 l0

ui

rn O C1 d

O 0

d•

d' 0

d•

u1 r O 0 O H

u1 H

d•

.r1 .rI Ei e

O i.fl u1 r-1

0 fr1

0 l0

0 N H

o L.r) to H

O ri

O l0

O N r-1

CuiI ACTIVATION OF ZINC SU

LPH

IDE

TABLE 4.17


-108-

Z •-, O.-. U I C N C c2 N (Jl O tT . U

o O

ln _.

0 O

In

0 O

t11

Cd I

M H

O

+1

M

,-1

O1 H

O

+I

01

H

<`

H

+1

h H

r-I H

O

+1

H

H

M O

O

+I O M

0 ... 1

N H

O

+1

N •

H

cl, H

O

+I

V

H

trl

O

+1

Ln

h O

O

+1 O t0

O

N O

O

+I N N

0

I II

[Cu ]exchanged

mole 27

1x10+5

1

C H

d,

O N

cr

N M

cr

In M

cr

tn M

[r I

01 N

M

01 N

M

h N

M

N H

M

N H

M

1 E detectable

1 'I

[Cu ]residual

mole 9 x10+5

N ,--I

tn

CO 01

O

V' CO

O

O CO

O

h h

O

n r- O

N H

tn

M CO

H

M CO

.•-I

tn CO .

,--I

O O

• N

O O •

N

N H

tn

O C

H H

O U

an au Onn C+ ni ,C r-t U X X r' O I

02 H H O [ H g

I

O N

M

H M

N

N N

O

CO 01

fr1 I

to

d' H I

I

N r`

N

H M

N

In LO

O

LO r

d' I

M

KI. H I

I

d, t0

O

V M

O

O In 0 I

01 M

(N I

H h

t0 I

[ZnII ] blank

mole

A,-lx10+5

I O O

H

cr M

N

tn 01

V

In CO

01

!` t0

O N

I h H

(NI

H O

M

1f1 O •

tn

r` 01

O H

H H

H N

1 CO 01

O

CO t0

H

0 N

M

0 O

LO

0 h

• H H

in +

H O Id H +-I X O., 4.) I

ot H H 41

C H N O `-' E

I

O N

d,

tn t0 d'

O N

tn

h CO

In

01 H

t0 I

01 CO

cr

N M

Ln

O r`

In

H N

t0

in CO

t0 I

N t0

• H

N O •

N

O r`

• N

H t0

• I''1

01 01

• [r

x C

x td C, H

.00

O 0

tI1

CO h

tn

01 M

t0

M t0

t0

O •cl,

h

N CO

r`

O O

`d'

In O

er

to H

V

O M

CN

O h

d'

CO d'

In

N N •

CO

In r`

r`

O 01

r`

h r`

r`

to CO

r`

0\ CO

r

â

O O Ln

t0 r` tn

O 01 In

O M

t0

O t0 l0

O t0

O O d'

H O er

N O cif

O H cr

O N

Tr

O M

er

0 N

CO

In r`

CO

N r`

CO

CO t0

CO

H t0 co

CO In O

(1) 01 E C '•'1 •,.I Ei E

O Ln tn H

O M

0 t0

0 N H

0 to tn H

O M

0 t0

O N H

O Ln tn H

O M

O t0

O (N H

'Q CO U

-109-

Zn

S

(S2

) an

d E

xpts

co

ncn

9

9-1

at in

itia

l p

H

ZnS

pretr

eate

d

at

pH

2

. 0,

thon

nM

at p

H 1

2.0

than

nN

rari

nr-

04 I

co 0 0 +1 m CO

O

rn 0 0

+1 co CO

O

rn 0 0

+1 O CO

O

rn 0 0 +1 01 CO

O

O 0 0 +1 .1 01

O

I

1/40 1/40 0 0 0 0

+1 +1 m CO

LO ul •• O O

lo 0 0

+1 01 Ul

O

1/40 0 0

+1 H O

O

VD 0 0

+1 O LO

O

I

o .-.1 0 +1 01 01

O

0) o 0

+1 cr 01

O

01 o 0 +1 rn 01

O

0 r-1 0 +1 lfl 01

O

01 0 0 +1 (N 0) O

1:: a) in 01+ Z O

X

IOH r-. ot H H 4) Zr-1 U O . E

I

-4.^

m

O

el

O O • rl

O 01

m

co o d' I

rl r-1 01 d' d' d'

N N N

O l0 N

o 0 N I

rn r` 1/4 0 N l0 in I co 0 H

m rn rn cr d'

II

[Cu

]

resid

ual

mol

e C

l x1

0+

5

N r-I •Ln

1/40 l0 • H 1/40 d' H

N M

H

d' d' r1 O

H H

N r1 to

H H r N N

M lp N

N ul N

r-I tn N

N r-1 ul

M in

H

ul m

H

x N

H O ri

H

l0 01

O

,0 a) iWl i+

O .g,--i X X r,

r,° â H H a) h HQ

C

I

r` H •

O r-1 d'

er M

d'

01 d' cr

l0 co

d' I

N m O H ri d'

01 H cr

r N

d'

N l0 d' I

N H 1/40 01 d' LO O r-1 H tn ri d' d' d' d'

II

[Zn

]blan

k

mol

e C

lx10

+5

I

in d' O H • . HH

1/4 0 d' H N • • H . - -I

O m . r1

I

rl rl m d' d' cr

. • • m m m

TIL d' . m

Ln ul • m

I O 01 • O

m H • H

1/44 H . H

1/4 0 N • r-I

O m •

H

II

[Zn

tota

l

mol

e

R7

lxl0

+5

I

N N •

tn

In O m ul

in in

el r tn

1/40 H

l0 I

in lO N Ln r` r`

N l0 1

H N. r`

C` rl O

I O l0 • cr

d' H Ln

(N ul d' m d' CO in tn in

x !; O fa ar A

O m

tn O d'

in

O r-I N en O O O O O 0 • tn ln tn to in

N H

In

O 0 ln

N O Ln

m 0 Ln

r-I H Ln

l0 N

in

0 m

in

x 0.,

O

tn N

10

Orl 1/40

O

1/4 0

m

1/4 0

en

1/40 O 1 m

rn • to in Lfl

H d • ul

ON Ln

in

tn O co • un tn

O 1/4 0

H 1/4 0

m 1/4 0

H N • 1/4 0

Ill N E 'ri -r1

Ea E O L to O H O O

H

0 tr Ln H

O m

O l0

O N r-I

0 n ui H

O O O N H

ACT

IVATION

OF ZINC

SULPH

IDE

H H U

TA

BL

E 4

.18

TABLE 4.18

Cu"


Zn S

(S2) and Expts

concn

ZnS pre-

Fraai-ori at

treated at

initial ng

treated at

initial OH

Cr

01 O O

+I

CO CO

O

01 O O

+ I

r O

O

01 O O

+ 1

1/1 CO

O

01 O O

+I

tn N

O

tn O O

+1

O tn

O

I

O H O

+I

01 O

ri

O H O

+I

OD O1

O

O H O

+ I

N O1

• O

O H O

+1

u1 01

• O

O H

O

+I

O 01 •

O

I

m N H

N N O 0

O +I +I

O ma 0D m N •

• • N r--I N

I

O O •

O 0

I N O

01 M • •

O O

... I I

SO C) 01 In G+ eli O

Û X X r-' C) I

à H H C)

.7 r--I U

a

I

V' 01 •

m

CO LO

• en

O h .

N

01 N .

cV

01 LO

• H I

I

H N .

M

O h

M

01 co

M

CO 0

• V'

H r-1 .

V' I

H r--I

• V'

O1 V' .

V'

r-I N

• cN

H r-1 •

to

H H

• tn

(Cu l1] residual

mole

R._l x10+5

N r-I

tn

co V' N M M ,-I V' V' CO V'

r-I H N N M

N ri

to

ri O

H

N V'

r-I

M N

r-1

cp O

r-I

r-I 0

r-i

N r-1

tf1

r-I O

r-1

M 1/40

0

r-1 V'

0

r-i O

0 V

H O

o V

S +OTxt_ô STOW

pabuet gox

a[

)

II

I

01 V' •

V'

in N

V'

CO r-1

M

N O

• M

l0 M

• M

I

r-I tl1 •

M

O1 n

• M

O O

V'

CD N

V'

CD N •

V I

O V' HO •.

M N

CO t,O

• H I

O tn •

to I

r-I •

V' r-1 1

In +

x0 r-1 X

H r+

r A o t H H C)

6 ra a

I I

O O

tn 01

N H

O

01 H

O H

l H

O r-I

1

u1 O

r-1

r-I r-I

r-I

V' H

ri

O (11

r-I

30 N

r-I

I 0 O1

r-I r--I

0 N

(+l r-1

o 01

W H

O I

r-i N

O N

O M

In +

r-I O RI r 4-) X OH A I

..2 H H N

h r-4 a

I

an ri O r-i

tn M

O r-I

h M O ,

n r-1 N r-I

M in M r-1

I

O In

V'

O 01 V'

V' H tn

CO cr In

O tn tn

4

O 0 to r--1

V' N

In r-I

N N

tn r-I

V' H

O r-I

N H

O r-I

X = g wri

A

O N

Ol r-I

ri • In

M •

to

r-i • r

g •

r

x a O O . N

co 01 . H

co 01 . r-1

co 01 .

r-1

CO 01 .

r--1

CO 01 . H

01 r--1 . to

O M .

l0

O M .

t,O

O M .

O

N M . tO

O M . O

to H .

I

01 tn . P`

01 to . h

0 V' .

h

0 N . r

0 N . I

N In B[ H

-E o to tn

H o M

0

0

N

o to to H

0 M

0 tp

0 N rl

o tn to r♦

0 re)

0 l0

0 N r-I

or more separate experiments (this includes the blank experiments).

4.2.2(b) Summary of major results

(I) In normal light

(i) At acid pH values, R is less than 1 (about 0.85) and is in-

dependent of time. R cannot be calculated at neutral and alkaline

pH, due to Cu" hydrolysis.

(ii) At initial pH values from 4.0 to 8.2 the pH of the system

increases during activation, whilst it decreases slightly at an init-

ial pH of 10.0. The pH Of the blank increases (but less so than during

activation) at acid initial pH values and decreases at neutral and

alkaline pH values.

(iii) The concentration of ZnII produced during activation is greatest

in the pH range 5.0 to 6.5 while it is reduced at pH 4.0 and at pH val-

ues above 7.0. The ZnII contribution from the blank exceeds that

expected from solubility considerations.

(II) In the dark

(i) R equals unity in the pH range 4.0 to 6.4.

(ii) The trends for ZnII release and pH changes are the same as those

in normal light.

(III) Under UV irradiation

(i) UV irradiation causes large quantities of ZnII to be released

-112-

from the blank at initial pH values in the range 4.0 to 8.2, exceeding

those found during activation after the first 15 minutes of reaction.

ZnII release during activation is similar to that found in the dark

and in normal light.

(ii) pH increases for the blank exceed those during activation at

acid initial pH values; a slight decrease is observed at an initial

pH of 8.2.

(iii) R ranges from positive (greater than unity) to negative values

(greater and less than unity).

(IV) Pretreatment experiments

in normal light

(i) Pretreatment of ZnS at an initial pH of 5.0 has no detectable

effect compared with untreated ZnS.

(ii) Pretreatment of ZnS at an initial pH of 2.0 followed by activ-

ation at pH 5.0 leads to reduced R values. [ZnII]exchanged

is the same

as for untreated ZnS; the pH changes during activation and for the blank

show the same trends but are smaller.

(iii) Pretreatment and activation at pH 2.0 yield R and [ZnII]exchanged

values which decrease with time.

(iv) Pretreatment at pH 12.0 followed by activation at pH 5.0 yields

R values slightly less than unity. [ZnII]exchanged

is less than for un-

treated ZnS whereas the pH changes are similar.,'

in the dark

Pretreatment in the dark followed by activation at 5.0 yields R

values of unity. [ZnII] exchanged is less than for untreated ZnS whilst

the pH changes are the same.

under UV irradiation

Pretreatment under W irradiation followed by activation at an

initial pH of 5.0 yields a similar pattern of results to that already

discussed for untreated ZnS under the influence of UV irradiation.

4.2.2(c) Subsidiary effects.

The concentration of ZnII released by the ZnS in the blank experiments

([ZnII] blank) increases very slowly with time after an initial increase. It

depends on pH, ZnS concentration, ZnS sample type, ionic strength and

incident light. [ZnII]blank values range from 0.64 x 10-5 to 2.84 x 10-5 M

which is much higher than the 10-6 M that would be expected on the basis of

simple ZnS solubility considerations (Figure 2.1, Chapter 2).

The formation of small amounts of oxidation products during the preparat-

ion and subsequent storage of ZnS is very likely. Their dissolution

and the consequent concentration of ZnII produced would depend both on ZnS

concentration and sample surface area. The high ZnII values also reflect the

fact that solubility frequently increases with increasing ionic strength (18).

The larger ZnII concentrations were recorded at acid initial pH (particularly

4.0) values, where the solubility of ZnS and, probably, its oxidation products

are expected to increase. The slow leakage of trace amounts of H2S (not detected

by the AgNO3 trap) from the system may also have contributed to the Zn II values

recorded in acid pH values. The effect of incident light on ZnS solubility is

discussed in Chapter 8.

4.3. RATE EQUATIONS.

4.3.1 Acid pH (4.0 to 6.5).

The experimental data have demonstrated that CuII uptake is

(i) independent of initial CuII concentration and of ZnS concen-

tration at the same initial pH value, provided that the initial rapid

-114-

activation step is complete.

(ii) independent of pH during the initial rapid step.

(iii) independent of ionic strength and of stirring speed.

(iv) dependent on both initial pH and incident light during the

second stage of activation.

(v) dependent, to some degree, on the nature of pretreatment

prior to the second stage of activation.

In Figure 4.13, CuII uptake has been plotted as a function of log10

(time) at initial pH values of 4.0 and 5.0. The data have been taken from

Figure 4.5 and Table 4.3 (A & B). CuII uptake shows a clear linear dependence

on lo910(time) for the initial, rapid activation step (0 to about 1.5 minutes)

After this there is a change in gradient with CuII uptake gradually becoming

independent of time. Each step is examined separately in the following

discussion.

Initial rapid activation step.

The logarithmic dependence during the initial step indicates that CuII

uptake may be expressed as

r = kilog10 (t) + r1 (1)

or dr = k1

dt 2.303 t

where r is the uptake of CuII by zinc sulphide expressed in mole g-1, t is

the time in minutes, r1 is the CuII uptake at 1 minute and k1 is a constant.

k1 has a value of 4.0 ± 0.6 x 10-6 mole g-1 (5.6 ± 0.9 x 10-6 mole m-2, aver-

age of the relevant sets of data). Other CuII uptake data, determined in the

dark and in normal light, including those in which all available CuII is con-

sumed in the initial step, exhibit similar behaviour to that shown in Figure

(2)

4.13. Figure 4.15 illustrates the case where all available copper is con-

sumed in the initial step. From equation (2),

log 10 lât) = -log 10 (time) + 1og10 k 1 2.303 (3)

i.e. for the initial step a of lo Cu" P plot g10(rate of Cu uptake) as a function

of 1og10(time) should be linear with a gradient of -1. This is confirmed in

Figure 4.14, where data from Figure 4.5 and Table 4.3(A + B) have been used

to construct the curves. Although it has not been possible to obtain data

for the initial activation step of pretreated ZnS, the very close correlat-

ion between fresh and pretreated ZnS shown in Figure 4.12 for the period 5-120

minutes suggests that "mild" pretreatment will not affect the initial step.

Second, slow activation step.

Examination of Figure 4.13 reveals that after the initial activation

step has been completed (i.e. after approximately 1.5 minutes) there is a

marked change in the uptake curve. All CuII uptake data obtained at acid pH

values show similar behaviour.

For a given initial pH and for constant light conditions the CuII uptake

data adhere fairly closely to the logarithmic dependence on time in the inter-

val 1.5 to 15 minutes i.e. T = k2log10t + constant. k2 is clearly different

(since the gradient alters) from that applying to the initial step. k2 has

a value in the range 1 to 2 x 10-6 mole g 1 (1.4 to 2.8 x 10-6 mole m-2;

average of relevant data).

After 15 minutes CuII uptake either becomes independent of time or in-

creases very slowly up to 120 minutes. In this period from 15 to 120 minutes

there is, apparently, no well-defined rate equation.

It is obvious from the results presented in Section 4.2.1(b) that the

type of pretreatment, initial pH and nature of incident light will affect CuII

Cut

rU

PT

AK

E (M

OL

E G

' Z

nS

x 1

0+6

)

10

9

8

7

6

5

4

3

2

1

1.5

10.1 100 101

102

TIME (MINUTES)

FIGURE 4.13 Coauptake as a function of Iogto (time) — varying add initial pH.

[ data from Figure 4.5, Table 4.3 (A & B) ]

Initial pH: A 4.00 0 5.00

RA

TE

OF

Cu4

UP

TA

KE

(M

OL

E G

" 1 M

in' 1

x 10+

e)

100

101

10°

FIGURE 4.14

Log10 (rate of CuII uptake) as a function of log10 (time) only for initial,

rapid phase of reaction. Data apply to initial pH values of 5.00 (o) and 4.00 (A )

and are obtained from Figure 4.5 and Table 4.3 (A & B).

il

10

CuII

UP

TA

KE

(MO

LE

0'1

Z

nS

x 1

0 +6

)

0

Cull undetectable in solution after this point

10-1

100

101

TIME (MINUTES)

FIGURE 4.15

Ci uptake as a function of log10 (time) where all Cua is consumed in

one minute. (Fig. 4.2 (C))

(i.e. Cuir in solution undetectable after this time)

Initial pH = 5.00; ZnS = 3.00 gl't (S1); [CA initial = 1.50 x 10'6 M

1.5

1.0

0.5

FR

AC

TIO

N O

F M

ON

OL

AY

ER

102 100 101

TIME (MINUTES)

10

9

2

1

0

1.5

cc W }

J o z o

1.0 2 u. o z o 5 cc LL

0.5

0

FIGURE 4.16 CUR uptake as a function of Iogio (time) at initial pH 7.2

(data from Table 4.3 (C)

-120-

uptake during this second stage. Any general rate equation which attempts to

express CuII uptake must, at least, be of the following form at constant

temperature:

r = f (t, pH, incident light, pretreatment).

4.3.2. Neutral and alkaline pH

Figure 4.16 shows the logarithmic time dependence of CuII

uptake at an

initial pH of 7.2. The behaviour is similar to that found at acid pH values,

however the magnitude of CuII

uptake is greater at pH 7.2.

Despite the fact that a proportion of the available CuII

was present in

the form of hydrolysis species, sufficient Cu2+ was present for the Cu2+

electrode to respond reasonably quantitatively towards Cu2+ abstraction.

Equation (1) appears to be obeyed during the initial step (k1 = 4.0 ± 0.5 x

10-6 mole g-1 or 5.6 ± 0.8 x 10-6 mole m-2) and during the period 1.5 to 15

minutes (k2 = 0.91 ± 0.20 x 10-6 mole g-1 or 1.1 ± 0.4 x 10-6 mole m-2).

After 15 minutes CuII uptake becomes virtually independent of time.

At more alkaline pH values, hydrolysis of CuII becomes much more pronoun-

ced and the Cu2+ electrode is useful only as a qualitative probe, responding

to changes in the very small concentration of Cu2+ present. Even if CuII

concentrations are kept below the precipitation edge for Cu(OH)2solid

form-

ation (5 x 10-6 M, see Figure 4.1(a)), Cu(OH)+, Cu(OH)2aq and Cu(OH)3 are

dominant solution species up to at least pH 12. The adsorption kinetics of

species such as these even on simple oxide surfaces is poorly understood (5,7),

let alone on a complex sulphide mineral. At this stage it is not possible to

develop a quantitative rate law for the CuII-ZnS activation reaction at alkal-

ine pH values.

4.3.3. Activation energy

The determination of an activation energy can frequently provide some

insight into possible reaction mechanisms. CuII

uptake was therefore meas-

ured at 298, 308, 318 K. Experiments were performed over the time span

0 to 10 minutes, in normal light at an initial pH of 5.0 with

[Cu].II nitial = 5.12 x 10-5 M and [ZnS] = 5.00 gA,-1 (S2).

The rate of CuII uptake is rather insensitive to changes in temperature,

as is shown by the rate constants tabulated below:-

Temperature (K) kl(mole m-2 x 10+6) k2(mole m-2 x 10+6)

298 6.5 ± 0.8 2.4 ± 0.4

308 5.6 ± 0.8 2.1 ± 0.4

318 6.0 - 0.8 2.4 ± 0.4

The insensitivity of CuII uptake towards temperature changes in the

range 298 to 318 K does not allow any calculation of activation energy to be

made.

-122-

CHAPTER 5. THE ACTIVATION OF ZnS WITH CdII

5.1 INTRODUCTION

This Chapter deals with the activation of ZnS with CdII. Although cad-

mium is not used in practice as an activator it is expected that it will ex-

change with ZnS due to the greater insolubility of CdS (Table 2.1 ). In

contrast to copper, the +2 oxidation state in cadmium is far more stable, a

fact which is reflected by the respective standard reduction potentials (E°):

E°, volt (ref 67,68)

+0.16

-0.40

Cut+/Cu+

Cd2+/Cd

Furthermore for equivalent concentrations of metal ion, hydrolysis of CdII

occurs at considerably higher pH values than for CuII (compare Figure 4.1(a)

with Figure 5.1) .

5.2. RESULTS

5.2.1. CdII uptake.

5.2.1(a) Preliminary work

(I) Planning. The objectives regarding electrode behaviour and time span

(0-2 hrs) for kinetic rexperiments, described in Chapter 4 for CuII, also

applied to CdII activation of ZnS.

(II) Hydrolysis of CdII.

The equilibrium diagram for CdII hydrolysis was calculated in a sim-

ilar fashion to that already outlined for CuII and ZnII in Chapter 4. The

relevant stability constants are tabulated in Table 5.1 and the resulting

diagram is shown in Figure 5.1.

2 4 8 8 10 17

-2

0

2

4

6

8

10

12

14

16

pH

FIGURE 5.1 Equilibrium diagram for Cdias a function of pH.

Cd(OH)2 wild

Cd(OH1+

IND

Cd1OH)2aq

NNW

Table 5.1 Thermodynamic stability constants for

CdII hydrolysis at 298K (7). Symbols

are those given by Sillen (6).

CdIl,logl0 K

K so

-13.66

K1 4.16

K2 4.23

K3 0.69

K4 -0.32

Figure 5.1 is referred to below.

5.2.1(b). Results

(I) General features.

The data are presented in Figures 5.2 to 5.6 and Tables 5.2 to 5.6

inclusive. The CdII uptake data represent the mean of two or more completely

separate experiments. Figures 5.2 to 5.6 show CdII uptake for the time spans

0-5 and 0-120 minutes. CdII uptake is expressed as mole per gram of ZnS and

as the fraction of monolayer covered. Monolayer coverage is based on an avail-

able area of 2O.8R2 for each CdII species (the origin of this value of 2O.8Â2

has been discussed in Chapter 4). For the zinc sulphide samples used, mono-

layer coverage corresponds to the following CdII uptake values:-

S3 (0.74 m2 g-1) . 5.9 x 10-6 mole g-1

Si (0.58 m2 g-1) : 4.6 x 10-6 mole g-1

S5 (0.57 m2 g-1) : 4.6 x 10-6 mole g-1

This allows an estimate of the depth of penetration of the zinc sulphide

crystal lattice to be made.

FR

AC

TI O

N O

F M

ON

OL

AY

ER

0.20

a (0 - 5 min .)

2.0

Cdl

UP

TA

KE

(MO

LE

G"1

ZnS

x 1

0+6

)

0.40

1.0

2 3 TIME (MINUTES)

b (0 - 120 min.)

4

5

0.40

0

1

20

60

120 TIME (MINUTES)

FIGURE 5.2 Cda uptake as a function of time — varying initial Cda concentration

Conditions: [ZnS] = 10.00 g1-1 (S4 ) Initial pH = 7.00 (Refer Table 5.3)

(CdII] initial = 2.92 x 10.5 M (i

=4.38x10-5 M A

0

4 5

FR

AC

TIO

N O

F M

ON

OL

AY

ER

2A

b(0-120 min)

2 3 TIME (MINUTES)

0 1

0.40

0.20


-126-

a (0 - 5 min.)

CdU

UP

TA

KE

(M

OL

E G

'1 Z

nS

x 1

0`8)

0.40

0.20

w

g 0 z

LL 0 Z 0

CC LL

FIGURE 5.3 CdII uptake as a function of time - constant [Cdr) /(ZnS) ratio

Conditions: Initial pH - 5.00 (Refer Table 5.2)

° [CdII] initial -1.48 x 10-6 M; (ZnS] - 5.00 gf 1 (S4 )

6 [CPU] initial - 2.92 x 10-6 M; [ZnS] - 10.00 gl-1 (S4 )


FIGURE 5.4 CdU uptake as a function of time — varying initial pH.

Conditions: [CdU] initial = 2.92 x 10-5 M [ZnS] - 10.00 g1-1 (S4 )

(Refer Tables 5.2 (B), 5.3 (A), 5.4)

Initial pH 5.00 0 7.00 A 9.00 Ll

2.0

CdIIU

PT

AK

E (M

OL

E G'1

Zn S

x 1

0+8

)

1.0

0.40

FR

AC

TIO

N O

F M

ON

OL

AY

ER

0.20

2.0 a (0 - 5 min )

0 1 2 3 TIME (MINUTES)

4 5

0.40

FR

AC

TIO

N O

F M

ON

OL

AY

ER

0.20

b(0-120min )

—128—

2.0 0- 120 mim

0.30

CdI

IUP

TAKE

(MO

LE G

'1 Z

nS

x 1

0 +6)

1.0

0.20

0.10 FR

ACT

ION

OF

MO

NO

LA

YE

R

60

120 TIME (MINUTES)

FIGURE 5.5 CA uptake as a function of time — varying initial pH.

initial pH = 5.00 0 (ZnS) = 10.00 gi' (S3 ), [Cd4] initial ° 2.92 x 10-6 M. (Refer Table 5.5)

Initial pH = 4.00 A [ZnS] = 5.00 gl-1 (S3), 1C" initial = 4.54 x 10-6 M;

-129-

a(0- 5 min .)

Cdn

UP

TA

KE

(MO

LE

G-1

Zn

S x

10+6

)

2.0

0.40

1.0 0.20

FR

AC

TIO

N O

F M

ON

OLA

YE

R

0

1

2 3

4

5 TIME (MINUTES)

b(0.120 min ,)

2.5

2.0

1.5

1.0

0.40

Cdff UP

TAK

E (

MO

LE G

"1 Z

nS

x 1

0 +6 )

W }

0.20 u.

z

LL

0.5

60

120

TIME (MINUTES)

FIGURE 5.6

CdW uptake as a function of time — varying light conditions.

Conditions:

[ZnS] - 10.00 gl't (S4, S5) (Refer Tables 5.3 (A), 5.8)

[Cdr] initial - 2.92 x 10'6 M

initial pH - 7.00

O in normal light ta in dark 0 under UV irradiation

-130- TABLE 5.2

(A)

CdII ACTIVATION OF ZINC SULPHIDE

pH initi = 5.00 al

Time minutes

[CdII] residual

mole R-1 x 10+5

(CdII] exchanged mole k-1 x 10+5

CdII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 1.46 - - 5.00 -

0.20 1.38 0.08 0.16 0.03

0.40 1.29 0.17 0.34 0.07

0.60 1.25 0.21 0.42 0.09

0.80 1.22 0.24 0.48 0.10

1.0 1.17 0.29 0.58 5.09 0.13

2.0 1.06 0.40 0.80 0.17

3.0 1.03 0.43 0.86 0.19

4.0 1.00 0.46 0.92 0.20

5.0 1.00 0.46 0.92 5.10 0.20

15.0 0.98 0.48 0.96 5.11 0.21

30.0 0.91 0.55 1.10 5.13 0.24

60.0 0.80 0.66 1.32 5.18 0.29

120.0 0.74 0.72 1.44 5.20 0.31

[ZnS] = 5.00 gß-1 (S4).nôrmal light298K


pHiniti = 5.00 al

Time minutes

[CdII residual mole k-1 x 10+5

(CdII]exchanged mole R 1 x 10+5

CdII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 2.92 - - 5.00 -

0.20 2.74 0.18 0.18 0.04

0.40 2.56 0.36 0.36 0.08

0.60 2.44 0.48 0.48 0.10

0.80 2.32 0.60 0.60 0.13

1.0 2.24 0.68 0.68 5.29 0.15

2.0 2.10 0.82 0.82 0.18

3.0 2.08 0.84 0.84 0.18

4.0 2.04 0.88 0.88 0.19

5.0 2.02 0.90 0.90 5.31 0.19

15.0 1.85 1.07 1.07 5.36 0.23

30.0 1.76 1.16 1.16 5.39 0.25 60.0 1.70 1.22 1.22 5.48 0.26 120.0 1.56 1.36 1.36 5.54 0.29

[ZnS] = 10.00 gp,-1(S4) normal light 298K

(B)

-131- TABLE 5.3

(A) CdII ACTIVATION OF ZINC SULPHIDE

pHiniti = 7.00 al

Time [CdII] residual

mole i- x 10+5 minutes

[CdII] exchanged+5 mole 2- x 10

CdII Uptake -1 +6 mole g x 10


0.0 2.92 - - 7.00 -

0.20 2.72 0.20 0.20 0.04

0.40 2.56 0.36 0.36 0.08

0.60 2.50 0.42 0.42 0.09

0.80 2.38 0.54 0.54 0.12

1.0 2.30 0.62 0.62 6.56 0.13

2.0 2.22 0.70 0.70 0.15

3.0 2.14 0.78 0.78 0.17

4.0 2.10 0.82 0.82 0.18

5.0 2.04 0.88 0.88 6.53 0.21

15.0 1.81 1.11 1.11 6.50 0.24

30.0 1.71 1.21 1.21 6.44 0.26

60.0 1.59 1.33 1.33 6.34 0.29

120.0 1.51 1.41 1.41 6.20 0.30

[ZnS] = 10.00 0-1 (S4) normal light 298K

(B) CdII ACTIVATION OF ZINC SULPHIDE

PH = 7.00

Time minutes

[CdII] resxd0 mole Q x 10l+5

(CdII] elcha0+5 mole R x

1 10

CdII Uptake mole g-1 x 10+6


0.0 4.38 - - 7.00 -

0.20 4.22 0.16 0.16 0.03

0.40 4.02 0.36 0.36 0.08

0.60 3.90 0.48 0.48 0.10

0.80 3.88 0.50 0.50 6.74 0.11

1.0 3.82 0.56 0.56 0.12

2.0 3.68 0.70 0.70 0.15

3.0 3.60 0.78 0.78 0.17

4.0 3.58 0.80 0.80 0.17

5.0 3.56 0.82 0.82 6.66 0.18

15.0 3.42 0.96 0.96 6.55 0.21

30.0 3.32 1.06 1.06 6.50 0.23

60.0 3.10 1.28 1.28 6.48 0.28

120.0 2.96 1.42 1.42 6.30 0.31

[ZnS] = 10.00 0-1 (S4) normal light 298K

TABLE 5.4


pH initi = 9.00 al

Time minutes

[CdII] residual

mole 1-1 x 10+5

[CdII] exchanged mole R,-1 x 10+5

CdII Uptake

mole g-1 x 10+6 pH Fraction of

Monolayer

0.0 2.92 - - 9.00 -

0.20 2.32 0.60 0.60 0.13

0.40 2.17 0.75 0.75 0.16

0.60 2.09 0.83 0.83 0.18

0.80 2.02 0.90 0.90 0.19

1.0 2.00 0.92 0.92 8.08 0.20

2.0 1.90 1.02 1.02 0.22

3.0 1.85 1.07 1.07 0.23

4.0 1.83 1.09 1.09 0.24

5.0 1.82 1.10 1.10 7.91 0.24

15.0 1.72 1.20 1.20 7.79 0.26

30.0 1.69 1.23 1.23 7.72 0.27

60.0 1.62 1.30 1.30 7.64 0.28

120.0 1.56 1,36 1.36 7.39 0.29

[ZnS] = 10.00 gR-1 (Sq) normal light 298K

-133-

TABLE 5.5

(A)

Cd11 ACTIVATION OF ZINC SULPHIDE

pH initi = 5.00 al

Time [CdI1] residual mole k- x 10+5 minutes

[CdI1] exchanged mole ß x 10+5

Cd11 Uptake -1 +6 mole g x 10+6


2.92 - - 5.00 -

5.0 1.99 0.93 0.93 6.41 0.16

15.0 1.84 1.08 1.08 6.52 0.18

30.0 1.70 1.23 1.23 6.52 0.21

60.0 1.51 1.42 1.42 6.51 0.24

120.0 1.40 1.53 1.53 6.49 0.26

[ZnS] = 10.00 gk 1 (S3) normal light 298K

(B)


pH initi = 4.00 al

Time [Cd 11) residual mole R- x 10+5 minutes

[Cd11] ]exchanged mole R- x 10+5

CdII Uptake -1 10+6 mole g x 10


0.0 4.54 - - 4.00 -

5.0 4.14 0.40 0.80 4.05 0.14

15.0 4.07 0.47 0.94 4.07 0.16

30.0 3.96 0.58 1.16 4.10 0.20

60.0 3.87 0.67 1.34 4.10 0.23

120.0 3.80 0.74 1.48 4.15 0.25

[ZnS] = 5.00 0,-1 (S3) normal light 298K

-134-

TABLE 5.6

(A)


pH initi = 7.00

al

Tine minutes

[CdI]residual mole k-1 x 10+5

[CdII]exchanged mole R-1 x 10+5

CdII Uptake

mole g-1 x 10+6


0.0 2.92 - - 7.00 -

5.0 1.97 0.95 0.95 6.86 0.21

15.0 1.39 1.53 1.53 6.70 0.34

30.0 0.93 1.99 1.99 6.60 0.44

60.0 0.62 2.30 2.30 6.82 0.51

120.0 0.44 2.48 2.48 6.80 0.55

[ZnS] = 10.00 g1-1(S5)under UV irradiation 298K

(B)


pH initial 7.00

al

Time minutes

[CdII]residual 1-1 x 10+5

[CdII]exchanged mole 1-1 x 10+5

CdII Uptake

mole g-1 x 10+6


0.0 2.92 - - 7.00 -

0.20 2.71 0.21 0.21 0.05

0.40 2.51 0.41 0.41 0.09

0.60 2.46 0.46 0.46 0.10

0.80 2.40 0.52 0.52 0.11

1.0 2.35 0.57 0.57 6.32 0.13

2.0 2.22 0.70 0.70 0.15

3.0 2.17 0.75 0.75 0.16

4.0 2.09 0.83 0.83 0.18

5.0 2.08 0.84 0.84 6.78 0.18

15.0 1.90 1.02 1.02 6.72 0.22

30.0 1.80 1.12 1.12 6.62 0.25

60.0 1.67 1.25 1.25 6.58 0.27

120.0 1.49 1.43 1.43 6.42 0.31

[ZnS] = 10.00 g1-1 (Ss) in dark 298K

CdII uptake values, when expressed as mole per gram are subject to an

error of

(i) ± 10% at or below 0.8 x 10-6 mole g-1,

and (ii) ± 5% above 0.8 x 10-6 mole g-1.

Tables 5.2 to 5.6 show CdII uptake (mole g-1 and fraction of monolayer)

as a function of time together with residual and exchanged CdII

concentrations.

The relevant pH values for the CdII-ZnS system are also tabulated as an aid

in commenting on CdII hydrolysis etc. (pH changes during activation and for

the blank are dealt with in Section 5.2.2).

The curves shown in Figures 5.2 to 5.4 and 5.6 show two distinct stages

in the uptake of CdII (Figure 5.5 shows only the second stage). An initial

rapid uptake of CdII occurs, is generally completed in one minute and

results in approximately 15% monolayer coverage. This is followed by a slower

step which, over 2 hours, never exceeds a CdII uptake equivalent to 60% mono-

layer coverage and is normally about 30% of a monolayer. A pH increase accom-

panies the CdII uptake. There was no detectable colour change during activat-

ion with CdII. X-ray diffraction of dry, CdII activated zinc sulphide did not

reveal any detectable changes in either lattice spacings or line intensity

from the untreated, cubic ZnS lattice. Changes in stirring speed had no effect

on reaction rate, provided that the ZnS remained in suspension.

(II) Influence of initial CdII concentration.

(Figures 5.2, 5.3 ; Tables 5.2, 5.3)

The results shown apply to ZnS concentrations of either 5.00 or

10.00 gt-1, initial CdII concentrations of 1.46 x 10-5 or 2.92 x 10-5 M and

initial pH values of 5.00 and 7.00. Within experimental error CdII uptake is

independent of initial CdII concentration and is the same for a fixed (Cd II

)initial ratio.At a ZnS concentration of 10.00 girl and an initial CdII (ZnS)

-136-

concentration of, say,0.50 x 10-5 M, all of the available CdII was consumed

in the initial phase of the reaction (i.e. residual CdII was less than

1 x 10-7 M). The behaviour was the same as for higher CdII concentrations

II up to the point where all of the CdII was consumed. At high

[Cd ]initial [ZnS]

ratios there was only a very slight change in electrode potential so that

calculation of CdII uptake was impracticable. It should be stressed that

these systems were not at equilibrium - analyzing for cadmium over a reaction

period of several days revealed that exchange was still occurring. For ex-

ample, under conditions corresponding to Table 5.2(B), CdII was undetectable

(pH = 5.8) after activation for 2 days.

For the systems shown in Figures 5.2 to 5.3 if, at the conclusion of

the 2 hour reaction period the CdII concentration was returned to its original

value (by titrating the Cd2+ electrode back to its initial potential reading),

there was no detectable alteration in reaction rate. Similarly there was no

effect on reaction rate when the pH was returned to its original value.

Opening several of the ports in the top of the reaction vessel just

prior to injecting the ZnS, and allowing these to remain open for up to 10

minutes of reaction time, showed that the introduction of some 02 and CO2

into the previously purged system had no detectable effect on CdII uptake.

This applied to a ZnS concentration of 5.0081-1,an initial pH of 5.0 and an

initial CdII concentration of 2.90 x 10-5 M.

(III) Influence of initial pH.

(Figures 5.4, 5.5 ; Tables 5.2(B), 5.3(A), 5.4, 5.5)

acid and neutral initial pH (4.0, 5.0, 7.0).

Instability of the Cd2+ electrode at pH 4 meant that kinetic information

during the first stage of activation could not be obtained at this initial pH.

CdII uptake values were therefore determined during the second stage of activ-

ation by AAS.

CdII uptake during the second stage of activation is independent of init-

ial pH in the range 4.0 to 7.0. CdII uptake is independent of initial pH in

the range 5.0 to 7.0 during the initial, rapid activation step. Although up-

take data could not be obtained for the rapid step at pH 4.0, the close corre-

lation between CdII uptake during the second stage of activation at initial pH

values of 4.0 and 5.0 suggests that CdII uptake is also independent of initial

pH for the rapid activation step in the pH range 4.0 to 7.0. After 1 minute

of reaction time, CdII uptake was generally about 0.6 x 10-6 mol g-1 or 15%

of a monolayer, increasing to about 1.4 x 10-6 mol g-1 or 30% monolayer cover-

age at 2 hours.

alkaline initial pH (9.0)

(Figure 5.4, Table 5.4)

CdII uptake at an initial pH of 9.0 is significantly greater in both the

rapid and slow stages of activation, in comparison with the data obtained at

acid and neutral pH values, for the first 15 minutes of reaction time. After

this the CdII uptake values are the same,within experimental error (during

activation at an initial pH of 9.0, the pH decreases to 7.8 at 15 minutes and

then 7.4 at 120 minutes).

(IV) Influence of incident light.

(Figure 5.6 ; Tables 5.3(A), 5.6)

Results are reported for experiments performed in the dark, in nor-

mal light and under W irradiation. Due to the noted instability of the Cd2+

electrode under W irradiation, CdII uptake could not be determined in the

initial stage of activation but was measured by AAS analysis during the second,

slower step.

At an initial pH of 7.0 there is no detectable difference between CdII

uptake measured in the dark or in normal light (note that the surface areas of

the two ZnS samples used are almost identical). Under UV irradiation CdII up-

take is significantly higher after 5 minutes of reaction time than for experi-

ments performed in the dark and in normal light (0.95 in comparison with 0.84 x

10-6 mole g-1). CdII uptake continues to be enhanced under UV irradiation,

reaching a value of 2.48 x 10-6 mole g-1 at 120 minutes. In contrast CdII

uptake in the dark and in normal light is about 1.4 x 10-6 mole g-1 at 120

minutes.

(V) Pretreatment experiments.

No pretreatment experiments were carried out for CdII activation of

ZnS since it was believed that the salient features of CdII activation could

be demonstrated with "fresh" ZnS. Furthermore pretreatment experiments perform-

ed during CuII activation of ZnS served mainly to confirm the results obtained

with "fresh" ZnS.

5.2.2 Exchange ratio, ZnII release and pH change.

5.2.2(a) Method of data presentation.

Tables 5.7 to 5.11 contain additional information which is complementary

to that already shown in Figures 5.2 to 5.6 and Tables 5.2 to 5.6. The inform-

ation may be readily correlated thus:-

Table 5.7 relatet to Figure 5.3, 5.4 ; Table 5.2

Table 5.8(A)relates to Figure 5.2, 5.4 ; Table 5.3(A)

Table 5.9 relates to Figure 5.4 ; Table 5.4


Table 5.11 relates to Figure 5.5 ; Table 5.5(A)

The data in Tables 5.7 to 5.11 show the pH during activation, pH of

the blank, total zinc, the zinc contribution from the blank, zinc exchanged,

residual cadmium, cadmium exchanged and the "exchange ratio" (R) as a

function of time (5, 15, 30, 60 and 120 minutes).

R is defined as

II - II

R = [Cd ] initial [Cd

II] residual

II _ II [zn ] total [Zn 'blank

The data represent the mean of two or more separate experiments.

5.2.2(b) Results.

(I) experiments in normal light. Tables 5.7 to 5.9, 5.11.

R. For the pH range 5 to 8, for two separate ZnS samples. and a range

of ZnS concentrations the value of R is equal to unity, within experimental

error. R ranges from a low of 0.82 to a high of 1.22; the most common value

is 1.0. Within experimental error, R is independent of time, pH in the range

5 to 8, ZnS concentration and initial CdII concentration.

[ZnZI

] released and pH changes.

At an initial pH of 5.0, the pH increases during activation. Within ex-

perimental error, this pH increase is the same for the blank. At neutral and

alkaline pH values the pH decrease during activation matches that of the blank.

This pH correlation is in marked contrast to the Cu" activation experiments.

The concentration of ZnII released by ZnS in the blank experiments

([Zn"I]bl

ank) increases slowly with time after an initial, rapid increase.


H H

TJ U

-140-

G U ..-+

U] C I a O cf) o2 N U -

o O O ,-1

O O

Ln

I

01 O

O

+I

tV 0• O

01 O

O

+I

Cr 01 O

Ori 01 O O

O O

+1 +1

00 O1 O1

O O

0 r-I

O

+I

L11 01

O

I

O r-1

O

+1

10 01

O

O r-I

O

+1

Cr O

r-I

N H

O

+1

N N

H

H H

O

+1

W O

H

H r-1

O

+1

h O

H

II

[C dI

exchanged

mole k-1x10+5

I

O

O1

O

N

0

r-1

O H

H

N N

H

W M

'-I I

10 d'

0

CO Cr

0

Ln tn

0

t0 W

0

N !`

0

[CdII]residual

mole k-1x10+5

N 01

N

N

O

N

Ln CO

r-1

t0 h

H

010 t` Ln

i-1 H

e0 cp

14

0 0

ri

CO 01

0

1 -1 01

0

O CO

0

C1' h

O

II

[ Zn ]exchanged

mole k-1x10+5

1 CO 01

0

v' r-I

r-I

ai N

H

Ln M

H

M Cr

H

1 CO d'

O

lo C1'

O

Ln Cr

O

N

e0

O

r LO

O

[ZnII ]bla nk

mole 271x10+5

I

O H

H

0\ H

H

r-1 N

r-I

N M

ri

CO Cl

H I

1.0 1t1

0

01 Ln

0

N h

0

19 h

O

Ln CO

O

II

[Zn ]

total

mole k-1x10+5

I

CO O

N

M M

N

O

Cr

N

h 10

N

H 00

N 1

Cr 0

r-I

ln O

r--I

h r-I

,-i*

CO Cl

H

N Ln

r•{

x

x g O aH

A O

Ln

M M

Ln

01 M

Ln

N Cl.

Ln

O Ln

Ln

CO Ln

Ln

0 O

Ln

0 ri

Ln

O H

Ln

.C1' ri

Ln

r'I

to

01 r-1 .

Ln

x al

o O

Ln

r-I M

tt1

10 M

In

01 M

to

a0 Cr

to

Cr Ln

to

0 0 OH

Ln Ln

1-1 M 00 O HH HN

Ln toton Ln

0 to

E-L ►i

O H Cl l0 n1

0 r-I Cl L,O N

o rn N

normal light

-141-

Zn S

concn 0

o

O

0 0 in

a• I

01 O

0

+1

N CO

O

0 ri

0

+1

n 01

O

r-i ri 0

+I

O rl

O

rI r-I

0

+1

l0 O

rl

.-I r-I

0

+I

1.11 0

r-I

I

CO 01 ri O O rl

0 0 0

+1 +I +1

N N o 0 01 o .

O O rl

r-I ri

0

+1

u1 o H

01 O

0

+1

t` o

O

II

[Cd ]exchanged

mole 2,

l x10+5

I CO CO

O

r-I r-1

rl

r-I N

r-I

fn M

rl

r-I C

rl I

r-I CO fv1 d' d' If1

O O O

d' to O

rl r-

0

[CdII]

residual

mole 27

1x10+5

N 01

N

d O

N

r-I CO

r-1

ri h

ri

01 in rl

r-I in

r-I

tO d'

r-I

t11 co fr1 o 01 01

. r-I 0 0

N co 0

111 r-

O

II

[ ]exc hanged

mole Q 1x10+5

i

n O

r-I

d' rl

rl

O rl

r♦

t0 N

rl

d' Cl rl

I

Ol N O d' IN tf1

O O O

r-1 10 O

r r-

• 0

to +

O

A X r-I r-r A 1

r+ d H H Q)

C r-i N g

1

I

O in

O

01 Lfl

O

r-I CO

O

N 01

O

O rl

rl I

r 01 N d' d• t11

• O O O

CO in

O

el C

0

in -I-

r--4 O ro r1 1) X 0 rr 4-) 1

d H H Q) h O — 6

I r- N

rl

fr1 r-

r-I

rl 01

r-I

co rl

N

d' d'

N I

t0 rl N 01 O O

0 r-I r-I

01 r-1

• ri

O in

• rl

1 X Ip a.-1

A

O O 1

O1 t0 t0

fr1 l0

ltl

d' tf1

10•

r1 d'

tfl

fr1 N

lC)

0 O

h O O CO CO h

l0 t0 l0

r-I tO • 10

O in

l0

O

O O !`

r1 Ln

t0

0 Ln

t0

d' d'

t0

dr el

tO

O N

t0

O O r

CO t` N 10 U1 111

t0 tO tO

t` d'

t0

N M

l0

a) CO E C

F -E

O in ir1 r-1

0 1r1

0 t0

0 .v-I

0 Ir1 Ill 0 r•-I Cl

0 0

0 r-1

W A H Z a

m

U Z H N

W O

z o H E

H H U a

H H b U

normal light 298K

-142-

a -, CO C) .7'-' 0 C: In I N O `-• c•

O O

• O .-i

a4

U CT

1

O O

+1

ON

0

O

+1

rn

0 0

chm ch O O O

O O O

+1 +I +I

O N

0 0 •

b a) 0) u) q+ raO 4 r-1

K .--k _ a1 1

ot H H (1)

'Ci ri U

R

I O

r-1

N

ri

N

r-1

M M

r-1 H

r-I rd in 0+ Ti O •rl ri N DC N'-+ s.lI

r-. à H H a)

"0 H 00 ..— û

N Q1 •

N

N CO

• ri

N n

• H

O1 %D

• H

N 10

lO 1[1 • •

H 1 r♦

[ZnI1]

exchanged

mole L-1x10+5

1

ch H

' r-I

vt. N r••1

ri

d' d'

r-1 r-I

In +

O iC

H r•

rA1ot H H a)

h r-i g

I

r- fn

O

N d'

O

r• d'

O

N ri In 10

O O

[znII)

total

mole k'

1x10+5

1 %0 111

• .-Ii

t0 111 1/40 01 %D000.) O .

r1 1-1 , r-1 N

r

â 1 •1 A

Ô

O

Ô O 1

CO h

CON

h

r0 CO

. N h

04 ô •

rn

N

°x •

r 0.1 N

•

N r̀

g 2

E •d

0 tr) 111 r-i

O fr1

O O U0 N

r•1

I •1

ACTI

VATION

OF

ZINC

SU

LPHIDE

H H

U

TABLE 5.10


-143-

1

Zn S

concn

(S5)

----

i

4!JI

(TS o •rl 0 7 1 T3

w b ô•o $4 a

r-1 + O

.•., -•i

o O ô

H

GI [ x 4.) 7.1

ro G r0

••i

a

o

r-I

0

I +1 N.

01

O

CO tD

o o

+1 +1 1/49 ri

h L0

O O

0 0

r- O

0

+1 01

l0 O

ra

r-I

0 0

}-1 r-I O it q U

I

01

O

0

+1

Ln Co O

H

O

r♦

0

+1

Co 01

O

o

r•1

0

+I

rn Cr)

O

O

. •-I

0

+1

t` 01

O

0

0

+I

r 01

O

II

[ Cd]exchanged

mole R lx10+5

U1 01

1

O

M O1 Ln 01

ri r-I

O M

N

co CI'

N

I

C7' CO

O

N O

r•1

N

rl

H

in N

ri

M

C1'

•-1

[CdII]

residual

mole

R,-1x10+5

N 1

01 01

N r•1

01 cr7

M 01

ri O

N

O

O

CN

C7'

O

N Co

ONO N N

O

Ol

H

O

OO

rl

t`

O

r•1

01

CN

r7

II

[

excha nged

mole L_ 1x10+5

MON

1 O N M

M

M 0

V

I O1

O

O

'-i

ri

H

N H

CN

H

+ x 0

g "XI

H •-'

A âI

H

H W

h O — E

01

CO I

0

M M

Ln N

r-I M

d'

N

N

O

N

I

CI'

in 0

H

Lc 0

H

0

0

H

0

0

Cr

lp

0

Ln +

r-1 O RS r-1 +.1 K O r•' 4-1 1

oi H

H W

?C

N. Co

I '••I

in O1 in CI' M W

co Ln

O1 O h H

I

tr7 in

ri

Ln l0 r•1

CN r H

O r-1 a) ,--I r-1 N

Eg r-I A

o O1

O Ol •r to

M M O O

h r

M O

r

O r•i

r

O O

r

00 r 1/40

N N s W

%.0 0

co Ln

W

N CN

W

x

a OW

O O

NW

00

S 1/40 •

WW

N O

00 00

WW

O

0

t`

H

Ln

O

in

H

CN

Lo

N

M

C1'

N

N Ln

E C ••i •r1 E1 F,

o in Ln O r1 M

O l0

O N

H 0 in in

H

O M

O ■.0

0 N

H

TABLE 5. 11


-144-

a ...... t!❑1 u <r]--r

O...d U tT

o O

O r-1

C4 I

r-I ri

O

+1

Ô • ,--i

r-1 r-I

O

+I

Ô

,-i

N r-I

O

+1

CO

r- I

N r-I

O

+I

• r-I

w

CI) O

O

+1

o

II

[Cd ]ex

changed

mole

2.'

1x10+5

I

M 01

O

Co O

ri

M N

r-I

N dr

ra

M tn

H [Cd

11]

residual

mole

2.-1x101

-5

ON .

N

01

H

CO

ri

ri

tf1

r--I

er

r-I

[II

Zn exchanged

mole .'

1x10+5

I

CO OD

O

C71

01

O

cl.

O

r-I

H

N

r-I

ri

N +

x o

g x

H H

A I r--r ot

H H Q) 0ri N

1

CO

q:1

H

rn

I's

r-1

rn

CO

H

0

0 N

0

N

[ZnII 'to tal

mole .7

1x10+5

I

,0 tn

N

CO h

N

M

01

N

ri

N i,

01

tn

CO •

M

.k

x a r-1 A

ô .

tn tNn

•

1/40

.

1/40

.

1/40

ho .

1/4.0

.

1/40

x p, ô . u1

.

1/4.0

N tn .

1/4.0

In .

1/40

t n .

1/4o

.

1/4.0

Q W

G

H

O u1 1/4r)

H

0

M

0

l0

0

N

r-I

[ZnII

]blank depends on pH, ZnS concentration, ZnS sample type etc. and has

been discussed in Chapter 4, Section 4.2.2(c)).

The release of ZnII in systems where CuII is present ([ZnIII total) is

similarly characterized by a rapid, initial release followed by a slow in-

crease with time. The [Znllltotal

is greatest at an initial pH of 5.0 and

decreases with increasing pH,other conditions being constant. The [Zn1I]total

is dependent on both the ZnS concentration and initial CdII concentration.

The [Zn1l]exchanged is independent of pH and is dependent only on ZnS concen-

tration and initial CdII concentration.

(II) experiments in the dark. (Table 5.10B)

R. R is equal to unity, within experimental error, and is independ-

ent of time.

[ZnII] released and pH changes

The effects are similar to those obtained in normal light.

(III) experiments under UV irradiation.(Table 5.10A)

Experiments were performed at an initial pH of 5.0. Release of zinc

from ZnS was enhanced under W irradiation during both the blank experiments

and activation compared to experiments performed in the dark and in normal

light. For the period 0 to 60 minutes the concentration of zinc released from

the blank is less than that found during activation. At 120 minutes the con-

centration of zinc released by the blank is slightly greater than during activat-

ion.

For blank experiments and during activation there is an enhanced pH in-

crease observed, compared to experiments performed in the dark and in normal

light.

- Calculation of R was only performed when MIII)total (ZnII)blank was

positive. R decreases from unity to 0.7 at 60 minutes.

5.3 RATE EQUATIONS

Experimental evidence has shown that CdII uptake is

(i) independent of initial CdII concentration in the pH range 4 to

8 provided that the initial, rapid activation step is complete.

(ii) independent of pH in the range 4 to 8.

(iii) independent of stirring speed.

(iv) enhanced under UV irradiation during the second stage of

activation.

In Figures 5.7 to 5.9, CdII uptake has been plotted as a function of log10

(time) at initial pH values of 5.0, 7.0 and 9.0. The data have been taken from

Figure 5.4 and Tables 5.2(B), 5.3(A) and 5.4.

CdII uptake is linearly dependent on 1og10(time) for the initial, rapid

activation step (0 to 1 minute). After this there is a change in gradient

with CdII uptake again following a logarithmic dependence on time up to 120

minutes.

5.3.1. Initial, rapid activation step.

The logarithmic dependence during the initial step indicates that CdII


I' = k11og10(t) + r1 (1)

dl' k1 dt 2.303t

where r is the uptake of CdII by zinc sulphide expressed in mole g-1, t is

the time in minutes, r1 is the CdII uptake at 1 minute and k1 is a constant.

or (2)


0.40

1.8

0.2

FIGURE 5.7 Cdd uptake as a function of logo (time)

Initial pH = 5.00 (data from Table 5.2 (B)

CdU

UP

TA

KE

(M

OL

E 0 Z

nS

x 1

0+

6)

0.6

0.2

1.0

1.4

1.8

2 0.20 ô

0.40

cc w } g o z o

z o

cc LL

—148—

10.1

100 101

102 TIME (MINUTES)

FIGURE 5.8 Cdauptake as a function of log10 (time)

Initial pH = 7.00 (data from Table 5.3 (A)

0.40 1.8

0.2

—149—

10"1 100 101

102

TIME (MINUTES)

FIGURE 5.9 CdIE uptake as a function of log10 (time)

Initial pH = 9.00 (data from Table 5.4)

5.00

7.00

9.00

1.1 ± 0.2

1.0 ± 0.2

0.84± 0.20

0.65 ± 0.10

0.60 ± 0.10

0.49 ± 0.10

k1 mole g-1 x 10+6 kl mole m-2 x 10+6 Initial pH

-150-

k1 has the following values, depending on initial pH:

Within experimental error, k1 is constant at initial pH values of 5 and 7,

decreasing slightly at 9. Other CdII uptake data, determined in the dark and

in normal light, including those in which all the available CdII is consumed

in the initial step, show similar behaviour to that shown in Figures 5.7 to

5.9 (the rate constants are an average of the relevant sets of data).

The uptake of CdII during the initial rapid step shows a similar pattern

of behaviour to CuII uptake at acid pH values.

5.3.2. Second, slow activation step.

Figures 5.7 to 5.9 show that CdII uptake during the period 1.0 to 120

minutes is clearly dependent on log10 time. i.e.

r = k2 logl0 (t) + r1 (3)

k2 has the following values, depending on initial pH (average of various sets

of data):-

Initial pH k2 mole g-1 x 10+6 k2 mole m-2 x 10+6

5.0 0.34 ± 0.06 0.59 ± 0.10

7.0 0.40 ± 0.06 0.69 ± 0.10

9.0 0.20 ± 0.05 0.34 ± 0.09

k2 is about the same at initial pH values of 5 and 7, within experimental error.

k2 at initial pH 9 is approximately half that recorded at the lower initial

pH values.

The uptake of CdII during the period 1 to 120 minutes is similar in

behaviour to CuII uptake at acid pH values during the time span 1.5 to 15

minutes. In the case of CdII uptake there is no plateau region during the

second stage of activation (compare CuII uptake at an initial pH of 4.0).

5.3.3. Activation energy.

CdII uptake was measured at 298, 308 and 318 K. Experiments were per-

formed over the time span 0 to 10 minutes, in normal light, at an initial pH

of 5.0 with [CdII]initial = 2.92 x ].0 5 M and [ZnS] = 10.00 g1-1 (S4).

The rate of CdII uptake is fairly insensitive to changes in temperature,

as is shown by the rate constants tabulated below:-

Temperature (K) kJ (mole m2 x 10+6 k2 (mole m 2 x 10+6

298 1.1 ± 0.2 0.59 ± 0.10

308 1.3 ± 0.2 0.57 ± 0.10

318 1.1 ± 0.2 0.67 ± 0.10

The noted insensitivity of CdII uptake towards temperature changes in

the range 298 to 318 K does not permit any calculation of activation energy

to be made.

CHAPTER 6. THE ACTIVATION OF ZnS WITH PbII

6.1 INTRODUCTION

This Cha_=er deals with the activation of ZnS with Pb". PbII is expected

to exchange with ZnS due to the greater insolubility of PbS (Table 2.1 ). The

+2 oxidation state in lead is intermediate in stability compared with copper

and cadmium, a fact which is shown by the standard reduction potentials (E0 ):

Eo, volt (ref 67, 68)

Cut+/Cu+ + 0.16

Pb2+/Pb - 0.13

Cd 2+/Cd - 0.40

Furthermore hydrolysis of PbII occurs at pH values intermediate between those

for CuII and CdII (compare Figures 4.1(a), 5.1 and 6.1).

6.2 RESULTS

6.2.1. PbII uptake.

6.2.1(a) Preliminary work

(I) Planning. The objectives regarding electrode behaviour and time span

(0-2 hrs) for kinetic experiments, described in detail in Chapter 4 for CuII,

also applied to PbII activation of ZnS.

(II) Hydrolysis of PbII.

The equilibrium diagram for PbII hydrolysis was calculated in a sim-

ilar fashion to that already outlined for CuII and ZnII in Chapter 4. The

relevant stability constants are tabulated in Table 6.1 and the resulting dia-

gram is shown in Figure 6.1.

—153—

-2

Pb1OH12 Solid

— Pb1OH12aq

2

4

8

8

10

12

pH

FIGURE 6.1 Equilibrium diagram for Pbir as a function of pH.

-154-

Table 6.1. Thermodynamic stability constants for

PbII hydrolysis at 298K(7).Symbols are

those given by Sillen (6)

PbI Ilog 10 K

K so

-17.0

K1 6.3

K2 6.3

K3 2.8

Kq 0.9

6.2.1(b) Results

(I) General features

The data are presented in Figures 6.2 to 6.6 and Tables 6.2 to

6.5 inclusive. The PbII uptake data represent the mean of two or more com-

pletely separate experiments. Figures 6.2 to 6.6 show PbII uptake for the time

spans 0-5 and 0-120 minutes. PbII uptake is expressed as mole per gram of ZnS

and as the fraction of monolayer covered. Monolayer coverage is based on an

available area of 20.8R2 for each PbII species (refer Chapter 4). For the zinc

sulphide sample used, monolayer coverage corresponds to the following PbII up-

take value:

S5 (0.57 m2 g-1) 4.6 x 10-6 mole g-1

This allows an estimate of the depth of penetration of the zinc sulphide crystal

lattice to be made.

PbII uptake values, when expressed as mole per gram are subject to an

error or

(i) ± 10% at or below 1.0 x 10-6 mole g-1.

and (ii) ± 5% above 1.0 x 10-6 mole 5-1.

Tables 6.2 to 6.5 show PbII uptake (mole g-1 and fraction of monolayer) as

PbU

UP

TA

KE

(M

OLE

G" 1

Zn

S x

10

+6)

0.40

0.40 2.0

P & U

PT

AK

E (

MO

LE

G'1

ZnS

x 1

0+6

)

1.0

2a

B

5 4 1 2 3 TIME (MINUTES) b(0-120 min.)

-155-

a (0 - 5 min )

0 60

120

TIME (MINUTES)

FIGURE 6.2 PO uptake as a function of time - varying initial PbII concentration

Conditions: (ZnS1 _ 5.00 gl-t (S5 ) Initial pH 5.00 (Reler fahles 6.2 (A) & (C)

[PbIII initial - 4.79 x 10-5 M

0.96x1(15M A

5 4 1

2 3 TIME (MINUTES)

b (0 - 120 min .)

-156-

a(0-5 min )

2.0 - 0.40

- 0.20

Mir

UP

TA

KE

(M

OL

E G

'1 Z

nS

x 1

0+e

)

2.0

1.0

- 0.40

0.20

FR

AC

TIO

N O

F M

ON

OL

AY

ER


FIGURE 63 PIF uptake as a function of time — varying ZnS concentration.

Conditions [PbI] initial' 4.79 x 10.6

M Initial pH = 5.00 (Refer Tables 6.2 (A) & (B) )

[ZnS] = 5.00 gI'1 0 (S5)

= 10.00 91-1 (S5)

0.40

—157—

a(0-5 min )

2.0

0.40

POI U

PT

AK

E (

MO

LE

G-1

Zn

S x

10+

6 )

1.0 0.20

FR

AC

TI O

N O

F M

ON

OL

AY

ER

1

2 3

4

5 TIME (MINUTES) b(0-120 min )

2.0

1.0

PbII U

PT

AK

E (

MO

LE

G' 1

Zn

S x

10+6

)

0


FIGURE 6.4

Foist uptake as a function of time — varying ZnS concentration.

Conditions: [Pb]4 initial = 4.79 x 10-6 M Initial pH -6.6 (Refer Tables 6.4 (A) & (B) I

[ZnS] = 5.00 gI-t 0 (S5)

10.00 g1-1 A

(S5)

1 2 3 4 5 TIME (MINUTES)

b (0.120 min)


FIGURE 6.5 Pb1r uptake as a function of time — varying initial pH Conditions: (ZnS] = 5.00 gl't (S5) (1)bal initial = 4.79 x 10'6 M

(Refer Tables 6.2 (A), 6.3 (A & B), 8.4 (A) 1

Initial pH = 4.00 0 = 5.00 0 = 6.60 0 = 7.77 +

PbIIU

PT

AK

E (

MO

LE

G"1

Zn

S x

10

+6

) 2.0

1.0 0.20

0.40

2.0

Pbr

UP

TA

KE

(MO

LE

G'1

Zn

S x

10+

6)

1.0

0.40

0.20

FR

AC

TI O

N O

F M

ON

OL

AY

ER

0.20

-159-

a (0 - 5 min )

0.40

-1 2 3 TIME (MINUTES) b(0-120 min )

4

5

Pbl

i U

PT

AK

E (

MO

LE

G'1

Zn

S x

10

+6I

10

2.0

0.50

60

120 TIME (MINUTES)

POI uptake as a function of time - varying light conditions.

Conditions: [ZnS] = 5.00 91'1 (S5) [P b 1 initial ° 4.79 x 10'6 M

Initial pH = 5.00 (Refer Tables 6.2 (A), 6.5 (A) & (B) 1

o in normal light A in dark O under UV irradiation

FIGURE 6.6

-160-

TABLE 6.2

(A) PbII ACTIVATION OF ZINC SULPHIDE

pHiniti - 5.00 al

Time minutes

[PbII] residual

mole JC-1 x 10+5

[PbII] exchanged mole R-1 x 10+5

PbII Uptake

mole g-1 x 10+6

pH Fraction of

monolayer

0.0 4.79 - - 5.00 -

0.20 4.58 0.21 0.42 0.09

0.40 4.48 0.31 0.62 0.13

0.60 4.42 0.37 0.72 0.16

0.80 4.31 0.48 0.90 0.21

1.0 4.24 0.55 1.00 5.18 0.24

2.0 4.10 0.69 1.29 0.30

3.0 4.07 0.72 1.44 0.31

4.0 4.05 0.74 1.48 0.32

5.0 4.03 0.76 1.52 5.24 0.33

15.0 4.01 0.78 1.56 5.33 0.34

30.0 3.94 0.85 1.70 5.41 0.37

60.0 3.97 0.82 1.64 5.56 0.36

120.0 3.92 0.87 1.74 5.74 0.38

[ZnS] = 5.00 g0,-1 (S5) Normal light 298K

(B) PbII ACTIVATION OF ZINC SULPHIDE

pHiniti = 5.00 al

Time minutes

[PbII] residual

mole R-1 x 10+5

[PbII] xchanged mole 9,-1 x 10+5

PbII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 4.79 - - 5.00 -

0.20 4.35 0.44 0.40 0.10

0.40 4.20 0.59 0.59 0.13

0.60 4.08 0.71 0.71 0.15

0.80 4.00 0.79 0.79 0.17

1.0 3.90 0.89 0.89 5.30 0.19

2.0 3.86 0.93 0.93 0.20

3.0 3.83 0.96 0.96 0.21

4.0 3.81 0.98 0.98 0.21

5.0 3.81 0.98 0.98 5.49 0.21 15.0 3.64 1.15 1.15 5.62 0.25 30.0 3.52 1.27 1.27 5.75 0.28 60.0 3.48 1.31 1.31 5.92 0.28 120.0 1,45 , 1.34 1.34 6.12 0.29

[ZnS] = 10.00 gR-1 (S5) Normal light 298K

TABLE 6.2

(C)

PbII ACTIVATION OF ZINC SULPHIDE

pH initi = 5.00 al

Time minutes

[PbII] residual mole R-1 x 10+5

[PbII] exchanged mole L-1 x 10+5

Pb Uptake _

mole g 1 x 10+6

pH Fraction of

Monolayer,

0.0 0.96 - - 5.00 -

0.20 0.82 0.14 0.28 0.06

0.40 0.71 0.25 0.50 0.11

0.60 0.67 0.29 0.58 0.13

0.80 0.64 0.32 0.64 0.14

1.0 0.63 0.33 0.66 5.16 0.14

2.0 0.62 0.34 0.68 0.15

3.0 0.63 0.33 0.66 0.14

4.0 0.63 0.33 0.66 0.14

5.0 0.63 0.33 0.66 5.21 0.14

15.0 0.58 0.38 0.76 5.24 0.17

30.0 0.46 0.50 1.00 5.28 0.22

60.0 0.46 0.50 1.00 5.30 0.22

120.0 0.43 0.53 1.06 5.33 0.23

[ZnS] = 5.00 0-1 (S5) Normal light 298K

-162-

TABLE 6.3

(A)


pH initial = 4.00

Time minutes

[PbII] residual

mole R-1 x 10+5


PbII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 4.79 - - 4.00 -

0.20 4.58 0.21 0.42 0.09

0.40 4.44 0.35 0.70 0.15

0.60 4.33 0.45 0.90 0.17

0.80 4.27 0.52 1.04 0.23

1.0 4.21 0.58 1.16 4.02 0.25

2.0 4.09 0.70 1.40 0.30

3.0 4.05 0.74 1.48 0.32

4.0 4.00 0.79 1.58 0.34

5.0 4.00 0.79 1.58 4.02 0.34

15.0 4.02 0.77 1.54 4.02 0.36

30.0 4.03 0.76 1.52 4.04 0.33

60.0 4.03 0.76 1.52 4.05 0.37

! 120.0 4.03 0.76 1.52 4.07 0.33

[ZnS] = 5.00 gR-1 (85) normal light

298K


pH initi = 7.77 al

Time minutes

[Pb"] residual mole 2-1 x 10+5

[PbII] exchang d mole R-1 x 10 5

PbII Uptake

mole g-1 x 10+6 pH Fraction

of Monolayer

0.0 4.79 - - 7.77 -

0.20 4.49 0.30 0.60 0.13

0.40 4.37 0.42 0.84 0.18

0.60 4.30 0.49 0.98 0.21

0.80 4.23 0.56 1.12 0.24

1.0 4.10 0.69 1.20 7.70 0.30

2.0 4.00 0.79 1.58 0.34

3.0 3.93 0.86 1.72 0.37

4.0 3.81 0.98 1.96 0.43

5.0 3.83 0.96 1.92 7.72 0.42

15.0 3.81 0.98 1.96 7.74 0.43

30.0 3.70 1.09 2.18 7.74 0.47

60.0 3.67 1.12 2.24 7.70 0.49

1120.0 3.63 1.16 2.32 7.65 0.50

[ZnS] = 5.00 gR 1 (S5) normal light

298K

-163-

TABLE 6.4

(A)


pHinitial = 6.60

Time minutes

[PbII] residual

mole k-1 x 10+5


PbII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 4.79 - - 6.60 -

0.20 4.48 0.31 0.62 0.13

0.40 4.31 0.48 0.96 0.21

0.60 4.22 0.57 1.14 0.25

0.80 4.16 0.63 1.26 0.27

1.0 4.12 0.67 1.34 6.65 0.29

2.0 4.07 0.72 1.44 0.31

3.0 3.99 0.80 1.60 0.35

4.0 3.96 0.88 1.66 0.36

5.0 3.89 0.90 1.80 6.75 0.41

15.0 3.83 0.96 1.92 6.79 0.42

30.0 3.78 1.01 2.02 6.79 0.44

60.0 3.70 1.09 2.18 6.75 0.47

120.0 3.71 1.08 2.16 6.68 0.47

[ZnS] = 5.00 g271 (S5) normal light 298K


pHiniti = 6.56

al

Time minutes

[PbII] residual

mole i-1 x 10+5

[PbII] exchanged mole Q-1 x 10+5

PbII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 4.79 - - 6.56

0.20 4.15 0.64 0.64 0.14

0.40 3.79 1.00 1.00 0.22

0.60 3.64 1.15 1.15 0.25

0.80 3.58 1.21 1.21 0.26

1.0 3.49 1.30 1.30 6.62 0.27

2.0 3.33 1.46 1.46 0.32

3.0 3.27 1.52 1.52 0.33

4.0 3.26 1.53 1.53 0.33

5.0 3.20 1.59 1.59 6.72 0.35

15.0 3.10 1.69 1.69 6.79 0.37

30.0 3.08 1.71 1.71 6.80 0.37

60.0 3.06 1.73 1.73 6.81 0.38

120.0 3.04 1.75 175 6.81 0.38

[ZnS] = 10.00 g9,-1 (S5) normal light 298K

-164-

TABLE 6.5

(A)


pH initial 5.00

Time minutes

(PbII ]residual

*yole i-1 x 10+5 [Pb22]exchanged mole Q-1 x 10+5

PbII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 4.79 - - 5.00 -

0.20 4.60 0.19 0.38 0.09

0.40 4.53 0.26 0.52 0.11

0.60 4.49 0.30 0.60 0.13

0.80 4.44 0.35 0.70 0.15

1.0 4.39 0.40 0.80 5.18 0.17

2.0 4.30 0.49 0.98 0.21

3.0 4.30 0.49 0.98 0.21

4.0 4.30 0.49 0.98 0.21

5.0 4.27 0.52 1.04 5.20 0.23

15.0 4.21 0.58 1.16 5.22 0.25

30.0 4.15 0.64 1.28 5.25 0.28

60.0 4.03 0.76 1.52 5.30 0.33

120.0 4.03 0.76 1.52 5.30 0.38

[ZnS] = 5.00 0,-1 (S5) in dark 298K

(B)


pH initi 5.00 al

Time minutes

[PbII] residual Q-1 x 10+5

[Pb"] exchanged mole 9,-1 x 10+5

PbII Uptake

mole g-1 x 10+6

pH Fraction of

Monolayer

0.0 4.79 - - 5.00 -

5.0 4.08 0.71 1.40 5.56 0.30

15.0 3.43 1.36 2.72 5.86 0.59

30.0 2.91 1.88 3.76 6.60 0.82

60.0 1.91 2.88 5.76 7.19 1.25

120.0 0.89 3.90 7.80 7.49 1.70

[ZnS] = 5.00 gi-1 (S5) under UV irradiation 298K

-165-

a function of time together with residual and exchanged PbII concentrations.

The relevant pH values for PbII - ZnS system are also tabulated as an aid in

commenting on PbII hydrolysis etc. (pH changes during activation and for the

blank are dealt with in section 6.2.2).

The curves shown in Figures 6.2 to 6.6 show two distinct stages in the up-

take of PbII. An initial rapid uptake of PbII occurs, is generally completed

within one minute and, depending on the experimental conditions, results in

from approximately 15 to 30% monolayer coverage. This is followed by a slower

step which, over 2 hours, never exceeds a PbII uptake equivalent to 17% mono-

layer coverage (under UV irradiation) and normally ranges from 23 to 50% of a

monolayer. A pH increase accompanies the PbII uptake.

During activation with PbII the initially white ZnS turns pale brown.

Under UV irradiation the activated ZnS turns dark grey, the colour becoming more

intense with time. X-ray diffraction of dry, PbII activated zinc sulphide did

not reveal any detectable changes in either lattice spacings or line intensity

from the untreated, cubic ZnS lattice. Changes in stirring speed had no effect

on reaction rate, provided that the ZnS remained in suspension.

(II) Influence of initial PbII concentration.

(Figure 6.2 ; Table 6.2(A) and (C)).

The results shown apply to a ZnS concentration of 5.00 g2C-1, initial

PbII concentrations of 0.96 x 10-5 or 4.79 x 10-5 M and an initial pH value

of 5.00. PbII uptake is dependent on initial PbII concentration, particularly

during the second stage of activation. The difference in PbII uptake values

recorded during the initial stage (up to 1 minute) may, in part, be due to a

slow electrode response following activation at an initial PbII concentration

of 0.96 x 10-5 M (refer Chapter 2).

At a ZnS concentration of 10.00 g2.-1 and an initial PbII concentration of

0.50 x10-5 M, all of the available PbII was consumed in the initial phase

of the reaction (i.e. residual PbII was less than 1.0 x 10-7 M). The

behaviour was similar to higher PbII concentrations, up to the point where II

all available PbII was consumed. At high [Pb ]initial ratios there was [ZnS]

only a very small change in Pb2+ electrode potential so that calculation of

PbII uptake was impracticable. These systems were not at equilibrium -

analyzing for PbII over a period of several days revealed that exchange was

still occurring. For example, under conditions corresponding to Table 6.2(A),

PbII was undetectable (pH = 6.2) after activation for two days.

Opening several of the ports in the top of the reaction vessel just

prior to injecting the ZnS, and allowing these to remain open for up to 10

minutes of reaction time, showed that the introduction of some 02 and CO2

into the previously purged system had no detectable effect on PbII uptake.

This applied to a ZnS concentration of 5.00 gR 1 , an initial pH of 5.0 and an

initial Pb T

concentration of 4.79 x 10-5 M.

(III) Influence of ZnS concentration.

(Figures 6.3, 6.4 and Tables 6.2(A) & (B), 6.4(A) & (B)).

Within experimental error, PbII uptake is independent of ZnS

concentrations during the initial, rapid stage of activation. PbII uptake

is clearly dependent on ZnS concentration during the second stage - for

a given initial PbII concentration, PbII uptake increases as the ZnS con-

centration is reduced.

(IV) Influence of initial pH.

(Figure 6.5 ; Tables 6.2(A), 6.3(A) & (B), 6.4(A)).

acid initial pH (4.0, 5.0)

PbII

uptake during the initial stage of activation at initial pH values

of 4.0 and 5.0 is independent of pH, within experimental error. After about

15 minutes of reaction time PbII

uptake is dependent on pH during the second

stage of reaction. In particular the PbII uptake curve at an initial pH of

4.0 reaches a plateau (Figure 6.5(b)) - PbII uptake stays constant at

1.5 x 10-6 mole-1 II g whereas Pb uptake at initial pH 5.0 continues to in-

crease.

near neutral and alkaline initial pH (6.6, 7.8)

PbII uptake during activation at initial pH values of 6.6 and 7.8 is in-

dependent of initial pH during the rapid stage of activation within experimental

error, and is enhanced at initial pH 7.8 during the second stage. In both

cases PbII uptake is enhanced compared to activation at acid initial pH values.

(V) Influence of incident light

(Figure 6.6 ; Tables 6.2(A), 6.5(A & B).

Results are reported for experiments performed in the dark, in normal

light and under UV irradiation. Due to the noted instability of the Pb2+

electrode under UV irradiation, PbII uptake could not be determined in the

initial stage of activation but was measured by AAS analysis during the second,

slower step.

At an initial pH of 5.0 there is no detectable difference between PbII

uptake measured in the dark and in normal light for the first 30 seconds of

reaction time. After this time PbII uptake is enhanced in normal light.

Under UV irradiation PbII uptake is dramatically increased after 5 minutes

of reaction time compared with experiments performed in normal light and in the

dark (e.g. at 120 minutes there is a fourfold increase under UV irradiation).

(VI) Pretreatment experiments.

No pretreatment experiments were carried out for PbII activation

of ZnS since it was believed that the salient features of PbII activation

could be adequately demonstrated with "fresh" ZnS. Furthermore pretreatment

experiments performed during CuII activation of ZnS served mainly to confirm

the results obtained with "fresh" ZnS.

6.2.2 Exchange ratio, ZnII release and pH change.

6.2.2(a) Method of data presentation

Tables 6.6 to 6.9 contain additional information which is complementary

to that already shown in Figures 6.2 to 6.6 and Tables 6.2 to 6.5. The in-

formation may be readily correlated thus:

Table 6.6 relates to Figures 6.2, 6.3, 6.5, 6.6 ; Table 6.2

Table 6.7 relates to Figures 6.4, 6.5 ; Table 6.4



The data in Tables 6.6 to 6.9 show the pH during activation, pH of the

blank, total zinc, the zinc contribution from the blank, zinc exchanged, res-

idual lead, lead exchanged and the "exchange ratio" (R) as a function of time

(5, 15, 30, 60 and 120 minutes).

R is defined as

R = - II [Pb"].nitial [Pb ]residual

[ZnII] - II total (Zn ) blank

The data represent the mean of two or more separate experiments.

CO rn N

normal lig ht

Zn S

concn

gR-1

(S5)

0 o

tri

o 0 ô H

O 0 Ln

a I

d' H .

O +1

01 d'

r-I

M r-1 .

O

+I

O M

H

r1 H .

O

+I

lS'i o

H

01 0 .

O

+ 1

r co

O

CO 0 .

O +I

l0 r

o

I

Ol 0 .

O +I

N rn

O

CO O .

O

+I (`9 co

O

CO O .

O

+1

in r

o

r O .

O

+I

d' r

0

(--- O .

O +I

r to

0

I

CO O .

O

+I

r r

0

' CO O .

O

+I

[r co

0

al O .

O

+1

H ch

o

01 O .

O

+I

t0 co

O

Ol O .

O

+I

H 01

O

't:1 a) Ou-) c+ ft A H O X X ,-+ (U 1

02 i-H r. CJ

A H W O

6

I tO r

O

co r

o

in CO

O

N co

O

r CO

O I

CO ch

O

to H

H

r N

H

H fM

H

d' M

H I

rn (y1

O

co (Y1

O

O Ln

O

O tn

O

M in

O

II

[Pb ]residual

mole 2

2 x10+5

O1 r

d'

M O

d'

H O

d'

d' 01

M

r 01

M

N 01

M

O1 r

d'

H CO M

d' 1/40

f+1

N tn

M

CO cr

M

Ln VI M

to Ol

O

(n W

O

CO v1

o

t0 'di

o

to d'

O

(M d'

o

II

[Zn ]exchanged

mole

2.-1x10+5

1

in to .

O

O to .

O

O CO .

O

d' O1 .

O

d' H .

H I

to O .

H

CO M .

H

O r

• H

co r .

H

O O •

N I

m d' .

O

tri cr .

O

in in .

O

co Lf1 .

O

CO Ln .

O

tn +

x O X

r-I r ,

rA o I 2 H H O

r-t N O @

r r d' to

0 O

O to

O

O to

O

tO tO

O 1

t0 r

O

OD r

O

0 co

O

W co

O

to 01

O I

s cr

O

r to

O

co to

O

O to

O

t0 to

O

S+OTxl_?f a

Tau

Tpqoq. [

uZ)

II

I

O O

H

r H

H

Cl) M

H

d «1

H

O co

H I

N co

H

tfl H

N

O to

N

d' to

N

Lf1 01

N I

O 01

O

(N O

H

(•1 r-(

H

CO H

H

d' N

H

X 2

A

O O

to

O N

Ln

r-I N

to C, r-I

N N

Ln

M N

to

(`') N

tn

O O

tr1

r1 d'

to

t0 d'

to

co d'

N

01 d'

tri

0 tf1

• Ln

0 O

in

0 N

in

H N

in

(N N

to

M N

tn

tY1 N

in

C

O O

to

d' N

in

(`1 f" 1

to

r-I d' in

tO Ln

tn

d' r

tn

O O

Lf1

01 d'

0

N to •

to

to r

in

N Ol

in

N r-1

to

O O to

H N

Ln

[r N

Ln

CO N

to

O M

to

('^ M

to

N N E C •r•I •ri E-' E

O Ln Lf1 r-I

O (•l

O t0

O N r-I

O Lo to H

O M

O tO

O N ,-1

O tn tn rl

O (`'1

O t0

O N rl

KC

Ca

U

Pb11 ACTIVATION OF ZINC SULPHIDE

-170-

z

[ O d U1 •

) ; H

O H -••. O O Ul I to O O N O O'r `- in o

1

1

: O O O O O O O O O 0

+ a I +I +I +I +1 +I I +I +I +I +I +I

H M M ul in M tn M 0 CT tfl

1 P M M N O M M N O O

r-I e-1 H H H H H H H H ■

1 M Vr f`'1 O V M N H H H H H H H H H H r••I H

[Mon

]

mole

R-1x 1

0+

5

I I O l0 H 0\ CO 01 01 H M ut 0) 01 O O O t11 t0 I I I O O H H H H H H H H

H Mil Z+ ro0 •r•I H

i•1 i H H W e H

G

W X 01 01 M CO 0 H 01 O 0 CO t0 Vr N r+ I CO CO I I I C N H O O O ot Vr M M M M M Vr M M M M fol

N <T C+ 4 O

H H W C: H N Q u E

U X N N u1 C` In CO I fY1 CO u1

N Î I to r r CO 0 • I H N Vr u1 tfl

e•-r 02 0 0 0 O H H • H H H r-1

+

x0 H

H H W 6

H O u E

Â I I Vr vs Vr Vr Vr I l0 t0 0 r CO X re) M fyl M tn r-•I r dr t0 M

à O O O o O O O O O O

tA +

H O b H

H H W {ÿ

N O — e

+ I I I d X In to Co 0 0 01 H t VI CO O-+ 01 H H M u1 r m o r1 er

ot O H H H H H H N N N

x O tn W O O DD O CO H N O tn

â O OD I W V' M O t0 t0 if) V N

• Â r tfl to 0

xL1.r t0 n h h r` t0 in r` n co co co O tn 01 al u1 DD to N O1 O H r-I

t0 t0 to t0 t0 l0 t0 t0 t0 t0 t0 t0 . . . •

N N

H• j O H to N

O H (0r1 W N

d H H

-171-

c O H .-. CI Z I 0

Nq )

Û 1T `--

O O •tn

O O ul

•

L4

h (N 01 h ln H H O O O

O O O O O

I +I + 1 +I +I +I

lf1 O CO lD O l0 H CO l0 Ln

H H O O O

d' O a0 lfl d' H r-i H H H

O O O O O

I +I +I +I +1 + I

re) 0 N H H d' s co tn d'

H H H H H

[Pb

II]

exchanged

mole R,-1 x10+5

Osr- lo l0 LO I r r t` •I O O O O O

LO CO 01 N lO 01 01 O H H 1 . • O O H H H

II]

[Pb

residual

mole

Q,-1x10+5

01 0 N l+1 r1 M r 0 0 0 0 o d' d. d. d' dt. d.

01 fn H O t` /r1 r CO CO r lD l0

d. m f`7 Cl f`'1 fh

II

[Zn exchanged

mole

R.-1x10+5

CO O l0 ln f 1 d' [ CO H lf1

I O O O H H

h O d' N ln O ( CO

• I O O O O O

tn + 21

H '-' r d

H H N

h O _ E

1/40 l0 l0 lD N I O O O O O

M V. d' 111 ID

I O O O O O

'

Lf) +

H O ro H +1 k O •-. 11 I d H

H Q) C H N 8

CO d' N N tn O frl to co (Ni

I • HHHH (N

H (N CO lfl W O O O N d'

I H H H H H

% A r-1 .A

O Ih (N d' d' d' O O O O 0 O

d' V. d' d' d d'

O r l0 O H (N O 1/40 ln d' N O

a) r r r r r

X 04

ON (N d' tl) r o 0 0 0 0 0 d, d' d' d' d' d'

r N d' d' O ln h r r r r lo

r r NNNN

v Cil . ., Ç:

•r1 -,-i E E

O ln ln 0 O O H m l0 (N

H

0 tn Ln 0 0 0 H M l0 N

H


—172—

ZnS

concn

g

Q-1

[ S5)

ô

Ln -,I a7

ô • in

I It

•I

•r1 4.)

sx

01 o Co O O r1 O H

0000

I +I + I +I +I

M ul N u1 Cr 01 rn C1

O O O O

o H

0

+I

U1 Ol

O

I

r` Co h in O O O N

0000

+I +I +I +I

M 1/40 N CO r h h cr

O O O N

Ln t0

0

+1

0 In

1/40

'0 W b'+ In

+ d O

Û X X.-. 0 Id

H H Q)

.fA H Aa p E

N CO V' l0 Ln in t0 r-

I O O O O

t0 r

O I

r1 O CO CO m W Co

0H .-I N

O C11

M

[Pb"

]] re

sid

ua l

m

ole t

-lx

10+

5

01 C` H ln M r` NNH O

V' cM V' cr V'

m O

V

01 r

V'

O M H H O cr 01 01

V' en N H

01 CO

O

II

[ Zn ]

exch

anged

m

ole

Q-lx

10+

5

1/40 H 0 0 in tfl r CO

I O O O O

0 CO

O I

h CO N 1..0 01 h tD H O H N r-4

O l0

O

II

[Zn 'b

lan

k m

ole

2.-1x

10+

5

O O r1 d' 1/40 t0' t0 t0

I • • O O O O

h t0

• O

I

M CO V V' 01 t0 H ul • .

O r1 0H 10

O W

• t0 H

ul +

rl O

43 K

. . I 02

H Hÿ

Q)

N H m

t0 r-I rl V' H N en V'

I H H rl H

s V'

r-1 I

O 10 t0 O 01 V' l C`

H M h O H

O V

N H

(xÿ x C a H

,q

O co 01 C11 CS1 o ri rl rl rl

. . . . .

Ul to In ul vl

C11 r .

U1

O o .

tn

M .--1 M r1

u1 [r N ul

. . . .

u1 Ln t0 t0

O 01 .

t0

.^". Q4

O O N Ul O ONNN f`1

u1 ul Ul U1 Ul

O M

U1

0 o

Ul

t0 t0 O 01 ul Co t0 H in Ln 1.0 h

C!1 V'

r`

C) Cn E s~ •.1 -rl E' E

O u1 in O O H m t0

O N r-1

O ut to O O H M t0

O N H

PbII

A

CTIV

AT

ION

OF

ZIN

C S

ULP

HID

E

II [ ]blank) I] blank

6.2.2(b) Results.

(I) experiments in normal light Tables 6.6 to 6.8.

R. For the pH range 4 to 8, for two ZnS concentrations and a range

of PbII

concentrations, the value of R ranges from 0.50 to 1.73. For a given

set of experimental conditions, R decreases with reaction time. At initial

pH values of 4.0 and 5.0, for a ZnS concentration of 5.00 g2,-1 and for an init-

ial PbII concentration of 4.79 x 10-5 M, R is greater than unity after 5

minutes of reaction time and steadily decreases to less than unity. There is

a similar decrease at initial pH values of 6.6 and 7.8, however R is always

greater than unity. R is clearly dependent on pH.

At a given initial pH (e.g. refer Table 6.6) R is dependent on both the

concentration of ZnS and on the initial PbII concentration.

[Zn11] released and pH changes.

At initial pH values of 4.0 and 5.0, the pH increases during

to a greater extent than for the blank. At initial pH values of 6

7.8, the pH during activation either increases slightly or remains

constant, while the pH of the blank decreases.

The concentration of ZnII

released in the blank experiments

increases slowly with time after an initial, rapid increase. [Zn I

activation

.5 - 6.6 and

essentially

depends on pH, ZnS concentration etc. and has already been discussed in Chapter

4 (section 4.2.2(c)).

The release of ZnII during activation `[ZnII]total

) is also characterized

by a rapid, initial release followed by a slow increase with time. [ZnII]total

is greatest at an initial pH of 4.0 and decreases with increasing pH, other

conditions being held constant. The [ZnII]total

is dependent on both the ZnS

concentration and initial PbII concentration. The [Zn1I]exchanged increases

with decreasing pH, other conditions being constant. Furthermore

-174-

[Zn1]exchanged is dependent on both the initial PbII concentration and on

the ZnS concentration.

(II) experiments in the dark. (Table 6.9(A)).

R. Within experimental error R is equal to unity and is independ-

ent of time.

[ZnII] released and pH changes.

Compared to an equivalent experiment carried out in normal light,

[SII] total

and [ZnII]

exchanged are reduced. [Zn

II]blank is approximately

the same as that obtained in normal light.

The pH increase during activation at an initial pH of 5.0 exceeds that

of the blank, however both these pH increases are less than those observed in

normal light.

(III) experiments under W irradiation. (Table 6.9(B))

Release of zinc from ZnS was enhanced under UV irradiation during

both the blank experiments and activation, compared to experiments performed

in the dark and in normal light. [ZnII]blank

is generally less than [ZnII]total'

however after 120 minutes of reaction time there is only a slight (tib%) diff-

erence between the two values.

During activation and for the blank experiments there is an enhanced pH

increase observed, compared to experiments performed in the dark and in normal

light. The pH increase during activation exceeds that of the blank.

R remained essentially constant for the first 30 minutes of reaction time.

Thereafter it increased to a value of 6.5 at 120 minutes.

6.3 RATE EQUATIONS

Experimental evidence has shown that PbII uptake is

(i) dependent on initial PbII concentration, particularly during the

second stage of activation.

(ii) independent of ZnS concentration during the initial, rapid stage of

activation but dependent on ZnS concentration during the second

stage.

(iii) independent of pH during the initial stage, following activation at

initial pH 4.0 and 5.0; similarly at initial pH 6.6 and 7.8, but

PbII uptake is enhanced compared with acid initial pH values.

(iv) dependent on pH during the second stage of activation.

(v) dependent on incident light after about 30 seconds reaction time.

(vi) independent of stirring speed.,

In Figures 6.7 and 6.8, PbII uptake has been plotted as a function of

lo410(time)for all of the sets of data (except those obtained under UV

irradiation) shown in Figures 6.2 to 6.6 and Tables 6.2 to 6.5.

PbII uptake is linearly dependent on 1og10(time) for the initial, rapid

activation step (0 to about 1.0 minute ). After the first stage has been com-

pleted PbII uptake does not follow any clear rate law up to 120 minutes of

reaction time.

6.3.1. Initial, rapid activation step

acid pH values (4.0, 5.0) Figure 6.7.

The logarithmic dependence during the initial step indicates that Pb


r = k1 logi0 (t) + t1 (1)

or dT � k1 (2) dt 2.3O3t

II

where t is the uptake )f PbII by zinc sulphide, expressed in mole g-1, t is

the time in minutes, r t is the PbII uptake at 1 minute and k1 is a constant.

In contrast to the behitviour of CuII and Cd II, kl depends to a limited extent

PbU

UP

TAK

E (M

OL

E G

-1 Z

nS

x 1

0+6

)

2.2

0.6

0.2

1.4

1.0

0.40

0.20

FR

AC

TIO

N O

F M

ON

OLA

YE

R

r

-176-

10'1 100 101

102 TIME (MINUTES)

FIGURE 6.7 PbT1 uptake as a function of log10 (time) (Tables 6.2, 6.3 (A), 6.5 (A) )

pHinitial - 5.0, [Pb]II initial - 4.79 x 1.0'5 M, [ZnS] _= 5.00 g1"1

pHiniti ! "- 5.0, [PbII1 initial = 4.79 x 10'5 M, [ZnS] = 10.00 g1'1

pHinitia! = 5.0, (Pb l initial = 0.96 x 10'5 M, [ZnS] = 5.00 g1-1

• pHinitial - 5.0, [PbII1 initial = 4.79 x 10'5 M, [ZnS] = 5.00 g1'1

+ p'` initial = 4.0, (P13311 initial '-4.79 x 10'5 M, ['ZnS] •-•5.00 g1-1

ZnS (S5 )

normal light

in dark

normal light

PbI

I UP

TA

KE

(M

OL

E G

.1 Z

nS x

10

+6)

2.2

0.6

0.2

2.4

1.8

1.4

1.0

0.40

0.20

FR

AC

TI O

N O

F M

ON

OLA

YE

R

101 100 101

102 TIME (MINUTES)

FIGURE 6.8 PbII uptake as a function of logt (time) (Tables 6.3 (B), 6.4)

° pH initial - 6.6, [Pbil] initial = 4.79 x 10"

5 M, [ZnS] = 5.0021"1 (S5)

pH initial = 6.6, (Pbil] initial = 4.79 x 10"5 M, (ZnS] = 10.0091"1 (S5)

PHinitial = 7.8, [PbU] initial = 4.79 x 10"5 M, [ZnS] = 5.OQjl"1 (S5)

on the particular experimental conditions. For the curves shown in Figure

6.7, kl varies from 0.51 ± 0.10 x 10-6 mole g-1 (0.89 ± 0.17 x 10-6 mole m-2)

to 1.1 ± 0.1 x 10-6 mole g-1 (1.9 ± 0.2 x 10-6 mole m-2).

near neutral and alkaline pH values (6.6, 7.8) Figure 6.8

The logarithmic dependence of r on time is again obeyed. k1 has a value

of 0.92 ± 0.08 x 10-6 mole g-1 (1.6 ± 0.1 x 10-6 mole m-2), representing an

average of the three sets of data shown in Figure 6.8 i.e. in the pH range

6.6 to 7.8, kl is constant, in contrast to the behaviour in the pH range 4

to 5.

6.3.2. Second, slow activation step.

During the second stage of activation during the period from about 1 to

120 minutes, PbII uptake does not follow any clear well-defined rate equation.

For some experiments there is an approximate linear dependence on loglp(time)

for the 1 to 10 minute period during the second stage, (e.g. at pH 5) but this

is not generally true. Any general rate equation which attempts to express

PbII uptake must, at least, be of the following form:

r II time, pH, MIS], [Pb ], incident light)

6.3.3. Activation energy.

PbII uptake was measured at 298, 308 and 318 K. Experiments were per-

formed over the time span 0 to 10 minutes, in normal light, at an initial pH

of 5.0 with [PbII]initial = 4.79 x 10-5

M and MIS) = 5.00 gk-1 (S5).

The rate of PbII uptake is relatively insensitive to changes in temper-

ature, as is shown by the rate constants tabulated below.

Temperature (K) k1 (mole m-2 x 10+6) k2 (mole m-2 x 10+6)

298 1.8 ± 0.2 0.47 ± 0.2

308 2.3 ± 0.2 0.51 ± 0.2

318 2.1 ± 0.2 0.70 ± 0.2

This demonstrated insensitivity of Pb11 uptake towards changes in temper-

ature in the range 298 to 318 K does not permit any calculation of activation

energy to be made.

CHAPTER 7 MASS SPECTROMETRIC DETERMINATION

OF ELEMENTAL SULPHUR ON ZnS SURFACES.

In this Chapter mass spectrometric evidence is presented concerning both

the qualitative and quantitative determination of elemental sulphur on zinc

sulphide surfaces.

7.1 IDENTIFICATION OF ELEMENTAL SULPHUR.

7.1(a) Identification of major ions.

(I) Preliminary comments. When an analysis sample containing elemental

sulphur (S°) is inserted into the ionization chamber of the mass spectrometer,

the S° evaporates from the ZnS surface. The evaporation rate is determined

by the solid probe temperature, vacuum pressure, nature of packing in the

capillary tube etc. The concentration of S° reaching the ionization chamber

normally increases to a maximum and then gradually decreases to zero (i.e. the

background level) (see Figure 7.4 for example). Careful checking showed that

ZnS did not contribute any peaks above the background level in the temperature

range studied (i.e. approximately 308 to 373K). This is in accord with the

mass spectrometric study by Goldfinger and Jeunehomme(87), who showed that sub-

stantial decomposition of ZnS only occurs at temperatures above 1000K at a

pressure of 5 x 10-7 mm Hg (6 x 10-5 Pa).

For good resolution in the mass range 10 to 220, a scan speed of 10 seconds

was required. This factor, coupled with the gradually increasing temperature

and evaporation behaviour of S° meant that identification of the major ions

could only be carried out on a qualitative basis.

(II) Major ions.

The mass meter was calibrated in the range 10 to 220. Figure 7.1 is

a line diagram of relative peak height (with background subtracted out) versus

150

100

50

+ S2

S3

I

+ S4+

S5+

S6+

I i t

150

0.01% S°/ZnS standard

150 _

CuII/ZnS pH 4 100

50

o

150

PbII/ZnS pH 4

RE

LAT

IVE

PE

AK

HE

IGH

T (

AR

BIT

RA

RY

SC

ALE

)

I

+ + + + S3 S4 S5 Se

I I I I I

+ + + + $3 $4 $6 se

I I L I I

S2+

2

+ $2

+

S+

I

S2

100

50

100

50

0

S+

1

—181—

0.02% S°/ZnS standard

S3+ S4+ S5+ S6+

I

100 200

M/e

Line diagram for elemental sulphur on ZnS surfaces prepared under various conditions. Mass range 10 — 200, Solid probe temperature 333 — 353 K.

FIGURE 7.1

-182-

M/e for elemental sulphur detected on various ZnS surfaces. The experiments

for copper and lead corresponded to the following conditions at an initial

pH of 4.0:

(i) [CuII]initial =

5.12 x 10-5 M, [ZnS] = 5.00 gR-1 (S2)

(ii) [FbII]initial = 4.79 x 10-5 M, [ZnS] = 5.00 gR-1 (S5)

Major peaks corresponding to S+, S2+, S3+, S4+, S5+ and S6+ were ident-

ified for both the standards and activated samples. S2+ at M/e 64 is the most

intense peak. Further work in the mass range 50 to 750 showed that small

(relative to S2+) peaks corresponding to S7+ and S8+ were also present.

7.1(b) Identification of elemental sulphur through isotope abundance.

Sulphur has 3 principal isotopes with the following isotopic abundances

(51) :

% abundance fractional abundance

S32 95.00 0.9500

S33 0.76 0.0076

S34 4.22 0.0422

G 99.98 0.9998

Since S2+ corresponds to the most intense peak in the sulphur mass spec-

trum, further identification of elemental sulphur on ZnS surfaces may be con-

veniently performed by comparing theoretical and experimental peak intensities,

arising from isotope combinations yielding peaks at mass numbers 64, 65, 66, 67

and 68.

The relative intensity of the isotope peaks in a molecule containing n

atoms may be expressed as (70, 82, 83):

-183-

(a + b + c n

where a, b and c represent the fractional isotopic abundances of sulphur.

Thus the isotopes of S give the terms

a2 + 2ab + 2ac + 2bc + b2 + c2

Hence the following theoretical relative intensities may be derived, taking

the 64 peak as 100%:

mass number

theoretical

relative intensity % relative intensity

64 a2 = 0.903 100

65 2ab = 0.0144 1.59

66 2ac+b2 = 0.0802 8.88

67 2bc = 0.00064 0.00078

68 c2 = 0.0018 0.002

Of these expected peaks only those occurring at mass numbers of 64, 65

and 66 are significant. Therefore comparison of theoretical peak intensities

with experimental values for peaks occurring at mass numbers 65 and 66, re-

lative to mass number 64, is a further step in the identification of elemental

sulphur.

Following extensive baking of the vacuum manifold, to ensure a very low

background, repeated scanning over the mass range 63 to 67 enabled the relat-

ive intensities shown in Table 7.1 to be determined.

TABLE 7.1 Relative peak intensities (based on 64 peak as 100%).

Mass Number Standards 0.02,0.01 wt% S0

CuII/ZnS pH 4

PbII/ZnS pH 4

Theoretical

64

65

66

100

2.0 ± 0.6

8.6 ± 0.5

100

1.8 ± 0.4

9.3 ± 0.5

100

2.0 + 0.5

8.7 + 0.4

100

1.59

8.88

The experimental conditions for copper and lead were as described above.

The results in Table 1 show that agreement between theoretical and ex-

perimental intensities for the sulphur 65 and 66 peaks compared with 64 is

good. Similar close agreement was obtained for other conditions where elemen-

tal sulphur was suspected.

The identification of the major ions S+ to S8+

in the mass spectrum,

combined with the close agreement between the theoretical and experimental

peak intensities, is conclusive evidence that elemental sulphur is present on

ZnS surfaces which have been "activated" under various conditions.

7.2 QUANTITATIVE DETERMINATION OF ELEMENTAL SULPHUR.

7.2(a) Analysis technique.

Since the S2+ peak at M/e 64 is the most intense peak in the mass spectrum

of S° it was decided, for the purposes of quantitative analysis, to carry out

scans in the mass range from approximately 62 to 67. The relationship between

the peak heights at mass numbers 64, 65 and 66 then provides an additional

check should any doubt arise concerning a particular 64 peak e.g. if it were

"off scale".

The peaks are symmetrical (see Figures 7.4 to 7.6), hence summation of

the peak heights for a complete scan will therefore be directly proportional

to the quantity of S° present. A complete scan corresponded to the time taken

for the sulphur peak at mass number 64 to return to the background level

after the analysis sample, contained in the solid probe, was inserted through

the vacuum lock directly into the ionization chamber. This generally took

place over a period of several minutes.

The background level was checked before and after each individual sample

(as was the background current) and generally increased only slightly during

analysis (more serious contamination, with a concomitant decrease in sensitiv-

ity, meant that extensive baking and cleaning procedures were required). The

peak heights mentioned above were determined with the background level "sub-

tracted out".

Standards were determined at the start, middle and end of an analysis run.

Unknowns were determined in duplicate at least.

A typical calibration curve for the standards is shown in Figure 7.2.

The curve is non-linear, probably due to the fact that the 70 V electrons used

undergo inelastic scattering from sulphur molecules. As the weight % of sul-

phur is increased, there are many more sulphur molecules in the vapour phase

leading to a greater number of scattered electrons. These scattered, lower

energy electrons evidently ionize the sulphur molecules more efficiently at

higher weight % sulphur, leading to an enhanced S2+ peak i.e. at the given

electron energy the efficiency of ionization and formation of S2+ increases

with sulphur concentration. Despite this observed non-linearity, sufficient

calibration points were always present for quantitative analysis to be perform-

ed and for very reproducible results to be obtained. Furthermore the standards

were chosen so that they normally fell above and below the S° levels in the

unknowns. The weight of S° in the unknowns was determined by taking the

appropriate relative intensity (sum of peak heights) and reading off the corr-

esponding weight % S° from the calibration curve.

0.01

0.02

WEIGHT PER CENT S°

FIGURE 7.2 Typical calibration curve for the mass spectrometric determination of elemental sulphur.

7.2(b) Detection limit.

(I) Weight %. A detection limit of 0.0005 wt % S° (1.6 x 10-7 mole

S° per gram ZnS) was set, representing a reading 3 times greater than the

background level. For a sample size of 8.0 mg this represents 40 nanogram

of S° ( 1.2 x 10-9 mole). This figure should not be taken as an absolute

limit With suitable refinements it should be possible to detect down to

about 4 nanogram of S°.

(II) Equivalent monolayer coverage.

For the purposes of discussion it is useful to calculate nominal

monolayer coverage (nominal meaning that no structural disposition of sulphur

atoms on the surface is implied). This calculation of monolayer coverage re-

quires a knowledge of the area occupied by a sulphur atom. Two cases may be

considered:

(i) Orthorhombic sulphur has 128 atoms per unit cell, the dimensions of

which are a = 10.44 R, b = 12.84 R and c = 24.37 R (8). In such a close packed array the void space is 25.9% of the total, hence each

sulphur atom occupies an area of 8.6 R2 . (Note that orthorhombic

sulphur is the stable form at STP (80)).

(ii) In cubic ZnS, cleavage along the 110 plane is normal. Each unit cell

(a = 5.42R) contains 4 sulphur atoms (8), therefore each sulphur atom

occupies an area of 29.4R2.

Both of these values have conventionally been used by various workers

(1, 76, 77) in discussions concerning elemental sulphur and monolayer coverage

on ZnS surfaces.

In this investigation ZnS sample S2 (0.68m2g-1) was used for mass spec-

trometric studies involving Cu" and ZnS sample S5 (0.57m2g-1) for CdII and

PbII. The weight % S° and mole of S° per gram of ZnS corresponding to mono-

layer coverage are shown in Table 7.2 for the two ZnS samples. Detection limits

-188-

are also shown and are expressed as percentage of monolayer coverage.

TABLE 7.2 Conditions for monolayer coverage and detection limits.

ZnS sample (A) Monolayer coverage

q ZnS

(B) Detection limits

wt% S° mole S0 per % monolayer

8.6Â2 per S atom

29.4Â2 per S atom

8.6A2 per S atom

29.4R2 per S atom

8.6R2 per S atom

29.4Â2 per S atom

S2

(0.68m2g-1) 0.042 0.012 1.3x10-5

-6 3.8x10 1.2 4.1

S5

(0.57m2g 1) 0.035 0.010 1.1x10-5

_6 3.2x10 1.4 4.9

From Table 7.2 it may be seen that, for a ZnS sample of surface area

0.68 m2 g-1, and assuming that each sulphur atom occupies an area of 8.6Â2,

the detection limit is 1.2% of a monolayer. This is some 10 to 20 times more

sensitive than other comparable published attempts (76,77,78).

-189-

7.2(c) Blank experiments.

S° was not detectable on the surface of

(i) "fresh" ZnS (i.e. direct from the storage container).

or (ii) ZnS which had been used in blank experiments, performed in the dark

and in normal light, over the pH range 4 to 10.5.

For ZnS (S2) conditioned under W irradiation at an initial pH of 4.0,

and for a blank experiment performed in normal light at pH 1.0, the following

results were obtained.

pH weight % S° mole S° per g ZnS % monolayer

8.6Â2/Satom 29.4,2/Satom

4.3(UV) 0.0016 ± 0.0004 5 ± 1 x 10-7 4 ± 1 13 ± 4

1.0 0.0006 ± 0.0003 2 ± 1 x 10-7

1.5 ± 0.7 5 ± 2

7.2(d) CdII activation experiments.

S° was not detectable on the surface of CdII activated ZnS in the pH

range 4 to 10.5 for experiments performed in the dark and in normal light.

For ZnS (S5) at a concentration of 5.00 g!t 1 ,activated with CdII (init-

ial concentration of 2.92 x 10-5 M) at an initial pH of 7.0 under W irradiat-

ion the following result was obtained:

pH weight % S° mole S° per g ZnS % monolayer

8.6Â2/Satom 29.42/Satom

6.6 0.0007 ± 0.0003 2± 1 x 10-7 2± 1 7± 2

7.2(e) CuII activation experiments.

CuII activation experiments were carried out under the following condit-

ions:

-190-

[Cu11]initial - 5.12 x 10-5 M

[ZnS] = 5.00 gR,-1 (S2)

There were no detectable differences in the amount of S° determined .per gram

of ZnS when thé ZnS concentration was increased to 10.00 g!C 1.

(I) Variation in sample drying conditions.

ZnS was activated with Cu" at an initial pH of 4.0, and after

solid-liquid separation was performed, the activated ZnS was washed and dried

in various ways (refer Chapter 3 for general details). The effect of these

various washing and drying procedures on the quantity of S° detected is shown

in Table 7.3.

TABLE 7.3 Effect of various washing/drying procedures

on S° determination (actual pH = 4.1).

Procedure wt% S°

1) Sample dried on filter paper over silica gel in desiccator (no washing or vacuum).

0.0048, 0.0052, 0.0056

2) Sample on filter paper washed with 3cm3 AR grade acetone and dried over silica gel in desiccator under water pump vacuum.

0.0053, 0.0050

3) Sample on filter paper washed with 3cm3 supernatant, then with 3cm3 AR methanol and dried as in (2).

0.0053, 0.0047

4) Sample on filter paper washed with 3cm3 conductivity water (pH 5.6), exposed to air for

0.0050

20 minutes, then dried as in (2).

Average 0.0052 ± 0.0003

Within experimental error, these various washing and drying procedures,

including limited exposure to air, had no significant effect on the amount of

S° detected. Procedure 3 (refer Chapter 3) was used as a standard method.

(II) Variation in Ionic Strength.

ZnS was activated with CuII at an initial pH of 4.0 at a background

ionic strength of 10-3 M KNO3. S° was detected at 0.0055 ± 0.0004 wt%. With-

in experimental error, variation in ionic strength had no detectable effect

on the amount of S° present.

(III) Variation in pH.

ZnS was activated with Cu" in normal light at various pH values

in the range 0.5 to 10.5. The amount of S° detected is shown in Table 7.4

and Figure 7.3. Under the particular experimental conditions the amount of

S° present decreases with increasing pH. Extrapolation of the curves shown

in Figure 7.3 suggests that S° is not detectable beyond about pH 7.5.

TABLE 7.4

Amount of elemental sulphur detected on the

surface of CuII activated ZnS (S2) as a

function of pH (normal light).

pH weight % S° mole S° per g ZnS

8.6Â2/Satom

% monolayer

29.4Â2/Satom

0.5 0.030 ± 0.003 9 ± 1 x 10-6 71 ± 7 244 ± 25

4.1 0.0052 ± 0.0003 1.6 ± 0.1x10-6 12 ± 1 42 ± 3

6.2 0.0016 ± 0.0004 5 ± 1 x 10-7 4 ± 1 13 ± 4

10.5 not detectable not detectable not detectable

not detectable

-192—

100

90

80

Ê 70

60 oa

co 50 a?

cc 4 }

40 g o z

30 *

20

10

0 2 4 6 8 10 12 14

FIGURE 7.3 % monolayer of elemental sulphur on the surface of

CO activated ZnS (S2 ) as a function of pH (normal light).

(IV) Variation in incident light.

ZnS was activated with CuII at an initial pH of 4.0 in the dark, in

normal light and UV irradiation. The results are shown in Table 7.5.

TABLE 7.5 Amount of elemental sulphur detected on the

surface of CuII activated ZnS (So) as a

function of incident light (actual pH 4.1).

Light weight % S° mole S° per g ZnS % monolayer 8.6R2/Satom 29.482/Satom

dark 0.0057 ± 0.0005 1.8 ± 0.2 x 10-6 13 ± 2 46 - 4

normal 0.0052 ± 0.0003 1.6 ± 0.1 x 10-6 12 ± 1 42 ± 3

UV irradiation

0.0058 ± 0.0005 1.8 ± 0.2 x 10-6 14 ± 2 47 4 4

Within experimental error, incident light has no significant effect on the

quantity of S0 formed.

(V) Stability of elemental sulphur at high pH.

ZnS was activated with CuII at an initial pH of 4.0 in normal light.

The pH was then raised to 9.5 and samples were taken at various times for S°

analysis. The results are shown in Table 7.6.

TABLE 7.6 Stability of elemental sulphur, initially formed

by CuII activation of ZnS at pH 4, at pH 9.5 as

a function of time.

Time, minutes

weight % S° mole S° per g ZnS

8.6Â2/Satom

% monolayer

29.4R2/Satom

0 0.0052 ± 0.0003 1.6 ± 0.1 x 10-6 12 ± 1 42 ± 3

1 0.0048 ± 0.0005 1.5 ± 0.2 x 10-6 11 ± 2 39 ± 4

10 0.0012 ± 0.0004 4 ± 1 x 10-7

3 ± 1 10 ± 3

30 not detectable not detectable not detectable

not detectable

° i S is unstable at pH 9.5 and is not detectable at 30 minutes, decreasing

to 25% of its initial value in 10 minutes.

7.2(f) Pbii activation experiments.

Pbii activation experiments were carried out under the following con-

ditions:

[Pb initial

.4.79 x 10-5 M

[ZnS] = 5.00 gß-1 (S5)

ZnS was activated with Pb" at initial pH values of 4.0 and 5.0 in normal

light, and at initial pH 5.0 in the dark and under UV irradiation.The quantity

of S° detected is shown in Table 7.7.

TABLE 7.7 Amount of elemental sulphur detected on the

surface of PbII activated ZnS (Ss) at various

initial pH and light conditions.

Conditions weight % S° mole S° per g ZnS

8.6â2/Satom

% monolayer

29.482/Satom

pH = 4.0 normal light

pH = 5.4 normal light

pH = 5.3 in dark

pH = 6.6

0.0030

0.0008

0.0008

0.0042

± 0.0005

± 0.0003

± 0.0003

± 0.0005

9 ± 2 x 10-7

3 ± 1 x 10-7

3 ± 1 x 10-7

1.3±0.2x10-6

8

2

2

12

± 2

± 1

± 1

± 2

30

8

8

41

± 5

± 4

± 4

± 5 UV irradiat-

ion

The amount of S° decreases with increasing pH under constant light con-

ditions. Within experimental error there is no detectable difference between

the amount of S° formed in normal light and in the dark, following activation

at an initial pH of 5.0. W irradiation causes a dramatic increase in the

quantity of S° formed on the surface of ZnS, following activation at an init-

ial pH of 5.0, compared with experiments performed in normal light and in the

dark.

7.3. SAMPLE MASS SPECTRA.

Due to the relative novelty of the analysis technique, sample "raw" mass

spectra are shown in Figures 7.4 to 7.6 . The mass range from approximately

62 to 67 has been scanned. The traces generally show background, sample

(denoted by two arrows showing insertion and removal of the sample) followed

by background. In some cases only a portion of the spectrum is shown (for

relative purposes only). In comparison with the standards, "tailing" was

frequently observed with the "activation" samples i. e. a longer time was

required to remove the S° from the surface, other conditions being constant.

-195-

(A) 0.01 wt% S° standard.

_-AM=NIM_ -=MMIN_ERVAMME t l ®

_= ===-_== _- == = _ _ a M= M I ,

= -__ iffiVffiEffiR_INEVE MIE i m='== =3.=0..... 3'e=== -- .-- _ B o === ----- - = .0i

mm

te

m

o® ® _ = =. .r=:= =- - - - --. . -_— = ==' = ==m= :ä ;. = ==== .asa= ____ _. , -._ ==;i â ® -. . .. =-:_ _. .`.._ =_._- =a 3 Ramora

sm

n -..------...=,=.-..3,E.-..==--._ -;_;.= ï -,-3. __ -- _: - - ': _=-......=-......:: n_ _:-:-.... ; =

Ea =a=: ° _ : i== m

a::=

1 3

s'IN ffi s Egte EM REM® ® =-= - 312 m- - = _-:_::;•--

-

:z' =g=:__

_ i : :: ;- r-=: =: i m ,- -, - : 5'-

, •

.■. .......... _ .,.. .... =--__ = s__ _

MgEEMa ffea ki

(B) CuII/ZnS pH 4.1 (background shown above).

(C) ZnS blank pH 4.0.

(D) Cuii/ZnS pH 6.2.

(E) ZnS blank pH 5.2.

Figure 7.4. Cuii/ZnS activation experiments. All spectra are on same scale.

[Normal light ; probe temperature 308-373K].

-196-

(A) 0.01 wt% S0 standard.

=s 1=1

= =_

2.—E ___ = =_ —__ =_ —.—.= âi== =__ ;_— _ = =_ =_ ,= —_ _= _ =_ : __ ;=_ _ ,- .-_

:__: ® __ _i _ âi.: ii: :z ::_ :== '__ ==3i -- "111-- -, -- S -

=_ :-- : : 1 _ : E __:: _• i i ®® i NiEligii....= .-_ =; =si= = sa: i ii , '3 := 9:= : _: =.:a ze

m= ;=_,?? :ë :;É s : - ,_;71:u. iz ;:o=:: _ : eE:a—••= --- — =. _ , . ._i:i.-i: =: ' ,=' " ? . ., m:1'; :: E;, i:E; =B!#?? 113 ; #p ; im== '

(B) portion of CuII/ZnS spectrum, pH 4.1 with background shown above.

(C) PbII/ZnS pH 4.0.

(D) CdII/ZnS pH 4.0 (background above, sample below).

_W_=_= _ EMB=_ E _

___= =_

â=- a==

_ =_= __E

e=111=_=IE l _ —__ i's—====1 ;x=====

° :3==91pimmis

_ ; __ EME1 ®

==_E__= _ __ _ _ _B _ 3I _ _-_

Activation experiments - CuII, PbII and CdII at pH 4. All spectra are on same scale, except (D) which is 2X. [Normal light ; probe temperature 308-373K].

Figure 7.5.

-197-

(A) 0.01 wt% S° standard.

(B) portion of CuII/ZnS spectrum - activated at pH 4.0, raised to pH 9.5 for 1 minute.

11 Izmir +i r ''1i1 GIIIII [: _ • :.. .?."'r..rs .r4'!'•1' ityWY' rr' i .7: :iTi IR i

(D) PbII/ZnS at pH 5.4.

! " ACCUCMART +-

! I

■■

ill:' NPR iiPIb11K I/! PRIM, i&" '.117.11 +A.1:► 1!!L'.tl'!'i7Wi1:' ` lLt;ii' "à1 Pl"iltiJllV" i!f l'l'r' ta

■!!ai ' ' d

1 it'al;I i tlait i II I tia rr r Rrrrrrÿrrst^rrmrnni rem !An rn

-4- + r- .-

IN aa ■ :: : __ E ri. .aa.a

iaw1

_—

a 1■ : ■

.E ■®1a a Mall

;: _ : ,. 1;

a■l''!1> ii ?1AaF7Rt1 1I It Mai I

1 ® PIAg m :w- r -1ti==,zJ:,: L ,â ill l l1" tl[!I te'illl,;ll.L.ai'i1gIIaRI: „IM ' r.' 1

(E) PbII/ZnS at pH 6.6 under UV irradiation.

r--F--I-1 I-

V VWY

i i 1 T

MINIM ■I■irl.

r Pia NaASI■Y r i ' ' Wii l' . WA4!iAäa iAlï!iA'4!a9'i" r ^!yJJ WvtiIL

■.a■. ■a■a 11111_111a ■i ■aa.aa■a as ■.oa y

■■■a■■a■aie . ■a■I ■ 111111111111111111111111 a■■ ■

a■■■■a■a■■ ■■aaa a■ iItttt IaNi i ■■a■a■■■■a■a■ ■ ■a MUM : il ltt® ■ti■it INN IMO ■l.a i i a■am » It ! Iia1 tfa! MI 1ai14üPEü !!il'#®i a■1 II11 ®I ■ ■ I ■ i 7I'raiKiril!' III I 'is i!N-'. R1hP9n1''lP1I lIA6! PII IIiP1111ii t /IIRrtA11i PI1!1 1®1 6/ 11(/®I r é:::+11C: 1;.:ii ils;i::a::A.,o3.i:Ctit'L'llû :üi :sl11'â7'i'C 17'11"IJ;tti'Ai.hYJ::n11J!I.:J1'l'L.l' ' ' 11'

(C) as in (B) but raised to pH 9.5 for 10 minutes.

Figure 7.6. Activation experiments - Cu", Pb". All spectra are on same scale, with bottom channel twice that of the top. [Normal light, except (E) ; probe temperature 308-373K].

, --Eh+

CHAPTER 8 INTERPRETATION.

8.1 ON THE STABILITY OF ELEMENTAL SULPHUR (S°).

Some 40 ionic and molecular species of sulphur in aqueous solution have

been isolated and identified (89,90). Valensi (15) has shown, however, that

only a few species are stable in appreciable quantity in the presence of water

i.e. S042-, HSO4-, S°, H2S, HS- and S2-.

The construction and use of equilibrium Eh-pH diagrams is well-established

(14). The Eh-pH stability fields for elemental sulphur and zinc sulphide in

equilibrium with dissolved sulphur and zinc are shown in Figure 8.1. The activ-

ity of dissolved zinc has been taken as 10-6, which effectively represents the

solubility of zinc sulphide (refer Figure 2.1). The activity of dissolved sul-

phur has been taken at 10-1 and 10-4(10-1 corresponds to that chosen by Garrels

(14), 10-4 is a convenient choice). The equations used in constructing this

diagram are shown in Table 8.1. Clearly the Eh-pH diagram for such a metal-

sulphur-water system is obtained by superimposing the respective diagrams for

the metal-water and sulphur-water systems. The field over which elemental

sulphur is stable is dictated by 4 equations - 10, 11, 13 and 14 in Table 8.1.

The S° stability field is critically dependent on the activity of dissolved

sulphur, contracting from an upper limit of pH 7.8 to pH 3, as the dissolved

sulphur activity is reduced from 10-1 to 10-4. According to Valensi (15) and

to Cloke (91), whilst polysulphide ions do not appear as principal species in

the Eh-pH diagram, they do assume appreciable concentrations near the alkaline

termination of the elemental sulphur field e.g. at a total dissolved sulphur

activity of 10-1, S52- has an activity of 10-4 at pH 7. The concentration of

polysulphides becomes increasingly important above pH 10 as pointed out by both

Valensi (15) and by Cloke (91).

Table 8.2 summarises the results of potential measurements conducted on

H2O ••■

H2

ZnS(s) HS- laa)

+0.6

+0.4

+0.2

0

- 0.2

- 0.4

- 0.6

Znls)' H2S1ag1

./".7:7 --

- 0.8

- 1.0

-199-

+1.0

+0.8

2 4 6

pH

8 10 12

FIGURE 8.1 Eh - pH stability fields for elemental sulphur and zinc sulphide in equilibrium with dissolved sulphur and zinc (298 K 1 atmosphere). The activity of dissolved zinc has been taken as 10-6 . Unbroken lines refer to dissolved sulphur activity of 10.1 , dotted lines to 10'4. Hatched region corresponds to experimental values (refer Table 8.2).

TABLE 8.1. Half-reactions and the relevant Eh equations which must

be considered in plotting the Eh-pH fields for elemental

sulphur and zinc sulphide in equilibrium with dissolved

sulphur and zinc in the pH range 0 to 14 (298 K, 1 atm)

Species are in aqueous phase unless otherwise specified.

(Data from ref. 14, 15, 68, 88).

Half-reaction Equation

1 02g + 4H+ + 4e -4- 2H2O Eh = 1.23 - 0.059 pH

2 2H+ + 2e — H2g Eh = -0.059 pH

3 H2S H+ + HS KH (H+](HS ]= 10-7.0 5 2 (H2S]

4 HS H+ + S2 KHS_ = (H+][S2 ] = 10-17.1 [HS-]

+ - 5 HSO4 + 9H + 8e t H2S + 4H2O Eh=0.290-0.066pH + 0.00741og10[HSO4 ]

[H2S]

6 5042 + 10H++ 8e 4 H2S + 4H2O Eh=0.303-0.074pH + 0.00741og10[5042 ] [H2S]

7 5042 + 9H+ + 8e 4 HS + 4H2O Eh=0.252-0.066pH + 0.O0741og10[SO42 ] (HS-]

8 SO42 + 8H+ + 8e t S2 + 4H2O Eh=0.148-0.059pH + 0.00741og10[S042 ]

[]

9 HSO4 f H+ + SO42 [H+][5042 ] KHSO -- - 10-1.9 4 [HSO4-]

10 SS + 2H+ + 2e + H2S Eh=0.142-0.O592pH-0.02961og10[H2S]

11 S + x+ + 2e f HS Eh=-0.0653-0.0296pH-0.02961og10[HS ]

12 Ss + 2e F S2 Eh=-0.476 - 0.02961og10[S2 ]

13 HSO4 + 7H+ + 6e f SS + 4H20 Eh=0.338-0.O691pH + 0.009871og10[HSO4 ]

14 S042- 042 + 8H+ + 6e { Ss + 4H20 Eh 0.357-0.0789pH + 0.009871og10[5042 ]

15 Zn(OH)2s --4- Zn2+ + 2OH KSO = 1.56 x 10-16

16 Zn2+ + Ss + 2e f ZnSs Eh=0.265 + 0.02961og10[Zn2+]

-

Continued

TABLE 8.1. (Continued)

Half-reaction Equation

r

17 Zn ++SO42 +8H++8e F ZnSs+4H2O Eh=0.334+0.00741og1p[Zn2+][5042 ]-0.0592pH

18 Zn2++HSO4 +7H++8e F ZnSs+4H2O Eh=0.320+0.00741og1p[Zn2+][HSO4 ]-0.0518pH

19 ZnSs + 2H+ * Zn2+ + H2S K = [Zn2+][H2S] = 6.82 x 10-5

[H+]2

20 ZnSs + H+ f Zn2+ + HS K = [Zn2+] [HS ] = 6.76 x 10-5

[H+]

21 Zn2+ + 2e t Zn s

Eh= -0.763 + 0.02961og10[Zn2+]

22 ZnSs + 2H+ + 2e F Zns + H2S Eh= -0.886 - 0.02961og1o[H2S]-0.592pH

23 Zn(OH)2s+SO42 +10InSe ; ZnSs+6H2O Eh=0.424+0.00741og10[S042 ]-0.0740pH

24 ZnS + Zn2+ + S2 KSO = 1.26 x 10-27

-202-

TABLE 8.2. Measured Eh potentials for selected activation systems.

Experiments performed at I = 10-1 M KNO3, in normal

light at 298K.

System Measured Eh

volt Comment

A. [ZnS]=5.0gQ 1, pH=4.0 0.495

B. [CuII]=5.12x10 5 M,pH=4.0 0.576 initial conditions

ZnS added at 5.091-1, after 1 min 0.556

5 min 0.513 separate test:

15 min

then pH increased to 5.0, after 5 min

pH increased to 7.0, after 5 min


0.484

0.450

0.396

0.327

UV irradiation has only a small, if any, effect at 15 min; irradiation for further 15 min results in pot-ential decrease of about 2 mV.

C. [PbII]=4.79x10_ 5 M,pH=4.0 0.524 initial conditions

ZnS added at 5.0 gt 1, after 1 min 0.512

5 min 0.495 separate test:

15 min 0.492 UV irradiation causes the potential to decrease by 20 mV over 15 min.

then pH increased to 5.0, after 5 min 0.484

The potential slowly (10 min) re-turned to within 5 mV of the pre-



0.418

0.397

irradiation value of 0.492 volt when the UV irradiation was shut off

D. [CdII] = 2.92x10-5 M,pH=4.0 0.540 initial conditions

ZnS added at 5.0 1M, separate test: gk after 5 min 0.540 UV irradiation has only a small, if

any effect at 5 min, irradiation then pH increased to 5.0,

after 5 min


0.514

0.451

for further 15 min results in potent-ial decrease of about 2 mV.

certain activation systems studied in this investigation. Similar measurements

are frequently performed in flotation pulps (1,2). It is common to find state-

ments regarding the stability of So being made in relation to the Eh-pH diagram

drawn at a dissolved sulphur activity of 10-1, when an activity of 10-4 or less

is more appropriate. The statements made in the preceding paragraph, taken in

conjunction with Figure 8.1, clearly illustrate the dangers involved in extra-

polating from a "standard" Eh-pH diagram to other ill-defined conditions.

The interpretation of measured Eh potentials is, moreover, fraught with

difficulties, the most serious of which are discussed briefly below:

(i) It has been established that platinum electrodes may be "poisoned"

by species such as H2S and S2 in flotation pulps (92). An electrode which

has been poisoned by H2S may register a potential which is independent of

the solution phase (92), due to the formation of a resistant film which

affects the charge transfer process at the electrode-solution interface.

In this present study, constant checking of the platinum-calomel pair

against ZoBell's solution (14, Chapter 3) suggested that poisoning of the

platinum electrode surface did not occur to any significant extent.

(ii) In the case of "slow" electrode reactions (5), the platinum elect-

rode may not respond to the reaction in question but rather to spurious

oxide potentials such as Pt(OH) + e Pt + OH- (standard reduction potent-

ial in acid solutions of 0.84 volt).

(iii) In the absence of well-defined redox couples having appreciable

(> 10-7 A cm-2) exchange currents and in sufficient concentrations, the

potential readings represent "mixed potentials" (5, 93-95). A mixed pot-

ential arises when the platinum electrode responds to a number of redox

reactions e.g. a trace of dissolved oxygen may have a marked effect on

the measured potential (5, 93, 94).

In view of these drawbacks, it is therefore highly unlikely that the

measured Eh potentials represent reversible, equilibrium potentials. They

cannot be used in quantitative discussions regarding solution composition.

Two further points must be considered:

(i) The Eh-pH diagrams are constructed from standard thermodynamic data

relating to reactions occurring at a platinum electrode. These reactions

may show rather different behaviour at a ZnS electrode and some caution

should be exercised in making comparisons. Williams (63) has made similar

comments in his study of the reactions occurring on the surface of GaAs

crystals in contact with aqueous electrolyte solutions.

(ii) The pH at the ZnS-solution interfade will probably be less than in

the bulk solution. The latter value is, of course, used in the Eh-pH

diagrams. Moiqnard, Dixon and Healy (9, also see Figure 8.5) have measured

negative mobilities of about 2.6 usec-1/volt cm-1 at pH 4 for CuII-ZnS

systems which corresponded very closely with the conditions used in this

investigation (i.e. I = 10-2 M KNO3 and (CuII]/ZnS area ratios above

and below those used here. Using this value, assuming that the simple

Sneoluchowski equation is valid, and that the shear plane coincides with the

Stern Plane, applying the Gouy-Chapman theory yields

pHStern Plane pHbulk ti -0.56

Allowing for the fact that the ionic strength in this present study is

higher (10-1 M KNO3), it is still likely that the pH at the Stern Plane is

about half a pH unit less than the corresponding bulk value, thus in part

contributing to the "stability" of the detected S0.

Without exception all of the measured Eh potentials are significantly higher

than the Eh-pH stability field for S°(at 10-1 or 10-4 dissolved sulphur activity)

predicts, assuming that S° is the stable phase. The effect of UV irradiation on

-205-

the PbII-ZnS system is significant. Taken in conjunction with the relevant

uptake data (Figure 6.6) and the enhanced formation of S° (Table 7.7) this

indicates that the rate of at least one reaction is increased in the presence

of UV irradiation.

The mass spectrometric evidence presented in Chapter 7 clearly shows that

° i S is present on the surface of ZnS under certain specified conditions used in

this investigation. Furthermore the stability of this S° decreases with in-

creasing pH (Tables 7.4 and 7.6). In this respect there is general agreement

with the equilibrium Eh-pH predictions. Bearing in mind the qualifications

outlined above with respect to the interpretation of measured Eh values and,

noting that pHStern Plane

is probably less than pHbulk, the experimental re-

sults nevertheless indicate that the S° formed on ZnS surfaces in the presence

of Cu, PbII and/or UV irradiation at acid pH values is unstable, slowly

decomposing to sulphur-hydrogen-oxygen compounds ( e.g. Tables 2.3 and 8.1):

i.e. S° SIQW H S O Reaction A

surface -4-- x y z

The results in Table 7.4 show that the amount of S° detected on the sur-

face of CuII or PbII activated ZnS in the pH range 4 to 6 decreases, suggesting

that the rate of decomposition increases with increasing pH (but is still slow

over the time span of 2 hours). The type of compound formed (and the values of

x, y and z) will depend on the prevailing Eh-pH conditions.

At alkaline pH values the mass spectrometric evidence (particularly Table

7.6) suggests that S° undergoes rapid decomposition (compared with acid pH

values).

o FAST

S surface Hxsy z Reaction B i.e.

8.2. THE SIGNIFICANCE OF THE LOGARITHMIC LAW.

8.2(a) Comparison with other activation studies.

The literature dealin• with •reviousl determined rate e•uations and act-

ivation energies has been criticised in detail in Chapter 2. Some brief com-

parisons are pertinent at this point

(I) Rate laws

CuII. The uptake of CuII during the initial activation step in the

pH range 4 to 7 has been shown to obey a logarithmic law, in agreement with the

work of Bazanova and Mitrofanov (37,39), whose results were rather inconclusive,

for they also suggested that first order kinetics might be involved. The re-

sults of this study show that the rate law during the initial step is clearly

not first order, as proposed by Mukai (25) and others (37,39). The logarithmic

law is also obeyed during the 1.5 to 15 minute period of the second stage of

activation, after which there is, apparently, no simple law up to a reaction

time of 2 hours. Certainly the parabolic rate law proposed by other workers

(25,26,36) is not obeyed. The data obtained by Gaudin et al (36), however,

apply to contact times of up to 63 hours, a far longer period than was studied

here and during which other mechanisms may be involved. Gaudin et al's study

applies only to a single pH value.. The work by Wada (26) and Mukai

(25) is similarly restricted. Only Mitrofanov and Bazanova (37) have shown that

CuZZ uptake at acid pH values is dependent on pH during the second stage of

activation.

The lack of any clear rate law at alkaline pH values is consistent with

other observations (25, 26, 36-39).

CdII

The uptake of CdII follows a logarithmic dependence on time in

the pH range 4 to 9. Beyond about pH 9 there is no clear rate law. As far as

can be ascertained no other rate laws relating to the uptake of CdII on ZnS

(or any other sulphide) have so far been proposed.

PbII

During the initial activation step in the pH range 4 to 8, PbII up-

take follows a logarithmic dependence on time, however the value of the rate

constant depends to a slight extent on the particular experimental conditions.

The data obtained in this study do not conform to the relationship

r = to (where n varies from 0.20 to 0.28) k

proposed by Fuerstenau and Metzger (97). There is general agreement, however,

between the current and latter studies in that

(a) the value of the rate constant depends to some degree on the

particular experimental conditions

and (b) the kinetics of PbII abstraction are considerably more complex

than for CuII and CdII.

During the second stage of activation the results of this study indicate

that there is no well-defined rate law (although the logarithmic law is obeyed

during at least part of the second stage during activation at an initial pH of

5), in contrast to Fuerstenau and Metzger's observations (97).

Of the metal ion activation studies performed on zinc sulphide by other

workers, only the uptake of AgI has been shown to obey the logarithmic law (98).

In this respect it shares a common characteristic with the activation behaviour

of CuII, CdII and PbII reported in this study.

(II) Activation energy.

The results of this present study have shown that the experimental

rate constants for Cu", CdII or PbII uptake are insensitive towards temperature

changes in the range 298 to 318 K during both the first and second stage of

activation. Activation energies could therefore not be evaluated. Wada and

Okada (26) have shown that the rate of uptake of CuII during the initial step

is insensitive to temperature changes in the range 293 to 323 K, although they

did determine an activation energy of 58 kJ mol-1 for the second step, assoc-

iating it with a parabolic rate law and structure sensitive diffusion processes.

Mukai and Nakahiro (25) determined an activation energy of 23 kJ mol-1 for the

first step, associating this with diffusion control in the solution phase and

first order kinetics. These two workers also determined an activation energy

of 36 kJ mol-1 for the second stage, during which they proposed that a parabolic

rate law was obeyed. No activation energies are reported for either CdII or

PbII uptake.

The results of this present study, notably the form of the rate laws coup-

led with the demonstrated insensitivity of the rate constants towards changes

in temperature, stirring speed, ionic strength etc., suggest that the kinetics

and mechanism of heavy metal ion activation is more complex than has previously

been realized.

8.2(b) The link with mechanism.

The logarithmic dependence of MII uptake on time shows a pattern of behav-

iour which is common in many types of "tarnish" reactions i.e. where a gas or

a liquid reacts with a solid surface (22 to 24). This logarithmic relationship

is particularly prevalent in the formation of very thin (about 30 R or less)

oxide films on metal and semiconductor surfaces e.g. in the low temperature

oxidation of aluminium (27, 109) copper (23a), germanium and silicon (24) etc.

The oxide thickness versus time curves obtained in these cases naturally bear a

striking similarly to the MII uptake-time data reported in this investigation.

The foundations of this logarithmic law for thin film oxidation processes have

been firmly laid by Landsberg (27), Mott (28) and Evans (23a) and it is gener-

ally expressed as

x = k1 in (k2t + k3)

-209-

where x is the film thickness formed in time t and kl to k3 are constants

(29). This logarithmic law is also often called the Elovich rate equation

(96) .

Establishing a rate law does not, however, distinguish a particular mech-

anism. Using an ideal oxide film as a model, Ritchie and Hunt (30) have clearly

shown that the presence of a slow surface step can lead to a logarithmic law.

The formation of a very protective space charge layer in oxides, due to positive

or negative ions moving more slowly than electrons, also results in a logarith-

mic law (19). It has also been found that the logarithmic relationship is

obeyed for charge transfer processes occurring at semiconductor surfaces during

chemisorption (96).

On the other hand Evans suggests that discrete mechanical breakdown of

the oxide film is characteristic of the logarithmic dependence of film thickness

on time (23). His theory, which is strongly supported by electron microscope

studies, is based on the existence of "flat gaps" or cavities between the film

and the metal. These form due to differences in the lattice dimensions of

the oxide and the metal; stress results, followed by subsequent mechanical

breakdown of the film. The flat gaps cause a reduction in the available sur-

face area and a consequent marked decrease in oxidation rate. In addition

Evans (23 a,b) reveals that a pore blocking mechanism can also lead to the

same logarithmic equation. In this case the volume of oxide formed is greater

than that of the metal destroyed, producing a lateral pressure which progress-

ively blocks the pores or paths of loosely arranged atoms along which oxidat-

ion proceeds.

Clearly further insight into the probable activation mechanisms of Cu",

CdII

and PbII with ZnS will best be obtained by comparing the relative rates of

uptake of the three metals, the effect of irradiation together with the presence

or absence of So on the activated ZnS surface. This comparison is carried out

in Section 8.4 and is preceded by some thermodynamic calculations in Section 8.3.

8.3. FREE ENERGY CALCULATIONS.

Extrapolating the results of calculations based on bulk thermodynamic

quantities to surface behaviour contains an element of risk. For example,

entropy terms in the bulk and at a surface are known to be different (50).

In the absence of precise surface thermodynamic information there is unfortun-

ately no alternative. The data contained in the following tables have been

determined using information extracted from references 5, 14, 51 and 53. The

bulk of this information was obtained from reference 14. The stable form of

the solid phase at 298 K and 1 atmosphere has been assumed present (note that

the conclusions reached remain unaffected even if amorphous, metastable solid

forms are present).

8.3(a) Oxidation.

The oxidation of sulphides has been reviewed in Chapter 2. The first step,

in the presence of dissolved oxygen, corresponds to (in acid solutions)

MSsolid + 02a

+ 2H+ i--- M2+ + Ssolid + H201i (1) Q Q Q Q

Thus the standard free energy of reaction `AGD may be calculated, yielding

the values shown in Table 8.3.

TABLE 8.3. Standard free energy of reaction for the oxidation

of various sulphides.

Compound AGr, kJ (298K)

ZnS -179.1

CuS -115.3

CdS -166.4

PbS -160.9

-211-

Additional oxidation reactions involving dissolved oxygen may also be formu-

lated. The important point, however, in all of these reactions is that Ad:

is always large and negative i.e. in equation (1), under standard state

conditions, the formation of elemental sulphur due to oxidation by dissolved

oxygen is thermodynamically possible. Even at initial, very low (e.g. 10-10M)

levels of dissolved oxygen, calculation shows that AGr is still negative, a

fundamental point which has frequently been overlooked and to which Gaudin (52)

has recently called attention. (N.B. AGr is the free energy of reaction under

specified experimental conditions away from the standard state and is related

to AGr through AGr = AGr + RT Eniknai (18)).

Whilst recognizing these thermodynamic implications the rate of oxidation

must also be considered. For synthetic and freshly formed sulphides, such as

those used in this present study, the rate of oxidation is generally quite slow

and should not occur to any significant extent during the two hours of reaction

time, under the particular experimental conditions used. The observations that

(i) the introduction of a small amount of oxygen into the system did

not affect the uptake of Cu", CdII or Pb"

and (ii),variations in drying conditions, including limited exposure to air

had no detectable effect on S° determination during CuII activation of

ZnS supports the published evidence that the rate of oxidation is slow.

(Note that the concentration of S° detected on the surface of CuII activated

ZnS at initial pH 4 is 8 to 10 times the level of 02 present, with both quan-

tities expressed as mole L-1). The concentration of dissolved oxygen initially

present in the systems studied does not exceed 10-6 M (0.05 ppm), so that the

quantity of oxidation products formed, if any, will be very small indeed.

8.3(b) Exchange reactions.

The exchange reaction between an aqueous metal ion, M2+ , and zinc sulphide aq

may be expressed as

2+ 2+ Maq + ZnSsolid É

MS solid + Znaq

leading to the AGr values in Table 8.4.

TABLE 8.4. Standard free energy of reaction for the

exchange reaction between M2+ and zinc sulphide.

M2+ aq

AGr, kJ (298K)

Cu2+ aq

Cd 2+ aq

Pb2+ aq

Under standard state conditions the exchange reaction between Cu2+, Cd2+aq aq

and Pb2+ is thermodynamically favourable. Moreover under the experimental aq

conditions used, calculation shows that AGr is negatitre in all cases hence the

exchange reactions are favourable once again.

8.3(c) Selected redox reactions producing elemental sulphur.

The reaction of Cu2+, Cdaq+ and Pb2+ with solid ZnS, H2Saq ,.HSaq andaqaq

S2 has been considered. The +1 oxidation state exists for copper, but is aq

unknown for either cadmium or lead (16,20,99) in bulk compounds or in aqueous

solution. The various reactions together with the relevant AGr values are

shown in Table 8.5.

(2)

2 Cuaq + 2 H2Sa q

Cd2+ +• H2S aq aq

Pb2+ +• H2S aq aq

2 Cut+ +• 2 HS- aq aq

Cd2+ +• HS aq aq

Pb2+ +• HS aq aq

2 Cut+ + 2 S2- aq aq

Cd2+ + S2- aq aq

Pb2+ + S2 aq aq

-213-

TABLE 8.5. Standard free energy of reaction for

selected redox reactions.

AGr, kJ (298 K) Group Reaction

A.

B.

C.

D.

2 Cuaq + 2 ZnSsolid ` ) Ssolid + Go2s soli d

+ 2Znâq

Cd2+ +• ZnS ----+So + Cdo + Zn2+ aq solid -4--- solid solid aq

Pb aq + Ssolid t-} Ssolid + Pbsolid + Znaq

o Ssolid

+ Cu2Ssolid+ 4H

aq

O 0 + � S

+ Cd solid

2H solid solid

aq

O 0 + � ---+Ssolid

+ Pbsolid + 2Haq

�

o + solid + Cu2so?.id

+ 2Haq

0 0 + �

so +

Cdsolid + Haq

O 0 + � Ssolid +

Pbsolid Haq

----+ Ssolid + Cu2Ssolid

--► S + Cd

o - solid solid

- So + Pbo solid solid

-113.9

+128.9

+ 75.4

-161.4

+105.1

+ 51.7

-241.3

+ 65.1

+ 11.7

-399.9

- 14.1

- 67.6

For reactions involving Cut+ , under standard state conditions the format-aq

ion of elemental sulphur is thermodynamically favourable in all cases, stressing

the instability of the +2 oxidation state in copper. 1Gr is also negative under

the prevailing experimental conditions.

In contrast, under standard state conditions,

reactions involving Cd2+ and Pb2+ is generally aq aq

exception of reactions involving S2 Reference aq

the formation

unfavourable,

to Figure 2.1

of Ssolid

from

with the

reveals that the

the activity of Sâq in equilibrium with solid ZnS in the pH range 2 to 8

is miniscule (i.e. much less than 10-16). Assuming metal ion activities of

10-4 and, say 10-16 for the activity of S2-, a simple calculation of AGr for

Cd2++ and Pb- results in AGr values of +100.1 and +46.6 kJ respectively. The

formation of Solid from either Cd2+ or Pb2+ (with the production of met- solid aq

allic Cd and Pb) is therefore thermodynamically unfavourable under the exper-

imental conditions used.

From the mass spectrometric data presented in Chapter 7 there is no

evidence to suggest that S is present as a separate (from ZnS), distinct

colloidal phase. If it were, variations in washing procedures for example

would be expected to have a marked effect on the level of S° detected and in-

consistent values would generally appear, in contrast to the observed results.

Reactions in Groups B to D in Table 8.5, which suggest that such a colloidal

phase might occur, therefore appear improbable. Furthermore, reactions in

Groups B and C predict that there would be a pH decrease during activation

(relative to the blank), which is quite the reverse of the observed pH

increase.

The mass spectrometric and kinetic evidence, when taken together, indicate

strongly that S° is formed on the ZnS surface.

A COMPARISON BETWEEN CuII, CdII and PbII.

8.4(a) Relative rate of uptake of CuII, CdII and PbII

(I) General Remarks

MII uptake as a function of time curves are shown in Figure 8.2 for

activation at acid to neutral pH values in the dark, in normal light and under

W irradiation. The rate constants k1 and k2, obtained as a function of

temperature during activation at an initial pH of 5.0 are shown in Table 8.6.

Temperature changes in the range from 298 to 318 K have no detectable influence

8.4.

(a)


(c)

IN DARK

10

8

6

4

2

(b1

i 0

I t i 0 30 60 90 120

TIME (MINUTES) (d) NORMAL LIGHT

10

10

0 1 2 3 4 . 5 TIME (MINUTES)

10

8

6

4

-215-

30 60 90 120 TIME (MINUTES)

UV IRRADIATION (e)

1 1 I 1

0 30 60 90 120 TIME (MINUTES)

FIGURE 8.2 Comparison between MIIuptake as a function of time for Cu (0), CdIIl0) and Ptf7t 111) in the dark (a,b), normal light (c,d) and under UV Irradiation (e).

CONDITIONS: [ZnS] = 5.00 g1-1(S2) (Cu initial [ZnS] = 10.00 sill 4S5) fCdII1lWitiel [ZnS] = 5.00 g1-1 (S5) [PbII] initial - 4.79 x 105 M (Tables 6.21A), 6.5)

pHlnitiel .. 5.00 pHinitiel - 7.00 pHinitial - 7'00

6.12 x 10-5 M (Tables 4.6(A), 4.8(A), 4.9(A) n• 2.92 x 105 M (Tables 5.3(A), 5.6)

TABLE 8.6. Comparison between rate constants for CuII,

CdII and PbII uptake as a function of temp-

erature at initial pH of 5.0.

MII rate constant mole m-2 x 10+6

298

Temperature

308 318

CuII

k1 6.5 ± 0.8 5.6 ± 0.8 6.0 + 0.8

k2 2.4 + 0.4 2.1 ± 0.4 2.4 ± 0.4

CdII k1 1.1 + 0.2 1.3 + 0.2 1.1 ± 0.2

k2 0.59± 0.10 0.57± 0.10 0.67± 0.10

PbII

k1 1.8 ± 0.2 2.3 t 0.2 2.1 ± 0.2

k2 0.47± 0.2 0.51± 0.2 0.70± 0.2

on the reaction rate constants.

During the initial stage of activation in normal light and in the dark

the rate of uptake of CuII is approximately 5 times faster than that of CdII

and about 3 times faster than PbII. This difference is maintained during the

second stage (at least during the period from approximately 1.5 to 15 minutes),

where the rate of CuII uptake is about 4 times as fast as either PbII or CdII.

Within experimental error the uptake rate of PbII and CdII is the same during

the second stage of activation. The magnitude and rate of uptake of all of

the metals is virtually unaffected by carrying out experiments in the dark or

in normal light. Under UV irradiation, however, the rate and magnitude of PbII

uptake is greatly enhanced during the second stage of activation. The amount

and rate of CdII uptake are also increased, but to a smaller extent than for

Pb". There is an increase in the actual uptake of Cu", but no significant

effect on the rate in the time period 5 to 120 minutes. The results suggest

that some minor rate enhancement occurs during the early (1 to 5 minutes) part

of the second stage of activation for Cu".

Ibis comparison between the rates of uptake of CuII, CdII and PbII in-

dicates that

(i) for CuII, the rate determining step(s) is (are) apparently

faster than for CdII and PbII

and,(ii) W irradiation accelerates at least one rate determining step,

for PbII in particular and, to a lesser degree, for CdII and Cu".

(II) Evidence from kinetic studies for the presence of S° on the

surface of activated ZnS.

During the initial activation step the rate of CuII uptake at init-

ial pH values of 4.0 and 5.0 is the same, within experimental error (refer

Figures 4.5(a) ; 4.6(a) ; 4.9(a). However the rate of CuII uptake and the

concentration of zinc exchanged is less at an initial pH of 4.0 compared with

5.0 during the second stage of activation in both normal light, in the dark

(refer Figures 4.5(b) ; 4.6(b) ; 4.9(b) ; 4.11(b) and Tables 4.14(A,B,E,F) ;

4.15 (A & C) ; 4.16 ; 4.17) and under UV irradiation. In particular the CuII

uptake curve reaches a plateau during the second stage. Similar behaviour is

observed for PbII (Figure 6.5), but not for CdII (Figure 5.5).

The observed results for CuII and PbII are consistent with a reaction pro-

duct retarding both the penetration of either CuII or PbII into the ZnS lattice

at pH 4 and the release of zinc. S° has been detected on the surface of ZnS

activated with either CuII or PbII at initial pH values of 4 and 5. The quan-

tity of S° detected on the activated ZnS surface is 3 to 5 times greater at an

initial pH of 4 compared with pH 5. It is likely that this S° retards the

movement of ions across the solid-solution interface (S° is non-ionic and diff-

usion of ions will be severely restricted). Retardation is enhanced at pH 4

in comparison with pH 5 presumably because the S° is present in greater quantity,

reducing the reaction rate.

-218-

8.4(b) Correlation between MII uptake, quantity of S° detected, R values

and pH changes.

The level of MII uptake after 30 minutes of reaction time, quantity

of S° detected, R values and associated pH changes are shown in Table 8.7

for Cu", CdII and PbII. These results have previously been presented but

are grouped together here for the sake of clarity.

The results listed in Table 8.7 clearly show that S° is not present on

the surface of CdII activated ZnS until it is subjected to UV irradiation,

when both S° is formed and the actual uptake of CdII is increased. In the

absence of W irradiation the pH changes during activation and for the blank

are essentially the same (i.e. ApH = 0)and R equals unity. For CuII and PbII

a different picture emerges - S° is present and, in the absence of UV irradiat-

ion, the pH rise during activation exceeds that of the blank (i.e. apH is

positive). R is clearly dependent on the incident light conditions. UV irrad-

iation dramatically increases both the amount of S° formed and the actual PbII

uptake during PbII activation but has no detectable effect on the quantity of

S° formed during CuII activation of ZnS. It should be realized that not all

of the S° formed may have been detected i.e. some portion will probably be

present as H2Saq, HSaq or as oxidation products, as discussed in Section 8.1.

Note, however that H2Sgas and SO4aq were not detected (Chapter 3).

In summary, experimental results showing the

(i) lack of detectable S° on the surface of CdII activated ZnS and

observation that R equals unity (except under UV irradiation)

and, (ii) presence of S° on CuII and PbII activated ZnS surfaces, associated

R values and pH changes,

when coupled with the fact that exchange reactions (Section 8.3(b)) are thermo-

dynamically favourable, strongly suggests that the reaction of a heavy metal ion

MII with ZnS can occur as a 2 stage process:-

-219-

Ln s H

ro .F.1 cli ro

O O

H g)

Ô

b •Fi Ld b

0 O n

o data

R

range rati

normal

light

l

co ô

rn

ô

Cr • Itt 4)

û

o 0

rn

ô

00 co ô ô

ô •

ô H

Ld4.)id b

o O

rn •

ô

(d 4.1 of '0

O 0 n

o data

°

• I

ô I

(d .p

'0

CO

• 0 I

id b

Â no data

stem— pH b]

normal

light ("1

O

O

to t0

O O

O H

O O

N O

O

an H

O

In O O +

tfl qv O + -

0.19

to data

1. 8±0.2

rd 4) r1:$

H • ô

N+1

O

b -I-) (d

C; no data

S?

detect

!nS

x 10+6

normal

light

1.6±0. 1

not detectable

0. 9±0. 1

1.8±0.2

N

g0

1/40

0o

I'd -1-)

27

O O

al H

b 4-)

•L)

O O n

o data

take at 3(

g-1

ZnS )

norma l

lig ht

N V'

o

d' N

r•••

t0 H

H

H N

H

t0 H

H

N Ln

H

O N

I H

t0 O (--

t0 Cr

r

ro 4.) Ld b

O O

N H

• H

rt 4) b b

O O n

o data

1.28

H id -Hx +.) a - H

o

0

Ln

•

0

v:11

0

N

0

Lri 0

CI'

0

Ln

H H

I

H H

Û

H H

Û

H H

W

H H

arison between

take at 30 minu tes reaction time,

b

es durin

from the relevan t Tables in Chapters 4 to 7.

4-)

0 •r•I

H HÛ

• g •rI

H

û

• 0 •rl

H

â

x N

r.

Â v › ro N 4

[". O Ld •r1 4) 4) ro •r1 b ro a v

8 U 0

(i) an exchange reaction

and (ii) a redox process involving the formation of S° which is strongly

influenced by the nature of the heavy metal ion and may be acceler-

ated by UV irradiation. This redox process is linked with the

solid-state properties of the zinc sulphide.

Obviously this second stage may not occur, as in the case of Cd.II, unless

UV irradiation is used.

The solid state properties of the sulphides and the relationship with a

surface redox process now require examination.

8.5. SOLID STATE PROPERTIES AND SURFACE REDOX BEHAVIOUR.

8.5(a) General remarks.

Comments on the structure of semiconductors, electron transfer between

semiconductors and oxidants/reductants in solution and the effects of illumin-

ation on processes occurring at the semiconductor-solution interface have al-

ready been made in Chapter 2. Aspects of this semiconductor theory are

applicable to the mechanism of S° formation on ZnS surfaces. The following

discussion excludes any consideration of the role of hydrolysis species and

concentrates on acid pH values. Effects in near neutral and alkaline media

are discussed in Section 8.8.

Insight into probable mechanisms linking solid state properties and sur-

face redox processes may initially be obtained from the band theory of inorganic

sulphides.

The structure, bonding and properties of inorganic sulphides,using the band

model, have been admirably reviewed by Jellinek (20) and by Schmidt and Siebert

(16). For transition metals, the (n + 1)s and (n + 1)p orbitals of the metal

combine with the 35 and 3p orbitals of sulphur to produce a valence band

("bonding", mainly due to sulphur) and a conduction band ("antibonding", main-

ly due to the metal). The nd orbitals of the transition metal will also overlap

with the sulphur orbitals, but much less so than the (n + 1)s and (n + 1)p

orbitals of the metal. Narrow bands are formed, however the nd orbitals may

still be treated as localized on the metal. This simplification generally

holds for oxides, sulphides and selenides, but frequently fails for tellurides.

For a cation Mm+, if the energy of the oxidation states m and (m - 1) lie be-

low the top of the valence band, the cation will be partly reduced to M(m-1)+

holes will be created in the valence band and the compound will be a broad-

band, p-type metallic conductor. The d orbitals of the Mm+ and M(m-1)+

ions are still localized, however, and the magnetic properties exhibited by

these sulphides are due to these cations.

8.5(b) Cuii.

Jellinek maintains (20) that if the energy of the M(m-1)+ ion is low

enough, complete reduction of the Mm + cation will occur and that in sulphides

the d9 Cu2+ ion is reduced to the stable, diamagnetic Cu+ ion. During this

reduction process,electrons are extracted from the valence band, (mainly due

to sulphur), creating a corresponding number of holes and causing metallic con-

duction. The formation of a hole in the valence band corresponds to oxidation

of a sulphide ion to S. If the concentration of ions becomes too large, pair-

ing of the S ions to form (S2)2- ions occurs, resulting in structure deform-

ation. Jellinek (20) and Schmidt and Siebert (16) maintain that this pairing

process occurs with part of the S- ions in CuS, which is a diamagnetic metallic

conductor. Jellinek states that the structure of the mineral covellite (CuS)

reflects the fact that all copper ions are present as Cu+. Bragg (8) supports

this, indicating that the formula is best expressed as Cu3SS2, rather than CuS.

The kinetics of this conversion process are not understood at the present time,

a situation which unfortunately applies to sulphur reactions in general (48).

These documented properties of CuS of course refer to bulk, crystalline

specimens. The processes involved in this study are linked to the surface

properties of ZnS and it is this feature which now requires attention.

8.5(c) The band gap at the surface of ZnS - the effect of Cu", CdII and PbII.

The +2 oxidation states in zinc, cadmium and lead are stable, in con-

trast to copier, and the bulk sulphides are always divalent (16,20) e.g.

attempts to synthesise lead disulphide have to date been unsuccessful.

However Sj is formed on the surface of CuII and PbII activated ZnS in the

dark and substantially increased quantities are formed on PbII activated ZnS

under UV irradiation, whereas S° is formed on ZnS and CdII activated ZnS

surfaces only under W irradiation. Bearing in mind the results presented in

this study and, noting that there is no evidence showing that the formation of

metallic Cd or Pb occurs during activation (nor is it feasible according to

bulk thermodynamic calculations), electron transfer processes involving the

reduction of PbII and Cdaq to Pbs and Cds at the solid-solution inter-

face appear extremely unlikely. Some alteration of the surface properties of

ZnS would appear to be far more probable.

A recent interesting approach to this alteration of surface behaviour

has recently been suggested by Maust and Richardson (101). These workers pro-

pose that activation of ZnS occurs due to a semiconductor "doping effect" at

the surface. The process is essentially the same as for the well established

bulk doping effects (102,103). From this viewpoint Maust and Richardson pro-

pose that, for a particular metal ion to be an activator, it must introduce

acceptor-like surface states far enough down in the band gap to make it

energetically favourable for the following reaction to occur to some degree:-

Mn+

+ e ± M(n-1)+ Reaction 1 surface surface

Here Mn+

surface is the activator metal ion that has exchanged with zinc on the

surface. For this reaction to be energetically favourable, the acceptor surface

state must lie below the Fermi level in the solid. This reduction process for

a semiconductor like ZnS corresponds to the removal of an electron from the

valence band. i.e. oxidation of S2- occurs. For a given surface state lying

initially below the Fermi level, the band bending that accompanies localization

of electrons raises the energy level of the state toward the Fermi level. The

equilibrium occupation corresponds to the surface state lying at or very close

to the Fermi level; the upward band bending at that point corresponds to the

equilibrium surface potential. If an acceptor state lies above the Fermi

level, occupation is not energetically favourable and, in the absence of other

effects, the reaction specified above does not occur to any significant extent.

With this central premise in mind, Maust and Richardson have carried out

a simple Madelung-Born calculation which enabled them to estimate the surface

state energies of various metal ions substituted at zinc sites on 110 ZnS

surfaces. Their calculation assumed that the surface concentration of the

activator was very small ( 1 activator ion per 1000 zinc ions), thus the sur-

face Madelung constant remained unaltered. They also assumed that the effect-

ive charge on the activator ion was the same as the effective charge on the

zinc it replaced. Furthermore their calculations were performed only for the

semiconductor-vacuum interface. Applied to AgI, CuII, CdII and Pb", these

calculations resulted in the following sequence of surface state energies:-

AgI < CuII < CdII < PbII

low energy < > higher energy

(near valence band) (near conduction band, but below Fermi level)

Mn particular AgI and CuII were found to lie below the valence band of ZnS.

Maust and Richardson consider that AgI, CuII, CdII and PbII are all activators,

so that Reaction 1 above is energetically favourable. They further maintain

that " ... the formation of a recognizable surface film of some second phase

can be viewed as the eventual and inevitable later stage of continued sub-

stitution of a surface impurity into the sphalerite lattice (101)". The

assumptions used in their calculations require modification - in this study

it has been found that after 1 minute of reaction time generally a minimum of

1 in 6 (e.g. for CdII) zinc ions have been replaced by the activator metal ion,

hence substantial alteration of the surface Madelung constant is likely. In

turn this will alter the energy of the induced surface state. The energy,

Ea, of the latter relative to the conduction band edge is given by Maust and

Richardson as (101) :

Ea = - e Ia - I2) + Vm (Y-1)

where Ea is the activator surface energy, Ia is its ionization potential, e is

the electronic charge, I2 is the second ionization energy for zinc, Vm is the

bulk Madelung potential and y is the ratio of surface to bulk Madelung con-

stants. Clearly any alteration in the surface Madelung constant will alter y

and thus the pcsition of Ea relative to the conduction band. Maust and Rich-

ardson's calculation does not allow for the presence of substantial quantities

of activator ion at the surface (e.g. 1 monolayer in the case of Cu/I) nor

does it account for the fact that surface structural changes may well occur

e.g. similar to those occurring in bulk copper sulphide. Such changes should

affect the value of the activator surface energy. Moreover as activation pro-

ceeds the nature of the surface will change, and probably, the position of Ea

i.e. some time dependent function will be involved. The position of the sur-

face acceptor states therefore cannot be accurately calculated unless this

theory is appropriately modified to account for processes occurring at the

semiconductor-aqueous solution interface.

For "shallow" bulk acceptor impurities the simplest calculation of ion-

ization energy is based on the hydrogen atom model (22, 102, 103). By analogy

-225-

with the hydrogen atom, the energy levels of an acceptor in a crystal are

given by (103):

2n2me4

En n2h2Ke2

where Ke is the dielectric constant of the crystal. When calculating ioniz-

ation energies in a crystal, the electron mass m, in the above equation must

be replaced by the effective electron mass. This is necessary because in a

semiconductor interactions occur between electrons and the periodic potential

of the crystal, frequently resulting in an enlarged (compared with a free

electron) mass when the electron occupies a quantum state near the edge of a

band (102,103). For cubic ZnS Ke is 8 and the effective electron mass is 1.1

(102). Hence the ionization energy of hydrogen (13.6 eV) is reduced by 1.1

eV above the top of the conduction band. Thus electrons at the top of the

valence band require relatively little energy to transfer to an occupied accept-

or level. At normal temperatures there will be sufficient thermal energy avail-

able for this transfer to occur. This simple model works remarkably well for

Group III and IV impurities in Si, Ge and GaAs (22,102), but cannot readily

account for "deep levels" which, due to their larger ionization energies result

in a smaller degree of thermal ionization (i.e. valence band to acceptor level

electron transfer) than do shallow impurities at normal temperatures. Measure-

ments of acceptor energy levels are sparse for ZnS. Bulk acceptor impurities

such as silver and copper produce levels between 0.5 and 1 eV above the edge

of the valence band (101, 104). Apart from this there is little other evidence

available regarding bulk impurity levels in zinc sulphide.

When acceptor (or donor) impurities are added to zinc sulphide (or any

other semiconductor), the total negative charges (electrons and ionized accept-

ors) must be equal to the total positive charges (holes and ionized donors) in

82

to 0.23 eV for ZnS (22,102). i.e. the acceptor states lie approximately 0.23

order that electrical neutrality is preserved. From this starting point Sze

(102) and Wclfstirn (105) have shown that when acceptor impurities of concen-

tration NA (per cm3) are added to a semiconductor crystal, the number of

ionized acceptors per cm3, NA-, is given by

NA- = NA

1 + 1 exp EA- EF g kT

where EA is the energy of the acceptor level, EF is the Fermi energy and g is

the ground state degeneracy factor (4 for acceptor levels). Clearly, at a given

temperature, both the semiconductor conductivity and the position of the Fermi

level will depend on the concentration of acceptor impurity, as pointed out

both by Sze (12) and Azaroff (103). Shockley (106), in particular, has dev-

eloped elegant graphical techniques for determining the position of the Fermi

level in precisely specified semiconductors.

Aside from these considerations of impurity levels, it is highly likely

that the actual band gap at the ZnS surface is altered in the presence of a

heavy metal ion activator. This will occur because the structure of the crystal

surface where, say, 1 in 6 or more zinc ions has been replaced by CuII, CdII

or Pb11 is very different from the unperturbed ZnS lattice. Experimentally

determined band gaps for various sulphides are shown in Table 8.8. Despite

the range of experimental values it is clear that the band gaps for PbS and

Cu2S are substantially less than that for ZnS whilst that for CdS is of a

similar magnitude. Therefore at an activated ZnS surface it is reasonable to

expect that the band gap of ZnS is narrowed in the case of CuII and PbII act-

ivation and remains essentially the same in the case of CdII. The degree of

narrowing will, of course, depend on the extent of substitution of zinc ions

Minimum room-temp. band gap

(ref. 51)

Band gap at 300K (ref. 108)

Range Average

Sulphides Band gap (ref.109)

ZnS (sphalerite,cubic)

CdS (hawleyite,cubic)

PbS (galena,cubic)

Cu2S (chalcocite, hexa-gonal)

CuFeS2(chalcopyrite, tetragonal)

FeS2 (pyrite,cubic)

3.54

2.42

0.37

0.53

3.69-3.88

2.41-2.43

0.29-0.41

1.05-1.93

3.76

2.42

0.36

1.57

0.20

0.90

0.10

* no data available for "CuS" (covelline)

-227-

TABLE 8.8. Band gaps in eV for sulphides.

by the activator ions and will become more pronounced as activation proceeds.

Such an effect occurs in the complex sulphide chalcopyrite, CuFeS2, whose band

gap reportedly ranges from 0.53 to 0.90 eV, intermediate between that of Cu2S

(1.57 eV) and FeS2 (0.10 eV) i.e. as substantial quantities of iron are sub-

stituted for copper, narrowing of the band gap of Cu2S occurs. A similar

narrowing in band gap occurs in the series ZnS — CdS — PbS i.e. for a cubic

sulphide, the band gap diminishes with increasing atomic number as the degree

of covalent bonding and metallic character increases(107).

Despite the deficiencies in Maust and Richardson's (101) calculated sur-

face acceptor levels and the relative paucity of experimental data for both

surface acceptor states and bulk acceptor levels, the concepts of surface

acceptor levels and band gap narrowing are nevertheless directly applicable to

this study. On this basis a surface mechanism may now be advanced to account

for the formation of S0 on the surface of activated/irradiated ZnS. An overall

mechanism for the activation process is then presented in Sections 8.7 and

8.8.

8.5(d) The formation of S° by a surface state mechanism.

A schematic diagram relating to processes occurring at the ZnS-solution

interface according to a surface state mechanism is given in Figure 8.3. The

position of the valence band is considered to be constant, whilst the Fermi

level is deleted for the sake of simplicity (for the same reason band bending

effects are not shown). The positions of the conduction band and acceptor lev-

el, relative to the valence band, are drawn for the systems ZnS blank, CdII-ZnS,

PbII-ZnS and CuII-ZnS. The diagrams should be taken to correspond to 30 minutes

reaction time i.e. the time at which samples were taken during activation for

mass spectrometric analysis.

(I) ZnS blank.

For ZnS in the absence of an activating heavy metal ion, S° is not

detected on the surface until the ZnS is irradiated with light of a wavelength

sufficient to promote electrons from the valence to the conduction band, result-

ing in the production of Zn22++ and Surface , as proposed in the stimulating

paper by Williams (63):

ZnS 4hv

) zn2+ + S° + 2e Reaction 2. aq surface

Reported band gaps for ZnS near room temperature range from 3.54 to 3.88 eV,

as may be seen in Table 8.8, thus W irradiation in the vicinity of 350

to 320 nm or less will cause reaction 2 to occur. The fact that S° has been

detected on the surface of UV irradiated ZnS in the present study is consistent

with this mechanism. Similar irradiation enhanced dissolution processes for

solids such as ZnO have been advanced by Gerischer (59), Morrison (56) and

-229—

(a) ZnS

solution

(i) DARK

solid

(ii) NORMAL LIGHT

solid solution

no detectable reaction

E (iii) UV IRRADIATION

solid solution

Zn2+ aq

ZnS h 2n 2+ +S + 2e

—

conduction band Zn2+ n2+

valence band

no detectable reaction

(b) CdII/ZnS

( + weak UV as2for ZnS; Cu2s replaces Zn

s FIGURE 8.3 Schematic diagram showing the formation of S° by a surface state mechanism

("S" denotes surface, "kT" denotes thermal energy)

Cu2s +a Cu

s

E

kT 1S +Y+SS 4kS

acceptor level

vb

r

cb

acceptor level

vb

(c) PbII/ZnS E E

(d) CuII/ZnS E

Pb S+ kT Pbs

75S S kT 14S.. + e

(+UV as for ZnS; Pb2s replaces Zn2s)

n2+ Zn2+ n2+

aq Pb2+

Pb s + e kT Pbs

XS2S kT ygs + e

E

b + e kT PbS

i4SS kT7154SS + e —

(weak h$— see UV)

E

cb acceptor level

vb

Zn2+ Zn2+

// no detectable reaction no detectable reaction

E

weak CdS +eÉ Cd S

i4SS4 14SS +e

(+ UV as for ZnS; Cd21.

replaces Zn2S)

Cu2s + e kT Cus

É 34S 2S _kT 34SS + e

F--

Cu2s + e kT Cus

'ASS kT —> S + e

(weak hti— see UV)

others (57, 58).

Pawlek (61) and Wadsworth (62) suggest that the following alternative

reaction occurs:

ZnS + 2Hâ f h > Znâ+ + H2Saq Reaction 3 q q

Pawlek (61) further identified sulphate and elemental sulphur in significant

quantities when ZnS was irradiated in acid solutions, but made no detailed inter-

pretation of his findings. with respect to the solid state properties of the ZnS.

Certainly large pH increases are observed when ZnS is irradiated in an aqueous

system (e.g. refer Table 4.17), however this need not mean that consumption of

Haq occurs according to Reaction 3. As has already been noted in Section 8.1

the S° formed is unstable - decomposition to H2Saq and HSaq will also consume

Ha q. Mass balances performed on zinc released compared with Haq consumed at

acid pH values reveal that the consumption of Ha+q stoichiometrically accounts

for not more than half the quantity of zinc released. Even allowing for buffer-

ing effects occurring in the H2S aq /HS aq

system, Reaction 3 cannot be confirmed

from the results of this present study. The presence of Ssurface on W

irradiated ZnS,as determined in this study, together with established behav-

iour of similar solids suggests that Reaction 2 is the most probable process.

The concentration of zinc found in solution during blank experiments per-

formed in normal light and in the dark (for the same ZnS sample and concentrat-

ion) is the same, within experimental error. Some UV radiation will be present

in normal light,but will be of extremely low'intensity:,so that Reaction 2 will

not occur to any detectable extent. Furthermore at normal temperatures (i.e.

around 298 K) there is insufficient thermal energy (denoted by kT in Figure 8.3)

available to promote electrons from the valence to the conduction band to any

appreciable extent.

(II) CdII-ZnS.

° i S is not detected on the surface of CdII activated ZnS until it

is subjected to UV irradiation. This suggests that the surface acceptor state

introduced by cd2+ is above both the valence band the the Fermi level so that

Reaction 1 is not energetically favourable for CdII-ZnS in the dark and in

normal light i.e.

2+ + Cd surface + e 4 surface

does not occur to any appreciable extent. In view of the similarly between

the band gaps of ZnS and CdS (refer Table 8.8), only a slight narrowing of the

ZnS band gap is expected to occur during activation by CdII. Hence there is

both insufficient radiation of the required intensity in normal light and not

enough thermal energy available to allow Reaction 2 to occur to any detectable

extent. This is demonstrated both by the fact that the exchange ratio is unity

and that S° is not detectable.

Under UV irradiation, Reaction 2 occurs and S0 is detected on the surface

of CdII activated, UV irradiated ZnS. Correspondingly enhanced uptake of CdII

is observed as Cd2+ ions replace the Zn2+ ions, produced according to Reaction

2, on the surface of the ZnS crystal lattice. Under UV irradiation it is also

probable that the following reactions occur to a small extent:

2+ by

Cd2+ + e --w surface Reaction 4

I S2 + e Reaction 5 surface surface

i.e. S° forms according to Reactions 2 and 5.

(III ) PbII-ZnS.

S0 is detected on the surface of PbII activated ZnS in the dark

and the exchange ratio, R, equals unity as is shown by the results in Tables

6.9(A), 7.7 and 8.7. This indicates that the surface acceptor state intro-

duced by Pb2+ is located close to the valence band and below the Fermi level,

so that thermal energy (kT) is sufficient to promote electrons from the val-

ence band to the surface acceptor state to an appreciable extent according to

the following reactions:-

2+ kT + Pb Pb surface + e {=---r surface Reaction 6

4Ssurface -<--Ssurface + e Reaction 7

On the basis of the respective band gaps for ZnS and PbS, narrowing of the ZnS

band gap is expected as Pb2+ substitutes for ZnS in the crystal lattice. How-

ever the band gap of PbII activated ZnS is sufficiently great so that thermal

excitation of electrons from the valence band to the conduction band will not

occur to any appreciable extent in the dark.

The quantity of S° formed on PbII activated ZnS surfaces increases sub-

stantially under UV irradiation. The results in Chapter 7 show that when ZnS

is subjected to UV irradiation (pH = 4.3), S° is detected on the surface at a

level of 5 ± 1 x 10-7 mole S° per g ZnS. For PbII activated ZnS at pH 6.6, the

level of S° detected under UV irradiation rises to 13 ± 2 x 10-7 mole S° per g

ZnS, whereas in the dark and in normal light at pH 5.3 to 5.4 S° is detected

at a level of 3 ± 1 x 10-7 mole S° per g ZnS i.e. enhanced S° formation occurs

in the presence of PbII and UV irradiation and cannot be accounted for simply

by adding the contributions resulting from Reactions 2 and 7. The variation

in pH simply stresses this difference in S° levels. However under UV irradiation

a large increase in PbII uptake occurs (Figure 8.2) so that, in reaction 7

more surface acceptors per unit volume are available for S0 formation through

Reaction 7. For PbII activation under UV irradiation, then, it appears that

Reaction 2 occurs, leading to S° formation and the production of Zn2+ . Pb2+ aq

ions replace Zn2+ in the crystal lattice, leading to significantly more surface

acceptors per unit volume (100 to 200% more) than in the dark and in normal

light (refer Table 8.7). An increase in surface acceptors per unit volume leads

to an increase in the number of ionized acceptors per unit volume as may be

seen from

NA- = NA

1+ g exp EA-EF

kT

assuming EA, EF and T are constant. In turn this will result in an increased

quantity of S° being formed through reactions 6 and 7. The increase in the

quantity of S° detected together with the enhanced PbII uptake and zinc release

support this mechanism.

In normal light at initial pH values of 4.0 and 5.0 the exchange ratio, R,

varies with ZnS concentration and initial PbII concentration. At the highest II [Pb ]initial ratios R is generally greater than unity after a reaction time of

[ZnS]

5 minutes and then becomes less than unity from about 30 minutes onwards (Tables

6.6(A), 6.8(A)). This effect does not occur in the dark - R is constant (unity)

with time (Table 6.9(A)). At lower [PbII

]initial ratios (Tables 6.6 (B,C)) (ZnS]

R is less than unity and either remains essentially constant, or decreases

slightly, with time. S° is detected on the surface of PbII activated ZnS in

normal light and, within experimental error,is the same as that found in the dark.

These results indicate that very low intensity radiation from normal light

does not contribute to any detectable extent to reactions 6 and 7 i.e. thermal

excitation is responsible for So formation. As Pb2+ substitutes for Zn2+ in

the crystal lattice, sufficient band narrowing occurs for reaction 2 to occur

to a limited extent. For example if the band gap in ZnS was reduced to 2 eV,

light of wavelength 620 nm or less would cause reaction 2 to occur. The extent

to which it occurs will depend on the intensity of light of the specified wave-

length. The slightly enhanced PbII uptake and zinc releases values (compare

Table 6.6(A) with 6.9(A)) for experiments performed in normal light compared

with those carried out in the dark supports the conclusion that band narrowing

occurs, leading to a weak reaction 2 with Pb2+ ions replacing Zn2+ in the

crystal lattice. The additional amount of So produced in this process is not

detectable (i.e < 1.6 x 10-7 mole S° per g ZnS).

At high [PbII initial ratios normal light has a subtle effect on PbII

[ZnS] uptake - it is enhanced, compared with the quantity of zinc released and R is

greater than unity. The effect occurs at initial pH values of 4.0 and 5.0,

but does not occur in the dark nor under UV irradiation. Normal light appar-

antly enhances an uptake step for PbII during the early stage of reaction -

as the reaction proceeds, band narrowing occurs and reaction 2 dominates this

effect, resulting in R values less than unity. Whilst this enhanced PbII

uptake is a minor effect it is interesting. The present results unfortunately

do not allow any further elucidation of the mechanism involved.

(IV) CuII-ZnS.

S° is detected on the surface of CuII activated ZnS and, within ex-

perimental error, the amount detected is independent of the incident light

conditions. This indicates that the surface acceptor state introduced by CuII

is located close to the valence band and below the Fermi level so that thermal

energy (kT) promotes electrons from the valence band to the surface acceptor

state to an appreciable extent according to the following reactions:

-235-

2+ kT + Cusurface + eCusurface Reaction 8

245surface ' S

sur face + e

Reaction 9

In the dark R equals unity. Certainly narrowing of the ZnS band gap

will occur as Cut+ replaces Zn2+ in the crystal lattice, however the band gap

is sufficiently large so that thermal excitation of electrons from the valence

band to the conduction band will not occur to any appreciable extent in the

dark i.e. reaction 2 does not occur.

In normal light R is consistently less than unity (0.85). Narrowing of

the ZnS band gap occurs, allowing reaction 2 to proceed to a very small extent

and releasing a small amount of zinc into solution. The total amount of zinc

released during activation experiments performed in normal light is very slightly

greater than in the case of experiments performed in the dark (Tables 4.15,

4.16), supporting this mechanism. Any additional S formed in normal light is

undetectable (i.e. < 1.6 x 10-7 mole S° per g ZnS).

Under UV irradiation only slightly enhanced zinc release is observed com-

pared with experiments performed in normal light and in the dark (Tables 4.15,

4.16, 4.17). Any additional S° formed is undetectable (i.e. < 1.6 x 10-7 mole

S° per g ZnS). After 1 minute of reaction time, CuII uptake is some 5 times

greater than for CdII or PbII and is in excess of monolayer coverage. It is

likely that this "film" of "CuS" itself absorbs at least some of the UV radiat-

ion, reducing the effective intensity and, at the same time impedes both further

uptake of CuII and release of zinc from the crystal lattice., A similar process

may be envisaged for both CdII and PbII. Figure 8.2(e) in particular shows

that PbII and CdII uptake level off with time while the Tables in Chapters 4,

5 and 6 show that zinc release is impeded during metal ion activation of UV

irradiated ZnS. In contrast, Figure 8.4 shows that zinc release occurs linearly

-236-

with time for unactivated, irradiated ZnS.

Cull uptake varies during activation of ZnS under varying light conditions

at acid initial pH values (e.g. Tables 4.15, 4.16, 4.17). The variations are

generally slight i.e. ±5%, which is within experimental error and are in accord

with both the similar small,or negligible, effects observed for zinc release

and S° formation discussed above and hence with the proposed mechanism for S°

formation.

(V) General remarks.

(i) It is pertinent to note that the proposed mechanism for S° form-

ation on Cull activated ZnS surfaces is consistent with the ESCA observat-

ions of Clifford et al (69, 78), who demonstrated that Cu' and S° are

present on such activated surfaces. The results of Storey and Platt (54),

showing that Cull is present are also in accord with the suggested mechan-

ism which is, in addition, consistent with the established structural

properties of bulk "CuS", reviewed by Jellinek (20) and discussed in Sec-

tion 8.5(b), although not all of the Cut+ is apparently reduced to the Cu+

state, according to the results of the present study. In bulk "CuS" pair-

ing of S ions occurs to form (S2)2- and structural deformation occurs due

to strain resulting from different S-S distances e.g. 2.76 Â in sphalerite,

2.05 R in the S22- units in covellite (8). Certainly structural strain

will occur on the activated surfaces examined in this study. At the very

least this will occur due to the different ionic radii of the substituted

metal ions causing strain i.e. Zn2+ = 0.74 R, Cut+ = 0.69 R , Cu+ =

0.96 R, Cd2+ = 0.97 R, Pb2+ = 1.20 R (51,132). It is not possible to

say whether or not pairing of S to form (S2-) occurs during Cull activat-

ion. This will probably only occur for a "distinguishable" film (see below).

The situation will be further complicated in systems where S° is present.

-237-

Sg+ was detected by the mass spectrometer-and, since recombination pro-

processes are highly improbable (83), this result indicates that S° was

present on the surface of ZnS as Sg rings which, upon evaporation into

the ionization chamber of the mass spectrometer, were fragmented by the

electron beam into mainly S2+ species. The disposition of these S° rings

upon an activated and/or irradiated surface is difficult to determine.

The "tailing" observed in S° determinations on activated and/or irradiated

ZnS surfaces, a feature which was absent in the standards, indicates that

the S° detected on the former surfaces is more firmly held. Beyond this

no further conclusions can be drawn with any certainty.

(ii) From the stoichiometric relationships between reactions8 and 9

and the MII uptake and S° concentrations per g of ZnS given in Table 8.7,

it appears that between 50 and 60% of the Cu2+ taken up by the ZnS is

present as Cu+ at pH 4 and is independent of the incidént light conditions.

For PbII at pH 4, in normal light a similar calculation shows that Pb'

has been entirely converted to Pb+. Both calculations assume, of course

that all of the S° formed on the ZnS surface is detected. At higher pH

values S° becomes increasingly unstable and calculation of the degree of

conversion of M2+ to M+ becomes uncertain.

Although these calculations are approximate the results imply that the

surface acceptor state introduced by Pb2+ into ZnS is much closer to the

valence band than is the corresponding state for Cu2+. This is in agree-

ment with the reported band gaps of 1.57 eV for Cu2S and 0.37 eV for PbS

i.e. the end product of continued substitution of M2+ into the ZnS lattice

is the formation of a "distinguishable" MS film, so that the initial sur-

face acceptor state - valence band difference eventually becomes the

band gap in the "distinguishable" MS film (N.B. "distinguishable" is

taken to mean "detectable" by XRD, ESR etc.). Note that the S° formed on

-238-

the surface of CuII activated ZnS at pH 0.5 (Table 7.4) is twice that

expected if all of the available CuII added was abstracted by the ZnS.

At this low pH, S° forms

(i) from ZnS (see Figure 8.1),

(ii) during CuII activation

and, (iii) probably, from decomposition of

Even at pH 2, the results shown

(c) demonstrate that the "CuS"

releasing copper to form Cua2++

the "CuS" film.

in Figure 4.12

film is unstable

so that at 120

minutes the CuII uptake is 40% of the value at

5 minutes. This behaviour is consistent with

the expected increase in solubility of sulphides

at low pH (e.g. Figure 2.1 for ZnS).

(iii) The results obtained in the present study indicate that, under

the conditions used, narrowing of the ZnS band gap at the solid-solution

interface occurs in the order

Cd2+ > Cu2+, Pb2+

When a "distinguishable" film is formed, the band gap at the surface should

resemble those of CdS, PbS and, probably,Cu2S.

The surface acceptor states introduced by Cu2+, Cd2+ and Pb2+ increase

in the order

Pb2+ < cu2+ « Cd 2+

[

low energy) [high energy

below Fermi level) above Fermi level

which is in disagreement with the order proposed by Maust and Richardson

(101), stressing the approximate nature of their calculations.

(iv) Differences in the relative rates of uptake of PbII and CdII under

UV irradiation are probably due to the fact that the wavelength of the

radiation from the UV lamp, necessary to cause reaction 2 to occur, is

different for PbII and CdII and has a higher intensity in the former case.

2- surface

surface + e 0

(v) It is relevant to note here that Sato (100) has detected, in his

electron diffraction study of the activation of ZnS by AgI, that some of

the abstracted silver is present as metallic silver. Further Gaudin,

Fuerstenau and Turkanis (98) have observed that subsequent deactivation

of such AgI activated ZnS with cyanide is favoured by the presence of an

oxygen atmosphere. These observations suggest that, at the surface of AgI

activated ZnS, reactions of the following form occur:-

+ o Ag surface + e

Agsurface

i.e. Ag+ introduces a low energy acceptor surface state into ZnS, resulting

in the formation of some metallic Ag and So. In the presence of a silver-

complexing agent like CN Agâq

may readily be controlled and hence Agsurface'

The observation by Gaudin et al (98) that an oxygen atmosphere was necessary

for deactivation is consistent with the known fact that So will oxidize to

hydrogen-sulphur-oxygen speciés in the presence of oxygen, particularly if

the latter is in excess., as in their study.

8.6. pH CHANGES.

8.6(a) Blank experiments.

At acid pH values a pH increase is observed when ZnS is placed into a

10-1 M KNO3, N2 flushed system at an initial pH values of 4.0 or 5.0. The mag-

nitude of this pH increase is dependent on the ZnS concentration and on the

incident light conditions.

In the dark and in normal light the increase in pH is due to:-

(i) adsorption of H at the ZnS-H2O interface,

-240-

(ii) probable adsorption of H by,and reaction,with soluble

ZnS oxidation products present on "fresh" ZnS.

(iii) production of OH due to the release of Sâq followed

by the formation of HS , H2S aq aq

i.e. Sâq + H2O2, -4-- HSaq + OHag etc.

Such pH changes are well established at oxide and sulphide-water interfaces (5,

9, 40). In normal light the pH increases are the same as those in the dark, with-

in experimental error, whereas massive pH increases are observed under UV irrad-

iation. As previously discussed, UV irradiation of ZnS causes Zn2+ and aq

Ssurface to be formed according to reaction 2.

hv 2+ o Ssolid aq + Ssurface + 2e Reaction 2

This reaction must be accompanied by a corresponding reduction process. The

most probable reactions are listed in Table 8.9.

TABLE 8.9

Reduction reactions.

Reaction Standard reduction potential, volt Reference Comment

I Haq

+ e 4- 4112 gas 0.00 67,68 suggested by tierischer for irradiation of Zn0 (59a)

II H201 +e F OHaq +4H2g -0.83 68 suggested by Dixon et al (43) for UV irradiation of ZnS.

III 3 02g+2Haq+2e- 4 H202, +1.23 67,68 [02 g]and [02aq ] very low.

IV 4-02g+H20R+2e 20Haq +0.40 67,68 [O2g] and [02aq] very low.

i.e.

-241-

These standard reduction potentials are obtained for reactions at a platinum

electrode and may well be different at a ZnS surface (Section 8.1 and reference

63). Furthermore the actual half-cell reduction potential, determined by the

relevant activities and fugacities, should be used in assessing which reaction

is most likely to occur.

All of these reactions in Table 8.9 will result in a pH increase when coupled

with reaction 2. III and IV are unlikely to be major contributors to the observed

pH increase due to the trace (10-6 M maximum) concentrations of dissolved 02

present. For this reason additional reduction reactions involving the format-

ion of oxy-anions or hydrogen peroxide from oxygen are not considered (42). I

and/or II appear most likely. Buffering effects in the H2Saq/HSaq system

probably obscure any stoichiometric link between Znâq formation and pH increase.

Unequivocal identification of the particular reaction(s) responsible for the

observed pH increase requires further work.

At neutral and alkaline pH values the pH of ZnS "blanks" normally decreases

during activation. These changes are generally quite small in terms of the net

consumption of OHaq (or production of Hâq) and are probably due to:

(i) adsorption of OH at the ZnS-H20 interface,

(ii) adsorption of OH by and reaction with soluble ZnS

oxidation products present on "fresh" ZnS.

Hydrolysis effects will also be important. These are discussed in Section 8.8.

8.6(b) Activation experiments.

During activation with Cu", CdII and PbII, pH changes are also observed.

In the case of CdII these are generally the same as for the blank and may be

attributed to the causes outlined in Section 8.6(a). For CuII and PbII activat-

ion a proportion of the pH change is undoubtedly due to the same reactions as

for the blank, however excess Haq consumption (or its equivalent) is generally

observed. A sample calculation for a CuII activation experiment performed in

-242-

the dark serves to illustrate the point (data from Table 4.16A).

dark (PHinitial = 5 0)

activation blank

pH at 120 mins 6.4 5.3

mole H+ consumed at 0.96 x 10-5 0.50 x 10-5

120 rains.

Even in the dark an excess of 0.46 x 10-5 mole of H+ is consumed during act-

ivation by Cu". Similar results are obtained for Pb".

These pH increases observed for CuII and PbII activation in the dark, at

least at acid initial pH values, are probably associated with decomposition of

S° to form H2S aq and HS aq

(equations 10, 11 in Table 8.1). In normal light

and under ûV irradiation the additional pH increases observed (i.e. above those

in the dark) will be due to reaction 2 coupling with I and II in Table 8.9.

Hydrolysis effects, discussed in detail in Section 8.8 , will complicate these

changes in neutral and alkaline solutions.

8.7. OVERALL MECHANISM AT ACID pH VALUES.

(N.B.the summaries in Tables 8.11 to 8.13 give a perspective view of the entire mechanism as a function of pH).

8.7(a) Exchange reaction.

Each of the heavy metal ion activators undergoes an exchange reaction with

ZnS. At acid pH values up to pH 6.5, M22q+

is clearly the dominant solution

species, as may be seen from Figures 4.1, 5.1 and 6.1. Several possible mech-

anisms may be envisaged (s denotes surface):-

(i) adsorption of M2+ , exchange between Zn2+ and M2+

aq lattice adsorbed with formation of MS on the ZnS surface then desorptiôn of

Zn2+ to form Zn+ . aq

M2++ + Zn Ss (Zn,M)Ss + Zn2+ Reaction 10 aq

(Note: (Zn,M)Ss denotes M2+ substituted for Zn2+ on the ZnS surface)

(ii) dissolution of ZnS to form Zn2+ and S2 aq aq

followed by precipitation of MS on to the

ZnS surface

M2+ + ZnS —} M2+ + ZnS + Zn2+ + S2 --4-ZnS • MS + Zn2+ ...Reaction 11 aq s aq s aq ag {-- s s aq

(Note: MSs denotes a separate, precipitated phase)

Similar mechanistic steps were tentatively proposed by Finkelstein and Allison

in a recent review (55).

The surface redox behaviour discussed in detail in Section 8.5 i.e. the

formation of surface acceptor states and band gap narrowing of ZnS, strongly

suggests that Reaction 10 occurs in preference to Reaction 11. Substitution of

M2+ for Zn2+, rather than precipitation of MS on to the ZnS surface, is consis-

tent with the mechanisms discussed in Section 8.5.

8.7(b) Surface redox behaviour.

This has been discussed in detail in Section 8.5.

8.7(c) Model for the logarithmic law.

(I) Initial stage of activation.

The exchange reaction between a heavy metal ion, Ma2++ and the ZnS sur-

face is represented by reaction 10 as:

M2+ + ZnS (Zn,M) S + Zn2+

aq — s aq

The following reaction scheme may be envisaged, in the absence of hydrolysis

steps ("s" denotes surface):

(i) diffusion of M2+ to the ZnS surface (fast) aq

fas M2+ M2+

aq aq,s Step (1)

(ii) adsorp ation of Mt+ on to a surface site, M2+ q.s ads

(fast)

(i.e. adsorbed site)

M fast} M2+ aq,s ads

Step (2)

(iii) a slow surface step involving transfer of M2+ from an adsorbed site,

yds. to a lattice site, Mlittice

(slow rate determining).

M2+ lattice 2+ ads slow DS Mlattice

Step (3)

(iv) transfer of Zn2+ from a lattice site to the ZnS surface (fast)

2+ fast 2+ lattice Zns Step (4)

(v) desorption of Zn2+ to form Zn2+ (fast) aq,s

`n2+ fast Zn2+ s aq,s Step (5)

(vi) diffusion of Zn2 away from the ZnS surface (fast) aq,s

Zn 2+ fast Zn 2+

aq,s aq

Step (6)

i.e. Step 3 is the slow, rate determining step in this overall mechanism.

The justification for choosing Step 3 as the slow step may now be discussed.

The solution diffusion steps (1 and 6) are not rate detemining as the ex-

perimental evidence has clearly demonstrated. Furthermore if either Step 4 or

5 were rate controlling, the same rate of uptake would occur for Cuis, CdII

and

Pb

r

II i.e. dt would be independent of the nature of the activator ion, in

marked contrast to the observed behaviour. Step 2 may also be ruled out as a

rate determining step for the following reasons:

(i) The standard enthalpies of hydration (zHoh) at 298 K for the

three metal ions studied in this investigation are (47):

AHoh kJ mole-1

Cut+ -2092

Cd2+ -1798

Pb2+ -1464

i.e. the hydration enthalpy is rather insensitive to the nature of the

cation and does not correlate with the significant differences in dl' dt

for copper, cadmium and zinc.

(ii) Bockris has investigated the rate of charge transfer reactions

at electrode-water interfaces in the presence and absence of electric

fields (115). For an ion "jumping" from a solution site to a site on the

metal surface

i.e. solution site -} site on metal surface

an important feature which emerges from the theory is that the reaction

rate is proportional to the concentration of the ions on the solution side

of the interface. In fact a first order dependence (coupled with other

terms) on ion concentration is observed, a feature which is not in evidence

in the present study.

The observed reaction rates are clearly dependent on the nature of the

metal ion, hence Step 3 appears to be the most likely rate-determining

step.

Suppose that a potential barrier exists for the transfer of Mads

lattice site (i.e. M2+ lattice and that this potential barrier is much larger than

for successive movements between lattice sites, so that diffusion in the sulphide

is not rate determining. The situation is then directly analogous to the oxidat-

ion of aluminium, where the rate determining step is the transfer at the metal-

oxide interface of Al into interstitial positions in the oxide as A13+ ions

to a

(22, 27, 28, 109). A very thin oxide film is formed and the logarithmic law

is obeyed.

The probability that an ion jumps from its adsorbed site (M2+ ) to a ads

lattice site is

v exp()

where v is the ionic vibration frequency, assumed to be the same at both the

adsorbed and lattice sites and AE is the barrier height or activation energy

(k is the Boltzmann constant, T the absolute temperature). This expression

was initially proposed by Mott (28) and elaborated on by others (110,116) for

the formation of protective oxide films on metals (particularly aluminium).

By analogy with the growth rate of oxide films, the rate of uptake of MII

is given by

dr -AE dt - Ay exp

here A is a constant.

Further, if the activation energy increases with increasing coverage or uptake

of MII, according to

DE = b1NoI' (3)

where bl is a constant, r is the uptake of MII in mole g-1 and No is Avogadro's

constant, then

(1)

(2)

dI' _ dt

Ay exp -b1NoI'

kT

(4)

Many gases exhibit an activation energy which increases with surface. coverage ac-

cording to equation 3(96).Activator ions, already present in lattice sites should

similarly impede incoming ions.

In the initial activation step, MII uptake is generally equal to, or less

than monolayer coverage. Space charge effects, which frequently enhance transfer

rates through oxide films (i.e. a few monolayers or more) may be neglected

(22, 28) .

From (4) ,

exp blNorldr = Av dt (5)

l kT )I

Integrating over the limits (t = 0, r = 0) and (t = t, r = r)

r t i.e. J exp bl dr = Av j dt

° kT } o

gives

b 11 (t

exp l

r kT oQrl1 [AvtJ

0

exp [biNor) = Avt + kT

kT b 1 No

taking natural logarithms yields

kT b 1 No

kT b1No

(6)

(7)

(8)

r = kT 2nt■ + kT b1No b1No

Zn biNoAv

kT

(9)

or

F= 2.303 kT log10t + 2.303 kT log10 (b1NoAv b1No b1No

II kT

which reduces to

r= kl logt + r1

where kl = 2.303 kT b1N0

and r1 = 2.303 kT log10 b1NoAv are constants at b1No tl kT

-248-

constant temperature for a given metal ion activator. Equation 11 is the

experimentally determined rate law shown in Chapters 4,5 and 6.

The MII uptake rates during the initial step increase according to the

following sequence

CdII < PbII < CuII

Moelwyn-Hughes (33) indicates that the frequency of vibration, vi, of an ion

of mass mi in a crystal is given by

vi = 2ffao pmi

where e is the electronic charge, ao is the equilibrium ionic distance, a is

the Madelung constant, and p is an empirical constant which varies by only

about 5 to 10%, on average, for crystals with a common anion, varying cation

(33). Generally variations in vi for a range of cations in crystalline hal-

ides, for example, are small (33), so that any variation in vi is unlikely to

account for the differences in the observed rates of uptake of CuII, CdII and

PbII on ZnS to any significant extent.

The observed differences in metal ion uptake strongly suggest that the

value of the activation energy, DE, determinines the reaction rate, decreasing

in the following sequence:

DECd2+ > AEPb2+ > AECu2+

This decrease in EE parallels the decrease in solubility for the sulphides which,

as discussed in Section 2.3, is determined by the properties of the d orbitals

possessed by the metal cation. The decrease in AE is also accompanied by an in-

crease in the electronegativity of the metal atom:Cu = 1.9, Pb = 1.8 , Cd = 1.7

(51,132). (Sulphur has an electronegativity of 2.5).

It is therefore highly likely that AE is smallest for Cu2+ because the

attraction exerted by the sulphide ions is greater than for Cd2+ and Pb2+,

(12)

causing a faster transition from M2+ to to M2+ttice'

Second stage of activation.

CdII uptake follows a logarithmic dependence on time until about 15% mono-

layer coverage then the rate constant, k1, alters to a new value,

logarithmic dependence continues over the 2 hour reaction period.

iour suggests that there is a critical surface coverage where the

energy increases sharply i.e. a critical quantity of CdII causes

activation energy and equation 3 becomes

= b2NoI'

k2, and the

This behav-

activation

an increased

(12)

where b2 > bl, thus k2 = 2.303 kT and k2 < kl, as observed experimentally. b2N0

A similar effect occurs for CuII during the reaction period 1.5 to 15 minutes.

For both CuII and PbII there is no clear rate law applicable over the en-

tire second stage of reaction. In both cases at acid initial pH values the

formation of a protective "film" of S0 on the surface retards further uptake of

either Cu2+ of Pb2+ (i.e. LE is very large), effectively leading to a limiting

uptake or "film thickness" at pH 4.

Further comments:

(i) It is unnecessary to invoke the "flat gaps" or "pore blocking"

mechaniGms proposed by Evans (23). The experimental results can be ade-

quately interpreted in terms of activation energies and the formation of

a protective film.

(ii) Activation energies for thin film oxidation processes where the rate

determining step involves a "jump" from a lattice site to an interstitial

position are large at room temperature (e.g. about 174 kJ mol-1 for alum-

inium, according to Mott (28)). Similarly the slow step 3 specified in

-250-

this proposed mechanism would be expected to have a large activation

energy, increasing with surface coverage. This is in accord with the ob-

servations that temperature changes in the range 298 to 318 K had no

detectable effect on the experimental rate constants.

(iii) On the basis of this model for the exchange reaction, the sur-

face redox processes outlined in Section 8.5 take place after the activator

ion reaches a lattice site i.e. M2+ The electronic transitions

lattice ce

caused by thermal energy are expected to be rapid. It should be realized

that this proposed model does not exclude the possibility that simultaneous

electron transfer may occur along with the exchange reaction, rather than

as a subsequent fast process. In the event that this occurs, some modific-

ation of the proposed mechanism will be necessary. The form of the rate

law should, however, remain substantially the same.

(iv) For the case of PbII activation, the mechanism outlined above is

generally in accord with the experimental, logarithmic dependence of Pb2+

uptake on time during the initial activation step. The slight concentrat-

ion dependence and varying kl values observed suggests that AE decreases

slightly with increasing Pb2+ concentration. aq

8.7(d) Magnitude of heavy metal ion uptake in other studies.

The literature regarding previous activation studies has already been

dealt with in Chapter 2.

The CuII and MII uptakes determined in the present study, expressed in terms

of monolayer coverage for synthetic, high purity, cubic ZnS are generally less

than those reported by other workers (36 to 39, 97, 124).

Generally natural sphalerite samples were used in these other studies.

Such samples commonly contain pores, cracks and grain boundary defects which may

lead to enhanced MII uptake. The criticisms made in Chapter 2 regarding the use

of very low surface areas and poor experimental precision are also valid here.

Furthermore natural samples of sphalerite frequently contain substantial

quantities of impurities (e.g. Fe in marmatite) which,if the incident light

conditions are not controlled, will probably lead to reaction 2 taking place

in normal light, due to narrowing of the ZnS band gap. Enhanced zinc release

will occur and will be accompanied by enhanced MII uptake, as discussed in

Section 8.5.

8.8 OVERALL MECHANISM AT NEAR NEUTRAL TO ALKALINE pH VALUES.

In neutral and alkaline pH regions, the formation of metal ion hydrolysis

species will complicate the reaction scheme outlined in Section 8.7. In addit-

ion modifications will be necessary due to the demonstrated instability of S°

in an alkaline environment (Table 7.6, Figure 8.1).

8.8(a) Blank experiments.

It is pertinent to comment briefly again on the effect of UV irradiation

on ZnS dissolution. The rate of dissolution is depressed at alkaline pH values,

an effect first detected by Dixon (42). Results obtained in this study are

shown in Figure 8.4.

20

10

-252-

0

30 60 90

120

TIME (MINUTES)

FIGURE 8.4 Znirrelease from ZnS under UV irradiation as a function of initial pH.

Initial pH o 4.0 0 5.0 0 8.2 (Table 4.17)

Dixon (42) and Dixon, James and Healy (43) maintain that this marked

reduction in rate is due to (a) the conversion of H2S into non-volatile

HS (i.e. accumulated reaction products

cause a decrease in rate)

and/or (b) the adsorption of Znii hydrolysis products

on the ZnS surface which subsequently retard

zinc release.

o i S is formed on the surface of ZnS under UV irradiation at acid pH values

(63,Chapter 7) but is unstable at pH 8 (14, Chapter 7,8), therefore soluble

sulphur species are very likely to be present at pH 8. The electrokinetic

studies of Moignard, Dixon and Healy (9) suggest that ZnII hydrolysis species

are present at the ZnS-H20 interface.

The slight pH decrease observed following irradiation at pH 8 (rather than

the massive increase observed at acid pH values) supports both the presence of

soluble sulphur species in solution e.g. HS aq

together with the formation of

ZnII hydrolysis species. Without further evidence it is likely that both (a)

and (b) contribute to a decrease in the rate of dissolution at alkaline pH val-

ues. Further investigation of W-irradiation enhanced dissolution of ZnS at

alkaline pH values is warranted.

8.8(b) Electrokinetic studies.

Electrokinetic studies of the MII-ZnS systems provide considerable insight

into possible activation mechanisms at near neutral and alkaline pH values,

particularly when they are combined with the results reported in this present

study. The relevant electrokinetic evidence is discussed below.

(I) 'uII-ZnS .

The concentration of CuII in relation to the available ZnS surface

area is of critical importance in interpreting electrokinetic behaviour.At low

CuII concentrations (less than that required for monolayer coverage) in neutral

and alkaline media, all of the available Cu" will be consumed in the initial

activation step although the reaction rate maY be significantly decreased. The

role of CuII hydrolysis may then be obscured. As the concentration of CuII is

increased above the amount needed to satisfy this initial step, CuII hydrolysis

will be of greater importance. Since it is common practice in flotation plants

to "overdose" with copper sulphate at alkaline pH values, resulting in a

"reservoir" of Cu(OH)2SOlid (1, 2, 44), electrokinetic information at high

u Concn (M) 00 O 21=10-5 m 51=10-5 6 21 .10"4

i

-254-

I Cu

I /ZnS surface area ratios is necessary. The most useful information is,

of course, obtained for varying CuII/ZnS surface area ratios. Of the studies

performed by Hukki et al (10), Salatic et al (11) and Moignard et al (9) only

the latter work achieves this aim. The relevant diagram, extracted from

Moignard, Dixon and Healy's study (9) is reproduced as Figure 8.5.

4 5 6 7 pH

8 9 10

FIGURE 8.5 Electrophoretic mobility-pH isotherms for 1.0 g t."1 ZnS in 10.2 M KNO3 in the

absence and presence of Cu(NO3)2. The ZnS was pre-equilibrated at pH 4.0 for

17 h before any Cu(NO3)2 was added. Added Cue+ concentrations are 2.05 x

103 ,5.125x 10-5, and 2.5 x 10-4 M

The pH range above 6.5 for the CuII-ZnS system is applicable here (the

mobility data obtained at acid pH values, which strongly suggest the presence

of So on the ZnS surface, have already been discussed in Chapter 2). The CuII

concentration/ZnS surface area ratios used by Moignard, Dixon and Healy ranged

from 1.2 to 12 x 10-5 mole Q 1/m2 ; the corresponding value in this study was

1.5 x 10-5 mole 2,-1/m2. All of these ratios were sufficient for the initial

activation step to be complete. Moignard et al's work at ratios of 12 and

8.7 x 10-5 mole 2.-1/m2 show the negative to positive (CR2) followed by an

extrapolated positive to negative (CR3) charge reversal which is characteristic

of the surface-nucleated metal hydroxide precipitation discussed in detail by

James and Healy (64). A surface coating of Cu(OH)2 forms on the ZnS surface

(similar charge reversals were observed by Hukki et al (10) and Salatic et al

(11)). The data obtained at a ratio of 1.2 x 10-5 mole 2.-1/m2, while not show-

ing charge reversal, certainly indicates that part of the available ZnS surface

is coated with Cu(OH)2 (the mobility becomes less negative). At the ratio of

1.5 x 10-5 mole 1-1/m2 used in this study charge reversal may not occur but the

ZnS surface should certainly nucleate Cu(OH)2 to a greater extent than for a

ratio of 1.2. Despite the fact that Cu is generally present in solution in

this study before ZnS is added, surface nucleation should occur but may be

accompanied by heterocoagulation of the ZnS and the suspended Cu(OH)2solid' for

they are apparently of opposite charge in at least the high pH region of the pH

range 6 to 10 - the reported pzc for high purity, synthetic ZnS ranges from

pH 5 to 8.5, depending on pre-equilibration conditions (9) whilst iep values

for Cu(OH)2 vary from pH 7.6 to 9.4 (72). (heterocoagulation may also occur

between Cd(OH)2solid ZnS and Pb(OH)2solid ZnS. depending on pH; according

to Parks (72) the iep for Cd(OH)2solid is > pH 10.5 whilst that for Pb(OH)2

ranges from pH 9.8 to 11).

(II) CdII-ZnS.

Although electrokinetic studies of the CdII-ZnS system have appar-

ently not been performed to date, adsorption/precipitation of CdII hydrolysis

species on the surface of TiO2 and a-FeOOH, for example, is known to occur

(71). There is therefore every reason to suppose that similar behaviour will

take place at the ZnS-H20 interface.

-256-

(III) Pb -ZnS.

Dixon (42) has studied the electrophoretic mobility of high purity

(99.99%), synthetic ZnS of surface area 1.2 m2 g-1 in the presence of Pb(NO3)2.

The ZnS concentration varied from 0.1 gß 1 to 0.5 gR 1 while the [Pb"] ranged _ II _

from 1C to ' M. At a [[ZnSj

ratio of 8.3 x 10-3 mole 2714m 2, there

is a positive oc negative charge reversal at pH 4 (CR1)followed by a negative

to positive cbarce reversal at pH 5.8 (CR2) with a final positive to negative

charge reversal at pH 7 (CR3). The behaviour is therefore similar to that

found in the ̀ u"-ZnS system. At lower [PbII] ratios similar to that used in [ZnS]

this _resent study i.e. 1.7 x 10-5 mole 9.-1/m 2, Charge reversal CR2 does not

occur, gut the mobility becomes less negative, indicating that the ZnS surface

has nucleated a cartial coating of Pb(OH)2.

8.8(c) Mecpan_ _em.

The detailed analytical evidence presented in Chapters 4, 5 and 6, support-

ed by electrokinetic studies, show that at near neutral and alkaline pH values,

enhanced DTII ,_take occurs, due to adsorption and precipitation of metal ion

hydrolysis s_ecies. For Cu" and Pb" in the approximate pH range 6.6 to 7.4,

the evidence indicates that activation is essentially dominated by the mechanism

operating at acid pH values. At alkaline pH values this mechanism is compli-

cated by hydrolysis effects. For CdII a simple exchange reaction occurs up to

pH 9, whereupon enhanced CdII uptake is observed. As pointed out in Section 8.1,

Sc is unstable in neutral and alkaline media according to reaction 8 i.e.

FA SsurfaceST HxSy z ... Reaction B

Precipitation cf bulk hydroxides such as Cu(OH)2, Cd(OH)2, Pb(OH)2 and Zn(OH)2

is known to be rapid (5, 7). At pH values above 7, with increasing pH, the

following scheme may be envisaged (it is particularly dominant above pH 8 for

Cuit and PbII, above pH 9 for CdII; s denoted surface):-

(i) rapid bulk precipitation of M(OH)2 followed by adsorption/

coagulation on to the ZnS surface

and/or

rapid surface-nucleated M(OH)2 precipitation resulting in

a ZnS surface partially or completely covered in M(OH)2.

Maq + 20H FAST

M(OH) 2 q

s ... Reaction 12

(ii) a slow surface reaction between M(OH)2surface and ZnS

in contact with an aqueous medium to yield (Zn,M)Ss,

Zn2+ and OHaq. Reactions involving S° formation/ aq decomposition may occur. Znâq is formed when the pH

is not high enough to form Zn(OH)2so1'd' though some

hydrolysis of Zn2+ to Zn(OH)+ is naturally expected. aq aq

M(OH)2s + ZnSs SLOW (Zn,M)Ss + Zn- + 20Haq ... Reaction 13

(N.B. (Zn,

(iii)

M)Ss

slow

- see under reaction 10, Section 8.7(a)).

dissolution of M(OH)2surface to yield

followed by Reaction 10.

M22++ and aq OH

a q

M(OH)2 s

SLOW

M2+ + 20H ... Reaction 14

aq aq

-258-

(iv) a slow surface reaction between M(OH)2surface and ZnS,

in contact with an aqueous medium, to yield (Zn,M)Ssurface . This occurs when the pH is high and Zn(OH)2 surface

enough for Zn(OH)2 to form. Reactions involving So form-

ation/decomposition may occur.

M(OH)2 s + ZnSs

SLOW (Zn,M)Ss + Zn(OH)2s .. Reaction 15

OR

(v) slow dissolution of M(OH)2surface

to yield M2++ and aq

OH aq M2+

then reacts according to reaction 10.

aq aq The Zn2+ produced rapidly precipitates on to the ZnS

aq surface as Zn(OH)2surface.

m(OH)2 s

SLOW M2+ ' aq

+ 20H aq

Reaction 14

Znaq + 20Hag

FAST Zn(OH)2 surface ... Reaction 16

In the situation where there is a large excess of MII such that M(OH)2 is present

as a bulk or suspended solid as well as on the ZnS surface, reactions similar to

those specified in reactions 13 to 16 but involving M(OH)-bulk,

will occur.

Certainly these metal hydroxide coatings are thermodynamically unstable

in the presence of ZnS, as the relevant pKso values indicate (see Tables 2.1,

4.1, 5.1 and 6.1). Evidence demonstrating that the reactions outlined above

are slow is sparse, however some information is available. Firstly the quantity

of zinc exchanged in a given time decreases as the initial pH increases above 7,as

is shorn by the data in Table 8.10 - i.e. the activation rate decreases with

i.e.

1.14 x 10-5

1.05 x 10-5

0.82 x 10-5

5

7

8

10

5.21 x 10-5

4.43 x 10-5

2.23 x 10-5

not detectable

[Zn]exchanged' (Cu - ZnS)

mole R-1

(PbII-ZnS) Initial pH

increasing pH. Secondly the observation that Cu(OH)2 coatings, deliberately

nucleated on sphalerite or marmatite surfaces, do not result in flotation with

xanthate under short conditioning times is added evidence that steps 13 to 15

TABLE 8.10 Concentration of zinc exchanged at 120

minutes as a function of initial pH.

Conditions: normal light, [ZnS] = 5.00 g2.-1

[Cu initial= 5.12 x 10-5 M [PbII]initial = 4.79 x 10-5 M

(Tables 4.14, 4.15 ; 6.6, 6.7, 6.8)

are slow (44,65). Longer conditioning to ensure conversion of at least some of

the Cu(OH)2 to "CuS" does lead to xanthate flotation. In contrast much shorter

conditioning times are required at acid pH values (44). Thirdly there is a

small decrease in the rate constant (k2) during the second stage of activation

for CuII at initial pH 7.2 and CdII at initial pH 9.0 compared with acid pH

values i.e. the presence of a partial surface coating reduces the rate constant.

Regarding the proposed mechanism outlined above, the obvious question to now

ask is:

"At alkaline pH values is the activation reaction subject to control by

slow surface reactions or by slow dissolution reactions?" i.e.

can reactions 13 and 15 be distinguished from reaction 14?

Copper is not quantitatively detectable (< 10-7 M) in solution at pH 8

-260-

or above in the presence of ZnS while zinc cannot be detected (< 10-7 M)

at pH 10. Reference to Figure 4.7 shows that the Cut+ electrode still res-

ponds, albeit qualitatively, to changes in [Cu24-] as activation proceeds - aq

the electrode potential decreases (corresponding to a decrease in [Cuâq])

with time. These "cycle" results of Figure 4.7 (i.e. ZnS addition followed

by two successive pH adjustments) taken together with electrokinetic and

analytical evidence, are consistent with progressive dissolution of part of a

Cu(OH)2 surface coating followed by activation of the newly exposed ZnS sur-

face. The marked decrease in rate following the second pH adjustment suggests

that monolayer coverage has been completed and that the second, slow activation

step is taking place. These results indicate that activation is proceeding, at

least in part, through Cuâq , thus reaction 14 appears viable. An immediate

criticism is that since a pH increase occurs during activation at pH 8, per-

haps this leads to removal of Cut+ (as hydrolysis species). However the aq

Cut+ electrode potential also decreases to lower [Cu2+] during activation at aq

pH 10 where it is accompanied by a slight pH decrease (Table 4.15 E). Removal

of Cuâq at pH 10 is therefore not due to hydrolysis of Cu2+ but rather to aq

reaction with ZnS. Similar experiments performed for PbII activation at pH 8

and CdII activation at pH 9.5 showed parallel behaviour to that of Cu". Clearly,

then, reaction 14 is viable.

The possibility that reactions 13 and 15 are taking place in parallel

with reaction 14 cannot be ruled out. Certainly solid-solid reactions are

frequently slow (22). There is no quantitative information available on

diffusion rates in the micro-crystalline or amorphous "thin films" that are

encountered here. The effect of UV irradiation at an initial pH of 8.2

(Table 4.17 C) is to enhance zinc release, although less so than for the

blank. This is certainly due to the occurrence of surface redox

processes, but W irradiation may also enhance the proposed surface reactions

and dissolution reactions at neutral and alkaline pH values. In the absence

of definitive evidence one if forced to conclude that reactions 13, 14 and 15

occur in parallel. Reaction 16 occurs, of course, when Zn(OH)2 formation is

feasible e.g. around pH 10.

The reasons for the pH changes occurring at near neutral and alkaline

pH values during heavy metal ion activation of ZnS are basically those discussed

in Section 8.6(b). They will obviously be complicated by reactions 12 to 16

and the formation/decomposition of sulphur species (e.g. reaction B).

8.9. SUBSIDIARY EFFECTS.

8.9(a) Influence of S0 on MII uptake - further comments.

The presence of S° on the CuII activated ZnS surface has the most pronounced

effect on CuII uptake at an initial pH of 4.0 after 10 minutes of reaction time

(Figures 4.5, 4.6, 4.9). Noting that the actual pH of the system at an initial

pH of 5.0 is 5.9 after 15 minutes reaction, increasing to 6.3 at 120 minutes,

enhanced adsorption due to CuII hydrolysis species might be suspected. The

electrokinetic studies of Moignard, Healy and Dixon (Figure 8.5) demonstrate

that, for similar CuII and ZnS concentrations to those used in the present study,

the electrophoretic mobility just starts to become less negative at pH 7. It

is therefore unlikely that adsorption/precipitation of CuII hydrolysis species

occurs in the pH range 5.9 to 6.3. The reduced CuII uptake at an initial pH of

4.0 compared to 5.0 is apparently due to the larger quantity of S0 reducing the

reaction rate. A similar argument would seem to apply in the case of Pb" act-

ivation of ZnS,

8.9(b) Effect of pretreatment on subsequent CuII activation.

Pretreatment of ZnS at an initial pH of 5.0, either in the dark or in normal

light, has no detectable effect on subsequent CuII activation, compared with

untreated ZnS. The reactions discussed above remain unaffected.

Pretreatment at pH 12.0 in normal light leads to slightly smaller amounts

of zinc being exchanged during subsequent activation at an initial pH of 5.0

(i.e. [ZnIi]excbangsd is smaller). The effect is a minor one and is probably

due to the presence of polynuclear ZnII hydrolysis species or solid phases such

as ZnS•HOHs , Zn2(OH)2SO4s and Zn(OH)2s on the ZnS surface, as originally

proposed by Moignard-, Dixon and Healy (9). Formed at pH 9.8 and above they are

metastable at lower pH values, as is shown by the marked similarity in electro-

kinetic behaviour between ZnS pre-equilibrated at pH 9.8 and ZnO, observed by

Moignard, Dixon and Healy (9) following the earlier studies of Healy and Jellet

(17). The stability constants of ZnS-HOHs and Zn(OH)2s are documented by

Sillen and Martell (6). The presence of these metastable species apparently

retards zinc release during activation at an initial pH of 5.0, but does not

affect CuII uptake.

Pretreatment under W irradiation at an initial pH of 5.0, followed by

activation, results in enhanced CuII uptake compared to untreated ZnS activated

in the presence or absence of W. Since the pH in this pretreatment experiment

was 7.2 prior to activation, the increased CuII uptake is undoubtedly due to

the precipitation of Cu(OH)2solid on to the ZnS surface, a process which is

known to occur at pH 7.2 and above as shown both in the study by Moignard,

Dixon and Healy (9) and by the results in this present investigation. The re-

lease of zinc in the blank and during activation again clearly demonstrates that

different rate determining steps apply during activation and UV enhanced dis-

solution. The "CuS layer" retards zinc release during activation.

Pretreatment in normal light at pH 2.0 should result in the formation of

elemental sulphur (S°) at the ZnS surface, since it is known to be thermodynam-

ically stable at low pH (14, Chapter 8). The markedly reduced CuII uptake

-263-

during activation at an initial pH of 5.0, following pretreatment at pH 2.0,

supports the model of a deposit of S retarding CuII uptake. [ZnII]total and

[Znii]blank are increased relative to the untreated ZnS due to the enhanced

solubility of ZnS at low pH (refer Figure 2.1). [ZnII]exchanged

is slightly

less than for untreated ZnS, again suggesting that zinc release is retarded by

the S° deposit. If activation is performed at pH 2.0, after pretreatment at

this pH CuII uptake is similar to untreated ZnS at 5 minutes and then steadily

decreases over 2 hours. At the same time zinc release from the blank steadily

increases and is 5 to 10 times greater than for untreated ZnS at pH 5'.0. Zinc

release during activation similarly increases but at a lower rate than for the

blank. At this low pH the solubility of both the ZnS and "CuS layer" is greatly

enhanced. Dissolution effects tend to obscure any further insight into the

activation mechanism at this low pH, although S° undoubtedly forms on the

activated ZnS surface.

8.9(c) Fundamental solubility considerations.

Obviously the fundamental solubility of ZnS and the "CuS layer" must be

considered, quite apart from any effects due to UV irradiation. This has al-

ready been dealt with for ZnS in Chapter 2 (Figure 2.1). "CuS", CdS and PbS are

more insoluble than ZnS (Table 2.1). Since the solubility of these sulphides is

very small until low (< 3) or high (> 12) pH values, their contribution to the

reaction scheme outlined above is negligible. While acknowledging their pres-

ence they play a sufficiently minor role to be excluded from mechanistic consid-

erations at this point.

8.10. SUMMARY.

In this thesis an attempt has been made to establish the rate laws and

mechanism by which the heavy metal ions CuII, CdII and PbII activate zinc sul-

phide, in an effort to clarify the conflicting and incomplete evidence which

exists at the present time.

It is clear that in considering the activation process, one must take

account of exchange, surface redox and adsorption/precipitation reactions. A

major feature which has emerged from this study is the connection between the

nature of the activator ion and the semiconducting properties of zinc sulphide.

This complex mélange of reactions has to date remained unidentified, with the

exception of the very recent preliminary work by Healy et al (9. 43, 66) and by

Maust and Richardson (101).

It has not been possible to link a particular result with a specific mech-

anism in all cases and it is sometimes necessary to consider that two or more

reactions are occurring in parallel. Much work remains to be done on determin-

ing the relative rate constants for these various reactions. Establishing a

rate law for the later stages of Cu" and PbII

activation at acid and near

neutral pH values has not been possible - it will necessarily be complex and

for practical flotation purposes is probably unnecessary. Similarly the dev-

elopment of a rate law at alkaline pH values will remain elusive until the

kinetics of the relevant precipitation, adsorption, surface and dissolution

reactions have been fully explored, not only for sulphides but for solid-water

interfaces in general.

There is also a need for the further application of spectroscopic techniques

to characterize sulphide surfaces, particularly the precise ratio of ionized

acceptor concentration to total acceptor concentration, the location of the

various surface energy states and the surface disposition of adsorbed species

which, in turn, must be blended with appropriate theoretical studies. Differ-

ential thermal analysis may well be useful in determining the strength of the

bond by which S° is linked to the sulphide surface.

"Collectorless flotation" has been discussed in detail in Chapter 2. With

the benefit of hindsight it is probable that the observed flotation behaviour

was due to undetected S0 in many cases. A similar reason may also apply to

cases where it has been reported that irradiation of sulphide surfaces en-

hances flotation (109). The recent paper by Heyes and Trahar (133) discussing

the collectorless flotation of chalcopyrite at pH 11 is important. Although

S0 is unstable at this pH, polysulphides (Sn2 ) assume appreciable concentrat-

ions at such strongly alkaline pH values and may contribute, through adsorption

on the chalcopyrite surface, to the observed behaviour. Further investigation

is clearly required.

The very least that the evidence presented in this thesis should do is to

raise questions concerning the use of collectors in sphalerite flotation and

the causes of natural hydrophobicity. The surface properties of sphalerite

can obviously be dramatically altered by the presence of an activator and sur-

face reactions can be controlled ("on" or "off") using light of a specified

wavelength and intensity. Perhaps collectors are unnecessary in sphalerite

(and other sulphides?) flotation!

The rate laws and proposed reaction mechanisms for Cu", CdII and PbII activ-

ation of ZnS are summarized in Tables 8.11, 8.12 and 8.13.

-266-

TABLE 8.11. CuII ACTIVATION OF ZnS (inert atmosphere, sealed system)

SUMMARY OF RATE LAWS AND REACTION MECHANISMS

T = 298 K

pH Rate Law Reactions ("s" denotes surface)

Major Minor

A. 4.0 to 6.5 0-approx 1.5 minutes

Cu2+ + ZnS —+ (Zn,Cu)S + Zn2+ ...1 — aq +--- aq

Cu2+ + e ktZ Cu+s

1132- kT

11So + e ...3 s s

o SLOW S H S O 4

Important under UV irradiation

_ = kl1og10(t) + r1

k1= 5.6±0.9x10-6mole m-2

approx 1.5 to 15 minutes

IF

INITIAL

STAGE

OF ACTIVATION IS

COMPLETE j

. = k 1og10(t) + constant

k2= 1.4 to 2.8x10-6mole m-2

15-120 minutes

1111+ + o

ZnS — Zn2 + S + 2e

— aq

+ s replaces Zn2 in lattice)

reduction reactions = f(t,pH,incident light,

pretreatment) Haq + e ! 11H2g

H2O2 + e OHaq + 11H2g

s y r— x z

8. near neutral

approx 6.6 to 7.4

generally as in

A. above

with k1 the same

and k2=1.1±0.4x10 6moole

Dependent on extent of CuII hydrolysis

(subsidiary to 1 to 4 unless [CuII]

very large)

2OH FAST- Cu2++ + • Cu(OH)2s ...5

S 9-'(Zn,Cu)Ss. +Znaq+ + 2OHaq Cu(OH)21ZnS ...6

OR

S Cu(OH)2s Cua

+4

+ 2OM ...7 then 1.

FAST S H S 0 ...8 s x y z

C. approx. 7.5 ;

and above

no quantitative

rate law

CuII hydrolysis important, 5 to 8 dominant with

9 and 10 if Zn(OH)2 precipitates ; 2 and 3

also occur.

Cu(OH)2s + ZnS S--W' (Zn,Cu)SS + Zn(OH)2s ...9

FA .T - Zn(OH)2s ..10 Znaq + 20Haq

-267-

TABLE 8.12. CdII ACTIVATION OF ZnS (inert atmosphere, sealed system)


T 298 K

pH

-

Rate Law Reactions ("s" denotes surface)

Major Minor

A. 4.0 to 8.0 0-1.0 minutes + + Cd2 + ZnS ------• (zd,Cd)S + Zn2+ ...1 Important under UV irradiation

ZnS 4 Znaq +S8 +2e

2+ 2+ (Cd replaces Zn in lattice

Cd g+ +e

h\4 Cds

ßf5

2- h-V+ 1 o

+ e +-- JSs

reduction reactions

r = kj1og1O(t) + r1

k1=1.1'_0.2x10 6mole m-2

(k1 is average value)

1.0 to 120 minutes

+-----aq s +----- s aq

IF INITIAL

STAG OF ACTIV-

ATION IS

COMPLETE

7 = k21og10(t) + r1

k2=0.64±0.10x10 6mole m-2

(k2 is average value)

aq + e , 4H2g

H2OR + e —• OH + )1H29

SLOW (acid pH) Ss HxSy Z

■

B.approx 8.1 to minutes Dependent on extent of CdII hydrolysis 9.0

r kllOglO(t) + r1

k1-0.84-0.20x10 6mo1e m-2

1.0 to 120 minutes

(subsidiary to 1 unless [CdII) very large)

Clog + 20H MET. Cd(OH)2s ...2

SLOW

. = k21o410(t) + r1

k2=0.34±0.09x10-6mole m-2

Cd(OH)2 + ZnS (Zrt,Cd)S +Zn2+ +20H aq

s s t s aq aq

...3

OR

Cd(OH)2s +LOWS Cdâq + 20Haq ...4 then.l

o FAST S —► H S O s*-- xyz

(neutral alkaline pH)

C. above pH 9 no quantitative

rate law

CdII hydrolysis important. 2 to 4 dominant

with 5 and 6 if Zn(OH)2 precipitates.

S Cd(OH)2s + ZnSs (Zn,Cd)SS + Zn(OH)2s...5

FAST Zn(OH)2s ...6 Zn2++ + 2OHaq

-268-

TABLE 8.13. II Pb ACTIVATION OF ZnS (inert atmosphere, sealed system)


T=298 K

pH Rate Law Reactions ("s" denotes surface)

Major Minor

A. 4 to 6.5

IF

INITIAL

STAGE

COMPLETE

0-1.0 minutes Pb2+ + ZnS -----+(Zn,Pb)S +Zn2+ ...1

Important under UV irradiation

r = k110410(t) + r1

k1=0.89 to 1.9±0.2x10-6mole m 2

> 1.5 minutes

aq

Pb2+ + e 1!Z+ Pbg .. .2 s

1152 kT+ 1fSO + e . 3

o stow Ss = HxSŸ z

...4

ZnS 11 ► Zn2+q s + S° + 2e a

(Pb2+ replaces Zn2+ in lattice

reduction reactions

r = f(t,pH,incident light

(Zn 51, (Pb"])

Nag + e *12g

H2O1 + e OH8 + 4N2 q 9

B. near neutral

approx 6.6 to 7.4

as in A. above with

k1 =1.6*0.1x10 6mole m-2

Dependent on extent of PbII hydrolysis

(subsidiary to 1 to 4 unless (PbII]

very large)

Pb2+ + 20Haq FAQ ST' Pb(OH)28

Pb(OH)2s+ZnSs S (Zn,Pb)SB+Znaq++20Haq ...6

OR

Pb(OH)2 LO-w+

Pb2+ + 201 ...7 then 1 Pb(OH)2s aq

o S FASTAST* HSÿ ...8 z

C. approx pH 7.5

and above

no quantitative

rate law

PbII hydrolysis important. 5 to 8

dominant with 9 and 10 if Zn(OH)2

precipitates. 2 and 3,also occur.

Pb(OH)28 +Zn5s Si- -:3 1-4 (Zn,Pb)Se+Zn(OH)2 ...9

HFA )Zn + 20 Zn(OH2s ..10q aq

-269-

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Minerva Access is the Institutional Repository of The University of Melbourne

Author/s:Ralston, John

Title:Heavy metal ion activation of zinc sulphide

Date:1978

Citation:Ralston, J. (1978). Heavy metal ion activation of zinc sulphide. PhD thesis, Department ofPhysical Chemistry, The University of Melbourne.

Publication Status:Unpublished

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Heavy metal ion activation of zinc sulphide: (1978)

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