aqueous oxidation of pyrrhotite and pyrite minerals - UNSWorks

Measurement of the rate of aqueous oxidation of sulphideminerals: aqueous oxidation of pyrrhotite and pyrite minerals

Author:Fabian, Cesimiro P.

Publication Date:1990

DOI:https://doi.org/10.26190/unsworks/5634

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MEASUREMENT OF THE RATE OF AQUEOUS OXIDATION OF SULPHIDE

MINERALS

AQUEOUS OXIDATION OF PYRRHOTITE AND PYRITE MINERALS

A thesis submitted for the degree of Master of Science in

The University of New South Wales.

Cesimiro P. Fabian

June, 1990

Candidate's Certificate:

This is to certify that most of the work presented in this

thesis was carried out by the candidate in the School of

Materials Science and Engineering of the University of New

South Wales and part at Aurotech N.L. and has not been

submitted to any other University or Institution for a

higher degree.

Cesimiro P. Fabian

Bachelor in Chemical Engineering Science and

Chemical Engineer (Peru)

ABSTRACT

The rate of consumption of oxygen by the aqueous oxida

tion of pyrrhotite minerals has been measured by an oxygraph

at constant fixed atmospheric pressure and about 31.5°c

temperature. The oxygraph, designed for the purpose of this

study, consisted of an stirred reactor and an electrolytic

cell for the production of oxygen interfaced to a microcom

puter. The internal pressure of the system, controlled by a

manometer containing two electrodes (contacts), indicates to

the computer the ON and OFF time of the electrolytic cell.

The leaching rates were measured in both air and pure

oxygen, in the absence of bacteria and with inoculation of

T. Ferrooxidans.

New potential-pH diagrams for the Fe-S-H2o system at

2s0 c temperature, neglecting pyrite, were constructed in

order to analyze the mineral chemistry of stoichiometric and

monoclinic pyrrhotite. New predominance area diagrams for

both ferric hydroxyl complexes, and ferric hydroxyl and

sulphate complexed species were derived.

The overall rate of consumption of molecular oxygen by

pyrrhotite minerals at an initial pH of 4.0, was found to be

as follows (grams of oxygen per minute): chemical oxida

tion, 0.00037; chemical oxidation with pure oxygen 0.00057

and bacterial oxidation 0.0020. However, the highest initial

rate of consumption of oxygen occurred in presence pure

oxygen.

These results, when compared with those of ''shake

flasks" experiments in terms of the extent of reaction, were

in good agreement and explained in detail in terms of sul

phide ions oxidation. It was also derived that the oxygen

formed in the sulphate ions (secondary reaction) comes from

the water rather than from the oxygen partial pressure. Thus

the rate of consumption of oxygen is a reliable method to

determine the rate of oxidation of pyrrhotite minerals at pH

values about 4.0. So the initial assumption on the stoichi

ometry of the aqueous oxidation of pyrrhotite minerals,

whether or not in the presence of microorganisms was con

firmed to form ferric hydroxide and elemental sulphur,

predominantly.

FeS + O.SH20 +0.7502 -----> FeOOH + s 0

Fe7s8 + 3.5H20 + 5.2502 -----> 7FeOOH + 8S0 •

The present knowledge on continuous bacterial oxidation

of sulphide minerals (6 tanks x 30 L) is discussed in terms

of flow diagram, pulp density, bacterial recirculation and

dissolved oxygen concentration. At 1.5 initial pH, the main

oxidations products are ferric ions and sulphate ions. The

concentration of ferric ions, hence the potential, increases

steadily throughout the leaching tanks. The bacterial degra

dation process was improved up to 50 percent iron oxidation

in 250 hours residence time by increasing the air sparge to

about 2.6 scc/min/L of slurry.

ACKNOWLEDGEMENT

The author wishes to acknowledge and to thank firstly

to Mr. Bruce Harris, Senior Lecturer in the School. With out

his frequent stimulation, including very frank expressions

of criticism and his painstaking reading, the thesis might

well never have been completed. I also extent my sincere

gratitude to Mr. Fred Scott, Professional Officer in the

School, for advising me on the interfacing of the computers

with the experimental equipment. Thirdly to the management

of Aurotech N.L. for releasing me from the confidentiality

agreement. Finally to my wife and daughters for their under

standing because the entire thesis was prepared after my

normal hours of work at different companies.

CONTENT

PART A: LITERATURE REVIEW

CHAPTER I

Introduction

CHAPTER II

Physical Properties of Iron Sulphide Minerals

CHAPTER III

Aqueous Oxidation of Sulphide Minerals

III.1 Aqueous Oxidation Of Chalcopyrite

III.2 Reduction of Chalcopyrite

III.3 Aqueous Oxidation of Pyrite

III.4 Summary

CHAPTER IV

Aqueous Oxidation of Pyrrhotite Minerals

IV.1 Introduction

IV.2 Acid Decomposition of Pyrrhotites

IV.2.1 Stoichiometry and Reaction Order

IV.2.2 Summary

IV.3 Oxidation of Ferrous Ions by Oxygen

IV.3.1 Oxidation of Ferrous Ions in Strong

Sulphuric Acid Solutions

IV.3.2 Oxidation of Ferrous Ions In Low

Concentrations of Acid Solutions

IV.4 Hydrolysis and Precipitation of Ferric

Page

1

5

14

14

18

19

22

24

24

25

27

36

38

39

43

Page

Ions 51

IV.4.1 Ferric Iron Species in Solution 56

IV.4.2 Hydroxyl Complexes Species 56

IV.4.3 Development of Predominance Regions

for Ferric Hydroxyl Species 57

IV.4.4 Sulphate Complexed Species 62

IV.4.5 Development of Predominance Regions

Ferric Hydroxyl and Sulphate Complexes

65

IV.4.6 Precipitation of Ferric Ions

IV.5 Aqueous Oxidation of Pyrrhotite

Minerals by Oxygen and Ferric Ions

IV.5.1 Aqueous Oxidation of Pyrrhotites

65

69

by Ferric Ions 76

IV.5.2 Aqueous Oxidation of Pyrrhotites

at Elevated Temperatures 83

IV.6 Bacterial Oxidation of Pyrrhotite

Minerals 87

IV.7 Literature Survey Conclusions 96

PART B:

CHAPTER V

Development of Potential-pH Diagrams for Pyrrhotite

Minerals

V.l Thermodynamic Properties of Pyrrhotite Minerals

V.2 Derivation of Potential-pH Diagram for the

99

99

Fe-s-H2o System at 2s0 c Neglecting Pyrite

Page

99

V.3 Analysis of the Oxidation Leaching of Pyrrhotite

Minerals

PART C: EXPERIMENTAL PROGRAM

CHAPTER VI

Introduction to the Experimental Program

CHAPTER VII

VII.I Initial Equipment Design

VII.1.1 Procedure for Determining the

Rate of Oxygen Consumption

103

112

120

120

VII.1.2 Evaluation of the First Oxygraph 123

VII.2 Second Equipment Design 124

VII.2.1 Procedure for Determining the

Rate of Oxygen Consumption 124

VII.2.2 Evaluation of the Second Oxygraph 126

VII.3 Third Equipment Design 127

VII.3.1 Procedure for Determining the 127

Rate of Oxygen Consumption

VII.3.2 Evaluation of the Third Oxygraph 130

VII.4 Consumption of Oxygen During Aqueous

Oxidation of Pyrrhotite Minerals

VII.4.1 Rate of Consumption of Oxygen

During Chemical and Bacterial

131

Oxidation of Pyrrhotite Minerals 132

VII.5 Justification of the Method for Determining

for Aqueous Oxidation Rate of Pyrrhotite

Page

Minerals 144

VII.5.1 Reproducibility of Reaction Rate 153

VII.5.2 Ferrous and Ferric Ions in Solution 154

VII.5.3 Chemical Oxidation of Pyrrhotite

with Pure Oxygen and Air

VII.5.4 Effect of Sodium Chloride and

Copper Ions

VII.5.5 Mass Balance: Shake Flasks

Experiments

VII.6 Discussion and Conclusions

VII.7 Present Knowledge on Continuous Bacterial

Oxidation of Pyrite Concentrate

VII.7.1 Introduction

VII.7.2 Experimental Program

VII.7.3 Experimental Results

VII.7.4 Discussion and Conclusions

CHAPTER VIII

Discussion and Conclusions

APPENDIX A

APPENDIX B

APPENDIX C

APPENDIX D

APPENDIX E

APPENDIX F

APPENDIX G

APPENDIX H

REFERENCES

157

161

163

179

184

184

185

189

196

198

CHAPTER I

INTRODUCTION

Research and development in hydrometallurgical proc

esses continues providing potential alternative methods for

some pyrometallurgical processes (e.g. bacterial oxidation

of refractory gold minerals) as a replacement for conven

tional roast-leach process. The kinetics and mechanism of

such hydrometallurgical processes are frequently unclear

making it difficult to operate under optimum conditions.

Basic studies of the equilibrium concentrations of chemical

species with complicated solution chemistry involving

simultaneous oxidation-reduction, metal-ligand complexation

and hydrolysis and precipitation are valuable in understand

ing these processes.

This study commenced by reviewing the mineralogical

properties of the various iron sulphide phases and minerals

and was followed by an analysis of the aqueous oxidation

process of chalcopyrite, pyrite and pyrrhotite. The oxida

tion of chalcopyrite and pyrite was reviewed in order to

understand the preferential oxidation product(s) of sulphide

ions whether to elemental sulphur and/or sulphate.

Hydrometallurgical processing of pyrrhotite minerals

to recover tin from cassiterite concentrate tailings from

Renison Bell, Tasmania, appears to be economically attrac

tive. Since the rate of aqueous oxidation of pyrrhotite

1

minerals is faster than that of some other sulphide miner

als (e.g. bacterial degradation of chalcopyrite and pyrite),

the determination of the kinetics and mechanism of the

oxidation of pyrrhotite minerals should improve the knowl

edge of their mineral chemistry and also assist in the plant

design.

Aqueous oxidation of the pyrrhotite minerals is

discussed thoroughly with particular reference to their acid

decomposition, the oxidation of ferrous ions and the hydrol

ysis and precipitation of ferric ions. Using the most

recent thermodynamic data, two new

(modified potential-pH diagrams) have

activity diagrams

been developed in

order to analyze the equilibrium between the various ferric

hydroxyl species in solution and their stability in equi

librium with ferric hydroxide. The second activity diagram

is for the sulphate complexing species in solution with

ferric hydroxyl species.

A new potential-pH diagram for the Iron-Sulphur-Water

system at 2s0 c, neglecting pyrite, was developed in order to

analyze the aqueous oxidation process of pyrrhotite miner

als. Recently published thermodynamic data (1982) was used

to construct these diagrams.

It is widely recognized in the literature that reduc

tion of dissolved molecular oxygen is the most important

cathodic process during the aqueous oxidation of sulphide

minerals. It is important because, at or near the rest

potential range from Oto 0.62 V (SHE), the only cathodic

reaction capable of maintaining a significant current is

2

the reduction of oxygen. Reduction of other dissolved oxi

dants, e.g. ferric ions, may occur but reduction of dis

solved oxygen still plays the important role of regenerating

reduced species. Ferric ions as oxidant is thoroughly ana

lyzed in Section IV.5. Thus measurements of rate of con

sumption of oxygen may help determine the kinetics of the

aqueous oxidation process of such sulphide mineral.

The higher rest potential value, 0.62 V, is referred

to pyrite which does not decompose at an appreciable rate on

standing in acid solutions and, consequently, pyrite is

defined as the more active sulphide mineral for oxygen

reduction.

In this study, three designs of oxygraphs were de

veloped, constructed and used in order to measure the con

sumption of oxygen during a chemical and bacterial oxidation

of pyrrhotite minerals. The third version oxygraph, which

far exceeded the performance to the first two oxygraphs,

works under a fixed atmospheric pressure and at constant

ambient temperature.

The oxygen consumed was produced electrolytically as

demanded by the aqueous oxidation process thus maintaining a

constant pressure in the system. Solutions of dilute sulphu

ric acid, sodium hydroxide and copper sulphate were used as

electrolytes for the production of oxygen for the three

oxygraphs, respectively.

Attempts have been made to obtain a reproducible

reaction rate. The results were about ninety per cent repro-

ducible. The reaction rates measured in these series of

3

experiments were compared with those obtained from 'shake'

flask experiments. The results indicated that the extent of

reaction obtained from the oxygraph were higher than that

from the 'shake' flasks experiments. This difference is

probably due to higher transference of oxygen, hydrogen ions

and ferric ions in the oxygraph forced by the vibramix.

The rate of oxidation of pyrrhotite minerals or any

other sulphide mineral could be measured by the use of the

third oxygraph; provided no gas (e.g. hydrogen sulphide)

is produced.

4

CHAPTER II

PHYSICAL PROPERTIES OF IRON SULPHIDES

Iron sulphides are the most abundant of the natural

metal sulphides. Pyrrhotite and pyrite constitute the binary

boundary of many multicomponent base metal sulphides of

economic importance, e.g. copper, nickel, zinc and pyritic

gold bearing minerals. New hydrometallurgical processes,

such as those for the recovery of gold and silver from

refractory sulphide minerals, are becoming more attractive

as many materials can not be treated economically by the

more conventional roast-leach processes. Thus an analysis of

the physical properties of pyrrhotite minerals may lead to

a better understanding their aqueous oxidation process.

Craig et al. (1) reviewed the mineral phases of nine

teen compounds in the Fe-S system. Table II-1 lists the

minerals and phases having compositions in the range from

FeS to FeS2 , gives their stoichiometric structure and indi

cates their relative stability. This table is not exhaustive

but it does indicate the extensive and, in some cases,

complex nature of the iron sulphide system.

A number of minerals and phase have not been fully

characterized because many natural mineral systems are more

complex than synthetic ones, because they have had a much

longer time to approach equilibrium than can be afforded by

experimentalists and because metastable phases are often

5

produced during the formation of synthetic materials.

6

TABLE II-1

MINERALS AND PHASES IN THE Fe-s SYSTEM(l)

MINERAL

NAME

Troilite

COMPOSITION STRUCTURE TYPE

(cell edges, A)

FeS Hexagonal 2C

THERMAL

STABILITY, 0 c

Max. Min.

140

Mackinawite Fes1 _x Tetragonal P4/nmm Often contains

0.07>x>0.04 Ni and Co

Hexagonal

pyrrhotite

Fe1_xs

44.9-50.0

at.% Fe

MC-Type Fe1 _xS

pyrrhotite 47.4-44.7

at.% Fe

NA-type Fe1_xS


at.% Fe

NC-Type Fe 1_xS


at. % Fe

SC

pyrrhotite

11C

pyrrhotite

6C

pyrrhotite

Hexagonal 1C

Hexagonal?MC

Hexagonal?

3C

Hexagonal?

NC

Hexagonal

Orthorhombic

11C

Orthorhombic?

6C

7

mp.1190 100

NiAs structure

308 262

"'266 209

"'213 "'100

"'100

"'100

Metastable Fe1_xs Hexagonal Metastable

pyrrhotite 0.06>X>0.03 Hexagonal?

4C Fe7+xsa Monoclinic 254

monoclinic 46.4-47.3

pyrrhotite at. % Fe

Anomalous Fe7+xsa Triclinic? ?

pyrrhotite 46.4 at. % Fe

Gamma iron Fe2s 3 Spinel ?

sulphide

Smythite Fe9S11 Pseudo-rhombohedral '''75

Hexagonal or monoclinic?

Greigite Fe3 s 4 Spinel Metastable(?)

Pyrite Fes2 Cubic 743

Marcasite FeS2 Orthorhombic Metastable

slightly s-

deficient

8

Table II-2 lists the pyrrhotite minerals which occur

naturally(3}. For the purposes of this study, characteriza

tion of the physical properties of iron sulphides will be

confined to this list, excluding pyrite.

Troilite should only be used to describe a polymorph

of stoichiometric FeS which occurs in meteorites (1} but it

has recently been frequently reported in terrestrial envi-

ronments usually associated with low-temperature

pyrrhotite (1,5}. Naturally occurring troilite

hexagonal

has been

found to have an iron-to-sulphur ratio varying from 1:0.995

to 1:1.007 (3}. Troilite and a number of other polymorphs,

are p-type semiconductors with the positive holes associat

ed with the sulphur irons as opposed to forming Fe3 + by

association with iron (3}.

The nomenclature used in the literature for pyrrho

tites is based on the cell edges of the niccolite (hexagonal

NiAs) structure. Troilite, FeS, which has a hexagonal cell,

has an 'a' axis equal to /3A and its 'c' axis equal to 2C,

where A and C refer to the cell edges of NiAs, thus troilite

is called '2C pyrrhotite'.

Another naturally occurring iron sulphide mineral of

near-FeS composition is mackinawite, Fel+xs (x = 0.057 to

0.064}, which has also been described as natural pyrrhotite

by Smith et al. (5), but it is not described as such by

Lambert (3).

9

TABLE 11-2

COMMON IRON-SULPHUR COMPOUNDS (3)

Compound Name Structure Atomic yin FeSy

FeS2 Pyrite Cubic 33.33 2.000 0.667

Fe7S9 Monoclinic 4C 46.67 1.143 0.993

pyrrhotite

Fe9S1o Hexagonal SC 47.37 1.111 0.947

pyrrhotite

Fe10S11 Intermediate llC 47.62 1.60 0.952

Fe11 5 12 pyrrhotites 6C 47.83 1. 091 0.957

FeS Troilite 2C 50.00 1.00 1.000

* N = mole fraction of FeS in the FeS-FeS2 system.

10

When examining the voluminous literature on pyrrhotite

mineralogy, it seems that monoclinic pyrrhotite, Fe7s 8 and

hexagonal pyrrhotite are more common and more frequently

studied. Monoclinic pyrrhotite is ferromagnetic having a

monoclinic superlattice with composition centered about

Fe7s 8 while anomalous monoclinic pyrrhotite and the other

pyrrhotite minerals are antiferromagnetic (1,3).

Anomalous monoclinic pyrrhotite and hexagonal pyrrho

tite appear to be widespread in low temperature sedimentary

environments. Naturally occurring monoclinic pyrrhotite has

a room temperature magnetization of 13.1 emu/g, but this

decreases with increasing temperature and at 320°c it be

comes paramagnetic. Its density is 4.6 g/cm3 and it is

called '4C pyrrhotite'.

In general, the most abundant natural pyrrhotites are

believed to consist of three different pyrrhotites mainly,

4C (monoclinic pyrrhotite), nC (intermediate pyrrhotites)

and 2C (troilite). The nC pyrrhotites have a composition

range approximately from Fe9 s 10 to Fe11 s 12 and show integral

types of superstructure (4,6,22).

Figure II.1 shows a summary phase diagram for the FeS

Fes2 system below 3so0 c reported by Craig et al. (1). It can

be seen that hexagonal pyrrhotite and troilite may coexist

at temperatures below so0 c. Monoclinic pyrrhotite may exist

independently or in conjunction with smythite below this

temperature.

The structure of pyrrhotite is thought to consist of

s2- ions arranged in a 3-dimensional hexagonal packing with

11

350 ,----,.-----,--~.----.---,----"T_;__..---,-----.---,.--..---.-----.---,.-----~

300

250

0 0

~200 Q) .... ::> -0 .... Q)

0. E 150

~

100

50

w II-0 ::i: 0:: c:: >-

w a.. >-I-(.)

w I-;:: 0 ::i: 0:: a:: >-a.. _J <t z 0 Cl <t X w ::i:

I

0.. _J '....,..,-,-A,.,..,.,.-;i + 6 z 0:: 0 IC> <t X w ::i:

HEXAGONAL PYRRHOTITE (1C) + PYRITE

"HEXAGONAL" PYRRHOTITE (MC) + PYRITE

"HEXAGONAL " PYRRHOTITE (NA) + PYRITE

("HEXAGONAL" PYRRHOTITE (NA ) + MONOCLINIC PYRRHOTITE )

("HEXAGONAL" PYRRHOTITE (NC) -+ MONOCLIN IC PYRRHOTITE )

MONOCLINIC PYRRHOTITE + PYRITE

(SMYTHITE + MONOCLINIC PYRRHOTITE )

I ,._ I

I I

SMYTHITE + PYRITE

308°

262°

254"

,.,_, 75•

Q L---'-----'--'-'-..,____._ _ __.__~.____. __ ........_ _ _._ _ __._ _ __. __ ....__ _ _.__ ______ ___;.___..,___-'----' 50 4 8 OW 46 W 44 42 4 0 38 36 34 zl- I- ~ ~ ::::i§ :r: c::; Atomic % Fe _l Ox ~ w 0 Oo:: ::E o:: 0: Zo:: Vl 2 I- ~li::

FIGURE 11.1: The Summary Phase0

Hiagram for the FeS - FeS 2 Segment below 350 C ( 1 ) .

the Fe2+ ions in the interstices between the sulphurs. The

interstices are of two types, octahedral and tetrahedral.

The octahedral sites, which are the larger, are equal in

number to the s 2 - ions and are occupied by Fe2 +; the smaller

tetrahedral sites are vacant (2).

Smith et al.(5) cited Lotgering (1956) who represented

the formula Fe7s 8 as Fet;fe3o;s8 . The arrows denote the

direction of electron spin and represents a cation vacan-

cy. Smith et al.(5) also cited Subbarono who together with

Lotgering (1964) maintain that Fe7s8 must contain ferric

ions to preserve electrical neutrality and they suggest an

ionic formula, Fe1122 +/Fec 112_3x)> 2+Fe2

3+a;s for Fe1_xs

pyrrhotite. However no evidence of ferric ions was detected

using a Mossbauer spectroscopy and, moreover, the ferric

iron-to-sulphide bonding is known to be extremely unstable

( 5) .

In 1970, Ward (6) proposed that the formula for Fe7s8

can be written as Fe52+Fe2

3 + Cl ;s8 in which stands for an

Fe vacancy and that the whole structure is electrically

neutral. But he also suggested that a better formula would

be "Fe72+as h+ 2", (it is believed that there is an error in

the number of atoms of sulphur given in this representation

and it is though that the correct formula could be

Fe72+os8h+ 2 ) where h+ stands for two electron holes, the

location of which are not specified.

The characterization of the chemical structure of

troilite, FeS, seems widely accepted as comprising ferrous

and sulphide ions, while that of the other pyrrhotite miner-

12

als is not clearly understood. However, it is not necessary

to clarify the existence of ferric ions in the atomic struc

ture of hexagonal or monoclinic pyrrhotite for the purpose

of this study, nevertheless Ward's (6) second representa

tion, Fe72+as8h+ 2 will be accepted as the most probable

because the literature in general considers only ferrous

ions, Fe2+ as the product of acid dissolution.

13

CHAPTER III

AQUEOUS OXIDATION OF SULPHIDE MINERALS

III.1 AQUEOUS OXIDATION OF CHALCOPYRITE

The impetus for the development of alternative hydro

metallurgical processes for the treatment of sulphide miner

als comes principally from the desire to realize a process

which would provide a viable alternative to the sulphur

dioxide and arsenic emissions associated with pyrometallur

gical processes. Many hydrometallurgical processes have been

developed in recent years, particularly for the treatment of

copper concentrates and refractory gold minerals. For exam

ple many methods have been proposed to process chalcopyrite,

CuFeS2 , the most abundant copper sulphide mineral. Among

these are The Sherrit-Cominco process, The CLEAR process,

The Minimet Rescherche process, The Dextec Process and The

Cymet Process.

Several attempts have been made to devise a process

based on the following reaction (Equation III.1.1)

CuFeS2 + H2S04 + 1.2502 ----> CuS04 + O.SFe203 + s 0 + H2o

(III.1.1)

Copper could be easily recovered from the ferric iron

precipitate and elemental sulphur, but the low yield of this

reaction appears to be unacceptable. Under optimum condi-

14

tions of a temperature of 115°c and oxygen partial pressure

of 200-500 psi. only 65% of the copper was dissolved in 2

1/2 hours (7).

The attraction and advantages of sulphuric acid leach

system is obvious for leaching sulphide minerals and the

recovery of the precious metals seems to be comparable with

that obtained by smelting practices (7).

Aqueous oxidation of chalcopyrite by ferric (Fe3 +)

ions and oxygen (Equation III.1.2 and III.1.3) has been

extensively studied (8,9).

CuFeS2 + 4Fe3 + -----> cu2 + + 5Fe2 + + 2s0 (III.1.2)

(III.1.3)

It has been proposed (8,9) that both of these reac

tions occur during the dump leaching of chalcopyrite ores,

while reaction (III.1.2) has been proposed for the leaching

of chalcopyrite concentrate in the presence of ferric ions

(10). When oxygen is present in the system,

ferrous ions, Fe2 + also occurs (8-10).

oxidation of

However, it has been pointed out that the oxidation of

ferrous ions should be prevented in a dump leaching, if it

were possible, because according to Liddle et al. (8)

"supplying dissolved oxygen increases the amount of

copper dissolved before Fe(OH) 3 precipitates only if the

initial ferric sulphate, Fe2 (S0 4 )3 concentration is 10.0g.L-

1 or lower. Chalcopyrite dissolution could be increased at

higher sulphate levels but oxygen must be excluded if pre-

15

cipitation of ferric hydroxide is to be prevented. The

importance of solution pH in preventing precipitation should

be emphasized".

More recently, studies of hydrometallurgical process

ing of chalcopyrite have concentrated on the electrochemical

phenomena involved (11-13). Warren et al. (11) state that

"electrochemical reactions of a mineral are a direct result

of the thermodynamic properties of the mineral, properties

of the electrolyte and their interactions at the mineral

electrolyte interface". Further, anodic polarization of

chalcopyrite is sensitive to pH at potentials higher than

0.7V (vs SHE) and insensitive to pH at lower potentials in

sulphate solution. Based on current and mass balance meas

urements, two intermediate sulphide phases appears to form

in the sequence chalcopyrite (CuFeS2 ) ---->bornite

(Cu5 FeS4 )----> covellite (CuS) which, mixed with the elemen

tal sulphur produced, form an electron conducting passive

layer on the mineral surface (11).

Below potentials (0.45-0.550 V), the reactions are

reported to be consistent with the following equations;

provided y>x (11).

CuFeS2 ----->

cu1 _xFe1 _yS2 _z + xcu2 + + yFe2 + + zs0 + 2(x+y)e-(III.1.4)

Cu1-xFe1-yS2-z ----->

(2-z)CuSn-s + (1-y)Cu2 + + (1-y)Fe2 + + 2(1-y)e- (III.1.5)

CuS( ) -----> cu2 + + s 0 + 2en-s (III.1.6)

It is noted that ferrous ions formation occurs in this

16

potential range.

Warren et al. (11) states that at potentials higher

than 0.7V, the "previously suggested processes still occur

but the amount of Cu 1_xFe1_yS2 _z formed is negligible com

pared with the amount of CuFes2 oxidized completely to cu2 +

and Fe2 +. This is due to the greater amount of material

forced to dissolve at higher current". Thus the

process is represented by the following reactions:

CuFeS2 -----> cu2 + + Fe3 + + 2s0 + Se-

overall

(III.1.7)

CuFeS2 + 8H2o ----> cu2 + + Fe3 + + 2so42 - + 16H+ + 17e

(III.1.8)

It is though that 88-90% of the chalcopyrite reacts via the

first reaction and remainder via the second reaction (11).

Independently, Hillrichs et al. (12) also studied the

kinetics of the electrochemical dissolution of chalcopyrite

and suggested slightly different electrochemical reactions,

at low potentials, in terms of the solid intermediate

products and ferric iron. They indicated that three solid

intermediate products are formed simultaneously, but appar

ently these were not characterized. Equations (III.1.9,

III.1.10 and III.1.11) describes these reactions. , The

production of ferric (Fe3 +) ion, which is included in these

equations, also disagrees with the reactions suggested by

Warren et al. ( 11 ) .

CuFeS2 -----> cu 1_xFeS2 + xcu2 + + 2xe- (III.1.9)

CuFes2 -----> CuFe1_xs2 + xFe3 + + 3xe- (III.1.10)

CuFes2 ----->

cu 1_xFe1 _yS2 + xcu2 + + yFe3 + + (2x+3y)e- (III.1.11)

17

However, Hillrichs et al. (12) states that "upon

immersing the electrode in the electrolyte, Fe(II) will be

dissolved preferentially" according to equation (III.1.10).

Thus both Warren et al. (11) and Hillrichs et al. (12) agree

on the nature of the electrochemical dissolution of chal-

copyrite at high potentials but they are not in full

agreement on the reactions which occur at lower potentials.

Some apparent anomalies, such as these, in the behaviour of

the mineral are probably due to the inherent physical and

chemical differences between samples from different sources.

III.2 REDUCTION OF CHALCOPYRITE

Electrochemical reduction studies of chalcopyrite in

sulphuric acid based electrolytes lead to the proposal that

the overall reduction can be represented as follows (13):

(2-x)CuFeS2 + (6-4x)H+ + (2-2x)e- ----->

Cu(2-x)S + (2-x)Fe2 + + (3-2x)HzS (III.2.12)

CuFeS2 + (3-2x)cu2 + + (4-4x)e- -----> 2Cucz-x)s + Fe2 +

(III.2.13)

The first reaction is appropriate in solutions con

taining no dissolved cupric ion, and the second to solutions

where it is present (13). It is likely that the intermediate

copper-iron sulphides may form initially and that thin

layers may even be present beneath the djurleite or chalco-

18

cite product. Moreover, additions of either cupric or fer

rous ions cause significant increases in the reduction

currents. The Cuc 2 -x)S product layer formed must be a porous

layer (13). This interesting area is not fully understood

and it should be considered for further studies.

III.3 AQUEOUS OXIDATION OF PYRITE

The aqueous oxidation of pyrite is also an important proc

ess in hydrometallurgically treatment of sulphide minerals.

Singer et al. (14) proposed a model in which both oxygen and

ferric ions, Fe3 +, play an important role in the aqueous

oxidation of pyrite:

Initial reaction:

+02

FeS2 (s) ------------> Fe2 + + s-compounds

Propagation cycle:

(slow)

Fe2 + + o2 (aq) ---------> Fe3 +

(III.3.14)

(III.3.15)

I\ I I I

Fe3 + + FeS2 (s) --------> Fe2 + + so42 -

L---------------------------------" (III.3.16)

This model consists substantially of three reactions:

1. The oxidation of pyrite by molecular oxygen or simple

dissociation to ferrous ions, Fe2 +, and s- compounds; reac-

tion (III.3.14). This has been proposed as a necessary

step and has also been supported by Mathews and Robins

( 15) ,

2. The oxidation of ferrous ions, Fe2+, to

19

ferric

ions, Fe3 + by molecular oxygen; reaction

(III.3.15). This is regarded as the rate limiting step,

3. The oxidation of pyrite by ferric ions, Fe3 +,

reaction (III.3.16) considered to be a fast reaction.

Singer et al.(14) also reported that the rate of

oxidation of pyrite, by ferric ions, is independent of the

presence or absence of oxygen but that, in the presence of

oxygen alone (e.g. no ferric ions) no oxidation was ob

served even after one week. Conversely, Mathews and Robins

(15) suggested that the presence of oxygen is likely to

retard the aqueous oxidation of pyrite. It appears that the

presence of oxygen increases the rest potential of pyrite

from 0.25V (in an inert atmosphere) to 0.63V in the presence

of oxygen (15), warranting the term of passivation to de

scribe the extreme irreversibility of the process. Moreover,

if oxidation by ferric ions is retarded, it could be due to

the oxygen being adsorbed on the pyrite surface and reducing

the number of active sites available for the ferric ions

attack.

Thus the overall reactions involved in the aqueous

oxidation of pyrite are( 14-15,18):

FeS2 + 3.502 + H2o -----> Fe2 + + 2S04 2- + 2H+ (III.3.17)

Fe2+ + 0.2502 + H+ -----> Fe3 + + 0.5Hz0 (III.3.18)

FeS2 + 14Fe3 + + 8H20 -----> 15Fe2 + + 2SO 2 -4 + 16H+

(III.3.19)

Fe3 + + 3H2o -----> Fe(OH) 3 + 3H+ (III.3.20)

It is noted that the formation of elemental sulphur

20

was not reported (14-16), however, the partitioning of

pyritic sulphur between sulphate and elemental sulphur,

during the aqueous oxidation of pyrite, has been shown to be

a function of the potential of the system (16,17). An

increase in the potential favours sulphate formation. More

over, aqueous oxidation of pyrite in sulphuric acid media at

110°c and partial pressure of oxygen of 0.4 MPa (18) and

6.28 MPa (16), slowly produce acid when the free acid con

centration is less than about 0.15-0.17 M; and slowly con

sume acid at higher concentrations (16,18,19). It is consid

ered that, at low acid concentrations, oxidation of ferrous

ions and the subsequent hydrolysis and precipitation of

ferric ions are fast reactions. Thus the over-all mechanism

of aqueous oxidation of pyrite has been stated to be a

combination of two competing electrochemical reactions,

equation (III.3.21) and (III.3.22) (16):

FeSz + 8Hz0 -----> Fe2 + + 2S042 - + 16H+ + 14e

(III.3.21)

and

Fes2 -----> Fe2 + + s 0 + 2e- (III.3.22)

and the overall anodic reaction of one mole of pyrite

could be:

FeSz + 4yH20 ----->

Fe2 + + yso42 - + (2-y)S0 + 8yH+ + (2+6y)e- (III.3.23)

21

the cathodic reduction of oxidants are:

Fe3 + + e- ------> Fe2+

(III.3.24)

(III.3.25)

Bailey et al. (16) states that "the overall stoichiom

etry of the pyrite decomposition reaction can be written in

the form:

FeS2 + (0.5 + 1.Sy + 0.25x)02 + (2+x-2y)H+ -----> (1-x)Fe2 + + xFe3 + + (2-y)S0 + yS04

2 - + (1-y+O.Sx)H20

(III.3.26)

in view of the absence of any other observable reaction

products."

Finally, it was shown (16) that during the aqueous

oxidation of pyrite (in sulphuric acid media, at tempera

tures of about 110°c and high oxygen partial pressure) the

sulphate formed from pyritic sulphur contains oxygen taken

from the water rather than from the high pressure oxygen

phase. This suggestion conflicts with the mechanisms previ

ously proposed, where the principal reaction is a heteroge

neous one between dissolved oxygen and the solid surface to

form the sulphate product.

III.4 SUMMARY

From this section, it is important to note that:

1. The leaching process of chalcopyrite and pyrite

22

are explained in terms of oxygen adsorption and electrochem

ical processes. Chalcopyrite produces two intermediate

products during its anodic dissolution, bornite and covel

lite. However, it is not clearly understood whether the

oxygen in the sulphate ion formed proceeds from the dis

solved oxygen or from the water.

2. The leaching kinetics of chalcopyrite tend to

maintain a constant acidity. This corresponds to a condition

where the rate of acid consumption approximately equals the

rate of acid production. It is understood that this buffer

ing action is the result of a dynamic equilibrium between

elemental sulphur and sulphate ion. At 90°c and 3.1 MPa

oxygen pressure the buffer pH during the leaching of chal

copyrite appears to be 2.2. Leaching of pyrite at 110°c by

oxygen under pressure slowly produce acid when the free acid

concentration was less than 0.17M H2so4 and slowly consume

acid at higher concentrations.

3. The leaching products of chalcopyrite and pyrite,

in terms of elemental sulphur or sulphate, depends on the

potential of the system. At low potentials (<0.5-0.7V), it

appears that elemental sulphur is the main product. At

higher potentials than 0.7V, sulphate ions is the main

product.

23

CHAPTER IV

AQUEOUS OXIDATION OF PYRRHOTITE MINERALS

IV.1 INTRODUCTION

In 1935 Mellor (20) defined the basic chemistry of

the iron sulphide minerals in aqueous solutions as a contri

bution to the general interest in iron and its compounds.

This interest, particularly in the basic chemistry of the

oxidation of iron sulphides, decreased after the beginning

of this century when new alternative sources and processes

for the production of sulphuric acid were found {e.g. zinc

and copper processes which produces sulphur dioxide), and

iron sulphides became to be considered a nuisance, producing

environmental hazards. However, since 1980 it appears that

research, at least in the field of copper hydrometallurgy,

is changing from process developments to fundamental re

search.

A number of chemical reaction sequences have been

proposed for the hydrometallurgical oxidation of pyrrhotite

minerals. Acid decomposition, with and without the addition

of oxidants, and alkaline decomposition, at low and high

temperature, are possible methods of processing pyrrhotite

minerals in aqueous solutions. Bacterial oxidation is usual

ly considered as a separate area of study.

Measurement of the rate of acid decomposition of

pyrrhotite minerals has been usually obtained by the deter-

24

mination of iron in solution; very often as ferrous ions,

Fe2+, but since hydrogen ions are consumed during the reac

tion, measurement of the changing pH of the solution has

also been used as a means of obtaining information about the

kinetics of the process.

IV.2 ACID DECOMPOSITION OF PYRRHOTITES

The simple and widely accepted chemical reaction of

pyrrhotite with acid solutions, is represented by the fol

lowing equation (20-30):

+ F 2+ FeS(s) + 2H -----> e (aq) + H2S(g) (IV.2.1)

According to this equation pyrrhotite, usually repre-

sented by the stoichiometric compound troilite (FeS) for the

sake of simplicity, yields iron in solution (most probably

as ferrous ions Fe2 +) and hydrogen sulphide, H2S(g)· Most of

the naturally occurring pyrrhotites are not stoichiometric

and it appears that they yield elemental sulphur propor

tional to their non-stoichiometry according to the following

equation (21,22):

FeS(l+x) + 2H+ -----> Fe2 + + H2S + xs0 (IV.2.2)

For example, acid decomposition of monoclinic pyrrho

tite has been reported to follow equation (IV.2.3) (21,22):

25

(IV.2.3)

If monoclinic pyrrhotite is leached the stoichiometry

suggests that 12.5% of the sulphur would appear in the

elemental form while the remaining 87.5% would be expected

to be evolved as hydrogen sulphide. Ingraham et al.(21)

obtained about 13.3% elemental sulphur during the leaching

of a mainly monoclinic pyrrhotite sample and 12% elemental

sulphur from a hexagonal pyrrhotite sample.

Another source of pyrrhotite, sometimes known as

'activated pyrrhotite' is produced by thermal decomposition

of pyrite. The purpose of this process, at least theorical

ly, is to eliminate one atom of sulphur from the pyrite

molecule, at temperatures of soo-aoo0 c, to produce the

artificial pyrrhotite.

heat

-----> FeS + s (IV.2.3)

Acid leaching of this product produces elemental

sulphur instead of sulphuric acid (22). Some investigators

(22-30) have also studied the acid decomposition of an

artificially prepared pyrrhotite.

In general, most authors have clearly identified the

mineral(s) present in their samples, while others have only

used the sulphur/iron ratio to characterize the sample.

These experimentally determined ratios do not usually fall

into the composition range of any known pyrrhotite mineral

or phase. Hamilton et al.(17) for example studied a natural

pyrrhotite having a sulphur to iron ratio of 1.13; although

26

this ratio is close to that of monoclinic pyrrhotite (1.143)

it does not fall into the composition range of any pyrrho

tite mineral. It must therefore be assumed that the samples,

whether they are natural or artificial, are very often a

mixture of minerals or phases.

The reactivity to acid decomposition of one mixture

is probably different to any another, the bulk chemical and

kinetic stability might also change. It has been reported

(23) that a low excess of sulphur in the structure of pyr

rhotite results in a higher rate of dissolution than a

mineral with a higher sulphur excess. Harris et al. (79)

found that different samples of pyrrhotite from the same

source show significantly different reactivities to aqueous

oxidation. Bugajski et al.(24) suggested that the chemical

composition of the sulphide has a large influence on their

properties, consequently their results (discussed further)

were proposed to be applicable only to monoclinic pyrrho

tite.

IV.2.1 STOICHIOMETRY AND REACTION ORDER

Table IV.1 summarizes the kinetic data which has been

reported for the acid decomposition of pyrrhotite. Although

this is not an exhaustive list, it does emphasize the varia

bility of kinetic observations made by various authors on

more or less similar materials.

27

TABLE IV.1

KINETIC DATA FOR ACID DECOMPOSITION OF PYRRHOTITE

Kind of Acid Reaction Temp. Activ. Rate Control Ref.

sample Mole Order

Precipit- HCl [HCl] 25

ate (0.003-0.1)

Natural H2so4 [H2so4 J1 · 3 30-80

co.2s-1)

Natural HCl [HCl]o. 9 30-90

(5-36%)

Natural HCl [HCl]

and artif. (0.01-1.8)

prepared

30-80

Artificial H2S04 [H2S04l 20-90

troilite (pH 2-5)

Natural 40-90

monoclinic NaClo4

Energy

Kcal/mol

13.2

7.0

9.8

chemical

diffusion

chemical

14.3 chemical

14.0 chemical

pyrrhotite {[H+]=0.1,[Na+]=0.9 [Clo4-]=1.0 mol/kg.}

28

26

21

23

30

24

The reported reaction order, with respect to acid

concentration, is unity in almost all cases. Applying the

general concept of mass transfer, the steps involved in

acid decomposition of pyrrhotite could be the transport

hydrogen ions from the bulk solution to the surface of

the

of

t~.

solid, adsorption at the interface, formation of an activat

ed complex, decomposition of the complex to products, de

sorption of the products at the interface, and the transport

of products from the surface. In this model hydrogen ions,

H+, adsorb at the surface on the anionic sites possibly

forming a hydrogen bond with sulphide ions, s2-, this then

decomposes to HS- which in turn reacts rapidly with hydrogen

ions, to form H2s (30).

During the last twenty years two suggestions have been

made about the rate limiting step during the acid decomposi

tion of pyrrhotite. In the first, Ingraham et al. (21) in

1972 suggested that the first half of the dissolution proc

ess was a diffusion controlled process. In the second, it

was thought that an activated complex, [FeS-2H+], was

formed (25). Moreover, it was re-stated by Tewari et al. in

1976 (30) and more recently by Bugajski et al. (24) in 1982

that the rate limiting step during the dissolution of pyr

rhotite was controlled by a chemical reaction. From stirred

reactor experiments, Ingraham et al. (21) stated that:

''the first half of the dissolution process is proba

bly controlled by the diffusion of some species through a

liquid layer that is adjacent to the mineral surface. The

latter part of the reaction is almost certainly influenced

29

by presence of the sulphur layer."

This suggestion is based on the small activation

energy obtained (7 Kcal/mol) and on the observation that the

reaction rate depended on the half power of the stirring

speed.

Yazawa et al. (26} suggested that the acid decomposi

tion of pyrrhotite is a chemical controlled reaction. Their

evidence for this conclusion includes:-

a} The activation energy is 12.3 Kcal/mol, which is

larger than that for a diffusion controlled process,

b} the dissolution rate is approximately proportional

to the surface area of mineral and the molarity of sulphuric

acid; although this would be true of both an interface and a

transport controlled reaction, and

c} the dissolution rate is not increased by faster

stirring of the solution.

Moreover, Bugajski et al. (24} also indicated that the

rate of dissolution of monoclinic pyrrhotite remained con

stant when the angular velocity of a disc electrode was

varied. This implies that the dissolution, according to

reaction (IV.2.1), occurs in the kinetic region, where the

rate of transport of reactants in the solution is at least

one order of magnitude greater than the rate of chemical

reaction, e.g. the reaction is chemically controlled or

controlled by diffusion through a solid reaction product.

It has to be pointed out that these conflicting

sults are not fully understood. However, it is noted

the conditions of the experiments were different in

30

re-

that

both

cases. While Bugajski et al.(24) used a rotating disc method

in an oxygen free aqueous solution, Ingraham et al. (21)

used a stirred reactor containing a minus 150-plus 200 mesh

particle size of natural pyrrhotite. Complete dissolution

of a 2.0 g. sample was obtained in 15 seconds in 500 ml. of

20 per cent hydrochloric acid at a stirrer speed of 600

r.p.m. and 9o0 c.

Ingraham et al. (21) are the only authors who have

reported elemental sulphur as a product formed during the

dissolution of monoclinic and hexagonal pyrrhotite in strong

acid conditions. Tewari et al. (30) considered that ferrous

ions and hydrogen sulphide are the major products during

the dissolution of artificial troilite in the pH range of

3.40-5.32. In this last study (kinetics of troilite dissolu

tion at one atmosphere pressure of hydrogen sulphide) it was

assumed that oxygen or any other oxidising agent present in

the solution was likely to oxidize dissolved hydrogen sul

phide to sulphur.

Aqueous oxidation of pyrrhotite at a pH of 4.6 may

also produce sulphate ions. Hamilton et al.(17) interpreted

the voltammograms of fresh surfaces of natural pyrrhotite

(Fes1 . 13 stoichiometry) at potentials both above and below

0.2V (SHE).

At potentials below 0.2V and at 4.6 pH, it was sug

gested that the anodic reactions form ferrous ions, elemen

tal sulphur and sulphate ions according to the following

equations (17).

31

Fes1 . 13 -----> Fe2 + + 1.13S0 + 2e

FeS1_13 + 4.52H20 ----->

(IV.2.4)

Fe2 + + 1.13S042 - + 9.04H+ + 8.78e- (IV.2.5)

At potentials higher than 0.2V and at 4.6 pH, the

stable iron species formed are ferric hydroxide, elemental

sulphur and sulphate ions. The inhibition to oxidation was

stated to be due to the formation of ferric hydroxide on the

surface of the mineral.

Fesl.13 + 3H20 -----> Fe(OH)3 + 1.13S0 + 3H+ + 3e

(IV.2.6)

Fes1 _13 + 7.52H2 ----->

Fe(OH) 3 + 1.13S042- + 12.04H+ + 9.78e- (IV.2.7)

However, it is pointed out (17) that at 4.6 pH, sul

phide ions, s 2-, are oxidised predominantly to elemental

sulphur rather than to sulphate ions. The formation of

ferric hydroxide is described as:-

Fe2 + + 3H20 -----> Fe(OH) 3 + 3H+ + e- (IV.2.8)

Thus it is noted that the aqueous oxidation of artifi

cial troilite in the pH range of 3.4-5.32 (30), in an hydro

gen sulphide environment, yields hydrogen sulphide rather

than elemental sulphur. Conversely, voltammographic studies

at potentials below 0.2V (SHE) and pH 4.6, in an inert

environment (nitrogen), showed that elemental sulphur is the

32

predominant product. It can be seen that under the same pH

conditions, the reaction products are affected by the envi

ronment in the system as would be expected.

It has been reported by Jibiki (23), G. van Weert et

al. (25) and Yazawa et al. (26) that acid decomposition of

pyrrhotite is characterized by an induction period during

which no hydrogen sulphide is evolved, although some iron

dissolved slowly. This induction period is reduced by in

creasing the temperature and the acid concentration of the

system. Jibiki (23) stated that:

"it was observed that the potential of the leaching

solution sharply decreases when the induction period ends."

Different suggestions have been made to explain the

cause of the induction period. For example it was suggested

by Yazawa et al. (26) that the induction period is a result

of prior oxidation of the mineral surface forming a rela

tively insoluble oxide film. They found a correlation be

tween the length of the induction period and the amount of

oxygen contained on the surface of the pyrrhotite. The

oxygen content was measured by weighing the amount of water

formed during the hydrogen reduction of the partially oxi-

dized pyrrhotite. G. van Weert et al.(25)

the preferential dissolution of magnetite,

proposed that

which existed

with their pyrrhotite sample, caused the induction period.

It should be noted that an induction period is not always

observed (23).

Scott et al.(27), Nicol et al.(29) and Nicol (30)

suggested that the induction period and the nature of the

33

rate limiting step during the dissolution of pyrrhotite can

best be interpreted by the use of an electrochemical model

rather than by the various adsorption theories previously

proposed. As these latter authors, except Ingraham et al.

(21), assumed that the dissolution rate of pyrrhotite is not

controlled by mass transfer processes, they reported that

the rate of dissolution must be limited by the transference

of the reactants species across the electrochemical-poten

tial barrier at the solid-solution phase boundary, a process

called 'ionic charge transference'. In other words, the rate

of dissolution must be controlled by a surface reaction in

the absence of protective surface-reaction products.

The electrochemical process stated by Nicol (30)

assumes that pyrrhotite minerals (non-stoichiometric slight

ly iron deficient compounds) dissolves through the following

mechanism:-

"only an FeS compound that is exactly stoichio

metric will dissolve in acid solutions by the transfer of

ferrous and sulphide (S2 _) ions across the

interface".

The following equations were found consistent with his

experiments:

Fe1-xS + 2xH+ + 2xe- -----> (1-x)FeS + XHzS (IV.2.9)

(1-x)FeS + 2(1-x)H+ -----> (1-x)Fe2 + + (1-x)H2s

(IV.2.10)

The rate of dissolution of various synthetic and

natural pyrrhotite minerals are dependent on the potential

34

of the sulphide surface. Figure IV.2 shows the rate of

dissolution of hexagonal and stoichiometric pyrrhotite, as a

function of potential, in 0.1 M HCl and 2s0 c as reported by

Nicol and Scott {29) and Nicol {30). It is noted that a

"triangular-wave potential sweep generator connected to a

potentiostat and a cyclic voltammogram" was used for the

experimentation. The small steady-state cathodic currents

were converted into equivalent rates {assuming two-electron

reduction reaction). The more important aspects of these

results were summarized by Nicol {30):-

"i. in each case, the rate increases rapidly with

decreasing potential with a maximum value at about

-0.3 to -0.4 {SCE) {-0.054V to -0.154V vs SHE).

Below this potential decreases slowly with decreasing poten

tial.

ii. The direct proportionality between the meas

ured dissolution rates and the cathodic currents is apparent

for both sulphides. The ratio of dissolution rate to

that of the two-electron reduction process was shown to be

consistent with equations 5 and 6 [{IV.2.9) and {IV.2.10)].

iii.The open circuit dissolution of the iron

sulphides was found to depend strongly on the potential

as would be expected. Thus, the addition of small amount of

reducing agents {including sulphide ions) resulted in in

creased rates whereas, the presence of some oxidising agents

significantly reduced the rate of dissolution.

iv. It was further demonstrated that under free

ly- dissolving conditions, the oxidation of hydrogen sul-

35

-6

FeS -7

-• • N

'n ... i -a .. ~

a ... Clthodlc current 0 ... FeS

-9

-100:-----:1..;---:i..:----:L.::--~~--:J~__,-__, -0. 1 -0.2 -0.3 -o... -0.5 -0.6 -o. 7

Potential V vs SCE

FIGURE IV.2. The rates of dissolution of iron sulphides in 0.1 HCl at 25 C as a function of potential. Cathodic currents were calculated from the two-electron reduction process shown in Equation (IV .2. 9) (29).

phide to elemental sulphur is the process that produces the

electrons required by reaction 5 [equation (IV.2.9)]"

An independent study (23), supports the observation

that the presence of oxygen, potassium dichromate, eerie

sulphate, potassium permanganate and hydrogen peroxide slows

the dissolution rate of pyrrhotite to an almost negligible

rate, though the critical concentration of each oxidant

(above which the effect of oxidant first appears) depends on

the particular oxidant, the acid concentration and the type

of pyrrhotite.

IV.2.2 SUMMARY

The acid decomposition of pyrrhotite is thought not to

be controlled by the transport of hydrogen ions, because

firstly the dissolution rate depends on the composition of

pyrrhotite, and secondly the estimated rate of hydrogen ion

diffusion controlled reaction is approximately

mole/cm2 .sec for a hydrogen ion concentration of 1 M (23)

and the maximum rate for pyrrhotite dissolution under the

same conditions is reported to be:

2.7 x 10-7 mole/cm2 .sec ( 1M HCl) Tewari et al. (30)

4 x 10- 7 mole/cm2 .sec [at -0.2 V (SCE), 1M HClO~] oJ

Nicol et al. (28)

10-9 mole/cm2 .sec (1M HCl) Jibiki (23)

1.095 x 10-6 mole/cm2 .sec (1M HCl) Ingraham et al.(21)

36

The transport of the reaction products, which are

predominantly ferrous ions and sulphide species of HS- or

H2S, also do not seem to control reaction rate of the proc

ess, since the ferrous ions concentration does not affect

the dissolution rate up to 0.1 M concentration and hydrogen

sulphide in solution does not slow the dissolution rate as

would be expected in a reaction controlled by the transport

of products from the reaction zone to the bulk solution.

Therefore, the dissolution of pyrrhotite must be controlled

by a chemical or an electrochemical reaction process at the

mineral surface.

When a sulphide has a very high decomposition rate, as

occurred in the study of Ingraham et al.(21), dissociation

of the solid may occur first with the subsequent reaction

between hydrogen ions and sulphide ions in the vicinity of

the solid surface (23). In this case, the reaction between

hydrogen ions and sulphide ions is, in general, fast because

of a homogeneous reaction. Therefore the transport of hydro

gen ions from the bulk solution to the sulphide ions or the

transport of reaction products from the vicinity of the

surface of the solid to the bulk solution will be the slow

est step for the reaction.

37

IV.3 OXIDATION OF FERROUS IONS, Fe(II), BY MOLECULAR

OXYGEN

Studies on the oxidation of ferrous sulphate

tions, made in the late 19th century and the early

century, were reviewed extensively by Mellor in 1935

solu-

20th

(20).

This

that

review and further work, discussed below, established

the rate of oxidation of ferrous ions depends on pH,

reaction media, ferrous ion concentration, molecular oxygen

concentration, temperature and the presence of certain

catalytic materials. Of these, the rate of oxidation of

ferrous ions seems to be more sensitive to pH and tempera

ture, thus the oxidation process may be classified, for

analysis, into two arbitrarily selected ranges of acid

concentrations:-

1. Strong concentrations of acid solutions of less

than pH 0, and

2. Acid solutions having a pH in the range Oto 7.

This analysis will concentrate mainly on solutions of

sulphuric acid and will attempt to highlight the variation

of the reaction rate when the acid concentration is changed

both at low and elevated temperatures. Oxidation of ferrous

ions at pHs greater than seven will not be discussed because

this study, the aqueous oxidation of pyrrhotite minerals,

will be conducted only in acid conditions.

38

IV.3.1 OXIDATION OF FERROUS IONS, Fe(II), IN STRONG

SULPHURIC ACID SOLUTIONS

The overall chemical reaction which takes place in

strong acid solutions, where Fe(III) remains soluble, is

described by the following equation (IV.3.1) (14,31-34):-

4Fe2+ + 02 + 4H+ -----> 4Fe3 + + 2Hz0 (IV.3.1)

The kinetics of oxidation of ferrous ions, Fe(II), in

this range of acid concentration has been reported to pro

ceed very slowly at ambient temperatures c~2s0 c) (14,31-34).

It has also been pointed out that the rate of reaction is

practically independent of acid concentration, second order

with respect to ferrous ions, Fe (II), and first order with

respect to the partial pressure of oxygen. In one molar

sulphuric acid solution the reaction, at 3o.s0 c, is de

scribed by the following rate equation (31):-

where Kt = 2.78 x 10-6 M- 1 atm.- 1 sec.-1 .

Since oxygen is only sparingly soluble in water,

supplying dissolved oxygen may be a problem in certain

aqueous oxidation processes involving sulphide minerals, if

oxygen is consumed by both homogeneous (e.g. oxidation of

ferrous ions) and heterogeneous processes (aqueous oxidation

of the sulphide mineral itself). Saturation with air at

39

ambient temperature provides only 6 or 7 mg.L- 1 of oxygen

(35).

The solubility of oxygen in water and in dilute acid

solutions (in moles of oxygen per liter of solution) may be

approximated by the use of Henry's Law:-

* [O laq. = 55.5 p02 / 14.7 (H - p02 > (IV.3.3)

H = Henry's Law constant value.

Biernat (32) (using this equation) calculated the

solubility of oxygen in sulphuric acid solutions and showed

that it decreases with increasing acid concentrations.

Similar decreases in the oxygen solubility have been shown

to result from increases in the concentration of ferrous

sulphate (32). The ratio of the solubility of oxygen in an

electrolyte solutions to that in pure water decreases as the

concentration of the electrolyte increases, but is essen

tially independent of both temperature (over the range of

298-348°K) and the oxygen partial pressure (over the range

O . 1-1 . o a tm . ) ( 36 ) .

The rate of oxidation of ferrous ions at elevated

temperatures is higher than that at ambient temperatures as

expected (31,33). Hoffman and Davidson (31), in 1956, point

ed out that the rate of oxidation, by molecular oxygen at

413 - 453°K, is very dependent upon the nature of the anions

present. At a given pH, the rate was found to increase in

the series perchlorate, sulphate, chloride, phosphate and

pyrophosphate. It is noted that the rate of oxidation of

40

ferrous ions, Fe(II), is higher in a chloride system than in

a sulphate system.

The oxidation of ferrous ions, at elevated tempera

tures, also proceeds by two independent paths (31,33). The

reaction rate at ferrous ions concentration of between 0.001

and 0.025 Min 1 M of sulphuric acid and at temperatures of

140-1So0 c can be expressed in terms of simultaneous bimolec

ular and termolecular paths (31):-

-d[Fe2 +] I dt = kb[Fe2 +].Po2·exp(-13,400/RT)

+ kt[Fe2 +] 2 .Po2·exp(-16,200/RT) (IV.3.4)

where Kb= 1.93 x 10-5 atm.- 1 sec.- 1 and

kt = 1.60 x 10-3m- 1 .atm.- 1 .sec-1 at 159°c.

and these "may be functions of sulphate and hydrogen

ions concentrations as well" (31).

Iwai et al.(33), in 1982, characterized more accurate

ly the oxidation of ferrous ions with dissolved molecular

oxygen identifying the effect of sulphate ions on the reac-

tion rate. This effect was identified in sulphuric acid

solutions containing 0.2 mol.L- 1 of ferrous sulphate "with a

sulphate ion concentration range below 0.01 mol.L- 1 ,

which approximately corresponds to a sulphuric acid

concentration above 0.4 mol.L- 1 ".

The oxidation reaction, in solutions of sulphuric acid

above 0.4 mol.L- 1 , was found to proceed through two parallel

paths. For the first path the reaction rate is independent

of the sulphate ion concentration, while both paths are

41

second order with respect to ferrous ions concentration and

first order with respect to the partial pressure of oxygen.

Thus, at 303.5°K, the kinetics in this range of acid concen

tration was described by

-d[Fe(II)] I dt = K1[Fe2 +1 2 .Po2·exp(-51,600/RT)

+ K2[S042-][Fe2 +1 2 .po2·exp(-144,600/RT) (IV.3.5)

where K1 = 9.1 x 10-10 .mol- 1 .L.Pa- 1 .min- 1

K2 = 1.61 x 10-7 .mo1- 2 .L2Pa- 1 .min- 1 .

The role of the sulphate ion, in the second reaction

path, is in controlling the actual concentration of the

ferrous ion reacting species and is related to other ferrous

sulphate complexes by their formation constants and the

dissociation constant of Hso4+ (33). Assuming that only Fe2 +

(aq), Feso4°, and FeHS04 are the principal species, the

predominance of ferrous sulphate species (Feso4°) was

calculated by Iwai et al.(33) from the following equations:

K, 3o.s0 c

Fe2 + + so 2 -4 = Feso4° K1 = 7.26 (IV.3.6)

Fe2 + + HS04 + = FeHS04

+ K2 = 1.67 (IV.3.7)

HS04 - H+ + so 2 - K3 = 0.0086 (IV.3.8) = 4

K1 = [Feso4°] I [Fe2 +] [S04 2+]'

K2 = [FeHS04 ] I [Fe2 +J [HS04 +],

K3 = [ H+] [So42-J I [HS04 -] .

Fe2+ (aq) = ferrous aquo-ion

42

These formation constants confirms the predominance

of ferrous sulphate, Feso4°, compared to the ionic species

under acid conditions.

Iwai et al.(33) suggested that the oxidation of fer

rous ions, in more dilute sulphuric acid solution, below 0.4

mol.L- 1 , also proceeds through a different reaction mecha-

nism. The reaction rate was second order with respect to

ferrous ions concentration and first order with respect to

the oxygen partial pressure. Furthermore, the reaction rate

was affected by sulphate and hydrogen ions. The rate clear

ly increased with a decrease in the hydrogen ion concentra

tion. However, Iwai et al. (33) stated that

"further detailed examination of the role of hydrogen

or sulphate ion in the oxidation of Fe(II) in dilute

sulphuric acid solution was not possible because of the

difficulty in manipulating experimental conditions''.

It is thought that the concentration of sulphate ion

is very sensitive to small changes of pH at concentrations

of sulphuric acid just below 0.4 mol.L- 1 containing 0.2

mol.L- 1 ferrous sulphate. Thus Iwai et al. (33) found it

difficult to maintain a constant concentration of sulphate

ions over a range of ferrous concentrations.

IV.3.2 OXIDATION OF FERROUS IONS, Fe(II) IN LOW

CONCENTRATIONS OF ACID SOLUTIONS

Oxidation of ferrous ions, Fe(II), in more dilute

solutions of sulphuric acid, (0-3.5 pH) and at temperature

43

of 2s0 c is slow (14,14,37,38) and the overall chemical

reaction proposed to take place up to a pH value about 3.5

appears to be the same overall reaction proposed in equation

(IV.3.1).

Although the reaction at pH values 0-2 also proceeds

along two paths (37), the reaction mechanism suggested in

this pH range differs from the reaction mechanism in strong

er acid solutions. The reaction mechanism suggested by Iwai

et al.(33), for the oxidation of ferrous ions, Fe(II), in

stronger solutions of sulphuric acid than pH zero is com

pared with the reaction mechanism proposed by Mathews and

Robins (37) at 0-2 pH.

Path 1 describes the proposed reaction mechanism when

the reaction rate is independent of sulphate ions in

stronger acid conditions than pH zero (33):

Path 1

Fe2 + + 02

Fe2 +o + Fe2 + 2

K1

<=====> Fe2 +o2

<=====>

fast

[Fe2 +-0=0-Fe2 +1 -----> [Fe3+-o-o-Fe3 +]

fast

(IV.3.9)

(IV.3.10)

(IV.3.11)

[Fe3 +-o-O-Fe3 +1 + 2H+ -----> [2Fe3 + + H202] (IV.3.12)

On the other hand, ferrous and sulphate ions may

associate themselves as the Lewis acid and base, respective-

44

ly, thus a Fe(II) sulphate-complex (Feso4°) may be the

reacting species in the rate determining step of the sul

phate ion dependent reaction path (31,33).

Path 2

Fe2 + +

ko

<=====> Fe2 +o 2

so 2 -4

kl

<======> Feso4° k'

(IV.3.13)

(IV.3.14)

fast

[Fe2+-0=0-FeS04°J----->[Fe3 +-o-O-Fe3 +.so42-] (IV.3.16)

fast

[Fe3 +-o-o-Fe3 +.so42-J + 2H+ ---->2Fe3 + + so4

2 - + H2o2

(IV.3.17)

Conversely, Mathews and Robins {37) suggested that

FeoH+ were the reacting species in the pH

range of 0-2. Although ferrous ions hydrolyzes to produce

FeoH+ and other arrays of mononuclear species at pH 7, their

stabilities are known with less precision than for other

common ions {49). The mechanisms proposed by Mathews and

Robins {37) are:

45

Path 1

k1

Fe2 + + o 2 <=====> Fe2 +.o22 +

k' 1

Fe2+.o22 + + Fe2 + +H2o----->

Path 2

Fe2 + + Oz <======> Fe2 +o2

Fe2 + + HzO -----> FeOH+

(IV.3.18)

2Fe3 + +OH-+ 0 H-2

(IV.3.19)

(IV.3.20)

(IV.3.21)

2Hz0 + Fe.022 + + FeOH+ -----> 2Fe3 + + SOH-

(IV.3.22)

(IV.3.23)

(IV.3.24)

The reaction paths suggested by Iwai et al.(33) and

Mathews and Robins (37) are not expected to follow the same

mechanism since the reactions are taking place at different

acid concentrations. The rate limiting step for both ranges

of pH values appears to be the formation of the Fe2 +o 2

complex. While the other rate limiting step in the pH range

below zero appears to be the formation of ferrous sulphate,

Feso4°, the other r.l.s. at pH values 0-2 appears to be the

formation of the Fe2 +.o22 + complex. Thus the rate equation

for the oxidation of ferrous ions, Fe(II), (Equation IV.3.1)

at pH values 0-2 appears to be more characterized by the

following rate expression (37,40):-

46

where K1 =

K2 = R =

-d [Fe2 +]

4dt

1.32 X 1011

1.76 X 104

d [Fe3 +]

4dt

M.min- 1

1.987 cal.gmol- 1 .°K- 1 .

(IV.3.25)

This kinetic expression is in good agreement with the

results obtained by Keenan (40). The main characteristics

of the rate equation is its approximately second order

dependence with respect to ferrous ions, first order depend

ence with respect to oxygen concentration and on hydrogen

ion concentration to an small negative exponent (-0.35).

The form of the oxygen dependence in the kinetic

expression reported by Mathews et al.(37) and Keenan (40) is

different to that of previous investigators. In all previous

studies, the rate was related to oxygen partial pressure

instead of the dissolved oxygen concentration. Mathews et

al.(37) and Keenan (40) suggested that dissolved oxygen

concentration should be used and further showed that under

conditions of high reaction rate, the actual dissolved

oxygen concentration is considerably below the equilibrium

saturation level implied by the use of oxygen partial pres-

sure.

Although only Singer et al.(14) has reported studies

on the kinetic of oxidation of ferrous ions in the pH range

from 1 to 7, the reaction rate has been measured for pH

47

2.0

Experimental points in low pH range

1.0 - d log I Fe(II) I

k" : dt

0 Po2 : 0.20 atm.

Temp. 2s·0c -1.0

7 - -2.0 >-n, "C 0 - I ... I ~ -3.0 I CTI I £ I

t:, I -4.0 6 / Extrapolation of

/ rate law:

-5.0 l d!Fe(II)] : k!F •2]!owJ2 R / dt e 02

I

-6.0 1 2 3 4 5 6 7

pH

Figure IV.3.1: Oxidation Rate of Ferrous Ions as a Function of pH (14).

values below 3.5 and for pH values greater than 4.5. The

reaction rate at pH values less than 3.5 is similar to that

formulated in equation (IV.3.2).

The mechanism of oxidation of ferrous ions in the pH

range from 3.5 to 4.5 does not appear to have been studied

by any investigator. The mechanism of reaction in this pH

range seems to be more complex since the reaction rate

increases rapidly. Figure IV.3.1 shows the dependence of

the rate of oxidation of ferrous ions on pH in the range of

1-7 as reported by Singer et al.(14).

At pH values greater than about 4.5, the overall

chemical reaction appears to be described by the following

equation (14,38,39):

Fe2+ + 0.2502 + 20H- + 0.5Hz0 -----> Fe(OH)3 (IV.3.28)

Minegishi et al.(38) pointed out that the overall

process comprises the sequential steps of the dissolution of

gaseous oxygen at the surface of rising bubbles and the

oxidation of ferrous ions by dissolved oxygen. Thus a better

approach would be to consider both processes:

0 2(g) ----> 0 2(aq)

Fe2+ + 0.25 Oz(aq) + 0.5 H2o =

(IV.3.29)

Fe(OH)~ ~

(IV.3.30)

Although the oxidation of ferrous ions at a constant

pH between 4.7-5.5 (38) and 6-7 (39) appears also to proceed

48

along two paths, one homogeneous reaction in the solution

and the other heterogeneous reaction on the surface of

ferric hydroxide precipitate; these reactions and the reac

tion rates were expressed differently,

a) homogeneous reactions:

Fe2 + + 0.25 Oz(aq) + 20H- + 0.5 HzO -----> Fe(OH)3

(38) (IV.3.31)

FeZ+ +Oz-----> Fe(III) + Oz- ( 39) (IV. 3. 32)

"the oxygen radical, o2-, is rapidly consumed to form

H02 by oxidising three additional ferrous ions" (39).

b) heterogeneous reactions:

Fe2+(ad) + 0.25 Oz(aq) + 20H- + 0.5 HzO -----> Fe(OH)3

( 38) (IV. 3. 33)

Fe2 +(ad) +Oz-----> Fe(III) + Oz- ( 39) (IV. 3. 34)

Where FeZ+(ad)• represents the adsorbed ferrous ions

on the surface of ferric hydroxide. Thus the catalytic

effect of ferric hydroxide, Fe(OH) 3 , on the oxidation of

ferrous ions was also expressed differently as follow,

Minegishi et al. (38): constant pH in the range 4.7-5.5

- d[Fe2 +] / dt = Ko[Fe2 +] [Oz(aq)J (IV.3.35)

Tamura et. al. ( 39): constant pH in the range 6-7

- d[FeZ+] / dt = K + k' [ Fe (III) ] [Fe2 +] (IV.3.36)

However, while the first rate equation, equation

IV.3.35, was obtained from studies in a sulphate system, the

second rate expression, equation IV.3.36, was obtained from

49

measurements in NaCl04-NaHC03 solution.

At pH values less than 5.0 and temperatures lower than

298°c the concentration of dissolved oxygen appears to reach

saturation immediately after the start of the oxidation

process. At higher pHs and temperatures, the oxygen concen

tration was found to be lower than the saturation value.

This observation led Minegishi et al. (38) to suggest that,

in the first case (pH<5, T<298°K), the overall rate was

likely to be controlled by chemical reactions; while at

higher pH and temperature, when the oxidation rate of fer

rous ions was higher, consequently controlled by both chemi

cal reactions and the rate of oxygen dissolution was in

volved. Figure IV.3.1 shows the dependence of the rate of

oxidation of ferrous ions on pH as found by Minegishi et al.

(38).

50

FIGURE (IV.3.~). Effect of pH on the Rate of Oxidation of

Ferrous Ions. Fe~·. after Minegishe et al. (38;.

IV.4 HYDROLYSIS AND PRECIPITATION OF FERRIC IONS

The hydrolysis and precipitation of ferric ions,

Fe(III), is an important process in the field of hydrometal

lurgy because it is a method for eliminating iron from leach

solutions.

Hydrolysis and precipitation of ferric ions is a

complicated process sensitive to a large number of varia

bles. It is recognized that this process has not been as

thoroughly studied in the sulphate system as it has in the

nitrate, perchlorate or chloride system (8,9,41). Although

ferric ions, Fe(III), complex strongly in sulphate and

chloride system, the chemistry in the sulphate system is

further complicated by the fact that the first dissociation

of sulphuric acid in aqueous solutions at 2s0 c proceeds

completely, [equation (IV.4.1)], while the second dissocia

tion is a moderately weak electrolyte, [equation (IV.4.2)],

(42). Thus it is worth to mention that the bisulphate ion

has a buffer action.

= +

= +

HSO + 4

so 2 -4

(IV.4.1)

(IV.4.2)

The pH at equilibrium between sulphate and bisul

phate ion is 1.99.

The composition and structures of iron (III)

(hydr)oxide precipitates depend on the ferric ion concentra

tion, the nature of the anion present, the pH, temperature

and period of ageing (45,47,48).

The polymers formed in nitrate solutions do not appear

51

to include this in the polymer chain. Whereas the polymers

formed in the chloride solution contain some chloride ions

in place of the hydroxyl ion (46,47)

Knight et al. (44) studied the hydrolysis and precipi

tation of ferric ions, Fe(III), in nitrate, perchlorate and

chloride systems. The time for a visible precipitate to form

during the addition of NaHC03 to iron solutions, having

concentrations in the range of O.Ol-0.3M, was measured at

2s0 c and the solid products formed were examined by x-ray

diffraction and electron microscopy. It was claimed that, in

the nitrate system, the precipitate is goethite, (alpha)

FeOOH, in perchlorate and chloride solutions the precipi

tates are lepidocrosite (beta)-FeOOH and (gamma)-FeOOH,

respectively. At high base concentrations (NaHC03 ), perchlo

rate solutions may produce (alpha)-FeOOH.

The hydrolysis and precipitation steps of iron (III)

at 90°c is followed by subsequent loss of water and internal

crystallization of (alpha)-FeOOH to (alpha)-Fe2o3 in nitrate

solution or by dissolution of (beta)-FeOOH and growth of

(alpha)-FeOOH in chloride solution (47).

Feitknecht et al. (48) measured the solubility con

stants of some metal oxides and hydroxides, including ferric

hydroxide, in aqueous solutions. Considering the case where

there exists only one modification of a certain crystalline

hydroxide or oxide and a sufficient amount of hydroxide ion

is added rapidly, a very fine crystalline precipitate is

formed with a disordered lattice, and this is an active

form of the compound Me(OH) 2 (active) or Me02 ; 2 (active). A

52

metastable equilibrium is established between this active

hydroxide or oxide and the solution, and this equilibrium

changes, more-or-less rapidly, approaching the limiting

value (of the solubility product) for an inactive form as

larger and ordered crystals are formed upon ageing. Thus

solubility product of a fresh precipitate is higher than

that of an aged, inactive form.

If there is only slight supersaturation of the metal

salt solution with hydroxide ions, nucleation and crystal

growth are often slow. Consequently, an inactive form of the

solid may settle out since equilibrium with the solution may

be established only gradually. Under such conditions short

term measurements do not give reliable values of the solu

bility product (48).

Hydroxides can occur in amorphous as well as various

crystalline modifications. The tendency to form amorphous

precipitates increases with the valency of the metal ion.

Amorphous precipitates may also show different activities.

Ageing of ferric hydroxide is strongly accelerated

when the temperature is increased and it may occur through

the following changes: either the active form of the unsta

ble precipitate becomes inactive, or a more stable modifica

tion of the precipitate is formed. Moreover, deactivation

of amorphous ferric hydroxide, Fe(OH) 3 may also be accompa

nied by condensation (formation of a more complex molecule)

and dehydration since ferric oxide, (alpha-Fe203), is more

stable than the primarily precipitated ferric hydroxide,

Fe(OH)~. Thus upon ageing of ferric hydroxide, non-homoge-~

53

neous solids may form and the solubility product, measured

under these circumstances, is usually referred to the most

active component (48):-

I

r----> (am.} Fe0n;2 (0H} 3 _n (inactive} I

(am.} Fe(OH} 3 (active~-----> (alpha}- FeOOH (active}

(am.}=amorphous

\ \

~----> (alpha}- Fe2o3

Feitknecht et al.(48} obtained the solubility

product for the various precipitates and they are reproduced

in Table IV.4.1.

54

TABLE IV.4.1

Precipitate Log S0

arn.Fe(OH) 3 (most active) -38

arn.Fe(OH) 3 (active) -38.7

arn.Fe(OH)? (inactive) -39.1 ~

(alpha)-Fe2o3 . -42.7

55

IV.4.1 FERRIC IRON SPECIES IN SOLUTION

Although a number of valence states of iron are well

known, the potential-pH diagram for the Fe-H2o system

(Figure IV.4.1) and other similar diagrams indicate that

only +2 and +3 oxidation states predominate at pH around 4.5

in the whole range of oxidation potential. Since it might be

assumed that the compounds precipitated are closely related

to the ferric complexes present at the instant of the pre

cipitation, it becomes very important to know which ferric

species exist in the solution.

IV.4.2 HYDROXYL COMPLEXED SPECIES

It is widely known that ferrous ions do not begin to

hydrolyze from an acid dissolution until the pH is in-

creased to around 7, thus an array of mononuclear species

such as FeOH+, Fe(OH) 2 , Fe(OH) 3 - and Fe(OH) 42 - are not

expected to be formed in solution. Conversely, ferric ions

begin to hydrolyze at pH>l (49).

Hedstrom (1953), quoted by Baes et al.(49), identified

various ferric hydroxyl species such as Fe(OH) 2 +, Fe(OH) 2 +

and Fe2 (0H) 24+ at 25°c in 3M sodium perchlorate. These

results have been widely referenced and confirmed by several

workers (18,41,43,44,63,46,50-52,54,55). Bierderman, also

quoted by Baes et al.(49), identified Fe3 (0H) 45 + as a minor

hydroxyl species. Moreover, ultracentrifugation of saturat

ed solutions led to an estimate of the equilibrium constants

56

-2 -1 a t 2 3 4 6 6 7 8 9 10 11 12 13 14 15 16 2,Z 2,2

Y)2 2

1,8 1,8

1,6 1,6

1,4 1,4

1,2 Fi O --? e " . 1,2

t t

0,8 ©----------.Q,8

0,6 0,6

0,4 --.... 0,4 -----0,2 0,2

0 ©-__ ++ 0 -- Fe - -0,2 0,2 -------- -0,4 --

-0,6

0,8 -0,8

-1 -1 I

Fe I

-1,2 cp -1,2

-1,4 -1,4

-1,6 -1,6

-1,s -1,8 -2 -1 0 1 2 3 -4 5 6 7 8 9 10 11 12 13 14 15pH16

F!GURF I\/ .4.1: Potential-pH Diagram for Lho Fe- Hz.0 S~stPm at 2':i°C (BS).

of Fe(OH) 3° and Fe(OH) 4 - hydroxyl species. Although other

hydroxyl species have been reported e.g., (Fe(H2o) 63+,

Fe(H2o) 5 (0H) 2 + and Fe2 (H2o) 8 (0H) 24+) there is still serious

disagreements about their exact chemical compositions and

equilibrium constants (49).

IV.4.3 DEVELOPMENT OF PREDOMINANCE REGIONS FOR FERRIC

HYDROXYL SPECIES

For the purpose of this study, predominance diagrams

showing the various ferric hydroxyl species and the solubil

ity of amorphous ferric hydroxide have been constructed on a

log-activity versus pH diagram. These diagrams were con

structed using software developed in the Department of

Mineral Processing and Extractive Metallurgy of The Univer

sity of New South Wales (95). In the first diagram, Figure

V.4.2, the predominance area of the various hydroxyl corn-

plexes is shown. In the second diagram, Figure V.4.3, the

stability of ferric hydroxide with respect to each ferric

hydroxyl species is given. Appendix D tabulates the recent

thermodynamic data used to construct these diagrams. The

hydrolysis reactions of ferric ions and the equations relat

ing their equilibrium constants and pH are presented in

Appendix E.

It can be seen from these diagrams that each complex

predominate under different conditions of pH and ferric ion

concentration. At low concentrations of hydroxyl species,

e.g. 1-5 mg.L- 1 , and pH values about 4.5, the predominant

57

··- --

0 1:298.15

' 6·· 7 . 1 l

21

3 l

4 l 1 3 4

2 I

5

6

~

::: 7 ..... ~

w ~8

0 1 2 3 4 5 6 7 8 9 10 11 pH

FIGLRE \V. 4 .2: Distribution of Ferric hydroxyl species as a function of pH at 25°C.

·-., 1 Fe+++ ( -1 • 1 l 12 Fel'JH++ (-54. 80) 3 Fe (CIH)2+(-106.74l 4 Fe (l'JHJ 3 CRQl C-154. 79) 5 FeCl'JHl4-C-198.39l 6 Fe2 CCIH) 2++++ (-111. 55)

17 Fe3Ce1Hl4+++++(-221.46l pH:0/0/.25 pHZel:0/0/0

lpFE!llll:0/.25/0

s

12 13 14

0 1:298.15

1 i I

~

3

4 1 5

6

~

::: 7 ...... - I w ~8

1

,.

B

2

3

1 Fe+++(-1.lJ

1

·2 FeClH++(-54.80) 3 Fe !ClHJ2+(-106.74J

14 Fe rnHJ 3 CRQJ C-154. 79) 5 Fe rnHJ 4- (-198. 39) 6 Fe2 (('.]HJ 2++++ (-111. 551

17 Fe3 Ct'lHJ 4+++++ (-221. 46) 8 f•t'lt'lH(RMJ (-109.B2J pH:0/0/.25 pH2t'l: 0/0/0 .

lpFE I I l I J: 0/. 25/0

0 1 2 3 4 5 6 7 B 9 10 11 12 13 14 pH

FIGURE lV .4.3: Distr ib.J tion of ferric hydroxyl species and their stability uith ferric hydroxide, Fe(OH~ (am.) at 25°C.

·-.,

ferric hydroxyl species would be Fe(OH) 2 +.

Table IV.4.2 summarizes the hydrolysis reactions for

ferric ions together with their equilibrium constants with 1

and 3M of sodium perchlorate solutions, or otherwise stated.

58

TABLE IV.4.2

-Log K, 25°c

lMa 3Ma Ref.

Fe(OH)~+ + H+ = Fe3 + + H2o 2.79 3.05 (49) (IV.4.3)

2.88b (52)

3 .16 3.06c (55)

2 .19 (58)

Fe(OH) 2 + + 2H+ = Fe3 + + 2H2o 5.85 6.31 (49) (IV.4.4)

6.85b 5.7c (52)

4.9 (55)

Fe2(0H)2 4+ + 2H+ = 2Fe3 + + 2Hz0 2.72 2.91 (49) (IV.4.5)

3.29b (52)

3.22c (55)

2.95 (58)

Fe3 (0H) 4 5+ + 4H+ = 3Fe3 + + 4Hz0 6.56 5.77 (49) (IV.4.6)

Fe(OH) 3 ° + 3H+ = Fe3 + + 3H2o <12 (49) (IV.4.7)

Fe(OH) 4 - + 4H+ Fe3 + + 4Hz0 <21. 6 (49) (IV.4.8) =

2Fe3 (0H) 4 5+ + 2H+ = 3Fe2 COH) 2

4+ + 2H2o (IV.4.9)

Fe2 COH) 2 + = 2Fe(OH) 2 + -3.19a (IV.4.10)

59

Although most of these species have been characterized

in the perchlorate system, it is assumed that they exist in

the sulphate system without incurring in great error (41).

This assumption was confirmed by Sapieszco et al. (1976)

(55) who studied the thermodynamics of aqueous hydroxo and

sulphate complexes and determined the equilibrium constants

for the three first hydrolysis reactions, given by equations

(IV.4.3), (IV.4.4) and (IV.4.5) over the temperature range

of 25-so0 c. The values of the equilibrium constants obtained

(included in Table IV.4.2) show a fair agreement with the

other values tabulated.

Music et al. (1982) (47) studied the formation of

Fe(III) oxyhydroxides and oxides formed by the hydrolysis of

nitrate, chloride and sulphate solutions and claimed that

"in sulphate solutions the formation of FeS04+

complex suppresses the polymerization process and the

formation of oxyhydroxides and oxides".

A Fe(OH)S04 precipitate is formed, at short times, by

hydrolysis of ferric solutions at 9o 0 c and transformation

occurs forming other relatively complex basic iron (III)

precipitates (brown precipitate) after long

times.

hydrolysis

There are still serious contradictions about the

composition and mechanisms of the hydrolysis process of

ferric iron, Fe(III). The difficulties involved in the

investigation of the composition and structures of iron

(III) hydroxide precipitates have been stated by Music et

al.(47) to be:

60

"i. small differences in the values of the parameters

change the composition, structure, and morphology of

the precipitate, and

ii.the experimental methods used in previous studies,

e.g. potentiometry, ultracentrifugation, electron

microscopy, visible spectrophotometry, IR spectropho

tometry and x-ray diffraction have limitations because

none can follow all the stages of the overall precipi

tation process.

Characterization of the precipitate in the colloidal

dimension range (where the hydroxy polymers, small iron

(III) (hydr)oxide particles of complex composition, and so

called amorphous iron (III) hydroxide form) is the least

susceptible to study. The use of 57Fe Mossbauer spectroscopy

does permit a study of each steps involved in the precipita

tion of iron hydroxides and their subsequent

transformation."

pre

and

In the absence of sulphate ions, hydrolysis and

cipitation of ferric ions, Fe(III), forms oxide

(oxy)hydroxide precipitates. The process is thought to take

place through the following steps (43-47).

i. The formation of low molecular weight species

(dimmers, trimmers),

ii. The formation of colloidal dispersions (sols.) of

ferric hydroxy polycations; and

iii.The formation of precipitates or colloidal

dispersions of various ferric oxyhydroxides.

On the basis of Mossbauer measurements, Music et al.

61

(47) schematically described the formation of solid phase formed

by hydrolysis of ferric nitrate [Fe(N03 ) 3 ] solution at go 0 c as

follows:

OH

Fe3 + <=====> Fe(OH) 2 + <=====> /\ [Fe Fe] 4 + <======> \/

OH

I II

OH OH OH OH

/ '- / n+ [Fe Fe] n/2

'\. / '\ OH OH

/\/ ---->[Fe Fe

\ / \ 0 0

Jn/2 + nH+ ------>

III IV

where "the first step there is formation of simple

hydrolysis products of iron (III), such as monomers

(I) and dimmers (II). The next step is the formation

of iron (III) hydroxy polymers (III). Oxybridges form

with prolonged time of heating or ageing and the

(alpha)-FeOOH structure (IV) develops. The final step

is loss of water and internal crystallization of

(alpha)-FeOOH to (alpha)-Fe2o3 (V)" (47).

IV.4.4 SULPHATE COMPLEXED SPECIES

The relative stability of sulphate and bisulphate ions

is determined by the pH of the solution and the potential-pH

diagram for the Fe-S-H2o system (Figure IV.4.4) shows that

bisulphate ions predominate up to 1.99 pH; thereafter,

62

sulphate ions. Thus ferric sulphate complexing reactions at

4-5 pH may take place predominantly at potentials above

about 0.6 V (SHE) (Figure IV.4.4).

In highly acidic solutions, where Fe(III) remains in

solution, the chemical reactions of ferric sulphate corn-

plexes can be described by the following equations (52,55-

57).

Fe3 + + SO 2- = 4 Feso4 +

Feso4 + + so42 - = Fe(S04 ) 2-

Fe3+ + HS04- = FeHS042 +

Log K,25°c Ref.

( 52) (IV. 4. 11)

1.92b (55)

(57) (IV.4.10)

(55) (IV.4.12)

If the reaction takes place in neutral sulphate solu

tions, the mechanism which describes the precipitation of

ferric hydroxide, would involve the following reactions

according to Sevryukov et al. (1981) (56).

(IV.4.13)

2Fe(OH)S04 + 2H20 = Fe2 (0H) 4 (S04 ) + so42 - + 2H+

(IV.4.14)

4Fe(OH)(S04) + 6H20 = Fe4S04(0H>10 + 3S042 - + 6H+

(IV.4.15)

4(H30)Fe(S04)2 + 6H20 = Fe4S04(0H>10 + 7S042 - + 14H+

(IV.4.16)

+ 2H20 = 4Fe(OH) 3 + so42 - + 2H+

(IV.4.17)

63

Thus it can be seen that the ferric sulphate complex

ing reactions are very dependent upon the pH of the system.

Formation of ferric hydroxide through this reaction mecha

nism implies a successive displacement of sulphate ions from

the inner sphere of the ionic complexes and the extent of

replacement increasing with the acidity. Moreover, the

above mechanism also suppresses the formation ferric hydrox

yl species as suggested by Music et al. (1981) (47). and

explains very well the formation of sulphuric acid during

the process. Goethite, detected in the precipitate, can be

formed by the loss of water according to the following

equation (56),

(IV.4.18)

64

V.4.5 DEVELOPMENT OF PREDOMINANCE REGIONS FOR FERRIC

HYDROXYL AND SULPHATE COMPLEXES

A third newly calculated diagram, Figure (IV.4.5),

showing the stability regions of both ferric hydroxyl and

sulphate complexes as a function of pH has also been con-

structed.

Feso4 +,

The ferric hydroxyl complexes

Feso4-, and Fe2 (S04 )3 in addition

considered are,

to the ferric

hydroxyl species. These are listed in Appendix D with their

thermodynamic data. It can be seen in Figure (IV.4.3) that,

at pH values around 4.5 and low ferric ion activity, the

predominant ferric hydroxyl species is Fe(OH) 2+.

IV.4.6 PRECIPITATION OF FERRIC IONS

Since the iron complexes described above are true

solution species which attain equilibrium with their sur

roundings in seconds (47); such species do not themselves

appear to precipitate although they may form polymers which

lead to iron precipitation. The mechanisms by which a solu

tion entity is transformed into a solid precipitate is

complex and not all steps are well understood. Moreover,

hydroxyl species can also contain sulphate, bisulphate or

sulphate-bisulphate ligands.

A number of precipitates of ferric ions, Fe(III), have

been reported to occur in leach dumps and laboratory column

leaching experiments. Unfortunately, all these precipitates

have not been characterized thoroughly. It appears that

65

0 T: 298.15 I • ' ' '9 )' ' ' ' ' ' ' ' ' [1 F • +++ (-1. I I

I: ·

0

4FoCOHl3CAQlC-154.79l

. . . . 2 FeCJH++ C-54. 80) . .• 3 Fe (l'JHl 2+ (-106. 74)

5 Fe CCJHl 4-C-198. 39) B 16 Fe2CCIH)2++++(-111.55l

21 , · 7Fe3CClHl4+++++C-221.46l 8 FeSC14+C-1B4.68l 9 Fe CSC4l 2- C-364. 36)

31 !10 Fe2(SCJ4l3CAQ) C-536.041 11 SC14--p H: 0/0/. 25

41---~..... 3 4 tpH2CJ:O/O/O · 5 pFECIIIl:0/.25/0

I I

5

;sl I 2 • • : 0...

,..... .... .... 7 ..... I.LI

~8

pSCl4:0/.25/0

o 1 2 3 4 s s 7 a s 10 11 12 13 14 \ pH

FIGURE I\/ .4.5: Distribution of Ferric hydroxyl and Sulphate complexes in

the Fe'+ -SO 2--H O system at 25°C. 4 2

various compounds precipitate simultaneously, and often only

the elemental composition of the precipitated material has

been determined, leaving its exact characterization uncer

tain. According to Sapieszko (55), in laboratory studies

"Fe(OH) 2+ and Feso4 + complexes seem to play the dominant

role in the precipitation of basic ferric sulphates. Feso4+

is the dominating species at all temperatures while Fe(OH) 2+

is the most abundant hydroxyl species at a pH less than 2,

at 25°C."

Thus the only chemical compound reported to precipi

tate was alunite type colloidal particles of hexagonal

crystal symmetry, having a chemical composition of

Fe3 (S04 )2 (0H) 5 .2H2o.

FeOH2+ + 2FeS04+ + 6Hz0 = Fe3(S04)z(OH)5.2H20 + 4H+

(IV.4.19)

Sabean (59) studied the formation of the jarosite

compound, KFe3 (S04 ) 2 (0H) 6 , and plotted a diagram showing the

role of temperature and pH in the Fe2 (S04 )3 - KOH system

(Figure IV.4.6). At temperatures up to 20°c and pH values

above 3.2, a nearly amorphous compound is formed (ferric

hydroxide). At higher temperatures, goethite and hematite

are formed. The hashed area represent the stability of the

jarosite compound, KFe3 cso4 ) 2 (0H) 6 . It can be seen that its

relative stability is increased as the temperature increases

from around 20°c and when the solution pH decreases from

around 3.2 to 1 pH.

Harvey and Linton (61), in 1981, speculated that fresh

precipitated amorphous ferric hydroxide, Fe(OH) 3 probably is

66

oc X Jarosite 200 A Hematite !'::I. /\!::,.

/ 0 Goethite 180 1/,

½ • nearly amorphous material 160 i

•XX b. b. 6 l). /1 140 % 120 i 100 •X,X/ 0 0

' 80

60 ~ . ~, 0 0 0 0 00 ,;o

~. 20 •,X·X. • • • • 0 00

0 1 2 J 4 5 6 1 8 9 10 11 12 13 1{

pi-I

FIGURE IV.4.6: Di2gram of Jarosi te Predominance in

a crystalline form of FeOOH but contains so much adsorbed

water that it appears to be noncrystalline. Their work was

done in nitrate media and whether this suggestion holds for

sulphate solution requires investigation.

Some potential-pH diagrams for the Fe-s-H2o system,

including the jarosite compound KFe3 cso4 ) 2 (0H) 6 have been

calculated at 25°c (19,60). It is seen in Figure IV.4.7

that this compound predominates at pH value less than about

3.0 and at potentials above 0.6 V (SHE). Moreover, this

stability partly supersedes the ferric ion predominance in

these diagrams. At higher pH values than 2.2, goethite,

FeOOH, (formulated as such by Lowson (19) and Brown (60)) is

the stable precipitate. Thus there is a correlation between

the temperature and pH stability plot derived by Babcan (59)

and the potential-pH for the Fe-s-H2o diagrams at 25°c where

jarosite compound is included.

Stumm et al. (54) stated that jarosite was a product

of the aqueous oxidation of pyrite. They emphasized that the

formation of jarosite is particularly favoured because

pyrite provides a source of iron and sulphuric acid. Thus

jarosite compounds are unlikely to form (unless sulphates

are already present) during the aqueous oxidation of pyrrho

tite minerals because the predominant sulphur species formed

is elemental sulphur and not sulphate ions.

The conclusions drawn from this literature are:

i. Jarosite is stable at low pH and moderately

oxidizing potentials,

ii. It is not clear how goethite forms. It may result

67

from ageing of Fe(OH) 3 or directly from the hydrolysis

process.

iii.Various crystalline or amorphous basic ferric

sulphate may precipitate under similar conditions.

Many of these compounds have not been well

characterized,

iv. Ferric hydroxide may form either from the

hydrolysis process or by ageing of basic sulphates.

68

, IGURE lV.4.7: SLabili ty Relati ons Among Jarosite- GoeLhite- Fe;.- -Fe:.;- at, 25°C, 1 aLm. as a funcUon of Eh, pH and activity of, total di ssolved iron. (Activi ty of total dissolved suiphur is 10~M and activity of dissolved pot,assium i s 10,'M) (60) .

IV.S AQUEOUS OXIDATION OF PYRRHOTITE MINERALS BY

MOLECULAR OXYGEN AND FERRIC IONS

In this section, the aqueous oxidation of pyrrhotite

minerals by molecular oxygen and ferric ions, Fe3 + will be

discussed in an attempt to characterize the chemical beha

viour of each pyrrhotite mineral whenever possible.

Table IV.5.1 lists the most common sulphide minerals

in order of their thermodynamic stability. It is noted

that, of the iron sulphides, monoclinic pyrrhotite, Fe7s8 ,

is the most thermodynamically unstable, followed by pyrite,

FeS2 , and troilite, FeS. However, it is known that pyrite

is the most widespread and abundant of all sulphide minerals

(62). This characteristic is not due entirely to its thermo

dynamic stability but as shown in Table IV.5.1, many more

reactive sulphides are theorically more stable than pyrite.

Thus other factors must contribute to the persistence of

pyrite or unstability of monoclinic pyrrhotite.

69

TABLE IV.5.1: FREE ENERGY OF FORMATION OF

FORMATION AT STANDARD TEMPERATURE AND PRESSURE (62)

SULPHIDES

MINERAL FORMULA G0 kJ/mole

Monoclinic Pyrrhotite Fe7s8 -748.52

Oldamite CaS -473.4

Tungstenite WS2 -298.0

Molybdenite MoS2 -297.6

Greigite (synthetic} Fe3s 4 -290.36 (63}

Alabondite MnS -218.1

Sphalerite ZnS -203.4

Chalcopyrite CuFeS2 -179.1

Stibnite Sb2s 3 -173.7

Orpiment As 2s 3 -168.7

Pyrite FeS2 -166.94 ( 7}

Greennockite CdS -145.7

Troilite Fe1.000S -100.42 ( 7}

Galena PbS -96.3

Mackinawite(synthetic} tetragonal FeS -93.30 (63}

Chalcocite cu2S -86.7

Cinnabar HgS -50.7

Covellite cus -48.98 (12}

Argentite Ag2S -40.2

70

Table IV.5.2 lists the equilibrium constant, K, for

acid decomposition of the common sulphide minerals. The

equilibrium constants are referred to the acid decomposition

of sulphides without the presence of an oxidant. The sul

phides may be attacked producing hydrogen sulphide and metal

salts. It is noted that the values vary widely from one

sulphide to another. While values of the equilibrium con

stant for monoclinic pyrrhotite, Fe7s8 , and troilite, FeS,

are large, the equilibrium constants for the copper sul

phides covellite, CuS, and chalcocite, cu2s, and silver

sulphide, argentite, Ag2S, are extremely small. Thus, it can

be seen that the acid decomposition of pyrrhotite minerals

is more favourable than that of copper and silver sulphides.

71

TABLE IV.5.2

CALCULATED EQUILIBRIUM CONSTANT VALUES FOR THE ACID DECOMPOSI

TION OF COMMON SULPHIDES AT 2s0 c (23)

SULPHIDE EQUILIBRIUM CONSTANT, K

9.95 X 107

5.13 X 103 (64)*

(artificially prepared)

FeS 3.91 X 102

cos 2.30

NiS(alpha) 1. 75

NiS(beta) 1.12 X 10-7

ZnS (wur.) 1.55 X 10-2

ZnS (sphal.) 7.4 X 10-S

CdS 7.08 X 10-7

PbS 7.95 X 10-8

CuS 1 X 10-15

Cu2s 3.16 X 10-28

Ag2s 1.26 X 10-29

HgS 6.3 X 10-33

72

Another fundamental property of conducting and semi

conducting minerals is the value of the rest potential. For

interfacial electrode processes, the rest potential corre

sponds to the equilibrium (no nett anodic cathodic current)

electrode potentials. Rest potential values for several

metal sulphides are shown in Table IV.5.3. It is important

to remember that a mineral electrode system will establish

and maintain a certain equilibrium potential that depends

not only on the solution composition but also on the compo

sition of the solid phase (65). Thus, when two phases of

different rest potentials are in electrical contact in an

electrolyte, they form a galvanic cell in which the current

will flow through the solution from the mineral having the

lowest potential. By such action, the phase with higher

rest potential in the electromotive series will behave

cathodically, acting as a site for the oxygen reduction

reaction and will be protected from dissolution, whereas the

phase with lower rest potential will undergo enhanced anodic

oxidation (66).

73

TABLE IV.5.3

REST POTENTIAL OF COMMON SULPHIDE MINERAL AT 2s0 c

Mineral Rest Potential (V vs. SHE)

pH 4 9K MEDIUM

(121) pH 2.5(109)

Pyrite FeS2 0.63 0.66

Marcasite (Zn,Fe)S 0.65

Chalcopyrite CuFeS2 0.52 0.56* 0.005

(0.450-0.550) (11)

Chalcocite cu2S 0.44

Covellite CuS 0.42** 0.45

Bornite Cu5 FeS4 0.42

Galena PbS 0.28 0.40

Sphalerite ZnS -0.24 0.46

Argentite Ag2S 0.28

Stibnite Sb2S3 0.12

Molybdenite Mos2 0.11

Pyrrhotite Fe1_xs -0.28 0.30(101)-0.065 to

-0.135

Pentlandite (FexNi1-x>9S0 -0.065 to

-0.145.

* anomalous

** 1.0 M HCl04

74

According to Table IV.5.3, pyrite has the highest rest

potential while pyrrhotite has the lowest. Consequently,

pyrite has been widely reported as the most active sulphide

mineral for oxygen reduction only slightly less active than

gold and platinum (67-69). Independently, Southwood (66)

presented an order of the common sulphide minerals according

to their rest potentials. Apparently, the acid and tempera

ture conditions are the same as such reported in Table

V.5.3. Although, both Hiskey et al. (65) and Southwood (66)

do agree in the following sequence, pyrite> chalcopyrite >

galena > sphalerite; the latter did not account for pyrrho

tite. Miller (121) reproduced the rest potential values of

various sulphide minerals reported by Majima (1969) at 4.0

pH value. It is seen that although pyrrhotite has also been

neglected the order, stated above, is maintained in this

medium. Natarajan et al.(109) measured the rest potential

of chalcopyrite, pentlandite and pyrrhotite in 9K medium and

2.5 pH and the values are reproduced in Table IV.5.3. In

this medium, pyrrhotite and pentlandite have similar rest

potentials which are approximately lOOmV lower than that of

chalcopyrite. Although it is not expected to have the same

rest potential value in these mediums (1M sulphuric acid, pH

4 and 2.5 pH), it is remarked that an accurate value of rest

potential for pyrrhotite may be difficult to determine since

the variable, heterogeneous composition of natural pyrrho

tite undoubtedly inhibits the reproducibility of the meas

urement. However, it might be concluded that pyrrhotite has

a lower rest potential value than the other common sulphide

75

minerals.

From the thermodynamic properties reported for pyrrho

tite minerals and their relationship with the other common

sulphide minerals, it is noted that monoclinic pyrrhotite

minerals are the most thermodynamically unstable sulphide

mineral followed by sphalerite, chalcopyrite, pyrite and

stoichiometric pyrrhotite. However, the highest equilibrium

constant for acid decomposition was reported for monoclinic

pyrrhotite, sphalerite and stoichiometric pyrrhotite. The

rest potential value for "pyrrhotite minerals" seems to have

the lowest value in the series. Thus pyrrhotite minerals, in

general, appears to be the most reactive common sulphide

mineral to acid decomposition and to an aqueous oxidation

process.

IV.5.1 AQUEOUS OXIDATION OF PYRRHOTITE MINERALS BY

FERRIC IONS

Ferric chloride and ferric sulphate are both impor

tant leaching reagents for sulphide minerals. Few investiga

tions have been conducted to elucidate the reaction kinetics

and to delineate the important leaching variables of pyrrho

tite minerals by these reagents. The stability of pyrrhotite

minerals in the presence of ferric ions, Fe3 +, can be pre

dicted by the oxidation potential of the ferrous-ferric

equilibrium (70) :-

76

Fe3 + + e- -----> Fe2 +

bG = -17.69 Kcal/mol

which at 25°c, is given by

[a.Fe3+]

Eh= 0.771 + 0.0591 log------

[a.Fez+]

a= activity.

(IV.5.1)

(IV.5.2)

Note that a solution of pure Fe3 + (without any

would have, in theory, an infinite oxidizing potential,

dilute solutions have virtually no oxidizing capacity since

even the slightest extent of reduction reduces the Fe3 +;Fe2 +

ratio quickly and hence the oxidizing potential.

The aqueous oxidation process of pyrrhotite minerals

whether as FeS, Fe7s8 or/and Fe9s 10 , by ferric ions, Fe3 +,

implies to possess a complicated model. This model must

explain the effect of the various chemical behaviours of the

pyrrhotite minerals, the media, acidity, and the concentra

tion of ferrous and ferric ions on the equilibrium and the

kinetics of oxidation. The process might be explained by

the following anodic reactions, (IV.5.3), (IV.5.4) and

(IV.5.5) where ferrous ions and elemental sulphur are

widely considered as the main products.

Fe1 _000s -----> Fe2 + + s 0 + 2e-

6G = -5.15 Kcal/mol, ( half-reaction)

Fe1 _000s + 2Fe3 + -----> 3Fe2 + + s 0

(IV.5.3)

bG = -30.35 Kcal/mol (overall reaction with ferric

ions)

77

Fe7s8 -----> 7Fe2+ + ss0 + 14e-

AG= -46.95 Kcal/mol,(half-reaction)

Fe7 S8 + 14Fe3 + -----> 21Fe2+ + s 0

(IV.5.4)

AG = -201.55 Kcal/mol(overall reaction with ferric

ions)

2+ 0 Fe9S1o -----> 9Fe + 10S + 18e- (IV.5.5)

(bG0 ) for Fe9s 10 appears to be not available)

It can be seen that the free energy change for the

overall reaction of monoclinic pyrrhotite with ferric ions

is more thermodynamically possible than that of stoichiomet

ric pyrrhotite and ferric ions. These reactions lead to a

decrease of ferric ions concentration and to an increase of

ferrous ions concentration if an oxidant, e.g. oxygen, is

absent from the solution.

Dissolved molecular oxygen may participate directly in

the aqueous oxidation process of pyrrhotite through a ca

thodic reduction process as well as indirectly by regenerat

ing ferric ions, Fe3 +. Oxygen is a strong oxidant if the

reduction occurs in a "more-or-less synchronous four-elec

tron step" (54). During the regeneration of ferric ions,

the maximum Fe3 +;Fe2+ ratio is limited to the equilibrium

value defined by the oxygen potential:-

78

=

AG= -113.376 Kcal.mol- 1 ., log K1 = 83.

Eh= 1.229168 - 0.05916 pH+ 0.0148 log a02

(IV.5.6)

(IV.5.7)

(IV.5.8)

Thus, this overall reaction can be subdivided in two

two-electron sequences:

02 + 2H+ + 2e- = (IV.5.9)

~G = -31.470 kcal.mol- 1 ., log K2 = 23.1

Eh= 0.6824 - 0.05916 pH - 0.0296 log [H202 ]

and

H202 + 2H+ + 2e- = 2H20 (IV.5.10)

AG= -81.906 Kcal.mol- 1 ., log K3 = 60

Eh= 1.776 - 0.05916 pH+ 0.0296 log [H202]

Oxygen is a much weaker oxidant if the two-electron

reduction sequence forming hydrogen peroxide, H2o2 (a stable

intermediate product in acid conditions) becomes operative

(54,71). However, the equilibrium potential for the reduc-

tion of hydrogen peroxide, [equation IV.5.10] 1.776V,

indicates that the hydrogen peroxide is one of the most

powerful oxidizing agents available and is unstable with

respect to both the oxidation of water and its own oxidation

and reduction.

Conversely, decomposition of hydrogen peroxide, H2o2 ,

the reverse reaction of equation (IV.5.9), is catalyzed by

many metal ions including ferric ions, Fe3 +, cupric ions,

79

cu2+, and cobalt ions, Co2+ of which ferric ions is likely

the most effective (58). Moreover, the presence of ferrous

ions [Fe2 +J consumes hydrogen peroxide rapidly to form

ferric ions and this decomposition is known to be promoted

by the presence of freshly precipitated ferric hydroxide in

solutions where pH values vary from 4.3 to 11.3 {58). Thus

if hydrogen peroxide, H2o2 , is formed during the aqueous

oxidation of pyrrhotite by molecular oxygen at pH around

4.5, it might be decomposed to molecular oxygen or oxidize

ferrous to ferric ions and affect the oxidation process of

pyrrhotite.

Investigators of conventional leaching processes,

involving acid ferric chloride and ferric sulphate solu

tions, have found that the overall reaction is best de-

scribed by the following stoichiometries (22,70,72,73)

whether the mineral is natural, artificial (from thermally

decomposed pyrite), or a stoichiometric compound or not:

[Ref.]

Thermally decomposed pyrite (22):

Fesl.15 + 2Fe3 + -----> 3Fe2 + + 1.15S0 (IV.5.11)

Natural pyrrhotite (70):

FeS+2Fe2 (S04 ) 3 ----->3FeS04 + s 0 (IV.5.12)

Natural monoclinic and hexagonal pyrrhotite (72):

FeS + 2FeC13 -----> 3FeC12 + s 0

Natural pyrrhotite (73):

(IV.5.13)

1.82Fe3 + ----->

2.71Fe2 + + o.013Ni2 + + o.011cu2 + + 0.342Co2+ + s 0 (IV.5.14)

80

Thus ferric ion leaching has several advantages over

direct acid decomposition of pyrrhotite. Firstly, sulphur is

produced in the elemental form eliminating the need for an

additional plant to recover sulphur dioxide (from a pyromet

allurgical process). Secondly, acid is not consumed during

the reaction and only a relatively low acid concentrations

is necessary to prevent the precipitation of an iron com

pound. Thirdly, the iron leaching media can be conveniently

regenerated, either by electrolysis or bacterial oxidation

(70). Moreover, accumulation of ferrous ions has no effect

on the rate of leaching (22,70).

In studies of sulphate leaching of pyrrhotite (70), at

temperatures below 5o0 c, linear kinetics were observed. The

apparent activation energy was reported to be 37.7 KJ/mol.

The rate was found to be independent of the ferric ion

concentration between 0.025 to 0.208 M of ferric ions. These

results, reported by Dutrizac from the unpublished work of

Lowe, suggest that the rate is controlled by chemisorption

on the surface of pyrrhotite. At temperatures above 50°c the

reaction rate decreased sharply because of the formation of

hydrogen sulphide by direct attack on the sulphide.

Subramanian et al. (22) also found that leaching of

thermally decomposed pyrrhotite with about 111 g/1 ferric

sulphate at 0.5 pH about 50% of the sulphur was released in

two hours, but five hours was needed to complete the reac

tion. The temperature at which the reaction took place was

not reported. Precipitation of any ferric iron compound was

prevented by maintaining the pH below 0.5.

81

Recent study (73) on the kinetics of leaching of

nickel, cobalt and copper from pyrrhotite-pendlandite miner

als by acidic ferric sulphate, indicate that the rates are

dependent on both the ferric ion and sulphuric acid concen

tration. The reactions took place in a packed bed reactor in

the absence of oxygen. Potassium permanganate was added

during the dissolution process to maintain a constant poten

tial of 550 V.

It was found that, at a given total iron concentra

tion, the rate of oxidation of pyrrhotite increases rapidly

if the potential of the solution is below 0.550V, but it is

insensitive to further increase in the potential above that

value. When the potential of the solution is increased to

the point where all soluble iron is oxidised to the ferric

state, the leaching rate becomes constant (73). It is also

reported that "the highest oxidation rate is obtained as the

mole fraction of ferric ion in solution is increased to

1.0". This latter statement seems to be in conflict with

that where the rate of oxidation is insensitive to the

potential above 550 V since the potential at 1.0 mole

fraction of ferric ions should be above 0.771 V.

82

IV.5.2 OXIDATION OF PYRRHOTITE AT ELEVATED

TEMPERATURES

Study of the aqueous oxidation of sulphide minerals at

elevated temperatures mainly involves the effects of oxygen

pressure, temperature, particle size, aeration characteris

tics of the autoclave and mixing intensity.

Table V.5.4 presents the stoichiometric reactions,

some experimental conditions for the aqueous oxidation of

pyrrhotite at elevated temperatures and the ferric iron

products reported to have formed under different experimen

tal conditions by a number of authors. It can be seen that

the stoichiometric reactions reported vary even under the

same apparent conditions. The prevalent ferric iron product

is a mixture of various forms of ferric compounds which

have been named differently by various workers. While Downes

(76) simply called it an "iron oxide", Goryachkin (77)

identified goethite and hematite {hydrogoethite), ferrofer

rite, hydromagnetite and basic iron sulphates. The stoichi

ometries for oxygen also varies but the oxygen consumption

rates were not reported. The reaction stoichiometries seems

to have been balanced only according to the iron compound(s}

reported.

83

TABLE IV.5.4:

SOME EXPERIMENTAL CONDITIONS AND STOICHIOMETRIC REAC

TIONS FOR THE AQUEOUS OXIDATION OF PYRRHOTITE AT ELEVATED

TEMPERATURES

1. Mineral: thermally decomposed pyrite with solid/

liquid ratio: 20-30%. Reaction temperature: 110-120°c.

"little sulphuric acid added". Reaction product reported,

iron oxide, ( 76)

4FeS + 302 (IV.5.16)

2. Mineral: nickel-pyrrhotite concentrate. Initial

pure oxygen pressure: 0.21 atm., solid/liquid ratio: 35-40,

reaction time: 3.8 hr., conversion of pyrrhotite: 93.6%,

+and ferric iron reaction product: goethite (hydrogoethite)

and hematite mixture, reaction temperature: 110°c (77).

5FeS + 502 + 2H20----> 4FeOOH + FeS04 + 4S0 (IV.5.17)

3. Mineral: pentlandite and hexagonal and monoclinic

pyrrhotite, solid/liquid ratio: 50. Initial pure oxygen

pressure: 0.5-1.0 atm., Some acidification was required.

reaction time: 3-4 hrs., pyrrhotite conversion: 90-95 % and

ferric iron reaction product: hydrated ferric oxide (78).

2FeS + 1.5 02 + nH20 -----> Fez03.nHzO + 2s0 (IV.5.18)

84

In 1955, Downes (76) suggested a reaction mechanism

for the aqueous oxidation of natural pyrrhotite with water

or for thermally decomposed pyrite with a "small amount of

sulphuric acid'' at temperatures of 110-120°c. The suggested

mechanism is:

FeS + 202 -----> FeS04

6FeS04 + 1.5 0 2 -----> 2Fe2 (S04 )3 + Fe2o 3

Fe2 (S04 ) 3 + 3H20 -----> Fe2o 3 + 3H2so4

(IV.5.19)

(IV.5.20)

(IV.5.21)

FeS + H2so4 -----> FeS04 + H2S (IV.5.22)

2H2S + 02 -----) 2Hz0 + 2s0 (IV.5.23)

H2S + Fez(S04)3 -----) 2FeS04 + HzS04 + s 0 (IV.5.24)

If these reactions are added, the final reaction is

given by equation (IV.5.25):

2FeS + 2FeS04 + 2HzS + 4 1/2 Oz+ HzO -----)

2Fe203 + 3H2S04 + s 0 (IV.5.25)

As this equation differs from the overall reaction

suggested by Downes et al. (76) (Table IV.5.4) it is pointed

out that this mechanism which is the only one published,

was suggested considering the most likely chemical reactions

that could take place under the conditions reported in Table

IV.5.4. Nevertheless the mechanism suggests independent

dissolution processes for pyrrhotite by oxygen and by sul

phuric acid, to form ferrous sulphate. This product is

oxidised to ferric sulphate and then hydrolyzed. On the

85

other hand, hydrogen sulphide is oxidised to elemental sul

phur by oxygen and ferric sulphate. It is believed that

Downes et al. (76) did not account that the oxidation of

hydrogen sulphide is a fast reaction in the presence on an

oxidant, e.g. oxygen, even at ambient temperature. Thus

this compound should not exist as a reactant in the overall

reaction, Equation IV.5.25. Although the oxidation of fer

rous ions and then hydrolysis of ferric ions are fairly well

understood during the aqueous oxidation common sulphide

minerals at ambient and elevated temperatures, the reaction

mechanism of aqueous oxidation of pyrrhotite is far from

clear.

86

IV.6 BACTERIAL OXIDATION OF PYRRHOTITE MINERALS

Pyrrhotite minerals, as other common sulphide miner

als, have been subjected to bacterial oxidation studies. As

pyrrhotite minerals themselves have little or no economic

value, bacterial oxidation has been used as a beneficiation

process to recover the valuable mineral or metals associated

naturally with them.

From the various Thiobacillus species, it is believed

that Thiobacillus ferrooxidans and Thiobacillus thiooxidans

are the major microorganisms directly responsibles for the

aqueous oxidation of sulphide minerals. T. ferrooxidans are

able to derive energy from the oxidation of acidic ferrous

iron as well as inorganic sulphur compounds. T. thiooxidans

are typically found in acidic environments where they ac

count for the production of sulphuric acid from reduced

inorganic sulphur compounds. These bacteria have a remarka

ble acid tolerance, superior to that of any other Thiobacil

lus species, but its optimum growth occurs near pH 4 (110).

Bacterial oxidation of a tin waste concentrate, con

taining about 50 per cent of pyrrhotite minerals, can sub

stantially destroy the pyrrhotite/cassiterite association by

aerial aqueous oxidation. While bacterial oxidation of

pyrrhotite/cassiterite concentrate at 1.8 initial pH favours

the formation of ferric sulphate and sulphuric acid; bacte

rial oxidation stated at 4.5 pH favours the formation of

iron oxides/hydroxides and sulphur as insoluble reaction

products (79).

87

Galvanic interactions, in the presence of microorgan

isms leading to a preferential oxidation of the active

sulphide mineral, play a significant role in a selective

bacterial oxidation of a mixture of sulphide minerals. The

microorganisms first oxidize sulphide minerals having a

lower electrode potential. Therefore, bacterial oxidation of

sulphide minerals is electrochemical in nature.

The rate of bacterial oxidation of a mixture of

pyrrhotite-chalcopyrite and pyrite-chalcopyrite minerals has

been studied by Rossi et al.(80) and Ahonen et al. (111).

They found that pyrrhotite, in a mixture with chalcopyrite,

may have either a positive or negative effect on the solubi

lization of copper, depending on the relative ratio of the

minerals. At an initial pH of 2.38, a pyrrhotite-to-chal

copyrite ratio of 0.5:1 appears the optimum for maximizing

both the rate and extent of the dissolution of copper. At

lower pyrrhotite-to-chalcopyrite ratios, the rate of disso

lution of copper is slower; whereas at higher ratios, copper

tended to be removed from the leaching solution (80). It is

not known why pyrrhotite minerals should improve this proc

ess however, it is though that the preferential dissolution

of pyrrhotite may rapidly increase the presence of ferric

ions in solution attacking the chalcopyrite mineral.

It has been shown (111) that pyrite has a positive

catalytic effect on the dissolution of chalcopyrite because

the electrical conductivity of these minerals allows the

formation of galvanic couple pyrite/chalcopyrite. It has

been reported in an earlier section, Section IV.5, aqueous

88

oxidation of pyrrhotite minerals by molecular oxygen and

ferric ions, that chalcopyrite has a lower rest potential

than pyrite, acts as an anode and is thus preferentially

solubilized.

The reported stoichiometric reactions which occur

during the bacterial oxidation of pyrrhotite minerals are

reproduced below. It can be seen that the amount of oxygen

reported by every investigator varies according to the

particular pyrrhotite mineral and the pH of the leaching

system. Bacterial oxidation of high purity hexagonal pyrrho

tite concentrate at 1.8 initial pH was reported to form

ferrous sulphate, sulphuric acid and elemental sulphur

(Equation IV.6.9). However, X-ray diffraction of the same

leach residues indicated that potassium and ammonio-

jarosite were the principal precipitation products (81). The

amount of iron, precipitated as jarosites, was not stated.

1. Stated at pH: 4.5 (79)

(bacteria) + 2+ 0 2FeS + 4H + o2 -----> 2Fe + S + 2H20 (IV.6.1)

(bacteria)

2Fe2+ + 2H20 + 0.502 -----> Fe2o3 + 4H+ (IV.6.2)

Summing equations (IV.6.1) and (IV.6.2):

(bacteria)

2FeS + 1.502 -----> Fe2o 3 + 2s0 (IV.6.3)

2. Initial pH: 2.38; Final pH: not reported (80)

(bacteria)

89

4FeS + 902 + 4H2o -----> Fe2 (S04 ) 3 + 2Fe(OH) 3 + H2so4

(IV.6.4)

3. Initial pH: 1.8; Final pH: 1.4 (81)

FeS1.10 + H2S04 -----) FeS04 + HzS + O.lS

4FeS04 + 2HzS04 + Oz

(bacteria)

-----> 2Fe2 cso4 )3 + 2H2o

HzS + Fez{S04)3 -----> FeS04 + HzS04 + so

FeS1.10 + Fe2(S04)3 ----> 3FeS04 + 1.1s0

(IV.6.5)

{IV.6.6)

(IV.6.7)

{IV.6.8)

Summing equations (IV.6.5), (IV.6.6), (IV.6.7) and

{IV.6.8):

2FeS1.10 + 2HzS04 + 02

----->2FeS04 + 2.2s0 + 2Hz0 {IV.6.9)

4. Initial pH: 1.91, Final pH: 1.04 (111)

2FeS + 4.502 + 3H+----->2Fe3 + + so42 - + HS04 -{IV.6.10)

2FeS + 1.502 + 6H+ ----->2Fe3 + + 2s0 + 3Hz0 (IV.6.11)

The stoichiometric reactions for the oxidation of

pyrrhotite are directly related to the pH of the system.

While Harris et al.(79) report one and half moles of oxygen

for every two moles of pyrrhotite (FeS/02 = 0.75/1) at 4.5

stated pH, Rossi et al. {80) report nine moles for every

four moles of pyrrhotite (2.25/1) at 2.3 pH and Kandemir

(81) one mole for every two moles of pyrrhotite (0.5/1) at

1.8 initial pH. Ahonen et al. (111) presented two simultane

ous reactions where sulphate/bisulphate ions and elemental

sulphur are main reaction products. This information will

90

be evaluated in the present study.

The performance of bacterial oxidation process is

directly related to the activity level of the microorgan

isms. This activity could be evaluated indirectly by analy

sis of the accumulation of metals in solution, the constant

decrease of pH and constant increase of potential in the

leaching

level of

system. Further, estimation of the microorganisms

activity from the slurry could be obtained by

protein determination followed by fluorescence microscopy

(83,110). Thus, since T. thiooxidans can produce negative pH

values, highly active microorganisms will constantly both

decrease the pH and increase the potential of the system

from the beginning of the process. Higher potential values

are probably due to higher ferric/ferrous ions ratio formed

during the oxidation process.

In general, bacterial oxidation of sulphide minerals

depends greatly in the acclimatization of the bacteria to

the particular mineral, e.g. attachment to the preferred

mineral sites resulting in direct attack, development of

metal tolerant strains of bacteria and temperature. Most

studies on bacterial oxidation have usually employed T.

ferrooxidans and T. thiooxidans which have the ability to

oxidize reduced sulphur compounds and reduced iron to obtain

energy for growth and other life processes. However, there

are other microorganisms which are currently evaluated e.g.

the thermophilic bacterium Sulfolobus acidocaldarious (115)

and moderately thermophilic Thiobacillus-like microorgan

isms (Leptospirillum ferrooxidans) (113). It is claimed that

91

these organisms may operate at temperatures of 70°c and

45°c, respectively, compared with 35-37°c favoured by T.

ferrooxidans and T. thiooxidans. Temperature increase might

remove the need to cool the reactor and improve the kinetics

of the process.

Although it is not understood clearly how the presence

of Thiobacillus ferrooxidans can affect the oxidation of a

sulphide mineral itself; Kandemir (81) suggested that the

dissolution could occur through an electrochemical mechanism

in which bacteria accelerate the cathodic reduction of

oxygen at the sulphide-bacteria interface. Moreover, disso

lution of pyrrhotite could firstly occur through a non

oxidative reaction (reaction IV.6.5), followed by bacterial

cyclic oxidation of ferrous ions (reaction IV.6.6). Ferric

ions then react in fast chemical reactions with hydrogen

sulphide and pyrrhotite itself (reactions IV.6.7 and

IV.6.8). Ahonen et al. (111) restated the above partial

equations and showed that at low pH values, e.g. 1.91 pH,

ferric ions precipitated mainly as jarosites. At higher pH

values, it precipitated as ferric hydroxide. This chemical

behaviour of ferric iron was fully described in section

IV.4.

However, it has to also pointed out that Ahonen et al.

(111) prepared the 9K solution with 0.4 g.L-1 each of

(NH4 ) 2so4 and MgS04 .7H2 which is different to that

suggested by Silverman et al. (114) g.L- 1 :

KCl 0.10, K2HP04 0.50, and MgS04 .7HzO 0.50). It is not

understood the effect of 9K media. Moreover, solubilization

92

of iron from pyrrhotite commenced without a lag period and

was completed only in three weeks. In contrast, only a minor

amount of pyrite was oxidized during the same period, where

soluble sulphate paralleled the iron solubilization. The

faster rate of oxidation of pyrrhotite is ascribable to its

mixed potential, which yields a self-sustained dissolution

below -0.15 V (SCE) (27).

Pyrrhotite minerals have been classified according to

their reactivity to aqueous oxidation as "more" and "less"

reactive species (79). Harris et al. (79) indicated that

during the bacterial oxidation of a "reactive" pyrrhotite,

the pH increases to a value where the oxidation of ferrous

to ferric species becomes kinetically favourable and the

overall rate of oxidation is determined by the rate of

chemical oxidation of ferrous ions. On the other hand,

during the oxidation of a "less active pyrrhotite", the pH

may decrease until a significant ferric ion concentration is

established and the overall oxidation will be determined by

the biologically assisted oxidation of ferrous ions to

ferric ions. This chemical behaviour may not be observed in

an larger operation scale since highly active microorganisms

will constantly both decrease the pH and increase the poten

tial of the system from the beginning of the process

(116,118). A more detailed analysis of continuous bacterial

oxidation of refractory gold minerals is presented in sec

tion VII.7.

It is also indicated by Harris et al. (79) that the

results of the bacterial leaching of pyrrhotite minerals

93

show some degree of lack of reproducibility. The principal

contributing factors to the low degree of reproducibility

are probably the wide range of variability of reactiveness

of pyrrhotite reaction and the as yet incomplete information

on the components of the indigenous microbial population.

Bacteria speciation changes in operations where long resi

dence time, e.g. 50 days, are maintained by enhancing the

growth of other Thiobacillus species than T. ferrooxidans

and T. thiooxidans, e.g. fungus (117). Once acclimatized,

they also become dormant under adverse conditions and revive

as soon as the conditions return to normal (118). It has

also been stated by Miller et al. (119) that "brief inter

ruption in the air supply resulted in a one-week decrease in

the metal extraction".

The principal bacterial oxidation products of sulphide

minerals, particularly those of pyrrhotite minerals such as

elemental sulphur, sulphate ions, ferric ions and other

metal associated with the substrates vary as a function of

pH, potential, leach time and the mineralogical properties

of each substrate (79). Brock et al. (120) reported the

conclusion of Sato who stated that sulphur released from the

crystal structure of pyrite is converted to unstable s 2

molecules, which could be instantly oxidised in the presence

of oxidising agents. From other sulphides (copper, lead,

silver and zinc) stable solid sulphur, s 0 , would be formed

which would eventually be oxidised to sulphate and reacts

slowly with ferric ions in the presence of bacteria. Al

though pyrrhotite minerals have not been mentioned by Sato,

94

it might behave predominantly as the "other sulphides" since

it yields mainly elemental sulphur rather than sulphate

ions at pH values less than 7 (79,80,101,111). Thus, Brock

et al. proposed that the genera Thiobacillus are able to

reduce ferric iron when grown on elemental sulphur as energy

source.

The rate of bacterial oxidation of pyrite was found

related to its type of conductivity, its real electron

structure. According to Karavaico et al. (121) bacterial

oxidation of pyrite with hole conductivity takes place at

higher rate and continuously compared to pyrite electron

conductive. Apparently, electron conductive pyrite is oxi

dised until all excess iron is oxidised and FeS2 becomes

stoichiometric in composition increasing its electrochemical

potential. Oxidation of hole pyrite, iron deficient, a hole

semiconductor remains uncompensated producing ferric iron

into solution for a long time. This kind of study has not

been found in the literature for pyrrhotite minerals.

95

IV.7 LITERATURE SURVEY CONCLUSIONS

The study of the literature indicated that at pH

values less than 7, reduction of oxygen and ferric ions are

the main cathodic reactions during the aqueous oxidation of

common sulphide minerals. However, there is very little

information for the rate of uptake of molecular oxygen

during their aqueous oxidations; particularly for the aque

ous oxidation of pyrrhotite minerals. Reduction of oxygen

and ferric ions are likely the main cathodic reaction wheth

er the process takes place in the presence or absence of

microorganisms, e.g., T. ferrooxidans and T. thiooxidans.

The reaction products produced during bacterial oxida

tion appear to be determined mainly by the pH, temperature

of the system and the actual sulphide mineral. At pH values

above 2 and at ambient temperature; pyrrhotite minerals tend

to yield ferric hydroxide and elemental sulphur predominant

ly.

Autoclave oxidation and bacterial oxidation of common

sulphide minerals are the most attractive alternative proc

esses, to conventional pyrometallurgical processes, for the

recovery of metal values. The kinetics of bacterial oxida

tion or aqueous oxidation at elevated temperature for a

particular sulphide mineral should indicate the residence

time, and, hence, the economic feasibility of the process.

The lack of knowledge of the oxygen consumption rate re

strict the understanding of both chemical and bacterial

oxidation.

96

The mechanism of the aqueous oxidation of pyrrhotite

minerals, whether as stoichiometric, hexagonal or monoclinic

pyrrhotite, is far from clear. Measurement of the rate of

consumption of oxygen; whether in the presence or absence of

microorganisms, may help clarify the process in more detail.

Thus determination of this rate must add to the knowledge

already available and to establish more efficient operating

conditions for chemical and/or bacterial leaching opera

tions. Moreover, measurement of the rate of uptake of

oxygen by pyrrhotite minerals has been unlikely published.

Thus an equipment, called an 'OXYGRAPH' is to be developed

in order to measure this rate.

Aqueous oxidation of sulphide minerals whether at

elevated temperatures with oxygen under pressure or bacteri

ally catalyzed is an electrochemical process in nature.

Sulphide minerals ordered according to their rest potential

values as: FeS<ZnS<PbS<CuS<Cu2S<CuFeS2 <Fes2 appears to

predict their preferential aqueous oxidation, where pyrite

seems the more refractory mineral, cathodically protected,

in the aqueous oxidation process.

A recent and specific potential-pH diagram for the

Fe-s-H2o system at 2s0 c for only pyrrhotite minerals, wheth

er as stoichiometric pyrrhotite and/or monoclinic pyrrhotite

has been scarcely found in the literature. Hamilton et al.

(101) derived a diagram for "FeS and pyrrhotite" using an

earlier thermodynamic data than 1977. Thus using recently

published thermodynamic data, new potential-pH diagrams are

to be developed. Potential-pH diagrams for the Fe-S-H2o

97

system at 2s0 c and higher temperature, where pyrite is

included, indicate that the direction of aqueous oxidation

process for pyrrhotite minerals should produce pyrite.

However, in an aqueous oxidation process of pyrrhotite

minerals; pyrite is not formed as main reaction product.

Thus, elimination of pyrite in this particular potential-pH

diagram should facilitate the understanding the aqueous

oxidation process of pyrrhotite minerals.

Finally, the aqueous oxidation of pyrrhotite minerals

is fully analyzed using these potential-pH diagrams.

Bacterial oxidation of pyrite concentrate at miniplant

scale (30 L x 6 leaching tanks) has not been found in detail

in the literature. Thus, a study performed at Aurotech N.L.

is presented where the effect of flow diagram and dissolved

oxygen concentration are discussed in detail.

98

PART B:

CHAPTER V

DEVELOPMENT OF A POTENTIAL-pH DIAGRAM FOR THE IRON

SULPHUR-WATER SYSTEM AT 2s0 c

V.1 THERMODYNAMIC PROPERTIES OF PYRRHOTITE MINERALS

Lowson (19), in 1982, reviewed the aqueous oxidation

of pyrite and listed the thermodynamic properties of a

number of the iron sulphides. As this seems the most com

plete list published recently it is reproduced in Table V.1.

The free energy of formation of the stoichiometric pyrrho

tite, troilite, Fe1 _000s, and monoclinic pyrrhotite, Fe7s8 ,

have values of -100.42 and -748.52 KJ/mol, respectively.

Values for intermediate pyrrhotites, such as from Fe9s10 , to

Fe11s 12 do not appear to have been published. All the above

thermodynamic data, reported by Lowson, was taken from a

1982 National Bureau of Standards publication (84).

V.2 DERIVATION OF THE POTENTIAL-pH DIAGRAM FOR THE

IRON-SULPHUR-WATER SYSTEM AT 2S°C NEGLECTING PYRITE

Potential-pH diagrams can be used to indicate whether

a particular reaction is thermodynamically possible. These

diagrams first provide a summary of the electron-transfer,

proton-transfer and combined electron-proton-transfer reac-

99

TABLE .. " .• 1 : Thermodynamic Properties of Iron Sulphide Minerals ( \'I )

H" A'" - Cp • a + bT + c'f'I, J de11-• mo1· 1 -G"(T)-Allr° 0 , Alft0

, A(°;t, H •• S",J Cp,J 1-/° O• J compouncl k,J mo1· 1 kJ mo1· 1 k,I moi- 1 kJ mo1· 1 clrll · 1 mol· 1 clr11· 1 mo1·• a b X 10·> C X 101 AT, K de1t· 1 mo1· 1

r,yrile, FrS 1 -174.56 -178.24 -166.9'1 9.63 52.93' 62.17 68.58 13.64· -9.05 200-780 -!0.61

pyrilr, FrS, -173.64 ~ 1 GZ.::1<1 62.09 74.81 5.52 -12.76 298-1000 pyril.r, P<'S, -17'1.0f, ~.162.7G 62.1 "/ pyril<', F<'S, -171.f,,1• mnrc.,~ilr, FrS, -1G9.'1f, -15R.57 9.74· 53.89· 62.43 67.67' 16.11 -8.77 200-70G 21.21 marl'nsilr, FrS, -l!iil.81 marr:isitr, FrS, -1 f10,Ci2 lroililt•, F1• 1_,.,S GR.95 t.roditc•, P<',_.,,S f,'1 .64 21.72 110.46· 29fHl 1 .troilitr. F" 1 •1111 S -100.4G -99.99 t00.'12' 9.35 G0.31· fiO.f,4 - 10.1:, 124.73 3.39 200-350 pyrrhotill' as Fr 0_,,.,S G0.79 49.87 38.58 46.82 -2.36 200-350 pyrrholilr ns Fr,S, -743.'I 1 -7:16.3H ·7•18.52· 73.72- 41:lf,.7G· 398.57

Fr,S, -279.91 -280.7!i 1,')2,2!1

proton-transfer and combined electron-proton-transfer reac

tions which are thermodynamically favourable when a mineral

is in a particular aqueous solution.

When a potential-pH diagram indicates that a particu

lar reaction is thermodynamically favourable it does not

mean that the reaction will take place at a significant

rate, or at all. Thermodynamics, therefore, only defines a

necessary precondition for a reaction and it determines the

direction in which an overall reaction will tend. Determina

tion of the rate and mechanism of the reaction can only be

obtained by a detailed study of the kinetics of the reaction

(85,86).

When pyrrhotite is leached by ferric sulphate or

ferric chloride solutions, or by oxygen under pressure in

acid solutions, iron enter the solution and the only new

solid phase formed is elemental sulphur. A study of the

potential-pH diagram for the Iron-Sulphur-Water system fails

to predict this chemistry. Pyrite occupies the region where

the products of the aqueous oxidation of pyrrhotite are

stable (87-91). Moreover, monoclinic pyrrhotite occupies the

region where the products of the aqueous oxidation of

stoichiometric pyrrhotite are stable. Pyrite can be formed

in hydrothermal systems involving hydrogen sulphide, e.g.,

it occurs through a dissolution, reprecipitation mechanism

rather than a transformation of the solid (17). Thus a

potential-pH diagram

pyrite should enhance

for the Fe-S-H2o system neglecting

the understanding of the mineral

chemistry of pyrrhotite minerals.

100

Apparently, nucleation and growth of a new mineral

phase is a process that does not take place readily at

ambient or elevated temperatures in these hydrometallurgical

processes. However, there is evidence to suggest that an

extremely thin layer of pyrite may form on the surface of

pyrrhotite preventing further reaction during aqueous oxida-

tion at ambient temperatures (or even autoclave) (88).

Moreover, after two hours of leaching, autoclave oxidation

of monoclinic pyrrhotite and troilite at 245°c, yielded

very small amount of pyrite and hexagonal

(92,93). •

pyrrhotite

Figures V.2.1, V.2.2 and V.2.3 show the potential-pH

diagrams for the Fe-s-H2o system constructed by various

authors. It is frequently reported that pyrite co-exist

with pyrrhotite minerals complicating the understanding of

their mineral chemistry. Aqueous oxidation of stoichiometric

pyrrhotite, formation of ferrous ions, Fe2 +, is pH dependent

rather than potential dependent. Contrarily, Aqueous oxida

tion of monoclinic pyrrhotite depends on both the potential

and pH.

The recent free energy of formation published and

used to construct the potential-pH diagrams are not only for

stoichiometric and monoclinic pyrrhotites, but also for the

other species considered; if they were available.

The potential-pH diagrams presented in Figures V.2.4

to V.2.9, were constructed by using software developed by

the Division of Mineral Chemistry of CSIRO (94). These

diagrams were forced to be consistent with the phase rela-

101

l·OrT -~-'- FccO,

O·!'> -

Fe ..

10• ~eS J

-o-s 1---'-"1 o'-·-) ~

I Fe

-.,

-1·0=

~L-ob--~1)-·-;--+---!:----:~----1-..:.H:s:·1.:s'....'_j 6 cl 10 12 14

HSO;;

0

Fez+

15a

-I Fe

-1· 5

0

pH

Unit activity for both dissolved sulphur and

dissolved iron.

Fe 0 2 - · ~--~

.... ----·---. ,_ -- -... __ ......... ..... --- ... T"::,-... ___ ....

I --I HFeo- >-

2

9b

HS lo S 2

-

. 8 12

pH

FIGURE V.2.1: Potential-pH Diagrams for the Fe-S-HzO System at 25°C. Reference (88) and (81), respectively.

0 2 4 6 8 10 12 14

2.0~----~---------------------~ 2

Uh

-0,0 ·

- :,0 . Fe -I

- 1,2 • 52-

-1,'l.

- 1,6

- 1,1)

- 2 '~ L< -_-1-.!.0 __ ...___2:,--,,,---'-4 -5:--cG::--, ---!7;---;:0--;;-9--;l;;--"0 °TI 1 2 13 l'l I~ 16 17 18-2

pH

FIGURE V.2.2: Potential-pH Diagrams for the Fe-S-He.o· System at 25°C. Reference (112) and (91), respectively.

O,!+ -, I I

0 -

0,2 ~so fy~ .?S FeqS10

-- ----- -- ----"""- ----- -- ------ -----Fe S

-02 ,·

0

r· l.SCE

-0,2,

-o 6 )

-OB I

-· -A t-lGLiRE V.2.3: Potential-pH Diagram for the FE:1-S System at 25°C for 10 M

cc,ncentraticm of dissol\,1ed speciPs (28).

tionship, actually observed during the oxidation-leaching

of pyrrhotite, by neglecting pyrite.

The potential-pH diagrams presented in the literature

for the Fe-S-H2o system usually include only stoichiometric

pyrrhotite and pyrite omitting monoclinic pyrrhotite (87-

91). Moreover, The stability of stoichiometric pyrrhotite at

a determined activity of iron dissolved species was present

ed differently by every author. This discrepancy mainly

stems from the value of the free energy of formation used.

Thus the mineral chemistry of pyrrhotite minerals will be

analyzed through the potential-pH diagrams presented for

this study drawn using the most recent thermodynamic data.

The species considered, and their free energy of

formation, are listed in Appendix A. Appendix B shows the

thermodynamic calculation involved in deriving the poten

tial-pH diagrams. Appendix C lists the equations relating

the potential and pH for the Fe-S-H2o system at 25°c involv

ing FeO, Fe3o4 , Fe2o3 and (alpha)-FeOOH species. However, it

is noted that the potential-pH diagrams being presented in

Figures V.2.4 to V.2.9 do not involve these species.

The temperature stability of the solid compound was

considered in estimating which solid is stable at 25°c.

The stability of a particular oxide, oxyhydroxide or hydrox-

ide was obtained from various sources (57,95). Ferric

hydroxide, Fe(OH) 3 , sometimes referred to as ferric oxy

hydroxide, FeOOH (95), is stable at 25°c. At temperatures

above about 90°c, well-defined crystals of goethite or

lepidocrocite are formed. On further heating, to above 130°,

102

2. 5.

2. 0~ s

1.5~

1. 0 ~

0 . 5 l::.__ Q

:r: 0~ w

-0.5~

- 1 . 0~

,._ - 1 . 5~

I -2.0

-2. ~ 2 0

FIGURE \/ .1.9: Eh-pH DIAGRAM FOR THE Fe-S-1120 SYSTEM AT 25°c Iron Sulphide Species: Fe

7s

8 _

1 Dissolved Species Activity: 10 M.

I R S + FE E-2 (RQ)

I~ S + FE? SB

I H2 S CRQl + FE E-2 D H2 S CAQJ + FE E

(RQ)

~I M IF H2 S CAQl + FE? SB H SE (RQ) + FE

G H SE CRQ) + FE7 SS

ILi ~ IH S E2 CRQl + FE ] S E2 CRQJ + FE [O H I 2 _J S E2 CAQJ + FE? SB

I ~ ~ IK S 04 E2 (AQ) + FE E-2 IRQJ L S 04 E2 CAQ) + FE E-3 ( AQ I

l~ S 04 E2 CAQ) + FE Qij E2 (RQ) S OLJ E2 CAQ) + FE [0 H 12

C S OLJ E2 CAQ) + FE CO H l 3

\ I I ~ IP S 04 E2 (AQ) + FE7 S8 =c:::: ~ Q H S OIJ E IRQJ + FE E-2 CAQJ

R H S OIJ E IRQ) + FE E-3 CAQ)

I ~ s H S OIJ E IRQ) + FE OIJ E:2 CRQJ T FE E-2 (AQJ + FE? SB

D I I I~ FE+ FE7 SB

H FE [OH 12 + FE7 SB F

I

2 LJ, 6 8 10 12 11.J 16 PH

2. 5.

2.0"-. T

1. 51-s

1. 0~

0.5l- A

:::r:: 0~ w

-0.5~

- 1. 01-

,,._ - 1 . 51-

-2. 0 1

-2. ~ 2 0

FIGURE \/ .1.i: Eh-pH DIAGRAM FOR THE Fe-S-H20 SYSTEM AT 25°c Iron Sulphide Species: Fe

1 000s_gnd Fe

7s

8 Dissolved S~ecies Activityi 10 M.

I A H2 S CAQJ + FE E-2 !AQ)

I~ H2 S CAQJ + FE

I H SE CAQl + FE E-2 CAQJ D H SE CRQ) + FE E H SE CRQl + FE CO H J 2

N F H SE CAQ) + FE 5 G S E2 CAQJ + FE 02 HE IAQJ H S E2 CAQJ + FE 1 S E2 CAQJ + FE CO H 12 .J S OLJ E2 CAQl -t- FE E-2 IAQJ

I "" ~ I t~ S 04 E2 CRQ) + FE (0 HI E-1 L S OLJ E2 CAQ) -+ FE 02 HE CAQJ

~,~ S OLJ E2 CAQl + FE E-3 IAQJ

-------- J ~ F S OLJ E2 CAQ) + FE 04: E2 IRQ) 0 S OLJ E2 CAQl + FE CO H 12 F S 04 E2 CRQ) + FE (0 H 13 Q S OLJ E2 CAQJ + FE S R H S 0~ E IRQ) + FE E-2 CRQl

I I~ Is H S 0~ E IRQ) + FE E-3 C AQ) T H S 0~ E IRQ) -t- FE 0~ E2 CAQJ

D . I ,~ FE E-2 lAQJ + FE S

8 I H FE CO H I E-1 IRQ) + FE S w FE+ FE S

I I ,x FE CO H 12 + FE 5

2 L! 6 8 10 12 14 16 PH

lrlt;

2. 51

2. 0~ w

1 • 5 f-

1 . 0 f-

0. 5~ u

::c 0~ w

-0.5

- 1 • 0f-

,._. - 1. 5f-

-2.0~

-2. ~'2 I 0

FIGURE \I .1.1: Eh-pH DIAGRA~ FOR THE Fe-S-H20 AT 2s 0 c Iron Sulphide Species: Fe

1 080s and Fe

7s

8 Dissolved Species Activity: .01 M.

I A S + FE E-2 CAQ)

I~ S + FE? SB

I H2 S CRQJ + FE E-2 IRQ) D H2 5 CAQJ + FE E H2 S CAQJ + FE S

~I p IF H2 S CAQl + FE7 58 G H SE CRQ) + FE

lul ~ IH H SE CRQ) + FE S H 5 E CRQ) + FE:7 58 ]

.J S E2 CAQJ + FE I ~ ~ IK S E2 CAQl + FE (0 H 12

L S E2 CAQl + FE S

l~ S E2 CRQl + FE7 S8 5 OY E2 CAQ) + FE E-2 IRQ) SOY E2 CAQ) + FE E-3 IAQJ

IP SOY E2 CAQ) + FE 04: E2 IRQ) SOY E2 CAQ) + FE (0 H l2 Q

R SOY E2 CRQ) + FE (0 H 13

I ~ s 5 OY E2 CAQ) + FE 5 T SOY E2 CAQ) + FE7 S8

D I u H S DY: E lAQ) + FE E-2 CAQ) V H S DY: E IRQ) + FE E-3 CAQ)

C .J w H S DY: E IRQ) + FE DY: E2 CRQJ

I I I~ FE E-2 lAQJ + FE 5 FE E-2 lAQJ + FE? SB

:z FE+ FE 5 I I I I I I I I I I RA FE [ 0 H l 2 + FE S 2 L! 6 8 10 12 1 lJ 1 ~B FE [ D H I 2 + F E7 S8

PH AC FE S + FE? S8

2. 5,

2. 0~ T

1 . 51-

1 . 01-

0. 5r- A

::r:: 0~ w

- 0. 51-

- 1 . 0~

r. - 1 . 51-

-2. 0 1

-2. ~ 2 IZl

FIGURE V .l.~: Eh-pH DIAGRAM FOR THE Fe-S-U20 SYSTEM AT 25°c Iron Sulphide Species: Fe

1 00 S_

6 Dissolved Species Activity: ?o M.

I A H2 S CAQl + FE E-2 IRQ)

I~ H2 S CAQl -t- FE

I H S E CRQl + FE E-2 CRQl D H 5 E CRQ) + FE E H S E CRQl + FE (0 H J 2

~I N IF H S E CAQ) + FE S G S E2 CAQJ + FE 02 HE IAQl

I ~ IH S E2 CRQl + FE M 5 E2 CAQl + FE (0 H I 2 I

.J S 04 E2 CAQl + FE E-2 IAQl I ~ ~ IK S 04 E2 C AQ l + FE [ 0 H I E- 1

L S 04 E2 CAQ) + FE 02 HE CAQl

~I~ S 04 E2 CRQ) + FE E-3 IAQl

-------- .J ~ f' 5 04 E2 CAQ) + FE 04: E2 IRQ) 0 S 04 E2 CAQ) + FE (0 H 12 F S 04 E2 CAQ) + FE (0 H 13 Q S 04 E2 CAQ) + FE S R H S 04: E IRQ) + FE E-2 CAQ)

I I~ I 5 H 5 04: E IRQ) + FE E-3 CAQ) T H S 04: E lAQ) + FE 04: E2 CAQl

0 I I~ FE E-2 (AQl + FE S B I H FE (0 H I E-1 IRQ) + FE S

w FE+ FE S I I iX FE [0 H 12 + FE 5

2 L! 6 8 10 12 14 16 PH

IRG

2. 5,

2. 0~ u

1. 5~ T

1 • 01

0.5t: s

:::c 0~ w

-0. 5~

-1.0~

.... -- 1. 5r

-2.01 -2.~2 0

f l G U l{ E 'I . 1 • ~ : Eh - p II U l AG l{ AM 1'' 0 l{ T II E i.,· e - S - II 2

U S Y ST EM AT :! :, " c

Iron SulphiQe Species: Fe1 0 0

s Dissolved Species Activity: 8.0001 M.

I A S + FE E-2 CRQ) ,~ H2 S CRQJ -t- FE E-2 IRQ)

I H2 S (RQl + FE ,

D H2 5 CRQJ + FE 5 E H SE CRQ) -t- FE

0 F H SE CRQ) + FE S G S E2 CRQJ + FE 02 HE !RQJ H S E2 CAQJ + FE 1 5 E2 CRQJ + FE CO H 12 .J S E2 CAQJ + FE S

I ~ ~ IK SOY E2 CRQ) + FE E-2 lAQJ L S OLJ E2 CRQJ + FE CO H l E-l

~~ S OLJ E2 CRQ) + FE 02 HE CRQJ

B ~ ~ Q 5 OY E2 CRQJ + FE E-3 IRQJ

SOY E2 ( AQ) -t- FE 04: E2 IRQ)

I u 1v ~~~ IP SOY E2 CAQ) + FE [0 H 12 Q S 04 E2 CRQJ + FE CO H I 3 R S QLJ E2 ( RQ) + FE S

I ~5 H 5 014 E IRQ) + FE E-2 CRQ) T H S DY: E IRQ) -t- FE E-3 CAQ)

C I I I~ H S 04: E lRQ) + FE 04: E2 CAQJ

H FE E-2 IAQJ + FE S E FE CO HI E-l IRQ) + FE S

X FE+ FE 5 1 FE CO H 12 -t- FE S

2 l! 6 8 10 12 14· 16 PH

(R[;

FIGURE V .1.4: Eh-pH DIAGRAM FOR THE Fe-s-112

0 SYSTEM AT 25°c Iron Sulphide Species: Fe

1 0 0s

Dissolved Species Activity: 8.01 M.

2. 5, I A S + FE E-2 CAQ) ,~ S ;- FE S 2. 0~

s I H2 S CRQJ + FE E-2 IRQ) 0 H2 5 CAQJ + FE E H2 S CAQJ + FE S

1.5~ ~I M IF H SE CAQ) + FE G H SE CRQ) + FE S

ILi ~ IH S E2 CRQJ + FE 1 . 01 1 5 E2 CAQJ + FE (0 H I 2 .J S E2 CAQJ + FE S

0. 5k._ Q ·1 ~ ~ IK S 04 E2 CAQ) + FE E-2 (RQl L S 04 E2 CAQJ + FE E-3 lAQl

~~ S 04 E2 CRQl + FE 04: E2 lRQ) ::c 0~ 5 04 E2 CAQJ + FE E o rt 1 2 w S 01.J E2 CAQ) + FE (0 H I 3 C .....

-0.5~ I I I . ~~ IP S 04 E2 CAQ) + FE S -c.: Q H S 01..! E lRQJ + FE E-2 CRQJ

A H S 01..! E lRQJ + FE E-3 CRQ) -1.01- I l 5 H 5 01..! E lRQJ + FE 01..! E2 CAQl

' T ' FE E-2 lAQ I ;- FE S 0 I I

u FE+ FE S -1.5~

H V FE (0 H 12 + FE S .....

F I I

-2.0

-2. ~ 2 0 2 lJ: 6 8 10 12 1 LJ 16 PH

these decompose to hematite, Fe2o3 . Ferrous hydroxide decom

poses to magnetite above 15o0 c (57,95). Thus, in this study,

only ferric hydroxide, stable at 25°c, was considered and

was referred to as FeOOH. Potential-pH diagrams for

stoichiometric pyrrhotite, FeS, were constructed for 10-2 ,

10-4 and 10-6 M dissolved iron species and these are pre

sented in Figures from V.2.4 to V.2.6. Stoichiometric pyr-

rhotite, (FeS) and monoclinic pyrrhotite,

superimposed for 10-2 and 10-6 M dissolved iron species.

These diagrams are represented in Figures V.2.7 and V.2.8.

Finally, a diagram for only monoclinic pyrrhotite at 10-2 M

was constructed, Figure V.1.9.

V.3 ANALYSIS OF THE CHEMISTRY OF PYRRHOTITE OXIDATION

LEACHING

The chemistry of pyrrhotite oxidation leaching is

analyzed in terms of the potential-pH diagrams derived for

the Fe-s-H2o system. At pH values less than 7, the stability

of pyrrhotite solid phase, whether as troilite, (FeS) or

monoclinic pyrrhotite, (Fe7s 8 ) is entirely located within

the predominance region of ferrous ions, Fe2 +. This thermo

dynamic behaviour might indicate that the initial reaction

of pyrrhotite takes place non-oxidatively.

1. Pyrrhotite can be dissolved non-oxidatively in acid

solutions producing ferrous ions and hydrogen

Stoichiometric pyrrhotite, monoclinic

sulphide.

pyrrhotite and

ferrous ions, predominate over the whole range of existence

103

of hydrogen sulphide.

(V.3.1)

Fe7Ss + 16H+ + 2e- -----> 7Fe2+ + 8H2S (V.3.2)

While dissolution of troilite may proceed independent

ly of the potential, dissolution of monoclinic pyrrhotite or

any other non-stoichiometric pyrrhotite mineral may depend

on the potential because it is a reduction process. A more

detailed discussion of this process has been given in the

section on acid decomposition of pyrrhotite (Section IV.2).

2. Aqueous oxidation of stoichiometric pyrrhotite at

pH values 4.0-4.5 may form only ferrous sulphate. Aqueous

oxidation of monoclinic pyrrhotite may also form an addi

tional molecule of elemental sulphur. This elemental sulphur

may oxidize to higher oxidation states such as sulphate ions

in the presence of sulphur oxidising bacteria.

FeS + 202 -----> FeS04 (V.3.3)

(V.3.4)

Additionally, it can be seen from Figures V.2.4,

V.2.6 and V.2.9 (0.01M activity of dissolved iron species)

that the aqueous oxidation of troilite and monoclinic pyr

rhotite at a pH less than about 4.2 and at potentials about

zero, may also form ferrous ions and elemental sulphur;

FeS + 2H+ + 0.502 -----> Fe2 + + s 0 + H2o (V.3.5)

104

Fe7S9 + 14H+ + 3.502 -----> 7Fe2 + + ss0 + 7Hz0 (V.3.6)

At pH greater than 4.2, at slightly negative poten

tials (Figures V.2.4, V.2.6 and V.2.9) stability of troilite

and monoclinic pyrrhotite with both elemental sulphur and

sulphate ions is thermodynamically possible. Further, it is

noted that the domain of stability of stoichiometric pyrrho

tite alone is congruent to that of monoclinic pyrrhotite at

the same activity of dissolved species, excepting at the

boundary line where stoichiometric pyrrhotite

potential independence.

exhibits

3. At potentials greater than zero and at pH values

around 4.0-4.5, aqueous oxidation of pyrrhotite may also

result in the production of ferric hydroxide and elemental

sulphur and/or ferric hydroxide and sulphuric acid:

FeS + 1.5H2o + 0.7502 -----> Fe(OH) 3 + s 0 (V.3.7)

Fe7S9 + 10.5H20 + 5.2502 ----->7Fe(OH)3 + ss0 (V.3.8)

and/or

FeS + 2.5H20 + 2.2502 -----> Fe(OH)3 + H2S04 (V.3.9)

Fe7S9 + 18.5H20 + 17.2502 -----> 7Fe(OH)3 + 8H2S04

{V.3.10)

These processes may occur through the

steps:-

a) An initial oxidation of pyrrhotite itself:

FeS + 2H+ + 0.502 -----> Fe2 + + s 0 + HzO

105

following

(V.3.11)

b)

hydroxide

Direct oxidation of ferrous ions to

rather than oxidation of ferrous ions

ferric

and

hydrolysis and precipitation of ferric ions:

Fe2 + + 2.5H20 + 0.2502 -----> Fe(OH)3 + 2H+ (V.3.12)

Adding equations (V.3.11) and (V.3.12), an identical

equation to (V.3.7) or (V.3.8) is obtained for stoichiomet

ric pyrrhotite and monoclinic pyrrhotite, respectively. Thus

these reaction products, ferric hydroxide and elemental

sulphur, seems to prevail at potential greater than zero and

at pH values around 4.0-4.5.

and

c) If oxidation of elemental sulphur to sulphur acid

occurs, the following equation may be added to equation

(V.3.9) and (V.3.10):

(V.3.13)

It is widely recognized that this reaction proceeds

very slowly at ambient temperatures and in the absence of

sulphur-oxidising microorganisms. Thus, adding equations a,

b, c, and d; and simplifying gives:-

FeS + 2.5Hz0 + 2.25 Oz -----> Fe(OH)3 + HzS04

(V.3.14)

The initial oxidation of pyrrhotite, step a, could be

confirmed with the following anodic reactions reported by

106

Hamilton et al.(17,101). Voltammograms for pyrrhotite at

4.6 pH and at 20 mV s- 1 were interpreted in terms of reac

tions forming elemental sulphur and sulphate. Below 0.2 V

(SHE), the reactions considered were (17,101):

Fes1 . 13 -----> Fe2 + + 1.13S + 2e

Fes1.13 + 4.52H2o ----->

Fe2 + + 1.13S042 - + 9.04H+ + 8.78e-

(V.3.15)

(V.3.16)

At higher potentials, the stable iron species at pH

4.6 was ferric hydroxide and hence the reactions would be:

Fesl.13 -----> Fe(OH)3 + 1.13S + 3e- (V.3.17)

Fes1 _13 + 4.52H2o ----->

Fe(OH) 3 + 1.13So42- + 9.04H+ + 9.78e- (V.3.18)

These reactions, equations (V.3.15) to (V.3.18) also

describe the potential-pH diagrams constructed by Hamilton

et al. and do agree with the chemical reactions discussed

above for this study. The potential-pH diagrams constructed

by Hamilton et al. (17) are reproduced in Figure (V.2.3).

Direct oxidation of ferrous ions to ferric hydroxide

at 4.6 pH value is referred to the study of Hamilton et al.

(17, 101) . A step on the voltammograph at 0.4 V (SHE) (17)

was assigned to the following reaction:

Fe2 + + 3H20----> Fe(OH) 3 + 3H+ + e- (V.3.19)

107

This half-reaction was added with the reduction of

oxygen to obtain equation (V.3.12). However, the second

publication of Hamilton et al. (101) indicates that this

reaction, equation (V.3.19), takes place at 0.2 V. Thus an

accurate potential at which this oxidation reaction occurs

is not yet determined.

4. Pyrrhotite minerals may also be leached by ferric

ions followed by re-oxidation of ferrous to ferric ions and

precipitation of ferric hydroxide, according to reactions:-

108

Stoichiometric pyrrhotite, FeS:

2FeS + 4 Fe3 + -----> 6Fe2 + +

2Fe2 + + 5Hz0 + 0.502 -----> 2Fe(OH)3 + 4H+

4Fe2 + + -----) 4Fe3 + + 2H 0 2

(V.3.20)

(V.3.21)

(V.3.22)

summing equations (V.3.20), (V.3.21) and (V.3.22)

2FeS + 3H20 + 1.502 -----) 2Fe(OH)3 + 2s0

Monoclinic pyrrhotite, Fe7s 8 :

-----> 21Fe2 + + ss0

4Fe2 + + 10Hz0 +Oz-----) 4Fe(OH)3 + 8H+

8Fe2 + + 202 + 8H+ -----> 8Fe3 + + H2o

(V.3.23)

(V.3.24)

(V.3.25)

(V.3.26)

similarly, summing equations (V.3.24), (V.3.25) and

(V.3.25)

-----> 252Fe2 + + 96S0 (V.3.27)

252Fe2 + + 126Hz0 + 6302 -----> 84Fe(OH)3 + 168Fe3 +

(V.3.28)

12Fe7S9 + 126Hz0 + 6302 -----) 84Fe(OH)3 + 96S0

(V.3.29)

simplifying:

Fe7S9 + 10.5Hz0 + 5.2502 -----> 7Fe(OH)3 + ss0

(V.3.30)

From these reactions, it can be deduced that oxidation

of ferrous ions, and hydrolysis and precipitation of ferric

ions, as well as the oxidation of sulphide ions play an

important role on the rate and extent of oxidation leaching

of pyrrhotite minerals themselves. Oxidation of hydro-

gen sulphide, H2s (S=-II), by oxygen may produce one or more

higher valence state sulphur compounds such as: s 0 (0),

s 2o 32-(+2), so3

2 - (+4) and so42 - (+6), from which only the

109

-II, 0 and +VI are "truly stable states" (96). A potential

pH diagram for the metastable sulphur-water system at 2s0 c,

Figure V.3.1, in which the +VI valence state is not consid

ered indicates the region of predominance of these metast-

able compounds (85,96). In the pH range of 4 to 6, the

species which might arise from the oxidation of hydrogen

sulphide consist of elemental sulphur, thiosulphate,

and thionates, HSo3 -.

SO 2 -2 3

Oxidation of sulphide ions, in solution, is affected

by several variables such as temperature, pH, sulphide ion

concentration, oxygen concentration, the presence of bacte

ria and neutral salts (96) as well as catalysts such as Pb,

Mno2 , etc.

Although there is disagreement on whether pyrite or

pyrrhotite produces more oxi-sulphur species (17,100) under

the same oxidation leaching conditions, the extent of oxida

tion of the sulphide ions depends on the ratio of molecular

oxygen to sulphide ion concentration, as well as the pH of

the system. In the concentration range of HS- lower than

2x10-3 M, the reaction product is elemental sulphur but at

concentration of 0.002M<[HS-]<0.003M a clear solution con

taining sulphur oxy-anions result. The oxidation is also

catalyzed by presence of metal ions even when a few parts

per million are present. The catalytic effect of metal ions

has been reported to be in the order Co>Ni>Mn>Cu=Fe (100).

Chen et al. (97) state that the rate of oxidation of

sulphide ion follows a complex relationship with pH. "The

rate increases from pH 6, reaching a maximum value at 8.5,

110

0·6

0·4

0·2

' ' 0 ' ...

IJJ

-0·2

-0·4

-0·6

-0·8

0

FIG." •?•l:

HS05

' ..........

........ .........

' ........ ........

HS-........

' ' ........

'

2 4 6 8 10 12

pH

Eh-pH Diagram for the Metastable S~lphuL 0

System at 25 C (8~.

14

declining to a minimum at 9.3 and attains a second maximum at

pH 11.5". It is also known (97) that the oxidation rate

between pH 2.2 and 6.5 is very slow.

The oxidation rate of elemental sulphur is greatly

dependent on the temperature and oxygen partial pressure.

Below the melting point of sulphur [120°c (52)) the oxida

tion rate is extremely slow and is chemically controlled

with an activation energy of 11.7 Kcal/mole (98).

Although potential-pH diagrams have proven to be very

useful in considering hydrometallurgical processes, they

have several shortcomings since the potential is considered

to be an independent variable (99). Angus et al.(99) argue

that potential is not a conserved quantity and consequently

it is impossible to locate a point on a potential-pH diagram

knowing a composition and pH. Secondly, small changes in

chemical composition can produce large potential differences

and, conversely, when two solid phases are present in equi

librium large composition changes do not produce a change in

potential. Thus Angus et al. (99) replaced the electrochemi

cal potential by its conjugate thermodynamic variable, the

average number of electrons per atom of the active element,

to produce a type of diagram similar to a conventional

metallurgical phase diagram. This type of diagram could be

developed further to study hydrometallurgical processes. A

diagram for the copper-water system is shown in Figure

V.3.2. At pH=O, the left hand solubility curve shows two

regions: one where cu1+ is the dominant species and the

other where cu2+ is dominant.

111

... -::, u

II

10

9

8

7

6

I l)l'i. o I

E • 0.102 V --------

E • 0.185 V

I

I I I

1I AQUEOUS 1

I

>I >1 ~· •I -1

~I Nf

WI i:.1 I

Q 5 ---------------.,.

.!2 4 I

3

2

Ir Cu o

8 AQUEOUS AQUEOUS

0 t__ _____ ---1, ______ ......_ _____ ___

0 2

ELECTRON NUMBER , Z

Figure 1/.3.Z. Electron number diagram for copper syste!T). Intersection of three dimensional figure ·with pH = 0 p 1 ane(.99).

PART C: EXPERIMENTAL PROGRAM

CHAPTER VI

VI.1 INTRODUCTION

Waste sulphide concentrate by Renison Bell, Tasmania

containing approximately 0.45% tin represents a loss of

about 19% of tin fed to the concentrator. This loss is

principally the result of an intimate association of cassit

erite pyrrhotite. The main objectives of the present study

are to modify the sulphide concentrate by chemical and/or

bacterial oxidation in order to release the tin mineral from

the cassiterite-pyrrhotite assemblage so that it becomes

amenable to recovery by physical separation processes, and

pyrrhotite minerals are converted to insoluble compounds

that are unlikely to produce environmental problems.

The mineralogical composition of this mineral was

reported in a previous study realized by Harris et al. (79).

The sulphide/cassiterite association comprise some 50 per

cent of mainly pyrrhotite but with minor amounts of chal

copyrite, pyrite, arsenopyrite, marcasite, sphalerite and

galena. Dolomite, siderite, quartz and other silicates

constitute the remaining gangue. The intimate association of

cassiterite mineral with sulphide minerals (and to a lesser

extent with silicates) essentially at particle size less

than 150 microns and the wide variability of the gange

minerals suite in the sill ores are the major factors influ-

112

encing to have up to 0.4% Sn in the waste concentrate.

Chemical analysis of this waste concentrate was reported in

an earlier study by Sawe (124) as follow,%: Fe 34.96, Cu

0.28, Ca 1.37, Mg 2.49, Pb 0.22, As 2.45, Zn 0.08, and Sn

0.40 (moisture 1.29%}.

Determination of the rate of aqueous oxidation of a

sulphide mineral is usually made by the measurement of the

accumulation of metallic and/or non-metallic cornponent(s} in

solution, e.g. iron, zinc and arsenic from bacterial oxida-

tion of pyrite-sphalerite-free arsenic mineral mixture

(116). It is also known that above pH 2, the rate of oxida

tion of ferrous ions and precipitation of ferric hydroxide

are enhanced. The use of this method to measure the rate of

aqueous oxidation of a sulphide mineral would not be suit

able to solutions where the pH value is higher than 2.

Another method used to determine the rate of aqueous

oxidation of a sulphide mineral is to measure the consump-

tion of oxygen. Vancelow (105) used an oxygraph to measure

the rate of aqueous oxidation of copper sulphide minerals

from bacterial oxidation. Bailey et al. (16) determined the

consumption of oxygen, produced electrolytically, during a

pressure leaching of pyrite. The use of this method assumed

that the reaction stoichiometry is fixed and known. Thus

the selection of the method to determine the rate of reac

tion of iron sulphide minerals depends on the pH.

At pH values higher than 2, determination of the rate

of aqueous oxidation of a sulphide mineral, e.g. pyrrhotite

minerals by the consumption of oxygen seems to be more

113

predictable than determination of any other reactant or

product. Moreover, the use of an oxygraph is easier and

gives continuous data. If determination of elemental sulphur

is used, there is also some concern about the solubility of

amorphous elemental sulphur in tetrachloroethylene. It has

been reported by Warren et al. (11) that amorphous elemental

sulphur, present from the aqueous oxidation of chalcopyrite,

is not soluble in carbon disulfide, cs2 . It is not known

whether some elemental sulphur formed during the aqueous

oxidation of pyrrhotite minerals is amorphous; if it is,

then if it is soluble in tetrachloroethylene.

In the present study, an attempt was made to measure

the rate of aqueous oxidation of a pyrrhotite concentrate

tailings from Renison Bell, Tasmania by continuously measur

ing the consumption of oxygen. In a preliminary study of

the aqueous oxidation of this mineral conducted by Harris et

al. (79); a series of stirred reactors and "shake flask"

experiments were tested in the presence of microorganisms.

At 4.5 controlled pH, the rate and extent of oxidation of

this reaction was determined by analysis of the leached

residue for total sulphur, sulphate sulphur and elemental

sulphur. The supernantant leachate was also analyzed for

ferrous and ferric ions.

The lack of reproducibility in terms of mass balance

and degree of pyrrhotite oxidation obtained by Harris et al.

(79) may be explained in terms of the oxidation of elemental

sulphur to sulphate ions and if the pH of the system is

controlled at 4.5 pH, jarosite compound may be formed which

114

has not been accounted for. Thus determination of the

consumption rate of oxygen during the aqueous oxidation of

sulphide minerals, e.g. pyrrhotite minerals seems a more

useful alternative at pH values above 2.0. Nevertheless, if

oxidation of elemental sulphur and/or oxidation of other

sulphides minerals are involved, the present method will

also account for the consumption of such oxygen.

The information about the kinetics and mechanism of the

aqueous oxidation of pyrrhotite minerals obtained, was

intended to add to the fundamental knowledge of the aqueous

oxidation of sulphide minerals and to assist in the develop

ment of a more viable process for the treatment of pyrrho

tite minerals containing valuable metals/minerals, such as

gold and cassiterite.

Particular attention was given to the design of the

equipment to determine the reproducibility of the measured

reaction rate and the range of oxidation reaction products

which are formed. Three equipment designs, called 'oxy

graphs' were developed to obtain a continuous measurement of

the reaction rate. Details of the designs are explained and

experimented in Chapter VII. However, it is pointed out

that the oxygraph only gives the rate of the aqueous oxida

tion. The assumption made is that a constant stoichiometry

stated as Equation VI.1.5 and/or Equation VI.1.12 will take

place constantly. If the actual reaction is stepwise, i.e.

oxidation of elemental sulphur to sulphate ions then precip

itation of jarosite compounds, which involve major consump

tion of oxygen, then the above assumption will not be cor-

115

rect.

As the leaching process takes place in a closed system,

isolated from the atmospheric pressure at constant tempera

ture; opening the system e.g. for sampling, will disturb the

internal pressure, as the atmospheric pressure could be

higher or lower than that at the beginning of the process.

Thus the number of samplings were reduced to two; the first,

on the third day of the process and last, at the termination

of the experiment. The progress of the reaction is indicated

by the rate of uptake of oxygen, as indicated by the time

the electrolytic cell is generating oxygen.

At 4.0 initial pH, the important partial reactions

which may take place during the aqueous oxidation of pyrrho

tite can be summarized as:

116

a. Acid dissolution of pyrrhotite:

FeS + 2H+ -----> Fe2 + + H2S (predominant)

FeS + H+ -----> Fe2 + + HS- (VI.1.2)

b. oxidation of hydrogen sulphide:

H2S + 0.5 02 -----> so + H2o

(metal catalyzed, cu2 + '

Fe2 + ( 97) )

c. oxidation of ferrous ions:

Fe2 + + 1.5H20 + 0.2502 -----> FeOOH + 2H+

Thus, the overall reaction would be, a+ b + c

FeS + 0.5H20 + 0.7502 -----> FeOOH + s 0

(VI.1.1)

(VI.1.3)

(VI.1.4)

(VI.1.5)

d. if oxidation of elemental sulphur takes place, the

additional consumption of oxygen would be:

s 0 + H2o + 1.5 Oz-----> HzS04 (VI.1.6)

However, it is known that this last reaction is very slow at

ambient temperature and it is unlikely to take place in

the absence of bacteria. Thus the overall reaction would

be; a+ b + c + d + e,

FeS + 1.5H20 + 2.2502 -----> FeOOH + H2so4 (VI.1.7)

It can be seen that stoichiometric pyrrhotite consumes

three times more oxygen and water when its reaction product

is sulphuric acid rather than elemental sulphur.

Similarly, aqueous oxidation of monoclinic pyrrhotite

may be considered as:

117

a. Fe7Ss + 14H+ -----> 7Fe2 + + 7H2S + s 0 (VI.1.9)

2+ + b. 7Fe + 1.7502 + 10.5H20 ----->7FeOOH + 14H

(VI.1.10)

c. (VI.1.11)

the overall reaction is

Fe7Ss + 3.5H20 + 5.2502 -----> 7FeOOH + ss0 (VI.1.12)

If sulphate ion is also formed, the overall reaction is

Fe7s8 + 11.5H2o + 17.2502 -----> 7FeOOH + SH2so4

{VI.1.13)

Thus, if the mass of the mineral is considered; one

gram of stoichiometric pyrrhotite and monoclinic pyrrhotite

consumes 0.2730 and 0.2595 grams, respectively of oxygen to

complete the reaction. This mass represents 208.8 and 198.5

cubic centimeters of oxygen for stoichiometric pyrrhotite

and monoclinic pyrrhotite, respectively (R=0.0821

L.atm/mol.°K)

It should be noted that no nett consumption or genera

tion of hydrogen ions is involved with either minerals when

ferric hydroxide and elemental sulphur are the only reaction

products. If oxidation of elemental sulphur takes place,

sulphuric acid may also be formed. These observations also

apply to hexagonal pyrrhotite, Fe9s 10 . However, elemental

sulphur has been shown to be the dominant product formed

from the oxidation of pyrrhotite at pH 4.6 although some

sulphate is also formed (17,101).

As the aqueous oxidation of pyrrhotite is heterogene

ous process, the slow steps in the series of reactions may

also be heterogeneous one involving solid-liquid and/or

118

solid-gas interfaces. However, in some cases, the slow step

in a heterogeneous process may be homogeneous, for example

the chemical oxidation of a soluble reactant to a soluble

product is involved, e.g. ferrous to ferric ions with dis

solved oxygen. It is not certain whether diffusion, as a

rate controlling step, can be eliminated even with the

vigorous agitation necessary to keep all the particles in

suspension. Diffusion across boundary layer of solution

associated with the solid surface may still be the rate

controlling step. Nevertheless, stirring as a variable, may

be eliminated provided great enough agitation is applied to

the leaching system and the rate does not increase with

increased stirring.

Continuous bacterial oxidation of pyrite concentrate

at miniplant scale {6reactors x 30 L capacity) developed at

Aurotech N.L. is discussed. In this study, the effect of

bacterial recycle whether with the slurry or liquor, and

dissolved oxygen on the rate of mineral degradation is

examined.

119

CHAPTER VII

VII.1 INITIAL EQUIPMENT DESIGN

The rate of consumption of oxygen can be used to deter

mine the oxidation rate of sulphide minerals at a given pH,

temperature and oxygen potential provided the stoichiometry

of the reaction is constant and known. Thus an 'oxygraph',

in principle, could be used to determine the rate of oxygen

consumption during the aqueous oxidation of pyrrhotite

minerals assuming the stoichiometry given in Equations

(VI.1.5) and/or (VI.1.12) take place.

FeS + 0.5H20 + 0.7502 -----> FeOOH + s 0

Fe7Sa + 3.5H20 + 5.2502 --->7FeOOH + as0

An 'oxygraph' would basically consist of:

1. a closed reactor for the aqueous oxidation

process itself,

2. a source of oxygen, which in this case was

produced electrolytically, and

3. an instrument to measure the oxygen consumed.

VII.1.1 PROCEDURE FOR DETERMINING THE RATE

CONSUMPTION OF OXYGEN

OF

The initial equipment design was based on maintaining a

constant oxygen pressure (and hence volume) in a closed (or

120

constant volume) reactor vessel. The oxygen consumed by the

aqueous oxidation was replaced almost instantaneously by

oxygen produced electrolytically in a separate reactor, in

order to maintain a constant pressure in the system.

The internal pressure in the leaching system was meas

ured with the aid of a manometer containing an electrical

contact. When the pressure in the cell decreased, the con

tact was 'broken' and this caused a constant current to flow

and generate oxygen in the electrolytic cell. When the

pressure increased, the contact was 'made' causing the

current, and hence oxygen generation, to cease. The rate of

oxygen consumption was therefore directly related to the

ON/OFF times of a constant current. Once the ON time start

ed, it was not turned off for a given time; thus uncertainty

in the measurement of small ON times or "noise" or "chatter"

were avoided. Details of the first oxygraph are shown in

Figure VII.1.

Solutions were prepared quantitatively in a volumetric

flask (lL) using analytical reagent grade chemicals. The

oxygen was produced electrolytically from acidified dis

tilled water. A "Towson and Mercer" hot-plate magnetic

stirrer was used to maintain the slurry in suspension.

A Rockwell AIM-65 computer recorded and stored the ON

and OFF times of the electrolytic cell. The data (maximum

1000 readings) was transferred to a PDP11 computer for

further calculations.

As the fresh waste concentrate samples contained in

fastened plastic bags and in drums of 25 Kg. were wet, a

121

FIGURE VII.1: First Oxygraph Design.

555 Timer

AIM-65

~-----------..

·-- ..... '

(B (A)

(D)

(A) Reactor Vessel (B) Electrolytic Cell (C) Manometer

(D) Magnetic Stirrer

sample of about 2 Kg. was taken to be washed with alcohol

and then several times with acetone and dried in a vacuum

for about 4 weeks. Once dried, the sample was ground in a

mortar and passed 100 % through a 150 mesh particle size.

These samples were kept in small plastic bags in a desicca

tor all the time.

A 5 per cent pulp density weight per volume of pyr

rhotite was used in each experiment, unless otherwise stat

ed. The sample studied in this experiment was identified as

Number 2 (79).

FIGURE VII.1

(A) Reaction Vessel: A vertical glass cylinder with

ground glass stopper (B 29/32 socket) with flat base was

used as the reaction vessel. A vertical glass rod was fixed

from the middle of the stopper with its end slightly bent

before a flattened section. The purpose of the rod was to

ensure good mixing of the slurry, which was stirred magneti

cally.

(B) Electrolytic Cell: Electrolysis of a dilute solu

tion of sulphuric acid produced the oxygen necessary to

continue the aqueous oxidation process. A platinum wire (the

anode) was installed inside the vertical glass tube (base

open). This vertical tube (which collected all the oxygen

produced) was connected to the reaction vessel by means of a

capillary tube. The cathode (platinum foil) was installed in

a 200 ml. beaker close against its wall. Hydrogen molecules

122

must be released quantitatively outside the system. A con

stant current of 80 mA was passed through the electrolytic

cell for a minimum period of 90 seconds per ON time, giving

0.0006 gr. oxygen/ON time.

(C) Manometer: The U tube attached to the reactor

vessel contained very dilute acidified water (2 drops of

concentrated sulphuric acid in 20 ml. distilled water) and

two platinum electrodes. In the left-side of the U tube, one

platinum electrode was immersed in the acidified water and

exited through a sealed socket (B 7/16 socket). This side of

the manometer was connected to the internal pressure of the

reactor vessel. The right-side electrode just touched the

water surface at the center of the glass tube; its vertical

position could be adjusted by a threaded screw connected to

the top of the socket (B 7/16 socket). This side of the U

tube was opened to the atmosphere.

The internal section of the system was completely

sealed from the external atmospheric pressure using high

vacuum silicone grease on ground glass joints.

VII.1.2 EVALUATION OF THE FIRST OXYGRAPH

A graph showing the rate of oxygen consumption of 1.5

grams of pyrrhotite at 4.5 pH is presented in Figure VII.1.2

and VII.1.3. As little information about the oxidation

behavior of pyrrhotite could be obtained from this design

due to the small reactor vessel and to the disturbance of

the manometer by stirring, a second larger 'oxygraph' and

123

1. ....... _-------------------------------,

-, L

0 H

f- .s

.8

Q. ~-J

~ 0

J en rJ) ~ .6 z ~ D a

u Z . 4 w Q )-

x D . 2

li Ll

w f-< er

0

FIGURE VII.1.2: Rate of Oxidation of 1.5 g. of Pyrrhotite at 4.5 pH.

1000 2000 3000 4000

Absolute Time (Min . )

5000

1------------------------------, FIGURE VII.1.3: Rate of Oxidation of Pyrrhotite at 4.5 pH.

(1.5 g. of mineral in 30 ml. of solution) Z I Moving Average: 5

.8 D H

~ ~ ~ J ~ z D u

C ·-E 'N 0

01

...:r .6 0 D D D

Z .4 w ~ > X D .2 ~ r, ~

,. ' UJ

~ <{ er:

0 10000 20000 30000 40000

ABSOLUTE TIME (Min.)

50000

electrolytic cell were designed.

VII.2 SECOND EQUIPMENT DESIGN

VII.2.1 PROCEDURE


FOR DETERMINING THE RATE OF

The second equipment design is shown in Figure VII.2. A

one liter quick-fit reactor vessel was used and the oxygen

was produced electrolytically from a solution of potassium

hydroxide. Two kinds of pyrrhotite mineral samples were

processed in this series of experiments: Sample O and Sample

2. Identification of these samples was stated by Harris et

al. (79). Thirty grams of pyrrhotite mineral in 600 ml. of

pH 4.0 distilled water (unless otherwise stated), represent

ing 5 per cent (w/v) of pulp density was used in every

experiment.

FIGURE VII.2

(A) Reactor Vessel: The quick-fit reactor vessel had a

spherical bottom and five necks. A variable speed "Vibramix"

stirrer, inserted through the middle neck, was used to

agitate the slurry and disperse air/oxygen. The "Vibramix"

operated at about 200V. A peristaltic pump was installed to

recirculate the air/oxygen from the internal surface reactor

vessel through the shaft of the "Vibramix" in order to

124

.. __

motor vibramix 1111

[ air pump

• . I.. 1 . . ~· . ' . ~ • C

• r _ r' ,~ • w

. ~ (A) Reactor Vessel

e e

r---+==+=1r-----====isss

,·

i t) ~ I (C) manometer .~ ....

~ I g Q

l ) (B) Electrolytic Cell

FIGURE VII.2: SECOND OXYGRAPH DESIGN

Timer

AIM-65

Rockwell

I PDP-11

increase the oxygen mass transfer to the slurry. The third

neck was used to connect the reactor vessel to the electro

lytic cell. The fourth neck was used occasionally to make

additions, such as copper ions and sodium chloride, through

a seal without affecting the internal volume or pressure of

the system.

(B) Electrolytic Cell: The cell consisted of a U shaped

glass vessel designed specifically for this system. The

electrolyte was a solution of 40 percent of

droxide with nickel wires as electrodes. The

potassium

left-side

hy

of

the electrolytic cell contained the anode and anolyte and

was connected to the reactor vessel. The right-side, con

taining the cathode, was open to the atmospheric pressure.

The cell was closed with a B 34/35 and a B7/16 socket sup

porting the electrodes. A constant current of about 150 mA

was used for the electrolysis. It was observed that an ON

time value of 110 seconds was satisfactory to supply the

oxygen necessary for the reaction giving 0.0014 grams of

oxygen per ON time.

A small glass 'U' tube containing slightly acidified

water and platinum electrodes was used as the controlling

manometer. The left arm of this 'U' tube was connected to

the leaching cell through the electrolytic cell; and the

right-side to atmospheric pressure through the electrolytic

cell. Both pairs of electrodes (electrolytic cell and

manometer) were connected independently to the controller

timer.

The whole system was covered by a plastic sheet and

125

heated, when necessary, by an infrared lamp. This lamp,

connected to a thermostat and a temperature controller,

heated the system, when required, up to 3o0 c during the

process.

T. ferrooxidans (BJR-Kl) isolated and acclimatized to

sulphide concentrate by Khalid (125) was inoculated, if

required, to the leaching system. The 9K media composition

was prepared for 600 ml. as described by Sawe (126) and

added in solid form (K2HP04 , 0.5g.; MgS04 .7H2o, 0.5 g.;

(NH4 )2so4 3.0g; Ca(N03 )2 .4H2o, O.Olg and KCl, O.lg for litre

of solution) The bacteria growth was examined through a

microscope at the completion of every experiment. The pH

was read by an Orion Research Ion Analyzer EA 940 at comple

tion of the leaching process. The number of samples was

limited to one, at completion of the experiment.

VII.2.2 EVALUATION OF THE SECOND OXYGRAPH

The rate of oxygen consumption obtained during experi

ments 1, 5, 6 and 7 can be seen from Figures VII.2.1,

VII.2.2, VII.2.3 and VII.2.4, respectively. The typical

leach curves still include cyclical variations which are

attributed to changes in ambient temperature and pressure

during the day. For example, Figure VII.2.4 shows the start

ing time of the process at 7.20 P.M. An apparent decrease,

then increase, in the oxygen consumption rate from about

7.20 AM, the following morning, until about 3.20 PM is noted

throughout the experiment.

126

~=~==~~==:::::::=====--:'.'.'-._::_ ==-------------------··--·----·---------------

1-~-~. r ~

-

0 z

~~

.µ

C

£ QJ E

...... c... QJ Q

. X

w

£ .. .,-

N

>-<

s w

? n::: .:J ~

<

>-<

lJ

_

~~

~

-CD

tO

(\J

0 (\.l

0 0 0 If) ..-1

0 0 0 0 ~

0 0 0 lf)

2 ·u1w

/ 0 · 6 80 JO· O

• N

D

L

.L d

W n

S N

D

::=J

N

3 8

A)<

D

-=:I D

3 _.L

\:I ?::I

---z H

~-L..

..._ __ ,,

UJ

L

H

t--

LU

1 r----, _

__

,;

_j

0 U)

OJ

<r

C ·g .9

-----cf'-J

OJ

-:r ['\J 0 0

0

. z 0 H

r-0.. ~ J (fJ

z 0 u z w ~ >x 0 lL 0 w r~ er:

.8

.7

.6

.5

.4

.3

.2

. 1

0 0

FIGURE VII.2.2: Experiment 5

""'· ... t \

T

: "'... :· .. . .:: ., : ·:.:~:-:-. .. ,_ ...... ·~·.,.~-: • .,,.:_JI'_··.: • ·"'-· ... -.~ ... . -,...,,... . ...... .. . . ..

1440

1

. · .. ....

2880

2

.

.. ·~·· ..... _ ·... •. . -..... · ..... ·.. . .... ...... ·· ................ _ ... _. __ ..... ,_ .. .. -............

4320 5760 7200

3 4 '.:>

. .': ..... .. ...

··.. ...

8640

6

. ... . , . .. .... ..-......... .

10080

7

ABSOLUTE TIME (MIN.)

--... ··· ...... .

11520

8

.. ... •' ..

..,. ...

12960 14.i:

9 DAY-

C

~~ . 9 Dis!

O')

~ .8 CJ CJ

0

. z 0 H

~ 0. ~ J G~ t_

0 r 1 \..,/

z w ~ > X 0 lL 0 w ~ <r: ('(

.7

.6

.5

.4

.3

.2

. 1

0

"·'· -. "\

. .. .......... .. . ..... - .. . .............. .

•,

0 1440

1

FIGURE VII.2.3: [xperiment !\Jo. 6

Bacteria Inoculation

7 .. ......... ' . .. .

-..... ...... · .. ··

2880

2

·····

4320

3

... ··..

······· .. . .

5760

4

ABSOLUTE TIME(MIN.)

7200 8640

s 6 DAYS

.s . 9 E -,

0~

OJ

co 0 0 0 0

. z 0

'. 8

.7

C .s I

0. ~ J /("\ UI

z 0 u z

.5

.4

W .3 ~

B . 2 / ~· _io f

7.2iJAtA..'-ro ~.2o fr'!

'~

~· )- t . !JT~ll.li ~ ,i""e

lL -~,_-

f. : f . ··1·- .: ..,..::--: .. , ... · ..... t,~ ... -..... "':' _ _. ___ ~-- -- -.:.--":- ·~---- -{·-------... __ } _____ ----t ··1 ·r--· --- .. ='. , . .,..•. ,, ·.. . ·'·· . . . .. . ,., . . - -0

w f-

:· . ~: '.· -, -··:' · .... -•.•· ·--:- . . . •... · . . 1 + , - . .. ··.,,

< er:

0 0 1440 2880

1 2

4320 5760

3 4

. . .

7200

5


Slope = m = -0.00C00744

. .': ... '\ ,. .. ·-. . . ....

;_

8640

6

-6 =-7.44 X 10

:... .. .. . . . ', -· ..... . . . . . .... ·. . .·. .. ......... . ···.. ·,

10080 11520

7 8

ABSOLUTE TIME {MIN.J ' ... ,

. ,.._ , . .. •,

- .... •,•

12960 14.i:

9 DAYS

Although the whole equipment was covered with plastic

sheets and heated by infrared lamp, when necessary, the

systematic variations are ascribed to the ambient tempera

ture rise in the mornings to above 30°c and, reaching about

3s0 c during the day. As the equipment was installed at the

east side window and top floor of the building, the ambient

temperature above 30°c was uncontrollable.

It is also known that atmospheric pressure (reported

daily by the Bureau of Meteorology) changes in the range

from about 996 to 1028 hectopascals. Because the right side

of the of the electrolytic cell was opened to the atmospher

ic pressure, it affected the manometer (small U tube). This

behavior was also uncontrollable on this 'oxygragh'. Never

theless, as Figures VII.2.1, VII.2.2, VII.2.3 and VII.2.4

still show consistently an overall linearity throughout the

experiments, these results will be discussed in detail

later.

A third equipment design was constructed in order to

eliminate these systematic variations due to temperature and

atmospheric pressure changes. But, as the Rockwell AIM-65

computer became unserviceable, a "Microbee" computer was

installed to store the data.

VII.3 THIRD EQUIPMENT DESIGN

VII.3.1 PROCEDURE FOR DETERMINING THE RATE OF


127

The third equipment design attempted to make the

'oxygraph' results independent of changes of atmospheric

pressure and ambient temperature. A constant atmospheric

pressure equivalent to the onset of the process and a con

stant temperature were maintained by sealing off the leach

ing system and enclosing it in a constant temperature

environment. Figure VII.3.1 shows the equipment in more

detail. It consisted of:

(A) Reactor Vessel: The same reactor vessel and agita

tor described for the second equipment was used. An air

sampler pump, Kimoto HS-6, was modified to recirculate air

or oxygen at about 9 1/min.

{B) Electrolytic Cell: Oxygen was produced by the

electrolysis of an acidified copper sulphate solution. It is

known that the electrode potential of copper and oxygen are

described by the following equations (110):

cu2 + + 2e- -----> cu0 E0 =0.3402 v

2H20 ----->Oz+ 4H+ + 4e- E0 =-1.229 V

(VII.3.1)

(VII.3.2)

(VII.3.3)

This reaction is forced to take place by a constant

current of 480 mA passed through the cell. In this process,

the anode reaction is the formation of oxygen and the

cathode reaction is the deposition of copper. Every experi

ment was started with a new copper sulphate solution con

taining at least 200 g/L of copper sulphate (solubility of

128

(11)

e

(5)

(4) (1) i

\ .. 0

(6) 0 •11 t, ~ .

(9) i.:..u--

FIGURE VII.3.1: THTRD OXYGRAPH DESIGN.

e

555 1 1 ~ Cu wire

p: / l'>Pt foil

1 @ (8)

(1) Reactor Vessel

(2) Electrolytic Cell

(3) Magnetic Stirrer

(4) Hg Manometer

(5) Air Pump

(6) Temperature Recorder·

Microbee Computer

IBM-PC

(7) Heating system

(8) fan

(9) Pt contacts

(10) Ammeter

(11) Vibramix

FlGURE VII. 3 .1: THIRD OXYGRAPH DESIGN

copper sulphate, cuso4 .5H2o in cold water is 24.3 in 100

parts (35)).

The anode was a platinum foil (2.5 cm x 6 cm, approx.)

and the cathode a copper wire. Therefore, hydrogen evolution

would not occur provided the copper concentration is high.

The electrical circuit was the same as that previously

described and it is detailed in Appendix H. The electrolytic

cell was connected to the leaching reactor by a short length

of thick walled high quality rubber tubing. The electrolyte

was gently agitated with a magnetic stirrer.

(C) Manometer: A conventional glass mercury manometer

(5mm internal diameter) was used to measure the pressure in

the system. The manometer was filled with mercury under

vacuum, the glass at the top was flame sealed. The bottom

part had three necks; two for the platinum electrodes used

to measure the height of the mercury interface, and the

third connected to the reactor system. One electrode was

immersed in the mercury and the other placed less than one

millimeter above the center of the mercury surface. These

electrodes were sealed at the sockets (B 7/16) with high

vacuum grease and black wax.

The ON/OFF time of the electrolysing current was

registered and stored in a "Microbee" computer,

able to save up to 2,500 readings. The data was

to a PDP-11 computer for processing, analysis and

which was

transferred

plotting.

As this computer became unserviceable, the data were trans

ferred to an IBM compatible PC.

Temperature of the environment was controlled by a

129

thermostat at 30.5(+/-1°c). Heat was supplied to the system

by means of an infrared heat lamp, and a fan was used to

ensure a uniform environmental temperature. It has to be

mentioned that the temperature in the environment increased

up to about 35°c due to the heat produced by the "Vibramix"

motor. To overcome this problem, a five centimeter hole was

made through the polystyrene wall, close to the motor, thus

the constant reaction temperature improved to about 31°C.

Ideally, the experimental equipment should have been in

stalled in an air conditioned environment.

Once the solution of sulphuric acid was placed in

leaching reactor, the system was operated for about a

hour in order to obtain a uniform temperature of 30.s0 c

the chamber. Then a sample of the mineral was added and

reactor system sealed off from atmospheric pressure.

VII.3.2 EVALUATION OF THE THIRD OXYGRAPH

the

half

in

the

The curves obtained for the rate of consumption of

oxygen from this 'oxygraph' are shown in Figures VII.3.2,

VII.3.3, VII.3.4, VII.3.5 and VII.3.6 and VII.3.7. It is

noted that the cyclic variations during the rate of oxygen

consumption were not eliminated fully. These cyclic varia

tions are directly related to the temperature variations

during the day. It was difficult to maintain a constant

temperature for extended periods without an air conditioned

environment. The temperature of the general environment

(outside of the system) and the heat produced by the motor

130

C ·g ,, ON

0)

...,. N n 0

D

. z D H

~ 0. ~ J G1 z 0 u z w [~

> y

0 u_ D w ~ er rr:

11 FIGURE Vll.3.2: Experiment 14 - .. ---1

l I

.8

.6

I .4 t . _.:.:., ........ ··.:.: ... . . . .

Total ON Time: 40,029 sec.

Total Reaction Time: 2,792 min ..

Extent of Reaction: 19.45 %

t .· .. ···, ..... ·:. ···,. . . ... : .:.-. . . ·. . ....... · .2 + ,. . . i ... .

tl • ---4--t----t-·-+-

... ··". . .... ·.. .... ......... .. .

··· .. ·. '

.· ... · ......... ·.· .. . . ___ :.:·. . .. · ·" .. · . .,.

o _· 1440 0 720 2160

1

ABSOLUTE TIME (Minutes)

I I

I I

2880

2 DAY~

it ----0,J 0

Ol

...:t N Cl C! D

~

z 0 H

rQ.

~ J (J)

z 0 u z w ~ )-x 0 lL 0 w r-~ er:

1 r

i T

I I

. 8 + I

I . 6 I

i

+ I i ,~ r_:._:. ,

f ':?ftJ, [ifL l T·· · · -.. -.... ~ .,_ · 1 e< . . ·'-\/ '\.,,,. : '/ .

....

2 +. .. ">; .. :..,t,,.,...r.. , ' .. ····;'·:·. . . . ,. <. ,:-:,:~ ..

0 0 1440

1

... '\.

2880

z

-----1 FIGURE VII.3.3: Experiment 15 Total ON Time: 114,617.4 sec. Total Reaction Time: 12,285.9 min. Extent of Reaction: 55.69 %

I I i

l

.,·,, . . ;:::: ...

. . t ,• ......... .

I' •'

.. .. ":· ....

:' /:? j f:'.jt'.? ,,. \:· .... .-.-::.:._!_._· .. :\·:·:_-_.:: ·-:.;·-· • .-..

..........

4320

3

. .....

5760

4

7200 5

8640 6

····:·-.- ... · .. -·. •.,.•,

10080 7

·1 I

11520 12960 8 DAYS

ABSOLUTE TIMt. (Minutes) \

C '§ 'N

0

OJ -;t N 0 0 0

. z D H

fQ.

~ J l8-z D u z w ~ > x D u. D w f<t cc

1

! .Bt.

I

I I

l I

i .GT I i

I . 4 + . •. ·-.:·.·.:· .. .. ' .

.2

I.· I k

I

·-------··------------- --·-1

FIGURE VII.3.4: Experiment 16 Total ON Time: 63,896.65 sec. Total Reaction Time: 4,212.7 min. Extent of Reaction: 31.04 %

··. !';,_::-::: .< :'._./ ·:_:;_,/·:··'': ._·., '•, :· : .:,. .-,./r '."'j:::.•i>'·'- _/i'·' , .. ,:-."·· ··'_::-, ..

. ...

!

:.-. . ,··. . :.:··.: ..

o L-+----+---~--1----t----t--+----'--~--+-_:_~· ·-+--0 1440 2aso 4320

1 2 .3 DAYS


C "§

" cI' OJ

-.:r N C) C)

C)

~

z 0 H

r-0.. ~ J '" Ui

z 0 u z w l'.l >x 0 u. 0

w r~ er

1 T---- --------' i !

I . 8 r

I I

j i

' 6 + . I

I

i I

\

.4 t i

... · ...

---

FIGURE VII.3.5: Experiment 17 Total ON Time: 88,474.27 sec. Total Reaction Time: 8,627.5 min. Extent of Reaction: 42.98 %

,~"'~.· .,. ~ . "•

:'":~~:; =:;)-.-:.

1• • I ,.. . : .... _

.. . 2 J/'{~:'.:}\//\,. :</'f /'.··.:/\\;:/LJ. I ·-

.. ···.· .. <:.:. ..... , .. ;.'\ ... ... ·· .... , .... .,. · .... ·/;_._~: ..... .-... :· .... :.:-.

·------·--··1

I I I

.;·. _::·~;.:.: ..

_\· ,,.

' ... ..

.--·.·. l --+---+-~~ 0 -- 2880

--+-----+----+------1-------l--0 1440 4320

1 2 3

ABSOLUTE TIME

5760 4

(Minutes)

7200 8640 5 6DAVS

C ·g ----... d'-1

OJ

-:r N Cl 0 0

. z D I-<

f-0.. ~ J ([)

Z' D u z w [j )-x D LL 0 w fer 0::

1 r-----------------------·--·-----·------------------- -------------·-----· -·----·------------ --------------· 1 I

I : j t FI~URE VII:3.~: Exp;riment 18 I 1 Total ON Time. 70,697.5 sec. 1 . 8 ~ Total Reaction Time: 4,523.3 min. I I Extent of Reaction: 34.35 % I + I

I t

.6 + \ : :, : .

+ -:>->. ... .... • ·.·:-,, i "?·'":.:·,:· .. I .. •· I···: ....

. 4 r . . ·" r '<\+~:::.<:f ::.t .. /);; /:\s;_.; ::.;; '

2 ), ., .. , .· ,.. .. ..... -

. T .. :\ ·:.: · ... :, .... . ,A•,,

+ i I

... ~·· . •, ..

:,\/',}·:·. :, . ·, ·.·

0 L--1 -1----1---t---+--- · I- I ---!-·--+--- I I ··-I----+-----!-- I I 0 1440

1

2880

2

ABSOLUTE TIME

4320 3

(Minutes)

5760

DAYS

C

~~ ON

OJ

..;r N 0 0 0

z D H

f-0. ~ J rJ)

z 0 u z w (.'.)

>x D LL D w f~ cc

1 ! /\,; -------------------- -----·--~---------- -- . ····-----·---·-·-·-----·-·--1 I . :,·.· T'· ...

I' · \. F;1GURE VIL 3. 7: Experiment 20 I

8 ·::· 1otal ON Time. 114,954 sec I · T ~\: Total Reaction Time: 7, 920 min. 1

, · X Extent of ReacU.on: 55. 85 % I I ;,;.. I

TI s

II , .. ··:~;. . I

..LI, . )\t ... . . ~ ...... :-:.:.

·~ • i • • • ··t.-t • ,.• ·. "·~ ·:,·. ·,·., ·:-~~-.

I ·.,. _. 4 -t- ilf t/,',':\}}: ~. ,

I ·: .. . . .. . : .· ..

. 2

. .. ..

:~?.i_: ·}· .. _:~::?-·.,_:. :::~::::~<,:~.---·.. ..,, ·.:

• .. ·

~-·'· ~. >

... ... · . .. : . ~."·:.

"lo • :: ...... ">.· ·,·: ..... . ..

........ · .. · ......

0 I I I I I -+-1 I 1. -t-+ I l-t------- +--+---t--i---- I I +--+---+- ·-r I I I I I 0 1440 2880 4320 5760 7200

1 2 3 4 5 DAYS

T I M E (MINUTES)

of "Vibramix" (constant) affected the system in the morn

ings.

However, it can be seen that all curves show consist

ently a general linearity, where the rate of consumption of

oxygen was faster at the beginning than that at the end of

experimentation, as would be expected.

The constant measurement of oxygen consumption during

the aqueous oxidation process allows us to determine the

extent of reaction. Thus the program "LEACH.BAS" also

calculates the extension of reaction and is presented for

each experiment in Figures VII.3.8, VII.3.9, VII.3.10,

VII.3.11, VII.3.12. These curves, tending to be parabolic,

also show the behavior of the extent of reaction, which

evidently is more aligned than the rate of oxygen consump

tion curves.

VII.4 OXYGEN CONSUMPTION DURING THE REACTION OF

PYRRHOTITE MINERALS

The rate of oxygen consumption will be analyzed in

three sections: Firstly, the results obtained from the

second oxygraph, where the rate of oxygen consumption was

determined for initial pH values of 5.5, 4.5 and 4.0. The

rate of oxygen consumption was compared when leaching took

place in the presence of bacteria and that against chemical

oxidation for two samples: Sample O and 2.

Secondly, an attempt was made to justify the fact that

131

·,o o'._

LU f-·f---i L--1

0 I :r-e---

'1 1 _ _,__

'~--/

o_

LL 0

~ L.--

0 ~ . . f--<(

D !--l X 0

2 a 1

--------------------------·------------------------------------------------------------·- __ 1

! ~~- i T ! t FIGURE VII.3.8: Experiment 14 I t Extent of Reaction I

1 ~ , 1 _ _, i I , I , I + I + ; i

10f

5

0

t i

+

t l

I

0

-·

720 1440 2160

A.BSOLUTE Titv1E (Minutes)

' ' i I I I

j

l ~

2880

60 --- ·---···

I ~

50 I LU FIGURE VII.3.9; Experiment 15 ------------r--Extent of Reaction r--l

L-I

0 I

40 a:: '.J LL

>-o_

30 ll 0

z 20 0 H ~-I

<t 10 0

ofL I H X

I 0 I I I I I --1 I I I I I I I I I I I 0 1440 2880 4320 5760 7200 8640 10080 11520 12960

A.BSOLUTE TIME (Minutes)

~

w fH f-0 I [[

IT >-0....

LL 0

z 0

40

30

H 10 f-<(

0 H X 0 0

0

FIGURE VII.3.10: Experiment 16 Extent of Reaction:

31.

04 %

720 1440 2160

-· --·--· I I

I

2880 3600 4320

ABSOLUTE TIME (tv1inutes)

'-._O o,

l .iJ 1-f --! ) r·-0 I (I [I

>Q_

LL 0

z 0 1--l !<(

D H X 0

50 .....----· --- ····-------- .... -·--·--··------- -. - -- ------- ·1

40 t

30 i 20 t 10

FIGURE VII.3.11: Experiment 17 Extent of Reaction: 42.98 %

I

0 w:c._-----•--i------t-----f------l 0 1440 2880 4320 5760 7200 8640


'-,0 0"-,

U_l L---., __ l I I

I ·--' 0 -T-__ ,:__

IT [[

>[l

I I L!., __

C.::J

"-7' ,:::_

0 l--i t-<t ii '·---~

l--l

X ,

CJ

5 0 -, · ---- --- ·- ---· I I

!

+ I

"O I .(4 ~

I 1 I I ' I

30 + I ! i

T I

I I

?OT l l

T I

10 t

0

I I

!

0

FIGURE VII.3.12: Experiment 18 Extent of Reaction: 34.35 %

-- ----- --·-·- --· ----- - ........ -1

/~

I I

--- 1--1----·+--+-----1---~----+--·-t----l-----+----t---+---1-----1

1440

~!SOI uT·r· /-. d L- j L

2880

T ·-1 J. MF

4320 5760

0v1inutes)

the rate of aqueous oxidation of pyrrhotite minerals could

be determined from the rate of oxygen consumption. There

fore, using the third 'oxygraph', the attempt consisted in

determining the reproducibility of the rate of oxygen con

sumption. The effect of air and pure oxygen on the rate of

oxygen consumption will be analyzed in terms of ferrous and

ferric ions in solution, pH in the system, the oxidation

products and extent of reaction.

Finally, the extent of oxidation obtained from 'shake'

flasks experiments was compared with those of both 'oxy

graphs' by determining the mass balance of the reaction

products.

VII.4.1 RATE OF CONSUMPTION OF OXYGEN DURING BACTERIAL

AND CHEMICAL OXIDATION OF PYRRHOTITE MINERALS.

Table VII.4.1 summarizes the results obtained from

seven experiments in terms of final pH, bacteria growth, if

inoculated, total hours of reaction and the number of read

ings, ON/OFF time of electrolytic cell (ON time fixed to be

constant).

Two preliminary investigations were conducted using

this sulphide

initiated the

microorganisms.

concentrate. Firstly, Sawe (124) in

experiments at pH 2.5 and the presence

In a second study, Harris et al. (79)

1980

of

in

1983 suggested 4.5 pH as the most favorable acidic condi

tion. It was thought that the soluble products, obtained

from the first study, are environmentally unacceptable. The

132

second experiment was aimed at producing both iron and

sulphur reaction products as insoluble and safe residues.

Thus, this study aims to determine the rate of consumption

of oxygen at 4.5 pH and, then, to evaluate the rate of

aqueous oxidation of pyrrhotite minerals, if the reaction

stoichiometries assumed are correct.

Experiment 1 and 2, both using Sample 2, took place to

validate the basic pH assumption at 5.5 and 4.0, respective

ly.

Experiment 1 (Figure VII.2.1) had a final pH of 3.85

and experiment 2, 3.1 after 237 (9.875 days) and 241 (10

days) hours of reaction, respectively. However, the number

of readings obtained for experiment 1 and 2 were 1,914 and

7,580. The number of ON times of the electrolytic cell is

determined by dividing the total number of readings by two.

This chemical behavior indicated that when the initial pH in

the system was decreased from 5.5 to 4.0, the oxygen con

sumption rate increased. As both experiments took place in

the absence of microorganisms, and under these conditions,

oxidation of elemental sulphur is unlikely to occur; it is

implied that the rate of consumption of oxygen increased due

to an increment in the rate of oxidation of pyrrhotite.

133

TABLE VII.4.1

Exper. Initial Final Bacteria Hours of Number

No. pH pH Growth Reaction Readings

Chemical Oxidation: Sample 2

1 5.5 3.85 negative* 237 1,914

2 4.0 3.10 negative* 241 7,580

Bacterial Oxidation: Sample 2

3 4.0 2.95 indicium** 456 5,864

4 4.5 3.88 indicium** 24 1,008


5 4.5 192 1,794

Chemical and Bacterial Oxidation: Sample 0

6 4.5 143 398

Chemical Oxidation: Sample No. 2

7 4.5 3.76 indicium 216 1,322

* not detected in the microscope.

** very small amount detected in the microscope.

134

Table VII.4.2 shows the oxygen consumed, extent of

oxidation and rate of oxygen consumption for every experi

ment. These values were determined on the assumption that

Equations (VI.1.5) and (VI.1.12) took place and Faraday's

Law applied, as stated in Appendix G.

Since the rate of oxygen consumption by the aqueous

oxidation of pyrrhotite minerals is slow under the present

conditions and seems to have a linear relationship with

time, the mean of the oxygen consumption rate might only be

used for explanation purposes.

It can be seen that the total oxygen consumed during

experiment 1 and 2 was found to be 1.309 and 5.1842 grams.

These values give 16.39 and 64.90 per cent of extent of

oxidation and 0.00009 g/min and 0.00036 g/min of rate of

oxygen consumption for experiment 1 and 2, respectively.

Faster consumption of oxygen at 4.0 initial pH than

that at 5.5 is probably ascribable to the initial reaction

of pyrrhotite that takes place non-oxidatively according to

the potential-pH diagrams. And, as stated by Yazawa (25),

dissolution of pyrrhotite is approximately proportional to

the surface area of the mineral and the molarity of sulphu

ric acid forming an [FeS-2H+] activated complex. Consequent

ly, higher oxygen usage was due to oxidation of ferrous to

ferric ions. At the final pH values obtained (above 2.0),

precipitation and hydrolysis of ferric ions is a fast reac

tion. Thus, the regeneration of hydrogen ions and continuity

of the reaction cycle seems to occur. On the other hand, as

oxidation of ferrous ions is fast at pH values around 4.5, a

135

direct attack of ferric ions on pyrrhotite minerals may also

occur.

It has been found in voltammographic studies that at

4.6 pH and potentials of 0.2V (101) ferric hydroxide dis

solves by the reverse of the following reaction:

Fe2 + + 3H20 ------> Fe{OH) 3 + 3H+ + e- {VII.4.4)

"This suggests that the oxide which gives rise to this

peak is formed directly and is not produced from species"

(17,101). As this work was done in nitrogen atmosphere, it

is not known whether this reaction would occur in the

presence of dissolved oxygen.

At the conditions of pH and potentials stated above,

the results were analyzed in terms of the formation of

ferrous ions, elemental sulphur and sulphate ions. At higher

potentials (0.4V), the stable iron species is "iron {III)

oxide", elemental sulphur and sulphate ions. Although the

potential in the present study has not been measured, it is

believed to be in this range, from 0.2 to 0.4 V(SHE). Thus,

the reaction products stated by Hamilton (101) (ferric

hydroxide, elemental sulphur and sulphate ions, where ele

mental sulphur is more predominant than sulphate ions) are

expected to occur here.

It was also stated by Hamilton et al. (101) that the

voltammetric study followed by X-ray photoelectron spectros

copy (XPS) results indicated that oxidation of pyrrhotite

proceeded through progressive removal of iron from the

136

lattice represented as:

Fes1 _13 ------> Fe1_xs1 _13 + xFe2 + + 2xe

and

(VII.4.5)

Fes1 _13 + 3xH2o -----> Fe1_xs1 _13 + xFe(OH) 3 + 3xH+

+3xe- (VII.4.6)

where the value of x will increase with an increase in

potential, and the rest potential of pyrrhotite in air

saturated pH 4.6 solution was found to be 0.3 V (SHE).

In experiments 3 and 4, Sample 2 was subjected to

bacterial oxidation at 4.0 and 4.5 initial pH values, re

spectively. Appendix G also shows the calculation by which

the oxygen consumption and oxygen generation could be ob

tained from the reaction stoichiometries and Faraday's Law.

It can be seen that 7.9877 grams of oxygen would be, in

theory, the amount of oxygen consumed by thirty grams of

pyrrhotite mineral sample.

Initially, the leaching system (containing only the

pyrrhotite sample) was sterilized with CIG "sterigas 27"

(12% w/w ethylene oxide in dichloro-difluoromethane) for

about 4 hours. Thereafter, the sterigas was flushed out by

sterile air. Distilled water (pH 4) was sterilized by auto

claving.

Experiment 3 started as a chemical oxidation, then

after 24 hours (770 readings were obtained already), the

system was inoculated with 5 ml, 109 bacterial suspension

acclimatized to this sulphide ore by Khalid (124). The 9K

137

nutrients were added in solid form equivalent to 600 ml of

solution. It was noted that the rate of oxygen consumption

was faster compared with that of experiments 1 and 2. It was

also observed that more froth was formed on the surface and

that the slurry showed to be more a brownish colour than

that of experiments 1 and 2. It has not been possible to

draw a graph for this experiment.

138

OXYGEN CONSUMED,

TABLE VII. 4. 2

EXTENT OF OXIDATION


AND RATE

Experiment Oxygen Extent of Rate of Oxygen

Consumed,g. Reaction,% Consumption,g/hr.


1 1.309 16.39

2 5 .1842 64.90

Bacterial Oxidation: Sample 2

3* 5.469 68.47

4 0.6894 8.63


5 1.227 15.36

Chemical and Bacterial Oxidation: Sample O

6 0.2722 3.41


7 0.9042 11.32

* ON time 150 seconds

139

0.00009

0.00036

0.0020

0.00048

0.0001

0.000032

0.00007

OF

To determine the oxygen to be consumed theoretically,

it was assumed that the reaction stoichiometries, stated in

Equations (VI.1.5) and (VI.1.12), for stoichiometric and

monoclinic pyrrhotite were taking place. As 30 grams of pure

stoichiometric pyrrhotite and monoclinic pyrrhotite would

consume 8.1905 and 7.785 grams of oxygen, respectively, the

mean value, 7.9877 grams, is thought to be the more accu

rate.

As it was reported in Chapter II, Physical Properties

of Iron Sulphides, the most abundant natural pyrrhotites are

believed to consist of three kinds of pyrrhotites, mainly:

monoclinic pyrrhotite (4C), intermediate pyrrhotites (nC)

and troilite (2C). However, as nC pyrrhotites having a

composition range from Fe9s 10 to Fe11s 12 have been scarcely

reported in the literature of aqueous oxidation, troilite

and monoclinic pyrrhotite are considered the predominant

minerals.

Table VII.2.2 also shows the amount of oxygen consumed

and the extent of reaction of pyrrhotite for experiment 3

and 4. An extent of 68.47 per cent was obtained for the

bacterial oxidation of Sample 2 in 456 (19 days) hours for

experiment 3. However, it was noted that after 200 hours of

reaction, only a very small amount of oxygen was consumed.

It is not understood whether the second half of the reaction

(after 200 hours) was controlled by a diffusion process by

the formation of ferric hydroxide and elemental sulphur

and/or the bacteria became inactive by the absence of opti

mal growth conditions.

140

Microscopic examination of the leached material indi

cated that bacteria had not grown; in the best cases (exper

iment 3 and 4) only very few microorganisms were detected,

although according to Kuenen et al. (110) (The Prokaryotes,

Chapter 81) pH 4.0 is the optimal acid level required for

growth of T. Thiooxidans, it is probable that the bacteria

have lessened their activity due to the absence of other

growth requirements. T. ferrooxidans are able to derive

energy from ferrous iron oxidation as well as inorganic

sulphur compounds while T. thiooxidans, which have been

shown to possess the capacity to chemelithotrophically

oxidize ferrous ions and reduced inorganic sulphur com

pounds, have been reported more recently to be highly active

under the following conditions (107):

- A pH range of 0.5-6 (125) with an optimum pH for growth

of 2,

- optimum temperature of 15-25°c for T.ferrooxidans and 10-

250c for T. thiooxidans (125),

- a source of nitrogen, phosphate and trace amounts of

calcium, magnesium and potassium, and

- oxygen and carbon dioxide from air used for organic

synthesis.

It is believed that some of these requirements may have

been suppressed in the 'oxygragh' (e.g. nitrogen, carbon

dioxide) since the equipment was sealed from the atmosphere

and only pure oxygen was produced.

In experiment 5 and 6, chemical and bacterial oxidation

of Sample o, named as the "less reactive" pyrrhotite by

141

Harris et al. (79) compared with Sample 2, was studied.

Experiment 5 and 6, Figures VII.2.2 and VII.2.3, with

an initial pH of 4.5 gave an extent of oxidation of 15.36

and 3.41 per cent in 192 (8 days) and 143 (6 days) hours of

reaction, respectively. From which the average rates of

oxygen consumption were 0.0001 and 0.000032 g/min, respec

tively. Experiment 7, Figure VII.2.3, took place under the

same initial 4.5 pH using Sample 2. In 216 hours, it pro

duced an extent of reaction of 11.32 per cent with an aver

age of oxygen consumption of 0.00007 g/min. These experimen

tal results are not completely in agreement. The discrepan

cies in the rate of oxygen consumption during experiment 5

(higher than experiment 7) and 6 (lower than experiment 7),

shown in Figure VII.2.2 (experiment 5) indicates that molec

ular oxygen has constantly been consumed during 8 days and

no experimental errors can be found. Nevertheless, although

Figure VII.2.3, derived from experiment 6 (Sample 0), shows

a lower rate of oxygen consumption than experiment 5, it

might be attributable to an experimental error.

The rate of oxygen consumption of Sample 2 with an

initial pH of 4.0 ([H+] = 10 x 10-5 mol/L) (experiment 2) is

about 3.6 times faster than that of Sample Oat 4.5 initial

pH ([H+] = 3.16 x 10-5 mol/L) (experiment 5). On the other

hand, the rate of consumption of oxygen of Sample 2 (experi

ment 1) with an initial pH of 5.5 ([H+] = 3.16 x 10-6

mol/L) and that of Sample O (experiment 5) at inital pH of

4.5 are approximately the same. Thus, it is deduced that the

ratio of concentration of hydrogen ions of 10 between pH 4.5

142

and 5.5 did not increase the rate of consumption of oxygen

of Sample 0. However, a ratio of concentration of hydrogen

ions of 3.16 between 4.0 and 4.5 pH did increase the rate of

consumption of oxygen of Sample 2 about 3.6 times faster

than that of Sample 0.

Aqueous oxidation of Sample 2, named "reactive pyrrho

tite" (79), wich was slightly faster than Sample 0, is

probably ascribable to the initial reaction which takes

place non-oxidatively. This is likely to be the behaviour of

stoichiometric pyrrhotite according to its potential-pH

diagram. This could be confirmed by the statement made about

this kind of pyrrhotite: "the pH will increase to a value

where the oxidation of ferrous ions becomes kinetically

favourable and the overall rate of oxidation is determined

by the rate of chemical oxidation of ferrous ions" (79).

On the other hand, with the aqueous oxidation of Sample

0, named "less reactive pyrrhotite" (79), the initial reac

tion seems to predominate with ferric ions rather with

hydrogen ions. This is likely to be the behaviour of mono

clinic (or/and hexagonal pyrrhotite) due to its stability

in the potential-pH diagram (it is stable at higher poten

tials than sotichiometric pyrrhotite). This may also be

confirmed by the statement that "the pH may decrease until a

significant ferric ion concentration is established and the

overall oxidation will be determined by the biologically

assisted oxidation of ferrous ions" (79). Thus, it appears

that Sample O was predominantly monoclinic pyrrhotite and/or

hexagonal pyrrhotite and Sample 2, stoichiometric pyrrho-

143

tite.

Thus, the rate of consumption of oxygen during chemical

oxidation of Sample O is slightly higher than that of

Sample 2. A clear cut conclusion if there is any difference

in the rate of oxygen consumption during bacterial oxida

tion of both samples can not be drawn from these experi

ments. Moreover, as the exact origin of Sample O is uncer

tain and some variability of the reactiveness of pyrrhotite

fraction was also found in the mill processing practice

(79), a full reproducibility of both samples might also not

be expected.

VII.5 JUSTIFICATION OF THE METHOD FOR DETERMINING THE

AQUEOUS OXIDATION RATE OF PYRRHOTITE MINERALS

It is widely known in the literature that there are

three areas of investigation of a reaction whether it is

homogeneous or heterogeneous: the stoichiometry, the kinet

ics and the mechanism. In general, the stoichiometry is

studied first, and when this is far enough along the kinet

ics is then investigated. With empirical rate expressions

available, the mechanism is then looked into. In this sec

tion, an attempt will be made to justify the method by which

it is shown that the rate of consumption of oxygen obtained

from the 'oxygraph' may be used to determine the rate of

aqueous oxidation of pyrrhotite minerals. The results ob

tained from these experiments are compared with those of

'shake' flasks in terms of elemental sulphur formed and

144

extent of reaction. Table VII.5.3 summarizes the results

obtained in these series of experiments which took place in

the third •oxygraph'.

145

TABLE VII.5.3

Experi Initial Final Hours Total Fe2 + Fe3 + Air/

ment pH pH Reaction Reading ppm. ppm. Oxygen

-----------------------------------------------------14 4.5 4.10 120 878 air

15 4.5 3.6 215 2,520 air

16 3.0 4.6 24 air

4.8 120 1,522

17 4.0 3.90 120 1,958 air

18 4.0 4.073 72 1,818 18.5 nil air

3.014 144 9 3.5

19 4.0 4.232 72 11.6 0.1 air

2.988 120 3.72 0.355

20 4.0 3.721 72 2,294 22 0.3 oxygen

146

The main reaction products obtained from the aqueous

oxidation of pyrrhotite minerals are likely to be amorphous

ferric hydroxide, FeOOH and elemental sulphur. However,

oxidation of some elemental sulphur to sulphate seems to

occur as ferric hydroxide accumulation takes place or/and as

the oxidative leaching process advances. This reaction,

oxidation of elemental sulphur, will be discussed later in

detail.

Case et al. (126) examined the products in the solution

phase as well as precipitates and starting material residues

of aerobic "hydrolysis" (pH variation from 2.5 to 5.8) of

ferrous sulphide by potentiometry, X-ray diffraction, X-ray

fluorescence, scanning electron microscopy, infrared spec

trometry and reported that lepidocrosite, (reported only as

FeOOH, it is usually known as (beta)-FeOOH) and sulphate are

the main reaction products. However, voltammographic

studies of surface oxidation of pyrrhotite minerals at 4.7

pH (17,101) reported Fe(OH) 3 (sometimes called only Fe(III)

or "iron

products.

extent of

oxide") and elemental sulphur

Harris et al. (79) reported that

reaction of bacterial oxidation

as the major

monitoring the

of pyrrhotite

minerals (the same concentrate as this study, at 4.5 stated

pH), by examining the reaction products by scanning electron

microscopy with back scatter detection and EDAX analysis and

X-ray diffraction, proved generally unsatisfactory. Case et

al. (126) also presented a balance of materials with plus or

minus 20 per cent precision after using all the experimen

tal methods indicated above.

147

Qualitative and quantitative analysis of the solid

reaction products from aqueous oxidation sulphide minerals

at pH values above 2.0 is complex to determine. It is com

plicated firstly, because ferric iron precipitates can

occur in amorphous and crystalline modifications and second

ly, ferric iron and sulphate ions may form many crystalline

or/and amorphous ferric basic sulphate precipitates under

similar conditions. Thirdly, these precipitates have not

been characterized thoroughly. According to Music et al.

(47) the instrumental methods indicated-above by Case et al.

(126) and also IR spectrophotometry, and ultracentrifugation

have their limitations. However, the use of 57Fe Mossbauer

spectroscopy does permit us, according to Music et al. (47),

to study the steps involved in the precipitation and their

subsequent transformations.

Since the hydrolysis and precipitation of ferric ions

generates hydrogen ions, some elemental sulphur may also

oxidize to sulphate ions reducing the pH during the aqueous

oxidation process. Generation of sulphuric acid may not be

due to the oxidation of pyrite which was present in the

sample (79) since pyrite would be inert in this system.

Nevertheless, voltammographic studies of the oxidation

products of pyrite at 4.7 pH showed elemental sulphur with

very little sulphate (101).

The shift of the final pH during the oxidation process

of pyrrhotite, whether to higher or lower value than the

initial pH depends on the initial value of the hydrogen

ions. It can be seen in Table VII.3.1 that from an initial

148

pH of 4.5, it decreases to a value slightly above 3.0; but

from a initial pH of 3.0, the final pH increased to a values

of 4.6-4.8. However, with an initial pH of 4.0, it is likely

to remain more constant during the process. The pH changes

during the process are probably due to the consumption of

hydrogen ions, e.g. small amount of hydrogen sulphide,

oxidation of ferrous ions (consumption) and to the produc

tion of hydrogen ions from hydrolysis and precipitation of

ferric ions. The rate of consumption and production of

hydrogen ions are likely to be the same at initial pH of

4.0.

Stability of 4.0 pH might be indirectly confirmed from

a mineralogical and electrochemical stability study of nick

el-iron sulphides (pentlandite and violarite) by Thornber

(127). It was stated that the acid leaching behavior of

violarite would be very similar to that of hexagonal pyrrho

tite (Fe9s10 ). Thus a potential-pH diagram for violarite

compiled from intermittent galvanostatic polarization (IGP)

and cyclic voltammetry data indicate that below pH 4.0,

anodic dissolution favors sulphur formation, and above pH

4.0 sulphate is formed. Thus a pH value slightly less than

4.0 will form elemental sulphur predominantly.

The rate of consumption of oxygen hence and, conse

quently, the rate of oxidation of pyrrhotite, is higher at

the beginning of the process than at the end. It has been

reported by Herbst (102) that small particles leach at a

high rate giving an increase in the initial rate of the

leaching curves, whereas the large particles leach at a slow

149

rate giving rise to the shape of the curves after extended

leaching. The initial high rate of oxidation of pyrrhotite

during these experiments may be due to the particle size

distribution of the feed used. However, the initial high

rate may also be due to the uncoated surface area of pyrrho

tite at the beginning of the process. As has been suggested

by Hamilton et al. (101), the rate of pyrrhotite oxidation

at 4.6 pH may be inhibited as the active surface becomes

covered with ferric hydroxide.

The kinetics of oxygen-acid leaching of chalcopyrite is

controlled by surface reaction and a plot of the integrated

rate expression for bach leaching "1-(1-a) 1 / 3 ., versus time

for nearly spherical monosize particles should yield

astraight line (102). If the chalcopyrite feed has a wide

distribution of particle sizes, the overall reaction behav

ior deviates strongly from linearity, demonstrating the

overall size distribution kinetics do not follow the "1-(1-

a)1/3•• law; even though the kinetics of the individual

particles are controlled by surface-reaction-rate (102).

Thus similar behavior may be occurring during the aqueous

oxidation of pyrrhotite since about 25 per cent of Sample 2

was far below 38 um (micrometer) particle size.

The actual pH in the leaching system, during the proc

ess, depends on both the consumption of hydrogen by the

dissolution reactions of pyrrhotite and the generation of

hydrogen ions by the hydrolysis and precipitation of ferric

ions. At the beginning of the process, consumption of hydro

gen ions appears to be predominant until the rate of hydrol-

150

ysis and precipitation of ferric ions increases. Yazawa et

al. (26) showed that the non-oxidative dissolution rate of

pyrrhotite is second order in the concentration (pulp densi

ty} of pyrrhotite up to about 50 per cent of the dissolu

tion, e.g.

- d[FeS]

dt

= [Fes] 2 mole

(l.min}

(VII.5.7)

where [FeS] means the concentration of pyrrhotite and K2 is

a constant. This rate equation was confirmed in the ratio of

acid normality to pyrrhotite molarity from 1.78 to 4.95 by

Jibiki (23). Thus, if regeneration of hydrogen ions does not

occur, acid depletion during the leaching may affect the

reaction rate in the latter stages. It was also reported

that the presence of oxidants (in the concentrations of acid

stated) inhibited the dissolution process (23,26).

Then, at a constant concentration of dissolved oxygen,

both reactions (consumption and production) may occur at

about the same rate. Firstly, if the rate constant for

oxidation

greater

of ferrous ions by oxygen at 25°c and pH values

than 4.5 is K=8.0 x 1013 liter2 .mole-2 .atm-1 .min- 1

and at pH below 3.5 K'=l.O x 10-7 atm- 1 .min- 1 , according to

Singer (14), then at 4.0 pH the rate constant will be ap

proximately above the mean of both values (e.g. slightly

above than 4.5 x 106 liter2 .mole-2 .atm-1 .min- 1 ). This value

shows that the oxidation of ferrous ions at 4.0 pH is a fast

reaction. Ferric hydroxide and any other oxide hydroxide of

151

Fe(III) precipitated in the leaching process is directly

related to their solubility product ([Fe3 +J[OH-J 3 = Kso).

According to Feitknecht et al. (48) fresh precipitates had

-log Kso values ranging from 38.0 to 42.7. Thus, the concen

tration of ferric ions is limited to the solubility product

of ferric hydroxide and to its reaction with pyrrhotite

minerals.

As previously discussed throughout the literature

review, hydrogen ions initially react with the sulphide

mineral to produce ferrous ions and some sulphur compounds,

which are predominantly elemental sulphur at 4.0 pH. The

oxidation of ferrous ions is fast at this pH giving ferric

ions, hydroxyl complexing species or may be even the sul

phate complexed species in solution. As the concentration of

ferric ions was found above the solubility of ferric hydrox-

ide, oxidation of pyrrhotite minerals by ferric ions may

also occur and this reaction is considered to be fast.

Otherwise, ferrous ions are directly oxidized to ferric

hydroxide producing hydrogen ions which could react with the

mineral. If the pH of the leaching solution is maintained

constant, the rate of consumption and production of hydrogen

ions must be the same.

Hydrogen Consuming Reactions:

FeS + 2H+ ------> Fe2 + + H2S

FeS + 2H+ + 0.502 -----> Fe2 + + s0

and/or

152

Equation

(V.3.1)

+ H20 (V.3.5)

Fe7S9 + 16H+ + 2e- -----> 7Fe2 + + 8HzS (V.3.2)

Hydrogen Producing Reactions:

Fe2 + + 2.5H20 + 0.2502 ------> Fe(OH) 3 + 2H+ (V.3.12)

If oxidation of elemental sulphur to sulphuric acid occurs,

the following equation may be added:

(V.3.13)

VII.5.1 REPRODUCIBILITY OF THE REACTION RATE

The reproducibility of the reaction rate measurement

was determined in experiments 14 and 15A, the results of

which are shown in Figures VII.5.13 and VII.5.14. It can be

seen that the reproducibility of the measured rate of reac

tion shows a satisfactory correspondence. This agreement

could be better appreciated by consulting the extent of

oxygen consumption curve, Figure VII.5.15.

It is noted that in order to determine the reproduci

bility of reaction experiments 14 and 15 were compared. As

only 878 readings were obtained from experiment 14 (Table

VII.3.1), the same number of readings were taken from the

results of experiment 15 and called experiment 15A. Table

VII.3.1 also shows the computerized total ON time values

obtained for all experiments. For the same number of read

ings, Experiments 14 and 15A indicate practically the same

153

C

--!:.::. ON 1 01

...:t N 0 0 0

. z D ~

r-0. ~ J (J)

z 0 u z w [j

>x 0 u. D w le:( Q:'.

.8 i

.6


Total Ol'J Time: 40,029 sec.

Total Reaction Time: 2,792 min.


4 1 .,- •• : •

. T . :-·;.-'_,·. · :·/ \' .. >:·:· :o._, .. . . - , :--> .. ·.

1 . ... .. "q• ·.····.,

0

'· ··· .. · .... -··,. . . . ; .. ' ....

720 1440 1

2160

ABSOLUTE TIME · (Minutes) \

. ... ··· ··: ..

2880 2 DAYS

C "§ '-,

u 0

0)

.:J N Cl 0

0

~

z D H

~ 0.. ~ J u1 z D u z w [j

> X D LL D w ~ I

1 er

1

.8

I

.6 t t I

. 4 t 1

' ....

.2

FIGURE VII.5.14: Experiment 15 A

Total ON Time: 40,026 sec.

_ Total Reaction Time: 2,436 min.


.... . _._..-··: :,,.-._._..... . ·• ·.:, .- ........ _· ... -.. .-._ .. ·· ..

...... -;··.-'·:·. ... . : :: .. , _:.'/ ,. .. .. , ... · ...... ··.' - .. ·. . . .- . .. .. ·.. . . .... . ... -. -:.,_.:.. . ..... ~ ....

·:·

l

I 0 ~--+-----f-----+----------~

0 720 1440 2160 1 2 DAYS

ABSOLUTE Tit~E (Minutes)

~

UJ j--

H l-... I

0 I a:: CI >Q ..

LL 0 --,-

/ ,c_

0 r-1

f<(

0 H ' ,, /'-.

0

20-.------- -------·---- ------------· ------, ~~1

i I

15-:

t -r I

10

5±

0 0

FIGURE VII. 5. 15: Extent of Reactions

Experiments 111 & 15 A: 19.45 % I

I l

_,.c________ l---------t----1-------1-------+-------1-----J

720 1440 1

2160


2880 2 DAYS

total ON time (differing by only three seconds) verifying

the high reproducibility of the 'oxygraph' results.

The 878 readings obtained from experiment 14 took in

2,792 minutes (46.5 hrs) residence time compared with the

same number of readings from experiment 15A which took in

2,436 minutes (40.6 hrs), representing about 87.25 per cent

of reproducibility. This variation of results, under the

same apparent conditions, is not understood. The variations

of the ambient temperature were slightly different for both

experiments (the peaks observed in Figures VII.5.13 and

VII.5.14) and the atmospheric pressure at the beginning of

each experiment may have been different and responsible for

this disagreement. The extent of reaction obtained for both

Experiments was 19.45 per cent.

VII.5.2 FERROUS AND FERRIC IONS IN SOLUTION

Ferrous and ferric ion content were determined by an

UV/Visible spectrophotometer (Unicam 6000), using o-

phenanthroline and thiocyanate as indicators, respectively.

Figures VII.5.16 and VII.5.17 show the standardization plots

prepared for the analysis of both states of iron.

In all the experiments, the concentrations of ferrous

and ferric ions were less than 25 ppm and 5 ppm, respective

ly. It is shown in Table VII.5.3 that experiments 18 and 19

indicate two values for ferrous and ferric ions and pH. The

first value was taken after 72 hours of reaction and the

second value at the end of the experiment (about 120 hours).

154

QI

u i:: <1l

,.0 I-< 0 1/)

,.0

<

FIGURE VII.5.16: Ferrous Ions Absorbance

1.5

1.0

0.5

.'

5

A= 0.18 [Fe 2+]

Procedure:

1. up to 1 ml. 0.36N H2so 4

depending on pH, 2. 5 ml. O.lM KH phthalate

buffering to 3.9 pH, 3. 4 ml. 0.2% W/W H2o

o-phenanthroline, 4. total volume 50 ml.

10

Ferrous Ions ConceQt~ation, ppm.

(1)

u C: <1l .0

1-1 0 tll

.0 <i:

FIGURE Vli.5.17: Ferric Ions Absorbance

1.5

- 1.0

o.s

5

3+ A= 0.142[Fe ]

Procedure: 1. 2 ml. SN HCl, 2. 5 m 1. 2 M KC N S

3. total volume 50 ml.

10

Concentration of Ferric Ions, ppm.

These results show that concentrations of both states of

iron in solution were higher at 72 hours than at the end of

the experiment. The decrease in the concentrations of both

iron species results from the decrease in the rate of aque

ous oxidation of pyrrhotite at constant concentration of

dissolved oxygen.

The initial pH value is slightly lower than that taken

at 72 hours and slightly higher than the final pH. It is

evident that there was more consumption of hydrogen ions,

molecular oxygen and ferric ions at the beginning of the

process than after the 72 hours of reaction. Concentration

of ferric ions, shown in Table VII.5.3, at 72 hours of

reaction was lower than that at 120 hours of reaction. Lower

concentrations of ferric ions at early stages of the process

than those at later stages may indicate the catalytic effect

of ferric ions. The aqueous oxidation of pyrrhotite minerals

was unlikely to be due to only hydrogen ions and dissolved

oxygen, but also to ferric ions. For example, it is thought

that stoichiometric pyrrhotite may react predominantly non

oxidatively with hydrogen ions (according to the potential

pH diagram) and monoclinic pyrrhotite with molecular oxygen

and ferric ions. However, it would be incorrect to discard

the possibility of cathodic dissolution of monoclinic pyr

rhotite according to the following reaction stated in Equa

tion V.3.2,

+ ~+ Fe7s 8 + 16H + 2e- -----> 7Fe~ + 8H2S (VII.5.8)

155

The extent of reaction due to each, ferric ions, hydrogen

ions or dissolved oxygen, is difficult to establish.

The gradual decline on the rate of aqueous oxidation

of pyrrhotite, observed after about 24 hours of reaction,

could be due to the fact that ferric hydroxide coated the

mineral decreasing the active surface area in contact with

hydrogen ions, ferric ions and/or dissolved oxygen in the

absence of bacteria, thus increasing the concentration of

ferric ions slightly.

Concentrations of both ferrous and ferric ions were

higher when the aqueous oxidation of pyrrhotite minerals

were conducted in pure oxygen. Concentration of ferrous ions

in the presence of pure oxygen was 22 ppm compared with that

12 ppm in the presence of air. Although the concentration of

ferric ions in pure oxygen was only 0.3 ppm at 72 hours of

reaction, it is higher than that when the reaction took

place with air. It is not understood how pure oxygen would

affect the kinetics of the reaction if dissolved oxygen is

assumed to be at its saturation level in both cases. Never

theless, aqueous oxidation of copper (13) and iron sulphide

minerals (77,78,127) at ambient and elevated temperatures is

faster with pure oxygen than with air.

Although the concentration of dissolved oxygen was not

measured during the leaching process, a recirculation rate

of air and/or oxygen was maintained at about 9 1/min. The

air and/or oxygen was pumped beneath the "Vibramix"

tion plate where it was dispersed as very small

Mineguishi et al. (55) reported the oxidation of

156

agita

bubbles.

ferrous

ions at 4.7 pH from rising bubbles of oxygen and nitrogen

without agitation. The concentration of dissolved oxygen

reached saturation immediately after the start of the oxida

tion. Thus, it is assumed that the concentration of dis

solved oxygen was at saturation level in all the experi

ments. Although Mineguishi et al. (55) did not determine the

ferric(assumed to be very low)/ferrous ratio, they reported

that the oxidation rate was about three times higher in a

solution containing ferric hydroxide precipitate than in a

solution in which no precipitate was initially present.

VII.5.3 CHEMICAL OXIDATION OF PYRRHOTITE WITH PURE

OXYGEN AND AIR

Table VII.5.4 summarizes the total ON time, total

reaction time, oxygen, water consumption (according to

stoichiometric reaction stated in Equation VI.1.5) and the

extent of reaction for the aqueous oxidation of pyrrhotite

obtained for these series of experiments.

It is noted that the program written in Basic Language

(IBM PC compatible) calculates directly the amount of oxygen

consumed in grams and the extent of reaction at every ON

time value. Both programs, named "CLOCK" for the Microbee

computer and "LEACH.BAS" for the IBM PC compatible are

listed in Appendix F. The Microbee computer was interfaced

with the electrolytic cell through the 555 timer chip to

retrieve the data.

The highest extent of reaction obtained when the pyr-

157

rhotite mineral was oxidized with air was 55.69 per cent in

about 205 hours (12,285 min. ,8.54 days) residence time. This

experiment (experiment 15, Figures VII.3.3 and VII.3.9) took

place at 4.5 initial pH and gave a rate of oxygen consump

tion of 0.0004 g. of oxygen per minute. Experiment 16

(Figures VII.3.4 and VII.3.10) at 3.0 initial pH gave 0.0006

g. of oxygen per minute and the extent of reaction obtained

was 31.04 per cent in 70 hours (4,213 min. ,3 days) residence

time.

Apparently, at initial pH of 3.0, the overall rate of

oxygen consumption is faster than that at 4.5. However, the

pH value of experiment 16 increased from 3.0 to 4.6 after 24

hours of reaction, and after 120 hours, it was still 4.8. It

is not understood why the pH value increased from 3.0 to 4.8

and was maintained by itself at this pH value. It is under

stood that the aqueous oxidation would be faster at early

stages of the process due to higher concentration of hydro

gen ions and free surface area of pyrrhotite minerals

(predominance of non-oxidative dissolution). However, when

the pH value reaches 4.8, the aqueous oxidation process

seems to decrease drastically (126).

Evidently, hydrogen ions are not being regenerated if

the pH value is self maintained at 4.8 (no formation of

ferric ions). Case et al. (126) with 5.05 initial pH and

"aerobic" conditions (without air sparging) obtained about

2 per cent of iron sulphide "hydrolysis" in 6 days. The

final pH was 5.8. Thus, concentration of hydrogen ions of

4.0 pH (sparging with air) seems to maintain the aqueous

158

oxidation rate of pyrrhotite minerals at optimum conditions.

As expected, the rate of oxygen consumption is higher

when the aqueous oxidation process takes place with pure

oxygen rather than with air. As shown in Table VII.3.2, the

amount of oxygen consumed in experiment 20, which took place

at initial pH of 4.0, was 4.57 grams of oxygen giving an

extent of oxidation of 55.85 per cent in about 132 hours

reaction (7,920 min., 5.5 days). Correspondingly, experiment

17 which took place with air and at 4.0 initial pH gave

42.98 percent of extent of reaction in 144 hours (8,628

min., 6 days) residence time. Although the rate of uptake of

oxygen showed some deviation from linearity in experiment

20, the total amount of oxygen consumed was divided by the

total reaction time giving 0.0006 and 0.0004 grams of oxy-

gen/minute for experiment 20 and 17, respectively. Thus

the rate of uptake of oxygen for experiment 20 was a bout 33

per cent faster than that for experiment 17 over 144 hours

residence time.

159

TABLE VII.5.4

EXTENT OF REACTION IN TERMS OF OXYGEN CONSUMED

Experi- Total Total Oxygen Water Extent

ment ON Time Reaction Consumed Consumed Reaction

Sec. Time,Min grams grams %

* 14 40,029 2,792 1.5929 0.5979 19.45

15 114,617 12,286 4.5609 1.7104 55.69

* 15A 40,026 2,436 1.5979 0.5973 19.45

16 63,897 4,213 2.5426 0.9535 31.04

17 88,474 8,628 3.5206 1.3202 42.98

18 70,698 4,523 2.8133 1.0550 34.35

20 114,954 7,920 4.5743 1.7051 55.85

20A ** 43,423 1,138 1.7279 0.6480 21.09

* For comparison of reproducibility of reaction with air(*)

and pure oxygen(**).

160

Similar conclusions were not made on the results of

experiment 18 (also reacted with an initial pH of 4.0),

since the curve shows unusual behavior. It is likely that

the contact at the mercury interface (manometer) became

dirty and unreliable.

The first 878 readings of experiment 20, conducted with

pure oxygen, were named experiment 20A. The rate of consump

tion of oxygen and extent of reaction calculated from these

results,

plots are

were compared with experiment 14 and 15A. These

shown in Figures VII.3.18 VII.3.19. It can be

seen that the reaction rate of experiment 20A with pure

oxygen was about 55 per cent faster than that of experiments

14 and 15A in the first 48 hours of reaction. Thus the rate

of consumption of oxygen is faster throughout the experi-

ment with pure oxygen than with air.

As the stoichiometric reaction, shown in Equation

VI.1.5, implies the consumption of water, its amount was

calculated and listed in Table VII.3.2. It is noted that

this value is very small.

VII.5.4 EFFECT OF SODIUM CHLORIDE AND COPPER IONS

Electrochemical reduction studies of chalcopyrite (13)

indicate that additions of cupric chloride or ferric ions

enhance significantly the reduction current. It is also

known that sodium chloride in solution improves the conduc

tivity of an electrolyte. On the third day, 5 grams of

sodium chloride and one gram of copper sulphate (0.004 M

161

~ 2 5 ~---------------------------·----· ------ ------- ---·-·-·-··--·-····-·--·--,

LJJ 1-H I-· 0 I CI CI >Q_

LL

20

± 15

0 10

z 0 1---j

f<(

C::J H >( CJ

T

f 5t

--0

0

Experiment 20A Experiments

14 15A

FIGURE VII. 5.18: Experiments 14, 15A and 20A. Extent of Reactions:

Experiments 14 & 15 A: 19.45 %

Experiment 20A: 21.09 ¾

+-------+-·----!1-----t-----720 1440 2160

1

ABSOLlJTE TIME (Minutes)

2880 2 DAYS

C .E

" d'" Ol

-::r N 0 0 0

i r--

f . z D H L I

0.. ~ J ([)

z

.8 t i

t .6 i

0 u

l .4

z w ~ > .2 X D lL

,. D .-~ .. -08 w f-

0 <{ n::

.. ·. .. . .-. · ..... : .. . . ........... '· .,·

·····

·-'.

Total ON Time: 43,423.29 sec. Total Reaction Time: 1,137.5 mi.n. Extent of Reaction: 21.09 %

....... '

------------------··

FIGURE VII.S.19: Experiment ;~OA

.. ' .. ·. .....

•.· .. : :

.. •: .. .... ..

l

-+--+---·+---t---+-----l'----t---+---t---+----1

180 360 540 720 900 1080 1260 1440

ABSOLLJTE TIME (Minutes)

copper ions) were added to Experiments 16 (Figure VII.3.4 and

VII.3.10) and 18 (Figure VII.3.6 and VII.3.12), respective

ly. The aim was to improve the rate of uptake of oxygen if

they could act as catalysts. However, an improvement on the

rate of consumption of oxygen was not observed in either

case. It is thought that cupric ions reacted according to

the following equations:

cu2 + + H2S----->CuS(c) + 2H+ K=2 x 10-15

and

FeS(c) + 2H+ -----> Fe2 + + H2S K=3 x 104

Where K is the equilibrium constant.

(VII.5.9)

(VII.5.10)

From which an overall reaction for monoclinic pyrrhotite

would be:

Fe7S0 + 7Cu2 + ----> 7Fe2 + + 7CuS + S (VII.5.11)

Electrochemical studies of the nature of the interac

tion of copper ions with galena and pyrrhotite in the pH

range of 3-5 was conduced by Nicol (128). It was found that

a surface layer with the electrochemical characteristic of

CuS is formed on these minerals, minimizing the oxidation of

the mineral. Contrarily to the aim of aqueous oxidation,

activation by copper ions at a pH value of 5 results in a

significant increase in recovery by flotation. It was also

reported that in the case of pyrite, a CuS layer is not

formed by the interaction with copper ions, alternatively

catalytic oxidation of the surface to elemental sulphur is

suggested.

162

VII.5.5 MASS BALANCE: SHAKE FLASKS EXPERIMENTS

A fourth series of experiment was performed in a shake

flask incubator in order to determine the extent of oxida

tion of pyrrhotite by measuring the quantity of elemental

sulphur formed and quantifying the unreacted mineral. The pH

values at which the oxidation took place and length of

reaction time were distributed as follows:

Initial pH: 1.5 2.5 3.5 4.0 4.5

Time, hrs.: 24

Number of Flasks: 1

24

1

72

3

120

5

120

5

This distribution allowed for the withdrawal of one

flask after 24, 48, 72, 96 and 120 hours of reaction. The

residue, a mixture of ferric iron precipitate, unreacted

pyrrhotite and elemental sulphur was filtrated and dried in

an airtight desiccator.

A laboratory size distillation column was installed in

a fume extraction cover to extract elemental sulphur using

tetrachlorethylene. The oxidized mineral was placed in an

extraction thimble and about 200 mls. of the extractant was

boiled

tilled,

thimble.

slowly overnight. The organic extractant, once dis

entirely covered the volume of the mineral in the

After the completion of the extraction process,

tetrachlorethylene was evaporated to dryness leaving the

elemental sulphur which was weighed.

From the residue, ferric hydroxide was separated from

163

the unreacted pyrrhotite manually by elution as the specific

gravities are 4.84 for ferrous sulphide and 3.4-3.9 for

ferric hydroxide (35). The limitations of this separation

are realized, since small particles of unreacted material

may remain with the ferric hydroxide. Table VII.5.5 shows

the results of this experiment.

164

TABLE VII.5.5

EXTENT OF REACTION IN TERMS OF ELEMENTAL SULPHUR PRODUCED

Residue Ferric Elemental Weight Reaction

!nit. Weight Hydroxide Sulphur Lost Extent as

pH grams + unreact. grams grams FeS Fe7s 8

pyrrhotite % %

grams

First day:

1.5 30.516 25.818 3.351 1.347 30.69 28.23

2.5 34.655 29.245 lost

3.5 35.090 29.050 5.246 0.794 48.04 44.19

4.0 34.765 29.009 5.103 0.653 46.73 42.98

4.5 33.876 28.361

Second day:

3.5 34.25 28.000 4.831 1.426 44.24 40.69

4.0 29.696 24.643 lost

4.5 34.958 28 .157 5.488 1.313 50.25 46.23

Third day:

3.5 34.728 28.285 5.346 1.097 48.95 45.03

4.0 35.560 28.239 5.741 1.580 52.57 48.36

4.5 35.386 28.203 6.140 1.043 56.22 51.72

Fourth day:

4.0 34.539 27.444 5.623 1.472 51.49 47.36

4.5 33.540 27.898 4.705 0.937 43.08 39.63

Fifth day:

4.0 32.960 28.240 4.720 1.327 43.22 39.76

4.5 33.838 27.535 4.200 2.103 38.46 35.38

* Weight lost after extraction of elemental sulphur.

165

From the results listed in Table VII.S.S two important

observations can be made. Firstly, the mass of the residue,

(mixture of ferric hydroxide, unreacted pyrrhotite and

elemental sulphur), is higher than the mass of the original

pyrrhotite sample. Secondly, the extent of reaction could be

determined from the amount of elemental sulphur extracted

from the residue and from a 30 gram sample of pyrrhotite

mineral. A value of 10.9411 and/or 11.8849 grams of elemen

tal sulphur could be theoretically obtained from 100 percent

of oxidation of 30 grams of mineral as stoichiometric pyr

rhotite and monoclinic pyrrhotite, respectively.

Figure VII.7.20, from (a) to (e), shows the extent of

oxidation versus pH obtained for each day. Plot (f) shows

the extent of reaction versus time (days) as a function of

pH. It can be seen that the extent of oxidation is lower at

low pH values, e.g. 1.5 pH than that at high pH values, e.g.

4.5 pH until the third day, whether stoichiometric or mono

clinic pyrrhotite is assumed. After the third day, the

extent of reaction is higher at 4.0 pH than that at 4.5 pH.

It can be seen in plot (f) that at 4.5 initial pH the

largest quantity of elemental sulphur, giving 53.97 per cent

of extent of oxidation, {mean value of 56.22 [FeS] and 51.72

[Fe7s8 ]), lessens after the third day to give 41.36 and

36.92 per cent as the aqueous oxidation continues. Similarly

at 4.0 initial pH, it lessens from 50.47 to 49.26 and 41.49

per cent. It is assumed that elemental sulphur

oxidized after the third day and this oxidation is

faster at 4.5 initial pH than that at 4.0.

166

is being

slightly

--··-

---

-------

.:t"

N

~----::----------~

...JO

...;__~

3 0

~

t<"

c-l

-·--------

~-------------~

..J 0

~o

o

o '

\ ...:r

\C"\

0

I Ct

r.n >

C

0 ·- .w

co D

·-X

0 q

_

0 .w

C

Q

.) .w

X

w

0 N

r--

- ...... >

~ ~

9 ~ ~

No

U. ~ 3c:) .::iO

.L N

3 .l. )(3

i ~

t

u, I..-

:z l.U

[ 0:: w

0

. X

w

-U

1 -::c U

1 c(

_J

lL

w

-::c c::( I :.(/1

::t

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-~

---:-

-----------_

,...JO

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C

r 0

It is not clearly understood why the formation of

elemental sulphur is lower at 1.5 pH than

er, this behavior is consistent with the

at 4.5 pH. Howev

results of the

toxygraph' at 3.0 initial pH. The pH after 24 hours in

creased to 4.6 and after 5 days to 4.8 and the leaching

process was slow. It was reported by Jibiki (23) and Yazawa

(26) that the presence of oxidants inhibits the non-oxida

tive dissolution of pyrrhotite minerals.

It is also not clear cut why elemental sulphur starts

to oxidize just after the third day of oxidation consistent

ly. Accumulation of oxidized sulphur species in solution may

had been taken place. It is agreed in the literature

(19,104) that the (main) overall oxidation reaction for

monosulphides (where sulphur is present as s2 - in the crys

tal lattice) produces Me2+, (Me=Fe, Zn, Pb), elemental

sulphur and to a lesser extent sulphate ions. Pyrite, in

which sulphur is present as s22- entities, is an exception

as virtually all its sulphur is converted to sulphate, and

s0 is generally reported to be absent or constitute a minor

product only.

Elemental sulphur

temperatures below its

is hardly oxidized, if at

melting point (119°c)

all, at

(98,104).

Therefore, it can be excluded as an intermediate in the

sulphate formation. At pH values about 4.0, the predominant

forms of reduced sulphur in water are H2s with 98 percent

and HS- with 2 percent (130). Chen and Morris (97) made an

extensive study on the effect of pH on the oxidation of H2s.

It was concluded that below a pH of 6, and in the absence of

167

biological activity, hydrogen sulphide can slowly be oxi

dized to sulphur which combines with the remaining sulphide

to form polysulphides. However, oxidation of sulphide to

sulphur by sulphur oxidizing bacteria can eliminate the rate

determining step on the initial sulphur formation.

Electrochemical studies in order to understand the

surface oxidation and/or floatability of pyrrhotite miner

als have been made by several investigators (17,101,128-

130). From these studies, it is generally established that

oxidation of pyrrhotite at pH about 4.0 would yield ferric

hydroxide, elemental sulphur and sulphate ions. Hodgson et

al. (129) from voltammographic studies stated that "strong"

oxidation would promote two sequential processes:

First Reaction: "this is considered as the partial

oxidation of Fe(II) to form a hydroxide and S(anionic) to

the polysulphide":

Fes1 _13 + H2 = [Fe(OH)aJ<s> 2 + s 22-(aq) + H+ + e

(VII.5.12)

Second Reaction: "Further oxidation of the solvatable

polysulphide formed and complete hydroxylation of the

Fe(OH>a site".

[Fe(OH)a](S2 ) +H2o= Fe(OH) 3 + S + H+ + e- (VII.5.13)

"Sulphur will continue to oxidize to so42-"

Adam et al. (130) indicated that in an oxygenated

solution at neutral pH the surface oxidation of pyrrhotite

to hydroxide or oxide and sulphate species of iron is

formed, accelerating electrochemical reactions. The electro

chemical reaction proposed are:

168

FeS = Fle

2+ + [_~ ::i::ociation)

Fe20(0HJ 3

4

or FeOOH (surface layer l formation)

Fe(OH)S04

(VII.5.14)

(VII.5.15)

"Ferrous ions released from pyrrhotite upon dissociation

react with hydroxyl ion generated by the reduction of oxygen

on the cathode surface of pyrrhotite and this results in a

stable iron hydroxide species" (130).

Thus from these electrochemical studies, it is general

ly agreed that the overall oxidation reaction products of

pyrrhotite minerals are ferric hydroxide and elemental

sulphur. Seemingly, some sulphide ion was oxidized to sul

phate at the surface of the iron hydroxide coating and

formed an iron basic sulphate.

The oxidation of elemental sulphur after the third day

may be explained by the effect of ferric hydroxide on sul

phide and ferrous ions. It was reported by Minegishi (38)

that the initial presence of 0.1 M of ferric hydroxide

enhanced the rate of oxidation of ferrous ions. Similarly,

Adam et al. (130) concluded that sulphide ion was oxidized

to sulphate at the surface of the ferric hydroxide. Thus, it

is proposed that ferric hydroxide also enhanced the oxida

tion of elemental sulphur after a certain value of ferric

hydroxide concentration. This process seems to proceed via

reduction of elemental sulphur to sulphide and then decompo

sition of H2s2o2 and H2so2 formed, according to equations

169

(VII.5.16 and VII.5.17 formulated by Lotens et al. ( 104) .

It is likely that the reduction of ferric hydroxide, report

ed from the electrochemical studies (17,101), takes place

due to this reaction. However, it is unlikely that the

consumption of oxygen increased from the third day of

oxidation, although some scatter is noted in the 'oxygraph"

results of experiment 15.

Bailey et al. (16) concluded from "oxygen-18 water

tracer tests" that "sulphate formed from pyritic sulphur

contains oxygen taken from water rather than from the high

pressure gas phase". It is assumed the same electrochemical

mechanism (oxidation of sulphur to sulphate) occurred during

the oxidation of pyrrhotite minerals. As no increase in the

consumption of oxygen is noted after the third day, the

sulphate formed from sulphur may contain oxygen from the

water rather than from the oxygen partial pressure.

In addition to elemental sulphur and sulphate, there

are indications that other sulphur species are present in

leach liquors. Lotens et al. referred to Brual et al., who

showed that not only s2o32-, but also s4o6

2 - exist during

the dissolution of pyrrhotite.

Independently, it is reported by Lotens et al. (104)

that oxidation of sulphide ion, s2 - at the crystal/leach

liquor interface to either s+ or s2+ depends on the oxida

tion agent used. The oxy-hydrolysis of those species to

H2s 2o2 and H2so2 , respectively, is followed by a "dispropor

tionation'' giving 75 and 50 per cent of elemental sulphur

yield, respectively.

170

2H2S202 ------> 3S + H2so3 + H2o

and

2H2S02 ------> S + H2so3 + H2o

(VII.5.16)

(VII.5.17)

"Higher sulphur yields can be achieved at pH level below a

certain value, which is specific for the particular

mineral,by direct formation of H2s which in turn reacts,

with oxidized sulphur species":

(VII.5.18)

It is proposed that this specific value for pyrrhotite

minerals is 4.0 pH. At this initial pH value, there is no

significant variation of pH during the oxidation process.

Table VII.5.6 shows the materials balance for experi

ments 18,19 and 20 conducted in the 'oxygraph' and in the

'shake' flasks. From the latter, the materials obtained from

5 days of leaching (120 hours) at initial pH values of 4.0

and 4.5 were selected.

The extent of reactions for experiments 18, 19, and 20

was derived from consumption of oxygen. For the 'shake'

flasks, it was derived from the elemental sulphur obtained.

171

TABLE VII.5.6

MATERIALS BALANCE

Base: 30 grams of Pyrrhotite Mineral (FeS)

Experiment 18 19 20

Extent of Reaction in Relation to:

- Oxygen Consumption, %:

34.35 40.79 55.85

- Elemental Sulphur, %:

25.61 36.66 26.09

4.0pH 5° 4.5pH 5°

Day Day

43.14 38.39

a: Final Weight, g:(FeOOH + s0 + FeS + Fe2 (S04 )3 .9H20)

33.285 32.534 32.700 32.960 33.838

b: Final Weight (theoretically), g:

33.869 34.592 36.291

Solution Phase:

H+ Liberated/consumed

Fe2 +, ppm: 9

Fe3 +, ppm: 3.5

3.72

0.355

22

0.3

Total Solution Phase, g: (b-a)

0.584 2.08 3.591

Leached Precipitated

c: Elemental Sulphur, g:

2.802 4.011 2.855

172

4.72* 4.20*

d: Elemental Sulphur to obtain (theoretically), g:

3.758 4.463 6.111 4.72* 4.20*

e: Residue (FeOOH + FeS + Fe2 (S04 ) 3 .xH20), g: (a-c)

30.483 28.523 29.845 28.240 29.638

f: Ferric Hydroxide to obtain (FeOOH, theoretically) g:

10.416 12.369 16.935 13.08 11.641

g: Unreacted FeS; g:

3.720 2.043 negli. 0.526 negli.

h: Unreacted FeS to obtain (theoretically), g:

19.695 17.76 13.245 17.058 18.483

i: Net Ferric hydroxide, FeOOH + Fe2 (s04 ) 3 .xH20:

10.788 10.763 16.6 11.182 11.155

Percentage of Elemental Sulphur Oxidized,%:

25.44 10.13 53.28

* considered the same value.

FeS Molecular Weight, g: 87.907.

FeOOH Molecular Weight, g: 88.8537.

173

The total value of oxygen consumed allows us to deter

mine the theoretical amount of elemental sulphur present in

the residue. If these values were compared with the practi

cal elemental sulphur obtained, it is noted that there is

variation. It is assumed that this variation is due to

oxidation of elemental sulphur of about 25.44, 10.13 and

53.28 (pure oxygen) per cent for experiment 18, 19 and 20,

respectively.

Similarly,

calculate the

the amount of oxygen consumed allows us to

theoretical quantity of ferric hydroxide

formed and of unreacted pyrrhotite. It is to be pointed out

that the extent of reaction of pyrrhotite minerals by

ferric ions is also included in the extent of reaction by

the consumption of oxygen since oxygen is consumed in the

oxidation of ferrous ions. If the theoretical amount of

unreacted pyrrhotite is compared with the practical value of

unreacted pyrrhotite, a ratio of discrepancies of theoreti

cal to practical values of 5.29, 8.69 and 13 is obtained.

Thus the separation by elution was not effective as expect

ed, but also these values appear appear to be to high.

There is a good agreement in the quantity of ferric

hydroxide obtained practically and theoretically. It is

expected that the practical value should be slightly higher

than the practical value since precipitation of basic ferric

sulphates is expected to occur as in experiment 18. The

practical value of experiment 19 is slightly lower than the

theoretical since collection of the slurry from the 'oxy

graph' was more difficult (due to splashes) than that from

174

'shake' flasks. Thus some experimental errors may have

occurred.

Theoretical values from the 'shake' flasks experiments

were derived from the practical elemental sulphur obtained.

Nevertheless, it is meanningless to calculate a comparison

between the extent of reactions obtained from the •oxygraph'

and those from the 'shake' flasks since the overall oxida

tion of elemental sulphur is about 18 per cent in the

•oxygraph' with air and 53 per cent with pure oxygen. The

percentage of overall oxidation of elemental sulphur ob

tained by Harris et al. (79) is from 2 to 20 five per cent

with T. ferrooxidans and in a non-sterile reactor.

(106) did not report elemental sulphur.

Sawe

Determination of the extent of oxidation by elemental

sulphur is also irrelevant since 0.902 grams of elemental

sulphur was extracted from 30 grams of unreacted pyrrhotite.

Moreover, apparently, the extent of oxidation was higher

from experiments conducted in the 'shake' flasks than from

those conducted in the •oxygraph'. This may not be correct

since oxidation of elemental sulphur could also depend on

the kind of stirring in the presence of ferric hydroxide and

unreacted pyrrhotite. The stirring was faster in the •oxy

graph' than in the 'shake' flasks, thus more elemental

sulphur was oxidized in the •oxygraph' than that in the

'shake' flasks.

If oxidation of elemental sulphur to sulphate has

occurred during the aqueous oxidation, the oxygen in the

sulphate was unlikely proceeded from the oxygen partial

175

pressure. It is shown in Table VII.5.7, 'oxygen distribu-

tion' that the oxygen consumed by the sulphur theoretically

oxidized for experiment 20 must be 6.4995 grams. The fact

that the actual oxygen consumed, 4.5743 grams, is less than

the oxygen to be consumed in the oxidation of sulphate

implies that the oxygen in the sulphate did not proceeded

from the dissolved oxygen. The fact that extent of reac

tion obtained for experiment 20 (conducted in pure oxygen)

is 55.85 and 26.09 per cent in terms of consumption of

oxygen and elemental sulphur, respectively; confirms that

the extent of reaction is more reliably determined by the

consumption of oxygen rather than by determining the elemen

tal sulphur formed. It was not possible to show the same

conclusion for experiment 18 and 19 since the oxygen to be

used in the sulphate is less than the actual oxygen con

sumed.

176

TABLE VII.5.7

DISTRIBUTION OF OXYGEN

18

Elemental Sulphur Obtained,g: 2.802

Experiments

19

4.011

20

2.855

Elemental Sulphur to be Obtained Theoretically, g:

3.758 4.463 6 .1·11

Elemental Sulphur Oxidized,g: 0.956 0.452 3.256

Sulphuric Acid to be formed,g:2.924 1.3827 9.96

Oxygen in Sulphuric Acid, g: 1.908 1.044 6.4995

Actual Oxygen Consumed, g: 2.8133 N.A. 4.5743

Final pH to be Obtained -1.095 -1. 043 -1.360

Final pH Obtained: 3.014 2.988 3.721

Stoichiometric Reactions: Equations VI.1.5 and VI.1.7

FeS + 0.7502 + 0.5H20 -----> FeOOH + s 0

FeS + 2.2502 + 1.5H20 -----> FeOOH + H2so4

(Molecular Weights: Sulphuric Acid, 98.0734; elemental

sulphur, 32.06; Oxygen, 15.9994; Hydrogen, 1.0079)

177

The fate of oxidized elemental sulphur products is not

clearly understood. The quantity of ferrous and ferric ions

in solution obtained does not balance the total amount of

sulphur variation between the actual and theoretical values

even as ferrous and ferric sulphate. No sulphur oxidized

compounds in solution such as H2s 2o 2 , H2so2 and so32 - has

been determined. Moreover, it was shown in the potential-pH

diagram (Figure V.3.1) for the metastable sulphur system

that at 25°c and 4.0 pH s 2o32 - and Hso3 - are stable. Thomp

son (125) mentioned thiosulphate, tetrathionate and tri

thionate which will be discussed later. However, precipita

tion of basic ferric sulphates such as defined by Adam et

al. (130) as Fe(OH)S04 , and/or Fe2 (S04 ) 3 .9H2o (coquimbite)

by Case et al. (126) can be expected from electrochemical

studies of pyrrhotite minerals at 3-5 pH values. More stud

ies need to be conducted to identify the sulphur species in

solution and 'basic ferric sulphate' precipitates, which are

beyond the aim of the present study.

178

VII.6 CONCLUSION AND DISCUSSION

Three designs of an 'oxygraph' were used in the deter

mination of the rate of consumption of oxygen during the

aqueous oxidation of a pyrrhotite concentrate tailing from

Renison Bell, Tasmania. The results of a fourth series of

experiments, using 'shake' flasks, were compared with those

of the 'oxygraph'.

The results obtained from the first oxygraph were not

reliable due to the reactor vessel size which could hold

only about 30 milliliters (1.5 grams of mineral

density of 5% w/v). Thus measurement of the low

at slurry

rate of

consumption of oxygen was affected by the vibration caused

by stirring the reactor vessel and by changes of temperature

and atmospheric pressure during the day.

The second oxygragh used, with a one liter reactor

vessel, was designed to determine the rate of consumption of

oxygen in the presence and absence of microorganisms. The

figures obtained from these series of experiments still

show some cyclic variations. The internal pressure of the

leaching

through

system was connected to the

the manometer. Thus, slight

atmospheric

variations

pressure

of both

ambient temperature and atmospheric pressure seemed to

disturb the electrolytic production of oxygen.

Nevertheless, as the rate of consumption

behaved with an overall linearity, it is posible

some conclusions. The higher consumption rate

179

of

to

of

oxygen

reach

oxygen

could be distinguished when the reaction process took place

in the presence of bacteria from a chemical process alone.

At 4.0 initial pH, the chemical oxidation of pyrrhotite

(Experiment 2) showed a rate of consumption of oxygen of

0.00036 grams per minute. Bacterial oxidation of the same

sample (Sample 2, Experiment 3) showed an improvement in the

rate of consumption of oxygen to 0.0020 grams/min.

The consumption of oxygen of Sample 2 was slightly

faster than that of Sample 0. It was not possible to state

accurately how much faster since 2 experiments at the same

conditions of initial pH were not conducted. This conclusion

is derived from the difference in the rate of consumption of

oxygen obtained by the difference in the ratio of hydrogen

ions concentrations at 5.5, 4.5 and 4.0 initial pH values

explained in the previous Section.

Under a constant fixed atmospheric pressure and tem

perature of about 31.5(+/- 5°) 0 c the rate of consumption of

oxygen during aqueous oxidation was controlled more accu

rately. Nevertheless, some cyclic variations due to ambient

temperature fluctuations were obtained. It was difficult to

maintain a constant temperature accurately (e.g. 31.s0 c

exactly) without an air conditioning system for long periods

(e.g. 8 days). However, the rate of consumption of oxygen in

the presence of air was 0.00037 grams per minute (experiment

15) and in pure oxygen 0.00057 grams/minute (experiment 20),

representing about 1.56 times faster in pure oxygen than

with air. The rate of uptake of oxygen in experiment 2

(second oxygraph) with air was 0.00036 g/min, which agrees

180

satisfactorily with experiment 15. Both experiments were

conducted for 9.5 days.

At an initial pH of 4.0, the overall chemical aqueous

oxidation of pyrrhotite minerals, was found to form ferric

hydroxide and elemental sulphur predominantly according to

Equations (VI.3.5) and/or (VI.3.12):

Stoichiometric pyrrhotite:

FeS + 0.5H20 + 0.7502 -----> FeOOH + s 0

Monoclinic pyrrhotite:

Fe7S9 + 3.5H20 + 5.2502 ---> 7FeOOH + ss0

Characterization and extent of formation of 'basic

ferric sulphate' compounds are difficult to determine with

out a 57Mossbauer sprectroscopy (47). Thus it is unkwon the

amount of Fe( 2 (S04 ) 3 .9H2o (coquimbite) (126), Fe(OH)S04

(130) and/or Fe2o3so3 .10H2o (56) that might have also

formed. These compounds were reported to form according to

electrochemical studies of ferrous sulphide and pyrrhotite

minerals in the pH range of 3 to 5.

The overall oxidation of elemental sulphur was found to

be about 18 per cent in a chemical oxidation process with

air and about 53 per cent with pure oxygen.

The source of oxygen consumed during the oxidation of

elemental sulphur is not clearly defined. However, it is

proposed that oxygen forming the oxidized sulphur species in

solution, e.g. sulphate, is provided by the water rather

than from the oxygen partial pressure. The fact that the

181

oxidation of elemental sulphur starts to occur after the

third day of the oxidation process implies that accumulation

of oxidized sulphur species and ferric hydroxide was taking

place. Moreover, oxidation of hydrogen sulphide and elemen

tal sulphur at 4.0 pH and ambient temperatures is very slow.

Lottens et al. (104) studied the behaviour of sulphur in the

oxidative leaching of sulphide minerals and described the

reaction mechanism as follows: "oxidation of s-2 at the

crystal/leach liquor interface to either s+ or s 2 + (depend

ing in the oxidizing agent used) and hydrolysis of s+ or s 2 +

sX+ + XH20 -----> H sXo + xH+ X X (VII.6.1)

disproportionate or react with hydrogen sulphide to form

elemental sulphur and so32-. The sulphite thus formed is

subsequently oxidized to sulphate or enters the "Wackenrod

er" reaction sequence with hydrogen sulphide generated from

the dissolution of the metal sulphide (104).

The proposition made in the present thesis may also be

confirmed by the electrochemical studies conducted by Bailey

and Peters (104) and Hamilton and Woods (17,101). Bailey et

al. (104) stated that "pyrite dissolution under acid oxygen

pressure leaching conditions was found to have an electro

chemical mechanism; i.e. the sulphate formed from pyrite

sulphur contains oxygen taken from water rather than from

the high-pressure gas phase". From the experimental condi

tions used to conduct the electrochemical studies, Hamilton

et al. (17,101) reported that in order to avoid any oxida-

182

tion of the electrode surface (pyrrhotite mineral) the prepa

ration was made in a nitrogen atmosphere (in a glove bag).

The cell solution was also purged with purified nitrogen.

Thus, it is assumed that oxygen was absent in the system and

the fact that formation of sulphate was determined also

confirms the possibility that the oxygen in the sulphate

formed from the pyrrhotite sulphur contains oxygen taken

from the water rather than from the oxygen partial pressure.

On the other hand, electrochemical studies on the reduc

tion of oxygen on sulphide minerals have shown that in the

limiting current region, all oxygen diffusing to the surface

is reduced to water (67-69).

Thus, it is concluded that if the oxygen forming the

sulphate ions comes from the water, the measurement of the

rate of consumption of oxygen is a reliable method to deter

mine the rate of oxidation pyrrhotite minerals.

More study needs to be conducted to identify the sulphur

species in solution and 'basic ferric sulphate' precipitates

to describe fully the reaction mechanism.

183

VII.7 PRESENT KNOWLEDGE ON CONTINUOUS BACTERIAL OXIDATION

OF PYRITE CONCENTRATE AT MINIPLANT SCALE

VII.7.1. INTRODUCTION

Neeling et al.(116) and Roberts (132) indicated some

results obtained from bacterial oxidation of pyrite concen

trate at pilot plant scale. Further evaluation of the

operation conditions of this pilot plant was needed in order

to reduce the processing time. Those operation conditions

were studied in the miniplant (the pilot plant was shut

down) and the results are presented in this study. Firstly,

it appeared that concentrations of dissolved oxygen higher

than about 3 mg.L- 1 have an adverse effect on the activity

of the microorganisms and hence on the rate of mineral

degradation. Secondly, in an attempt to enhance the rate of

bacterial growth and to prevent "washing out" of bacteria in

a continuous flow system, the recycling effect of some

concentrate and/or liquor was studied. Thirdly, the beha

viour of the microorganisms, T. ferrooxidans and T. Thiooxi

dans, in sulphur oxidizing 9K media was studied.

In this study, a continuous multistage bacterial

leaching plant, using the same concentrate as Neeling et al.

(116), was examined in order to clarify the effect of both

concentrate and liquor recycle, the effect of dissolved

oxygen and the effect of the 9K media on the bacterial

activity and the rate of mineral degradation.

184

VII.7.2 EXPERIMENTAL PROGRAM

The continuous bacterial oxidation of pyrite concen

trate from Hellyer, Tasmania in a small scale system con-

sisting of six stirred reactors in series, was studied.

system was operated continuously (24 hours per day - 7

per week) for the duration of all the experiments.

The

days

The

equipment used is shown in Figures VII.7.21 and VII.7.22.

The reactor tanks were constructed of PVC and used

stainless steel agitators with double radial impellers,

actuated by air motors. Diaphragm and peristaltic dosing

pumps were used to control the flow rates of slurry and

bacterial nutrients (9K solution). A pump installed between

the slurry feed tank (first tank) and Tl (second tank)

controlled the residence time throughout the system. Flow of

slurry through the reactors Tl to TS was by gravity.

The pyrite concentrate had an average chemical composi

tion of: gold 3 mg.L- 1 , silver 80 mg.L- 1 , iron 42.5 %, zinc

7.5 % and arsenic 0.5%. Mineralogical examination indicated

that the concentrate was predominantly pyrite with some

sphalerite and minor amounts of arsenic mostly in solid

solution with pyrite. The particle size was 90% passing 38

micrometers.

Results of the bacterial oxidation process were evalu

ated by analysis of iron, zinc and arsenic accumulated in

solution and by measuring the slurry potential (Eh) and pH.

The elements in solution were analyzed using an AA-1275

Varian Atomic Absortion Spectrophotometer (AAS); Eh and pH

185

H.C.

':llvQ.Q.'( Fi:f:.I) 16"'/0 v->/v l'D. (i-!EL.0?l) FIGURE VII. 7.21: CONTINUOUS BACTERIAL OXIDATION OF PYRITE CO,.\JCE:\/TRA TE

$I=

(_ Pver~,1 v-,1;111._ nrn1ro.li..,}c...( w£1.t.u-)

Aia, 5,11(2.bE: \\C..

12..rr,.l/r.,-,~

1\

PU ~P".J -0- peri'.:>foltcc.. -$- ~t).phrdm

\-\. c.. -= \.le.o.T..,0 (,O..t

II q~

11 I "" .,,,., 0 2. m\/m;" '::::,\vrr~ rec.it'<!.. 1---_

1'

T2

'T 0-\-c,.,\ \JolvMe./f11-n k. .. ,:.3 Z.4 L · \(.Jcrk:.;.,.,d \Jo\vMe/10."1<..-::. 28,4 L TvJc 1-0..J..:,:.\ ivv,~e.\\Qrs/10.,,,\:.....

a

TZ.

k"t.nu, \:ietwee"' h., pc.Iler~-:: 16cm. ,Di'bti:nc.e- be.."tweet'\ hdt0m'i,,,,.~el/t!.f"o..n.A ~o,...,4-e.nk: l,,-{5c...,... AiQ.. '!>fA!l.6'1::: '50 fhj\;>'12 = 345 ~h ... (flow \Uk ~sio). 5TIU.i.llJ~. 12-f'\1:. 220--2~0. Alfl MOTOl-S · 4AI,,\- H.Y~24 MC!).EL

C..FM (.j--re,e.. /.\,.() ~m~~ '& x (p (ht01ar:i)-=- 4'6 ci::-t1. 'IDLUM\: OF Tr\\C.1<.(IJ&(l,-:. \l.'2'l L.

0

T4 .1:.

"'T5

To d''"' 11,ic,

FIGURE VII.7.22: MINIPLANT FOR COI\JTJI\JUOUS BACTERIAL OXIDATION OF IRON SULPHIDE Ml!\iERALS

Varian Atomic Absortion Spectrophotometer (AAS); Eh and pH

by an Orion Research Ion Analyser EA940.

Samples from each stirred reactor and the thickener(s)

(overflow and underflow) were taken for analysis and the

pulp density was measured every Monday, Wednesday and Fri

day. Flow rates, pH, Eh and dissolved oxygen (D.O.) were

monitored every day. Procedure for the determination of

iron, arsenic and zinc is given in Appendix H.

A N.A.T.A. recognized external laboratory determined

the gold and silver contents of the original concentrate

and of the samples taken from slurry feed, Tl and TS.

Although cell counts could be made on the liquor and on

the surface of the mineral samples, the cell number determi

nation on the surface of the ore were usually used as a

measure of effectiveness of the leaching process. These cell

counts were usually higher that those obtained from the

liquor when the relative bacteria activity is high. Bacteria

activity and cell morphology, determined by the proportion

of green, orange and red cells, obtained from nitrogen

determination by the Kjeldahl Method and by Fluorescence

Microscopy, were examined by a microbiologist.

Water, at a controlled temperature of about 38°c,

circulated through independent coils in each stirred tank.

Additions of pyrite concentrate, bacterial nutrients (9K),

cell recycle, whether as slurry or solution, were made to

the slurry feed tank. The level of dissolved oxygen was

controlled by sparging compressed air (40 p.s.i.) into each

stirred reactor independently. A dissolved oxygen meter was

186

used to measure the oxygen in each tank and the sparged air

rate was adjusted manually to keep the oxygen level approxi

mately constant.

In this section, three flow diagrams designated as

MINIP 03, MINIP 04 and MINIP 05 are discussed. Table 1

shows the more important operating conditions under which

each system was run. Whether the slurry recycling took place

from TS to SF tank or from T2 to Tl, detailed in each mini

plant, it is reported simply as slurry recycle.

187

TABLE 1

OPERATION CONDITIONS

SYSTEM

Number of Reactors

Reactor Size, L.

Pulp Density, % (w/v)

MINIP 03

6

30

15

Initial pH in SF Tank 1.6-1.75 1.2

MINIP 04 MINIP 05

6 6

30 30

15 then 10 12-13

1.65,Tl:l.5

Rate Slurry Feed, ml/min 12.2 17-18 17-18

Slurry Recycle, ml/min 2.1 7-8 7-8

Dissolved Oxygen,mg.L- 1 . 2-3 3 3-5 then >6

Stirring Speed 210-220 200-210 215-220

9K Solution* standard standard modified

Temperature, 0 c 36-38 36-38 36-38

Flow Rate Air Sparge,scc/min/L:1.5

Residence Time, Hrs. 250

Total Time in Operation,Hrs.860

1.5

500

816

1.5 then 2.6

300

1780

* "Standard" 9K was prepared according to the formulation

described by Silverman et al.(114). For 1 L of solution

without ferrous ions as follow, g: (NH4 )2so4 3.0, KCl 0.10,

K2HP04 0.50, and MgS04 .7H2o 0.50. "Modified" indicates that

3 g of K2HP04 , instead of 0.50 g, was used.

188

VII.7.3 EXPERIMENTAL RESULTS

The rate of degradation of refractory gold minerals,

pyrite, by bacterial oxidation seems directly related to

the level of activity of the genera Thiobacillus, from which

T. Ferrooxidans and T. Thiooxidans seem to have more impor

tance. The level of activity may be indirectly determined by

the rate of accumulation of dissolved species in solution,

by the rate of reduction of pH values and by the increase

of slurry potential throughout the stirred reactors system.

The level of activity of microorganisms may be enhanced by

selecting a correct system pattern, level of dissolved

oxygen, residence time, pulp density and temperature among

other variables.

As would be expected, stirring speed and pulp density

influence dissolved the oxygen concentration in the system.

At constant air sparge flow rate, the concentration of

dissolved oxygen increased because the stirring speed in

creased automatically when the pulp density decreased. In

MINIP 03 and MINIP 04, the dissolved oxygen was maintained

at low level (<3 mg.L- 1 ) by decreasing the air sparge flow

rate. When the air sparge rate was increased, a reduction

in the ratio of active to inactive cell population from

about 85-90 per cent was noted. This reduction of activity

was accompanied by excessive formation of froth for several

hours, reduction in cell size and a change in cell shape.

The ferric ions concentration was found to increase

189

steadily from the slurry feed tank to TS. Table II shows the

typical values obtained from MINIP 05. The ferric ion

concentration is calculated as the difference between

"soluble" iron concentration (as measured by the AAS) and

the ferrous iron concentration as measured by potassium

permanganate titration.

The increase of slurry potential throughout the stirred

tanks is probably due to the increasing ferric/ferrous

ratio.

190

TABLE 2

FERROUS IONS TITRATION BY POTASSIUM PERMANGANATE IN MINIP 05

(Data Extracted from the Spreadsheet, Date: 5-2-88)

Fe2 + g/L:

Fe* g/L:

Fe3 +;Fe2 +:

Eh,rnV(SHE):

pH

SF

4.87

4.90

0.01

582

1. 77

Tl

1.51

1. 62

0.07

579

1. 79

T2

1.46

1.67

0.24

584

1.80

* "soluble" iron determined by AAS.

191

T3

1. 79

1.97

0.10

593

1. 73

T4

1.40

2.88

1.01

684

1.66

T5

1.18

4.54

2.85

678

1. 54

In MINIP 03 system described by flow diagram Figure

VII.7.20, the slurry feed rate was kept at about 12 ml/min

(SF to Tl) and bacteria recirculation took place from TS to

Tl at 2 ml/min as slurry. During the leaching run it was

found that the concentration of iron, zinc and arsenic

increased up to a maximum of 9500, 12000 and 350 mg.L- 1

after about 300 hours and then decreased to about 6000,

9500 and 280 mg.L- 1 , after 860 hours when the experiment was

terminated. Based on the "total" dissolved iron analysis,

the average extent of pyrite oxidation was only 12.5%.

In MINIP 04, Figure VII.7.22, because of the low oxida

tion rate it was decided to stop the bacterial recycle,

decrease the pulp density from 15 to 10 per cent and in

crease the residence time from 260 to about 500 hours main

taining about the same level of dissolved oxygen. Bacterial

recycle from T5 to Tl was suspended because it had been

found that inactive bacteria were occasionally pumped back

to the system. A net slurry flow of 5 ml/min was maintained

by pumping 15 ml/min of slurry with an initial pulp density

of 10 % (w/v) from SF to Tl and recirculating 10 ml/min from

Tl to SF tank. An automatic pH controller (Prominent dulcom

eter dosing pump, Type PHO) was installed to maintain the pH

at 1.2 in Tl. Using an air sparge of about 1.5 scc/min/L of

slurry, the dissolved oxygen was maintained at about 3

L-1 mg. .

Figures VII.7.23, VII.7.24 and VII.7.25 show the

variation of slurry potential (Eh), "total" dissolved iron

192

FIGURE \/II.7.22

MINI-PLANT 04, NO SLURRY OR SOLUTION RECYCLE SLURRY FLOW 5 ML/MIN AT 10% W/V SOLIDS

FEED

'- ~ ~

r""'\ ' ,, ·" / \../ .,.,..

15 'rr.\ lr,,·..-.

10 \"(111,~·.:1,. T2 T3 T4 TS

..

Tl

> E

...... 0 a 0

..:,

K .:. .,; 0 a:. Q lJ.i

~ 0 V) !Q. Q

700

690

680

670

660

650

64-0

630

620

610

600

690

580

570

1560

tltlO

a .a

9

8

7

6

5

..

3

2

D Sf

0 200

MTNTPT,Al\J'T' 04 SLURRY POTENTIAL

PER T.A.NK

400 600

cmAULATI\IE TIME, Hr51.

800

SF Tank + Tank 1 X fonk +

<> Tank 2 Tr;mk 3 V Timk 5

FIGURE VII.7.23: Variation of Potential in MINIP 04 at 1.5 Initial pH.

MlNIPLANT 04-IRoN OISSOWTION

0 200 400 600 800

CUMULATIVE TIME, HP.S.

+ Tt C> T2 t.. T3 X H V T5

FIGURE VII.7.24: Dissolved Iron in MINIP 04 per Tank

17

Hi

15

14

1.3

12 K

z 11

0 10 et:

Cl g

~ 8 Cl Li 7 Li

t;i u.

5

+ J

2 1

0 a

,-.r , . .., , .'lJ•::a.. J I

I RON DJ.S SOLU TION

r I t-~-.~

I \

PER TANK

/ ~

I /

v-J

/ u\

' l \ ' i \ F \ / \/ g C'J

.oaa

),'. Trinf; 5 ·

'"'"

V TOTAL

Q CH\

FIGURE VII.7.25 : Total and Pe r Tan k Dissolved Iron.

( MINIP 04).

17

Hi

15

14

1.3

12 ),:

11 z 0 10 It:

0. 9

~ 8 Q 7 ~, ~,

ti Cl.

5

4

:3

2

0

JRON DJ:SSOLUTION PER TANK

0 / \I // \\ ,/f \, I <4\

'· , ~ \ \ ~ ci ..... \ \.

(\ ,.aa "'"' '""' CUMUL'-.T!VE TIME, Hr.i.

r·, r..-1.,·, .:1r•XI I

V TOTAL

FIGURE VII.7.25: Total and Per Tank Dissolved Iron.

(MIN IP 04).

and the percentage of iron dissolution in each tank, respec

tively. At 1.5 pH in the SF tank, the slurry potential

initially increased then decreased steadily in all tanks.

This behaviour is not understood. However, it appears that

when the potential increased, the concentration of dissolved

species also increased.

Increasing the residence time from about 250 hours in

MINIP 03 to about 500 hours in MINIP 04 and decreasing the

pulp density from 15 to 10 % (w/v) resulted in the concen

tration of dissolved species in TS increasing to a maximum

of about 8,250 mg.L- 1 , 10,000 mg.L- 1 and 350 mg.L- 1 of iron,

zinc and arsenic, respectively. This indicated an extent of

oxidation of about 16 per cent of iron.

In contrast, the system described by the flow diagram

of MINIP 05, Figure VII.7.26, was run at a higher dissolved

oxygen concentration 3-5 mg.L- 1 . simply by slowly increas

ing the air sparge rate over a period of about 100 hours. In

the MINIP 05 system the leach solution, containing ferric

ions, was recycled via the overflow of thickener (2) to the

slurry feed tank. Ferric acid leaching took place in SF tank

and the resultant ferrous ions were eliminated via the

overflow of thickener 1. Some of the slurry from T2 was

recycled to Tl, in order to obtain longer residence time and

to enhance growth of the bacteria.

Figures VII.7.27, VII.7.28, VII.7.29 and VII.7.30 show

the variation of pH, Eh, "total" iron and zinc concentration

in each tank, respectively. After about 650 hours of opera

tion, the pH of the "modified" solution was maintained at

193

FIGURE \/II.7.26:

r·10 ollfied 9K

feed WO..

WO-t 0

( D

L. er . , '

r-) - ~ ., '- - - -, I t--

~~\ -n T3 T4 T5

•.

1 t . l 1 t soUd

.. '

Ill r--ecycle

'

- -Tr\'(,1 L) '- Tf-l'lz.

12.

~ //

wa.ste ,11 u (D

sotu-tlon recycle \

\V

oold lea.eh

It

z 0 ~ c,.

~ ...l 0 VI !!'.!. LI.

" 0 ~ N

c,.

~ 0 VI VI zs:

rJ

26

24

22

2(l

18

16

1+

12

10

B

fi

4

2

0

0

+ Tl&:2

0.2

MIN1PLANT 05 IRON DISSOLUTION

0.4 O.fi (Th0u50ndii)

CUMULATIVE TIME, HRS.

t.. T4 l( T5

0.1'1

V Toti;il

FIGURE VII.7.29: Total and per Tank Dissolved Iron.

·MIN1PLANT 05 ZINC DISSOLUTION

140

130

120

110

10(l

9~,

80

70

60

50

4-0

30

2(l

10

0

0 0.2 0.4 O.!i 0.8 ( Thou IIQn d~)

CUMULATIVE TME, HRS.

SF + Tl,i:2 <> T3 )( T5 V foli;il

FIGURE VIU.30: Total and per Tank Dissolved Zinc.

:c u.

/

> E £. 11.1

2.1

.R -2 ,/

'Ei I

1.9

1.8

1.7

1.6

1.5

1.4

1.J

1.2

1. 1

0

Li SF funk A T,1n,,. ,

.)

I I I

\ I I.

0.2

pH MJNIPLANT :> PER Tli.NK

-"1 __ .-----,,

\ \ C:r'

0.4 O.fi 0.8 ( Trou ,,,..n Qt1)

. CUMULATI\IE TIME, Hl"ll.

+ Tank I <> Tank 2 X fonk 4 V T,;mk 5

FIGURE VII.7.27: Variation of pH per Tank.

700

61)0

680

1570

660 fillO

6+i.1

6~0

6:'!0

610 fiOO

5!il0

!!BO

570

!ilSO

550 540

tiJO

0 0.2

D '.:iF TUl,r.

...:. 1nm: )

FIGURE VII. 7 .28:

SLURRY .POTENTIAL PER Tli.NK

0.4 0.6 (Thouer.in9")

CUMULATI\.£ Tl~4E, Hl"ll.

O.B

+ Tunk I X fon,: 4

<> Tank 2 V Trmk 5

Variation of Potential per Tank. (MINIP 05)

1.5. The potential in each tank did not vary greatly and

was almost constant in T4 and T5. Zinc was completely dis

solved in the first three tanks. The concentration of iron

in T5 improved after the pH of the "modified" 9K solution

was kept at 1.5.

Table 3 shows the results obtained when the air sparge

flow rate was increased to about 2.6 scc/min/L.slurry in

terms of dissolved iron, pH and Eh, after the oxygen concen-

tration was raised to >6 mg.L- 1 . The "total" iron concentra

tion in T5 was found to be about 20.6 compared with 27 mg.L-

1 in T4. It is not understood how the "total" iron concen-

tration in T5 can be lower than that in T4 but it could be

that, because of some lack of air sparge in TS, an iron

precipitation formed which is not soluble in the hydrochlor

ic acid leach in the chemical analysis procedure. No attempt

was made to identify any such material. A slight increase in

pH and decrease in slurry potential (Eh) in the same tank

was noted.

Dissolution of pyrite increased substantially up to

about 50 per cent related to iron oxidation. It was also

noted that dissolved oxygen concentration increased well

above 7 mg.L- 1 (7-9 mg.L- 1 )

194

TABLE 3

RESULTS FOR MINIP 05

For Dissolved Oxygen Concentration: >6 mg.L- 1

SF

pH 1. 35

Eh (mV) 677

Fe, g/L. 14. 1

Tl

1. 06

755

16.9

T2

0.97

783

21. 2

T3

0.88

807

24.7

T4

0.85

814

27.0

For Dissolved Oxygen Concentration: <6 mg.L- 1

At 1000 Hours of Operation (MINIP 05)

Fe,g/L 3.2 3.3 3.7 2.82 4.57 5.14

At 1440 Hours of Operation (MINIP 05)

Fe,g/L 4.74 2.66 2.84 3.08 3.63 4.15

195

T5

0.88

728

20.6

VII.7.4 DISCUSSION AND CONCLUSIONS

The rate of degradation of pyrite concentrate from

Hellyer, Tasmania by bacterial oxidation can be enhanced by

acclimatization of the microorganisms, F. ferrooxidans and

F. thiooxidans to high concentrations of dissolved

c-10 mg.L- 1 ). This acclimatization is a slow process.

oxygen

Under

the operation conditions of MINIP 05, it took about 4 weeks.

Bacterial oxidation of pyrite concentrate in a continuous

system was enhanced up to about 50 per cent based on iron in

pyrite oxidation (260 hours residence time) after the air

sparge rate was slowly increased to about 2.6 scc/min per

liter of slurry.

The effect of the 9K solution whether as "standard" or

"modified" on the rate of degradation of pyrite concentrate

at low air sparge rate was found to be non-existent. In

creasing the concentration of potassium ortho-phosphate in a

standard 9K solution from 0.5 to 3.0 grams/L was thought to

enhance the growth of sulphur oxidizing bacteria; hence the

oxidation of elemental sulphur and reduced sulphur compounds

formed during the process.

Bacterial oxidation of sulphide minerals is an electro

chemical process in nature as widely discussed in the liter

ature. The mineral sphalerite found in the pyrite concen

trate was oxidized completely in the first three tanks.

Thus sulphide minerals with lower rest potentials will be

oxidized before those sulphides with higher rest potentials.

196

This electrochemical behaviour seems to follow the next

pattern in a 9K medium and at 2.5 pH (109,124):

FeS<ZnS<CUS<Cu2S<CuFeS2<FeS2.

The effect of stirring speed up to 300 r.p.m. seems to

have no effect on bacterial oxidation as long as settlement

is avoided. In the equipment used for this study and with

about 15 per cent pulp density (w/v), 220 r.p.m with double

radial impellers and diffusing aeration kept the slurry

thoroughly suspended.

It is not known whether the flow diagram of MINIP 05 is

the optimum. However, new flow systems without bacterial

recycle should be studied. Growth of the microorganisms

might be enhanced in an independent batch reactor at low

pulp density (e.g. 5 per cent w/v} in order to be fed to the

slurry feed tank and thus avoid bacteria "wash out" and/or

the size of the first tank (SF} should be from 2 to 3 times

greater than the size of the following tanks.

Bacterial population of green cells found in the mini

plant (from the solids} varied from 7 x 108 to 10 x 1012 per

milliliter of slurry. Once the microorganisms are acclima

tized, the number of inactive cells increase (red and orange

cells) when the environment is changed (e.g. dissolved

oxygen, presence of adverse conditions). But when the condi

tions return to normal, it seems that the microorganisms

revive or/and its rate of growth is very fast that the

number of green cells could be restored in about 24 hours.

No severe loss of microorganisms has been observed in the

miniplant at any time.

197

CHAPTER VIII

DISCUSSION AND CONCLUSION

Bacterial oxidation of iron sulphide minerals, with

particular reference to pyrrhotite and pyrite, is an at

tractive alternative to recover the metal values associated

with them. However, bacterial oxidation of pyrite is still

industrially impractical today due to its slow kinetics.

The electrochemistry of these minerals, as well as of

other sulphide minerals, is being explained more often by

using their potential-pH diagrams and electrochemical meth

ods. Firstly, because analysis of the lines in the poten

tial-pH diagrams may indicate the path in the leaching

mechanisms. Secondly, because electrodissolution of slurry

could play an emerging role in hydrometallurgy, e.g. to

process pyrite.

There is a dramatic difference in behaviour between the

aqueous oxidation of pyrite and pyrrhotite. The main differ

ence, found from this study and electrochemical studies

(16,101,104, 129,130), occurs in the production of elemental

sulphur and/or sulphate ions. Although it is known that

increasing the potential favours the formation of sulphate

in both cases, pyrite predominantly produces sulphate, and

pyrrhotite produces elemental sulphur. These electrochemical

behaviours are also found in the bacterial oxidation proc

ess. As the chemical oxidation of hydrogen sulphide and

198

elemental sulphur are very slow processes at ambient temper

ature, predominance of sulphur pyrite oxidation to sulphate

and sulphur pyrrhotite to elemental sulphur is not known

accurately. However, it seems that the state of oxidation of

sulphur present in pyrite and pyrrhotite as s 22- and s 2-,

respectively, their electronic structure, affects the anodic

electrochemistry.

The bacterial oxidation of pyrite concentrate in the

miniplant with initial and final pH values of 1.5 and ~1.0,

respecitively, was found to produce sulphuric acid predomi

nantly with small formation of elemental sulphur (less than

four per cent). Chemical oxidation of pyrrhotite minerals at

initial pH values of 4.0 was found to produce elemental

sulphur predominantly. Oxidation of reduced sulphur com

pounds from the leaching of pyrite and pyrrhotite is en

hanced by the presence of T. ferrooxidans and T. thiooxidans

and by chemical aqueous oxidation with pure oxygen.

Bacterial oxidation of pyrite concentrate at miniplant

scale was found to be faster than that at laboratory scale.

Although it is impossible to provide a complete explanation,

one reason could be an increased concentration of dissolved

oxygen in the system, a thorough acclimatization of the

microorganisms to the pyrite. At the beginning of any new

flow diagram, the rate of degradation is slow, then it

increases steadily by itself. The process increases to an

optimum level when a steady-state condition in each stage

of the process (acclimatization of the microorganisms to the

flow diagram) is reached. This process of acclimatization

199

could take 2 or 3 residence times.

The rate of degradation of pyrite by bacterial oxida

tion by T. ferrooxidans and T. thiooxidans needs to be

improved to become an economical process. It is expected

that the development of genetic engineering techniques may

greatly improve the bioprocessing of pyrite by the year

2000. However, under the conditions studied and the extent

of reactions obtained, bacterial oxidation of pyrrhotite

minerals is faster than that of pyrite: It could be said to

be approximately three times faster.

Thus if it is decided to recover the tin encapsulated

in the pyrrhotite mineral, chemical aqueous oxidation with

air or preferable with pure oxygen is an alternative. Howev

er, bacterial oxidation at initial pH value of 4.0 could be

more suitable. The next stage, after this study, should be

the investigation of this process at miniplant scale and/or

bacterial heap leaching.

200

APPENDIX A

THERMODYNAMIC DATA AT 25°c (10)

Formula Gf0 (Kcal/rnol)

HS- 2.88

s2- 20.5

HSO -4 -180.69

so 2 -4 -177.97

Fe2 + -21.75 (51)

-18. 95;1-.

FeOH+ -66.3 (7)

HFeo2 - -90.3 (7)

Fe3 + -4.06 (51)

-1. 1 *

FeoH2 + -54.83

Fe2 (0H) 2 4+ -111.68

Fe(OH)2 + -104.68

Feo42 - -111.7

H+ 0.0

Fe(s)(alfa) 0.0

FeO -58.59

Fe3o 4 -242.7

Fe203 -177.4

Fe(OH)2 -116.3

Fe(OH)3 -166.5

(alfa)-FeOOH -116.77(18)

H2S(aq) -6.66

S(s) 0.0

Fe1.000S -24.0

Fe7s8 -178.9(7)

H2 (g) 0.0

02 (g) 0.0

H20(l) -56.688

* Data used to calculate the equations in Appendix C.

--o--

APPENDIX B

DERIVATION OF POTENTIAL-pH DIAGRAMS

Equilibria in aqueous solutions can be represented by

the general reduction reaction:

aA + xH+ + ze- = bB +

for which the standard reaction isotherm:

AG - A.G 0 = RT ln K

can be treated using the following substitutions and

assumptions:

T = 298.15°K

~G = - zFE (in Kcals)

where activitity of water is assumed to have a value

of one and pH= -log aH+

F = 23.0609 Kcal/ volt,

R = 1.98717 cal / deg. mol.

to give the potential - pH relationship:

zE / 0.05916 = -!G0 /l.3642 - xpH - log a b1a a B I A

(in Kcal. l

APPENDIX C

EQUATIONS RELATING Eh, pH AND ACTIVITY FOR THE IRON-WATER

SULPHUR SYSTEM AT 25~c

A. 2H+ + 2e- -----> H2

Eh= 0 - 0.05916pH 0.0295 log a.H2

B. 02 + 4H+ + 4e- -----> 2H20

Eh= 1.229 - 0.05916 pH+ 0.0148 log a.02

Two Dissolved Species:

la. Fe(OH)+ + H+ = Fe2 + + H20

pH= 6.772 - log (Fe2 + / Fe(OH)+)

lb. HFeo2- + 2H· = Fe(OH)+ + H20

pH= 11.98 - 1/2 log (Fe(OH)+ I HFeo2-)

2a. Fe(OH) 2 • + H+ = Fe3 + + H20

pH= 2.168 - log [Fe3 • / Fe(OH) 2 +]

2b. Fe(OH)2+ + H+ = Fe(OH) 2 + + H20

pH= 4.983 - log [(Fe(OH) 2 + / Fe(OH)2•]

2c. Fe2(0H)24 + + 2H+ = 2Fe3 + + 2H20

pH= 1.421 - 1/2 log [(Fe3 +) 2 / Fe2(0H)24 +]

2d. 2Fe(OH)2· + 2H+ = Fe2(0H)24+ + 2H20

pH= 5.745 - 1/2 log [(Fe2(0H)24•) / (Fe(OH)2•) 2 ]

3a. Fe3 • + e- = Fe2 •

Eh= 0.770 - 0.0591 log (Fe2 • / Fe3 +)

3b. Fe(OH) 2 + + H+ + e- = Fe2 • + H2o

Eh= 0.898 - 0.0591pH - log [Fe2 • I Fe(OH) 2 +]

3c. Fe(OH)2+ + 2H+ + e- = Fe2• + 2H20

1

Eh= 1_193 - 0_1182pH - log [Fe2 • / Fe(OH)2•]

3d. Fe(OH) 2 • + e- = Fe(OH)•

Eh= 0.497 - 0.0591 log [Fe(OH)• / Fe(OH) 2 •]

3e. Fe(OH)2• + H· + e- = Fe(OH)• + H20

Eh= 0.793 - 0.05916pH - 0.05916 log (Fe(OH)• / Fe(OH)2+)

3f. Fe(OH) 2 • +H2o+ e- = HFeo2- + 2H·

Eh= -0.920 + 0.1183pH - 0.05916log [HFeo2- / Fe(OH) 2 ·]

3g. Fe(OH)2· + e- = HFeo2- + H·

Eh= -0.624 + 0.05916pH - 0.05916log [HFeo2- I Fe(OH)2•]

3h. Fe2(0H)24 • + 2e- = 2Fe(OH)•

Eh= 0.453 - 0.0296 log [(Fe(OH)•) 2 / Fe2(0H)24 •]

3i. Fe2(0H)24 • + 2H• + 2e- = 2Fe2 • + 2H20

Eh= 0.854 - 0.05916pH - 0.0296log [(Fe2 •) 2 / Fe2(0H)24 •]

4a. Fe042 - + 7H+ + 4e- = Fe(OH)• + 3H20

Eh= 1.352 - 0.104 pH - 0.0148 log [Fe(OH)· / FeQ42 -1

4b. Fe042 - + SH· + 4e7 = HFeo2- + 2H20

Eh= 1.000 - 0.0739pH - 0.0148log(HFe02- / Fe042 -)

Eh= 1.679 - 0.1576pH - 0.0197 log(Fe3 •/FeQ42 -)

Sb. Fe04 2 - + 7H• + 3e- = Fe(OH) 2 • + 3H20

Eh= 1.636 - 0.1379pH - 0.0197log[Fe(OH) 2 • I FeQ42 -1

Sc. FeQ4 2 - + 6H• + 3e- = Fe(OH)2• + 2H20

Eh= 1.537 - 0.1182pH - 0.0197log(Fe(OH)2•/Fe042 -)

Sd. 2Fe042 - + 14H+ + 6e- = Fe2(0H)24• + 6H20

Eh= 1.650 - 0.1379pH - 0.0098log[Fe2(0H)24 •/(Fe042 -) 2 ]

Two Solid Substances:

2

6a. FeO + 2H+ + 2e- =Fe+ H20

Eh= -0.041 - 0.05916pH

6b. Fe3Q4 + SH+ +Se-= 3Fe + 4H20

Eh = -0.086 - 0.05916pH

6c. Fe203 + 6H+ + 6e- = 2Fe + 3H20

Eh = -0.053 - 0.05916pH

7a. Fe304 + 2H+ + 2e- = 3Fe0 + H20

Eh = -0.212 - 0.05916pH

7b. Fe203 + 2H+ + 2e- = 2Fe0 + H20

Eh = -0.069 - 0.0591pH

7c. 3Fe203 + 2H+ + 2e- = 2Fe3Q4 +

Eh = 0.214 - 0.05916pH

H20

Sa. Fe(OH)2 + 2H+ + 2e- = Fe + 2H20

Eh = -0.063 - 0.05916pH

Sb. FeOOH + 3H+ + 3e- = Fe +

Eh = 0.003 - O.Os'916pH

Sc. Fe(OH)3 + 3H+ + 3e- = Fe

Eh= 0.051 - 0.05916pH

9a. FeOOH + H+ + e- = Fe(OH)2

Eh= 0.136 - 0.05916pH

2H20

+ 3H20

9b. Fe(OH)3 + H+ + e- = Fe(OH)2 + H2o

Eh= 0.281 - 0.05916pH

10a. 3FeOOH + H+ + e- = Fe3Q4 + 2H20

Eh= 0.719 - 0.05916pH

10b. 3Fe(OH)3 + H+ + e- = Fe3Q4 + SH20

Eh= 1.155 - O.G5916pH

Solid and dissolved species:

3

11a. FeO + 2H+ = Fe 2 + + H20

pH = 6 .155 - 1/2 log(Fe2 +)

llb. FeO + H+ = Fe(OH)+

pH = 5.68 - log [Fe(OH)+)J

llc. HFeo2- + H+ = FeO + H20

pH= 18.19 + log (HFeo2-)

2pH = 11.68 - log(Fe2 +)

lle. Fe(OH)2 + H+ = FeOH+ + H20

pH= 5.05 - log[Fe(OH)+J

llf. HFeo2- + H+ = Fe(OH)::z

pH= 18.82 + log (HFeo2-)

12a. Fe(OH)3 + 3H+ = Fe3• + 3H20

3pH = 3.42 - log(Fe3•)

12b. Fe(OH)3 + 2H ... = ,.Fe(OH) 2 • + 2H20

2pH = 1.25 - log[Fe(OH) 2 +J

12c. Fe(OH)3 +H ... = Fe(OH)2+ + H20

pH= -3.75 - log[Fe(OH)2+]

12d. 2Fe(OH)~ + 4H+ = Fe2(0H)2 4 + + 4H20

4pH = 4.00 - log[Fe2(0H)24 +]

13a. FeOOH + 3H+ = Fe3 + + 2H20

3pH = -1.682 - log(Fe3 +)

13b. FeOOH + 2H+ = Fe(OH) 2 + + H:20

2pH = -3.85 - logFe(OH) 2 +

13c. FeOOH + H ... = Fe (OH) 2+

pH = -8.86 - logFe(OH):2+

13d. 2FeOOH + 4H ... = Fe2(0H)2 4+ + 2H20

4

4pH = -6.2190 - log[Fe2(0H)2 4 +]

14a. Fe203 + 6H+ = 2Fe3 + + 3H20

6pH = -3.77 - log[(Fe3 •) 2 J

14b. Fe203 + 4H+ = 2Fe(OH) 2 + + H2o

4pH = -8.10-log[Fe(OH) 2 +] 2

14c. Fe203 +H2o+ 2H+ = 2Fe(OH)2+

2pH = -18.10 - log[Fe(OH) 2 +] 2

14d. Fe203 + 4H+ = Fe2(0H)24 + + H2o

4pH = -6.61 - log [Fe2(0H)2 4·]

15a. Fe2 + + 2e- = Fe

Eh= -0.409 + log(Fe2 •)

15b. FeOH· + H· + 2e- =Fe+ H2o

Eh= -0.213 - 0.0296pH + 0.0296log[Fe(OH)+J

15c. HFeo2- + 3H· + 2e- =Fe+ 2H20

Eh= 0.493 - 0.0887pH + o.u296log(~Feo2-)

15d. Fe3• + 3e- = Fe

Eh= -0.0159 + log(Fe3 +)

Eh= 0.880 - 0.2365 pH - 0.0296log[(Fe2 •) 3 J

16b. Fe3Q4 + SH+ + 2e- = 3Fe(OH)• + H20

Eh= 0.292 -0.148pH - 0.0296log[(Fe(OH)•J 3

16c. Fe3Q4 + 2H20 + 2e- = 3HFeo2- + H·

Eh= -1.82 + 0.0296pH - 0.0296log[(HFeo2-) 3]

16d. Fe203 + 6H• + 2e- = 2Fe2 • + 3H20

Eh= 0.658 - 0.177pH - 0.0296log[(Fe2 •) 2 ]

16e. Fe203 + 4H+ + 2e- = 2Fe(OH)• + H20

Eh= 0.266 - 0.118pH - 0.0296log[(Fe(OH)•) 2 ]

16f. Fe203 + H20 + 2e- = 2HFeo2-

5

Eh= -1.145 - 0.0296log((HFe02-) 2 ]

17a. Fe(OH)3 + 3H+ + e- = Fe2 + + 3H20

Eh= 0.972 - 0.177pH - 0.05916 log(Fe2 +)

17b. Fe(OH)3 + 2H+ + e- = Fe(OH)+ + 2H20

Eh= 0.580 - 0.118pH - 0.05916 log[Fe(OH)+J

17c. Fe(OH)3 + e- = HFeo2- + H20

Eh= -0.832 - 0.05916 log(HFeo2-)

17d. FeOOH + JH+ + e- = Fe2 + + 2H20

Eh= 0.67 - 0.177pH - 0.05916 log(Fe2 +)

17e. FeOOH + 2H+ + e- = Fe(OH)+ + H20

Eh= 0.2697 - 0.1183pH - 0.05916 log[Fe(OH)+)

17f. FeOOH + e- = HFeo2-

Eh = -1.1479 - 0.05916 log(HFeo2-)

18a. 2Fe04 2- + lOH+ + 6e- = Fe203 + 5H20

Eh= 1.716 - 0.098pH + 0.0098 log [(Fe042 -) 2 )

18b. FeQ42 - + 5H+ + ~- = FeOOH + 2H20

Eh= 1.712 - 0.099 pH+ 0.0197 log (Fe042 -)

18c. FeQ42 - + SH+ + 3e- = Fe(OH)3 + H20

Eh= 1.611 - 0.099 pH+ 0.0197 log(Fe042-)

Iron - Sulphur - Water System

19a. Fe2 + + H2S = FeS + 2H+

2pH = 1.106 - log[(Fe2 +) (H2S)aq]

19b. Fe2 + + Hs- = FeS + H+

pH= -5.89 - log [(Fe2 +) (Hs-)J

19c. FeS + 2H20 = HFeo2- + s 2 - + 3H+

3pH = 49.33 + log [HFeo2-)(s2 -)J

6

2pH = 36.41 + log [HFeo2-J

19e. FeS + 2H20 = HFeo2- + H2S + H+

pH= 29.42 + log [(HFe02-)(H2 S<aq>}]

19f. FeS + 2H20 = Fe(OH)2 + Hs- + H+

pH= 17.57 log(Hs-)

19g. FeS + 2H20 = Fe(OH)2 + 5 2 - + 2H+

2pH = 30.49 + log (s2 -)

20a. FeS + 2H+ + 2e- =Fe+ H2S

Eh= -0.376 - 0.05916pH - 0.0296 log[(H2S)<aq>J

20b. FeS + H+ + 2e- =Fe+ Hs-

Eh = -0.583 - 0.0296 pH - 0.0296 log(Hs-)

20c. FeS + 2e- =Fe+ s 2 -

Eh = -0.965 - 0.0296 log(s2 -)

21a. Fe3Q-4 + 3H2S + 2H+ + 2e- = 3FeS + 4H20

Eh = 0.782 - 0. 0)3916 + 0.0296 log [ ( H2S) -.q)

21b. Fe3Q-4 + 3Hs- + 5H+ + 2e- = 3FeS + 4H20

Eh = 1.402 - 0.148 pH + 0.0296 log[HS-) 3 ]

21c. Fe:304 + 3s2 - + 8H+ + 2e- = FeS + 4H20

Eh= 2.548 - 0.236 pH+ 0.0296 log [(s2 -) 3 )

21d. Fe(OH)3 + s 2 - + 3H+ + e- = FeS + 3H20

Eh= 2.084 - 0.177 pH+ 0.05916 log (s2 -)

21e. FeOOH + s2 - + 3H+ + e- = FeS + 2H20

Eh= -0.3857 - 0.1775 pH+ 0.05916 log (s2 -)

21f. Fe203 + 2s2 - + 6H+ + 2e- = 2FeS + 3H20

Eh= 0.5416 - 0.1775pH - 0.0296 log[(s2 -) 2 ]

22a. Fe7Se + 2H+ + 2e- = 7FeS + H2Scaq>

Eh= -0.0919 - 0.05916 pH - 0.02958 log [(H2S) (aq,J

7

22b. Fe,Se + H+ + 2e- = 7FeS + HS-<aq>

Eh= -0.2988 - 0.02958 pH - 0.02958 log [(HS-)]

22c. Fe,5e + 2e- = 7Fe5 + s 2 -

Eh = -0.6808 - 0.02958 log{S2 -)

Eh= 0.1373 - 0.4733pH - 0.02958 log[(Fe2 +) 7 (H2S)e<aq>J

23b. Fe,5a + 14H20 + 2e- = 7HFe02- + 852 - + 21H+

Eh = -10.938+0.6212pH-0.0296 log[(HFeo2-)'(S2-)

0 J

23c. Fe,5e + 14H20 + 2e- = 7Fe(OH)2 + 852- + 14H+

Eh= -6.99 + 0.414 pH - 0.02958 log ((52 -) 0]

23d. 3Fe,5a + 28H20 + 8e- = 7Fe3Q4 + 2452 - +56 H+

Eh= -4.97 + 0.4141 pH - 0.0074 log [(S2 -) 24 ]

24a. 7Fe2 + + 85 + 14e- = Fe,Se

Eh= 0.1454 + 0.0042 log (Fe2 +) 7

Eh= 0.3058 - o.~611 pH+ 0.0009 log [(Fe2 +)'(S042 -) 0 J

25b. 7HFe02- + 8S042 - + 85H+ + 62e- = Fe,Se + 46H20

Eh= 0.511 - 0.0811pH + 0.0010 log [(Fe2 +) 7 (504 2 -) 9 )

Eh= 0.29 - 0.053 pH+ 0.0010 log [(Fe2 +) 7 (HS04-) 8]

25d. 7HFe02- + 8H504- + 77H+ + 62e- = Fe,5e + 46H20

Eh = 0.4959-0.0735pH+0.0010 log[HFeo4-)'(HFeo4·· ~8]

25e. 7Fe(OH)2 + 85042 - + 78H+ + 62e- = Fe,Se + 46H20

Eh= 0.3838 - 0.0744 pH+ 0.0009 log [SQ42 -) 0)

25f. 7Fe(OH) 2 + 8HS04- + 70H+ + 62e- = Fe,5e + 46H20

Eh= 0.3686 - 0.0668 pH+ 0.0010 log [HS04-) 8 )

26a. 7Fe3 0 4 + 245042 - + 248 H+ + 200e- = 3Fe,Se + 124H20

Eh= 0.3460 - 0.0733 pH+ 0.0003 log [504 2-)

24)

8

Eh= 0.3319 - 0.0662 pH+ 0.0003 log [(HS04-) 24 ]

26c. 7Fe(OH)3 + 8S042 - + 85H+ + 69e- = Fe,Se + 53H20

Eh= 0.3734 - 0.0728 + 0.0009 log [S04 2 -) 8 J

26d. 7Fe(OH)3 + 8HS04- + 77H+ + 69e- = Fe,Se + 53H20

Eh= 0.3597 - 0.0660pH + 0.0009 log [HS04-) 8]

26e. 7Fe203 + l6S042 - + 170H+ + 138e- = 2Fe,Se + 85H20

Eh= 0.3416 - 0.07288 pH+ 0.0004 log [(S042 -) 16]

26f. 7Fe203 + 16HS04- + 154H+ + 138e- = 2Fe,Se + 85H20

Eh= 0.32789 - 0.0602 pH+ 0.0004 log [(HS04-) 16]

26g. 7FeOOH + 8S042 - + 85H+ + 69e- = Fe,Sa + 46H20

Eh= 0.3586 - 0.0729 pH+ 0.0008 log [(S042 -) 0]

26h. 7FeOOH + 8HS04- + 77H+ + 69e- = Fe,Sa + 46H20

Eh= 0.3449 - 0.0660 pH+ 0.0008 log [(HS04-) 8]

,,

9

APPENDIX D

FREE ENERGY OF FORMATION, G0 f, OF FERRIC IRON

HYDROXYL AND SULPHATE COMPLEXES AT 25°c.

Ferric Species G0 f, Kcal/ mol.

Robins (18) Naumov et. al. (70)

Fe3 + -1.1 -4.27

Fe(OH) 2 + -54.80 -58.0

Fe(OH) 2 + -106.74 -108.42

Fe(OH) 3 0 -154.79 -161.9

Fe(OH) 4 - -198.39 -201.7

Fe2 (0H) 2 4+ -111.55 -111.63 ( 7 )

Fe3 (0H) 4 5+ -221.46

FeOOH(am) -109.82 ( 7 )

Fe(OH)-:,(am.) -166.48 ...,

FeS04 + -184.68 ( 7)

Fe(S04 ) 2 - -364.36 ( 7)

Fe2 (S04 )3 -536.04 ( 7)

H2o (liqui'd) -56.688

H+ 0.0

--C1--

APPENDIX E

EQUILIBRIUM DATA FOR IRON (III)-WATER SYSTEM AT 25°C

1. Fe(OH) 2 + + ff+ = Fe3 + + H:20

pH = 2.19 log [Fe3 +] I [Fe(OH) 2 +]

pH = 2.16 log [Fe3 +] I [Fe(OH) 2+]

2. Fe(OH)2+ + 2H+ = Fe3 + + 2H::z0

pH = 2.835 - 1/2 log [Fe3 +J I [Fe(OH):2+]

pH = 3.367 1/2 log [Fe3 +] I [Fe(OH)2+]

3. Fe(OH):::;, 0 + 3H+ = Fe3 + + 3H20

pH= 4.00 - 1/3 log [Fe3 +] / [Fe(OH):::;,0]

4. Fe(OH)4- + 4H+ = Fe3 + + 4H20

pH= 5.399 - 1/4 log [Fe3 +] I [Fe(OH)4-J

5. Fe2(0H}:;;i:4 + + 2H+ = 2Fe3 + + 2H20

pH= 1.475 - 1/2 ;l,og [Fe3+]2 / [Fe2(0H)24 +]

6. Fe:::;,(OH).,._5 + + 4H+ = 3Fe 3 + + 4H20

pH= 1.575 - 1/4 log [Fe3+J::. I Fe:::;, (OH).,.. ,:5+

7. FeOOH<am> + 3H+ = Fe3 + + 2H20

pH= 1.136 - 1/3 log [fe3 +]

8. Fe(OH)2+ + H+ = Fe(OH) 2 + + H20

pH= 3.48 - log [Fe(OH) 2 +] / [Fe(OH)2+]

9. Fe(OH):::, 0 + 2H+ = Fe(OH) 2 + + 2H20

pH= 4.90 - 1/2 log [Fe(OH) 2 +] / [Fe(OH):::;, 0]

10. Fe(OH)4- + 3H+ = Fe(OH) 2 + + 3H20

pH= 6.469 - 1/3 log [Fe(OH) 2 +] / (Fe(OH)4-J

11. Fe2(0H)24 + = 2Fe(OH) 2 +

-1.429 = log [Fe(OH) 2 +] 2 / [Fe2(0H)24 +]

1

...

12. Fe3(0H)4 5+ + H+ = 3Fe(OH) 2 + + H20

pH = -0.273 - log [Fe(OH) 2 +] 3 I Fe::;,(OH)4 3+

13. FeOOHcam> + 2H+ = Fe(OH) 2 + + H20

pH = 0.611 - 1/2 log [Fe(OH) 2 +]

14. Fe(OH)::3° + H+ = Fe(OH):2+ + H20

pH = 6.332 - log [Fe(OH)2+] I [Fe(OH)3<:>)

15. Fe(OH)4- + 2H+ = Fe(OH):z+ + 2H20

pH= 7.96 - 1/2 log [Fe(OH)2+] / [Fe(OH)4-J

16. 2Fe(OH)2+ + 2H+ = Fe2(0H)24 + + 2H20

pH= 4.195 - 1/2 log [Fe2(0H)24 +] / [Fe(OH)2+] 2

17. 3Fe(OH)2+ + 2H+ = Fe3(QH)43 + + 2H20

pH= 5.357 - 1/2 log [Fe3(0H)43 +] / [Fe(OH)2+] 3

18. FeOOHcam> + H+ = Fe(OH)2+

pH= -2.25 - log [Fe(OH)2+]

19. Fe(OH)4- +ff+= Fe(OH)3° + H20

pH= 9.594 - log~[Fe(OH)3°) / [Fe(OH)4-J

20. 2Fe(OH)3° + 4H+ = Fe2(0H)24 + + 4H20

pH= +5.263 - 1/4 log[Fe2(0H)2 4 +] / [Fe(OH)::3°) 2

21. 3Fe(OH)3° + 5H+ = Fe3(QH)43+ + 5H20

pH= 5.942 - 1/5 log[Fe3(0H)4 5 +] / [F~~OH)::,,=] 3

22. FeOOHcam> + H20 = Fe(OH)::3°

-8.589 = log[Fe(OH)3°]

23. 2Fe(OH)4- + 6H+ = Fe2(0H)24 + + 6H20

pH= 6.706 - 1/6 log [Fe2(0H)2 4 +] / [Fe(OH)4-J 2

24. 3Fe(OH)4- + 8H+ = Fe3(QH)4 3 + + 8H20

pH= 7.31 - 1/8 log[Fe3(0H)45 +] / [Fe(OH)4-) 3

25. Fe(OH) 4 - + H· = FeOOHcam> + 2H20

pH= 18.183 + log[Fe(OH)4-]

2

26. 2Fe3(0H)45 + + 2H+ = 3Fe2(0H)24 + + 2H20

pH= 1.871 - 1/2 log [Fe2(0H)24 +] 3 / [Fe3(0H)45 +] 2

27. 2FeOOH<am> + 4H+ = Fe2(0H)2 4 + + 2H20

pH= 0.969 - 1/4 log [Fe2(0H) 2 4 +]

28. 3FeOOH<-m> + SH+ = Fe3(0H)4 5 + + 2H20

pH= 0.788 - 1/5 log[Fe3(0H) 45 +]

0000000000

* Equations using Naumov's data (70)

3

APPENDIX F

EXPERIMENTAL DATA CALCULATIONS

A representation of the ON and OFF sequence times of

the electrolytic production of oxygen may be represented

as:

0 1 3 4 5 .. I

!--------------!------!----------!------!------!-----)

OFF ON OFF ON OFF ON

The ON time was fixed at 140 seconds, but it could be

varied according to the consumption of oxygen in the system.

RATE OF UPTAKE OF OXYGEN

(a) ON Time: I - (I

(b) OFF Time: I - (I

1), FOR I= 2 to N step 2

1), FOR I= 1 to N - 1 Step 2

(c) Cycle Time: ON Time+ OFF Time

(d) Real Time: N - O (seconds)

I= actual reading of ON and OFF times of the

electrolytic cell.

N = Total number of readings during the experiment.

RATE OF UPTAKE OF OXYGEN= ON TIME/ ON TIME+ OFF TIME

= ON TIME/ CYCLE TIME

Thus, the rate of consumption of oxygen is plotted against

the real time, the rate of consumption of oxygen during the

aqueous oxidation of pyrrhotite could be obtained. A program

written in Basic Language showing the procedure to calculate

this rate is reproduced below.

PROGRAM

wo qE~ r11ia11titiriiriiitttistiiiit1iiiiia111is UO ?.Ef i LEACH. BAS ¥ 120 REn HHHUUUtUUUHHUHUHtUHUHt 130 c~s : ?RHi: : PRINT : PRINT 140 PRINT TAB(20)~P~JGRA~ME LEACHIN9 QF ?YRRHGTITE!' 150 ?R!N~ : ?RI~T : PRINT

200 FnirH "ANJ L:/\CH RATES IN DECIMAL FORriAT, THIS DATA FILE CAN BE PLOTTED'' 110 ?RlNT ·AS IS, QR IT CAN BE SMOOTHED £YTHE iiOvING AVERAGE METHOD USIN5 •; 120 PRINT 'LEACH, BAS;' 130 PRINT : PRINT : PRINT !40 PRINT "ENTER NAME OF DATA FILE CONTAINING ON i!MES AND OFF TIMES" :50 PRINT • IN DECIMAL FORnAT "; : INPUT I$ :60 CLOSE 2 : OPEN '1 ",2, I$+" .D65"

'.BO N:-1 90 IF EDFl2J THEN 310 00 N:Ntl : INPUT #2,D(N) : SOTO 290 10 PRINT : PRINT "ENTER NAME OF DATA FILE TO CONTAIN ABSOLUTE TIMES"; 20 PRINT "AND LEACH RATES"; : INPUTS$_ 30 CLCSE 3 : OPEN "0",3,S$t",D65" 40 PR!!4T : PRINT "HHER NAME OF DATA FILE TO CONTAIN EXTENT OF OX IDATHJN" !5 PRINT "OF PYRRHOTITE 1

; : INPUT S$ 10 CLOSE 4 : OPEN •o• ,4,S$t" .D65" 10 Q=O : DS=O : 06=0 : A=O lO rRINT : PRINT "DATA FILE •u•,D65 contains"D(O)"readings• : PRINT !O PRINT "How itany readings rio you want to process •; : INPUT N !1 CLS 12 FU=" 14 F2$="·

ABSOLUTE TINE

ON TINE

CYCLE TIHE

'6 F3$= 1 (minutes) (seconds) (seconds) ¼

LEACH RATE

8 F4$="11111 HHI.I Htl,1 IHI.I i,HHI 9 PRINT F1$: PRINT.F2$ : PRINT F3$ : PRINT 0 FOR I=2 TON STEP 2 0 REH CALCiJLAT:uN OF "REAL Til'iE" 5 IF Di!){D(i-2) THEN G=G+lOOOOO: IF D(I-2)=D{O) THEN G=O 0 Dl=(D(I}+A+QJ/60 3 N=INT{D1t10+::.;i10 0 REM CALCULATi'JN OF "ON THIE" J D2=D{!I-D(!-!):!F D(IJ<D(I-1) THEN D2=D(Il+100000!-D{I-11 ) D5=D5+D2 i REM C~LCULATIC'i OF "CYCi_E TIME" D3 ) F D(i-2):D(:;) •-iEN D(;-2)=(;

EHENTION OF" PYRRHDT ITE • OXIDATION, x• · H.HH"

) D3=D(I)-D(l-2;: IF !l(:) < D(I-2} THEN D3=D(l)+i0)00ti-D(I-2j I REi~ CALCULAT I QN OF nLEACH RATE '1

( D2iD:)

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APPENDIX G

CALCULATION OF OXYGEN CONSUMPTION AND PRODUCTION

The consumption of oxygen to be determined is based on

Equations (VI.1.5) and (VI.1.12) previously discussed and

thirty grams of sample of mineral.

Equation (VI.1.5): FeS + 0.5H20 + 0.7502 ----->

Equation (VI.1.12):

FeOOH +

Stoichiometric Monoclinic Oxygen

Pyrrhotite Pyrrhotite

Molecular Weight, grams: 87.907 647.409

Oxygen consumed, gr.: .75*32*30/55.847 5.25*32*30/647.409

8.1905 7.785

Average: 7.9877 grams of oxygen.

If this average value of both results is considered as

the theoretical amount of oxygen consumed by thirty grams of

the pyrrhotite mineral sample, the incurred error could be

minimal (2.48 per cent).

ELECTROLYTIC PRODUCTION OF OXYGEN

These calculations are based on Faraday's Laws

[M=(A/F)*I*t] and determined for experiments No.1 and 2 from

the second equipment design.

EXPERIMENT 1 EXPERIMENT r, .::...

Total Number of Readings: 1,914 7,580

Total Number of ON times: 957 3,790

ON Time Value, Sec.: 110 110

Current, mA. : 150 150

From Faraday's Law, it is known that in order to pro-

duce one mole of molecular oxygen from water it is necessary

to pass through the electrolytic cell four faradays. Thus,

M1 =(0.15coul/sec*110*957sec*32 g.02 *FJ/(4F*96500 coul]

M1 = 1.309 g of oxygen

M2 =(.15coul/sec*110*3790 sec *32 g.02 *FJ/(4F*96500 coul]

M2 = 5.1842 g. of oxygen

Thus the percentage of conversion of thirty grams of the

pyrrhotite mineral sample will be:

Experiment No.l: 1.309 * 100/7.9877 = 16.39 %

Experiment No.2: 5.1842 * 100/7.9877 = 64.90 %

RATE OF OXYGEN CONSUMED, g/hr.

Experiment No.1 = 1.309 g 02/237 hours

= 0.0055 g/min.

Experiment No.2 = 5.1842 g 02/241 hours

= 0.0215 g/min.

Note: ON time for experiments 14-20 was 480 mA.

APPENDIX H

PROCEDURE FOLLOWED TO DETERMINE IRON, ARSENIC AND ZINC

DURING THE CONTINUOUS BACTERIAL OXIDATION OF PYRITE

Samples of slurry from each stirred reactor and the

thickener(s) (overflow and underflow) were taken for analy

sis every Monday, Wednesday and Friday.

From each slurry sample, 1ml aliquot was taken into a

test tube and named "total" sample. Another 1 ml aliquot was

taken into another test tube and named "soluble" sample. To

the "total" sample, 1ml of SN HCl was added, stirred manual

ly and left for about 30 minutes. After this time, 8 ml of

pH 2 distilled water was added (dilution 10). The "soluble"

sample was diluted with 9 ml of pH 2 distilled water direct

ly. Both set of samples, "total" and "soluble" were centri

fuged at 1500 r.p.m for 20 minutes. The solution after

centrifugation was diluted further to 100 and/or 1000 in

order to read in the AAS.

It is believed that the 1ml of SN HCl added to the

"total" sample only redissolved the precipitation products

without attacking the mineral.

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6 Aurotech N. L.

1 June 1989

MrC Fabian 18 Broadhurst Crescent BATEMAN WA 6155

DearCes

L•,d floo~. 34 C(1l11, Sireel. \Vc;1 P,·,th. Wesiern 1\u,1rali,, 6()():, GP() Rox M9+1. Perth 6001

CONFIDENTIALITY COVENANT

Telephone lnt: +6193214400 Fax lnt: +619321 4646

We write in response to your request seeking a release from the Confidentiality Covenant you entered into with Aurotech N.L.

As Aurotech is changing its emphasis from bacterial leaching research and development to gold exploration/production we believe it to be fair and reasonable to release you from the Confidentiality Covenant.

We wish you well in your endeavours and thank you for the service given to Aurotech during the period of your employment.

Yours sincerely

BWRIDGEWAY Chairman/Managing Director

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