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University Microfilms International 300 North Zeeb Road Ann Arbor, Michigan 48106 USA St. John's Road. Tyler’s Green High Wycombe, Bucks, Engiand HP10 8HR 77- 18,655

TH(M, Gary Carlisle, 1943- THE CHARACTERIZATION OF THE METAL SULFIDE SULFUR DIOXIDE REACTION IN AQUEOUS MEDIA. The American IMiversity, Ph.D., 1977 Oianistry, physical

Xerox University Microfilms, Ann Arbor, Michigan 48ioe

© 1977

GARY CARLISLE THOM ALL RIGHTS RESERVED THE CHARACTERIZATION OF THE METAL SULFIDE

SULFUR DIOXIDE REACTION

IN AQUEOUS MEDIA

by

Gary Carlisle Thom

Submitted to the

Faculty of the College of Arts and Sciences

of The American University

in Partial Fulfillment of

the Requirements for the Degree

of

Doctor of Philosophy

in

Chemistry

Signatures of Committee;

^ / V ' Cha±rman:^^î^ li S

Dean of the College

Dat‘='/ “ 1977

The American University Washington, D.C. 20016

THE AMERICAN UNIVERSITY LIBRARY 535-0 Man, being the servant and interpreter of nature, can do and understand as much and so much as he has observed in fact or in thought of the course of nature: beyond this he neither knows anything nor can do anything.

Francis Bacon "Novum Organum"

iii ACKNOWLEDGEMENT

This work could not have been completed without the guidance and

generous assistance given to me by Drs. Paul F. Waters, Robert T. Foley,

and Albert F. Hadermann. I am most grateful to my mother, Ethel B. Thom, who spent many long hours typing the manuscript. Finally, I am indebted

to The American University and its Department of Chemistry which has

provided me with both my graduate and undergraduate training. This

training, in which I was able to freely explore my academic interests,

has laid the groundwork for a productive and rewarding career.

iv TABLE OF CONTENTS

Page

ABSTRACT

DEDICATION...... H i

ACKNOWLEDGEMENT...... iv

LIST OF TABLES...... vii

LIST OF FIGURES...... viii

CHAPTER

I. INTRODUCTION...... 1 A. Current Status of Flue Gas Desulfurization Processes.. 3 B. The FGD Sludge Disposal Problem...... 4 C. Research Objectives...... 10

II. BACKGROUND...... 13 A. The Metal Sulfide-Sulfur Dioxide Reaction in Aqueous Media...... 13 B. The Oxidation of Metal Sulfides in Hydrometallurgy.... 18 C. The Hydrogen Sulfide-Sulfur Dioxide Reaction in Aqueous Media...... 29 D. The Chemistry of Sulfur...... 34 E. Analyses for Thiosulfate and Sulfane Sulfonates...... 51 F. The Kinetics and Mechanism of the Acid Decomposition of Thiosulfate...... 55

III. EXPERIMENTAL...... 59 A. Materials...... 59 B. Apparatus...... 60 C. Procedure...... 62 D. Methods and Instrumentation...... 65 E. Kinetic Studies...... 66 F. Data Analysis...... 67 TABLE OF CONTENTS (CONT'D)

Page

IV. RESULTS AND DISCUSSION...... 69 A. Characterization of the Ferrous Sulfide-Sulfur Dioxide Reaction...... 70 B. Characterization of the Sulfide Ore-Sulfur Dioxide Reaction...... 101 C. Comparison of Sulfide Reactivities...... 115 D. Kinetics of the Acid Decomposition of Thiosulfate.... 124 E. Mechanism of the Oxidative Dissolution of Sulfide Ores...... 136

V. CONCLUSIONS...... 147

VI. FUTURE RESEARCH...... 149

APPENDIXES

1. DESCRIPTION OF FIRST AND SECOND GENERATION FLUE GAS DESULFURIZATION PROCESSES...... 151

2. HYDROMETALLURGICAL LEACHING REACTIONS...... 159

3. KINETIC DATA...... 169

REFERENCES...... 175

vi LIST OF TABLES

Page

I. Domestic Reserves of Sulfide Ores...... 11

II. Sulfur Oxide Emissions from Non-Ferrous Smelters...... 19

III. Sulfur Allotropes...... 36

IV. Sulfur Molecular Species...... 38

V. UV Spectra of the Sulfite, Thiosulfate, and Sulfane Disulfonate Anions...... 54

VI. Summary of Results for Base-line FeS Experiments...... 80

VII. Summary of Results for FeS-Reactant Variation Experiments...... 97

VIII. Qualitative Assessment of the Differences Between the Reactant Variation and Base-line Experiments...... 99

IX. Summary of Results for Sulfide Ore Experiments...... 109

X. Qualitative Assessment of the Differences Between the Sulfide Ore and Base-line Experiments...... Ill

XI. Summary of Effects of Ore Surface Area on the Reaction Rates at 75 °C and the Adsorption of SOg at 50 “C 116

XII. Summary of Kinetic Results for the Decomposition of Thiosulfate at 40.0 °C...... 135

vii LIST OF FIGURES

Page

1. Pourbaix diagram of the Sg/S^^/S^^/HgO system...... 23

2 . Pourbaix diagram of the Sg/S2 0 3 ^ /S^ /S^^/H^O system. 24

3. Pourbaix diagram of the Cu2 S/CuS/Sg/Cu/H2 0 system.... 25

4. Mole fraction of S , 2 < n < 8 , in a saturated vapor above the liquid phase...... 39

5. Viscosity-temperature curve for liquid sulfur...... 41

6. Sulfite reduction apparatus...... 61

7. Heating curve for a reaction temperature of 70.8 °C. 64

8. FeS-S0 2 reaction characterization at 31.1 °C experiment 95 ...... 73

9. FeS-SOg reaction characterization at 40.2 °C experiment 96 ...... 74

10. FeS-S0 2 reaction characterization at 58.3 °C experiment 94 ...... 75

11. FeS-S0 2 reaction characterization at 69.0 “C experiment 90 ...... 76

12. FeS-S0 2 reaction characterization at 70.8 °C experiment 8 8 . 77

13. FeS-S0 2 reaction characterization at 71.9 “C experiment 89 ...... 78

14. FeS-SOg reaction characterization at 82.0 °C experiment 91 ...... 79

15. Composite of the SO2 pressure characterization curves for the FeS-S0 2 reaction at five temperatures..... 81

viii LIST OF FIGURES (CONT’D)

Page

16. Relationship between -dP/dt and ^[SgOg "]/dt for the base-line FeS-S0 2 experiments (#), r^ = 0.98; points for the reactant variation experiments (a ), and the sulfide ore experiments (■) are identified by experiment number...... 83

17. Relationship between -dP/dt and d[Fe^*l/dt for the base-line FeS-SOg experiments (•), r = 0.93; points for the reactant variation experiments (a ), and the sulfide ore experiments (■) are identified by experiment number...... 84

18. Relationship between d[Fe ]/dt and d[S 2 Û 3 ]/dt for the base-line FeS-S0 2 experiments (#), r^ = 0.92; points for the reactant variation experiments (a ), and the sulfide ore experiments (■) are identified by experiment number...... 85

19. Arrhenius plot for the -dP/dt data from the base-line FeS-SOo experiments; In k = -16,783/RT + 22.77, E„îa = 16.783 kcal mole“^, r^ = 0.98...... 87

20. Arrhenius plot for the d[S2 0 g^"]/dt data from the base-line FeS-S0 2 experiments; In k = -15,569/RT + 17.56, Eg = 15.569 kcal mole"^ r^ = 0.97...... 8 8

21. Arrhenius plot for the d[Fe. 2 + ]/dt data from the base-line FeS-S0 2 experiments ; In k = -12,182/RT + 12.15, Eg =12.182 kcal mole"^ r^ = 0.99...... 89

22. Potential energy diagram for the FeS-SÜ2 reaction showing the activation energies for the dissolution of FeS (El), and for the formation of thiosulfate (E2 ) ...... 91

23. Relationship between the pressure drop, -AP, and the maximum thiosulfate concentration for the base-line FeS-S0 2 experiments (•), r^ = 0.99; points for the reactant variation experiments (a ), and the sulfide ore experiments (■) are identified by experiment number...... 92

24. Effect of using 0.1 mole H 2 SO 4 instead of 0.2 mole HCl in the FeS-S0 2 reaction at 70.9 °C, experiment 92...... 94

ix LIST OF FIGURES (CONT'D)

Page

25. Effect of using 0,05 mole instead of 0.10 mole FeS in the FeS^SOg reaction at 90,2 °C, experiment 99 ...... 95

26. Effect of using calciner feed in place of Na^SOg in the FeS^S0 2 reaction — r a simulated . sulfite reduction process at 69.6 "C, experiment 106...... 96

27. PbS-S0 2 reaction characterization at 73.2 °C, experiment 101...... 103

28. FeS (pyrrhotite)-S0 2 reaction characterization at 77.4 °C, experiment 104...... 104

29. ZnS-S0 2 reaction characterization at 75.0 ®C, experiment 98 ...... 105

30. CuFeS2 ~S0 2 reaction characterization at 76.5 °C, experiment 105...... 106

31. FeS2 -S 0 2 reaction characterization at 75.3 °C, experiment 107...... 107

32. CU2 S-SO2 reaction characterization at 75.0 “C, experiment 103...... 108

33. SO2 pressure vs. temperature characterization for the base-line FeS, reactant variation, and FeS (pyrrhotite) experiments showing the adsorption of SO2 on the ore surface prior to the reaction...... 119

34. SO2 pressure V£. temperature characterization for the PbS experiment showing the adsorption of SO 2 on the ore surface prior to the reaction...... 1 2 0

35. SO2 pressure vs. temperature characterization for the CuFeS2 and ZnS experiments showing the adsorption of SO2 on the ore surface prior to the reaction...... 1 2 1

36. SO2 pressure vs. temperature characterization for the FeS 2 and Cu^S experiments showing the adsorption of SO2 on the ore surface prior to the reaction...... 1 2 2

37. Decomposition of thiosulfate in the FeS/Na2 S0 3 /HCl reaction product supernatant from experiment 96 at 40 °C; k = 1.067 liter^ mole ^ hr”^, n = 2.38, r^ = 0 . 9 5 ...... 126

X LIST OF FIGURES (CONT'D)

Page

38. Decomposition of thiosulfate in the FeS/Na2 S0 3 /HCl reaction product supernatant from experiment 94 at 40 ®C; k = 1.012 liter^ mole"^ hr“^, n = 1.57, r^ = 0.76...... 127

39. Decomposition of thiosulfate in the FeS/NagSOg/HCl reaction product supernatant from experiment 90 at 40 °C; k = 1.006 liter^ mole~^ hr~^, n = 1.22, r^ = 0 . 8 8 ...... 128

40. Decomposition of thiosulfate in the FeS/Na^SOg/HCl reaction product supernatant from experiment 89 at 40 “C; k = 1.008 liter^ mole~^ hr”^, n = 1.43, r^ = 0.94...... 129

41. Decomposition of thiosulfate in the FeS/Na^SOg/HCl reaction product supernatant from experiment 91 at 40 °C; k = 1.010 liter^ mole“^ hr“^, n = 1.46, r^ = 0.68...... 130

42. Decomposition of thiosulfate in the FeS/Na2 S0 3 /H2 S0 ^ reaction product supernatant from experiment 92 at 40 °C; k = 1.008 liter^ mole~^ hr” ^, n = 1.39, r^ = .0.92...... 131

43. Decomposition of thiosulfate in the FeS/calciner feed/HCl reaction product supernatant from experiment 106 at 40 “C; k = 1.013 liter^ mole~^ hr“^, n = 1.45, r^ = 0.96...... 132

44. Decomposition of thiosulfate in the PbS/Na^SOg/HCl reaction product supernatant from experiment 1 0 1 at 40 ®C; k = 1.008 liter^ mole~^ hr~^, n = 1.58 r^ = 0.87...... 133

45. Decomposition of thiosulfate in the FeS (pyrrhotite)/ Na^SOq/HCl reaction product supernatant from experiment 104 at 40 °C; k = 1.003 liter^ mole ^ hr n = 1.03, r^ = 0.82...... 134

46. Dissociation of thiosulfuric and the sulfane sulfonic acids on the surface of colloidal Sg...... 139

47. Chemical model for the oxidative dissolution of metal sulfides in acid and alkaline solution showing the acid decomposition of thiosulfate and the surface deposition of Sg...... 141

xi CHAPTER I

INTRODUCTION

The uncontrolled release of sulfur dioxide into the atmosphere by fossil-fueled electric generating plants has created one of the major air pollution problems facing our energy-intensive society. This problem will continue to be of major concern since our demand for electricity is doubling approximately every ten years. Currently, over

50 per cent of the nation's 172,000 megawatt (MW) electrical generating capacity is provided by 450 coal-fired steam electric generating plants which consume 85 per cent of the 390 million tons of combustion coal produced annually.^ In 1970, coal-fired power plants accounted for over 60 per cent of the 34.3 million tons of sulfur dioxide released into the atmosphere. As a result of their large contribution to sulfur dioxide emissions, these plants have been the subject of considerable emission control actions by the federal government.

The Clean Air Act Amendments of 1970 established the Environmental

Protection Agency (EPA) and gave it the responsibility for promulgating ambient air quality standards and establishing New Source Performance

Standards (NSPS) for a number of industrial activities, including steam electric generating plants. Accordingly, in 1971 EPA issued national ambient air quality standards for sulfur dioxide and established NSPS for power plants. Under these regulations, sulfur dioxide emissions from new coal-fired power plants was limited to 1.2 lb/million BTU. To 2 ensure the achievement of the ambient air quality standard for sulfur dioxide, the necessary reductions in power plant emissions were to be attained by May 31, 1975.

Initially, reductions in sulfur dioxide emissions from coal-fired power plants relied on one or more of the following options; the use of low-sulfur coal, the substitution of oil or natural gas, the modernization of equipment, the construction of new power plants outside urban areas, and the installation of tall stacks to disperse pollutants more effectively. During the period from 1970 to 1974, these measures led to a reduction in total sulfur dioxide emissions from 34.3 to 31.4 million tons and a 25 per cent reduction in urban levels of sulfur dioxide.3 However, further progress has been threatened by the increasingly scarce supplies of low-sulfur coal and the cleaner burning oil and natural gas. Furthermore, those plants using tall stacks or other "supplementary control strategies", and those relocated in rural areas still failed to meet the emission standards set forth in the NSPS.

To fully meet the NSPS requirements it has been estimated that some form of flue gas desulfurization (FGD) would be required on approximately one-fourth of the 450 coal-fired power plants operating in 1974. By 1980, when the annual consumption of coal by utilities is expected to increase by 60 per cent to 625 million tons, the number of plants requiring FGD will increase substantially. This increase is mainly the result of the Energy Supply and Coordination Act of 1974.

This act, which was passed as a result of the 1973 Arab oil embargo, prohibits any power plant with the capability to burn coal from burning natural gas or fuel oil. A. Current Status of Flue Gas Desulfurization Processes

Flue gas desulfurization (FGD) involves removing the SOg from

combustion gases either by adsorbing the SOg in an alkaline solution or

slurry thereby producing a sulfite salt, i.e., wet scrubbing; or

adsorbing the SOg on a suitable solid adsorbent, i.e., dry adsorption.

Depending on what is done with the spent adsorbing medium, a. flue gas

desulfurization (FGD) process may be either regenerable or non-regenerable. Regenerable, or closed loop processes, regenerate the

adsorbing medium and recover the SO2 for further processing. Thus, the

adsorbing medium may be continuously recycled and the SO2 may be

converted into a useful product such as sulfuric acid, sulfur, or

gypsum. Non-regenerable, or once-through processes, require a

continuous supply of new adsorbent and produce a waste sludge by-product

(e.g., CaSOg) which must be disposed of in an environmentally acceptable

manner.

The active FGD research and development program undertaken by the

EPA during the past decade has resulted in the development of more than

50 FGD p r o c e s s e s . After extensive testing and evaluation, ten

processes, which have SOg removal efficiencies from 70 to 90 per cent,

were scaled up for commercial application. Although FGD systems have

been installed or are planned in 109 power plants having a generating

capacity of 42,128 MW,^^ by late 1976 only 8,577 MW FGD capacity were

operating. The operational FGD capacity will increase to approximately

24,416 MW by 1980, but this is still only about one-fourth of the

estimated 90,000 MW needed to meet the ambient air quality standard for

SO2 . Thus, much work remains to be done. 4

The ten commercial FGD processes have been the subject of several special symposia^where a detailed status and process description of each has been discussed. For the purposes of exemplifying FGD process chemistry a description of each of these processes is given in

Appendix 1. These processes are listed below with the extent to which each is currently used in the 8,577 MW capacity:

Non-regenerable (7569 MW - 8 8 per cent)

• Lime and/or limestone scrubbing,i.e., three processes (7549 MW)

• Double alkali scrubbing^* (20 MW)

Regenerable (1008 MW - 12 per cent)

• Sodium carbonate scrubbing^^ (375 MW)

• Magnesium oxide scrubbing^®" (365 MW)

• Wellman-Lord p r o c e s s ^ (115 MW)

• Cat-ox process^** (110 MW)

• Dilute acid scrubbing^° (23 MW)

• Activated carbon adsorption^^ (20 MW)

B. The FGD Sludge Disposal Problem

Of the ten processes mentioned above, the non-regenerable lime/limestone processes provide a large majority of the FGD capacity

( 8 8 per cent). By 1980, they will provide FGD to 85 per cent or 20,677

MW of the projected total of 24,416 MW. The greatest problem with these processes is the disposal of enormous volumes of sludge. For a power plant burning high sulfur coal (3 to 5 weight per cent sulfur), eight

to nine cubic feet of sludge is generated per ton of coal burned. This

translates into 80,000 cubic feet of sludge per day for a 1000 MW power p l a n t . Thus in 1980, when 20,677 MW will have sludge-generating FGD, as much as 1 . 6 million cubic feet of sludge will have to be disposed of per day. And for the 90,000 MW FGD capacity required to meet the SO2 ambient air quality standard, about 6 . 1 million cubic feet per day will be generated unless a larger percentage of FGD is performed by regenerable processes. On a weight basis, EPA estimates that by 1980 approximately 131 million tons of sludge will be generated per annum.^ ^

The magnitude of the sludge problem may be illustrated by the disposal plans for a new power plant at Shippingport, Pa., where a valley is being dammed and will, over the next 25 years, be filled to a depth of 400 feet with sludge pumped through a special seven-mile-long pipeline.

Although some commercial usage of sludges is feasible from a technical standpoint, the potential outlet is so small that the vast majority of sludge will not be marketed.As a result, disposal of sludges by ponding and landfilling appears to be the only important near-term alternative. Disposing of large quantities of sludge in this manner requires that their environmental impact be considered. The composition of sludge is somewhat variable but typical sludge from a limestone scrubbing process consists of an aqueous slurry of CaS0 3 "l/ 2 H 2 O,

CaS0 4 *2 H 2 0 , CaC0 3 and fly ash (mainly Si0 2 , AI 2 O 3 , and Fe 2 0 3 , with small amounts of alkaline earth and heavy metals, and trace amounts of heavy metal oxides). Since some of these constituents are highly toxic and sludge is high in total dissolved solids (TDS), adequate steps must be taken to prevent the pollution of ground and surface waters in both ponding and landfill operations. 6

Ponding of untreated waste sludge requires the installation of an impermeable pond liner such as clay, plastic, or rubber. Although ponding provides a short-term solution to the waste disposal problem, pond reclamation has not been demonstrated to be cost effective.

Furthermore, little is known about the possible formation of secondary pollutants due to the long-term biological action on ponded sludge.

Landfilling requires that the sludge be dewatered to greater than

70 per cent solids and/or undergo chemical fixation to give the sludge sufficient physical stability to be handled by earth-moving equipment.

Dewatering of untreated sludge, which contains less than 50 per cent solids, may be accomplished by mechanical filtration, addition of dry solids (i.e., fly ash), addition of flocculating agents, conversion of the sludge to gypsum, or a combination of these methods.Following dewatering, chemical fixation may be used to further increase the stability of the sludge and to chemically and/or physically bind up the soluble constituents. Commercial processes for chemical fixation are available^®’ but, of course, they increase the overall cost of FGD.

Other sludge disposal methods currently under consideration include disposal in underground mines, surface mines, oceans, and deep-well injection.^®

The use of sludges in various commercial products has also been investigated. As a result of these investigations, technology has been developed which can utilize sludge in a variety of products and applications including mineral wool, bricks, sintered concrete products, soil amendment, mineral recovery, road base materials, parking lot materials, artificial aggregate, lightweight aggregate, and aerated 7 concrete.^** However, due to the inability of these new technologies to economically compete with existing ones, little or no utilization of sludges in these applications is expected.

The sludge disposal problem would seem to provide a strong incentive for power plants to use regenerable scrubbing processes which produce a potentially marketable by-product. However, this has not been the case for several reasons: 1 ) the reluctance of the power industry to enter into chemical marketing; 2 ) the lack of a demand for by-products, particularly in remote areas where many power plants are located; and 3) the problem of increasingly short supplies of petroleum feedstocks required to generate hydrogen sulfide for those regenerable

FGD processes which produce elemental sulfur as a by-product.

Historically, the generation and distribution of electricity has been the only concern of the electric utility industry. Thus, industry executives are naturally hesitant to enter into the chemical marketing of FGD by-products, particularly when dealing with low-demand goods.

However, in the future the power industry will, either through regulatory and/or economic pressures, be forced into FGD by-product marketing. The alternative FGD by-products include elemental sulfur, sulfuric acid, gypsum, sodium sulfate, ammonium sulfate, and liquid sulfur dioxide. Of these, only sulfuric acid and elemental sulfur have significant market potential. Elemental sulfur is the most environmentally and economically desirable product since it has low solubility, substantially reduced storage requirements, low shipping

costs, and wide applicability as a raw material. Currently, the market

for sulfur is demand-limited and, as a result, FGD derived sulfur cannot

compete with the sulfur derived from the Frash and Claus processes. 8

However, research is underway which will lead to the use of sulfur in a variety of products including fertilizers, foams, coatings, construction materials, pavements, and batteries.^® Once these products reach

commercial production the demand for sulfur will increase substantially

and the sulfur derived from FGD can be more readily marketed.

Because elemental sulfur is the most desirable FGD by-product, most regenerable processes have been designed to recover sulfur. A

recent status report on these processes^ found that of the 14 currently

in operation or under development, 7 produced elemental sulfur as the

final product, 6 produced a concentrated (85 per cent) stream of SO2 which required further processing to obtain elemental sulfur or H 2 SO4 ,

and one recovered gypsum. However, 9 of the 14 processes require a

reducing gas (H2 S or CO and H 2 ) to produce the final elemental sulfur product. Associated with the generation of these reducing gases are

two very critical problems. First, the severe shortages of natural gas

and other petroleum-based feedstocks used to produce the gases have

resulted in the hydrocarbon feedstocks being used for end products more valuable than sulfur. Although coal is the logical alternative source

of hydrogen, the technology for generating reducing gas from coal is not adequately developed. Second, the generation of the highly toxic

and corrosive H 2 S requires that special precautions be taken to prevent

leakage or accidental release of the gas and that corrosion-resistant materials be used in the construction of the high temperature H 2 S

synthesis reactor.

In summary, there are three major problem areas associated with both non-regenerable and regenerable FGD processes: disposal or 9 utilization of waste sludges and the reclamation of sludge disposal ponds or landfills (non-regenerable processes); and the generation of reducing gases (regenerable processes). Significantly, it is also these three areas to which the research reported here is applicable.

In a previous study,designed to identify alternatives to sludge disposal, it was found that elemental sulfur could be recovered from waste sulfite sludges using an autoclave type process which involves the aqueous reduction of acidified sludge with ferrous sulfide. The process chemistry may be represented by the following overall equations;*

CaSOg + H 2 SO4 — ► CaSO^ 4- + H 2 SO3 (1)

2 H 2 SO3 + FeS —► l/4Sg + FeSO^ + ZH^O (2)

Actual yields of elemental sulfur from the process were as high as 95 per cent under the following conditions: 25 psig SO2 , 90 ®C, pH < 1.

Based on the results of this study, a preliminary economic evaluation showed that a 33 per cent annual gross return would be realized from a

2.7 million ton per year sulfite sludge reduction plant. The study also showed that, according to the overall reactions taking place, the reduction of sulfite by sulfide proceeds by a series of reactions involving the formation of thiosulfate and sulfane sulfonic acid intermediates which lead to the formation of elemental sulfur.

This process has potential application to the solution of each of the three problems (sludge disposal, pond reclamation and reducing gas generation) associated with FGD processes. Utilizing the sulfide

*A11 reactions take place in the aqueous phase; FeS and Sg are solids. 10 contained in the enormous reserves of natural occurring sulfide ores

(Table I), sulfur can be recovered from sulfite sludges either prior to or after disposal in ponds or landfills. These ores can also be used as a direct source of sulfide reductant for the in situ recovery of sulfur from regenerable FGD process streams, thus avoiding the use of scarce petroleum feedstocks to generate H^S. Furthermore, the sulfite reduction process enables the metal values contained in the ores to be recovered using conventional hydrometallurgical and electrochemical techniques.

C. Research Objectives

The overall purpose of the research performed in this study was to further elucidate the chemistry of the sulfite reduction process under various experimental conditions and to determine the general applicability of natural-occurring sulfide ores to the reduction process.

In fulfilling this purpose four specific objectives were identified:

• Confirm, through the use of spectrophotometric analysis, the

formation of the thiosulfate and sulfane sulfonic acid

intermediates as postulated in the previous study^® and to

determine their concentrations as a function of reaction time

• Determine the effect of temperature on the reaction rate of the

process

• Determine the suitability of six naturally-occurring sulfide ores

as reductants and compare their reactivity with that of FeS used

in base^hne studies 11

TABLE I

Domestic Reserves of Sulfide Ores®® (Estimated millions of tons)

Chalcopyrite

Chalcocite ) 150

Covellite

Galena 65

Sphalerite 45

Total 260 12

• Determine the reaction order and rates of decomposition of the

thiosulfate and sulfane sulfonic acids formed in each metal

sulfide-sulfur dioxide reaction; compare the reaction order to

that obtained in previous kinetic studies; and determine, by

comparing the rates, if the presence of particular ions affects

the rate constants.

The attainment of these objectives will not only provide the basis for further research into identifying the optimum conditions for the sulfite reduction process but also lead to a more complete understanding of sulfur-producing FGD processes and the development of more direct, efficient means of controlling sulfur dioxide emissions and recovering sulfur as a commercial powder. 13

CHAPTER II

BACKGROUND

A complete understanding of the chemistry of the sulfite reduction process and its relationship to the hydrometallurgical leaching of sulfide ores and the development of the methodology needed to conduct this research requires that several subject areas be reviewed. The following disucssion covers six major areas of aqueous solution sulfur chemistry applicable to the problem; a) the metal sulfide-sulfur dioxide reaction; b) the oxidation of metal sulfides in hydrometallurgy; c) the hydrogen sulfide-sulfur dioxide reaction; d) the chemistry of sulfur as it relates to these reactions; e) analytical determinations for thiosulfate and sulfane disulfonates; and f) the kinetics of the decomposition of thiosulfate into elemental sulfur. Since the production of elemental sulfur is the primary purpose of both the sulfite reduction and the hydrometallurgical processes, particular attention will be given to those reactions which yield elemental sulfur and the mechanisms (or lack of them) which have been proposed for these reactions.

A. The Metal Sulfide-Sulfur Dioxide Reaction in Aqueous Media

Several investigations®^”®® of the aqueous metal sulfide-sulfur dioxide reaction have been conducted, each in relation to the 14 development of new hydrometallurgical pressure leaching processes for sulfide ores. In 1970, Kunda and Rudyk^^ studied the aqueous reduction of sulfur dioxide by FeS (pyrrhotite), the reductant, at a pH < 1.0 using SO2 overpressures of 0 - 90 pounds per square inch gage (psig) and at reaction temperatures between 38 and 120“C. Their results showed that SO2 reacts exothermically with pyrrhotite to form elemental sulfur and ferrous sulfate according to the proposed reaction:*

2 SO2 + FeS —► 28 + FeSO^ (3)

By monitoring the ferrous ion concentration as a function of time, the effects of reductant concentration, SO2 overpressure, agitation rate, temperature and presence of impurities on the kinetics of the reaction were determined.

The primary mechanism they proposed for the aqueous reduction of sulfur dioxide to elemental sulfur (reaction 3) is a stepwise process given by the reactions:

2 SO2 (g) + 2 H 2 O (1) 2 H 2 SO 3 (aq) (4)

2 H 2 SO3 + 2FeS — ► 2 FeS0 3 + 28^8 (5)

SO2 + 2 H 2 8 — ► 38 + 2 H 2 O (6 )

8 O 2 + 2Fe803 — ► 8 + 2FeS04 (7)

In addition, Kunda and Rudyk proposed that a secondary mechanism, catalyzed by the addition of coal, produced elemental sulfur via the disproportionation of sulfurons acid by the reactions:

*These and all subsequent reactions take place in the aqueous phase unless otherwise noted; Fe 8 , other metal sulfides, and elemental sulfur ( 8 or 8 3 ) are solids. 15

6 H 2 SO 3 + 2S — ► H 2 S 2 O3 + 2 H 2 S3 O 6 + 3 H 2 O (8 )

H 2 S2 O 3 + 3 /2 SO2 — ► H 2 S3 O 6 + 1/2S (9)

3H2S3O6 — ► 3H2SO4 + 3802 + 38 (10) giving, for the secondary mechanism, an overall reaction:

9/2802 + 3 H 2 O — ► 3 H 2 8 O 4 + 3/28 (11)

The first two steps of the proposed primary mechanism (4 and 5) represent straightforward reactions. However, the next two steps

( 6 and 7) fail to adequately describe the mechanism by which elemental sulfur is formed primarily because at temperatures below 160 “C, elemental sulfur exists as 8 g and not "8 ". (This very important point will be discussed extensively in 8 ection D). The secondary mechanism, which leads to the formation of elemental sulfur in the colloidal state, was found to predominate at longer reaction times and higher temperatures. In addition, elemental sulfur formed after removal of the reductant — a process which led to the suggestion of the disproportionation mechanism. Although "disproportionation" reactions do occur in these systems,they are much more complex than those proposed here.

Kunda and Rudyk also observed two interesting phenomena associated with the formation of elemental sulfur by the primary mechanism:

1 ) the elemental sulfur formed on the surface of the reductant, and 2 ) the presence of coal catalyzed the formation of elemental sulfur.

Although no explanation was offered for these phenomena, it appears certain that these observations indicate a surface catalyzed reaction.

A proposed mechanism for this reaction will be discussed in Chapter IV,

8 ection E. 16

The formation of elemental sulfur on the surface of the reductant initially reduces the leaching rate and eventually prevents complete dissolution of the FeS. Thus, for a given amount of SO 2 ,either excess reductant must be used or the occluded reductant must be withdrawn from the reaction vessel, treated to remove the elemental sulfur, and then recycled for further reaction. Several methods may be used to remove the elemental sulfur from the surface of the reductant. These include solvent extraction (ammonium sulfide or tetralin) and heat induced volatilization of the elemental sulfur. The problem of surface occlusion is not unique to Kunda and Rudyk’s leaching process but is associated with all sulfur-producing hydrometallurgical processes. In fact, the surface occlusion of ores during leaching represents a potentially major technological roadblock to commercialization of sulfur-producing hydrometallurgical processes because of the increased costs associated with longer leaching times and sulfur removal. However, the complete understanding of both the chemistry of the formation of elemental sulfur in these processes and the mechanism of surface catalysis should lead to the development of chemical process control methods which will eliminate or reduce surface occlusion.

Two other studies of the metal sulfide-sulfur dioxide reaction will be mentioned briefly. Sobol and Frash^^ used SO2 as an oxidant in the leaching of aqueous suspensions of several metal sulfide ores to yield elemental sulfur and sulfates. By monitoring the leaching rate as a function of time, temperature, metal sulfide concentration and SO2 pressure, they determined the optimum leaching conditions and also found that the addition of H^SO^ increased the amount of ore dissolved. 17

However, they offered no mechanism to explain the formation of elemental sulfur.

Finally, Meyers, Hamersma and Kraft^^ outlined a process for the pressure leaching of copper concentrates with 100 psig SO2 in3.6 M HCl at 180 °C. For chalcopyrite, they proposed the following reaction:

4 CuFeS2 + 3802 + 12HC1 — ► 4CuCl + 4 FeCl 2 + U S + 6 H 2 O (12)

Their initial experiments resulted in greater than 98 per cent dissolution of chalcopyrite which indicates little or no surface occlusion. However, they offer no mechanism for the formation of elemental sulfur.

In very closely related studies, Meyers and co-workers^** have been working for several years on a regenerable process to remove pyritic sulfur from coal by leaching with FeClg. Although extensive kinetic studies have been conducted on this process, a mechanism is proposed which again does not adequately explain how elemental sulfur,

Sg, is formed. The stoichiometry they propose for this process is:

FeS 2 + 4 .6 1 ^ 2 (9 0 4 ) 3 + 4 .8 H 2 O — ►

1 0 .2 FeS0 4 + 4 .8 H 2 SO4 + 0.88 (13)

Leaching is conducted under the following conditions: 50 - 150 psig air (or 35 - 135 psig O 2 ), 110 - 130 “C, pH 1.3 - 1.9. Approximately

60 per cent of the pyritic sulfur is converted to sulfate and 40 per cent to elemental sulfur. The latter is removed from the coal either by toluene extraction or by heat-induced volatilization. 18

B. The Oxidation of Metal Sulfides in Hydrometallurgy

Sulfur oxide emissions from the nation’s 25 non-ferrous smelters account for about 15 per cent or 5.1 million tons of the total emissions of this pollutant (Table II).EPA has estimated that these emissions could be reduced by 75 per cent, to less than 1 million tons by 1979 with the use of SO2 control processes similar to those discussed in

Chapter I for coal-fired power plants. With the specter of federally enforced emission reductions looming over the pyrometallurgical process industry, increasing attention is being given to hydrometallurgical processes which eliminate the SO2 emissions associated with the smelting of sulfide concentrates. In addition, hydrometallurgical leaching techniques may provide a relatively simple means for recovering metals from deep, low-grade ore deposits. The increased activity in hydrometallurgical research is reflected in the convening of several international conferences on hydrometallurgy,the introduction of a new international journal, "Hydrometallurgy", by Elsevier-Amsterdam in 1975, and the recent appearance of annual reviews on hydrometallurgy in the periodical literature. ^ ° Several excellent general reviews have also been published.**

The hydrometallurgical extraction of metals from sulfide ores involves the leaching of pulverized ore concentrates in the presence of an oxidant at temperatures between 20 and 150 “C. Subsequently, the dissolved metal ions are recovered by cementation, electrorefining, hydrogen reduction, selective precipitation, or solvent extraction.

Depending on the type of process employed, either atmospheric or autoclave pressure leaching with air or oxygen may be employed. 19

TABLE II

Sulfur Oxide Emissions from Non-Ferrous Smelters^ ® (thousands of tons)

Smelter 1971 1979 (with controls)

Copper 4730 858

Lead 109.2 23.3

Zinc 289.3 34.1

Total 5128.5 915.4

3 20

Furthermore, the leaching solution may be acidic (pH < 3) or alkaline

(pH > 10). Acidic leaching under mild oxidizing conditions, low pressures (< 150 psig O 2 ), and moderate temperatures (< 120 °C), converts the sulfide in an ore nearly quantitatively to elemental sulfur. However, with increasingly stronger oxidizing conditions

(e.g., higher temperatures) the yield of elemental sulfur decreases and sulfate becomes the major product.**’ On the other hand, alkaline leaching under either mild or strongly oxidizing conditions has been found to produce thiosulfate, sulfane disulfonates and small amounts of elemental sulfur, all of which are eventually oxidized to sulfate.**^' **’• **®

Thus, the oxidation of metal sulfides in hydrometallurgy may be summarized in the following general reactions where O 2 may be oxygen or other suitable oxidant:

Acid Leaching

MeS + I/2 O2 + 2H'^ — ► Me^+ + l/8 Sg + H^O (14)

MeS + 2 O 2 + 2H"^ —► Me^+ + HgSO^ (15)

Alkaline Leaching

2MeS + 40g + 60H" — ► 2MeO^^“ + SgOgZ"+ 38^0 (16)

®2®3^~ + 20g + 20H" — ► 2S0^^" + H 2 O (17)

Reaction 14 is identical to the overall reaction given for the sulfite reduction process,

FeS + 2 H 2 SO3 — ► l/4Sg + FeSO^ + 28^0 (2) with the exception that the oxidant is sulfurons acid (i.e., aqueous sulfur dioxide). A comparison of the reaction conditions also reveals similarities : low pressures, moderate temperatures, and low pH. 21

Finally, since elemental sulfur is formed in both of these reactions under very similar conditions, one must conclude that the mechanism of its formation is the same.

This conclusion prompted the author to conduct a thorough search of the literature on hydrometallurgy to determine how the problem of the mechanism of Sg formation was dealt with in the leaching of sulfide ores. The following discussion and Appendix 2 contain a summary of the information obtained in this literature search.

1. Pourbaix Diagrams

The two most important thermodynamic variables in the dissolution of minerals are voltage (i.e., oxidation potential or free energy) and pH. Pourbaix**® has provided a very useful graphical tool for the presentation of thermodynamic data in the form of pH vs. potential diagrams. These diagrams extend from very high oxidizing to very high reducing conditions and make it possible to visualize, for those substances whose electrochemical potential has been determined, the stable and metastable phases which can exist in the various gas, solid, and aqueous solution phases. Thus, these diagrams have been used extensively by hydrometallurgists to describe reaction paths and phases which influence the surface reactions occurring during leaching. The following description of the construction of Pourbaix diagrams has been taken from a recent discussion by Wadsworth®® on the theoretical aspects of the leaching of sulfides.

All reactions involving aqueous solution equilibria can be expressed in the form: 22

aA + cHgO + ne- bB + inH'*' (18) where, for n 0, A represents the reactant in the oxidized state while

B is the reduced state. According to the Nerst equation:

E = - -Ol U lI - _ n _ I V, In a, (19) 0 nF nF where Z ° = AG“, the standard free energy change, and p, °, , and a, refer, respectively, to the chemical potential, the stoichiometric coefficient, and the activity of the ith component. At room temperature, equation 19 applied to reaction 18 becomes:

E„ = En° + S0.0591pH - P.--QA21log u " n n - a ^A

If voltages are referred to the hydrogen half-cell potentials, E q = E^, then a Pourbaix diagram can be plotted for a particular set of equilibria at various values of E^ and pH to show the thermodynamic relationships between the solid, gaseous, and dissolved components.

The three Pourbaix diagrams which are most relevant to hydrometallurgical leaching and this research are shown in Figures 1 to

3. Figures 1 and 2 show two different interpretations of the sulfur-oxygen-water system. Wadsworth®® has used Figure 1 as the basis for stating that the only stable sulfur species in this system are

HSO^", S0^^“ , HgS, HS", S^“ , and elemental sulfur. Thus, according to

Figure 1, under the acid leaching of sulfide ores, elemental sulfur often forms but metastable sulfur intermediates, such as thiosulfate and sulfite, are not observed. On the other hand, Wadsworth contends

that in alkaline leaching the metastable intermediates, such as

thiosulfate, and sulfane disulfonates do form, as shown in Figure 2.

However, he does not show these in his diagram (Figure 1). 23

0.5

so.

G

•0.5

HS"

■1.0 2 6 10 14

pH

5 0 Figure 1. Pourbaix diagram of the Sg/S^“ /S®'*'/H2 0 system. 24

CO 0.5 CM HSO/

V.X,

-0.5

HS'

• 1.0 2 6 10 14

pH

47 Figure 2. Pourbaix diagram of the 8 3 /8 2 0 3 ^“ /S^“ /S'*'*'/H2 0 system. 25

Dissolution

.2+ Cu CuO

Passivation 0.2 Jtz UJ

CuO

Cu

0.4

•0.6 2 6 10 14

pH

Figure 3. Pourbaix diagram of the CugS/CuS/Sg/Cu/HgO system 50 26

The formation of thiosulfate and sulfane disulfonates in the alkaline leaching of molybdenite®^ (M0 S2 ), galena®^ (PbS), pentlandite

(NiPeS), chalcopyrite (CuPeS2 ) , pyrrhotite (FeS), and pyrite®® (FeS2 )

is well established. This indicates that Figure 2 gives a more accurate representation of the equilibria present in alkaline solution. In addition. Figure 2 applies to acid leaching also, since, as will be shown later, thiosulfate and sulfane disulfonates are the precursors to

the formation of elemental sulfur, Sg. Thus, the construction and

interpretation of Pourbaix diagrams fail to give a complete picture of metal sulfide oxidation. Indeed, Peters®** recognized the fact that

Pourbaix diagrams failed to predict the formation of the various sulfane disulfonates and, therefore, put these species into the category of metastable species occurring along the reaction path during leaching.

With regard to metal sulfide-oxygen-sulfur-water Pourbaix diagrams

(Figure 3), Peters noted that similar diagrams, which describe the

leaching of CuFeS 2 with Fe 2 (8 0 ^ ) 3 or FeClg, failed to predict the

formation of elemental sulfur.

2. The Chemistry of Oxidative Leaching

The chemistry of hydrometallurgical leaching processes has been

thoroughly reviewed by Burkin,®® and Forward and Warren.®® As a part

of this investigation, a review of the most recent periodical

literature®®"®^ on hydrometallurgical chemistry was conducted and is

summarized in Appendix 2. This review shows the wide variety of

oxidants which can be used to leach sulfide ores. For acid leaching

the following oxidants have been used: 1 ^ 2 (8 0 4 )3 , FeCl 3 , CI2 , 0 3 , HNO 3 , 27

K 2 Cr 2 0 y, and anodic oxidation. In alkaline leaching only O 2 and

NaOCl have been used as oxidants.

However, more importantly, the review given in Appendix 2 points out the complete lack of an adequate mechanism describing the formation of elemental sulfur, Sg. This is primarily because all studies have investigated the kinetics of leaching as determined by monitoring the concentration of dissolved metal ions as a function of time. Several studies,’^' some of which have been reviewed by Wadsworth,**®' ®®'®® on acid leaching under the mild oxidizing conditions favorable to the

formation of elemental sulfur, have led to the development of linear, parabolic, mixed kinetic equations, and electrochemical models to explain the decrease in the rate of ore dissolution with time. These equations are not applicable under severe oxidizing conditions or in alkaline leaching where the formation of sulfate or thiosulfate is favored and linear kinetics are followed.

The retardation of leaching kinetics in acid is due to the formation of a non-reactive layer of elemental sulfur on the ore surface which reduces the ore surface area and the dissolution rate. That this is clearly the case has been shown by direct microscopic examination of the ore surface®® and the fact that the removal of the sulfur layer by solvent extraction or volatilization "reactivates" the ore.®^' ®^' ®®

Yet, the major question remains unanswered: What is the mechanism by which elemental sulfur forms in acid leaching, and, more specifically, how and why does it form on the surface of the ore? Such questions can

only be answered by realizing that elemental sulfur exists as the Sg

ring and is formed through a series of stepwise reactions involving 28 thiosulfuric, sulfane monosulfonic, and sulfane disulfonlc acids.

Ginzburg et al., in a series of extensive studies on the oxidation of sulfide ores in acid media, appear to be the only investigators who have observed surface occlusion, the formation of elemental sulfur, and the fact that the oxidation of sulfides to sulfates takes place in several stages, producing first thiosulfuric acid, then sulfane disulfonic acids, and finally sulfate. However, they did not correlate the formation of elemental sulfur with the occurrence of the sulfane disulfonic acids. Furthermore, although Dresher,®^ Forward,®® and

Vizsolyi et al.®® have provided important information on the formation of thiosulfate and sulfane disulfonates in alkaline leaching, the mechanism is still not understood.

Indeed, in 1963, Pawlek, in commenting on the work of Vizsolyi et al.,®® stated that, "Great efforts should be made to continue studies that will explain the mechanisms of sulfur oxidation, and only then will the pressure leaching processes for sulfides be fully understood." Yet, as of 1977, the understanding of these mechanisms had not been realized.

As indicated by Burkin**® in his review lecture on the winning of non-ferrous metals: "It is not understood at present what kind of mechanism occurs when a sulfide ion in a solid structure changes into a sulfate ion in acidic solutions containing an oxidizing agent, or into an anion of a lower oxyacid of sulfur in alkaline solutions The mechanisms, on the atomic scale, by which solids do not change in composition or structure during leaching, breakdown to release ions into solution and form solid products such as sulfur, are also not understood." 29

It was the stated®® purpose of this research to provide an explanation of the role of sulfane sulfonic acids in the formation of elemental sulfur in aqueous media and to provide evidence for the existence of these anions. It is hoped that this research will serve to stimulate further investigations into the complicated but intriguing field of sulfur chemistry.

C. The Hydrogen Sulfide-Sulfur Dioxide Reaction in Aqueous Media

This reaction should, as expected, be analogous in some respects to both the metal sulfide-sulfur dioxide and metal sulfide-oxidation reactions. All three reactions produce sulfur in the elemental form under mild oxidizing conditions, temperatures below 120 °C, and low pH.

The interaction of hydrogen sulfide with excess sulfur dioxide in aqueous solution is known to produce an extremely complex mixture of elemental sulfur and sulfane disulfonic acids (polythionic acids) called

Wackenroderfe liquid. Since the chemistry of this reaction will be discussed in the next section, it will only be said here that the mechanisms involved in the formation of Wackenroder's liquid still represent a classical problem in inorganic chemistry. The remainder of this section will deal with three regenerable flue gas desulfurization

(FGD) processes which use a reaction exactly analogous to the

Wackenroder reaction in producing elemental sulfur.

The three processes, APCI/IFP ammonia scrubbing, Stauffer's

"Aquaclaus" process, and the Citrate process, are described in Appendix

1. Each process produces elemental sulfur by reacting H 2 S with an

S0 2 -laden scrubbing solution thereby regenerating the SO2 adsorbent. 30

Although the chemistry involved in such a reaction is extremely complex,

the developers of these three processes have, to varying degrees,

ignored this complexity. Each has either referred to the formation of

elemental sulfur as occurring directly by the Claus reaction in aqueous media:

ZHgS + SO2 — ► 3S + 2H 2O (21) or has used this reaction in giving an inadequate description of the process chemistry. At best, terms such as "aqueous" Claus and

"Aquaclaus" are oversimplistic and cannot be used interchangeably with descriptions for gas phase reactions.

This can be seen by examining the chemistry of the Claus process.

In the original gas phase Claus process,HgS is reacted with a

stoichiometric amount of oxygen in a combustion chamber (Claus vessel): 400 °C 2 H 2 S + O 2 ------► l/4Sg + 2H20 (22)

Since the yield of this reaction is only about 70 per cent, the

remaining SOg and HgS are passed through two reaction ovens loaded with

catalysts where additional elemental sulfur is formed according to: 300 *C 2H2S + SOg------► 3/8Sg + 2H2O (23)

giving a 92 to 94 per cent overall yield for the process. The product

streams from these two reactions are cooled to condense sulfur in the

liquid state. Subsequently, the sulfur may be either stored or shipped

as a liquid or a solid.

Of course, reactions 22 and 23 just give the starting materials

and final products and do not indicate the mechanism by which elemental

sulfur forms. Although not yet shown experimentally, the mechanism

must involve the formation of gas phase sulfane free radicals which. 31

through a series of bimolecular reactions, form catena-^Sj^flt cyclo-Sn

(2 < n < 10), of increasing size, n. As the gas is cooled, the chain length or ring size increases (cf p. 39), When the sulfur vapor is condensed, the liquid exists as chains (4 < x < 10®) and rings

(5 < X < 33) which,on further cooling below 120 °C, form solid cyclo-octasulfur. The reactions involved in this mechanism definitely do not occur in aqueous media where gas phase free radicals cannot exist.

Despite the dissimilarity of the Claus process and the Wackenroder reaction, the developers of the APCI/IFP process®®' ® ® indicate that sulfur forms directly by reaction 21. On the other hand, the inventors of Stauffer's "Aquaclaus" process®^' ®® are aware that other reactions occur which involve the formation and degradation of thiosulfate and sulfane disulfonates. They give the following reactions:

HSO 3 - + S — ► 8 2 0 3 ®' + H+ (24)

6 S 2 O 3 ®- + 9 H 8 O3 " + 9H’*' — *► 4 8 3 0 6 ®" + 8 4 O6 ®" + SgOg®" + gHgO (25)

5 H 2 8 + 8 4 O 5 ®" + 2H+ — ► 98 + 6 H 2 O (26)

None of these reactions occur as written. Reaction 24 does not occur in acid solution; the products of reaction 25 are present in the product solution but do not form by this reaction; and in reaction 26, H 2 8 does not react directly with any species to form elemental sulfur. (See

Section D.)

The developers of the Citrate process®®"®® have performed in depth literature and chemical research of their process.®® However, after reviewing a considerable amount of the excellent work which has been done to elucidate the chemistry of sulfane disulfonates and the mechanisms involved in Wackenroder's reaction, they incorrectly concluded 32 that in their reaction system there "was no apparent reaction of thiosulfate or of the sulfane disulfonates to form sulfur." Furthermore, the authors refuse to accept the fact that the formation of elemental sulfur, Sg, must occur through a series of stepwise reactions involving the sulfane sulfonic acids. Instead, they give the following equations to account for the formation of elemental sulfur, "S";

SOg®- + 2H 2 S + 2H+ —► 3S + SHgO (27)

S20g®" + 2H 2 S + 2H+ — ► 4S + 3H 2O (28)

Finally, they draw a distinction between the colloidal sulfur formed in

Wackenroder's reaction and the "crystaffine sulfur" obtained in the

Citrate process. This distinction lacks scientific foundation and contradicts observations made in a previous study®** conducted by the original developers of the process. That study exhibited two photographs, one showing an elemental sulfur sol and the other showing how the sol can be "precipitated" by heating the solution above the melting point of sulfur (119.6 °C); this process is used to remove the sulfur from the reaction solution. A sol is necessarily the first form in which elemental sulfur appears visibly in Wackenroder’s solution because it consists of flocculated non-ionic cyclo-Sg molecules.®®

Subsequently, the sol can be destroyed by heating, thereby forming what could be interpreted as "crystalline sulfur," but is merely another allotropie form of cyclo-Sg.

A concluding note must be added concerning these three processes.

Although each has the advantage of producing elemental sulfur in a regenerable scrubbing process, there are two major disadvantages associated with using hydrogen sulfide as a reductant. First, the 33 generation of hydrogen sulfide requires the use of increasingly short supplies of natural gas;

CH4(g)+.l/2Sg(s)+2H20(g) C02(g) + 4H2S(g) (29)

Second, because of the extremely toxic, corrosive, and explosive nature of hydrogen sulfide, special safety precautions must be taken in its generation, storage, and use. Other reductants, such as CO or Hg, may be used, but these generally have the disadvantages similar to those of hydrogen sulfide and, in addition, produce undesirable side products in the reaction with SO2 .

With natural gas shortages becoming increasingly critical, EPA recently conducted a study® ® to examine alternate reductant sources for these processes. This study concluded that the most viable alternative is the gasification of coal or residual fuel oil to produce a hydrogen-carbon monoxide-rich gas which may be used as a reductant.

However, this solution merely shifts the problem from the use of a scarce natural resource, methane, to the use of other not so scarce natural resources, coal and oil. One obvious option the report failed to identify or investigate is alternative sources of sulfide.

As indicated in the previous chapter, the enormous reserves of natural-occurring sulfide ores can be used as a direct source of reductant for the in situ recovery of sulfur from FGD process streams.

Indeed, a process using this approach has already been proposed.®® It would integrate a regenerable FGD scrubbing system with a hydrometallurgical pressure leaching and metals recovery process for sulfide ores. In this process the SOg-laden scrubbing liquor would serve as the lixiviant for the ores, recovering the elemental sulfur 34 and dissolved metals by conventional hydrometallurgical techniques.

Thus, the need for generating hydrogen sulfide in FGD processes and sulfuric acid in hydrometallurgical processes would be eliminated.

D. The Chemistry of Sulfur

The occurrence of elemental sulfur and its oxy-acids (i.e., thiosulfuric, and sulfane sulfonic), and the mechanisms by which these compounds form in the various processes discussed up to this point, can be correctly explained by thoroughly examining the published literature on the chemistry of these compounds. The following discussion includes a review of the physical and electronic structure of elemental sulfur which have been used to correctly predict its degradation by nucleophilic agents such as bisulfite and cyanide ions. The nucleophilic degradation and rearrangement of sulfane sulfonates will then be discussed in some detail followed by a review of the origin of sulfane sulfonic acids and elemental sulfur in the aqueous reaction between hydrogen sulfide and sulfur dioxide (Wackenroder's liquid).

Several excellent texts and comprehensive reviews of sulfur chemistry are available. The properties and reactions of elemental sulfur have been reviewed in texts by Pryor,®* Meyer,^ ® Nickless,^ and Meyer and Schmidt.Meyer has also reviewed the most recent developments in elemental sulfur chemistry.The chemistry of both elemental sulfur and its oxy-acids has been reviewed by Senning,^ ® ® and

Schmidt and Siebert.^®**

Elemental sulfur, as a result of its occurrence in the free state, was one of the first chemical substances to be discovered. Its broad 35

reactivity led to its early association with mysticism and alchemy, and

later to its industrial importance. Sulfur occurs in a variety of forms

some of which have more than one crystalline structure. Table III

summarizes the various allotropie forms of sulfur. Table IV gives a

description of the several sulfur molecular species which exist in the

solid, liquid, and gaseous states. The temperature distribution of the

species present in saturated sulfur vapor are shown graphically in

Figure 4. The most stable sulfur molecules are cyclo-Sg, -Sj^2 » -Sg,

in this order. The unique stability of these molecules is due to

favorable symmetry and non-neighbor interaction between atoms across

the rings. In addition, experimental ionization potentials,

thermodynamic considerations, and theoretical considerations indicate

that all molecules, cyclo-S^, 6 < n < 12, including Sy, Sg, and S^Q, which suffer from unfavorable bond distances and contain non-equivalent

atoms, exist in the solid, liquid, and vapor phases.^®®

At room temperature the thermodynamically stable form of sulfur is

Sqj. The Sq( has an orthorhombic crystalline structure and consists of

Sg molecules in the form of staggered eight-membered crown-shaped rings:

S— S ^ ^ ^ S •s—

This structure allows for considerable cross-ring interaction between

nonbonded atoms thereby enhancing its stability.

When the orthorhombic form is heated, a sequence of changes takes

place beginning with the enantotropic transition to the 3 form at

95.3 °C. As this conversion is relatively slow, it is possible to 36

TABLE III

Sulfur Allotropes®®' ^®®

Name Description a Orthorhombic crystals; cyclo-Sg; stable to 95.3 °C where it slowly converts to 6 -Sg; melting point: 112.8 “C

3 Monoclinic I crystals; cyclo-Sg; stable from 94.4 °C to its melting point of 119.6 “C

Y Monoclinic II crystals; cyclo-Sg; melting point: 106.8 “C; above 75 ®C, converts rapidly to 3-Sg

6 Monoclinic III crystals; cyclo-Sg; very unstable; transforms to y-» Ç-» 3-, or a-Sg r) Monoclinic IV crystals; cyclo-Sg; lifetime: 10 minutes; observed during the preparation of 6 -Sg from rubber solutions containing sulfur and selenium

Ç Monoclinic V crystals; cyclo-Sg; lifetime: 12 hours; crystallizes together with 6 -Sg to which it spontaneously converts

0 1 , K Unstable allotropes of cyclo-Sg; identification inconclusive 0, T or ambiguous

X Designation for the free or bound cyclo-Sg molecule; the lattice unit in the polymorphs of cyclo-Sg; stable in solution and occurs in liquid and frozen sulfur ir Designation for all ring components, other than cyclo-Sg, contained in liquid or frozen liquid sulfur p Polymeric sulfur; designation for the catena-sulfur contained in liquid or frozen liquid sulfur i|j Fibrous sulfur; consists of a long helical structure containing 3-1/2 atoms per turn

(|) Fibrous sulfur which has been stretched; consists of long sulfur helixes (ip) that are packed regularly with Y-Sg 37

TABLE III (CONT'D)

Name Description w Designation for that portion of sublimed sulfur which is insoluble in CS 2 ; also called "white" sulfur; identification inconclusive

E, p Rhombohedral crystals; cyclo-Sg; occurs in the equilibrium liquid and vapor; reacts 10** times faster than cyclo-Sg 38

TABLE IV

Sulfur Molecular Species*®'

Form Description cyclo-Sg Exists in at least three stable solid crystalline forms; a, 3 » and y» colloidal sulfur is mainly a; exists in the liquid as the major species from about 120 to 159 °C (X-sulfur) cyclo-Sg Rhombohedral or trigonal crystals; Engel's or Aten's sulfur; produced by the acidification of NSgSgOg catena-S X Polymeric sulfur; forms long helixes with ten atoms comprising each three turns of the helix; molecular weight; 1 0 ** to 1 0 ®; present in the following metastable mixtures of allotropes: ip,

Becomes increasingly important species in sulfur vapor above 1800 °C 39

1.0

0e 0.1 1a a E

0.01

400 800 1200

Temperature, *K

Figure 4. Mole fraction of S^, 2 < n < 8, in a saturated vapor above the liquid phase. 40 maintain the form up to Its melting point of 112.8 °C. The monoclinic crystals are stable to their melting point of 119.6 “C where

a pale yellow, low viscosity liquid forms. This liquid consists primarily of Sg rings and shows no unusual behavior up to 159.4 °C.

At this temperature almost all properties of liquid sulfur suffer a

discontinuity. These Include the density, heat capacity, velocity of

sound, polarlzablllty, compressibility, molar polarization, electric

conductivity, surface tension, and many other properties. However, the most striking effect at this temperature, which Is often called the X

temperature. Is the abrupt 100-fold reversible Increase in the viscosity

(Figure 5). These phenomena, which have been extensively studied by

MacKnlght and Tobolsky,^°® are due to thermal scission of the Sg rings

into linear dlradlcals which attack other Sg rings In a chain reaction

resulting In the formation of polymers. The polymerization of sulfur

Is an equilibrium process and may be represented as follows:

cyclo-Sg *S-Sg-S* Initiation (30)

•S-Sg-S* + cyclo-Sg "S-Sg-S-Sg" Propogatlon (31)

At higher temperatures, the viscosity decreases dueto thermal

cleavage of the polymer. Rapid thermal quenching of molten sulfur

freezes the above equilibria and results In the formation of a

transparent elastic substance known as plastic sulfur (S^). This

substance hardens when kept at room temperature, with the formation of

visible orthqrhomblc sulfur crystals. Subsequent extraction with

carbon disulfide removes the Sg crystals leaving polymeric sulfur (S%).

The Sg molecule, cyclooctasulfur, although frequently shown with

single sulfur-sulfur bonds, may be represented by two resonance forms: 41

% e

160 240 320

Temperature, *C

Figure 5. Viscosity-temperature curve for liquid sulfur 99 42

SI IS SI IS 1! li = il It Si IS SI IS '^S=s<^ '^s—rS^

Schmidt^®® has explained these structures by considering the electronic configuration of sulfur:

Is^ 2s^ 2p^ 3s^ 3p2 3p^ 3p^

Thus, each sulfur atom has two lone pairs of electrons (3s% and 3p%) and two bonding electrons (3Py and 3p^), the latter forming covalent O bonds with adjacent sulfur atoms. The fact that sulfur, being a third-period element, has the tendency to expand its octet to a decet, thus including its five empty 3d orbitals, results in the formation of sulfur-sulfur multiple bonds from the (3p -► 3d)TT interactions.

However, recent considerations regarding the atomic wave functions, radial functions from self-consistent field theory, and orbital energies do not justify strong d-orbital participation in S-S bonds. Other calculations on cyclo-Sg, -Sg, -S^, and catena-Sg ions have confirmed that the contribution of d-orbitals to the ground-state energies of divalent sulfur compounds is negligible. On the other hand, values calculated for excited state energy levels fit observed spectra better when d-orbitals were included.Thus, at the very least there appears to be some d-orbital participation in S-S bonding, although it may be restricted to the higher energy states occurring during the transition period of chemical reactions.

The delocalized electronic system resulting from d-orbital participation, however weak it may be, indicates that elemental sulfur and its analogues should react as electrophilic centers or Lewis acids. 43 undergoing nucleophilic attack by Lewis bases via an S^2 mechanism.

This reaction results in the formation of sulfane mono- and disulfonates and their acids which bear a close resemblance to elemental sulfur in both chemical and physical properties.These similarities increase rapidly with increasing chain length of the sulfane. Thus, in the sulfane disulfonic acid + 4 -2 +4 HOgS-Sjj—SOgH each of the sulfur atoms of the chain terminating sulfonic acid groups has an oxidation number of +4 , whereas the compensating oxidation number of the sulfane, -2, is distributed along the sulfur chain. The longer the chain the higher becomes the average oxidation number of each sulfur atom — increasing from -2, asymptotically approaching zero with the extension of the chain and, thereby, approaching elemental sulfur itself.

The most marked difference, of course, must be found between elemental sulfur and the first member of the sulfane disulfonic acid group, that is, monosulfane disulfonic acid with n = 1. This is indeed the case as recently shown experimentally by Meyer, et al.^°^

This interpretation has enabled a large number of apparently diverse mechanisms to be rationally explained and others to be predicted.

A typical example of the Lewis acid-Lewis base concept is the stepwise degradation of cyclooctasulfur in neutral or slightly alkaline media by bisulfite.

It has long been known that elemental sulfur dissolves in boiling aqueous sodium sulfite solutions with the formation of sodium thiosulfate. Indeed, this reaction has been used for the preparation of

Na2S20g. For a long time the literature did not offer an indication as 44 to the actual course of this reaction. In most textbooks on inorganic chemistry it still is formulated as

Na2S0g + S —► Na2S20g (33)

That this formulation of the familiar redox reaction does not describe the true reaction course (the ideal aim of every mechanistic study) , but merely represents the starting materials and the end product, follows from the fact that sulfur, under the conditions of the thiosulfate synthesis, is present not as atomic sulfur but in the Sg ring form. A more realistic formulation of the thiosulfate formation therefore has to be written as

Sg + BNa^SOg — ► 8Na2S20g (34)

If the thiosulfate formation actually did proceed according to this reaction, then it would be a ninth order reaction, which, of course, is impossible. It is clear, therefore, that this "simple" reaction must proceed through intermediate stages. The first step is a nucleophilic

2 — attack of the base SOgH” (or SO3 , depending on the pH of the solution) on the "acid" cyclo-octasulfur (all eight sulfur atoms in this ring are, by nature, electronically equivalent; there is no electrophilic center for this first reaction step). Thereby, a sulfur- sulfur bond of the ring is broken and a new S-S bond between the attacking SO3H anion and the attacked sulfur atom is formed. This bimolecular reaction between the sulfur ring and the reductant or nucleophile is rate-determining for the overall reaction. It results in the formation of an open sulfur chain anion, terminated at one end by a SO3H group, that is the anion of octasulfane monosulfonic acid,

H 2 S9 O 3 : 45

SI IS'-.^ II It + SOoH" "S-Sg-S-SOoH (35) SI IS

The octasulfane monosulfonic acid anion is then degraded in a rapid series of steps via an S^ 2 mechanism — the bisulfite ion acting as the nucleophile attacking the electrophilic center of the anion, e.g., the terminal sulfur atom with its formal octet. This process results in the ejection of the moreweakly nucleophilic sulfane monosulfonic acid group and formation of thiosulfate, the final product:

"HOgST'S^g-SO^H — HO3S2" + “S-S5-SO3H (36)

"H0 3 sT"S^g-S0gH -fast» ROgSg" + “s-S^-SOgH (37) and so forth until:

“H O g S ^ ’s-ft-SOgH — 2HO3S2” (38)

The overall reaction is thus given by:

8HSO3" + Sg —► 8HS2O3" (39)

An identical first step and similar reaction mechanism has been shown for the reaction of sulfur with cyanide:^®®

t| ir^CN" NCS-S-S5-S" (40) Si is

NCS^S^Sg-S" T"CN" — NCS-S5-S" + SON" (41)

NCS^S^4-S" CN“ NCS-S4-S" + son” (42) and so on until:

NCS-S + CN — ► 2SCN (43) 46 with the overall reaction given by;

8 CN" + Sg — ► 8 SCN" (44)

In their elucidation of the reaction mechanism, Bartlett and

D a v i s , i n d i c a t e that two factors contribute to reactions 41 to 43 being much faster than the rate determining reaction 40. First, the sulfur-sulfur bond in the chain is more easily broken than in the eight-membered sulfur ring, and second, the thiocyanate group is a much better leaving group than the polysulfane group.

Similar effects due to leaving group stability are encountered in the sulfite-induced degradation of sulfane disulfonic acids and their anions in aqueous solution:

H 0 3 S-Sn-S0 3 H + (n-l)H2 S0 g HO3 S-S-SO3 H + (0 -1 X8 2 8 2 0 3 (45)

This reaction takes place spontaneously and quantitatively only in the presence of excess sulfurous acid. Stoichiometric and lesser amounts lead to mixtures of different sulfane disulfonic acids as indicated by the equilibrium nature of the reactions.

Two mechanisms have been proposed for this reaction. Both lead to the same end products: thiosulfate and monosulfane disulfonate which is not degraded by sulfite. On the basis of both radionuclide studies and the existing possibilities for the point of heterolytic cleavage of the sulfur-sulfur bonds, Foss^^® has proposed that the reaction leading to the displacement of thiosulfate is the one that is greatly favored.

For trisulfane disulfonate, the reaction proceeds by a series of displacements of the weaker nucleophilic thiosulfate ion by the stronger nucleophilic sulfite ion: 47

■O3 S-S-S-S-SO3 " + S0 3 ^~ ^ "O3 S-S-S-SO3 " + 8 3 0 3 ^ (46)

-O3 S-S-S-SO3 " + 3 0 3 ^" "O3 S-S-SO3 " + 8 2 0 3 ^“ (47)

Schmidt^®® proposes a slightly different mechanism based on his interpretation of the reaction between labeled sulfite and disulfane disulfonate which results in the entire activity being present in the disulfonate (see reaction 47). Extrapolating this result to higher sulfane disulfonates leads to the prediction that the leaving group is the sulfane monosulfonate group:

O3 S—S—S—S—SO3 + 3 0 3 ^ O 3 S-S—3 + O 3 3 —3—SO3 (48)

"033-^3"^303^" 2 8 2 0 3 ^” (49)

3ulfane disulfonates also undergo disproportionation reactions which are catalyzed by thiosulfate. These reactions ultimately result in the formation of elemental sulfur. The mechanism as proposed by

Foss^^®' ^^^ for trisulfane disulfonate and higher disulfonate ions, involves an initial nucleophilic attack of the thiosulfate on the sulfane disulfonate resulting in the formation of the next lower sulfane disulfonate and a sulfane monosulfonate ion. For trisulfane disulfonate the initial reaction:

O 3 3 —3 + O3 S—S—3—3—SO3 . O3 3 —3—3—3 O 3 + S— 8 3 0 3 (50) is followed by a building-up of the sulfane monosulfonate ions into polysulfane monosulfonate ions through a series of consecutive nucleophilic displacements :

"0 3 3 2 -8 '^^+^5^2°3~ ^ "0 3 8 2 - 8 2 " + 8 2 0 3 ^" (51)

O 3 8 2 - 8 2 + 3 ^ 2^)3 ^2^3^ (52) 48 and so forth, until Internal displacement produces ring closure and elemental sulfur:

S-Sg—S—S—SO3 — ► Sg + S2O3 (53)

Foss^^^ also points out that a thiosulfate-catalyzed rearrangement of trisulfane into disulfane and tetrasulfane disulfonate may occur as a result of an initial displacement (reaction 50) followed by:

0382—8 + O3S—S—8—S—SO3 . ■ O3S—S—S—S—S—SOg + S2O3 (54)

Similar rearrangements occur in still higher sulfane disulfonates. In fact, any sulfane monosulfonate occurring in reactions 50 to 53 should be expected to attack a disulfonate chain to give another sulfane disulfonate, preferably a higher one, since the thiosulfate is thebetter leaving group relative to a mono-, di-, or polysulfane monosulfonate group.

The thiosulfate-catalyzed disproportionation of disulfane disulfonate has been shown by Fava and Bresadola^^^ to proceed by the following mechanism:

O3S—S + O3S—S—S—SO3 ■. " O3S—S—S—S—SO3 + SOg^ (55)

SOg^" + “038-8-8-803“ “O3S-S-SO3" + "8-803“ (56) giving an overall reaction:

2 O38—8—8—8O3 " O38—8—8—8—8O3 + O38—8—8O3 (57)

If significant amounts of thiosulfate are present in solution it will undergo an acid-catalyzed conversion to elemental sulfur and sulfur dioxide. Davis^^® has interpreted the decomposition of acidified sodium thiosulfate in aqueous solution as proceeding by a mechanism which involves a series of displacements of sulfite by thiosulfate: 49

+ H*’' HSSO3 " (58) ___ ^ ^ JJ+ “O3SS" + HS-S03~ — HSSSO3" + HSOg" (59)

“038-3“ + 1^8^803” — ► 8888-803“ + 803^“ 88O3" (60) and so on, until

883803 — ► 8g + 88O3 (61)

Because the sulfite ion is immediately converted* by protonation, to sulfurous acid in the quite acid medium of the reaction, the reverse degradation reactions involving the displacement of thiosulfate from the polysulfane monosulfonates by sulfite do not occur.

The oxy-acids of sulfur are the primary products of the reaction of hydrogen sulfide and sulfur dioxide in aqueous media. This reaction results in the formation of a mixture of polysulfane disulfonic

(8 2 8 ^0 3 ) and polysulfane monosulfonic acids (HgSj^Og), and elemental sulfur known as Wackenroder's liquid. Wackenroder^^performed the first detailed investigation of this reaction in 1846 when he identified the existence of trisulfane disulfonic acid (pentathionic acid H^S^Og).

A comprehensive review of the literature prior to 1952 on the chemistry of sulfane disulfonic acids in general, has been given by Goehring.^^®

Although the complete elucidation of the mechanisms leading to the formation of these acids still represents a classical problem in inorganic chemistry, considerable progress has been made.

Stamm et al.^ ^ ® have proposed the following steps for the formation of di- and tetrasulfane disulfonic acids from hydrogen sulfide. 50 sulfurous acid and thiosulfuric acid in aqueous solution at pH = 0.5:

HgS + SOg ^ HgSgOg (62)

H2S2O2 + 2SÜ2 — H^S^Og (63)

H2S2O2 + 2H2S2O3 —► H2SgOg +2H2O (64)

In a more extensive study, Blasius and Burmeister^^^ employed radio-paper chromatography to determine the products of Wackenroder's reaction. They found that the major product of the reaction is disulfane disulfonic acid, HgS^Og, and proposed the following mechanism:

H2S + 2H2SO3 —► 3H2SO2 (65)

H2SO2 + 2H2SO3 — ► H^SgOg + 2H2O (66)

H 2S30g + H2S — ► 2H2S2O3 (67)

H+ + S^Og^" + SgOgZ- ^ S^Og^" + HSOg" (68)

The monosulfane disulfonate and thiosulfate then react to form higher sulfane disulfonates (n > 3) via the equilibria given in reaction 45.

Schmidt^reports that the product of reaction 62 actually represents an intermediate which reacts with sulfurous acid to form the sulfoxylic acid produced in reaction 65:

H 2 S + SO2 H 2 S2 O 2 (69)

H 2 S2 O2 + H 2 SO3 + H 2 O 3 H 2 SO2 (70) to give an overall reaction:

8 3 8 + 2 8 2 8 O3 3 8 2 SO2 (71) as shown in reaction 65.

Finally, in a recent study of the complex Wackenroder reaction by

8 telnle,i^® he proposed the following reaction sequence. The first step in the reaction is the formation of disulfurous acid (8 2 S2 O3 ) 51

anion, which is produced from the interaction of sulfur dioxide with

water:

SO2 + H 2 O HSOg" + h'*' (72)

HSO]" + SO2 HS2 O5 " (73)

The disulfurous acid anion then reacts with hydrogen sulfide to produce

thiosulfate and the unstable sulfoxylic acid which leads to the

formation of disulfane disulfonic acid:

HgS + HS2 O5 " HS 2 O3 - + H 2 SO2 (74)

HS2O3" + H 2 SO2 HO 3S-S-SOH + 0H“ (75)

HO3 S-S-SOH + HSO3 " HO3S-S-S-SO3H + OH" (76)

The formation of trisulfane disulfonic acid and elemental sulfur would

then occur by the following steps:

HO3 S-S-SOH + H 2 S2 O3 HO3 S-S-S-S-SO3 H + H 2 O (77)

H 2 SO2 + 8 3 8 8 3 8 3 0 + 8 3 0 (78)

X8 2 S2 O — ► S2 JJ + X8 2 O (79)

Although the last two steps contradict the previously described mechanisms for the formation of elemental sulfur, reactions 74 to 76,

in combination with the mechanisms proposed by Stamm, Blasius, and

Schmidt, have provided the basis for the interpretation of the results

obtained in this work.

E. Analyses for Thiosulfate and Sulfane Sulfonates

Quantitative methods for the determination of soluble thiosulfate

and sulfane disulfonates include gravimetric, volumetric, polarographic,

chromatographic, and spectrophotometric techniques. A description of

these methods may be found in the comprehensive reviews by Blasius 52 et and Karchmer.^ ^ Newly developed techniques and improvements in established methods include high pressure liquid chromatography/^^ coulometric titration,and thin-layer high voltage electrophoresiscoupled with coulometric titration,^ radiochemical analysis,or remission spectrophotometry.^^® However, most of these methods are either not ion-specific and time consuming, or they require special equipment. For example, the most common method for thiosulfate is iodometric titration.This method is not satisfactory for the research described here because the other oxidizable species in solution, e.g., sulfite, or dissolved SO2 and the sulfane monosulfonic acids are also titrated.

Similar problems arise in the quantitative analysis of sulfane disulfonates. Since individual disulfonates are very similar in their chemical and physical properties and they are present with other oxy-acids of sulfur, their direct analysis is not practicable. Indirect methods may be used to determine individual sulfane disulfonate concentrations, but a large number of analysis equations is; necessary, which must be obtained through correspondingly different analyses. In addition, these analyses are time consuming and they give erroneous results due to the changes in sample composition with time. New methods, such as high pressure liquid chromatography,eliminate these problems but require special equipment not available to us.

In selecting an analytical method for use in this study, a prime consideration was that a relatively large number of samples was anticipated. Thus, a method was sought which could be performed rapidly, thereby avoiding sample decomposition, and at a very minimum. 53 eliminate the need for separate analyses for sulfite and thiosulfate.

After a series of test runs, it was found that ultraviolet absorption spectrophotometry met these requirements.

The ultraviolet (uv) absorption spectra of the various sulfur oxy-acids have been known for some time. For example, in 1937, Ley and

Konig^®° found that the uv absorption spectrum of thiosulfate in solution appears as a single intense band with a peak at 215 nm. This peak has been assigned to a charge-transfer-to-solvent (c.t.t.s.) type of excitation.^^ There appears, however, to be only limited use of the spectra of the sulfur oxy-acids in quantitative analysis, although each known species follows Beer's law very closely.

The pertinent absorption parameters for the several species of interest in this study, S0 2 (aq),^®^' and sulfane disulfonates,^®® are given in Table V. The high molar absorbtivities of

the thiosulfate and sulfane disulfonate species proved to be very valuable in the real-time analysis of the reaction solutions: It meant that reasonable absorbance values (i.e., 0.3 to 1.0) could be

obtained at low species concentrations (10~^ to 10” ** M). Thus, the

reaction solution, which was approximately 10” ^ M in thiosulfate had

to be diluted by a factor of 1000 prior to analysis. When these high

dilutions were made they effectively "froze" the reaction, enabling the

thiosulfate concentration to be determined without significant

decomposition.

Finally, it will be noted that although uv analysis distinguishes

between S0 2 (aq) and thiosulfate, it does not allow for individual

sulfane disulfonate analysis due to the proximity of their values. 54

TABLE V

UV Spectra of the Sulfite, Thiosulfate, and Sulfane Disulfonate Anions

log e max

Species X , nm 1 0 “® liter mole“^ cm“^ max

SOg 276 2.78

S2 O 3 "- 215 3.57

-195 -

216 3.85 SA" 214 3.92

217 4.19 55

In fact, when the peak absorbance is read at 215 nm all species are included. However, this does not affect the data interpretation with respect to either the formation of these precursors of elemental sulfur or, as will be seen in the next section, their subsequent decomposition into elemental sulfur.

F. The Kinetics and Mechanism of the Acid Decomposition of Thiosulfate

The first extensive and reliable kinetic studies on the decomposition of acidified solutions of thiosulfate were conducted by

La Mer and co-workers.^®® They studied the production of monodispersed sulfur hydrosols prepared by the reaction between very dilute solutions of strong acids (HCl or H 2 SO4 ) and sodium thiosulfate:

^2^3^ + H"^ — ► HSO3 + S (80)

Reactions were performed at 25 °C using the following reactant concentration ranges : 0.001 to 0.003 M thiosulfate, 0.001 to

0.006 M H"*", and ionic strengths from 0.007 to 0.03. Upon mixing the reactants the solution remained transparent for a period of time which was inversely proportional to the initial reactant concentrations. After this induction period, discrete particles of sulfur formed as indicated by the appearance of a Tyndall beam. However, during the induction or homogeneous reaction period, the formation of molecularly dispersed sulfur could be followed by transmittance measurements at 300 nm. Thus, the increase in optical density as a function of time was measured in

10 cm quartz cells with a Beckmann DU spectrophotometer.

A comparison of the reaction rates at constant ionic strength showed that the rate of fomnation of molecularly dispersed sulfur could 56 be expressed as a function of the initial thiosulfate and hydrogen ion concentrations :

[H+]q1 / 2 (81)

The fractional order of this equation and the shape of the optical density vs. time plots were interpreted as an indication of an autocatalytic reaction, probably of the chain type. It was found that the reaction 80 exhibited a positive Bronsted primary salt effect as shown by the increased reaction rate in the presence of lanthanum salts.

This effect was attributed to the fact that the La®* species set up very large local deviations from the average bulk anion concentration by attracting the negative ions thereby promoting more rapid reaction among them. This interpretation was supported by the observation that the calalytic effect of La®* was greatly reduced in the presence of sulfate since the doubly charged anion competed with the reacting species for loci near the La®*. Thus, the rate-controlling step of the thiosulfate decomposition was predicted to be a reaction between negatively-charged ionic species, either 8 2 0 3 ^” or HS2 O3 ".

Keller^®® has conducted a kinetic study of the reaction between

H 2 S and H 2 S2 O3 in aqueous media (pH 3.5 to 5.5), as it relates to the

formation of elemental sulfur in a proposed flue gas desulfurization process. He found that the reaction rate was independent of the H 2 S

concentration and that the rate of disappearance of thiosulfate was

given by the equation; -d[S203^~] = k [8203^" ] 3/2 [H+] 1/2 (82) dt A comparison of this equation with equation 81 shows that the reaction between H 2 S and H 2 8 2 O3 , and the acid decomposition of thiosulfate 57 apparently proceed by a very similar mechanism. Thus, when these two equations are combined it can be seen that the rate of formation of elemental sulfur may be determined by monitoring the rate of disappearance of thiosulfate; d[S] ^ -d[S203^~ 3 (83) dt “ dt Based on La Mer's observations, Davis^^® has interpreted the decomposition of acidified thiosulfate solutions on a mechanistic basis.

The rate expression (Equation 81) was explained using the following series of bimolecular nucleophilic displacement reactions at the sulfur atom;

8203^" + H'* H82O3- (84)

88203“ + 8303^ — ► 118303“ + 803^ (85)

H8303- + 8303^ — H84O3" + 803®- (86)

HS403- + 8303^- — H83O3- + SOgZ- (87)

H8303- + 8303^ — H8g03 + SOgZ" (88)

H8g03“ + 8303^ - ► HSyOg" + SOgZ" (89)

88^03“ + S203Z- — 88303- + SOgZ- (90)

88303“ + 8303^ — ► H8903— + SOgZ- (91)

889O3- — ► 8g + H8O3- (92)

To derive the overall rate equation Davis represented each of these reactions by a first-order differential equation. Then, through a series of successive evaluations and substitutions (see reference 113 for the complete derivation), he obtained the following expression;

= [H+]q 1/2 [SgOgSlg* 2 > n > 1 (93)

Thus, reactions 84 to 92 are reasonable representations of the mechanism by which elemental sulfur, Sg, is formed. These reactions also account 58 for the positive salt effect observed by La Mer, since all reacting species are anions. Finally, this mechanism shows thiosulfate participating in each of the reactions leading to the formation of Sg.

Thus, the spectrophotometric determination of thiosulfate in the presence of the transient sulfane sulfonate species, which absorb at the same wavelength (Table V, Section E), gives absorbance values which are only directly proportional to those for thiosulfate. However, for the purposes of this study in applying equations 81 and 83, the absorbance values were assumed to be totally due to thiQsulfate.

The elegant derivation by Davis gives further proof that the formation of elemental sulfur in aqueous solution must proceed through a series of stepwise reactions involving thiosulfate and polysulfane monosulfonic acid intermediates. In light of the evidence provided here and in previous sections, one must discard as untenable and oversimplistic those mechanisms which have, for too long a time, shown elemental sulfur, "S", as being formed by a one-step bimolecular reaction. 59

CHAPTER III

EXPERIMENTAL

A. Materials

Three reagents were used in each experiment: sodium sulfite (or calciner feed), a pulverized metal sulfide ore, and 4.0 M hydrochloric acid or 2.0 M sulfuric acid. Anhydrous sodium sulfite powder (Baker

Analyzed) was used without further purification. Stock solutions of

4.0 M hydrochloric acid and 2.0 M sulfuric acid were prepared from the concentrated acids purchased from Mallinckrodt. Seven different sulfide ores were used. The ferrous sulfide (laboratory grade: 62.9% Fe,

37.1% S), used for base-line and comparative studies, was obtained from Fisher Scientific Co. The other six naturally-occurring sulfide ores were purchased from Ward's Natural Science Establishment, Inc.: galena (PbS), pyrrhotite (FeS), sphalerite (ZnS), chalcopyrite (CuFeSg),

chalcocite (CugS), and pyrite (FeS2 ). Prior to use, each ore was pulverized with a mortar and pestle, and then screened through a

standard 200 mesh sieve. To minimize differences in the surface area, an identical screen distribution was selected for each ore:

14.6% +200 and 85.4% -200. The calciner feed or sulfite sludge (63.24%

MgSOg, 12.20% MgSO^, 5.37% MgO, 14.00% HgO, 0.47% C) was supplied by

the Environmental Protection Agency. The feed was pulverized (not

screened) before using. 60

Two standard solutions were prepared for use in determining the

Beer law relationships for thiosulfate and the ferrous-phenanthroline complex. A 0.1007 N NagSgOg solution, prepared using reagent grade crystals, was standardized^**® using primary grade K2^^2^7" Similarly, the standard Fe(NH4 )2 (S0 4 ) 2 solution, containing 5.683 mg Fe^*/ml, was prepared from reagent grade crystals and standardized^**^ against 0.1 N

K 2 Cr2 0 j prepared from the primary standard grade reagent.

B. Apparatus

Experiments were conducted using the apparatus shown in Figure 6 .

The major components of this apparatus are a Fisher-Thermix heater-stirrer, a 500 ml Parr high pressure reaction bottle, a 0 - 30 psig Monel pressure gage (Air Products), a 0 - 100 °C thermometer calibrated to ± 0.05 °C, and a sampling port consisting of a rubber septum inserted in a stainless steel tubing fitting. The sampling port provided the means of removing samples from the reaction mixture with

Hamilton Gas-Tight syringes equipped with 24 ga 5.5 in. stainless steel needles (Glenco Scientific, Inc.). The leakage rate of this apparatus at 17.0 psig was < 0.1 psig per hour.

Constant temperature kinetic studies were conducted in two Lauda

K-2/R (Brinkmann Instruments) water baths, maintained at 25 “C ± 0.05 and 40 “C ± 0.05. A Forma Scientific, Inc., Walk-In Environmental Room was used to maintain other samples at 5 “C ± 0.05. 61

Thermometer Monel Pressure Gage

Pressure Relief Valve

/ Sampling Port Stainless Steel Tubing

Paar High Pressure Reaction Bottle

Heater-Stirrer

Figure 6 . Sulfite reduction apparatus. 62

C. Procedure

The quantities of reactants used in each experiment were similar

to those used in previous work^® and were based on a one-tenth stoichiometry of the following reaction:

NagSOg + l/2FeS + 2HC1 — ► l / 8 S g + l/2FeS0^ + 2NaCl + H^O (94)

Thus, in all cases the Na2 S0 g : HCl mole ratio was 1 : 2 or 0.1mole

NagSOg and 0.2 mole HCl. (In one case 0.1 mole H^SO^ was used.) This insured the complete conversion of the SO^ to SO2 (aq) according to:

NagSOg + 2HC1 — ► SO^ (aq) + 2NaCl + H^O (95)

The quantities of metal sulfides used depended on the experiment being run and the reactivity of the sulfide. For the more reactive ferrous sulfide, used in base-line experiments, 0 . 1 mole or two times the FeS stoichiometry of reaction 94 was employed. For comparison, one experiment with ferrous sulfide was conducted at the one-tenth reaction stoichiometry of 0.05 mole FeS. The less reactive naturally-occurring sulfide ores were used in 0 . 2 mole quantities, i.e., four times the reaction stoichiometry. An exception was the more reactive pyrrhotite which was used in 0 . 1 mole quantities.

Reactions were performed in the Parr high pressure reaction bottle

(Figure 6 ). Reactants were weighed to the nearest mg and added to the reaction bottle in the same order for each experiment. First, the sodium sulfite (or calciner feed) was quantitatively transferred to the bottle and completely dissolved, with stirring, in 1 0 0 ml of distilled water. Then the pulverized ore (reductant) was added followed by

thorough mixing. Finally, 50.0 ml of 4.0 M HCl or 50.0 ml of 2.0 M

HgSO^ was mixed with 50 ml of distilled water and added to give a total 63 volume of 200 ml. The reaction bottle was then immediately attached to the apparatus. During the course of the reaction the mixture was heated and stirred (720 rpm) using the Fisher Thermix apparatus. A typical heating curve is shown in Figure 7. Reactions were run at various temperatures (30 to 90 °C) for periods of 120 to 315 minutes.

Temperature and pressure readings were taken at intervals of 5 or 10 minutes. In addition, at 10 to 30 minute intervals, 100 or 200 ]il samples were removed from the reaction mixture using the gas-tight syringes. Each sample was diluted immediately by a factor of 125, 200,

250, 500, or 1000 using Class A volumetric flasks. Aliquots from these dilutions were subsequently analyzed for SOg^ or 8 2 0 3 ^", and ferrous ion. Thus, the progress of each reaction was followed as a function of time, pressure, temperature, [SO^^" ] or [8 2 0 3 ^]' and [Fe^*].

Upon completion of an experiment any remaining 80^ pressure was released, the reaction bottle removed from the apparatus, stoppered, and cooled to about 10 “C in an ice bath. This "froze" the reaction and allowed the excess ore to settle out. [In experiments run during the latter part of this research it was found that it was necessary to centrifuge the reaction supernatant to ensure complete removal of the excess ore.] Subsequently, 25 ml aliquots of the reaction supernatant were transferred to 25 ml volumetric flasks for kinetic studies on the decomposition of thiosulfate. 64

75

u e

50

a>E

25

50 100

Time, minutes

Figure 7. Heating curve for a reaction temperature of 70.8 °C. 65

D. Methods and Instrumentation

1. Sulfite and Thiosulfate

The samples withdrawn from the reaction mixture were analyzed first for SO^^" and . Since there was some sample decomposition with time, these analyses were performed within 30 minutes of sampling; some samples were stabilized by the addition of one or two drops of

4.0 M HCl. Both S0g2" and SgOg^" were determined by their uv absorption spectra using a Cary 14 Recording Spectrophotometer. One cm quartz cells

(distilled water blank) were used for spectral scans between 2 0 0 and

350 nm. Sulfite was determined qualitatively by its absorbance at

276 nm (e = 600 liter mole cm "^ ).

Thiosulfate was determined quantitatively by its absorbance at

215 nm (e = 3.68 x 10^ liter mole”^ cm"^). The Beer’s law relationship for thiosulfate was determined by using the 0.1007 N NSgSgOg at dilutions from 1:100 to 1:1000. A linear least squares regression analysis of the concentration-absorbance data gave a coefficient of determination (r^) of 0.9984 and the following equation which was used to calculate the thiosulfate concentration in the reaction mixture:

[SgOgZ-] = (2.728 X 10-4 A - 9.473 x 10"®) x D.F. (96)

where : A = Absorbance at 215 nm

D.F. = Dilution Factor of the syringe-dilution method.

The molar absorbtivity (e) of thiosulfate calculated from this equation is Eq r» 2- == 3.43 x 10® liter mole cm . This compares Û2 W 3 favorably with the literature value of 3.68 x 10® liter mole"^ c m ^ ®®

Since the absorbance for thiosulfate is a factor of about six greater 6 6 than that of sulfite, to obtain suitable absorbance values for the two species, dilutions for the thiosulfate analysis had to be six times greater than those for sulfite. At this higher dilution the sulfite absorption peak did not appear in the spectra and thus did not interfere with the thiosulfate analysis,

2. Ferrous Ion

Subsequent to the sulfite/thiosulfate analysis all samples taken in experiments using iron-containing ores were analyzed for ferrous ion by the ferrous phenanthroline m e t h o d . This method involves the formation of the red ferrous-phenanthroline complex whose absorbance is read at 510 nm against a water blank. Absorbance values were determined using a Beckmann Spectronic 20. A Beer's law plot was made using the standard 5.683 mg Fe^* per ml solution at five dilutions between 1:100 and 1:1000. A linear least squares regression analysis on the concentration-absorbance data gave a coefficient of determination

(r^) of 0.9998 and the following equation which was used to calculate the ferrous ion concentration in the reaction mixture :

[Fe%+] = (0.1121A + 1.265 x 10"®) x D.F. (97)

where A = Absorbance at 510 nm

D.F. = Dilution Factor of syringe-dilution method.

E. Kinetic Studies

The kinetics of the decomposition of thiosulfate were studied for

those metal sulfide-sulfur dioxide reactions in which greater than

0.0600 M SgOg^" was formed. Using the syringe-dilution method, the 67 thiosulfate concentration in the simples held at constant temperature was determined approximately every 24 hours by uv analysis. These studies were conducted for a period of 8 to 24 days. Although kinetic measurements were initially made at 5, 25 and 40 °C, the rates of thiosulfate decomposition at 5 and 25 °C were too slow to be meaningful.

Therefore, only the measurement at 40 °C are discussed in Chapter IV,

Section D.

F. Data Analysis

Characterization curves for each reaction were drawn by plotting the pressure,?,[8 2 0 3 ^ ], and [Fe®*] data vs, time. Using these curves and the original data, maximum values for -dP/dt, dfSgOg^j/dt, and d[Fe®*]/dt were calculated for the time period during which two or three data points gave a straight line. The peak reaction pressure invariably coincided with the maximum value for d[Fe^*]/dt and was followed shortly thereafter by the maxima in dtSgOg^-j/dt and -dP/dt.

The temperature of each reaction was determined by taking the average temperature during -dP/dt; the average increase in temperature during this period was 2.7 °C. The total decrease in pressure, -AP, was calculated as the difference between the maximum pressure and the pressure at the time of the maximum [SgOg^"]; the [Fe^*] was also noted at the maximum [8 2 ,0 3 ^] ' These various parameters were interrelated through the use of least squares regression analysis.

The kinetic data were initially displayed in linear plots of time (hr) molar thiosulfate concentration. Subsequently, the order of the thiosulfate decomposition reaction with respect to time was 68 determined by plotting the logarithm of the rates, U = dc/dt, at various times against the logarithm of the corresponding reactant concentration, c.^^z The slope of the line, corresponding to the order of the reaction with respect to thiosulfate, and the y-intercept, corresponding to the logarithm of the rate constant, were determined by log-log least squares regression analysis. 69

CHAPTER IV

RESULTS AND DISCUSSION

The experiments conducted as a part of this research were designed to elucidate the chemistry of the sulfite reduction process by characterizing the metal sulfide-sulfur dioxide reaction under a variety of experimental conditions. While the results from these experiments have achieved the four specific research objectives identified in

Chapter I, Section C, additional information has been derived through an in depth analysis of the recorded experimental data. The complete results and their interpretation are presented in the following discussion under five major subject areas: a) the characterization of the ferrous sulfide-sulfur dioxide reaction at five temperatures and three reactant variations; b) the characterization of the sulfide ore-sulfur dioxide reaction for six naturally-occurring ores; c) a comparison of sulfide reactivities using induction period pressure- temperature relationships; d) the kinetics of the acid decomposition of thiosulfate in nine reaction solutions; and finally in summarizing the results, e) a chemical model is proposed which describes the mechanism of the oxidative dissolution of sulfide ores in both acidic and alkaline media.

Prior to conducting a systematic study on the seven sulfides, a series of preliminary experiments was performed. These experiments were necessary to develop the required analytical methodology to determine 70

the most suitable sulfide : SOg mole ratios, and to define the

approximate ranges of the various reaction parameters, i.e., the

reaction time, the pressure change, and the thiosulfate and ferrous ion

concentrations. These experiments and the subsequent investigation

showed that, in general, because ferrous sulfide was more reactive than

the six naturally-occurring sulfide ores, the experiments with ferrous

sulfide involved; a) a lower quantity of sulfide (0 , 1 vs. 0 . 2 mole); b) less reaction time ( 1 2 0 vs. 2 1 0 minutes); c) lower reaction pressures but higher pressure changes; and d) higher reaction rates accompanied by higher thiosulfate and ferrous ion concentrations. However, with

regard to the kinetics of the acid decomposition of thiosulfate, the reaction product solutions from both the ferrous sulfide and the naturally-occurring sulfide ore experiments showed the same reaction order and approximately the same reaction rates. These observations are discussed quantitatively in the following sections.

A. Characterization of the Ferrous Sulfide-Sulfur Dioxide Reaction

Two types of experiments were employed to characterize the FeS-SOg reaction. The first type was a series of base-line experiments designed to determine the effect of temperature on the reaction. These base-line experiments, which were conducted at five temperatures between

31.1 and 82.0 °C, established several important relationships among the reaction parameters. In the second type of experiment, three reactant variations were made in the base-line FeS experiments. These variations included the use of: a) 0.1 mole of HgSO^ instead of 0.2 mole of HCl; b) 0.05 mole instead of 0.1 mole of FeS; and c) calciner feed in place 71 of Na^SO^ as the source of SOg. The first two variations determined

the suitability of using an alternate acid and a decreased reductant

concentration in the reaction. The third and most important variation enabled a direct comparison to be made between the idealized base-line experiments and an experiment which represented the actual sulfite reduction process.

Prior to discussing the results of these experiments in detail, it is of interest to note two observations which gave the first qualitative evidence for the formation of Sg via a thiosulfate intermediate. In all reactions using FeS, a green color and a yellowish colloid appeared at the top of the reaction mixture within five minutes after the addition of the acid. Five to ten minutes later, depending on the reaction temperature, the colloid disappeared and the solution became light blue-green — a condition which persisted throughout the remainder of the reaction.

These phenomena may be explained by considering the inherent similarity between the FeS- 8 0 2 and H 2 S-SO2 reactions, and the mechanisms which have been proposed for the HgS-SOg reaction (Chapter II, Section D)

Thus, the following mechanism is proposed to account for the initial products formed in the FeS-S0 2 reaction. Prior to the complete

conversion of the sulfite to SOg via reaction 95, there is sufficient

free acid in solution to react with the FeS to form a small amount of

H 2 S which subsequently reacts with SO2 to form thiosulfuric acid:

FeS (s) + 2H+ — ► HgS + Fe%+ (98)

H 2 S + SO2 — ► [H2 S2 O 2 ] [0]» H2 S2 O3 (99)

The thiosulfate then associates itself with the ferrous ion to form the 72

green-colored metastable ferrous thiosulfate pentahydrate. ^^3 As the temperature of the reaction solution increases, this compound dissociates thereby releasing the thiosulfate which subsequently undergoes acid catalyzed decomposition to form colloidal Sg, Thus, the green color gives way to the blue-green color of ferrous chloride tetrahydrate or ferrous sulfate heptahydrate.^**3 subsequent disappearance of the

Sg colloid may occur through the nucleophilic degradation of the Sg ring by SOg or it may simply deposit on the surface of the FeS, thereby being removed from solution. A proposed mechanism for the surface deposition of Sg will be discussed in Section E.

1. Base-line FeS-S0 2 experiments

The seven base-line FeS experiments were conducted at 31.1, 40.2,

58.3, 69.0, 70.8, 71.9, and 82.0 °C. The characterization curves for each of these experiments are shown in Figures 8 to 14. The summary of results for the various experimental parameters is given in Table VI.

The most useful measure of the progress of the reaction was the change in pressure with time. Figure 15 shows how the pressure-time function gives a unique characterization for five different reaction temperatures.

The pressure-time function also enables the rate of the reaction to be determined through the application of the ideal gas law. Since the pressure, P, is proportional to the amount of SO2 above the reaction solution, [SO2 ], we may write:

PV “ [SÛ23RT (100) where V is the volume of the reaction bottle, R is the gas constant, and T is the reaction temperature. All reactions were conducted at 73

0.3

20

0.2 15 n

g

CM O co 10

5

50 100 150 200

Time, minutes

Figure 8 . FeS-S0 2 reaction characterization at 31.1 “C, experiment 95. 74

0.3

20

IS CO •a

CM O CO .2+ CD 10

5

SO 100 150 200

Time, minutes

Figure 9. FeS-S0 2 reaction characterization at 40.2 °C, experiment 96. 75

0.3

20

0.2 15 •W .2+

e. § 3 c/9 %■ 10

5

50 100

Time, minutes

Figure 10. FeS-S0 2 reaction characterization at 58.3 °C, experiment 94. 76

0.3

20

0.2 15 .2+

& §

CM 3 o CO 10

0.1

5

50 100

Time, minutes

Figure 11. FeS-SOg reaction characterization at 69.0 °C, experiment 90. 77

0.3

20

0.2 15 •a F e‘ 9 c. S sT s I .=» COJ 3 o CO 5" 10 ÏÏ

5

50 100

Time, minutes

Figure 12. FeS-S0 2 reaction characterization at 70.8 “C, experiment 8 8 . 78

20

0.2 15 C3 .2+ e

e

CM 3 O o c/3 n 10 §■

s

50 100

Time, minutes

Figure 13. FeS-SOg reaction characterization at 71.9 “C, experiment 89, 79

20

0.2 rs 15 e Fe‘ I.

3 w §• 10 i

5

50 100

Time, minutes

Figure 14. FeS-SOg reaction characterization at 82.0 “C, experiment 91. 80 ,Ct g a Ü ÜÜ e •H m o o o O o o •H S o> r4 eg o O o VO H H CM tH rH rH rH

X

1 CM + O ON iH m CM VO CM CO CM m VO O CO <" CO <" O 0) CM o o rH rH rH iH rH cn £l

CM rH tH 00 CO CO + m CO m CO e- VO VO CM s o eg 00 00 0) tH rH tH tH H iH H o O OO o O O

i O CO m VO H CM o> 00 tH ON sf OJ m 00 m cn VO CO o O rH CM CM CM CM O a eg o OOO o OO 03

bO (U •H as CO 00 CO rH < CO 1 A o rH m m m

u na 1 c ■P •H CO rH o H VO d + s CO m rH CO CO CJN 0) eg B 0) a o o CM CO CO CO •H I CM U 'd o O (U iH

KM 1 O C • c •H 03 o 1 8 H 00 VO iH ON CO o U •H eg tH <" O CO rH O CM d CO a Ü rt O o o eg U1 m Ü C eg m o 03 O r4 P 'd o 1 0) •H eg 4J CO O 4J rH 2 eg cd c/3 Q) I—I CO 1 n3 i u 4J 00 CO m O CO o o B u d •H •H o CO r4 tH o VO 00 cu rH CM rH CM w % 1 eg CMI O CO 1—I CO ■§ •r4 °CM 03 !—1 CM CO O 00 as O H Ü o rH O 00 ON OH CM CO m VO 00 CA p, O m VO <" O 00 ON rH X Î3 ON ON ON ON 00 00 ON IH M I I 81

20

15 & of

CM e 82.0 CO 10 71.9

58.3

5 40.2

31.1

50 100 150 200

Time, minutes

Figure 15. Composite of the SO2 pressure characterization curves for the FeS-SÛ2 reaction at five temperatures. 82 constant volume and imposing the restriction of constant temperature, this relationship may be simply expressed as;

^V,T “ [80%] (101)

After the peak pressure, the reaction begins as indicated by the decrease in SOg pressure.* Since the initial linear portion of the pressure time curve occurs at nearly constant temperature we may write:

-dP/dt = -dESOgj/dt (102)

Now, the overall reaction may bewritten in general terms as:

SÛ2 (g) + FeS(s) — ► products (i.e., and Fe^*) (103) whose rate equation may be expressed as :

Rate = -d[S02]/dt (104)

Taking into consideration the two reaction products whose concentrations were monitored, we may also write:

-d[S02]/dt oc +d[S2 0 3 ^“ ]/dt

Finally, substituting equation 105 into equation 102, we see that the following relationships between the various rates should exist:

-dP/dt = +d[S203=-3/dt (106)

-dP/dt = +d[FeZ+]/dt (107)

d[Fe=+]/dt = d[S2 0 3 2 -]/dt (108)

Using the data in Table VI, these relationships were plotted and are shown in Figures 16 to 18. A linear least squares regression analysis for each pair of parameters showed that the coefficient of determination, r^, was greater than 0.92 for each relationship. Thus,

The slight reaction prior to the peak pressure did not affect the results. 83

7

99 6

5 106

4

CM CO 3 101

104 2

1

10 20 30

-dP/dt, 10^ psigmin'^

Figure 16. Relationship between -dP/dt and d[S2 0 3 ~ ]/dt for the base­ line FeS-S02 experiments (•), r^ = 0.98; points for the reactant variation experiments (a ) , and the sulfide ore experiments (■) are identified by experiment number. 84

5

4

M 3 92

2

106 99

1

104

■ 105

10 20 30

■dP/dt, 10^ psigmin'^

Figure 17. Relationship between -dP/dt and d[Fe ]/dt for the base-line FeS-S02 experiments (•), r = 0.93; points for the reactant variation experiments (a) , and the sulfide ore experiments (■) are identified by experiment number. 85

7

G

E 5 S 106 en 4

3 S-

104 2

1

105,

1 2 3 4 S

d{Fe2+]/dt, 10^ M min'1

Figure 18. Relationship between d[Fe^*]/dt and dfS^Cy/"]/dt for the base-line FeS-SO^ experiments (•), r^ = 0.92; points for the reactant variation experiments (a ) , and the sulfide ore experiments (■) are identified by experiment number. 8 6 it may be assumed that these parameters are accurate indicators of the reaction rate. As such, they may be used to determine the energy of activation, E^, for the reaction through the application of the integrated form of the Arrhenius equation:

In k = - + In A (109)

The Arrhenius plots of the data for the three rate parameters given in

Table VI are shown in Figures 19 to 21. The corresponding energies of activation were determined by performing a logarithmic least squares regression analysis on the data points,*

Comparing these activation energies, it is seen that those for

-dP/dt and d[S2 0 2 ^"]/dt are very similar, 16.783 and 15,569 kcal mole"^, respectively. This result is not unexpected in view of the high correlation between these two rate parameters and the fact that, according to reaction 99, SO2 reacts directly to form thiosulfate. The lower activation energy of 12.182 kcal mole"^ for the d[Fe^*]/dt parameter indicates that the dissolution of FeS is not the rate controlling step in the overall reaction and occurs first since it requires less energy. Similarly, the higher activation energy of

15.569 kcal mole"^ for the thiosulfate indicates that its formation occurs after the dissolution of the ore. These two facts lead to the conclusion that the overall reaction involves the formation of an intermediate metastable species prior to the formation of thiosulfate.

This is consistent with the mechanism outlined in reaction 99, which

*The results from experiment 95 (31.1 °C) were not included in the Arrhenius plots of -dP/dt and d[S202^']/dt. 87

50

20

10 CL CM

5

2

2.8 2.9 3.0 3.1 3.2

1/T, 1Q3 K 1

Figure 19. Arrhenius plot for the -dP/dt data from the base-line FeS-S0 2 experiments; In k = -16,783/RT + 22.77, = 16.783 kcal mole”\ r^ = 0.98. 8 8

1 S g c

■a

0.5

2.8 2.9 3.0 3.1 3.2

1/T , 10^ K '1

Figure 20. Arrhenius plot for the ]/dt data from the base-line FeS-SOo experiments; In k = -15,569/RT +17.56, E = 15.569 kcal mole- 1 r^ ' = 0.97. - -- 89

5.0

2.0

I S 1.0 *o

^ 0.5

0.2

2.8 2.9 3.0 3.1 3.2 3.3

1/r, 10^ K 1

Figure 21. Arrhenius plot for the d[Fe^*]/dt data from the base-line FeS-SO, experiments; In k = -12,182/RT + 12.15, E = 12.182 kcal mole"', r^ = 0.99. 90 shows the formation of the sulfane sulfoxylic acid intermediate. This mechanism and the activation energies of the FeS-S0 2 reaction may be schematically summarized in the potential energy/mechanism diagram shown in Figure 22.

The activation energies for the dissolution of FeS (12,182 kcal mole~^) and the formation of thiosulfate (15.569 kcal mole”^) compare favorably with those reported in the literature. Recently, Tewari and

Campbell^reported an activation energy of 14.3 ± 1.7 kcal mole"' for the acid dissolution of FeS. Since the formation of thiosulfate is the rate-controlling step in the overall reaction, its energy of activation may be compared with that for similar reaction processes occurring in the leaching of sulfide ores. A recent review of these processes’*^ shows the energies of activation ranging from 6 to 20 kcal mole"'.

Finally, Table VI contains values for the parameters -AP,

[SgOq/"] , [Fe^"*"], [S„0»^“ ] : [Fe^*], and the time at which these L 3 max 2 3 max values occur. Figure 23 shows a very high correlation (r^ = 0.99) between -AP and [S«0_^~] . This relationship is not unexpected — / 3 max as discussed earlier, the pressure is a direct measure of the SO2 in the gas phase and, according to equation 99, SO2 reacts directly to form thiosulfate. Thus, the total pressure drop, -AP, should be directly proportional to the amount of SO2 consumed and the amount of thiosulfate formed.

Additional support for the potential energy/mechanism (Figure 22) may be obtained by examining the ratios given in

Table VI. The data show that this ratio increases with reaction temperature up to 69.0 °C. By comparing the individual thiosulfate 91

B B LU I

S2 O 3 '

Reaction Coordinate

Figure 22. Potential energy diagram for the FeS-S02 reaction showing the activation energies for the dissolution of FeS (E^), and for the formation of thiosulfate (E2 ). 92

101*

0.3

106

104 0.2

i So E S E N PO

0.1

105

98

1 2 3 45 6 7

A P, psig

Figure 23. Relationship between the pressure drop, -AP, and the maximum thiosulfate concentration for the base-line FeS-S0 2 experiments (#), r^ = 0.99; points for the reactant variation experiments (a ) , and the sulfide ore experiments (■) are identified by experiment number. 93 and ferrous ion concentrations, it is seen that the lower ratios are principally due to the lower than expected [SgiO^^"] for the corresponding

[Fe^*]. Thus, it is reasonable to assume that reaction temperatures below 69 °C are sufficiently high to form the intermediate products,

Fe^* and H 2 S2 O2 , but not to form the final product, thiosulfate.

Consequently, at the termination of these experiments, the reaction was incomplete with a large portion of the products remaining in the intermediate state.

Similar qualitative evidence for the two tier nature of the reaction may be seen by noting that, although the ferrous ion concentration is nearly the same at the beginning of each reaction, the time at which thiosulfate first appears decreases as the reaction temperature increases (Figures 8 to 14). The time-temperature dependency of thiosulfate formation is also observed in the reactions conducted above

69.0 °C (Figures 10 to 14) where it is seen that the thiosulfate concentration is initially lower than that of ferrous ion. However, as the reaction proceeds, the thiosulfate concentration increases rapidly and eventually becomes about 1.4 times greater than the ferrous ion

concentration before it begins to decompose into S g . The decomposition of thiosulfate will be discussed in Section D.

2. FeS - Reactant Variation Experiments

The three reactant variation experiments were conducted using

H^SO^, less FeS, and calciner feed in place of the corresponding reactant used in the base-line FeS-S02 experiments. The results of these three experiments are shown in Figures 24 to 26 and the reaction parameters are summarized in Table VII. The first experiment, in which 94

0.3

20

0.2

15 .2+ 0

i 1 D. CM I 8 O 10

0.1

5

50 100

Time, minutes

Figure 24. Effect of using 0.1 mole H2S0^ instead of 0.2 mole HCl in the FeS-SOg reaction at 70.9 “C, experiment 92. Concentration, mole liter 95 CM

o 0 CM O Oi 4-1 (d B O •H 4J O rt lU w CM O % V3 CA 0) P<4 (U Æ 4J B •H en CM (U 0) rH O 8 O rH O (S «W a» 0 in E TJ CO 01 4-1 (0 B •H 0) I-) O 8 m o CD bO B •H (T\ (0 o\ B 4J 04 B O (U 8 4-1 •H U |4 to 0) (U 04 a 04 X M 0)

in CM m )4 a 00 •H k lO o m 8

Djsd 'ajnssajd Zgs 96

0.3

20

0.2 15 'Cl .2+ £a [

(M e V) 10

0.1

5

50 100 ISO

Time, minutes

Figure 26. Effect of using calciner feed in place of Na2S0o in the FeS-S0 2 reaction — a simulated sulfite reduction process at 68.6 “C, experiment 106. 97

V 0) u e G OO o •H•HCMCM o HS r H i H H

X fO S 1 + m r H rs. est 0 0 m Q) O i H r H r H CsJ CO

CM O + 1-4 r H 0 0 X rs. m i n 4) c 4 r H r H CM k O O O O G

I g ë J O r*. N N i n CM m •a* m m I O s CMCM CM CM g c/3 O O o i PO P x •H i n m m < CD 1 CL < T \ o U 1 :

4J g g I •O T e g ‘ + B \G m C M N \o \ o so g 0) X 'm Pw Ow C M r H H o u •OO O w

4J 1 I • O e •H e O r~. • CM ON so X e n m I o M CM « 0) CO o tH • o » I 5 CO * c 4-1 x H e 0 4J (0 » o 0 0 %» m i 0 0 e •fM & 4 m rx- % o o : CL f H H 7 «g O I f 4 ' c g c/3 ON CMVO : H KJ O O CO o ON sO I u Ôl o 0 0) 4J G 0> r H U S e O p 4 O o .o I (0 • H Q < r g C * 0 4J 4.» 8 O CO«H 01 B Ü « CO m 0) U O O I es r H C M o P L r H k CO 0) W • X es i OS o d C g g 4J CO

p , Ô CM ON SO B K z Os O M H J I 98

0.1 mole of HgSO^ was substituted for 0,2 mole of HCl, was designed to determine the suitability of using the more readily available and less expensive sulfuric acid. In the second experiment, the effects of an increased reaction temperature and using 0.05 mole instead of 0.10 mole of FeS were determined. Finally, and most importantly, the third experiment simulated the sulfite reduction process by using calciner feed in place of ^^2 8 0 ^ as the source of sulfur dioxide. This experiment enabled direct comparisons to be made between the mechanism and kinetics of the base-line experiments and those of the sulfite reduction process.

To facilitate the comparison of the results from these three experiments with those of the base-line experiments, each of the parameters is plotted on the appropriate base-line figure. In addition.

Table VIII gives a summary of the qualitative differences between the parameter values for the variation and base-line experiments. Several factors contributed to the differences in the parameter ratios shown in this table. The following discussion gives an interpretation of how these factors led to the observed differences in the reaction parameters for each experiment.

H^SO^. The presence of S0^^“ in the reaction mixture has the overall effect of decreasing each of the rate parameters and increasing the reaction time (see Table VII). However, the dES^Og^^j/dt : d[Fe^*]/dt ratio shows no change and, therefore, the increase in the ratios of each of these rates to -dP/dt must be due to a low -dP/dt, i.e., a decrease in the rate of consumption of SOy,. This decrease in -dP/dt is an indication of the competition between S0 ^^“ and SO2 in the reaction 99

.C(U o BI u g CM PL + 1 O o CO < 4) #» O 1 PL CSJ S i CO M CO ■M d (U •H« .O i s U 3 n) •H 'd u M d ■u " CO d) d) + + u CM + o T + d M en CM + t$ eu O 0) M (U CSJ pL cd 0) (U d CO u > CM •H d tfl o CM H d T3 /a •H 1 cO Q (U CO (U cd r d m u w d m c T3 4J o cd M CM d d H o cd + + 1 1 d -H CM Ph 1 o CM M d (D d •H d q PL 1 M (U •pH •H î i e 4J d (0 CO Cd o CO |3 CO•H A I d) M ,3 CO Cd m 0) CO > 4J < d 4J d) d 4J > cd 1 d +J • r i +j CM + + + ’^ 1•H ' ^ 4J CJ en u g Cd cd O % d) d) u 4J (U CSJ 1 •H Pd CO tH çd d) d d x t ^0) "0) 1 OO O ' 4J 4J 0 > CN| C D d) CO d •XJ eu H u O O 00 . "O o r». OC CD (ü a i a ; M CO cd u d) cd (U & .o •> m u d eu a o eu H M cd •H tH o eu is^î u 4J o 8 d o Cd 8 0 *i-j Td cd "r4 CO CO c/a ü eu Q) M H CH o eu iH eu î.. ël cd • Cd ' h cd IH d) CH d) > o o o CM /a CM d)

S* C3 CH 0 \ vO m g CTi CTi O rH llî cd CO 4J CO ^ “ H cd 00 cd 0 ) (U0 1 0 0 with FeS to form thiosulfate. The net result is reflected in the higher ;-AP ratio which arises from the fact that the same amount of thiosulfate is formed but less SO^ is consumed.

0.05 mole FeS. The effect of using one-half of the 0.10 mole FeS base-line quantity is to substantially reduce the rate of appearance of

Fe^"*", d[Fe^*]/dt. However, the higher reaction temperature and pressure (Figure 25) favor the formation of thiosulfate, thereby increasing dfSgOg^^j/dt. These effects combine to greatly increase the ratio of the two rates. The lower ; -AP ratio is due to the fact that at the higher reaction temperature there is some decomposition of thiosulfate to Sg. Finally, an examination of the data in Table VII shows that the higher ; [Fe^"*"] ratio is due to the lower [Fe^*]. This indicates an increased efficiency in the conversion of reactants to products as would be expected at the higher reaction temperature.

Calciner feed. In most respects, this simulation of the sulfite reduction process gave results nearly identical to those obtained in the base-line experiments. This similarity is exhibited in the

[S,0o^”^ ]__„ nicix : -AP ratio which does not change (Table VIII), and in the same values for the ; [Fe^*] ratio, and the reaction time

(Table VII). These results show that the process proceeds by the same mechanism as that of the FeS-SOg reaction. However, with regard to the kinetics there are two differences: d[Fe^*]/dt and -dP/dt are low when compared to the values obtained in base-line reactions conducted at the same temperature. Thus, an examination of Figures 18 and 16, respectively, shows that in qualitative terms, d[Fe^^]/dt is 1 0 1 considerably lower (— ), and -dP/dt is lower (-) than would be expected, while dfSgOg^^j/dt remains unchanged. These differences are reflected in the parameter ratios shown in Table VIII, Although the reason for the decrease in the two rate parameters was not fully investigated, it is probably due to the Mg^"*" contained in the calciner feed which greatly increases the ionic strength of the reaction solution. The result is a reduced solubility of SOg in the solution which decreases both the rate of FeS dissolution and the rate of consumption of SOg. The fact that dESgOg^^j/dt remains unchanged may be attributed to the presence of a considerable amount (12%) of S0^^“ in the calciner feed which reacts to form thiosulfate.

In summary, the results from the three reactant variation experiments indicate that the development of the most economic and technically feasible reaction conditions for the sulfite reduction process can include the use of HgSO^, a reduced quantity of FeS, and a higher reaction temperature.

B. Characterization of the Sulfide Ore-Sulfur Dioxide Reaction

The potential utilization of the sulfite reduction process in recovering elemental sulfur from spent flue gas desulfurization sludges depends on both the availability and reactivity of the sulfide ores which could be used as reductants in the process. Thus, it is necessary to determine the reactivity of those ores whose abundance is relatively large and from which essential metal values could be recovered through the application of hydrometallurgical techniques. As a first step toward the selection of suitable candidate ores, the following discussion 1 0 2 examines the reactivities of six naturally-occurring sulfide ores

through an analysis of the reaction rates and mechanism as obtained from the characterization of the MeS-SOg reaction.

The six ores evaluated in these experiments were: galena (PbS), pyrrhotite (FeS), sphalerite (ZnS), chalcopyrite (CuFeSg), pyrite

(FeSg), and chalcocite (CugS). Since preliminary experiments showed

that, with the exception of pyrrhotite, these ores were less reactive than the ferrous sulfide used in base-line experiments, 0.2 mole of each ore was used instead of 0.1 mole. However, 0.1 mole of the pyrrhotite was used to facilitate a direct comparison with the base-line

FeS experiments. All experiments were conducted at a reaction temperature near 75 °C. This temperature was chosen for two reasons:

1) it exceeded the temperature required for the formation of thiosulfate; and 2) it was the same temperature at which a majority of the base-line data were collected thus enabling the two sets of results to be easily related.

The results from the six MeS-SOg experiments are shown in Figures

27 to 32 and summarized in Table IX. Experiments are arranged in order of decreasing ore reactivity, based on the -dP/dt parameter. However, it should be noted that this order does not apply to FeS (pyrrhotite) since only half as much ore was used in this experiment. In fact, the results from this experiment indicate that if 0.2 mole were used, pyrrhotite would be the most reactive ore. This is indeed the case as will be discussed in Section D.

As seen in Table IX, not all of the rate parameters and

concentrations were determined. For PbS, ZnS, and CUgS this was due to 103

0.3

20

0.2 n 15 e s CD S

3 CM O S 10

5

50 100 150 200

Time, minutes

Figure 27. PbS-SOg reaction characterization at 73.2 ®C, experiment 101. 104

0.3

20

0.2

IS n •a e 2 » Ia. I 3 s & 10

0.1

.2+

5

SO 100 ISO 200

Time, minutes

Figure 28. FeS (pyrrhotite)-S0 2 reaction characterization at 77.4 °C, experiment 104. 105

0.3

20

0.2 IS CO

2

00 CD

10 CD

S

50 100 150 200

Time, minutes

Figure 29. ZnS-S0 2 reaction characterization at 75.0 °C, experiment 98. 106

0.3

20

0.2 15 o n e s

C/3 10

5

.2+

50 100 150 200

Time, minutes

Figure 30. CuFeS2 “S0 2 reaction characterization at 76.5 °C, experiment 105. 107

0.3

20

0.2 15 n

2 s

CM 8 10

5

.2+

50 100 150 200

Time, minutes

Figure 31. FeS2 “S0 2 reaction characterization at 75.3 “C, experiment 107. 108

20

IS & I CM e V» ID

5

50 100 ISO 200

Time, minutes

Figure 32. CugS-SOg reaction characterization at 75.0 “C, experiment 103. 109

u Oi •ovO *a fH •d *o •O E ■He rHo o\ rH rHq MOOrH H E CM rH CM CM CM CM

Se rH 1 + 00 m 0) ( I 1 OCN Cx CM rH CO

CM tn m + nO vO CO Z 1 r~. 1 m o 1 0 O o o w o o d

X J vOm rHo OOS mas c m rH O VO vO o Z m CM o o 1 1 d P4 «H o CM m o m O < O 2 1 & \D CM CM o O CO o 0 ) 0) & 3 Q) X u P 9 •o 7c •OQ) 0 U3 M •H m œ 0 ) + 8 O

PbS. An examination of the data points in Figure 23 shows that this experiment produced both the highest and greatest pressure drop. In addition, although fluctuations in the thiosulfate concentrations occurred during the experiment (Figure 27), no uniform decrease in concentration was observed. Instead, the concentration was still increasing at the end of the 210 minute experiment. These results are indeed odd in view of the data obtained in similar base-line experiments (e.g., Figure 11) and the fact that the dESgOg^^j/dt : -dP/dt ratio increased only slightly while the t^2^3^~^max ' ratio showed no change. One might at least expect that the higher values for

^^2^3^~^max -AP would be accompanied by an increase in the two rates associated with these parameters. But the data show that both of the rates were less than would be expected for a reaction temperature of

73.2 °C.

These unusual observations are due to a positive salt effect which is exerted on the reaction by the large doubly-positive plumbous ion. Ill

eu rd )4 X 44 o CO 8 O d 1 ' 44 o CM p 44 p < 1 U o 1 eu + CH P C/ p eo (U eu 14 > o Xi X 44 CO CO •H CO 4J > r 4J T3 •u G "d 44 -4- CO (U d + 0) 8 1 •H u 'H CM + O d M P CM + P g) eu (U O (U > 14 p. p p cO O (U K W 44 X CH W T3 CO CO m -d ■d •H (U p d cO rH •H 44 0) iH C*4 d Xi 1 o CO w m 4-1 U CO 'd d •H • % «44 CO 44 o 44 o p ■d •H •H 1 + 44 d 1 4J t 3 CM p 1 •H M -rX d d eu 9 CO •H eu CO p 1 o CO & 8 p O i CO (U •d rH + CO >4 > QJ rH 44 •H 4J P 8 44 o 0) C44 T3 •H d > rH ■d M CO eu •H d 44 (U (U U u 4J en 1 T3 d P *H d CO CM \ •H X 44 eu 4J (U P P o s eu *H eo •H fl O g M d P rH 4J CH 1 (U eu 00 cO CO P 44 d ' H d d (U •H eo eu o* (U 'd ■d rH X 1 d p g eu (U d 4J w eo CO o (U (U CO p p & P « 0) eu "d CH eo eu 1 d •H eu en P eu X ^ d 1*4 V4 (U ü 44 rH o g 'È P d » d d d 0) M O eu en u %4 O H d 0) 44 eu C*H P 1*4 eu •H d - eu d •H u H O eu H 00 p eu d OO o > d rH •H o d p 44 •H o O CO eo d 44 eo •rH eu H eu rH U U d çg 00 cO rH d (U eu cO O' 14 d > CO HO 1 1 2

This effect and the presence of excess SOg prevent the decomposition of thiosulfate into elemental sulfur via the following mechanism. First, as a result of the positive salt effect, the negatively charged thiosulfate is attracted to the vicinity of the formed from the dissolution of PbS. The resulting local deviations from the average bulk anion concentration are similar to those observed by La Mer^^^ for

La^* (Chapter II, Section F). In the present case, the anions are probably more loosely associated to Pb^* than they would be to La^* due to the lower charge and larger size (1.20 vs. 1.016 A) of Pb^* compared to La^*. However, the association is sufficiently strong to overcome the weaker attractive forces between the ore surface and thiosulfate which lead to the deposition of Sg on the ore surface (see Section E of this chapter).

Second, the presence of excess SOg (i.e., high pressure) prevents the build-up of the long chain sulfane disulfonic acids which are precursors of Sg:

H0gS-S^-S0gH+ (n-l)H2S0g:#: HOgS-S-SOgH + (n-iyHgSgOg (110)

Since the equilibrium of this reaction is shifted to the right, no thiosulfate decomposition occurs (i.e., -dfSgOg^'j/dt = 0); this is reflected in the slightly increased ^ESgOg^^j/dt :-dP/dt ratio. Finally, since no Sg forms, the ore surface does not become occluded by the deposition of Sg on its surface and, therefore, the reaction continues unimpeded until the termination of the experiment. This is shown in

Figure 27, where it is seen that at the end of the experiment d[S20g^“ ]/dt is positive and dP/dt is negative. These unique results provide the basis for the observed high values of [SgOg^"]^^ and -AP. 113

FeS (pyrrhotite). This experiment provides another example of the effect of high SOg pressure on the equilibrium of reaction 110. By comparing the three rate parameter ratios in Table X, it can be deduced that since dESgOg^^j/dt : -dP/dt remains unchanged, the high dESgOg^^j/dt : dEFe^*]/dt ratio is due to a decreased FeS dissolution rate (i.e., low dEFe^*]/dt). As a result the overall rate of reaction is reduced as reflected in the lower values for both dESgOg^'J/dt and

-dP/dt (Figure 16). Despite the low reaction rates, the high SOg pressure shifts the equilibrium of reaction 110 to the right, thereby contributing to the increased dESgOg^^j/dt : dEFe^*]/dt ratio and the high ES?Og^"]maY" However, there is some decomposition of thiosulfate as reflected in the decreased ESIgOg^^l^^x • "AP ratio. Finally, the high ESgOg^" ' EFe^"*"] ratio (Table IX) is due to the combined effects of a low EFe^*] and high ESgOg^"]^^^.

ZnS and CuFeS2 . Both of these ores were relatively unreactive, giving dESgOg^^j/dt : -dP/dt ratios very similar to those obtained in base-line experiments conducted at 31.1 and 40.2 °C. The lower

ESgOg^"]^^^ : -AP for the ZnS experiment is due to adsorption of SOg on the ore surface without reaction to form thiosulfate (see Section C).

This increases the total pressure drop relative to the amount of thiosulfate formed.

The parameter ratios for the CuFeSg experiment show changes which are similar to but, due to lower reactivity, less than those for the

FeS (pyrrhotite). Thus, because dEFe^*]/dt is low, dESgOg^^j/dt : dEFe^*]/dt increases, and dEFe^*]/dt : -dP/dt decreases, while the product of these two ratios remains: unchanged (Table X). The high 114 pressure prevents the decomposition of thiosulfate thereby increasing the [8 2 ,0 2 ^"] : [Fe^*] ratio shown in Table IX. Finally, the adsorption of SOg on the ore surface without reaction results in a lower

FeSg and CugS. These ores were the least reactive of all the ores tested as indicated by the high SO^ pressures attained during the reaction: 20.2 and 19.9 psig, respectively (see Figures 31 and 32).

Due to their low reactivity, the thiosulfate concentrations in these two reactions were not monitored. Therefore, the only conclusion which can be made is that, based on the -dP/dt values shown in Table IX, these ores were very unreactive toward SO^.

In summary, the naturally-occurring sulfide ores, which are less reactive than ferrous sulfide, may be arranged in the following order of reactivity based on the -dP/dt rate parameter and the use of 0.2 mole of each sulfide: FeS (pyrrhotite) > PbS > ZnS > CuFeS2 > FeSg > CUgS.

Furthermore, although the results of these experiments show that the metal sulfide-sulfur dioxide reaction occurs by the same mechanism as the FeS-S0 2 reaction, low reactivity, a high SO2 pressure during the reaction, and the adsorption of SO^ on the ore surface produce changes in both the rate and concentration parameters. Finally, the positive salt effect observed in the PbS-S0 2 reaction has provided some indication as to how the occlusion of the ore surface by Sg may be prevented. This will be discussed further in Section E. 115

C. Comparison of Sulfide Reactivities

One of the major rate controlling factors in a heterogeneous

reaction, such as the MeS-SOg, reaction, is the surface area of the

reacting solid phase. Other factors which may affect the rate include

the mass transfer of the gaseous reactant into the aqueous phase and the

transport and diffusion of the reactant and product species to and from

the solid/solution interface. These various processes will be used in

the development of a reaction model given in Section E; the remainder of

this section will discuss the effects of ore surface area on the

MeS-SOg reaction rates and the adsorption of SOg on the ore surface.

In order to minimize the effect of surface area as a variable in

the comparison of reaction rates, each ore used in this research was

pulverized and screened to the same screen distribution (see Chapter III)

Assuming this provided ore samples with comparable surface area, then

the ore volume can be substituted as a direct measure of its surface

area.* Accordingly, the volume of ore used in each experiment was

calculated from the actual mass of ore used and its bulk density.

This enabled values for two additional parameters to be calculated which aided in the interpretation and verification of the experimental

results: 1) a reaction rate normalized to the amount of ore used; and

2) the amount of SOg adsorbed per unit volume of ore.

A tabulation of ore volumes, normalized reaction rates, and SOg

adsorption data is given in Table XI. The normalized rate for each

*A rigorous treatment would require the accurate determination of surface area by BET low temperature nitrogen adsorption; however, the use of the ore volume as a substitute parameter proved more than adequate for the semi-quantitative interpretations given here. 116

0 o 0) O -2 IP o g CO 00 Mf CO o VO CO G\ 00 r4 CO O' tH CM o iH H &0 Id < •H r4 r4 H CM o o o o O O (0 (U (N P. O 3 W CO CO (U eu

(U CO U Q) O SJ O m CO a 1 CO Q) 00 VO CM ov r4 VO UO OV CÜ m *H I % M 0) CO CO r4 CM O o iH O O d g PM 44 p . •H a csj 44 o o CM»H œ O O ctf CO cu 0) M-4 ei (N 0 0) o O *H Id 4J IW p . 5 ë a CO CO 0) 03 4J OJ < 0) cd 1 X N Pd 4J o0) *H 13 00 H O 00 CO 00 O -û tH a e cr» CO 00 ov CO CM CM O o 44 4J cd o pü Ü 1 O •H *3 00 m CO CO tH o O O o CO 03 g 4J 1 r4 CO CO a o Ü I U cd tz; cd 2 O o A 0) a 44 o 44 UO W 1 u g) VO CO VO VO O VO <" 00 CO (U gm 00 o\ CO C30 o> 00 VO g •H ë a g rH O rH rH rH VO 00 m 2

acd Ü 0) CM -w o •H CO CO T3 4J ■s (U g I 0) § H & cs •H i i CM CO IP o 0 CO CO CO ! I

(Tables VI, VII, and IX) by the corresponding ore volume. A comparison can then be made between the base-line rate and the rates for the reactant variation and sulfide ore experiments. For the reactant variation experiments, the use of 0.05 mole of FeS at a reaction temperature 10 ®C higher than the base-line experiment, which used

0.10 moles of FeS, results in a doubling of the normalized rate.

Although this result is unexpected in view of the fact that one-half as much FeS was used in this experiment, it indicates that the 10 °C increase in temperature and higher SO^ pressure (see Figure 25) were the rate determining factors in this experiment.

The lower rates for the calciner feed and HgSO^ experiments, as discussed in Section A, are due, respectively, to the lower solubility of SOg in the calciner feed reaction mixture and the formation of thiosulfate from both S0^^“ and SO^ in the HgSO^ medium. The net effect in both cases is to reduce -dP/dt and the normalized rates by amounts proportional to the magnitude of the cause; thus, the presence of sulfate inhibits the reaction to a greater extent than does the solubility of SOg in the reaction mixture (Table XI).

For the sulfide ore experiments the data show the same order of reactivity as given in the previous section; FeS (pyrrhotite) > PbS >

ZnS > CuFeSg > FeSg > CugS. The normalized rates may also be used to determine the quantity of ore required to obtain a particular reaction rate. Thus, if one wished to obtain a reaction rate equal to that of the base-line experiment, the following multiplicative increases in each ore volume would be required; 2.7, 7.1, 19.1, 39.0, 112, <».

Correspondingly, the reciprocal of these factors gives a reactivity 118 scale with the FeS base-line reactivity equal to one. If the base-line reactivity is set equal to 100, then the following numerical reactivity scale applies:

FeS FeS (pyrrhotite) PbS ZnS CuFeSg FeSg CUgS

100 37.5 14.1 5.2 2.6 < 0.1 0

This scale shows that only FeS (pyrrhotite) and PbS are sufficiently reactive at 75 °C to be considered as candidate ores for the sulfite reduction process. However, an increase in the reaction temperature by

30 °C should increase the reaction rates of ZnS and CuFeSg by approximately eight-fold, thereby making these two ores suitable candidates at higher temperatures than those used here in the comparison study. FeSg and CugS appear to be insufficiently reactive

to be useful at any temperature compatible with low-pressure aqueous solution reactions.

The adsorption of SO^ on the ore surface can be estimated from the

SOg pressure vs. temperature characterization curves shown in Figures 33

to 36. Each figure shows an SO^ calibration curve which was determined

from the reaction of Na^SO^ with HCl under the same experimental

conditions as those used for all reactions. The curves for each

experiment show how the SOg pressure increases during the period when

the reaction solution is being heated to the reaction temperature. As

such, these curves characterize each experiment during the induction period, i.e., the time prior to the beginning of the reaction, at which

point the SO^ pressure begins to level off.

If it is assumed that there is no reaction between the metal sulfide

and SOg during the induction period, then the lower reaction curve

pressures relative to the SOg calibration curve pressures may be 119

SOo Calibration 20

FeS ( pyrrhotite ): 77.4 *0

0.05 mole FeS 90.2 *C 15

oT

Base-line: 82.0 *C CM e CO 10 Caiciner feed : 68.6 *0 i-line: 71.9 *0

5

25 50 75

temperature, "C

Figure 33. SO2 pressure vs. temperature characterization for the base­ line FeS, reactant variation, and FeS (pyrrhotite) experiments showing the adsorption of SO2 on the ore surface prior to the reaction. 1 2 0

SO- Calibration 20

IS PbS

g o ' Vi 10

5

25 SO 7S

Temperature, "C

Figure 34. SO2 pressure temperature characterization for the PbS experiment showing the adsorption of SO2 on the ore surface prior to the reaction. 1 2 1

SO- Calibration 20 CuFeS. ZnS

15

CM O CO 10

5

25 50 75

Temperature, "C

Figure 35. SOo pressure temperature characterization for the CuFeS2 ana ZnS experiments showing the adsorption of SO2 on the ore surface prior to the reaction. 1 2 2

SO- Calibration 20 FeS, CUjS

15 â fi-

CM O C/S 10

5

25 50 75

Temperature, °C

Figure 36. SO2 pressure vs. temperature characterization for the FeS2 and CU2 S experiments showing the adsorption of SO2 on the ore surface prior to the reaction. 123

attributed to the adsorption of SO2 on the ore surface. Using the ideal

gas law this may be expressed by the following relationship:

" % , T “ (111) Thus, at constant volume and temperature, the difference between the

SO2 calibration curve pressure and reaction curve pressure is proportional to the amount of SOg adsorbed on the ore surface. For the

current analysis, pressure differences were determined at 50 °C; all

experiments were conducted at a constant volume of 200 ml. It should be noted that the assumption that no reaction occurs during the

induction period is not entirely correct. This can be seen in Figure 33 where the slightly skewed nature of the reaction curves relative to the

SO2 calibration curve is apparent. However, at 50 ° C this effect is minimized and only a slight, constant error is introduced into the

analysis.

The SO2 pressure difference for each experiment and the

corresponding values for the SO2 adsorbed per unit volume of ore are

given in Table XI. An examination of these data reveal several

important relationships which provide additional support for the

differences in the reaction parameters discussed in the previous two

sections.

A direct, unqualified comparison can be made between the data from

the reactant variation and the base-line experiments since the FeS used

in these experiments was taken from the same batch. Thus, it is not

unexpected that the use of one-half the amount of FeS (experiment 99)

results in one-half the SOg pressure difference and the same amount of

SO2 adsorbed per cm^. The lower pressure difference observed in the

calciner feed experiment reflects the reduced solubility of SOg in 124 the calciner feed reaction mixture; thus, the amount of SOg adsorbed per cm^ is reduced. Finally, the higher pressure difference for the

HgSO^ experiment is due to the increased solubility of SO2 in HgSO^, thereby producing an apparent increase in the surface adsorption of SOg.

The less reactive sulfide ores exhibited a considerably reduced

SOg adsorption as shown in the decreased pressure differences and in the SOg adsorbed per cm^. Although FeS (pyrrhotite) and PbS produced

SOg adsorptions which were proportional to their reaction rates, both

ZnS and CuFeSg exhibited larger pressure differences than would be expected from their low reaction rates. For ZnS this was due to the relatively large amount of SO^ adsorbed per cm^; for CuFeSg, the large ore volume used in the experiment accounted for the increase in SOg adsorbed. Thus, in both of these ores, SO^ adsorption occurs but very little reaction takes place as indicated by their low reaction rates. For FeS2 and Cu^S, the data show that FeS2 adsorbs less SO2 despite the fact that it is more reactive. Finally, it is interesting to note that, as shown in Figure 36, at higher temperatures SO2 is desorbed from the surface of these ores as indicated by the convergence of the calibration and reaction curves.

D. Kinetics of the Acid Decomposition of Thiosulfate

In those reactions conducted above a reaction temperature of 68 °C

(Figures 11 to 14, 25, 26 and 28), the thiosulfate characterization curves show a peak in the concentration followed by a decrease which is approximately logarithmic with time. The results discussed below show that this decrease is due to the acid decomposition of thiosulfate into 125 elemental sulfur. This interpretation was anticipated from an observation made in a previous study^® where it was found that heating the FeS-S02 reaction-product filtrates resulted in the formation of Sg.

Although the determination of reaction rates was not a part of that study, as discussed in Chapter II, Section F, the combined rate equations of La Mer^^^ and Keller^show that:

- = k[S20g=-]3/2 (112)

Thus, by monitoring the thiosulfate concentration in the reaction supernatant as a function of time, it can be determined through a comparison of reaction orders if the mechanism given by Davis^^® to explain the above equations, applies to the MeS-SOg reaction.

Accordingly, constant temperature (40 ± 0.05 ®C) kinetic studies were conducted on the reaction supernatant from all reactions in which greater than 0.0600 mole of thiosulfate was formed.

The results of nine experiments which satisfy the above concentration criteria are shown in Figures 37 to 45; the rate data from which the figures were constructed are given in Appendix 3. Table

XII gives a summary of the kinetic results. For each experiment, both the rate constant, k, and reaction order, n, were determined by log-log least squares regression analysis. The results show that, although there is some variation in the reaction order, the average reaction order is 3/2. Furthermore, comparing this average with the individual values, particularly those obtained in the reactant variation experiments (106 and 92) and the sulfide ore experiments (101 and 104), shows that most experiments follow the 3/2 order reaction kinetics.

Consequently, the kinetics are similar to those observed by La Mer and 126

I O)

I

0.5

2 5 10 20 SO

Thiosulfate concentration, 10^ mole liter

Figure 37. Decomposition of thiosulfate in the FeS/Na2 S0 3 /HCl reaction product supernatant from experiment 96 at 40 "C; k = 1.067 liter^ mole“^hr”^, n = 2.38, r^ = 0.95. 127

E

QC

0.5

2 510 20 50

Thiosulfate concentration, 10^ mole liter '

Figure 38. Decomposition of thiosulfate in the FeS/Na2 S0 g/HCl reaction product supernatant from experiment 94 at 40°C; k = 1.012 liter^ mole“^hr”\ n = 1.57, r^ = 0.76. 128

20

10

5 a e E

2 K

1

0.5

2 5 10 20 50

Thiosulfate concentration, 10^ mole liter'^

Figure 39. Decomposition of thiosulfate in the F^:S/Na2 S0 g/HCl reaction product supernatant from experiment 90 at 40°C; k = 1.006 liter^ mole“^hr“^, n = 1 .2 2 , r^ = 0 .8 8 . 129

20

ID

u 5

E

K 2

1

0.5

2 5 10 20 50

Thiosulfate concentration, 10^ mole1 liter

Figure 40. Decomposition of thiosulfate in the FeS/Na2S0g/HCl reaction product supernatant from experiment 89 at 40° C; k = 1.008 liter^ mole“ ^hr“^, n = 1.43, r^ = 0.94. 130

Q)

Si S

0.5

2 S 10 20 50

Thiosulfate concentration, 10^ mole liter'^

Figure 41. Decomposition of thiosulfate in the FeS/NaoSOo/HCl reaction product supernatant from experiment 91 at 40 C; k = 1.010 liter^ mole“ ^hr ^, n = 1.46, r^ = 0.68. 131

O) e E

I

0.5

2 5 10 20 50

Thiosulfate concentration, 10^ mole liter'^

Figure 42. Decomposition of thiosulfate in the FeS/Na2 S0 3 /H2 S0 ^ reaction product supernatant from experiment 92 at 40°C; k = 1.008 liter^ mole“^hr"^, n = 1.39, r^ = 0.92. 132

01

01 O E

I

0.5

2 5 10 20 50

Thiosulfate concentration, 10^ mole liter'^

Figure 43. Decomposition of thiosulfate in the FeS/calciner feed/HCl reaction product supernatant from experiment 106 at 40 °C. k = 1.013 liter^ mole“^ hr”^, n = 1.45, r^ = 0.96. 133

20

10

JS 1 5

E

I 2

1

0.5

2 5 10 20 50

Thiosulfate concentration, 10^ mole liter'^

Figure 44. Decomposition of thiosulfate in the PbS/Na2S0g/HCl reaction product supernatant from experiment 101 at 40 °C, k = 1.008 liter^ mole“^hr“^, n = 1.58, r^ = 0.87. 134

10 •

1

I

2 S 10 20 SO

Thiosulfate concentration, 10^ mole liter'^

Figure 45. Decomposition of thiosulfate in the FeS (pyrrhotite)/Na2S0g/ HCl reaction product supernatant from experiment 104 at 40 °C; k = 1.003 liter* mole“*hr“\ n = 1.03, r* = 0.82. 135

m VO 00 00 CM VO CH as 00 as VO OV Ov 00 00 o o o o o o O o o

00 CNJ CO VC G\ m 00 CO o cn m CvJ CO O m CVJ rH rH rH rH rH rH rH rH Pi o

1^ C\| VO 00 00 CO 00 CO m o 4J O VD H o O O rH O O rH CM CO iH o o o O O O O O O O rH I—I rH I-N 0) u 4J 0) I MCÜ T-j 4J

rH rH CJ CJ X s, g 'd rH CO 2 0) CJ o O 0) Pd CO rH rH rH rH rH CO PM CVJ O O o o U CVJ CO Cd I rd w « I u o a p rH 0) CO CO CO CO CO CO CO 0 CVJ 0) o o o o o o •fH Cd JJ § CO CO CO CO CO CO Ü z : •H 4J (N CNJ CVJ CVJ CvJ CNJ rH 4J I•H CJ Cd cd cd Cd Cd cd cd O Cd W Cd s 53 % 53 % CJ 0 X O Q) u 14J (S CO CO CO CO CO CO CO rH u CO 0) 0) 0) 0) (1) a) Q) cd to k Pm Pm Pm Pm PM O PM

VO

Keller, and it may be concluded that the mechanism is the same as that derived by Davis (Chapter II, Section F). According to this mechanism, the decomposition of thiosulfate into Sg involves intermediate polysulfane monosulfonic acids (or polythionates). This mechanism will be used in the next section to develop a model for the oxidative leaching of sulfide ores.

Finally, due to the large variation in the rate constants for the base-line FeS/Na2 S0 g/HCl experiments, no conclusions can be made regarding the effects of particular ions on the reaction rate. Thus, additional experiments will be required to determine why the positive salt effect, previously observed in the PbS-S0 2 reaction (Section B), was not observed in the decomposition of thiosulfate (experiment 1 0 1 ).

One would expect that, in the absence of SO^, the Pb^* would increase the rate constant relative to the other experiments.

E. Mechanism of the Oxidative Dissolution of Sulfide Ores

The results of this study may be summarized in a chemical model which describes the mechanism by which metal sulfides undergo oxidative dissolution in acid solution. Since a similar mechanism takes place in alkaline solution, the model described here applies equally well to both acid and alkaline leaching processes. The importance of this model to both the sulfite reduction process and the hydrometallurgical leaching processes can be seen by the fact that heretofore no explanation has been offered for the mechanism by which elemental sulfur forms and subsequently occludes the surface of sulfide ores undergoing oxidative dissolution in acid solutions. Since the occlusion of the ore surface 137 retards the dissolution rate, it presents an economic and technological roadblock to the development of any oxidative leaching-type process which produces elemental sulfur. Thus, the understanding of the leaching mechanism is necessarily the first step in controlling surface occlusion and defining the most favorable conditions under which leaching should be conducted.

The typical heterogeneous reaction involves a series of simultaneous and successive steps. These include the following:

• Mass transfer of the gaseous reactant from the gas phase into

the aqueous phase

• Transport of the dissolved reactants to the solution/solid

interface of the sulfide bearing particles

• Adsorption of the reactants on the solid interface

• Reaction on the surface

• Diffusion of various species through the particle matrix

• Desorption of the reaction products from the surface

• Transport of the reaction products into the body of the

solution.

These steps have provided the basis for several leaching models.

Forward and Mackiw®^ proposed a model for the alkaline dissolution of sulfide ores based on their determination that thiosulfate and sulfane disulfonates were the major products of the ammonia oxygen-pressure leaching of pentlandite (FeNiS). Wadsworth®® has reviewed several models, which have been developed for the acid dissolution of sulfide ores. Although these models fit the experimentally observed nonlinear

(i.e., parabolic) and mixed (linear and nonlinear) kinetics which result 138 from the occlusion of the ore surface by elemental sulfur, they do not explain the mechanism by which elemental sulfur forms and subsequently occludes the ore surface.

The model presented here explains both of these mechanisms. It is based primarily on the model developed by Forward,®® and the results obtained in this and previous research.^® In addition, two assumptions were made in developing the model. These are: 1) the surface of colloidal Sg is negatively charged; and 2) the surface of the ore is positively charged. Regarding the first assumption, surface charge on colloids is a common phenomenon which results from an intrinsic charge

(not predicated here), the adsorption of ions from solution, or the dissociation of molecules at the colloid surface. In the case of colloidal Sg, Pauli®**® has observed that the colloid is not composed entirely of Sg but also of thiosulfuric and polysulfane sulfonic acids which dissociate at the surface giving rise to a negative charge as shown in Figure 46. Thus, colloidal Sg, designated here as Sg", would be attracted to a positively charged species.

With regard to the second assumption, in a previous study on the sulfite reduction process,*® it was observed that the Sg product was formed or deposited on the surface of the FeS during the reaction. In fact, the increase in weight of the FeS during the reaction was used to quantitatively determine the amount of Sg formed. Whether the adsorption of Sg on the surface of the FeS is electrostatically induced or is the result of other surface phenomena remains to be proven.

However, for the purposes of this discussion, the surface of the ore is assumed to have a net positive charge, in order that colloidal Sg" species adsorb strongly on its surface. 139

H2 Û H*

H'

Figure 46. Dissociation of thiosulfuric and the sulfane sulfonic acids on the surface of colloidal 140

The model for the oxidative dissolution of metal sulfides in acid and alkaline solution is shown in Figure 47. Although the model is based on O 2 as the oxidative agent, the same mechanisms apply to SOg as well as to any other oxidant used in the presence of oxygen, including those discussed in Appendix 2, i.e., Fe2 (S0 ^)g, FeClg, CI2 , HNOg,

K2 Cr2 Üy and NaOCl. The following discussion centers mainly on how this model applies to FeS. However, a brief description is given of the differences encountered in the generalized case for both the acid and alkaline oxidative dissolution of metal sulfides.

The first step in the oxidative dissolution process for the

FeS-S0 2 system is the mass transfer of SO2 from the gas phase into the aqueous phase. Sulfur dioxide dissolves readily in water (at 15 °C,

45 vol. SO2 in 1 vol. H 2 O) to give a rather complicated reaction mixture. This mixture, which is commonly called "sulfurous acid", consists primarily of physically dissolved SO^, loosely hydrated SO2 ,

"disulfurous acid" (H2 S2 0 ^) , and the ions H^O"*", HS2 0 ^” and HSOg“ .

However, undissociated sulfurous acid, HgSOg, and disulfurous acid do not exist in detectable amounts in aqueous solutions of sulfur dioxide.

Thus, the dissolution of SOg in HgO may be represented by the following equilibria:

SO2 (g) ^ SO2 (aq) (113)

SOg (aq) + HgO SO^'HgO "H2 SO3 " (114)

"HgSOg" :;±: HSO^" + H"*" (115)

HSO3 - + SO2 HS 2 O5 - (116)

Since each of these species has an electronegative character, any one could be adsorbed on the surface of the ferrous sulfide. For the 141

§

/\ 0 CO 1 jS CO IS'” \/ T3 o ^ O

Ü Ü CO_ o CO a a « H *o

0 ] CO iH CJ 4JCO ^4J s ,

M -l CO O C S O CO • H M -l 4J iH 3 3 H CO O O CO 1 -1 iHCO ,W 3 t) 0) % ^ g

CO g 0 01 5 1 S ’® M -l "i4 O H V TS0) cO O (U H - ilCO 0 0 6 m

si- 0) I 142 present discussion sulfur dioxide will be used; however, similar reactions may be written for the other species.

The adsorption of SOg on the surface of the FeS and the subsequent reaction between SO^ and FeS results in the dissolution of the FeS with the formation of ferrous ion and the sulfane sulfoxylic acid intermediate. This intermediate subsequently reacts with HSO^" to form thiosulfate which may diffuse out into solution or remain within the reaction zone and react with sulfoxylate to form disulfane disulfonic acid:

FeS + SOg (aq) FeS'SOg (ads) — ► Fe%+ + [SgOg^"] ( 117)

[SgOgZ"] + 2H+ HgSgOg (118)

HgSgOg + HSOg- H 2 S2 O3 + HSOg" (119)

HgSgOg + HSO2 " HO3 S-S-SOH + 0H“ (120)

HOgS-S-SOH + HSOg" HO^S-S-S-SO^H + OH” (121)

Subsequently, the chain length of the disulfonic acids is increased by the nucleophilic displacement of a sulfite group by thiosulfuric acid:

HOgS-Sg-SOgH + HS-SO3H HO3S-S3-SO3H + H2SO3 ( 122)

HO3S-S3-SO3H + HS-SO3H HO3S-S4-SO3H + H2SO3 ( 123) and in general:

H0 3 S-S^-S0 3 H + HS-SO3 H H 0 3 S-S^+3 -S0 3 H + H 2 SO3 (124)

Of course, these equilibria are not shifted to the right until the thiosulfuric acid concentration becomes high enough to counter the "k degrading effect of the sulfurous acid.

*Sulfurous acid is used here to designate physically dissolved or loosely hydrated SO2 . 143

The increasing concentration of polysulfane disulfonic acids in

the presence of high concentrations of thiosulfuric acid leads to the production of elemental sulfur via two mechanisms:

• The thiosulfate-catalyzed decomposition of the higher sulfane

disulfonates into lower disulfonates and the sulfane

monosulfonates, followed by a building up of the monosulfonate

ions into polysulfane monosulfonates through a series of

consecutive nucleophilic displacements

e The acid-catalyzed conversion of thiosulfate to sulfur and

sulfurous acid by a series of displacements of sulfite ion by

thiosulfate.

The first of these mechanisms is given by the following set of equations

(for trisulfane disulfonic acid):

HOgS-S^^f HOg^S^-S-SOgH :;±: HO3 S-S-S-SO3 H + ”8 -8 2 0 3 %: (125)

HO3 8 2 -S- + H 8 -8 2 O3 H HO3 8 2 -S2 ” + H 8 -8 O 3 H (126)

^°3®2“^2~"'^^^^2°3® ^ HO3 S2 -S3 ” + H 8 -8 O3 H (127)

and so forth until internal displacement produces ring closure and

elemental sulfur:

'S^sps-S-SO^^ — » 8 g + “8 2 O3 H (128)

The second mechanism involves the acid catalyzed conversion of

thiosulfate to elemental sulfur and sulfurous acid by a series of

displacements of sulfite by thiosulfate:

+ H'*' HSSO3 ” (129)

“0 o8-8”' T ^ 8 ^ 0 o~ ^ H8-8-80q” + 80.=” » H80,“ (130) ■J J j j fast d 144 and so on, until

HSgSOg” — ► Sg + HSO3 " (132)

The subsequent occlusion of the ore surface occurs as shown in

Figure 47. The ore surface, which has a net positive charge, acts as a large positive ion would act in the positive salt effect discussed in

Section B. Thus, the negatively charged species are attracted to the vicinity of an ore particle where, through the previously described reactions, Sg is formed. Of course, Sg also forms the body of the reaction solution with the resulting production of colloidal Sg” . In either case, the negatively charged Sg” colloid is attracted to the surface of the positively charged ore surface where it deposits, thereby resulting in surface occlusion. As a result, the surface area of the ore is reduced and dissolution takes place at a reduced rate.

Eventually, the ore surface becomes completely occluded and dissolution

terminates.

For the general oxidative dissolution in acid solution shown in

Figure 47, similar reactions take place:

MeS + O2 — ► Me^+ + [S0 2 ^"] ^ ^ M e ^ + + [SgiO^^"] (133)

[SgiO^Z”] + 2H'*' H 2 S2 O 2 (134)

H 2 S2 O2 + I/2 O2 — ► H 2 S2 O 3 (135)

The subsequent formation of the polysulfane sulfonic acids and elemental

sulfur proceeds by the same reactions as given for the dissolution of

ferrous sulfide (reactions 120 to 132). Finally, the occlusion of the

ore surface by Sg occurs by the same process as described for FeS

(Figure 47). 145

The prevention of surface occlusion is exemplified by the PbS-SOg reaction (Experiment 101, Figure 27). As noted in the discussion of this experiment (Section B), the presence of the large doubly-charged plumbous ion exerted a positive salt effect on the reaction. Thus, thiosulfate and the other negatively charged polysulfane sulfonates were attracted principally to the Pb^* (aq) and to a lesser degree to the ore surface. Furthermore, due to the high concentrations of SOg in the solution, elemental sulfur did not form. Similar effects undoubtably occur with other large inorganic or organic cations under a variety of conditions and concentrations of all interacting species. Consequently, exploitation of the positive salt effect observed by La Mer and apparently active here, appears to be a viable method for preventing surface occlusion by Sg, thereby maintaining the linear dissolution kinetics required for the commercial scale application of hydrometallurgical leaching processes.

The alkaline dissolution of sulfides proceeds by a series of reactions analogous to those which occur in acid media. The following reaction scheme shows the formation of thiosulfate and disulfane disulfonate which is followed by the formation of the polysulfane sulfonates :

MeS + O 2 — Me=+ + [802^"]Me^+ +[ 8 2 0 2 ^"] (136)

[802^"] + 2 H 2 O — ► H 2 SO2 + 20H" (137)

[SgiOg:-] + I/2 O2 S2 O 3 :- (138)

SgOjZ” + H 2 SO2 — ► "O3 8 - 8 -8 -OH + oh" (139)

“O3 8 -8 -8 OH + SOgZ" — ► “O3 8 -8 -8 -8 O 3 " + OH" (140)

"0 3 8 - 8 "'^ "0 3 8 ^ 8 ^ 0 3 " “O 3 8 -8 -8 -8 -8 O3 " + 8 0 3 %" (141) 146

ZSgOjZ" tpi: “O3 S-S-S" + SO]:" (142)

" O g S g - S ^ "OgS^S^S-SOg" "O3 S-S-S-S-SO3 " + SgOjS" (143)

^3 ^2 "^ + 8 ^ 2 0 3 tpi: ^3^2"^2 ^2^3^ (144)

~°3®2~®2~^^5^2°3~ ~°3^2"^3~ ®2°3^~ (145) The major difference between alkaline and acid dissolution is that in alkaline solution the products cannot decompose into elemental sulfur. Instead, in the presence of increasing amounts of sulfite formed in reactions 141 and 142, all products are degraded to thiosulfate:

(n—1) O3 8 + 0 3 8 —8 ^ — 8 0 3 O 3 8 —8 —8 O 3 + (n—1) 8 —8 O 3 (146)

O3 8 + 8 —8 j^—8 O 3 1 O 3 8 2 + ®~^n—1~^^3 (147)

Finally, in the presence of heat and oxygen, sulfate is formed:

^2°3^~ ^°2 + 20H" — ► 280^2- + (148)

8 0 3 ^" + I/2 O2 S0 ^^“ (149)

Thus, all products in the alkaline oxidative dissolution of sulfides are soluble and there is no occlusion of the ore surface by elemental sulfur. As a result, dissolution follows linear kinetics, that is, the rate of ore dissolution is constant with time and is a simple function of ore surface area, oxygen pressure and/or other oxidant. This is indeed the case, as observed in several studies."^®' 147

CHAPTER V

CONCLUSIONS

The characterization of the metal sulfide-sulfur dioxide reaction under a variety of experimental conditions has shown that the sulfite reduction process for recovering elemental sulfur from spent flue gas desulfurization sludges takes place by a mechanism involving the formation of thiosulfate and sulfane sulfonic acid intermediates which subsequently decompose into elemental sulfur. The most favorable process conditions include: a reaction temperature near 100 °C, SOg pressures between 15 and 20 psig, and a stoichiometric amount of ferrous sulfide. In addition, the less expensive sulfuric acid may be used in place of hydrochloric acid.

Several naturally-occurring sulfide ores were found to be suitable for use in the reduction process, although each was less reactive than ferrous sulfide. Four ores were sufficiently reactive to warrant consideration as suitable candidates for use in the process: FeS

(pyrrhotite), PbS (galena), ZnS (sphalerite), and CuFeSg (chalcopyrite).

Accordingly, each of these ores is also suitable for use in an integrated regenerable FGD-hydrometallurgical pressure leaching process which would eliminate the need for the hydrogen sulfide and the other methane-based reducing gases currently used in regenerable FGD processes.

Finally, a chemical model has been developed for the oxidative dissolution of sulfide ores in both acid and alkaline solutions. This 148 model explains the mechanism of ore dissolution as applicable to both

the sulfite reduction process and all oxidative hydrometallurgical leaching processes. In addition, the model has provided the basis for explaining the mechanism by which elemental sulfur forms and subsequently occludes the ore surface thereby preventing complete dissolution. This mechanism has revealed a relatively simple method by which surface occlusion may be controlled thereby significantly improving the technical and economic outlook for hydrometallurgical pressure leaching processes. 149

CHAPTER VI

FUTURE RESEARCH

This work has disclosed many, yet unexplored, research avenues in the complex but intriguing field of sulfur chemistry. Several research studies which would aid in the further elucidation and verification of the mechanisms involved in the metal sulfide-sulfur dioxide reaction are summarized below:

• An identification of individual sulfane sulfonates and the

determination of their concentrations as a function of time

in the acid decomposition of thiosulfate

• A determination of the effect of ore surface area, as measured

by BET low temperature nitrogen adsorption, on the MeS-SOg,

reaction rates and surface adsorption of SOg

• A study of the MeS-SOg reaction rates and thiosulfate

decomposition rates at temperatures up to 120 °C

• A determination of the SOg concentration as a function of

MeS-SOg reaction time and its effect -on the rates of thiosulfate

decomposition

• A characterization of the MeS-SOg reaction at constant pressure

• A comprehensive study of the influence of the positive salt

effect on ore occlusion and reaction kinetics 150

# A determination of the yield of elemental sulfur in the sulfite

reduction process using sludges from other non-regenerable FGD

processes.

In addition to providing data directly related to the metal sulfide- sulfur dioxide reaction, each of these studies will provide much needed information on the chemistry of those sulfide-sulfur dioxide reactions which are closely related to the methods currently used for controlling sulfur dioxide emissions into the atmosphere. 151

APPENDIX 1

DESCRIPTION OF FIRST AND SECOND

GENERATION FLUE GAS DESULFURIZATION PROCESSES

Although several reviews and descriptions of flue gas

desulfurization (FGD) processes can be found in the literature, no single

source gives a general description of the ten processes which comprise

the first generation of non-regenerable and regenerable FGD. This

review fulfills this need and, in addition, exemplifies the chemistry

involved in FGD processes. A similar review is given for those second

generation regenerable FGD processes which use hydrogen sulfide as a

reducing agent to convert adsorbed sulfur dioxide to elemental sulfur

in aqueous media.

A. Non-Regenerable Flue Gas Desulfurization Processes

1. Lime and/or Limestone Scrubbing^®

Stack gases are scrubbed with a recirculating slurry of lime

(calcined limestone), limestone, or a lime/limestone mixture. The SOg

reacts with the slurry forming calcium sulfite and calcium sulfate

(sludge) which is drawn off for disposal in a settling pond. Since

limestone is less reactive than lime, more limestone is required to

obtain high SO2 removal efficiencies. However, the use of lime results 152 in the increased formation of deposits (scaling) in the scrubber and higher costs for the construction and operation of a lime kiln. The major problems with the lime/limestone processes are: process control, scrubber scaling, corrosion-erosion of various parts of the system, elimination of mist from the scrubbed stack gases, reheating the stack gases to provide adequate plume bouyancy, and sludge disposal. Some of these problems have been solved through ongoing test and evaluation programs. More extensive programs are underway to eliminate the remaining problems. Although the lime/limestone processes do not recover SOg in a useful form, they have the advantages of relative simplicity, low investment, and no by-product marketing. For these reasons lime/limestone FGD is favored by the electric power industry as the best first generation SO2 control technique.

2. Double Alkali Scrubbing^®

This three-step wet scrubbing process involves adsorption, precipitation, and filtration; two alkalis are used, one for adsorption and one for precipitation. In the first step, SO2 is adsorbed in a recycled solution of NaOH, 0 ^ 2 8 0 3 and Na 2 C0 3 to yield a spent liquor containing Na2 S0 3 , NaHS0 3 and Na2 8 0 ^. The spent liquor is sent to two chemical mix tanks where it is mixed with a slurry of calcium hydroxide

(slaked lime) or limestone. This precipitates calcium sulfite and sulfate (sludge) and regenerates the NaOH. The sludge/slurry is then sent to two reactor clarifiera where a high molecular weight anionic polymer is added to aid precipitation. Excess calcium ion is precipitated in the second reactor clarifier by adding Na2 C0 3 or COg. 153

The clarified alkaline liquor is then recycled to the scrubber. Sludge is removed from the bottom of both reactor clarifiera, and, prior to disposal, it is vacuum filtered to remove water and any soluble salts which are returned to the scrubber. Although this process produces a sludge waste product, no sludge disposal pond is required. Furthermore, the scrubber scaling and corrosion-erosion problems inherent in the lime/limestone processes are greatly reduced.

B. Regenerable Flue Gas Desulfurization Processes

1. Sodium Carbonate Scrubbing^®

In this process a spray-dryer serves as the reaction chamber for the removal of SO2 from the hot flue gases. Here SO2 reacts with an atomized solution of Na2 C0 3 to form a dry powder containing NagSOg, NagSO^, and unreacted NagCO^. The powder is subsequently removed from the gas using a cyclone separator followed by an electrostatic precipitator, thereby producing a clean flue gas. The powdered salts are combined,mixed with coke, and conveyed to a high-temperature, ceramic-lined molten salt reducer where, at 2000 °F, the coke reacts with the molten sulfite and sulfate salts to produce HgS and COg. The molten mixture is then passed to a quench tank where it is dissolved in H^O. After the aqueous solution is cooled and filtered, it is precarbonated with the COg evolved from the decomposer. Additional recarbonation occurs in a bicarbonate crystallizer using CO^ from the reducer off-gas. The resulting aqueous

NagCOg/NaHCOg solution is sent to a decomposer where COg is thermally stripped from the solution and a pure solution of NSgCOg is regenerated 154 for return to the scrubber. Gases evolved from both the precarbonator and crystallizer are rich in H 2 S and are combined for subsequent recovery of elemental sulfur in a Claus plant. Tail gases from the Claus plant are returned to the spray-dryer for final clean-up. Although the total process eliminates the scaling, plugging, and sludge disposal problems encountered in the lime-limestone processes, several problems have been encountered in the operation of the molten salt reducer. These problems are currently under study.

Magnesium Oxide Scrubbing 19

Sulfur dioxide is adsorbed in a slurry of magnesium oxide,forming insoluble magnesium sulfite. The scrubbing liquor is then centrifuged to separate the sulfite crystals which are dried and calcined in the presence of a reducing agent (CO or Hg). This regenerates the MgO and produces a concentrated ( 1 0 to 16 per cent) stream of SOg which may be processed into sulfuric acid, liquid sulfur dioxide, or elemental sulfur. Since the magnesium oxide is recycled, it must be protected from contamination by efficiently (> 99 per cent) removing fly ash from the flue gas prior to the desulfurization step. A major advantage of this process is that the dried magnesium sulfite may be shipped to a regeneration facility for processing into the final product; as a result,

the power plant is not required to deal in chemical processing and marketing. However, the process is not without its problems, some of which are similar to those encountered in lime/limestone scrubbing.

They include; the erosion of pumps, valves, and piping; excessive wear on the internal parts of the centrifuge; and the need for stack gas reheat. 155

3. Wellman-Lord Process^**

Sulfur dioxide is adsorbed in a solution of sodium sulfite,

converting the sulfite to bisulfite. The spent adsorbing solution is

then sent to a separate vessel where thermal stripping (evaporation)

removes the SOg from the solution and regenerates the NagSOg. The

resulting dilute SOg gas stream can be used as feed to a sulfuric acid

plant. Small quantities of sodium sulfate, a potentially marketable waste product, are also produced. However, to obtain the benefits of

the Wellman-Lord process, a power company must either go into the

sulfuric acid business by building its own plant or it must be located

near one. The equipment for this process is also more expensive to

install than a lime or limestone scrubber.

4. Cat-Ox Process^**

Flue gases at about 500 °C are passed through a hybrid mechanical

filter/electrostatic precipitator. The cleaned gas is then sent to a

vanadium pentoxide catalyst bed where the SO2 is oxidized to SO3 .

Following cooling the SO3 laden gases enter an adsorbing tower where the

SO3 is adsorbed in a 75-80 percent sulfuric acid solution, a small

portion of which is periodically drawn off, cooled, and sent to storage

tanks. A mist eliminator removes any entrained sulfuric acid mist prior

to discharging the cleaned gas to a stack. Although this process

requires large capital expenditures and produces sulfuric acid which is

of marginal commercial value, no raw material other than the flue gas

is required. 156

5. Dilute Acid Scrubbing^®

This process, also kpown as the Chiyoda Thoroughbred 101 process, is a three-step wet scrubbing operation involving adsorption, oxidation, and crystallization. The SOg is adsorbed in a dilute acid solution where it is oxidized to sulfuric acid in the presence of oxygen and ferric sulfate catalyst. The sulfuric acid product stream is split: some of the acid returns to the scrubber and the rest goes to a crystallizer where it is reacted with pulverized limestone to precipitate the sulfate as hydrated calcium sulfate (gypsum). The slurry is centrifuged to separate the gypsum which is then dried, stored, and subsequently sold for use in wallboard and as a retarder for Portland cement. The overflow from the crystallizer is sent to a clarifier and then recycled to the adsorption and oxidation steps. This process was developed in Japan where a market exists for gypsum. However, no such market exists in the U.S., and, by

1978, the Japanese market is expected to become saturated.

6 . Activated Carbon Adsorption^^

This process involves three basic steps: 1) adsorption of dilute

SOg on char (activated carbon); 2 ) regeneration of the char with the resultant production of SOgi and 3) reduction of SOg to elemental sulfur.

In the adsorption step, flue gases are passed through a tower where a counterflowing bed of char adsorbs SOg, Og and HgO vapor. The adsorbed

SO2 then reacts with the Og and HgO to form HgSO^ which is retained in the pores of the char pellets. As the char becomes saturated it is removed from the adsorbing tower and sent to a regenerator vessel where it is heated to 1200 “F in an inert atmosphere. At this temperature. 157 the reduction of HgSO^ to SO^ yields a concentrated ( 2 0 per cent) stream of SOg. The SOg-rich gas is passed through a vessel containing crushed coal where, at 1500 °F, SOg is reduced by carbon to gaseous sulfur.

Finally, the product sulfur is condensed from the reactor effluent and

the tail gas is recycled to the adsorber to remove any residual sulfur,

thus providing a closed-loop system. This system has several advantages over the conventional wet scrubbing operations. These are: increased reliability as it handles no slurries, requires no pH control, and there is no scaling or plugging; no stack gas reheating is required, and a marketable by-product (pulfur) is produced.

7. APCI/IFP Ammonia Scrubbing®®’ ®°

This process, developed jointly by Air Products and Chemicals, Inc.,

(APCI) and the Institut Français du Petrole (IFP), involves the

countercurrent scrubbing of SO2 laden flue gases with a solution of ammonium sulfite. The resulting ammonium bisulfite is processed in a

forced circulation evaporator where the SO2 is liberated. The SO2 is

then sent to a high temperature catalytic reactor where it is reduced

to HgS by hydrogen or carbon monoxide. The reactor effluent gas stream, which contains H 2 S and SO 2 in a 2:1 mole ratio at 160 °C is sent to an

IFP "aqueous Claus" reactor. Here the gases enter near the bottom of

the reactor vessel and countercurrently contact a special solvent/catalyst

system to form elemental sulfur. The molten sulfur product is drawn off

from the bottom of the unit while any ammonia in the exit gases is

recovered and returned to the scrubber. 158

8 . Stauffer's "Aquaclaus" Process®^'

This process is another adaptation of the "aqueous Claus" reaction.

The SO2 laden flue gas is fed to a scrubber containing a solution of sodium hydrogen phosphate (NaOH and HgPO^ at pH = 4) which adsorbs the

SO2 as sodium bisulfite while forming sodium dihydrogen phosphate. The spent scrubbing solution is then sent to a reactor where the adsorbed

SO2 is reacted with HgS at pH = 2.8 to 4.5 to form elemental sulfur and regenerate the sodium hydrogen phosphate. After the elemental sulfur is separated from the phosphate solution, the latter is recirculated to the scrubber.

9. Citrate Process®®"®®

The S0 2 ~bearing gas is first cooled to between 45 and 65 °C and cleaned to remove sulfuric acid mist and particulate matter. The gas is then sent to a countercurrent scrubber where the SO2 is adsorbed in a solution of sodium citrate, citric acid, and sodium thiosulfate. Here the citrate solution acts as a buffer, maintaining the pH around 4, thereby increasing the SO 2 adsorption capacity; the thiosulfate forms a complex with the SO 2 preventing its oxidation to sulfate. The SO2 laden solution is then pumped to a stirred reactor, where at atmospheric pressure and 65 °C, it is reduced to elemental sulfur by H 2 S, thereby regenerating the adsorbing solution. Prior to recycling this solution the sulfur is removed by oil flotation and melting. The H 2 S requirement is met by reacting natural gas, sulfur and steam at 700 “C in the presence of a catalyst. Other reducing gases, such as hydrogen and/or carbon monoxide, may be used in place of methane. 159

APPENDIX 2

HYDROMETALLURGICAL LEACHING REACTIONS

The following review is not intended to give a detailed description

of hydrometallurgical processes — this may be obtained by consulting

specific references. Rather, its purpose is to demonstrate the variety

of oxidants which may be used to oxidize the sulfide in ores to elemental

sulfur under acidic conditions or thiosulfate under alkaline conditions.

In conjunction with the reviews in the literature,this review

represents a current state-of-the-art of leaching chemistry. If

comparisons are made between some of the reactions given here and others

reported in the literature, some inconsistencies will be found. This is

apparently due to the lack of understanding of the mechanisms involved

in leaching. This is especially true for the mechanism involved in the

formation of elemental sulfur, Sg. No author has referred to elemental

sulfur in his leaching equations as Sg, let alone attempted to explain how it is formed. Thus, much remains yet to be understood about the mechanisms involved in leaching chemistry.

A. Acid Leaching

Unless otherwise noted all of the following reactions were conducted

under mild oxidizing conditions favorable to the formation of elemental

sulfur: use of a mild oxidant, low pressures (< 150 psig Og), moderate

temperatures ( < 120 °C), and pH < 3. 160

1. Ferrie Sulfate

Thomas®® has investigated the kinetics of dissolution ofchalcocite in aqueous acidic ferric sulfate solutions and assumed thefollowing reactions to be occurring;

CugS + Fe 2 (3 0 ^ ) 3 — *► CUSO4 + CuS + 2FeS0^ (IB)

CuS + Fb 2 (3 0 ^ ) 3 —► CuSO^ + 2FeS0^ + S (2B)

In a process described by Pearson,®® simultaneous leaching and electrowinning of copper from sulfide ores is performed in an electrolytic cell:

Cu 2 3 + 2 2 ^ 2 (3 0 ^ ) 3 + H 2 O + 3 /2 O 2 — 2CuS0^ + 4FeS0^ + H^SO^ (3B)

0 0 3 8 + Fe2(30^)3 —► CuSO^ + 2FeS0^ + CuS (4B)

CuS + Fe 2 (S0 ^ ) 3 —► CuSO^ + 2FeS0^ + S (5B)

CuS + I/ 2 O 2 + H 2 SO4 — ► CuSO^ + H 2 O + S (6 B)

Cu^+ + 2e“ —► Cu (7B)

In a study of chalcopyrite leaching, Dutrizac®® found that

stoichiometric yields of sulfur and ferrous iron were obtained according

to the reaction:

CuFeSg + 2Fe2(30^)3 —► CuSO^ + 5FeS0^ + 2S (8 B)

Finally, Subramanian®® and co-workers have developed a process for

leaching pyrrhotite according to:

FeS + 2 ^ 2 (0 0 4 ) 3 —► 3FeS0^ + S (9B)

Sulfur separation after leaching was performed in a Soxhlet apparatus

using CS2 . However, the authors indicate that sulfur separation could

also be performed by solvent extraction with perchloroethylene or by

volatilization. 161

2. Ferrie Chloride

The acid ferric chloride leaching of CuFeSg and PbS has been studied by Haver®® ® and co-workers who report that over 99 per cent of the sulfur can be extracted from the leach residue with an ammonium sulfide solution:

CuFeSg + SFeClg — ► CuCl + 4 FeCl 2 + 2S (lOB)

PbS + 2FeCl3 — ► PbCl2 + ZlFeClg + S (IIB)

Subramanian,®^ Van Weert,®^ and their co-workers have developed processes to leach pyrrhotite based on the reaction:

FeS + 2 FeCl 3 — ► 3 FeCl2 + S (12B) although, at a pH near 0, as much as 80 per cent of the sulfide sulfur may be converted to HgS.

3. Cupric Chloride

The use of cupric chloride as a leachant has only been recently investigated. Demarthe®® has proposed a new process in which CuFeS2 is leached in an electrochemical cell at a potential of 600 mV :

CuFeS2 + 3 CuCl2 — ► 4CuCl + FeCl 2 + 2S (13B)

Cathro®® has proposed a somewhat different process in which CuFeS2 is first "activated" with elemental sulfur to convert the chalcopyrite into more readily soluble sulfides:

5CuFeS2 + 4S — ► Cu^FeSg + 4FeS2 (14B)

Cu^FeSg + yCuClg —*► 12CuCl + FeClg + 68 (15B) 162

4. Chlorine

Groves and Smith®^ have studied the leaching of several sulfide ores with chlorine produced by reacting sodium hypochlorite with hydrochloric acid. Although they observed elemental sulfur as an intermediate product, it was oxidized to sulfate during the leach:

NaOCl + 2HC1 — Clg + NaCl + HgO (16B)

CuS + Clg — ► CuClg + S (17B)

CUgS + Clg — ► 2CuCl + S (18B)

CuFeSgf S/ZClg-^CuClg + FeCl^ + 2S (19B)

2 FeS 2 + CI 2 — ► FeCl2 + FeS + 3S (2GB)

S + 3 CI2 + 4 H 2 O — ► H^SO^ + 6HC1 (21B)

5. Sulfuric Acid/Oxygen

Downes and Bruce®® have described a process for the recovery of elemental sulfur from pyrite and pyrrhotite which involves the thermal decomposition of pyrite followed by leaching the pyrrhotite under 150 psig O 2 . The following overall reactions occur:

FeS 2 FeS + S (22B)

4FeS + 3 O 2 — ► 2 2 ^ 2 0 3 + 4S (23B)

To account for the formation of elemental sulfur in reaction 23B they proposed the following mechanism :

FeS + 2O2 — ► FeSO^ (24B)

6FeS0^ + 3/2O2 — *► 22^2(80^)3 + Fe2Ü3 (25B)

Fe2 (304)3 + 3H2O — 2^ 20^ + 38^ 30^ (26B)

FeS + H2S0^ — ► FeSO^ + H 2 S (27B) 163

ZHgS + Og ZHgO + 2S (28B)

HgS + FegCSO^y^ —► 2FeS0^ + HgSO^ + S (29B)

Kunda et al.®® proposed a slightly different mechanism for reaction 23B to explain the formation of the various iron compounds in the leaching process :

FeS + HgSO^ — ► FeSO^ + H^S (30B)

FeS + FegCSO^Xg — ► 3FeS0^ + S (31B)

2FeS0^ + l/ZOg + H^SO^ —^ FegCSO,). +H^O (32B)

FegCSO^)^ + 2 H 2 O — ► 2Fe(0H)S0^ + 8 2 8 0 ^ (33B)

Fe2 (S0 ^)g + 6 H 2 O — 2Fe(0H)g + 3 H 2 S0 ^ (34B)

H 2 S + I/2 O2 — S + H 2 O (35B)

Forward and Veltman have reported^® the extraction of 99 per cent of the metals from galena and sphalerite concentrates according to the following equations:

PbS + H^SO^ -f- I/2 O 2 — ► PbSO^ + S + H 2 O (36B)

ZnS + H^SO^ + I/ 2 O 2 — ► ZnSO^ + S + H^O (37B)

Veltman and co-workers^® have developed an acid pressure leaching process in which copper is extracted from chalcocite and the sulfide sulfur is converted into elemental sulfur:

CugS + ZHgSO^ + O 2 — ► 2CuS0^ + S + 28^0 (38B)

Finally, Vizsolyi et al. have developed avery similar process for the leaching of chalcopyrite concentrates:

CuFeS2 + H 2 SO4 + 5 /4 O 2 + I/2 H 2 O — ► CUSO4 + FeCOH)^ + 2S (39B)

The elemental sulfur is separated from the concentrate by hot filtration, solvent extraction, or distillation. 164

6 . Hydrochloric Acid/Oxygen

Exploratory tests conducted by Kunda and co-workers^^ have shown that sulfidic copper, iron, and zinc concentrates can be oxidized in the presence of various chloride salts which appear to accelerate ore dissolution. Although most of the sulfide sulfur was oxidized to the elemental form, as much as 30 per cent was oxidized to sulfate: The most important reactions taking place are:

CuFeSg + 3HC1 + iVAOg — ► CuCl + FeClg + 2S + 3 /2 H 2 O (40B)

CuS + HCl + 1/40% — ► CuCl + S + I/2 H 2 O (41B)

FeS + 2HC1 + I/2 O 2 — ► FeCl2 + S + H^O (42B)

ZnS + 2HC1 + I/2 O 2 — ► ZnCl2 + S + H 2 O (43B)

7. Nitric Acid/Sulfuric Acid

Prater,^** Bjorling,’® and their co-workers have developed a process for the dissolution of copper concentrates in nitric acid. Sulfide is converted to elemental sulfur and iron is precipitated as basic iron sulfate. Although part of the elemental sulfur is oxidized to sulfate and some NO 2 is formed, the major products may be represented by the equation:

eCuFeSg + lOHNO^ + lOHgSO^ — ►

6 CUSO4 + lONO + 1 2 s + 3Fe202'4S02*9H20 + SHgO (44B)

8 . Oxygen

The reactions reported for oxygen leaching in acidic solutions®^ are identical to those for 8 :2 8 0 ^/0 2 » This is because H 2 S0 ^ is formed in oxygen pressure leaching (see reactions 24B to 39B). However, one study 165 which takes only oxygen into consideration in developing a kinetic model

is of interest. Bartlett^® has developed a computer model for

predicting the rate of oxidative leaching of copper and iron from

low-grade sulfide ores based on the following equations:

CuFeS2 + II/4 O 2 — ► CUSO4 + l/2Fe2Û3 + S (45B)

FeS 2 + 31/120% + 1/2H%0 — ►

1/3F62(S0^)3 + 1/9[1 ^ 3 (8 0 4 )2 (0 8 )3 -2 8 2 0 ] + 7/98 (46B)

9. Bichromate

In a study to identify alternative oxidizing agents suitable for

leaching,Baur et al.^^ have reported that dichromate oxidized

chalcopyrite according to:

3CuFe82 + + 288'*’ — ►

3Cu%+ + 3Fe%+ + 4CrS+ + 68 +I 4 8 2 O (47B)

10. Anodic dissolution/Electrolysis

Habishi^® has reviewed the electrometallurgy of metal sulfides,

emphasizing the importance of anodic dissolution reactions. 8e reported

that sulfides produce elemental sulfur according to the generalized

reaction:

Me8 —► Me^* + 8 + 2e“ (48B)

with deposition of the metal at the cathode:

Me^* + 2e“ —► Me (49B)

These general reactions havebeen applied tothe anodicdissolution of pyrrhotite,®^ chalcopyrite,**^ and to the electrolysis of galena suspensions.^* In addition, these reactions have been used as an electrochemical model for the dissolution of metal sulfides. 166

B. Alkaline Leaching

Unlike acid leaching, the leaching of sulfide ores in alkali has been conducted under a variety of temperatures, pressures and pH values.

Thus, the leaching conditions are reported separately for each oxidant.

It is important to note that the following studies established that thiosulfate and sulfane disulfonates are formed as intermediates in the leaching process.

1. Ammonia/Oxygen

In a series of leaching experiments conducted at 50-105 °C, 10-40 psig Og, and pH - 9, Forward and Mackiw®* were the first to observe the formation of thiosulfate, S„0_^“, and the sulfane disulfonate, S 0 / j 3 6 in the ammonia/oxygen pressure leaching of a mixture of pentlandite, chalcopyrite, pyrrhotite, and pyrite ores. They recognized the fact that sulfide is not converted in a single step to sulfate, but follows through a series of reactions in which the first product is SgOg^", this being further oxidized to S^O^^” , which in turn is oxidized to sulfate. Furthermore, they found that by increasing the temperature or oxygen pressure the oxidation of and is accelerated thereby decreasing their maximum concentrations, the time period during which high concentrations exist, and the ratio. These phenomena were summarized in the followingequations:

NiS-FeS + 3FeS + 70^ + lONH^ + 48^0 —

N i (8 8 3 ) ^ 3 0 4 + ZFegO^'HgO + 2 (8 8 4 ) 2 8 3 0 3 (50B)

2 (8 8 4 ) 2 8 2 0 3 + 2 O 2 — («8 4 )2 8 3 0 ^ + (8 8 4 ) 2 8 0 4 (51B)

(8 8 4 )2 8 3 0 ^ + 2 O 2 + 4 NH 3 + HgO — ► 8 8 4 .8 O 3 . 8 8 2 + 2 (8 8 4 ) 3 8 0 4 (52B) 167

Consistent with these observations. Forward and Mackiw also proposed the first model to describe the mechanism by which sulfide ores dissolve. This model is discussed in Chapter IV, Section E.

Vizsolyi et al.®^ observed very similar results under the same leaching conditions in their study of the ammonia/oxygen pressure leaching of galena. They gave the following overall reaction:

2PbS + 40^ + 2 NH 3 + HgO — ► PbO.PbSO^ + (8 8 4 ) 3 8 0 4 (53B)

In addition, the leaching of galena in ammonium sulfate solutions was investigated at 110 “C and 20 psig 0^. The p8 of the leaching solution increased from an initial value of 4.9 to 8 . 8 after 30 minutes and finally decreased to 7.7 after 3 hours of leaching. Under these conditions a small amount of elemental sulfur was initially formed and later oxidized to thiosulfate when the p8 increased to 8 .8 . The following reactions were given for the dissolution mechanism and to account for the increase in p8 with time:

PbS + 2 O 3 + 2 8 3 O — ► P b (08) 3 + 8 3 SO4 (54B)

Pb(08)3 + (8 8 4 ) 3 3 0 4 PbS04 + 288^ + 28^0 (55B)

Although the authors recognized the formation of thiosulfate and sulfane disulfonates, they incorrectly stated that during dissolution of the PbS, "the sulfur, released as S^“ , must first oxidize to S, then to and only after a considerable time to S^^." Theinaccuracy of this oxidation sequence is proven in Chapter II, Section D.

2. Potassium 8ydroxide/0xygen

Dresher et al.®^ investigated the kinetics of the leaching of molybdenite in the temperature range 100 to 175 °C, 0 to 700 psig 168 oxygen, and pH > 10. Under these conditions molybdenite was found to leach according to linear kinetics. The authors also monitored the thiosulfate as a function of time and found, as Vizsolyi et al. had, that it increases rapidly and then decreases slowly. However, they offered no explanation for this phenomenon but indicated that thiosulfate was directly oxidized to sulfate, thus refuting the proposition that sulfane disulfonates were present in the leaching solution. The following reactions were given to explain their observations:

MoSg + 5 /2 O 2 + 40H“ — ► MoO^z- + SgOgZ- + (56B)

SgOgZ" + 20^ + 20H" — ► 2 3 0 4 ^" + H 3 O (57B)

3. 3odium Hypochlorite/3odium Hydroxide

Cox and 3chellinger®° developed a process for leaching molybdenite from low grade ores at 65 °C, atmospheric pressure, and pH = 10. They gave the following reaction, based on the fact that thiosulfate was observed as a product:

Mo3g + 7NaOCl + 4e" — ► 8 0 0 4 ^" + 8 2 0 3 ^“ + 7NaCl (58B)

Bhappu and co-workers®^ have essentially refuted the results of all investigators who report the presence of thiosulfate in alkaline leaching solutions. They show that molybdenite leaches according to the reaction:

Mo32 + 9C10" + 60H" — *► 8 0 0 4 ^“ + 2 3 0 4 ^" + 9 Cl" + IHgO (59B) in which the oxidation of sulfide goes to completion, with sulfate as the only end product. They go further and state that thiosulfate and other intermediate oxidation states of sulfur probably do not exist in basic media, or, if so, only for infinitesimal time periods. 169

APPENDIX 3

KINETIC DATA

The following tables give the kinetic data collected in the constant temperature (40 °C) thiosulfate decomposition studies. These data were used in the construction of Figures 37 to 45 and in the calculation of the rate constants (k), reaction orders (n), and coefficients of determination (r^) given in Table XI. The reaction velocity, v) = dc/dt, is given for each concentration.

A. Base-line Experiments

Experiment 96 (FeS/NagSOg/HCl)

c, mole liter”^ t, hr U, 1 0 ** mole liter ^ hr”^

0.1441 140.82 5.492

0.1279 170.32 5.834

0.1136 194.83 3.616

0.0884 264.52 2.226

0.0823 291.92 1.817

0.0789 310.63 1.322

0.0757 334.83 - 170

Experiment 94 (FeS/NagSOg/HCl)

c, mole liter"^ t, hr U, 1 0 ** mole liter” ^ hr” ^

0.2045 53.18 8.406

0.1832 78.52 7.283

0.1497 124.52 6.667

0.1333 149.12 6.767

0.1171 173.06 3.866

0.0973 224.28 4.870

0.0759 268.22 2.958

0.0688 292.22 0.965

0.0641 340.95 -

Experiment 90 (FeS/NagSOg/HCl)

0.2236 8.70 9.102

0.2115 32.98 8.865

0.1766 72.35 7.970

0.1567 97.32 5.602

0.1269 150.52 4.244

0.1076 196.00 4.986

0.0948 221.67 3.410

0.0853 249.53 1.748

0.0523 438.33 1.840

0.0433 487.23 171

Experiment 89 (FeS/NagSOg/HCl)

c, mole liter”^ t, hr U, 1 0 ** mole liter”! hr”!

0.2189 19.30 6.897

0.1965 51.78 8.109

0.1766 76.32 7.105

0.1594 100.53 6.449

0.1179 164.88 2.826

0.1103 191.77 3.646

0.1007 218.10 2.991

0.0871 263.57 3.197

0.0789 289.22 2.366

0.0723 317.12 1.536

0.0433 505.88 0.777

0.0395 554.81 - 172

Experiment 91 (FeS/NagSOg/HCl)

c, mole liter” ! t, hr V), 10** mole liter” ! hr” !

0.1437 5.77 5.821

0.1281 32.57 4.837

0.1173 54.90 3.717

0.1072 82.07 2.466

0.1018 103.97 2.901

0.0958 124.65 4.333

0.0847 150.27 4.524

0.0741 173.70 3.651

0.0651 198.35 1.712

0.0578 241.00 0.972

0.0529 291.40 1.284

0.0496 317.10 - 173

B. Reactant Variation Experiments

Experiment 92 (FeS/Na2S03/H2S04)

c, mole liter” ! t, hr U, 10** mole liter”! hr”!

0.2519 4.07 12.37

0.2183 31.23 10.71

0.1949 53.08 7.923

0.1785 73.78 7.109

0.1436 122.87 4.663

0.1321 147.53 5.663

0.1184 171.72 5.231

0.0961 214.35 2.844

0.0745 290.30 -

Experiment 106 (FeS/Calciner Feed/HCl)

0.2033 11.25 14.42

0.1878 22.00 8.776

0.1698 42.51 8.276

0.1444 73.20 8.910

0.1256 94.30 6.914

0.0802 159.96 4.379

0.0566 213.85 2.000

0.0336 328.85 0.811

0.0296 378.28 - 174 c. Sulfide Ore Experiments

Experiment 101 (PbS/NagSOy/HCl)

c, mole liter”! t, hr U, 10" mole liter”! hr”!

0.3220 7.03 17.17

0.2767 33.42 9.028

0.2557 56.68 8.075

0.2385 77.98 8.270

0.1790 149.93 5.193

0.1665 174.00 6.092

0.1504 200.43 3.739

0.1417 223.70 -

Experiment 104 (FeS-pyrrhotite/Na2 S0 3 /HCl)

0.1657 74.96 4.908

0.1553 96.15 4.283

0.1354 142.61 3.806

0.1242 172.04 4.106

0.1155 193.23 2.951

0.1004 244.40 3.602

0.0928 265.50 2.092

0.0883 287.01 2.298

0.0659 384.48 1.996

0.0378 525.30 - 175

REFERENCES

1. U.S. Congress, Senate, "Control of Sulfur Oxides," S. Doc, 94-59, 94th Cong., 1st sess., 1975, p. 1.

2. U.S. Congress, Senate, "Progress in the Prevention and Control of Air Pollution in 1975," S. Doc. 94-228, 94th Cong., 2nd sess., 1976, p. 80.

3. Council on Environmental Quality, "Environmental Quality — The Sixth Annual Report," U.S. Government Printing Office, Washington, D.C., 1975, p. 306.

4. U.S. Environmental Protection Agency, "Standards of Performance for New Stationary Sources," Federal Register, 36, 24876 (1971).

5. U.S. Code, ]J, sec. 793 (1974).

6. H.S. Rosenberg, R.B. Engdahl, J.H. Oxley, and G.M. Genco, Chem. Eng. Progr., 71 (5). 66 (1975).

7. K.S. Murthy, H.S. Rosenberg, and R.B. Engdahl, J. Air Pollut. Contr. Ass., 851 (1976).

8. A.V. Slack, Environ. Sci. Technol., 2» H O (1973).

9. R.E. Harrington, Int. J. Sulfur Chem., Part B, %, 57 (1972).

10. A.V. Slack, Int. J. Sulfur Chem., !_■, 67 (1972).

11. "SO 2 removal technology enters growth phase," Environ. Sci. Technol., 6, 688 (1972).

12. "Stack Gas Cleaning: A 1976 Update," Environ. Sci. Technol., 10, 417 (1976).

13. "Proceedings: Symposium on Flue Gas Desulfurization - Atlanta, November 1974," EPA 650/2-74-126a/b (NTIS No. PB 242-572/3), December 1974.

14. U.S. Environmental Protection Agency, "Proceedings: Symposium on Flue Gas Desulfurization, New Orleans, March 1976," EPA 600/2-76- 136a/b (NTIS No. PB 255-317/8), May 1976. 176

15. J.B. Pfeiffer, Ed., "Sulfur Removal and Recovery from Industrial Processes," Advances in Chemistry Series 139, American Chemical Society, Washington, B.C., 1975.

16. A. Saleem, J. Air Pollut. Contr. Ass., 22, 172 (1972).

17. A.V. Slack, H.L. Falkenberry, and R.E. Harrington, Environ. Sci. Technol., 22, 159 (1972).

18. A.V. Slack, Chem. Eng. Progr., 72 (8), 94 (1976).

19. R.L. Gall and E.J. Piasecki, Chem. Eng. Progr., 71 (5), 72 (1975).

19a. U.S. Environmental Protection Agency, "Conceptual Design and Cost Study— Sulfur Oxide Removal from Power Plant Stack Gas— Magnesia Scrubbing— Regeneration: Production of Sulfuric Acid," EPA-R2- 73-244, May 1973.

20. A. Tamaki, Chem. Eng. Progr., 71 (5), 55 (1975).

21. W.F. Bischoff and Y. Habib, Chem. Eng. Progr., 71 (5), 59 (1975).

22. "Coal Research (IV): Direct Combustion Lags Its Potential," Science, 194, 172 (1976).

23. F.T. Princiotta, "Sulfur Oxide Throwaway Sludge Evaluation Panel (SOTSEP), Volume I: Final Report - Executive Summary," EPA 650/2-75-010-a (NTIS No. PB 242 618),U.S. Environmental Protection Agency, April 1975.

24. F.T. Princiotta, "Sulfur Oxide Throwaway Sludge Evaluation Panel (SOTSEP), Volume II: Final Report - Technical Discussion," EPA 650/2-75-010-b (NTIS No. PB 242 618), U.S. Environmental Protection Agency, April 1975.

25. J.W. Jones, "Alternatives for Disposal of Flue Gas Cleaning Wastes," Presented at the American Institute of Chemical Engineers 82nd National Meeting, Atlantic City, N.J., September 1976.

26. "Putting industrial sludges in place," Environ. Sci. Technol., 874 (1972).

27. "The Stabilization Game," Environ. Sci. Technol., 9^, 622 (1975).

28. J.R. West, Ed., "New Uses of Sulfur," Advances in Chemistry Series No. 140, American Chemical Society, Washington, D.C., 1975.

29. G.C. Thom, M.S. Thesis, The American University, Washington, D.C., 1974. 177

30. U.S. Department of the Interior, Bureau of Mines, "Commodity Data Summaries 1976," U.S. Government Printing Office, Washington, D.C., 1975.

31. W. Kunda and B. Rudyk, Can. Met. Quart., 551 (1970).

32. S.I. Sobol and T.M. Frash, Tsvet. Metal., (2), 14 (1974).

33. R.A. Meyers, J.W. Hamersma, and M.L. Kraft, Environ. Sci. Technol., 9, 70 (1975).

34. R.A. Meyers, J.W. Hamersma, J.S. Land, and M.L. Kraft, Science, 177, 1187 (1972); E.P. Koutsoukos, et.al., "Meyers Process Development for Chemical Desulfurization of Coal," EPA-600/2-76-143, U.S. Environmental Protection Agency, May 1976.

35. U.S. Congress, Senate, "The Cost of Clean Air." S. Doc. 93-122, 93rd Cong., 2nd sess., 1974, pp. 94-109.

36. M.E. Wadsworth and F.T. Davis, Eds., "Unit Processes in Hydrometallurgy," The Metallurgical Society, AIME, Gordon and Breach Science Publishers, New York, N.Y., 1963.

37. D.J.I. Evans and R.S. Shoemaker, Eds., "International Symposium on Hydrometallurgy." The Metallurgical Society, AIME, New York, N.Y., 1973.

38. J.C. Yannopoulos and J.C. Agarwal, Ed., "Extractive Metallurgy of Copper," An International Symposium, The Metallurgical Society, AIME, New York, N.Y., 1976.

39. M.E. Wadsworth, J. Met. (JOM), 28 (3), 4 (1976).

40. B.R. Palmer and R.B. Sudderth, Min. Eng., 28 (2). 72 (1976).

41. V.N. Mackiw, Can. J. Chem. Eng., 46 (2), 3 (1968).

42. M.E. Wadsworth, Miner. Sci. Eng., 4 (4), 36 (1972).

43. K.N. Subramanian and P.E. Jennings, Can. Metall. Q., 11, 387 (1972).

44. J.C. Paynter, J. South Afr. Inst. Min. Metall., 158 (1973).

45. A.R. Burkin, Proc. Roy. Soc. Lond. A, 338, 419 (1974).

46. J.M. Toguri, Can. Metall. Q., 14, 323 (1975).

47. M.L. Jansen and D.A. Milligan, J. Met., 27 (1). 13 (1975).

48. R.B. Bhappu, D.H. Reynolds, and W.H. Stahmann in "Unit Processes in Hydrometallurgy," M.E. Wadsworth and F.T. Davis, Eds., The Metallurgical Society, AIME, Gordon and Breach Science Publishers, New York, N.Y., 1963, pp. 95-113. 178

49. M. Pourbaix, "Atlas of Electrochemical Equilibria in Aqueous Solutions," Pergamon Press, New York, N.Y., 1966.

50. M.E. Wadsworth in "Physical Chemistry an Advanced Treatise, Volume VIl/Reactions in Condensed Phases," H. Eyring, Ed., Academic Press, New York, N.Y., 1975, pp. 413-470.

51. W.H. Dresher, M.E. Wadsworth, W.M. Fassell, Jr., Trans. AIME, 206, 794 (1956).

52. A. Vizsolyi, H. Veltman, F.A. Forward in "Unit Processes in Hydrometallurgy," M.E. Wadsworth and F.T. Davis, Eds., The Metallurgical Society, AIME, Gordon and Breach Science Publishers, New York, N.Y., 1963, pp. 326-343.

53. F.A. Forward and V.N. Mackiw, Trans. AIME, 203, 457 (1955).

54. E. Peters in "International Symposium on Hydrometallurgy," D.J.I. Evans and R.S. Shoemaker, Eds., The Metallurgical Society, AIME, New York, N.Y., 1973, pp. 205-228.

55. A.R. Burkin, "The Chemistry of Hydrometallurgical Processes," E. and F.N. Spon, Ltd., London, 1966.

56. F.A. Forward and I.H. Warren, Metall. Rev., 137 (1960).

57. I.I. Ginzburg, Ya. I. Ol'shanskii, and V.V. Belyatskii, Trudy Inst. Ged. Rudnykh Mestorozhden., Petrog., Mineral, i Geokhim., No. 59, 1 (1961); Chem. Abs., 180i (1962).

58. G. Thomas, T.R. Ingraham, and R.J.C. MacDonald, Can. Metall. Q., 6, 281 (1967).

59. E.W. Pearson, Min. Eng., 26 (12), 40 (1974).

60. J.E. Dutrizac, R.J.C. MacDonald, and T.R. Ingraham, Trans.Met. Soc. AIME, 955 (1969).

61. K.N. Subramanian, E.S. Stratigakos, and P.H. Jennings, Can. Metall. Q., U , 425 (1972).

62. F.P. Haver, R.D. Baker, and M.M. Wong, U.S. Bureau of Mines Report of Investigations 8007, 1975.

63. F.P. Haver, and M.M. Wong, U.S. Bureau of Mines Report of Investigations 8105, 1976.

64. G. Van Weert, K. Mah, and N.L. Piret, CIM Bull., 78 (1), 97 (1974).

65. J.M. Demarthe, L. Gandon, and A. Georgeaux in "Extractive Metallurgy of Copper," An International Symposium, The Metallurgical Society, AIME, New York, N.Y., 1976, pp. 825-848. 179

66. K.J. Cathro in "Extractive Metallurgy of Copper," J.C. Yannopoulos and J.C. Agarwal, Eds., An International Symposium, The Metallurgical Society, AIME, New York, N.Y., 1976, pp. 776-792.

67. R.D. Groves and P.B. Smith, U.S. Bureau of Mines Report of Investigations 7801, 1973.

68. K.W. Downes and R.W. Bruce, Can. Min. Metall. Bull., 48, 127 (1955).

69. W. Kunda, B. Rudyk, and V.N. Mackiw, C.I.M. Bull., 819 (1968).

70. F.A. Forward and H. Veltman, J. Met., IJL, 836 (1959).

71. H. Veltman, S. Pellegrini, and V.N. Mackiw, J. Met., 19 (2), 21 (1967).

72. A. Vizsolyi, H. Veltman, I.H. Warren, and V.N. Mackiw, J. Met., 19 (11), 52 (1967).

73. W. Kunda, R. Hitesman, and H. Veltman in "Extractive Metallurgy of Copper," J.C. Yannopoulos and J.C. Agarwal, Eds., An International Symposium, The Metallurgical Society, AIME, New York, N.Y., 1976, pp. 793-813.

74. J.E. Prater, P.B. Queneau, and T.J. Hudson, Transactions of the American Institute of Mining, Metallurgy and Petroleum Engineers (Society of Mining Engineers), 254, 117 (1973).

75. G. Bjorling, I. Fait, E. Lindgren, and I. Toromaov in "Extractive Metallurgy of Copper," J.C. Yannopoulos and J.C. Agarwal, Eds., An International Symposium, The Metallurgical Society, AIME, New York, N.Y., 1976, pp. 725-737; G. Bjorling in "International Symposium on Hydrometallurgy," D.J.I. Evans and R.S. Shoemaker, Eds., The Metallurgical Society, AIME, New York, N.Y., 1973, pp. 701-717.

76. R.W. Bartlett in "International Symposium on Hydrometallurgy," D.J.I. Evans and R.S. Shoemaker, Eds., The Metallurgical Society, AIME, New York, N.Y., 1973, pp. 331-374.

77. J.P. Baur, H.L. Gibbs, and M.E. Wadsworth, U.S. Bureau of Mines Report of Investigations 7823, 1974.

78. F. Habashi, Miner, Sci. Eng., 3 (3), 3 (1972).

79. I.F. Mazalov, Z.B. Rozhdestvenskaya, T.F. Simon, and G.D.Medvedeva, Z. Prik. Khim., 2691 (1974).

80. H. Cox and A.K. Schellinger, Eng. Min. J., 159 (10), 101 (1958).

81. R.B. Bahppu, D.H. Reynolds, and W.S. Stahmann in "Unit Processes in Hydrometallurgy," M.E. Wadsworth and F.T. Davis, Eds., The Metallurgical Society, AIME, Gordon and Breach Science Publishers, New York, N.Y., 1963, pp. 95-113. 180

82. P.H. Yu, C.K. Hansen,and M.E. Wadsworth In "International Symposium on Hydrometallurgy," D.J.I. Evans and R.S. Shoemaker, Eds., The Metallurgical Society, AIME, New York, N.Y., 1973, pp. 375-402.

83. M.E. Wadsworth in "Annual Review of Physical Chemistry," H. Eyring, C.J. Christensen, and H.S. Johnston, Eds., Annual Reviews, Inc., Palo Alto, Calif., Vol. 23, pp. 355-383, 1972.

84. R.K. Mishra, Ph.D. Thesis, University of Utah, Salt Lake City, Utah, 1973.

85. J.J. Jan, Ph.D. Thesis, University of Denver, Denver, Col. 1974.

86. D. Verhulst, Eng. Sc. D., Columbia University, New York, N.Y., 1974.

87. R.K. Clifford, Ph.D. Thesis, University of Utah, Salt Lake City, Utah, 1972.

88. G.C. Thom, Ph.D. Thesis Proposal, The American University, Washington, D.C. 1975.

89. C.E. Ennis in "Proceedings: Symposium on Flue Gas Desulfurization, New Orleans, March 1976," U.S. Environmental Protection Agency, EPA 600/2-76-136a/b (NTIS No. PB 255-317/8), May 1976, pp. 865-875.

90. H. Hirai, R. Odello, and H. Shimamura, Chem. Eng., 79 (8), 78 (1972).

91. J.S. Hayford, L.P. VanBrocklin, and M.A. Kuck, Chem. Eng. Progr., 69 (12), 54 (1973).

92. F. G. Sherif, J.S. Hayford, and J.E. Blanch, Ger. Offen., 2,250,959, April 1973.

93. W.I. Nissen, D.A. Elkins, and W.A. McKinney in "Proceedings: Symposium on Flue Gas Desulfurization, New Orleans, March 1976." U.S. Environmental Protection Agency, EPA 600/2-76-136a/b (NTIS No. PB 255-317/8), May 1976, pp. 843-864.

94. J.B. Rosenbaum, W.A. McKinney, H.R. Beard, L. Crocker, and W.I. Nissen, U.S. Bureau of Mines Report of Investigations 7774, 1973.

95. S. Vasan, Chem. Eng. Progr., 71 (5), 61 (1975).

96. W.J. Moore, "Physical Chemistry," Fourth Edition, Prentice-Hall, Inc., Englewood Cliffs, N.J., 1972, p. 513.

97. D.W. Hissong, K.S. Murthy, and A.W. Lemmon, Jr., "Reductant Gases for Flue Gas Desulfurization Systems," EPA 600/2-76-130 (NTIS PB No. 254 168), U.S. Environmental Protection Agency, May 1976. 181

98. G.C. Thom, "An Investigation of the in situ Recovery of Elemental Sulfur from Spent Flue Gas Desulfurization Sludges," Unsolicited Research Proposal, The American University, Washington, D.C., June 1976.

99. W.A. Pryor, "Mechanisms of Sulfur Reactions," McGraw-Hill Book Co., Inc., New York, N.Y., 1962.

100. B. Meyer, Ed., "Elemental Sulfur," Interscience Publishers, New York, N.Y., 1965.

100a. G. Nickless, Ed., "Inorganic Sulfur Chemistry," Elsevier Publishing Co., Ltd., Amsterdam, 1968.

101. B. Meyer and M. Schmidt in "International Review of Science, Inorganic Chemistry Series Two, Main Group Elements, Groups VI and VII," Volume 3, V. Gutmann, Ed., Butterworth and Co., (Publishers) Ltd., London, 1975.

102. B. Meyer, Chem. Rev., 76, 367 (1976); Chem. Rev., 429 (1964).

103. A. Senning, Ed., "Sulfur in Organic and Inorganic Chemistry," Marcel Dekker, Inc., New York, N.Y., 1971.

104. M. Schmidt and W. Siebert in "Comprehensive Inorganic Chemistry," J.C. Bailar, Ed., Pergamon Press, Ltd., Elmsford, N.Y., 1973, Vol. 2, pp. 795-933.

105. W.J. MacKnight, and A.V. Tobolsky in "Elemental Sulfur," B. Meyer, Ed., Interscience Publishers, N.Y., 1965, pp. 95-97.

106. M. Schmidt in "Sulfur in Organic and Inorganic Chemistry." A. Senning, Ed., Marcel Dekker, Inc., New York, N.Y., 1971, Vol. 2, pp. 71-112.

107. B. Meyer, L. Peter, and K. Spitzer, Inorg. Chem., 16, 27 (1977).

108. J.L. Kice in "Sulfur in Organic and Inorganic Chemistry." A. Senning, Ed., Marcel Dekker, Inc., New York, N.Y., 1971, Vol. 1, pp. 153-207.

109. P.D. Bartlett, and R.E. Davis, J. Amer. Chem. Soc., 80, 2513 (1958)

110. 0. Foss, Acta Chem. Scand., ^2, 959 (1958).

111. 0. Foss, Acta Chem. Scand., 1610 (1961).

112. A. Fava, and S. Bresadola, J. Amer. Chem. Soc., 77, 5792 (1955).

113. R.E. Davis, J. Amer. Chem. Soc., 80, 3565 (1958).

114. H. Wackenroder, Liebigs Ann. Chem., 60, 189 (1846). 182

115. M. Goehring, Fortschr. Chem. Forsch., 2» 444 (1952).

116. H. Stamm, W.W. Magers, and M. Goehring, Z. anorg. u. allgem.Chem., 244, 184 (1940).

117. E. Blasius and W. Burmeister, Z. Anal. Chem., 168,1 (1958).

118. M. Schmidt, Z. anorg. u. allgem. Chem., 289, 141 (1957).

119. K. Steinle, Thesis, University of Munich, 1962.

120. E. Blasius, G. Horn, A. Knochel, J. Münch and H. Wagner in "Inorganic Sulfur Chemistry," G. Nickless, Ed., ElsevierPublishing Co., Ltd., Amsterdam, 1968, pp. 199-239.

121. J.H. Karchmer, Ed., "The Analytical Chemistry of Sulfur and Its Compounds," Wiley Interscience, New York, N.Y., 1970.

122. J.N. Chapman and H.R. Beard, Anal. Chem., 2268 (1973).

123. E. Blasius and J. Munch, Z. Anal. Chem., 261, 198 (1972).

124. E. Blasius, A. Knochel, N. Spannhake, and H. Wagner, J. Chromatogr., 70, 109 (1972).

125. E. Blasius, H. Augustin, and G. Klemm, J. Chromatogr., 108, 53 (1975).

126. E. Blasius and J. Munch, J. Chromatogr., 137 (1973).

127. E. Blasius and J. Munch, Z. Anal. Chem., 272, 353 (1974).

128. E. Blasius and H. Wagner, Chem. Zeitung, 98, 154 (1974).

129. E. Blasius and J. Munch, Z. Anal. Chem. 272, 346 (1974).

130. H. Ley and E. Konig, Z. Physik. Chem., B 41, 365 (1937).

131. R. Sperling and A. Treinin, J. Phys. Chem., 68, 897 (1964).

132. D.A. Ratkowsky and J.L. McCarthy, J. Phys. Chem., 66, 516 (1972).

133. T.E. Eriksen and J. Lind, Acta Chem. Scand., 2 ^ , 3325 (1972).

134. C.R. Blandamer and M.F. Fox, Chem. Rev., 70, 59 (1970).

135. R.M. Golding, J. Chem. Phys., 33, 1666 (1960).

136. V.K. La Mer and A.S. Kenyon, J. Colloid Sci., 2^, 257 (1947).

137. E.M. Zaiser and V.K. La Mer, J. Colloid Sci., 3, 571 (1948). 183

138. R.H. Dlnegar, R.H. Smellie, and V.K. La Mer, J. Amer. Chem. Soc., 23, 2050 (1951).

139. J.L. Keller, U.S. Patent 2729543, January 3, 1956.

140. I.M. Kolthoff, F.B. Sandell, E.J. Meehan, and S. Bruckenstein, "Quantitative Chemical Analysis," Macmillan Company, Toronto, Canada, 1969, p. 847.

141. G.H. Ayres, "Quantitative Chemical Analysis," Second Edition, Harper and Row, Publishers, New York, N.Y., 1968, pp. 625, 644.

142. K.J. Laidler, "Chemical Kinetics," Second Edition, McGraw Hill Book Company, New York, N.Y., 1965, p. 16.

143. R.C. Weast, Ed., "Handbook of Chemistry ofPhysics," 56th Edition, The Chemical Rubber Co., Cleveland, Ohio, 1975, p. B-104.

144. P.H. Tewari and A.B. Campbell, J. Phys. Chem., 80, 1844 (1976).

145. R.C. Weast, Ed., "Handbook of Chemistry of Physics," 56th Edition, The Chemical Rubber Co., Cleveland, Ohio, 1975, p. B-196.

146. Wo. Pauli, J. Colloid Sci., 2» 33 (1947).

147. B. Jirgensons and M.E. Straumanis, "A Short Textbook ofColloid Chemistry," Second Revised Edition, The Macmillan Co., New York, N.Y., 1962, p. 149.