Electrochemical Studies of Catalysed Aqueous Sulphide ......5.2 Voltammogram of Vanadium (V) in Borate Buffer at pH 9.2. 8 5 (First Scan, commenced at 0.35 V vs. SHE. 50 mV s_1.) 5.3

1

Electrochemical Studies of Catalysed Aqueous Sulphide Oxidation

A thesis submitted for the degree of

Doctor of Philosophy of the University of London

and

The Diploma of Imperial College

by

Ian Thompson

Department of Mineral Resources Engineering Sept. 1987

Im perial College of Science and Technology

University of London

LONDON

SW7 2BP

But it's all right now,In fact it's a Gas Gas Gas!

Mick Jagger (1968)

Abstract 3

Abstract


This thesis concerns the mechanism of oxidation of aqueous sulphide solutions in the

British Gas Stretford Process, which uses atmospheric oxygen to achieve the partial

oxidation of hydrogen sulphide, producing elemental sulphur and water. Hydrogen

sulphide is absorbed in an alkaline solution (pH 8.5-9.5) containing vanadium (V)

salts and anthraquinone derivatives which act as oxidation catalysts.

The important methods of removing hydrogen sulphide from fuel gases were reviewed,

and a detailed description of the Stretford Process was provided.

The thermodynamic data on sulphur species were presented in the form of Eh/pH

diagrams, and the literature relating to the oxidation of sulphide solutions was

surveyed. The redox behaviour of sulphide and polysulphide solutions were

investigated using electrochemical techniques such as cyclic and pulse voltammetry at

gold ring-disc electrodes. It was shown that polysulphide species were important

intermediates in the oxidation of HS" ions.

The aqueous chemistry of vanadium was described and the electrochemical behaviour

of vanadium (V) and (IV) solutions at pH 9.2 were investigated at mercury, carbon and

gold electrodes. The electrochemical reduction of vanadium (V) was shown to be

irreversible and to lead to vanadium oxide films, rather than to solution species.

The redox chemistry of anthraquinone was reviewed and electrochemical studies were

made of the compound anthraquinone 2,7-disulphonate. The reduced species and

intermediates were identified using UV-visible spectrophotometry and ESR

spectroscopy.

The reduction pathways of oxygen in alkaline solution were reviewed, and the role of

hydrogen peroxide as a possible reactive intermediate was investigated.

The Stretford Process chemistry was examined using stopped-flow spectroscopic

methods; these enabled the courses of the redox reactions between sulphide solutions

and solutions containing the catalysts to be followed.

A mechanism was proposed for the Stretford Process, and possible process

improvements were discussed.

Acknowledgements 4

Acknowledgements

I would like to thank Dr. G. H. Kelsall for his supervision during the course of this

work. Thanks must also go to Dr. T. Ritter from the British Gas London Research

Station and to the other staff there who have given me help; Dr. D. Keene, Dr R.

Mounce, Dr R. Gibbons, Roy Lowry, Lucien Anthony, and Susan Mahony.

From Imperial College I would particularly like to thank Gordon "the glass" as well as

the other technical and academic staff in the Mineral Resources Engineering and the

Chemistry Departments. Research is not carried out alone. The present and past

members of the research group have both aided my studies and, through their

company, made them more enjoyable. Thank you.

I would also like to acknowledge the financial help of British Gas, the Science and

Engineering Research Council and last but definitely not least, Corina Thompson.

Contents 5

Contents

Abstract 3

Acknowledgements 4

Contents 5

List of Figures 9

List of Tables 12

1 . Introduction: The Importance of Sulphide Oxidation 141 .1 The Removal of Hydrogen Sulphide 1 4

1.1.1 Absorption by liquids 15

1.1.2 Adsorption by solids 16

1.1.3 Electrochemical oxidation of hydrogen sulphide 17

1.1.4 Aqueous oxidation of hydrogen sulphide 1 8

1 .2 The Stretford Process 1 9

1.2.1 Historical development 2 0

1.2.2 Operational Problems 2 1

1.2.3 Mechanistic studies 2 3

1 .3 Objectives of the Present Study 2 5

1.3.1. Research Approach 2 5

2 . Review of Sulphide Oxidation 2 62 .1 The Oxidation States of Sulphur 2 9

2 . 1.1 Sulphide (-II) 2 9

2.1.2 Poly sulphides (-1 to 0) 3 0

2.1.3 Elemental Sulphur (0) 3 2

2.1.4 Polythionates (0 to IV) 3 3

2.1.5 Thiosulphate (II) 3 3

2.1.6 Sulphite (IV) 3 4

2.1.7 Sulphate (VI) 3 4

2 .2 Electrochemical Studies of Sulphide Oxidation 3 4

2 .3 Chemical Oxidation of Sulphide using Oxygen 3 7

2.3.1 Rate of Reaction of Sulphide Solutions with Oxygen 3 82.3.2 Effect of Temperature and pH on Reaction Rate 3 9

2.3.3 Catalysis of Sulphide Oxidation 3 9

2.3.4 B acterial Action in S ulphide Oxidation 4 0

2 .4 The Production of Elemental Sulphur 4 0

3 . Sulphide Electrochemistry 413.1 Thermodynamic Calculations 4 3

3 .2 Experimental 4 4

3.2.1 Solution Preparation 4 5

3.2.2 Electrochemical Instrumentation 4 6

3.2.3 Electrode Pretreatment 4 7

3.2.4 Experimental: Ion chromatography 4 8

3 .3 Sulphide Voltammetry: Results and Discussion 5 0

3 .4 Thiosulphate Voltammetry: Results and Discussion 5 1

3 .5 Polysulphide Voltammetry: Results and Discussion 5 3

3 .6 Ring-Disc Studies: Results and Discussion 5 6

3 .7 Calculated Polysulphide Concentrations vs. Potential 6 2

3 .8 Detection of Polysulphides Using Ion Chrom atography 6 4

3.8.1 Ion Chromatography: Results and Discussion 6 4

3 .9 Summary 6 5

4. Vanadium 664 .1 Vanadium (V) 6 7

4 .2 Vanadium (IV) 7 0

4 .3 Vanadium (V)/(IV) Compounds 7 2

4 .4 Vanadium (III) 7 4

4 .5 Vanadium (II) 7 4

4 .6 Vanadium Electrochemistry 7 5

4.6.1 The Vanadium (V)/(IV) Couple 7 5

4.6.2 Vanadium (TV) reduction 7 6

4.6.3 The Vanadium (ffl)/(II) Couple 7 6

4 .7 Oxidation of Vanadium (IV) solutions using Oxygen 7 6

4 .8 Vanadium Sulphides 7 7

4.8.1 V3S,V 5S4,VS 7 7

4.8.2 V2S5 7 8

4.8.3 VS2 andVS4 7 8

4 .9 Vanadium -Sulphur Complexes 7 84 .1 0 Summary 8 0

5. Vanadium Electrochemistry 8 2

5 .1 Vanadium Electrochemistry: Experimental 8 2

5.1.1 Solution Preparation 83

5 .2 Vanadium (V) Voltammetry: Results and Discussion 8 4

5 .3 Summary 9 1

Contents 6

Contents 7

6. Review of Anthraquinone Redox Chemistry 9 26 .1 Anthraquinone Reduction 9 2

6.1.1 Substituent effects 9 4

6.1.2 Photo-reduction 9 5

6 .2 Anthraquinones in the Production of Hydrogen Peroxide 9 6

7. Redox Chemistry of Anthraquinone 2,7-disulphonate 9 7

7 .1 Purification of Anthraquinone 2,7-disulphonate 9 7

7.1.1 Analysis of the Purified Material 9 7

7 .2 Experimental: Voltammetry 9 87 .3 Experimental: Exhaustive Electrolysis 9 9

7.3.1 Calculations: Exhaustive Electrolysis 1 0 0

7.3.2 Calibration of Exhaustive Electrolysis Apparatus 1 0 1

7 .4 Voltammetry: Results and Discussion 1 0 3

7 .5 Exhaustive Electrolysis: Results and Discussion 1 0 8

7 .6 UV-Visible Spectrophotometry: Experimental 1 1 1

7 .7 Results and Discussion: UV-Visible Spectrophotometry 1 1 3

7.7.1 Spectral Assignments 1 1 4

7 .8 ESR Spectroscopy: Experimental 1 1 6

7 .9 ESR Spectroscopy: Results and Discussion 1 1 7

7 .1 0 ESR Spectral Structure 1 1 9

7 .1 1 Summary 1 2 0

8 . Oxygen Reduction 1228 .1 The Oxygen / W ater Couple 1 2 4

8.1.1 The Evolution of Oxygen 1 2 6

8 .2 Hydrogen Peroxide 1 2 7

8 .3 Superoxides 1 2 9

8 .4 Experimental 1 2 9

8 .5 Oxygen Reduction: Results and Discussion 1 3 0

8.6 Summary 13 2

9 . The Redox Chemistry of the Stretford Process 1 3 3

9 .1 Experimental 1 3 3

9.1.1 S topped Flow Apparatus 1 3 4

9.1.2 Experimental: Measurement of Solution Potential 1 3 5

9.1.3 Experimental: Preparation o f51V NMR Samples 1 3 6

9 .2 Reaction of AQ27DS and HS“: Stopped Flow Results 1 3 7

9.2.1 Rate Studies 1 3 8

9.2.2 S olution Potential Measurements 1 4 0

9 .3 Reaction of V(V) and HS“: Stopped Flow Results 1 4 2

9.3.1 Vanadium (V) Reduction 1 4 4

9 .4 Interaction of AQ27DH" ions with Oxygen 1 4 7

9 .5 Stretford Solution Chemistry: Electrochemical Results 1 4 8

9 .6 The Stretford Process: Possible Mechanism 1 5 0

1 0 . Conclusions 1 5 2

1 0 .1 The S(-II)/S(0) Redox Couple 1 5 2

1 0 .2 The V(V)/V(IV) Redox Couple 1 5 2

10.3 The Anthraquinone/Anthraquinol Redox Couple 1 5 3

1 0 .4 The C ^/O H - Redox Couple 1 5 4

1 0 .5 The Redox Chemistry of the Stretford Process 1 5 4

1 0 .6 The Mechanism of the Stretford Process 1 5 5

1 0 .7 Concluding Remarks 1 5 6

Appendix: Thermodynamic Data Used in Eh-pH Diagrams 1 5 8

R eferences 1 6 1

Contents 8

Figures 9

List of Figures

1 .1 The Stretford Process. 19

2 .1 Possible valence states of sulphur in aqueous media. 2 62 .2 Efo-pH diagram for the sulphur/water system at 298 K. 2 7

2 .3 . Efo-pH diagram for metastable sulphur system at 298 K. 2 8

2 .4 E^-pH diagram for the sulphur/water system at 298 K. 2 9

(Oxy-sulphur anions not considered.)

3 .1 Eh-pH Diagram for the Au/Cl/S System. 4 2

3 .2 Eh-pH diagram for the sulphur/water system at 298 K. 4 3

3 .3 Metastable Eh-pH diagram for the S/H20 system at 298 K. 4 4

3 .4 A Rotating Ring Disc Electrode. 4 4

3 .5 Ion Chromatography Apparatus. 4 9

3 .6 Voltammogram of HS~ on Gold Plated Disc Electrode. 5 0

([HS_] = 10 mol m-3, pH = 9.2, nth. cycle, 20 mV s_1.)

3 .7 Cyclic Voltammograms of Sodium Thiosulphate. 5 2

([Na2S2C>3] = 10 mol m~3. 1st Scans. 100 mV s-1. pH = 8.2.)

3 .8 Voltammograms of Polysulphide Solution at a Gold Disc. 5 3

([Sx] = 1 mol m-3. xav = 2. pH = 8.2. Scan rate 50 mV s-1.)

3 .9 E^-pH Diagram of the Sulphide/Polysulphide System. 5 4

3 .1 0 Voltammograms of Poly sulphide Solution at a Gold Disc. 5 5

([Sx2~] = 1 mol m~3. xav = 2. pH = 8.2. Scan rate 50 mV s-1.)

3 .1 1 Ring-Disc Voltammetry of Sulphide Solution at Au RRDE. 5 7

([HS~] = 10 mol m~3. co = 9 Hz. Scan rate = 100 mV s_1.)

3 .1 2 Ring-Disc Voltammetry of Sulphide Solution at Au RRDE. 6 0

([NaOH] = 1 kmol m“3, [HS"] = 1 kmol m"3, co = 4 Hz)

3 .1 3 Ring-Disc Potential Pulse Study. Au RRDE. 6 1

([HS“] = 10 mol n r 3, co = 9 Hz. pH 9.3.)

3 .1 4 Poly sulphide Distribution vs. Potential (pH = 14) 6 2

3 .1 5 Polysulphide Distribution vs. Potential (pH = 9) 6 3

3 .1 6 Ion Chromatography Results. 6 4

4 .1 E^-pH Diagram for the Vanadium-Water System. 6 7

4 .2 Structure of the V10O286" i°n- 6 9

4 .3 Vanadium (V) Speciation. 7 0

4 .4 Structure of V1804212’- 7 1

4 .5 Vanadium (IV) Speciation in Solution. 7 2

4 .6 Vanadium (IE) Speciation in Solution. 7 4

Figures 10

5 .1 Hanging Mercury Drop Electrode. 8 2

5 .2 Voltammogram of Vanadium (V) in Borate Buffer at pH 9.2. 8 5

(First Scan, commenced at 0.35 V vs. SHE. 50 mV s_1.)

5 .3 Efo-pH Diagram for the V-H2O System at 298 K. 8 65 .4 Voltammogram of Vanadium (V) in Carbonate Buffer at HMDE. 8 7

(1st. Scan, 50 mV s '1, pH 9.3, [V(V)] = 10 mol n r3, T = 40 °C.)

5 .5 Cyclic Voltammogram of Vanadium (V) on a Gold Electrode. 8 8(1st. Scan, 50 mV s’1, pH 9.3, [V(V)] = 10 mol n r3, T = 19 °C.)

5 .6 Cyclic Voltammogram of V(V) on a Vitreous Carbon Electrode. 8 9

(1st. Scan, 50 mV s’1, pH 9.3, [V(V)] = 10 mol n r3, T = 20 °C.)

5 .7 Cyclic Voltammogram of VS43", HS' on a Gold Disc. 9 0

([VS43-] = 0.2 mol n r3, [HS-] = 0.36 kmol n r 3.)

6 .1 9,10-Anthraquinone. 9 26 .2 Possible Intermediates in the Reduction of Anthraquinones. 9 2

6 .3 Anthraquinone 2,7-disulphonate (AQ27DS). 9 4

6 .4 Na4 NN'-disulphomethylanthraquinone-2,6-disulphonamide. 9 5

7 .1 Electrochemical Cell Design for Voltammetry Experiments. 9 87 .2 Exhaustive Electrolysis Apparatus. 9 9

7 .3 Plot of Log it vs t during the reduction of Fe(CN)g3". 1 0 2

7 .4 Cyclic Voltammogram of AQ27DS. 1 0 3

([AQ27DS] = 1 mol m“3, Sweep Rate 5 mVs"1, pH 9.3.)

7 .5 Cyclic voltammetry of AQ27DS at a rotated gold disc electrode. 1 0 6

(Scan rate = 20 mVs"1. pH = 9.23. C0 = 0.357 mol m-3.)

7 .6 Plot of i vs. co1/2 for reduction of AQ27DS. 1 0 7

7 .7 Plot of Reduction Potential vs. pH for AQ27DS. 1 0 8

7 .8 Plot of Charge vs. Time During the Electrolysis of AQ27DS. 1 0 9

(Electrolysis potential = -0.6 V vs. SHE. pH = 9.3.)

7 .9 Plot of Current vs. Time for Electrolysis of AQ27DS. 1 1 0

(Electrolysis potential = -0.6 V vs. SHE. pH = 9.3.)

7 .1 0 Electrolysis with Linked UV-Visible Spectrophotometry. I l l

7 .1 1 Spectra at 15 C Charge Intervals during AQ27DS Reduction. 1 1 3

7 .1 2 Absorbance(330 nm) vs Charge during AQ27DS Reduction. 1 1 4

7 .1 3 Electrochemical ESR Apparatus. 1 1 6

7 .1 4 Flow Profile accross a Tube. 1 1 7

7 .1 5 Normalised ESR signal (S/iiim) vs. V f 2̂ 3. 1 1 8

7 .1 6 Structure of AQ27DS-" 1 1 9

7 .1 7 Actual and Simulated ESR Spectra of AQ27DS-" 1 2 0

Figures 11

8 .1 Efo-pH diagram of the O2/H2O System at 298 K. 12 4

8 .2 The Structure of Hydrogen Peroxide. 1 2 7

8 .3 E^-pH Diagram for the H20 2 /H 20 System at 298 K. 1 2 8

8 .4 Cyclic Voltammograms Showing Oxygen Reduction 1 3 0

8 .5 Experimental and Calculated O2 Reduction Currents at a RDE. 1 3 2

9 .1 Stopped Flow Apparatus. 1 3 4

9 .2 Gold Indicator Electrode for Measuring the Solution Potential. 1 3 5

9 .3 Spectra Taken During Reaction between AQ27DS and HS“. 1 3 7

9 .4 UV-visible Spectrum of Sodium Polysulphide. pH 9.3. 1 3 8

9 .5 Plot of ln(Abs 330 nm) vs. Time During Reduction of AQ27DS. 1 3 9

9 .6 Evans Diagram Showing the establishment of a mixed potential. 1 4 0

9 .7 Measured and Theoretical Solution Potentials vs Time. 1 4 1

9 .8 Spectra Taken During Reaction betweenV(V) and HS~. 1 4 3

9 .9 Spectra of 10 mol V(IV) m"3, before and after aeration. 1 4 5

9 .1 0 E^-pH Diagram for the V-S-H2O System at 298 K. 1 4 6

9 .1 1 Voltammetry of Stretford Solution During Reduction. 1 4 9

Tables 12

List of Tables

1 . 1 Typical Stretford Solution Composition. 2 1

2 . 1 The pKa values of Polysulphides. 3 2

4 .1 51V NMR Chemical Shifts of V(V) Species. 6 S

4 .2 Some Known Vanadium Sulphides. 7 7

4 .3 Spectral Summary of Thiovanadates. 8 0

7 .1 UV-Visible Spectral Summary of AQ27DS Reduction. 1 1 5

A .l AGf° Values for Vanadium Compounds at 298 K. 1 5 8

A .2 AGf° Values for Sulphur Compounds at 298 K. 1 5 9

A .3 AGf° Values for Vanadium Sulphides at 298 K. 1 6 0

/

13


Introduction 14

1. The Importance of Sulphide Oxidation

The oxidation of aqueous sulphide species (H2S, HS', S ") is of considerable

technological importance, and sulphide oxidation processes have been devised for

many applications. The gas industry has developed methods of removing hydrogen

sulphide from fuel gases based on dissolving the gas in an aqueous solution and then

oxidising this solution by aeration. Pollution control and effluent treatment processes

must decrease the concentration of aqueous sulphide species, which would otherwise

cause a loss of the dissolved oxygen in rivers and lakes. Many industries evolve

hydrogen sulphide, controlled oxidation of which can produce elemental sulphur; in

this way a toxic pollutant can be converted into a saleable by-product.

1.1 The Removal of Hydrogen Sulphide from Fuel Gases

When the North Sea gas fields start to become depleted, coal gasification processes to

produce methane (Synthetic Natural Gas or SNG) will need to be developed in the

United Kingdom. All coal deposits contain some sulphur, present in organic and

inorganic forms. Organic sulphur occurs within the coal matrix, in organic compounds

such as thiols. Inorganic sulphur occurs predominantly as inclusions of the mineral

pyrite (FeS2). Although improved mineral processing techniques can reduce the amount

of pyritic sulphur, it is impossible to remove the organic sulphur by physical

processing. Attempts have been made to extract the organic sulphur using chemical

methods; these methods were reviewed by Eliot [1] who concluded that they have only

been partially successful at removing the organic sulphur content and are unlikely to be

implemented on an industrial scale in the near future.

Though coal deposits vary greatly, organic sulphur commonly comprises from 30 to

70 % of the coal's total sulphur content. Thus all coals, and gases derived from coals,

are likely to contain sulphur compounds for the foreseeable future.

In America the oil price rise of the mid-seventies increased interest in developing natural

gas fields containing a high proportion of hydrogen sulphide - sour fields. This led to

developments in gas desulphurisation processes.

In 1984 the Gas Research Institute (Chicago, Illinois) initiated an investigation into

aqueous sulphide oxidation processes that are used to purify both sour natural gases,

SNG and other fuel gases produced from coal. America has large reserves of coal, but

some deposits, especially those in the South West, have a high sulphur content. Coal

gasification with subsequent gas desulphurisation [2] is seen as a means of avoiding

the atmospheric pollution that would otherwise result from direct combustion of these

high sulphur coals; recently the KILnGAS process for producing clean gas from high

sulphur coal has been demonstrated on an industrial scale [3]. The resulting fuel gas

could be used, for example, for electricity generation.

Introduction 15

Combined cycle electricity generation schemes have been proposed. Coal is gasified

and the resulting gas is purified. This gas is then burnt in a gas turbine which

generates electricity, and the exhaust gas temperature is still high enough to raise steam

to drive a conventional steam turbine. As well as offering reduced atmospheric

pollution resulting from sulphur removal, this scheme offers increased generating

efficiencies [4].

During coal gasification, the sulphur in coal is converted into hydrogen sulphide. If

methane is to be produced, this has to be removed, since it would cause poisoning of

the methanation catalysts that are used later in the process. Even if sulphur resistant

catalysts were to be developed, its toxicity, the objectionable odour of hydrogen

sulphide, and its detrimental effect on steel pipelines [5] would still necessitate its

complete removal.

Although other methods have been suggested, such as selectively permeable

membranes [6,7], there are at present three basic ways of removing hydrogen sulphide

from fuel gases:

1. Absorbing the gas in a liquid.

2. Adsorbing the gas on the surface of a solid.

3. Chemically converting the gas to a less toxic product.

1.1.1 Gas Absorption by a Liquid

Gas absorption processes are usually followed by regeneration of the absorbing

solution and release of the hydrogen sulphide, which still requires further processing.

For example, alkanolamines are widely used for absorbing the 'acid gases' hydrogen

sulphide and carbon dioxide [8]. Simplified reactions are:

HORNH2 + H2S -> HORNH3+ + HS- ( 1 .1)

HORNH2 + C 02 + H20 HORNH3+ + H C03- (1.2)

Heating the solutions reverses the above reactions and regenerates the hydrogen

sulphide and carbon dioxide.

The many other liquid absorption processes, which remove hydrogen sulphide and

carbon dioxide from gases, have been reviewed recently [9]. The Rectisol process

absorbs hydrogen sulphide into cooled methanol under high pressure, and then releases

the gas when the pressure is decreased [10], and the Potash Vacuum and Benfield ‘

processes rely on the absorption and desorption of hydrogen sulphide by aqueous

solutions of potassium carbonate [10,1 1 ].

All the above processes suffer from the disadvantage that the acid gases are only

concentrated, and not converted into non-toxic compounds. Although carbon dioxide

can be vented safely to the atmosphere, hydrogen sulphide cannot, and so it must

undergo further chemical treatment. The most common processes involve partial

oxidation with atmospheric oxygen to produce elemental sulphur.

Introduction 16

There are various ways of achieving this oxidation, the earliest of which was

developed last century by Claus, and is still in use today. Hydrogen sulphide is split

into two streams. One stream, consisting of a third of the gas, is combusted with air to

produce sulphur dioxide:

H2S +3/2 0 2 -> S02 + H20 (1.3)In a subsequent reaction chamber the sulphur dioxide that is produced acts as an

oxidising agent for the remaining hydrogen sulphide, and forms elemental sulphur:

2 H 2S + S 0 2 <-> 3/2 S2 + 2 H20 (1.4)

Reaction (1.4) is an equilibrium; at the temperature occurring in the hydrogen sulphide

combustion furnace, the equilibrium lies to the left. The combustion products must be

cooled to about 650 °C in order to shift the equilibrium to the right.

The sulphur vapour is condensed, and after further purification the sulphur can be sold.

However, at 650 °C, the equilibrium mixture still contains appreciable quantities of

hydrogen sulphide. Repeating the reaction scheme can reduce this quantity further but

the thermodynamics of the process militate against the complete removal of hydrogen

sulphide.

Increasingly stringent environmental regulations have meant that conventional Claus

processes now require tail gas purification units before the off gases can be vented.

This has meant that Claus Units have become more complex and expensive, although

they are still widely used [10].

1.1.2 Gas Adsorption on a Solid

Processes involving gas adsorption on solids (followed by chemical reaction), have

been widely used to remove the hydrogen sulphide from coal gases. One historically

important method used dry iron (III) oxide; for which the adsorption reaction that is

quoted by Kohl and Riesenfeld [10] is:

^ e2^3 3 H2S —> Fe2S3 3 H20 (1.5)

However, an iron (1H) sulphide phase has never been identified, and it is likely that the

adsorption takes place with simultaneous reduction of the iron(III). Thus the phase

Fe2S3 may be better regarded as a mixture of FeS2 and FeS:

Fe20 3 + 3 H 2S -» FeS2 + FeS + 3 H20 (1.6)

Periodically, air is blown through the sulphidised bed. This oxidises the 'Fe2S3',

producing elemental sulphur and regenerating the iron (III) oxide:

F̂ 2̂ *3 + 3/2 0 2 —> Fe203 3 S (1.7)The sulphur forms around the iron (III) oxide particles and eventually prevents further

reaction. Fouled beds contain between 40 and 50 % sulphur, which in principle can be

recovered. However, the beds were commonly discarded, or combusted to yield

sulphur dioxide for sulphuric acid manufacture.

Introduction 17

Zinc oxide filters [12] have also been used to remove trace amounts of hydrogen

sulphide:

ZnO + H2S -> ZnS + H20 (1.8)

The sulphidised bed cannot be easily converted back to zinc oxide, so zinc oxide filters

are used only when complete elimination of sulphur is required; they are used, for

example, as 'guard tubes' to protect catalysts.

Recently silica gel has been suggested as a selective adsorbent to remove hydrogen

sulphide from biogas [13], but the process has not been demonstrated on a large scale.

Activated carbon filters [14] have also been proposed; oxidation of the adsorbed

species produces elemental sulphur, which eventually deactivates the surface. The

carbon can be reactivated by contact with steam, but it is difficult to make the process

continuous. Again, no industrial applications of this principle have yet been

implemented.

1.1.3 Electrochemical Oxidation of Hydrogen Sulphide

Hydrosulphide ions can be oxidised at an anode to form either free sulphur, as in

reaction (1.9), or polysulphide ions (S22',S32_,S42“ and S52") by reactions such as

(1.10):HS‘ —> S + H+ + 2e" (!-9 )

2 HS- -> S22‘ + 2 H+ + 2 e- ( 1 .10)

At the cathode, the reduction of protons produces hydrogen; a by-product which can

create income to offset the cost of the electrical energy required.

2 H+ + 2e- -> H2 (1.11)

A process for the direct electrolysis has been proposed by Bolmer [15], but the

sulphur produced can passivate the anode surface; addition of a sulphur solvent at

85 °C has been suggested to prevent this deactivation. A second problem with direct

electrolysis is that polysulphide ions can diffuse to the cathode where they can undergo

reduction, thus decreasing the current efficiency.

Dandapani, Sharifker and Bockris [16] showed that the use of elevated temperatures

(85 °C) and high sodium hydrosulphide concentrations enabled the electrolysis to

proceed without passivation; polysulphide solutions were produced that increased in

concentration until elemental sulphur precipitated. Cation exchange membranes

prevented the polysulphides from reaching the cathode, and high current efficiencies

were reported.

Lim and Winnick investigated the electrolysis of hydrogen sulphide when it was

dissolved in a molten potassium sulphide/sodium sulphide mixture [17,18]. The

electrolysis was operated at a temperature of around 800 °C, and high current densities

were achieved on graphite electrodes. The anode compartment was purged with

hydrogen, and the sulphur vapour produced reacted with this to re-form hydrogen

sulphide. The cell thus acted as an electrochemical hydrogen sulphide concentration

Introduction 18

device. They proposed that this process would be suitable for desulphurising gases

that would be subsequently fed to molten carbonate fuel cells, which also operate at

high temperatures. However, one problem with the process was the absorption of

carbon dioxide by the sulphide electrolyte (forming alkali metal carbonates).

Indirect electrolysis has also been suggested; at the anode an oxidant is generated that is

capable of oxidising HS' ions in a subsequent chemical step. Kalina and Maas [19]

used electrochemically generated iodine (present as 13“ in the iodide solution) as the

oxidant:

At Anode: 31- —» I3- + 2e* ( 1 .12)

At Cathode: 2 H+ + 2 e" h 2 (1.13)

In Electrolyte: Is' + H2S —» 2H+ + 31- + S (1.14)

Overall reaction: h 2s —̂ H2 + S (U 5 )

Olson [20] proposed a similar scheme based on the electrochemical production of an

iron (HI) complex; hydrogen sulphide was absorbed and oxidised in one vessel and the

iron (II) produced re-converted to iron (III) in an external electrochemical cell.

All electrochemical routes suffer from the cost penalty of utilising electrical energy

rather than using oxygen as the oxidant. Efficient electrolysis has been claimed using a

cell voltage of only 0.5 V on a laboratory scale [16], but this still represents an energy

requirement of 837 kWh per tonne of sulphur produced. To date no electrochemical

process for H2S removal has been operated on an industrial scale.

1.1.4 Aqueous Oxidation of Hydrogen Sulphide

Aqueous oxidation processes ensure that the hydrogen sulphide is converted into a

non-toxic product, and not merely concentrated. These methods offer the advantage

that, unlike the Claus Process, they can remove the hydrogen sulphide completely. In

contrast to the solid adsorption processes, they are easily adapted for continuous use,

and they can utilise the oxidising power of atmospheric oxygen rather than consuming

expensive electrical energy.

Aqueous oxidation processes employ solutions which contain oxidising agents capable

of producing elemental sulphur from hydrogen sulphide; the reduced solutions are then

re-oxidised with air and recycled, so that the oxidising agents complete a redox cycle

and are not consumed. The overall reaction for all these process is given by equation

(1.16):

H2S + 1/2 0 2 -» S + H20 (1.16)

Various oxidising agents have been used to catalyse this reaction: iron (III) salts, with a

suitable sequestering agent, are used in the Low -Cat process [21]; arsenic (V)

compounds are used in the Thylox and Vetrocoke processes [10]; and organic

oxidants such as quinones are used in the Perox and Takahax processes [10].

Introduction 19

Amongst these alternatives, the Stretford Process is one of the most commercially

successful; the process uses an absorbing solution containing vanadium(V) salts and

soluble anthraquinone derivatives.

1.2 The Stretford Process

The Stretford process was developed in the 1960's to oxidise hydrogen sulphide in coal

gas to sulphur using an aqueous solution. The process was developed at the North

West Gas Research Laboratories at Stretford, near Manchester. The process is shown

schematically in Fig. 1.1 There are three main components: the absorber, the reactor,

and the oxidiser.

The absorber is a gas/liquid contacting device which ensures that any hydrogen

sulphide in the gas stream is absorbed into the solution. Since the pKaj of hydrogen

sulphide is about 7, and the absorbing solution is buffered at around pH 8.5, the gas is

absorbed according to reaction (1.17):

H2S(g) <-» H2S(aq) —» HS" + H+ (1-17)

The carbonate/bicarbonate buffer solution prevents the protons released from lowering

the pH.

Fig. 1.1 The Stretford Process.

In the reaction tank and the oxidiser, the Stretford process achieves the oxidation of this

hydrosulphide ion to sulphur:

HS“ +l/2 0 2 -> S + OH- (1.18)

Air is blown through the solution in the oxidiser and the sulphur produced, which is

naturally hydrophobic, is carried to the surface by the rising air bubbles. Thus,

aeration serves the dual purposes of oxidising the solution and carrying the sulphur

particles to the surface, where the sulphur-containing froth can be skimmed off and

Introduction 2 0

filtered. However, reaction (1.18) proceeds slowly using atmospheric oxygen without

a catalyst, and higher oxidation products (such as thiosulphate, sulphite, and sulphate)

tend to be produced. This constitutes an effluent problem which would otherwise be

absent.

It was found that using a solution containing vanadium (V) salts and anthraquinone

derivatives increased the rate of reaction greatly. It is thought [2 3 ] that the

vanadium (V) salts are responsible for the hydrosulphide oxidation, according to

simplified reactions such as (1.19):

HS- + 2V5+ + OH- -> S + 2V4+ + H20 (1.19)

If this is the case, then the vanadium(V) salts should not strictly speaking be termed

catalysts, since they are consumed stoichiometrically according to equation (1.19);

however, they are regenerated in a subsequent aeration tank, reaction ( 1 .20), and so

take part in a catalytic redox cycle:

2V4+ + 1/2 0 2 + H20 -> 2V5+ + 2 OH- (1.20)

The anthraquinone derivatives that are added are said to catalyse the reoxidation of

vanadium (IV) to vanadium (V) by atmospheric oxygen [23].

The overall reaction occurring in the process is the same as that given for other aqueous

oxidation processes:

H2S + 1/2 0 2 -> S + H20 (1.16)

Since this reaction does not involve the production or consumption of protons, no

permanent pH change will occur. However, since gas absorption, hydrosulphide

oxidation and vanadium (V) regeneration occur in different vessels, local changes in pH

would be expected.

1.2.1 The Historical Development of the Stretford Process

The development of the Stretford process of aqueous sulphide oxidation has been

reported in some detail by Vasan [24], Moyes and Wilkinson [23], and Nicklin and

Holland [25]. The major stages in the history of the process are as follows:

In 1963 the Stretford Process was developed jointly by the North Western Gas Board

and the Clayton Aniline Company Limited; it was intended for use on streams of coke

oven gas, and utilised a solution containing sodium anthraquinone disulphonate in an

alkaline sodium carbonate/bicarbonate buffer. The original process did not use a

solution containing vanadium (V) salts. Three plants were built to this design and

worked satisfactorily on a feed of gas containing 0.08 % (by volume) hydrogen

sulphide, but the hydrosulphide ion concentration in the working solution could not

exceed 1.25 mol m“3.

In an attempt to increase the maximum HS" concentration that could be oxidised, and to

reduce the reaction time, which was about half an hour for the first generation of plants,

various oxidising agents were investigated for use in the process. Vanadium (V) was

Introduction 2 1

chosen for study since it did not to form a sulphide precipitate under plant operating

conditions. It was found that the vanadium (V) was an effective oxidising agent for the

hydrosulphide, but that the vanadium (IV) species produced could not easily be re

oxidised by aeration alone. However, if anthraquinone disulphonate salts were

present, regeneration of the vanadium (V) was rapid. Using this system,

hydrosulphide solutions of concentration 30 mol m-3 could be oxidised in several

minutes. This enabled a much smaller volume of liquid to be re-circulated in order to

achieve the same gas throughput rates as obtained in the first generation of plants. In

this basic form, the Stretford process has remained in use to the present date and there

are now over 100 installations worldwide [26]. The composition of a typical Stretford

solution is given in Table 1.1 (The Data is taken from Murin et A1 and Mallot

[2 6 ,2 7 ]):

Com pound kg n r 3 mol m

Sodium anthraquinone 2,7-disulphonate (Na2AQ27DS) 3.00 7.3

Sodium vanadate (NaV03) 1.70 32.5

Sodium citrate (Na3C 02CH2C (0H )C 02CH2C 02) 10.00 52.1

Sodium carbonate (Na2C 03) 6.25 59.0

Sodium hydrogen carbonate (NaHC03) 18.75 223.0

Sodium thiosulphate (Na2S20 3) (variable) 89.50 360.0

Sodium sulphate (Na2SC>4) (variable) 40.2 283.0

Sodium thiocyanate (NaSCN) (dependent on HCN in feed)

Table 1.1 Typical Stretford Solution Composition.

Although the Stretford Process has been one of the most successful methods of

hydrogen sulphide removal, several problems still exist.

1.2.2 Operational Problems Experienced by Stretford Plants

Sometimes elemental sulphur can form in the absorbing vessel, which can build up so

as to restrict, and eventually block the gas flow [26]. If the solution is over-oxidised,

the production of soluble sulphoxy compounds, primarily thiosulphate (S20 32")

results. The thiosulphate concentration can build up in the solution, and if left

unchecked, sodium thiosulphate would precipitate. To prevent this happening, a

portion of the solution must be discarded or treated (for example in a fixed salt recovery

unit), and fresh liquor added. Thiosulphate production also consumes hydroxide ions:

2 H2S + 2 0 2 + 2 OH" S20 32" + 3 H20 (1.21)

Therefore, alkali has to be added continually to the solution to prevent the pH from

decreasing.

Fixed salt recovery units prevent the loss of sodium vanadate; working solution that

has been withdrawn from the re-circulating circuit is incinerated in a reducing

atmosphere. This converts the thiosulphate and thiocyanate into gaseous hydrogen

sulphide, ammonia and carbon monoxide; these gases are fed to the sour gas input.

Introduction 2 2

Solid sodium vanadate and sodium carbonate are also produced; which can be used to

make up fresh solution. Sodium citrate and sodium anthraquinone disulphonate

(Na2AQDS) are both destroyed by the reductive incineration. The Na2AQDS is

expensive and this loss constitutes one of the main operating costs of the process.

In several plants a black solid, containing vanadium, sulphur and oxygen precipitated

from the solution. Chemical analyses revealed that samples originating from different

plants had different compositions. It was unclear whether the sulphur was chemically

combined in the compound, or whether it was physically entrained in the precipitate. It

was claimed that the compound precipitated from solution when the pH was allowed to

rise above 9, and that a high concentration of carbonate ions promoted precipitation

[26]. One solution to this problem was to increase the flow rate to keep particles

suspended until they entered the aeration vessel. It was noted that on prolonged

aeration the compound redissolved. Complexing reagents have also been added to a

number of Stretford solution in the hope of preventing the vanadium from precipitating.

Tartrate ions, citrate ions, and di-sodium ethylenediamine tetraacetate (Na2EDTA) have

all been used industrially. These compounds are claimed to form complexes with the

vanadium (V) species, but in a recent study Haley [28] contradicted Malott [27], by

asserting that no complex was formed between vanadium (V) and citrate ions, and only

a weak complex was formed with tartrate ions.

If hydrogen cyanide is present in the feed gas, it can react with elemental sulphur to

form the thiocyanate ion:

HCN + S + OH- SCN- + H20 (1.22)

This reaction consumes alkali, and the presence of thiocyanate is also thought to

increase the conversion of hydrogen sulphide to thiosulphate. To combat this, pre

washing the gas to absorb hydrogen cyanide (for example by contacting the gas with

polysulphide solutions) was introduced.

It was found that bacteria present in the Stretford solutions could affect the plant's

performance adversely. Analyses showed that most solutions contained 1012 bacterial

cells m-3. The bacteria included autotrophic sulphur bacteria of the genus th io b a c illu s;

these can oxidise thiosulphate in solution to produce sulphuric acid, which consumes

the sodium carbonate and bicarbonate. Bacteria cells are also encapsulated in a slime

layer which can break away and cause the solution to foam when it is aerated. These

problems caused British Gas to investigate the addition of various biocides, which

proved effective at controlling the bacterial population. Microbial problems were not

apparent in the early coke oven gas desulphurisation plants; this is thought to be due to

the biocidal concentrations of thiocyanate (SCN- ) formed from the hydrogen cyanide

in the feed [29].

Introduction 23

1.2.3 Mechanistic Studies

Working from free energy of formation data by Israel and Meites [30], the calculated

value of Eo' for the vanadium (V)/(IV) couple, at pH 8.5, is +0.05 V vs SHE. At this

pH the indicated vanadium (V) species are HV2O73- ions, and the vanadium (IV)

species V4O92" ions. The standard reduction potential of the sulphate/hydrosulphide

couple at this same pH is only -0.20 V, [31]. This means that vanadium (V) is

thermodynamically capable of oxidising hydrosulphide ions to sulphate.

It is clear from the literature that disagreement exists as to the nature of the vanadium

(V) and vanadium (IV) species existing in the Stretford solution. Malott [27] argued

that the vanadium (V) species HjjV O ^3'11)-, HnV207(4"n)“, V3093" and V4O124" may

all be present in the solution. Habayeb and Hileman [32] suggested V2094",V3093‘

and V4O124" are the major species, whilst many workers have simply assumed that the

oxidising species is V5+. Malott suggested that V2052- may be the predominant

vanadium (IV) species, but Pope [33] claims that V jg C ^12" is the major vanadium

(TV) form in alkaline solution.

Aeration of alkaline sulphide solutions can result in a variety of reaction products, as

Kuhn and Kelsall [34] pointed out in their review. The rate of reaction and the

oxidation products are highly dependent on the pH, the solution potential (determined

by the dissolved oxygen content) and the temperature. Catalysts can further change the

product mixture.

Andrzheevskii [35] undertook a study of the mechanism of the oxidation process

occurring in the Stretford Process. He studied the oxidation of sodium sulphide

solutions at pH 9.0 using three solutions containing: 0.1 kmol m-3 sodium

anthraquinone 2,6-disulphonate (Na2AQ26DS), 0.1 kmol m-3 sodium metavanadate

(NaVC>3 ), and a mixture of composition similar to that used in working Stretford plants

(12.5 mol Na2AQ26DS m-3, 40 mol NaVC>3 m"3)- A set of experiments were

conducted in the absence of air or dissolved oxygen, adding sulphide to the above

solutions to make the concentration 50 mol HS_ m-3. The reaction was followed by

monitoring the remaining hydrosulphide concentration. He analysed the samples using

two methods: a potentiometric titration using mercury (II) nitrate; and a 'chemical'

method using cadmium acetate in acetic acid. Unfortunately, no experimental details of

either analysis method were given. Andrzheevskii found that the hydrosulphide

concentrations as measured by the two methods were identical initially, but as the

reaction proceeded, the cadmium acetate method indicated a rapid removal of HS-

(decreasing from 50 mol m-3 to zero in 10 minutes), whereas the potentiometric

titrations suggested a much slower rate of reaction (after 10 minutes 40 mol HS- m-3 remained).

Introduction 2 4

Andrzheevskii argued that this was evidence for the HS- being present in two forms,

one free in solution, the other in a complex with vanadium. He proposed that the

lowering of the sample pH during the cadmium acetate analysis caused rapid oxidation

of the complexed hydrosulphide, while the potentiometric titration recorded the total

[HS-] in both complexed and uncomplexed forms.

Without details of Andrzheevskii's analysis techniques it is difficult to interpret his

results. However, Boulege [36] utilised mercury (II) chloride solutions to titrate

solutions containing HS-, Sn2_, S2O32-, and SO32-. He states that at pH 13.0 the first

end point (detected by a sulphide selective electrode) corresponds to the completion of

the two reactions:

Hg2+ + HS- -> HgS + H+ (1.23)

Hg2+ + Sn2- -> HgS + (n-l)S (1.24)

This end point is detected by the sharp decrease in potential as the sulphide selective

electrode responds to the decrease in [S2-]. If the pH is then adjusted to 7-8, two

further end points can be detected corresponding to equations (1.25) and (1.26):

Hg2+ + 2 S 20 32- -> Hg(S20 3 )22- (1.25)

Hg2+ + 2 S 0 32- -> H g(S03)22- (1.26)

These end points are detected because free Hg2+ ions in solution are known to interfere

strongly with the response of the electrode, which causes a further decreases in

potential.

Andrzheevskii carried out his potentiometric titration at pH 9.0, and so it is not clear

which reactions occurred. Certainly the total sulphide and polysulphide

concentrations will be recorded, and he may also have titrated dissolved S2032- and

SO32-. His analysis method using acidified cadmium acetate almost certainly involves

precipitation of cadmium sulphide. Without experimental details, it is unclear how

polysulphide ions, thiosulphate and sulphite ions will react; it may be that they are

oxidised during the analysis. Thus these species may account for the discrepancy

between his two analyses, and his claim that some of hydrosulphide must be present in

the form of a complex must be regarded as speculative; however, Harrison and

Howarth have recently obtained 51V NMR evidence for complex formation between

HS“ and vanadium (V) [37].

Andrzheevskii's work reveals some interesting observations: oxidation of HS* using

stoichiometric amounts of sodium metavanadate or Na2AQ26DS (assuming one

electron reduction in both cases) proceeds relatively slowly, in around 90 minutes; a

mixed solution (containing a third molar excess oxidising power) oxidised the solution

much more rapidly, in about 15 minutes, and when air was admitted oxidation

proceeded even more rapidly. If the same solution was used repeatedly to oxidise

samples of hydrosulphide solution, with oxygenation used to restore the oxidising

power, as in the industrial process, Andrzheevskii noted that the solution lost its ability

Introduction 25

to oxidise the HS~ ions in the absence of air. Over 60 minutes, no loss of HS~ was

detected (using the potentiometric titration). However, if oxygen was then admitted,

complete oxidation was effected in 10-12 minutes. It was stated - without the

supporting evidence, that prolonged oxygenation of these solutions for 40 minutes did

not regenerate detectable amounts of vanadium (V), although this contradicts other

workers [26]. No detailed product analyses were given in Andrzheevskii's work, but

the following generalisations were offered:

1. Using solutions of NaV03 or Na2AQ26DS with oxygen as the oxidising

agent, thiosulphate was the main oxidation product.

2. Using a mixed solution with oxygen, elemental sulphur was the main

oxidation product.

In 1984 a Research programme was started by the Gas Research Institute

(Chicago: USA) into the Stretford Process [26]; they have built a bench-scale

circulating flow unit in the hope of determining the optimum operating conditions for a

Stretford Plant. They also acknowledged that the complex chemistry of the process is

not well understood, and considered that the key to improving the process performance

lies in a better understanding of the basic chemistry.

Thus, previous workers have not demonstrated unambiguously the mechanism of

oxidation of hydrogen sulphide in the Stretford Process, although many studies of the

process have been undertaken. Since 1963, some operating problems have been

overcome by using a practical experimental approach, but little of the fundamental

chemistry has been elucidated.

1.3 Objectives of the Present Study

The aim of the present study is to elucidate the reaction mechanism, in the hope that this

can help both to solve operational problems (such as the formation of thiosulphate and

the precipitation of vanadium salts) and to point the way towards future process

improvements.

1.3.1. Research Approach

The approach taken was to investigate the redox couples involved in the process

separately and then to consider the interactions between these couples. The couples;

S(-II)/S(0), V(V)/V(IV), anthraquinone/anthraquinol and 0 2/H20 were studied using

electrochemical techniques such as cyclic and pulse voltammetry. Products were

identified using UV-visible and esr spectroscopy. Finally, the redox chemistry of the

process was investigated by using stopped flow spectrophotometry and conducting

small scale batch experiments on solutions containing two or more of the redox

couples.

Review of Sulphide Oxidation 2 6

2. Review of sulphide OxidationSulphur has a range of oxidation states from -II to VI:

-II

h 2s

HS-

0s S20 32-

II IV

S 0 32-

h s o 3"

VIS 042-

h s o 4_

Fig. 2 .1 Possible valence states of sulphur in aqueous media,

('per' compounds are omitted)

Sulphide solutions represent the lowest oxidation state of sulphur, and in theory they

can be oxidised to any of the higher states. However, only the -II, 0, and +VI states

are thermodynamically stable in aqueous solution at normal temperatures and pressures.

Balanced redox reactions between the various species can involve the production or

consumption of protons, e.g.

Thus at high pH's the forward reaction is favoured, i.e. the oxidation can be achievedbe

using a relatively low potential. Certain acid-base equilibria may also^mportant, e.g.

The values for the equilibrium constants of these equations are not always known with

great certainty. The second acid dissociation constant K2 for equation (2.3) has values

reported as far apart as 10-13 and 10-19 [38]; however, at a pH of around 9 there is no

doubt that the predominant solution sulphide species are hydrosulphide ions.

The available thermodynamic information can be summarised in the form of an E^-pH

diagram, which can be used to predict the most stable species at any given E^ and pH.

The solution potential, E^, can be applied electronically at an electrode surface, or by

adding a redox couple to the solution. In the latter case the potential at equilibrium will

be give by the familiar Nemst equation, and can be measured with a suitable indicator

electrode (e.g. a platinum wire):

E = reversible potential vs. SHE / V.

E° = standard reduction potential vs. SHE / V.

z = number of electrons transferred; F= Faraday's constant / C (mol electrons)-1 R = Gas Constant / J mol-1 K_1; T = Temp. / K.

ar, a0 = activities of reduced and oxidised species respectively.

H2S S + 2H+ + 2 e - (2 .1)

H2S(aq) <-> HS- + H+

HS- <-> S2- + H +

(2.2)

(2.3)

E = E° + RT ln (a 0 )

zF ar

(2.4)

Review of Sulphide Oxidation 27

The thermodynamics of the sulphur - water system were first summarised in the form

of an E^-pH diagram by Valensi [39]. His diagram is shown in Fig. 2.2. Notice

that there are only three stable oxidation states.

Fig. 2.2 Eh-pH diagram for the sulphur/water system at 298 K.

Sulphur Species present at unit activity [39].

Oxidation of hydrosulphide at pH 9.0 from Eh = -0.4 V to +0.1 V would be predicted

to form sulphate. This prediction that sulphate will be the main oxidation product is in

clear disagreement with the findings of numerous workers who have studied the

oxidation of sulphide solutions in these potential and pH regions.

The reason for this apparent contradiction is that E^-pH diagrams are based solely upon

the assumption that equilibration between species is taking place under thermodynamic

control, and no account is taken of the rate of the possible reactions. Reactions

producing sulphate ions, for instance, are known to proceed at a very slow rate during

the atmospheric oxidation of sulphide solutions. Some account of these kinetic factors

can be made by excluding from the diagram species which are known to form very

slowly. This can be considered equivalent to adding to their free energy of formations

in order to compensate for their large activation energies. Peters [40] produced a

diagram from which sulphate species are excluded, Fig. 2.3 (overleaf).


Fig. 2.3. Eh-pH diagram for metastable sulphur system at 298 K. [40].

A potential change from = -0.4 V to -0.1 V, again at pH 9, would now predict that

sulphide will be oxidised to thiosulphate, and that a further increase to + 0.1 V would

yield sulphite:

2HS- + 8 OH- -+ S20 32- + 5 H 20 + 8 e- ( 2 . 5 )

S20 32- + 6 OH- -> 2 S 0 32- + 3 H20 + 4e- (2.6)

Hamilton and Woods [41,42] studied the electrochemical oxidation of sulphide

solutions at a gold electrode, and found that under mildly oxidising potentials (+0.2 V

vs. SHE at pH 9.2) the rate of production of all sulphoxy compounds was negligible.

They produced an E^-pH diagram with all such species excluded, Fig. 2.4 (overleaf),

and used it to explain their experimentally observed oxidation products: polysulphide

ions and elemental sulphur.


Fig. 2.4 Eh-pH diagram for the sulphur/water system at 298 K

Oxy-sulphur anions not considered [42].

Thus, three Eh-pH diagrams can be drawn for the sulphur/water system: one which

considers all thermodynamically stable sulphur species; one which considers

metastable sulphur and sulphoxy species; and one which only considers metastable

sulphur species. These diagrams would predict the sulphide oxidation product to be

sulphate, thiosulphate and sulphite, or sulphur respectively.

2.1 The Oxidation states of Sulphur

The redox chemistry of sulphur is a rich and complex area of study, and complete

textbooks have been written on this subject alone [43]. This review will be confined

to the aqueous sulphur redox chemistry, and will concentrate on those studies which

have been carried out in alkaline solutions.

2.1.1 Sulphide (-II)

Below pH 7, sulphide (-II) exists as the species H2S(aq), which is moderately soluble:

H2S(g) H2S(aq) K = 0-101 (2.7)This means that under a partial pressure of hydrogen sulphide of one atmosphere,

aqueous concentrations of up to 100 mol n r3 can be achieved. Above pH 7 the

dominant species are the HS' ions:

^ 2^(aq) ^ HS“ + H+ pK ai = 6.99 (2.8)


Using free energy data from Zhdanov [31], the [HS-] in solution can be calculated

according to equation (2.9):

H2S(g) + OH- <-> HS- + H20

log[HS-] = pH - 7.995 + lo g P ^ s (2.9)

For instance at pH 8.5 a partial pressure of hydrogen sulphide of only 0.01 atm. would

be in equilibrium with a solution containing 30 mol HS- m"3.

The HS- ion can deprotonate further to form the sulphide ion:

HS- <-> S2- + H+ p K ^ = 13-19 (2.10)

Until recently the value for pKa2 was accepted at around 13, implying that strongly

alkaline sulphide solutions contained the S2_ species. However, evidence has now

accumulated to suggest a value for pKa2 of 19 ± 2 [38, 44-46]. An early calculation

of pKa2 [47] relied on UV-visible spectrophotometric measurements; assigning an

absorbance band at 230 nm to the HS" ion and a band at 360 nm to the S2_ ion.

Giggenbach [45] pointed out that the absorbance band at 230 nm due to HS" was

subject to a blue shift on increasing the hydroxide ion concentration, which would

produce an apparently falling absorbance value if measurements were made at a single

wavelength. He suggested that the weak absorbance at 360 nm was due to the presence

of polysulphide ions, which had been formed through the partial oxidation of the

sulphide solution. He failed to observe this band, previously assigned to the S2_ ion,

when oxygen was completely excluded, and from his own measurements suggested a

value of p K ^ of around 17.1. This result was largely ignored for the next decade, but

Meyer [38] later confirmed the higher value using IR Raman spectroscopy. He

identified the presence of the HS" ion by its peak at 2570 cm"1, and found that it was

present even in 16.9 kmol m-3 NaOH.

It now seems that a high value for p K ^ will have to be accepted, which has many

implications [44]. The S2_ ion will not be a predominant species in aqueous solutions,

even in strongly alkaline media. The value for AGf° for S2_ must be revised, together

with any associated thermodynamic values calculated from this, which means that the

solubility products of many metal sulphides will be even lower than previously

thought. The formation of mercury sulphide has limited the value of polarographic

investigations into the oxidation of sulphide solutions; HgS is formed at lower

potentials than those at which the sulphide is oxidised [48].

2.1.2 Polysulphides (-1 to 0)

Polysulphides are the low chain length sulphur di-anions; S22", S32-, S42-, S52-. They

have formal oxidation states intermediate between -I and 0. Alkali metal polysulphides

(e.g. Na2S4) can be prepared [49,50] and they will dissolve readily in water to form

bright yellow solutions.


Polysulphides can also be formed by the dissolution of elemental sulphur in alkaline

sulphide solution:

(n -l)S + HS- + OH- -> Sn2' + H20 (2.11)

Polysulphides are more easily oxidised by atmospheric oxygen than sulphide species

[5 1 ], forming S, S2C>32-, SO32- and SC>42_. Polysulphides are also formed as

intermediates during the oxidation of sulphide solutions, especially around pH 7.

Partially oxidised solutions develop a yellow-green colouration which is due to

polysulphides.

It has been known for a long time that polysulphides in solution are always present in

equilibrium with each other, even when solutions are prepared from solids with a fixed

stoichiometry [50,52-53]. Giggenbach [53], working at an ionic strength of 2,

derived a set of equilibrium constants for the reactions:

3 S 52- + HS- + OH- —̂ 4 s 42' + h 2o Kcq == 2 x 10-4 (2.12)

2 S42' + HS- + OH- —̂ 3 S 32- + h 20 Keq == 1.8 x 10-2 (2.13)

S32’ + HS- + OH- —) 2 S22- + h 20 Keq == 4 x 105 (2.14)

These equilibrium constants can be used to determine the concentration of an individual

polysulphide ion in a solution of known sulphur to sulphide ratio and pH (see section

3.5.1). Nevertheless Power and Richie [49] ignored the above equilibria and simply

assumed that a solution made up from Na2S4 contained only the anion S42-.

Using these equilibrium constants enables the concentrations of individual polysulphide

ions to be calculated; the results show that tetrasulphide is usually the major species in

aqueous solution and that HS" ions are present at a comparable concentration.

Schwarzenbach and Fischer's earlier study is in agreement with these findings [50].

The average chain length in poly sulphide solutions that are saturated with sulphur is

never greater than five.

Since polysulphides are thermodynamically unstable in alkaline solution, there exists

the possibility of spontaneous disproportionation to produce thiosulphate:

4 S42' + 8 OH" + H20 <-» 3 S20 32- + 10 HS- (2.15)

The equilibrium constant and the rate of this reaction have been measured [54]; the

forward rate is slow, but increases significantly with temperature. However,

polysulphide solutions at room temperature show no noticeable change in their UV-

visible spectra even after months of storage.

Polysulphides protonate to form the polysulphanes, which are reported to be yellow

solids. There is an experimental difficulty in determining the first and second

dissociation constants of the polysulphanes, since in the course of an acid titration

polysulphide species can disproportionate to form HS" and elemental sulphur:

Sn2- + H+ -> (n -l)S + HS- (2.16)


Schwarzenbach and Fischer [50] used a continuous flow technique to achieve the

mixing of acid and polysulphide solutions, and determined the pH downstream from

the mixing vessel, after the poly sulphide solution had only been acidified for 10 ms.

They claimed that this time was too short for the polysulphide to disproportionate, and

from the resulting titration curves determined the following pKa values:

Species P K a l P K a2H2S2 5 .0 9 .7

h 2s 3 4.2 7 .5

h 2s4 3 .8 6 .3

h 2s 5 3 .5 5 .7

Table 2.1 The pKa values of Polysulphides [50].

The above values imply that at pH 8.5 all the polysulphides would be present in

aqueous solution as their dianions except the disulphide, which would be present as

HS2~ Disulphides normally constitute only a minor component in a polysulphide

mixture.

Lessner et al [55] studied the electrochemical redox behaviour of polysulphide

solutions at pH 12. They found that using slow sweep voltammetry at platinum and

cobalt electrodes, sulphur was deposited during the positive going scans. This current

peak and the associated electrode passivation masked the oxidation reactions that

produce higher polysulphides, eg:

5 S 42- -» 4 S 52- + 2e- (2.17)

Upon negative going potential scans they found that hydrogen was evolved before a

well defined diffusion limited current peak due to polysulphide reduction was

observed.

Using voltage pulse methods, the same authors found that the resulting currents were

much smaller than those expected from the calculated concentrations of tetrasulphide.

They concluded that the electroactive species were a minor component in the

equilibrium mixture, and suggested that they were supersulphide ions, S2", which are

known to be produced from tetrasulphide ions at elevated temperatures [54]:

S42- 2 S2‘ (2.18)

Using the equilibrium constant of this reaction Lessner et al [55] predicted

supersulphide concentrations that were consistent with their observed currents over the

temperature range 25-80 °C.

2.1.3 Elemental Sulphur (0)

Pure sulphur exists at room temperature in the crystalline orthorhombic form,

consisting of stacked layers of puckered Sg rings. Above 368 K it transforms into the

monoclinic phase, which is stable up to the melting point of 392 K.


When molten sulphur is heated to 457 K, the Sg rings break and chain molecules result;

if this liquid phase is then quenched rapidly amorphous sulphur is formed. This form

of sulphur deforms plastically and can be stretched to several times its original length.

It has also been suggested that amorphous sulphur can also result from the

electrochemical oxidation of metal sulphides [56]. Sulphur is insoluble in water and

has a high electrical resistance (resistivity 1.9 x 1015 Q m); sulphur coatings therefore

passivate electrodes. Colloidal sulphur is a poorly defined material [43]; it can contain

polythionates of the type SO3" -Sn-S03~, where n has a value of 10 to 20.

It is clear that all reactions of Sg must first require ring scission, which demands a

considerable activation energy; the S-S single bond strength is 226 kJ mol-1 [57].

Therefore, elemental sulphur is resistant to further oxidation, and in many systems the

formation of sulphur is irreversible. Habashi and Bauer [58] found that elevated

temperatures and a high pressure of pure oxygen were required to effect the complete

oxidation to sulphate.

2.1.4 Polythionates (0 to IV)

Polythionates have the general formula (OgS-Sn-SOg)2-, the best characterised ions are

those having n = 1-4. They can be prepared by reducing sulphurous acid with

hydrogen sulphide, a process which produces a complex mixture: Wackenroder’s

solution. Tetrathionate is produced quantitatively by the oxidation of thiosulphate with

iodine:

2 S2032" + I2 —> 2 T + S40 62- (2.19)

In acid solution the polythionates disproportionate to give S, SO2 and SO42'.

2.1.5 Thiosulphate (II)

The thiosulphate ion has the structure S-SOg2- [59], so the two sulphur atoms have

differing chemical environments. Thiosulphate solutions disproportionate in acid

solution to give elemental sulphur and sulphur dioxide:

S20 32- + 2 H+ -> H20 + S + S 0 2 (2.20)

In alkaline solution the reverse reaction can occur, and thiosulphate can be prepared by

heating sulphur with sulphite solution. Many metals form soluble complexes with

thiosulphate, particularly silver and mercury. Thiosulphate solutions are used to

dissolve the light sensitive silver bromide in photographic emulsions to 'fix' the image.

The fact that mercury forms a complex ion means that the oxidation of thiosulphate

cannot be studied using polarography, as the the mercury surface is oxidised

preferentially [48]:

Hg + 2 S2032- —> Hg(S203)22“ + 2 e" (2.21)

Thiosulphate is also difficult to reduce, and no reduction waves are observed at a

dropping mercury electrode. At a platinum electrode, potentials o f-1.75 V vs SHE are


required before HS" is produced. It is this kinetic inertness that has led to thiosulphate

being named a m etastable sulphide oxidation product. Though not

thermodynamically stable, thiosulphate solutions can nevertheless be kept for weeks

without appreciable disproportionation or oxidation.

2.1.6 Sulphite (IV )

Although aqueous solutions of sulphur dioxide have been termed sulphurous acid, it

has now been established that the free acid does not exist, and aqueous solutions

contain SO2 (aq). Aqueous sulphur dioxide solutions can give rise to two series of

salts, the sulphites containing S032‘, and the bisulphites containing HSO3". The

bisulphite ion can react with itself to form the metabisulphite ion S2O52":

2 H S 0 3- <-» S20 52- + H20 (2.22)

S2O52" ions exist in dehydrated solid salts and in concentrated aqueous solutions.

Sulphite ions can be oxidised to sulphate, and in alkaline solution they will act as

reducing agents, for instance slowly removing dissolved oxygen. Samec and Weber

[60] studied the electrochemical oxidation at a gold electrode. They found that the

oxidation rate was much slower than that expected for a diffusion controlled process,

and was characteristic of an adsorption process followed by an irreversible two electron

transfer to yield sulphate.

2.1.7 Sulphate (VI)

Sulphuric acid, and its two series of salts, the sulphates and bisulphates, represent

sulphur in its highest normal oxidation state. They are thermodynamically stable in

aerated aqueous solutions at all pH s, and are the ultimate oxidation product of all other

sulphur salts. Theoretically, the sulphates may be reduced to sulphur (0) or sulphide

(—II), (see Fig. 2.2). In practice these reactions are highly irreversible; sulphates are

not normally reduced in aqueous media even in the presence of powerful reducing

agents. In fact, sulphates are so resistant to reduction that they are commonly used as

background electrolytes in electrochemical studies. However, sulphate can be reduced

to HS" by the bacteria v ib rio d e s u lp h u r ic a n s which can achieve this in cold aqueous

solution [61].

2.2. Electrochemical Investigations into Sulphide Oxidation

In an investigation of the oxidation and reduction of hydrosulphide ions on a rotating

gold electrode at pH 6.8 and 9.2, Hamilton and Woods [41] concluded that at low

potentials, first sub-monolayers, then multilayers of sulphur were produced. They

found the ratio of anodic charge to cathodic charge was greater than 1 , and furthermore

that this charge imbalance increased with increasing rotation rate. They concluded that

this must be due to a soluble intermediate which was dispersed at high rotation rates.

Polysulphides are known to be soluble, giving yellow-green solutions, and so they

postulated that the reaction proceeded through a poly sulphide intermediate. Sulphur (0)

is known to exist in polymeric form; rhombic sulphur consists of stacked Sg rings.


Therefore, polysulphide ions are reasonable intermediates to propose for the oxidation

of sulphide to sulphur. Indeed, Allen and Hickling suggested a similar mechanism in

1957 [62].

In a later paper, Buckley, Hamilton and Woods [42] showed that the initial sub

monolayer coverage, which formed in the underpotential region (~ -0.2 V vs. SHE at

pH 9.2), showed an X-ray photoelectron spectrum that was consistent with a gold

sulphide type structure. If this potential was maintained for extended periods (~ 10

mins.), then multilayers of sulphur were formed which passivated the electrode

surface. They confirmed that the oxidation and reduction of the sulphur proceeded via

soluble polysulphide species by detecting them using a rotating ring-disc electrode

(RRDE) [63].

The above authors [42] first plated the disc with sulphur in a positive going potential

scan, then reduced the adsorbed sulphur in a negative going scan; they found that

polysulphide species were produced:

nS + 2 e- -> Sn2- (2.23)

By holding the ring at a highly negative potential (-0.92 V vs. SHE) the polysulphide

ions were further reduced to HS~ according to equation (2.24):

Sn2- + 2 (n -l)e - + nH+ -> n HS‘ (2.24)

where n = 2,3,4 or 5.

Polysulphides can also be produced by the chemical dissolution of sulphur in

hydrosulphide solutions:

(n -l)S + HS- -> Sn2- + H+ (2.25)

By comparing the charges passed due to equations (2.23) and (2.24), and allowing for

the chemical production of polysulphides via chemical dissolution (2.25) they

calculated the mean chain length in the polysulphide intermediates, according to

equation (2.26):

n = 1 + (Qr -qr)/N Q d (2.26)where Qr = charge due to polysulphide reduction at the ring / C.

Qd = charge due to sulphur reduction at the disc / C.

N = the RRDE collection efficiency, a constant for a given geometry.

qr = charge due to the reduction of chemically produced polysulphide / C.

They found that at pH 9.2, n = 3.3, indicating that a mixture of different polysulphides

was produced.

At potentials above +0.25 V, the adsorbed sulphur layer can be oxidised to form

sulphate (2.27). Sulphate can also be formed directly through the oxidation of

hydrosulphide to sulphate, equation (2.28); this reaction can occur in parallel with

sulphur formation.

S + 8 OH- -» S 042- + 4 H20 + 6 e-

HS- + 9 OH- -> S 042- + 5 H20 + 8 e-

(2.27)

(2.28)


In a comprehensive review series, Zhdanov [48] reported that in alkaline solution

under mildly oxidising potentials, sulphide was oxidised to yield polysulphide ions,

which he ascribed to the dissolution of an initial deposit of sulphur in the sulphide

solution, equation (2.25). At much higher potentials, 1.0-1.7 V vs. SHE,

thiosulphate, and some sulphate, were formed in addition, and it was not until a

potential of above 1.7 V was applied that the oxidation product was predominantly

sulphate.

Moscardo-Leveist and Plichon [64] studied the electrochemical oxidation of sodium

sulphide in an equimolar sodium hydroxide/water melt at 100 °C and again found that

two oxidation steps were involved; the first step yielded elemental sulphur and di- and

tri-sulphides and the second, at higher anodic potentials, produced sulphite ions.

Remick and Camara [65] studied the electrochemistry of the sulphide/polysulphide

couple. They prepared their polysulphide solutions by dissolving elemental sulphur in

alkaline sulphide solutions according to equation (2.25). Polysulphide solutions

prepared in this way contain a number of polysulphide species; the average length of

the polysulphides being determined by the ratio of sulphur(O) to hydrosulphide(-II)

used. Giggenbach [53 ] studied such solutions by UY-visible spectroscopy and

determined the concentration of each polysulphide species present from their

absorbances. The analysis is complicated by the fact that the separate polysulphide ions

show absorbance maxima at very similar wavelengths. Nevertheless, he calculated the

absorbance maxima and the extinction coefficients for a range of polysulphides.

Remick and Camera [65] confirmed Allen and Hicklin’s earlier mechanism [62 ]

concerning the oxidation of polysulphide solutions. Their results were consistent with

the following scheme:

1. Adsorption of polysulphide ion onto the metal (M) surface:

Sn2- + M —> M~Sn2- (2.29)

2. Oxidation of the adsorbed poly sulphide by solution polysulphide:

M—Sn2- + Sn2- —> M—Sn.j + Sn.j2- + 2 e- (2.30)

3. The adsorbed layer of polysulphide can then be regenerated by reaction with

solution hydrosulphide ions:

M—Sn_! + HS- + OH- -> M~Sn2- + H20 (2.31)

In this way, a surface layer of polysulphide acts as an electrocatalyst for the oxidation

of polysulphide solutions. This mechanism explains the intermediate formation of

poly sulphides (e.g. S42-) observed by Hamilton and Woods [42]. The first step may

be oxidation of hydrosulphide to form tetrasulphide ions:

4HS- + 4 OH- -> S42- + 4 H 20 + 6 e- (2.32)

These can then adsorb onto the electrode surface:

M + S42- —> M—S42- (2.33)

and there undergo oxidation to produce adsorbed sulphur:

M—S42- + HS- + OH- —> M—S + S42- + H20 + 2 e- (2.34)


The polysulphide produced in equation (2.34) could re-adsorb and repeat the reaction

scheme. In this way a sulphur monolayer could be built up, with polysulphide ions

produced as intermediates close to the electrode surface. These could either adsorb

onto the electrode or, if the electrode is rotated, be dispersed into solution.

Remick and Camera [65] studied the electrocatalytic activity of several electrode

materials. They found that platinum was a poor electrocatalyst, especially for the

cathodic reduction of polysulphide to sulphide. They attributed this to the removal of

the adsorbed polysulphide layer on cathodic sweeps. As expected, conducting metal

sulphides were found to be better electrocatalysts for both the oxidation and reduction

reactions. Their work, and the study by Hodes and Joost [66], concluded that CoS,

NiS, and M0S2 were the most effective electrocatalysts. Platinum and carbon were less

effective.

2.3 Chemical Oxidation of Sulphide Solutions using Oxygen

Many workers have studied the air oxidation of sulphide solutions, and recently a

review of the topic was published by Kuhn, Kelsall and Chana [34]. When the

oxidation is achieved electrochemically, the applied potential governs the extent of

oxidation. In the same way, with sulphide solutions which are chemically oxidised, the

extent of oxidation is largely determined by the molar ratio of dissolved O2 to HS'.

Studzinska [67] reviewed the reaction and reported that a low ratio of O2 to HS-

favoured sulphur production, whereas a high ratio resulted in S2O32-, SO32-, and

SO42- formation.

The solution pH also has an important role to play; sulphur is thermodynamically stable

only in acidic or neutral solutions. Therefore, it would be expected that under mild

oxidation in this pH region, sulphur would be the predominant oxidation product, and

sulphoxy species would predominate at a higher pH. A substantial pH change can also

occur on oxidation, and it must be ensured that the solution is adequately buffered to

prevent this altering the reaction course. Alferova and Titova [68] carried out the

oxidation of sulphide solutions at various pH values by aerating them for 24 hours. At

pH 7, they found that most of the starting sulphide was converted into elemental

sulphur. As the pH was increased, the proportion of sulphide that was converted into

thiosulphate and sulphite rose, until at pH 15 conversion to thiosulphate was almost

complete. O'Brien and Birkner [22] listed the reaction products from the results of a

number of studies, including their own. In alkaline solutions, thiosulphate and sulphite

were the main products, although polysulphides, sulphur, and sulphate were also

reported.

It is well known that, in elemental sulphur, Sg rings can be broken by ultra-violet light.

Since polysulphide ions absorb in the near ultra-violet wavelength regions, they too are

likely to be decomposed by strong light sources, forming reactive free radicals.

Therefore, it is possible that ambient light conditions can affect the reaction rate of


aqueous sulphide oxidation and alter the reaction products. Cox and Sandalls [69]

showed that light of wavelength 300 to 400 nm caused photo-oxidation of gaseous

mixtures containing oxygen and traces of hydrogen sulphide. Sulphur dioxide was the

major product, and this was thought to be produced from the reaction between HfjS and

O* or OH* free radicals. Pelizetti [70] showed that hydrogen sulphide can be cleaved

in aqueous solution by visible light, producing hydrogen and elemental sulphur. He

added colloidal cadmium sulphide particles, which acted as photocatalysts: They

absorbed photons to generate electron-hole pairs, and the HS“ ions then acted as hole

scavengers, so becoming oxidised to form sulphur (or polysulphides).

Many workers studying the oxidation of aqueous sulphide solutions did not monitor the

product distribution, following the reaction instead by the decrease in the sulphide

concentration. In many industrial applications, e.g. waste water treatment, it is

relatively unimportant to determine this distribution (providing the products are not

toxic). The possibility of using as a measure of the extent of oxidation was ignored,

and no attempts were even made to measure it with a suitable indicator electrode.

2.3.1 Rate of Reaction of Sulphide Solutions with Oxygen

The reaction rate is usually defined as the rate of loss of sulphide ions, taking no

account of the products. Thus plots of 'rate' against pH can be misleading, as various

reactions are known to be predominate in the different pH regions.

It is clear that the uncatalysed oxidation of sulphide solutions using oxygen alone

proceeds slowly. At 25 °C, ti/2, the time taken for the sulphide concentration to reach

half its initial value, is several hours. Many workers noted that the reaction was

preceded by an induction period varying from 15 minutes to two hours [71-74]. Such

an induction period is characteristic of an autocatalytic reaction. Bowers [7 2 ]

suggested that the catalytic products were polysulphide ions, whilst Cline and Richards

[75] proposed that the catalysts were free radicals. They found that a high surface area

to volume ratio in their glass reaction vessels reduced the induction time.

Bhaskarwar and Kumar [76] studied the oxidation of hydrosulphide solutions using a

foam bed contactor operating at 75 °C. They suggested that the reaction proceeded

through a Sg2- intermediate which could either undergo mild oxidation and ring closure

to form elemental Sg, or further oxidation to form S2O32-. The foam was stabilized

using dodecyl sulphate or octyl phenoxy polyethoxyethanol surfactants.

Several workers [22,73,77] found an approximate first order dependence of the

reaction rate on sulphide concentration. O’Brien and Birkner [22] determined that the

reaction was first order with respect to the oxygen partial pressure, although Chen and

Morris [77] quote this order to be 0.56.


2.3.2 Effect of Temperature and pH on Reaction Rate

Selmeczi [78] reported that the reaction rate increased substantially with temperature,

and Bowers [72] showed an Arrhenius plot of reaction rate vs. temperature in the

range 20 °C to 50 °C. The change in rate over this range was approximately 20 fold.

The pH was reported by all authors to have a distinct effect, which is not surprising in

view of the varying reactions occurring in differing pH regions. Chen and Morris

reported two rate maxima at pH values of 8.5 and 11.5 [77]. Snavely and Blount

[74 ] found a substantial rate increase at pH 11.5, and Alferova and Titova [68]

reported maximum rates at extremes of pH.

2.3.3 Catalysis of Sulphide Oxidation

Snavely and Blount found that just 5 mg Co2+ dm-3 effected complete oxidation of

aerated 6 mol m-3 sulphide solution in only 60 seconds [74]. Interestingly, they also

noted that their own early results on uncatalysed systems were in error because a

chromium plated oxygen probe had catalysed the reaction. It may be that other studies

purporting to be on the uncatalysed system have also been affected by trace

contaminants (e.g.,in the chemicals used). This may explain the varying rates and

reaction products observed by different workers under apparently similar experimental

conditions.

Transition metals, noble metals, activated carbon, and organic compounds are all

known to catalyse the oxidation. Studzinska claimed that the most effective catalysts

were the transition metals [67] and she ranked their effectiveness in the order:

Ni2+ > Co2+ > Mn2+ > Cu2+ > Fe2+

Organic catalyts, such as phenols and hyroquinones, were ranked less effective still,

but still increased the rate of oxygen uptake 10 to 20 fold. Activated carbon is also

known to catalyse hydrogen sulphide oxidation both in the gas phase [79], and in

aqueous solution [73]. Oeste [80] proposed that the hydrogen sulphide oxidation

over activated carbon took place via an electrochemical mechanism; Kuhn and Kelsall

made a similar suggestion about the catalytic effect of the transition metals [34].

Metal sulphides have extremely low solubility products. For example, for copper(II)

sulphide:

CuS + H20 = Cu2+ + OH' + HS- (2.35)

K sp = [Cu2+] [OH-] [HS-] = 6 x 10-37This means that heterogeneous metal sulphide particles are usually produced when

solutions containing transition metal catalysts are added to aqueous sulphide solutions.

Surprisingly, workers have consistently ignored the presence of these colloidal

particles, which is unusual in view of their possible role as redox catalysts.


Since many of these metal sulphides are known to be electrically conducting and to

catalyse oxygen reduction [81], Kuhn and Kelsall proposed the following mechanism

for transition metal catalysis of sulphide oxidation [34]. They suggested that the

overall equation (2.38) was split into two half reactions, equations (2.36) and (2.37),

each of which could occur on the metal sulphide surface:

A similar mechanism has been proposed to explain the catalytic activity of colloidal gold

particles [82]. Since CoS, NiS, and M0S2 are now known to be electro-catalytically

active for sulphide oxidation [65], this mechanism explains their effectiveness as

catalyts. However, an eight electron transfer is unlikely to occur in one step, and the

reaction is likely to proceed via a series of intermediates, such as soluble polysulphides.

Electrochemical hydrosulphide oxidation at the mineral sulphide surfaces, galena (PbS)

and pyrite (FeS2), has shown that polysulphide ions can be formed [56]. These metal

sulphides were found to be more effective electrocatalysts than the noble metals,

platinum and gold.

Vanadium does not precipitate a solid sulphide phase when vanadium (V) salts are

added to hydrosulphide solutions. Vanadium (V) is a moderately powerful oxidising

agent and is capable of directly oxidising the hydrosulphide to produce sulphur [83].

This can produce solid vanadium oxide phases, which are noted for their catalytic

abilities during gas phase oxidations at high temperatures; for instance in the well

known contact process which oxidises SO2 to SO3 [84]. Recently 51V NMR

studies [37] have shown that vanadium (V) can also form oxy-sulphur complexes in

solution, which may assist the catalytic action of vanadium(V).

2.3.4 Bacterial Action in Sulphide Oxidation

Experience in working gas desulphurisation plants shows that bacterial oxidation can

change the product distribution of the sulphide oxidation [29]. However, a culture time

of several days is required before the bacterial population can significantly change the

oxidation pathway. Therefore, it is unlikely that experiments on the 'chemical'

oxidation of sulphide solutions have been affected by bacterial action, despite the fact

that no special precautions appear to have been taken to exclude bacteria.

2.4 The Production of Elemental Sulphur

It is clear from the above discussion that elemental sulphur is not a thermodynamically

stable oxidation product, except under conditions of mild oxidation in acidic solutions.

Although sulphate (oxidation state +VI) is the preferred product, it is observed only

after prolonged oxidation, and more commonly metastable sulphur species are formed;

these include sulphur (0), thiosulphate (II), and sulphite (IV). Concentrations of

intermediates such as polysulphides and polythionates can also accumulate. Many

workers have observed that the available reaction pathways are followed in parallel,

HS- + 9 OH- -> S 0 42- + 5 H 20 + 8 e-

4 H20 + 2 0 2 + 8e- -» 8 OH'

HS- + 2 0 2 + OH" —» S 042- + H20

(2.36)

(2.37)

(2.38)


resulting in the simultaneous production of thiosulphate and sulphur from

hydrosulphide solutions for instance.

Sulphur is the desired product from many industrial oxidation processes because it is

non-toxic, easily handled and a saleable by-product. Complete conversion to sulphur

can be anticipated only under oxidation conditions where it is thermodynamically stable;

i.e.,low solution potentials of around 0.0 V vs SHE and a pH of around 5. These

conditions can be provided if the oxidation is carried out electrochemically, and

industrial processes based on this principle have been proposed [15,16].

Operating at an acidic pH retards the absorption of hydrogen sulphide into the aqueous

phase. Industrial processes operating at this pH must compensate by ensuring the rapid

oxidation of the H2S once it is in solution. Although there is an industrial process

which uses an acidic solution containing an Fe (III) complex [21], most processes

have utilised alkaline working solutions. Vanadium (V), iron (III), and arsenic (V)

have all been used industrially as the oxidising agents [29,85,10]. These processes

all produce a range of higher oxidation state sulphur products.

Maximum production of elemental sulphur can be achieved by removing the sulphur

from the reaction system as soon as possible. In the Stretford Process the aeration

which occurs in the oxidiser also serves to remove the sulphur by froth flotation [35].

It has also been suggested that de-oxygenating the absorbing solution prior to it

contacting the gas stream containing the hydrogen sulphide can decrease the rate of

production of thiosulphate [86].

Sulphide Electrochemistry 4 2

3. Sulphide ElectrochemistryElectrochemical studies of hydrosulphide ions in aqueous solution are made difficult

because most metals form their sulphides when oxidising potentials are applied. Of the

noble metals, gold is known to form a sulphide coating less readily than platinum [87],

and so this metal was chosen for the working electrode material. Gold can dissolve in

sulphide solutions to form the gold (I) complex; AuS". Garrels and Christ [88]

produced an E^-pH diagram for the Au/S/Cl system (Fig. 3.1), which shows this

complex to be thermodynamically stable in alkaline solution:

Fig. 3.1 Eh-pH Diagram for the Au/Cl/S System.

Dissolved concentrations / kmol n r 3: Au as marked; Cl 1; S 0.1.

However, in practice, corrosion of gold in sulphide solutions is not a serious problem;

in the present study it was found that several hours continual potential cycling from

-0.9 V to +0.3 V vs. SHE. (at 20 mV s-1 ), in a solution containing 1 kmol Na2S m"3 at pH 14, was required to strip a gold layer only 1 Jim thick.

Recent evidence from Buckley, Hamilton and Woods suggested that gold electrodes

that were immersed in sulphide solutions (at pH 9.2) and held at potentials higher than

-0 .5 V vs. SHE became coated with a gold sulphide phase [42]. However, such a

phase is likely to be semiconducting and has been shown not to passivate the electrode

towards further sulphur deposition.


3.1 Thermodynamic Calculations

Fig. 3.2 shows an Eh-pH diagram for the sulphur/water system at 298 K. It was

created using the program PPE produced by Angus [89,90] running on the Apple He

microcomputer. The thermodynamic data, in the form of the free energies of formation

of the species considered, were taken from a review by Zhdanov [31]. Notice that

there are only three stable oxidation states. Oxidation of hydrosulphide at pH 9.0 from

Eft = -0.4 V to Ejj = +0.1 V would predict that sulphate would be formed as the

predominant product.

Fig. 3.2 Eh-pH diagram for the sulphur/water system at 298 K.

[S species] = 10 mol n r 3.

Zhdanov [31] quotes a value of 86.31 kJ mol-1 for the AGf° of S2-, which implies that

the value for the p K ^ of H2S is 13. If Zhdanov had accepted the new, higher value for

the pKa2 of around 19 (see section 2.1.1) then the value for AGf° (H2S) would be

119 kJ mol-1 (assuming that AGf° (HS-) = 12.05 kJ mol-1). However, utilising this

higher value does not change the above diagram significantly, merely eliminating the

area of predominance for S2' ions.

A m etastable E^-pH diagram can be produced by eliminating all sulphur (VI)

compounds from the calculations (see section 2). Such a diagram is shown in

Fig. 3.3. and this indicates that the oxidation of a sulphide solution at pH 9.3 can

proceed to form thiosulphate (at E^ > -0.3 V vs. SHE), sulphite (E^ > -0.175 V) or

dithionate (E^ > +0.12 V). It is also worth noting that, although the polysulphide

species do not appear on the diagram, the area of predominance of the ion S52" is

masked only by that of elemental sulphur. Elemental sulphur can only form when

suitable nuclei are available, and often a high degree of supersaturation is required


before such nuclei are formed. Thus, in mildly alkaline solution, yellow coloured

polysulphide solutions can result from the atmospheric oxidation of hydrosulphide

ions, despite the fact that they are not thermodynamically stable.

Fig. 3.3 Metastable Eh-pH diagram for the Sulphur-water system

298 K.S(VI) species excluded. [S species] = 10 mol m“3.

3.2 Experimental

Cyclic voltammetry and potential pulse studies were conducted at room temperature

(~ 20 °C) using hydrosulphide solutions at pH 9.3 with a rotating ring- disc electrode

(RRDE). A schematic diagram of this electrode is shown in Fig. 3.4.

crztr̂

Fig. 3.4 A Rotating Ring Disc Electrode

Sulphide Electrochemistry 45

The ring and disc potentials can be controlled independently using a bipotentiostat, and

the potential of the ring can be adjusted, such that any metastable oxidation products

passing to the ring become reduced. As the electrode is rotated, these soluble species

produced at the disc are spun out to the ring, and the reduction current at the ring can be

used to obtain a measure of their concentration in solution. The collection efficiency,

N, is defined as that proportion of soluble product, produced at the disc, which is

transported to the ring. The collection efficiency is dependent on the electrode

geometry and can be predicted theoretically [63 ], or experimentally checked by

monitoring the transport of a material which is known to be oxidised and reduced

reversibly (such as Fe(CN)64").

3.2.1 Solution Preparation

A carbonate buffer solution of pH 9.3, containing 0.059 kmol Na2C03 m"3, 0.223

kmol NaHCC>3 m"3 and 0.10 kmol Na2S04 m-3, was prepared by dissolving the

appropriate mass of analytical grade materials (BD H ) in triply distilled water.

Similarly a borate buffer, having a pH of 9.2, was made up containing 12.5 mol

Na2B4O7.10H2O m-3, 0.9 mol NaOH m-3 and 0.1 kmol Na2S04 m-3. A stock

solution containing 0.1 kmol HS" m"3 was prepared by dissolving an accurately

weighed amount (about 12 g) of transparent, dried crystals of Analar sodium sulphide

(BDH) in 500 cm3 the appropriate deoxygenated buffer solution.

The molarity of this stock solution was checked by conducting an iodate titration:

exactly 1 cm3 of the stock solution was taken and mixed with a 15 cm3 aliquot of

potassium iodate solution (0.025 kmol m“3). The mixture was made highly alkaline by

adding 10 cm3 of sodium hydroxide (10 kmol m“3) and boiled for 10 minutes. Under

these conditions the sulphide was oxidised into sulphate:

4 I 0 3- + 3HS- + 3 0H - 41- + 3 S 0 42- + 3 H20 (3.1)

The unused iodate was then back titrated. Excess potassium iodide solution (5 cm3 of

5 % KI by mass) was added to the cooled solution which had been made acidic by

adding 20 cm3 of H2SO4 (4 kmol m"3). This converted the unused iodate to iodine:

I0 3- + 51- + 6 H+ -» 3 I 2 + 3 H 2 0 (3.2)

The liberated iodine was then titrated with thiosulphate (0.1 kmol m-3):

6 S20 32" + 3 I 2 -> 6 S40 62- + 61- (3.3)

As the end point neared, the solution became a pale yellow colour and the solution was

diluted to 150 cm3 with distilled water. Several drops of sodium starch glycollate

(BDH) were added and the end point was detected when the characteristic deep blue

colour of the starch-iodine complex was discharged. The concentration of the original

hydrosulphide solution was calculated from equation (3.4):

C = 7.5 x 105 (3.75 x 10"4 - 16.666 Vt) (3.4)

where C = Concentration of stock hydrosulphide solution / mol m“3 Vt = Volume of thiosulphate solution taken / m3


Freshly opened sodium sulphide was found to contain about 32 % Na2S, which

corresponds to the formula Na2S.(H20)9 23.

Stock sulphide solutions could be kept for several weeks without degradation in a

septum-stoppered bottle with a nitrogen atmosphere over the liquid. Measured volumes

were withdrawn using a glass syringe and needle, and injected into a larger volume of

the nitrogenated buffer solution, to make solutions containing 10 mol HS" m '3.

Electrochemical studies were carried out in a three compartment glass cell of

conventional design, (see Chapter 7, Fig. 7.1).

Sodium tetrasulphide (Na2S4) was prepared according to the method given by

Schwarzenbach and Fischer [50]. Under an inert atmosphere 12.588 g of sodium was

dissolved in 400 cm3 Qf absolute ethanol, forming sodium ethoxide:

Na + C2H5OH C2H5ONa + 1/2 H2 t (3.5)

Dry hydrogen sulphide was then bubbled through the solution until it became saturated,

whereupon sodium hydrosulphide was formed:

C2H5ONa + H2S -> NaHS + C2H5OH (3.6)

Excess hydrogen sulphide was absorbed in Dressel bottles containing solutions of

NaOH (10 kmol m '3) and CUSO4 (1 kmol m-3), to prevent its escape into the

atmosphere.

To this, 26.386 g of elemental sulphur was added; this dissolved with the evolution of

hydrogen sulphide, producing a deep red solution of sodium tetrasulphide:

2 NaHS + 3S -> Na2S4 + H2S T (3.7)

To ensure that reaction (3.7) proceeded to completion the solution was refluxed under

an inert atmosphere for one hour. It was then cooled to below 40 °C and the solvent

evaporated under vacuum so that the solution was reduced to 1/10 of its original

volume. A yellow crystalline product was obtained, which was filtered under vacuum

in an inert atmosphere and dried for one week over P2O5. A yield of 33.64 g was

obtained (71 %).

Polysulphide solutions were prepared by either dissolving the appropriate mass of

Na2S4 in an oxygen-free buffer solution, or by dissolving elemental sulphur in

sulphide solution. A stock solution, of average polysulphurisation index of two, was

prepared by adding elemental sulphur to a solution of Na2S.9H20 in the molar ratio

1:1. The total sulphur concentration was 0.1 kmol n r 3 and the sulphur took several

days to dissolve, forming a transparent, bright yellow solution. Such a solution will

contain not only S22“, but also HS', S32", S42" and S52".

3.2.2 Electrochemical Instrum entation

Gold and platinum ring-disc electrodes of were mounted in a motor unit (Oxford

Electrodes) which allowed the rotation speed to be continuously varied up to 50 Hz.

The disc areas were 0.3848 cm2 and the electrode dimensions were rj = 0.35,


T2 = 0.375, 13 = 0.4 cm. This geometry provided a theoretical current collection

efficiency of 0.17, which was verified experimentally using a solution containing

1 mol K4Fe(CN)6 m-3. The hexacyano iron (II) was oxidised at the disc and the

product reduced at the ring, the ratio of the two currents was 1: 0.17.

The potentials of the ring and the disc, relative to a saturated calomel reference electrode

(EIL), were controlled independently using a bipotentiostat built at Imperial College,

based on conventional operational amplifier design. The control potentials were

provided by two Hi-Tek PPR1 waveform generators, and bright platinum counter

electrodes were used. The ring and disc currents were passed through resistors and the

resulting voltages applied to the inputs of separate J J PL4 chart recorders. The

thiosulphate cyclic voltammetry was carried out using one channel of the bipotentiostat,

and the polysulphide voltammetry was conducted using the Solartron 1286

Electrochemical Interface.

3.2.3 Electrode Pretreatm ent

Initial experiments were carried out using platinum electrodes that had been coated with

gold. This coating was achieved by polishing the platinum surface until a mirror finish

was obtained, then electroplating the gold from aqueous solution under potentiostatic

control, whilst the electrode was rotated at 40 Hz. The electroplating solution contained

1 mol AUCI4'm -3 and 0.1 kmol HC1 m-3. During the electroplating, the potential was

maintained at 0.287 V vs. SHE (0.045 V vs. SCE) which caused a current of

approximately 1.3 mA to flow. After 1 minute the electrode was disconnected and the

gold surface gently polished with a lint-free tissue to prevent the deposit from becoming

dendritic. The electroplating was then continued until the desired thickness of gold had

been achieved. Gold coatings prepared in this way were bright and adherent, and

voltammograms recorded on pure gold and gold-plated platinum surfaces were

identical.

Gold electrodes that have been exposed to the atmosphere adsorb oxygen on their

surfaces, which gives rise to a reduction current on the first negative-going potential

scan. Standing the electrode in nitrogenated buffer, or potential cycling in the same

media desorbs this oxygen. During rotated-disc experiments, care was taken to

maintain a nitrogen atmosphere above the solution surface; the disc rotation had the

effect of aerating the solution which could cause large oxygen reduction currents to

flow (280 |iA at -0.5 V vs. SHE).

After polishing with 0.3 qm alumina powder until a mirror finish was obtained, the

electrode was introduced into the working solution. Cathodic polarization and potential

cycling were investigated as possible methods of electrode activation. Holding the

electrode at a highly negative potential (-1.7 V vs SHE) removed the adsorbed oxygen,

but subsequent voltammograms recorded in sulphide solutions showed current

densities lower than those which were obtained after the electrode had undergone


potential cycling, suggesting the presence of adsorbed sulphur. It was found that

potential cycling at 10 V s-1 between -1.25 V and +1.75 V vs. SHE produced an active

gold surface. If the anodic limit of the potential scans was reduced to -0.2 V vs.

SHE, the electrode surface was not activated; adsorbed sulphur on the gold electrode is

not oxidised to sulphate until a potential of 0.5 V vs. SHE is exceeded [41]. After

potential cycling, the electrode was held at the cathodic limit, prior to commencement of

a potential scan or pulse.

Current densities were calculated from the geometrical surface area unless otherwise

stated. The real surface area of the gold disc electrode was determined according to the

method of Dickertmann et al. [91], which relies on the integration of the charge

passed when a layer of gold oxide is formed. The gold electrode was placed in a

solution containing 1 kmol HCIO4 m"3, and the potential scanned from 0.5 V to 1.7 V

vs. SHE at 10 mV s"1; polycrystalline gold with a roughness factor of one forms a

monolayer oxide coating with the passage of 0.40 mC cm-2. Using this method the

roughness factor of the polished gold electrode was found to be 1.3.

3.2.4 Experimental: Ion chromatography

The experimental apparatus was assembled as shown in Fig. 3.5 (overleaf). A glass

reservoir held the eluent, 0.1 kmol Na2C03 above m-3, which was constantly sparged

with nitrogen. The eluent was pumped through a Kontron 414T pump, through a

pressure damper and injection port and into a Dionex AG3 guard column. This

guard column prevented strongly adsorbing ions from poisoning the main ion exchange

column. The main column was made from the same material and achieved the

separation of the sulphur anions. Two detectors were provided: a LKB 2238

Unicord SII fixed wavelength UV detector (at 254 nm) and a Dionex ECD

electrochemical detector. The latter consisted of a silver working electrode (held at

-0.1 V vs Ag/AgCl), a gold counter electrode and an Ag/AgCl reference electrode.

Any sulphide or polysulphide passing to this detector produced an oxidation current

and formed Ag2S at the working electrode.

Complete exclusion of air from the working solutions was found to be essential to

prevent oxidation of the polysulphide species. Aqueous polysulphide solutions were

prepared by diluting stock solutions (see section 3.2.1) with deoxygenated purified

water to form solutions in the concentration range 0.1-1 mol m-3. Samples were

withdrawn into a glass syringe, and analysed immediately.


The concentrations of poly sulphide ions in the injected samples were calculated from

the rate constantsAthe equilibration between species given by Giggenbach [53] (see

section 2.1.2). From the initial concentrations of S(0) and S(-II), and knowing the pH,

the equilibrium concentrations of the polysulphide species were calculated. The

program TKSOLVER,run on a Digital 350 microcomputer, was used to solve

numerically the resulting set of simultaneous equations. Corrections were made for the

effect of ionic strength by calculating the activity coefficients according to the method of

Albert and Serjeant [92].


3.3 Sulphide Voltammetry: Results and Discussion

A voltammogram of a gold plated platinum electrode sulphide solution recorded at pH

9.2 is shown in Fig. 3.6. The main oxidation peak at +0.09 V vs. SHE is due to the

oxidation of HS" ions producing layers of elemental sulphur. This non-conducting

layer inhibits further oxidation and passivates the electrode. On the reverse potential

scan the corresponding reduction was not observed until a potential of -0.43 V vs. SHE

was reached. This large peak separation is evidence that sulphur formation is a highly

irreversible process, as would be expected for a phase formation reaction.

Fig. 3.6 Voltammogram of HS" on Gold Plated Disc Electrode.

[HS‘] = 10 mol m-3, pH = 9.2 (Borate Buffer), nth. cycle, 20 mV s_1.

The integrated charge under the oxidation peak was found to be approximately

20 C m-2 (based on the real surface area). A monolayer coverage of sulphur,

assuming 2 e" discharge, has been calculated to correspond to charge densities of 3.5 and 2.3 C m r2 [41 ,93 ] (depending on the assumptions made about the sulphur

packing). It is therefore apparent that several monolayers of sulphur were formed, and

this conclusion is in agreement with those of previous workers [41,42,94].


There was a charge imbalance over a complete cycle, more charge being passed on the

positive-going scan. There are several possible explanations for this imbalance:

1. The sulphur layer was not completely reduced. This would imply that the sulphur

layer would build up after repeated cycles. However, prolonged potential cycling

did not completely passivate the electrode, nor could any visible sulphur deposits be

seen.

2. A reaction that produced soluble sulphur oxidation products had occurred in parallel

with that of sulphur formation. Possible alternative oxidation products include

polysulphides, thiosulphate, sulphite and sulphate.

3. The sulphur layer was reduced to form polysulphide rather than hydrosulphide

ions. Polysulphide ions could diffuse into the solution before they were further

reduced.

Since the sulphur layer does not build up to completely passivate the electrode after

repeated cycles, the reason for the charge imbalance must be either (or both) of the

second and third possibilities.

Hamilton and Woods [41] suggested that sulphate production in the positive-going

scan, and polysulphide production in both the positive and negative-going scans, were

the reasons for the charge imbalance. To investigate the possibility that polysulphide

intermediates were formed, ring-disc electrochemical studies were conducted. The gold

ring was held at a reducing potential, in order to detect any polysulphide ions, and the

disc was subjected to either a triangular potential waveform or a potential pulse.

3.4 Thiosulphate Voltammetry: Results and Discussion

Thiosulphate is known to be a metastable oxidation product from sulphide oxidation.

Theoretically, it can be oxidised to tetrathionate (at potentials above 0.18 V vs. SHE),

or reduced to form hydrosulphide ions (below a potential of -0.3 V vs. SHE). The

actual redox behaviour at a gold electrode was investigated to determine whether

thiosulphate can, in reality, be reduced in the potential range that is required to ensure

polysulphide reduction in a ring-disc experiment.

A gold disc electrode was cycled between the potential limits -0.75 V and 0.3 V vs.

SHE in a solution containing 10 mol Na2S2C>3 m-3 at pH 8.2; no reduction currents

above those obtained with the buffer solution alone were observed.


If the potential range was increased, voltammograms such as those shown in

Fig. 3.7 were observed:

Potential vs. SHE / V

Fig. 3.7 Cyclic Voltammograms of Sodium Thiosulphate.

[Na2S203] = 10 mol m-3. 1st Scans. 100 mV s-1. pH = 8.2.

In the negative-going scan, it can be seen that no reduction currents were observed until

the cathodic limit was reached, when hydrogen was evolved at a potential of -0.7 V vs.

SHE. In the positive-going scan, an oxidation peak at 0.6 V vs. SHE can be seen

which was due to the formation of gold oxide. At higher potentials, 1.1 V and 1.25 V

vs SHE, further oxidation peaks can be seen. However, the magnitudesof these current

peaks are approximately an order of magnitude lower than the diffusion limited current

calculated from the Levich equation (7.14), even assuming only a one electron

oxidation:

^ 2^ 3^" 1/2 S ^ g ^ - + e~ (3.8)

This reaction occurred only after an overpotential of almost 1 V was applied. In fact it

has been suggested that thiosulphate oxidation proceeds via a chemical reaction with

hydrogen peroxide, which is evolved at an anode at these potentials [95].

It can be concluded that thiosulphate is electrochemically inactive at a gold electrode in

the potential range of interest to the present study.


3.5 Polysulphide Voltammetry: Results and Discussion

Polysulphide solutions were prepared by diluting measured volumes of stock solutions

(see section 3.2.1) in nitrogenated buffer. Potential scans were commenced from the

cathodic limit, or the electrode rest potential (-0.17 V vs. SHE); similar results were

obtained in both cases. Typical results are shown in Fig 3.8:

Fig. 3.8 Voltammograms of Polysulphide Solution at a Gold Disc.

[Sx] = 1 mol n r 3. xav = 2. pH = 8.2. Scan rate 50 mV s-1.

Aqueous polysulphide solutions always contain a proportion of free HS“ ions, and so

voltammograms of polysulphide and hydrosulphide solutions are very similar.

Commencing at -0.8 V vs. SHE, the first positive-going scan showed a oxidation pre

wave at around -0.5 V vs SHE. The integrated charge density under this peak (based

on the real surface area) was 0.5 C m-2, which corresponds to the discharge of a sub-

monolayer (~ 0.2 monolayers) of sulphide ions. Similar peaks have been observed in

the voltammetry of dilute hydrosulphide solution by Hamilton and Woods [41,42],

and were attributed to the formation of a gold sulphide phase at the electrode surface.

This oxidation peak disappeared after prolonged potential cycling provided the positive

limit was kept below 0.5 V, under these conditions an adsorbed sulphur layer is likely

to be permanently present. If the [HS“] were to be increased, this would have the effect

of decreasing the potential at which the phase forms; this explains why this oxidation

peak only appeared only as a shoulder on the hydrogen evolution current in

concentrated HS‘ solutions.

The main oxidation peak, which is due to the formation of multilayers of elemental

sulphur, was observed at 0.05 V vs SHE. When the electrode was rotated, this had


little effect on the peak current density, which was much lower than the expected

diffusion limited value (ip = 1.04 A m"2, iijm = 6.5 A m-2); moreover the current

decreased as the potential was increased above 0.1 V vs. SHE. All this is consistent

with sulphur passivation of the electrode surface. However, the fact that there was

some increase in the oxidation peak upon rotation indicated that soluble oxidation

products were also produced.

In the negative-going scan, at a stationary electrode, the reduction currents at -0.5 V vs.

SHE were seen to consist of two waves. These are due to the reduction of multilayer

sulphur (at -0.4 V vs. SHE) and the gold sulphide layer (at ~ -0.5 V vs. SHE).

An Eft-pH diagram showing the sulphur polysulphide system is shown in Fig. 3.9.

Fig. 3.9 E^-pH Diagram of the Sulphide/Polysulphide System.

[S Species] = 10 mol n r 3.

The diagram was produced using the program PPE [89,90], running on an Apple He

microcomputer, and the thermodynamic data was taken from Zhdanov's review [31].

The Eft-pH diagram was not adjusted to account for the new value of AGf°(S2-), since

the free energies of formation of the polysulphides were calculated from equilibrium

potential measurements on sulphide/polysulphide system, and rely on the lower value

for pKa2(H2S) [52]. Despite the uncertainties in the thermodynamic data, it is

apparent that at potentials lower than -0.45 V vs SHE, all poly sulphide species are

thermodynamically unstable and can be reduced to form S (—11).


In Fig. 3.8 it can be seen that the reduction of the polysulphide solutions at a rotated

electrode resulted in an increased reduction current below -0.4 V. However, no clear,

diffusion-limited plateau was seen in either the positive- or negative-going potential

scans. The diffusion-limited current density calculated from the Levich equation

(7.14), assuming a diffusion coefficient of 5.2 x 10' 10 m2 s-1 [42] and a rotation

speed of 20 Hz, is around 4.2 A m-2 for a one electron transfer. The current density at

-0 .6 V vs. SHE was only 1.8 A m-2 . If the lower potential limit was decreased, scans

such as Fig. 3.10 were obtained:

Fig. 3.10 Voltammograms of Polysulphide Solution at a Gold Disc.

[Sx2-] = 1 mol n r 3. xav = 2. pH = 8.2. Scan rate 50 mV s-1.

There was a current plateau at a potential of -0.95 V vs. SHE; it is conceivable that this

current might be due to the diffusion-limited reduction of poly sulphide ions:

Sn2- + 2 (n -l)e “ + nH+ nHS- (3.9)

However, this current density is too high to be attributed solely to this reaction; the

calculated diffusion-limited value is 8.4 A m-2 for the complete reduction of all

polysulphide species, whereas the observed value was 65 A m- 2 . Thus, the major

proportion of the observed reduction current at -0.95 V vs. SHE must be due to

hydrogen evolution (E0' H+/H2 = -0.48 V vs. SHE at pH 8.2). Rotating the electrode

increased the hydrogen evolution current, and adsorbed sulphur is known to have a

dramatic effect on the hydrogen overpotential [96].


Therefore, there is a problem in deciding which potential should be applied to detect

polysulphide species. Oxidising potentials cannot be used, since elemental sulphur is

formed and the electrode passivates. At a potential of -0.6 V vs. SHE, the

polysulphide species may not be completely reduced. If the potential were to be

lowered to around -0.95 V, a substantial hydrogen evolution current would flow;

which may be modified by the local [HS~] (and [H+] ). Buckley et al. [42] showed a

clearly defined current plateau at -0.6 V vs. SHE (at pH 9.2), and claimed that the

magnitude of this reduction current was consistent with the complete reduction of all the

polysulphide species. Nevertheless, they chose a substantially more negative potential

in order to detect polysulphides at a gold ring: -0.92 V vs. SHE (unless they have

erroneously quoted the potential vs. SCE rather than vs. SHE). At this potential a

substantial hydrogen evolution current is likely to flow, and their assumption that this

background current is constant (irrespective of surface sulphide concentration and pH

changes) is questionable. In the present ring-disc study, a detection potential of

-0.75 V vs. SHE was applied.

3.6 Ring-Disc Studies: Results and Discussion

To determine whether polysulphide ions were produced from the oxidation of HS"

ions, and from the reduction of elemental sulphur, ring-disc electrode studies were

undertaken. An experimental problem was that the electrochemical activity of the ring

(even when held at -0.75 V vs. SHE) decayed with time. The experiment was

conducted after the initial decline in activity had stabilised, and periodically the electrode

could be reactivated by potential cycling. Buckley et al.[42] noted a similar problem in

their studies, and reactivated the electrode by pulsing to a highly positive potential.

This deactivation suggests that sulphur was still adsorbing onto the electrode surface,

even at these low potentials, implying that that the polysulphides may not be reduced

under mass transport control at the ring. The potential of the ring was maintained at

-0.75 V vs. SHE, while the disc potential was swept from -0.75 V to 0.3 V vs. SHE,

returning to -0.75 V. The resulting ring and disc currents are shown in Fig. 3.11

(overleaf).

At the disc, similar results to those described previously were observed; a small

oxidation peak at -0.53 V vs. SHE was seen in the positive-going scan, as the gold

sulphide phase was formed, and the major oxidation peak was seen at 0.065 V vs.

SHE. The integrated charge under this peak was 358 qC, which corresponds to 2-3

monolayers of sulphur.


Fig. 3.11 Ring-Disc Voltammetry of Sulphide Solution at Au RRDE.

[HS-] = 10 mol m"3. co = 9 Hz. Scan rate = 100 mV s-1.

Ring potential = -0.75 V vs. SHE.

At the ring, there was little response in the positive-going (disc potential) scan. A

slightly increased reduction current was seen, which reached to a maximum value of

—2 jiA. Assuming that this was due to polysulphide reduction, and that the

polysulphide ions were produced at the disc, the corresponding disc oxidation current

would be 11.8 |iA. This represents only about 5 % of the peak oxidation current at the

disc, and supports the theory that polysulphides are rapidly oxidised at a gold electrode

to form elemental sulphur [42]. Thus, only a small proportion of the poly sulphides

that are produced can diffuse into solution and reach the ring.

On the negative-going sweep, a reduction current began to flow at the disc when the

potential reached -0.3 V vs. SHE and this increased in magnitude as the potential was

decreased further. The integrated charge that was passed in the negative-going scan,

after subtraction of the hydrogen evolution charge, was 200 jiC. This includes a


component which is due to the reduction of the gold sulphide surface phase; Buckley et

al. [4 2 ] quoted a value of 0.9 C m"2 for the reduction of this layer , which

corresponds to 45 pC on the experimental surface area. Therefore the charge that was

passed in reducing the multilayer sulphur, Q^, was 155 JJ.C (which was 43 % of the

charge that was passed to form the sulphur).

At the ring, during the negative-going (disc potential) scan, a reduction current was

seen which reached a maximum when the disc potential was -0.535 V vs. SHE. This

current arose from the reduction of polysulphide ions which were swept out from the

disc; if the electrode was stationary, no reduction currents were observed. There are

two possible ways in which these polysulphide ions can be produced: from the

electrochemical reduction of adsorbed sulphur (3.10), or from the chemical dissolution

of sulphur in HS" solution (3.11).

n S + 2 e" —» Sn2' (3.10)

nS + nHS" + nOH" —» Sn2" + n H 2 0 (3.11)

At a disc potential of -0.15 V vs. SHE, no current flowed at the disc, yet a reduction

current of about 2 pA below background was seen. This must have been due to the

production of polysulphide ions by chemical dissolution of the sulphur layers (3.11).

An estimate of the current due to polysulphide production due to the electrochemical

reduction of sulphur can be gained by subtracting this 2 qA from the ring reduction

currents observed at a lower disc potential. The integrated charge due to the reduction

of electrochemically produced poly sulphide, Qr, was found to be 21.2 |iC.

If it is assumed that:

1. At the disc, all the sulphur was reduced to form polysulphide ions according to

equation (3.10)

2. At the ring, all the polysulphide reaching this electrode was fully reduced to HS":

Sn2~ + 2(n-l)e" + nH+ -> n HS" (3.12)

Then the ratio of the two charges will be given by:

N Qd/ Qr = 1 / (n-1) where N = the ring collection efficiency (3.13)

From equation (3.13) the average poly sulphide chain length (n) can be calculated.

Substituting in the values; = 155 |iC, Qr = 21.2 jiC and N = 0.17 gives n = 1.8.

In a similar analysis, Buckley et al. [42] calculated that the polysulphide had n = 3.3

(at pH 9.2 and [HS‘] = 0.2 mol m~3). It can be concluded that in both cases the

reduction product is likely to contain a mixture of polysulphide species. This is not

unexpected, bearing in mind the predominance of different polysulphide species as the

potential range was scanned (see Fig. 3.9).


The reduction of the sulphur layers was found to correspond to 155 p.C, and if this

resulted in the formation of a poly sulphide of average stoichiometry 82-, the sulphur

must have been deposited with the passage of 1.8 x 155 |iC, i.e. 279 fiC. As the

sulphur was deposited, it was also chemically dissolved to form more polysulphide

ions. Over the time span of the deposition of the sulphur (about 8 s) a reduction current

at the ring of 2 (lA was observed due to the reduction of these polysulphide ions.

Thus, at the ring a further charge of 16 |iC was passed. Assuming a collection

efficiency of 0.17, the amount of sulphur dissolved would have required the passage of

16 / 0.17 = 94 {iC for its production at the disc. Thus, in total, the 279 + 94 = 373 JJ.C

would be expected to have been passed on the anodic scan, which is is approximate

agreement with the observed anodic charge of 358 (iC. Therefore, the charge

imbalance at the disc can be attributed to the production of polysulphides in both the

positive- and negative-going scans; there must have been little or no direct oxidation to

produce sulphoxy species.

One criticism of the above approach, is that it is only legitimate to use ring and disc

charges rather than currents when the collection efficiency and the composition of the

intermediate species remain constant as the disc potential that is scanned. A constant

collection efficiency requires that the reduction of polysulphide ions at the ring always

operates under mass transport control, irrespective of the polysulphide flux over the

ring. The nature of the polysulphide ions that are produced from sulphur reduction are

likely to vary with potential, and so the above calculations will only lead to an average

value for the chain length of the polysulphide. Voltage pulse studies can provide a

more accurate estimate of the polysulphide that is produced under a particular reduction

potential.


In a highly alkaline, concentrated solution of HS', the chemical dissolution of sulphur

is favoured (3.11). A ring-disc electrode study in a solution containing 1 kmol

Na2S.9H20 m' 3 and 1 kmol NaOH m-3 showed that the polysulphide species could be

detected in the positive-going scan, as shown in Fig. 3.12:

Fig. 3.12 Ring-Disc Voltammetry of Sulphide Solution at Au RRDE.

[NaOH] = 1 kmol n r 3, [HS“] = 1 kmol m“3, co = 4 Hz, nth* scan.

Scan rate = 20 mV s"1, ring potential = -0.90 V vs. SHE.

Although a polysulphide reduction current was still detected on the negative-going (disc

potential) scan, the largest current was seen during the positive-going scan. Once

sulphur layers have been formed, they can be dissolved by the high concentrations of

HS" flowing over the electrode surface, to produce polysulphide ions which can be

reduced at the ring.

Potential pulse studies also confirmed that polysulphides were produced during the

formation and reduction of elemental sulphur. The experiments were conducted in

deoxygenated solutions containing 10 mol HS" m-3 at pH 9.3. The ring potential was

maintained at -0.7 V vs. SHE throughout, and the disc potential was stepped from this

to a more positive value for 4 s, and then stepped back to -0.7 V vs. SHE. The


experiment was repeated, with and without electrode rotation, for different values of the

disc potential step. The resulting current response is shown in Fig. 3.13:

Disc

Potential-0 7 V

0 0 Vvs. SHE

Us.-----

Fig. 3.13 Ring-Disc Potential Pulse Study. Au RRDE.

[HS“] = 10 mol n r 3, co = 9 Hz. pH 9.3.

The potential step to 0.0 V vs. SHE was sufficient to form multilayers of elemental

sulphur, and at the ring a small reduction current was seen due to the detection of

polysulphides. When the potential was pulsed back to -0.7 V vs. SHE, the sulphur at

the disc was reduced, and an increased reduction current was observed at the ring. If

the electrode was stationary, the ring current decreased to zero (apart from the

capacitive current spikes which were produced when the potential was pulsed). The

ratio of ring to disc currents were lower than those expected for the production of a

poly sulphide on average chain length 1.8. This suggested that either that less

polysulphide was produced (a calculation of the average chain length gave n = 1.1) or

that the ring electrode had become deactivated. The ratio of ir to i^ rose rapidly in the

first 250 ms following the potential pulse, thereafter remaining approximately constant

at about 0.025. The reduction of the sulphur layers, produced when the potential was

pulsed to -0.7 V vs. SHE, resulted in the production of a smaller proportion of

polysulphide ions than was observed in the potential scan study.

If the disc potential was pulsed to potentials below that at which multilayers of sulphur

can form, i.e. no higher than -0.1 V vs SHE, no current response was seen at the ring.

When the disc potential was stepped to 0.1 and 0.2 V vs. SHE, higher ring currents

were seen while the disc was held at these potentials. This indicated that more

polysulphides were formed, either through direct oxidation of HS" ions or by

dissolution of the sulphur layer.


3.7 Calculated Polysulphide Concentrations vs. Potential

A knowledge of the equilibrium concentrations of polysulphide ions as a function of

potential can be used in two ways. It can enable the chain length of a polysulphide

produced under a particular electrode potential to be predicted, and it can be used to

calculate the composition of a polysulphide solution from the solution potential.

The concentrations were calculated by considering the following equilibria:

S22" + 2e- 2 S2' Ei (3.14)

2 S32- + 2 e- *-» 3 S 22- e2 (3.15)

3 S42" + 2 e ‘ <-> 4 S32" e3 (3.16)

4 S52- + 2 e- 0 5 S42" e4 (3.17)

S2- + H+ HS- p K - ^ S ) (3.18)

The standard electrode potentials (Ej0- E40) and pK^CH^S) were calculated from the

free energy of formation data in Zdhanov’s review [31]. The redox potentials were all

set equal. Since these potentials are dependent on the species concentration via the

Nemst equation, a set of simultaneous logarithmic equations are produced, which were

solved iteratively using a numerical method and the program TKSOLVER. The

process was repeated for a range of potential values, and in this way a profile of the

polysulphide concentrations was built up. The results are shown in Fig. 3.14 and

Fig. 3.15:

Fig. 3.14 Polysulphide Distribution vs. Potential (pH = 14)


Fig. 3.15 Polysulphide Distribution vs. Potential (pH = 9)

These figures indicate that the potential ranges of predominance of the individual

polysulphide species are small. A range of only 100 mV separates the areas of stability

of S52- and S22-. This means that using the solution potential to calculate the

polysulphide concentration is insensitive; a deviation of only a few millivolts in the

recorded potential would seriously alter the calculated composition. Furthermore,

when gold indicator electrodes were used to measure the potential of polysulphide

solutions, values outside the theoretically predicted likely potential range were obtained;

e.g. -0.14 V vs. SHE for a solution containing 0.1 mol Na2S4 m~3 at pH 9.3.

Solutions made from Na2S4 disproportionate, forming a range of polysulphide species

in equilibrium. If these species were not equilibrating reversibly at the gold electrode

surface, the observed solution potential would be outside the theoretical range.

Inaccuracies in the published thermodynamic data may also account for this difference.

Certainly there is a conflict between equilibrium constants that are calculated from the

thermodynamic data, and those which have been determined experimentally [53].

However, the above figures do indicate that there will always be more than one polysulphide present in significant amounts at any particular solution potential (if the species can equilibrate under thermodynamic control).


3.8 The Detection of Polysulphides Using Ion Chromatography

Because of their possible importance as an oxidation intermediates in the Stretford

Process, it was decided to investigate the possibility of detecting poly sulphides using

the technique of ion chromatography. Although polysulphides absorb in the

UV-visible spectral range, spectrophotometry can not be used routinely since other

components of a Stretford solution absorb in the same spectral region.

3.8.1 Ion Chromatography: Results and Discussion

Using the calculations outlined in section 3.1.4, it was determined that only three

solution species were important in polysulphide mixtures in the pH range 10-14; S42",

S52" and HS". An injection of polysulphide solution resulted in the detection of three

peaks, as can be seen from Fig. 3.16:

20-

Electrochemical Detector , Response

1-5-

UV Detector Response

1 0-3 mM HS"

2 10 mM Sj?"

3 10 mM H$ Saturatedwith elemental Sulphur

cJj Oc_OIS)

4

£ 00005

O l

O3

Fig. 3.16 Ion Chromatography Results


The peak heights of the second and third peaks (those two peaks having the longest

retention times) correlated well with the calculated concentrations of S42' and S52".

However, the response of the first peak, which was assigned to the HS' ion, was

larger than expected. This suggested that disproportionation had occurred as the

sample traversed the column.

3.9 Summary

Oxidation of HS" ions at pH 9.3 has been shown to produce a sub-monolayer of

adsorbed sulphur on a gold electrode at low potentials (-0.4 V vs. SHE), and

multilayers of sulphur at higher potentials (0.05 V vs. SHE). Associated with the

formation of elemental sulphur is the production of polysulphide anions, Sn2" (n = 2 to

5), which can also be produced by the dissolution of the initial sulphur layer. The

production of such polysulphide species accounted for the difference in charge between

the positive and negative-going scans.

Upon reduction of the sulphur layers, polysulphide ions were produced, which were

detected at a ring electrode in a rotating ring-disc electrode study. By a comparison of

the charges passed in the production of these ions from elemental sulphur, and their

reduction to HS" ions, an estimate of the average polysulphide chain length could be

gained. This was calculated to be 1.8, which suggests that a mixture of poly sulphides

was produced. This is consistent with thermodynamic predictions which show that a

number of polysulphides can exist in solution at comparable concentrations at any given

potential.

Ion chromatography was investigated as a means of detecting polysulphide ions in

solution, but disproportionation of the polysulphide species as they traverse the ion

exchange column and the air-sensitive nature of the solutions, made the method

unsuitable for routine use.

Vanadium Review 6 6

4. VanadiumVanadium is a lustrous, corrosion-resistant metal which is used in large quantities to

make steel alloys tougher. There are only a few concentrated deposits of vanadium

minerals, and most vanadium is generated as a co-product from the processing of iron,

phosphorus or uranium ores. The vanadium content is dissolved from these ores under

oxidising and acidic conditions, forming a solution containing the VC>2+ ion; the pH is

then raised and vanadium anions are formed (see section 4.1) which can be separated

from the aqueous solution using solvent extraction. Recently ion-exchange columns

have also been suggested for this purpose [97]. The vanadium is then stripped from

the organic solvent or ion-exchange resin to form a concentrated aqueous solution, and

precipitated as the oxide. Conventional pyrometallurgical reduction of the oxide is not

facile, since vanadium reacts with oxygen, nitrogen and carbon at high temperatures;

instead, most vanadium is produced as the iron alloy, f e r r o v a n a d i u m . Very pure

vanadium can be produced by the reduction of the chloride VCI4 with magnesium, the

reduction of the pentoxide V2O5 with aluminium, or using the Van Arkel method which

relies on the decomposition of the iodide VI3 at high temperatures.

Vanadium has atomic number 23 and mass number 51, and is in the first row of group

VB in the transition series; it can exist in oxidation states between II and V in aqueous

solution. Vanadium, like other metals in this region of the periodic table (Nb, Ta, Cr,

Mo and W) shows a marked tendency to form polymeric ions in aqueous solution,

which makes the chemistry of vanadium both interesting and complex.

Various authors have produced E^-pH diagrams showing the thermodynamically most

stable species [61,90,97-100]. These diagrams differ from each other for a variety

of reasons. Firstly, AGf° values for the polymeric vanadium species are not always

available; Pourbaix [6 1 ], for instance, did not consider the formation of any

decavanadate species, although their existence and AGf° values have now been well

established. Secondly, lines are drawn between two solution species under different

criteria, Pourbaix used the criteria of equal activities, whereas Post and Robins [99]

drew the lines to show when the total dissolved vanadium was distributed equally

between the two species. When polymeric species are involved, this can lead to a

considerable shifting of the equilibrium lines. Finally, diagrams are drawn under

different total vanadium concentrations; Zipperian and Raghavan produced a diagram

using a very low vanadium concentration of 0.2 mol m-3 and as a result show only

monomeric species, whereas Post [98] considered the equilibria at a total vanadium

concentration of 1 kmol n r 3 and showed the polymeric species V2074- and

V10O286" had areas of predominance.

Post [99] also noted that some of the available thermodynamic data on vanadium had

been misprinted in the past, which had passed unnoticed by other authors; he corrected

these values where necessary: Fig. 4.1 (overleaf) shows a diagram using his

corrected data [98]:

Vanadium Review 67

p H

Fig. 4.1 Ejj-pH Diagram for the Vanadium-Water System [98]

1 kmol V m-3 T = 298 K.

4.1 Vanadium (V)

Vanadium (V) solutions can be prepared by dissolving sodium metavanadate (NaVC^),

ammonium metavanadate (NH4VO3), or vanadium (V) oxide (V2O5). It is the only

oxidation state that is stable in air throughout the entire pH range. Babel et al. recently

reviewed the use of vanadium (V) as an oxidising agent in acidic media [101]; they

showed that it could be used to oxidise many organic and inorganic compounds (e.g.

alcohols to aldehydes and SC^2- to SO42'). Under these conditions, the predominant

species are VC>2+ ions. If such a solution is made alkaline, tetrahedral VO43- ions are

produced. As dilute (< 0.01 mol V(V) m“3) alkaline solutions are made more acidic,

VO43- ions protonate to form HVO42- and H2VO4- (at pH 13 and 8 respectively).

Hydrated vanadic acid (usually written as HVO3) precipitates at about pH 5 [102].

The situation is more complex with higher V(V) concentrations; consider the

acidification of a moderately concentrated vanadium (V) solution (10 mol m-3) at pH

14. Initially, only VO43- ions are present and the Raman spectrum is simple. As the

pH falls below 14, new bands appear at 810, 503 and 228 cm-1, which have been

assigned to V-O-V stretches [103]. This is consistent with the presence of the dimeric

V2O74- ions. Similar changes are seen in the 5iV NMR spectrum [104]; in strongly

alkaline solution a single absorption at 536.2 ppm (relative to VOCI3) is observed, due

to VO43- ions, whereas below pH 14 this peak shifts slightly to 533 ppm and a second

peak is seen at 556.2, due to the V2O74- species. If the pH is further lowered to below

Vanadium Review 6 8

11 this line broadens and shifts to 562 ppm, which is consistent with the protonation of

these ions to form HV2O73-. The presence of dimeric species is confirmed by the

precipitation of Na4V207 at concentrations in excess of 1 kmol V(V) m"3 .

In the pH region 7-8 there were early disagreements as to whether the vanadium was

present in trimeric or tetrameric species. However, Ingri and Brito [105] showed that

below 20 mol V(V) m"3, the V3093" ions predominate, and at higher concentrations

V4O124" ions are formed. Habayeb and Hileman studied vanadium (V) speciation

using 51V NMR [32] and confirmed the presence of V40i24"- They also collated the

chemical shift data for the known vanadium (V) species; this compilation is shown as

Table 4.1:

Species 51V NMR Shift / ppm

vo43- -536.2

hvo42- -533

v 2 o 74- -556.2

HV20 73- -562

V30 93- -573

V3O105- -410, -500, -516

O 1 -577

h2v 4o 134- -582

V50 155- -586

1OI> -582

V 1 0 ° 2 8 6' -419,-495,-510

Table 4.1 51V NMR Chemical Shifts of V(V) Species.

Around pH 6.5, the species V50 i5^“ and VÔjg6- have been reported [33], and they

are seen as possible precursors to the decavanadate structure. It must be noted that acid

titrations are difficult to interpret in this pH region, since some of the equilibria,

especially those involving polymeric anions, are achieved very slowly. This may

account for some of the early disagreements regarding vanadium (V) speciation.

However, there is now no doubt that between pH 2 and 6, and at vanadium

concentrations greater than 1 mol m-3, the orange decavanadate ions are formed;

V io028^“> and ^ V jqC^s4'- The presence of these species has now been

established by potentiometric [105 ], cryoscopic [106 ], and spectroscopic [107 ]

methods.

Vanadium Review 69

Solid salts containing decavanadate anions have been isolated, and the minerals pascoite

(Ca3Vio028-7H20 ) and hummerite (K2Mg2Vio028-16H20 ) have been shown to

contain them [108]. The structure of the V^gC^s6- ion consists of ten linked VOg

octahedra [109]:

Fig. 4.2 Structure of the Viq0 286' ion.

If decavanadate solutions are allowed to stand for several weeks, or if they are warmed,

sparingly soluble orange salts precipitate. Although they are termed trivanadates, and

have stoichiometries such as KVgOg, they do not contain V30g" ions, instead having a

structure consisting of layers of linked VOg octahedra separated by layers of cations

[ 110].

Vanadium Review 70

The available information on vanadium (V) speciation has been summarised in the

form of an activity-pH diagram by Post and Robins [99]:

Fig. 4.3 Vanadium (V) Speciation [99].

4.2 Vanadium (IV)

Vanadium (IV) is stable in aqueous solution provided that oxygen is excluded. Post

[98] and Rossotti and Rossotti [111] showed that the predominant species below

pH 3 was the blue vanadyl ion, V 0 2+. Above this pH they suggested that the two

complexes, VO.OH+ and (VO)2(OH)22+, were formed. Rohrer et al noted that a series

of solid sulphate complexes were obtained between V 0 2+ and concentrated sulphuric

acid [112]. At about pH 4 a solid precipitates from aqueous solution which has the

composition VO(OH)2 (which could be regarded as V2O4. 2H2O). They noted that this

was soluble in excess alkali, forming brown, air sensitive "vanadite" solutions.

Vanadium Review 71

Pope has recently reviewed the vanadium (IV) speciation in such alkaline solutions

[33]. Crystalline alkali metal salts can be precipitated from vanadium (IV) solutions,

and although they have the empirical formula M2V3C>7.nH2C), they have been shown to

contain the polyanion Vig04212"[113 ]. This anion consists of an almost spherical

shell of linked VO5 square pyramids, surrounding a central cavity about 0.45 nm in

diameter:

Fig. 4.4 Structure of Vi804 212‘.

In the solid salts the central cavity is occupied by a potassium ion or a water molecule.

The anion appears to be stable in vanadium (IV) solutions between pH 9 and 13, and at

concentrations above 2 mol V(IV) m"3 [33]. As it has only recendy been identified, no

value for AGf° has yet been proposed, and for this reason it does not appear on any of

the published E^-pH or activity-pH diagrams.

Vanadium (IV) has a magnetic moment of 1.73 Bohr Magnetons, which means that,

unless the spins are completely paired, the 51'V NMR spectra are poorly resolved. This

has limited the information available concerning the vanadium(IV) speciation in

solution, and much less is known about vanadium (IV) anions than vanadium (V)

anions. Post and Robins [99], although omitting any mention of the V ^g C ^12- ion,

do show four vanadium (IV) species: V02+ (low pH), HV2O 5- (high pH),

(V O )2(O H )22+ and V4O92- (at high V(IV) concentrations). The situation is

summarised in Fig. 4.5 (overleaf).

Vanadium Review 72

1 2 3 4 5 6 7 8 9 10 11 12 13 14pH

Fig. 4.5 Vanadium (IV) Speciation in Solution [99].

4.3 Vanadium (V)/(IV ) Compounds

Several mixed-valence soluble polyvanadates have been reported in the literature; they

have been formulated as partially reduced decavanadate structures,

e.g.HV3IVV7V0286“. Ostrowetsky reported six ions in the pH range 4 to 6.5 with

y iv :y v ratios ranging from 2:8 to 7:3, of which the green 3:7 and 7:3 ions were the

most stable [114]. Solid mixed-valence alkali metal vanadium oxides are also known

[115]. Many of these compounds are semiconductors, and some of them show a

metallic lustre; they have been termed "vanadium oxide bronzes" e.g. K2V3O8.

It is possible to prepare a range of 19-nucleate blue-violet anions having y lv:VV ratios

from 5:14 to 7:12 [33] (e.g K g H V ^ V V ^ O ^ .l l^ O ). The basic structure of these

anions consists of an ellipsoidal cluster of 18 VOn polyhedra. These formulae are

almost twice that of those proposed by Ostrowetsky [114] and it is possible that his

assumption that the ions are based Y \ q clusters is incorrect..

Hayek and Pallasser [116] obtained the mixed-valence crystalline decavanadate

structures Na6VIV8Vv20 24.8H2 0 and K6V lv8VV20 24 .5H 20 ; they acidified

thiovanadate solutions containing V:S ratios in the range 1:1 to 1:4 using acetic acid.

Vanadium Review 73

The pH was lowered to about 8.5, and after 12 hours heating the brown-black crystals

precipitated from solution. The sulphide solution had effected the partial reduction and

itself been oxidised to elemental sulphur, which was removed by washing with carbon

disulphide. The mixed-valence ammonium salt (NH4)2V30g.l/2 H2O was prepared in

a similar manner, thought it did not to contain a decavanadate ion, but instead consisted

of linked VIVC>5 square pyramids and VV207 di-tetrahedral units. It is interesting to

note that even in the presence of excess reducing agent, the vanadium (V) is not

completely reduced to vanadium (IV) compounds, but instead forms a mixed-valence

precipitate.

Post [98] studied the atmospheric oxidation of vanadium (IV) solutions in acid

solutions, and found that at pH 2.5 and in the presence of sodium ions, blue or brown

mixed oxidation state solids were obtained. He reported one solid with a Vv: V ^ ratio

of 1:4 which had similar properties to the mineral corvusite (~V2C>4 8 .1/2 H2O).

Oxides with intermediate V/TV stoichiometries such as VgO^ can also be produced by

heating the appropriate masses of the oxides V2O5 and V2O3 at 600 °C for 10 hours

[117]. V gO ^ can also be prepared by reducing V2O5 in a stream of hydrogen [118].

Vanadium Review 74

4.4 Vanadium (III)

The oxide V2O3 is not amphoteric, unlike the vanadium (IV) and (V) oxides. It is

insoluble in alkaline solutions, but dissolves in acid forming green V(H20) 63+ ions.

Above pH 1 these hydrolyse to form VOH2+ and V2(OH)24+, and if the pH is further

raised to 7, V2O3 precipitates. The vanadium (III) speciation is summarised in Fig.

4.6:

Fig. 4.6 Vanadium (III) Speciation in Solution [99].

4.5 Vanadium (II)

Vanadium (II) represents the lowest accessible oxidation state of vanadium. The

electronic configuration is cfi, which confers upon the aqueous species a kinetic

inertness, and the ligand substitution reactionsAthe purple V(H20)62+ ion are slow. It is

a powerful reducing agent, and is oxidised by water. Because of this instability little is

known of its hydrolysis behaviour [98]. It is readily oxidised by atmospheric oxygen,

forming the green V(H20)g3+ ion, and it has been used to remove trace amounts of

oxygen from inert gases [100].

Vanadium Review 75

4.6 Vanadium Electrochemistry

Most studies on the electrochemistry of vanadium have been carried out in highly

acidic solution, where each of the oxidation states V(II), V(III), V(IV), and V(V) can

be produced by controlled potential electrolysis [100]. A comprehensive review of the

work done up to 1976 is provided by Israel and Meites [30]. In acidic solution

polarography of V(V) solutions is difficult because V(V) is capable of oxidising a

mercury surface. However, in alkaline solution mercury metal should still be stable at

potentials high enough to oxidise V(IV) to V(V) [61 ]. One problem with the

interpretation of polarographic results is that vanadium coatings may catalyse other

reactions; vanadium (V) has been shown to be an electrocatalyst for carbon oxidation

[119] , and vanadium alloys reduce the overpotential required for hydrogen production

[ 120] .

4.6.1 The V(V)/V(IV) Couple

In acidic solution, vanadium (V) causes oxidation of mercury and platinum electrodes,

and consistent pre-treatment is required to obtain meaningful results. In the pH range 7

to 10, of most relevance to the present study, relatively few investigations have been

attempted.

Below pH 2, the reduction of a vanadium (V) solution proceeds in two stages; the

reduction from VC>2+ to V 02+ is reversible, but a further decrease in electrode potential

causes the production of V(H20 )g2+ [100]. Magri-Elouadseri and Vittori [121],

studied the electrochemical behaviour of carbon paste electrodes incorporating

vanadium (V) solids at pH 0. They found that the vanadium (V)/(IV) couple was

reversible and gave a half wave potential close to the expected standard potential

(0.944 vs. SHE), but that at more negative potentials subsequent reduction proceeded

directly to produce vanadium (II).

Van den Berg and Huang [122 ] conducted polarography on vanadium (V) at

concentrations of 0.02 mol m-3 and at pH 7. They found that the vanadium (V)

underwent reductive adsorption at potentials of -0.678 V vs. SHE, forming

vanadium (IV) on the mercury surface, which was further reduced to vanadium (II) at

a potential of —1.0 V vs SHE. Between pH between 2 and 9, up to four reduction

waves were observed by Filipovic et al. [123]. The first two they assigned to the

adsorption and reduction of hydrogen polyvanadate ions, the third was attributed to the

reduction of dissolved vanadium (V) to vanadium (IV), and the fourth to a further

reduction to form vanadium (II). From pH 9 to 12.5 they observed only the third and

fourth waves. They noted that the reduction to vanadium (IV) occurred only after an

overpotential of about 0.8 V had been applied, (e.g. when a potential of —1.16 V vs

SHE was reached at pH 9.3). At solutions with a pH higher than 12.5, Filipovic

[123] and other workers [30] have found a single irreversible reduction wave which

has been attributed to the reduction of V(V) to form V(II).

Vanadium Review 76

Stromberg et al. [124] found that vanadium oxide films were obtained when reducing

potentials were applied to platinum or carbon electrodes in alkaline vanadate solutions.

They proposed that the initial films (on carbon electrodes) consisted of V2O3, and that

this was converted to V2O2 at potentials lower than -1.15 V vs SHE. However, they

had no direct evidence of the film compositions, and relied solely upon thermodynamic

predictions.

4.6.2 V(IV) Reduction

It appears that the reduction of V(IV) at the dropping mercury electrode in acid media is

totally irreversible and proceeds directly to V(II) [100,123]. At a carbon paste

electrode, the reduction of V(IV) to V(III) was found to be very slow [121], and the

potential had to be lowered until a reduction process forming V(II) occurred.

Gala et al. studied the deposition of vanadium from vanadium (IV) solutions onto steel

cathodes at pH 10 [120]. They noted that in the metal could not be deposited from

solutions containing vanadium alone, but suggested that elemental vanadium could be

co-deposited with nickel, forming an alloy. To achieve this, the potential had to be

lowered to such a level that hydrogen was also evolved. However, their analysis

techniques did not distinguish between vanadium in a nickel alloy, and entrained grains

of vanadium oxides. Indeed, they did not consider the possibility of entrained phases.

4.6.3 The V(III)/V(II) Couple

The reduction of vanadium (III) in acid solution is reported to be reversible on mercury

and carbon paste electrodes [100 ,121]. Filipovic et al. [123 ] report a half wave

potential of -0.29 V vs. SHE, which is close to the expected standard potential of

-0.263 V vs. SHE. In alkaline solution vanadium (III) forms solid V2O3.

4.7 Oxidation of Vanadium (IV) Solutions using Oxygen

The oxygen/water half cell can apply sufficient potential to oxidise vanadium (IV)

solutions:

E °(02/H20 ) = 1.23 V, E°(V(V)/V(IV)) = 0.944 V (at pH 0).

Since the potential of the vanadium(V)/(TV) couple decreases with pH at a greater rate

than the O2/H2O couple, the thermodynamic driving force for vanadium (IV) oxidation

using oxygen increases with pH. This explains why acidic vanadium (IV) solutions

can be handled without taking any special precautions to exclude air, whilst alkaline

solutions are air sensitive.

Post [98] studied the oxidation of vanadium (IV) by oxygen under acidic conditions.

Working at a temperature of 90 °C, he noted that the oxidation proceeded with a

decrease in pH, due to reactions such as:

4 V 02+ + 2 H2O + O2 —̂ 4 V02+ + 4 H+ (4.1)

This change in pH may alter the predominant vanadium solution species, and so alter

the reaction mechanism.

Vanadium Review 77

Dean and Herringshaw [125] looked at the air oxidation of vanadium (IV) in alkaline

solution. They showed that oxidation to V(V) was rapid and complete (0.8 mol

V(TV) m-3 being completely oxidised by the dissolved oxygen in air saturated solutions

in 10 s at pH 14). They noted that under conditions of excess oxygen, hydrogen

peroxide was produced as the oxygen reduction product, and that this itself was capable

of oxidising more V(IV).

4.8 Vanadium Sulphides

Vanadium can form a number of solid sulphide phases with varying ratios of V:S,

some of which have not been fully characterised [126]. Like the oxide phases, mixed-

valence compounds are known, and most vanadium sulphides possess sulphur-sulphur

bonds. Mixed metal Mo-V sulphides are also known [127]. Vanadium sulphides are

electrically conducting, and show paramagnetic behaviour owing to the presence of

unpaired electrons. Table 4.2 shows some of the known vanadium sulphides

(thermodynamic data are from Mills [128]):

Compound Comments AG f° / k j mol- 1

v 3s Can exist in two metallic forms.

V5S4 Metallic structure

VS Non-stoichiometric solid -192v 7s8 Hexagonal Structure.

V3S4 Layered structure

V2S3 Prepared by direct reaction -518

v 5s8 Monoclinic structure.

V2S5 Prepared by decomp, of (NH4)3VS4v s 2 Non stoichiometric

vs4 Exists as mineral Patronite -413vs5 Amorphous semiconductor

Table 4.2 Some Known Vanadium Sulphides.

4.8.1 V3S, V5S4, VS

The compound with the highest V:S ratio is V3S, which can exist in two forms, both of

which are metallic. V5S4 is also reported to have a metallic structure [126]. As the

sulphur content is increased the metallic character is lost and the compounds become

semiconducting. Stoichiometric VS is unstable at room temperature and dis-

proportionates to form cation-deficient V7Sg and cation rich V9Sg; V7Sg has a

hexagonal NiAs-type structure.

Vanadium Review 78

A non-stoichiometric range of compounds with formula from Vq̂ sS to Vq̂ S are

also known. Within this range, the compounds V2S3 (Vq̂ S ) and V3S4 (V0/75S) have

been prepared. V2S3 can be prepared by heating vanadium pentasulphide to 300 °C in

an inert atmosphere [129]:

V2S5 -> V2S3 + 2S (4.2)

V3S4 is prepared by combination of the elements at 800-1000 °C; it has a monoclinic

unit cell and is thought to consist of alternate layers of V2+ and V3+ ions. V3S4 absorbs water at room temperature, and loses H2S when it is heated, forming an oxide

phase. If V3S4 is heated in air or oxygen, it oxidises to form V2O3, V2O4 and V2O5 successively, evolving SO2 [130].

4.8.2 V2S 5Vanadium pentasulphide is the sulphur analogue of vanadium pentoxide. It can be

prepared by heating ammonium tetrathiovanadate (see section 4.9) to 100 °C in an inert

atmosphere, whereupon ammonia and hydrogen sulphide are evolved [129]:

2 (NH4)3VS4 -> V2S5 + 6 NH3 + 3 H2S (4.3)

V2S5 is a black amorphous powder, insoluble in water, alcohol, ether or carbon

disulphide. If it is heated above 290 °C in the absence of air it decomposes to form

V2S3; if air is present, it oxidises readily at 100 °C forming vanadium pentoxide:

2 V2S5 + 15 0 2 -> 2 V20 5 + 10SO2 (4.4)

4.8.3 VS2 and VS4Stoichiometric vanadium disulphide is not known, although a compound of

stoichiometry Vj 2^2 has been reported. If the proportion of sulphur is further

increased, VS4 is produced. VS4 exists in nature as the mineral patronite, and can be

prepared in the laboratory by heating the elements together at 400 °C for several weeks.

The structure is monoclinic,the vanadium atoms sitting in between S22- pairs; in this

respect the mineral is similar to pyrite (FeS2).

4.9 Vanadium -Sulphur complexes

If hydrogen sulphide is passed into an alkaline solution containing V, Mo or W anions

a range of colours are produced. These colours are due to the thioanions of the

transition metals and depending on the metal, pH, and metal to sulphide ratio, virtually

any colour can be produced. Muller [131] recently reviewed the transition metal

thiometalates. He noted that, as hydrogen sulphide was passed through an aqueous

oxometalate solution, changes in the UV-visible and IR raman spectra were consistent

with the successive formation of MO411", MC^S11-, MO2S211", MOS3n_, and MS4n‘

(M = V, Mo, W or Re). The rate of formation of these thiometalates was governed by

the polarising power of the central metal atom; the lower the polarising power of the

metal, the greater the electron density on the oxygen atoms and hence the faster the rate

of complex formation.

Vanadium Review 79

In aqueous solution thiometalates are not very stable, especially at low pH. They can

be hydrolysed to form oxometalates, they can form solid metal sulphides, or they can

undergo intramolecular redox processes:

MfS2’̂ -> Mr-2(S22-) (e.g. M = Mo, r = 6) (4.5)

Thiometalates can be attacked by nucleophiles to give a reduced metal centre, and this

process of sulphur abstraction is more apparent in vanadium than molybdenum

complexes:

Mr-S + Nu Mr_2 + NuS (e.g. Nu = CN‘) (4.6)

This reaction may explain why when VO43- was reacted with H2S in the presence of

CN“ a polyvanadate with a low vanadium valence was obtained [132].

Ranade et al.[133] prepared a series of thiovanadate complexes by passing H2S

through weakly buffered ammoniacal solutions containing 0.1 mol V(V) m-3 at 5 °C.

They found that initially a complex was formed which absorbed at 360 nm in the UV-

visible spectrum (and weakly at 305 and 460 nm). This spectrum was similar in

structure to the isoelectronic complex M0O2S22", and so they attributed it to V02S23-.

(Because of the high affinity of vanadium for sulphide, the monothiovanadate (VO3S3')

could not be produced in aqueous solution, although it could be formed in a methanolic

solution). As further H2S was passed through an aqueous solution the trithiovanadate

and tetrathiovanadate complexes were formed:

VO2S23- + H2S VOS33" + H20 (4.7)

VOS33- + H2S VS43" + H20 (4.8)

Yatsimirskii and Zakharova [134] studied the hydrolysis of VS43-. They concluded

that on dissolution in sodium hydroxide solutions at pH 13-14, solid ammonium

tetrathiovanadate dissolved with hydrolysis to form V02S23-, and that this rapidly

hydrolysed further to form the vanadate ion, VO43".

From concentrated solutions, solid salts containing the tetrathiovanadate ion can be

prepared. Busine and Tridot prepared the ammonium salt [129] by passing hydrogen

sulphide through 5.9 kmol m~3 ammonium sulphide solution containing 37 mol

V(V) m-3; after several days at O °C intensely-coloured violet crystals were produced.

If a more concentrated vanadium (V) solutions was used, or if the temperature was

raised, the vanadium (V) became reduced and vanadyl hydroxide (VO(OH)2)

precipitated. This precipitate would redissolve in excess ammonium sulphide,

suggesting that vanadium (IV) thiosalts can also be produced.

Vanadium Review 80

Harrison and Howarth [37] folloy&d Busine and Tridot's method and prepared a

range of thiovanadate complexes. -Whey determined the 51V NMR shifts (relative to

VOCI3) of the free anions and their.Trotonated forms. Their results, together with a

summary of the UV-visible spectr&/(from [133]), are summarised in Table 4.3.

51V NMR

Species Colour UV-visible Absorbances Chemical Shift

X / nm (e / m2 mol*1) ppm vs. VOCI3VO43- Colourless -541

V 03(0H)2- Colourless -539

V03S3- (orange) 305,442 (in methanol) -250

HV03S2' -121V 0 2 s 2 3 - Yellow/red 305, 360 (8-400), 460 (weak) 184

HV02S22- <230

VOS33- Red 295, 325, 459 (e~600), 521 740

HVOS32- 748

VS43- Violet 267, 351, 394, 538 1395

HVS42- 1392

Table 4.3 Spectral Summary of Thiovanadates [37,133].

4.10 Summary

Vanadium, in common with other transition metals in the same region of the periodic

table, shows a tendency to form polymeric anions in alkaline solutions. The degree of

condensation in these species is highly dependent on the total vanadium concentration.

Efo-pH diagrams are only of limited value in determining the predominant vanadium

species under particular solution conditions. This is partly due to the above dependence

of the vanadium speciation upon concentration, and partly because thermodynamic

values are still not available for key polymeric species. However, in S tretford

Process solutions it is likely that the dimeric and tetrameric species H V ^C^" and

V4O124" ^ present.

Hydrogen sulphide can interact with; vanadium (V) in two ways; as a reducing agent

and as a complexing agent. Complete reduction of Stretford Process solutions to

the V(IV) oxidation state is likely to produce the brown polyanion, V ig C ^ 12". On

prolonged exposure to reducing environments it is conceivable that further reduction

will occur, forming a precipitate of.vanadium (III) oxide, V2O3. Mildly reducing

conditions, or re-oxidation of vanadium (IV) solutions, can produce mixed-valence

(V)/(IV) compounds (e.g.VIvgVX2024^~)* The sodium salts of these ions may precipitate if the sodium ion concentration is high.

Vanadium Review 81

Thio complexes are known to be produced when (ammoniacal) vanadium (V) solutions

contact H2S. As sulphur is substituted for oxygen in the vanadate ion (VO43-) the

complexes VO2S23-, VOS33- and VS43- are formed.

Since it is known that isoelectronic thiomolybdate complexes can undergo

intramolecular redox processes, it is possible that vanadium (V) catalysis of sulphide

oxidation proceeds via thio-complex formation followed by an intramolecular redox

reaction. Thus, initially a complex such as Vv02S23- might be formed, which is

converted to Vm02S23-; the disulphide ion so produced may then desorb from the

complex. In this way vanadium (V) could oxidise sulphide solutions producing

polysulphide solutions and reduced vanadium species.

The electrochemical reduction of vanadium (V) in alkaline solution is slow, and large

overpotentials are required to obtain measurable currents; the formation of thio-

complexes may offer reaction pathways with a lower activation energy, and so allow a

higher rate of sulphide oxidation.

Vanadium Electrochemistry 82

5. Vanadium ElectrochemistryThe reduction kinetics of vanadium (V) in solution at pH 9 was investigated at a variety

of electrode surfaces. In the absence of specific chemical interactions, oxidising agents

that show reversible behaviour at electrode surfaces are reduced rapidly by chemical

means, whereas oxidising agents which show irreversible reduction at an electrode

react only slowly with chemical reductants.

5.1 Vanadium Electrochemistry: Experimental

Cyclic voltammetry and potential pulse studies were carried out using a Tacusel

hanging mercury drop electrode (HMDE). The electrode consisted of a mercury

reservoir connected to a glass capillary. By turning a micrometer drive, a drop of

mercury was made to hang from the capillary tube. The bore was cleaned prior to use

with 6 kmol HNO3 m~3 and triply distilled water, then rendered hydrophobic by

treating it with a solution of dimethyldichlorosilane (2 % in 1 ,1,1-trichloroethane,

BDH). A diagram of the HMDE is shown in Fig. 5.1:

/

Micrometer thread

/

\

Electrical contact

Mercury resevoir

f . Silicone rubber seal

Glass capillary

iH M ercury bead

Fig. 5.1 Hanging Mercury Drop Electrode.


The surface area of one drop was calculated by making 25 complete revolutions of the

micrometer drive, and measuring the mass of mercury ejected. From this, the average

mass of mercury ejected by one revolution was derived; assuming that the drop had a

spherical shape and knowing the density of mercury, enabled the surface area to be

calculated (1.974 x 10~6 m2). The capillary bore was 100 |im in diameter, and it was

shown that the area of attachment corresponded to only 0.4 % of the total drop area.

Gold, platinum and vitreous carbon rotating discs (see section 3.2) were also used as

working electrodes. Bright platinum counter electrodes were used in all cases, and the

potentials were controlled relative to saturated (KC1) calomel reference electrodes

(EIL). All potentials are reported versus the standard hydrogen electrode (SHE),

assuming that the potential of the saturated calomel electrode was 0.242 V vs. SHE.

The electrochemical studies were carried out in a three compartment cell (see Chapter 7,

Fig. 7.1) using a Thompson Ministat MP81 potentiostat. The control potentials

were provided by a Hi-Tek PPR1 waveform generator and the currents were passed

through a standard resistor. The resulting voltages were then applied to the inputs of a

J J PL4 chart recorder.

Most studies were undertaken at room temperature (~20 °C), but a series of

experiments were recorded at 40 °C using a jacketed electrochemical cell which

contained heating water maintained at 41 °C by a Grants thermostatted water bath.

Voltammograms were commenced from the positive potential limit (0.331 V vs. SHE

for a mercury electrode) for vanadium (V) solutions, or from the rest potential for

vanadium (IV) solutions (-0.193 V vs SHE).

5.1.1 Solution Preparation

A carbonate buffer of pH 9.3 was prepared by dissolving the appropriate mass of

analytical grade chemicals (BDH) in triply distilled water to produce a solution

containing 0.059 kmol Na2C03 0.223 kmol NaHCC^ n r 3 and 0.1 kmol

Na2S04 m~3. A borate buffer of pH 9.2 was similarly prepared, containing 12.5 mol

Na2B4Oy.lO H20 m"3,0.9 mol NaOH m"3 and 0.1 kmol Na2S04 m-3.

A stock solution of 0.1 kmol V(V) m~3 was prepared by dissolving the appropriate

mass of NaVC>3 (BDH) in the carbonate buffer. The white crystals dissolved slowly

with conventional stirring, but the dissolution rate could be increased by placing the

flask in an ultrasonic bath. V(V) solutions were also prepared by dissolving vanadium

pentoxide (BDH) in dilute sodium hydroxide, according to reaction (5.1):

V2O5 + 3 NaOH —> HV20 73- + H20 + 3 Na^* (5.1)


The colourless stock solutions could be kept for many months without degradation, and

they were diluted with the appropriate buffer solution before use. All solutions were

thoroughly deoxygenated by sparging with White Spot grade nitrogen (BOC) for at

least an hour before any electrochemical investigations were commenced. Identical

results were obtained from vanadium (V) solutions which had been prepared from

sodium vanadate and vanadium pentoxide starting materials.

A stock solution containing 10 mol vanadium (IV) m-3 was prepared by dissolving

0.635 g of blue vanadyl sulphate, VOSO4.6H2O (BDH), in 250 cm3 of oxygen-free

carbonate buffer. Since the predominant V(IV) species are thought to be V ig C ^12"

ions (see section 4.2), 6.7 cm3 of 1 kmol NaOH m-3 solution was added to allow for

the hydroxide ion consumption during reaction (5.2):

I 8 VOSO4 + 48 OH' -> V180 4212- + I 8 SO42- + 24 H20 (5.2)

The resulting dark brown solution was diluted tenfold with an oxygen-free buffer

solution before electrochemical studies were made.

A solution of the complex VS43- was prepared according to the method of Harrison and

Howarth [37]. 10 cm3 of aqueous ammonia "0.880" (BDH) was added to 90 cm3 of

distilled water and the resulting solution was saturated with hydrogen sulphide. 1 cm3 of stock vanadium (V) solution (prepared from V2O5 as detailed above) was then

added; this produced a deep purple solution containing 1 mol VS43- m"3. If this

solution was allowed to contact air it turned orange initially and after a longer time

became colourless, as elemental sulphur was precipitated. A solution for

electrochemical studies, initially containing 0.2 mol VS43- m"3, was prepared by

diluting the above solution five fold with an oxygen-free carbonate buffer.

5.2 Vanadium Voltammetry: Results and Discussion

A voltammogram of a 1 mol V(V) m"3 is shown in Fig. 5.2. A sharp reduction peak

was observed on the negative going scan at 0.3 V vs. SHE. If the potential was

maintained at 0.3 V and a new drop of mercury expelled an oxidation current was seen

to flow for a short time. These peaks were peculiar to the mercury electrode and the

charge under them corresponded to the passage of 1.1 C m"2. The value was

independent of the sweep rate, the stirring rate and the vanadium (V) concentration

(providing it was above 10"3 mol m-3). This is consistent with the process responsible

being the reduction of a monolayer of mercury (I) vanadate. The formation of a

mercury (I) salt with vanadium anions has been reported [135], and used a means of

determining the vanadium concentration in solution. In the concentration range 10"4 to

10"2 mol m"3, less than a monolayer of mercury (I) vanadate is formed at a HMDE

when it is held at an oxidising potential for 60 s [135]. The vanadium concentration

determines the fraction of the surface that is covered, and hence the charge that is

passed reducing this layer in a subsequent cathodic stripping potential scan. In this way

cathodic stripping voltammetry can be used to determine the vanadium concentration in


solution. Calculations show that the close packing of mercury (I) ions results in a

charge density of 2.9 C m-2, so it is likely that the monolayer coverage is determined

by the packing of the larger vanadium (V) ions (e.g. HV2O73-).

Fig. 5.2 Voltammogram of Vanadium (V) in Borate Buffer at pH 9.2

First Scan, commenced at 0.35 V vs. SHE. 50 mV s"1.

Below 0.2 V vs. SHE, no further reduction was observed until a potential of -1.0 V vs.

SHE was reached. This potential is considerably lower than the reversible potential

required to reduce V(V) to V(IV) (-0.1 V vs. SHE), as can be seen from the E^-pH

diagram for the vanadium-water system (Fig. 5.3). This diagram was produced

using the computer program POURB, which was re-written in FORTRAN 77 from a

listing provided by Froning et al [136]. The thermodynamic data, in the form of AGf°

values, were taken from a recent review by Israel and Meites [30]. These values are

shown in the Appendix.

A potential of -1.0 V vs. SHE at pH 9.3 is sufficient to produce vanadium (II) oxide.

The peak current density at -1.2 V was about -2.8 A m -2 (see Fig. 5.2). This

compares a value of -1.4 A m"2, which can be calculated for the peak current during a

reversible one electron transfer (using equation 7.11 and assuming: x> = 0.05 V s"1,

C0 = 1 mol m"3, D0 = 5 x 10' 10 m2 s-1 and r = 3.96 x 10-4 m). The fact that the

observed reduction current was double the reversible one electron value suggests that

the reaction may proceed to form vanadium (HI) oxide (Vj Oj ) or V3O5 (oxidation state

31/3) rather than vanadium (IV) ions.

5 ■

^ -K H2 -10 -0-8 -0-6 -04 -02 6 02Potential /V vs. SHE

HV2O73" + 4 e" + 3 H2O —> V2O3 + 7 OH-

3 HV20 73' + 10 e- + 8 H20 -> 2 V 30 5 + 19 OH'

(5.3)

(5.4)

Vanadium Electrochemistry 8 6

Fig. 5.3 Eh-pH Diagram for the V-H20 System at 298 K.

Activity of V species = 0.01.

However, the formation of solid phases often proceeds at lower current densities than

those predicted by equation (7.11), and it may be that vanadium (II) oxide (VO) is

formed:

HV20 73- + 6 e" + 4 H20 -> 2 VO + 9 OH- (5.5)

If the reduction products were soluble species, they would be dispersed away from the

electrode surface as the solution was stirred and would not be available for re-oxidation

on a subsequent positive-going scan. Therefore, stirring the solution would have the

effect of suppressing the re-oxidation peak at 0.05 V vs. SHE. In fact, stirring did not

suppress this peak, which implied that the reduction product was a solid film which

was adsorbed on to the electrode surface.

The reduction of water to form hydrogen has an extremely high overpotential on

mercury. However, the presence of the vanadium oxide phase facilitated hydrogen

evolution, as can be seen from Fig. 5.2; this is consistent with the known catalytic

activity of vanadium [120].


At higher V(V) concentrations, similar results were obtained. A voltammogram

recorded at a concentration of 10 mol m“3 is shown in Fig. 5.4 (this was taken at

40 °C, but the increased temperature did not substantially affect the voltammogram).

Fig. 5.4 Voltammogram of Vanadium (V) in Carbonate Buffer at HMDE.

1st. Scan, 50 mV s"1, pH 9.3, scan commenced 0.33 V vs. SHE.

[V(V)] = 10 mol m-3, T = 40 °C.

The peak reduction current was shifted to a less negative potential than in Fig. 5.2, to

-0.9 V vs. SHE. A small reduction current was also observed at -0.17 V vs. SHE, a

potential which is accessible using H2S as a reducing agent. The reversible potential

for the V(V)/V(TV) couple according to equation (5.6) at this pH is -0.10 V vs. SHE.

Therefore, it is possible that reduction of vanadium (V) to (IV) may be responsible for

this peak, producing V jg C ^12- or V4092" ions, as shown in equations (5.6) and (5.7).

As yet, no thermodynamic data has been published for the V18O4212" moiety, so it does

not appear on the E^-pH diagram shown in Fig. 5.3.

2 HV20 73' + 4e- + 3H 20 V4O92- + 8 OH“ (5.6)

9 HV20 73" + 18 e- + 12H 20 V180 4212' + 33 OH“ (5.7)

From the stoichiometry of both of the above equations it can be seen that neither

reduction is likely to proceed in a single step, since in both cases a large structural

rearrangement is required. This explains why the reduction current was less than an

order of magnitude lower than that predicted for a reversible one electron transfer.

At gold and platinum electrodes, a smaller potential range was available due to the low

overpotential required for hydrogen evolution. Fig. 5.5 shows a voltammogram

recorded on a gold flag electrode in a carbonate buffer solution (pH 9.3) containing

Vanadium Electrochemistry 8 8

10 mol V(V) m-3. The peaks at 0.4 V vs. SHE on the positive going scan, and 0.3 V

on the negative going scan were due to gold oxide formation and reduction,

respectively. Using vanadium concentrations of 10 mol m~3 and above, a small

reduction current (approximately 1/10 of the magnitude of a reversible one electron

reduction), was observed at a potential of -0.55 V vs. SHE. This current may have

been due to the reduction of vanadium (V) to (IV) in solution. The re-oxidation peak at

0.02 V vs. SHE was observed only when a potential of -0.7 V vs. SHE was exceeded

on the negative going scan, which suggested than a film of reduced vanadium oxide

was again formed at these lower potentials.

Fig. 5.5 Cyclic Voltammogram of Vanadium (V) on a Gold Electrode.

1st. Scan, 50 mV s-1, pH 9.3, scan commenced 0.245 V vs. SHE.

[V(V)j = 10 mol nr3, T = 19 °C.


Direct evidence that reduction of vanadium (V) results in the production of a solid film

was provided by voltammetry using a vitreous carbon disc electrode. A typical cyclic

voltammogram is shown in Fig. 5.6:

Fig. 5.6 Cyclic Voltammogram of V(V) on a Vitreous Carbon Electrode.

1st. Scan, 50 mV s"1, pH 9.3, [V(V)] = 10 mol n r 3, T = 20 °C.

This showed a pattern of vanadium (V) reduction and re-oxidation similar to that

observed in Fig. 5.4 and Fig. 5.5. Large reduction currents were observed only at

highly negative potentials (-1.0 V vs. SHE in Fig 5.6) and there was an extremely

large potential separation between the reduction and re-oxidation peaks. Repeated

potential scans resulted in decreasing current densities. An inspection of the electrode

after such scans revealed that it had become coated with an iridescent layer. Further

experiments showed that the thickness of this layer could be increased by holding the

electrode at a potential of -1.0 V vs. SHE. The vanadium oxide layers appeared blue,

green or purple depending on their thicknesses; this behaviour is characteristic of the

optical interference patterns produced by thin layers, and can only occur when the layer

thickness is at least quarter of the wavelength of the incident light (i.e. > 0.1 |im).


Voltammetry of a solution containing 1 mol V(TV) m~3 on a gold disc electrode revealed

that no reduction currents could be detected above the background currents that were

seen in the carbonate buffer, in the potential range -0.6 V to +0.3 V vs. SHE. Since

the reversible potential for V(V)/(IV) at pH 9.3 is about -0.10 V vs. SHE, based on

equation (5.6), an anodic limit of 0.3 V represents an oxidising overpotential of

400 mV.

The above evidence demonstrates that the reduction of V(V) to V(IV) and the oxidation

of V(IV), are irreversible processes at a variety of electrode surfaces. If the

vanadium (V) phase is HV2073" and the vanadium (IV) phase is V18O4212" (see

section 4.1 and 4.2) then it is not surprising that V(V) reduction is slow, since the

formation of V1g0 4212“ requires a considerable structural rearrangement. The

overpotential for this reduction is so high that the potential has to be lowered to values

where other reduction reactions can occur, producing oxide phases such as V3O5,

V2O3 and VO.

This suggests that the reduction of V(V) by hydrogen sulphide in the Stretford Process

proceeds via a specific chemical interaction between the two species. The thiovanadate

complexes (see section 4.9) are well known and are likely reaction intermediates. The

complex VS43- was prepared, and the cyclic voltammogram of this species was

recorded (Fig. 5.7).

Fig. 5.7 Cyclic Voltammogram of VS43-, HS" on a Gold Disc.

100 mV s '1, pH 9.8, [VS43"] = 0.2 mol n r 3, [HS“] = 0.36 kmol n r 3.


As can be seen from Fig. 5.7, oxidation and reduction currents were observed with a

large peak separation. This voltammogram is very similar to those observed on gold

electrodes in hydrosulphide solution (see section 3.3). The saturation of a solution of

ammonia with hydrogen sulphide results in the production of ammonium

hydrosulphide:

NH3 + H2S NH4+ + HS" (5.8)

The 10 % aqueous ammonia solution, as used in the preparation of the VS43" complex

(section 5.1), contained 1.8 kmol NH3 m'3. When this was saturated with hydrogen

sulphide, a solution containing 1.8 kmol HS' m' 3 was formed. This is about 1000

times greater than the VS43" concentration, and explains why the voltammogram is

essentially that of a hydrosulphide solution. If the hydrosulphide concentration was

lowered, the complex decomposed. Thus, the redox behaviour of VS43' was masked

by the large background currents due to the presence of HS' ions.

5.3 Summary

The reduction of vanadium (V) was found to be irreversible on a variety of electrode

surfaces, and led to the formation of solid oxide films (V3O5, V20 3 and VO) rather

than to V(IV) solution species. Irreversible behaviour is commonly observed when a

large structural rearrangement is necessitated as the reactant is reduced. In the present

case, HV2073" is the probable V(V) species and V1g04212'is the likely V(IV) species;

it is clear that such a rearrangement will be required.

The fact that vanadium (V) is an effective oxidising agent for hydrogen sulphide in the

Stretford Process suggests that there is some specific chemical interaction between

them that facilitates V(V) reduction, such as the formation of thiovanadate complexes.

An attempt was made to investigate the redox chemistry of the thiovanadate ion VS43',

using cyclic voltammetry, but large background currents due to the oxidation of HS'

ions obscured any currents that might have been due to the reduction of VS43' ions.

Anthraquinone Review 92

6. Review of Anthraquinone Redox ChemistryAnthraquinones contain two carbonyl groups on an anthracene backbone. Fig. 6.1

shows the structure and nomenclature of 9,10-anthraquinone.

2

3

Fig. 6.1 9,10-Anthraquinone.

In the following discussion the 9,10- prefix should be assumed.

6.1 Anthraquinone Reduction

Each of the two carbonyl groups in the anthraquinone can be reduced to a hydroxy

group. This reduction can be regarded as electron transfer followed by protonation. If

only one carbonyl group is reduced the product is a semiquinol, if both are reduced a

quinol is produced. Reduction to the quinol is shown in equation (6.1).

0 ^ 0 + 2 H+ + 2 e o CM

(6.1)

This equation can be written in an abbreviated form:

AQ + 2 H+ + 2 e- —̂ AQH2 (6.2)

AQ and AQH2 represent the anthraquinone and anthraquinol respectively. This

reduction could proceed through any one of seven intermediate species. Fig. 6.2

shows the possible reaction pathways:

AQH22+ <-> AQH + <-> AQH+ H+

T le - 1U e- Tvl e-

a q h 2.+ <-> AQH- AQ:

H+ H+

t i e - U e - T i e-H+ H+

a q h 2 <-> AQH- <-> AQ2-

Fig. 6.2 Interm ediates in the Reduction of Anthraquinones [137].


In acidic aqueous solution, Bailey and Ritchie [138] studied the reduction of a variety

quinones, and found that this proceeded to invariably produce the corresponding

quinol. Quershi [1 3 9 ] studied the electrochemical reduction of 18 hydroxy-

anthraquinones derivatives and found a two electron reduction in all cases.

This reduction occurs reversibly at the dropping mercury electrode (DME), and the

polarographic studies up to 1974 were reviewed by Chambers [137]. Since the

reaction occurs reversibly, and the diffusion coefficents of the quinone and quinol

forms are similar, the half wave potentials ( E ^ ) are good approximations to the formal

standard potentials (Eo').

Heyrovski and Kuta [1 4 0 ] noted that the electrochemical reduction of an

anthraquinone is dependent on the stability of the corresponding semiquinones; this

stability can be measured by the value of the semiquinone formation constant Ksq,

defined by:

AQ + AQ2' 2 AQ*“ (6.3)

K[AQ-12

[AQHAQ2-]

Polarographic and voltammetric results are dependent on the kinetics of the above

reaction, as well as the value of its equilibrium constant. Only if the equilibrium is

established rapidly,compared to the time taken to complete a potential scan, will the

value of Ksq affect the polarographic reduction wave. Heyrovski and Kuta [140]

found that if Ksq« 1 a single polarographic wave corresponding to a two electron

reduction was observed; if K sq» 16 then two separate waves were observed,

corresponding to two consecutive one electron reductions. However, when Ksq had a

value between 1 and 16, a single wave was observed with a slope corresponding to

anything between 2/3 and 2 electrons per molecule.

This explains why, in many instances, single reduction waves can be observed,

indicating electron transfer numbers close to one while coulometry always results in a

value close to 2. Exactly this kind of behaviour was observed in the present study

when disodium 2,7-anthraquinone disulphonate (Na2AQ27DS) was reduced. The

reduction of anthraquinones may or may not be accompanied by the uptake of protons,

depending on the first and second acidity constants of the corresponding anthraquinols.

From the pH dependence of the reduction potential the number of protons consumed

can be determined.

Savenko and co-workers [141,142] noted that the polarography of anthraquinone-

1,5-disulphonic acid was affected by adsorption of the oxidised and reduced forms on a

mercury electrode, and this resulted in the inhibition of the reduction reaction.


6.1.1 Substituent Effects

Zuman [143] reviewed the effect of substituent groups on the Ej/2 potentials of

quinone/quinol couples and found that substituent groups increased the reduction

potentials according to equation (6.4):

AE1/2 = C 5 (6.4)

where:

S = Log fKa (Subst. Benzoic acid)1 Ka (Benzoic acid)

AEj/2 = the difference in the half wave potentials between the substituted and

unsubstituted quinones.

C = proportionality constant (for each group of quinones)

The term S , the total polar substituent constant, is based on the ratio of the acidity

constants of substituted and unsubstituted benzoic acids. Electron-withdrawing

substituents on the benzene ring help to delocalise the negative charge on the benzoate

anion, through polar and resonance effects. This same charge stabilisation occurs in

substituted AQH*" and AQ2- anions. Thus, upon substitution, the acid-base equilibria

of the quinol will be altered; this will decrease the concentration of free quinol, and so

increase the reduction potential. Therefore, it would be expected that the nature and

position of the substituent groups on an anthraquinone affect the molecule's ability to

become reduced.

It is believed that the anthraquinone salts act as oxidation catalysts and regenerate the

V(V) species in solution in the Stretford Process. Reduced quinols can be re-oxidised

by bubbling air through a solution containing them. Randell and Phillips [144 ]

investigated the effectiveness of various substituted anthraquinones in catalysing the

oxidation of Stretford solutions, which had been reduced previously by HS- ions.

Their results appear to indicate that at pH 9 many substituted anthraquinones show a

catalytic ability as good as, or superior to, that achieved by AQ27DS.

Fig. 6.3 Anthraquinone 2,7-disuIphonate (AQ27DS).


Randell and Phillip measured the time taken for the oxygen concentration (initially zero)

to rise to 20 % and 80 % of saturation. Using AQ27DS these times were 9 and 16

minutes respectively. Simultaneously, a platinum electrode measured the solution

potential, which rose from -0.188 V to +0.045 V (versus SHE). Two materials in

particular appeared to enable the oxygen content to rise in about half the time taken

using AQ27DS: a mixture of tetra-sodium disulphomethyl AQ2,6 and 2,7-

disulphonamides; and tetrasodium disulphomethyl AQ 1,5-disulphonamide:

0

Na0oS-CHo-N-S0p J J Na c

Fig.6.4 Na4 NN'-disuIphomethylanthraquinone-2,6-disulphonamide.

No explanation of their increased catalytic activity was offered, except to state that the

compounds showed a greater solubility than AQ27DS.

6.1.2 Photo-reduction

The photochemistry of anthraquinones has been widely studied. Anthraquinones have

been suggested as photocatalysts for solar energy storage and for the splitting of water

[145]. In alcoholic solution, photo-reduction occurs and anthraquinols are the main

products, but in water the photolysis becomes more complicated and reduced and

hydroxylated products are formed (mainly a-hydroxy anthraquinone sulphonates).

Moore [146] studied the UV and Raman spectra of the species generated by the

irradiation of 2,6 anthraquinone disulphonate using laser light at 351 nm. Triplet Tj

2,6 anthraquinone disulphonate (3n7t*) was initially produced; in the presence of

reducing agents such as sodium nitrite the triplet state was reduced to form the radical

anion:

hv NaNC>2AQ26DS -» 3AQ26DS -» AQ26DS-' (6.5)

The radical anion was extremely long lived in the absence of oxygen, and showed UV

absorbances at 400 and 510 nm [146,147]; the protonated form only showed the

absorbance at 400 nm.


When oxygen was present, the radical was quickly quenched, a process which resulted

in the production of superoxide ions:

AQ26DS-" + 0 2 -> AQ26DS + 0 2 ~ (6.6)

The rate constant for the above reaction was found to be 9 x 108 M' 1 s_1; calculations

show that a solution containing 1 mol m_3 of AQ26DS-" would completely deoxygenate

an oxygen-saturated aqueous solution in less than 1 s. The superoxide ions that are

produced are powerful oxidising agents and will react with water to produce hydrogen

peroxide [148].

2 0 2 - + H20 -» 0 2 + H 02- + OH- (6-7)

H 02- + H20 -> H20 2 + OH- (6-8)

The AQ26DS •" radical anion can also be attacked by hydroxyl radicals to form

hydroxyl substituted anthraquinones.

Kano and Matsuo [147] showed that the AQ26DS-" radical anions could be stabilised

by adding surfactants that produced micelles; ions that were bound to these micelles

were stable for several weeks, even in aerated solutions.

6.2 Anthraquinones in the Production of Hydrogen Peroxide

The industrial preparation of hydrogen peroxide relies on the hydrogenation of

anthraquinone derivatives in non aqueous media, and the subsequent re-oxidation with

oxygen to produce hydrogen peroxide [149]:

AQ + H2 -» AQH2 (6.9)

catalyst

AQH2 + O2 —> AQ + H2O2 (6.10)

An alkyl substituted anthraquinone (such as 2-ethyl anthraquinone) is used, dissolved

in a solvent of methyl cyclohexyl acetate or alkyl benzene [150]. The reoxidised

solution is contacted with distilled water, and the hydrogen peroxide partitions itself

into the aqueous phase.

Keita and Nadjo [151] showed that analogous reactions could occur in aqueous

solution; they reduced the water-soluble sodium salt of 2,6 anthraquinone disulphonate

(Na2AQ26DS) electrochemically, and then re-oxidised it with oxygen to produce

hydrogen peroxide with 100% efficiency. They suggested that this method could be

used to produce relatively dilute hydrogen peroxide for immediate local use.

Therefore, it is feasible that, on re-oxidation, the reduced AQ27DS in the Stretford

Process solutions generates hydrogen peroxide in-situ. Hydrogen peroxide in

alkaline solution is a powerful oxidising agent which is capable of oxidising sulphide

solutions and re-oxidising vanadium (IV) solutions.

Anthraquinone Electrochemistry 97

7. Redox chemistry of anthraquinone 2,7-disulphonateThe anthraquinone disulphonate (AQDS) used industrially in the Stretford Process

is an isomer mix, but the most active isomers are believed to be the 2,7 and 1,5

disulphonates, of which AQ27DS is the most active. The redox chemistry was

investigated using cyclic voltammetry, at stationary and rotating disc electrodes, and

controlled potential coulometry was conducted. The reduction products and

intermediates were analysed by UV-Visible spectrophotometry and e s r spectroscopy.

7.1 Purification of 2,7 anthraquinone disulphonate

A 15 g sample of the crude di-sodium 2,7 anthraquinone disulphonate (L.B.

Holliday and Co. Ltd., Huddersfield, England) was dissolved in 30 cm3 of distilled

water, and this solution was placed on top of a column packed with alumina. The

column was then eluted with distilled water, and a mobile pale-yellow band collected;

an orange band remained adsorbed at the top of the column.

The AQ27DS was recrystallised from an 80:20 acetoneiwater mix using the following

procedure: 200 cm3 of the above solution was added to 800 cm3 of boiling propanone

(acetone), the solution was reboiled and then filtered hot. It was then cooled in an ice

bath and the pale-yellow crystals were recovered by filtration. The solid was dried

overnight in an oven at 110 °C; a yield of 7.9 g was obtained (52 %).

7.1.1 Analysis of the purified 2,7 anthraquinone disulphonate

It has been shown [152] that liquid chromatography can separate the isomers of

anthraquinone sulphonates. Using equipment at the British Gas London Research

Station, the solid was analysed using High Pressure Liquid Chromatography (HPLC).

A sample of the pure 2,7 isomer was kindly provided by Dr. M. Bruce (Department of

Chemistry, Manchester University), this had been prepared by a regio-selective

synthesis route and was over 99% pure. Using this as a standard, it was shown that

the material which had been purified as above contained 99.1 % AQ27DS.

The 13C NMR spectrum was also recorded, and showed six distinct carbon resonances;

this is consistent with the structure of 2,7 anthraquinone disulphonate. Unpurified

material showed resonances which could be attributed to the presence of other isomers.

Industrial AQDS contains a broad range of isomers. The material that is used currently

in operating Stretford Plants (Elvada) contains only 22 % of the 2,7 isomer; with the

2,6, 1,5, 1,6, 1,7, and 1,8 isomers all present in significant quantities. It is believed

that the 2,7 and 1,5 isomers are the most effective catalysts.


7.2 Experimental: Voltammetry

Solutions of AQ27DS (in the range 1-5 mol m~3) were prepared by dissolving the

appropriate mass of the sodium salt in conducting carbonate buffer (0.059 kmol

Na2C 0 3 m-3, 0.223 Kmol NaHC03 n r 3, 0.10 Kmol Na2S 0 4 n r 3: pH 9.3). All

working solutions were freshly prepared on the day of the experiment and were

nitrogenated for two hours before use with white spot grade nitrogen (BOC pic).

Cyclic voltammograms were recorded using a conventional electrochemical cell design

and a saturated calomel reference electrode (SCE electrode), as shown in Fig. 7.1

below. Either a hanging mercury drop electrode (HMDE), a gold flag or a platinum foil

was used as the working electrode.

Counter Electrode

Fig. 7.1 Electrochemical Cell Design for Voltammetry Experiments.

A Hi-Tek PPR1 waveform generator provided control potentials for the

Thomson MP81 potentiostat, and were also applied to the x-input of a J J PL4 chart

recorder; the currents were passed through a suitable standard resistor and the resulting

voltage was applied directly to the y-input of the chart recorder. At sweep rates above

100 mV s_1 the voltammograms were recorded on a N icolet E x p lo re r I

oscilloscope.


7.3 Experimental: Exhaustive Electrolysis

In order to determine the number of electrons involved in the reduction of AQ27DS and

to provide a supply of the reduced compound, exhaustive electrolysis was performed,

using the apparatus shown in Fig. 7.2. This incorporated a cation exchange

membrane (Nafion 425, D uPont) which prevented the diffusion of the reduced

species to the anode where they would otherwise have been re-oxidised.

Counter electrode

Test solution

Mercury working electrode

Referenceelectrode

Ion exchange membrane

Fig. 7.2 Exhaustive Electrolysis Apparatus.


The AQ27DS solution was reduced at a stirred mercury pool electrode, using a

catholyte containing 3.54 x 10"4 moles dissolved in 35 cm3 of carbonate buffer (pH

9.3). The anolyte contained 1.27 x 10“2 moles of potassium hexacyano iron(II)

solution (K4Fe(CN)6) dissolved in the same buffer; using this solution meant that

hexacyano iron(III) was formed at the platinised titanium mesh anode. Preliminary

experiments had shown that if oxygen was allowed to be evolved at the anode, it could

diffuse through the ion exchange membrane and chemically re-oxidise the reduced

solution in the catholyte compartment. Both compartments were nitrogenated with

Zero Grade nitrogen (BOC pic) prior to electrolysis, and a nitrogen atmosphere was

maintained above the working solution throughout the reduction.

A Luggin probe placed close to the mercury surface was connected to a saturated

calomel electrode (EIL); the liquid film around a closed ground glass joint provided an

electrical connection whilst minimising diffusion from the reference electrode

compartment into the working solution. The working electrode potential was controlled

using a Solartron 1286 Electrochemical Interface, and the current was passed

through an internal resistor, the resulting voltage was then fed to a Hi-Tek DIBS

digital integrator.

7.3.1 Calculations: Exhaustive Electrolysis

If it is assumed that the overpotential for reduction is sufficient to achieve complete

reduction of the test solution, and there are no competing reactions, the integrated

charge would be expected to rise asymptotically with time towards a value of znF C.

(z = number of electrons transferred, n = number of moles of reactant present and

F = Faradays constant).

Furthermore, if the electron transfer at the electrode is extremely rapid the current will

be limited by the mass transport of reactant to the electrode :

it 00 Ct (7.1)where Ct = concentration of reactant at time t

The constant of proportionality will depend on the number of electrons transferred (z),

the electrode area (A), the thickness of the Nemst diffusion layer (8) and the diffusion

coefficient of the reactant (D0) according to the equation:

it = zFD0ACt

5(7.2)


If the system is closed, and of volume V, then the concentration of oxidised species at

any time is given by:

Ct = C0(l - x)

where x is the fractional conversion and will be given by:

x = Charge passed _ f i 3t

Total Charge Required zFC0V

.-. Ct = C0 - l i 2 t - (7.3)zFV

Substituting for Ct in equation 7.2 gives:

it = zFD0AC0 - D0A f i 3t (7-4)

5 V5

Since the initial current, it=o, is given by:

it=o = zFDqACq (7.2)

6

Equation (7.4) becomes:

it = it=o - D0A J i 3t (7.5)

V5

Differentiating with respect to t:

dit = - DqA it

dt V5

Rearranging,

l d i t = - DqA dt

it V5

and integrating gives:

In it = ln it=0 - ^ 0A t 7̂ '5^

5V

Thus a plot of In it vs. t would be expected to be a straight line with an intercept of

In (it=o) and a gradient of -D0A/5V.

7.3.2 Calibration of Exhaustive Electrolysis Apparatus

The apparatus was calibrated using the reduction of potassium hexacyano iron (III),

(potassium ferricyanide), which is known to undergo a reversible one electron

reduction. A potential of -0.2 V vs. SHE (-0.442 V vs. SCE) was applied to the

electrode, which was in contact with a catholyte solution containing 4.9 x 10"4 moles

of K3Fe(CN)6. Care was taken not to allow the hexacyano iron (III) to contact the

mercury cathode before potential control was established, since the solution is capable

of oxidising the mercury surface:

2 Fe(CN)63- + 2 Hg + 2 OH’ -» Hg20 + H20 + 2 Fe(CN)64- (7.6)


Electrolysis was allowed to proceed for 2 x 104 seconds, after which time the

reduction current had fallen from 28 mA to 22 pA. The charge passed after this time

was 45.15 C; a value in reasonable agreement with the theoretical value of 46.98 C

expected for the one electron reduction of the K3Fe(CN)g.

A plot of log it vs. t is shown below:

Fig. 7.3 Plot of Log it vs t during the reduction of Fe(CN)<53_.

The slope of this graph gives a value of -DoA/2.303SV; the solution volume (V) and the

electrode area (A) were measured directly, and a value for D0 of 1.02 x 10“9 m2s_1 was

taken from the literature [153,154], and corrected for the viscosity (r |) of the working

solution according to the Stokes-Einstein relationship (D0 = K^T / 67tqa). These

values enabled the mean thickness of the Nemst diffusion layer (8) to be calculated to

be 5.53 Jim (under the particular stirring conditions employed).


7.4 Voltammetry: Results and Discussion

A typical voltammogram for AQ27DS in aqueous alkaline solution is shown in

Fig. 7.4. The reduction was found to comply with many of the requirements of a

reversible electrode reaction [155]: the peak separation was independent of the voltage

sweep rate, the ratio of cathodic to anodic peak heights was one - independent of the

voltage sweep rate - and the peak reduction current was found to be directly

proportional to the square root of the sweep rate.

Fig. 7.4 Cyclic Voltammogram of AQ27DS.

[AQ27DS] = 1 mol n r 3, Sweep Rate 5 mVs"1, pH 9.3.

The reduction was found to occur at a half wave potential of -0.25 V vs. SHE at

mercury, gold and platinum electrodes. Since the diffusion coefficients of the oxidised

(quinone) and reduced (quinol) forms are likely to be equal, the half wave potential

provides an estimate of the formal standard potential at this pH. This potential is

slightly greater than that required to oxidise HS" to elemental sulphur (E0' = -0.28 V

when [HS-] = 10 mol m“3).

Only a single reduction peak was observed, even though potentials down to -1.4 V

vs. SHE were accessible using a mercury electrode, before hydrogen was produced.

This behaviour suggested that the reduction was proceeding directly to form the quinol

in a two electron step.


The current at a hanging mercury drop electrode can be split into components due to

planar and radial diffusion of the reactant to the electrode surface. It can be shown

[155] that the peak current at a planar electrode is given by:

ipl = 0.4463 zFAC0 ( zF f 2 o 1/2 D 0 1*2 (7.7)RT

where v = voltage sweep rate / Vs"1 and other symbols are as defined previously.

At 20°C in aqueous solution this simplifies to:

ipl = (2.71 x 105) z3/2ACod1/2 Dq1/2 (7.8)

For a spherical electrode of radius r m, an extra term due to radial diffusion must be

added:

isp = ipl zFADqCq (})(Gt) (7.9)r

(j)(at) is a dimensionless constant which is dependent on the applied overpotential. At the peak current obtained during a reversible reduction its value is 0.7516 [155].

Substituting this value into equation (7.9) gives:

isp = ipl + (7-25 x IO ^zADqCq (7.10)r

Subtituting in for ipi*.

isp = (2.71 x IO ^z^ A C qD ^ D o1/2 + (7.25 x 104)zACoJ2o (7.11)r

Equation (7.11) constitutes a quadratic equation in D01/2, which can be solved at a

given peak current and sweep rate providing z, A and r are known. Following the

methods described previously (section 5.1) the drop area and radius were calculated to

be 1.97 x 10"6 m2 and 3.96 x 10"4 m respectively. From the peak current densities

shown in Fig. 7.4, and assuming that z = 2, the value of D0 was found to be

3.74 x 10"10 m2 s"1.

The above calculations assume that the reduction occurs in a two electron process. The

peak separation was found to be about 40 mV, in between that expected for one

electron and two electron processes (59 and 29.5 mV respectively). This may be due to

the two electron quinol product being in equilibrium with the semiquinone, as

discussed in section 6.3.


Richardson and Taube [156] extended the theory first proposed by Polcyn and Shain

[157] which relates the observed peak separation (AEp) to the comproportionation

constant (Kg), where:

[AQ27DS] [AQ27DS2-]

If the first electron is transferred at a standard potential of E^j, and the second at E ^ ,

then they showed that the conproportionation constant was equal to:

Their analysis assumes that both charge transfers are reversible, that the reaction rate is

sufficient to maintain Nemstian concentrations at the electrode surface and that the

reactant diffuses linearly towards the electrode surface. In practice the assumption of

linear diffusion applies well to planar electrodes, and even the hanging mercury drop

electrode used shows only about 10 % deviation due to it being spherical (see equation

(7.11) above).

When (E°i - E02) is greater than 120 mV, two reduction peaks can be seen and the E°

values estimated directly from the half wave potentials; however, in the present case the

peaks are superimposed so that only a single reduction peak can be seen. Richardson

and Taube [156] produced a table of values and a working curve to enable the value of

(Eî - E°2) to be estimated from the peak separation between the negative and positive

going scans (AEp). The observed value of was about 40 mV, which corresponds to

(E°i - E ^ ) = 0. Substituting this value into equation (7.13) gives Kc ~ 1, although

the error in AEp is such that K̂ . could lie in the range 0.2 to 4.

Kc = [AQ27DS-~]2 (7.12)

(7.13)RT


Cyclic voltammetry at a rotated gold disc electrode is shown in Fig. 7.5. Current

crossovers can be seen, which may be due to adsorption of the quinol on the electrode

surface.

Potential vs. SHE / V

Fig. 7.5

Cyclic voltammetry of AQ27DS at a rotated gold disc electrode.

Scan rate = 20 mVs"1. pH = 9.23. C0 = 0.357 mol n r 3.

The currents appear to be mass transport controlled, and would be expected to follow

the Levich equation:

ilim = 1.554 z F A D o ^ c o ^ n '^ C o (7.14)

where co = rotation rate / s_1

o = kinematic viscosity / m2 s-1

and other symbols are as defined previously.

A plot of current vs. (rotation rate)1/2 would therefore be expected to be a straight line

with a gradient of 1.554 zFAD02/3 o"1/6 C0. Experimental results are shown in

Fig. 7.6 (overleaf).


Fig. 7.6 Plot of i vs. co1/2 for reduction of AQ27DS.

The gradient was found to be 2.07 x 10"5 A s1/2. Assuming two electron reduction

enabled a value for the diffusion coefficient (D0) of 3.73 x 10“ *0 m2 s_1 to be

calculated; this is in good agreement with the value reported previously. For

comparison, Compton [158] found a value of 4.7 x 10“10 m2 s_1 for the similarly

sized compound, 1,8-dihydroxyanthraquinone.


The pH dependence of the reduction potential was determined by conducting cyclic

voltammetry at a hanging mercury drop electrode, after the pH had been adjusted by

sparging the solution with carbon dioxide for a short time. The pH was monitored after

each adjustment by withdrawing samples and measuring the pH with a Corning 150

pH meter. The pH could be lowered in this way from 9.3 to 7.1. A plot of the half

wave potential vs. pH is shown in Fig. 7.7.

Fig. 7.7 Plot of Reduction Potential vs. pH for AQ27DS.

The slope of the graph was found to be -31.6 mV pH"1; it follows from the Nernst

Equation that the reduction potential should decrease at a slope given by 59h/z, where

h = the number of protons consumed in the reduction. Since z = 2, h must be equal to

one; i.e. the reduction must proceed in a two electron, one proton process:

AQ27DS + 2e- + H+ AQ27DSH" (7.15)

7.5 Exhaustive Electrolysis: Results and Discussion

Electrolysis was performed at potentials of -0.382 V and -0.6 V vs. SHE; providing

132 and 350 mV overpotential respectively. Assuming Nemstian conditions apply,

132 mV is sufficient to ensure that the equilibrium concentration of the oxidised

AQ27DS is reduced to less than 0.01 % of its initial value. At both these potentials

identical results were obtained; the resulting plot of charge vs. time is shown in

Fig. 7 .8 (overleaf).


Fig. 7.8 Plot of Charge vs. Time During the Electrolysis of AQ27DS.

Electrolysis potential = -0.6 V vs. SHE. pH = 9.3.

As expected, the charge rose to reach a maximum value which corresponded to

complete, two electron reduction; experimental values ranged from 94 % to 107 % of

the theoretical charge. When complete reduction had been achieved, the potential of the

mercury pool electrode could be stepped to 0.0 V vs. SHE and the solution re-oxidised

with the passage of 86 % of the cathodic charge.

The reduction current was found to decay logorithmically, as expected for a mass

transport limited reaction (section 7.3.1). Using a value for the Nemst diffusion layer

of 5.53 pm (section 7.3.2) enabled an estimate of the diffusion coefficient (D0) to be

calculated.

The slope of the Log(i) vs. t plot (Fig. 7.9 overleaf) was found to be -1.91 x 10' 4 (correlation coefficient 0.998), this results in a calculated value of the diffusion

coefficient of 8.9 x 10' 10 m2 s-1 . This is only in moderate agreement with the values

obtained from rotating disc experiments and hanging mercury drop electrode cyclic

voltammetry (3.74 x 10" ̂ ) . However, the stirred mercury pool electrode does not

provide a well defined hydrodynamic regime and diffusion coefficients calculated from

such results are likely to be less accurate.


Fig. 7.9 Plot of Current vs. Time for Electrolysis of AQ27DS.

Potential = -0.6 V vs. SHE. pH = 9.3.


7.6 UV-Visible Spectrophotometry: Experimental

Observation of the electrode surface during cyclic voltammetry revealed that a deep

red-brown colour was produced on the negative going scan; this colour was discharged

on the return scan. This suggested that the production of the quinol could be followed

spectrophotometrically,

An apparatus was assembled to continuously monitor the UV-Visible spectrum of a

Na2AQ27DS solution throughout its reduction, by pumping the solution through a

flow-through UV cell. The apparatus is shown in Fig. 7.10.

Fig. 7.10 Electrolysis with Linked UV-Visible Spectrophotometry.

The anolyte contained 0.5 kmol m-3 potassium hexacyano iron (II) (K4Fe(CN)g)

dissolved in carbonate buffer (pH 9.3) and the catholyte contained about 0.5 mol m-3

Na2AQ27DS dissolved in the same buffer. The two solutions were separated by a

Nafion cation exchange membrane (DuPont) and were nitrogenated with oxygen-free

(CP) grade nitrogen (BOC pic) before the reduction was commenced. Throughout

the experiment a nitrogen atmosphere was maintained in the anolyte, catholyte and

reference compartments.

A titanium working electrode of large surface was prepared by dissolving the oxide

coating from a slotted titanium plate in hot 4 kmol m-3 sulphuric acid. The potential of


this electrode was then maintained at -0.78 V vs. SCE (-0.538 V vs. SHE). The

reversible potential for hydrogen evolution at pH 9.3 was -0.638 V vs. SHE, and a

preliminary voltammogram had shown that hydrogen evolution was not significant until

a potential of -0.658 V were reached.

The potentials were controlled by a Solartron 1286 Electrochemical Interface,

and the charge passed was monitored by a Hi-Tek DIBS digital integrator. Once the

solution had been deoxygenated, 3.24 x 10"4 moles of Na2AQ27DS were added, to

form a solution of concentration 1.392 mol m-3 . When the working potential was

applied, an initial current of -3 mA flowed, which decreased as the AQ27DS was

reduced, and eventually reached a "residual" value of -240 |iA. This was assumed to

be due to oxygen diffusion into the apparatus, causing re-oxidation of the quinol

formed.

Preliminary experiments had shown that diffusion of oxygen through plastic tubing and

peristaltic pumps was a serious problem, and so PTFE-lined stainless steel tubing and

an enclosed diaphram pump were used in the apparatus shown in Fig. 7.10. The

reduced quinol was extremely oxygen-sensitive and the problem of oxygen diffusion

was most apparent when the solution was essentially reduced. Data given by Esco

(rubber) Ltd. showed that silicone rubber has a permeability of 3.16 x 10"4 moles of

0 2 m-2 s"1 (1 mm thickeness, AP accross membrane 1/5 atmosphere). Even the short

lengths used to connect the inflexible metal tubing to the flow-through UV cell could,

in theory, allow enough oxygen diffusion sufficient to sustain currents of -300 |iA. A

correction was made for oxygen diffusion by subtracting the charge passed due to the

experimentally observed "residual" current.

At regular charge intervals the UV-visible spectrum was recorded using a Hewlett

Packard HP8451A diode array spectrophotometer, with the solution flowing

through a Hellma 170.004Q quartz cell (path length 1 mm). The spectra were

referenced against the carbonate buffer solution.

ABSO

RBAN

CE


7 .7 UV-Visible Spectrophotometry: Results and Discussion

The UV-visible spectra taken at approximately 15 C charge intervals are shown in

F ig. 7 . 11 .

Fig.7.11

Spectra taken at 15 C Charge Intervals during AQ27DS Reduction.

Concentration = 1.392 mol m-3. Path Length = 1mm. pH = 9.3

It can be seen that the peak at 330 nm, due to AQ27DS, decreased and two new peaks

appeared as the reduction proceeded: a sharp peak at 410 nm and a broad shoulder at

520 nm. By measuring the absorbances at 330 nm of solutions containing different

concentrations of AQ27DS, it was shown that the Beer-Lambert law (Equation 7.16)

was followed, and so the absorbance at 330 nm could be used to monitor the AQ27DS

reduction.

A = eC 0l (7.16)

A = optical absorbance

8 = extinction coefficient / m2 mol-1 C0 = concentration / mol m-3 / = optical path length / m

The extinction coefficient was calculated from a graph of absorbance (330 nm) vs.

concentration, and was found to be 460.3 m2 mol- 1 . (e=4,603 in non-SI units;

cm-1 mol-1 dm 3).


A plot of the absorbance(330 nm) vs. charge (after correction for oxygen diffusion)

during the reduction of AQ27DS is shown below:

Fig. 7.12 Absorbance(330 nm) vs Charge during AQ27DS Reduction

Concentration = 1.392 mol n r 3. Path Length = 1mm. pH = 9.3

The absorbance at 330 nm decreased linearly with charge passed as the AQ27DS was

reduced (deviations were seen as the reduction neared completion, when the charge

correction due to oxygen diffusion became very important). The charge taken to reduce

the absorbance to half its initial value was 32.3 C, corresponding to 1.04 F mol-1 .

This suggests that a two electron process is required to achieve complete reduction, a

conclusion in agreement with earlier results.

7.7.1 Spectral Assignments

The starting material, AQ27DS, shows absorbances at 330 nm (e = 460 m2 mol-1 ) and

258 nm (s = 4450 m2 mol- 1 ) in the UV spectral region. The reduction product at

pH 9.3 is believed to be AQ27DSH", and so the absorbances at 410 nm (e = 980

m2 mol-1 ) and 277 nm (e = 104 m2 mol-1) are attributed to this species.

Work done by McQuillan [159] has shown that the AQ27DS*' radical anion can be

produced by controlled potential electrolysis of AQ27DS at pH 13.2 in a solution

containing 0.5 kmol m-3 tetraethyl ammonium hydroxide. In such a solution two

reduction waves can be seen in a cyclic voltammogram, corresponding to two

successive one electron reductions:

AQ27DS + e--- > AQ27DS-” (7.17)

AQ27DS-- + e----> AQ27DS2' (7.18)

The tetraethyl ammonium hydroxide stabilises the radical anion, possibly by

incorporating it within a micelle; Kano and Matsuo [147] found that micelles of


sodium laurate or sodium lauryl sulphate could stabilise the radical anion AQ26DS-".

Alternatively, an ion-pairing interraction between the tetraethyl ammonium cation and

the radical anion may be present.

By applying a potential sufficient to achieve only one electron reduction, McQuillan

produced a solution containing AQ27DS-" as the major species, and therefore was able

to measure its UY-Visible spectrum [159]. He found absorbances at 403 and 525 nm.

It is likely that the shoulders at around 390 nm and 520 nm observed in Fig. 7.11 are

due to AQ27DS-".

At more negative potentials, McQuillan produced the di-anion AQ27DS2" which

absorbed at 454 and 540 nm. The pKa of AQ27DSH- is 10.8 [159], and so at pH 9.3

approximately 3 % o f the AQ27DSH- would be present as the di-anion. Therefore the

shoulder observed at 450 nm in Fig. 7.11 is likely to be due to AQ27DS2-, present as

a minor component. As the reduction proceeds, protons are consumed and, despite the

buffering of the solution, the pH is likely to rise slightly. Calculations show that a

change of only 0.3 of a pH unit would double the equilibrium concentration of the di

anion. This explains why the shoulder at 450 nm became more pronounced as the

reduction progressed.

The spectral assignments are summarised in Table 7.1:

Species ^max / nm 8 / m2 mol-1

AQ27DS 330 460

258 4450

AQ27DS-- 525 -660

403 -480

A Q 27D S2" 540 -300

454 -1200

AQ27DSH- 410 -980

277 - 104

Table 7.1 UV-Visible Spectral Summary of AQ27DS Reduction

Products


7.8 E S R Spectroscopy: Experimental

The peak separation from voltammetry (section 7.4) and the UV-visible results

(above) indicated that the radical anion, AQ27DS-", was produced during the

reduction of AQ27DS, even though the major product was AQ27DSH". If the

radical anion were present, then the solution would show a strong esr signal. The

following experiment was designed to detect any esr signal produced during the

reduction of the anthraquinone.

Solutions of 1 and 5 mol AQ27DS n r3 were prepared by dissolving the appropriate

mass of the sodium salt in conducting carbonate buffer (0.059 kmol Na2C03 m-3,

0.223 kmol NaHC03 m~3, 0.1 kmol Na2S04 m“3: pH 9.3). The solutions were

deoxygenated by passing purified nitrogen through the solution. The nitrogen was

purified by passing it through a series of Dressel Bottles containing: reduced

anthraquinone-2-sulphonate (AQ2S) to remove oxygen, water to remove spray,

and finally drying agents to remove traces of moisture. The AQ2S was reduced by

contacting an alkaline solution with a zinc/mercury amalgam.

The deoxygenated solution was loaded into the drive syringe of the electrochemical

esr apparatus that is shown in Fig. 7.13.

Tube cross-section

- to waster Platinum tube

counter electrode

Referenceelectrode

Solution

esr davify

LL___

$&(

Platinum half -cylinder working electrode

Porous plug

Fig. 7.13 Electrochemical E S R Apparatus.


A stepper-motor driven syringe enabled the flow of solution through the apparatus

to be carefully controlled. A potential of -0.358 V vs SHE (-0.6 V vs SCE) was

applied to the platinum half-cylinder working electrode; the potential was controlled

using a laboratory-built potentiostat incorporating a conventional operational

amplifier circuit. The reduced solution flowed into the esr cavity of a B ruker

ER200 TT esr spectrophotometer, through a 1 mm diameter quartz tube, and the

esr spectrum was recorded using a field setting of 3395 Guass and microwave

frequency of 9.54 GHz. The field was swept slowly over a 5 Guass range, using a

0.5 s integrating time constant. The integrated esr signal strength was proportional

to the magnitude of the central resonance; the cavity response was calibrated by

placing in it a MgO crystal containing a known number of spins (200 ppm Mn2+)

and measuring the resulting esr signal.

7.9 Electrochemical E S R : Results and Discussion

The flow-through tube electrode provides a well defined mass transport rate, and

the mass transport limited current at a half-cylinder electrode should be given [158]

by:

iIim = 2.75 zF Xe2/3 D02/3 VfW C0 (7.19)

where Xe = length of electrode / m (see Fig. 7.13)

Vf = volumetric flow rate / m3s_1 and other symbols are as previously defined.

Thus the limiting current should be directly proportional to the third root of the flow

rate; the slope of a graph of iiim vs. Vf1/3 gives an estimate of the diffusion

coefficient (D0), although it is acknowledged that rotating disc results (see section

7.4) are more accurate.

It was found that the above plot was linear, the value of the slope being 8.8 x 10-3 A (m 3 s '1)"1/3; the value of D0 calculated from this was 1.0 x 10-10 m2 s' 1 - somewhat lower than the value of 3.73 x 10_1° m2 s' 1 obtained using the rotating

disc electrode.

Radical species produced at the tubular electrode are carried into the esr cavity, but

the velocity profile across a section of tube is parabolic and not uniform, with the

flow slowest near the walls (see Fig. 7.14).

Fig. 7.14 Flow Profile Across a Tube.


The radicals, produced at the tube edge, diffuse towards the centre of where the

flow is fastest. Thus the number of radicals within the esr cavity is dependent on

the flow velocity profile across the tube and the rate of radial diffusion towards the

centre. This combination has been considered by Compton [158] who found that

the resulting esr signal strength was given by equation (7.20).

S = S0 (tt/4)2/3 j2/3 r2 inm V (7.20)

zFD0l/3 Vf2/3 IK

where S = esr signal strength

S0 = esr signal strength for one mole of spins within cavity.

1 = length of esr cavity / m

r = radius of the tube / m

IK and IK' are sensitivity factors, and are constant for a given geometry.

Thus the normalised signal strength (S/ixim) should be directly proportional to V f2/3.

Such a plot is shown if Fig. 7.15:

Fig. 7.15 Normalised ESR signal (S/ium) vs. Vf~2/3.


As can be seen the resulting plot is not linear, showing a signal enhancement at low

flow rates. This is consistent with a mechanism of radical production as follows:

AQ27DS + 2e" + H+ -> AQ27DSH" (7.21)

AQ27DSH- <-> AQ27DS2' + H+ (7.22)

AQ27DS2" + AQ27DS <-> 2AQ27DS-- (7.23)

The electrochemical reduction results in the formation of the protonated di-anion

(7.21) which exists in equilibrium with the unprotonated form (7.22); this then

reacts with the starting material to produce the radical anion (7.23).

If the flow rate is slow, the reduced material undergoes greater radial diffusion,

towards a region of high AQ27DS concentration, which will shift the equilibrium

(7.23) to the right, and result in an increased esr signal.

7.10 ESR Spectral Strucure

Some difficulty was encountered in obtaining a fully resolved esr spectrum; a high

concentration of AQ27DS-" increased the spin exchange owing to reaction (7.23),

which had the effect of broadening the esr signals, and a low concentration meant

that the spectrum was hard to resolve from the instrumental 'noise'.

Nevertheless, using a concentration of 5 mol AQ27DS m-3 enabled a partially

resolved esr spectrum to be obtained. This was consistent with the radical structure

shown below:

Fig. 7.16 Structure of AQ27DS*-

There are three groups of equivalent protons, labelled and H3. This would

be expected to produce 27 resonances, but spectral overlap and line broadening will

reduce this number.


A spectrum was simulated using the program EPSIM77 assuming the following

values:

Spectral lineshape Lawrentzian

Linewidth 0.19

Splitting Constants 0.21 Gauss (2 protons)

0.55 Gauss (2 protons)

1.025 Gauss (2 protons)

The actual and simulated esr spectra are shown below:

Fig. 7.17 Actual and Simulated ESR Spectra of AQ27DS*"

7.11 Summary

Anthraquinone 2,7 disulphonate (AQ27DS) was reduced in aqueous solution at pH

9.3 to give a deep red coloured air-sensitive solution. Cyclic voltammetry and

exhaustive electrolysis indicated that the anthraquinone was reversibly reduced in a

two electron, one proton process at a variety of electrode surfaces. From limiting

current results at a rotating disc electrode the diffusion coefficient of AQ27DS was

calculated to be 3.73 x 10"10 m2 s"1.


UV-visible spectrophotometry confirmed that AQ27DSH" was the major reduced

species, but also indicated that the di-anion (AQ27DS2-) and radical species

AQ27DS-" were also present. A strong esr signal confirmed the presence of the

radical and the spectral features were consistent with the chemical structure of

AQ27DS-".

The esr signal strength variation with flow rate showed that the radical was not

produced directly at the electrode in a one electron process, but was formed via a

comproportionation reaction between the di-anion and the AQ27DS starting

material. The peak separation from voltammetry enabled the comproportionation

constant (Kc) to be estimated, and it was found to be in the range 0.2 to 4.

Oxygen Reduction 122

8. Oxygen ReductionOxygen and sulphur both lie in group VI of the periodic table and there are certain

similarities in their chemistry: for example they both form compounds in the -II

oxidation state and sulphur can often replace oxygen in its compounds (e.g. SO42" and

S20 32-).

However, oxygen is the first member of the group and many aspects of oxygen

chemistry differ dramatically from those of sulphur The smaller size of oxygen means

that the element is more electronegative than sulphur, and oxygen tends to form bonds

with a high degree of ionic character. Thus, water is a highly polar molecule, which

leads to its hydrogen-bonded liquid structure at room temperature, whilst hydrogen

sulphide is non-polar and exists as a gas.

The fact that oxygen is restricted to eight electrons in its outer shell limits its maximum

co-ordination number to four, and in practice it rarely exceeds two. Whereas sulphur,

because of its d orbitals at available energy levels, is able to form compounds with

higher co-ordination numbers. Commonly four co-ordinate compounds are formed

(e.g. SO42") and co-ordination number of up to six are possible (e.g. SFg). This

expansion of the octet rule allows sulphur to form a range of compounds with a formal

oxidation state greater than II (see section 2.1) whilst oxygen exhibits only the

oxidation states -II, -I and zero.

The bond energy of the oxygen-oxygen double bond is more than three times that of the

single bond (498 and 142 kj mol-1 respectively [160]), whilst for sulphur it is less

than twice (425 and 226 kJ mol-1 [161,57]). This results in a tendency for sulphur,

unlike oxygen, to form catenated molecules (e.g. the polysulphides and polythionates).

Oxygen can exist in two allotropes; ozone (O3) and oxygen (O2). Although oxygen has

a bond order of two, it is paramagnetic and contains two unpaired electrons. This has

been explained by molecular orbital bonding theory, which predicts that two tc*

orbitals should be singly occupied. Because of this, oxygen has been termed a

di-radical and many of the reactions of oxygen proceed via radical mechanisms (e.g.

combustion).

By contrast ozone is diamagnetic. The molecule is bent with a central bond angle of

116 °. Ozone is prepared by the action of a silent electrical discharge in a stream of

oxygen, or by the electrochemical oxidation of sulphuric acid at low temperatures.

Ozone is an endothermic compound (AHf° = 142 kJ m ol'1) and because of this

thermodynamic instability it can decompose to form oxygen. However, this

decomposition is slow in the absence of ultra violet light or transition metal catalysts,

because of the high activation energy that is required. Like many sulphur compounds,

ozone is an example of a metastable compound.


Ozone is a more powerful oxidising agent than oxygen, as can be seen from a

comparison of their standard reduction potentials:

0 3 + 2H+ + 2e- -> 0 2 + H20 E° = 2.07 V (8.1)

0 2 + 4H+ + 4e- -» 2 H 20 E° = 1.23 V (8.2)

In fact few oxidants are as effective in acid solution; ozone is capable of oxidising

sulphide to form sulphate [83].

Oxygen itself should also be a good oxidant. Certainly at high temperatures it is

extremely effective, being reduced to form water or carbon dioxide in the oxidation of

fuels, for instance. However, in aqueous solution at room temperature, oxygen is

often a much less effective oxidising agent than its standard electrode potential would

indicate. There are two reasons for this; the low solubility of oxygen in aqueous

solution and the kinetic inertness of the 02 molecule towards reduction.

Oxygen is only slightly soluble in water; a saturated solution in equilibrium with pure

oxygen at one atmosphere pressure contains only 1.25 mol 0 2 m"3. Like other gases,

the solubility of oxygen decreases as the temperature is increased. This is why

increasing the temperature in an attempt to increase the rate of an oxidation can

sometimes have the reverse effect. However, the diffusion coefficient (which can limit

the transport of oxygen to a reacting surface) increases with temperature, so it follows

that the product of the diffusion coefficient and the solubility must go through a

maximum. This maximum is found at 60 °C in aqueous solution. Gold dissolution in

cyanide solution (8.3) is limited by the transport of oxygen to the metal surface.

4A u + 0 2 + 8 CN- + 2H 20 -> 4Au(CN)2‘ + 4 OH- (8.3)

The fastest rate of gold dissolution is found to occur at 60 °C.

The kinetic inertness of oxygen is due to the high strength of the oxygen-oxygen bond

(498 kJ mol-1 ) and the need for four consecutive electron transfer steps. Oxygen

reduction can sometimes proceed to form hydrogen peroxide (H20 2) rather than water,

since this does not require the cleavage of the oxygen-oxygen bond.


8.1 The Oxygen / W ater Couple

Water represents the lowest oxidation state of oxygen and is produced on the complete

reduction of molecular oxygen. Theoretically, this can be achieved by applying a

potential lower than line (b) in Fig. 8.1:

Fig. 8.1 Eh-pH diagram of the 0 2/H20 System at 298 K [61].

However, complete reduction of gaseous oxygen (8.4) involves a four electron

transfer, and is a highly irreversible process.

0 2 + 4H+ + 4 e “ -> 2 H 20 E°= 1.23 V (8.4)

The standard electrode potential for reaction (8.4) has proved difficult to achieve in

practice; exchange current densities for the reaction on Pt and other noble metals are

typically 10"10 - 10-11 A [62]. Side reactions, which would otherwise be considered

slow, can compete with reaction (8.4) in determining the rest potential. Ordinary

platinum electrodes in pure acid and in the presence of 1 atm. 0 2 usually achieve a

potential of around 1.0 V vs. SHE.


A general scheme for oxygen reduction for oxygen reduction has been reproduced in

reviews by Tarasevich et al. [162] and Schiffrin [163]:

I 1O j <-» ( 0 2 ) Sur ^ (C ^ a d s ^ ( ^ Q ^ a d s “ > H2 0 (8 .5 )

u

(H 2°2)sur H 2°2

In this scheme 0 2 , ( 0 2)sur, and ( 0 2)a(js correspond to molecular oxygen in the bulk

solution, at the electrode surface, and in the adsorbed state, respectively. There exists

two basic reaction pathways; the direct reduction to produce water, or a consecutive

reaction pathway proceeding through hydrogen peroxide (which can also

disproportionate chemically, to produce oxygen and water). At the reversible potential

of equation (8.4), nearly all metal electrodes are covered in an oxide film, the nature of

which is potential and time dependent. These films have a pronounced effect on the

oxygen reduction kinetics, and make the observed reduction waves particularly

complex to interpret. Vesovic et al. [164] recently pointed out that despite years of

intensive study, the mechanism of oxygen reduction at many electrode surfaces has not

yet been established.

Rotating ring-disc electrode studies have been made on the reduction of oxygen, and

were recently reviewed by Tarasevich [162]. He pointed out that the method was

incapable of yielding rate constants for all the reactions in scheme (8.5), but drew the

following qualitative conclusions:

i) On gold, mercury, pyrite (FeS2) and carbon electrodes the consecutive reaction

pathway is operative, hydrogen peroxide is formed as an intermediate and can be

desorbed slowly from the electrode surface.

ii) On platinum, palladium and silver electrodes direct, four-electron reduction

predominates.

A simple explanation of this distinction between the two groups was suggested in the

differing affinities of the materials for molecular oxygen and hydrogen peroxide; the

first group possess a low affinity for these species, which is insufficient to break the

oxygen-oxygen bond.

When oxygen reduction is carried out a rotated electrode, any hydrogen peroxide

intermediate that is formed can be swept away from the electrode surface. This tends

to occur at high rotation rates and low overpotentials. At pH 5, with a slow negative

going potential sweep at a rotated pyrite (FeS2) electrode in an oxygen-saturated

solution, Biegler et al. [165] observed two distinct reduction waves. They attributed

the first wave to reaction (8.6) and the second wave to reaction (8.7).

0 2 + 2 H+ + 2 e" —> H20 2 (8.6)

0 2 + 4 H+ + 4 e- -> 2 H20 (8.7)

Vesovic et al. [164] noted a similar change from a two to a four electron process at a

gold electrode as the overpotential was increased.


In fact, hydrogen peroxide can be produced by the electrochemical reduction of oxygen

at porous carbon electrodes. This has been used as an industrial method of producing

hydrogen peroxide, but has now been largely superceded by methods based on the

hydrogenation and re-oxidation of anthraquinones (see section 6.2).

Behret et al [81] studied the oxygen-reduction activity of transition metal sulphides.

They found that the sulphides of the metals cobalt, iron, and nickel showed the greatest

activity. Since the sulphides of these metals are known to be catalytically active for

sulphide oxidation (section 2.3.3) and oxygen reduction, it is likely that their known

catalytic action on sulphide oxidation proceeds via a coupled electrochemical

mechanism.

The fact that the four electron reduction of oxygen requires a high overpotential at most

electrode surfaces has important technological implications. Only a limited number of

metals are suitable for use in the construction of metal-air batteries, and efficient

oxygen-consuming cathodes in fuel cell systems have remained reliant on expensive

catalysts. As a consequence, the great potential of fuel cells for the efficient conversion

of fuels to electricity remains largely untapped. Oxygen reduction has not yet replaced

hydrogen evolution in electrolytic processes, such as those employed in the chlor-alkali

industry, despite a possible saving of 0.8 V in the cell voltage if this practice could be

adopted.

8.1.1 The Evolution of Oxygen

The oxidation of water to produce oxygen is also irreversible: an overpotential of 0.6 to

0.8 V is required to produce oxygen from acid solution at a current density of

20 0 A m“ 2 on platinum or platinum / iridium anodes [1 6 6 ]. Even higher

overpotentials are developed on the lead oxide anodes that are used industrially for

electrowinning metals, and this voltage loss represents a considerable waste of electrical

energy. On the other hand, the high overpotential is advantageous in aqueous batteries

with positive electrodes which develop potentials greater than the reversible potential

for oxygen evolution. The spontaneous generation of oxygen would rapidly discharge

such batteries if the overpotential were low.

According to Tarasevich [162], the origin of the kinetic hindrance towards oxygen

evolution lies in the nature of the electrode surface at the high potentials required.

Relatively few metals are resistant to corrosion, and in those that are, this can usually

be attributed to the formation of a passivating layer of metal oxide. Thick surface layers

form a barrier towards electron transfers (although electrons are believed to be able to

tunnel through the thin surface layer that forms on a platinum electrode in acidic

solution). Despite years of research, the mechanism of oxygen evolution, even on

platinum (the most intensively studied material), remains speculative.


8.2 Hydrogen Peroxide

Hydrogen peroxide is industrially produced by the reduction of an anthraquinone

derivative with hydrogen, followed by re-oxidation with oxygen (see Section 6.2). It

can also be prepared by the electrochemical oxidation of water, via a peroxodisulphuric

acid intermediate. Sulphuric acid is oxidised at low temperatures and using a platinum

anode; under these conditions it is oxidised to form peroxydisulphuric acid:

2 H 2S 04 -> H2S20 8 + 2H+ + 2e- ( 8.8)

This can be hydrolysed to form hydrogen peroxide (8.9), which can be removed by

vacuum distillation.

H2S20 8 + 2 H20 -> H20 2 + 2 H2S 04 (8.9)

Thus the overall reaction involves the oxidation of water:

2 H20 -> + 2H+ + 2e- (8.10)

Hydrogen peroxide can also be prepared by the electrochemical reduction of oxygen

(reaction 8.6) as was discussed above.

Hydrogen peroxide is also evolved when the solid peroxide salts react with water.

These salts contain the 0 22' ion and are prepared by reacting group I or II metals with

oxygen. Hydrogen peroxide is used as the starting material to form a range of organic

peroxides, which are useful as oxidants or a source of free radicals (e.g. benzoyl

peroxide).

Hydrogen peroxide has the structure shown in Fig. 8.2:

94°

Hx?97° 1.49 A

< Q " “

H

Fig. 8.2 The Structure of Hydrogen Peroxide.

Pure hydrogen peroxide is a pale-blue viscous liquid, which possesses a structure

containing a network of three-dimensional hydrogen bonds. However, it is not used as

a solvent because of its oxidising nature and its ready decomposition.

Hydrogen peroxide is another example of a metastable compound; it is

thermodynamically unstable, yet solutions can be stored for months without

decomposition, because of the large activation energy required. If it is exposed to light

(which can provide the large activation energy) or traces of transition metals (which

provide a lower activation energy mechanism) then it decomposes rapidly.


An Efo-pH diagram for the hydrogen peroxide / water system is shown in Fig. 8.3:

Fig. 8.3 Eh-pH Diagram for the H20 2 / H20 System at 298 K [61].

Below lines (2) and (3) hydrogen peroxide can be reduced to form water:

H20 2 + 2 H+ + 2 e" -> 2 H 20 (8.11)

Above lines (4) and (5) it can be oxidised to form oxygen:

H20 2 -> 0 2 + 2H+ + 2e- (8.12)

Between these family of lines hydrogen peroxide is doubly unstable and can

decompose according to reaction (8.13):

2H 20 2 -> 2 H 20 + 0 2 (8.13)

Thus, if hydrogen peroxide contacts a metal surface having an electrode potential within

this region, it will spontaneously decompose; this is an example of the electrochemical

catalysis of a chemical reaction.

Since hydrogen peroxide can be reduced in the region below lines (2) and (3) in

Fig. 8.3, it follows that it will act as an oxidant towards redox couples which have

their solution potentials in this region. In this way hydrogen peroxide acts as a

moderately powerful oxidising agent, both in acid and alkaline solution.


Conversely, towards redox couples having their potentials above lines (4) and (5),

hydrogen peroxide can act as a reductant; manganate (VII) solutions can be reduced to

manganese (II), for example. Because of the slope of these lines with pH, it is possible

for hydrogen peroxide to act as an oxidant at a low pH and a reductant at a higher pH.

The reduction of hydrogen peroxide is catalysed by the presence of transition metal ions

in solution. Mo(VI), for example, forms a complex with hydrogen peroxide [167]:

Mo0 42' + H20 2 Mo0 52' + H20 (8.14)

This complex is more readily reduced than hydrogen peroxide:

Mo0 52- + 2H+ + 2e- Mo042- + H20 (8.15)

Many of the reactions of hydrogen peroxide can also proceed via free radical

mechanisms.

8.3 Superoxides

The action of oxygen on potassium, rubidium and cesium gives rise to yellow

crystalline solids of the formula MO2. They contain the superoxide ion (C>2~) which is

an extremely powerful oxidising agent. In aqueous solution the superoxide ion will

react with water to form hydrogen peroxide:

2 02" ■+■ 2 H20 —̂ 0 2 + H20 2 (8.16)

The superoxide ion is formed as the first intermediate during oxygen reduction, and its

production in an adsorbed form was thought to be the rate-limiting step for oxygen

reduction at a gold electrode in acidic solution [168].

8.4 Experimental

A carbonate buffer of pH 9.3, containing 0.059 kmol Na2CC>3 irf3, 0.223 kmol

NaHC03 m-3 and 0.1 kmol Na2S04 n r 3, was prepared by dissolving the appropriate

masses of analytical grade materials (BDH) in triply distilled water. This buffer

solution was saturated with oxygen by sparging with pure oxygen (BOC) for one hour

before the commencement of electrochemical measurements. According to the Kent

oxygen meter handbook [169], such a saturated aqueous solution at 20 °C will contain

1.35 x 10-6 mol 0 2 m-3. Electrochemical measurements were made in a glass, three

compartment electrochemical cell of conventional design (see Fig. 7.1).

A Hi-Tek PPR1 waveform generator provided the control potentials for the

potentiostat, which was built in Imperial College using a conventional operational

amplifier circuit design. A gold or platinum rotating disc electrode (see section 3.2)

was used as the working electrode, a bright platinum flag as the counter electrode and a

saturated calomel electrode (EIL) as the reference electrode. All potentials are quoted

relative to the standard hydrogen electrode (SHE), assuming that the potential of the

saturated calomel electrode was 0.242 V vs. SHE. The working electrodes were spun

using a motor unit (Oxford Electrodes) which allowed the rotation speed to be

continuously varied up to 50 Hz. The current flowing at the working electrode was


passed through an internal resistor in the potentiostat and the resulting voltage was

applied to the y-plates of a Nicolet 5091 storage oscilloscope. The potentiostat

control voltage was applied to the x-plates which enabled a voltammogram to be

recorded on the oscilloscope. Permanent copy was obtained on a Gould 60000 x-y

plotter, by connecting this to the plotter output terminals of the oscilloscope.

The gold and platinum electrodes were pretreated by potential cycling from -0.8 V to

+1.2 V vs SHE at 10 V s-1 (see section 3.2.3). Slow potential scan voltammograms

were then recorded, starting from either the positive or negative potential limit, at a

number of different rotation rates.

8.5 Oxygen Reduction: Results and Discussion

Cyclic voltammograms showing oxygen reduction at gold and platinum electrodes are

shown in Fig. 8.4:

Fig. 8.4 Cyclic Voltammograms Showing Oxygen Reduction

Gold and Platinum Rotating Disc Electrodes, co = 9 Hz. pH 9.3.

[ 0 2] = 1.35 mol n r 3. Scan rates, Pt = 10 mV s_1, Au = 1 mV s"l.


The reversible potential for the O2 / H2O couple at pH 9.3 is 0.681 V vs. SHE.

However, no reduction currents flowed at platinum or gold electrodes until the potential

was reduced to below 0.4 V and 0.2 V vs. SHE respectively. This demonstrates that

the direct reduction of oxygen to water is a highly irreversible process. The reversible

potential for the O2 / H2O2 couple (in equilibrium with 1 mol H202 m-3) at this pH is

0.221 V vs. SHE. Thus, it can be seen that oxygen reduction at a gold electrode does

not commence until this potential is reached.

At neither electrode surface was a clear, diffusion limited current plateau seen before

hydrogen evolution commenced. Both electrode surfaces showed some degree of

hysterisis, but this was most noticeable in the case of gold. This behaviour suggested

that the gold surface was deactivated on the positive-going scan, at a potential of

-0.35 V vs SHE. Hoare [168] noticed a similar deactivation after the first scan and

suggested that it was due to a change in the electrode surface owing to the activity of a

Au-0 layer, which changes with time. The same author [170] also noted that the

presence of platinum sites on a gold surface can alter the electrocatalytic activity of a

gold surface.

It is apparent from Fig. 8.4 that oxygen reduction requires a substantially lower

overpotential on a platinum electrode, but the current still does not show a perfectly

defined diffusion limited plateau (cf. to Chapter 7, Fig. 7.5). Nevertheless, the

reduction current at a potential of -0.5 V vs. SHE did show an approximately linear

dependence on the square root of the rotation rate, (co)1/2. The reduction currents that

would be expected from the Levich equation (7.14) , assuming that the diffusion

coefficent D0(02) = 1.8 x 10"9 m2 s-1 [171], were also calculated. The experimental

results (on platinum), together with the theoretical two and four electron reduction

currents, are shown in Fig. 8.5 (overleaf).


□ i / mA ♦ 2e n 4e

Fig. 8.5 Experimental and Calculated O2 Reduction Currents at a RDE.

Pt RDE area = 3.85 x 10'5 m2, [ 0 2] = 1.35 mol m'3. T = 293 K.

As can be seen from the above figure, the experimentally observed values fall in

between those expected for two and four electron reductions. At very low rotation rates

the current was close to that expected for a four electron process; as the rotation rate

was increased the current tended towards the two electron limit. This was consistent

with the suggestion that hydrogen peroxide is produced as a metastable intermediate.

At a low rotation speed H2O2 remains on or close to the platinum surface and can be

further reduced to water, whereas at high speed more H2O2 is dispersed into solution.

8.6 Summary

Oxygen reduction was shown to be a slow reaction at gold and platinum electrodes.

Platinum was a more effective electrocatalyst for oxygen reduction than gold; however,

it did not show reduction currents large enough to be attributed to the complete four

electron reduction of oxygen to water. The reduction currents that were observed were

intermediate in magnitude between those expected for two and four electron processes.

This behaviour suggested that hydrogen peroxide was formed as a metastable

intermediate in a two electron reduction, and was then reduced further to form water.

This conclusion that hydrogen peroxide is an important intermediate is in agreement

with the results of previous workers; it is due to the high strength of the oxygen-

oxygen bond. A direct four electron reduction would involve the cleavage of this bond

at an early stage during reduction, whilst a two electron reduction leaves the bond

intact.

Stretford Process Chemistry 133

9. The Redox Chemistry of the Stretford Process The Stretford Process achieves the oxidation of hydrogen sulphide to elemental

sulphur. The gas containing the hydrogen sulphide is contacted with an alkaline

solution containing vandadium (V) salts and anthraquinone disulphonates; the hydrogen

sulphide dissolves and deprotonates in the alkaline solution and reacts with the two

oxidising agents. The reduced solution is then passed to an oxidising vessel, where air

is passed through the process solution. This serves to re-oxidise the solution and to

recover the sulphur produced; sulphur is naturally hydrophobic and concentrates in the

froth at the liquid surface, where it can be skimmed off and filtered. The oxidised

solution is recycled to the gas absorber where it contacts more hydrogen sulphide. A

more complete description of the Stretford Process is given in section (1.2).

In the process there are four linked redox couples; S(-II)/S(0), V(V)/V(IV),

anthraquinone/anthraquinol and 0 2/H20 . In the preceding chapters these redox

couples have been investigated separately using electrochemical techniques. This

section is concerned with the interaction between the redox couples, in order to

determine the reaction mechanism that occurs in the Stretford Process.

A variety of techniques have been applied to study the chemical reactions involved:

i) Stopped flow spectrophotometry has been used to follow the course of reactions

that involve species which absorb in the UV-visible region of the spectrum.

ii) The solution potential has been measured by the use of a suitable indicator

electrode, in order to determine the extent of reduction that has occurred.

iii) Small scale batch experiments have been conducted and the reaction products have

been identified using 51V NMR spectroscopy, cyclic voltammetry and conventional

chemical analyses.

9.1 Experimental

A carbonate buffer solution of pH 9.3, containing 0.059 kmol Na2CC>3 m-3, 0.223

kmol NaHC03 m-3 and 0.10 kmol Na2SC>4 was prepared by dissolving the

appropriate masses of analytical grade materials (BDH) in triply distilled water.

Similarly a borate buffer, having a pH of 9.2, was made up containing 12.5 mol

Na2B4O7.10H2O m-3, 0.9 mol NaOH m' 3 and 0.1 kmol Na2SC>4 m-3. A stock

solution containing 0.1 kmol HS" m"3, was prepared by dissolving an

accurately weighed amount (about 12 g) of transparent, dried crystals of A nalar

sodium sulphide (BDH) in 500 cm3 the appropriate deoxygenated buffer solution.

The molarity of this stock solution was checked by conducting an iodate titration as

detailed in section (3.2.1), and it was diluted with the appropriate volume of oxygen-

free buffer before use.


Polysulphide solutions were prepared either by dissolving the appropriate mass of

Na2S4 in an oxygen-free buffer, or by dissolving elemental sulphur in a sodium

sulphide solution. Stock solutions were made up containing 0.1 kmol S m"3 and were

diluted for use with oxygen-free buffer solution.

Vanadium (V) solutions, containing 0.1 kmol V(V) m-3, were prepared by dissolving

NaV03 or V2O5 (BDH) in a carbonate buffer solution or a dilute sodium hydroxide

solution respectively. The colourless stock solutions could be kept for many months

without degradation, and they were diluted with the appropriate buffer solution before

use.

Vanadium (IV) solutions containing 10 mol vanadium (IV) m"3 were prepared by

dissolving 0.635 g of blue vanadyl sulphate, VOSO4.6H2O (BDH), in 250 cm3 of

oxygen-free carbonate buffer. 6.7 cm3 of 1 kmol NaOH n r 3 solution were added to

allow for the hydroxide ion consumption during reaction (9.1):

I 8 VOSO4 + 48 OH- -» V180 4212- + I8 SO42- + 24 H20 (9.1)

The resulting solutions were dark brown, but became green and eventually colourless if

they were exposed to the atmosphere.

9.1.1 Stopped Flow Apparatus

A diagram of the Stopped flow apparatus is shown in Fig. 9.1:

Fig. 9.1 Stopped Flow Apparatus.


A Hi-Tech SFA-11 stopped flow attachement was modified for use with oxygen

sensitive solutions by using glass syringes and PTFE-lined stainless steel tubing

throughout. The attached quartz cell had four optical faces, so that it could be used

with a 2 or 10 mm path length. The two reservoir syringes were filled with the

reactants and the optical cell was placed in a Hewlett Packard 8451A diode array

spectrophotometer. The two reactants were loaded into the drive syringes from the

reservoir syringes. Then, when the syringe pistons were simultaneously depressed by

the drive plate, they were mixed within the optical cell.

The diode array spectrophotometer was capable of recording a full UV-visible spectrum

in 100 ms, and such spectra were recorded after fixed time intervals following the

mixing of the reagents. The spectra were stored (as digital data) in the memory of the

machine and could be transferred onto a magnetic disc for permanent storage. Using

the Hewlett Packard program KINETICM , the spectra could be recalled and the

absorbance values extracted at a fixed wavelength.

9.1.2 Experimental: M easurement of Solution Potential

The solution potential was measured throughout the reaction between 2,7,

anthraquinone disulphonate (AQ27DS) and sodium sulphide solution using a gold bead

electrode. A diagram of the apparatus used is shown in Fig. 9.2:

seal electrodeSolution

flow

Fig. 9.2 Gold Indicator Electrode for Measuring the Solution Potential


The potentials was measured relative to the saturated calomel electrode (EIL ) and

converted to the SHE scale assuming that the potential of the latter was 0.242 V vs.

SHE. The indicator electrode assembly was connected to the stopped flow apparatus in

place of the stop syringe, so that the chamber containing the indicator electrode was

flushed with the reaction mixture at the same time as the optical cell was filled. Care

was taken to completely expel all the air bubbles at this stage. The reaction mixture

initially contained 0.16 mol AQ27DS m~3 and 50 mol Na2S n r 3 in deoxygenated

carbonate buffer. The potential between the gold bead and the reference electrodes was

measured at 600 s intervals, as the reaction between AQ27DS and HS“ proceeded.

9.1.3 Experimental: Preparation of Samples for 5 iV NMR

Four samples were prepared for 51 v Nuclear Magnetic Resonance (NMR) spectroscopy

containing:

1 . 100 mol NaV03 m~3 bi carbonate buffer.

2 . 10 mol NaV03 m-3 in carbonate buffer.

3. 500 mol NaV03 m"3 and 500 mol Na2S m"3 in carbonate buffer.

4. 5 mol NaV03 m"3 and 500 mol Na2S m-3 in carbonate buffer.

All the samples were thoroughly deoxygenated before they were mixed, and were

sealed into glass NMR sample tubes under a nitrogen atmosphere. Samples one and

two were colourless, sample four became yellow as the reagents were mixed and

sample three became a dark-brown colour and a black, hydrophobic solid precipitated

from the solution.

Spectra were obtained using a Bruker 200 NMR spectrometer, using an exciting

radiation frequency of 52.6 MHz. Liquid VOCI3 was used as a reference material and

all chemical shifs are quoted relative to it.


9.2 Reaction Between AQ27DS and HS": Stopped Flow Results

When equal volumes of solutions containing 0.32 mol AQ27DS m-3 and 100 mol

Na2S m-3 were mixed in the stopped flow apparatus, the series of spectra shown in

Fig. 9.3 were obtained:

Fig. 9.3 Spectra Taken at 600 s Intervals During Reaction between

AQ27DS and H S \ [AQ27DS]0 = 0.16, [HS"]0 = 50 mol n r 3.

T = 17 °C. Cell Path Length = 1 cm.

There above spectra are very similar to those shown in Fig. 7.11, which were

obtained during the electrochemical reduction of AQ27DS. This suggests that the

reduction product (which has an absorbance peak at 410 nm) was the same in both

cases. In chapter 7, evidence was presented that suggested that this reduction product

was AQ27DSH". There was no visible deposition of elemental sulphur during the

reaction, so it is likely that the HS" ions are oxidised to form polysulphide ions (e.g

S42-):

3 AQ27DS + 4 HS- + OH" 3 AQ27DSH" + S42' + H20 (9.2)


Polysulphides ions also absorb in the UV-visible region, and the spectrum of a

polysulphide solution (prepared by dissolving Na2S4 in a carbonate buffer) is shown in

Fig. 9.4:

VAVEIENGTH (rut)

Fig. 9.4 UV-visible Spectrum of Sodium Polysulphide. pH 9.3.

However, complete reduction of all the AQ27DS, according to equation (9.2), would

only produce a S42' concentration of 0.053 mol m'3. Since emax at 380 nm is 112.5

m2 mol-1, the increase in absorbance at this wavelength due to the production of the

polysulphide ions would amount to only 0.06 absorbance units. This is less than

10 % of the optical absorbance at 380 nm due to AQ27DSH". Thus, the presence of

polysulphides would be expected to produce a shoulder at around 380 nm on the

absorbance peak at 410 nm. An inspection of Fig. 9.3 shows that such a shoulder is

present.

9.2.1. Reaction of AQ27DS and HS“: Rate Studies.

Since there was a large excess (approximately 200 fold) of HS~ ions over the AQ27DS

in the above experiment, the [HS'] was assumed to remain constant throughout the

reaction. This enabled the rate order with respect to AQ27DS to be calculated; a zero-

order reaction would cause a linear decrease in [AQ27DS] with time, a first-order

reaction would produce a logarithmic decrease with time, and a second-order rate

would produce a linear decrease of [AQ27DS]"1 with time.


In fact, a logarithmic decrease in [AQ27DS] (as monitored by its absorbance at 330 nm)

with time was observed, demonstrating that the reaction was first-order with respect to

AQ27DS. The plot of In (Abs 330 nm) against time is shown in Fig. 9.5:

Oi

-1

Ln(Abs330 nm)

-2

'30 1000 2000 3000 WOO 5000 6000 7000 8000

Time / s

Fig. 9.5 Plot of ln(Abs 330 nm) vs. Time During Reduction of AQ27DS

[AQ27DS]0 = 0.16, [HS‘]0 = 50 mol m-3. T = 17 °C. 1 = 1 cm.

It follows from themathematicsof first-order kinetics, that the concentration of reactant

R (in this case AQ27DS), remaining after time t will be given by:

In (R) = In (R0) - kt (9.3)

where R0 = the initial concentration of reactant

The concentration of AQ27DS is related to the absorbance at 330 nm (A) through the

Beer-Lambert law (9.4):

A = e R l (9.4)

e = extinction coefficient / m2 mol"1; 1 = path length / m

By substituting (9.4) into (9.5) it follows that:

In (A) = In (A0) - kt (9.6)

Therefore, the first-order rate constant (k) is given by the slope of Fig. 9.5, which

was found to be 2.53 x 10"4 s"1.

This value means that at 17 °C, in the presence of 50 mol HS' m-3, half the

anthraquinone would be reduced after 45 minutes (i.e. t1/2 = 45 mins.). Assuming that

the rate of reaction doubles for each 10 °C rise in temperature, means that at 40 °C (at

which the Stretford Process operates) t1/2 will be reduced by a factor of four.


Nevertheless, in order to achieve 75 % AQ27DS reduction, a residence time of 23

minutes would still be required. Early Stretford Plant liquors contained only

anthraquinone disulphonates; these plants were characterised by long residence times

(in the absorber and reactor vessels) and low H2S throughputs.

9.2.2. Reaction of AQ27DS and HS-: Solution Potential M easurements

Placing an inert metal indicator electrode in a solution containing an oxidising agent and

a reducing agent which are reacting chemically, enables the extent of reaction to be

monitored; as the reduction proceeds, the potential decreases. Indicator electrodes can

be used in industrial processes; for instance, they can be used to follow the extent of

oxidation during the oxidative leaching of uranium ores.

When there are two redox couples present in non-equilibrium conditions, the observed

solution potential will lie between the reversible potentials that each couple would attain

separately (given the concentrations of its oxidised and reduced forms). However, this

value will lie closer to potential of the redox couple which shows the most reversible

behaviour at the electrode surface. This situation is summarised in the "Evans

Diagram" shown in Fig. 9.6:

Fig. 9.6 Evans Diagram Showing Anodic and Cathodic Polarisation

Curves During the Establishment of a Mixed Potential.


If separate polarisation curves were drawn for the electrochemical reduction of

AQ27DS and oxidation of HS" they would appear as shown in Fig. 9.6, with the

reduction reaction polarising cathodically and the oxidation reaction polarising

anodically. When the two couples are allowed to react chemically, the reaction at the

indicator electrode surface gives rise to a "short circuited" reaction current, which is

dependent on the reaction rate. At this particular value of the current, the two potentials

are both equal to the mixed potential. The oxidation of sulphide at a gold electrode

has been shown to be highly irreversible (see Fig. 3.6) and so the anodic polarisation

curve rises rapidly. Conversely, the reduction of AQ27DS at a gold surface was

reversible (see section 7.4), and the cathodic polarisation curve falls gently. Therefore,

the observed mixed potential of a gold bead electrode in a reacting mixture of AQ27DS

and HS‘ ions will lie close to the equilibrium potential of the AQ27DS / AQ27DSH"

couple; its value will depend on the relative concentrations of the quinone and quinol.

If it is assumed that only the AQ27DS and AQ27DSH- concentrations determine the

potential of the gold indicator electrode, that this potential is attained rapidly compared

to the rate of change of concentrations and that the reduction is a first order process,

then the solution potential will be given by a modified form of the Nemst equation:

g _ Eo _ RT { In ([AQ27DS]0 - exp(ln[AQ27DS]0 - k t )} (9.7)

zF [AQ27DS]0

Since the value of k has been determined in section (9.2.1) and values of E° = -0.273 V

vs. SHE and z = 2 can be estimated from section (7.4), the variation of the potential

with time can be theoretically predicted. The calculated and experimentally observed

potential measurements are shown in Fig. 9.7:

Fig. 9.7 Measured and Theoretical Solution Potentials vs Time.

Solution Conditions as in Fig. 9.3.


From Fig. 9.7 it can seen that there is a reasonable agreement between the theoretical

and experimentally observed values. The potential fell as the reduction of the AQ27DS

proceeded, and this decrease in potential as the reaction progressed was of the correct

magnitude as that predicted when z = 2 in equation (9.7). If z were to equal 1, a drop

in potential twice that observed would be predicted; therefore, these measurements

provide evidence that AQ27DS is reduced in a two electron process.

However, the agreement between theory and experiment is not sufficiently close to

allow the potential measurements to be used to predict the concentrations of the reduced

and oxidised forms, nor to determine the first-order rate constant. The discrepencies

are largest at the start of the reaction, when the assumption that the potential is attained

rapidly compared to the rate of change of [AQ27DS] is most suspect.

9.3 Reaction between V(V) and HS_: Stopped Flow Results.

The application of UV-visible spectrophotometry to the study of the reactions of

vanadium (V) was limited by the high optical absorbance of V(V) solutions in the UV

region; a 10 mm cuvette containing a solution of 10 mol NaVC>3 m~3 showed complete

light absorbance below 370 nm (dilution showed that ^ max = 270 nm , £ = 320 m2 mol-1). For this reason, it was not possible to study the reactions of concentrated

vanadium (V) solutions, such as those used in the Stretford Process ([V(V)] = 32

mol m"3), using the Hi-Tech SFA-11 stopped flow apparatus. However, the optical

path length (2 mm) was short enough to allow a study of the reaction between a

solution containing 0.5 mol V(V) m-3 and excess HS' (44 mol m~3).


Spectra were taken at two second intervals during the reaction between the two above

solutions, and the resulting series is shown in Fig. 9.8:

Wavelength / nm

Fig. 9.8 Spectra Taken at 2 s Intervals During Reaction between

V(V) and HS“. [V(V)]0 = 0.5, [HS-]0 = 44 mol n r 3.

T = 17 °C. Cell Path Length = 2 mm.

The reaction was extremely rapid, and an absorbance peak at 360 nm appeared within

the first two seconds. An attempt to decrease the rate of reaction by reducing the [HS~]

by a factor of ten succeeded only in decreasing the magnitude of the absorbance

maxima. This behaviour suggested that an equilibrium was rapidly established between

the HV2O73' ions and the HS" ions. An examination of the spectral properties of the

known thiovanadate complexes (see Table 4.3) revealed that the complex ion

V02S23- possessed an absorbance maxima at 360 nm. This suggested that an

equilibrium involving this ion may have been established:

HV20 73- + 4HS- 2 V 02S23" + 2 H 20 + OH" (9.8)

After the rapid formation of the absorbance maxima at 360 nm, there was a small

increase in the optical absorbance in the spectral region 300-380 nm with time.

Polysulphide species absorb in this spectral range, and they are likely to have been

responsible. Indeed, separate experiments had shown that such an increase could also

be seen when the HS" solution was allowed to react with an aerated buffer.


Vanadium (IV) solutions did not show an absorbance maximum at 360 nm. Instead,

these brown solutions showed a broad absorbance throughout the whole UV-visible

spectral range. This complex and intense spectrum is consistent with vanadium (IV)

existing as the complex polyanion V jg C ^12-. This spectral pattern was not observed

in the above experiment, which shows that V ^g C ^12- was not formed under these

conditions.

The 51V NMR spectra of 10 mol V(V) m-3 showed peaks at -547, -562 and -573 ppm

vs. VOCI3; these were attributed to the species H2V2072-, HV2O73" and VgC^3-

respectively. Upon reaction with the sulphide solution, all these peaks disappeared,

which is consistent with the formation of thio complexes. However, no peak at

184 ppm (attributed by Howarth [37] to V02S23- complex) could be detected,

although it was not clear whether the detection limits of the machine would be exceeded

at this relatively low concentration (5 mol V m"3). The spectrum was not scanned

above 300 ppm, so it is not possible to rule out the presence of other thio complexes

(see Table 4.3).

9.3.1 Vanadium (V) Reduction

At higher vanadium (V) concentrations it was not possible to follow the reaction

between V(V) species and HS" ions using UV-visible stopped flow spectrophotometry,

because of the strong absorbance of the V(V). However, the following observations

could be made:

i) When a solution containing 10-100 mol V(V) m' 3 was reacted with an equal

volume of equimolar Na2S, the mixture instantly turned a green-brown colour. If a

large excess of sulphide was added, after several minutes a brown-black solid

began to precipitate from the solution.

ii) If this brown-black solid was separated and dissolved in hydrochloric acid, it

dissolved to form a blue solution.

iii) When hydrogen peroxide was added to the green-brown solutions, the solution

instantly turned turbid, and a yellow solid (which was identified as sulphur) could

be separated from a clear solution.

These observations imply that when more concentrated solutions of vanadium (V) were

used, reduction of the vanadium (V) to vanadium (IV) was achieved. Vanadium (IV)

solutions (prepared from VOSO4) appeared brown, but when they were exposed to air

they became a blue-green colour. After prolonged exposure to the atmosphere they

turned colourless. The UV-visible spectra of vanadium (IV) solutions before and after

exposure to the atmosphere are shown in Fig. 9.9 (overleaf).


Fig. 9.9 Spectra of 10 mol V(IV) m“3, before and after aeration.

pH = 9.3, T = 17 °C. Cell Path Length = 10 mm.

This behaviour suggests that vanadium (IV) solutions, which contain V 180 4212"

anions, can be oxidised initially to form a mixed valence V(V)/V(IV) ion in solution.

Mixed valence anions that are blue and green are now known, though still poorly

characterised (see section 4.3). Prolonged aeration formed the V(V) ion, HV2O73".

It is likely that the brown-coloured solutions that result when V(V) (at vanadium

concentrations > 10 mol m"3) and HS" are mixed, contain Vig04212“ ions:

12HS“ + 9 HV20 73“ -» V180 4212- + 3 S42" + 21 OH" (9.9)

This reduction appears to produce predominantly polysulphides, rather than elemental

sulphur, since the precipitation of sulphur was not observed unless an large excess of

vanadium (V) was used.

The brown solid that precipitates when vanadium (V) is in prolonged contact with

sulphide solutions must contain vanadium in the (IV) oxidation state or lower; the blue

solutions that were produced when the solid was dissolved in acid are characterisitic of

V 0 2+ ions. The possibility that the solid was a vanadium sulphide is remote, since an

examination of the E^-pH diagram for the vanadium-sulphur-water system

(Fig. 9.10) shows that there is no area of thermodynamic stability for a sulphide

phase at pH 9.3.


Fig. 9.10 E^-pH Diagram for the V-S-H20 System at 298 K.

Dissolved S species not shown. Activity of Dissolved species = 0.01.

Fig. 9.10 was produced using the computer program POURB, which was re-written

in FORTRAN 77 from a listing provided by Froning et al [136]. The thermodynamic

data, in the form of AGf° values, are shown in the Appendix and were taken from

Israel and Meites [30] and Mills [128].

From Fig. 9.10 it can be seen that the first solid phase to be encountered as the

solution potential is decreased at pH 9.3 is V3O5; this would be expected to dissolve in

acid to form the blue vanadyl cation (V 02+), and the green V3+ ions (which would be

oxidised on exposure to the atmosphere forming more vanadyl ions):

V3O5 + 8 H+ V 02+ + 2 V 3+ + 4 H20 (9.10)

It is also possible that the solid may have been a mixed-valence salt (see section 4.3).

Thermodynamic data are not yet available for these species and so they cannot be

included in Fig. 9.10.


Hydrogen peroxide reacted rapidly with polysulphide solutions, producing elemental

sulphur:

2 H 202 + 2 S42" -> S8 + 4 OH' (9.11)

By contrast, when hydrogen peroxide was added to solutions containing HS" at pH 9.3

no sulphur was produced; the oxidation proceeded instead to form sulphoxy

compounds, such as thiosulphate:

4 H 20 2 + 2HS- -> S20 32- + 5 H 20 (9.12)

However, when H20 2 was added to the green-brown solution, produced by mixing

equal volumes of V(V) and HS" (both at concentrations of 10 mol m"3), elemental

sulphur was again produced. This suggested that polysulphides were produced during

the reaction between V(V) and HS- and that these were oxidised to sulphur by H20 2.

Hydrogen peroxide was also a sufficiently strong oxidant to convert vanadium (IV) to

vanadium (V), forming a colourless solution:

v l8°4212' + 9 H 20 2 + 15 o h - 9H V20 73- + 12H zO (9.13)

Therefore, the addition of hydrogen peroxide to a reduced solution, containing

vanadium (IV) and polysulphide ions, can produce elemental sulphur and vanadium (V)

ions.

9.4 Interaction of AQ27DSH" Ions with Oxygen.

Hydrogen peroxide is a metastable intermediate in the electrochemical reduction of

oxygen at noble metal electrodes (see section 8.1), and can also be produced when

reduced anthraquinones react with oxygen in aqueous solution (see section 6.2).

Because of its possible role in the Stretford Process it was decided to analyse for the

presence of hydrogen peroxide during the oxidation of AQ27DSH" solutions.

Hydrogen peroxide was detected by titrating with As(III) [172 ]. H20 2 oxidises

As(III) to As(V) and unreacted As(III) was then back-titrated in acid solution with

iodate:

As02" + H20 2 —> A s03" + H20 (9.14)

I0 3- + 2H3As03 + 2H + + Cl- -> IC1 + 2H3As04 + H20 (9.15)

Preliminary checks were made to ensure that As(III) could not be oxidised either

through prolonged aeration or by direct reaction with anthraquinone 2,7 disulphonate

(AQ27DS).

A solution of AQ27DS was reduced electrochemically using the exhaustive electrolysis

described in section (7.3). The reduced solution was then re-oxidised externally by

mixing the solution with pure oxygen in a gas syringe, and recording the volume of gas

absorbed. The re-oxidised solution was then titrated with As(III) to detect the presence

of hydrogen peroxide.


It was found that the reduced solution absorbed oxygen in the molar ratio of O2 to

AQ27DSH", 1:2. No hydrogen peroxide was detected in the re-oxidised solution; this

suggested that the oxygen was reduced to form hydroxide ions in a four electron

process:

2AQDSH- + 0 2 -> 2AQ27DS + 2 OH" (9.16)

However, if the AQDSH" was injected into an aerated solution of As(III), under

conditions of excess oxygen at all times, then hydrogen peroxide was detected

quantitatively according to equation (9.17):

AQ27DSH- + 0 2 + H20 AQ27DS + H20 2 + OH“ (9.17)

In a solution containing As(III), any hydrogen peroxide produced would react

immediately according to reaction (9.14).

This behaviour suggested that hydrogen peroxide was produced as an intermediate

during the reduction of oxygen. When a reduced species was available in solution,

(such as As(III) or AQ27DSH"), the hydrogen peroxide reacted with it immediately

(producing AQ27DS and As(V) respectively). In the Stretford Process, the reaction

between AQ27DSH- and oxygen is likely to give rise to the in-situ production of

hydrogen peroxide; which is capable of producing elemental sulphur from polysulphide

solutions and oxidising V(IV) to V(V).

9.5 Stretford Solution Chemistry: Electrochemical Results

When a S tretford Process solution (containing 33 mol V(V) m~3 and 8 mol

AQ27DS m_3 in a carbonate buffer) was reacted with HS~ ions (10 mol m“3) the

solution instantly turned a turbid dark brown colour. Over the succeeding twenty

minutes this turbidity disappeared, leaving a dark-brown coloured solution. Both

vanadium (IV) and AQ27DSH" ions absorb strongly in the visible region, and together

would produce such a brown colouration. The turbidity may have been due to reduced

vanadium oxide (e.g. V203).

In the absence of oxygen, no precipitation of elemental sulphur was observed, even

after standing for twelve hours, although the solution contained a stoichiometric excess

of oxidising agents. However, when air or oxygen was bubbled through the solution,

sulphur was formed immediately and the red-brown colouration was discharged

simultaneously.


Continual voltammetry of the above solution during these reactions was conducted

using a gold flag electrode, with the potential limits set at +0.3 V and -1.2 V vs. SEE.

Initially a reduction wave was seen at around -0.285 V on the negative-going scan,

which was attributed to the reduction of AQ27DS. The re-oxidation peak was partially

suppressed, and occurred at 0.05 V vs. SHE. Similar patterns were seen during the

voltammetry of vanadium (V) solutions at gold electrodes (see Fig. 5.5), and were

attributed to the re-oxidation of vanadium oxide films. This suggested that the

formation of such films deactivated the gold electrode towards the re-oxidation of

AQDSH-.

When the sodium sulphide solution was added, the series of voltammograms shown in

Fig. 9.11 were obtained. Two new reduction peaks at -0.65 and -0.8 V vs. SHE

appeared, and increased in intensity during the 50 minutes following the addition.

Fig. 9.11 Voltammetry of Stretford Solution During Reduction

[AQ27DS] = 8 mol m-3, [V(V)] = 33 mol m-3, [HS-] = 10 mol n r 3.

Au Flag Electrode, pH = 9.3, Scan Rate = 200 mV s"1.

When the solution was oxygenated, these same peaks disappeared in about 20 minutes.

A possible explanation for this behaviour is that the species responsible for the

reduction peaks are polysulphide ions; cyclic voltammetry of polysulphide solutions at

a gold electrode revealed two reduction peaks (see section 3.5). The potentials of the


observed peaks in Fig. 9.11 differ from those seen in section (3.4), (-0.5 V and

-0.95 V vs. SHE), but this may be due to the state of the gold electrode surface; which

is likely to have been covered by films of elemental sulphur then vanadium oxide as the

negative-going scan proceeded.

9.6 The Stretford Process: Possible Mechanism

The results are consistent with the following mechanism occurring in the Stretford

Process:

Hydrogen sulphide dissolves in the alkaline solution producing hydrosulphide ions:

In the absorber and reactor vessels, the hydrosulphide ions are oxidised to form

poly sulphides. In practice, a mixture of polysulphides is likely to be produced, since

there is always more than one polysulphide present in significant concentrations in an

equilibrated solution (see Fig. 3.15). However, the following equations will be

written assuming that the predominant polysulphide species is S42":

3AQ27DS + 4H S- + OH" 3 AQ27DSH" + S42' + H20 (9.19)

12 HS“ + 9H V 20 73- -> V180 4212- + 3 S42" + 21 OH" (9.20)

The reaction between AQ27DS and HS" ions has been shown to be slow, whereas the

reaction with vanadium (Y) is more rapid. The thermodynamic driving force is greater

for the V(V)/HS" reaction, since the reversible potential values at pH 9.3 are -0.28,

-0.25 and -0.10 V vs. SHE for the HS"/S(0) AQ27DS/AQ27DSH" and V(V)/V(IV)

couples respectively. However, it is likely that the oxidation of hydrosulphide

solutions with vanadium (V) proceeds via the formation of a thiovanadate complex; it

has been shown that dilute vanadium solutions (< 1 mol V(V) m“3) may form the

complex ion V02S23" when contacted with excess HS":

HV20 73- + 4HS" <-» 2 V 02S23" + 2 H 20 + OH" (9.21 )

Transition metal thiovanadates are known to undergo intramolecular redox reactions

which can give rise to poly sulphides and reduced vanadium phases:

Colloidal V2O3 can dissolve to form vanadium (IV):

9 V 20 3 + 9H V 20 73- -4 2 V 180 4212- + 3 OH- + 3 H20 (9.25)

In this way, the formation of thiovanadate complexes can increase the rate of oxidation

of hydrosulphide solutions^

H2 S + OH" HS- + H2 0 (9.18)

vvo2s23- -> vmo2s23-2 V m 0 2S23- + H 2 0 - > V 20 3 + 2 S 22- + 2 OH"

(9.22)

(9.23)


In the oxidiser, the reduction of oxygen with AQ27DSH" gives rise to the in-situ

production of hydrogen peroxide:

A Q D SH - + 0 2 + H2 0 -> 2 A Q 2 7 D S + H 2 0 2 + O H ' (9 .2 6 )

This hydrogen peroxide can achieve the oxidation of the polysulphide ions to form

elemental sulphur (9.27), and assist in the oxidation of the vanadium (IV) species

(9.28) and colloidal V2O3.

2 H2 0 2 + 2 S42“ S8 + 4 O H ' (9 .2 7 )

V 18O4212' + 9 H 2 0 2 + 15 O H ' - > 9 H V 20 73- + 12 H 2 0 (9 .2 8 )

V 20 3 + 2 H2 0 2 + 3 OH- - » H V 2O 73- + 3 H 2 0 (9 .2 9 )

The brown vanadium (IV) species (V1804212-) will react slowly with oxygen to form

mixed-valence V(V)/V(TV) compounds and eventually form vanadium (V):

2 V i80 4212" + 9 0 2 + 30 OH" - > I 8 H V 2O 73- + 6 H 2 0 (9 .3 0 )

The rising air bubbles in the oxidiser also serve to recover the sulphur produced in

equation (9.23) by froth flotation, because of the naturally hydrophobic nature of

elemental sulphur. The re-oxidised solution is then re-cycled to the absorbing vessel

where it can undergo another redox cycle.

Conclusions 152

10. ConclusionsThe redox couples that are involved in the Stretford Process were studied using

electrochemical techniques, the interactions between them were investigated and a

process mechanism was proposed. The following results emerged:

10.1 The S(-II)/S (0) Redox Couple

The literature concerning the oxidation of sulphide solutions was reviewed. It revealed

that previous measurements of p K ^ for H2S had been in error and that a new value of

19 ± 2 will have to be accepted. This means that the S2" ion will not be the

predominant species in aqeous solution, even in highly alkaline media. At pH 8.5-9.5,

at which the Stretford Process operates, HS" ions are present. The published

Efo-pH diagrams revealed that elemental sulphur was not a thermodynamically stable

oxidation product at the process pH; sulphur’s only region of stability was below pH 7.

Operating the process below this pH would not only decrease the dissolution kinetics of

hydrogen sulphide, but would also alter the vanadium (V) speciation, producing

decavandate anions. These large anions are known to have slow reaction kinetics and

upon reduction they can precicipitate the sodium salts of mixed valence (V/IV)

vanadates.

Therefore, the Stretford Process is required to operate at a pH where the desired

product is thermodynamically unstable. Nevertherless, elemental sulphur is still

produced in high yield. However, this thermodynamic instability means that the

formation of some higher oxidation state products, such as thiosulphate, is inevitable.

Previous studies on the atmospheric oxidation of sulphide solutions showed that a wide

range of reaction rates and oxidation products could be observed, depending on the

particular temperature, pH and catalyst used.

The electrochemical oxidation of HS“ ions on gold electrodes at pH 9.3 was shown to

produce a sub-monolayer of a gold sulphide phase at low potentials (-0.4 V vs. SHE)

and multilayers of sulphur at higher potentials (0.05 V vs. SHE). Associated with the

formation and reduction of this sulphur layer, was the production of polysulphide ions,

Sn2“ (n = 2 to 5). The poly sulphide ions were detected at a ring electrode in a rotating

ring-disc electrode study. By comparing the charges passed during their production

and reduction, the average polysulphide chain length was calculated to be 1.8.

10.2 The V(V)/V(IV) Redox Couple

The literature relating to the aqueous chemistry of vanadium in alkaline solutions was

reviewed. This showed that vanadium, in common with other transition metals in the

same region of the periodic table, displays a marked tendency to form polymeric

anions. Early disagreements about V(V) speciation have been largely resolved, but

uncertainty still exists as to speciation of the lower vanadium oxidation states. V(V)

exists as the ions HV2O73' and V40 j2^" in the Stretford Process solutions; upon

reduction these are likely to form the brown V(IV) polyanion V180 4212”. However,

Conclusions 153

prolonged exposure to reducing environments can produce a precipitate of

vanadium (III) oxide (V2O3), and mild reduction may also produce mixed-valence

(V)/(IV) compounds, such as VjqC ^ 6".

The electrochemical reduction of vanadium (V) was found to be irreversible on a variety

of electrode surfaces. This reduction led to the production of vanadium oxide films

(V3O5, V2O3 and VO) rather that to a solution species. Irreversible electrochemical

behaviour and high overpotentials are commonly associated with processes that require

a large structural rearrangement, as is the case with the reduction of HV2073- to

V 18O4212". The effectiveness of V(V) as an oxidising agent in the S tre tfo rd

Process suggested that there was a specific chemical interaction occurring which

facilitated V(V) reduction. A range of vanadium-sulphur complexes (the thiovanadates)

are known, which are likely to formed under the chemical conditions prevailing in the

absorbing vessel of the Stretford Process. It is possible that these thiovanadate

complexes can undergo intramolecular redox reactions, producing polysulphide ions

and a reduced oxidation state vanadium complex; in this way the formation of

thiovanadate complexes may offer reaction pathways with lower activation energies

than would otherwise be the case. This mechanism may explain the catalytic role of

V(V) in the process.

10.3 The Anthraquinone/Anthraquinol Redox Couple

The reduction of the anthraquinone 2,7-disulphonate (AQ27DS) can produce either the

semiquinol in a single electron process, or the fully reduced quinol in a two electron

process. An equilibrium can be established between these two species:

AQ27DS + AQ27DS2" <-> 2 AQ27DS-' (10.1)

The comproportionation constant for this reaction can determine the voltammetric

behaviour of the compound. If the equilibrium constant is high, two consecutuve one

electron reductions are observed, whilst if it is low, a single wave is seen

corresponding to a two electron reduction.

The values of pKaj and p K ^ for the quinol (AQ27DSH2) are ~7 and 10.8 respectively.

As would be expected from these values, the electrochemical reduction of AQ27DS at

pH 9.3 was found to proceed in a two electron, one proton process:

AQ27DS + 2e- + H+ AQ27DSH- (10.2)

However, the presence of a strong esr signal from the reduced solution showed that

AQ27DSH" existed in equilibrium with the radical species, AQ27DS*". From an

analysis of the voltammetric peak separation, the comproportionation constant was

estimated to be in the range 0.2 to 4.

Conclusions 154

The radical anion reacts rapidly with oxygen to form the superoxide ion 0 2*", which

can further react with water to form hydrogen peroxide:

AQ27DS-- + 0 2 -> AQ27DS + 0 2-‘ (10.3)

2 0 2*" + 2 H20 -» 0 2 + H20 2 + 2 OH- (10.4)

Reactions (10.3) occurs extremely rapidly, which has the effect of shifting equilibrium

(10.1) to the right. In this way the oxygen can become reduced to form hydrogen

peroxide, while the reduced AQ27DSH" is oxidised to form AQ27DS:

AQ27DSH- + 0 2 + H20 -> AQ27DS + H20 2 + OH- (10.5)

When a solution of AQDSH" was allowed to react with oxygen, hydrogen peroxide

was detected quantitatively according to equation (10.5).

10.4 The 0 2/OH_ Redox Couple

The direct electrochemical reduction of oxygen at platinum and gold electrodes gave rise

to currents at a rotating disc electrode which were intermediate in value between those

expected for mass transport limited two and four electron processes. At low rotation

rates on platinum, the current approached the value predicted for a four electron

process, and at higher rates it tended towards the two electron value. This suggested

that hydrogen peroxide was formed as a metastable intermediate, via a two electron

process (10.6). At low rotation rates H20 2 was further reduced at the electrode surface

to form hydroxide ions (10.7), whilst at higher speeds it was dispersed into solution.

0 2 + 2 e - + H2 0 - > H 2 0 2 + 2 OH- (1 0 .6 )

H2 0 2 + 2 e " - » 2 OH- (1 0 .7 )

Previous workers have also found that hydrogen peroxide can be produced in alkaline

solutions from the two electron reduction of oxygen, and this has been attributed to the

high strength of the oxygen-oxygen bond. The direct four electron reduction of oxygen

to form hydroxide ions involves the cleavage of this bond, whilst it remains intact

during the two electron reduction to form hydrogen peroxide.

10.5 The Redox Chemistry of the Stretford Process

Stopped flow spectrophotometric studies indicated that AQ27DS reacted with HS" ions

to form the fully reduced quinol and polysulphide ions:

3AQ27DS + 4HS" + OH" 3 AQ27DSH- + S42“ + H20 (10.8)

The reaction was found to be first order with respect to AQ27DS under conditions of

excess HS" and the first order rate constant was determined to be 2.53 x 10~4 s"l (at a

temperature of 17 °C and a concentration of 50 mol HS" m-3).

Conclusions 155

The high optical absorbance of V(V) and V(IV) solutions prevented the reaction

between concentrated vanadium (V) solutions (> 1 mol m"3) and HS" ions being

investigated. However, the reaction with a dilute vanadium (V) solution (0.5 mol m"3)

led to the formation of the thiovanadate complex VO2S23':

HV20 73- + 4HS- <-> 2 V 02S23' + 2 H 20 + OH' (10.9)

The VO2S23' ion was identified by its optical absorbance at 360 nm, although the 51V

NMR peak that had been attributed these species could not be detected.

Vanadium solutions at a concentration greater than 10 mol V(V) m"3 reacted rapidly

with HS' ions, forming polysulphide ions rather than elemental sulphur:

12 HS- + 9H V20 73' -> V180 4212- + 3 S 42' + 21 OH' (10.10)

Solutions containing the V i804212- anion were re-oxidised slowly when they were

exposed to air. This re-oxidation produced a blue-green mixed valence compound

initially; vanadium (V) was regenerated only after prolonged oxygenation.

10.6 The Mechanism of Sulphide Oxidation in the Stretford Process

The above evidence points to the following process mechanism:

Hydrogen sulphide dissolves in the alkaline solution producing hydrosulphide ions:

H2 S + OH' -> HS- + H2 0 (10.11)

In the absorber and reactor vessels, these are oxidised to form polysulphides (such as

S42") and the two oxidising agents in solution are reduced:

3AQ27DS + 4H S- + OH' -> 3 AQ27DSH' + S42“ + H20 (10.12)

12 HS- + 9 HV20 73' -> Vi80 4212- + 3 S 42- + 21 OH' (10.13)

The reduction V(V) to V(IV) is likely to proceed via a mechanism of thiovanadate

formation, followed by an intramolecular redox reaction, and the desorption of the

resulting polysulphide.

In the oxidiser, oxygen is reduced by its reaction with the anthraquinol, which leads to

the in-situ production of hydrogen peroxide:

AQDSH' + 0 2 + H20 -> 2AQ27DS + H20 2 + OH' (10.14)

This hydrogen peroxide can react with the polysulphide ions in solution, producing

elemental sulphur:

2 H20 2 + 2 S42" —> Sg + 4 OH' (10.15)

Hydrogen peroxide is also capable of re-oxidising the vanadium:

V ig042^2” + 9 H20 2 + 15 OH“ —> 9 HV20 73' + 12 H20 (10.16)

although vanadium (IV) can be re-oxidised more slowly by its direct reaction with

oxygen.

2 V180 4212" + 9 0 2 + 30 OH' -> 18HV20 73' + 6 H20 (10.17)

Conclusions 156

The rising air bubbles in the oxidiser also serve serve to recover the sulphur produced,

which is naturally hydrophobic and concentrates in the froth. This froth is filtered and

the sulphur produced can be further purified for sale. The re-oxidised solution is then

returned to the absorbing vessel where it can commence another redox cycle.

10.7 Concluding Remarks

There are several process implications arising from the above mechanism:-

i) Polysulphide solutions tend to be oxidised to form thiosulphate when they contact

oxygen directly. Therefore, maintaining an oxygen-free environment in the

absorbing tower and reaction vessel (where the polysulphides are produced) would

be expected to decrease the rate of thiosulphate production. Since sulphur is

thermodynamically unstable at the process pH, some production of thiosulphate is

inevitable.

ii) If vanadium (V) is exposed to highly reducing conditions (i.e. when too much H2S

enters the absorber) a number of reduced-valence state vanadium salts can

precipitate. The oxides V3O5 and V2O3 are thermodynamically stable at low

solution potentials, and the production of the alkali metal salts of poorly

characterised mixed-valence V(V)/(IV) anions (e.g N ag V io C ^.B Ô ) is also

possible. The vanadium (V) concentration in the process solution is likely to be

critical; dilute solutions do not oxidise the HS" ions (merely forming thiovanadate

complexes) whereas concentrated solutions are more likely to precipitate the

vanadium from solution.

iii) To be an effective oxygen-reduction catalyst, the anthraquinone must produce the

maximum yield of hydrogen peroxide in the oxidising vessel. This is governed by

the stability of the radical semiquinol, which exists in equilibrium with the fully

reduced quinol. If the equilibrium concentration of the semiquinol is low, it might

not be sufficient to reduce the oxygen as fast as the gas enters solution. Conversely,

if the semiquinol were highly stable, it would also be unreactive towards oxygen

and so would be of no use as a re-oxidation catalyst. The stability of the

semiquinol will be affected by the nature and position of any substituent groups;

this explains the widely different catalytic activity of anthraquinone disulphonate

isomers.

iv) The rate-determining step in the Stretford Process will depend upon the

operating conditions; in a poorly optimised plant, the transport of H2S into solution

in the absorber, or of oxygen into solution in the oxidiser, may limit the plant's

throughput. However, in most operating plants it is the oxidation of HS" to form

polysulphide in the absorbers and reaction vessels is likely to be the rate-limiting

step. Early Stretford plants used solutions that contained only anthraquinone

salts, and they suffered low throughput rates because of the slow reaction kinetics

between anthraquinone and HS" ions (the first order rate constant between AQ27DS

Conclusions 157

and excess HS" at 17 °C is only 2.53 x 10"4 s_1). The addition of vanadium (V)

salts to the process solution has greatly improved the oxidation kinetics of the

process, which is probably due to the role of the thiovanadates (since the

electrochemcial reduction of vanadium (V) at a number of electrode surfaces has

been shown to be slow). The search for a more effective sulphide oxidation

catalyst may be best focused on those transition metals which are also known to

form thiometalates; for example molybdenum, tungsten, and rhenium.

Appendix 158

Appendix: Thermodynamic Data Used in Eh-pH Diagrams

The following thermodynamic data were used in the calculation of the E^-pH diagrams.

Table A.1 was taken from the recent review by Israel and Meites [30], these values

were in close agreement with the earlier compilation by Post [98] (from which some of

the values were taken).

Species AGf° / k j mol-1V 0.0y 2+ -218

VO -404.2y3+ -251.3

V20 3 -1139

VO+ -451.8

VOH2+ -471.9

v 305 -1816

V407 -2473

v 204 -1318.6

V 02+ -446.4

VOOH+ -657

(VOOH)22+ -1331

V4092- -2784

V60i3 -4109

v 205 -1419.4

vo 2+ -587

VOJ- -783.7V 043- -899.1

h v o 42- -974.9

H2vo 4- -1020.9

H3V04 -1040.3

v 2074- -1720

h v 2o 73- -1792

h 3v 2o 7- -1864

V30 93- -2356

V40124- -3202

V10°286' -7675

HV1o0285' -7708

H2Vio0 284- -7729

v o 2.h 2o 2+ -746.4

v o .h 2o 23+ -523.4

Table A .l AGf° Values for Vanadium Compounds at 298 K.

Appendix 159

Table A.2 shows the thermodynamic data that were used in the calculation of the

Sulphur / water E^-pH diagrams. These were taken from the recent compilation by

Zhdanov [31].

Species AGf° / k j mol

S 0.0S2- 86.31

S22" 79.5

S32’ 73.6

S42- 69.0

S52’ 65.7

SQ2(aq) -300.7

S0 32- -486.6

S 042- -744.6

S20 32’ -518.8

S20 42- -600.4

S20 52- -791

S20 62" -966

S20 82- -1110.4

s 3 062 - -958

s 4062 - - 1022.2s 5 o 62 ’ -956.0

HS' 12.1H2S(aq) -27.9

HS03- -527.8

h s2o4- -614.6

HSO4- -756.0

H2S 0 3 -537.9

h 2s o 4 -744.6

H2S204 -616.7

H2S20g -1110.4

Table A.2 AGf° Values for Sulphur Compounds at 298 K.

Appendix 160

Table A.3 shows the AGf° values for the vanadium sulphides. AHf° and S° values

were given for these compounds by Mills [128], and the AGf° values were calculated

using S° values of 28.93 and 31.82 J k-1 mol-1 for elemental vanadium and sulphur

respectively [173].

Species AGf° / k j mol-1VS -191.3

V2S3 -516.2

VS4 -286.0

Table A.3 AGf° Values for Vanadium Sulphides at 298 K.

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Documents

Electrochemical Studies of Catalysed Aqueous Sulphide ......5.2 Voltammogram of Vanadium (V) in Borate Buffer at pH 9.2. 8 5 (First Scan, commenced at 0.35 V vs. SHE. 50 mV s_1.) 5.3