electrochemical studies of catalysed aqueous sulphide ......5.2 voltammogram of vanadium (v) in...
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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
Electrochemical Studies of Catalysed Aqueous Sulphide Oxidation
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
Electrochemical Studies of Catalysed Aqueous Sulphide Oxidation
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).
Review of Sulphide Oxidation 28
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.
Review of Sulphide Oxidation 2 9
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)
Review of Sulphide Oxidation 30
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.
Review of Sulphide Oxidation 31
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)
Review of Sulphide Oxidation 3 2
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.
Review of Sulphide Oxidation 3 3
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
Review of Sulphide Oxidation 3 4
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.
Review of Sulphide Oxidation 35
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)
Review of Sulphide Oxidation 3 6
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)
Review of Sulphide Oxidation 37
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
Review of Sulphide Oxidation 38
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.
Review of Sulphide Oxidation 3 9
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.
Review of Sulphide Oxidation 4 0
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)
Review of Sulphide Oxidation 4 1
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.
Sulphide Electrochemistry 4 3
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
Sulphide Electrochemistry 4 4
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
Sulphide Electrochemistry 4 6
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,
Sulphide Electrochemistry 4 7
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
Sulphide Electrochemistry 4 8
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.
Sulphide Electrochemistry 49
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].
Sulphide Electrochemistry 50
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].
Sulphide Electrochemistry 51
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.
Sulphide Electrochemistry 52
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.
Sulphide Electrochemistry 53
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
Sulphide Electrochemistry 54
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).
Sulphide Electrochemistry 55
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].
Sulphide Electrochemistry 56
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.
Sulphide Electrochemistry 57
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
Sulphide Electrochemistry 58
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).
Sulphide Electrochemistry 59
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.
Sulphide Electrochemistry 60
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
Sulphide Electrochemistry 61
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.
Sulphide Electrochemistry 62
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)
Sulphide Electrochemistry 63
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).
Sulphide Electrochemistry 64
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
Sulphide Electrochemistry 65
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 V^Ojg6- 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.
Vanadium Electrochemistry 83
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)
Vanadium Electrochemistry 84
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
Vanadium Electrochemistry 85
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].
Vanadium Electrochemistry 87
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.
Vanadium Electrochemistry 89
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).
Vanadium Electrochemistry 90
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.
Vanadium Electrochemistry 91
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].
Anthraquinone Review 93
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.
Anthraquinone Review 94
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).
Anthraquinone Review 95
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.
Anthraquinone Review 96
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.
Anthraquinone Electrochemistry 98
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.
Anthraquinone Electrochemistry 99
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.
Anthraquinone Electrochemistry 100
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)
Anthraquinone Electrochemistry 101
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)
Anthraquinone Electrochemistry 102
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).
Anthraquinone Electrochemistry 103
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.
Anthraquinone Electrochemistry 104
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.
Anthraquinone Electrochemistry 105
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
(E^i - 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
Anthraquinone Electrochemistry 106
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).
Anthraquinone Electrochemistry 107
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.
Anthraquinone Electrochemistry 108
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).
Anthraquinone Electrochemistry 109
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.
Anthraquinone Electrochemistry 110
Fig. 7.9 Plot of Current vs. Time for Electrolysis of AQ27DS.
Potential = -0.6 V vs. SHE. pH = 9.3.
Anthraquinone Electrochemistry 111
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
Anthraquinone Electrochemistry 112
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
Anthraquinone Electrochemistry 113
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).
Anthraquinone Electrochemistry 114
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
Anthraquinone Electrochemistry 115
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
Anthraquinone Electrochemistry 116
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.
Anthraquinone Electrochemistry 117
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.
Anthraquinone Electrochemistry 118
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.
Anthraquinone Electrochemistry 119
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.
Anthraquinone Electrochemistry 120
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.
Anthraquinone Electrochemistry 121
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.
Oxygen Reduction 123
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.
Oxygen Reduction 124
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.
Oxygen Reduction 125
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.
Oxygen Reduction 126
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.
Oxygen Reduction 127
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.
Oxygen Reduction 128
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.
Oxygen Reduction 129
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
Oxygen Reduction 130
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.
Oxygen Reduction 131
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).
Oxygen Reduction 132
□ 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.
Stretford Process Chemistry 134
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.
Stretford Process Chemistry 135
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
Stretford Process Chemistry 136
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.
Stretford Process Chemistry 137
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)
Stretford Process Chemistry 138
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.
Stretford Process Chemistry 139
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.
Stretford Process Chemistry 140
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.
Stretford Process Chemistry 141
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.
Stretford Process Chemistry 142
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).
Stretford Process Chemistry 143
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.
Stretford Process Chemistry 144
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).
Stretford Process Chemistry 145
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.
Stretford Process Chemistry 146
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.
Stretford Process Chemistry 147
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.
Stretford Process Chemistry 148
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.
Stretford Process Chemistry 149
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
Stretford Process Chemistry 150
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)
Stretford Process Chemistry 151
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 ^O ) 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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