The University of New South Wales

Faculty of Applied Science

School of Chemical Engineering and Industrial Chemistry

A Study of Molybdenum Redox Couples

by

PAUL PENNISI

A Thesis Submitted as Part of the Requirements for the Degree of

Master of Science

1995 U N S W 1 7 JUL 1997

LIBRARY CERTIFICATE OF ORIGINALITY

I hereby declare that this submission is my own work and that, to the best of my knowledge and belief, it contains no material previously published or written by another person nor material which to a substantial extent has been accepted for the award of any other degree or diploma of a university or other institute of higher learning, except where due acknowledgment is made in the text.

I also declare that the intellectual content of this thesis is the product of my own work, even though I may have received assistance from others on style, presentation and language expression.

£maUM Paul Pennisi ACKNOWLEDGMENTS

I would like to thank Professor Maria Skyllas-Kazacos for her endless help during this project and the Vanadium Battery Group whose various members have shared their knowledge and expertise, they have been a great help.

My gratitude also goes out to the Australian Research Council for their support with this project.

The support and knowledge gained from the staff and friendships made, in the School of Chemical Engineering and Industrial Chemistry was greatly appreciated.

Much credit goes to my Mother and Father whose support saw me through my undergraduate degree.

Most of all I would like to thank my wife for her patience over the last two years. ABSTRACT

Evaluation and optimisation of a molybdenum redox cell was the main objective of this work. There is demand for new energy storage systems. Redox flow cells are an attractive alternative because of their long life, simplicity, low long term cost, wide range of applications, high energy efficiency, as well as being electrically rechargeable.

Using a single metal ion in different oxidation states for both the positive and negative half cell electrolytes, the cross mixing, low energy density, irreversibility, and cross contamination which most redox systems suffer from can be largely eliminated.

Many electrolytes, electrode materials, membranes, and molybdenum salts were screened to optimise the molybdenum redox cell. The methods used to screen the different solutions and electrodes was solubility tests of molybdenum and cyclic voltammetry. The primary criteria required for a good combination is both significant solubility (more than one mole per litre) and significant electroactivity shown by cyclic voltammetry at an electrode surface.

Solubility studies of several molybdenum salts in a range of supporting electrolytes have shown that the maximum solubility of a molybdenum compound which also had significant electroactivity at a graphite electrode, was 1.6M, the salt being disodium molybdate dihydrate in 4M sulphuric acid. The kinetics of four molybdenum couples Mo(VI)/Mo(V), Mo(V)/Mo(IV), and two for Mo(IV)/Mo(III) was investigated at a glassy carbon electrode. The heterogeneous rate constants for these couples were 3.7E-4, 3.9E-5, 1.8E-4, and 2.5E-5 respectively. However, the Mo(VI)/Mo(V) reduction reaction could not be observed at the glassy carbon electrode, it was found that Mo(VI) is reduced directly to Mo(IV).

The redox couples initially proposed for a Mo - Mo redox flow cell were Mo(VI)/Mo(V) and Mo(III)/Mo(IV) for the positive and negative half cell electrolytes respectively. The molybdenum cell using 1M sodium molybdate in 4M sulphuric acid as the electrolyte, with 25cm2 graphite felt electrodes, and a current density of 20mA/cnr gave an energy efficiency of more than 70% showing that it can operate as a redox cell. The coulombic efficiency was high although further improvements in the total energy efficiency would be achieved by better cell design, and a more electroactive electrode material to catalyse the Mo(VI)/Mo(V) reaction and reduce the voltage losses. Unfortunately, however, the discharge voltage was found to be 0.4V which could limit its applications. Further work is thus required to achieve higher energy efficiencies with the molybdenum redox cell.

Overall the molybdenum redox cell has shown some promising features which could make it feasible as a renewable energy source although the low discharge potential does imply that practical applications would be very limited. TABLE OF CONTENTS

ABSTRACT Table of Contents CHAPTER 1 INTRODUCTION CHAPTER 2 THEORETICAL BACKGROUND 2.1. Redox Flow Cell 11 2.1.1. Redox Flow Cell Advantages 11 2.1.2. Basic Requirements for a Redox Cell 12 2.2. Why molybdenum? 13 2.3. Safety, Health, and the Environment 14 2.4. Molybdenum Chemistry 15 2.4.1. Different Oxidation States of Molybdenum 17 2.4.2. Molybdenum Half Cell Reaction Standard Reduction Potentials 19 2.4.3. The Effect of pH on Molybdenum Chemistry 22 2.5. Molybdenum Blue 25 2.6. The Effect of Acid Type on Molybdenum Chemistry 29 2.6.1. Hydrochloric Acid 29 2.6.2. Sulphuric Acid 32 2.6.3. Molybdenum Salt Solubilities 34 2.7. Redox Fuel Cell Application of Mo Redox Couples 37 2.8. Battery Definitions 38 2.9. Electrochemical Kinetic Theory 39 2.9.1. Polarisation Losses 40 2.9.2. Cyclic Voltammetry 42 2.9.3. Summary of Heterogeneous Reaction Kinetics 51 2.9.3.1. Irreversible Reactions 51 2.9.3.2. Reversible Reactions 52 2.9.3.3. Quasi-reversible Reactions 53 2.9.4. Instrumentation Used for Cyclic Voltammetry 55 CHAPTER 3 SUMMARY OF LITERATURE REVIEW AND RESTATEMENT OF OBJECTIVES CHAPTER 4 EXPERIMENTAL 4.1. Chemicals and Materials 58 4.2. Instruments 59 4.3. Experimental Procedures 60 4.3.1. Solubility Tests 60 4.3.1.1. Direct Dissolution 60 4.3.1.2. Electrolytic Reduction and Dissolution 61 4.4. Analytical Methods for Mo Concentration Determinations 62 4.4.1. Potentiometric Titrations 63 4.4.2. Atomic Absorption Spectroscopy 64 4.5. Preparation of Solutions of Different Oxidation States 66 4.6. Temperature Stability Tests 67 4.7. Cyclic Voltametric Studies 67 4.7.1. Electrodes Used for Cyclic Voltammetry 68 4.7.1.1. Counter and Reference Electrodes 68 4.7.1.2. Working Electrodes 68 4.8. Molybdenum Redox Flow Cell Performance Tests 70 4.8.1. Flow Cell Components and Setup 70 4.8.2. Cell Electrode Fabrication 72 4.8.3. Cell Resistance Calculation 72 4.8.4. Efficiency Calculation 74 CHAPTER 5 OBSERVATIONS, RESULTS, AND DISCUSSION 5.1. Solubility Studies of Molybdenum Compounds 75 5.1.1. Solubility Studies of Mo(VI) Compounds 75 5.1.2. Solubility of Electrolytically Generated Mo(VI) Solutions 84 5.1.3. Molybdenum Solution Reduction Observations 86 5.1.4. Mo(III) Stability to Atmospheric Oxygen 91 5.2. Thermal Stability Studies of Molybdenum Electrolytes 94 5.3. Cyclic Voltammetry Studies 98 5.3.1. Effect of Acid Concentration 98 5.3.2. Screening the Molybdenum Salts 107 5.3.3. Electrode Selection 109 CHAPTER 6 CYCLIC VOLTAMETRIC STUDY OF THE KINETICS OF MOLYBDENUM REDOX COUPLES 6.1. Mo(III)/Mo(IV) Couple 118 6.2. Mo(V)/Mo(VI) Couple 120 6.3. Mo(IV)/Mo(V) Couple. 121 6.4. Kinetic Parameters of Concentrated Molybdenum in Dilute Sulphuric Acid at a Glassy Carbon Electrode 123 6.5. Electrode Area 129 CHAPTER 7 EFFECT OF COMPLEXING AGENTS, ELECTROLYTE ADDITIVES AND ANION CONCENTRATION ON THE CYCLIC VOLTAMETRIC BEHAVIOUR AND SOLUBILITY OF MOLYBDENUM 7.1. Effect of Sulphate Concentration on Mo Blue Formation 132 7.2. Effect of Complexing Agents 135 CHAPTER 8 MOLYBDENUM REDOX CELL PERFORMANCE TESTING 8.1. Design and Operation 137 8.2. Cell Resistance 138 8.3. Molybdenum Cell Charge-Discharge Testing 139 8.4. Electrode and Membrane Stability to Mo(VI) and Mo(III) 145 8.5. Mo(III) and Mo(VI) Stability Under Paraffin Oil 146 8.6. Molybdenum Source 146 CONCLUSION BIBLIOGRAPHY APPENDICES 1 Mo Standard Solution AAS Data 156 2 Potentiometric Titrations 157 3 ln(ip) vs (E1(-E0’) plots 160 4 ip vs v0-5 plots 168 5 Cyclic Voltammograms at Different v 176

7 CHAPTER 1 INTRODUCTION

There is demand for new energy storage systems for a wide range of applications. Redox flow cells which employ two soluble redox couples in the positive and negative half cell electrolytes are an attractive alternative because of their potentially long life, reliability, low maintenance, simplicity, low cost, and wide range of applications.

The aim of a redox flow cell system is to have an electrically rechargeable bulk energy storage system with a high overall efficiency, extended cycle life, high reliability, that can operate at ambient temperatures, and at the same time be economical. Earlier work with the Fe-Cr redox couple revealed that an inherent problem with such systems is the cross mixing of the two electrolytes across the membrane (67).

Using a single metal ion in different oxidation states for both the positive and the negative half cell electrolytes, the cross mixing, and capacity loss, due to cross contamination which most redox systems suffer from can be largely eliminated. The advantages of the vanadium redox flow cell which employs V(V)/V(IV) anc* V(III)/V(II) have already been demonstrated (15,16,73-78). Like vanadium, molybdenum also exists in a number of oxidation states and therefore has the potential to be used in a redox flow cell.

8 There are primary cells which use molybdenum as the anode (13), and examples of chemically regenerative redox cells can be found (12,14). A molybdenum / vanadium redox cell has been described by Kummer (5). The couples were Mo(III)/(IV) versus V027V02+. The cathodic and anodic reactions respectively being:

VO,+ + 2H+ + e- = V02+ + H20 E° = 1.0V

[(Mo)23+Mo4+] = [Mo4+]3 + 2e‘ E° = 0.0V

In an earlier study by Trinh (15), a Mo - Mo redox flow cell system was proposed which employed the Mo(III)/Mo(IV) and Mo(VI)/Mo(V) redox couples in dilute sulphuric acid as the negative and positive electrolytes respectively. The molybdenum cell using one molar molybdic acid in 5M sulphuric acid, gave an energy efficiency of around 70%.

Further work was, however required to fully understand the chemical and electrochemical behaviour of these redox couples and to optimise the Mo - Mo system.

The two redox couple systems Mo(III)/(IV) and Mo(V)/(VI) were thus studied as potential couples for the negative and positive electrolytes respectively.

9 Many electrolytes, electrode materials, membranes, and molybdenum salts were thus screened to identify the optimum components for a molybdenum redox cell. The method used to screen the different solutions was solubility tests of molybdenum salts, while the electrochemical reversibility was studied by cyclic voltammetry. The primary criteria required for a good redox couple is both significant solubility ( greater than 1M ) and high electroactivity at the electrode surface as shown by cyclic voltammetry

10 CHAPTER 2 Theoretical Background

2.1. Redox Flow Cell

A redox flow cell converts chemical energy stored in its active material directly into electrical energy in an electrochemical oxidation - reduction reaction. The redox cells have fully soluble redox couples that react at inert electrodes. Each half cell electrolyte contains an electroactive species in a different oxidation state. The redox couple electrolytes are stored in separate tanks and are pumped through the cell stack where the oxidation - reduction reactions take place. The two electrodes on each side of the cell are separated by an ion selective membrane. Being chemically reversible reactions, they can be recharged. Self discharge and loss of system capacity by diffusion of the cations across the membrane is still possible to some extent. The main components of the redox flow cell are illustrated in Figure 2.1

2.1.1. Redox Flow Cell Advantages

The redox flow cell has many advantages over conventional battery systems. These include:

* Long storage life * Couples with low viscosity even at high concentration * No problem from cross mixing as in the vanadium redox system * Two fully soluble redox couples * Absence of solid state reactants and their morphological changes at the electrode

11 * No solid state slow processes which can lower voltage efficiencies are involved * No cycle life limitation * No limit on electrode life as they are inert and serve only as current collectors and the reaction sites for the redox reactions. There are no plating reactions limiting electrode life * Discharge does not damage the battery so replacement costs are low * Capacity is determined by the solution concentration and volume, so it can be increased by using a larger reservoir * May be instantly recharged by replacing electrolyte * Since all cells are fed from the same electrolyte reservoirs, all cells are at the same SOC (state of charge) * It is possible to charge a battery at 2V and discharge at 100V without affecting the life or performance of the battery * Continuous SOC monitoring is possible with the use of open circuit cells

2.1.2. Basic Requirements for a Redox Cell

* The anodic reactant has at least one higher oxidation state and can lose an electron in the reaction.

* The cathodic reactant has at least one lower oxidation state and can gain an electron.

* The sum of the free energy changes for the two reactions is negative -dG = nFE (E = open circuit voltage)

* The potential should be at least one volt.

12 * The electrolyte must be a good conductor to avoid voltage drop within the cell during discharge.

* Resistance losses need to be minimised (see 2.7. Polarisation Losses).

* For long storage at least one component is able to be isolated. (19)

— +

Conducting Membrane Plastic Substrate Carbon Fibre Electrode

Negative Positive Half-Cell Half-Cell Electrolyte Electrolyte

Figure 2.1. Redox Flow Cell

2.2. Why Molybdenum?

The chemistry of molybdenum is similar to that of vanadium. While a number of elements exist in several oxidation states, a close examination of the standard reduction tables reveals that there is only a handful of redox couple systems which satisfy the following basic requirements for redox cell application. Both oxidised and reduced forms must be highly soluble species (that is, should not involve gaseous or solid compounds). Secondly, there should be sufficient separation of the standard reduction potentials of the two redox couples to allow a reasonable cell voltage (preferably > IV). Superficially, molybdenum appears to meet these requirements.

13 Also, preliminary studies showed promising results of around 70% efficiencies could be achieved without optimisation. Trinh’s (15) cell consisted of 200mls of a 1M molybdic acid solution in 5M sulphuric acid with the charged positive and negative electrolytes being Mo(VI) and Mo(III) respectively. The electrode was a carbon felt, and the cell was charged and discharged at 20mA/cm2 for about an hour each. An overall energy efficiency under these conditions of nearly 70% was achieved. (15)

2.3. Safety, Health, and the Environment

The molybdenum compounds pose relatively little threat as toxic substances. The metabolism of molybdenum in mammals is linked to that of copper and iron. It is an integral part of certain enzymes, particularly xanthine oxidase.

Excretion is mainly in the urine. Levels of molybdenum in the body can be increased in the presence of inorganic sulphate in the diet. Usually excretion is so rapid and efficient that there is practically no storage in the body.

Most animal experiments have indicated that molybdenum has a relatively low toxicity. Lethal dose for rats is (LD50) 125mg/kg for the trioxide, and 333mg/kg for ammonium molybdate(17), and for CaMo04 101mg/kg(3). For vanadium pentoxide the LD50 is just 23mg/kg.(3) No occurrence of significant toxicity of molybdenum compounds has been observed in humans, nor has any cumulative effect been noticed. (3)

14 2.4. Molybdenum Chemistry

The principal use of molybdenum in industry is for catalysts, pigments, lubricants, and high temperature stainless steel. Because it is only about 1-1.5 x 10'4 % of the earth’s crust it is relatively expensive. The main ore of molybdenum is molybdenum disulphide (21), and it is also a product of the copper industry. (44)

Table 2.1. Comparison of the Terrestrial Abundance of Molybdenum with Other Elements (44)

Element Parts Per Million

Chromium 122

Vanadium 136

Chlorine 126

Molybdenum 1.2

Tungsten 1.2

Holmium 1.4

Terbium 1.2

15 Molybdenum is commercially available as disodium molybdate, the ammonium molybdate salt, molybdic acid, phosphomolybdic acid, molybdenum disulphide, and both molybdenum dioxide and trioxide.

There are many compounds of molybdenum in the oxidation states Mo(V) & Mo(VI), which have one or more oxygen atoms as terminal ligands or as bridging ligands. Strong multiple bonding between the oxygen atom and the metal atom results in_ a range of compounds. (22)

Molybdenum compounds may disproportionate to mixtures of compounds in which molybdenum occurs in different oxidation states. The most stable species was thought to be Mo(VI) but it is readily converted to a Mo(V)/(VI) divalent blue species which seems to be more stable.

The following reactions of molybdenum, some of which occur simultaneously, hold for the pH range of 0-14. (13)

H2Mo04 + 2H+ + e — Mo02+ ■+■ 2H20

Mo02+ + e‘ = Mo02

H2MoO, + e + H+ = MoChOH + H20

16 Molybdenum forms a series of oxides, Mo-oxygen systems are complex and not completely defined. Two oxides of definite composition are Mo03 and MoO,, with the first being the more stable oxide. Mo03 is insoluble in water but soluble in alkaline solutions and mineral acids forming molybdates. Molybdenum compounds may disproportionate to mixtures of compounds in which molybdenum occurs in different valence states.

2.4.1. Different Oxidation States of Molybdenum

Mo(II)- Molybdenum (II) species are seldomly mentioned in the literature and at first it appears that they do not exist at all. Bowen (34) describes ’New Molybdenum Species’ including some which have molybdenum in an oxidation state of two. No monomeric compounds of Mo(II) with saturated ligands are known. Mo(II) complexes are usually stabilised by metal to metal bonds or by unsaturated ligands.

Mo(III)- This species is always found hydrated and not as free Mo(III). The monomeric [Mo(H20)6]3+ has a characteristic green colour, but Mo(III) also exists as the dimeric [Mo2(OH)2(H20)8]4+, and as the trimeric [Mo3(OH)4]5+.

Mo(IV)- This oxidation state exists as the oxomolybdenum(IV) ion

[Mo304(H20)J4+ with absorption at 310 nm. This species can be reduced at a potential of - 0.25V(vs NHE) giving the trimeric oxomolybdenum(III) species via an intermediate of Mo2(III)Mo(IV). The Mo(IV) species is more stable to air oxidation than the Mo(III) or the Mo(V) species.

17 Mo(V)- The dimeric oxo species exists as [Mo204(H20)6]2+. The following quote describes the chemistry of this and most molybdenum species well, "The chemistry of Mo(VI) is not clear in aqueous solutions, the chemistry of Mo(V) is even more obscure."(7,53)

Mo(VI)- At low pH cationic molybdenyl species such as Mo022+ form. Mo(VI) can be chemically, electrochemically, and photochemically reduced. It is possible that other impurities could aid the partial reduction process. "The polymerisation of acidified solutions of Mo(VI) yields the most complicated of all the polyanion systems. It is often difficult to explain why, under given circumstances, a particular degree of aggregation or a particular structure is preferred." (44)

Elwell reports that molybdenum displays basic characteristics in the lower oxidation states and acidic properties in the higher states. This could help explain solubility differences.

Molybdenum solutions in the literature have been studied extensively but the molybdenum concentrations used have been usually in the millimole range whereas for redox cell application the desired concentration range would be over 1M. As a result of this, much of the literature on molybdenum chemistry is not immediately applicable.

18 2.4.2. Molybdenum Half Cell Reaction Standard Reduction Potentials

The following half cell reactions involving molybdenum ions have been reported in the literature together with their corresponding reduction potentials.

E° MoOp- + 4H20 + 6e < = = > Mo + 80H 0.91(31,57) 2Mo03 + 2H+ + 2e' < = = > Mo,05 + H20 0.76(31) Mo(VI) + e- (8F HC1) < = = > Mo(V) 0.70(56) Mo(VI) + e- (2F HCL) < = = > Mo(V) 0.53(56,26,57) Mo(VI) + e (8N sulphuric) < = = > Mo(V) 0.53(59)

Mo03(s) + 4H+ + 4e- < = = > Mo03+ + 2H>0 0.48(26) H2Mo04(aq) + 2H+ + e' < = = > Mo02+ + 2H20 0.40(57) H2Mo04(aq) + 6H+ + 6e‘ < = = > Mo + 4H20 0.00(26,57) Mo042' + 8H+ + 6e‘ < = = >Mo + 4H2Q (E = E° -0.0788V/pH) - 0.15(31)

Mo(V) + 2e* (2F HC1) < = = > Mo(III)red 0.11(56,57) Mo02+ + 4H+ + 2e < = = > Mo(III) + 2H20 0.00V(31,57)

0.5Mo2O5 + H+ + e' < = = > Mo02 + 0.5H2O - 0.23V(31) Mo(V) + 2e- (2F HC1) < = = > Mo(III)green - 0.25V(56,57,59)

Mo(IV) + e' (4.5F Sulphuric) < = = > Mo(III) 0.10(56,59,26)

MoQ2 + 4H+ + e- < = = > Mo(III) + 2H2Q - 0.31V(31)

Mo(III) + 3e (acidic) < = = > Mo 0.20(31,57,59)

1/2 Mo203 + 3H+ + 3e- < = = > Mo + 3/2 H2C> 0.12(31)

19 The following three tables give the standard potentials for the half cell reactions of Mo(IV) < = = = > Mo(III), in equal concentrations, at 25°C as measured in HC1 (with H3PO., or CH3COOH) as a function of H+ versus NHE(31). Note that Mo(III) has two forms, one is deep green and the other is a red/green. From the potentials given, Table 2.4 illustrates that at lower hydrogen ion concentrations the green Mo(III) form is preferred. This is confirmed in Table 2.2 and 2.3 up to phosphoric acid and acetic acid concentrations of 0.34 and 0.99 mol/L respectively.

Table 2.2. Mo(IV)/(III) Half Cell Potentials (vs NHE) in 0.3N HC1 and Varying the Phosphoric Acid Concentration (31)

h3po4 Mo(III) Mo(III) green green/red (Mol/L) (Volts) (Volts)

0.17 -0.042 -

0.34 -0.029 -

0.50 - 0.084

1.0 -0.002 0.105

2.0 - 0.123

3.9 0.002 0.173

12.0 - 0.307

20 Table 2.3. Mo(IV)/(III) Half Cell Potentials (vs NHE) in 0.3N HC1 and Varying the Acetic Acid Concentration (31)

CH3COOH Mo(III) Mo(III) green green/red Mol/L (Volts) (Volts)

0.495 0.034 -

0.99 0.063 -

1.98 0.082 0.161

3.96 - 0.193

11.87 - 0.203

Table 2.4. Mo(IV)/(III) Half Cell Potentials at 20°C Varying the Hydrochloric Acid Concentration (vs NHE) (31)

HC1 Mo(in) Green (Mol/L) (Volts)

0.03 0.03

0.08 0.11

0.17 0.25

21 2.4.3. The Effect of pH on Molybdenum Chemistry

Electrochemistry of Mo(VI) molybdate aqueous solutions experience pH dependent equilibria involving oxo-ligand protonation, oxo-bridge formation, and di, oligo, or poly-merisations, by disproportionation reactions. (41)

At a pH above 7 or 8 Mo(VI) occurs as the tetrahedral monomeric molybdate ion MoO/', but polymerisation occurs at concentrations in excess of 0.0001M at lower pH values. Molybdates form insoluble precipitates with most cations with the exception of the alkali metal ions. (40)

Addition of strong mineral acid to.molybdate produces a precipitate of molybdic acid which then dissolves in an excess of the acid to form molybdenyl compounds of the type MoO.SO.,. The molybdic acid formed has a very strong tendency to condense and polymerise, not only with itself but with a variety of other acids to form isopoly and heteropoly complex ions. Most molybdates are polymolybdates of the formula M6Mo(Mo602_,).4H20 (M a heavy metal). Their solubilities vary in water. The alkali salts are readily soluble, whereas the alkaline earth and heavy metal molybdates are relatively insoluble. (6,52)

The MoO/' ion is stable above pH 5. Acidification leads to the first polymeric species the heptamolybdate ion as:

7Mo042' + 8H+ = [Mo70Mf

22 This reaction is complete at pH 4.5. Then at pH 2.9-1.5 the octamolybdate ion forms [Mo8026]4 , and at pH 0.9 molybdic acid precipitates. At pH < 0.9 the precipitate redissolves to form species such as Mo02 .

The formal potentials of Mo(IV)/(III) and Mo(V)/(VI) are influenced by the acid concentration, the values increase with increasing acid concentration. (23) It has been concluded also, that potential is directly related to the anion

concentration and is independent of H+ concentration, but for HC1, H3P04, and

CH3COOH potential increases with an H+ ion concentration increase. (4) The sequence of the acid’s anion, to form increasingly stable complexes is CIO4-, NO3, Cl, so;-. (3)

Many intermediate reactions occur, especially hydration and protonation. Adjustment of acidity, concentration, and temperature can produce solids of many species which are apparently not present in solution. Changing the pH changes the equilibrium and changes the species present in solution, as shown by the following examples. (44)

7[Mo04]2' + 8H+ < = = > [Mo70,,r + 4H20

8[Mo04]2" + 12H+ < = = > [Mo802J4- + 6H20

36[Mo04]2' + 64H+ < = = > [Mo36On2]8- + 32H.O (44)

23 The following figure is an equilibrium diagram of potential versus pH for aqueous molybdenum. The data used to generate this figure was originally obtained by Baes and Mesmer from NBS Tables (40). The upper dashed line represents 1 Atm oxygen and the lower dashed line 1 Atm hydrogen. The concentrations of dissolved species are 10'6M.

Figure 2.2 suggests that there are no soluble Mo(IV) species and does not consider Mo(V) species. This figure does not take into account or predict which species would exist in the presence of any anions like sulphate which complexes in various ways with molybdenum depending on the pH.

pH

Figure 2.2. Potential - pH Diagram for Molybdenum

24 2.5. Molybdenum Blue

Molybdenum solutions are often a strong blue colour known as molybdenum blue of which the composition is uncertain. Molybdenum blue is formed readily by mild partial reduction of Mo(VI) to a divalent Mo(VI)2Mo(V) species. There are probably several distinct compounds included as Mo(blue). This is one example which highlights the complexity of molybdenum chemistry. (6)

Molybdenum blue (divalent molybdenum species of V and VI oxidation states), is formed in either acidic or basic solutions by partial reduction of molybdates or by oxidation of lower valency state species. (6)

Mild reduction of acidic molybdate gives a strong blue colour. This molybdenum blue has an uncertain composition, which is probably a mixture of valencies, of hydrous molybdenum oxides. (3) The polymerisation of isopolymolybdates increases with increases in acid concentration, but these higher polymers are more difficult to characterise. (42) The Gmelin Handbook series (9) gives a very comprehensive list of molybdenum species, including those included as molybdenum blue as well as how they are formed.

It is well established in the literature that molybdenum blue arises from partial reduction of Mo(VI). Sidgwick explains more specifically that if molybdenum in sulphuric acid solutions is exposed to dust it is ’superficially’ reduced and turns blue. Dust then includes impurities. It is claimed that molybdenum blue contains two valence states Mo(VI) and Mo(V), which are only stable when containing water. (43)

25 These species are also suggested elsewhere (44). Molybdenum blue is made up of a large number of distinct and accurately stoichiometric phases proposed to be compounds such as Mo4On, Mo17047, and MogO^. As oxygen is eliminated a series of Mon03n.i species are feasible. Intermediate phases are by no means fully understood however their non stoichiometric ratios of Mo and oxygen imply mixed valence states. (44)

Heteropoly molybdates such as [Mn(IV)Mo9032]6' and [Ni(IV)Mo9032]6' also exhibit the typical molybdenum blue intense colours upon mild reduction. (44)

The phenonemon of molybdenum blue has been studied by Khan. The study confirms the instability of Mo(VI) solutions, which form molybdenum blue. It was found that many species influence the formation of molybdenum blue. Among those studied were bismuth, phosphate, acid concentration, and even molybdenum concentration. The formation of blue over a wide pH range suggests that more than one species is involved. It is not surprising that molybdenum blue has such an intense colour, since it has an element which is in two oxidation states, due to the electron transfer bands. The solid molybdenum blue is quoted to vary between the formula MoO,5and MoO,. In a closed system the amount of oxygen in the complex will depend on the initial ratio of molybdenum and oxygen. (7)

Molybdenum blue is formed when the molybdenum salt dissolves, reacts with a proton, and is slowly transformed to a new pentavalent species. It is proposed that molybdenum blue is formed by reaction of the polymerised form of hexavalent molybdenum (say Mo40132') with the positively charged species of pentavalent molybdenum.

26 When acidic Mo(VI) is mixed with a solution of Mo(V) in a ratio of 2:1, molybdenum blue is certainly formed. The rate of formation and the intensity of the final blue colour is dependent on the acid concentration of the solution. It has been shown that the intensity of the colour, and rate of formation, decreases with a decrease of pH. (7)

It has been reported that Mo(VI) in the presence of organics and ultraviolet light results in the formation of molybdenum blue as shown below. Many examples of this can be found in the literature. The relevance to this study is that organic impurities will contribute to molybdenum blue formation. (45,48,49,50)

• 1

Mo(VI)=0 Mo(VI ,=°

hv O -—> Charge ~> 0 Transfer •

Mo(VI) Mo(V)(49)

The electrochemical properties of a isopolymolybdic acid thin film modified carbon fibre microelectrode, prepared by dip coating, gives three couples of surface redox waves between 0.70V and 0.1V vs SCE in 2M sulphuric acid. The stronger the acidity, the better is the stability and reversibility of the electrode, the reactions taking place can be expressed as:

27 Mo80264- + mH+ + 2ne' < > [HmMo,2n(VI)Mo2n(V)026]^^ n = l,2,3 m=2,5,7

The cyclic voltammogram of a molybdenum anion film has three redox waves. With an increase in scan rate the reduction potential slightly shifts towards positive potentials, while the oxidation towards negative. The valence state of the metal cations can be examined by X-ray photo electron spectrum (XPS). The binding energy values for Mo(VI) are 235.5 and 232.4eV, whereas for Mo(V) they are 234.4 and 231.3eV. It was confirmed that the molybdenum blue species does contain both oxidation states.(51)

It has been shown that the potential/pH dependence is not 59mV/pH as usual, but 82.5mV/pH. This is because of the interaction between the protons and redox molybdenum species. Peak potential separations increase with the decrease of scan rate. This is due to the bridging oxygen atoms being reactive and consumed in the early stages of the reduction. The addition of an electron will result in a weakening of the bridge-oxygen bond, and an increase of the basicity of the anions. At low scan rate there is enough time to weaken this oxygen bond, while at the faster sweep rate, it does not. This is why the redox reversibility of a Mo80264‘ film at fast scan rates is better. (51)

Further text worth consulting include reference 87.

28 2.6. The Effect of Acid Type on Molybdenum Chemistry

2.6.1. Hydrochloric Acid

In less than 3M HC1, Mo(V) produces a brown colour, while at HC1 concentrations above 5.75N the colour is emerald green. Other unspecified ionic species exist at intermediate acidities. (37) It has been shown that green Mo(III) solutions contain cationic species and the orange-red solutions of Mo(III) contain anionic species depending on the starting species of the Mo(III), the nature of the oxidant, and the nature and concentration of the acid used. The cationic aquomolybdenum (III) chloride is more susceptible to aerial oxidation and is more reactive. (38,39)

Electrochemically prepared Mo(III) in 0.1M HC1 slowly reoxidised to Mo(V) probably with reaction with H+ even if stored under oxygen free conditions. It seems the Mo(III) species (probably MoCl6 3J is slowly hydrolysed to a more reactive species which can be oxidised by H+ at low Cl' concentrations.

Also, in HC1 Mo(III) reacts instantaneously with Mo(VI) to give Mo(V), and Mo(III) can react in several ways. So Mo(III) can be oxidised by Mo(VI), H+, or electrochemically at high voltages. Note that molybdenum potentials are more positive in HC1 than in H2S04 . In HC1 concentrations of 2-7M (and higher), Mo02C12, Mo02C13', and Mo02C14', are probably important species. (9,25)

29 Mo(VI) in an acidic chloride solution, at low concentrations (below 0.001- 0.0001M Mo) avoids polyacid formation, and Mo042' is converted to monomeric molybdic acid. In this case polymers such as HMoXV in dilute acid and (Mo02C12)2 in more concentrated acid are probable. As the HC1 concentration increases the hydroxyl groups of Mo(OH)6 are replaced by chloride ions. Mo02C12 forms at HC1 concentrations of 2-7M, and then

Mo02C13', Mo02C142', is formed at higher concentrations.

The two following reactions were studied in HC1 and were found to be zero order with respect to the hydrogen ion concentration as determined in solutions containing 2, 4, 8, and ION H+; ION CT, and 0.005M in Mo(VI), Mo(V), and Mo(III) ions. Additionally, it was found that the first reaction was first order with respect to Mo(VI), and zero order with respect to Mo(V). Similarly, the second reaction was found to be first order with respect to Mo(V), and zero order with respect to Mo(III). (13)

Mo(VI) + e' < = = > Mo(V) Mo(V) + 2e- < = = > Mo(III)

The kinetics of the first reaction above was studied and confirmed by a second group, and it was found that the half cell potential was independent of the total molybdenum concentration, but decreases with an increase in Mo(V). At Mo(VI)/Mo(V) = 1, the potential was found to be 0.61, 0.48, and 0.43 vs NHE in 6.2, 3.0, and 1.1N HC1 respectively, the reaction being:

(Mo03+)2 + 2H20 < = = > (Mo022+)2 + 4H+ + 2e‘ in 2-4N HC1 Mo03+ + H20 < = = > MoO;2+ + 2H+ + e- in 6-8N HC1 (31)

30 Chalilpoyil (11) also confirms that Mo(V) is reduced to a dimeric Mo(III) product. The electrochemistry of molybdenum in trifluoromethanesulphonic (HTFMS) acid was examined. HC1 and H2S04 as electrolytes often produce mixtures of complexes. Whereas, HTFMS allows study without complex formation. (11)

Reduction of Mo(IV) at -0.25V vs NHE in HPTS with the hydrogen ion concentration ranging from 0.5-4.0M gives a Mo(III) ion. This ion is oxidised by ClOf as well as oxygen. (30)

In polarographic studies of Mo(VI)/(V) the amount of molybdenum precipitate decreased as either the HC1 or the Cl' ion concentration was increased. (25) In the equimolar system of Mo(VI)/(V), in dilute HCI, an addition of thiocyanate ions caused a shift of potential from 0.5V to 0.64V at high thiocyanate concentration. The potential difference is independent of molybdenum concentration and proportional to the logarithm of thiocyanate. The potential of this system in a solution containing equal amounts of Mo(VI) and Mo(V) in 1.0M HCI increased when phosphoric acid was added. Mo(V) forms a more stable complex with phosphoric acid then with Mo(VI). (31)

The reduction of Mo(VI) to Mo(V) in the presence of gluconic acid is reversible from pH 2.3 to 6.8. Mo(VI) is reduced to Mo(V) over the entire pH range while Mo(V) is reduced to Mo(IV) or Mo(III) depending on pH. The Mo(VI) and gluconic acid form a complex as it does with many ligands. (32)

31 Mo(III) chloride solutions can be obtained either by electrolytic reduction or by zinc or aluminium. The formal potentials of Mo(IV)/(III) and Mo(V)/(IV) are influenced by the acid concentration, the values increase with an increase of acid concentration. (23)

2.6.2. Sulphuric Acid

Mo(VI) reacts with H2S04 to form complex molybdenyl Mo022+, and molybdyl Mo04+ cations which form soluble compounds. In equilibrium with 1-4N FI2S04 there is molybdic acid monohydrate, in the range 6-7N Mo02S04, then at 9-

12N H2S04 Mo02S04 may precipitate. (3)

For a solution which is 2.4g/L Mo, 1-18.5N sulphuric acid, and has a Mo(VI)/Mo(V) ratio varied from 1:20 to 20:1, the potential measured between a platinum and a reference electrode was found to increase with both an increase in the Mo(VI) concentration, and the sulphuric acid concentration. (31)

32 Table 2.5. Redox Potential Versus Acid Concentration at 20°C for a Mo(VT)/(V) Ratio of 1 (31)

Sulphuric Acid Potential Concentration (vs NHE) Normality volts

18.5 0.53

9.7 0.47

4.7 0.43

1.0 0.41

The redox potentials of a solution of a Mo(VI)/(V) ratio of 1, at a molybdenum concentration of 0.01M were measured at 25°C in 0.1 to ION HC1, H2S04, H3P04, tartaric acid, oxalic acid, and citric acid. Additionally, the anion concentration was kept constant while the proton concentration was varied. The conclusion from these experiments was that the potential is directly related to the anion concentration and independent of the hydrogen ion concentration. (31)

33 Table 2.6. Redox Potentials of Mo(V)/(III) at 20°C Varying the Acid Concentration (vs NHE) (31)

Sulphuric Acid Voltage (g/L) (vs NHE) volts

0.9 -0.01

4.5 0.03

9.4 0.10

17.5 0.23

Three polarographic reduction E1/2 waves are found in 0.1 M sulphuric acid plus 0.2M sodium sulphate with values of 0.06, -0.29, and -0.60V. The first is attributed to the reduction of Mo(VI) to Mo(V), the other two to the reduction of Mo(V) to Mo(III). The presence of nitrate ions causes a catalytic current to be produced. There have been polarographic methods developed for the determination of molybdenum concentrations from these waves. (6)

2.6.3. Molybdenum Salt Solubilities

The maximum solubility of molybdic acid in H2S04 has been studied by two groups, Table 2.7 and 2.8 agree well in that the maximum solubility is in 8N sulphuric acid. Note that it is reported that the solubility of molybdic acid can be increased to 280g/l (from 224g/l in sulphuric acid) by using 11N HC1. Table 2.8 also compares the solubility in 3M sulphuric acid at two temperatures, clearly showing that the solubility is reported to decrease with increase in temperature, this will be discussed further later.

34 Table 2.7. Maximum Solubility of Molybdic Acid in Various Acids (4)

Acid Solubility g/1 20°C

11N HC1 280

4N HN03 154

8N H2S04 224

Table 2.8. Solubility of Molybdic Acid in Vaiying Concentrations of Sulphuric Acid (5)

Acid Solubility g/1 20°C

3M H2S04 140

4M H2S04 200

3M H2S04 90 @ 50°C

Solubility values for some more unusual molybdenum complexes in various electrolytes have also been published. (36) Not only are the solubilities low, but the electrolytes are reported to often form non-conductive solutions (36) (see Table 2.9).

35 Table 2.9. Solubilities of Three Molybdenum Salts at 20°C in Various Solvents (36)

Salt Solvent Solubility g/L

MoBr3 50% Tartaric Acid 51.5

MoBr3 50% Dextrose 51.5

MoBr3 Ammonium Nitrate 10

MoC13 Ethyl Chloride(lOml) 10

Benzene(20ml)

MoC15 Carbon 51.5

Tetrachloride

MoC15 Ethyl Bromide 16

MoC15 Aniline (10ml) 50

Ethyl Bromide(5ml)

36 2.7. Redox Fuel Cell Application of Molybdenum Redox Couples

In a redox fuel cell at least one couple can be regenerated chemically with a renewable source. For example Kummer (5) studied the reduction of (0.5M) sodium molybdate to Mo(IV) and Mo(III) by hydrogen in concentrated (3M) sulphuric acid. It was found that highly acidic (>3M sulphuric acid) molybdate solutions can be readily reduced by hydrogen using a platinum catalyst, to a trinuclear Mo ion. So, it would be theoretically possible to have a molybdenum redox cell where the negative Mo(III)/(IV) couple was continuously recharged by hydrogen.

Pt/H, Mo(IV)3 + 2e- < = = = > Mo(III)2Mo(IV) E° = 0.0 V vs NHE (5)

The main advantage being that there would only be the need for a small volume of negative electrolyte to be used and thus large cost savings in terms of energy density and capacity. Relevant publications about fuel cells include that by Larsson (33). Kummer (5) used vanadium for the cathodic couple and molybdenum for the anodic as follows.

Anode: Pt/H2 (230cm3/min flow by) 45-50°C

Mo(IV)3 + 2e- < = = = > Mo(III)2Mo(IV) E° = 0.0 V vs NHE (5)

Cathode: 75-80°C

V02+ + 2H+ + e" <===> V02+ + H20 E° = -1.0 V vs NHE (5)

37 2.8. Battery Definitions

Batteries can be classified as either primary or secondary. If the reactions which are to provide the current are not readily reversible, the battery is referred to as a primary battery. If they are reversible, the battery is called a secondary or storage battery.

The important and significant characteristics for a battery include all those of the redox system as listed in 2.1.1. Additionally the following points must be considered :

Available capacity Available energy Power it can deliver Cycle life

The theoretical capacity of a cell is given by Qt = xnF. Where x is the theoretical number of moles of reactants associated with complete discharge of the cell and n is the number of electrons involved in the reaction. The coulombic efficiency is given by Q(/Q„ where Qp is the actual number of coulombs delivered.

The energy efficiency is defined as the ratio of the actual and theoretical energy delivered. The power delivered by a cell is the product of the current and the associated cell voltage produced during discharge. The cycle life of a battery is the number of times a cell can undergo charge discharge tests before its performance has been diminished below some arbitrary limit.

38 2.9. Electrochemical Kinetic Theory

For the general electron transfer reaction where O is reduced and R is oxidised:

O + ne' —> R

The equilibrium potential is given by the Nernst equation:

[R] Eeq = E° - RT/nF In — Equation 2.1.

[O]

where E° = Standard Reduction Potential

R = Universal Gas Constant n = number of electrons involved in the reaction F = Faradays Constant T = Temperature in Kelvin

When a current flows through the cell, however, the potential deviates from the equilibrium or zero current potential because of polarisation or ohmic losses.

39 2.9.1. Polarisation Losses

An electrode through which a current passes has a potential different from its zero current or equilibrium value. This difference is called overpotential (rj), associated with voltage losses. It can be defined as the difference in potentials when the electrode is at equilibrium and when it is sustaining a net anodic or cathodic reaction. Overpotential contains ohmic, concentration, and activation overpotentials.

Ohmic or IR drop results from a potential of any working electrode in a cell of finite resistance. Concentration polarisation exists with sustained electrolysis, where concentrations at the electrode surface differ from that in the bulk of the solution. Activation overpotential is the potential in charge transfer reactions at the electrode - electrolyte interfaces required to drive the reaction. Activation overpotential is related to the rate of the overall electrochemical reaction which is dependent on the activation energy (Ea). Ea depends on the amount of energy required to transfer an electron through the double layer established between the electrode and its ions in solution. So a low Ea is desirable for kinetic reasons. (19,20)

Within a cell there are various significant causes of polarisation. These have to be minimised in order to maximise cell efficiencies. Practically, applying an electric current results in a certain degree of irreversibility. This is due to losses of some of the electrical energy in the form of heat in the internal resistance of the cell.

40 For a galvanic cell, or a battery undergoing discharge:

Discharge : Ecell = Ecq - EI>ol

For a galvanic cell, or a battery undergoing charge:

Charge : E„„ = E„, + E,„,

Epo, is the potential loss due to polarisation in the cell.

Concentration overvoltage is associated with the depletion or accumulation of electroactive material near the electrode surface. The polarisation curve presented in Figure 2.3 initially shows a sharp fall in cell voltage due to electrode polarisation overvoltage (region (i)). The middle region (ii) is linear exhibiting ohmic polarisation, associated with the cell resistance. Finally, at the high current drain, (region(iii)) IR polarisation is combined with concentration polarisation at the electrode surfaces.

Current (mA)

Figure 2.3. Typical Polarisation Curve for an Electrochemical Cell

41 2.9.2. Cyclic Voltammetry

Cyclic voltammetry is an electrochemical technique for studying electroactive species rapidly and over a wide potential range. It involves cycling the potential of an electrode which is immersed in an unstirred solution of the electroactive species and measuring the resultant current as a function of the applied potential. The potential of the working electrode is controlled versus a reference electrode such as a standard calomel electrode (SCE) as illustrated in Figure 2.4.

The voltage applied to the working electrode is scanned linearly with time from an initial value E; to a preset limit Es (switching potential), where the direction of the scan is reversed (refer to Figure 2.5). Differences between initial and successive scans or with different scan rates are important clues to discovering information about reaction mechanisms.(8)

After the potential has passed the potential where one or more electrode reaction takes place the direction of the scan is reversed and the electrode reactions of intermediates and products, formed during the forward scan, often can be detected.(64)

42 A typical cyclic voltammogram is shown in Figure 2.6. In general the current increases as the rate of reduction increases at more negative potentials until a maximum is reached and then the current decreases steadily. The cathodic peak in the cyclic voltammogram results from the competition of two factors, the increase in the rate of reduction as the potential is made more negative and the development of a thickening depletion layer across which reactant must diffuse. At potentials more than 100/n mV away from the peak, the reactant concentration at the electrode surface is small compared to that far from the electrode, and the current is then controlled by the rate of diffusion of reactant through the depletion layer.(64)

Important parameters of measurement in cyclic voltammetry are the anodic (ipa) and cathodic (ipc) peak currents and the peak potential separation dEp = Epa - Epc. Where Epa and Epc are the anodic and cathodic peak potentials respectively. Results can be used to calculate the formal potential, the standard heterogeneous rate constant and the diffusion coefficients of the redox couples, and thus give important information on the suitability for the redox flow cell application.

Inert Go« Inlet ---- Working Electrode

Reference Electrode-

— Auxiliary Electrode

Figure 2.4. Electrochemical Cell for Voltammetiy

43 A + <• —

Switching time. X

Fig 2.5. (a) Potential sweep (b) Cyclic Voltammogram (15)

POTENTIAL (Volls) Fig 2.6. Cyclic Voltammogram Showing Anodic and Cathodic Peak Currents and Potentials

44 For a Reversible Reaction, the peak current ip is given by:

ip = 0.4463nFACo*(nF/RT)1/2Do1/2v1/2 2.2.

While for a Irreversible Reaction:

ip = 0.227nFACo*k° exp[-(anaF)(Ep-E°')/RT] 2.3 ip = 299000n(a na)1/2AC0‘D01/2v1/2 2.4

A = Surface area of working electrode (cm2) CG* = Concentration of active species (mol/cm3) k° = Heterogeneous rate constant (cm/s) a = Transfer coefficient Ep = Peak potential (V) E°’ = Formal Potential (V) n = Number of electrons in the reaction na = Number of electrons in the rate determining step

D° = Diffusion coefficient (cm2/s) v = Scan rate (mV/s) K = Equilibrium constant i = Current (A)

45 If the reaction is completely reversible, using for the baseline the cathodic curve, which would have been obtained if there had been no change in direction of the potential scan, the anodic peaks are the same, independent of the switching potential and identical in height and shape to the cathodic wave. The ratio of ia/ic = 1, is independent of switching potential. This behaviour is diagnostic of the absence of various coupled chemical reactions, as is the peak potential being proportional to the scan rate. (79)

In practice, the position of the anodic wave on the potential axis is a function of the switching potential as illustrated in Figure 2.5. This is due to the fact that on the cathodic scan, the surface concentration of substance R does not quite equal CG* at potentials close to the peak, and if the anodic scan is started under these conditions, the concentration of R at the surface of the electrode is much less than the concentration of O at the corresponding potentials for the cathodic scan. This causes an anodic shift in the wave, which decreases as the switching potential is made more cathodic (for reversible charge transfer). (79,85)

Generally the anodic portion of the cyclic voltammogram is not affected as much by the preceding chemical reaction as is the cathodic portion. For a chemical reaction preceding a reversible charge transfer, generally the height of ia increases as the switching potential is made more cathodic (refer to Figure 2.7). For a chemical reaction preceding an irreversible charge transfer, the curves are more drawn out because of the effect of the electron transfer coefficient. (79)

46 For an irreversible reaction succeeding a chemical reaction, if a very rapid reaction is involved in experiments with a very slow scan rate the curve will reflect the chemical step almost entirely. If the scan rate is fast compared to the rate of reaction the curves correspond to uncomplicated charge transfer.(79)

Cyclic Voltammetry can be used to detect when a chemical reaction takes place prior to electron transfer, following electron transfer, or between electron transfer steps. For example if the product of an electrode reaction is lost via chemical reaction then the peak will be reduced in magnitude, and it will be completely absent if the reaction half-life is much less than the scan duration.(64)

For base line correction using Figure 2.7, Nicholsans equation (8) is:. i„a = (ipa)o + 0.485(isp)o + 0.086i1>c 2.5. isp = current at the switching potential

Instead of using Nicholsarfs equation, if the anodic sweep is stopped and the current is allowed to decay to zero the resulting cathodic current-potential curve is identical in shape to the anodic one, but is plotted in the opposite direction on both axis. This happens because allowing the anodic current to decay to zero results in the diffusion layer being depleted of R and populated with O at a concentration near CD* such that the cathodic scan is virtually the same as that which would result from an initial cathodic scan in a solution containing only species R.

47 Other similar methods for correcting for the non zero base line are given elsewhere in the literature. For example Sawyer (56) gives a formula which has similar constants to Nicholson. Another author gives an even more complicated formula involving no less than 8 variables within it, all of which come from a single voltammogram. (62)

-100 -200 -300

Figure 2.7. Parameters Required for Base Line Evaluation for the Reverse Scan in a Cyclic Voltametric Experiment (8)

48 Some other important definitions include:

Ep/2 = The cathodic half peak potential

E1/2 = The half wave potential

E° = The standard potential or the potential when all the activities are equal to unity at standard conditions of 25°C, latm.

E°’= The formal reduction potential or the experimental potential which is observed in a given medium when the concentrations are in the ratios prescribed by the stoichiometry of the reaction. The prime signifies that the effect on the free energy of the reactants and products embodied in activity coefficients has been combined with the thermodynamic reduction potential to form a term that is directly measurable but subject to solution conditions.

D0 = The diffusion coefficients of the oxidised species, D0 and DR are usually equal and so E1/2 is usually close to the formal potential.(64) k° = The standard heterogeneous electron transfer rate constant. It is a property of the reaction between the particular compound and the electrode surface used. The heterogeneous rate constant k°, gives a measure of the kinetic ease and speed of a redox couple, large k° means that equilibrium is reached faster, whereas a small k° means that a longer time is required to reach equilibrium. kf and kb can be made large by applying a large potential relative to E°’. At large overpotentials the current levels off due to mass transfer limitations rather than because of the heterogeneous kinetics.(8)

a = The transfer coefficient. It arises because only a fraction of the energy (applied potential) that is put into the system lowers the activation energy barrier. (65)

49 One criterion for reversibility is Epa-Epc = 57/n mV, which is independent of scan rate and concentration. The criteria for diffusion control is that ipc/v1/2 must be constant (v is sweep rate). If the reaction is reversible and diffusion controlled then ipc/[nFACo*(D0nFv/RT)1/2] = 0.446. A quasi-reversible or irreversible reaction will give a much smaller value. For a quasi-reversible reaction the heterogeneous rate constants for the forward and backward reaction are about equal. For an irreversible cathodic reaction the forward constant is much bigger than the reverse.(64)

Reversible means that the reaction is fast enough to maintain the concentrations of the oxidised and reduced forms in equilibrium with each other at the electrode surface. Redox couples whose peaks shift further apart with increasing scan rate are quasi-reversible. (65)

The current depends on two steps, the movement of electroactive material to the surface and the electron transfer reaction. The electron transfer rate constant for a reduction process is a function of potential and can be described by (65):

k,= k°exp[(-a nF/RT)(E-E°')] 2.6

50 2.9.3. SUMMARY OF HETEROGENEOUS REACTION KINETICS

2.9.3.I. IRREVERSIBLE REACTIONS

Irreversible reactions typically exhibit no significant current flow except at high overpotential

k° < 0.00003 cm/sec k° < 0.00002v1/2

Ep-Ep/2 = 48/(ana) mV at 25°C ip proportional to C0‘, v1/2 but Ep is a function of v (Ep changes by -30/(ana) for each 10 fold increase in v)

Ep occurs beyond E0’ by an activation overpotential related to k° by ip = 0.227nFACo*k° exp[-(anaF)(Ep-E°’)/RT] k° and a are treated as phenomenological parameters evaluated by experiment.

(8)

51 2.93.2. REVERSIBLE REACTIONS

Reversible reactions typically exhibit fast charge transfer processes.

The potential and surface concentrations are in equilibrium.

Current flows due to the electrodes’ surface concentration not being in equilibrium with the bulk such that the mass transfer to the electrode surface occurs continuously.

kG > 0.02 cm/sec

Ep is independent of v ip is proportional to v1/2

Ep-Ep/2 = 56.5/n mV at 25°C

If Ep is proportional to v1/2 then it is diffusion controlled ipa/ipc = 1 independent of v, Es, D0 k° > 0.3v1/2 cm/sec

52 2.933. QUASI-REVERSIBLE REACTIONS

For quasi-reversible reactions ia and ic contribute to the current measured in the overpotential range where mass transfer effects are not important. While opposing charge transfer reactions must be considered, a noticeable activation over potential is required to drive the reactions.

Peak heights for the anodic and cathodic reactions at the surface of the electrode are also an indication of reversibility. Over a small potential range and with a reasonably fast scan rate, the same ions which are first oxidised are then reduced at the surface of the electrode. Generally, for a given system, if the anodic and cathodic peaks are similar in size and occur at very close potentials then the system is said to be quasi reversible.

0.00003 < k° < 0.02

ip is not proportional to v1'2 such that it shows electron transfer kinetic limitations where the reverse reaction has to be considered

0.3v1/2 > k0 > 0.00002v1/2 cm/sec

Note that any system may appear reversible, quasi-reversible or totally irreversible depending on the time scale of the demands made on the charge transfer kinetics.

53 Fig 2.8. Cyclic Voltammogram corresponding to (a) Reversible (b) Quasi- Reversible (c) Irreversible Systems

54 2.9.4. Instrumentation Used for Cyclic Voltammetry

The electrochemical cell used for cyclic voltammetry involves a three electrode system. The counter electrode provides the current that is needed at the working electrode. This way nearly no current flows through the reference electrode and its potential remains constant. This three electrode cell also helps to minimise voltage errors due to ohmic loss through the solution by placing the reference electrode close to the working electrode surface. The instrumentation required for cyclic voltammetry is illustrated in Figure 2.9. and includes a potentiostat, current to voltage converter, and a recorder.(65)

Some of the things which will influence the current produced at an electrode surface are: electrode surface area, reactant concentration, temperature, viscosity, scan rate, and applied potential. Usually all of these properties are kept constant so that the current is proportional to concentration. Current interference can come from the electrolysis of impurities, solvents, and the electrode itself.

WAVEFORM POTENTIOSTAT

controlled

CURRENT TO VOLTAGE CONVERTER

Eiectrode designation: O—working, |— auxiliary, reference.

Figure 2.9. Instrumentation for Cyclic Voltammetry

55 CHAPTER 3 SUMMARY OF LITERATURE REVIEW AND RESTATEMENT OF OBJECTIVES

For optimum performance of the molybdenum redox battery the electrolyte should have the following characteristics:

~ not decompose on contact with mild oxidants ~ not be reduced on contact with mild reductants ~ stable to decomposition over long periods of use ~ high solubility in each oxidation state to maximise energy density ~ be stable over the temperature range of -5°C to 40°C ~ high electrolyte conductivity for minimum efficiency losses ~ high electrochemical reaction rates for high energy efficiency ~ low permeation through the ion selective membrane

Published solubility data for molybdenum compounds indicate that molybdenum is soluble in different oxidation states, while an earlier study by Trinh (15)(refer to section 2.2) of a molybdenum redox cell, showed that such a system would be promising. This preliminary information was the starting point for this study.

The aim of this project was thus to fully assess the feasibility of a molybdenum redox cell by evaluating the electrochemical reversibility of the molybdenum redox couples in various supporting electrolytes so that the two most suitable couples could be chosen for the positive and negative half cell electrolytes. Other important factors that required evaluation and optimisation were:

56 Molybdenum solubility Molybdenum source Availability Cost Compatibility of electrolyte and salts Supporting electrolyte choice and concentration Electrolyte stability in various oxidation states to temperature, air, and light Study the various redox couples of soluble molybdenum salts Electrochemical activity at various electrode surfaces, especially carbon Electrode selection Selective membranes for battery operation Electrolyte additives Cell efficiencies via charge discharge tests

While a comprehensive study of each of these factors would be beyond the scope of the present thesis, their study should provide an overall indication of the feasibility of a molybdenum redox cell.

57 CHAPTER 4 EXPERIMENTAL

4.1. Chemicals and Materials

The following chemicals were employed in evaluating the molybdenum redox cell.

Molybdic acid (A.R. grade Aldrich & BDH) Sodium molybdate (VI) dihydrate(A.R. grade Ajax) Ammonium molybdate (A.R. grade Ajax) Molybdenum disulphide (Technical Grade) Dodeca phospho molybdic acid(A.R. M & B) Molybdenum trioxide (A.R. grade M & B) Molybdenum dioxide (A.R. grade Stern Chemicals) .

Paraffin Oil (Ajax)

Nitrogen gas (CIG)

Sulphuric acid 98% (A.R. grade Ajax) Hydrochloric acid 36% (A.R. grade Ajax) Nitric acid 70% (A.R. grade Ajax) Phosphoric acid 85% (A.R. grade Ajax) Acetic acid (A.R. grade Ajax) Potassium hydroxide (A.R. grade M & B) Sodium hydroxide (A.R. grade Ajax)

58 Nafion ion selective membrane (Dupont) Selemion ion selective membrane (Asaki Glass Co. Japan) AMV ion selective membrane (Asaki Glass Co. Japan) CMV ion selective membrane (Asaki Glass Co. Japan)

Additives :

Sodium Sulphate (A.R. grade M & B) Potassium Sulphate (A.R. grade Ajax) Ammonium sulphate (A.R. grade Ajax) Calcium sulphate (A.R. grade Ajax) Sodium citrate (A.R. grade Ajax) Tartaric acid (A.R. grade Ajax) EDTA sodium salt (A.R. grade Ajax)

4.2. Instruments

Princeton Applied Research - PAR Model 173 Potentiostat RDE 4 Pine Instruments Potentiostat RDE 3 Pine Instruments Potentiostat PAR Model 175 Universal Programmer Escort EDM-1133 Multimeter Computerised Electrochemical System (Model 270 PAR) X-Y recorder (Model D-72BP Riken Denshi) Stabilised DC power supply (Bremi BRS 41 Italy) DC power supply (GPR 1820H Goodwill Instruments Taiwan) Portable chart recorder(YEW type 3057) Magnetic pumps (EPDM MD10 Tokyo Iwaki Co. Japan)

59 4.3. EXPERIMENTAL PROCEDURES

4.3.1. Solubility Tests

4.3.1.1. Direct Dissolution

The saturation solubility of the various molybdenum salts was determined by the following procedure. Firstly a 1M solution was prepared of each of the compounds, molybdic acid, sodium molybdate, and ammonium molybdate, in 3M, 4M, 5M, and 7M sulphuric acid. Then gradually as the solid dissolved, more solid was added to the volumetric flasks. This was repeated until no further dissolution occurred, indicating that the saturation point had been reached.

The various molybdenum salts were first directly dissolved in the various electrolytes. The addition of the salt was done gradually by adding around 0.05 moles per litre per day until the solution was saturated. These solutions were then left to equilibrate for a month. This way a direct calculation of the solubility could be done. The solubilities of the saturated solutions in sulphuric acid solutions were then reconfirmed by using atomic absorption spectroscopy with standard solutions of known concentration.

60 4.3.I.2. Electrolytic Reduction and Dissolution

Electrolytic reduction and dissolution of the molybdenum salts was performed for two reasons. Firstly, for the solubility of the lower oxidation states (III, IV, V) to be estimated, a method was required to obtain these species. Secondly, it has been found that vanadium pentoxide has a very low solubility in dilute sulphuric acid but by electrolytic dissolution relatively stable supersaturated solutions of more than 2M can be prepared. Similarly, electrolytic dissolution was used to determine whether more concentrated molybdenum solutions could be prepared by this method.

The procedure for electrolytic reduction and dissolution involved the use of a PVC cell shown in Figure 4.1. The two cells were separated by an AMV ion selective membrane and lead or graphite plate electrodes were employed. A current density of 20mA/cnr was applied from a power supply. Either a solution of Mo(VI) can be reduced to Mo(III), or a quantity of molybdenum salt can be gradually introduced into the negative haf cell with nitrogen bubbling through the solution to keep the salt suspended as it dissolves and is reduced. The nitrogen also improves mass transport in the cell and prevents molybdenum concentration build up at the electrode surface as well as helping to prevent air oxidation of the Mo(III) formed.

When the molybdenum salt has been electrolytically dissolved and reduced to a lower oxidation state, it is then possible to reoxidise the solution to Mo(VI) by reversing the polarity of the cell. The appearance of a precipitate at any point during the electrolysis indicates that the saturation point of one of the intermediate species has been passed. After electrolysis the original volume is made up with distilled water to compensate for any water losses due to the decomposition or evaporation of water during the process.

61 Figure 4.1. Dissolution Cell Apparatus

Solution PVC CELL

Lead Electrodes

Gasket Membrane Gas Inlet

4.4. Analytical Methods for Molybdenum Concentration Determinations

It is important to be able to calculate accurately the concentration of the molybdenum species in solution. These tests should only be done after the solutions have had time to equilibrate. Methods of analysis used in this study include potentiometric titration, Atomic Absorption spectroscopy (AAS), and Inductively Coupled Plasma (ICP) emission spectroscopy. Other standard methods found in old text books involving complicated titrations, which are subject to interferences, are simply not as practical. (28,29)

62 4.4.1. Potentiometric Titrations

Any oxidation reduction system that gives rise to a reversible potential can be titrated potentiometrically. The potential of a solution changes as an oxidant such as potassium permanganate is added. This potential variation is described by the Nernst equation and arises from the change in the ratio of the oxidised and reduced forms of the analyte during the titration. Not only can the total molybdenum concentration be determined but this method also gives important information on oxidation states. The method is fast and overall simple.

The solution was reduced in several ways, electrochemically, by zinc and by aluminium. The result was not improved by any of these. The titrations were actually worse when zinc or aluminium were used. They seemed to interfere with the potentials of oxidation for the molybdenum.

All possible variables for the titration were tried including varying the potassium permanganate rate of addition, the speed of the pen on the chart recorder, the scale of the graph, the sample size, and different sample sizes diluted with both the supporting electrolyte and water.

Since molybdate polymers, which are highly negatively charged species, may be anions of relatively weak acids, and partially protonated anions will also be present, the measuring of potentiometric endpoints with the necessary precision is difficult as the end point is not sharp (42). As can be seen in Appendix 2. by the graphs of potential versus permanganate added, the results are inconclusive. The end points are not at all clear.

63 After exhaustive trials it was thus concluded that this method was not valid for determining the concentration of molybdenum when the oxidation state was not known. Similarly, it was not useful for determining the oxidation state unless the concentration was known.

Nevertheless, while this was the case, potentiometric titrations were found to be useful in two cases. Firstly, if the molarity of a solution was known, then from the total volume of titrant, the oxidation state of the solution could be calculated. Secondly, if the oxidation state of the solution was known then the molarity could be calculated. Examples of typical calculations are given in Appendix 2.

4.4.2. Atomic Absorption Spectroscopy

A second analytical method was needed as a backup to the potentiometric method so as to be able to compare and confirm results. AAS was selected since it is another convenient method of analysis. Care must be taken with samples, however because of the many interferences which are possible. (86) Also the samples must be diluted several thousand times (5000 in this case) and so care must be taken when diluting the samples to minimise the errors.

64 As with many other elements, molybdenum analysis can be affected by the solvent, complexing agents, other ions present, pH, and the type of mineral acid used. To overcome the interferences, especially caused by sulphate ions from the sulphuric acid, the standards prepared were treated in exactly the same manner as used to treat the unknown sample. The samples were diluted from 4M H2SCh with distilled water to keep the pH relatively constant and the sulphate ion concentration low. The wavelength used for molybdenum was 313.3nm. (13)

Atomic Absorption Spectroscopy (AAS) was used to measure the concentrations of molybdenum in the sulphuric acid solutions made electrolytically. It is useful as it does not matter what oxidation state the molybdenum is in.

The concentration range of the standards prepared was 10 to 30 milligrams per litre (PPM) as recommended in the AAS suppliers’ handbook. As can be seen in the standard graph of concentration versus absorption, a linear response is obtained at least up to an absorbance of 0.26 corresponding to a molybdenum concentration of 30PPM (Figure 4.2). This confirms that Beers law (which states absorbance equals the product of the absorptivity constant, path length, and the concentration of the species absorbing) is obeyed within this concentration range.

65

of or for

various 10ml

together solutions

in

positive

The the

mixing mixed

Absorption example by

solution. were (SOCs). For vs

4 simulate

to S0 2

States

H accurately ratios.

charge

and

Mo(IV) in

Curve

of of

most

Molybdenum

Oxidation Mo(V)

30ml states different

Mo(VI)

PPM in

in and

and 66 prepared

various Different

Concentration resulted at of were

Mo(IV)

solutions Mo(III)

of

of

charge Mo(lII)

Solutions

of of Mo(III) electrolyte Standard Molybdenum

of solutions

solutions cell

and

20ml states

half obtain and

4.2.

to

prepared Mo(VI) Preparation

0.050 0.100 0.150 0.200 0.250

0.300 different

Figure 4.5. negative ratios of

Mo(VI) the Absorban The 4.6. Temperature Stability Tests

Solutions of various molybdenum compounds were studied in various concentrations of sulphuric acid at different temperatures to determine the optimum ion concentration and solution composition for different applications.

The solutions were sealed and placed in temperature baths at temperatures of 20, 30, 45, and 60°C, then left for seven days to equilibrate, before samples were taken and analysed for molybdenum concentration. The filtered samples were then analysed by AAS. It was also important to record the time when the first precipitate was observed, if it takes six months for a precipitate to form, then obviously this result would not be as serious in the redox cell as one where the solution showed a precipitate after just six hours.

4.7. Cyclic Voltametric Studies

Many supporting electrolytes (of different concentrations), electrode materials, additives, and molybdenum salts, were screened to identify the optimum conditions for a molybdenum redox cell. The primary criteria required for a good combination of electrolyte and electrode material, are significant solubility and electroactivity.

67 For the cyclic voltametric studies a 50ml beaker was employed as the electrochemical cell, containing around 20ml of solution. The cell was covered by a teflon lid which had three holes in it for each of the electrodes. The three electrodes were connected to a potentiostat and the resulting cyclic voltammograms were recorded on an X-Y recorder (refer to Figure 2.9). The scan rate was kept constant at 4 V/minute for these studies. The open circuit voltage of the solution measured between the reference and working electrodes was the starting potential for the scans as no reaction takes place at this potential. Unless otherwise specified, all cyclic voltammograms were recorded at room temperature and all potentials were measured against an SCE reference electrode.

4.7.1. Electrodes Used for Cyclic Voltammetry

4.7.1.1. Counter and Reference Electrodes

The counter electrode was an ultra pure, high grade graphite rod, 6mm in diameter. The reference electrode used was a standard calomel electrode (SCE).

4.7.I.2. Working Electrodes

With the molybdenum solutions, working electrodes of platinum, palladium, lead, bismuth and carbon were investigated.

68 Graphite electrodes were made by bonding a piece of graphite rod about 15mm long to a stainless steel rod with a three parts to one silver epoxy glue. The end 30mm was then encased for insulation by an epoxy resin. Once the epoxy dried the graphite was wax impregnated to fill the pores. After polishing, the graphite should conduct with the stainless steel rod very well with less than one ohm resistance between them. This also leaves a small known surface area of graphite to be used as the working electrode. A similar method was used to fabricate working electrodes of Pt, Pd, Pb, and Bi. In the case of the graphite working electrode, the method was improved on by eliminating the stainless steel rod and having a continuous graphite rod insulated by epoxy to eliminate the contact resistance thereby reducing the overall resistance of the electrode.

Figure 4.3. Working Electrodes for Cyclic Voltammetry

Steel Rod

Graphite

69 4.8. Molybdenum Redox Flow Cell Performance Tests

4.8.1. Flow Cell Components and Setup

The cell apparatus employed two graphite felt electrodes either pressed against a graphite plate or bonded to a conducting plastic which was reinforced with a bronze mesh on the back to allow easy current collection by a copper plate. The electrode surface area was 25cm2. The membranes evaluated were CMV, AMV, Selemion, and Nafion. The total volume of 1M molybdenum solution was 75ml in each half cell. The charge discharge current density was kept constant at 20-40mA/cm2. The upper and lower voltage limits were between 1.2 and 0.3 volts respectively. The tests were carried out at ambient temperatures with 4 molar sulphuric acid as the supporting electrolyte.

+

Copper Membrane _ Current Collector Carbon Fibre Electrode

Graphite Plate

Negative Positive Half-Cell Half-Cell Electrolyte Electrolyte

Figure 4.4. Redox Flow Cell Apparatus

70 During operation the cell was started with both half cells containing either a solution of both 0.5M Mo(IV) and 0.5M Mo(V) (Mo(4.5+)) uncharged state, or charged solutions of Mo(VI) and Mo(III). The pumps are switched on circulating the solutions through the cell with the current controller automatically and continuously controlling the charging and discharging currents. A two pen chart recorder was used to graph the cell current and the cell voltage as a function of time during the charge and discharge cycles.

The following reactions occur when charging a Mo(4.5+) solution.

At the Positive electrode : Mo(IV) = Mo(V) + e' -

followed by Mo(V) = Mo(VI) + e'

At the Negative electrode : Mo(V) + e' = Mo(IV)

followed by Mo(IV) + e‘ = Mo(III)

A range of readily available electrodes and membranes were evaluated in the above test cell. In order to compare the various membranes the same electrode material Sigri carbon fibre felt was employed. When comparing the different electrode materials AMV membrane was employed throughout the experiments.

71 4.8.2. Cell Electrode Fabrication

Carefully controlled fabrication of the electrode is required to minimise resistances associated with poor bonding of the graphite felt to the conducting plastic. The type of polymer and felt used will make a difference. A specially formulated conducting plastic material was employed as the substrate (74). The granulated polymer was spread out into a mould. Typically, the press temperature is raised to 180°C, with no pressure for 20 minutes. Then at 180°C 40 ton of pressure is applied. First the bronze mesh is placed underneath the substrate at 40 ton pressure at 180°C for 20 minutes. Then a mould with a window is placed on top so that the graphite felt can be pressed into the polymer and bonded with only a slight pressure for 20 minutes. The electrode is then left to cool.

4.8.3. Cell Resistance Calculation

The lower the resistance of the cell the higher the voltage and overall efficiencies. The electrical resistance of each cell component can be measured individually, for example the resistance of the copper current collector, or the dry carbon felt electrode. Alternatively the total electrical resistance of the assembled cell excluding the membrane, can be measured.

All of the resistances are measured in similar ways by varying the applied current and measuring the potential between two points. Then the resistance (R) can be calculated simply by plotting this current versus voltage and then R = (slope)/Area.

72 To measure the overall cell resistance which is the sum of the ohmic and polarisation resistances, the cell is assembled with a membrane in place and solutions corresponding to 50% SOC (Mo(VI)/(V) and Mo(III)/(IV) ratios of 1) are circulated through the positive and negative half cells respectively. A constant current is applied for one minute to charge the cell and the voltage recorded. Then the cell is discharged for one minute and the voltage is recorded. This procedure is repeated for a range of charge and discharge current densities. The results are plotted as polarization plots similar to that shown in Figure 4.5. The slope of each plot gives the cell resistance for the charge and discharge processes respectively at 50% SOC.

Discharge Charge

Figure 4.5. Typical Polarization Plot for Calculating Cell Resistance

73 4.8.4. Efficiency Gilculation

The empirical calculations given below in association with Figure 4.6 are valid when a constant current is used for both charging and discharging as was the case in this study. Note that c and d are measured by first finding the midpoint of the curves. If the charge and discharge curves are symmetrical, the mid point represents the average charging and discharging voltage respectively. A limitation exists when the curves are not symmetrical however, making this assumption invalid. In general however, symmetrical charge and discharge curves are obtained at low current densities.

Charge

> '—* (1.1 o < _/i- o >

Time (h) Figure 4.6. Charge - Discharge Curve

Coulombic Efficiency: Ec = a/b 4.1

Voltage efficiency: Ev = c/d 4.2

Energy Efficiency: E = Ec x E* 4.3

where ^discharge ^discharge charge^charge J Vdischarge charge

74 CHAPTER 5 OBSERVATIONS, RESULTS, AND DISCUSSION

5.1. Solubility Studies of Molybdenum Compounds

5.1.1. Solubility Studies of Mo(VI) Compounds

The solubility of the molybdenum compounds is important because the higher the concentration the greater the energy density of the electrolyte, which also means a smaller total volume required for a given capacity, reduced storage tank costs and reduced weight.

The maximum working solubility needs to be established. That means that all conditions experienced by the solution should not affect its ability to perform as a redox cell at maximum capacity and efficiency. Obviously, for redox battery application, the redox species should not precipitate in any of the oxidation states which are likely to occur during charge or discharge of the cell.

The maximum solubilities of the Mo(VI) compounds molybdic acid, disodium molybdate, and ammonium molybdate, in 4M sulphuric acid, were measured directly by salt addition until saturation at room temperature, and are shown in Table 5.1. These results are confirmed by the results reported by Kummer and Oei (5). Published data is listed in Tables 5.4 and 5.5 for comparison. It was noted that the sodium molybdate dissolved much faster than the other two salts. This is significant when making large volumes of solution. The experimental solubilities of these three salts in sulphuric acid concentrations of 3, 4, 5, and 7 moles per litre are shown in Figure 5.1.

75 Table 5.2 is a list of experimental solubilities of molybdic acid in various electrolytes. Concentrated hydrochloric acid gives the highest solubility for molybdic acid. Table 5.3 is a list of various molybdenum salts in 4M sulphuric acid and shows that their solubilities are not as great as the salts listed in Table 5.1.

Figure 5.1. Experimental Mo(Vl) Solubility vs Sulphuric Acid Concentration

Molybdenum (VI) Solubility

Sulphuric Acid Concentration

_H_ Ammonium Molybdate Molybdic Acid Sodium Molybdate

76 Table 5.1. Experimental Saturation Solubilities of Mo(VI) Salts in 4M H2S04 at Room Temperature

Salt' Molybdenum (VI) Solubility mol/1

H2Mo04 1.3

(NH4)6Mo7024.4H20 1.4

Na,Mo04.2H20 2.1

Table 5.2. Experimental Solubility of Molybdic Acid in Various Electrolytes

Electrolyte Solubility mol/L

Hydrochloric Acid (11M) 1.6

Phosphoric Acid (4M) 0.5

Nitric Acid (4M) 0.9

Acetic Acid (4M) 0.4

Sodium Hydroxide (1M) 0.3

Potassium Hydroxide (1M) 0.3

77 Table 5.3. Experimental Solubility of Various Molybdenum Salts in 4M Sulphuric Acid

Salt Solubility mol/L

Molybdenum 0.1 Disulphide

Molybdenum Trioxide 1.1

Molybdenum Dioxide 0.2

Phosphomolybdic Acid 1.0

Molybdenum in a solution of concentrated phosphoric acid forms the complexes Mo02(P04)' and Mo02(P04)>4', and in a dilute solution of molybdic acid the protonated species HMo02P04 and H4Mo02(P04)2 form. (9)

Table 5.4. Published Solubility Data for Comparison

Salt Mo03.2H20 Na2Mo04 H2Mo04

Temp 20°C 20°C 20°C

mol/L 0.008 1.6 1.7

g/lOOml 0.138 39.38 28.0

Solvent water water 11N HC1

Reference 2 2 4

78 Table 5.5. Published Solubility Data

Solvent Temp H2Mo04 Ref

mol/1

11N HC1 20°C 1.7 4

4N HN03 20°C 0.95 4

8N H2S04 20°C 1.38 4

6N H2SO., 20°C 0.86 5

8N H2S04 20°C 1.24 5

6N H2S04 50°C 0.56 5

As discussed in the Gmelin Handbook B 3b, the cation has a real influence on solubility. In fact it is suggested that the species differ with each cation. (9) The degree of molybdenum polymerisation increases with the hydrogen ion concentration until precipitation occurs (pH 1.1-1.8), but then dissolves again at higher acid concentrations. (4)

79 The solubility in this study for molybdic acid in 11N HC1 as shown in Table 5.2. is lower from that quoted in the literature (refer to Table 5.5 and Figure 5.2). It was found that initially a more concentrated solution could be formed in HC1 but that it was not stable. The rate of precipitation was quite remarkable once the solution was seeded. This is one reason why sulphuric acid was chosen as the preferred electrolyte over HC1, and as will be seen later it is not the only reason when electroactivity is considered.

Kirk-Othmer states that SOA2' complexes are comparatively stable despite the reactions being reversible. This is one good reason for sulphuric acid to be used as the major electrolyte in the molybdenum redox system.

The equilibrium solubility, quoted from reference 4 and illustrated in Figures 5.2 and 5.3, were determined at 20°C with daily periodic stirring of the solutions for two hours. The time for establishing equilibrium was determined from the equilibrium curve. With an increase in the acidity from 1-1 IN HC1 the solubility of the molybdic acid went from 3.6 to 260 g/L. The equilibrium was reached after 15 days as shown in Figure 5.3.

At low acidities 2-5N FIC1 the solubility of molybdic acid at 20°C is higher than at 50 and 100°C. With an increase in the acidity of the HC1 above 6N the solubility is higher for the two higher temperatures (refer to Figure 5.4). The higher solubility of molybdic acid in HC1 in the colder state is probably explained by partial formation of colloids. When these solutions are heated, coagulation of the colloidal species takes place with precipitation of molybdic acid, this phenonemon is also known as thermal precipitation (4).

80 The solubility curve shown in Figure 5.3 for molybdic acid in nitric acid has a maximum in 6N nitric acid. At equilibrium at 20°C the maximum solubility of molybdic acid in nitric acid is 14.3 g/L. Equilibrium is reached after 18 days. Figure 5.4 shows that in 4N nitric acid at 20°C and at 100°C the initial solubility drops from 154.0 to 7.6 g/L.

Equilibrium, according to Shapiro, of molybdic acid in sulphuric acid at 20°C is reached after 20 days. The solubility curve shown in Figure 5.3 has a maximum of 224 g/L in 8N sulphuric acid. Figure 5.4 shows that with an increase in temperature the solubility decreases and the solubility maximum is shifted to a higher acid concentration. (4)

Kummer (5) also studied the solubility of molybdic acid in sulphuric acid, and confirms the results presented. For a 3M sulphuric acid solution the solubility is given as 140 g/L. The solubility decreases with increase in temperature and is 90 g/L for this 3M solution at 50°C.

Acid Normality-

Figure 5.2. Isotherms of the Solubility of I I2MoO^ in 1) FIC1, 2) H2S04, and 3) IIN03 at 20°C (ref4)

81 Time (Days)

Time (Days)

Time (Days)

Figure 5.3. Effect of Time on the Establishment of Equilibrium in I-I2MoO„at 20°C: a)HCl, b) IIN03, c) II2SO„ Acid Concentrations N: 1)3, 2) 6, 3) 9, 4) 11 (ref4)

82 HCL (N)

HN03 (N)

Figure 5.4. Effect of Temperature on the Solubility of H2Mo04: a) MCI, b) HN03, c) H2SO<, Temperature °C: 1) 20, 2) 50, 3) 100. Dissolution time 4hrs. (ref4)

83 5.1.2. Solubility of Electrolytically Generated Mo(VI) Solutions

Solutions of Mo(VI) were also prepared electrolytically by first reducing a suspension of the relevant Mo(VI) salt to Mo(III) and then reoxidising it to the Mo(VI) state. The Mo(VI) solutions prepared by this procedure were always blue however, suggesting the presence of the Mo(blue) species.

Saturation solubility results, in moles per litre, of these Mo(blue) solutions in (4M H2SO.,) made electrochemically, are shown in Table 5.6. as measured by AAS and confirmed by ICPES.

Table 5.6. Saturation Solubility Results of the Mo(blue) Solutions Made Electrolytically

Salt Solubility mol/L

H2MoO, 1.0

(NH4)6Mo702,.4H20 1.3

Na,MoO.,.2H20 1.6

84 The most important point is the contrast between maximum solubilities of the solutions produced by the two different methods as shown in Tables 5.1 and 5.6. The addition of Mo(VI) salts to an acid solution results in a solution of Mo(VI) which has a much higher saturation concentration than the solution made by electrochemical reduction and re-oxidation.

The reason for this difference in solubilities therefore is that during oxidation of Mo(III) or Mo(IV) solutions, the stable Mo(blue) intermediate is formed. The solutions produced by the electrochemical method give erroneous solubility values therefore since they are not 100% Mo(VI) solutions but contain the molybdenum blue mixed valence species. This will be described in more detail in the following section. Overall therefore it is not the Mo(VI) species which is concentration limiting but the molybdenum blue species. This would have significant consequences for the operation of a molybdenum redox cell. This mixed valence species, while being able to be oxidised and reduced, will precipitate during cycling on the positive half side of the cell if the molybdenum concentration exceeds the saturation solubility of this species. This would lead to blockages in the solution channels, thus debilitating the pumps and reducing the flow of electrolyte through the electrode felt.

The solubilities of the electrolysed solutions containing molybdenum blue are the critical solubility values for this application as these species will exist during cycling of a cell.

85 5.1.3. Molybdenum Solution Reduction Observations

Molybdenum solutions of different oxidation states were prepared by electrolytic reduction of various Mo(VI) salts.

The theoretical time required for electrolytic dissolution can be calculated from the following equation.

t = mnF/I

t = time seconds m = moles of electroactive material F = Faraday’s constant n = number of equivalents per mole of molybdenum to achieve the desired oxidation state I = electrolysis current (amps)

The actual electrolysis time required was longer however due to the fact that the current efficiency of the process is less than 100%. During electrolytic dissolution the reactions taking place would be:

At the anode : FLO = 2H+ + 1/2 02 +2e'

At the cathode : Mo(VI) + e‘ = Mo(V)

86 Mo(V) + e* = Mo(IV)

Mo(IV) + e- = Mo(III)

Note though that there are intermediate complex species containing more than one oxidation state during electrolysis. It has been shown that Mo(VI) reduced to Mo(III) results in a dimeric product Mo,(III). Instead of Mo(VI) reducing in a single electrolysis it has two steps such that Mo2(III) is formed. (11)

2Mo(VI) + 2e‘ = Mo2(V)

Mo>(V) + 4e- = Mo2(III)

It was found that the clear Mo(VI) solutions made from molybdic acid, disodium molybdate, and ammonium molybdate when reduced to the lower oxidation states changed colour. The colour for the Mo(III) species was distinctly dark green. The colours for the other two oxidation states were brown/orange for Mo(V) and dark red for Mo(IV). These observations are summarised below in Table 5.7.

Note also that the Mo(blue) species occurs between the Mo(V) and Mo(VI) oxidation states. This species is an intense blue.

87 Table 5.7. Colours of the Molybdenum Oxidation States

Oxidation State Colour

cm) Dark Green

(IV) Dark Red

(V) Brown/Orange

(V)/(VI) Dark Blue

(VI) Clear Yellow

The unbalanced equations of the proposed molybdenum redox system are most probably :

Mo(III)/(IV) : Mo3(OH)4 5+ = Mo30(OH)3 5+ + H+ + e = [Mo304(H20)9] 4+

Mo(VI)/(V) : H2Mo04(aq) + 2H+ + 2e‘ = Mo02+ + H20 = [Mo204(H20)6] 2+

(i)

88 It was found that Mo(III) prepared by reduction in the cell described in Figure 4.1, regardless of the parent Mo(VI) salt used, was soluble to concentrations greater than two molar in 4M sulphuric acid. From these observations it has been found that Mo(VI) is more stable than Mo(III) but it is less soluble. More stable again is the Mo(blue) divalent species to which Mo(VI) is readily converted.

Mo(VI) was found to dissolve very slowly and the white powder (of molybdic acid, sodium molybdate, and ammonium molybdate) dissolves to give a clear light yellow solution. However upon sitting for a few days to a few weeks the solutions gradually turn blue which is characteristic of the molybdenum blue species. Obviously the molybdenum blue being produced is due to some partial reduction to Mo(V) or a mixed Mo(V)/Mo(VI) species.

It has been reported that impurities including dust are capable of reducing the Mo(VI) to an intense blue Mo(V)/(VI) complex. The occurence of this partial reduction to molybdenum blue is known to also be light catalysed in the presence of organics, and other impurities such as phosphates (7). In this case impurities are most likely to be responsible for the Mo(blue) formation in the yellow Mo(VI) solutions.

Mo(IV) and Mo(V) are less easy to distinguish than Mo(VI) and Mo(III) due to their similar dark colours. Furthermore, the intensity of the colours change with changing acid and changing molybdenum concentration. For this reason colour is not the most reliable method for distinguishing oxidation state.

89 A suggested reaction for the "foul smell" referred to here is given as: S042 + 4H+ + 2e _ = H2S03 + H20 H2S03 + 4H+ + 4e~ = S + 3H20 S + 2H+ + 2e = H2S

Mo(III) solutions were found to produce a foul odour after a few days, even if kept under nitrogen. The odour was like that of hydrogen sulphide. Decomposition of the sulphuric acid may be responsible for this odour which would not be desirable if the cell were to be operated without being closed to atmosphere.

It was reported that Mo(III) in dilute HC1 was oxidised by hydrogen ions even under oxygen free conditions. The Mo(III) is slowly hydrolysed to a more reactive species which can then be oxidised by H+. Perhaps a similar reaction happens in the sulphuric acid solutions. (25)

It is also noted that the higher the concentration of molybdenum the more likely is the solution to turn a deeper blue. It was found that all solutions turned at least a pale blue. After electrochemical reduction the molybdenum solutions were electrochemically reoxidised to Mo(VI) to check if saturation had been reached. The blue colour only disappeared to form the yellow solution after several hours of extra electrolysis at a current density of 20mA/cm2 for a 50 ml aliquot.

Despite the blue colour however, the half cell potential of the blue solutions was about the same as that of the original yellow solutions. This suggests that Mo(VI) is the corresponding predominant species and that only a small proportion of it had formed the molybdenum blue complex. It should be noted however that in general the potential (measured between a reference electrode and a graphite or glassy carbon electrode) is often quite insensitive to slight changes in condition variables.

90 Even a solution of only 1M molybdic acid showed signs of instability on standing. Mo(III) is readily oxidised by air resulting in a red solution. More detailed studies of Mo(III) stability in air are presented in Section 5.1.4. One molar solutions of Mo(IV) and Mo(V) (from molybdic acid in 4M sulphuric acid) were prepared by mixing appropriate ratios of Mo(III) and Mo(VI). Mo(V) is oxidised by air to molybdenum blue. Also Mo(IV) is oxidised from its red colour to a more brown colour indicating oxidation by air. It was found that Mo(4.5+) (being 50% of each Mo(IV) and Mo(V)) precipitates significantly over a period of a month at just one molar concentration. The stability of the molybdenum solutions would thus be a problem at higher concentrations than one molar.

5.1.4. Mo(III) Stability to Atmospheric Oxygen

This study was performed to observe the stability of a Mo(III) electrolyte if exposed to atmospheric oxygen. The conditions of a redox half cell were simulated in a 100ml volumetric flask and the oxidation state was monitored with time.

To simulate the agitation of the solution in a redox system, which is caused by the pumps, the solution was stirred with a magnetic stirrer. This action exposes the solution to the air. A volumetric flask was chosen as the vessel since the narrow neck does not enhance ventilation and convection of air to the solution surface.

91 75ml of a 100% Mo(III) solution produced by the reduction of one molar sodium molybdate in 4M sulphuric acid solution, was used for the experiment. Samples of the solution were taken at 10 minute intervals and analysed by potentiometric titration.

open

100% Mo(lll)

------magnet

magnetic stirrer

Figure 5.5. Mo(III) Oxidation Apparatus

As seen in Figure 5.6, the amount of oxidation was significant over a short time period. Within 90 minutes the concentration of Mo(III) had dropped from 100% to 67%. At this time the rate of oxidation decreased, so that after a total of 580 minutes, the solution was 57% Mo(III).

It has been reported that Mo(III)2 in HPTS (p-toluenesulphonic acid) is oxidised by oxygen (and other oxidants) to the Mo(V)2 dimer, Mo2042+. In this electrolyte it is claimed that Mo(IV) does not react with oxygen (35). More details for this oxidation are given in a paper by Hills(46), and the results presented agree with those of the previously mentioned author. The hexa-aqua molybdenum (III) species with oxygen in HPTS gives the di-oxo-Mo(V) ion [Mo20.,(H20)(l]2+ according to the following reaction sequence which shows intermediate species.

92 Mo(III) + 02 < = = = > MoO,3+

Mo023+ + Mo(III) < = = = > Mo02Mo6+

Mo02Mo6+ + 2H20 < = = = > MoA2+ + 4H+ overall : 2Mo(III) + 02 + 2H,0 < = = = > Mo2042+ + 4H+ (46)

The reaction to Mo(V) is a two electron reaction, further oxidation to Mo(VI) is slow. It is concluded that stable monomeric and dimeric Mo(III) ions are unlikely. (46,47,54)

Richens and Sykes explains that at high acid concentrations (>4M) Mo(III)3 is electrochemically oxidised firstly to a mixed valent species namely

Mo(III)2Mo(IV). (30) This shows that the chemistry of molybdenum is highly influenced by the method of reduction or of oxidation. Cyclic voltammetry suggests minor structural changes prior to the formation of Mo(III)3.

In 2M HPTS (30):

Mo(IV)3 + 2e‘ < = = = > Mo(III)2Mo(IV) -0.10V vs NHE

Mo(III)2Mo(IV) + e' < = = = > Mo(III)3 -0.18V vs NHE

93 110

100 II-

100 300 Time (Minutes)

Figure 5.6. Molybdenum (III) Oxidation Rate

5.2. Thermal Stability Studies of Mo Electrolytes

Thermal stability studies of Mo(VI) and Mo(III) were performed at various temperatures. Solutions of sodium molybdate, molybdic acid, and ammonium molybdate were reduced to the Mo(III) oxidation state and exposed to constant temperatures of 25, 30, 45, and 60°C as were the original Mo(VI) solutions. The solutions were placed in a water bath, at concentrations below that of their room temperature saturation points, that is initial concentrations of 1.5, 1.0, and 1.3M respectively.

94 Precipitates were observed in the Mo(VI) solutions after just one day, even at the relatively low temperature of 30°C. The precipitate was always blue despite the starting Mo(VI) solutions being clear yellow. The stability of Mo(III) is not such a problem at the low temperatures however, some precipitation was observed at the higher temperatures. As the Mo(III) solutions are very dark green, it was difficult to determine exactly when the precipitation began although it was suspected that this occured after 2-3 days. All solutions were equilibrated for at least seven days.

Figure 5.8. shows that while the solubilities dropped with increase in temperature, for Mo(III) the solubility levelled off above 30°C. In most cases the Mo(VI) concentration levelled off after 45°C as shown in Figure 5.7. The results in Figures 5.7 and 5.8 also show that when the solutions which had been kept at 60°C for seven days, were allowed to re-equilibrate at room temperature, the Mo(III) salts mostly redissolved, but the Mo(VI) salts did not.

This experiment was duplicated to ensure the phenonemon was real. The cited literature value (see Table 5.5) for the solubility of Mo(VI) at 50°C confirms the thermal precipitation observed although it had not been previously quantified to any extent.(3)

The important point to note is that some precipitation will occur at these concentrations if the temperature fluctuates by more than 10°C. The Mo(VI) solutions form the molybdenum blue species which does not completely redissolve on cooling to room temperature. The molybdenum blue mixed valence species is likely to be the most thermodynamically stable species and this explains why it is formed so easily.

95 This instability was further confirmed after examining the electrolyte of a cell which had been subjected to 12 charge - discharge cycles especially if the molybdenum concentration was greater than one molar. Especially noticeable was the positive half cell electrolyte which stained with a dark blue film everything that it came in contact with. This included pipes, joints, reservoir, and even the membrane.

As described later, to avoid the problem of the molybdenum blue species clogging the system it was important to not fully charge the electrolyte. Studies were also undertaken to find a method to suppress its formation.

Temperature °C Ammonium Molybdate Molybdie Acid — Sodium Molybdate

* Footnote: The second set of readings at 20°C resulted from leaving the solutions which had been held at 60°C to cool and equilibrate at room temperature. Figure 5.7. Temperature Stability of Mo(VI) Salts

96 6 1.3

Temperature °C H_ Ammonium Molybdate Molybdic Acid -J>— Sodium Molybdate

* Footnote: The second set of readings at 20°C resulted from leaving the solutions which had been held at 60°C to cool and equilibrate at room temperature.

Figure 5.8. Temperature Stability of Mo(III) Salts

97 5.3. Cyclic Voltammetry Studies

Further screening of suitable supporting electrolytes, sulphuric acid concentration, electrode composition, and molybdenum source was undertaken using cyclic voltammetry. A typical cyclic voltammogram obtained at a graphite electrode in a solution of 1M Mo(VI) in 3M H2S04 is shown in Figure 5.9. From the peak currents, and peak potentials of the anodic and the cathodic curves, qualitative and quantitative kinetic information is obtained and used to determine the kinetic constants for the reactions as detailed in section 2.9.2.

5.3.1. Effect of Acid Concentration

The one molar solutions of molybdic acid were studied by cyclic voltammetry in 3, 4, 5, and 7M sulphuric acid solutions at a graphite electrode and the curves are shown in Figures 5.9. to 5.12. for a scan rate of 4V/min. While the principal peaks look similar, important differences in peak heights and separations can be observed with varying acid concentrations.

The cyclic voltammograms show the complex nature of molybdenum under these conditions. There are many more small peaks, indicating a large number of reactions. This indicates that there are many equilibria between the ions in solution, forming different complexes which are reduced or oxidised at different potentials.

98 The variations observed with changing acid concentration arise because of the different pH of each of the solutions. As reported in the literature on the chemistry of molybdenum (refer to sections 2.4 and 2.5), the different acid concentrations leads to the formation of different complexes. The various species in solution will give different peaks at different potentials, due to variations in equilibrium potentials. The sulphate ions form different complexes with molybdenum depending on the hydrogen ion concentration (refer to section 2.4).

In choosing the optimum acid concentration, and cell conditions, it is important that the peaks corresponding to the charging reactions for Mo(VI) and Mo(III) are away from the oxygen and hydrogen evolution reactions respectively. Operating a cell near these potentials would lead to poor coulombic efficiencies due to gassing side reactions during charging.

The approximate anodic and cathodic peak potentials at a graphite electrode as identified from their positions in the cyclic voltammograms, are listed in Table 5.8.

99 Figure Current (mA) Figure

5.10. 5.9.

from Cyclic from Cyclic -0.5

Molybdic

Molybdic

Voltammogram Voltammogram

Acid, Acid, 100

in in

at

at

4M 3M

a a

Graphite Graphite II HjSO^, 2 S0 4 ,

Scan Scan

Electrode Electrode

Rate Rate

4V/min

4V/min

in in

1M 1M E

Mo(VI)

Mo(VI) vs

SCE

E

vs

SCE E vs SCE -0.5.

Figure 5.11. Cyclic Voltammogram at a Graphite Electrode in 1M Mo(VI) from Molybdic Acid, in 5M II2S04, Scan Rate 4V/min

E vs SCE

Figure 5.12. Cyclic Voltammogram at a Graphite Electrode in 1M Mo(VI) from Molybdic Acid, in 7M II2S04, Scan Rate 4V/min

101 Table 5.8. Average Half Cell Reaction Potentials as Measured by Cyclic Voltammetry at a Graphite Electrode for 1M Mo(VI)

cv Reaction Voltage Peak (vs SCE)

A Mo(III) —> Mo(IV) 0.0

B Mo(IV) —> Mo(V) 0.7

C Mo(V) — > Mo(VI) 1.6

D Mo(VI) — > Mo(V) 1.0

E Mo(V) — > Mo(IV) -0.1

F Mo(IV) —> Mo(III) -0.7

As can be seen in Figures 5.9 - 5.12 however, most of these major peaks actually consist of several multiple peaks. In Section 6 these peaks are dealt with in more detail, and as will be shown later, the peaks labelled here as A and F for example, actually consist of two separate couples.

102 At a constant scan rate the peak separations and peak heights vary for each solution. Comparing Figures 5.11 and 5.12 for the 5 and 7M solutions, it is seen that the peaks are separated more distinctly than for the 3 and 4M sulphuric acid solutions. In all of the solutions, the peak C corresponding to the oxidation of Mo(V) to Mo(VI), is close to the oxygen evolution reaction. The peaks in 5 and 7M sulphuric acid are also shifted to slightly more positive potentials, but otherwise they are all similar in shape. The important feature to note however, is that peak D, corresponding to the reduction of Mo(VI) to Mo(V) is not clearly evident in all cyclic voltammograms suggesting that Mo(VI) is reduced straight to Mo(IV) which then reacts with the remaining Mo(VI) to form Mo(V).

Peak heights are also of considerable importance. The greater the peak height the greater is the current produced by the half cell reaction. The relative peak ratios vary for each solution and so comparison is made difficult.

Various other electrolytes were screened by cyclic voltammetry including KOH, NaOH, acetic acid, perchloric acid, hydrochloric acid, and nitric acid as seen in Figures 5.13 to 5.18, however, the response is poor for most electrolytes except nitric acid. This indicates that the molybdenum species formed in these supporting electrolytes do not have acceptable electrochemical reversibility. Acetic acid has a relatively poor conductivity and this is reflected in the cyclic voltammogram by drawn out peaks.

103 Figure Current (mA) = Current (mA) c

5.14. 5.13.

Mo(VI) Cyclic Mo(VI) Cyclic

Voltammogram Vollammogram

from from

Molybdic Molybdic 104

at

at

Acid Aeid

a a

Graphite Graphite

in in

1M 1M

NaOII,

KOII,

Electrode, Electrode,

Scan Scan

in

in Rate Rate

0.1 E 0.1M E

vs

M

4V/min vs 4V/min

SCE

SCE Figure Figure

5.16. 5.15. Current (mA)

Mo(VI) Cyclic 4V/min Mo(VI) Cyclic

Voltammogram Voltammogram

in from

4M

Molybdic Perchloric 105

at at

Acid

a a Acid,

Graphite Graphite

in

Scan 1M

Acetic

Electrode, Rate Electrode,

4V/min Acid,

E in

in Scan

vs

E 1.0M 0.5M

SCE vs

Rate

SCE

E vs SCE

Figure 5.17. Cyclic Voltammogram at a Graphite Electrode, in 1.0M Mo(VI) in 11M Hydrochloric Acid, Scan Rate 4V/min

E vs SCE

Figure 5.18. Cyclic Voltammogram at a Graphite Electrode, of 0.7M Mo(VI) in 1M Nitric Acid, Scan Rate 4V/min

106 5.3.2. Screening the Molybdenum Salts

The solubility of the salts is a critical factor in redox cells as concentration is proportional to energy storage capacity. The electroactivity of the salts is important as well. It was found that there were no great differences in the cyclic voltammograms in 4M sulphuric acid at a graphite electrode for each of the compounds molybdic acid, disodium molybdate and ammonium molybdate. While there was no great difference in peak potential, the cyclic voltammograms obtained for sodium molybdate and ammonium molybdate solutions did show slightly higher peak heights at the same concentration of one mole per litre of molybdenum. This could be attributed to differences in pH, and the ability of the anions to carry charge more effectively through solution than a molybdic acid solution, although no conductivity measurements were performed to verify this.

Two other salts which were screened for electroactivity were phosphomolybdic acid and molybdenum dioxide (Figures 5.19 and 5.20). However at the graphite working electrode, multiple peaks are evident for each of the redox couples showing that a large number of species are present in the acid solutions. Considering sodium molybdate has the greatest solubility and dissolves the fastest therefore, it was chosen as the preferred source of molybdenum for the redox application.

107 E vs SCE

Figure 5.19. Cyclic Voltammogram at a Graphite Electrode, in 1.0M Phosphomolybdic Acid Mo(VI) in 4M Sulphuric Acid, Scan Rate 4V/min

E vs SCE

Figure 5.20. Cyclic Voltammogram at a Graphite Electrode, in 0.1M Molybdenum Dioxide Mo(IV) in 4M Sulphuric Acid, Scan Rate 4V/min

108 5.3.3. Electrode Selection

The cyclic voltammograms presented in the previous section show that the molybdenum reactions are diffusion controlled and quasi-reversible at a graphite electrode. A number of other electrode materials were also investigated and Figures 5.21 to 5.25 show the cyclic voltammograms for platinum, palladium, bismuth, glassy carbon, and lead.

The palladium electrode shows practically no peaks at all. The platinum electrode had about 20% of the surface area of the graphite electrodes used, and the resulting peaks are nearly half the size of the ones observed at graphite. This shows that the current densities of the molybdenum reactions at platinum electrodes are greater than at graphite. The bismuth electrode shows two large peaks over the potential range -1.0 to 2.0 V but no others which most likely involves the oxidation of bismuth. The lead electrode only shows two small peaks over the entire potential range at negative potentials, which would involve lead being oxidised to lead sulphate. The glassy carbon electrode shows a good response and is used in section 6 for kinetic studies. The graphite electrodes give about 20% more current for a given surface area compared to the glassy carbon electrode, although this is probably due to the larger than geometric surface area arising from its porosity. While glassy carbon shows good electrocatalytic response for the molybdenum reaction, these appear at different potentials compared to graphite. This is discussed further in Chapter 6.

109 Figure 5.21. Cyclic Voltammogram at a Platinum Electrode (Area 5mm2), in 1M Mo(VI) from Molybdic Acid, 4M Sulphuric Acid, Scan Rate 4V/min

< c E vs SCE ■*—> a u.a Uc u3

Figure 5.22. Cyclic Voltammogram at a Palladium Electrode (Area 0.7mm2), in 1M Mo(Vl) from Molybdic Acid, 4M Sulphuric Acid, Scan Rate 4V/min

110 Figure 5.23. Cyclic Voltammogram at a Bismuth Electrode (Area 28mm2), in IM Mo(Vl) from Molybdic Acid, 4M Sulphuric Acid, Scan Rate 4V/min

E vs SCE

Figure 5.24. Cyclic Voltammogram at a Glassy Carbon Electrode (7mm2), in IM Mo(VI) from Molybdic Acid, 4M Sulphuric Acid, Scan Rate 4V/min

111 Figure

C u r r e n(mA) t 5.25. Rate Cyclic 1M

Mo(VI) 4V/min

Voltammogram

from 112

Molybdic

at

a

Lead

Acid,

Electrode

4M

Sulphuric

(Area E

Acid,

vs 28mm

SCE

Scan 2 ),

in

CHAPTER 6 CYCLIC VOLTAMETRIC STUDY OF THE KINETICS OF Mo REDOX COUPLES

The cyclic voltammograms obtained in molybdenum solutions (of different oxidation states) are very complex with many peaks apparent over the large potential range of -1.0 to 2.0V. The present series of experiments involved varying the scan rates, the molybdenum concentration, the sulphuric acid concentration, as well as the molybdenum salt, to determine any effect on the peak currents, peak potentials and thus the effect on the heterogeneous rate constants and diffusion coefficients for the reactions of interest.

The significant redox couples for the molybdenum reactions in sulphuric acid were separated, by narrowing the potential range scanned, and studied by cyclic triangular wave voltammetry. The molybdenum salts used were molybdic acid, disodium molybdate, and ammonium molybdate. The sulphuric acid concentration, the molybdenum concentration, and the scan rates were varied (from 5 to 200mV/sec) in order to evaluate the kinetics of the reactions involved. All experiments were performed at a glassy carbon electrode (area of 7mm2) polished with a fine p!200 grit polishing paper between scans.

Due to its higher solubility compared with the other molybdenum salts tested, the cyclic voltammetry study focused on sodium molybdate. The Mo(III)/Mo(IV) couple was studied in a one molar Mo(III) solution prepared from sodium molybdate, while the Mo(V)/Mo(VI) reactions were studied in one molar Mo(V) prepared by mixing Mo(III) and Mo(VI) solutions in a ratio of 1:2,The Mo(IV)/Mo(V) reaction was also studied using a Mo(IV) solution prepared from either molybdenum dioxide (0.1M) or by mixing Mo(III) and Mo(VI) solutions.

113 Nitrogen was purged through the solutions to minimise electrolyte oxidation by atmospheric oxygen. The electrode was polished between scans to eliminate interferences. Experiments were repeated on solutions chosen at random so that errors could be calculated and reproducibility could be estimated.

As shown earlier, a 6mm graphite electrode gave about 20% greater peak current densities compared to glassy carbon, over the potential range -1.0 to 2.0V, however the resulting cyclic voltammograms were not very reproducible. This was especially the case at the higher potentials (at 1.1V), where the graphite surface was visibly oxidised and significant carbon peaks were observed. The carbon peaks can be seen on the scan obtained at the graphite electrode in a 4M sulphuric acid solution with no molybdenum present (Figure 6.1).

In contrast, a 3mm glassy carbon electrode gave results which were much more reproducible. It was also noted that the carbon oxidation and reduction peaks were absent at the glassy carbon electrode suggesting that glassy carbon is also more stable in the solutions over the potential range used.

114 Figure Current (mA)

6.1.

Rate LOWER:Glassy Cyclic

4V/min

Voltammogram

Carbon, 115

in at

4M a

Sulphuric TOPrGraphite

Acid

E Only,

vs Electrode,

SCE

Scan

Glassy carbon is relatively cheap compared to say platinum, and has increased reversibility for redox couples and reactions which involve subsequent proton transfer. Glassy carbon electrodes have a negligible concentration of surface acid-base groups and comes closest to an ideal inert redox electrode. (56)

After just a few cycles the glassy carbon electrode gave an equilibrium curve which no longer changed with consecutive scans, whereas graphite behaved rather unpredictably especially at positive potentials (due to carbon surface reactions).

From the peak potential separations observed in the cyclic voltammograms of Figures 5.9-5.12 the reactions appear to be irreversible at a graphite electrode and so the equations used to calculate the relevant kinetic constants were those for the irreversible reactions as set out in Bard and Faulkner (8) as discussed below.

It was found necessary in some cases to make peak height corrections for cathodic non zero base lines. To calculate the associated peak currents a few methods are available. In the present study Nicholsaris equation was used as was defined in Figure 2.7 (8). This being given by:

V - (v)o + 0.485(is)o + 0.086i1>a

The heterogeneous rate constant k° and the transfer coefficient (a), were determined from plots of ln(ip) vs (Ep-E°) as given by Equation 2.3 for irreversible reactions. Thus, for the conditions used:

116 ip = 0.227nFACo*k° exp[-(anaF)(Ep-E°')/RT]

To solve for k° and a, the logarithms of both sides were taken and the constant

numerical values for n, na, A, F, R, and T were substituted into the equations. These values being 1, 1, 0.070686cm2, 96487CmoF, SAMJK^mol'1, and 293K respectively. The concentration of molybdenum C0*molcnV3, was varied and so was not incorporated into the constants shown in Equations 6.1 and 6.3. Thus, by equating the slope and intercept of the plots of ln(ip) vs (Ep-E°) as in Equations 6.1 and 6.2, the parameters can be calculated.

k° = exp[intercept]/(1548.20Co') 6.1 a = -slope/39.609 6.2

The diffusion coefficient for the molybdenum ions was- determined from plots of ip vs v1/2 as given by Equation 2.4 for irreversible reactions. Thus, for the conditions used:

ip = 299000n(a na)1/2AC0*D01/2v1/2

To solve for D0 the constant numerical values for n, na, and A were substituted into the equation.

D0 = [slope/(21135.114C0(a122)]2 6.3

117 The actual plots are shown in Appendix 3 and 4. From equations 2.3 and 2.4, it can be seen that for an irreversible reaction the peak potential separations (Ep-E°) increase with increasing scan rate. A series of cyclic voltammograms at various scan rates is shown in Appendix 5 for the molybdenum redox couples. This shows that peak potentials do change with scan rate inferring that the reactions are indeed electrochemically irreversible. The molybdenum reactions can be better described as being quasi-reversible. A criteria for a reaction being quasi-reversible is given in Section 2.9.3.3. as 0.00003 < k° < 0.02. It will be shown in Section 6.4 that this is true for molybdenum at a glassy carbon electrode.

In the present study, the following three molybdenum redox couples were studied: Mo(III)/Mo(IV), Mo(V)/Mo(VI), and Mo(IV)/Mo(V). Each couple is discussed separately in Sections 6.1 to 6.3.

6.1. Mo(III)/Mo(IV)

As seen in Figure 6.2 at the negative potential range there are two sets of molybdenum peaks associated with the oxidation and reduction of two different species. The anodic and cathodic peak potentials for the most negative couple are at around -0.5 and -0.6V respectively, which is for the reaction between Mo(IV) and Mo(III), while the second couple occurs at 0.0 and -0.2V. The more positive of the two also corresponds to a change in colour from dark green to a dark red solution during oxidation. Both a green and a red Mo(III) species has been reported, in fact the potentials at which they occur indicate that they actually involve a divalent species Mo(III),Mo(IV), which is well documented.

118 Figure it Two the the the C u r r e n(mA) t In

is

this M M Mo(III)>Mo(IV) reduction one Mo(III)/(IV)

proposed groups o o

(III) (III) 6.2.

section at

red 3 2 the

M

Richens

o

of

that more (IV) the

—

Mo(IV) and Rate Mo(III) Cyclic couple

> these

two

— positive M B and

— 50mV/sec

o in

Voltammogram pairs

> >

(III) will reactions red in

from

Sykes the Mo(IV),

M

acidic

2 be potential

M of o single (IV)

Sodium

o

referred

(30), peaks ( (the

1V) 119 media A

3

+

scan) +

and

and

2e figure

being + which 2e

Molybdate,

‘

at to

which

e'

B V

Paffett

as a can

vs

A. 0.0V to

were Glassy

Reaction

SCE

explains the be

(24), vs

described in found REACTION REACTION

left NHE

Carbon 4M

report A contains these

to Sulphuric and

be

as: (5)

Electrode,

observations, a Reaction

associated mechanism A B both

Acid,

peaks

B in

,

Scan

with with

and 1M

for A

6.2. Mo(V)/Mo(VT)

As seen in Figure 6.3 the Mo(V)/Mo(VI) couple is not readily observed. While the oxidation of Mo(V) to Mo(VI) occurs at 1.0V, the reduction of Mo(VI) to Mo(V) that is expected to occur around 0.8V vs SCE (26) is not visible. The reduction of Mo(VI) is reported to be a 2 electron mechanism producing Mo(IV), this explains why no Mo(V) reaction peak is observed. The Mo(IV) in this reaction would then disproportionate with remaining Mo(VI) to form Mo(V).

Mo(VI) + 2e‘ —> Mo(IV)

Mo(IV) + Mo(VI) —> Mo(V)

Dong and Wang (80-82) have studied Mo(VI) in 2M H2S04 by cyclic voltammetry at carbon fibre microelectrodes at 200mV/s. The response in the potential range of -0.1 to 0.7V vs SCE showed three distinct peaks at potentials of about 0.0, 0.2, 0.3V. They showed that no increase in peak current is observed if the concentration of Mo(VI) is elevated. This is evidence that the rate of adsorption of the molybdenum species on the carbon fibre microelectrodes is very rapid, belonging to an irreversible strong adsorption process. The X-ray photoelectron spectrum showed that the Mo(VI) reduction products include both Mo(VI) and Mo(V).

120 V vs SCE

Figure 6.3. Cyclic Voltammogram at a Glassy Carbon Electrode, in 1M Mo(V) prepared from Sodium Molybdate, in 4M Sulphuric Acid, Scan Rate 50mV/sec

6.3. Mo(IV)/Mo(V)

This couple was studied in a Mo(IV) solution. The oxidation and reduction peaks were found to be approximately at 0.3 and -0.1V vs SCE respectively. The oxidation and reduction potentials for the Mo(III)/Mo(IV) couple ’A’ given in Section 6.1 was -0.2 and 0.0V vs SCE respectively. Thus some difficulty arises in separating these two couples at a glassy carbon electrode. The multiple peaks which are shown in Figures 5.9 to 5.12, at a graphite electrode, shown as peaks A and E, are due to the overlap of these two reactions.

121 Figure

6.4. Current (mA) Cyclic Rate Mo(IV)

50mV/sec

Voltammogram

prepared 122

from - -

- at 0.1 0.1

MoO a

Glassy z ,

in

Carbon 4M

Sulphuric vs

Electrode,

SCE

Acid,

in

0.1M Scan

6.4. Kinetic Parameters of Concentrated Molybdenum in Dilute Sulphuric Acid at a Glassy Carbon Electrode

A detailed kinetic study of the four molybdenum redox couples discussed in 6.1 to 6.3, was undertaken. The Mo(IV)/Mo(III) and Mo(VI)/Mo(V) couples were studied at three molybdenum concentrations (as disodium molybdate) and three sulphuric acid concentrations. In addition molybdenum solutions made from ammonium molybdate and molybdic acid were studied at three different molybdenum concentrations in 4M sulphuric acid. The Mo(V)/Mo(IV) couple was also studied in 4M sulphuric acid at three different molybdenum concentrations.

Using Equation 2.4, from plots of ln(ip) vs (Ep-E°), values of the heterogeneous rate constants were obtained for different molybdenum and sulphuric acid concentrations and are summarised in Tables 6.1 to 6.4 for each of the molybdenum couples. The heterogeneous rate constant is seen to be independent of molybdenum concentration, molybdenum source, and sulphuric acid concentration in the ranges studied. Similarly, the diffusion coefficients of the various Mo ions (Table 6.5) were independent of these parameters. This data is also represented in the graphs of Appendix 3 and 4.

123 Table 6.1. Heterogeneous Rate Constants for Mo(III)/Mo(IV) (Reaction A), (Anodic and Cathodic Values given Respectively (cm/s))

Heterogeneous Rate Constant [Mo] (cm/s) Solution mol/1 3M 4M 5M

h2so4 h2so4 h2so4

1.0 1.2E-4 2.2E-4 1.1E-4 1.6E-4 1.0E-4 1.1E-4

Sodium 0.5 2.1E-4 3.3E-4 0.6E-4 Molybdate 1.6E-4 2.7E-4 0.7E-4

0.25 1.4E-4 5.0E-4 0.7E-4 1.7E-4 3.2E-4 0.7E-4

1.0 - 1.9E-4 - 1.5E-4

Ammonium 0.5 - 2.4E-4 - Molybdate 2.8E-4

0.25 - 2.3E-4 - 2.6E-4

Average Anodic : 2.0(±1.4)E-4 Average cathodic : 1.6(±0.9)E-4

124 Table 6.2. Heterogeneous Rate Constants for Mo(III)/Mo(IY) (Reaction B), (Anodic and Cathodic Values given Respectively (cm/s))

[Mo] Heterogeneous Rate Constant (cm/s) mol/1 Solution 3M 4M 5M

h2so4 h2so4 h2so4

1.0 3.1E-5 4.1E-5 0.5E-5 1.3E-5 4.7E-5 0.5E-5

Sodium 0.5 4.6E-5 4.1E-5 0.3E-5 Molybdate 4.3E-5 4.1E-5 0.3E-5

0.25 1.5E-5 5.0E-5 0.5E-5

0.9E-5 4.5E-5 0.4E-5

1.0 _ 0.3E-5 _ Ammonium 0.5E-5 Molybdate

0.5 - 0.7E-5 - 1.0E-5

0.25 _ 0.8E-5 _

1.4E-5

1.0 _ 10E-5 H2Mo04 5.9E-5

0.5 _ 11E-5 _

8.3E-5

0.25 _ 3.1E-5

2.4E-5 Average Anodic : 2.6(±1.9)H-5 Average Cathodic: 2.3(±2.0)E-5

125 Table 6.3. Heterogeneous Rate Constants for Mo(VI)/Mo(V), Anodic Case (cm/s)

Heterogeneous Rate Constant [Mo] (cm/s) Solution mol/1 3M 4M 5M

h2so4 n2so4 h2so4

1.0 3.0E-4 4.5E-4 3.7E-4

0.5 2.7E-4 6.1E-4 4.4E-4 Sodium Molybdate 0.25 2.2E-4 2.5E-4 4.5E-4

1.0 - 3.4E-4 - Ammonium Molybdate 0.5 - 2.3E-4 -

0.25 - 2.0E-4 -

1.0 - 2.2E-4 -

H2MoC>4 0.5 - 2.6E-4 -

0.25 - 2.1E-4 -

Average Anodic : 3.7(±1.3)E-4

126 Table 6.4. Heterogeneous Rate Constants for Mo(V)/Mo(IV), (Anodic and Cathodic Values given Respectively (cm/s))

Solution 0.1M Mo 0.05M Mo 0.025M Mo

4M H2S04 5.0E-5 3.8E-5 3.8E-5 1.2E-5 2.3E-5 7.3E-5

Average Anodic : 4.2(±0.7)E-5 Average Cathodic : 3.6(±3.3)E-5

The errors in the calculated means for the heterogeneous rate constants are relatively large, but this is expected as the value is very sensitive to very slight shifts of the slope. Further error is introduced considering that the geometric and active surface areas are different. As seen in the plots of ln(ip) versus (Ep- E0) in Appendix 3, in all cases the lines fall very close together, thus showing that the value of the heterogeneous rate constant is independent of the acid and molybdenum concentrations.

The various plots of ip versus the square root of the scan rate all have very similar slopes, as can be seen in Appendix 4. From the slopes of these plots the diffusion coefficient was calculated according to Equation 6.3 and the averages given in Table 6.5 are derived from the experiments involving three different acid and molybdenum concentrations as for the heterogeneous rate constant.

127 T;iblc 6.5. Calculated Average Molybdenum Diffusion Coefficients Anodic and Cathodic Values given Respectively (cm2/scc)

Reaction A 5.9(±4.3)E-9 Mo(III)/Mo(lV) 7.5(±4.7)E-9

Reaction 13 4.2(±2.7)E-10 Mo(III)/Mo(lV) 2.2(±2.1)E-10

Reaction C 3.5(±4.S)E-9 Mo(VI)/Mo(V)

Reaction D 7.9(±4.2)E-9 Mo(V)/Mo(IV) 4.5(±3.3)E-9

The values of the diffusion coefficients shown in Table 6.5 are much lower than would be expected for solution diffusion. This discrepancy is discussed further in Section 6.5.

6.5 Anomalous Diffusion Coefficients

The anomalous diffusion coefficients reported above for the molybdenum couples could result from a number of factors. Firstly, the estimated surface area of the electrode could be inaccurate.

Surface roughness results in the true surface area of a solid electrode being greater than the geometric area. As shown in Equations 2.3 and 2.4 for calculation of k° and Du it is the total active surface area which is required as an important parameter. In this study the extremely small values obtained for Dc(of 10'9 cnij/sec), suggest that the actual active surface area of the electrode may be much lower than the geometric surface area used in the calculations. It is well known that many redox reactions occur only at specific active sites on graphite or glassy carbon electrodes (76,77). The glassy carbon electrodes surface then is not 100% covered by active sites for the molybdenum redox reactions. Note that some -CO, -COM, -COOH groups will exist on the surface, and that they may or may not be the active sites for reaction.

Tabic 6.6. Calculated Geometric Surface Area Depending on the Diffusion Coefficient

Diffusion Actual Coefficient Area (cm/sec) (cm2)

10 s 0.0005

10- 0.002

107 0.005

io8 0.017

109 0.071

128 Assuming various values of D0 therefore, Equation 2.4, was used to calculate the expected effective surface area and the values are given in Table 6.6.

Since for most solution species, the diffusion coefficient is expected to be of the order of lO'6 cm2/sec, these calculations suggest that the effective surface area of the glassy carbon electrode may be only 3-7% of the total geometric area.

If this were the case however, one would expect the glassy- carbon electrode to behave like a microporous electrode array. The glassy carbon electrode was polished between scans potentially giving slightly different electrode configurations but with an overall common surface area. If all of the area is not active, the electrode must have areas of activity separated by relatively inert sections. These areas then in themselves, being very small, may be influenced by similar forces as a micro electrode.

Compact micro electrodes have been used to take advantage of enhanced diffusion, low charging currents, and reduced solution resistance effects. Radial diffusion at the edges of microelectrodes contribute to the overall diffusion and results in quasi steady state currents for moderate sweep rates with reversible redox couples. This diffusion flux affects various properties but the "most striking" effect of an array of closely spaced electrodes where the diffusion layers overlap is that it becomes possible to detect the electrogenerated products at the adjacent electrodes. (89) This effect could be a contributing factor for molybdenums complicated reactions which lead to lower measured currents.

These characteristics are related to the electrode and not specific to molybdenum. If reaction products were defusing to nearby active sites and reacting further then this could help explain the complicated voltammogram results.

Also, the current at one electrode can effect that at its neighbour when both are at the same potential, because the diffusion layers overlap. The current is reduced. This reduction in current is called a shielding factor. (89)

For example, with a conventional 3 electrode system for the reduction of a 5mM Ru solution, and 8 micro electrodes together, the resulting cathodic current was only 28% of 8 times the current at a single electrode. The importance of radial diffusion of redox species to the electrodes is clear. The effective radial diffusion to the combined electrode area was reduced by using the closely spaced electrodes together. (89) 129 The glassy carbon electrode used was not in itself a random array of microelectrodes but a random array of various sized electrodes which may still be influenced by similar effects.

At very short times the electrodes behave independently but with the scan rates used here, the times are relatively long and the diffusion layers have time to merge and eventually the entire active surface behaves as one electrode (as would be expected). (90)

For micro electrodes the equation which can be used to calculate the diffusion coefficients should have another term included:

ip = 29 9000n(ana) 1 /9 ACo★ Do1 /91v /? + (constant) nADoCo ★ /r. o

compared to:

ip = 299000n(ctna)1/2ACo*Do1/2v1/2 (91,45)

The plot of i versus v°'5 is thus not expected to go through the zero. For voltammetry with micro electrodes the dependence of the current on potential may not exhibit a peak especially at low scan rates. (91) For the molybdenum system peaks are usually found. Plus it has been shown that the peaks are more defined at lower scan rates then high scan rates. (91) So based on this information it does not appear that the molybdenum glassy carbon system is acting like a micro electrode array despite the lower than expected diffusion coefficients.

Additionally a wave like potential dependence may only be observed if the radius of the electrode is well below 10 cm. Only then does the steady state resulting from the domination of the second term on the right hand side of the above equation exist. (91)

From the above, it can be concluded that since micro electrode behaviour is not observed in the cyclic voltammograms for the molybdenum system, the low diffusion coefficient must be due to other complications.

For instance, it may be due to the complexity in the electrochemistry and the overlapping of reaction peaks. The diffusion coefficients are calculated from plots of i versus v ' which originates from equation 6.3. Now, this calculation depends on the data plot being linear and to pass through the origin. The majority of plots do approximately go through the origin but in some cases the plot does not pass through the origin suggesting that equation 2.2 is inappropriate for this system.

130 CHAPTER 7 EFFECT OF COMPLEXING AGENTS, ELECTROLYTE ADDITIVES AND ANION CONCENTRATION ON THE CYCLIC VOLTAMETRIC BEHAVIOUR AND SOLUBILITY OF MOLYBDENUM

In an effort to increase the concentration of molybdenum and increase the energy density of a molybdenum redox cell, the use of complexing agents and additives was considered. Electrolyte additives can be used for two purposes. Firstly additives can be used to stabilise a species in solution and even increase solubility. This is done by either a dispersing agent or a complexing agent. Secondly additives can be used in order to increase electroactivity and reversibility of the reactions of interest.

Solution additives have also been used to increase the reversibility of redox couples such as the vanadium system by depositing on the electrode and forming active sites for the redox reaction. Other additives such as Pb and Cd have been used as inhibitors for the hydrogen evolution reaction during charging for the vanadium redox battery (77). In the molybdenum system, the formation of the Mo(blue) species limits both the reversibility of the Mo(VI) reduction as well as the molybdenum concentration. The solution used to test the additives was 1M sodium molybdate in 4M sulphuric acid.

The various complexing agents used included sodium citrate, tartaric acid, EDTA sodium salt, and D-gluconic acid sodium salt. Sodium lauryl sulphate, sodium hydrogen sulphate, sodium sulphate, potassium sulphate, calcium sulphate, and ammonium sulphate were used to vary the sulphate concentrations of the electrolyte.

131 7.1. Effect of Sulphate Concentration on Mo (blue) Formation The anion concentration, in a 1M Mo(VI) sodium molybdate solution in 4M sulphuric acid, was changed by adding various sulphate salts (as listed in Table 7.1). Cyclic voltammetry was then performed in each of these solutions. Usually when a clear yellow Mo(VI) solution is reduced at a graphite electrode (at positive potentials) a blue product is formed. It was found that at a certain anion concentration, while reducing a solution of Mo(VI) at a graphite electrode, the Mo(blue) species did not form. This was only found to occur under two conditions, these being with an additional 2M K2S04 or 1M (NH4)2S04. The sulphate concentration was then varied to be able to see what was the minimum sulphate concentration necessary for this phenonemon to occur. As shown in Table 7.1 only these two high concentrations of sulphate gave this result. Cyclic Voltammograms of molybdenum solutions containing additives are shown in Figures 7.1 and 7.2. The cyclic voltammograms are not significantly different to the ones which have no additives shown in earlier Chapters.

It was proposed in Chapter 5 that the molybdenum solubility is limited by the solubility of the molybdenum blue species. In the presence of the excess sulphate which was discussed above, to inhibit the formation of this molybdenum blue species, the solubility of molybdenum (VI) was reassessed. It was found however, that the solubility of molybdenum (VI) in these solutions was not enhanced at all but reduced and limited to just above one molar in sodium molybdate in 4M sulphuric acid with 2M potassium sulphate. The molybdenum concentration was also lower for a solution containing 4M sulphuric acid and 1M ammonium sulphate. This can be explained by considering that the sulphate concentration in these solutions is 5 to 6 molar. Thus, while the increase in sulphate ion concentration inhibits the formation of Mo(blue), the solubility of the Mo(VI) sulphate formed decreases due to a ’common ion effect’.

132 Table 7.1. Effect of Sulphate Concentration on Mo(blue) Formation During Reduction of Mo(VI) at a Graphite Electrode, in 1M Sodium Molybdate and 4M Sulphuric Acid

Salt Molarity Mo(blue) Added YorN

Na2S04 1.0 Y

- 2.0 Y

NaHS04 2.0 Y

k2so4 0.5 Y

- 1.0 Y

- 1.5 Y

- 2.0 N

(NH0£O. 0.5 Y

- 1.0 N

CaS04 Sat’d Y

133 Figure Current (mA) Figure

7.2. 7.1.

Ammonium Mo(VI) Cyclic Potassium Mo(VI) Cyclic

10

Voltammogram Voltammogram - from from

Sulphate,

Sulphate,

Sodium Sodium 134

Scan

Scan

Molybdate, Molybdate, at

at Rate

Rate a a

4V/min Graphite Graphite

4V/min

1.0 4M 4M 1.0

E Sulphuric E Sulphuric Electrode,

Electrode, vs vs

SCE SCE

Acid, Acid,

in in

1M 1M 2M 1M

7.2. Effect of Complexing Agents

The additives EDTA sodium salt, tartaric acid, D-gluconic acid sodium salt, and sodium citrate were evaluated as potential complexing agents for the molybdenum ions. The additives which were used were found to have little effect if not a negative one on the peak potentials and peak, heights at a graphite electrode as can be seen in Figures 7.3 to 7.5. For example, the curves for sodium citrate and tartaric acid as additives show flat peaks indicating less reversible reactions than seen previously for the molybdenum reactions. The solubility of molybdenum in the presence of these additives was also reduced to less than one molar as precipitates were formed with additive concentrations of 0.5M sodium citrate, 0.5M tartaric acid, 0.5M EDTA sodium salt, and 0.1M D- gluconic acid sodium salt 97%.

E vs SCE

Figure 7.3. Cyclic Voltammogram at a Graphite Electrode, in 1M Mo(VI) from Sodium Molybdate, 4M Sulphuric Acid, 0.5M EDTA Sodium Salt, Scan Rate 4V/miri

135 E vs SCE

Figure 7.4. Cyclic Voltammogram at, a Graphite Electrode, in 1M Mo(VI) from Sodium Molybdate, 4M Sulphuric Acid, 0.5M Tartaric Acid, Scan Rate 4V/min

E vs SCE

- -10

Figure 7.5. Cyclic Voltammogram at a Graphite Electrode, in 1M Mo(VI) from Sodium Molybdate, 4M Sulphuric Acid, 0.5M Sodium Citrate, Scan Rate 4V/min

136 CHAPTER 8 MOLYBDENUM REDOX CELL PERFORMANCE TESTING

8.1. Design and Operation

As with any battery system, redox cells exhibit gradual loss in capacity with continuous cycling. In redox cells, this is due to cross mixing of electroactive species and thus some self discharge as a result of differential transfer of active species across the membrane. An imbalance occurs in the molybdenum ion concentration so that the solutions will not charge up to 100% at the preselected upper charging voltage. This leads to reduced efficiencies and cell capacity. In the molybdenum redox cell, as in the case of the vanadium system, the capacity can be fully restored by simply remixing the two solutions at required intervals.

Before any cell studies could be undertaken however, considerable effort was required in achieving a suitable cell design which would be relatively easy to assemble, while minimising leakage and flow problems. The final cell design developed for the charge discharge cycling of the molybdenum redox cell is illustrated in Figure 8.1.

+ Copper Current Collector Carbon Fibre — Electrode Passing the Electrolyte Graphite Plate

Negative Positive Half-Cell Half-Cell Electrolyte Electrolyte

Figure 8.1. Final Redox Flow Cell Apparatus 137 8.2. Cell Resistance

Before undertaking any charge - discharge cycling of the molybdenum electrolytes, each cell was first tested for its overall electrical resistance. A good electrode assembly produced from carbon felt on a conducting plastic substrate should have a resistance of less than 1 Ohm/cm2 for good performance. Obviously the lower the resistance, the lower the ohmic energy losses.

For an electrochemical cell, the electrode polarisation losses are equal to the sum of the individual overpotentials associated with the charge transfer kinetics (activation overpotential) at the electrode, and mass transport of reactant species (concentration overpotential). In an operating cell therefore, the cell voltage is given by:

EcELL " Ec - Ea - T|a " qc " IRcell

Where;

Ec° = cathodic half cell potential Ea° = anodic half cell potential qA = activation overpotential x)c = concentration overpotential

IRcell = ohmic resistance losses

138 The ohmic resistance of a particular cell is made up of the electrical resistances of the electrodes, the membrane, the electrolyte, and any contact resistances. For example the two electrodes used with the molybdenum redox cell consisting of Sigri carbon felt heat bonded onto a conducting plastic (SEA Ltd Austria) had a resistance of 0.6 Ohm/cm2 each. When these electrodes were placed in the test cell without a membrane, but with each electrode touching, an overall electrical resistance of 1.4 Ohm/cm2 was measured. Depending on the type of membrane used therefore, an overall ohmic resistance of between 2 and 4 Ohm/cm2 would be expected, not including any reaction polarisation at the electrodes.

Various combinations of electrodes (of individual resistances less than 1 Ohm/cm2) and membranes were tested and the cell resistance values determined for charging and discharging. Cell efficiencies were also determined over a range of current densities as detailed in the following section.

8.3. Molybdenum Cell Charge - Discharge Testing

The molybdenum cell with 75mls per half cell of a 1M sodium molybdate solution in 4M sulphuric acid was cycled with a current density of 20mA/cm2. A typical charge - discharge curve is shown in Figure 8.2. One of the most obvious features of this curve is the presence of two steps in both the charge and the discharge cycles.

139 Charge Discharge

Figure 8.2. Typical Charge - Discharge Curve for the Molybdenum Redox Cell Employing the Mo(III)/Mo(lV) vs Mo(V)/Mo(VI) Redox Couples

Charging Discharging

Time (hours)

Figure 8.3. Charge - Discharge Curve for the Molybdenum Redox Cell From Reference (15) (idiarcc = i^^ = 20 mA/cm2)

140 Preliminary studies by Trinh(15) indicated that the molybdenum redox couple as described here, Mo(VI)/(V) and Mo(IV)/(III), under the same conditions had a discharge potential of around 0.8V as illustrated in Figure 8.3. In the present study however, it has been found that the cell discharges at a voltage of approximately 0.8V for only a small part of the full discharge curve and then discharges at an average of 0.4V for the remainder of the cycle. For the experimental conditions used by Trinh, the theoretical capacity (in time) can be calculated as being 10.72 hours for full discharge at 20mA/cm2 for the conditions of 200ml of 1M Mo per half cell at 20mA/cnr on a 25cm2 electrode. The theoretical discharge time is given by:

t = mnF/A = (l*0.2)*l*96500/(0.5*3600) =10.72 hours for full discharge

Examining Figure 8.3 (15) therefore, reveals that the discharge time of only one hour, represents only 10% of the theoretical capacity of the system. This explains therefore the difference in the discharge curves obtained here compared with those reported by Trinh (15).

The average coulombic efficiency (Ec) of a cell is affected by self discharge due to diffusion of the molybdenum ions across the membrane, as well as any gassing at the electrodes during charging. The voltage efficiency (Ev) is associated with polarisation. This polarisation is due to ohmic losses associated with the electrode, membrane, and resistance of the electrolyte itself, as well as concentration and reaction overvoltage losses.

141 Several different membranes were tested in the molybdenum cell using Sigri graphite felt electrodes, and a solution of 1M sodium molybdate in 4M sulphuric acid. The volume on each side of the cell was 75ml and a current density of 20mA/cm2 was employed for both charging and discharging.

As shown in Table 8.1, all of the membranes gave similar results except that Nation exhibited more solution transfer than the others (of about 15ml out of 75ml) after 12 cycles. The coulombic efficiency (Ec) is high for all of the membranes indicating that the membrane component in a molybdenum cell would not be a problem. The coulombic efficiency was greater than 90% and remained high for up to thirty cycles. The voltage efficiency is fairly constant for each type of membrane showing that the area resistance values for each membrane are similar (2-2.5Q/cnr). An overall energy efficiency (Ej.) of up to 76% was obtained with the AMV membrane.

142 Table 8.1. Effect of Membrane Type on Cell Efficiencies with Sigri Felt Electrode (Cycles 1 & 12 respectively), 20mA/cm2, 1M Sodium Molybdate in 4M H2S04

Membrane Ec Ev Er

AMV 96 78 76 96 78 76

CMV 96 78 75 95 74 70

Nafion-112 95 77 73 90 72 65

Selemion 95 77 73 91 72 66

Several different graphite felt electrode materials were also tested using the AMV membrane, 1M sodium molybdate in 4M sulphuric acid, and a volume on each side of the cell of 75ml. The current was varied from 20 to 40 mA/cm2. The various carbon felts bonded to the conducting plastic electrodes available (made as in section 4.8.2) often gave high area resistances and poor reproducibility. To avoid this problem, most of the charge discharge tests were performed with the various carbon fibre felts being pressed directly onto a graphite plate which was in turn connected to a copper current collector. This way, lower cell resistances were attained and this removed the effect of the variable of the conducting plastic substrate.

143 Table 8.2. Effect of Graphite Felt Electrode Materials On Cell Efficiencies with an AMV membrane (Cycles 1 & 12), 20mA/cm2, 1M Sodium Molybdate in 4M Sulphuric Acid

Electrode R-ci-ll Ec Ev Er Charge

Sigri 4.1 97 78 76 97 78 76

Toray 4.9 97 74 72 97 70 68

FMI 5.0 94 71 67 92 69 63

LeCarbone 5.1 94 70 66 91 65 59

Carbon 5.9 93 70 65 Foam 91 63 57

As seen in Table 8.2, the Sigri felt was found to give the best performance with the total energy efficiency remaining fairly constant at 76% for 14 cycles at 20 mA/cm2. At 30mA/cm2 the overall efficiency was 67% dropping to 60% after just 20 cycles. At this current, the coulombic efficiency was 97% however, the voltage efficiency decreased from 69% to 62%. At 40mA/cm2 the energy efficiency was only 51%. This shows that low current densities are required for the present molybdenum redox cell system. However, a novel cell design, with lower resistances, could yield higher efficiencies.

144 During cell cycling, the negative half cell showed no problems, in that there was no precipitation formed after cycling. In the positive half cell however, the molybdenum blue, which forms readily upon reduction of the Mo(VI) positive half cell solution, precipitated during discharge and even deposited on all surfaces. Experiments where the sulphate comcentration was increased were not performed due to molybdenums’ reduced solubility (refer to section 7.1). The problem of the Mo(blue) precipitation may be eliminated by employing the Mo(IV)/Mo(V) couple in the positive half cell, however this would be impractical as the discharge voltage would then be even lower.

8.4. Electrode and Membrane Stability in Mo(VI) and Mo(III)

During cell cycling tests, it was observed that the Mo(VI) electrolyte stained the AMV membrane an intense dark blue colour, corresponding to a precipitate of Mo(blue) on the surface. There was no sign of physical deterioration, however. Pieces of AMV were left in a solution of Mo(VI) for over six months with similar results. Similarly, the graphite fibre felts exposed to the same conditions did not show signs of deterioration.

Furthermore, pieces of AMV membrane were exposed to Mo(III) solutions for more than six months. During this time the solutions had turned completely red (indicating molybdenum oxidation). Apart from the red colouration, the AMV membrane showed no sign of physical deterioration, indicating that membrane stability would not be a problem with the Mo redox cell.

145 8.5. Mo(III) and Mo(VI) Stability Under Paraffin Oil

By covering these solutions with paraffin, atmospheric oxygen is excluded. The yellow 1M Mo(VI) sodium molybdate in 4M sulphuric acid solution, stored under paraffin for more than a year had only turned a light blue. Indicating that only a very slight reduction to molybdenum blue had occurred.

The green Mo(III) solution (produced by electrolytically reducing the above Mo(VI) solution) was also stored under paraffin oil for six months. The characteristic dark green colour changed to a browner appearance. However, when the 1M molybdenum solution was potentiometrically titrated the oxidation state had not changed within experimental error. Paraffin would thus be a useful barrier for excluding atmospheric oxygen from Mo(III) for long term storage.

8.6 Molybdenum Source

The experimental results obtained for solubility and electrochemical behaviour of the molybdenum redox couples have shown that while the performance and energy density of a molybdenum redox cell would be lower than the vanadium system the availability and cost of molybdenum may be a further limitation of the system. Molybdenum is available in a few forms but mostly as Mo(VI) salts. Approximate costs not including delivery are summarised in Table 5.9.

146 Figure 8.4. Comparative Costs of Molybdenum Salts vs V205

Ammonium Molybdate Vanadium Pcntoxidc Sodium Molybdate Molybdic Acid

As can be seen in Figure 8.4, molybdenum is significantly more expensive than vanadium which is a further disadvantage of a molybdenum redox cell.

147 CONCLUSION Solubility studies of several molybdenum salts in a range of supporting electrolytes have shown that the maximum solubility of a molybdenum compound which also had significant electroactivity at a graphite electrode, was 1.6M, the salt being disodium molybdate dihydrate in 4M sulphuric acid.

The kinetics of the four molybdenum couples Mo(VI)/Mo(V), Mo(V)/Mo(IV), and two for Mo(IV)/Mo(III) (denoted earlier as A and B) was investigated at a glassy carbon electrode. The heterogeneous rate constants for these couples were 3.7E-4, 3.9E-5, 1.8E-4, and 2.5E-5 respectively. However, the Mo(VI)/Mo(V) reduction reaction could not be observed at the glassy carbon electrode, Mo(VI) being reduced directly to Mo(IV).

The redox couples initially proposed for a Mo - Mo redox flow, cell were Mo(VI)/Mo(V) and Mo(IV)/Mo(III) for the positive and negative half cell electrolytes respectively. The molybdenum cell using 1M sodium molybdate in 4M sulphuric acid, as the electrolyte gave an energy efficiency of more than 70% showing that it can operate as a redox cell. The coulombic efficiency was high although further improvements in the total energy efficiency would be achieved by better cell design, and a more electroactive electrode material to catalyse the Mo(VI)/Mo(V) reaction and reduce the voltage losses. Unfortunately, however, the discharge voltage was found to be 0.4V which could limit its applications. Further work is thus required to achieve higher energy efficiencies with the molybdenum redox cell.

Overall the molybdenum redox cell has shown some promising features which could make it feasible although the low discharge potential does imply that practical applications would be very limited.

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155 APPENDIX 1.

Molybdenum standard solution AAS data for Figure 4.2.

PPM Mo Absorbance

10 0.100 15 . 0.138 r = 0.9979 20 0.187 a = 0.0124 25 0.220 b = 0.00856 30 0.273

156 APPENDIX 2.

Potentioractric Titration (by Permanganate) of Molybdenum Solutions the x-axis is time and the y-axis is voltage., note that the end points are not clear

Figure 1. 1M Mo(III) Reduced by Zinc

Figure 2. 1M Mo(III) Reduced by Aluminium

157 Figure 3. Mo(III) Unknown Concentration Reduced Electrolytically

Figure 4. Mo(III) Unknown Concentration Reduced Electrolytically, In this Case, the Rate of Addition of the Permanganate is Changed, as is the x-y Scale

158 APPENDIX 2 Examples of Molybdenum Solutions Potentiometric Titration Calculations

Case 1.

For an aliquot of solution of 4ml known to be 1M in molybdenum, the expected volume of oxidising titrant (0.2M KMn04) can be calculated. Assuming the oxidation state of the molybdenum is III, (and the final oxidation state is VI) this volume is 12ml. Then from the actual volume of titrant the actual oxidation state can be calculated. For example if it was 8ml then the oxidation state of the molybdenum would be (IV).

Case 2.

For an aliquot of solution of 4ml known to be in an oxidation state of III, the expected volume of titrant (0.2M KMn04) required to oxidise it to the VI state can be calculated as being 12ml, assuming the molybdenum concentration is 1M. Then from the actual volume of titrant the actual molarity of the molybdenum can be calculated. For example if it was 8ml then the molarity of the molybdenum would be 0.66M.

159 Appendix 3.

The natural logarithm of the anodic and cathodic peak currents verses the difference between the anodic and cathodic peak potential and the formal potential are plotted as they are used to calculate the heterogeneous rate constants.

Ea-Ec versus ln(ia)

Reaction B, 1M Sodium Molybdate

Ea-Ec (mV) "■Na-4 •■■NH3-4 *«»H2-4 ***Na-3 —Na-5

Na-Sodltim Molybdate, NH3*Ammonlum Molybdate, H2**H2Mo04. -3,4.or 5M H2S04

160 Ea-Ec versus; ln(ic)

Reaction B, 1M Sodium Molybdate

' Ea-Ec (mV) ■ Na-4 -NH3-4 *»H2-4 »»Na-3 — Nn-5

Na°Sodlum Molybdate, NH3=Ammonlum Molybdate, H2=H2Mo04. -3,4,or 5M H2S04

Ea-Ec versus. In(ia)

Reaction B, 0.5M Sodium Molybdate

100 Ea-Ec (mV) — Na-4 — NH3-4 — H2-4 **Na-3 — Na-5

Na»Sodlum Molybdate, NH3»Ammonlum Molybdate, H2=H2Mo04. -3,4,or 5M H2S04

161 Ea-Ec versus ln(ic)

60 80 100 Ea Ec (mV) 'Na-4 "“NH3-4 «»*H2-4 »«Nn-3 "-Na-5

Na-Sodlum Molybdate. NH3-Ammonlum Molybdate, H2-H2Mo04. -3,4.or 5M H2S04

Ea-Ec versus In(is)

Reaction B, 0.25M Sodium Molybdate

Ea-Ec (mV) *"Na-4 »«"NH3-4 «««H2-4 **'

Na»Sodlum Molybdate, NH3=Ammonlum Molybdate, H2*H2Mo04. -3.4,or 5M H2S04

162 Ea-Ec versus In(ic)

Reaction B, 1M Sodium Molybdate

Ea-Ec (mV) ■■Na-4 •"NH3-4 »*H2-4 «**Na-3 *“Na-5

Na^Sodlum Molybdate, NH3“Ammonlum Molybdate, H2*M2Mo04. -3,4,or 5M M2S04

Ea-Ec versus ln(ia)

Reaction A, 1M Sodium Molybdate

Ea-Ec (mV) *—Na-4 »«NH3-4 ***:Na-3 «"Na-5

Na=Sotllum Molybdate, NH3=Ammonlum Molybdate, H2=H2Mo04. -3,4,or 5M H2SQ4

163 Ea-Ec versus ln(ic)

Reaction A, 1M Sodium Molybdate

Ea-Ec (mV) Na-4 ■“•NH3-4 ««Na-3 "«Na-5

Na»Sodluin Molybdate, NM3»Ammonlum Molybdate, H2=H2Mo04. -3.4,or 5M H2SQ4

Ea-Ec versus ln(ia)

Reaction A, 0.5M Sodium Molybdate

Ea-Ec (mV) — Na-4 -NH3-4 «*Na-3 — Na-5

Na^Sodlum Molybdate, NH3“Ammonlum Molybdate, H2=H2Mo04. -3,4,or 5M H2SQ4

164 Ea-Ec versus ln(ic)

Reaction A, 0.5M Sodium Molybdate

Ea-Ec (mV) — Na-4 NH3-4 «»Na-3 — Na-5

Na»Sodlum Molybdate, NH3«Ammonlum Molybdate, H2»H2Mo04. -3,4,or 5M H2S04

Reaction A, 0.25M Sodium Molybdate

Ea-Ec (mV) Na-4 — NH3-4 *«»Na-3 -Na-5

Na = Sodlum Molybdate, NH3=Ammonlum Molybdate, H2=*H2Mo04. -3,4,or 5M H2S04 Ea-Ec versus |n(ic)

Reaction A, 0.25M Sodium Molybdate

Ea-Ec (mV) — Na-4 — NH3-4 »**Na-3 — Na-5

Na<*Sodltim Molybdate, NH3"Ammonlum Molybdate, H2»H2Mo04. -3,4,or 5M H2S04

Ea-Ec versus, ln(ia)

Reaction C, 1M Sodium Molybdate

Ea-Ec (mV) — Na-4 — NH3-4 »**H2-4 »»Na-3 — Na-5

Na»Sodlum Molybdate, NH3*Ammonlum Molybdate, H2»H2Mo04. -3,4,or 5M H2S04

166 Ea-Ec 'versus ln(ia)

Reaction C, 0.5M Sodium Molybdate

Ea-Ec (mV) ■■Na-4 ““NH3-4 »»H2-4 *«*Na-3 -"Na-5

Na“Sodlum Molybdate, NH3>Aminonlum Molybdate, H2«H2Mo04. -3,4,or 5M H2S04

Ea-Ec versus, |n(ic)

Reaction C, 0.25M Sodium Molybdate

Ea-Ec (mV) — Na-4 —NH3-4 **H2-4 *®:Na-3 -Na-5

Na»Sodltim Molybdate, NH3=Ammonlum Molybdate, H2=H2Mo04. -3,4.or 5M H2SQ4

167 Appendix 4.

Anodic and cathodic peak currents versus the square root of the scan rates have been plotted here, as they were used to calculate the diffusion coefficients

(ic) vs Square Root Scan Rate

Reaction A, 0.25 M Molybdenum

1 1.5 2 2.5 Square Root Scan Rate (mV/sec) ~ 0.5 'Nn-4 *»*NH3-4 «»Na-3 «-Na-5

Na=sodium molybdate, NH3=ammonitim molybdate, H2=H2Mo04, -3,4,5M H2S04

168 (ia) vs Square Root Scan Rate

Reaction C, 1.0 M Molybdenum t

CO e <

LU r—

C U> t— o

(0 o CL .y XJ o 5

Square Root Scan Rate (mV/sec) ~ 0.5 — Na-4 *«** NH3-4 **»H2-4 **Na-3 — Na-5

Na-sodium molybdate, NH3=ammonium molybdate, H2=H2Mo04, -3,4,CM H2S04

(ia) vs Square Root Scan Rate

Reaction C, 0.5 M Molybdenum

CO E <

'Cf LU >—

0) 5 o -V

Square Root Scan Rate (mV/sec) ~ 0.5 ■■Na-4 •‘“NH3-4 ®«H2-4 *:*¥Na-3 ■■Na-5

Na=sodium molybdate, NH3=ammonium molybdate, H2=H2Mo04, -3.4.5IVI H2S04

169 (ia) vs Square Root Scan Rate

Reaction B, 0.25 M Molybdenum

CL E < S i—

"c OJ 3 O

«VJ Q_

Na=sodiuin molybdate, NH3=ammonium molybdate, H2=H2Mo04, -3.4.5M H2S04

(ic) vs Square Root Scan Rate

Reaction B, 0.25 M Molybdenum

'Tn tt. E <, 2 c Q- .0 *0 o XZ Id O

Square Root Scan Rate (mV/sec) ~ 0.5 — Na-4 w* NH3-4 <*»H2-4 ***Na-3 — Na-5

Na=sodium molybdate, NH3=ammonium molybdate, H2=H2Mo04, -3,4,5M H2S04

170 Na=sodium Cathodic Peak Current/1 E-4 (Amps) 2 . Anodic Peak Current/1 E-4 (Amp: =soc!ium

molybdate, molybdate, (ic) (ia)

NH3=ammonium

NH3=ammonium

vs vs “ “ Reaction »Na-4

Reaction Square Na-4 Square

Square Square

«»«NH3-4 *««NH3-4

Root

Root

B,

molybdate, B, molybdate,

0.5

0.5

Scan

Scan Root

Root

M M »«*H2-4 171 »**H2-4

Molybdenum

Molybdenum H2=H2Mo04, Rate Rate H2=H2Mo04,

Scan

Scan

(mV/sec) ***Na-3 (mV/sec) » : *Na-3

-3,4,

-3,4,5M

Rate

Rate «

~ “ 5M » * “

0.5 Nn-5 0.5

Na-5 H2S04 H2S04

Na=sodium Cathodic Peak Current /1E-3 (Amps) j 2 Anodic Peak Current /1E-3 (Amps)

molybdate, (ic) (ia)

NH3=ammonium

vs vs “ Reaction •Na-4

Reaction Square

Square Square

»»NH3-4

Root

B. molybdate, B,

1.0 1.0

Scan Root

Root

M m 172 ««H2-4

Molybdenum

Molybdenum H2=H2Mo04, Rate

Scan

Scan

(mV/sec) *«Na-3

-3,4,

Rate Rate

"•Na-5 ~ 5M

0.5

H2S04

(ic) vs Square Root Scan Rate

Reaction A, 0.5 M Molybdenum

(/) ° CL £ < 2.5 CO LlJ S 2 c (1) § 1.5 o _v;

! 0-5 £ o«U o 0 0.5 1 1.5 2 2.5 3 3.5 Square Root Scan Rate (mV/sec) ~ 0.5' [»-Na-4 *»NH3-4 *:«Na-3 — Na-5

Na=sodiiim molybdate, NH3=ammonium molybdate, H2=H2Mo04, -3,4,5M H2S04

(ia) vs Square Root Scan Rate

Reaction A, 0.25 M Molybdenum

Square Root Scan Rate (mV/sec) ~ 0.5 Na-4 .*** NH3-4 ^Na-3 "-Na-5

Na=sodium molybdate, NH3=ammonium molybdate, H2=H2Mo04, -3,4,5M H2S04

173 Na=sodium Ancdic Peak Current/1 E-3 (Amps) I z Cathodic Peak Current/1 E-3 (Amps)

molybdate, (ia) (ic)

NH3=nmmoniwm

vs vs Reaction Reaction

Square

Square Square Square ■*Na-4 • “ ■Na-4

Root

Root

A,

A, molybdate, » »«NH3-4

“

0.5 *NH3-4 1.0

Scan

Scan Root Root

M M

Molybdenum

174

Molybdenum Rate H2=H2Mo04,

Rale

»«Na-3 ***Na-3

Scan

Scan

(mV/sec) (mV/cec)

» *-Na-5 —

-3.4.5M

Na-5 Rate Rate

~ ~

0.5 0.5

H2S04

Na=sodium Anodic Peak Current/1 E-3 (Amps) j 2 Anodic Peak Current/1 E-4 (Amp

molybdate, (ia) (ia)

NH3=ammonium

vs vs Reaction Reaction

Square

Square Na-4 Square Square “ -Na-4

*»*NH3-4

Root Root C, A,

molybdate,

»«NH3-4

0.25 1.0

Scan Scan Root

Root

M

M »»H2-4

Molybdenum

175

H2=H2Mo04, Molybdenum Rate

Rate **»Na-3

Scan Scan

(mV/sec) (mV/sec) **»Na-3

—

.

-3.4.5M

-3.4.

Na-5

Rate Rate

*-Na-5 ~ ~ 5M

0.5 0.5

H2S04 H2S04

APPENDIX 5.

Cyclic voltammograms of the molybdenum solutions at different scan rates (20, 50, 100 mV/sec), at a glassy carbon electrode, using 1M sodium molybdate in 4M sulphuric acid, showing the four molybdenum redox couples.

Mo(III)/Mo(IV)

-0.5

-0.5‘

V vs SCE

176 Current (mA) Mo(V)/Mo(VI) Mo(IV)/Mo(V) 177 V

VS

SCE V

vs

SCE