The Electrodialysis of Lithium Sulphate to Lithium Hydroxide ENG470

MURDOCH UNIVERSITY

The Electrodialysis of Lithium Sulphate to Lithium Hydroxide ENG470 – Engineering Honours Thesis

Thesis submitted to the School of Engineering and Information Technology, Murdoch University, to fulfil the requirements for the degree of Chemical and Metallurgical Engineering Honours.

Written by: Hollie Harrison Unit Coordinator: Professor Parisa Arabzadeh Bahri & Dr. Gareth Lee Thesis Supervisors: Dr Aleks Nikoloski

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Author’s Declaration I declare that this thesis is my own account of my research and contains as its main content work which has not previously been submitted for a degree at any tertiary education institution.

Hollie Harrison

I Acknowledgements

Firstly, I would like to sincerely thank my supervisor, Dr. Aleks Nikoloski for his unwavering support and encouragement throughout this project. I would also like to acknowledge and wholeheartedly thank Kwang-loon Ang (Allan) for his hours of work in helping me to complete this research, without whom I would not have been able to complete this thesis. I would also like to thank Jacqueline Briggs for all of her vital assistance and patience with the ion-chromatography machine.

To my friends within the university, I would like to thank you for your help, support and advice, without which this project and university life would have been the most stressful experience of my life. Each one of you has helped me to the best of your ability, and I am so very grateful for that.

To my family for their complete and unconditional love, support and encouragement throughout my entire student career. I would not have had the ability to complete this degree without you, and for that I cannot thank you guys enough.

Lastly to my amazing partner, Bevan Green, thank you for all of the lunches you brought me when I had no time to get food. Thank you for your unconditional support and all of the encouragement you’ve given me over the past year.

II Abstract

There is currently an increasing demand for lithium-ion batteries, and therefore a push within the industry to produce lithium hydroxide. Electrodialysis has been shown to be a promising new technology for producing lithium hydroxide.

A three-compartment batch electrodialysis cell was constructed, utilising an anionic exchange membrane and a cationic exchange membrane. This cell was constructed in order to produce lithium hydroxide from lithium sulphate salt. The cell was run under multiple different conditions to observe the effect that they would have on the recovery of lithium within the lithium hydroxide of the catholyte compartment within the cell. The initial pH of the solution, the temperature of the system, the initial concentration of lithium sulphate and the residence time within the cell were all tested in separate experiments in order to observe how they would influence the system and the production of lithium hydroxide.

The results of this study indicated that by decreasing the initial concentration of the lithium sulphate within the cell, the lithium recovery is dramatically increased, at 30 wt.% lithium sulphate, 18.3% of the lithium is recovered within 4 hours into the catholyte solution as lithium hydroxide. At 5 wt.% lithium sulphate, 81.2% of the lithium is recovered within 4 hours into the catholyte as lithium hydroxide.

The results also suggest, the rate of production of lithium hydroxide is fastest when the residence time within the cell is reduced, however, a longer residence time within the cell will increase the lithium recovery. A 4-hour test at 30 wt.% of lithium sulphate yielded a 23.1% lithium recovery within the catholyte solution. When this residence time was doubled, the recovery was increased to 37% lithium within the catholyte as lithium hydroxide.

III

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IV Table of Contents Author’s Declaration ...... I Acknowledgements ...... II Abstract ...... III Table of Tables ...... VI Table of Figures ...... VII 1. Chapter 1- Introduction ...... 1 2. Chapter 2- Literature Review ...... 3 2.1. Introduction to Lithium ...... 3 2.1.1. Uses for Lithium ...... 3 2.1.2. History of Lithium ...... 5 2.1.3. Lithium Hydroxide (LiOH) ...... 6 2.2. Electrodialysis ...... 7 2.3. Membranes ...... 11 2.3.1. Types of membranes explored ...... 11 2.4. Electrolytic Solutions ...... 20 2.4.1. Catholyte ...... 21 2.4.2. Anolyte ...... 22 2.5. Factors That Can Affect Efficiency ...... 22 2.6. Conclusions and Recommendations ...... 25 3. Chapter 3- Materials and Methods ...... 28 3.1. Solution Preparation...... 28 3.2. Analytical Methods ...... 29 3.3. Experimental Materials and Set-up ...... 30 3.4. Experimental Method...... 35 3.4.1. Preliminary Experiment ...... 35 3.4.2. pH Alteration ...... 37 3.4.3. Temperature ...... 37 3.4.4. Initial Concentration of Li2SO4 ...... 37 3.4.5. Residence Time ...... 38 4. Chapter 4- Results and Discussion ...... 39 4.1. Preliminary Experiment ...... 39 4.2. Effect of pH...... 43 4.3. Effect of Temperature ...... 48 4.4. Effect of Starting Concentration ...... 54 4.5. Effect of Residence Time ...... 62 4.6. General Discussion ...... 70 5. Chapter 5- Conclusion and Recommendations ...... 73 5.1. Future Work ...... 74 6. References ...... 75 7. Appendix ...... 80

V Table of Tables

Table 1: Calculated lithium demand (Basic Scenario) forecast and share in 2020 for different applications Modified from (Martin et al., 2017)...... 4 Table 2: Operating conditions for each experiment ...... 35 Table 3: Mass transfer and recovery, experiment 1: preliminary ...... 43 Table 4: Mass transfer and recovery, experiment 2: pH 11 ...... 47 Table 5: Mass transfer and recovery, experiment 3: pH 7 ...... 48 Table 6: Mass transfer and recovery, experiment 4: 40°C,...... 54 Table 7: Mass transfer and recovery, experiment 5: 60°C,...... 54 Table 8: Mass transfer and recovery, experiment 6: 15 wt.% ...... 61 Table 9: Mass transfer and recovery, experiment 7: 10 wt.% ...... 61 Table 10: Mass transfer and recovery, experiment 8: 5 wt.% ...... 62 Table 11: Mass transfer and recovery, experiment 9 ...... 69 Table 12: Mass transfer and recovery, experiment 10 ...... 69 Table 13: Mass transfer and recovery, experiment 11 ...... 69 Table 14: System mass balance, experiment 1: preliminary...... 84 Table 15: System mass balance, experiment 2: pH 11 ...... 85 Table 16: System mass balance, experiment 3: pH 7 ...... 86 Table 17: System mass balance, experiment 4: 40°C ...... 87 Table 18: System mass balance, experiment 5: 60°C ...... 88 Table 19: System mass balance, experiment 6: 15 wt.% ...... 89 Table 20: System mass balance, experiment 7: 10 wt.% ...... 90 Table 21: System mass balance, experiment 8: 5 wt.% ...... 91 Table 22: System mass balance, experiment 9: 2 hours ...... 92 Table 23: System mass balance, experiment 10: 4 hours ...... 93 Table 24: System mass balance, experiment 11: 8 hours ...... 94

VI Table of Figures

Figure 1: The movement of ions within an electrodialysis cell. Modified from (Mroczek et al. 2015)...... 7 Figure 2: Separation test of Li and Cl ions by electrodialysis. Modified from (Hoshino. 2013)...... 9 Figure 3: Electrodialysis process for turning lithium sulphate into lithium hydroxide and sulphuric acid. Modified from (Ying et al., 2008)...... 10 Figure 4: Recovery ratio (% of lithium recovered) achieved by electrodialysis (a) IL-i- OM and (b) High-durability IL-i-OM. Modified from (Hoshino, 2013)...... 12 Figure 5: Nanofiltration membrane with monovalent ions permeating through the membrane wall. Modified from (Ge et al., 2015)...... 14 Figure 6: nanotube membrane technology in the desalination of water. Modified from (Das et al., 2013)...... 15 Figure 7: Flow sheet for the conventional production of ultra-pure water. Modified from (Xu and Huang, 2008)...... 17 Figure 8: Flow sheet for the production of ultra-pure water utilizing conventional electrodialysis. Modified from (Xu and Huang, 2008)...... 18 Figure 9: Flow sheet for the production of ultra-pure water utilizing electrodialysis with a bipolar membrane. Modified from (Xu and Huang, 2008)...... 19 Figure 10: Electrodialysis cell for the production of lithium hydroxide. Modified from (Ying et al., 2008)...... 21 Figure 11: Electrodialysis cell with cationic and anionic movement of particles. Modified from (Valero et al., 2011)...... 23 Figure 12: Production of lithium hydroxide at varying current densities. Modified from (Ying et al. 2008)...... 24 Figure 13: Energy consumption (squares) and the current efficiency (circles) as related to the current density. Modified from (Ying et al. 2008)...... 25 Figure 14: Front View of Electrodialysis cell ...... 31 Figure 15: Top view of Electrodialysis cell ...... 32 Figure 16: Cathode ...... 33 Figure 17: Experimental Setup ...... 34 Figure 18 Concentration of lithium within the salt and catholyte chambers, experiment 1: preliminary ...... 39 Figure 19: Concentration of sulphate within the anolyte and salt chambers, experiment 1: preliminary ...... 40 Figure 20: Cell voltage, anodic and cathodic potentials, experiment 1: preliminary .. 41 Figure 21 Concentration of lithium within the catholyte and salt chambers, experiment 2: pH 11...... 44 Figure 22: Concentration of lithium within the catholyte and salt chambers, experiment 3: pH 7...... 44 Figure 23: Concentration of sulphate within the anolyte and salt compartments, experiment 2: pH 11...... 45

VII Figure 24: Concentration of sulphate within the anolyte and salt compartments, experiment 3: pH 7...... 45 Figure 25: Cell voltage, anodic and cathodic potentials, experiment 2: pH 11 ...... 46 Figure 26: Cell voltage, anodic and cathodic potentials, experiment 3: pH 7 ...... 46 Figure 27: Concentration of lithium within the catholyte and salt chambers, experiment 4: 40°C ...... 49 Figure 28: Concentration of lithium within the catholyte and salt chambers, experiment 5: 60°C ...... 49 Figure 29: Concentration of sulphate within the anolyte and salt compartments, experiment 4: 40°C ...... 50 Figure 30: Concentration of sulphate within the anolyte and salt compartments, experiment 5: 60°C ...... 51 Figure 31: Cell voltage, anodic and cathodic potentials, experiment 4: 40°C ...... 52 Figure 32: Cell voltage, anodic and cathodic potentials, experiment 5: 60°C ...... 52 Figure 33: Concentration of lithium within the catholyte and salt chambers, experiment 6: 15 wt.% ...... 55 Figure 34: Concentration of lithium within the catholyte and salt chambers, experiment 7: 10 wt.% ...... 56 Figure 35: Concentration of lithium within the catholyte and salt chambers, experiment 8: 5 wt.% ...... 56 Figure 36: Concentration of sulphate within the anolyte and salt compartments, experiment 6: 15 wt.% ...... 57 Figure 37: Concentration of sulphate within the anolyte and salt compartments, experiment 7: 10 wt.% ...... 58 Figure 38: Concentration of sulphate within the anolyte and salt compartments, experiment 8: 5 wt.% ...... 58 Figure 39: Cell voltage, anodic and cathodic potentials, experiment 6: 15 wt.% ...... 59 Figure 40: Cell voltage, anodic and cathodic potentials, experiment 7: 10 wt.% ...... 60 Figure 41: Cell voltage, anodic and cathodic potentials, experiment 8: 5wt.% ...... 60 Figure 42: Concentration of lithium within the catholyte and salt chambers, experiment 9: 2 hours...... 63 Figure 43: Concentration of lithium within the catholyte and salt chambers, experiment 10: 4 hours...... 63 Figure 44: Concentration of lithium within the catholyte and salt chambers, experiment 11: 8 hours...... 64 Figure 45: Concentration of sulphate within the anolyte and salt compartments, experiment 9: 2 hours...... 65 Figure 46: Concentration of sulphate within the anolyte and salt compartments, experiment 10: 4 hours...... 65 Figure 47: Concentration of sulphate within the anolyte and salt compartments, experiment 11: 8 hours...... 66 Figure 48: Cell voltage, anodic and cathodic potentials, experiment 9: 2 hours ...... 67 Figure 49: Cell voltage, anodic and cathodic potentials, experiment 10: 4 hours ...... 67 Figure 50: Cell voltage, anodic and cathodic potentials, experiment 11: 8 hours ...... 68

VIII Figure 51: Current efficiency and the lithium recovery of experiments 1-11...... 71 Figure 52: Lithium hydroxide and sulphuric acid production rates from each test ..... 71 Figure 53 Middle compartment of electrodialysis cell ...... 81 Figure 54: Dimensions of Electrodialysis Cell ...... 82 Figure 55 Apparatus placement within the electrodialysis cell ...... 83

1. Chapter 1- Introduction

At the beginning of 2017, the largest global use for lithium was batteries (Unites

States Geological Survey, 2017; Martin et al., 2017). These batteries are used in handheld devices, computers and other products where a lead based battery is heavy and impractical. Currently there is a major push in the research field to investigate new or improved ways to increase our ability to have portable power. Lithium is also being looked at in order to develop car batteries for electric and hybrid cars in the future (Hoshino, 2014; Hwang et al., 2016; Tahil, 2007). Their appeal comes from their ability to store more energy within handheld devices with fewer charges (Tahil,

2007), and this innovation could potentially be transferred into the new car battery technology.

Lithium recovery from lithium chloride salts has been researched extensively.

However lithium chloride resources are becoming limited and other means of lithium recovery must be looked into (Hoshino, 2014). Sulphate brines are in abundance, but the technology to produce lithium hydroxide from sulphates has not yet been properly established (Hoshino, 2014). It has been suggested that electrodialysis could possibly be a relatively simple and cost effective method for producing lithium hydroxide from lithium sulphate.

The aim of this study is to improve technology to produce lithium hydroxide. In order to do this, electrodialysis will be carried out on lithium sulphate salt to determine:

1. That electrodialysis is a possible method for producing lithium hydroxide from

lithium sulphate.

1 2. The effect that the pH of the solution has on the production of lithium hydroxide.

3. The effect of temperature on the production of lithium hydroxide.

4. The effect of the initial concentration on the production of lithium hydroxide

5. The effect on the production of lithium hydroxide in relation to residence time

within the cell.

2 2. Chapter 2- Literature Review

2.1. Introduction to Lithium

Lithium is a resource that is increasingly becoming more popular as new technologies are being developed to incorporate lithium ion batteries into their functionality. To decrease the cost these new technologies, and therefore make them more appealing to consumers, the cost of production of lithium needs to be reduced, it has been suggested that this could be done through electrodialysis (Ying et al., 2008;

Hoshino, 2013; Hoshino, 2014). Currently, the electrodialysis of lithium chloride has been investigated and extensively researched, making the process of extracting lithium from chloride brines efficiently methodized, however the chloride brine resource is becoming increasingly limited in its natural economical supply, whereas sulphate brines are plentiful, but the research towards the extraction of lithium from these brines is yet to be determined fully (Hoshino, 2014).

Lithium is found in hard rock, such as pegmatites. These igneous rocks are formed by crystalized magmatic fluid, forming minerals that contain lithium such as spodumene, lepidolite and petalite (Evans, 2008; Tahil, 2007). Lithium is also present in the form of brines in salt lakes and additionally can be found in seawater. This occurs when hard rocks are leached and the concentration of the brine and seawater can vary greatly depending on the location of the sample (Evans, 2008; Tahil, 2007).

2.1.1. Uses for Lithium

In 1976 a National Research Council Panel estimated that the demand for lithium was approximately 3,200 tonnes per year. In 2008, the demand equated to approximately 16,000 tonnes of elemental lithium (Evans, 2008). In 2015, 35% of lithium consumed that year went towards batteries, with ceramic and glass

3 applications being the second biggest use for lithium, consuming 32%. The demand for lithium in this year was calculated to be approximately 173,000 tonnes and in

2020 the basic demand for lithium is forecast to be around 270,000 tonnes (Martin et al., 2017).

Table 1: Calculated lithium demand (Basic Scenario) forecast and share in 2020 for different applications Modified from (Martin et al., 2017).

Demand 2015 Share 2015 Demand 2020 Share 2020 [t] [%] [t] [%] Batteries 53,629 35 76,673 34 Glasses and 0,549 32 86,717 38 Ceramics Lubricating 13,840 9 14,507 7 greases Polymers 8,650 4 2,315 1 Air 8,650 5 5,325 2 conditioning Aluminium 1,730 1 0 0 Continuous 10,380 5 18,478 8 casting Other 15,570 9 22,235 10 Sum 172,998 100 226,250 100

Today lithium ion batteries are being used in a rage of technologies, lithium is being used as opposed to lead due to its lower density, making it a much lighter battery, therefore a more attractive alternative to the lead acid battery. Lithium-ion batteries are also greatly attractive due to their ability to store more energy, allowing portable devices to last longer with fewer charges (Tahil, 2007). Lithium is also being used in a range of medical devices as lithium micro-batteries (Hwang et al., 2016).

As a means of combatting the issue with the depleting fossil fuels resource, and the growing environmental crisis of global warming, lithium-ion batteries are increasingly being developed and utilized in hybrid cars (Hoshino, 2014; Hwang et

4 al., 2016; Tahil, 2007). Lithium is also widely used in glass and ceramics production, this application of lithium is the second largest next to batteries in the industry

(Martin et al., 2017). In glasses and ceramics, lithium improves their durability when temperature is involved (Dakota Minerals, 2017). Not only does lithium increase the performance of glass and ceramics in terms of their thermal qualities, it also enhances the mechanical strength of ceramics and the colourfastness of glasses (Martin et al.,

2017).

2.1.2. History of Lithium

Lithium was first discovered in the 1790s in the form of the mineral petalite (Royal

Society of Chemistry, 2017). Traditionally, Lithium has been recovered from salt lakes or in other words, brines. South America produced a large amount of the world’s lithium, with two of its salt lakes, one in Argentina, the other in Chile, these

Salt Lakes produced approximately 70% of the world’s lithium (Hoshino, 2013;

Hoshino, 2014).

Lithium Reserves are part of lithium reserve bases in which lithium can be economically produced from at the time of production, they denote the holding of realistic recoverable lithium (Tahil, 2007). The term, reserve base is the identified source of lithium; it includes the lithium reserves, i.e. the economically recoverable lithium, the marginal reserves, the lithium that is only marginally economical to produce and the sub-economic reserves of lithium. These reserves will only be economical to produce in the event that new technologies are developed or the global price of lithium rises sufficiently in order to extract the lithium in these reserves economically (Tahil, 2007). However, the industry cannot rely on lithium prices to increase in order to have the ability to economically extract from the currently sub-

5 economic reserves, therefore the industry must find alternative, economical methods of production.

2.1.3. Lithium Hydroxide (LiOH)

Lithium hydroxide (LiOH), in its anhydrous form (containing no water) is a white crystalline (solid) substance that is soluble in water and will produce an alkaline

(basic) liquid. Lithium hydroxide is insoluble in ether and only slightly soluble in ethanol. When in its crystalline form, lithium hydroxide has a melting point of 450°C and will decompose at 924°C (Daintith, 2008). Reacting lithium salts or lithium ores with lime can make this particular compound. Lithium hydroxide can also be produced by reacting lithium metal or lithium hydride with water, however this reaction is exothermic and therefore quite aggressive (Daintith, 2008).

Lithium hydroxide is one of the materials used in the production of electric vehicle batteries. As such, as the demand for electrically run vehicles increases, so does the demand for lithium hydroxide (Warburton, 2016). Lithium hydroxide was previously produced by aqueous causticisation reactions between lime, which is produced by hydrating calcium oxide with water, and lithium carbonate. However, as one of the main applications of lithium hydroxide is in the production of lithium batteries, in which the lithium hydroxide is used to produce the cathode material within the battery. The lithium hydroxide produced needs to be of battery grade, meaning that it needs to be very pure, and almost completely free of contaminants. Producing lithium hydroxide through causticisation means that obtaining a battery grade lithium hydroxide product is problematic (Buckley et al., 2011).

Usually, the lithium hydroxide produced through causticisation is obtained from spodumene ore or brine water in which the lithium is present as a salt, quite usually lithium chloride or lithium sulphate (Sharma, 2016). Lithium Hydroxide can also be

6 produced through converting lithium chloride into lithium carbonate utilizing soda ash (Sharma, 2016).

2.2. Electrodialysis

A physical process would be a process in which lithium ions are physically or mechanically separated from a substance or compound, these methods are typically not the most accurate of separation methods, allowing for other metallic ions to be recovered alongside lithium. However, chemical processes in comparison, are a lot more selective than physical processes (Hwang et al., 2016).

Figure 1: The movement of ions within an electrodialysis cell. Modified from

(Mroczek et al. 2015).

Electrodialysis was invented in the 1950s in order to desalinate brackish water

(Valero et al., 2011; Reahl, 2006). It is essentially an extension of electrolysis, which

7 is an electrochemical process in which ions in solution are passed to an anode or cathode, oxidation of the solution occurs at the anode and reduction occurs at the cathode.

The system incorporates ion exchange membranes; cationic membranes and anionic membranes can be used. Usually these membranes are alternated with 3 fluid streams, the dilute stream, that contains the substance to be extracted, the concentrate and the electrolyte that can be described as the catholyte and the anolyte. The electrolytic liquid provides the ions with a mode of transport between the semi- permeable membranes (Mroczek et al., 2015). Figure 1 illustrates the movement of anions through the anionic membranes and the movement of cations through the cationic membrane.

The chlor-alkali industry had adapted diaphragm cells that were used to produce chlorine and caustic soda, to produce electrodialysis cells. Instead of having a diaphragm, an ionic membrane known as an ion-exchange membrane is used

(O’Brien et al., 2005). Diaphragms were originally made of asbestos; the anode of the cell would be placed between the two diaphragms and a copper gauze after each diaphragm acted as the cathodes of the system (O’Brien et al., 2005).

In earlier years, electrodialysis was initially used to produce sodium hydroxide from rock salt, sodium chloride. Initially, sodium hydroxide was produced by electrolysis, which had been experimented with. Diaphragms and mercury cathodes were explored in order to produce other products rather than just sodium hydroxide and chlorine gas (Mazrou et al., 1997). Thus electrodialysis utilizing anionic and cationic exchange membranes was able to produce not only sodium hydroxide, but hydrochloric acid as well (Mazrou et al., 1997).

8 The pH, voltage, flow rate and the number of membranes in the electrodialysis cell are the numerical factors, along with the electrolyte used, that will greatly influence the recovery of lithium ions. Therefore, these factors need to be optimised in order to efficiently and economically recover lithium ions through electrodialysis (Hwang et al., 2016). For example, Figure 2 shows the separation of lithium and chloride utilizing cationic and anionic membranes, with an anolyte of water and a catholyte of hydrochloric acid in an electrodialysis cell.

Figure 2: Separation test of Li and Cl ions by electrodialysis. Modified from

(Hoshino. 2013).

As can be seen in Figure 2, lithium ions will penetrate through to the cathode through the cation exchange membrane, while the chloride ions will permeate towards the anode through the anion exchange membrane.

Figure 3: Electrodialysis process for turning lithium sulphate into lithium hydroxide

and sulphuric acid. Modified from (Ying et al., 2008).

Figure 3 illustrates a continuous electrodialysis cell in which lithium sulphate salt is put into solution to produce lithium hydroxide and sulphuric acid. The dilute stream in this particular example would be the lithium sulphate solution, while the concentrate would be the lithium hydroxide that is being produced, and the electrolytes used in this cell were lithium hydroxide and sulphuric acid. The Li2SO4 is pumped into the compartment between the cationic membrane (CEM) and the anionic membrane (AEM). The lithium ions within the solution will them permeate through the cationic membrane towards the anode to produce lithium hydroxide along with hydrogen gas. The sulphate ions will pass through the anionic membrane to produce sulphuric acid and oxygen gas. Figure 10 depicts a simpler schematic of this cell

(Ying et al., 2008).

The reactions taking place within the cell are as follows:

Overall reaction: Li2SO4 (aq) + 2H2O  H2SO4 (aq) + 2LiOH (aq) (1)

Half-cell reactions:

+ - Anode: H2O (aq)  2H (aq) + 1/2O2 (g) +2e (2)

- - Cathode: 2H2O (aq) +2e  H2 (g) +2OH (aq) (3)

10 2.3. Membranes

Membranes can be used in a wide range of technologies; they are not limited to their use in electrodialysis. Other uses include gas separation and simple physical separation of particles. Gas separation, employs specialty robust and highly selective membranes that are used in order to economically separate certain gasses. Gas separation membranes can be used in fuel cells for cars or other vehicles, and reactors that utilize membranes for the production of hydrocarbons (Koros. 2002).

The industrial application of membranes first started in 1950, this was when artificial membranes were invented (Tanaka et al., 2012). Membranes become

‘stacked’ alternating between cation and anion specific membranes. Depending on the membrane, only certain ions will be able to permeate through the membrane. The way the membranes are layered in the stack will also determine which ions will be extracted from the dilute stream into the concentrate (Mroczek et al., 2015).

Different ion-exchange membranes have different permselectivities, this means that the membrane is selective in the cations or the anions that can pass through. This broadly has to do with whether they are monovalent or multivalent cations and anions

(Mroczek et al., 2015; Ball and Boatang., 1987). Lithium, being a monovalent cation can be separated from multivalent cations by using a permselective membrane.

2.3.1. Types of membranes explored

Although ion-exchange membranes are the conventional membranes used when running an electrodialysis cell, bipolar membranes can be used in electrodialysis in place of or in addition to ion-exchange membranes.

2.3.1.1. Ion-Exchange Membranes

Ion-exchange across membranes had initially been investigated through the use of biological membranes prior to the invention of artificial ion-exchange membranes

11 (Tanaka et al., 2012). This specific type of membrane is now the primary membrane used in the purification of water and demineralisation industries. Ion-exchange membranes can also be used in the treatment and recycling of sewage water for reuse within households and membrane reactors (Tanaka et al., 2012).

Figure 4: Recovery ratio (% of lithium recovered) achieved by electrodialysis (a) IL-i-

OM and (b) High-durability IL-i-OM. Modified from (Hoshino, 2013).

Ion-exchange membranes are a type of polymeric membrane in which the polymer matrix has charged groups attached (Rottiers et al., 2015). Hoshino, (2013) used ionic liquid impregnated organic membranes, IL-i-OM (Gore-TexTM) and high-durability

IL-i-OM (Nafion 324) membranes in order to recover lithium from sea water, the recovery of unwanted minerals and the recovery of lithium was then calculated and recorded. In Figure 4 it can be seen that the high-durability membrane recovered more lithium in the same amount of time as opposed to the normal IL-i-OM. Nafion membranes tend to have a high durability and have the ability to be subjected to harsh environments while retaining their ion-exchange properties (O’Brien et al., 2005).

Anionic exchange membranes allow negatively charged ions to permeate through the membrane as the groups attached to the polymer matrix within the membrane are positively charged (Rottiers et al., 2015).

Mroczek et al., (2015) had originally purchased Nafion membranes, however they had to cut them to shape and found that procedure to be tedious and inaccurate.

Instead, they were able to use a PCCell electrodialysis system that had an anionic and cationic membrane provided.

Nie et al., (2017) utilized and Asahi Glass Selmion ASA anionic membrane to allow for the anions in the feed solution to migrate through a 40 cell stack to the anolyte.

Cationic exchange membranes allow positively charged ions to permeate the membrane as the groups attached to the polymer matrix within the membrane are negatively charged (Rottiers et al., 2015).

Hoshino, (2013) recovered the lithium in the form of lithium chloride. Later it was stated that a cation exchange membrane was used in order to allow the lithium ions to

13 permeate through to the concentrate. An anon exchange membrane was used for the chloride ions to permeate through using a 0.1M HCl solution.

A SELMIONTMCMV membrane allows cations, such as lithium to permeate through to the cathode side of the cell. At the same time, it prevents ionic liquid and water from permeating through to the concentrate (Hoshino, 2014).

2.3.1.2. Non-ionic membranes

Non-ionic membranes are membranes that do not require charged particles to function, their selectivity is non-ion specific. Instead their permselectivity is based on other traits such as particle or molecule size, organic or inorganic etc. Nanofiltration membranes are a form of membrane that do not work in the same way as ion- exchange membranes. Instead of allowing an ion of a specific charge to permeate the membrane, monovalent ions are instead allowed to permeate through the membrane while other multivalent ions are unable to permeate (Ge et al., 2015).

Figure 5: Nanofiltration membrane with monovalent ions permeating through the

membrane wall. Modified from (Ge et al., 2015).

Ultrafiltration membranes are another type of non-ion-exchange membrane that have been used in electrodialysis cells. Serre et al., (2016) utilized an ultrafiltration

14 membrane to neutralize the organic acids that are retained in cranberry juice in an attempt to reduce the acidity of the juice. This was explored, as raw cranberry juice is too acidic to be deemed consumable by the market.

Carbon nanotube membranes have been used to purify saline water. These membranes are being explored due to the depleting amount of fresh water that is in existence and accessible at this point in time. Global warming is a big factor in the ever-increasing contamination of fresh water with salts, as fresh water is vital in order to produce food and other commodities such as lithium, it is important that other means of desalination be explored (Das et al., 2013).

Figure 6: nanotube membrane technology in the desalination of water. Modified from

(Das et al., 2013).

The nanotubes are made from sheeted graphite that are subsequently rolled up into a tube. When the sheets are rolled, they are either rolled singularly or rolled up with

15 multiple sheets to produce a nanotube with multiple layers. The water molecules will pass through the nanotube membranes while the salts in the water will be retained within the membrane (Das et al., 2013). However, as this type of membrane is permselective only to water, its application in industry is limited to only water purification until such time wherein other potential uses could be further investigated.

2.3.1.3. Bipolar Membranes

Bipolar Membranes comprise of two layers, a cationic- exchange and an anionic- exchange layer. These features give the bipolar membranes the ability to split solvents into their sub-part. For example, water can be split into H+ and OH- (Xu and Huang,

2008). However bipolar membranes are not limited to electrodialysis, they can be utilized in food processing, food control and chemical or biochemical synthesis (Xu and Huang, 2008).

Hwang et al., (2016), using a bipolar membrane, Neosepta BP-1 together in alternating stacking with a cation-exchange membrane Neosepa CMX, was able to produce hydroxyl and hydrogen ions. This is due to water splitting in the catalytic layer occurring when voltage was applied to the system. The hydroxyl ions together with the lithium ions within the feed solution then produce lithium hydroxide.

Ultrapure water production is a prime example of the ability for electrolysis and electrodialysis to simplify a conventional process. See Figure 7, Figure 8 and Figure

Figure 7: Flow sheet for the conventional production of ultra-pure water. Modified

from (Xu and Huang, 2008).

Figure 7 illustrates the conventional method in which ultra-pure water is produced.

Initially, the feed water is fed into microfiltration that is then passed through to the softener and into a storage tank. The water then undergoes reverse osmosis before going into a degasifying column. The water is then put through another round of reverse osmosis and then into UV-sterilization. Once it has been sterilized, it’s put through a mix-bed ion exchange and subsequently undergoes ultrafiltration to then be stored and in-situ filtered to produce ultra-pure water.

Figure 8: Flow sheet for the production of ultra-pure water utilizing conventional

electrodialysis. Modified from (Xu and Huang, 2008).

Figure 8 depicts the process that is used to produce ultra-pure water utilizing electrodialysis. Compared to Figure 7, this process is condensed, requiring fewer unit operations such as ultrafiltration and in-situ filtration. The feed is pre-treated before going into microfiltration and de-gassing. It is then UV- sterilized and put through reverse osmosis, entering the final step of the process where the water is purified and

18 stripped of any salts by electrodialysis. The feed comes into the middle of the cell. As can be seen in Figure 8, the chloride ions in the solution migrate towards the anode and the sodium ions migrate towards the cathode. Within this particular cell, as shown in Figure 8, there are 2 cationic-exchange membranes and 2 anionic-exchange membranes. By having a second set of membranes within the cell, the chances of impurities permeating the membranes into the concentrate are minimized.

Figure 9, fits into the flow sheet of Figure 8, however a bipolar membrane is present in this particular cell. By having the bipolar membrane present, the voltage drop in the system is minimized resulting in better energy efficiency (Xu and Huang.,

2008).

Figure 9: Flow sheet for the production of ultra-pure water utilizing electrodialysis

with a bipolar membrane. Modified from (Xu and Huang, 2008).

19 When using bipolar membranes in an electrodialysis system (BMED), gas production is reduced, energy consumption of the system has added efficiency. The installation and performance of the system are also increased. In addition, their compact size makes them convenient and versatile in their application (Xu and

Huang., 2008). While bipolar membranes improve the performance of electrodialysis, they also increase the cost of the system (Wang et al., 2010).

2.4. Electrolytic Solutions

The electrolytic solution in electrodialysis is the medium in which the ions travel from the feed solution to the anode and the cathode through the membranes within the electrodialysis cell. The electrolytic solution can be broken into two solutions known as the catholyte and the anolyte.

In order for lithium ions to be liberated from lithium manganese oxide, they need to be replaced with hydrogen ions, therefore hydrochloric acid was used in this particular process. However, because of the characteristics of the bipolar membrane used, water can also be used to produce lithium hydroxide. This is because the water splitting within the catalytic layer of the membrane produced the hydroxyl ion and hydrogen (Hwang et al., 2016).

Figure 10: Electrodialysis cell for the production of lithium hydroxide. Modified from

(Ying et al., 2008).

As can be seen in Figure 10, water was used along with a cationic and anionic membrane. A 1.0 mol/L solution of lithium sulphate salt is placed in the cell between the two membranes. The anolyte and the catholyte are both water, the sulphate ions in the lithium salt permeates through the anionic membrane into the anolyte to produce sulphuric acid and oxygen through oxidation. While the lithium ions will permeate the cationic membrane to produce lithium hydroxide and hydrogen gas through reduction

(Ying et al., 2008).

2.4.1. Catholyte

In the case of a 3 or more compartment cell, the catholyte resides in the compartment of the cell that the positively charged ion has permeated through the membrane into. This is where the cation will migrate to and become most commonly a hydroxide. Figure 10 illustrates this; the catholyte resides in the same compartment of the electrodialysis cell as the cathode. This is also true of a 2-compartment electrodialysis cell if the membranes being used are cationic. However, if the

21 membranes being used are anionic, then the cathode will be submerged in the original salt solution, with an anionic membrane either side of the catalytic compartment. For example, lithium sulphate is being used to produce sulphuric acid and lithium hydroxide, with sulphuric acid being the main product of the process. The lithium ions will stay within the catalytic compartment as they cannot permeate through the anionic membranes and the sulphate ions will permeate through to the anolyte to react and produce sulphuric acid (the concentrate).

2.4.2. Anolyte

The anolyte resides in the compartment of the anode; usually the anolyte within the cell is the feed containing the lithium to be extracted from the lithium salt. For a 3- compartment cell, the anolyte, the catholyte, the dilute and the concentrate would be

3-4 different solutions. The anolyte and the catholyte can be the same solution that will react to produce two different solutions. For example, in Figure 10, the catholyte and the anolyte are both H2O. However, in order for the water in the cell to become sufficiently ionised to carry a charge in order to allow the process to eventuate, there must be some lithium hydroxide already present in the catholyte. In this case some sulphuric acid present in the anolyte. If this were a 2-compartment cell however, the sulphuric acid would be produced within the dilute solution, in Figure 10, this solution is labelled salt, as the anode would be submerged in the lithium sulphate solution, thus making it the anolyte and already conductive.

2.5. Factors That Can Affect Efficiency

The voltage applied to the electrodialysis cell is the driving force of the entire process. Direct voltage is applied to the system to drive the anions to the anode and the cations to the cathode (Valero et al., 2011), as illustrated in Figure 11.

Figure 11: Electrodialysis cell with cationic and anionic movement of particles.

Modified from (Valero et al., 2011).

Direct current (DC) is used as this means that the current only flows in one direction. Alternating current (AC) would result in the current being supplied to the system would reverse periodically, therefore reversing the voltage in the system, and there would not be a continuous current running through the system. Therefore, this would have an effect on the transfer of the lithium ions.

While increasing the voltage of the system will increase the transfer rate of the lithium within the cell, the membranes will be detrimentally affected by this increase in power supply (Mroczek et al., 2015). Therefore, a balance between the voltage

23 supplied to the system and the degradation rates of the membranes must be optimized in order to economically produce lithium ions.

Hoshino (2014) utilized a voltage of 2 V to concentrate the lithium ions to the cathode side of the cell, allowing the concentration to increase with time. After 2 hours of this applied dialysis voltage, the concentration of lithium ions had reached

24.5% recovery. Another study found that a voltage 6.5 V per membrane, using a bipolar membrane, with all flow rates of 0.44 mL/(cm2min) yielded extraction rates of lithium manganese oxide to be approximately 70% (Hwang et al., 2016).

The pH of the initial salt solution is one of the factors that may have an effect on the rate at which lithium will be transferred from the feed. With a lower pH, the rate of transfer of lithium from the lithium salt was increased (Hwang et al., 2016).

Mroczek et al., (2015) observed that at a pH range of approximately 2-4, the optimal transfer rates were achieved using a 3-membrane stack in an electrodialysis unit. The highest transfer rate obtainable under these parameters was 0.28 mg/(hour.cm2).

1.2 700 A/m^2

0.9 1000 A/m^2 0.6

1400 A/m^2 0.3

LiOH Concentration LiOH Concentration (mol/L) 0 1800 A/m^2 0 100 200 300 400 Time (min)

Figure 12: Production of lithium hydroxide at varying current densities. Modified

from (Ying et al. 2008).

24 100 15

80 12

60 9

Current efficiecy 40 6 Energy consumption

Current Current eficiency (%) 20 3

0 0 Energy Consumption (kWh/kg(LiOH)) 400 800 1200 1600 Current Density (A/m^2)

Figure 13: Energy consumption (squares) and the current efficiency (circles) as

related to the current density. Modified from (Ying et al. 2008).

Figure 12 illustrates that higher current densities will improve the rate at which lithium hydroxide is produced. However, increasing the current density results in increasing the energy consumption, which will ultimately increase the cost of production. As can be seen in Figure 13, 500 A/m2 was found to have the optimal current density as there was a current efficiency of 80%.This current density also had the lowest energy consumption for this particular experiment at 6 kWh/kg of lithium hydroxide (Ying et al., 2008).

2.6. Conclusions and Recommendations

Whilst electrodialysis has been used and proven to work for multiple salts and other processes, the electrodialysis of lithium sulphate to lithium hydroxide has not been extensively covered. Lithium hydroxide has been produced in a 3-compartment electrodialysis cell using lithium sulphate and water; see Figure 3 and Figure 10.

However, as the reaction between sulphate and water will produce sulphuric acid, this

25 adds some complication if a 2-compartment cell is being used. Sulphuric acid will be produced within the anolyte compartment where the dilute solution is residing.

Considering this, the compartment that will retain the sulphuric acid must have the ability to withstand the acidity of the sulphuric acid in the concentration at which it will be produced. As the sulphuric acid would be produced while the dilute solution is being split into lithium and sulphate ions, the process would need to be a batch process as it would be burdensome to attempt to remove the sulphuric acid without affecting the amount of lithium ions within the solution from permeating a cationic membrane.

While water can be used as an electrolyte, more specifically, a catholyte, the conductivity will be increased if a dilute solution of lithium hydroxide is used in order to drive the process forward to produce lithium hydroxide from the lithium sulphate.

Membrane selection is a factor that will have an impact on the production of lithium hydroxide, to what degree is unknown. Nafion membranes have been used in the past and have been shown to be durable and effective (O’Brien et al., 2005;

Hoshino, 2013). However, bipolar membranes have also been becoming increasingly popular in the electrodialysis field of study, their ability to increase the efficiency of the system is an attractive feature that they offer, however they do not seem to eliminate the need for any other ion-exchange membrane and may not work in a 2- compartment electrodialysis cell.

The temperature at which the process is run will have an effect on the kinetics of the system, however this has not been a factor researched to a great extent.

Temperature should therefore be varied in ongoing experiments to investigate whether it has any effect on the production of lithium hydroxide and/or the efficiency of the system. The temperature will also have an effect on the corrosive nature of the

26 sulphuric acid, which could detrimentally affect the membrane durability and the analytic compartment within the cell.

Current density through the electrodes will also influence the production of lithium hydroxide. Investigating how optimal current densities coincide with the optimal temperature and voltage of the system and their effect on the current efficiency can also be taken into consideration in future research of the production of lithium hydroxide from lithium sulphate through the means of electrodialysis.

Future research should consist of lithium hydroxide being produced from lithium sulphate. The temperature of the system should be explored in order to observe how it influences the efficiency of the system, whether it has a positive or a negative effect on the transfer of lithium ions through the cationic membranes and the production of lithium hydroxide. The current density alongside the voltage supplied to the system can also be manipulated to observe the effects of ion transfer across the membrane and the production of lithium hydroxide. Different membranes can also be tested and compared once the optimal parameters have been set in order to see if different membranes produce better results in terms of the production of lithium hydroxide.

27 3. Chapter 3- Materials and Methods

Eleven experiments were carried out in this study. The first experiment was to determine whether the electrodialysis cell would produce any lithium hydroxide when run under normal conditions. Subsequent experiments were then carried out to observe the effects of pH, the effect of the temperature, initial concentration and the residence time within the cell.

3.1. Solution Preparation

Three solutions were prepared before each experiment. The first of the solutions is the anolyte, which was made up of 5 wt.% of sulphuric acid. Measuring 102 gram of

98% concentrated sulphuric acid solution into a 2 L volumetric flask does this. The solution is then made up with deionised water to the mark on the flask and thoroughly mixed.

The next solution to be made is the catholyte, which also a 5 wt.% solution, although this solution is made up of 5 wt.% lithium hydroxide monohydrate. 100 grams is measured into a 2 L volumetric flask and deionised water is then added to sufficiently dissolve the material. Once dissolved, the solution is then made up with deionised water to the mark on the volumetric flask and thoroughly mixed.

The final solution is the salt solution, which was made of 30 wt.% lithium sulphate. This is made using either 349.17 g of lithium sulphate monohydrate, or 300 g of lithium sulphate. The lithium sulphate monohydrate is easier to dissolve than the lithium sulphate.

When the materials were hard to dissolve, a stirrer bar was placed inside the volumetric flask and left to dissolve on a hot plate set to stir. This could take multiple days in the case of the lithium sulphate.

28 3.2. Analytical Methods

The concentration of the lithium present in each of the salt and catholyte samples was determined using an Inductively Coupled Plasma Mass Spectrometry (ICP-MS) machine. This was done by carrying out a 1,000,000 times dilution, 0.1 mL was diluted into 10 mL for each sample, this was then repeated twice more.

The sulphate concentration both within the anolyte and the salt samples was determined using Ion Chromatography. This was done by carrying out a 10,000 times dilution by diluting 10 μL into 100 mL and then transferred into the instrument.

An acid-base titration was also carried out in order to determine the concentration of acid produced within the anolyte. This was done using 0.1 M sodium hydroxide

(NaOH), a 1 mL sample of the anolyte (H2SO4) was placed in a beaker and diluted with approximately 10 mL of deionized water. A pH meter was placed in the diluted sample to read the pH. As the NaOH is added, the pH increased, when the pH changed to read above 7, the amount of titrant was recorded. This was then used to calculate the concentration of acid, the equations for this can be found in the appendix.

Both the cell voltage and the cell current were measured and recorded at regular sampling intervals with a multimeter that was clipped to the anode and cathode at the top of the cell. The current and voltage on the power supply were also recorded at these intervals.

An Hg/HgSO4 in 3 M K2SO4 reference electrode was used connected to a multimeter to measure the cathodic and anodic potentials. The reference electrode was immersed into a Luggin capillary, which was held close to the surface of the immersed cathode and anode. The measurement was taken and recorded at every sampling interval.

29 Three separate thermometers were submerged within each compartment of the cell and the temperature was read and noted at every sampling interval.

3.3. Experimental Materials and Set-up

Each of the 11 experiments carried out using an electrodialysis cell consisting of 3 compartments all made of acrylic that is held together by silicone glue. The cell was equipped with 2 membranes a Fumasep FAB-PK-103 anionic membrane and a

Fumasep FKB-PK-130 cationic membrane. The membranes are attached to the middle compartment in circular windows on either side of the centred compartment, the membranes are then held in place with nylon screws. Figure 53 (in Appendix) shows the dimensions of the middle compartment of the cell. In order to prevent leakage, or damage to the membrane, a rubber washer was used to separate the membrane from the acrylic of the cell, and then another rubber washer was used between the membrane and a silicone washer. This was all screwed into place tightly to ensure no leakage. These membranes must be kept wet at all times and need to be stored in deionised water when not in use.

Figure 54 (in Appendix), which can be found in the appendix, illustrates the dimensions of the cell and the fixtures for the apparatus that is to go into the cell. The height of the cell without the top covers is the same as the height of the middle cell,

17.3 cm. The middle compartment is capable of holding approximately 0.7 L while the left and right compartments are capable of holding approximately 1.9 L each.

Figure 14 shows the experimental setup of the electrodialysis cell, the catholyte resides in the left compartment of the cell, the salt resides in the middle compartment, and the anolyte resides in the right compartment.

Figure 14: Front View of Electrodialysis cell

Figure 15: Top view of Electrodialysis cell

Figure 15 shows the top view of the electrodialysis cell and where all of the apparatus is placed. Figure 55 shows a schematic of Figure 15. The apparatus consists of:

 Three condensers, which are placed in the back left opening of the left

compartment, the back opening of the middle compartment and the back right

opening of the right compartment.

 Two aquarium heaters, which are placed in the openings below the condensers in

the left and right compartments of the cell.

32  Two temperature sensors, which are placed to the right and the left of the

aquarium heaters in the left and right compartments respectively.

 Three thermometers, which are placed in the bottom left opening of the left

compartment, the bottom opening of the middle compartment and the bottom right

of the right compartment.

 Two aeration tubes, placed in the back right of the left compartment and the back

left of the right compartment.

 One N2 gas sparging tube placed into the middle opening of the middle chamber to

provide a cooling effect.

 Two luggin capillaries, which are placed in the remaining openings beside the

anode and the cathode.

The cathode resides in the left compartment with the catholyte, while the anode resides in the compartment with the anolyte, left and right compartments respectively as seen in Figure 15 The cathode and the anode are made of titanium (Ti) mesh as can be seen in Figure 16. The anode however, is coated in iridium oxide (IrO2).

Figure 16: Cathode

The entire experimental setup can be seen in Figure 17. The power supply, placed below the fume hood is connected to the anode and cathode; the anode is connected to the positive output while the cathode is connected to the negative output. The aquarium heaters are both connected to the controllers to the right of Figure 17 with the control for the pump to the aerator sitting on top of the aquarium controllers. The pump to the aerators can be seen in Figure 14 and Figure 15 behind the cell. Tubing is run from the bath around the cell to syphon water out from the ice bath to maintain the temperature of the cell.

Figure 17: Experimental Setup

34 3.4. Experimental Method

Table 2 summarises all of the operating conditions for all experiments 1-11.

Experiment 1 was the preliminary experiment in which the experiment was conducted under the standard conditions as a base line. This involved a pH of 2, a temperature range of 20-30 °C, an initial lithium sulphate concentration of 30 wt.% with a residence time of 4 hours at a current of 3 A.

Experiments 2 and 3 were conducted with the initial salt solution’s pH altered to

11 and 7 respectively. Experiments 4 and 5 were conducted with elevated temperatures, while experiments 6, 7 and 8 were conducted with the initial salt concentration decreased. Finally experiments 9, 10 and 11 were conducted to determine what affect the residence time within the cell had on the recovery of lithium hydroxide.

Table 2: Operating conditions for each experiment

Experiment pH Temperature Salt (Li2SO4) Residence Current ID No. range (°C) concentration time (hour) (A) (wt.%) 1 2 20-30 30 4 3 2 11 20-30 30 4 3 3 7 20-30 30 4 3 4 2 35-45 30 4 3 5 2 50-60 30 4 3 6 2 20-30 15 4 3 7 2 20-30 10 4 3 8 2 20-30 5 4 1.8-2.7 9 2 20-30 30 2 3 10 2 20-30 30 4 3 11 2 20-30 30 8 3

3.4.1. Preliminary Experiment

The electrodialysis cell was set up as shown in Figure 17. For each experiment, 1 L of 5 wt.% LiOH solution was added to the left compartment of the cell and 1 L of 5

35 wt.% H2SO4 solution was added to the right compartment. The Li2SO4 salt solution was added to the middle compartment in a quantity of 0.4 L at 30 wt.%. Before setting the power source, the water was turned on to run through the condensers and the pump was turned on for the aeration to provide sufficient mixing of the solutions.

The power supply, a galvanostat, was then set to output a constant current at its maximum of 3 amperes. The power supply and the timer were simultaneously started and the first samples were taken. Samples of 3 mL are taken at 30-minute intervals from each compartment of the cell. Measurements are also taken and recorded at these intervals; the current and the voltage being displayed on the power supply were recorded along with the measured value taken from a multimeter. The anodic and cathodic potentials were also recorded along with the temperature from each compartment of the cell.

Throughout the duration of the experiment the temperature needed to be monitored and kept within the 20-30°C range. In order to do this, ice was placed in a bath around the cell. To prevent the cell from possible cooling, the aquarium heaters were being set to 25°C. As the ice melted within the bath, the water was syphoned out. The level of the middle salt compartment also needed to be monitored, as the level of the solution had to sit above the membrane window. This was done by adding deionised water to top up the compartment at each sampling interval.

When the cell had run for 4 hours the final samples and measurements were taken and recorded before the power supply is turned off. Once the power supply had been turned off, along with all of the pumps, water and heaters, the apparatus, apart from the aerators, was removed from the cell and washed. The remaining solution was then taken out by reversing the pumps on the aerators and bottled. The cell is then thoroughly rinsed and filled with deionised water to maintain the membranes.

36 3.4.2. pH Alteration

A 2 M solution of LiOH-H2O was made up by dissolving 8.4 g of LiOH-H2O into

100 mL of deionised water in a volumetric flask. This solution was then used to increase the pH of the Li2SO4 salt solution while it was in the cell and being mixed by the N2 gas sparging and monitored by a pH meter. Once the pH reached 11 for experiment 2 and pH 7 for experiment 3, the experiments were run in the same way as experiment 1 under the conditions stated in Table 2.

3.4.3. Temperature

For experiment 4, before the power supply output was set, the aquarium heaters were set to 40°C and the solution was left to heat to this temperature. Once the outer compartments of the cell had reached this temperature, the power supply was set and the experiment was run under the conditions stated in Table 2.

Experiment 5 required additional heating in order to get the anolyte and catholyte to 60°C. The solutions were first placed in beakers on hotplates and heated to 60°C as it would take the aquarium heaters too long to reach this temperature at that volume.

Once the solutions had reached the desired temperature they were placed in the cell in their corresponding compartments and the experiment was run under the conditions stated in Table 2.

At these temperatures the lithium sulphate would precipitate into the end of the N2 gas sparging tube. A pair of tweezers and deionised water was used to flush out the blockage and keep the salt within the salt compartment of the cell.

3.4.4. Initial Concentration of Li2SO4

Solutions of 15 wt.%, 10 wt.% and 5 wt.% Li2SO4 were made up by dissolving 75 g of Li2SO4 salt into 500 mL of deionised water, 50 g of Li2SO4 salt into 500 mL of deionised water and 25 g of Li2SO4 salt into 500 mL of deionised water respectively

37 for experiments 6,7 and 8. These experiments were then run under the conditions stated in Table 2. Experiment 8 was unable to run at a current of 3 amperes, and the current needed to be reduced as the concentration of the Li2SO4 salt reduced within the middle compartment.

3.4.5. Residence Time

Experiments 9, 10 and 11 were run in the same way as the initial experiment, as stated in Table 2 however the sampling times were different. In experiment 9, samples and measurements were taken and recorded every 20 minutes in order to provide enough data points for a 2-hour test. Experiment 10 had samples and measurements taken every half an hour as per the previous tests, and experiment 11 samples and measurements were taken every hour in order to avoid too many samples.

38 4. Chapter 4- Results and Discussion

4.1. Preliminary Experiment

The initial experiment was conducted in order to identify whether lithium hydroxide could be produced from lithium sulphate by bath electrodialysis. Ying et al., (2008) previously found that using a continuous method of electrodialysis, lithium hydroxide and sulphuric acid can be produced from lithium sulphate salt.

40 35 30 25 20 Catholyte 15 Salt 10 5 0

Lithum iConmcentration LithumiConmcentration (g/L) 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 18 Concentration of lithium within the salt and catholyte chambers,

experiment 1: preliminary

Figure 18 illustrates that the concentration of lithium within the salt solution gradually decreases over time while the lithium concentration within the catholyte chamber is increasing. This indicates that lithium hydroxide is being produced within the catholyte chamber, confirming that electrodialysis is a potential method for producing lithium hydroxide. The lithium present in the salt decreased from 36.84 g/L to 30.11 g/L and the catholyte lithium concentration increased from 7.68 g/L to 10.01 g/L.

39 35

15 Anolyte Salt 10

5 Lithum iConmcentration LithumiConmcentration (g/L) 0 0 1 2 3 4 Time (hours)

Figure 19: Concentration of sulphate within the anolyte and salt chambers,

experiment 1: preliminary

Figure 19 shows that there is a steady increase of sulphate within the anolyte chamber. As for the salt chamber, it appears that there were some disturbances during the sample assay, but the overall trend depicts that the sulphate within the salt chamber has decreased while the anolyte chamber has increased. The anolyte consists of sulphuric acid, a by-product of this particular system that could later be concentrated and sold in practical application of this system. The sulphate within the anolyte increased from 2.55 g/L to 4.32 g/L while the salt decreased from 28.45 g/L to 22.35 g/L.

40 18 16 14 12 10 Cell Voltage 8 Cathode Potential

Potential Potential (V) 6 Anode Potential 4 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 20: Cell voltage, anodic and cathodic potentials, experiment 1: preliminary

Figure 20 shows that over time, the anodic and cathodic potentials remained similar for the duration of the experiment. The cell voltage decreased throughout the course of the experiment, while throughout the experiment the cell ran at a consistent current this would imply that the resistance within the cell is decreasing in keeping with Ohm’s law. This is due to the increase in ionic conductivity as a result of the increased concentrations of lithium and sulphate ions in both the anolyte and catholyte chambers.

The current density of this system can be calculated using the surface area of the anode. The anode having dimensions of 64 mm x88 mm x1 mm has a surface area without the diamond voids of 11,504 mm2. The diamonds are 9 across and 9 down on the anode and give a total space of 166 mm2. The total of the diamond shaped voids was then calculated to be 2,988 mm2. The rest of the empty area of the anode that did not have a total diamond shape to it (refer to Figure 16 is 28 mm2 therefore the total

41 void within the mesh of the anode is 3,016 mm2. The total active anode area is then calculated by:

∑ 퐴 = ∑ 퐴푎푛표표푑푒 푠푢푟푓푎푐푒 − 2 ∑ 퐴푣표𝑖푑 𝑖푛 푚푒푠ℎ + ∑ 퐴𝑖푛푠𝑖푑푒 푑𝑖푎푚표푛푑푠 − ∑ 퐴 푏푎푠푒

= 11,504 – (3,016 × 2) + (166 × 26 × 1) – (64 × 0.5 × 1)

=8,096 mm2

Using this, a current of 10 A would result in a current density of 125 mA/cm2. And applied current of 3 A gives 37.1 mA/cm2, which converts to 371 A/m2.

The amount of salt transferred in terms of lithium during this experiment was 0.36 mole this was calculated using equation (9) (Ying et al., 2008).

푛퐿𝑖(푇푟푎푛푠푓푒푟푟푒푑) = (퐶퐿𝑖,푡 × 푉퐿𝑖,푡) − (퐶퐿𝑖,0 × 푉퐿𝑖,0) (9)

Where CLi,t is the concentration of lithium in the final catholyte sample, CLi,0 is the concentration of lithium in the initial catholyte sample, VLi,t is the final volume of the catholyte and VLi,0 is the initial volume of the catholyte.

Using the moles of lithium transferred and equation (10), the current efficiency (ŋ) can be calculated:

푛 퐹 퐶푢푟푟푒푛푡 퐸푓푓푖푐푖푒푛푐푦 (ŋ) = 퐿푖 × 100 (10) 퐼푡

Here F is Faraday’s constant, I is the current applied to the cell and t is the duration of the experiment in seconds. For experiment 1, the current efficiency in relation to lithium transfer from the salt to the catholyte was 80%.

Table 14, which can be found in the appendix, shows the mass balance for experiment 1. The concentration of lithium was determined by Inductively Coupled

Plasma Mass Spectrometry (ICP-MS), while the concentration of sulphate was determined by Ion Chromatography (IC). 98.74% of the lithium was accounted for within the experimental data, the remaining 1.26% could be due to lithium hydroxide and lithium sulphate precipitating onto the equipment in small quantities or

42 evaporation of lithium sulphate and lithium hydroxide solution from the catholyte and the salt compartments of the cell. The salt compartment tended to run slightly hotter than the other two compartments within the cell. This is due to its smaller volume of solution held within the compartment.

Table 3: Mass transfer and recovery, experiment 1: preliminary Lithium Sulphate g % g % Mass transferred from salt 3.14 19.2 2.77 22.8 Li Recovery in catholyte 2.48 17.3 - - Recovery in anolyte - - 1.67 14.9 g/h g/h Average Transfer Rate 0.66 0.53

Table 3 depicts the transfer and recovery of both lithium and sulphate within the system. 17.3% of the total lithium within the original salt solution was recovered into the catholyte. 14.9% of the sulphate was recovered in the anolyte from the original salt solution. The transfer rates in Table 3 refer to the rate of transfer of the lithium and sulphate through the cationic and anionic membranes into the catholyte and anolyte respectively.

4.2. Effect of pH

Experiments 2 and 3 were run in order to observe the impact that the pH of the lithium sulphate solution would have on the recovery of lithium in the form of lithium hydroxide from the catholyte compartment. Experiments 2 was run at a pH of 11 and experiment 3 was run at a pH of 7 under the conditions as stated in Table 2.

43 35 30 25 20 15 Catholyte 10 Salt 5 0 Lithum Conmcentration LithumConmcentration (g/L) 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 21 Concentration of lithium within the catholyte and salt chambers,

experiment 2: pH 11

40 35 30 25 20 Catholyte 15 Salt 10 5

0 Lithum Conmcentration LithumConmcentration (g/L) 0 1 2 3 4 5 Time (hours)

Figure 22: Concentration of lithium within the catholyte and salt chambers,

experiment 3: pH 7

Figure 21 and Figure 22 illustrate the increase in concentration within the catholyte chamber in experiments 2 and 3, respectively, and the decrease in the lithium concentration within the salt chambers. Experiment 2 saw the salt concentration reduced from 30.53 g/L to 18.44 g/L while the lithium concentration increased from

6.74 g/L to 8.22 g/L. Experiment 3 shows that the concentration within the salt

44 chamber went from 33.40 g/L to 27.85 g/L and the lithium concentration in the catholyte chamber increased from 7.75 g/L to 9.74 g/L. This data suggests that a change in pH does not overwhelmingly influence the increase in lithium concentration within the catholyte.

15 Anolyte 10 Salt

Sulphate Conmcentration SulphateConmcentration (g/L) 0 0 1 2 3 4 5 Time (hours)

Figure 23: Concentration of sulphate within the anolyte and salt compartments, experiment 2: pH 11

30 25 20 15 Anolyte 10 Salt 5 0 Lithum Concentration LithumConcentration (g/L) 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 24: Concentration of sulphate within the anolyte and salt compartments, experiment 3: pH 7 Figure 23 and Figure 24 illustrate the increase in concentration within the anolyte chamber in experiments 2 and 3, respectively, and the decrease in the sulphate concentration within the salt chambers. Experiment 2 saw that the salt concentration

45 decreased from 26.75 g/L to 19.49 g/L while the sulphate concentration increased from 2.91 g/L to 5.05 g/L within the anolyte chamber. Experiment 3 shows that the concentration within the salt chamber changed from 23.58 g/L to 19.87 g/L and the sulphate concentration in the anolyte chamber increased from 4.04 g/L to 4.94 g/L.

18 16 14 12 10 Cell Voltage 8 Cathodic potential Potentil Potentil (V) 6 Anode potential 4 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 25: Cell voltage, anodic and cathodic potentials, experiment 2: pH 11

18 16 14 12 10 Cell Voltage 8 Cathodic potential Potentil Potentil (V) 6 Anode potential 4 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 26: Cell voltage, anodic and cathodic potentials, experiment 3: pH 7

46 The cell voltage and the electrode potentials differ very slightly between the experiments run at different pH. This is shown in Figure 25 and Figure 26. There was a greater voltage drop across the cell throughout the duration of the test run at pH 7.

Again this data suggests that there is a drop in resistance over time within the cell in compliance with Ohms law. However, there was a greater drop in resistance when running at a pH of 7 than there was when running at a pH of 11.

The number of moles of lithium transferred in experiment 2, 0.22 moles is less than the number of moles transferred in experiment 3, 0.29 moles. Both of these tests also had a lesser transfer of moles of lithium across the membrane than experiment 1, which had 0.36 moles transfer across the membrane from the salt chamber into the catholyte. These numbers also correspond to the current efficiency, the lower the number of moles that transfer, the lower the current efficiency. Experiment 2 had a current efficiency of 49% and experiment 3 had a current efficiency of 64%, both lower than the 80% current efficiency that was found in experiment 1.

Table 4: Mass transfer and recovery, experiment 2: pH 11

Lithium Sulphate g % g % Mass transferred from salt 5.07 39.3 3.15 27.1 Li Recovery in catholyte 1.53 13.0 - - Recovery in anolyte - - 1.98 19.1 g/h g/h Average Transfer Rate 0.77 0.60

Table 5: Mass transfer and recovery, experiment 3: pH 7

Lithium Sulphate g % g % Mass transferred from salt 2.70 17.9 1.82 17.2 Recovery in catholyte 2 15.4 - - Recovery in anolyte - - 0.75 8.1 g/h g/h Average Transfer Rate 0.54 0.29

The mass balance and accountability of lithium and sulphate for experiments 2 and

3 can be found in Table 15 and Table 16 in the appendix. Table 4 shows that 13.0% if the lithium within the salt solution at pH 11 was recovered within the catholyte and

Table 5 shows that 15.4% of the lithium was recovered at a pH of 7, comparing these values to experiment 1, which was run at the salt solutions unaltered pH of 2, there was a 17.3% recovery of lithium. This suggests that, together with the current efficiency, increasing the pH of Li2SO4 salt solution has a detrimental effect on the lithium recovery from the salt compartment into the catholyte.

Sulphate recovery was of 19.1% in experiment 2 and 8.1% in experiment 3, while experiment 1 had a recovery of 14.9%. The recovery of sulphate was highest in experiment 2, this would suggest that increasing the initial pH of the starting salt solution would be beneficial if sulphuric acid was the main product of this process.

Experiment 2 had a higher rate of transfer for both lithium and sulphate compared to experiment 1 and experiment 3, this means that lithium would be produced faster if the pH were to be elevated but at a cost to recovery and current efficiency.

4.3. Effect of Temperature

Experiments 4 and 5 were run to determine what would happen if the temperatures within the cell were increased and how it would influence the recovery of lithium.

48 Experiment 4 was run at a temperature range of 35-45°C while experiment 5 was run at a temperature range of 50-60°C. Table 2 outlines the rest of the conditions that these experiments were conducted under. The cell was closely monitored in experiment 5, as the temperature threshold on the membranes is 60°C, and possible damage could ensue if they were to be run over 60°C.

35 30 25 20 15 Catholyte 10 Salt 5

Lithum Concentration LithumConcentration (g/L) 0 0 1 2 3 4 5 Time (hours)

Figure 27: Concentration of lithium within the catholyte and salt chambers,

experiment 4: 40°C

40 35 30 25 20 Catholyte 15 Salt 10 5 Lithum Concentration LithumConcentration (g/L) 0 0 1 2 3 4 5 Time (hours)

Figure 28: Concentration of lithium within the catholyte and salt chambers,

experiment 5: 60°C

49 Figure 27 and Figure 28 show the change in concentration within the salt and catholyte chambers in experiments 4 and 5 respectively. Experiment 4 saw the lithium concentration decrease in the salt cell from 32.75 g/L to 23.48 g/L and experiment 5 saw the salt chamber change from 33.92 g/L to 24.99 g/L. The catholyte chamber in experiment 4 increased in the lithium concentration from 7.18 g/L to 8.23 g/L and experiment 5 had the lithium concentration increase from 8.54 g/L to 11.00 g/L in the catholyte chamber. This suggests that elevating the temperature from 40°C to 60°C increased the concentration slightly in the catholyte.

10 Anolyte Salt 5

Sulphate Concentration SulphateConcentration (g/L) 0 1 2 3 4 5 Time (hours)

Figure 29: Concentration of sulphate within the anolyte and salt compartments,

experiment 4: 40°C

50 25

10 Anolyte Salt 5

Sulphate Concentration SulphateConcentration (g/L) 0 1 2 3 4 5 Time (hours)

Figure 30: Concentration of sulphate within the anolyte and salt compartments,

experiment 5: 60°C

Figure 29 and Figure 30 illustrate the increase in concentration within the anolyte chamber in experiments 4 and 5 respectively and the decrease in the sulphate concentration within the salt chambers. Experiment 4 saw the salt concentration reduce from 22.59 g/L to 19.42 g/L while the sulphate concentration increased from

23.78 g/L to 5.06 g/L. Experiment 3 shows that the concentration within the salt chamber went from 22.63 g/L to 18.27 g/L and the sulphate concentration in the catholyte chamber increased from 3.83 g/L to 4.89 g/L. In comparison to the graphs, there is not much effect in the increase in concentration of sulphate within the anolyte.

51 16

8 Cell Voltage

6 Cathode Potential Potential Potential (V) Anode Potential 4

0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 31: Cell voltage, anodic and cathodic potentials, experiment 4: 40°C

16 14

12 ) 10 8 Cell Voltage

6 Cathode Potential Potential Potential (V 4 Anode Potential 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 32: Cell voltage, anodic and cathodic potentials, experiment 5: 60°C

Figure 31 and Figure 32 show that the anodic and cathodic potentials for experiments 4 and 5 are relatively similar. However, there is a greater voltage drop in experiment 5 at the higher temperature of 60°C, suggesting that there is less resistance within the cell at a higher temperature. Ionic conductivity between the chambers would naturally increase with elevated temperature.

52 The number of moles transferred from the salt to the catholyte was greater in experiment 5 than in experiment 4. Experiment 5 has 0.21 moles of lithium transfer across the membrane into the catholyte while experiment 4 had 0.15 moles transfer.

When compared to experiment 1 though, with a transfer of 0.36 moles, it seems to suggest that increasing the temperature of the system is detrimental. Experiment 4 had a current efficiency of 34% while experiment 5 had an increased current efficiency of

48%, however these current efficiencies are not great compared to experiment 1 where the current efficiency was 80%. This would suggest that the increase in temperature is uneconomical as it hinders the transfer of lithium ions yet also decreases the current efficiency.

Table 17 and Table 18 in the appendix, show the mass balances for experiments 4 and 5 respectively. The accountability for lithium in both of these experiments is lower compared to that of the other experiments. This is due to the lithium sulphate beginning to precipitate at 40°C, some of the solid salt was then lost when flushed out of the N2 gas sparging tubing, where the bulk of the precipitation occurred, as it was preventing the gas from cooling and mixing within the camber and had to be removed. Lithium sulphate precipitating in the salt chamber is an issue, as the solid cannot pass through the membrane, therefore making the process inefficient.

Table 6: Mass transfer and recovery, experiment 4: 40°C,

Lithium Sulphate g % g % Mass transferred from salt 4.38 31.7 1.82 18.1 Recovery in catholyte 1.05 8.2 - - Recovery in anolyte - - 1.05 12.0 g/h g/h Average Transfer Rate 0.64 0.33

Table 7: Mass transfer and recovery, experiment 5: 60°C,

Lithium Sulphate g % g % Mass transferred from salt 4.41 30.7 2.36 24.2 Recovery in catholyte 1.48 11.2 - - Recovery in anolyte - - 0.901 10.2 g/h g/h Average Transfer Rate 0.69 0.38

Table 6 and Table 7 show that when the cell is run at 40°C 8.2% of the lithium within the salt is recovered into the catholyte, while at 60°C this recovery is increased to 11.2% suggesting that an increase in temperature from 40°C to 60°C is beneficial when trying to increase recovery of the lithium material. These two recoveries are both still lower, however, than that of experiment 1 at 14.9% which seems to suggest a detrimental effect from operating at elevated cell temperature. Table 6 and Table 7 also show that when the temperature is increased from 40°C to 60°C, that the rate of transfer is increased, however the rate of transfer is only lightly faster in experiment 5 than it is in experiment 1.

4.4. Effect of Starting Concentration

The impact that the initial concentration of the salt solution has on the recovery of lithium is tested in experiments 6, 7 and 8. Experiment 6 had the salt concentration

54 reduce from 30 wt.% to 15 wt.%, in experiment 7 the concentration was further reduced to 10 wt.%, and in experiment 8 the concentration was again reduced to 5 wt.% lithium sulphate. Table 2 outlines the rest of the operating conditions that experiments 6, 7 and 8 were conducted under.

20 18 16 14 12 10 Catholyte 8 6 Salt 4

Lithum Concentration Concentration Lithum (g/L) 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 33: Concentration of lithium within the catholyte and salt chambers,

experiment 6: 15 wt.%

55 14

6 Catholyte Salt 4

2 Lithum Concentration Concentration Lithum (g/L) 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 34: Concentration of lithium within the catholyte and salt chambers,

experiment 7: 10 wt.%

10 9 8 7 6 5 Catholyte 4 3 Salt 2

Lithum Concentration LithumConcentration (g/L) 1 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 35: Concentration of lithium within the catholyte and salt chambers,

experiment 8: 5 wt.%

Figure 33, Figure 34 and Figure 35 show the concentration of lithium within the catholyte and the salt compartments in the cell over time. Experiment 6 shows that the lithium centration within the salt decreases from 17.07 g/L to 10.78 g/L while the

56 catholyte increases from 7.60 g/L to 9.29 g/L. Experiment 7 shows that the lithium concentration within the salt decreases from 11.90 g/L to 6.28 g/L and the catholyte increased from 7.32 g/L to 8.67 g/L. Finally, in experiment 8, the lithium concentration decreases in the salt from 5.92 g/L to 1.08 g/L and increases from 7.21 g/L to 9.01 g/L in the catholyte. From these graphs, it is clear that the test run at 5 wt.% has the largest increase in concentration within the catholyte chamber.

6 Anolyte 4 Salt 2

Sulphate Concentration SulphateConcentration (g/L) 0 1 2 3 4 5 Time (hours)

Figure 36: Concentration of sulphate within the anolyte and salt compartments,

experiment 6: 15 wt.%

57 8

4 Anolyte 3 Salt 2

1 Sulphate Concentration SulphateConcentration (g/L) 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 37: Concentration of sulphate within the anolyte and salt compartments,

experiment 7: 10 wt.%

5 4.5 4 3.5 3 2.5 Anolyte 2 Salt 1.5 1

Suphate concentration Suphate concentration (g/L) 0.5 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 38: Concentration of sulphate within the anolyte and salt compartments,

experiment 8: 5 wt.%

Figure 36 shows the concentrations of the anolyte and salt chambers for sulphate in experiment 6. The salt chamber in experiment 6 decreases from 10.57 g/L to 7.43 g/L while the anolyte chamber increases from 3.91 g/L to 5.29 g/L. Figure 37 shows the concentrations of the anolyte and the salt chambers for experiment 7, where the

58 original concentration of the salt was 10 wt.% Li2SO4. The figure shows that the initial salt concentration was 6.92 g/L that then decreased to 3.82 g/L. The anolyte chamber initially started at 3.54 g/L but increased to 4.56 g/L. Figure 38 illustrates the increasing concentration of the anolyte chamber, 3.69 g/L to 4.55 g/L, and the decreasing concentration of the salt chamber, 3.08 g/L to 1.07 g/L in experiment 8.

From these graphs, the increase in sulphate concentration within the anolyte is relatively similar in each experiment. This suggests that decreasing the initial concentration of the lithium sulphate salt solution could have a positive effect on the increase in concentration of lithium while there is not an obvious negative effect on the sulphate concentration.

18 16 14 12 10 Cell Voltage 8 Cathode Potential

Potential Potential (V) 6 Anode Potential 4 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 39: Cell voltage, anodic and cathodic potentials, experiment 6: 15 wt.%

18 16 14 12 10 Cell Voltage 8 Cathode potential

Potential Potential (V) 6 Anode Potential 4 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 40: Cell voltage, anodic and cathodic potentials, experiment 7: 10 wt.%

15 Cell Voltage

10 Cathode Potential Potential Potential (V) Anode Potenial 5

0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 41: Cell voltage, anodic and cathodic potentials, experiment 8: 5wt.%

Figure 39, Figure 40 and Figure 41 illustrate the cell voltages and the potentials of experiments 6-8. Experiments 7 and 8 initially decreased in cell voltage, but later on began to increase, more so for experiment 8. This suggests an increase in resistance, or a restriction of ion manoeuvrability therefore creating resistance within the cell. As the ions begin to decrease in the salt compartment, there is less conductivity within that compartment. In experiment 8, the applied current had to be repeatedly lowered

60 as the concentration decreased within the salt cell to prevent the voltage passing through the cell from increasing uncontrollably.

The molar transfer of lithium for experiments 6, 7 and 8 were 0.24 moles, 0.2 moles and 0.26 moles respectively. The lower concentration allowed for more lithium molar transport across the membrane from the salt compartment to the catholyte.

These values correspond to a current efficiency of 54% for experiment 6, 45% for experiment 7 and 75% for experiment 8.

Table 19, Table 20 and Table 21 show the mass balance and accountability for lithium and sulphate from experiments 6, 7 and 8 respectively. The accountability for each of these experiments is relatively high. The small loss of lithium and sulphate could be due to precipitation.

Table 8: Mass transfer and recovery, experiment 6: 15 wt.% Lithium Sulphate g % g % Mass transferred from salt 2.70 37.7 1.38 30.7 Recovery in catholyte 1.18 17.8 - - Recovery in anolyte - - 1.23 30.0 g/h g/h Average Transfer Rate 0.46 0.31

Table 9: Mass transfer and recovery, experiment 7: 10 wt.% Lithium Sulphate g % g % Mass transferred from salt 2.34 47.7 1.29 45.3 Recovery in catholyte 1.40 30.4 - - Recovery in anolyte - - 0.59 22.1 g/h g/h Average Transfer Rate 0.45 0.23

61 Table 10: Mass transfer and recovery, experiment 8: 5 wt.%

Lithium Sulphate g % g % Mass transferred from salt 1.95 81.6 0.74 58.2 Recovery in catholyte 1.83 80.8 - - Recovery in anolyte - - 0.58 49.1 g/h g/h Average Transfer Rate 0.46 0.16

Table 8, Table 9 and Table 10 show that the recovery of lithium from the salt solution to the catholyte drastically increased from 17.8% in experiment 6 to 30.4% in experiment 7 and finally to 80.8% in experiment 8, as the concentration of the initial salt solution is decreased. However, the mass transfer rates are relatively the same within these three experiments, they are however, slower than that of experiment 1.

These concentrations are a more accurate representation of what might be used in industry as initial concentrations.

4.5. Effect of Residence Time

Experiments 9, 10 and 11 were conducted in order to determine what the impact residence time within the cell has on the recovery of lithium. Experiment 9 was run for 2 hours, experiment 10 was run for 4 hours and experiment 11 was run for 8 hours, the conditions at which these tests were run can be found in Table 2.

62 40 35 30 25 20 Catholyte 15 Salt 10 5

0 Lithum Concentration LithumConcentration (g/L) 0 0.5 1 1.5 2 2.5 Time (hours)

Figure 42: Concentration of lithium within the catholyte and salt chambers,

experiment 9: 2 hours

40 35 30 25 20 Catholyte 15 Salt 10

5 Lithum Concentration LithumConcentration (g/L) 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 43: Concentration of lithium within the catholyte and salt chambers,

experiment 10: 4 hours

63 40 35 30 25 20 Catholyte 15 Salt 10

5 Lithum Concentration LithumConcentration (g/L) 0 0 1 2 3 4 5 6 7 8 Time (hours)

Figure 44: Concentration of lithium within the catholyte and salt chambers,

experiment 11: 8 hours

Figure 42, Figure 43 and Figure 44 depict the concentrations of lithium within the salt and the catholyte in experiments 9, 10 and 11 respectively. In experiment 9, the lithium concentration decreased in the salt chamber from 33.25 g/L to 29.68 g/L and the catholyte increased from 6.89 g/L to 8.29 g/L. Experiment 10 ran for 2 hours longer than experiment 9 and the salt went from 33.25 g/L to 2.64 g/L and the catholyte went from 6.89 g/L to 9.12 g/L. Finally, experiment 11 ran for a total of 8 hours and the concentration within the salt compartment went from 33.25 g/L to 19.03 g/L while the catholyte increased from 6.89 g/L to 10.93 g/L. From these graphs it is evident that there is an increase in lithium concentration within the catholyte.

64 25

Anolyte 10 Salt

5 Sulphateconcentration

0 0 0.5 1 1.5 2 2.5 Time (hours)

Figure 45: Concentration of sulphate within the anolyte and salt compartments, experiment 9: 2 hours

Anolyte 10 Salt

5 Sulphate Concentration SulphateConcentration (g/L) 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time(hours)

Figure 46: Concentration of sulphate within the anolyte and salt compartments, experiment 10: 4 hours

65 25

10 Anolyte Salt 5

Sulphate Concntration SulphateConcntration (g/L) 0 0 1 2 3 4 5 6 7 8 Time (hours)

Figure 47: Concentration of sulphate within the anolyte and salt compartments, experiment 11: 8 hours

Figure 45, Figure 46 and Figure 47 show the concentration of sulphate in experiments 9, 10 and 11 respectively. In experiment 9, the salt concentration within the salt compartment decreased from 22.78 g/L to 20.61 g/L and the sulphate within the anolyte increased from 3.72 g/L to 4.17 g/L. Experiment 10, which is depicted in

Figure 46, shows that the concentration of sulphate further decreases from 22.78 g/L to 19.01 g/L and that the sulphate within the anolyte further increases from 3.72 g/L to 4.83 g/L. Finally, Figure 47, which shows the sulphate concentration for experiment 11 that was run for 8 hours. The salt decreased more in this experiment from 22.78 g/L to 17.86 g/L, and the anolyte increased further from 3.72 to 5.65 g/L.

These graphs clearly show that with an increase in residence time within the cell, there is a notable increase in the concentration of sulphate.

66 20 18 16 14 12 10 Cell Voltage

8 Cathode Potential Potential Potential (V) 6 Anode Potential 4 2 0 0 0.5 1 1.5 2 2.5 Time (hours)

Figure 48: Cell voltage, anodic and cathodic potentials, experiment 9: 2 hours

20 18 16 14 12 10 Cell Voltage

8 Cathode Potential Potential Potential (V) 6 Anode Potential 4 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 Time (hours)

Figure 49: Cell voltage, anodic and cathodic potentials, experiment 10: 4 hours

67 20 18 16 14 12 10 Cell Voltage

8 Cathode Potential Potential Potential (V) 6 Anode Potential 4 2 0 0 0.5 1 1.5 2 2.5 3 3.5 4 4.5 5 5.5 6 6.5 7 7.5 8 8.5 Time (hours)

Figure 50: Cell voltage, anodic and cathodic potentials, experiment 11: 8 hours Figure 48, Figure 49 and Figure 50 show the cell voltage and potentials for experiments 9, 10 and 11 respectively. The cell voltage initially decreases in experiment 9 and then gradually increases over the 2 hours, while the voltage overall decreases in experiments 10 and 11.

With an increase in residence time, there is an increase in molar transfer.

Experiment 9 had 0.20 moles transfer with a current efficiency of 90%, experiment 10 had 0.32 moles transfer at a current efficiency of 72% and experiment 11 had 0.58 moles transfer at a current efficiency of 65%. It can be deduced, that with an increase in time, there is a decrease in current efficiency.

68 Table 11: Mass transfer and recovery, experiment 9

Lithium Sulphate g % g % Mass transferred from salt 1.82 13.2 1.14 11.7 Recovery in catholyte 1.40 10.6 - - Recovery in anolyte - - 0.33 3.7 g/h g/h Average Transfer Rate 0.79 0.17

Table 12: Mass transfer and recovery, experiment 10

Lithium Sulphate g % g % Mass transferred from salt 4.26 29.8 1.76 18.1 Recovery in catholyte 2.24 17.4 - - Recovery in anolyte - - 0.95 10.7 g/h g/h Average Transfer Rate 0.76 0.32

Table 13: Mass transfer and recovery, experiment 11

Lithium Sulphate g % g % Mass transferred from salt 5.93 42.2 2.20 20.9 Recovery in catholyte 4.05 31.7 - - Recovery in anolyte - - 1.76 20.1 g/h g/h Average Transfer Rate 0.59 0.45

Table 11, Table 12 and Table 13 show the mass transfer of lithium and sulphate with their respective recoveries for experiments 9,10 and 11. These tables indicate that, with an increase in time, there is an increase in recovery. Experiment 9 had a lithium recovery of 10.6%, experiment 10 had a recovery of 17.4% for lithium and experiment 11 had a lithium recovery of 31.7%. But when comparing the lithium mass transfer rates, over time the mass transfer decreases significantly.

69 In terms of the sulphate, the recovery roughly doubles each time with a doubling in residence time within the cell. In experiment 9, the recovery for sulphate was 3.7%; this is after 2 hours within the cell. In experiment 10, the residence time doubles from

2 hours to 4 hours and the recovery for sulphate more than doubles to 10.7%. Finally, in experiment 11, the residence time within the cell doubles again and the recovery for sulphate is 20.1%, roughly doubled again.

4.6. General Discussion

Figure 51 shows the recovery of lithium with the corresponding current efficiency from each experiment. From this graph is can be deduced that experiment

8, where the initial starting concentration of the Li2SO4 salt was lowered to 5 wt.%, had the highest recovery of lithium within the catholyte chamber. This particular experiment also had a significant current efficiency in the transfer of lithium from the salt compartment to the anolyte.

The highest current efficiency was found in experiment 9, where the cell was only run for 2 hours. This indicates that the lithium transfer from the salt chamber to the catholyte chamber was fast during the run time of the cell, however, there is a poor lithium recovery at this retention time.

Experiment 11 was run for 8 hours, and Figure 51 shows that this test had the highest recovery for lithium out of all of the tests that began at a concentration of 30 wt.% Li2SO4. When compared to experiment 1 and 10, the recovery was more than doubles with only a minor loss in current efficiency within the cell.

70 100 90 80 70 60

% 50 Current efficiency 40 Lithium Recovery 30 20 10 0 1 2 3 4 5 6 7 8 9 10 11 Experiment no.

Figure 51: Current efficiency and the lithium recovery of experiments 1-11.

2.5

1.5

g/h LiOH production rate 1 H2SO4 production rate

0.5

0 1 2 3 4 5 6 7 8 9 10 11 Experiment no.

Figure 52: Lithium hydroxide and sulphuric acid production rates from each test

Figure 52 shows the production rates of each experiment for both lithium hydroxide and sulphuric acid. It can be seen, that experiment 9, where the cell was run for 2 hours under conditions outlined in Table 2, had the fastest rate of production for lithium hydroxide at 2.29 g/h of lithium hydroxide being produced. Experiment 1, conducted under conditions outlined in Table 2, also yielded a relatively high rate of production and from Figure 51 had a higher recovery of lithium than experiment 9, however, the current efficiency for experiment 9 was greater than that of experiment

Experiments 10 and 11 also produced relatively high rates of lithium hydroxide production, these experiments were run for 4 and 8 hours respectively under the conditions outlined in Table 2.

72 5. Chapter 5- Conclusion and Recommendations

This study aimed to determine whether it was possible to produce lithium hydroxide from lithium sulphate salt through the means of electrodialysis using a 2 membrane, 3 compartment cell. It was found, that in reduced concentrations of the initial lithium sulphate salt, that the recovery of lithium as lithium hydroxide within the catholyte compartment of the cell was maximised. It was also found that the rate of production of lithium hydroxide was at its fastest when the cell was run for a shorter time period. However, in running the cell for longer, the recovery of lithium was increased.

The residence time within the cell was found to have a significant impact on the recovery of lithium as lithium hydroxide within the catholyte compartment of the cell.

The longer the cell ran for, the recovery of lithium increased, however the current efficiency deteriorated, and the transfer of lithium ions through the cationic membrane slowed over time. However, it was observed that a doubling of residence time leads to a doubling in the recovery of lithium as lithium hydroxide, with less than a 10% decrease to the current efficiency.

It is recommended, that changing the pH of the initial pH solution of the lithium sulphate should not be done, as it has detrimental effects on both the recovery of lithium as lithium hydroxide and the current efficiency of the system. Similarly, this process is recommended to be carried out at room temperature in order to achieve maximum results. At elevated temperatures, technical issue with the lithium sulphate precipitating in the salt chamber results in less lithium ions available to migrate through the cationic membrane to produce lithium hydroxide. Likewise, there is less sulphate ions available to pass through the anionic membrane to produce sulphuric

73 acid. Running the cell at elevated temperatures also significantly reduce the current efficiency of the system.

5.1. Future Work

It is recommended that different membrane pairs be experimented with in order to observe the effects of their permiselectivity and the resulting purity of the lithium hydroxide produced. Different membrane pairs may also influence the recovery of the lithium within the catholyte. It may be necessary to install more than one membrane pair to obtain greater purity of the final product.

The addition of impurities into the lithium sulphate salt should also be experimented with in order to observe their effects on the production of lithium hydroxide in situation when the salt is not pure. The migration of impurities to the catholyte should, ideally, be minimised in the production of lithium hydroxide for commercial use.

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79 7. Appendix

Equations for determining the concentration of sulphate from the acid-base titration:

Moles of NaOH used:

푁푁푎푂퐻 = 푇푖푡푟푎푛푡 푣표푙푢푚푒 (퐿) × 푀표푙푎푟푖푡푦 표푓 푇푖푡푟푎푛푡 (4)

2− 퐻2푆푂4 + 2푁푎푂퐻 → 2퐻2푂 + 푆푂4 + 2푁푎 (5)

∴ 1 푚표푙 표푓 퐻2푆푂4 푡표 2 푚표푙 표푓 푁푎푂퐻

Moles of H2SO4:

푁 푁 = 푁푎푂퐻 = 푁 (6) 퐻2푆푂4 2 푆푂4

Equation 3 depicts that the number of moles of sulphate within the acid is a stoichiometric ratio of 1:1 so the number of moles of acid is equal to the number of moles of sulphate within the 1 mL sample.

푁 퐶 (푚표푙⁄퐿) = 푆푂4 (7) 푆푂4 0.001

퐶푆푂4(푔⁄퐿) = 퐶푆푂4(푚표푙⁄퐿) × 푀푆푂4 (8)

80 6.5 cm 5 cm

17.3 cm

4 cm

10.4 cm

Figure 53 Middle compartment of electrodialysis cell

81 Figure 54: Dimensions of Electrodialysis Cell

Figure 55 Apparatus placement within the electrodialysis cell

83 Table 14: System mass balance, experiment 1: preliminary.

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 36.84 400 14.74 Catholyte 7.68 1000 7.68 Total 22.41 Final Salt 30.12 385 11.59 Catholyte 10.01 1015 10.16 Total 21.75 Mass lost due to sampling 0.38 Mass unaccounted for 0.28 Deviation (%) 1.26 Accountability (%) 98.74 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 28.45 400 11.38 Anolyte 2.55 1000 2.55 Total 13.93 Final Salt 22.35 385 8.61 Anolyte 4.32 978 4.22 Total 12.83 Mass lost due to sampling 0.23 Mass unaccounted for 0.879 Deviation (%) 6.41 Accountability (%) 93.58

Table 15: System mass balance, experiment 2: pH 11

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 30.53 400 12.21 Catholyte 6.74 1000 6.74 Total 18.95 Final Salt 18.44 387 7.14 Catholyte 8.22 1007 8.27 Total 15.41 Mass lost due to sampling 0.45 Mass unaccounted for 3.09 Deviation (%) 16.70 Accountability (%) 83.30 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 26.75 400 10.70 Anolyte 2.914 1000 2.914 Total 13.61 Final Salt 19.49 387 7.54 Anolyte 5.05 969 4.89 Total 12.44 Mass lost due to sampling 0.35 Mass unaccounted for 0.83 Deviation (%) 6.24 Accountability (%) 93.76

Table 16: System mass balance, experiment 3: pH 7

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 33.40 400 13.36 Catholyte 7.75 1000 7.75 Total 21.11 Final Salt 27.85 383 10.67 Catholyte 9.74 1001 9.75 Total 20.41 Mass lost due to sampling 0.36 Mass unaccounted for 0.34 Deviation (%) 1.62 Accountability (%) 98.38 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 23.58 400 9.43 Anolyte 4.04 1000 4.04 Total 13.47 Final Salt 19.87 383 7.61 Anolyte 4.98 962 4.79 Total 12.41 Mass lost due to sampling 0.24 Mass unaccounted for 0.83 Deviation (%) 6.28 Accountability (%) 93.71

Table 17: System mass balance, experiment 4: 40°C

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 32.75 400 13.10 Catholyte 7.18 1000 7.18 Total 20.28 Final Salt 23.48 371.5 8.27 Catholyte 8.23 1000 8.23 Total 16.95 Mass lost due to sampling 0.33 Mass unaccounted for 3.00 Deviation (%) 15.03 Accountability (%) 84.97 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 22.59 400 9.04 Anolyte 3.78 1000 3.78 Total 12.82 Final Salt 19.42 371.5 7.21 Anolyte 5.05 952 4.83 Total 12.05 Mass lost due to sampling 0.23 Mass unaccounted for 0.53 Deviation (%) 4.21 Accountability (%) 95.78

87 Table 18: System mass balance, experiment 5: 60°C

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 33.92 400 13.57 Catholyte 8.54 1000 8.54 Total 22.10 Final Salt 24.99 366.5 9.16 Catholyte 10.23 979 10.02 Total 19.18 Mass lost due to sampling 0.35 Mass unaccounted for 2.57 Deviation (%) 11.82 Accountability (%) 88.18 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 22.63 400 9.05 Anolyte 3.83 1000 3.83 Total 12.88 Final Salt 18.27 366.5 6.69 Anolyte 4.88 968 4.73 Total 11.42 Mass lost due to sampling 0.21 Mass unaccounted for 1.24 Deviation (%) 9.79 Accountability (%) 90.21

88 Table 19: System mass balance, experiment 6: 15 wt.%

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 17.08 400 6.83 Catholyte 7.60 1000 7.60 Total 14.43 Final Salt 10.78 383 4.13 Catholyte 8.79 999 8.78 Total 12.91 Mass lost due to sampling 0.20 Mass unaccounted for 1.32 Deviation (%) 9.30 Accountability (%) 90.70 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 10.57 400 4.23 Anolyte 3.91 1000 3.91 Total 8.13 Final Salt 7.43 383 2.85 Anolyte 5.29 971 5.14 Total 7.98 Mass lost due to sampling 0.13 Mass unaccounted for 0.03 Deviation (%) 0.33 Accountability (%) 99.66

89 Table 20: System mass balance, experiment 7: 10 wt.%

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 11.90 400 4.76 Catholyte 7.32 1000 7.32 Total 12.08 Final Salt 6.28 385 4.76 Catholyte 8.67 1006.5 7.32 Total 11.14 Mass lost due to sampling 0.14 Mass unaccounted for 0.8 Deviation (%) 6.71 Accountability (%) 93.29 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 6.92 400 2.77 Anolyte 3.54 1000 3.54 Total 6.3 Final Salt 3.82 385 1.47 Anolyte 4.56 971 4.42 Total 5.89 Mass lost due to sampling 0.08 Mass unaccounted for 0.33 Deviation (%) 5.32 Accountability (%) 94.68

90 Table 21: System mass balance, experiment 8: 5 wt.%

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 5.92 400 2.37 Catholyte 7.21 1000 7.21 Total 9.57 Final Salt 1.08 384.5 0.42 Catholyte 1.08 384.5 0.42 Total 9.45 Mass lost due to sampling 0.11 Mass unaccounted for 0.02 Deviation (%) 0.19 Accountability (%) 99.81 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 3.08 400 1.23 Anolyte 3.69 1000 3.69 Total 4.922 Final Salt 1.0667 384.5 0.41 Anolyte 4.55 978 4.45 Total 4.86 Mass lost due to sampling 0.05 Mass unaccounted for 0.01 Deviation (%) 0.11 Accountability (%) 99.89

91 Table 22: System mass balance, experiment 9: 2 hours

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 33.25 400 13.30 Catholyte 6.89 1000 6.89 Total 20.19 Final Salt 29.68 387 11.49 Catholyte 8.29 1000 8.29 Total 19.77 Mass lost due to sampling 0.07 Mass unaccounted for 0.34 Deviation (%) 1.71 Accountability (%) 98.29 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 32.78 400 9.11 Anolyte 3.72 1000 3.72 Total 12.83 Final Salt 20.61 387 7.97 Anolyte 4.18 970 4.05 Total 12.02 Mass lost due to sampling 0.08 Mass unaccounted for 0.73 Deviation (%) 5.71 Accountability (%) 94.29

92 Table 23: System mass balance, experiment 10: 4 hours

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 33.25 400 13.30 Catholyte 6.89 1000 6.89 Total 20.19 Final Salt 23.36 387 9.04 Catholyte 9.13 1000 9.13 Total 18.17 Mass lost due to sampling 0.42 Mass unaccounted for 1.61 Deviation (%) 8.12 Accountability (%) 91.88 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 22.78 400 9.11 Anolyte 3.72 1000 3.72 Total 12.83 Final Salt 19.01 387 7.36 Anolyte 4.83 970 4.68 Total 12.03 Mass lost due to sampling 0.13 Mass unaccounted for 0.66 Deviation (%) 5.23 Accountability (%) 94.76

93 Table 24: System mass balance, experiment 11: 8 hours

Lithium Initial Concentration (g/L) Volume (L) Mass of Lithium (g) Salt 33.25 400 13.30 Catholyte 6.89 1000 6.89 Total 20.19 Final Salt 19.04 387 7.37 Catholyte 10.94 1000 10.94 Total 18.31 Mass lost due to sampling 0.54 Mass unaccounted for 1.34 Deviation (%) 6.81 Accountability (%) 93.19 Sulphate Initial Concentration (g/L) Volume (mL) Mass of Sulphate (g) Salt 22.78 400 9.11 Anolyte 3.72 1000 3.72 Total 12.83 Final Salt 17.86 387 6.91 Anolyte 5.65 970 5.48 Total 12.39 Mass lost due to sampling 0.37 Mass unaccounted for 0.07 Deviation (%) 0.58 Accountability (%) 99.42