The Electrodialysis of Lithium Sulphate to Lithium Hydroxide€¦ · The Electrodialysis of Lithium Sulphate to Lithium Hydroxide ENG470 – Engineering Honours Thesis ... History

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

This page was intentionally left blank

I

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

II

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.

III

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.

IV

This page was intentionally left blank

V

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

VI

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 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

VII

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

VIII


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


experiment 4: 40°C .............................................................................................. 49


experiment 5: 60°C .............................................................................................. 49


experiment 4: 40°C .............................................................................................. 50


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


experiment 6: 15 wt.% ......................................................................................... 55


experiment 7: 10 wt.% ......................................................................................... 56


experiment 8: 5 wt.% ........................................................................................... 56


experiment 6: 15 wt.% ......................................................................................... 57


experiment 7: 10 wt.% ......................................................................................... 58


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


experiment 9: 2 hours........................................................................................... 63


experiment 10: 4 hours......................................................................................... 63




experiment 9: 2 hours........................................................................................... 65





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

IX

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

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.

2

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.

3

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

4

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

[t]

Share 2015

[%]

Demand 2020

[t]

Share 2020

[%]

Batteries 53,629 35 76,673 34

Glasses and

Ceramics 0,549 32 86,717 38

Lubricating

greases 13,840 9 14,507 7

Polymers 8,650 4 2,315 1

Air

conditioning 8,650 5 5,325 2

Aluminium 1,730 1 0 0

Continuous

casting 10,380 5 18,478 8

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

5

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-

6

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

7

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

8

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).

9

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.

10

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)

11

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

12

(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).

13

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

14

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

15

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

16

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

9.

17

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.

18

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

19

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).

20

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).

21

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


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


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

22

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.

23

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

24

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).

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

from (Ying et al. 2008).

0

0.3

0.6

0.9

1.2

0 100 200 300 400

LiO

H C

on

cen

tra

tio

n (

mo

l/L

)

Time (min)

700 A/m^2

1000 A/m^2

1400 A/m^2

1800 A/m^2

25

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

0

3

6

9

12

15

0

20

40

60

80

100

400 800 1200 1600

En

erg

y C

on

sum

pti

on

(k

Wh

/k

g(L

iOH

))

Cu

rre

nt

efi

cie

ncy

(%

)

Current Density (A/m^2)

Current efficiecy

Energy consumption

26

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

27

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.

28

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.

29

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.

30

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.

31

Figure 14: Front View of Electrodialysis cell

32

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.

33

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

34

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

35

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

ID No.

pH Temperature

range (°C)

Salt (Li2SO4)

concentration

(wt.%)

Residence

time (hour)

Current

(A)

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

36

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.

37

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

38

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.

39

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.

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.

0

5

10

15

20

25

30

35

40

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5Lit

hu

m i

Co

nm

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

40

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.

0

5

10

15

20

25

30

35

0 1 2 3 4

Lit

hu

m i

Co

nm

cen

tra

tio

n (

g/

L)

Time (hours)

Anolyte

Salt

41

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

0

2

4

6

8

10

12

14

16

18

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potential

42

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

calculated by:

∑ 𝐴 = ∑ 𝐴𝑎𝑛𝑜𝑜𝑑𝑒 𝑠𝑢𝑟𝑓𝑎𝑐𝑒 − 2 ∑ 𝐴𝑣𝑜𝑖𝑑 𝑖𝑛 𝑚𝑒𝑠ℎ + ∑ 𝐴𝑖𝑛𝑠𝑖𝑑𝑒 𝑑𝑖𝑎𝑚𝑜𝑛𝑑𝑠 − ∑ 𝐴 𝑏𝑎𝑠𝑒

= 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).

𝑛𝐿𝑖(𝑇𝑟𝑎𝑛𝑠𝑓𝑒𝑟𝑟𝑒𝑑) = (𝐶𝐿𝑖,𝑡 × 𝑉𝐿𝑖,𝑡) − (𝐶𝐿𝑖,0 × 𝑉𝐿𝑖,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:

𝐶𝑢𝑟𝑟𝑒𝑛𝑡 𝐸𝑓𝑓𝑖𝑐𝑖𝑒𝑛𝑐𝑦 (ŋ) = 𝑛𝐿𝑖𝐹

𝐼𝑡× 100 (10)

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

43

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.

44

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

experiment 2: pH 11


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

0

5

10

15

20

25

30

35

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5Lit

hu

m C

on

mce

ntr

ati

on

(g

/L

)

Time (hours)

Catholyte

Salt

0

5

10

15

20

25

30

35

40

0 1 2 3 4 5

Lit

hu

m C

on

mce

ntr

ati

on

(g

/L

)

Time (hours)

Catholyte

Salt

45

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.


experiment 2: pH 11


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

0

5

10

15

20

25

30

0 1 2 3 4 5

Su

lph

ate

Co

nm

cen

tra

tio

n (

g/

L)

Time (hours)

Anolyte

Salt

0

5

10

15

20

25

30

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Anolyte

Salt

46

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.

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

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

0

2

4

6

8

10

12

14

16

18

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

til

(V)

Time (hours)

Cell Voltage

Cathodic potential

Anode potential

0

2

4

6

8

10

12

14

16

18

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

til

(V)

Time (hours)

Cell Voltage

Cathodic potential

Anode potential

47

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 %


Li Recovery in catholyte 1.53 13.0 - -


g/h g/h


48

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

Lithium Sulphate

g % g %


Recovery in catholyte 2 15.4 - -


g/h g/h


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.

49

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.


experiment 4: 40°C


experiment 5: 60°C

0

5

10

15

20

25

30

35

0 1 2 3 4 5

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

0

5

10

15

20

25

30

35

40

0 1 2 3 4 5

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

50

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.


experiment 4: 40°C

0

5

10

15

20

25

0 1 2 3 4 5Su

lph

ate

Co

nce

ntr

ati

on

(g

/L

)

Time (hours)

Anolyte

Salt

51


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.

0

5

10

15

20

25

0 1 2 3 4 5Su

lph

ate

Co

nce

ntr

ati

on

(g

/L

)

Time (hours)

Anolyte

Salt

52

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

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.

0

2

4

6

8

10

12

14

16

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potential

0

2

4

6

8

10

12

14

16

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potential

53

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.

54

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

Lithium Sulphate

g % g %


Recovery in catholyte 1.05 8.2 - -


g/h g/h


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

Lithium Sulphate

g % g %




g/h g/h


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

55

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.


experiment 6: 15 wt.%

0

2

4

6

8

10

12

14

16

18

20

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

56





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

0

2

4

6

8

10

12

14

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

0

1

2

3

4

5

6

7

8

9

10

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

57

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.



0

2

4

6

8

10

12

0 1 2 3 4 5Su

lph

ate

Co

nce

ntr

ati

on

(g

/L

)

Time (hours)

Anolyte

Salt

58





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

0

1

2

3

4

5

6

7

8

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Su

lph

ate

Co

nce

ntr

ati

on

(g

/L

)

Time (hours)

Anolyte

Salt

0

0.5

1

1.5

2

2.5

3

3.5

4

4.5

5

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Su

ph

ate

co

nce

ntr

ati

on

(g

/L

)

Time (hours)

Anolyte

Salt

59

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.

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

0

2

4

6

8

10

12

14

16

18

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potential

60

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

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

0

2

4

6

8

10

12

14

16

18

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode potential

Anode Potential

0

5

10

15

20

25

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potenial

61

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 %




g/h g/h



Lithium Sulphate

g % g %




g/h g/h


62


Lithium Sulphate

g % g %




g/h g/h


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.

63


experiment 9: 2 hours



0

5

10

15

20

25

30

35

40

0 0.5 1 1.5 2 2.5

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

0

5

10

15

20

25

30

35

40

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

64



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.

0

5

10

15

20

25

30

35

40

0 1 2 3 4 5 6 7 8

Lit

hu

m C

on

cen

tra

tio

n (

g/

L)

Time (hours)

Catholyte

Salt

65





0

5

10

15

20

25

0 0.5 1 1.5 2 2.5

Su

lph

ate

co

nce

ntr

ati

on

Time (hours)

Anolyte

Salt

0

5

10

15

20

25

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Su

lph

ate

Co

nce

ntr

ati

on

(g

/L

)

Time(hours)

Anolyte

Salt

66



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.

0

5

10

15

20

25

0 1 2 3 4 5 6 7 8

Su

lph

ate

Co

ncn

tra

tio

n (

g/

L)

Time (hours)

Anolyte

Salt

67

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


0

2

4

6

8

10

12

14

16

18

20

0 0.5 1 1.5 2 2.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potential

0

2

4

6

8

10

12

14

16

18

20

0 0.5 1 1.5 2 2.5 3 3.5 4 4.5

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potential

68


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.

0

2

4

6

8

10

12

14

16

18

20

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

Po

ten

tia

l (V

)

Time (hours)

Cell Voltage

Cathode Potential

Anode Potential

69

Table 11: Mass transfer and recovery, experiment 9

Lithium Sulphate

g % g %




g/h g/h



Lithium Sulphate

g % g %




g/h g/h



Lithium Sulphate

g % g %




g/h g/h


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.

70

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.

71

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

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

0

10

20

30

40

50

60

70

80

90

100

1 2 3 4 5 6 7 8 9 10 11

%

Experiment no.

Current efficiency

Lithium Recovery

0

0.5

1

1.5

2

2.5

1 2 3 4 5 6 7 8 9 10 11

g/

h

Experiment no.

LiOH production rate

H2SO4 production rate

72

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

1.

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.

73

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

74

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.

75

6. References

Ball, Donald L and Daniel A.D. Boatang. 1987.Method for the Recovery of Lithium

from Solutions by Electrodialysis. US4636295A.

Buckley, David, J. David Genders and Dan Atherton. 2011. Method of Making High

Purity Lithium Hydroxide and Hydrochloric Acid. US2011/0044882A1.

Daintith, John, ed. 2008. Oxford Dictionary of Chemistry. 6th ed. Oxford University

Press

Dakota Minerals. 2017. Lithium uses.

https://www.dakotaminerals.com.au/lithium/lithium-uses (Accessed:

15/6/2017)

Das, Rasel, Md. Eaqub Ali, Sharifah Bee Abd Hamid, Seeram Ramakrishna and Zara

Zaman Chowdhury. 2013. “Carbon nanotube membranes for water purification:

A bright future in water desalination.” Desalination 336: 97-109

http://dx.doi.org/10.1016/j.desal.2013.12.026.

Evans, R. Keith. 2008. An Abundance of Lithium.

http://citeseerx.ist.psu.edu/viewdoc/download?doi=10.1.1.363.1242&rep=rep1

&type=pdf.

Ge, Liang, Bin Wu, Qiuhua Li, Yaqin Wang, Dongbo Yu, Liang Wu, Jiefeng Pan,

Jibin Miao and Tongwen Xu. 2015. “Electrodialysis with nanofiltration

membrane (EDNF) for high-efficiency cations fractionation” Journal of

Membrane Science 498: 192-200.

http://dx.doi.org/10.1016/j.memsci.2015.10.001.

Hoshino, Tsuyoshi. 2013. Preliminary Studies of Lithium Recovery Technology from

Seawater by Electrodialysis using ionic Liquid Membrane. Japan: Elsevier B.V.

https://www.dakotaminerals.com.au/lithium/lithium-uses

76

Hoshino, T. 2014. “Lithium Recovery from Seawater by Electrodialysis using Ionic

Liquid-based Membrane Technology.” The Electrochemical Society 58

(48):173-177. http://doi.org/10.1149/05848.0173ecst

Hwang, Chi Won, Min Ho Jeong, Young Joong Kim, Won Keun Son, Kyung Suk

Kang, Chang Soo Lee and Taek Sung Hwang. 2016. “Process design for lithium

recovery using bipolar membrane electrodialysis system.” Separation and

purification 166, (March) :34-40.

http://dx.doi.org/10.1016/j.seppur.2016.03.013

Koros, William J. 2002. “Gas separation membranes: needs for combined mterils

science and processing approaches.” Special Issue: Polymer Membranes 188 ,

no. 1(November): 13-22. http://doi.org/10.1002/1521-

3900(200211)188:1<13::AID-MASY13>3.0.CO;2-W

Liu, Chaofeng , Zachary G. Neale and Guozhong Cao. 2016. “Understanding

electrochemical potentials of cathode materials in rechargeable batteries.”

Minerals today 19, no.2 (March):109-123.

https://doi.org/10.1016/j.mattod.2015.10.009

Martin, Gunther, Lars Rentsch, Michael Höck and Martin Bertau. 2017. “Lithium

market research- global supply, future demand and price development.” Energy

Storage Materials 6, (November): 171-179.

http://dx.doi.org/10.1016/j.ensm.2016.11.004

Mazrou, S., H. Kerdjoudj and A.T. Cherif. 1997. “Sodium hydroxide and

hydrochloric acid generation from sodium chloride and rock salt by electro-

electrodialysis.” Journal of Applied Electrochemistry 27 no. 5, (May): 558-567.

https://doi.org/10.1023/A:1018498612326

http://doi.org/

https://doi.org/10.1016/j.mattod.2015.10.009

77

Mroczek, Ed, Gaetano Dedual, Duncan Graham and Lew Bacon. 2015. Lithium

Extraction from Wairakei Geothermal Fluid using Electrodialysis. Melbourne:

Proceedings World Geothermal Congress.

https://pangea.stanford.edu/ERE/db/WGC/papers/WGC/2015/39000.pdf

Nie, Xiao-Yao, Shu-Ying Sun, Xingfu Song and Jian-Guo Yu. 2017. “Further

investigation into lithium recovery from salk lake brines with different feed

characteristics by electrodialysis.” Journal of Membrane Science 530, (May):

185-191. http://dx.doi.org/10.1016/j.memsci.2017.02.020

O’Brien, Thomas F., Tilak V.Bommaraju and Fumio Hine. 2005. Handbook of Chlor-

Alkali Technology Volume I: Fundamentals. New York: Springer Science +

Business Media Inc.

Reahl, Eugene. R. 2006. Half a Century of Desalination With Electrodialysis.

https://pdfs.semanticscholar.org/5391/c71537ced99f87e457413e9fb38776dc4ab

c.pdf

Rottiers, T., G De la Marche , B.Van der Bruggen and L Pinoy . 2015. “Co-ion fluxes

of simple inorganic ions in electrodialysis metathesis and conventional

electrodialysis.” Journal of Membrane Science 492, (October): 263-270.

http://dx.doi.org/10.1016/j.memsci.2015.05.066

Royal Society of Chemistry. 2017. Periodic table, Lithium.

http://www.rsc.org/periodic-table/element/3/lithium (Accessed: 20/4/2017)

Serre, Elodie, Elodie Rozoy, Karine Pedneault, Stella Lacour and Laurent Bazinet.

2016. “Deacidification of cranberry juice by electrodialysis: Impact of

membrane types and configurations on acid migration and juice

http://www.rsc.org/periodic-table/element/3/lithium

78

physicochemical characteristics.” Separation and Purification Technology 163,

(May):228-237. http://dx.doi.org/10.1016/j.seppur.2016.02.044

Sharma, Yatendra. 2016. Production of lithium hydroxide. CA 2966525 A1.

https://encrypted.google.com/patents/CA2966525A1?cl=en

Sociedad Minera Salar de Atacama.1999. Production of Lithium Carbonate from

Brines. US5993759 A. https://www.google.com/patents/US5993759

Tahil, William. 2007. The Trouble with Lithium. http://www.meridian-int-

res.com/Projects/Lithium_Problem_2.pdf

Tanaka, Yoshinobu, Seung-Hyeon Moon, Victor V.Nikonenko and Tongwon Xu.

2012. “Ion-Exchange Membranes.” International Journal of Chemical

Engineering 2012, (August): 1-3.http://doi.org/10.1155/2012/906952

Unites States Geological Survey. 2017. Mineral Commodity Summaries: Lithium.

https://minerals.usgs.gov/minerals/pubs/commodity/lithium/mcs-2017-lithi.pdf

Valero, Fernando, Angel Barcelo and Ramon Arbos. 2011. Electrodialysis

Technology- Theory and Applications. Desalination, Trends and Technologies.

Michael Schorr (ed.), ISBN:978-953-307-311-8, InTech. Available from:

http://www.intechopen.com/books/desalination-trends-and-

technologies/electrodialysistechnology-theory-and-applications

Ying, Xu, Wang Xiaolin, Chen Feiguo and Yu Lixin. 2008. “Preparation of lithum

hydroxide from lithium sulphate in a three- compartment membrane

electrodialysis system.” ProQuest Dissertations & Theses Global (1026937088),

http://libproxy.murdoch.edu.au/login?url=http://search.proquest.com/docview/1

026937088?accountid=12629

Wang, Yaoming, Xu Zhang and Tongwen Xu. (2010). “integration of conventional

electrodialysis and electrodialysis with bipolar membranes for production

http://www.intechopen.com/books/desalination-trends-and-technologies/electrodialysistechnology-theory-and-applications

http://www.intechopen.com/books/desalination-trends-and-technologies/electrodialysistechnology-theory-and-applications

http://libproxy.murdoch.edu.au/login?url=http://search.proquest.com/docview/1026937088?accountid=12629

http://libproxy.murdoch.edu.au/login?url=http://search.proquest.com/docview/1026937088?accountid=12629

79

of organic acids.” Journal of Mmbrane Science 365 (1-2): 294-301

https://doi.org/10.1016/j.memsci.2010.09.018

Warburton, Simon. 2016. “FMC triples lithium hydroxide production.” MLA 8.

http://www.todaysmotorvehicles.com/article/fmc-lithium-expansion-

electric-vehicles-battery-052416/ (Accessed 13/7/2017)

Xu, Tongwen and Chuanhui Huang. (2008). “Electrodialysis- Based Separation

Technologies: A Critical Review.” AIChE Journal 54, (12) : 3147-3159.

http://doi.org/10.1002/aic.11643

https://doi.org/10.1016/j.memsci.2010.09.018

http://www.todaysmotorvehicles.com/article/fmc-lithium-expansion-electric-vehicles-battery-052416/

http://www.todaysmotorvehicles.com/article/fmc-lithium-expansion-electric-vehicles-battery-052416/

80

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

Moles of NaOH used:

𝑁𝑁𝑎𝑂𝐻 = 𝑇𝑖𝑡𝑟𝑎𝑛𝑡 𝑣𝑜𝑙𝑢𝑚𝑒 (𝐿) × 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑜𝑓 𝑇𝑖𝑡𝑟𝑎𝑛𝑡 (4)

𝐻2𝑆𝑂4 + 2𝑁𝑎𝑂𝐻 → 2𝐻2𝑂 + 𝑆𝑂42− + 2𝑁𝑎 (5)

∴ 1 𝑚𝑜𝑙 𝑜𝑓 𝐻2𝑆𝑂4 𝑡𝑜 2 𝑚𝑜𝑙 𝑜𝑓 𝑁𝑎𝑂𝐻

Moles of H2SO4:

𝑁𝐻2𝑆𝑂4=

𝑁𝑁𝑎𝑂𝐻

2= 𝑁𝑆𝑂4

(6)

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.

𝐶𝑆𝑂4(𝑚𝑜𝑙 𝐿⁄ ) =

𝑁𝑆𝑂4

0.001 (7)

𝐶𝑆𝑂4(𝑔 𝐿⁄ ) = 𝐶𝑆𝑂4

(𝑚𝑜𝑙 𝐿⁄ ) × 𝑀𝑆𝑂4 (8)

81

Figure 53 Middle compartment of electrodialysis cell

10.4 cm

4 cm

17.3 cm

5 cm

6.5 cm

82

Figure 54: Dimensions of Electrodialysis Cell

83

Figure 55 Apparatus placement within the electrodialysis cell

84

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

85

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




86

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

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



Initial


(g) Salt 23.58 400 9.43 Anolyte 4.04 1000 4.04

Total 13.47




87

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

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



Initial


(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



88

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

Lithium

Initial


(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




(g)





89

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

Lithium

Initial


(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




(g)





90


Lithium

Initial


(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




(g)





91


Lithium

Initial


(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




(g)





92

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

Lithium

Initial


(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




(g)





93


Lithium

Initial


(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




(g)





94


Lithium

Initial


(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




(g)





Documents

The Electrodialysis of Lithium Sulphate to Lithium Hydroxide€¦ · The Electrodialysis of Lithium Sulphate to Lithium Hydroxide ENG470 – Engineering Honours Thesis ... History