Recovery of Lithium from Geothermal Brine with Lithium-Aluminum

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Recovery of Lithium from Geothermal Brine with LithiumAluminum Layered Double Hydroxide Chloride Sorbents Mariappan Parans Paranthaman, Ling Li, Jiaqi Luo, Thomas Hoke, Huseyin Ucar, Bruce A. Moyer, and Stephen Harrison Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b03464 • Publication Date (Web): 27 Oct 2017 Downloaded from http://pubs.acs.org on October 28, 2017

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Environmental Science & Technology

Recovery of Lithium from Geothermal Brine with Lithium-Aluminum Layered Double Hydroxide Chloride Sorbents

1 2 3 4 5 6 7 8 9 10 11

Mariappan Parans Paranthaman1*, Ling Li1, Jiaqi Luo1, Thomas Hoke1, Huseyin Ucar1, Bruce A. Moyer1, and Stephen Harrison2** 1 2

Chemical Sciences Division, Oak Ridge National Laboratory, Oak Ridge, TN 37831, USA Rakehill Technologies, LLC, Benicia, CA 94510, USA

ABSTRACT

12

We report a three-stage bench-scale column extraction process to selectively extract lithium

13

chloride from geothermal brine. The goal of this research is to develop materials and processing

14

technologies to improve the economics of lithium extraction and production from naturally

15

occurring geothermal and other brines for energy storage applications. A novel sorbent, lithium

16

aluminum layered double hydroxide chloride (LDH), is synthesized and characterized with X-ray

17

powder diffraction, scanning electron microscopy, inductively coupled plasma optical emission

18

spectrometry (ICP-OES), and thermogravimetric analysis. Each cycle of the column extraction

19

process consists of three steps: 1) loading the sorbent with lithium chloride from brine; 2)

20

intermediate washing to remove unwanted ions; 3) final washing for unloading the lithium

21

chloride ions. Our experimental analysis of eluate vs. feed concentrations of Li and competing

22

ions demonstrates that our optimized sorbents can achieve a recovery efficiency of ~91% and

23

possess excellent Li apparent selectivity of 47.8 compared to Na ions and 212 compared to K

24

ions, respectively in the brine. The present work demonstrates that LDH is an effective sorbent

25

for selective extraction of lithium from brines, thus offering the possibility of effective

26

application of lithium salts in lithium-ion batteries leading to a fundamental shift in the lithium

27

supply chain.

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28 29

INTRODUCTION

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The global demand for lithium salts, the raw material for lithium-ion batteries, has increased

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significantly in recent years as a result of technological advancements in electronic devices and

32

electric cars requiring high energy-density batteries.1-3 The brine lake deposits, pegmatite

33

minerals, and clays currently provide the major resources for the near-critical element lithium.4

34

Until the late 1990’s, lithium production was dominated by industries utilizing spodumene and

35

pegmatite mineral deposits found in the United States. However, South America's brine lakes in

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Chile and Argentina together with Australia's spodumene currently account for about 90% of

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global lithium production estimated to be 135,000 tpa in 2016.5 Production from Australian

38

based spodumene concentrate is converted into lithium products in China. Not yet exploited,

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geothermal brines in particular are of interest for a variety of reasons.6 First, certain geothermal

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brines provide a source of geothermal power due to the fact that hot geothermal lakes are stored

41

at high pressure underground, which when released to atmospheric pressure, can provide a flash-

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steam of nearly 325 °C. The flash-steam can be used to generate power. In certain geothermal

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waters and brines, associated binary processes can be used to heat secondary high temperature

44

fluids, which can provide steam for the generation of electricity without the flashing of the

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geothermal brine. Additionally, geothermal brines contain a variety of salts, which can be

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recovered and utilized for high value by-products of energy production. It is known that

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geothermal brines can include various metal ions, particularly alkali and alkaline earth metals, as

48

well as transition metals, such as manganese, zinc, in varying concentrations, depending upon

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the source of the brine. Recovery of these metals is potentially important to the chemical,

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agricultural and pharmaceutical industries. 2 ACS Paragon Plus Environment

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As lithium has gained importance as a near-critical element for use in various energy storage

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applications,7 there are continuing efforts to develop simple and inexpensive methods for the

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recovery of lithium. For example in a typical 50 MW geothermal plant in the Salton Sea, USA,

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15,000 tons of lithium carbonate or hydroxide salts can be produced annually by recovering and

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converting of lithium chloride from the geothermal plant waste solutions.8 In addition, there are

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about 390 MW of geothermal power currently produced in the Salton Sea Known Geothermal

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Region (SSKGR), USA from the deep geothermal resource.8 This is approximately equivalent to

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120,000 tons of lithium carbonate production. Hence, the lithium salts produced from geothermal

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brine solutions will be potentially low-cost and provide a significant supply of lithium resources.

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Compared to recovering lithium from hard-rock mining, conventional recovery from brine is

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much cheaper, eco-friendly and simple, as it only involves evaporating the brine under solar heat

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to achieve the lithium product.9 However, this evaporation process has drawbacks such as low

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lithium grades, high dispersions of compositions, uncertainty of recovery rate, and a series of

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time-consuming steps.1 Various alternative methods such as ion exchange (adsorption),10-13

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hydrometallurgy,4 and solvent extraction14 have been developed to extract lithium from naturally

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occurring brine solutions. For example, An et al.4 used a two-stage hydrometallurgy process to

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recover lithium from brine collected from Salar de Uyuni, Bolivia. On the other hand, ion

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exchange utilizing resins or sorbent materials is suitable for the recovery of lower concentrations

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of Li ions from brines. Several sorbents such as manganese oxides,10-11, 13 and layered hydrogen

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titanate, H2TiO312 have been evaluated for lithium recovery. Nevertheless, the quality of the

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current developed sorbent materials is not completely satisfactory in terms of recovery

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efficiency, Li capacity, selectivity relative to other competing bulk ions, and stability. More

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efficient sorbent materials and associated process technology for lithium extraction are therefore



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needed to be developed. Previous studies have found that lithium salts can be intercalated into

75

gibbsite (γ-Al(OH)3), yielding a layered lithium aluminum double hydroxide chloride (LDH).15-16

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LDH crystallizes in hexagonal symmetry with a space group P63/mcm at room temperature. X-

77

ray and neutron powder diffraction patterns have revealed that the structure of the LDHs consists

78

of [LiAl2(OH)6]+ layers separated by water molecules and hydroxide ions.17 In fact, LDHs have

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captured attention for a while owing to a variety of applications in catalyst, drug delivery agents,

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adsorbents, and ion scavenging.16,

81

characterization of the lithium aluminum double hydroxide chloride ((nominal composition

82

LiCl.mAl(OH)3.nH2O) with varying Li:Al ratios), and evaluate the efficiency and Li selectivity

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of these sorbents for lithium recovery from geothermal brine in column experiments.

18-23

In this study, we focus on the synthesis and

84 85

EXPERIMENTAL METHODS

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Sorbent synthesis

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Lithium aluminum layered double hydroxide chlorides (LDHs), LiCl.mAl(OH)3.nH2O with Li:Al

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molar ratio of 1:3 have been synthesized as follows.24 Lithium hydroxide (LiOH.H2O, Alfa

89

Aesar, 99.9% pure) and aluminum hydroxide (gibbsite Al(OH)3, Strem Chemicals, 99% pure)

90

were used as the starting materials. LiOH.H2O and Al(OH)3 were mixed in a stainless steel

91

container with approximately 75 wt.% of de-ionized (DI) water. The mixture was stirred for 48 h

92

in an oil bath held in individual preparations at 18 °C, 21 °C, 30 °C, 40 °C, and 90 °C.

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Additional DI water was added to the mixture to prevent the samples from drying out. Equivalent

94

mole ratios of HCl to initial LiOH.H2O were slowly added to the gel-like samples drop by drop

95

with stirring to replace the OH– with Cl– ions associated with Li+. The pH of the solution was

96

always kept above 5.5. After the ion exchange step, the mixture was stirred for 1–2 h to allow for


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97

a complete reaction. The mixture was then washed with DI water and vacuum filtered overnight.

98

The resultant filtrate had a pH of 6. Finally, the sample was ground into granulated form with

99

average particle size of 150 µm for lithium extraction. After identifying the optimal synthesis

100

temperature of 90 °C, we also synthesized LDH with varied Li:Al molar ratios of 1:1.25; 1:1.5;

101

and 1:2 to optimize the composition and achieve a single LDH phase.

102 103

Bench scale column extraction setup

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A schematic picture of the column is shown in Figure 1. Three solutions namely the initial brine,

105

intermediate washing, and final strip solutions are loaded into the column successively. The

106

brine typically contains various salts including LiCl, NaCl, KCl, CaCl2, etc. (more details of the

107

brine compositions are provided in Table 1 and also in reference (2)). The wash solution contains

108

mainly NaCl with a small amount of LiCl to remove any unwanted ions trapped in the sorbents,

109

while the strip solution contains a low concentration of purely LiCl. The system is controlled by

110

a computer automatically through two pinch valves, and is constantly kept at 95 °C. The eluate

111

solutions are collected manually for each bed volume (BV). One bed volume is defined as BV =

112

Lπr2 (BV = Bed Volume, L = Bed length of the column in cm, and r = radius of the column = 0.5

113

cm). In the experiments, the column radius was 0.5 cm, with a bed height of 13 cm, with a

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starting mass of LDH (Li:Al = 1:1.25) of 8 g. Flow rates were adjusted to 1.33 mL/min for load.

115

Typically, the flow rates for brine solutions are 8 BV/h followed by 8 BV/h for wash and 2 BV/h

116

for strip solutions.

117



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118 119

Figure 1. Schematic of the bench scale column extraction setup.

120 121

Characterization

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Phase formation of all sorbents was investigated by X-ray powder diffraction (XRD) with a

123

PANalytical X`Pert Powder diffractometer with CuKα1 radiation (λ = 1.54059 Å). The feed and

124

eluate solutions as well as the sorbents were analyzed by inductively coupled plasma optical

125

emission spectroscopy (ICP-OES) at Simbol Materials, Inc. to obtain the elemental

126

concentrations. A scanning electron microscopy (SEM) fitted with an energy dispersive

127

spectrometer (Merlin ZEISS XG-3262) was used to investigate the microstructure of the

128

sorbents. Thermogravimetric analysis (TGA) were performed from room temperature up to 800

129

°C with a heating rate of 5 °C /min under flowing dry air to determine the water percentage

130

present in each sorbent.

131 132 6 ACS Paragon Plus Environment

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133

RESULTS AND DISCUSSION

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Optimizing the reaction temperature

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Figure 2 shows the XRD patterns of the LDH sorbents with nominal Li:Al = 1:3 synthesized at

136

various temperatures. Almost all peaks in each pattern can be indexed with a hexagonal

137

P63/mcm model. The data indicates a nearly single phase of layered LDH with a very small

138

amount of unreacted gibbsite γ-Al(OH)3. Furthermore, sharper peaks can be clearly observed for

139

the 90 °C sample compared to those for samples synthesized at lower temperatures, which

140

indicate that the 90 °C sorbent is of best crystallinity. This can be readily explained, from the

141

previous study of time-resolved in-situ X-ray diffraction experiments involving LiCl and

142

gibbsite.16 The temperature dependence of the intercalation reaction data of LiCl and Gibbsite

143

was fitted into Avrami-Erofe’s ev equation. By assuming Avrami-Erofe’ev kinetics, a two

144

dimensional diffusion-controlled growth model following instantaneous nucleation was

145

proposed. Hence, Avrami-Erofe’ev equation becomes

146 147

α = 1 - exp(-kt)

(1)

148 149

The validity of this equation for the whole process was confirmed by a Sharp-Hancock analysis

150

and the plot of ln k versus 1/T yielded Ea, activation energy of 27 kJ mol-1.16 This value is in

151

agreement with the activation energy for lithium intercalation in TiS2 giving Li0.2TiS2. Hence, a

152

highly crystalline sample is expected to be superior in that the well-defined layered structure

153

would allow for more efficient insertion/extraction of the LiCl ions.



*

Unreacted Gibbsite

90 °C

* *

Intensity (arb.u)

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40 °C

* * *

30 °C 21 °C 18 °C LDH Gibbsite Al(OH)3

10

20

30

40

50

60

70

80

90

2 theta (degrees)

154 155 156 157

Figure 2. XRD patterns of LDHs with nominal Li:Al molar ratio of 1:3 synthesized at 18 °C, 21 °C, 30 °C, 40 °C, and 90 °C. The SEM images in Figure 3 show that the morphology of the LDHs with Li:Al = 1:3

158

changes with temperature as synthesized at 18 °C, 30 °C, 40 °C, and 90 °C. The magnifications

159

for all the images are 100 kX. It can be observed that the sorbents synthesized at 18 °C, 30 °C,

160

and 40 °C have a needle/rose petal like morphology, whereas the morphology of the 90 °C

161

sorbent is transformed to flat and platelet. A platelet morphology could likely enhance the

162

lithium adsorption efficiency since it results in a higher surface area.25 Sorbents were not pre-

163

treated before ICP-OES analysis. To better evaluate the stoichiometry of the sorbent, we

164

dissolved the sorbents with nominal Li:Al = 1:3 synthesized at various temperatures in 5% nitric

165

acid solution, and determined the chemical composition via ICP-OES. As shown in Table S1, the

166

actual value of Li:Al ratio is very close to 1:3 irrespective of the synthesis temperatures.

167

Furthermore, TGA was employed to determine chemical composition of the sorbents. Figure 4

168

shows a total weight loss of 44% for the Li:Al = 1:3 sample as a function of temperature. The

169

weight loss is mainly due to water evaporation during heating, from which the H2O content is

170

estimated to be 0.7 moles. Similar weight loss was observed for other LDH compositions as well. 8 ACS Paragon Plus Environment

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171

To prepare single phase LDH sorbent and eliminate unwanted gibbsite impurities, we

172

synthesized the LDHs with varying Li:Al ratios. Figure 5 shows the XRD patterns of LDHs with

173

nominal Li:Al ratios of 1:3, 1:2, 1:1.5, and 1:1.25 synthesized at 90 °C. The nominal Li:Al = 1:3

174

sample had a small amount of unreacted gibbsite Al(OH)3 as indicated by the extra peak marked

175

by a star in Figure 5 and also in Figure 2, and the rest of the sorbents are single phase Li-Al

176

LDH. We also determined the chemical composition via ICP-OES of these sorbents, and the

177

results were reported in Table S2. We used single-phase nominal Li:Al = 1:1.25 sample

178

nominally LiCl.1.25Al(OH)3.mH2O as the sorbent for the bench-scale column evaluation run.

179

Note that the sorbents pre-treated by reacting with water at 95 °C to strip part of LiCl and make

180

the LDH sorbents deficient with cation vacancies suitable for LiCl intercalation before the

181

column evaluation runs.

182 183 184 185

Figure 3. SEM images showing the morphology of LDHs with nominal Li:Al ratio of 1:3 synthesized at 18 °C, 30 °C, 40 °C, and 90 °C.



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186

105 100 95

Al/Li = 3:1

90

Weight (%)

85 80 75 70 65

weight loss

60 55 50 0

100

200

300

400

500

600

700

800

900

o

Temperature ( C)

187 188 189 190

Figure 4. TGA data of LDH with nominal Li:Al ratio of 1:3 synthesized at 90 °C.

Intensity (arb.u)

*

Unreacted Gibbsite

Li:Al = 1:1.25 Li:Al = 1:1.5 Li:Al = 1:2

*

Li:Al = 1:3 LDH Gibbsite Al(OH)3

10

20

30

40

50

60

70

80

90

2 theta (degrees)

191 192 193

Figure 5. XRD patterns of LDHs with nominal Li:Al ratio of 1:3, 1:2, 1:1.5, and 1:1.25 synthesized at 90 °C. 10 ACS Paragon Plus Environment

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194 195 196

Bench scale column extraction

197

6 shows the metal-ion concentrations for different cations in both the eluate and feed solutions in

198

the course of the experiment. As discussed above, the sorbent had been synthesized through LiCl

199

intercalation into the hexagonal layers of gibbsite, Al(OH)3 as shown in the equation below:

200

Column experiments demonstrated the recovery of LiCl from simulated geothermal brine. Figure

LiOH + HCl + 3 Al(OH)3 + 0.7H2O

LiCl.3Al(OH)3.0.7H2O

(2)

201

While the three solutions are passing through the column, the sorbent composition changes

202

dynamically with LiCl ions coming in and out of the layered LDH structure. When brine is

203

loaded into the column, LiCl intercalates into the sorbent.

204

x LiCl(brine) + (1-x) LiCl.3Al(OH)3.0.7H2O ≈ LiCl.3Al(OH)3.0.7H2O

(3)

205

The washing step employs an aqueous intermediate washing solution to wash away unwanted

206

salts, particularly KCl, CaCl2, MgCl2 and B residing in the inter-particle volume of the sorbent

207

after initial loading of the brine step without removing lithium intercalated in the sorbent.

208

Finally, the strip eluent, which contains a small concentration of LiCl in aqueous solution,

209

simply reverses the LiCl uptake owing to the now drastically reduced chloride concentration in

210

the aqueous phase. The process cycle thus operates by chloride "swing". Compared with ion

211

exchange, chloride swing obviates the consumption of chemical reagents, such as HCl for

212

stripping, that would otherwise add cost and secondary waste issues.

213

Although the loading-elution cycle was not optimized, the concentration profiles shown in

214

Figure 6 show the uptake and elution of LiCl. It can be seen from Fig. 6(b) that while loading the

215

brine, the Li output concentration in the eluate solution decreases down to ~31 mg/L until Bed

216

Volume, BV6, and gradually breaks through at 360 mg/L at BV9 when the capacity saturation of

217

the sorbent is reached. Stripping was effected with 250 mg/L LiCl rather than pure water,



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because over-depletion of LiCl from the sorbent material causes irreversible degradation of the

219

LDH presumably back to gibbsite (as shown in Figure S1). We evaluated the sorbents through

220

X-ray powder diffraction analysis after each cycle of passing the water strip solutions containing

221

various concentrations of LiCl (0; 125; 250; 500 mg/L) through the column to determine the

222

stability of sorbents. Strip solutions containing 0 and 125 mg/L LiCl caused the degradation of

223

the LDH phase revealed by the formation of extra Gibbsite Al(OH)3 phase in the XRD patterns.

224

However, LDH phase is retained through the strip solution containing 250 and 500 mg/L.

225

Typical XRD patterns of LDH sorbents loaded with brine is shown in Figure S2. It may be seen

226

that on stripping the Li output concentration increases to a peak concentration of 2.3 g/L between

227

BV11 to BV13 as the Li ions temporarily stored in the sorbent diffuse into the low-chloride

228

aqueous stream. It should be noted that the Li ions present in the strip solution would also

229

contribute somewhat to this large increase. We have taken that into account for calculating the

230

separation factor for Li over other metal ions.

232 233 234 235 236 237 238 239 240 241 242 243 244 245 246 247 248 249 250 251

4

10

3

10

2

10

Li out Na out Ca out

1

10

4

2500

6

8

K out B out Mn out 10

12

14

2000 L i o ut (mg/ L)

231

Metal ion concen tr ati ons (mg /L)

218

1500

Li out Li in Li adsorbed

1000 500 0 Lo ad 2

4

Wash

6 8 10 12 Bed Vol ume (coun ts)

Strip 14

16

Figure 6. The metal-ion concentrations of the feed (in) and eluate (out) solutions as a function of bed volume. 12 ACS Paragon Plus Environment

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252

The recovery of Li in the loading step approached 91% in our run without optimization. Li

253

recovery efficiency was calculated from the ratio of total Li mass in the strip compared to the

254

total Li mass present in the brine solution. It is a fact that, since the sorbent material must retain a

255

certain amount of LiCl to maintain structural integrity, the eluent will always contain a small

256

level of Li at the minimum. We are currently addressing the question of the optimization of the

257

load-wash-strip cycle conditions as well as minimization of the LiCl concentration in the strip

258

solution to maintain the integrity of the sorbent. Selectivity is also very good for purposes of

259

treating geothermal brine. As shown in Table 1, the Li concentration at the peak strip

260

concentration at BV13 was much higher than that of any of the competing metals. Assuming a

261

fraction is collected at the peak concentration, the grade of LiCl would be excellent in relation to

262

its small concentration relative to the competing ions. Defining the separation factor for lithium

263

over another metal M as SFLi/M = ([Li]strip/[Li]brine)/ ([M]strip/[M]brine), where [Li]strip and [Li]brine

264

denotes the lithium concentration in the strip and brine solutions, respectively, where [M]strip and

265

[M]brine denotes the other metal cation concentrations in the strip and brine solutions, very

266

promising values are obtained, especially for sodium at SF = 34.3 (at BV=13), where BV = Bed

267

Volume; and for potassium at SF = 146.2 (at BV=13). Taking the combined bed volumes of

268

BV13-16, the separation factors improve because of the rapid decrease of concentrations of

269

competing ions. Possible tailing of the lithium concentration at BV16 is evident on stripping, and

270

we will pursue this question in future efforts. Efforts are also being made to modify the LDH

271

sorbent compositions or develop alternate extraction methods to improve Li selectivity and

272

capacity and also remove Na and K efficiently from the eluate solution so that high purity battery

273

grade lithium carbonate products can be successfully developed.

274



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275

CONCLUSION

276

A series of LDH sorbent materials with optimum process conditions have been prepared and

277

examined for potential use in lithium sorption from geothermal brine using a chloride

278

concentration swing principle. A Li recovery efficiency of ~91% and a high Li selectivity of 47.8

279

compared to Na ions and 212 compared to K ions from the brine have been achieved. Column

280

results, though not optimized for process purposes, indicate the potential for excellent recovery

281

with rejection of major competing ions, including Na, K, Ca, Mn, and B. Owing to the sorbent’s

282

low cost, ease of preparation, environmentally friendly nature, and straightforward load-wash-

283

elution behavior without consumption of reagent chemicals nor secondary waste, the LDH

284

sorbent seems promising for treating geothermal brine and for creating a new supply of lithium

285

capable of meeting a significant fraction of the future growth in lithium demand for energy

286

storage applications.

287 288 289 290

291 292 293

Table 1. Metal-ion concentrations and separation factors in a typical load wash strip run in a column experiment.

Metal Li

Conc. in brine (mg/L) 360

Conc. @ BV13 (mg/L) 2340

SFLi/M @ BV13*

Conc.@ BV13-16 (mg/L) 5079

Average conc. @ BV 13-16 1269.8

SFLi/M @ BV13-16*

Na

44000

7470

34.3

10474

2618.5

47.8

K

16500

657

146.2

886

221.5

212.0

Ca

30400

1660

106.6

2410

602.5

143.6

Mn

1420

199

41.5

361

90.25

44.8

B

390

19.5

116.4

35

8.75

126.9

*Contribution from the Li ions present in the strip solution was taken into account while determining the separation factor, SFLi/M = ([Li]strip/[Li]brine)/ ([M]strip/[M]brine). 14 ACS Paragon Plus Environment

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294 295

ASSOCIATED CONTENT

296

Supporting Information

297

Tables S1 and S2, Figures S1 and S2. The supporting information is available free of charge on

298

the ACS Publications website.

299

Keywords

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Li extraction; brine solution; LDH sorbents, synthesis, bench scale extraction

301

AUTHOR INFORMATION

302

Corresponding Author

303

∗

304

**Present address: Alger Alternative Energy, LLC formerly with Simbol Materials (Stephen

305

Harrison).

306

Notes

307

The authors declare no competing financial interest.

308

ACKNOWLEDGEMENTS

309

This work was supported by the Critical Materials Institute, an Energy Innovation Hub funded

310

by the U.S. Department of Energy, Office of Energy Efficiency and Renewable Energy,

311

Advanced Manufacturing Office. The research on the synthesis of sorbents (JL, TH) was

312

supported by the U.S. Department of Energy, Office of Science, Office of Workforce

313

Development for Teachers and Scientists (WDTS) under the Science Undergraduate Laboratory

314

Internship program. Thanks are due to Fred Sloop and Daejin Kim (ORNL) for ICP analysis.

Phone (865) 574-5045; e-mail: [email protected] (M. P. Paranthaman).



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Additional information

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This manuscript has been authored by UT-Battelle, LLC under Contract No. DE-AC05-

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00OR22725 with the U.S. Department of Energy. The United States Government retains and the

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publisher, by accepting the article for publication, acknowledges that the United States

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Government purposes. The Department of Energy will provide public access to these results of

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sponsored

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Table of Contents (TOC)

This is a schematic representation of the recovery of Li from the used geothermal brine that is coming out of a geothermal power plant and the application of extracted lithium salts for producing lithium-ion batteries.


Recovery of Lithium from Geothermal Brine with Lithium-Aluminum

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