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Hydrometallurgy 117–118 (2012) 64–70
Contents lists available at SciVerse ScienceDirect
Hydrometallurgy
journal homepage: www.elsevier.com/locate/hydromet
Recovery of lithium from Uyuni salar brine
Jeon Woong An a, Dong Jun Kang a, Khuyen Thi Tran b, Myong Jun Kim b, Tuti Lim c, Tam Tran b,⁎
a
b
c
Technology Research Institute, Korea Resources Corp, Seoul, South Korea
Department of Energy & Resources Engineering, Chonnam National University, Gwangju, South Korea
School of Life Sciences and Chemical Technology, Ngee-Ann Polytechnic, Singapore
a r t i c l e
i n f o
Article history:
Received 27 December 2011
Received in revised form 4 February 2012
Accepted 8 February 2012
Available online 16 February 2012
Keywords:
Uyuni salar brine
High purity lithium carbonate
Removal of magnesium and sulphate
Calcium
Boron
a b s t r a c t
A hydrometallurgical process was developed to recover lithium from a brine collected from Salar de Uyuni,
Bolivia, which contains saturated levels of Na, Cl and sulphate, low Li (0.7–0.9 g/L Li) and high Mg (15–
18 g/L Mg). Unlike other commercial salar brines currently being processed, the high levels of magnesium
and sulphate in Uyuni brine would create difficulties during processing if conventional techniques were
used. A two-stage precipitation was therefore first adopted in the process using lime to remove Mg and sulphate as Mg(OH)2 and gypsum (CaSO4.2H2O). Boron (at 0.8 g/L in the raw brine), a valuable metal yet deleterious impurity in lithium products, could also be mostly recovered from the brine by adsorption at a pH
lower than pH11.3 in this first stage. The residual Mg and Ca (including that added from lime) which were
subsequently precipitated as Ca–Mg oxalate could be roasted to make dolime (CaO ∙MgO) for re-use in the
first stage of precipitation. Evaporation of the treated brine up to 30 folds would produce 20 g/L Li liquors.
The salt produced during evaporation was a mixture of NaCl and KCl, containing acceptable levels of sulphate,
Mg, Ca, etc. The final precipitation of lithium at 80–90 °C produced a high purity (99.55%) and well crystalline
lithium carbonate.
© 2012 Elsevier B.V. All rights reserved.
1. Introduction
Lithium has found application in many industries, from the
manufacturing of glass, ceramics, rubbers and pharmaceuticals to
production of lithium-ion batteries. The global market share of lithium used in batteries has grown significantly over the last few years
and has reached 23% in 2010 (USGS, 2011). The demand for lithium
is forecast to increase by ~ 60% from 102,000 t to 162,000 t of lithium
carbonate equivalent in the next 5 years, with application in batteries
taking a large percentage (40,000 t) of this growth (Hykawy, 2010;
Siame and Pascoe, 2011). Yaksic and Tilton (2009) completed an extensive survey and reported that the current resource of lithium in
continental/salar brines is approximately 52.3 million tonnes of lithium equivalent, mostly in Chile, Argentina and Bolivia, of which
23.2 million tonnes is recoverable. On the other hand, lithium from
mineral resources accounts for 8.8 million tonnes, where large deposits have been located in the USA, Russia and China. The reserves
and recoverable resources of lithium were estimated by Evans
(2008) as 29.79 million tonnes Li.
Several commercial projects have been recently considered or developed to process lithium carbonate from lithium minerals. A mine
treating 1 million tonne spodumene ore per year to produce
131,000 tpa lithium concentrate (6% Li2O) was recently built by
⁎ Corresponding author. Tel.: + 82 62 530 1726; fax: + 82 62 530 0462.
E-mail address: [email protected] (T. Tran).
0304-386X/$ – see front matter © 2012 Elsevier B.V. All rights reserved.
doi:10.1016/j.hydromet.2012.02.008
Galaxy Resources in Western Australia (Galaxy Resources, 2011).
The concentrate was sent to its lithium carbonate plant in Jiang Su,
China to produce 17,000 tpa lithium carbonate of 99.9% grade. Other
projects are also considered for the processing of wastes from China
clay operations (Siame and Pascoe, 2011; Sitando and Crouse, 2012)
and zinnwaldite wastes from tin–tungsten mines in Cinovec, the
Czech Republic (Jandova et al., 2008, 2009, 2010). However, up till
now lithium has been mostly produced from salar brines which contain 0.06–0.15% Li due to the low cost of production as shown in
Table 1 below for major operations around the world (Evans, 2008;
Roskill, 2009). Brine production until 2013 by Chemetall, FMC and
SQM is forecast to increase by an annual rate of 4–5% per year, whereas up to a growth rate of 35% per year is expected for the lithium carbonate production from China, due to the significant expansion of the
lithium battery industry (Roskill, 2009).
Salar de Uyuni in the highland of Bolivia, ranked as one of the richest in the world, has a lithium reserve estimated at 5.4 million tonnes,
(USGS, 2011). Apart from the fact that the Uyuni salar is located well
above 3600 m altitude, the processing of this brine would present a
challenge as it has a high Mg/Li ratio (16–22:1) while the Li content
only averages b0.08% Li compared to other geothermal or salar brines
(Table 1). Apart from other by-products such as boric acid, potash,
sylvite (KCl), sylvinite (KCl ∙NaCl) commonly recovered from salar
brines, the value of Mg in the Uyuni brine is as high as for lithium if
a chemical grade Mg oxide or hydroxide could be produced.
The recovery of lithium from salar brine usually involves continuous solar pond evaporation of the brine in several stages until its
J.W. An et al. / Hydrometallurgy 117–118 (2012) 64–70
65
Table 1
Compositions of various brines of commercial value around the world (Boryta et al., 2011; Evans, 2008; Roskill, 2009).
Source
Clayton Valley, USA
Salton Sea, USA
Salar de Atacama, Chile
HombreMuerto, Argentina
Salar de Uyuni, Bolivia
Searles Lake, USA
Great Salt Lake, USA
Dead Sea, Israel
Sua Pan, India
Bonneville, USA
Zabuye, China
Taijinaier, China
Na
K
B
Li
Mg
Ca
Cl
SO4
(wt.%)
(wt.%)
(wt.%)
(wt.%)
(wt.%)
(wt.%)
(wt.%)
(wt.%)
4.69
5.00–7.00
9.1
9.9–10.30
7.06
11.08
3.70–8.70
3.01
6
8.3
7.29
5.63
0.4
1.30–2.40
2.36
0.24–0.97
1.17
2.53
0.26–0.72
0.56
0.2
0.5
1.66
0.44
0.005
0.039
0.04
0.0163
0.01–0.04
0.157
0.068–0.121
0.0321
0.0054
0.0018
0.0012
0.002
0.0057
0.0489
0.031
0.019
0.07–0.57
0.965
0.018–0.14
0.65
0.4
0.0026
2.02
0.0057
0.0106
0.02
7.26
14.20–20.90
18.95
15.80–16.80
5
12.3
7.00–15.60
16.1
7.09
14
9.53
13.42
0.34
42–50
1.59
0.53–1.14
0.5–0.97
3.09
0.045
2.26–3.9
0.045
0.019–0.09
0.0306
0.0016
0.026–0.036
1.29
0.071
0.007
0.003
0.007
content reaches 6% Li. Products of 99.5% to 99.99% lithium carbonate
can now be manufactured using brine evaporation, precipitation of
K and Mg, solvent extraction or ion exchange to remove other impurities (Boryta et al., 2002, 2011).
In salar brines containing Mg, carnallite (KCl.MgCl2.6H2O) or
bishoffite (MgCl2.6H2O) starts precipitating when the brine is concentrated to ~ 4.4% Li (Atashi et al., 2010; Boryta et al., 2011;
Wilkomirsky, 1999). The solar evaporation reaches 5.5–6.5% Li before
the concentrated brine is processed to lithium carbonate or chloride.
By then lithium carnallite, LiCl.MgCl2.6H2O, also precipitates thus reducing the lithium recovery from the process. The brine by then will
contain ~ 6% Li, 1–4% Mg, 0.5–1% boron (as borate). The separate recovery of Mg is therefore necessary for the production of high purity
lithium carbonate or lithium chloride. Boron (B) should also be recovered as apart from improving the process economics its removal from
the concentrated brine is essential for the production of low B feedstock required for advanced battery manufacturing or Li metal production (Boryta et al., 2011).
A new process scheme therefore needs to be developed for high
Mg salar brines such as that found in Uyuni, Bolivia. A different approach has to be implemented to remove Mg and B as by-products
before lithium carbonate can be recovered.
Mg can be removed from sea water, bittern or process liquors as
Mg(OH)2 using slaked dolomite (CaMgCO3) or lime (Ca(OH)2)
(Carson and Simandl, 1994; Karidakis et al., 2005). In the presence of
sulphate, however, gypsum (CaSO4.2H2O) is normally formed. The
Mg(OH)2–CaSO4∙2H2O mixture has all properties (BET area, oil adsorption, particle size, whiteness, etc.) required to be used as fire
retarding fillers (Hull et al., 2011; Karidakis et al., 2005). However, further refining of this mixture is required before Mg can be recovered as
a chemical grade product or as feedstock for Mg metal production. The
use of NaOH to precipitate Mg(OH)2 would no doubt produce a high
purity Mg product. However the process yields very fine and poorly
crystalline Mg(OH)2, creating difficulty during the solid/liquid separation stage required for its recovery from the highly viscous brine
(Baird et al., 1988; Henrist et al., 2003). Attempts have been made to
promote the “ripening” of nano-sized primary crystals of Mg(OH)2
to improve its settling and filterability by aging the precipitate in
24 h in non-stirring conditions (Turek and Gnot, 1995). Alamdari et
4.61
0.94–2.00
0.061
0.83
3.41
al. (2008) could precipitate Mg(OH)2 of a larger particle size range
(5–60 μm) within 60 min from a sea bittern containing 30 g/L Mg by
adding seeds (5–40 μm) and adopting slow NaOH addition.
After Ca, Mg, B and sulphate are recovered, the concentrated brine
can then be purified to remove other residual impurities. The clean
brine is then subjected to carbonation to precipitate lithium carbonate using sodium carbonate. The conventional process used by most
plant operations can produce lithium carbonate of >99.5% grade.
This study therefore focuses on the development of a general process
to treat the salar brine from Uyuni, Bolivia and the optimization of the
unit operations used to recover a high purity lithium carbonate. The
process considered also has to recover Mg, B and K as by-products.
2. Experimental
2.1. Materials
Approximately 15 m 3 of the salar brine from Salar de Uyuni was
sampled for the experimental testwork conducted at our laboratories
in Korea. The brine was analysed by ICP-AES and its composition is
shown in Table 2. Hot plates or temperature-controlled water baths
were used in experiments where heating was required. All chemicals
used in the study were of analytical grade.
2.2. Equipment used
In this study the chemical compositions of the brines were determined by ICP-AES (Optima-5300DV, Perkin-Elmer) or IC (ICS-2000,
Dionex). The structure of lithium carbonate product was determined
by X-ray diffraction analysis (Cu-tube60kv 50 Ma, Philips). The zeta
potential of the Mg(OH)2 precipitate was measured using an instrument from Otsuka Electronics (ELZ-Z2).
2.3. Experimental techniques
The precipitation of Mg and Ca salts was conducted at ambient
temperature (20–22 °C) using various reactors incorporating different modes of stirring. Different precipitants were added (to known
volumes of brine, either 2 L or 4 L) in solid form to avoid the dilution
Table 2
Compositions (in g/L) of the Uyuni salar brine before and after treatment.
Brine before/after treatment
Na
K
Ca
Mg
Li
B
Cl
SO4
Raw Brine 1
Raw Brine 2
After lime precipitation
After oxalate precipitation
After evaporation to 20 g/L Li
After evaporation to 30 g/L Li
105.4
95.01
92.77
115.2
56.44
37.61
15.7
16.9
16.3
15.9
52.5
41.3
3.33
0.34
13.96
b0.05
b0.05
b0.05
16.7
17.4
0.30
0.17
0.35
0.65
0.84
0.76
0.82
0.82
19.8
30.3
0.70
0.68
0.36
0.20
2.91
4.22
203.7
187.5
189.5
196.3
201.4
215.6
21.3
22.6
0.67
0.76
20.2
13.3
66
J.W. An et al. / Hydrometallurgy 117–118 (2012) 64–70
Ca(OH)2solid
Brine
Mg, B Removal
L
Na2C2O4solid
S
Mg(OH)2
CaSO4.2H2O
Adsorbed B
Re-processing
to recover
Mg(OH)2 and B
Ca, Mg Removal
L
S
Roasting of Ca-Mg oxalate
CaO.MgOrecycling
Solar evaporation
Residual Ca, Mg, etc…
Purification
20 g/L Li
Li
Na2CO3
solid
Evaporation
Carbonation
80-90 oC
1.5 g/L Li
Carbonate treatment
before recycling to
solar evaporation
99.55%
Li2CO3
Fig. 1. Flowsheet developed for the recovery of lithium as carbonate from Uyuni brine.
of the original brine, which hindered the mass balance calculation
and the brine concentration afterwards. The precipitate suspension
was based on either (a) magnetic bar stirrer, (b) impeller-type stirrer
set at 200 rpm and (c) impeller plus ceramic balls for grinding the
precipitant added as solid. The use of ceramic balls (5 mm diameter)
added to the reactor at 5 g/L would help the reaction to reach equilibrium quickly within 1 h. Samples were taken during the precipitation
to confirm the reaction has reached completion. Brine evaporation
tests were also conducted by continuous boiling and the concentrated
liquors were separated from the precipitates by vacuum filtration or
centrifuging.
3. Results and discussion
Concentration (g/L)
This study investigates conditions for, (a) recovery of Mg, B and Ca
as by-products, (b) purification to remove deleterious impurities
including sulphate which causes gypsum formation during processing and (c) concentrating the brine by evaporation to achieve
20–60 g/L Li for the precipitation of lithium carbonate. Fig. 1 outlines
the major steps involved in the recovery of lithium and other byproducts. Each unit operation of this flowsheet was tested and the results are presented and discussed as follows.
3.1. Recovery of magnesium and boron
Different alternatives were evaluated to recover Mg, B and to reduce the levels of Ca and sulphate from the raw brine. Stabcal software (Huang, 2008) was first used to predict the conditions for
selectively precipitating Mg and Ca using different precipitants. A typical example of the results obtained from Stabcal modelling is shown
in Fig. 2 for the addition of lime to the raw brine at the Lime/Mg molar
ratio of 1:1 and 0.1:1. Using lime, Mg could be removed as Mg(OH)2
Ca
Mg
Na (1:1)
B
Li
Na (0.1:1)
pH
Fig. 2. Stabcal modelling showing the stability of different Ca, Mg, Li and B species at different pHs and lime additions of 1:1 lime/Mg molar ratio. For Na, the precipitation of Na2SO4∙10H2O was predicted at 0.1:1 lime/Mg addition at a low pH range, whereas at 1:1 lime/Mg molar ratio the precipitation takes place at pH > 12.
J.W. An et al. / Hydrometallurgy 117–118 (2012) 64–70
67
Fig. 3. Measured concentration profiles of Brine 2 after precipitation of Mg at different lime/Mg molar ratios. Continuous or dotted lines are those predicted by stoichiometric calculations (for Mg) or from Stabcal, except pH which was measured throughout the test.
at a pH > 8.6, whereas the precipitation of gypsum (CaSO4.2H2O) and
Na2SO4 takes place over a range of pH from pH2 to 12. The equilibrium concentrations of Ca, Mg, Na, and SO4 ions are therefore controlled by these precipitates during lime addition. Lithium is stable
as soluble Li + ions over a wide range of pHs whereas solid
(CaO)2∙ B2O3 and (CaO)3∙B2O3 could be formed at a pH > 12.
Several other precipitants (containing Na or Ca to avoid contaminating the brine) were also tested, including sodium carbonate, sodium
phosphate, caustic (sodium hydroxide) and sodium oxalate. The precipitation of Ca/Mg carbonate, phosphate, hydroxide and oxalate took
place rapidly and residual Ca/Mg concentration could reach steady
state within 1 h. All carbonate, phosphate and oxalate precipitates
could be recovered easily by filtration. However the Mg(OH)2 precipitate formed from NaOH addition was fine and non-crystalline (particle
sizes b10 μm) and could not be easily separated from the viscous brine.
Our preliminary studies based on precipitation tests and Stabcal
modelling showed Ca could be removed selectively in the first stage
using phosphate. Mg could also be removed as pure Mg3(PO4)2 salts
in the subsequent stage. These phosphates are valuable fertilizer
products. However even at additions of phosphate at PO4/Mg stoichiometric ratios >2:1 the residual Mg concentration in the treated
brine was still very high (at ~4 g/L Mg). This excessive use of phosphate would contaminate the liquor and cause further complications
in the following stages. When sodium carbonate was used as precipitant, the double salt of CaCO3.MgCO3 was formed in the neutral pH
range, negating the complete recovery of pure Mg products.
Lime is the best precipitant that can be used for removing Mg from
the brine, although it has to be added in slight excess for complete Mg
removal. The precipitation of gypsum (CaSO4.2H2O) also reduces the
sulphate to b1 g/L in the treated brine. The precipitate formed (at
pH > 8.6) is a mixture of gypsum and Mg(OH)2 which needs to be further processed to separate Mg from the gypsum waste. As shown
later, this precipitate also contained boron which was recovered by
adsorption during the first-stage precipitation. Further processing of
this mixed product is required to recover high purity Mg and B byproducts from the mixed precipitate. The results of this study will
be reported later. Residual Mg after lime precipitation and almost
all soluble Ca can then be subsequently removed by adding sodium
oxalate. This second-stage precipitate can be converted to MgO ∙CaO
(dolime) by roasting for re-use in the lime precipitation stage.
Stabcal modelling (Fig. 2) was also used to predict accurately the
equilibrium concentrations of Na, K, Li and Cl and other minor elements in the brine as shown in Fig. 3. At a lime addition of 0.1:1
Ca(OH)2/Mg molar ratio, the pH rises from pH7.1 to pH8.6, effecting
the precipitation of Mg(OH)2. Stabcal however could not predict the
Mg equilibrium conditions accurately as there was large variation of
Mg concentration with respect to pH along the steep precipitation
line of Mg(OH)2, as seen in Fig. 2. The predicted Mg equilibrium concentration therefore was stoichiometrically calculated from the lime
additions. Sodium sulphate started precipitating also (Fig. 2) within
this range of pH, causing slight dips in both Na and SO4 concentrations as measured by chemical analysis. As more lime is added,
Na2SO4 re-dissolves, releasing SO4 ions for the formation of the less
soluble gypsum. The steady state concentrations measured after one
hour however showed lower precipitated Mg compared to stoichiometric predictions (i.e. higher Mg concentration remaining in the
brine) indicating incomplete reaction of lime and Mg ions. This is
expected with lime neutralization as the formation of gypsum coating
0.7
Boron Concentration (g/L)
Zeta Potential (mV)
60
40
20
PZC
0
0.6
0.5
0.4
0.3
0.2
0.1
0
0
-20
7
8
9
10
11
12
13
0.2
0.4
0.6
0.8
1
Ca(OH)2/Mg2+ Ratio
pH
Fig. 4. Zeta potential measurements at different pHs showing the point-of-zero-charge
(PZC) at pH11.3.
Fig. 5. Removal of boron from Brine 1, at different consecutive additions of Ca(OH)2
(diamonds: addition at 0.1× stoichiometry each time, circles: at 0.2 × mole ratio, triangles: 0.3 × mole ratio).
68
J.W. An et al. / Hydrometallurgy 117–118 (2012) 64–70
120
220
Concentration (g/L)
Na
20
80
210
40
200
15
10
Cl
K
0
0.7
0
190
0
10
SO4
5
20
30
40
50
60
32
0
10
20
30
40
50
60
4.5
0
10
20
30
40
50
60
50
60
0.6
24
0.5
3
0.4
16
0.3
Mg
0.2
Li
B
1.5
8
0.1
0
0
10
20
30
40
50
60
0
0
0
10
20
30
40
50
60
0
10
20
30
40
Concentration Factor
Fig. 6. Concentration profile of concentrated brines after evaporation of Brine 1.
3.2. Recovery of residual magnesium and calcium
The residual Mg and Ca, including added amounts from lime,
could be removed using sodium oxalate (added at 1:1 sodium oxalate/[Ca + Mg] molar ratio to b50 ppm Ca and 170 ppm Mg). The concentrations of other ions after sodium oxalate treatment are also
shown in Table 2. The precipitate obtained is a mixed Ca–Mg oxalate
which can be converted to dolime (CaO∙MgO) by roasting for re-use
in the first stage of lime precipitation.
3.3. Concentration of lithium by evaporation
After Mg, B and Ca were removed the brine contained Li, Na and K
in a chloride brine of low sulphate. The loss of lithium throughout the
removal of Mg and Ca by lime and oxalate precipitation is minimal
(within 3% of the analytical analysis error) as indicated in Table 2.
Following conventional techniques, lithium could be concentrated
to 6% Li by solar evaporation. The operation was simulated in the laboratory in several trials to concentrate the brine several folds from the
initial volume of 4 L. In this context the degree of evaporation was
measured as the ratio between the initial volume and the volume of
brine recovered after evaporation and pressure filtration or centrifuging, defined as Concentration Factor, representing the number of folds
the brine was concentrated by. Heavy losses of Li due to evaporated
brine caught in NaCl, KCl and other complex salts precipitated during
evaporation were observed. As shown in Fig. 7, to achieve 20 g/L and
30 g/L Li, the maximum recovery of Li from evaporation and pressure
filtration is ~ 80% and 60%, respectively. In practice, using pond solar
evaporation the recovery is expected to be lower as the solid/liquid
separation operation would not be as effective as in laboratory
studies.
The cumulative losses of major ions are shown in Fig. 8. In the first
period when the concentration factor is b5, the precipitation of NaCl
is more predominant. After that (concentration factor: 5–20) KCl
co-precipitates with some loss of Mg measured, possibly as Mgchloride salt or by adsorption on NaCl and/or KCl. By mass balancing
the cumulative number of moles losses, i.e. Δ(mol), from the liquor
obtained after evaporation, it was confirmed that most of the salt
100
80
Li recovery (%)
would prevent complete dissolution of lime particles for the reaction.
The pH rose steadily from pH8.6 to pH9.2 at further additions of lime
until the 1:1 Ca(OH)2/Mg molar ratio was reached. After this, the pH
would rise to pH > 11.5 at a further addition of lime. Above pH12,
Stabcal predicts that sodium sulphate would re-precipitate at 1:1
lime/Mg molar ratio (as shown Fig. 2).
The recovery of boron during lime addition corresponding to the
lime/Mg molar ratio of 0.1:1 to 1:1 is mainly due to adsorption by
the Mg(OH)2-gypsum mixed precipitate. This was confirmed by the
measurement of zeta potential of the Mg(OH)2 precipitate. Fig. 4
shows the zeta potential Ezeta measured as 0–60 mV for the pH
range of pH8–11.3, with the zero-point-of-charge at pH11.3. Above
this pH, the surface of Mg(OH)2 is only lightly negatively charged. As
shown in the Stabcal stability diagrams (Fig. 2) boron exists as
HB4O7− or B(OH)4− cation within the range pH8–12, and therefore is
easily adsorbed onto Mg(OH)2 below the point-of-zero-charge. The
Stabcal modelling however could not predict the adsorption process
as shown in Fig. 3, where large discrepancies exist between the
model prediction and boron concentration measured. The solution
pH increased steadily to pH ~ 11.5 at further additions of lime until
1:1 stoichiometry. After this, the adsorbed boron was released from
the Mg(OH)2 back to the brine. However with lime addition past 1:1
stoichiometry, (CaO)2B2O3 and (CaO)3B2O3 started precipitating as indicated by Stabcal, causing incomplete recovery of boron at high pHs.
The adsorption process was reproducible at different modes of lime
addition as shown in Fig. 5 (at 0.1×, 0.2× or 0.3× lime/Mg stoichiometric ratio). Typical compositions of the brine after lime treatment,
added at 1:1 Ca(OH)2/Mg molar ratio are listed in Table 2.
60
40
20
0
0
10
20
30
40
50
Achieved Li concentration (g/L)
Fig. 7. Li recovery after consecutive evaporation of Brine 1 to achieve 10, 20, 30, 40 and
50 g/L Li targets.
J.W. An et al. / Hydrometallurgy 117–118 (2012) 64–70
25
Run 1
Run 2
Run 3
Run 4
Run 5
10
Δ(Mol)
20
1.5
15
Na
5
0
10
20
30
40
50
K
0.5
60
0
15
Run 1
Run 2
Run 3
Run 4
Run 5
1
0
0.025
Run 1
Run 2
Run 3
Run 4
Run 5
0.015
0.01
Mg
0.005
0
0
10
20
30
40
50
0
60
10
20
30
40
50
20
30
40
50
60
0.04
Li
0.05
0.03
Run 1
Run 2
Run 3
Run 4
Run 5
0.1
0
60
10
0.05
0.02
10
20
30
40
50
Run 1
Run 2
Run 3
Run 4
Run 5
B
0.01
0
0
Cl
5
0.15
0.02
Run 1
Run 2
Run 3
Run 4
Run 5
10
0.2
0.03
0
25
2
20
69
0
60
0
10
20
30
40
50
60
Concentration Factor
Fig. 8. Cumulative losses (as number of moles) of various ions after evaporation of Brine 1.
precipitated during the early stage of evaporation (at b20 folds) was a
mixture of NaCl and KCl (i.e. Δ(mol) (Cl −) = Δ(mol)(Na +) + Δ(mol)
(K +)). Up to the concentration factor of 20 the loss of lithium and
boron is minimal. This loss increases at further evaporation of the
brine. For example at the concentration factor of 30, the lithium concentration in the liquor was 20 g/L whereas Δ(mol)(Li +) is 0.1 mol,
representing losses of ~ 20% (of ~0.4 mole Li contained in 4 L of the
original brine of ~0.1 mol/L Li). Sulphate concentration shows a
one-to-one linear relationship with the concentration factor up to
20 folds, highlighting the fact that no sulphate salt will precipitate
up to this level of evaporation. A concentration ratio of 14 and 26
was measured for boron and sulphate at 20-fold evaporation confirmed minimal losses of these two ions (within experimental error
by division). At a higher concentration factor (>30) the loss of sulphate was also observed (Fig. 6) indicating sulphate salts started
precipitating.
Evaporative concentration up to 30-folds can therefore be conducted in future commercial operations to produce concentrated
brines of ~20 g/L Li.
Table 3
Composition of lithium carbonate (99.55%) product.
Component
Ca
Mg
Na
K
Li
B
Sr
Content (%)
Component
Content (ppm)
0.04
Fe
b 10
b 0.01
Cu
b 10
0.15
Al
b10
0.06
Ni
b10
18.7
Zn
b 10
0.06
Cr
b 10
0.04
Pb
b 10
in the range 80–90 °C, where the maximum solubility of lithium carbonate is between 1.5 and 1.8 g/L Li (corresponding to 7.8–8.2 g/L
lithium carbonate). Under these conditions, the recovery of lithium
in the final stage is expected to be higher than 90% using a feed liquor
of 20 g/L Li.
The product produced from this precipitation was a wellcrystalline materials as confirmed by its XRD pattern. After washing
in hot water, a high purity product (99.55% lithium carbonate)
could be recovered. The composition of this product is shown in
Table 3.
4. Conclusions
3.4. Lithium carbonate precipitation
At 20 g/L Li concentration, precipitation of lithium carbonate to recover the product can be performed using sodium carbonate precipitant. As shown in Fig. 9, the solubility of lithium carbonate decreases
at temperatures higher than ambient (Garrett, 2004). The carbonation stage therefore is usually conducted at elevated temperature
Li2CO3solubility (g/L)
16
14
12
10
A hydrometallurgical process was developed to recover lithium
from the Uyuni salar brine containing 15–18 g/L Mg, 0.7–0.9 g/L Li
and saturated with Na, Cl and sulphate. A two stage precipitation
was employed to remove Mg, Ca, B and sulphate from the brine
using lime first and then sodium oxalate. The mixed Mg(OH)2, gypsum and adsorbed boron from the first stage of precipitation needs
to be re-processed to produce high purity Mg and B products. Residual Mg and Ca can then be totally removed by sodium oxalate in the
second stage of precipitation. Following a 30-fold evaporation the
brine could be concentrated to achieve 20 g/L Li. Further polishing
purification can be conducted to remove residual Ca and Mg before
the treated brine was subjected to carbonation at 80–90 °C using sodium carbonate. A high purity (99.55%) and well crystalline lithium
carbonate could then be produced from this process.
Acknowledgement
8
6
0
20
40
60
80
100
Temperature (oC)
Fig. 9. Lithium carbonate solubility at different temperatures.
Data from Garrett (2004).
This work was supported by a grant from the Energy Efficiency &
Resources of the Korea Institute of Energy Technology Evaluation
and Planning (KETEP), Ministry of Knowledge Economy, Korea (no.
2010 T100100408).
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