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Extraction of lithium with functionalized lithium ion-sieves
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Extraction of lithium with functionalized lithium ion-sieves
Xin Xu a,b, Yongmei Chen a, Pingyu Wan a, Khaled Gasem b, Kaiying Wang a,b, Ting He c, ⇑ Hertanto Adidharma b, Maohong Fan b,d,e, a National Fundamental Research Laboratory of New Hazardous Chemicals, Beijing University of Chemical Technology, Beijing 100029, China b Department of Chemical and Petroleum Engineering, University of Wyoming, Laramie, WY 82071, USA c Idaho National Laboratory, Idaho Falls, ID 83415, USA d School of Energy Resources, University of Wyoming, Laramie, WY 82071, USA e School of Civil and Environmental Engineering, Georgia Institute of Technology, Atlanta 30332, USA article info abstract
Article history: Due to the technology advancement and the large-scale application of lithium-ion batteries Received 10 June 2016 in recent years, the market demand for lithium is growing rapidly and the availability of Received in revised form 30 August 2016 land lithium resources is decreasing significantly. As such, the focus of lithium extraction Accepted 18 September 2016 technologies has shifted to water lithium resources involving salt-lake brines and sea Available online 19 September 2016 water. Among various aqueous recovery technologies, the lithium ion-sieve (LIS) technol- ogy is considered the most promising one. This is because LISs are excellent adsorbents Keywords: with high lithium uptake capacity, superior lithium selectivity and good cycle perfor- Water lithium resources mance. These attributes have enabled LISs to separate lithium effectively from aqueous Lithium ion-sieve technology Lithium adsorbents solutions containing different ions. The present work reviews the latest development in Lithium adsorption/desorption mechanisms LIS technology, including the chemical structures of ion-sieves, the corresponding lithium LIS batteries adsorption/desorption mechanisms, the ion-sieves preparation methods, and the chal- LIS composite materials lenges associated with the lithium recovery from aqueous solutions by the LIS batteries. Besides, some common LIS composite materials forming technologies, including granula- tion, foaming, membrane and fiber formation, and magnetization, which are used to over- come the shortcomings in industrial column operations, are also explored. Ó 2016 Elsevier Ltd. All rights reserved.
Contents
1. Introduction ...... 277 2. LIS effect ...... 278 3. LISs overview ...... 279 3.1. LMO-type LISs ...... 279 3.1.1. Li-Mn-O ternary phase diagram ...... 279 3.1.2. Spinel structures of LMO precursors ...... 280 3.1.3. Doping modifications ...... 282 3.2. LTO-type LISs...... 283
3.2.1. Layered H2TiO3 ...... 283 3.2.2. Spinel titanium oxides ...... 284
⇑ Corresponding author at: Department of Chemical and Petroleum Engineering, University of Wyoming, Laramie, WY 82071, USA. E-mail addresses: [email protected] (X. Xu), [email protected] (Y. Chen), [email protected] (P. Wan), [email protected] (K. Gasem), [email protected] (K. Wang), [email protected] (T. He), [email protected] (H. Adidharma), [email protected], [email protected] (M. Fan). http://dx.doi.org/10.1016/j.pmatsci.2016.09.004 0079-6425/Ó 2016 Elsevier Ltd. All rights reserved. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 277
4. Lithium intercalation/de-intercalation mechanisms ...... 285 4.1. In spinel structures ...... 285 4.1.1. Redox mechanism...... 285 4.1.2. Ion exchange mechanism...... 286 4.1.3. Composite mechanism ...... 286 4.2. In layered structures ...... 286 5. Preparation methods of LIS precursors ...... 287 5.1. Solid-phase synthesis methods ...... 287 5.1.1. Solid-phase combustion method ...... 287 5.1.2. Microwave combustion method ...... 288 5.2. Soft-chemical synthesis methods ...... 288 5.2.1. Sol-gel method ...... 288 5.2.2. Hydrothermal method ...... 290 5.2.3. Co-precipitation method ...... 291 5.2.4. Molten-salt synthesis method ...... 292 5.2.5. Other methods ...... 292 6. Batteries for lithium recovery from aqueous solution ...... 294
6.1. k-MnO2 – Pt battery ...... 295
6.2. k-MnO2 – Ag battery ...... 295 6.3. The electrostatic-assisted lithium desorption process ...... 296 6.4. Outlook of the electrochemical method for recovering lithium ...... 297
6.4.1. Porous LiMn2O4 electrode ...... 297 6.4.2. Cost-effective counter electrodes...... 298 6.4.3. Coating modifications...... 298 6.4.4. The metal-doped LMO electrodes ...... 299 7. Forming of LISs for industrial applications ...... 299 7.1. Granulation ...... 300 7.2. Foaming ...... 302 7.3. Membrane formation ...... 303 7.4. Fiber formation and magnetization ...... 304 7.4.1. Fiber formation ...... 304 7.4.2. Magnetization ...... 304 8. Concluding remarks ...... 305 Acknowledgements ...... 307 References ...... 307
1. Introduction
Lithium is a critical energy material and a strategic resource for the 21th century. This lightest alkali metal plays a grow- ing role in numerous processes such as rechargeable batteries, thermonuclear fusion, medical drugs, lubricant greases, cera- mic glasses, air conditioning, dyes, cements, adhesives, and electrode welding [1–25]. Consequently, the market demands for lithium resources are increasing across the world [26–28]. Generally, lithium is obtained from two major resources: the lithium mineral ores including spodumene and petalite ores, and the lithium water resources including salt-lake brines and sea water [29–32]. Currently, the former is well exploited while the latter is being developed by industries with relatively low efficiency. Nonetheless, according to a reported research [29], more than 60% of the total lithium amount (about 26.9 Mt) exists in brines and sea water, especially those located in Chile, Bolivia, Argentina, China and the USA (Fig. 1). Therefore, a great potential exists for obtaining lithium from aqueous sources, if an efficient lithium recovery technology can be developed. Thus far, several technologies have been exploited for extracting lithium from the brines and sea water [33–44]. The most commonly used is the solar evaporation process [33–36], involving several stages to precipitate and crystallize NaCl, MgCO3, Mg(OH)2, and Li2CO3. However, this process takes several months to concentrate the brines and the Li2CO3 product is mixed with other minerals such as magnesium salts. An alternative method is co-precipitation process [37], where agents such as
AlCl3 are added to the brines followed by adjusting the pH with caustic alkali to obtain lithium aluminum hydroxide. This method, however, is highly constrained by the operating temperature, pH, Li/Al molar ratio. Another one, a solvent extrac- tion method, uses a specific chelating agent dipivaloylmethane [38,39] to recover lithium in the presence of various metal ions in aqueous solutions. However, this method could only be applied in low Mg/Li brines. More seriously, the corrosive additives may damage the process equipment. A hydrometallurgical method [40,41] that is designed for extracting lithium from brines containing high level of magnesium and sulfate, employs a two-stage precipitation to recover lithium. During the process, lime is added to the brines to precipitate Mg and Ca followed by solar evaporation. Lithium ion-sieve (LIS) is comprised of a series of lithium selective adsorbents with unique chemical structures cap- able of separating lithium ions from briny solutions and seawater. In the past several decades, the LIS technology is con- 278 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Fig. 1. Map of lithium resource availability and geostrategic impacts [29]. Reproduced from Ref. [29].
sidered the most promising technology for lithium recovery from aqueous solution [42–49]. The LIS technology provides a new effective method to recover lithium from solutions involving multiple coexisting ions (Na+,Mg2+,K+,Ca2+, etc.). Due to the superior lithium selectivity of LISs, high purity lithium products are readily accessible. As such, this technol- ogy exhibits high lithium separation efficiency in comparison to other technologies mentioned above. The main advan- tage is that no additional purification is needed in industrial operations. Additionally, LIS technology offers high theoretical lithium uptake capacity, low regeneration loss of raw materials, relatively less energy consumption, and an environmentally friendly lithium adsorption/desorption process. However, since its inception in 1980s, the practical applications in industry have been developing slowly due to poor virtual lithium uptake capacity, water pollutions by manganese dissolution, and poor powder forming process. Recently, their lithium adsorption/desorption efficiency have been greatly improved by exploring the electrochemical method, and they have become a significant research topic again. In this review, details on the chemical structures, the mechanisms for lithium intercalation and de-intercalation, and the preparation methods of several types of LISs will be discussed. In all, emphasis is placed on lithium uptake capacity, regen- eration performance, and the advantages and challenges of lithium recovery from aqueous solutions by the electrochemical method. Furthermore, some commonly applied material forming technologies (e.g., granulation, foaming, membrane and fiber formation, and magnetization) for LIS composites are illustrated to highlight the challenges in industrial column oper- ation, followed with some future prospects regarding LISs.
2. LIS effect
Since first prepared by Volkhin et al. [50] in 1971, ion-sieve oxides have received increasing attention in the past sev- eral decades due to special properties and performance as metal ions [51–55]. Ion-sieve oxides are adsorbents for specific metal ions with effective ion-screening performance. Such adsorbents are derived from the corresponding precursors that contain the target metal ions. Typically, the precursors present stable molecular frameworks, even if the target ions are stripped from their crystal sites, the vacant crystal sites still can be retained. Consequently, the formed vacant crystal sites could only accommodate ions the ionic radii of which are smaller than or equal to that of the target ions. For LISs, only lithium ions could reenter the vacant sites because lithium has the smallest ionic radius among all metal ions. This is evidenced by the fact that when LISs are placed in aqueous solutions containing different kinds of metal ions, only the lithium ions can be adsorbed. The working principle of LIS is shown in Fig. 2. The first step is the formation of LIS with hydrogen filled state [LIS (H)] by stripping the lithium ions from the lithium filled state [LIS (Li)], mainly through Li-H ion exchange, followed by the steric-effect based selective adsorption of lithium ions of LIS from Li+-containing solutions. Then, the spent LIS (H) is regenerated for formation of LIS (Li) through adsorption of the lithium ions. The overall process can be denoted as the ‘‘LIS effect” [56–60]. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 279
Fig. 2. Schematic representation of LIS process.
3. LISs overview
Generally, LISs can be classified into two types by chemical elements: the lithium manganese oxides-type (LMO-type) and lithium titanium oxides-type (LTO-type). The LMO-type LISs are the most popular selective lithium adsorbents cur- rently because of high lithium uptake capacities, excellent regeneration performance, and superior lithium selectivity. Additionally, the recovery of lithium from aqueous solutions has been greatly improved recently by applying electro- chemical techniques. However, this type of LISs suffers from manganese dissolution in aqueous solution, which may cause serious water pollution in industrial operations. In comparison, the LTO-type LISs can overcome this challenge because the titanium compounds are not harmful to the water environment and they can be removed from aqueous solution easily [61–63]. Furthermore, compared to the LMO-type LISs, the LTO-type LISs have much stable molecular structures, which is attributed to the large titanium-oxygen bond energy. However, the LTO-type LISs are of limited use in lithium recovery from aqueous solution when electrical potential is applied. This limitation may hinder future industrial applica- tions of LTO-type LISs. In conclusion, both the LMO-type and LTO-type LISs have their unique advantages and challenges. As such, future work will likely focus on minimizing their corresponding drawbacks to meet large-scale industrial applications.
3.1. LMO-type LISs
3.1.1. Li-Mn-O ternary phase diagram Since the discovery [64], some of the LMO-type LISs have been well developed [42,59,65]. Typically, their precursors pre- sent spinel structures. However, due to manganese multiple valence states, a number of lithium manganese oxides with dif- ferent crystal structures could be formed. Fig. 3 illustrates an isothermal cross section of the Li-Mn-O phase diagram at 25 °C 1 [66]. The stoichiometric defect spinel phases are defined by the Mn3O4 Li4Mn5O12 k-MnO2 triangle (the blue area in Fig. 3a). The stoichiometric spinel phases lie on the tie line between Mn3O4 and Li4Mn5O12, which can be represented by the general formula LixMn3 xO4 (0 6 x 6 1.33). The manganese-oxides defect spinels, located between Mn3O4 and k-MnO2, are represented by the general formula Mn3 xO4 (0 6 x 6 1). The lithium-manganese-oxide defect spinels are expressed by the general formula Li2O yMnO2 (y > 2.5) and lie on the tie line between Li4Mn5O12 and k-MnO2. Currently, the blue triangle of LiMn2O4 Li2Mn4O9- Li4Mn5O12 in Fig. 3b is the active area for preparing the precursors of LMO-type LISs. Therefore, it is possible to obtain high Li- + Mn precursors such as Li5Mn4O9 and Li7Mn5O12 in principle, implying that high Li capacity LISs may be obtained in the future. + Currently, only a few LMO-type LIS precursors with high Li adsorption capacities such as k-MnO2, MnO2 0.31H2O, and MnO2 0.5H2O, which are derived from LiMn2O4,Li4Mn5O12, and Li1.6Mn1.6O4 have been prepared respectively. Chitrakar et al. [67] built a phase diagram consisting of additional proton-type manganese oxides as a function of manganese valence state, Li/Mn, and H/Mn mole ratios, as shown in Fig. 4. In this figure, the LMO-type LIS precursors can be classified by two reaction types and exhibited in two perpendicular planes: the vertical plane is the redox reaction region and the horizontal
1 For interpretation of color in Fig. 3, the reader is referred to the web version of this article. 280 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Fig. 3. (a) An isothermal cross section of the Li-Mn-O phase diagram at 25 °C and (b) an expanded region of the Li-Mn-O phase diagram.
Fig. 4. Phase diagram of LMO and their delithiated products [67]. Reproduced from Ref. [67]. plane represents the ion exchange region. Their main lithium uptake properties from the aqueous solutions are primarily summarized in Table 1.
3.1.2. Spinel structures of LMO precursors Invariably, chemical structures determine chemical properties, thus the lithium recovery by LMO precursors is attributed to their unique chemical structures. In fact, the synthesized LMO precursors all present spinel structures [71–79]. Among these, the LiMn2O4 structure is the most representative one, as shown in Fig. 5. The LiMn2O4 spinel has a cubic crystal structure, which belongs to Fd3m space group. In this structure, the lithium ions occupy the 8a sites of tetrahedrons. The Mn3+ and Mn4+ ions are randomly distributed over the 16d sites of octahedrons with the mole ratio 1:1, and the oxygen anions occupy the 32e sites of face-centered cubes. Thus, the LiMn2O4 spinel can be rep- resented by the formula of (Li)8a[Mn(III)Mn(IV)]16dO4, which belongs to the most typical spinel formulas (AB2O4). The approximately cubic close-packed array of oxide ions incorporates the MnO6 octahedral connected to one another in three dimensions by sharing edges. On the other hand, the LiMn2O4 unit cell could be considered as a complicated cubic structure: 32 oxygen atoms and 16 manganese atoms occupy the half sites of octahedral interstice (16d) while the other half sites (16c) are vacant. Here, 8 lithium atoms occupy 1/8 sites of the tetrahedral interstice (8a). The Li+ could intercalate/de-intercalate in 3D networks of vacant octahedral and octahedral interstices along 8a-16c-8a-16c channel, which is the structural basis of Li+ intercalation/de-intercalation in spinel LiMn2O4 [76]. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 281
Table 1 Classification for some LMO type LISs.
Materials Li+ uptake Regenerations Li+ selectivity Ref. capacity Precursors Ion sieves
1 LiMn2O4 k-MnO2 16.9 mg g (LiCl – Equilibrium distribution (Kd): [43] solution, Li+ Ca2+ >Mg2+ >K+ >Na+ pH = 9.19) 1 + 2+ 2+ + + 1D nano 1D nano k-MnO2 ca.20mgg – Kd:Li Ca >Mg >Na >K [47] 1 LiMn2O4 (10 mmol L LiCl solution, pH = 10.1) 1 + Li4Mn5O12 MnO2 0.4H2O 39.6 mg g The Li adsorption capacity of the The adsorption capacities of other ions were [68] (10 mmol L 1 LiCl spherical ion sieve after 55 cycle’s almost zero except for Mg2+; the ratio of Mg/ solution, adsorption-desorption was Li was reduced to less than 1 from 746 pH = 10.1) 0.4 mmol g 1 1 + 2+ 2+ + + Li4Mn5O12 H4Mn5O12 49.6 mg g – Kd:Li Mg >Ca >K >Na [69] (0.1 mol L 1 LiCl + LiOH solution) 1 + + 2+ 2+ + + Li1.6Mn1.6O4 MnO2 0.5H2O 42.1 mg g The Li adsorption capacity Kd:Li Mg >Na >K >Ca [70] (initial Li+ reduces from 4.08 mmol g 1 to concentration 3.62 mmol g 1 after six times 10 mmol L 1, pH = 10.1) 1 + + 2+ + 2+ + Li1.6Mn1.6O4 MnO2 0.5H2O 40.9 mg g The Li adsorption capacity is Kd:Li Cg >K >Mg >Na [65] (seawater) 35.4 mg g 1, and Li+ desorption rate still reaches 96% at the second cycle operation 1 + Li1.5Mn2O4 H1.5Mn2O4 ca. 15.3 mg g – 400 times higher concentration of Li could be [60] (pH = 8.1) achieved while most of Na+ remains in artificial seawater that contains 10 ppm of Li+ and 10,000 ppm of Na+ by chromatographic separation 1 Li1.51Mn1.63O4 H1.36Li0.07Mn1.65O4 34.07 mg g –– [71] (pH = 12.01) 1 Li1.57Mn1.65O4 H1.39Li0.01Mn1.65O4 ca.37mgg –– [48] (pH = 12)
Fig. 5. Perspective view of (a) a cubane core in the unit cell of spinel LiMn2O4, (b) an extended three-dimensional framework structure of LiMn2O4, and (c) k-MnO2 with void spaces after removal of Li ions. Li, Mn, and O atoms are shown as green, pink, and red, respectively [77]. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.) Reproduced from Ref. [77].
In spinel LiMn2O4, the two metal cations Li and Mn are shown in 1:2 ratio; however, this stoichiometric proportion can be relaxed somewhat under certain conditions. For instance, a sixth of manganese ions in 16d sites can be substituted by lithium ions without changing the whole crystal framework, which is illustrated in Fig. 6. The corresponding LIS of the replaced precursor Li1.33Mn1.67O4 (or Li4Mn5O12) presents theoretically higher lithium capacity than that of k-MnO2, because more lithium ions can be extracted from it or inserted into it. However, the neutron diffraction studies by Ammundsen et al. [75] showed that the lithium reinsertion progress is just for the tetrahedral sites, not for the octahedral sites, indicating the lithium extraction/insertion reaction may be expressed by the following equation:
þ þ ðLiÞ½Li0:33Mn1:67 O4 þ H $ðHÞ½Li0:33Mn1:67 O4 þ Li ð1Þ
Other typical lithium-rich LMO precursor is Li1.6Mn1.6O4 (or Li2Mn2O5), whose corresponding LIS is MnO2 0.5H2O. The MnO2 0.5H2O exhibits the highest theoretical lithium capacity among all the available manganese-type LISs (ca. 72.3 mg g 1). In this compound, the proportion of cations and anions (4:5) is different from that of the typical spinel 282 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Fig. 6. (a) Comparison of two quadrants of cubic spinel lithium manganese oxides and (b) polyhedral structure model and Rietveld refinement of XRD powder pattern for Li1.33Mn1.67O4 (or Li4Mn5O12) recorded under 8.6 GPa. C, cubic spinel phase (space group Fd3m); W, tungsten gasket [78]. Reproduced from Ref. [78]. compounds (3:4), implying extra lithium ions are probably in the interstitial sites of the spinel structure with unequivocal location [74]. Unfortunately, there is not any published structural model for Li1.6Mn1.6O4 so far. Chitrakar et al. [67] proposed three hypothetical models by the preliminary Rietveld analysis: (1) a model with excess Li in the 16c site (Li)8a[Li0.2]16c[Li0.4- Mn1.6]16dO4; (2) a model with oxygen deficiency (Li)8a[Li0.5Mn1.5]16dO3.75, and (3) a model of a hexagonal lattice with cation deficiency (Li0.8h0.2)3b(Mn0.8h0.2)3aO2 (‘‘h” represent the vacant sites in spinels). The simulation results indicated that all the models traced the XRD peaks of the heat-treated sample, but the third model (hexagonal lattice with cation deficiency) clo- sely traced the relative intensity of the XRD peaks. The X-ray absorption spectroscopy of Li1.6Mn1.6O4 samples by Ariza et al. [74] showed that no overall shift of the manganese absorption edge after lithium extraction/reinsertion. Further, the struc- tural arrangement and manganese oxidation state remained unchanged when lithium was extracted and reinserted, which confirmed the ion exchange mechanism for lithium extraction and reinsertion. In summary, there are still some controver- sies on the crystal structure of Li1.6Mn1.6O4. Perhaps, future studies should be focused on this issue to affiliate the develop- ment of LIS with superior lithium uptake properties.
3.1.3. Doping modifications 3 1 3+ Due to the specific electronic configuration of 3d orbit, t2g-eg,Mn can induce Jahn-Teller effect which can cause a serious distortion in the octahedral MnO6 structure. Such distortion will be accompanied by reduction in the stability of LMO and a decrease in the efficiency of Li+ intercalation/de-intercalation process [80–84]. More seriously, a large amount of manganese dissolution in water can lead to water pollution in industrial operations. Therefore, some doping modifications were pro- posed to substitute Mn3+ with other more effective metal ions. In lithium-ion battery field, a great variety of cation (including Co2+,Ni2+,Cr3+,Mg2+,Al3+,Fe3+, and rare earth element ions) substitution has been applied to inhibit capacity fading and improve the electrochemical performance [85–108]. Sim- ilarly, LIS modifications through metal ions doping were proposed to improve the lithium uptake properties in aqueous solu- tions. Chitrakar et al. [109] investigated the impact of LimMgxMn(III)yMn(IV)zO4 (0 6 x 6 0.5) on the dissolution of manganese during acid treatment, the results showed that lithium adsorptive capacity and chemical stability of the proto- nated samples increased with increase in Mg/Mn ratio. Tian et al. [110] synthesized the Mg2+ doped lithium manganese spi- nel by a soft chemical method. The batch experiment suggested that sorption of Li+ showed a high pH and initial concentration dependent profile. Furthermore, the kinetic experiments showed that the adsorption process followed a + pseudo-second-order model. Feng et al. [111,112] studied the Li extraction reactions in both LiMg0.5Mn1.5O4 spinel and + LiZn0.5Mn1.5O4 spinel. They found that the Li extraction and insertion proceeded topotactically by ion-exchange type mech- + anisms. In addition, Liu et al. [113] verified that the Li extraction/insertion reactions with LiAlMnO4 and LiFeMnO4 spinels in the aqueous phase also followed the ion-exchange mechanisms. Ma et al. [114] prepared a series of LiMxMn2 xO4 (M = Ni, Al, Ti; 0 6 x 6 1) spinels and compared their lithium recovery properties in aqueous solutions. The results demonstrated that the LiAl0.5Mn1.5O4 exhibited a relatively high Li extraction ratio and relatively low Mn and Al extraction ratios during acid + treatment, but the LiNi0.5Mn1.5O4 and LiTi0.5Mn1.5O4 spinels did not present satisfactory Li extraction and adsorption prop- erties due to the significant cell contraction or expansion. The Sb-doped LMO spinel was first synthesized by Chitrakar et al.
[115]. The obtained samples Li1.16Sb(V)0.29Mn(III)0.77Mn(IV)0.77O4 presented a well crystallized spinel-type structure, with an exchange capacity reaching 5.6 mmol g 1 for Li+ and the following affinity order K < Na Li. In a follow-up study, Ma et al. [116] prepared series of Li-Sb-Mn composite oxides with different Sb/Mn molar ratios using solid-state reactions. The results indicated that the Sb/Mn molar ratio of Li-Sb-Mn composite oxides is a crucial factor in determining their structure and Li+ X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 283 extraction and adsorption properties. Further, the acid treated spinel Li-Sb-Mn composite oxide with Sb/Mn molar ratio of 0.05 displayed a high Li+ adsorption capacity of 33.23 mg g 1 in lithium solution. The ion-exchange property of iron-doped lithium manganese oxides Li1.33FexMn1.67 xO4 (x = 0.15, 0.30, and 0.40) in Bolivian brine was studied by Chitrakar et al. [31]. The results demonstrated that the adsorbent with Fe/Mn ratio of 0.1 obtained from calcination of the precursor at 450 °C presented the highest lithium extractability with HCl solution. The adsorbent showed lithium uptake of 18.1 mg g 1 from the raw brine at a final pH of 2.0 (initial pH was 6.7), and the uptake increased to 28 mg g 1 at a final pH of 7.2 after an addi- tion of 1 mol L 1 NaOH solution to the raw brine (initial pH was 8.2). From the above description of the doped-LMO spinels, it is clear that doping modifications can improve the lithium adsorptive property effectively. However, comparing with the great progress of ion-doped manganese oxide spinels in elec- trochemistry field, little attention has been paid to LIS refinement. Currently, only a few studies of single metal-doped LISs were explored. The Li+ adsorptive property of multiple ions-doped LISs, including multiple cation-doped, multiple anion- doped and cation-anion-doped LISs in aqueous solution, is still a virgin area for exploration. Previous research of multiple ions-doped LiMn2O4 exhibited high capacity retention, high discharge capacity, and good cycling performance in the lithium ion batteries. This is because the multiple ions-doped LiMn2O4 presented enhanced structural stability [117–126]. In addi- tion, the co-doping has a synergistic effect on the improvement of the cycling life of the materials as cathodes, which can counteract all the factors responsible for the capacity loss of the single ion-doped LiMn2O4 [127–130]. In a similar way, it is convincible that multiple-ion doping is beneficial in improving the regeneration performance and lithium uptake capac- ities of LISs in aqueous solutions. Future research needs to focus on the synergistic effect of different ions on lithium recovery property.
3.2. LTO-type LISs
Currently, two categories of LTO-type LISs exist: H2TiO3 with a layered structure and H4Ti5O12 with a spinel structure. Although the quantity of LTO-type LISs is limited, a great potential exists for developing these green lithium adsorbents for industrial applications benefitting from a minimal risk of water pollution.
3.2.1. Layered H2TiO3 Fig. 7 presents the chemical structure of layered H2TiO3. Layered H2TiO3 is derived from layered Li2TiO3 precursor. The crystal structure of this precursor could be better described as Li[Li1/3Ti2/3]O2; specifically, the structure can be repre- sented as a cubic close packed oxygen atoms, with metal atoms placed in octahedral voids. Li and Ti form two types of layers in the Li2TiO3 structure. One (Li) layer is occupied only by Li atoms, whereas another (LiTi2) layer is occupied by 1/3 of Li and 2/3 of Ti. Meanwhile, lithium ions in the (Li) layers constitute 75% of the total number of lithium in the Li2TiO3 structure, whereas the remaining 25% are located in the (LiTi2) layers [131]. Correspondingly, all the lithium ions are replaced with the protons in the layered H2TiO3 structure. Correspondingly, some authors considered the struc- ture of H2TiO3 was transformed from the layered Li2TiO3 by topotactically replacing lithium ions with protons in early studies. However, Yu et al. [132] discussed the structure of H2TiO3 by analyzing the differences between Li2TiO3 and
Fig. 7. Crystal structure of Li4Ti5O12 (Yellow tetrahedra represent lithium, and green octahedra represent disordered lithium and titanium) [150]. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.) Reproduced from Ref.[150]. 284 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
H2TiO3 and simulating the XRD patterns of HxLi2 xTiO3 (0 6 x 6 2), they pointed out that the structure with a layered double hydroxide type with the 3R1 sequence of oxygen layers is more reasonable for H2TiO3 and the H2TiO3 could be described as a stacking of charge-neutral metal oxyhydroxide slabs [(OH)2OTi2O(OH)2]. In the future studies, more exper- imental verifications are required to confirm the well refined structure.
Since Li2TiO3 was first prepared by Onodera et al. [133] in 1988, numerous studies were reported on its electrochemical applications [134–138] and photocatalytic applications in pollutant degradation [139–141]. Recently, the ion exchange behavior in salt-lake brines was investigated by Chitrakar et al. [131] Although the lithium adsorption rate was relatively slow (requiring one day to attain equilibrium at room temperature), the Li+ capacity could reach 32.6 mg g 1 (4.7 mmol g 1) at a pH of 6.5, which is the highest value among the lithium adsorbents studied to date in acid solution. In addition, H2TiO3 was found capable of efficiently adsorbing lithium ions from the brine containing competitive cations such as Na+,K+,Mg2+ and Ca2+. This is because the exchange sites are so narrow that Na+ (0.102 nm), K+ (0.138 nm) and Ca2+ (0.100 nm), which have ionic radii larger than Li+ (0.074 nm), cannot enter these sites. As to Mg2+, although its ionic radius (0.072 nm) is close to Li+, a high energy is required for the dehydration of magnesium ions to enter the exchange sites because the free energy of D 0 1 D 0 1 hydration for Mg ( Gh = 1980 kJ mol ) is four times of that for Li ( Gh = 475 kJ mol ) [142]. Besides, Shi et al. [143] found that the separation coefficient of Li to Mg reached 102.4 at the 8th adsorption cycle, which constitutes an excellent separation of Li+ and Mg2+ in salt-lake brines. He et al. [144] investigated the maximum lithium adsorptive capacity of 1 H2TiO3 reached 57.8 mg g at the optimum condition by designing an orthogonal test.
3.2.2. Spinel titanium oxides
Spinel titanium oxides present another kind of LTO-type LISs, which are derived from the spinel Li4Ti5O12 precursors. In lithium-ion batteries field, Li4Ti5O12 spinel has been considered one of the most promising anode candidates for the next-generation large-scale lithium-ion batteries to be used in powering electric vehicles (EVs) or hybrid electric vehicles (HEVs). This favorable rating is due to a high potential of around 1.55 V (vs. Li/Li+) during charge and discharge, excellent cycle life due to the negligible volume change, and high thermal stability and safety [145–147]. In the field of lithium recovery from aqueous solutions, a great potential exists for developing the spinel Li4Ti5O12. The corresponding LIS (H4Ti5O12) has a large lithium capacity and a cycling performance better than that of manganese-type LISs, mainly due to the stronger TiAO bond. Additionally, Li4Ti5O12 has similar chemical structure as Li4Mn5O12 (Fig. 8). However, to the best of our knowledge, to date there are very limited reports on the lithium recovery property of spinel
H4Ti5O12. Dong et al. [148] developed three-dimensionally ordered nano Li4Ti5O12 precursor by using PMMA colloidal crystal templates. The corresponding ion sieve exhibited a high selectivity for Li+, 56.81 mg g 1 saturated lithium capac- ity and an excellent acid stability. In the last year, Moazeni et al. [149] reported the H4Ti5O12 with nanotube morphology synthesized via a simple two-step hydrothermal process presented 39.43 mg g 1 lithium capacity in 120 mg L 1 of lithium solution.
Fig. 8. Schematic representation in spinel manganese oxides by the composite mechanism (a) Li+ extraction reactions and (b) Li+ insertion reactions. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 285
4. Lithium intercalation/de-intercalation mechanisms
4.1. In spinel structures
In the past several decades, three main mechanisms were proposed to explain Li+ intercalation/de-intercalation reaction in spinels. They are the redox mechanism, the ion-exchange mechanism, and the composite mechanism.
4.1.1. Redox mechanism + 4.1.1.1. Disproportionation model. In the early study of Hunter [64],Li de-intercalation/intercalation process in LiMn2O4 was + considered as a pair of redox reactions. The author pointed out that the Li de-intercalation in LiMn2O4 was accompanied by the disproportionation of Mn (III) in an acid environment, as seen in Eq. (2). In this process, only the surface Mn (III) is con- verted to Mn (II) and Mn (IV). With the dissolution of Mn (II), the conversion of surface Mn (IV) and interior Mn (III) takes place. Meanwhile, the lithium ions diffuse from the interior to the surface in crystal form and eventually dissolve into solu- tion. This process can be explained by invoking a mechanism somewhat analogous to that proposed by Kozawa and Powers
[151] for the cathodic reduction of c-MnO2 in alkaline electrode. Similarly and as shown in Eq. (3), the Mn (IV) is reduced to + + Mn (III) and O2 with Li reentering the 8a tetrahedron sites in the Li intercalation process, which was verified by Ooi et al. [152]. Li+ de-intercalation:
þ þ 4ðLiÞ½MnðIIIÞMnðIVÞ O4 þ 8H ! 3ð Þ½Mn2ðIVÞ O4 þ 4Li þ 2MnðIIÞþ4H2O ð2Þ Li+ intercalation:
þ 4ð Þ½Mn2ðIVÞ O4 þ 4Li þ 4OH ! 4ðLiÞ½MnðIIIÞMnðIVÞ O4 þ 2H2O þ O2 ð3Þ
(In above equations, the (), [], h represent 8a tetrahedron sites, 16d octahedral sites and cavity sites of spinel structures, respectively.) In conclusion, the driving force of Li+ de-intercalation/intercalation process is the disproportionation of surface Mn (III), with Mn (IV) remains in the crystal structure, and Mn (II) dissolves in the water solution. This theory can explain both (a) the Li+ capacity of spinel-type lithium ion sieves gradually declines in recycling operations and (b) the lattice parameter of the spinel structure turns smaller after Li+ de-intercalation since the radius of the hydrogen atom is smaller than that of the lithium atom.
4.1.1.2. Stage model. Ooi et al. [153] proposed a stage model to describe the process of the topotactic insertion of Li+ in k-
MnO2. They pointed out that the mobility of lithium ions and electrons in k-MnO2 structure is independent. Because of the alternative distribution of Mn (III) and Mn (IV), the migration of electrons can be achieved easily in the spinel structure. Thus, the Li+ insertion process can be divided into the following two steps: first, the insertion of lithium ions into tetrahedral vacant sites of the k-MnO2 framework, accompanied by a reduction of Mn (IV) to Mn (III)
þ k-MnO2 þ xLi þ xe ! LixMnO2 ð4Þ and second, a migration of excess positive charge ( xe) to the surface of k-MnO2 with the oxidation of hydroxide ions in the aqueous phase
xOH ðsÞ!ðx=2ÞH2OðsÞþðx=4ÞO2ðsÞþxe ð5Þ Accordingly, the overall reaction could be expressed by Eq. (6):
k-MnO2 þ xLiOH ! LixMnO2 þðx=2ÞH2O þðx=4ÞO2 ð6Þ
In the above equations, x is the number of moles of inserted Li+ ions per mole of Mn atoms in the solid, (s) represents the + surface of particle. Eq. (6) indicates that the electron deficiency of k-MnO2, which is produced by the insertion of Li ions, can be compensated for by a supply of electrons due to the oxidation of OH to O2. According to this model, the sites for + + OH oxidation are not necessary to be neighboring to the sites for Li insertion. Therefore, the Li insertion in k-MnO2 can proceed without destruction of the spinel structure.
In the case of lithium de-intercalation process of LiMn2O4, the redox mechanism could be verified by a variety of exper- iments. A pH titration study on k-MnO2 by Ooi et al. [154] showed that the lithium ions adsorption consumed an equivalent + amount of LiOH, implying the k-MnO2 with vacant tetrahedral sites was formed after Li extracted from LiMn2O4. The same titration curves were well analyzed on the basis of redox mechanism with Eh (Mn (IV)/Mn (III)) of 1 V (vs. SCE) in their fol- lowing studies [155]. Mosbah et al. [156] obtained spinel related compounds LixMnO2k (0.02 6 x 6 0.50) from spinel type + LiMn2O4 by solid-liquid reaction with various oxidizing agents. Similarly, Tang et al. [157] studied Li extraction from the 3 orthorhombic LiMnO2 by using a solution containing an oxidizing agent (0.5 mol dm (NH4)S2O8). The redox mechanism + could well explain the Li extraction process from the spinel LiMn2O4. 286 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
4.1.2. Ion exchange mechanism In an early study of Shen and Clearfield [158], manganese oxides prepared electrolytically in the temperature range 10– 95 °C presented ion-exchange behavior: the oxides prepared at 25 °C or lower exhibited a high ion exchange capacity equiv- alent to one proton per two Mn atoms, those prepared at higher temperatures exhibited correspondingly lower capacities + and the oxides prepared at 95 °C showed only surface exchange. Furthermore, the Li exchanged solid formed a spinel LiMn2- O4 when heated and the authors considered the HMnO4 could be converted from the spinel LiMn2O4 by treating with dilute acid. Sato et al. [44] suggested that the valence state of the surface Mn ions remained unchanged during the Li+ extraction and insertion by X-ray photoelectron spectroscopic (XPS) analysis, implying an occurrence of the Li+-H+ ion exchange reac- tion at the surface irrespective of the solution pH. Furthermore, the Fourier transform infrared photoacoustic spectra of the Li+ extracted spinel demonstrated the existence of the surface hydroxyl groups, which was considered to be associated with the vacant 8a tetrahedral sites in the spinel structure. The results suggested hydrogen existed in Li+ extracted spinel in the + form of hydroxyl groups, the Li extracted spinels might not be k-MnO2. Zhang et al. [47] demonstrated the MnO2 ion-sieve has two kinds of exchange groups by pH titration experiment. When pH is lower than 4.46, the tetrahedral 8a sites show stronger acidity and they could be easily exchanged, while pH is higher than 4.46, the octahedral 16d sites demonstrate weaker acidity. In addition, Koyanaka et al. [159] verified that the lithium adsorption capacity was proportional to the con- tent of hydrogen ions by analyzing the correlation between the composition and the adsorption capacity of several spinel- type k-MnO2. However, they believed that the reason for selective adsorption of lithium ions in k-MnO2 was not based upon the so-called ion-sieve effect, rather a proper ion exchange reaction occurring only between lithium ions and protons. + Another application of ion exchange is to illustrate the Li extraction process of spinel type Li4Mn5O12 (Li1.33Mn1.67O4), as shown in Eq. (7).
Li4Mn5O12 $ H4Mn5O12 ð7Þ
+ + A neutron diffraction studies by Ammundsen et al. [75] on H exchange of Li1.33Mn1.67O4 and the reverse reaction (Li + exchange of H1.33Mn1.67O4) have clearly shown the ion exchange mechanism during the Li extraction and re-insertion. Kim et al. [160] studied the bonding natures between tetrahedral and octahedral Li sites of the Li1.33Mn1.67O4, showing Li in manganese oxides is highly ionized in both sites but the net charge of Li is greater for tetrahedral sites than octahedral. The results suggested that the tetrahedral sites have higher Li+/H+ exchangeability than the octahedral sites, and are prefer- able for the selective adsorption for Li+ ions. According to this mechanism, lithium ions in the solid can be fully replaced by protons, the Mn (III) and Mn (IV) sites in the crystal remain the same during the Li+-H+ ion exchange, and the protonated solid presents a perfect reversibility and high specificity for Li+ because the spinel structure is retained. As such, this model explains the improved lithium capacity when the pH of the solution increases since the high concentration of OH can promote the Li+ intercalation in LISs. In a recent study of the lithium intercalation/de-intercalation in LISs, it was found that the ion-exchange mechanism is the dominant mechanism; albeit, it is still imperfect. For instance, it cannot be used to explain the observations pertaining to (a) man- ganese loss in solution and (b) Li+ storage capacity gradual decline during recycling.
4.1.3. Composite mechanism Although the redox mechanism and the ion-exchange mechanism could explain many phenomena in lithium recovery from aqueous solutions, both have limitations. Therefore, a composite mechanism was proposed based on the two mechanisms. Ooi et al. [161] investigated the insertion reactions of lithium ions in different kinds of spinel-type manganese and divided the lithium insertion sites into three groups: a redox-type site, a Li+-specific ion-exchange site, and a nonspecific ion-exchange site. The proportion of each site varies depending on the preparation conditions. The oxidation state of man- ganese in the heat-treated precursor plays an important role in the formation of the insertion sites. Feng et al. [162] con- cluded that the lithium ions were extracted/inserted by both redox-type mechanism and ion-exchange type mechanism, and Li+ extraction/insertion sites could be classified into redox-type site and ion-exchange-type site. Generally, Li+ ions were preferentially extracted from/inserted into the ion-exchange sites, unless the number of Mn (III) in the crystal is increased. The reason is that the numbers of redox-type and ion-exchange-type sites are correlated well with the amounts of Mn (III) and Mn (IV), respectively. Furthermore, they found that the proportions of the two types of sites varied depending on prepa- ration conditions (i.e., heat treatment temperature and Li/Mn mole ratio of starting material) of the lithium manganese oxide spinels. The schematic representation in Fig. 9 shows the Li+ extraction/insertion by the composite mechanism in spinel manganese oxides. The composite mechanism can illustrate the lithium adsorption/desorption processes in aqueous solution well, but it is too complicated to verify and experimental verifications are needed to prove its validity.
4.2. In layered structures
In comparison with lithium intercalation/de-intercalation in spinel structures, the process of lithium intercalation/de- intercalation in layered H2TiO3 structure is deemed simple according to Chitrakar et al. [131]. Specifically, due to the narrow + + exchange sites that can only accommodate Li and H , the layered H2TiO3 was found capable of efficiently adsorbing lithium X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 287
Fig. 9. The process of lithium recovery by the layered H2TiO3 (a) Crystal structure of layered Li2TiO3, (b) H2TiO3 and (c) H2TiO3 upon lithium exchange (TiO6 as octahedra, Li as blue ball, H as light pink ball) [131]. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.) Reproduced from Ref. [131]. ions from the brine containing competitive cations such as Na+,K+,Mg2+, and Ca2+ in extremely large excess; meanwhile, the layered H2TiO3 can be regenerated by substituting lithium ions with protons. The whole process is shown in Fig. 9. + In a previous study, Hosogi et al. [163] treated Li2TiO3 with molten AgNO3 at 300 °C and found that Ag cannot be + exchanged with Li at (LiTi2) layers, but it can be exchanged at (Li) layers forming a layered material Ag[Li1/3Ti2/3]O2. The + + results demonstrated that the activity of Li at the (LiTi2) layer was lower than that at the (Li) layer. Thus, the Li in the + (Li) layers is first exchanged to form H[Li1/3Ti2/3]O2, followed by further exchange of Li in the (LiTi2) layers to form the fully + exchanged phase H[H1/3Ti2/3]O2 [164]. Recently, He et al. [144] pointed out that H inserted in the (HTi2) layers of H[H1/3Ti2/3] + O2 cannot be re-exchanged for Li . Furthermore, the practical lithium adsorptive capacity is much lower than the theoretical adsorptive lithium capacity of H2TiO3 even at the optimal conditions. The lithium intercalation/de-intercalation processes could be expressed as follows.
þ þ Li½Li1=3Ti2=3 O2 þ H ! H½Li1=3Ti2=3 O2 þ Li ð8Þ
þ þ H½Li1=3Ti2=3 O2 þ 1=3H ! H½H1=3Ti2=3 O2 þ 1=3Li ð9Þ
þ þ H½H1=3Ti2=3 O2 þ xLi ! H1 xLix½H1=3Ti2=3 O2 þ xH ð10Þ
5. Preparation methods of LIS precursors
As mentioned previously, LISs are transformed from the corresponding precursors by replacing H+ with Li+ in the crystals. Therefore, the essence of obtaining LISs is to synthesize the LIS precursors. In the past decades, several preparation methods were developed, including solid-phase synthesis and soft-chemical synthesis methods. Among them, the solid-phase synthe- sis methods contain the solid-phase combustion method and the microwave combustion method; and the soft-chemical synthesis methods consist of the sol-gel method, the hydrothermal method, the co-precipitation method, the molten-salt synthesis method, the spray-drying method and the microemulsion method.
5.1. Solid-phase synthesis methods
5.1.1. Solid-phase combustion method The solid-phase combustion method is the most widely used approach to prepare LIS precursors. This method begins with mixing lithium salts (usually Li2CO3, LiOH, CH3COOLi, LiNO3) and manganese salts (usually MnO2,Mn2O3,Mn3O4, MnO, c- MnOOH, MnCO3, Mn(CH3COO)2, Mn(NO3)2 or titanium salt (TiO2) evenly with a defined stoichiometric ratio, followed by cal- cination at a suitable temperature for an appropriate period.
Hunter [64] prepared LiMn2O4 by using Li2CO3 and several manganese oxides, such as Mn2O3,Mn3O4, and MnO, as raw materials. The k-MnO2 spinel was then obtained by acidizing the LiMn2O4. Yang et al. [165] synthesized well-crystallized Li1.33Mn1.67O4 spinels at 400 °C by using LiNO3 and different manganese sources, including c-, b-MnO2, hollandite, and H+-form birnessite manganese oxides. The pH titration results illustrated that the lithium adsorbents prepared by the man- ganese sources with relatively smaller pore size (tunnel size) had more uniform acidic sites and less spinel lithiation steric hindrance. Takada et al. [166] prepared a cubic spinel Li4Mn5O12 with lattice parameter a = 8.1616(5) Å by using lithium acetate CH3COOLi and manganese nitrate Mn(NO3)2 as raw materials. Chitrakar et al. [167] synthesized a low-crystalline orthorhombic LiMnO2 (o-LiMnO2) by reacting either c-MnOOH or Mn2O3 with LiOH H2O under steam atmosphere at 120 °C, followed by heating the samples at 400 °C in air to convert to cubic Li1.6Mn1.6O4. The acid-treatment samples 1 1 (H1.6Mn1.6O4) showed that the lithium uptakes could reach 4.75 mmol g (33 mg g ). Shi et al. [143] obtained the Li2TiO3 by a solid-phase reaction between TiO2 and Li2CO3, the acid-modified adsorbent precursor H2TiO3 presented an excellent adsorption ability to Li+—the equilibrium adsorptive capacity reached 39.8 mg g 1 in LiOH solution. Kim et al. [168] prepared 288 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
a spinel-type LiM0.5Mn1.5O4 (M = Ti, Fe) by solid-state reaction. The raw materials were Li2CO3, MnCO3 and TiO2 or FeOOH. The products showed a topotactic extraction of Li+ in aqueous phase, mainly through an ion exchange mechanism. Similarly, the doped lithium manganese oxides such as Sb-doped LiMn2O4 and Fe-doped Li1.33Mn1.67O4 could also be prepared by this method [31,116]. Although the solid-combustion method is simple and easy to apply, it still has some technical challenges; namely, the non-uniform contact of powders and insufficient chemical reactions may lead to the formation of impurities and uneven dis- tributions of particle size. Therefore, some methods were proposed to improve it. Kosova et al. [169] prepared highly dis- persed stoichiometric and nonstoichiometric LixMn2O4 spinels by a mechanochemical method. Here, the lithium and manganese compounds were mixed using a ball mill in raw materials blending stage, followed by roasting the mixtures at certain temperatures. In this process, evenly mixed starting materials by ball milling could lead to mutual diffusion of reactant phases homogeneously after melting, and an increase in the contact area between different reactants could be achieved, thus forming fine particles with high specific surface area. Huang et al. [170] first synthesized LiMn2O4 yBry nanoparticles by a room-temperature solid-state coordination method. In their study, the lithium acetate, manganese acet- ate, and lithium bromide were used as the starting materials, and the citric acid worked as the chelating reagent. After grind- ing them into powders, the polyethylene glycol was applied to mix the powders and make them react to reach the best possible homogeneity. Then the mixture was thermally treated by programmed-temperature calcination, and the products could be finally obtained. Compared to the conventional solid-state method, this method possesses the advantages of simple manipulation and better prospects for commercialization. Further, it could realize evenly mixed raw materials, a low syn- thesis temperature, and small-grain sized powders.
5.1.2. Microwave combustion method An alternative solid phase synthesis method is the microwave combustion method, which was developed in recent years. Due to the in situ mode of energy conversion of microwave heating, the energy will be generated internally within the mate- rial, instead of stemming from external sources of the conventional heating processes. By this method, the nucleation rate and crystal growth rate could be greatly improved, resulting in shorter reaction times and higher preparation efficiency. Fu et al. [171] investigated the differences between the microwave-induced combustion method and the conventional solid- state reaction method in the properties of the LiMn2O4 powder synthesized. They found that the LiMn2O4 powder prepared by microwave-induced combustion method presented excellent capacity and reversibility, which was attributed to the small and uniform particle size produced. Cui et al. [172] verified that the rheological phase-assisted microwave synthesis method was superior in terms of shorter reaction time and regular cubic spinel shapes. Nakayama et al. [173] used the microwave synthesis technique to control the grain size of LiMn2O4. The results demonstrated that the grain size of the samples obtained were fine (ca. 0.2–0.5 lm) compared to traditional solid-state reactions. Chitrakar et al. [174] prepared semicrys- talline orthorhombic LiMnO2 (o-LiMnO2) by microwave irradiation within 30 min at 120 °C. The extractability of the final product Li1.6Mn1.6O4 reached 99%, and the dissolution of manganese during acid treatment was less than 2%.
5.2. Soft-chemical synthesis methods
The soft-chemical synthesis methods are aimed at improving the shortcoming of heterogeneous mixture of solid-phase synthesis methods. Usually, in these methods, atomistic-level mixtures of the starting materials are prepared by dissolving the soluble lithium compounds and manganese or titanium compounds into aqueous solution. The final products could be obtained by roasting the homogenous mixture in specific environment. Herein, we introduce some most commonly used soft-chemical synthesis methods for preparing LIS precursors.
5.2.1. Sol-gel method It is well known that sol-gel method is typically used to prepare nanoparticles with excellent chemical homogeneity, high purity, and uniform phase distribution [175–177]. The schematic presentation of fabricating LIS precursors by sol-gel method is illustrated in Fig. 10. This procedure refers to a process in which the raw materials (soluble lithium source and manganese/titanium source) are dispersed in a sol, followed by agglomeration to form a gel with continuous three-dimensional network. The product LIS precursors could be then obtained by heating the aerogel at a proper condition. In this process, the starting materials are mixed at an atomic scale in a multi-component system, resulting in a perfect crystallization of the LIS precursor.
Fig. 10. A schematic presentation of LIS precursors synthesized by sol-gel method. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 289
As an example, Sun et al. [178] synthesized spinel LiMn2O4 powders by using an aqueous solution of metal acetates containing polyacrylic acid (PAA) as a chelating agent. The results showed that the sol-gel method required much lower calcination temperature and shorter calcination time than conventional solid-sate reactions, mainly because the mixed materials derived from the homogeneous gel presented high sinterability. Moreover, it was found that many physicochem- ical properties of LiMn2O4 such as particle size, specific surface area, and microcrystalline morphology could be controlled by varying the pyrolysis conditions and chelating agent quantity. Fig. 11 showed typical scanning electron micrographs (SEM) images of LiMn2O4 prepared in different conditions by sol-gel method.
Fig. 11. Scanning electron micrographs of the LiMn2O4 powders by sol-gel method in different conditions: (a) 300 °C, x = 1.67, (b) 500 °C, x = 1.67, (c) 750 °C, x = 1.67, (d) 800 °C, x = 1.67, (e) 650 °C, x = 1.67, and (f) 650 °C, x = 0.5 (x indicates the molar ratios of PAA to the total metal ions) [178]. Reproduced from Ref. [178]. 290 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
On the other hand, there are many reported studies on preparing conditions and mechanisms of the sol-gel method.
Seyedahmadian et al. [179] explored the optimum conditions to prepare a pure spinel LiMn2O4 by sol-gel method: They sug- gested that (a) pH should be controlled in the range of 4–6, (b) the mole ratio of citric acid and metal ions should be 1, (c) the lithium nitrate and manganese nitrate should be applied as the source of metal ions, and (d) ethanol should be used as the solvent. This research offers significant instructional details for preparing perfect crystalline spinel LiMn2O4 by the sol-gel method. Kushida and Kuriyama [180] investigated the crystallization mechanism of sol-gel synthesized spinel LiMn2O4. They found that the crystallization into the spinel LiMn2O4 began with combustion of C and H atoms that were released from the Li-Mn gel at 280 °C, and the rearrangement of Li, Mn, and O atoms for the crystallization that occurred subsequently in the process of non-aqueous sol-gel synthesis at 290 °C. The results indicated that the minimum temperature of annealing the Li-
Mn gel precursor should be above 290 °C. Chu et al. [181] prepared spinel Li4Mn5O12 particles via a facile sol-gel method. Their method called for stoichiometric ratio of LiOH 2H2O (AR) and Mn(CH3COO)2 4H2O (AR) to be dissolved in de- ionized water, followed by adding citric acid to the mixed solution at 75 °C in water bath with constant stirring. Subse- quently, the aqueous ammonia was added to adjust the pH to 6.5. Finally, the gel was formed by drying the solution at
100 °C under vacuum environment. Xie et al. [182] synthesized the Ni and Fe dual-doped Li4Mn5O12 spinels by the sol- gel method in a similar way. In their research, the Ni(CH3COO)2 4H2O and Fe(NO3)3 were used as the extra raw materials added in the mixed solution and the final products Li4Mn5 x yNixFeyO12 (0 6 x 6 1, 0 6 y 6 0.75) spinels exhibited perfect consistency with pure Li4Mn5O12 as evidenced by the XRD patterns (Fig. 12). Hao et al. [183] prepared Li2TiO3 using different stoichiometry involving lithium nitrate (LiNO3) and tetrabutyl titanate [Ti(C4H9O)4] as raw materials and glacial acetic acid as the chelating reagent. The average particle size calculated from the transmission electron microscope (TEM) images of the obtained powder was 6–12 nm by the sol-gel method, which is much finer than that prepared by the conventional solid- state method.
5.2.2. Hydrothermal method The hydrothermal method can overcome several disadvantages in preparing LIS precursors, including unevenly mixed raw materials, complicated operation, and high cost on equipment [184]. Moreover, it can create new compounds that could
Fig. 12. Powder XRD patterns of pristine Li4Mn5O12 and doped Li4Mn5 x yNixFeyO12 with various doping amounts [182]. Reproduced from Ref. [182]. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 291 not be synthesized by the other methods because of its unique homogeneous nucleation mechanism [185,186]. For instance,
Li1.6Mn1.6O4, was first synthesized by Chitrakar et al. [67] by heating LiMnO2 that was obtained by the hydrothermal method 1 at 400 °C. The corresponding LIS material MnO2 0.5H2O yielded the maximum lithium uptake value in theory (72.29 mg g ). This material could not be obtained by the conventional solid-state method, because various impurities such as MnO2, LiMn2O3, and Li2O may be obtained when calcinating the raw metal source mixture with an Li/Mn stoichiometric ratio of 1. The other advantage of hydrothermal method is that different morphologies such as nanorods, nanowires, cubic shape, and spheres of LIS precursors could be prepared by changing the hydrothermal conditions. Table 1 lists some different LIS precursors obtained in different hydrothermal environment. It is well known that the nano-adsorbents with specific mor- phologies have a great diverse impact on the adsorption behaviors [187–189]. For LISs, several properties such as lithium uptake capacity, Li+ adsorption equilibrium time, lithium selective performance, and reproducibility may be affected by their different specific morphologies. However, currently there are no published works on the systematic comparison of the dif- ferences in lithium adsorption behaviors among diverse LIS precursors with different morphologies. In general, the smaller the adsorbent particle size is, the larger the lithium uptake capacity can reach and the shorter the Li+ adsorption equilibrium time requires. The reason is that smaller particles usually produce larger specific area, which benefits the contact of Li+ and LIS. Furthermore, the Li+ intercalation/de-intercalation processes in LIS crystal structure would be much easier (see Fig. 13 and Table 2).
5.2.3. Co-precipitation method The co-precipitation method is a widely used soft-chemical method to prepare doped LIS precursors. In a typical process, the required metal sources are dissolved in the aqueous solution, followed by adjusting the pH of the solution to a specific value until the precipitates are formed. The final products are then obtained by calcinating the precipitates at an appropriate temperature for a proper length of time.
Feng et al. [112] prepared the LiMg0.5Mn1.5O4 by using LiOH, Mg(NO3)2, and Mn(NO3)2 as the metal sources. The acid-treat samples presented a high selectivity and a large capacity for Li+ among alkali metal ions. However, the Al-doped and Fe- doped LiMn2O4 samples, LiAlMnO4, and LiFeMnO4 prepared by the co-precipitation method presented worse lithium capac- ities than that of Mg-doped sample [113]. This reduction in capacity may be attributed to the fact that the acid extraction rate of Mg (about 50%) is higher than that of Al and Fe (both of them are less than 10%). To study the effect of tetrahedral substitution products on lithium capacity, Feng et al. [111] synthesized the Zn-doped sample LiZn0.5Mn1.5O4 by the same method. The results showed poor Li-H exchange behavior, because the Zn2+ at tetrahedral sites could not be separated from the spinel structure, which has a significant impact for lithium de-intercalation. Liu et al. [199] synthesized a spinel LiNi0.5- Mn1.5O4 with the modified oxalate co-precipitation method by controlling pH value of the precursor solution and introduc- ing excessive Li source in the precursor. They found that the pH value of the precursor solution could affect the morphology, stoichiometry, and crystallographic structure of the target material. Naghash and Lee [200] prepared uniformly sized, fine particles of spinel LiMn2O4 by the co-precipitation method. In their process, LiNO3 and MnSO4 were dissolved in the solution
Fig. 13. Different morphologies of LIS precursors synthesized by hydrothermal method. 292 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Table 2 Some common LIS precursors obtained in different hydrothermal environment by hydrothermal method.
Precursors Metal sources Hydrothermal Morphology Size Ref. condition
LiMn2O4 LiOH + Mn(NO3)2 70–180 °C, 8–48 h Mono-dispersed nanorods 10–20 nm in diameter, several [47] micrometer in length
LiMn2O4 LiOH H2O+b-MnO2 150 °C, 12 h Agglomerated nanowires ca. 400 nm [69]
LiMn2O4 LiNO3 + MnSO4 H2O 150 °C, 12 h Large polygons 200 500 nm in diameter and length [43]
Li4Mn5O12 LiNO3 + MnSO4 90–180 °C, 12 h One-dimensional nanorods 20–140 nm in diameter, 0.8–4 lmin [46] length
Li4Mn5O12 LiOH + Mn(NO3)2 110 °C,144 h Cubic shape 40–50 nm [190]
Li1.27Mn1.73O4 LiOH + c-MnOOH 170 °C, 24 d – – [55]
Li1.6Mn1.6O4 LiOH + Mn2O3 210 °C, 10 h Agglomerated particles with 0.1–0.3 lm [191] looser surface
Li1.6Mn1.6O4 LiNO3 + Nano-MnO2 120 °C, 12 h Urchin-like spheres – [192]
Li1.6Mn1.6O4 LiOH H2O+Mn 120 °C, 24 h One-dimensional nanowires 50–200 nm in diameter, 0.5–2 lmin [193]
(NO3)2 4H2O length
Li1.6Mn1.6O4 LiOH 130–160 °C, Fine particles <0.2 lm [70]
+ KMnO4 + MnCl2 6–24 h
Li1.6Mn1.6O4 LiOH + c-MnOOH 120 °C, 24 h Needle particles ca. 0.1 lm [67]
Li4Ti5O12 LiOH H2O + Ti(SO4)2 100 °C, 20 h Regular spheres with rough 0.3–1.0 lm [194] surface
Li4Ti5O12 LiOH H2O + anatase 180 °C, 24 h Hierarchically spheres ca.4lm [195]
TiO2
Li4Ti5O12 LiOH H2O+Ti 180 °C, 24 h Spheres with anomalous surface 1–5 lm [196]
(OC4H9)4
Li4Ti5O12 LiOH + Ti[OCH 170 °C, 36 h Flower-like agglomerated 300–500 nm [197]
(CH3)2]4 spheres
Li4Ti5O12 LiOH + Ti[OCH 130 °C, 12 h Sawtooth-like agglomerated 400–600 nm [198]
(CH3)2]4 spheres
first, and then the solution was added drop-wise under a nitrogen atmosphere to a stirred stearic acid in tetramethylammo- nium hydroxide (TMAH). After the pH of the solution was adjusted to 8–10 by using acetic acid, the formed precipitation was aged for 1 h before filtration. Finally, the LiMn2O4 was obtained by calcinating the precipitation in the given environment at a heating rate of 1 °C/min. In conclusion, the co-precipitation method presents the advantages of simple operations, low cost, and easily controlled conditions; however, the added precipitators may result in the high local ion concentrations, which may produce uneven particle size distribution.
5.2.4. Molten-salt synthesis method In early studies, the molten-salt synthesis method was mainly used to prepare ceramic materials [201]. Recently, some studies have been reported on using the molten-salt synthesis method to prepare LIS precursors [202–207]. In this process, the low melting point molten salts, such as LiCl, LiNO3,Li2SO4, and LiOH, do not only act as the reactants but also as the sol- vents [157,203,207]. The obtained samples presented high purity in single-phase powders.
Helan et al. [206] prepared LiMn2O4 powder by a molten-salt method, using an eutectic mixture of LiCl and MnO2 salt at 900 °C. The final products showed good crystallization property, implying this method could be extended to an economic large-scale preparation of LiMn2O4. In their following studies, LiMn2O4 prepared using LiCl-Li2CO3-Mn(CH3COO)2 4H2O and LiSm0.01Mn1.99O4 prepared using LiCl-MnO2-SmCl3 6H2O were investigated [204,205]. Yang et al. [203] tried a variety of lithium-containing fluxes to synthesize the lithium manganese oxide using a needle-like c-MnOOH raw material as the manganese source. Based on the obtained results, their research group explored a series of plate-form spinels with different
Li/Mn ratios, which were obtained by a topotactic reaction of Li-birnessite in LiNO3 flux at 400 °C. A well-ordered high- crystalline spinel LiNi0.5Mn1.5O4 synthesized by a molten-salt method was obtained by Kim et al. [202] using a mixture of LiCl and LiOH salts as the lithium source. The results showed that a pure LiNi0.5Mn1.5O4 phase could be obtained only in the covered condition, indicating the limited amount of oxygen was needed in the synthesis of LiNi0.5Mn1.5O4. Further, it was found that the LiCl salt content could affect the crystal growth of LiNi0.5Mn1.5O4 effectively; i.e., large-particle-size prod- ucts could be obtained if a large amount of LiCl salt was added. Compared to the other methods, the molten-salt synthesis method presents four main advantages: (1) this method pro- vides a way to achieve low calcination temperature in the crystallization process, which is caused by the high diffusion rates of the reaction components in the molten media; (2) the molten medium facilitates better reaction conditions with high ion concentrations of reactants; (3) the operation is simple; and (4) the final products show high purity performance [208]. How- ever, a certain degree of risk in the experiment may limit its application in industry.
5.2.5. Other methods There are many other methods for preparing LIS precursors, but with limited applications. A summary of these follows. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 293
Sinha and Munichandraiah [209] prepared submicron-sized particles of LiMn2O4 spinel via the microemulsion method. The obtained powder exhibited pure crystalline spinel structure by annealing at 900 °C. TEM and HRTEM of the samples in Fig. 14 shows that the particle size was about 200 nm or smaller and the lattice fringed with a separation of 4.759 Å, which corresponds to the (1 1 1) plane of the cubic lattice of LiMn2O4. Kim et al. [210] synthesized Li4Mn5O12 by the solution-combustion method, using LiNO3, Li(CH3COO) 2H2O, and Mn(CH3- COO)2 4H2O as raw materials. The preparation process was outlined as follows. All materials were first totally dissolved in de-ionized water with an Li/Mn stoichiometric ratio of 1/4, followed by drying the solution at 80 °C overnight, and then the dried material was combusted at 300 °C. After combustion, the material was calcined at various temperatures (300, 400, 500
Fig. 14. TEM (a) and (b) and HRTEM (c) of LiMn2O4 powder prepared at 900 °C by the microemulsion method [209]. Reproduced from Ref. [209].
Fig. 15. SEM images of the precursor and spinel LiMn2O4 powders prepared by the spray-drying method: (a) precursor powders and (b) spinel LiMn2O4 powders [214]. Reproduced from Ref. [214]. 294 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
and 600 °C) for different period of times (3, 5, 10 h) in air. The results showed that the Li4Mn5O12 could be obtained above 400 °C. The spray-drying method is another alternative way to prepare LIS precursors, the basic idea of which is to produce of highly dispersed powders from a fluid feed by evaporating the solvent [211–213]. This process is achieved by mixing a heated gas with an atomized (sprayed) fluid of high surface-to-mass ratio droplets, ideally of equal size, within a vessel (dry- ing chamber), causing the solvent to evaporate uniformly and quickly through direct contact. Wu et al. [214] synthesized spinel LiMn2O4 via sintering of a precursor, which was obtained from a solution of the corresponding metal acetate at 600 °C for 15 h by the spray-drying method. The obtained sample particles presented fine, narrowly distributed, and well crystallized property, as shown in Fig. 15.
6. Batteries for lithium recovery from aqueous solution
In the past several decades, the traditional lithium extraction process by LIS has not been widely successful in industry. The main reason is that the lithium adsorption/desorption cycle is inefficient, because it takes a relatively long time to reach equilibrium. Tanaka et al. [215] proposed a valid apparatus to recover Li+ from sea water, i.e., granular lithium adsorbent stored in a bottom valve of sailing ships for allowing sea water to pass through the lithium adsorbent. After 20 days of sailing, the lithium adsorbent reaches saturation. The lithium then is recovered, following an acid treatment to affect desorption. This method has a low recovery efficiency; albeit, it is environmentally friendly and economical. The lithium separation membrane (LSM) is a type of selective lithium membrane with a complex constitution, which permits only Li+ to permeate from one side of the membrane to the other, as shown in Fig. 16. By applying the LSM device, the lithium separation process is more environmentally friendly and more cost-effective than the conventional lithium recovery process. Furthermore, due to the differential Li+ concentration between the two sides of the membrane, the lithium separation process could generate electricity [216]. This technology could be suitable for industrial-scale mass production of lithium from brine/sea water, as well as for the recycling of used lithium ion batteries. However, this device still suffers from a low Li+ recovery efficiency, even worse than the conventional LIS recovery technology. Future work should focus on developing new lithium separation membranes with higher lithium recovery efficiency and better eco- nomic applicability. As a matter of fact, there are two major factors limiting the lithium recovery efficiency: (1) low diffusion rates of Li+ in aqueous solutions and (2) low ion-exchange rates of Li+ in adsorbent particles. The former can be improved by increasing the adsorption temperature, pH of initial solution, and lithium concentration in aqueous solution. The latter, which is the dom- inant limiting factor, could be improved by adding an internal driving force. Recently, an electric driving force to augment the lithium recovery system was proposed [217–219]. The added driving force is intended to accelerate the speed of Li-H exchange on the surface of the LIS. Similar to the rechargeable lithium ion batteries, batteries for recovering lithium ions from aqueous solutions contain both lithium capturing and lithium releasing processes, which is the basis of lithium regeneration.
Fig. 16. World-first dialysis system for Li recovery from seawater by using LSM [216]. Reproduced from Ref. [216]. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 295
6.1. k-MnO2 – Pt battery
Kanoh et al. [220] were the first to propose the electrochemical method to recover Li+ from aqueous solutions. In their experimental device, the k-MnO2-modified platinum electrode, which was prepared by a thermal-decomposition method, was regarded as the working electrode. The Pt-wire electrode and the saturated calomel electrode (SCE) were the counter and reference electrodes, respectively. They investigated the possibility of extracting Li+ from sea water by cyclic voltamme- + try method, briefly described as follows. During Li insertion into MnO2, the Mn (IV) was reduced to Mn (III), and the water was oxidized to oxygen gas. When the sweep reversed, the Mn (III) was oxidized to Mn (IV), and the water was reduced to hydrogen gas. The chemical reactions for this process are expressed as follows. Li+ intercalation:
þ Working electrode : 2k MnO2 þ Li ! LiMn2O4 e ð11Þ þ Counter electrode : 2H2O ! 4H þ O2 þ 4e ð12Þ
Li+ de-intercalation: þ Working electrode : LiMn2O4 ! 2k MnO2 þ Li þ e ð13Þ Counter electrode : 2H2O ! 2OH þ H2 2e ð14Þ
Although this system can recover Li+ from aqueous solutions with a high selective performance and efficiency, it still has a drawback. Specifically, this cycle process is accompanied with the electrolysis of water, which is energy consuming. There- fore, the solution (including pH and ion concentration) always needs to be controlled in a stable environment for recovering Li+ by constant potential electrolysis method (see Fig. 17).
6.2. k-MnO2 – Ag battery
To reduce the energy consumption in lithium recovery process, several LiFePO4 batteries [218,219,221] based on the mix- ing entropy battery and the desalination battery [222,223] were well explored with perfect selectivity between Li+ and Na+ 2+ [224–228]. However, the LiFePO4 batteries suffer from severe interference from the significant presence of coexisting Mg in + brines and sea water. This shortcoming indicates that the LiFePO4 batteries are not fit for recovering Li from salt-lake brines and sea water. Therefore, based on the energy consumption of LiFePO4 batteries, Lee et al. [217] proposed a k-MnO2—Ag rechargeable battery to extract Li+ from solutions containing various metal ions, including Li+,Na+,Mg2+,Al3+,K+, and 2+ + Ca . Due to the excellent selective performance of k-MnO2, only Li could enter the tetrahedral sites of k-MnO2 spinel struc- ture. The schematic of the lithium ion capturing process is shown in Fig. 18 and a complete cycle reactions are shown in Eqs. (15)–(18). Li+ intercalation:
þ Working electrode : 2k MnO2 þ Li ! LiMn2O4 e ð15Þ Counter electrode : Ag þ Cl ! AgCl þ e ð16Þ
Li+ de-intercalation: þ Working electrode : LiMn2O4 ! 2k MnO2 þ Li þ e ð17Þ Counter electrode : AgCl ! Ag þ Cl e ð18Þ
Fig. 17. Schematic of the lithium ion capturing process in source water (Step 1, discharging process) and the lithium ion releasing process in reservoir solution (Step 2, charging process) [217]. Reproduced from Ref. [217]. 296 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Fig. 18. The battery cell voltage (V) vs. charge (Q) at each discharging-charging cycle of two consecutive processes: (a) 1st stage and (b) 2nd stage, the charge-discharge current density of 0.5 mA cm 2); selective lithium ion (c) capturing in source water (1st step, discharging process) and (d) releasing in reservoir solution (2nd step, charging process) (Li+:d,Na+:j,K+:◆,Mg2+:▲,Ca2+:.) [217]. Reproduced from Ref. [217].
A complete cycle of lithium recovery by applying k-MnO2—Ag system consists of two steps. Step 1 describes the discharg- + ing process of the rechargeable battery, which generates energy. In this step, Li is selectively inserted into the k-MnO2 pos- itive electrode from salt water comprised of different metal ions, such as Na+,K+,Ca2+, and Mg2+. This insertion is accompanied by the capture of chloride ions on the negative electrode. Step 2 describes the charging process of the recharge- + able battery, which consumes energy. Here, Li is released from k-MnO2 positive electrode into the reservoir solution under the same experimental conditions of Step 1, accompanied by the release of chloride ions from the negative electrode. The energy consumed per cycle could be given by the circular integral of the voltage with respect to the charge, as shown in the following equation: I W ¼ DEdq ð19Þ c where W is work energy (J) consumed per cycle, DE is the cell voltage (V), and q is the charge (C). The results showed that this system was feasible for converting simulated actual brine into lithium-rich brine consuming only about 1.0 Wh of energy per mole of lithium recovered. Further, it was found that the excellent selective performance for lithium ions of the k-MnO2 electrode is achievable even in the presence of magnesium ions. This device is highly successful in selectively recovering lithium ions from the salt-lake brines and sea water with low energy consumption and a short cycling period. These attributes could make it a good candidate for industrial application. In future studies, efforts should be directed to (a) increase the cyclic ability and (b) find more economical material substitute for Ag in the chloride-ion-capturing electrodes.
6.3. The electrostatic-assisted lithium desorption process
The ion exchange membrane (IEM) technology applied on the batteries has been developed rapidly in recent years [229– 234]. Due to the excellent ion selectivity of the IEMs, the deionization efficiency has been greatly improved when combined with an electrical potential. Using this technology, the lithium ions recovery process is more efficient, more economical, and more environmentally friendly than the conventional acidification method. Fig. 19 illustrates the working principle of the electrostatic-assisted recovery process of lithium ion [235]. In this device, the selective lithium adsorbent electrode and the activated carbon electrode are coated on the current collector, and the anion exchange membrane is attached to the surface of the activated carbon electrode. The operation of this system contains X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 297
Fig. 19. Schematic diagram of the electrostatic-assisted recovery process of lithium ion: (a) graphite sheet as a current collector, (b) selective lithium adsorbent electrode, (c) anion exchange membrane, and (d) activated carbon electrode [235]. Reproduced from Ref. [235].
three steps. First, the lithium ions are adsorbed on the selective lithium adsorbent layer (k-MnO2) when the sea water flows past the electrodes. Second, the other ions that are not adsorbed on the adsorbent layer are removed by flushing them with effluent water. Third, the lithium ions are desorbed from the selective lithium adsorbent layer electrode when a potential is applied between the electrodes. The lithium releasing process reaches equilibrium when the conductivity of effluent water reduces approximately to the conductivity of de-ionized water. This method modified the principle of membrane capacitive deionization, and the experiment results verified that it was feasible to desorb lithium from the lithium adsorbent. Although the desorption efficiency by this method is only approxi- mately 45% compared to that of the conventional desorption process, this method avoids the use of acidic solution to release lithium ions. This means the lithium desorption process is more environmentally friendly and the manganese dissolution ratio can be decreased effectively.
6.4. Outlook of the electrochemical method for recovering lithium
Though the battery technologies for lithium recovery from aqueous solutions grew rapidly, the performance of regener- ation, the lithium capacity, the lithium recovery efficiency, and the economic applicability of these technologies still require further development. Some available methods for improving these electrochemical technologies are summarized in the fol- lowing sections.
6.4.1. Porous LiMn2O4 electrode Similar to the problems of aqueous rechargeable lithium batteries (ARLBs), which include short cycling life of the elec- trode materials, low reversible capacity and inefficient charging performance [236], the batteries for lithium recovery from aqueous solutions technology also suffers from poor lithium regeneration performance, insufficient lithium capacity, and low lithium adsorption/desorption rate. In previous studies, Qu et al. [237] published a paper on using porous LiMn2O4 as cathode material with high power and excellent cycling for aqueous rechargeable lithium batteries. Similarly, this material could be applied to the electrochemical technology for lithium recovery from aqueous solutions.
Porous LiMn2O4, consisting of nano grains was prepared using polystyrene (PS) as the template. In comparison with the LiMn2O4 obtained by traditional solid-state reaction, it is observed that the porous LiMn2O4 exhibited a three-dimensionally ordered porous structure (Fig. 20), which is similar in structure to LiCoO2 and the like materials prepared from template methods [237–240]. The porous LiMn2O4 presented not only an excellent rate behavior and a reversible capacity (108 and 90.0 mA h g 1 at the charge/discharge current densities of 5000 and 10,000 mA g 1, respectively), which is related to the smaller electrochemical resistance and reduced polarization of the porous LiMn2O4, but also an excellent cycling perfor- mance (with no more than a 7% capacity loss after 10,000 full cycles at 9 °C under the 1000 mA g 1 condition) in ARLBs. According to the explanation of Tang et al. [236], there are four reasons for the excellent cycling performance of the porous
LiMn2O4: (1) the high crystallinity of porous LiMn2O4 ensures the structure stability during cycling operations; (2) a short diffusion pathway for lithium ions provided by the porous LiMn2O4 improves the electrode/electrolyte contact, facilitates the de-intercalation/intercalation of lithium ions, enhances the utilization efficiency of the material, and reduces the polar- ization; (3) the porous structure can accommodate the strain/stress during the charge and discharge process from Jahn- Teller effects; and (4) there is no acid, such as HF, which is usually present in the organic electrolytes.
Although the porous LiMn2O4 electrode exhibits an excellent performance in ARLBs, to the best of our knowledge, there is no report on using the porous LiMn2O4 electrode in lithium recovery from aqueous solutions. In this regard, it is very likely that the lithium regeneration performance and the lithium recovery efficiency of the lithium recovery process could be greatly improved by using the porous LiMn2O4 electrode. 298 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Fig. 20. SEM micrographs of the as-prepared (a) solid and (b) porous LiMn2O4, and (c) TEM micrographs of the as-prepared porous LiMn2O4 [237] Reproduced from Ref. [237].
6.4.2. Cost-effective counter electrodes Previous studies indicate that the materials used for counter electrode are usually expensive metals (Ag, Pt), thus render- ing these methods economically unacceptable. More recently, Trocoli et al. [221] proposed the nickel hexacyanoferrate as II the suitable alternative to silver for electrochemical lithium recovery in the FePO4-M(1/n)KNiFe (CN)6 battery. Nickel hexa- II cyanoferrate, an abundant and environmentally friendly material, is derived from the M(1/n)KNiFe (CN)6 (M = Li, Na, K, Mg, Ca; n = 1, 2). Its chemical reactions during lithium recovery process are presented below: Li+ intercalation:
IIð Þ ! IIIð Þ þð = Þ nþ þ ð Þ Mð1=nÞKNiFe CN 6 KNiFe CN 6 1 n M e 20 Li+ de-intercalation: IIIð Þ þð = Þ nþ ! IIð Þ ð Þ KNiFe CN 6 1 n M Mð1=nÞKNiFe CN 6 e 21
As mentioned previously, the lithium selectivity of FePO4/LiFePO4 electrode is worse than the LiMn2O4/k-MnO2 electrode, II implying that the k-MnO2—M(1/n)KNiFe (CN)6 battery may present better lithium selectivity and high lithium regeneration efficiency. Moreover, this battery for lithium recovery from aqueous solutions may have low energy consumption, which is mainly caused by the resistance of the electrolyte and diffusion-limiting currents in lithium recovery process, as compared with the experimental results of the FePO4—Ag battery. In future studies, more cost-effective materials for counter electrode should be explored in LISs battery systems to reduce the industrial cost involved.
6.4.3. Coating modifications In ARLBs research area, coating modifications are commonly used for improving the electrode conductivities, mechanical properties of active materials, and cycling performance. This is because the coating materials can block the direct contact of X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 299
Fig. 21. The mechanism model of the effects of the PPy coating on a-MoO3 nanoplates during the charge and discharge processes [244]. Reproduced from Ref. [244]. the active materials and the electrolyte, which can avoid the dissolution of the active materials [241–245]. Moreover, coating modifications for cathode materials with spinel structures are beneficial in inhibiting their Jahn-Teller effect [92,246], imply- ing it can improve the Li+ regeneration performance to some degree. For instance, the multi-walled carbon nanotubes
(MWCNTs) coated LiMn2O4 can achieve some successes in improving the cycling stability of LiMn2O4 in aqueous solutions. 1 For example, when the discharge capacity of the hybrid LiMn2O4/MWCNTs was maintained at 112.8 mA h g after 1000 1 cycles, the capacity loss was only 3.59% while that of the virginal LiMn2O4 decreased to 50.1 mA h g after 1000 cycles. This performance was attributed to the fact that smaller particle size and the larger surface area of the hybrid material might have reduced the Jahn-Teller effect during the charge/discharge process [247].
Coating is also applicable to anode materials. For instance, MoO3, which is a hot anode material in ARLBs, has only a cou- ple of redox peaks at 0.72/ 0.6 V (vs. SCE) in CV curves. This anode can be changed to two couples of redox peaks at 0.39/ 0.32 V and 0.75/ 0.59 V (vs. SCE) by coating with polypyrrole (PPy), indicating that the nanostructuring improves the electrochemical activity of MoO3. Tang et al. [244] developed a nanocomposite of the PPy-coated MoO3 nanoplates (PPy@MoO3) with perfect cycling performance and excellent rate capability for the ARLBs. In this case, the PPy coating not only prevents or inhibits the dissolution of active Mo ions, but also it can buffer the possible volume changes during the cycling process. Meanwhile, the PPy coating can partially act as a conductive binder to increase the contact between par- 2 ticles; therefore, the particle-to-particle resistance is decreased. In addition, anions such as SO4 exist at the surface of MoO3. + Thus, Li ion transference or moving between the MoO3 host and the aqueous electrolytes could be favored since there will be enough Li+ ions in the PPy coating due to adsorption from the anions. The schematic representation of the mechanism is shown in Fig. 21. In future work to improve the efficiency of LISs battery systems, some coating modifications of ARLBs can be used as ref- erence, because the lithium recovery processes from aqueous solutions based on LISs by the electrochemical method are very similar to the lithium cycles in ARLBs. However, coating may also cause the decrease of the lithium interface diffusion rate between aqueous solution phase and the LIS particle phases, resulting in reducing the lithium intercalation/de-intercalation rates. How to keep high lithium recovery efficiency without reducing the lithium intercalation/de-intercalation rates by coating modifications is a viable research topic that requires further study.
6.4.4. The metal-doped LMO electrodes
The other alternative for increasing the lithium regeneration using LIS batteries can be by replacing the LiMn2O4/k-MnO2 electrode with a metal doped-LMO electrode. As mentioned above, the metal-doped LMO presented excellent lithium regen- eration performance in aqueous solutions [31,109,110]. In current studies, all the working electrodes of the lithium recovery batteries based on LISs chose the LiMn2O4/k-MnO2 electrode as the cathode [217,220,235]. However, serious manganese dis- solution caused by the Jahn-Teller effect of Mn3+ can reduce the lives of LIS-batteries. Using the metal-doped LMO electrode can inhibit this phenomenon effectively because the spinel structures consist of less amount of Mn3+. However, there are no published data currently on applying metal-doped LMO electrodes using LIS batteries for recovering lithium from aqueous solutions. Thus, developing metal-doped LMO electrodes has a realistic potential to improve the lithium regeneration per- formance, and as such it is worthy of further study.
7. Forming of LISs for industrial applications
Usually, LISs are present in powder form, which produce poor liquidity, insufficient permeability, and low recycling effi- ciency in industrial applications. More seriously, powder form of LISs can cause large pressure drop, loss of powder, and high 300 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 energy consumption in column operations. Therefore, some available methods were proposed to improve the form of LISs, such as granulation, foaming, membrane and fiber formation, and magnetization.
7.1. Granulation
The granulation technology is the most commonly used method for forming LISs adsorbent powders. By applying this technology, high mechanical stability and excellent water permeability of LISs could be obtained.
Onodera et al. [248] first prepared a new composite material that was based on k-MnO2 and macroporous silica. The com- posite can adsorb lithium ions selectively and has merit for use in column operations. Xiao et al. [249] prepared a spherical
PVC-MnO2 ion sieve of 2.0–3.5 mm diameter by an anti-solvent method using ultrafine Li4Mn5O12 powder as the precursor, poly (vinyl chloride) as a binder, and N-methyl-2-pyrrolidone as a solvent (Fig. 22). The lithium adsorption behavior of this material has been systematically studied. The findings indicate that the lithium uptake process was controlled by intraparti- 5 cle diffusion mechanism, according to the calculation results of the film mass-transfer coefficient [kf = (1.8–2.5) 10 - 1 10 2 1 ms ] and the pore diffusivity [Dp = (1.6–1.8) 10 m s ] by the corresponding mathematical model. Further, this composite material has high Li+ equilibrium adsorption capacity (3.38 mmol g 1), superior lithium ions selectivity and excel- lent regeneration performance. These characteristics indicate a promising potential for its use in industry when dealing with lithium recovery from salt-lake brines and sea water.
In recent year, Xiao et al. [250] proposed a polyacrylamide (PAM)-MnO2 ion sieve, which was prepared by using the inverse suspension polymerization method. In their study, the lithium adsorbent was Li4Mn5O12, the binder was acrylamide, the cross-linker was N,N0-methylenebisacrylamide (MBA), and the initiator was ammonium persulfate (APS). Interestingly, the lithium adsorption process of PAM-MnO2 was also considered as an intraparticle-controlled diffusion process. This was evidenced by calculating the film mass-transfer coefficient (kf) and pore diffusivity (Dp), which were estimated in the range of (4.0–6.0) 10 5 ms 1 and (2.8–4.5) 10 10 m2 s 1, respectively. Moreover, the experimental results showed that the 1 + PAM-MnO2 ion sieve had a lithium equilibrium adsorption capacity of 2.68 mmol g at 303 K, and the Li loading amount of PAM-MnO2 ion-sieve remained almost constant after 30 cyclic regenerations in a packed column. Hong et al. [251] designed a composite lithium adsorbent based on alpha-alumina bead (AAB). Their design procedure called for immersing the AAB in Li/Mn acetate solution, calcinating the sample followed by acid treatment after drying, and then preparing the final product in the form of spinel-type hydrogen manganese oxide immobilized on AAB. Due to the superior stability of AAB and good lithium selective property, this composite material presented excellent overall per- formance, as characterized by less than 2% Li+ adsorption capacity loss after 14 repeated use and 8.87 mg g 1 Li+ adsorption capacity in real seawater, even if the lithium concentration in seawater is low. On the other hand, their research group also used the chitosan as a binder material to prepare the chitosan Li1.33Mn1.67O4 granule with meso-porous structure [252],
Fig. 22. SEM photomicrographs of spherical PVC-MnO2 ion sieve: (A) whole particle, (B) profile image of the particle, (C) exterior particle surface, and (D) a section of the particle [249]. Reproduced from Ref. [249]. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 301
Fig. 23. Cross-linking of chitosan monomer by sulfuric acid [252]. Reproduced from Ref. [252]. which was caused by the cross-linking reaction of chitosan, as Fig. 23 illustrates. In this process, the sulfuric acid, the most widely used cross linker, ionically crossed-linked the chitosan molecule, leading to the morphology change of chitosan LMO. As shown in Fig. 24, the number of meso-pores increases after sulfuric acid treatment as compared to the original morphol- ogy of chitosan LMO. The increasing number of meso-pores is beneficial to Li+ uptake because the diffusion of seawater inside a granule is closely related to the number pores in the adsorbent granule. In addition, the cross-linking could also enhance the mechanical stability of the granules.
In early studies by Ryu et al. [253], a cylinder-type LMO was prepared employing water glass (Na2SiO3) that could be used as the binder at high temperature. Owing to the off-gas emissions from the reactants during the heat treatment and lithium extraction process, this composite material produced a porous structure, which allowed liquid to flow for the Li-H ion- exchange reaction. Although the cylinder-type LMO showed a lower adsorption capacity (15.06 mg g 1) than the LMO pow- der (27.62 mg g 1), the structure of cylinder-type LMO remained in good condition after lithium recovery process for several times, thus implying a promising regeneration performance could be obtained in industrial applications. In their later studies
[254], they used Na2CO3 as the additive for preparation of the LMO-cylinder. This was motivated by its (a) perfect thermal stability, which suppressed side reactions with the spinel structure during the heat treatment and (b) high solubility in acid solutions, which generated additional internal channels and pores in the structure during the acid treatment. They also 1 reported that the maximum adsorption capacity of the composite adsorbent (19.19 mg g ) obtained when the Na2CO3 concentration was 7 wt%. The improvement was the result of the enhanced surface area and pore volume, which was possibly achieved by the fast diffusion into the LMO structure of the aqueous solution and good contact of Li+ and the solid surface.
Fig. 24. SEM image of chitosan LMO granule: (a) 30 , (c) 160 , chitosan HMO granule (after Li+ extraction); (b) 30 , (d) 160 , chitosan LMO granule before and after Li+ extraction [252]. Reproduced from Ref. [252]. 302 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
7.2. Foaming
In recent years, foam-type lithium adsorbents have been developed. These adsorbent are a kind of composite materials, in which the LIS is the core adsorbent, and the organic templates are the modified materials. Usually, they present excellent mechanical property, good regeneration performance, and superior lithium selectivity in industrial uses. Ma et al. [255] prepared a foam-type lithium adsorbent of spinel manganese oxide (MO) by the polyurethane template method. The results demonstrated that the MO foam was composed of lithium ion sieve of MO and oxygen-containing cross- linked pitch that plays a role of a binder for closely bonding MO together into a homogeneous three-dimensionally interpen- etrating network. The SEM and TEM images of MO foam at different magnifications are presented in Fig. 25. This three- dimensionally interpenetrating network is beneficial for the ions migration in the lithium recovery process and their pene- tration of the working solution. The cavities within the MO foam shown in Fig. 25e and f could be functioned as many mini reaction rooms for the solutions penetrating through the pitch foam and accommodating the solutions. Therefore, the MO located on the surface of the pitch particles is able to adsorb Li+ from the solutions easily, indicating that the lithium uptake process is more efficient. However, the combination of the pitch support skeleton and the MO has the tendency to become loose; albeit, the MO foam can keep the structure and morphology during acid treatment and lithium adsorption process. Therefore, further studies should focus on optimizing its preparation and application methodology to improve the overall performance, especially its mechanical repeatability. Nisola et al. [256] fabricated a macroporous polyvinyl alcohol (PVA) foam composites with high loading of uniformly distributed LISs. The corresponding preparation schematic diagram of this material is shown in Fig. 27. In their work, the hierarchical porosity composed of macro-/mesopores was effectively achieved by surfactant blending combined with cryo-desiccation. Using this procedure, insoluble LIS/PVA foams with high water absorbency and flexibility was obtained by the glutaraldehyde cross-linking process. Compared with the LIS powder, the LIS/PVA foams produced minimal reduc- tions in Li+ capacity and kinetic properties because of its high total porosity and surface area, hydrophilicity of the PVA matrix, and high LIS loading. Further, it was found that the overall performance was closely related to the LIS loading
Fig. 25. SEM and TEM images of MO foam at different magnifications: (a, d) 50 ; (b, e) 100 ; (c, f) 5000 [255]. Reproduced from Ref. [255]. X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 303
Fig. 26. Schematic diagram of the preparation of LIS/PVA foam via surfactant blending combined with cryo-desiccation [256]. Reproduced from Ref. [256]. amount. That is increasing the LIS loading, caused the Li+ selectivity of the LIS/PVA foams to approach closely that of the LIS powder. In this regard, 250 wt% LIS/PVA foam produced the highest adsorption performance but showed low mechanical property. As such, this material could be used practically for Li+ mining from secondary sources, since it does not need extra energy intensive processes like packed-bed and membrane systems (see Fig. 26). Han et al. [257] investigated the lithium recovery process of the millimeter-sized spherical ion-sieve foams (SIFs) with hierarchical pore structure, which was prepared by the mechanical foaming method along with drop-in-oil and agar gela- tion. The results showed that the lithium adsorption capacity in the LiOH solution decreased as the agar content increased during the fabrication process. Probably in this case, the SIFs fabricated with lower agar content possessed more open pores, which in turn increased the contact probability with the hierarchical structure developed in the inner parts of the SIFs. The maximum lithium uptake amount of the SIFs could reach 3.4 mg g 1 in natural seawater. In addition, the lithium adsorption efficiency was over 95%, and the lithium desorption efficiency was maintained at about 86% even after five cyclic operations.
7.3. Membrane formation
Membrane forming is another typical approach for developing composite lithium adsorbents for industrial application. Research in this area is of particular interest to the scientific community, because these membranes have several advantages. These include high mechanical property, high specific surface area, and excellent lithium recovery performance. Conse- quently, with greater development of the lithium recovery industry, more lithium adsorbent-embedded-membranes will be used.
Umeno et al. [258] prepared a membrane-type adsorbent of spinel-type Li1.33Mn1.67O4 by the solvent-exchange method using PVC as a binder. In this study, a parallel-flow cell with suitable spacers around the membranes was designed. As shown in Fig. 27, this system is advantageous for the effective lithium recovery from seawater utilizing natural flow. The properties of precursor membranes with different initial PVC concentrations and binder additive contents were further explored. The results showed that the membrane effective surface area increased with an increase in the PVC content, which agreed with an increasing order of lithium adsorption rate. This indicates that the diffusion in solid phase is the rate-determining step, in analogy to the situation of lithium adsorption on granular adsorbents. However, the lithium adsorption rate correlated only to manganese oxide content when the lithium adsorption rate was normalized per surface area, implying that it decreased with PVC content. Therefore, it can be concluded that the outer surface area plays an important role in controlling the lithium adsorption rate, no matter what the adsorbent form exists, be it granular or membrane.
Based on the above studies, Zhu et al. [259] developed a PVC-H1.6Mn1.6O4 lithium ion-sieve membrane by the solvent- exchange method, followed by the HCl treatment using spinel-type manganese oxide powder (Li1.6Mn1.6O4) as the precursor, PVC as the binder, and N,N-dimethyl acetamide (DMAc) as the solvent. The working principle is shown in Fig. 28. It was found that the thickness of membrane affects the adsorption capacity. They determined that the optimum thickness of the liquid film to prepare the membrane-type adsorbent was 0.3 mm, which accommodates the adsorption capacity, mechanical strength, and economical point of view. Further, the lithium ions could be regenerated effectively without sig- nificant loss in the adsorption capacity, indicating that PVC-H1.6Mn1.6O4 lithium ion-sieve membrane has good mechanical property and excellent lithium recovery performance.
Chung et al. [260] proposed a polymeric membrane reservoir system that contained LIS derived from Li1.33Mn1.67O4 for recovering Li+ from seawater. The polymeric membrane reservoir was prepared from non-woven fabric, polysulfone (PSf) membrane, PSf/non-woven fabric composite membrane, and KimtexÒ. Among these materials (Fig. 29), the composite mem- 304 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Fig. 27. Schematic representation of parallel flow adsorption equipment [258]. Reproduced from Ref. [258].
Fig. 28. Working principle of PVC-H1.6Mn1.6O4 lithium ion-sieve membrane [259]. Reproduced from Ref. [259]. brane based on KimtexÒ showed the best results of mechanical property, Li+ extractability, and Li+ lithium adsorption/des- orption efficiency. It was considered as the most promising system for lithium recovery from seawater. However, the phe- nomenon of severe leakage of the inorganic particles limited its industrial application.
7.4. Fiber formation and magnetization
Compared with previous forming methods, fiber formation and magnetization are the newest forming technologies in industrial applications of LISs. Currently, the research on fiber formation and magnetization is quite limited since it is more challenging.
7.4.1. Fiber formation Due to the unique structural properties of nanofibers such as high specific surface area, favorable morphology, high porosity, and unique mechanical strength [261–265], the composite nanofiber materials can resolve the persisting limita- tions observed in currently available LIS composites. Recently, a composite poly (acrylonitrile) (PAN) nanofiber with
H1.6Mn1.6O4 (HMO) lithium ion-sieve was prepared via electrospinning [266]. As shown in Fig. 30, this nanofiber LIS com- posite produced minimal adsorption capacity reduction, improved handling, recyclability, and longevity of the LISs com- pared with currently available LIS composites. Accordingly, the nanofiber LIS composites like HMO/PAN should be further explored as static membrane filters for different types of Li+ source streams and a potential industrial-scale application.
7.4.2. Magnetization An effective magnetic-separation method for lithium employing magnetic particles was introduced this year [267]. Com- pared with other forming methods, it has the advantages of non-toxicity, low cost, and facile synthesis. The preparation of X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 305
Ò Fig. 29. SEM image of inorganic absorbent and membranes. (a) Li1.33Mn1.67O4, (b) surface image of non-woven fabric, (c) cross-sectional image of Kimtex , (d) surface image of PSf membrane, (e) non-woven fabric side of composite membrane, and (f) cross-sectional image of PSf membrane [260]. Reproduced from Ref. [260].
the magnetically separable and reusable adsorbent, as shown in Fig. 31, was a result of growing magnetite (Fe3O4) on LMO with various lithium precursor contents. Among the composite adsorbents, M-HMO-2.5 (magnetic H1.33Mn1.66O1.67) gave the highest chemical stability and superior Li+ selectivity. Specifically, the Li+ adsorption was above 86% even after 6 adsorption- + 2+ + + + desorption cycles, and the order of the equilibrium distribution coefficients (Kd) was Li >Mg >Na >K . Further, the Li adsorption of M-HMO-2.5 was 1.2 mg g 1 in concentrated seawater in spite of these harsh environmental conditions com- pared with the real briny solution or seawater. The M-LMO-2.5 (magnetic Li1.33Mn1.66O1.67) adsorbent can be easily sepa- rated from a liquid under an external magnetic field within a few minutes. The results demonstrated that the magnetic lithium adsorbents were promising candidates for recovering Li+ from aqueous Li resources. In future work, decreasing the LMO destruction by magnetite growth and testing the adsorption performance in other aqueous lithium solution should be the main research focus.
8. Concluding remarks
LISs represent a new kind of cost-effective green lithium-ion adsorbents. They attract great interest in lithium extraction from aqueous solutions. This is partly attributed to the excellent lithium selectivity, high theoretical lithium uptake capacity, and superior cycle performance. Additionally, the lithium recovery process has excellent safety, reliability, and high effi- ciency compared to other conventional methods. It is expected that the development of LIS technologies to recover lithium resource from salt-lake brines and seawater is a growing trend, especially since land lithium resources are diminishing. Such developments will undoubtedly relieve the market pressure for lithium resources across the world. 306 X. Xu et al. / Progress in Materials Science 84 (2016) 276–313
Fig. 30. (a) SEM images of HMO/PAN prepared with various HMO loadings; (b) mechanical testing of pure PAN and HMO/PAN nanofiber composites; (c) equilibrium Li+ adsorption experiments using HMO powder, pure PAN and HMO/PAN nanofiber (20 wt% loading); and (d) recyclability test of HMO/PAN (60 wt%) [266]. Reproduced from Ref. [266].
Fig. 31. Outline of synthetic route to M-LMO composite adsorbents [267]. Reproduced from Ref. [267].
Though nano structured LISs have been prepared by conventional soft-chemical synthesis methods, they have been mainly confined to Li1.6Mn1.6O4,Li1+xMn2 xO4 (0 6 x 6 0.33), Li4Ti5O12 spinels, and layered Li2TiO3. Nevertheless, the Li- Mn-O phase diagram predicts that LISs with high Li+ capacity derived from precursors with high Li/Mn ratios is possible. This finding implies that developing new methods for synthesizing high-efficiency nano-structured LISs is a promising research direction. Moreover, exploring multiple-element-doped LMO can effectively improve the regeneration and the lithium uptake capacity of LISs. Similarly, it is important to study synergistic effects of different ions on lithium recovery property.
For LTO-type LISs, more attention should be focused on exploring the Li1+xTi2 xO4 (0 6 x 6 0.33) spinels since their chemical X. Xu et al. / Progress in Materials Science 84 (2016) 276–313 307
structures are very similar to the Li1+xMn2 xO4 (0 6 x 6 0.33) spinel. In LISs battery research, four research directions can be undertaken to improve the lithium recovery from aqueous solution: (1) using porous LiMn2O4 working electrode to improve the regeneration performance and lithium recovery efficiency; (2) applying new cost-effective counter electrode to reduce the cost; (3) coating electrodes to mitigate dissolution of LISs; and (4) employing doped LMO electrodes to increase lithium regeneration. In addition, some fundamental research is urgently needed to gain a better understanding of the lithium adsorption/desorption behavior, including adsorption kinetics, isothermal adsorption performance, and mechanisms of lithium intercalation/de-intercalation under an electric driving force. Perhaps, developing new forming technology is one of the biggest challenges encountered currently in transitioning LIS laboratory technologies to industrial-scale applications. Noteworthy, the current available LIS forming technologies reduce not only the lithium capacity but also the lithium uptake efficiency compared to raw LISs. Therefore, future studies should be focused on identifying better adhesives, optimizing lithium adsorbents, and improving preparation conditions for forming composite materials.
Acknowledgements
This project was supported by the National Natural Science Foundation of China (Grant No. 51374016), Beijing University of Chemical Technology (BUCT) Fund for Disciplines Construction and Development (Project No. XK1525); the Fundamental Research Funds for the Central Universities (YS1406) and Wyoming Engineering Initiative in the U.S.
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