Article pubs.acs.org/IECR
Recovery of Lithium Ions from Seawater Using a Continuous Flow Adsorption Column Packed with Granulated Chitosan−Lithium Manganese Oxide Taegong Ryu,†,⊥ Yuvaraj Haldorai,‡,⊥ Arunkumar Rengaraj,§ Junho Shin,† Hye-Jin Hong,† Go-Woon Lee,∥ Young-Kyu Han,‡ Yun Suk Huh,*,§ and Kang-Sup Chung*,† †
Mineral Resources Research Division, Korea Institute of Geoscience and Mineral Resources, Daejeon 305-350, Republic of Korea Department of Energy and Materials Engineering, Dongguk UniversitySeoul, Seoul 04620, Republic of Korea § Department of Biological Engineering, Biohybrid Systems Research Center (BSRC), Inha University, Incheon 402-751, Republic of Korea ∥ Testing and Certification Center, Korea Institute of Energy Research (KIER), Daejeon 305-343, Republic of Korea ‡
ABSTRACT: We demonstrated an efficient recovery of lithium ions (Li+) from seawater using a continuous flow column packed with an adsorbent, chitosan−lithium manganese oxide (LMO). The effects of Li+ concentration, contact time, and recyclability were investigated. The adsorbent showed good removal efficiency of Li+ from seawater. The maximum adsorption capacity was calculated to be 54.65 mg/ g. The adsorption process fit well with the Freundlich isotherm with a correlation coefficient of 0.9924. Kinetics studies showed that the adsorption process was consistent with the pseudo-second-order model. The recyclability was tested after extraction of Li+ from the adsorbent using sulfuric acid. The adsorption capacity decreased slightly after recycling the adsorbent three times. This may be due to the dissolution of Mn2+ and deformation of the chitosan structure. A study of the selectivity of Li+ in seawater showed that the selectivity increased in the order Li+ > Mg2+ > Na+.
1. INTRODUCTION In recent years, there has been increasing interest regarding the recovery of resources from seawater because of the large amounts of dissolved ions, but most rare metals exist at much lower concentrations than other metals in conventional sources. Among the 112 known elements, lithium (Li) is one of the 31 rare metal elements and has a great importance due to its wide range of applications such as lithium ion batteries, portable electronic devices, and the automotive industry.1,2 Currently, “Li is recovered from the salt lakes and mines which contains 14 million tons of Li. Although the amount of Li in those resources is quite insufficient at this point, alternative resources should be found to satisfy the increasing demand in the near future”.3 Seawater is a huge reservoir of Li+ (although the concentration of Li+ is 0.17 mg/L),4,5 and the concentrated seawater obtained from desalination plants or salt farms is a potential source of Li+. Several methods have been investigated for the recovery of Li+ from seawater, such as adsorption,6 solvent extraction,7 and coprecipitation.8 In the above methods, adsorption plays an important role because of the selectivity of metal ions, low cost, and environmentally friendly nature. Several studies have reported the preparation of lithium manganese oxide (LMO), such as LiMn2O4, Li1.33Mn1.67O4, and Li1.6Mn1.6O4, for the recovery of Li+ from seawater.9−13 LMO © 2016 American Chemical Society
provides a three-dimensional spinel structure to facilitate the insertion of Li+ with a reversible reaction in the aqueous media. These inorganic materials display high selectivity to lithium, as well as low toxicity and chemical stability, which is desirable for applications to metal ion recovery from seawater. The lithium adsorbent is formed by the topotactic extraction of lithium from LMO using an acid treatment, in which an ion exchange (H+/ Li+) reaction occurs, maintaining the spinel structure. Therefore, LMO can have high selectivity to lithium when used as an adsorbent. The ion-exchange capacity of the adsorbent depends on the composition of LMO. Wang et al.13 reported that the replacement of the manganese site with lithium in the spinel structure affects the lithium uptake in the adsorption process. In their study, the lithium uptake increased with increasing Li/Mn ratio in the spinel structure. Chitrakar et al.14 also reported that the lithium adsorbent (H1.6Mn1.6O4) derived from Li1.6Mn1.6O4 showed the highest lithium uptake compared to that derived from LiMn2O4 and Li1.33Mn1.67O4. On the other hand, Li1.6Mn1.6O4 cannot be obtained by conventional solid state Received: Revised: Accepted: Published: 7218
April 27, 2016 June 3, 2016 June 14, 2016 June 14, 2016 DOI: 10.1021/acs.iecr.6b01632 Ind. Eng. Chem. Res. 2016, 55, 7218−7225
Article
Industrial & Engineering Chemistry Research reactions using a mixture of lithium and manganese compounds.14,15 Ryu et al.16 prepared a cylinder-type LMO adsorbent using water glass as a binder for the recovery of Li+ from seawater. Recently, Li+ adsorption/desorption properties of lithium ion sieves in aqueous solutions and the recovery of lithium from borogypsum were reported.17 Although many studies have evaluated LMO powders for Li+ recovery by an adsorption process, the powder form of an adsorbent limits its industrial applications. To solve these problems, compositing LMO with a polymer is a suitable approach. Umeno et al.18 produced a polyvinyl chloride (PVC)−LMO based membrane-type adsorbent for Li + recovery from salt lakes. Chung et al.19 prepared a polymeric membrane reservoir system that contained the ion exchange adsorbent, Li1.33Mn1.67O4, for the recovery of Li+ from natural seawater. Ma et al.20 synthesized polyurethane−LMO foam for the recovery of Li+ from aqueous solutions. Han et al.21 fabricated a spherical foam with a hierarchical pore structure using a lithium ion sieve powder and agar for the excellent recovery of Li+ from seawater. Recently, the use of a granulated chitosan−LMO complex for the recovery of Li+ from seawater using a batch process was reported.22 The water permeability and hydrophilicity of chitosan are advantageous when it is applied as a binding material compared to other polymer binders, such as poly(vinyl alcohol) or PVC. To the continuation of our previous work, this study examined the recovery of Li+ from seawater using a continuous flow adsorption column packed with the granulated chitosan− LMO adsorbent. The effects of initial Li+ concentration, contact time, and recyclability of the adsorbent for Li+ recovery were investigated. In addition, the adsorption isotherm and kinetics models were analyzed.
Figure 1. Schematic diagram of the continuous flow column packed with granulated chitosan−LMO adsorbent for Li+ recovery.
double jacketed vessel. The experiments were carried out by varying the initial concentrations of Li+ (30, 60, 120, 180, 240, and 360 mg/L) for different times (5 min−6 days) at room temperature. The samples were collected from the reservoir at different time intervals to analyze the amount of Li+ recovered. 2.4. Characterization. Fourier transform infrared spectroscopy (FTIR) was performed using a Bruker VERTEX 80 V spectrometer, and X-ray diffraction (XRD; RIGAKU, DMAX 2500) was performed using Cu Kα radiation. The morphology of the granulated chitosan−LMO was examined using a HITACHI S-4300 scanning electron microscope (SEM). The amount of Li+ recovered from seawater was analyzed by inductively coupled plasma mass spectrometry (ICP-MS; PerkinElmer ELAN 6100).
3. RESULTS AND DISCUSSION Figure 2 presents the XRD patterns of the LMO and granulated chitosan−LMO. For both samples, the positions and relative intensities of all the XRD peaks corresponding to (111), (311), (222), (400), (331), (511), (440), and (531) planes were in good agreement with those of LMO [JCPDS No. 35-0782]. Chitosan peaks were not observed in the granulated chitosan− LMO sample because of the small amount of chitosan and amorphicity. The above results confirmed that the LMO powder and chitosan had been mixed thoroughly with each other in the granules and their individual characteristics were unaffected by the granulation. The recovery of Li+ from seawater was carried out using a continuous flow adsorption column packed with the granulated chitosan−LMO. Figure 3 shows the FTIR spectra of chitosan− LMO granule before Li+ adsorption, after Li+ adsorption, and after Li+ extraction. The Li+ was extracted from the adsorbent using 0.2 M sulfuric acid. The peaks observed at 505, 630, 1070, and 1427 cm−1 corresponded to the LMO. The chitosan had a band at 3100−3400 cm−1 due to the stretching vibrations of −OH and −NH groups.23 The band at 1650 cm−1 was assigned to the amide I group (C−O stretching along the N−H deformation), 1374 cm−1 was attributed to the COO− group in carboxylic acid salt, and 1157 cm−1 was considered to the special peak of the β(1−4) glucosidic bond in the polysaccharide unit. The intensities of chitosan peaks were weak due to the small amount of chitosan. On the other hand, the intensities of the LMO peaks were shifted after the acid treatment. The peak at 1670 cm−1 was assigned to the NH3 deformation band of chitosan. The decreasing peak intensity (1670 cm−1) may be due to the cross-linking of −NH2 functional groups with the acid.24 In addition, the intensities
2. MATERIALS AND METHODS 2.1. Materials. Chitosan, lithium carbonate (Li2CO3), manganese(II) carbonate (MnCO3), lithium chloride, lactic acid, and sulfuric acid were obtained from Sigma-Aldrich and used as received. 2.2. Preparation of Chitosan−LMO Granule. A solid state synthetic method was adopted to synthesize LMO according to the reported procedure elsewhere.22 In a typical experiment, the starting materials Li2CO3 and MnCO3 with suitable molar ratios of 1.33 (Li) and 1.67 (Mn) were ball milled and heated in a furnace for 4 h at 500 °C to mix rigorously. “The binder material was prepared by dissolving chitosan (6 wt %) in 4% lactic acid at 80 °C. After complete dissolution, a yellowish viscous binder was obtained by cooling the solution to room temperature. The LMO and chitosan binder was then mixed completely using a mechanical mixer with a ratio of 4 (w/v). The chitosan−LMO adsorbent was granulated into cylinder-shaped material (diameter of 0.7 mm) using an extruder”.22 2.3. Preparation of Column for Efficient Recovery of Li+. The experimental setup was constructed for the continuous adsorption of Li+ from seawater using chitosan−LMO granules. The setup consisted of three major parts: (i) reservoir, (ii) pump, and (iii) double jacketed vessel. The double jacketed vessel was 14 cm in height and 2.25 cm in width. Inside the vessel, a 1 cm wide pipe was used to hold the adsorbent of around 1.5 g. Two liters of seawater spiked with various concentrations of Li+ was circulated (flow rate 0.9 L/min) using an external pump (Figure 1). The temperature of the system was controlled at room temperature by flowing water through a 7219
DOI: 10.1021/acs.iecr.6b01632 Ind. Eng. Chem. Res. 2016, 55, 7218−7225
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Figure 2. XRD patterns of (a) LMO and (b) chitosan−LMO granules.
Figure 3. FTIR spectra of chitosan−LMO granules before Li+ adsorption and after Li+ extraction using acid. The inset shows the IR spectrum of adsorbent after Li+ adsorption.
of peaks at 1539 and 902 cm−1 increased due to the ionic interaction between NH3 of chitosan and SO42− of the acid.25 These results clearly showed that chitosan was cross-linked with acid after Li+ extraction. The FTIR spectrum of chitosan− LMO granule after Li+ adsorption is also shown in Figure 3, inset. The result revealed that, when compared to the spectrum of adsorbent before Li+ adsorption, no new peaks were observed indicating the physical adsorption of Li+ onto the adsorbent. Figure 4 represents SEM images of the chitosan− LMO granule before and after Li+ adsorption. After adsorption, one can see the Li+ ions on the chitosan−LMO surface. The adsorption capacity of chitosan−LMO granules as a function of the Li+ concentration was examined by varying the initial concentration of Li+ from 30 to 360 mg/L and keeping the adsorbent concentration constant for different time intervals. Because the Li+ concentration in seawater is too low (0.17 mg/L), LiCl was used to artificially spike the Li+ concentrations to 30, 60, 120, 180, 240, and 360 mg/L. The spiked samples were circulated into a continuous flow setup (as discussed in section 2), and the adsorption process was carried out for 5 min−6 days at room temperature. Samples were collected at different times to evaluate the adsorption capacity, and all adsorption experiments were conducted in duplicate. The recovery percentage of Li+ and the adsorption capacity of Li+ at equilibrium (qe) were calculated using the following equations: % recovery =
qe =
C0 − Ce × 100 C0
(C0 − Ce)V m
Figure 4. SEM images of chitosan−LMO granules before (a, b) and after (c, d) Li+ adsorption.
where C0 and Ce are the initial and final Li+ concentrations (mg/L), respectively, V is the volume of solution (L), and m is the mass of the adsorbent (g). Figure 5 shows the recovery of Li+ versus the contact time for different initial concentrations of Li+. The amount of Li+ adsorbed per unit mass of adsorbent increased with increasing Li + concentration. This might be because, at higher concentrations, the driving force for ion migration between
(1)
(2)
Figure 5. Recovery of Li+ using chitosan−LMO adsorbent. 7220
DOI: 10.1021/acs.iecr.6b01632 Ind. Eng. Chem. Res. 2016, 55, 7218−7225
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the bulk solution to the particle surface due to the weaker driving force.26,27 The breakthrough points for 30, 60, 120, 180, 240, and 360 mg/L Li+ inlet concentrations were observed as 75, 45, 26, 13, 9.7, and 9 min, respectively. Although the breakthrough time was shorter at high concentration, the Li+ uptake was higher due to the availability of binding sites for Li+. Adsorption isotherm models are used widely to describe the adsorption progress and examine the mechanism of adsorption. The adsorption isotherms are basic requirements for the design of an adsorption system. The adsorption equilibrium data provide the information for the capacity of an adsorbent on the amount required to adsorb a unit mass of target ion under the system conditions. The Langmuir and Freundlich models were applied to the experimental data. The Langmuir adsorption isotherm has been applied to many adsorption processes and is based on the assumption that adsorption takes place at specific homogeneous sites on the adsorbent. The linearized Langmuir isotherm is represented by the following equation:28
the aqueous phase and solid phase increases. The recovery of Li+ increased with increasing Li+ concentration from 30 to 360 mg/L. The adsorption of Li+ reached equilibrium at 3 days at all Li+ concentrations tested. After this time, no appreciable increase in the Li+ recovery was observed because there were no active sites available for Li+ adsorption. The recovery of Li+ after 6 days for initial Li+ concentrations of 30, 60, 120, 180, 240, and 360 mg/L were calculated to be 16.04, 22.81, 37.68, 46.76, 58.50, and 62.94 mg/L, respectively. The packed-bed setup with granulated chitosan−LMO adsorbent provided excellent feasibility in operation. During this process, the granules showed good mechanical stability and constantly provided a small back pressure due to the flow of water into the adsorbent material. The breakthrough curve analysis was determined using the experimental data obtained from the continuous adsorption. Figure 6 illustrates the effect of inlet concentration of Li+ on the
Ce C 1 = + e qe KLqm qm
(3)
where Ce is the Li+ concentration at equilibrium (mg/L), qe is the adsorption capacity at equilibrium (mg/g), KL is the Langmuir adsorption constant (L/mg), and qm is the maximum adsorption capacity (mg/g). Figure 7a presents the Langmuir isotherm obtained by plotting Ce as a function of Ce/qe. The qm and KL values obtained from the slope and intercept of the linear plot were 54.65 mg/g and 0.009, respectively, with a correlation coefficient of R2 = 0.9640. The Freundlich isotherm model assumes a heterogeneous adsorption surface with sites that have different energies of adsorption and provides no information on the monolayer adsorption capacity. The linear form of the Freundlich adsorption isotherm is represented by the following equation:29
Figure 6. Effect of initial concentration on breakthrough curve for Li+ adsorption on chitosan−LMO adsorbent.
log qe = log KF +
breakthrough curves at a bed height of 1 cm. It was clear from Figure 6 that as the initial concentration increased from 30 to 360 mg/L the breakthrough point time decreased. Also, the breakthrough curves became steeper and the breakthrough volume decreased because of the lower mass-transfer flux from
1 log ce n
(4)
where KF is a constant related to the adsorption affinity and n is the heterogeneity coefficient. The value of n varies with the heterogeneity of the adsorbent and gives an idea regarding the favorability of the adsorption process.
Figure 7. Adsorption isotherms: (a) Langmuir and (b) Freundlich. 7221
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Figure 8. Kinetic models: (a) pseudo first order, (b) pseudo second order, and (c) intraparticle diffusion plots.
The linear plot of log Ce as a function of log qe (Figure 7b) shows that the adsorption process obeys the Freundlich isotherm with a KF value of 1.93 (mg/g (L/mg)1/n) and n = 1.98 obtained from the intercept and slope of the curve with an R2 = 0.9924. The n value indicates the degree of nonlinearity between the solution concentration and adsorption as follows: if n = 1, the adsorption is linear; if n < 1, it is chemisorption; if n > 1, it is physisorption. Therefore, the adsorption of Li+ occurred as a multilayer physical adsorption process. Among the two isotherms, the Freundlich model was found to represent the equilibrium data with a much better fit and higher R2 value. The mechanism of adsorption and the rate-controlling steps, such as mass transport and chemical reaction processes, can be elucidated by the kinetic data of adsorption. The adsorption kinetics was evaluated using pseudo-first-order, pseudo-secondorder, and Weber intraparticle models. The linear form of the pseudo-first-order equation can be expressed as30 log(qe − qt ) = log qe − (k1t )/2.303
adsorbent and adsorbate as covalent forces. The pseudosecond-order model can be expressed as31 t 1 t = + qt qe k 2qe 2 (6) where qe and qt are the adsorption capacities at equilibrium and at time t, respectively, and k2 (g/mg min) is the rate constant. The values of qe and k2 were calculated from a linear plot of t/qt versus t (Figure 8b) with a slope 1/qe and an intercept of 1/ k2qe2. Table 1 lists the parameters along with the correlation Table 1. Kinetic Equilibrium Parameters of Li+ Recovery Using Granulated Chitosan−LMO Adsorbent pseudo second order
(5)
where qe and qt are the adsorption capacities at equilibrium and at time t, respectively, and k1 (min−1) is the rate constant. The values of k1 and qe were calculated from the slope and intercept of the log(qe − qt) versus t plot (Figure 8a). The low R2 value suggested that the pseudo-first-order model was unsuitable for the adsorption of Li+ onto the chitosan−LMO granule. The pseudo-second-order model is based on the assumption that the rate-limiting step may be chemisorption involving the valence forces by the sharing of electrons between the
concn (mg/L)
qe (mg/g)
K2 (g/mg·h)
R2
30 60 120 180 240 360
7.08 15.44 24.92 31.68 41.04 45.63
0.062733 0.007877 0.002202 0.001665 0.001199 0.001104
0.9429 0.9497 0.8222 0.9211 0.9582 0.9862
coefficient values. The higher correlation coefficients indicated that the pseudo second order was predominant and the adsorption was controlled largely by the chemisorption process. The adsorption mechanism of the adsorbate onto the adsorbent is a multistep process involving the transport of solute ions from the aqueous phase to the surface of solid 7222
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mesopore diffusion and the second represents micropore diffusion. The kid values showed that diffusion was very fast in the first linear part and became stable in the second linear part. An extrapolation of the linear portions of the plots back to the y-axis gave the intercepts, which provide a measure of the boundary layer thickness. The deviation of the straight lines from the origin may be due to differences in the rate of mass transfer in the initial and final stages of adsorption. Furthermore, such deviation of the straight line from the origin indicates that pore diffusion is not the sole ratecontrolling step. Figure 8 c shows the curvature at the initial period, which is generally attributed to the boundary layer diffusion effects.35 The slope of the plot is defined as a rate parameter, which is characteristic of the rate of adsorption in the region where intraparticle diffusion is rate controlling. To examine the selectivity of Li+, the competitive adsorption of Na+ and Mg2+ from seawater was investigated (Figure 9a). The adsorbent recovered Li+ effectively from seawater, even though the Li+ concentration was extremely low compared to those of Na+ and Mg2+ in seawater.4 The metal ions adsorbed on the adsorbent were extracted using 0.2 M sulfuric acid, and the adsorption process was repeated for three cycles. The recovery of Li+, Na+, and Mg2+ from seawater was 96, 3.1, and 3.3%, respectively, in the first cycle. The amount of Li+ recovered decreased to 88 and 85% in the second and third cycles, respectively. No appreciable changes in the recovery of Na+ and Mg2+ were observed in the second (2.7 and 3.7%) and third (3.2 and 4.1%) cycles. The selectivity of Li+ adsorption arises from the LMO spinel structure, which is known for its specificity toward Li+. The crystal sites of LMO are so narrow that metal ions other than Li+ are excluded from the sites because of the large steric effect in both the dehydrated and
particulates followed by the diffusion of solute ions into the interior of the pores, which is a slow process and rate determining.32 Weber and Morris proposed an empirical relationship that if intraparticle diffusion is the rate-controlling factor, the uptake varies with the square root of time. The intraparticle diffusion model is expressed as33 qt = k idt 1/2 + C
(7)
where kid is the intraparticle diffusion constant (mg/g min1/2) and C is a constant (mg/g). The kid and C values can be obtained from a linear plot of qt vs t1/2 and are listed in Table 2. Table 2. Intraparticle Diffusion Model first linear part
second linear part
C0 (mg/L)
kid (mg/g min1/2)
R2
kid (mg/g min1/2)
R2
30 60 120 180 240 360
0.177 0.392 0.129 0.465 0.329 0.286
0.980 0.974 0.958 0.975 0.977 0.963
1.145 2.548 2.699 3.997 4.157 3.872
0.634 0.997 0.932 0.942 0.758 0.821
The C value indicates the thickness of boundary layer: the larger the C value, the greater the boundary layer effect.34 If the plot of qt vs t1/2 gives a straight line, the adsorption process is controlled by intraparticle diffusion only. On the other hand, if the data exhibits multilinear plots, then two or more steps influence the adsorption process. The intraparticle diffusion plot obviously showed the trend of diffusion at different times (Figure 8c). Figure 8c exhibits two straight lines: the first straight portion indicates macropore and
Figure 9. (a) Selectivity, (b) recyclability, and SEM images of (c) chitosan−LMO granules before Li+ adsorption and (d) after three repeated recycles of Li+ extraction using acid. 7223
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Industrial & Engineering Chemistry Research hydrated forms. Because of the larger ionic radius of Na+,36 it suffers large steric hindrance at the entrance of the spinel sites. On the other hand, although Mg2+ has the same ionic radius as Li+, it has an approximately 4 times higher free energy of hydration than that of Li+.37 Therefore, it requires more energy to dehydrate and enter the adsorption sites. The characteristic feature of an adsorbent from a practical application point of view is its lifetime because longer periods of time lead to a significant decrease in the cost of the process. Therefore, it is essential to evaluate the stability and reuse of the adsorbent for the practical implementation. The amount of Li+ recovery using a Li+ concentration of 360 mg/L for 6 days was repeated three times with the same adsorbent, and after each experiment, the Li+ was extracted from the adsorbent using sulfuric acid and recycled. The results indicated that the adsorbent exhibited good activity even after three cycles of reuse. The Li+ adsorption decreased slightly with increasing cycle number due to deformation of the chitosan structure and the dissolution of Mn2+.16 The degradation of the chitosan integrity was confirmed by FTIR spectroscopy (Figure 4). The dissolution of Mn2+ was validated by ICP-MS, and the results are shown in Figure 9b. The dissolution amount of Mn2+ from LMO after the acid treatment was calculated as 7.2, 8.5 and 9.2% for the first, second, and third cycles, respectively. Parts c and d of Figure 9 display SEM images of a chitosan− LMO granule before and after three cycles of reuse. Before the acid treatment, no specific morphology was observed (Figure 9c). The granule showed a globular cluster of chitosan and LMO. After the acid treatment, the apparent physical nature of the granule changed remarkably due to cross-linking. The chitosan binder was deformed like a spider web, including numerous 20−40 μm mesopores (Figure 9d). The mesoporous structure is believed to influence the Li+ uptake because the diffusion of seawater inside a granule is closely related to the number of pores in the adsorbent granule.
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ACKNOWLEDGMENTS
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REFERENCES
This research was supported by the national research project entitled “The Development of Technology for Extraction of Resources Dissolved in Seawater” of the Korea Institute of Geoscience and Mineral Resources (KIGAM), funded by the Ministry of Oceans and Fisheries. We acknowledge Jun Yeong Kim for his kind help during manuscript preparation.
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4. CONCLUSIONS In this study, Li+ was recovered successfully using a continuous flow column packed with chitosan−LMO granules. SEM analysis revealed no changes in the surface morphology of the adsorbent after Li+ recovery. The experimental results indicated that the recovery efficiency initially increased rapidly, and the maximum recovery was reached within 3 days. A further increase in reaction time did not show significant changes in the equilibrium concentration. The maximum uptake of Li+ was calculated to be 54.65 mg/g. The continuous packed bed adsorption process provided excellent feasibility for the recovery of Li+. The adsorption and kinetic studies showed that the process was in good agreement with the Freundlich model and second-order kinetics. The adsorbent exhibited high selectivity for Li+ compared to Na+ and Mg2+, which was attributed to the spinel structure of LMO.
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AUTHOR INFORMATION
Corresponding Authors
*Fax: +82-32-872-4046. E-mail:
[email protected] (Y.S.H.). *Fax: +82-42-868-3645. E-mail:
[email protected] (K.S.C.). Author Contributions ⊥
T.R. and Y.H. contributed equally to this work.
Notes
The authors declare no competing financial interest. 7224
DOI: 10.1021/acs.iecr.6b01632 Ind. Eng. Chem. Res. 2016, 55, 7218−7225
Article
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DOI: 10.1021/acs.iecr.6b01632 Ind. Eng. Chem. Res. 2016, 55, 7218−7225
