Process synthesis: Selective recovery of lithium from lithium ion battery

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Process synthesis: Selective recovery of lithium from lithium ion battery cathode materials Quan Li, Ka Yip Fung, Lingda Xu, Christianto Wibowo, and Ka Ming Ng Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/acs.iecr.8b04899 • Publication Date (Web): 25 Jan 2019 Downloaded from http://pubs.acs.org on January 27, 2019

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Industrial & Engineering Chemistry Research

Process synthesis: Selective recovery of lithium from lithium ion battery cathode materials Quan Li 1, Ka Yip Fung 1, Lingda Xu 1, Christianto Wibowo 2, Ka Ming Ng 1* 1

Department of Chemical and Biological Engineering, The Hong Kong University of Science and Technology, Clear Water Bay, Hong Kong 2

ClearWaterBay Technology, Inc., 671 Brea Canyon Road, Suite 5, Walnut, CA 91789, USA A manuscript submitted to Ind. Eng. Chem. Res. Original submission: October 2018 Keywords: Process synthesis; Selective dissolution; Oxalic acid; Lithium recovery; Anti-solvent precipitation

Abstract A selective dissolution process to recover lithium from cathode materials by oxalic acid was investigated. The chemical reaction responsible for dissolution was identified, and the effects of operating parameters including temperature, acid concentration as well as solid-to-liquid feed ratio on lithium recovery were studied using LiNi0.5Mn0.3Co0.2O2 (NMC-532) as the base case cathode materials, leading to a recovery of 96.3% lithium from cathode materials. However, 2.16% of manganese was also dissolved with lithium. An integrated process based on chemical and antisolvent precipitation was synthesized to separate and recover manganese and lithium with high recovery and high purity from the liquid after dissolution. This process was also shown to work for on other cathode materials including LiNixCoyMn1-x-yO2 of various metal ratios, LiMn2O4 and LiCo0.95Mn0.05O2. *Correspondence concerning this article should be addressed to Ka M. Ng. Tel.: +852 2358 7238 and Fax: +852 2358 0054, E-mail address: [email protected]

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Introduction Lithium ion batteries (LIBs) are widely used as energy storage devices especially in portable electronics and electric vehicles (EVs) due to their high energy density, long cycle life, and light weight. Among the different types of cathode materials used in EVs, LiNixMnyCo1-x-yO2 is widely recognized as the first choice for EVs.1 Driven by the expanding EV market, the LIB demand is projected to reach $77.42 billion by 2024.2 The lithium from mining simply cannot meet the need and the price of lithium carbonate in China escalated from $15,000 to $24,000 per ton in 2017 alone.3 It is clear that lithium has to be recycled for sustainability. At present less than 1% of lithium from spent LIBs was recycled in the world.4 Various pyrometallurgical and hydrometallurgical processes have been developed for recovering metals from spent LIBs. For example, high temperature smelting has been used to produce a slag that contains Li, Al, Mn and an alloy that contains Co, Ni, and Cu. The alloy is further purified by chemical precipitation to remove Cu, followed by solvent extraction to separate Co and Ni. However, lithium is not recovered in this process.5 In the hydrometallurgical processes, lithium can be recovered by first dissolving the metal ions from the cathode current collector using acids such as HCl6-8, HNO39, 10, and H2SO411-14, followed by separating the metals by techniques such as solvent extraction14, and selective chemical precipitation15. For example, Tedjar and Foudraz16 first separate cathode powder from other cell components by crushing, screening, magnetic separation, and flotation. The cathode powder is then dissolved in acid, followed by chemical precipitation to remove Cu. Co is then recovered by oxidation or electrolysis, and lithium is recovered as lithium carbonate by pumping CO2 into the solution. In most of these separation processes, lithium is recovered last because of the high solubility of lithium salts. However, to simplify the recycling process, it is beneficial to selectively recover lithium before other metals


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from the cathode materials, especially for those that do not contain expensive metals such as Co. There are reports on the use of chemical agents to selectively dissolve lithium from cathode materials. For example, sodium fluoride can be used to recover lithium as lithium fluoride at low pH, but the possible generation of HF complicates the process. Sodium persulfate can oxidize Ni, Co, or Mn in LiNixCoyMn1-x-yO2 to insoluble oxides while dissolving lithium in the aqueous solution17. However, lithium is not fully recovered and the process is economically unattractive for commercial applications. Instead of using a mineral acid as a dissolving agent, an organic acid also can deliver complete dissolution through complexation with transition metal ions18-22. Further, Chen et al.23 used tartaric acid to dissolve Li from lithium cobalt oxide with a recovery of 98% of lithium along with 3% of Co in the solution. A better selectivity between Li and Co was achieved by Sun et al.24, who used oxalic acid to achieve a recovery of 98% of lithium, and to completely precipitate out Co as cobalt (II) oxalate as shown in the following reaction: 2LiCoO2(s) + 4H2C2O4(aq) 2CoC2O4(s) + Li2C2O4(aq) + 4H2O(aq) + 2CO2(g)

(1)

Gao et al.25 followed the same approach, and proposed a selective dissolution process to treat the LiNi1/3Mn1/3Co1/3O2 (NMC-111) using formic acid and H2O2, about 40% of transition metals (Ni, Co, and Mn) were dissolved along with lithium. Considering the fact that all the corresponding transition metal (Ni, Mn, and Co) oxalate salts are highly insoluble, oxalic acid would be a better choice than other organic acids for selective lithium recovery from various cathode materials containing Ni, Mn, and Co. Most processes are developed based on trial and error, leading to the possibility to recover the wrong product or excessive usage of chemicals. A deeper understanding of the phase behavior is essential to choose the right flowsheet. The conceptual design of precipitation / crystallization


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processes is governed by the solid-liquid equilibrium (SLE) phase behavior. Techniques such as fractional crystallization26, 27, extractive crystallization28, reactive crystallization29, and drowningout crystallization30 are well-established in the literature. Wibowo and Ng31 developed a unified procedure, which represents basic operations such as chemical addition and solvent removal as movements on the phase diagrams, for synthesizing a separation process based on precipitation. This procedure has been applied to synthesize a separation process for cathode materials such as LiFePO4 and LiCoxMn1−xO232. The same approach is used here to synthesize a process to selectively recover lithium by chemical precipitation. LiNi0.5Mn0.3Co0.2O2 (NMC-532) is used as the base case to illustrate the separation process and to show that Li2CO3 can be recovered in high purity and high yield. SLE phase behavior can be conveniently visualized using a phase diagram, which is usually high-dimensional in nature as the system of interest consists of a number of molecular and ionic species. The representation of such a high-dimensional phase diagram for an ionic system is wellestablished33, 34. The number of independent coordinates (F) required for a complete graphical representation of the isothermal-isobaric phase behavior of electrolytic systems is given by Eq. 2: F=m+n+s–2

(2)

where m is the number of simple cations present in the system, n the number of simple anions, and s the number of non-dissociating molecular species or solvents. Among these independent coordinates, the system contains m – 1 cationic coordinates, n – 1 anionic coordinates, and s solvent coordinates. As it is impossible to visualize the phase diagram in its entirety when its dimensionality is larger than 3, various projections and cuts can be made to reduce its dimensionality. A projection is produced by not explicitly plotting the effect of one or more intensive variables on the phase behavior, whereas a cut is generated by plotting the phase behavior


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at fixed values of one or more intensive variables. Each variable that is fixed or not explicitly considered reduces the dimensionality of the phase diagram by one. This paper aims at synthesizing separation processes to selectively recover lithium from various cathode materials, namely LiNixMnyCo1-x-yO2, LiMn2O4, and LiCo0.95Mn0.05O2. After removing metals that are partially dissolved with lithium by chemical precipitation, lithium can be recovered as lithium carbonate, a common precursor for cathode material synthesis, by chemical precipitation. The phase behavior of different systems was investigated to elucidate the experimental results and rationalize the process design, based on which the best flowsheets can be identified. General approach to synthesize the lithium recovery process Figure 1 depicts a three-step conceptual process for selective lithium dissolution from cathode materials, including acid dissolution (step 1), impurity removal (step 2), and lithium recovery (step 3). The objective of the process is to recover lithium with high purity, and high yield. Organic acids such as citric acid18, succinic acid19 or lactic acid20, are not considered for step 1 as the corresponding transition metal (Co, Ni, Mn) salts are also highly soluble. In step 2, even though most transition metals are expected to stay in the solid phase, some of them would inevitably come along with lithium, giving stream F1. These divalent metal impurities are removed in step 2 by chemical precipitation. Other than manganese, this step also works for other impurities such as Cu, and Fe. Depending on the chemical used in this step, different streams (F3 or F3’) is generated. In step 3, lithium is recovered as Li2CO3 either by chemical precipitation or anti-solvent precipitation at room temperature. This general recovery process was created by performing experiments guided by the established crystallization separation techniques. Design decisions and additional experimental work were performed iteratively. To provide the reasoning behind how decisions were made and how possible


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optimization was identified, relevant phase diagrams are provided as needed. The experimental details are discussed next. Experimental Materials and reagents All cathode materials, including NMC cathode materials with different metal ratios (111, 532 and 811), LiMn2O4, and LiCo0.95Mn0.05O2, were provided by Zhengyuan Material Co. (Guizhou, China). H2C2O4·2H2O, KOH, NaOH, K2CO3, and 100% ethanol were purchased from VWR. LiOH·H2O was purchased from Sigma-Aldrich. All chemicals were used without further purification. Experimental procedure Selective dissolution of lithium by using oxalic acid (step 1) Selective lithium dissolution from cathode materials was carried out in a 100 mL three-neck round-bottom flask with a magnetic stirrer, a heating mantle, and a reflux condenser. The base case dissolution experiment was conducted with 0.25 g of cathode powder and 25 mL of 1 M oxalic acid, giving a solid-to-liquid ratio of 10 g/L. The dissolution process was carried out at 95 °C for 12 hours. The effect of oxalic acid concentration, temperature, and solid-to-liquid ratio on dissolution performance were investigated. Removal of manganese as manganese hydroxide by chemical precipitation (step 2) As part of the manganese was dissolved along with lithium, it had to be removed from the liquid mixture by chemical precipitation. Chemical precipitant (NaOH, KOH, or NH4OH) was added to


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F1 in a 20 mL glass vial, and the mixture was agitated with a magnetic stirrer at room temperature. After equilibration for 4 hours, the solution and the precipitated solids were separated by vacuum filtration. Collected solids were washed with DDI water and dried in a vacuum oven overnight at 70 oC. Recovery of lithium as lithium carbonate by chemical precipitation (step 3) The chemical precipitation process for lithium recovery was conducted in a 100 mL culture bottle with a magnetic stirrer. Chemical precipitant (Na2CO3 or K2CO3) was added to F3. The mixture was placed in a water bath at 80 °C. After equilibration for 4 hours, the liquid solution and the precipitated solids were separated by vacuum filtration. Collected solids were washed with DDI water and dried in a vacuum oven overnight at 70 oC. Recovery of lithium carbonate by anti-solvent precipitation (step 3) Anti-solvent precipitation was conducted in a 100 mL conical flask with a cap to minimize solvent evaporation. Chemical precipitant (K2CO3, NH4OH) was added to F3’. A certain amount of ethanol was then added as anti-solvent. The mixture was stirred for 4 hours to attain equilibrium. The solids collected after filtration were washed with DDI water and dried in a vacuum oven overnight at 70 oC. Analytical methods The pH of the solution was measured with a pH/mV meter (Model Mettler- Toledo AG SevenGo2) and the metal concentrations were determined by inductively coupled plasma optical emission spectrometry (ICP-OES) (Model PerkinElmer Optima 7300 DV). The standard solutions were provided by High-Purity Standards, 99.998%, and diluted in double deionized (DDI) water


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with a resistivity of 18.2 MΩ cm at room temperature. The typical instrumental detection limits are 0.3 ppm for lithium, 0.5 ppm for nickel, 0.1 ppm for manganese, and 0.2 ppm for cobalt. The chemical composition of the solids was confirmed with Fourier transform infrared spectroscopy (FTIR) using a Bruker Vertex 70 Hyperion 100 FTIR spectrometer, and thermal gravimetrical analysis (TGA). The crystal structure was characterized by an X-ray diffraction (XRD) system (Model PW1830 Philips, 2KW, Cu anode, graphite monochromator). Results and discussion The dissolution chemical reaction (step 1) Selective lithium dissolution was carried out to investigate the chemical reaction between oxalic acid and NMC-532. Lithium and manganese were found in the liquid after selective lithium dissolution as measured by ICP-OES. Nickel, manganese, and cobalt were present in the solid residue in a ratio of 5: 2.3: 2. The original particles of cathode material and those after selective dissolution are shown in Figure 2. After selective dissolution, the secondary particles broke into smaller pieces. The original spherical primary particles disappeared and parallelepipedic particles constituted by an aggregation of the primary particles formed. The XRD peaks of the solid residue could not be indexed to a single oxalate, due to the coupling effect of three phases,35 indicating the formation of mixed oxalates. The presence of oxalates was further confirmed by FTIR and TGA. Figure 4 shows the FTIR spectroscopic results of the solid residue. The bands at 1315 and 1361 cm-1 can be assigned to O-C-O symmetric stretching, and the band at 1622 cm-1 is the result of asymmetric stretching. The bands at 827 and 491 cm-1 are related to the O-C-O and MO2 ring bending vibration, respectively. The broad band at 3376 cm-1 is the fingerprint of hydration in the residue.36-39 Figure 5 shows the TGA results obtained for the solid residue in air at a heating rate


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of 10 K/min. No weight loss was observed below 150 oC, implying that the powder was completely dry. From 150 to 230 oC, the weight loss was about 20%, which corresponded to the release of two structural water molecules40. The weight loss matched well the theoretical value of 19.79% when H2O was released from a mixture of NiC2O4·2H2O, MnC2O4·2H2O, and CoC2O4·2H2O in the ratio mentioned above. The second stage weight loss resulted from oxalate decomposition.41-43 Taken together, it could be concluded that the liquid phase after selective lithium dissolution was made up of Li+, Mn2+, HC2O4- and H2C2O4 (stream F1 in Figure 1). The solid residue was a mixture of MnC2O4·2H2O, CoC2O4·2H2O, and NiC2O4·2H2O. Thus, the dissolution chemical reaction follows: 10LiNi0.5Mn0.3Co0.2O2(s) + 25H2C2O4(aq)  2CoC2O4(s) + 5NiC2O4(s) + 3MnC2O4(s) + 10LiHC2O4(aq) + 20H2O(aq) + 10CO2(g)

(3)

The effects of various operating parameters including oxalic acid concentration (C), temperature (T), and solid-to-liquid feed ratio (S/L ratio) on this reaction with 12 hours of dissolution is discussed next. A summary of the results is presented in Table 1. The second row shows the base case: C = 1 M, T = 95 oC, and S/L= 10 g/L. The third to fifth rows show how the wt% of metal dissolved in F1 change in response to the change in a single operating parameter while keeping the remaining two parameters constant. Since only trace amounts of Ni and Co were dissolved in F1, they are neglected and the following discussion is focused on the concentration change of lithium and manganese. Effect of oxalic acid concentration The oxalic acid concentration was varied from 0.75 to 2.0 M while keeping temperature and S/L ratio constant (third row of Table 1). As shown in Table 1, 84.0% of lithium was dissolved from


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the cathode material when the oxalic acid concentration was 0.75 M. It increased to 95.4% for 1 M oxalic acid (base case), and a similar efficiency was obtained for higher acid concentrations. The data accuracy was checked by repeating some of the measurements and deemed to be sufficient for process conceptualization. As shown in Table 1, 20.1% of manganese, which has a relatively high solubility in oxalic acid, was also dissolved at an oxalic acid concentration of 0.75 M. It increased to 32.9% when the acid concentration was increased to 2 M. An oxalic acid concentration of 1 M was selected for further experiments, as it provided a high lithium recovery, but only a moderate amount of manganese was dissolved in F1. Effect of temperature The temperature was varied from 80 to 95 oC while keeping acid concentration and S/L ratio constant (fourth row of Table 1). The experimental results show that the lithium recovery increased from 81.7% to 95.4% when the temperature was increased from 80 oC to 95 oC. Similarly, F1 contained more manganese at a higher temperature. As lithium is the target compound to be recovered, 95 oC was selected for further experiments at the expense of a higher energy cost. Effect of solid-to-liquid feed ratio The S/L ratio was varied from 5 to 35 g/L (fifth row of Table 1). Experimental results show that 95.4% of lithium was dissolved for an S/L ratio of 10 g/L (base case). It dropped to 86.6% for an S/L ratio of 15 g/L, and was further reduced to 79.0% at 35 g/L. As the amount of lithium being dissolved from the cathode material depends on the dissolution time, additional experiments were conducted at two loadings (10 g/L and 35 g/L) for a prolonged period (35 h). Figure 6 shows that 95.4% of lithium was dissolved within 12 hours for an S/L ratio of 10 g/L, whereas 30 hours was needed to achieve a lithium recovery of 96.3% at an S/L ratio of 35 g/L. The S/L ratio was


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gradually increased to 50 g/L and the results are summarized in Table 2 for a dissolution time of 35 hours. The lithium recovery was above 95% at all S/L ratios. However, more Ni and Co were dissolved with Li when the S/L ratio increased. At an S/L ratio of 50 g/L, the concentration of Ni, Mn, and Co were 417.1 ppm, 729.3 ppm, and 239.0 ppm, respectively, and the pH of the solution increased from 1.63 to 8.56. In the solution, H+, HC2O4-, and C2O42- were present as an equilibrium system. C2O42- species are dominant at basic pH44 and form complexes with metal ions. This explains why more metal was dissolved at a higher pH. F1 obtained from 1 M of oxalic acid, at an S/L ratio of 35 g/L, a temperature of 95 oC and a dissolution time of 35 hours was further considered below. Mn removal by chemical precipitation (step 2) As part of the manganese is dissolved with lithium, it has to be removed by chemical precipitation for subsequent lithium recovery. The experiment was performed using 5 mL of F1 solution containing 2341 ppm of Li+, 120.7 ppm of Mn2+, and 3.96 ppm of Co2+. For simplicity, the small amount of Co was neglected in the experiments. KOH (1 M), NaOH (1 M), and NH4OH (5 M) were used for Mn precipitation, and the results are summarized in Table 3. Experimental results show that only 2.28% of Mn was precipitated as Mn(OH)2 when 3.2 mL of 1 M KOH was added (second row). This increased to 100% when 3.4 mL of 1 M KOH was used. Similarly, 2.8 mL of 1 M NaOH and 4 mL of 5 M NH4OH precipitated all Mn (Third and fourth row, respectively). Note that complete Mn recovery was achieved only when the pH of the solution reached around 10.6 after adding the precipitant. The precipitated Mn(OH)2 contained impurities from the precipitant, and its metal contents were analyzed by ICP-OES. The solids contained 0.27 wt% potassium oxalate and 28.3 wt% sodium oxalate when 3.4 mL KOH or 2.8mL NaOH was used as the precipitant, respectively. More oxalate


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was precipitated when NaOH was used, mainly because the solubility of sodium oxalate (34.1 g/L) is much lower at room temperature than that of potassium oxalate (364 g/L). When NH4OH was used as precipitant, the amount of ammonium oxalate that was coprecipitated with Mn(OH)2 cannot be measured by ICP-OES. Giving its low solubility (44.5 g/L), a substantial amount of ammonium oxalate was expected to precipitate with Mn(OH)2 and this option was abandoned. Let us illustrate the above results with our understanding of the phase behavior. Consider using KOH as the chemical precipitant, the overall feed for step 2 contains Li+, H+, K+, Mn2+, OH-, and C2O42-. Complex ion such as HC2O4- is not considered, since it is assumed to fully dissociate into simple ions for plotting purpose. According to Eq. 2, the corresponding phase diagram has four independent coordinates, including three cationic coordinates and one anionic coordinate. There is no solvent coordinate for water as H+ and OH- are both present in the system. To visualize the high-dimensional phase diagram, its dimensionality can be reduced by lumping Li + into H+, as lithium salts have relatively high solubility and would not precipitate in this process. Figure 7a shows a triangular prism illustrating the vertices and co-ordinates of such a three-dimensional phase diagram. Projected components are labeled on the corresponding vertices of the phase diagram with a smaller font size. The saturation surfaces are not included in Figure 7a for simplicity purpose, as we only need the saturation surface of Mn(OH)2 to illustrate the separation process and there is also not enough information to estimate the saturation surfaces of other salts. As K2C2O4 is much more soluble than Mn(OH)2, we roughly know that the compartment for precipitating Mn(OH)2 is much larger than that for K2C2O4. By taking a cut (dashed line in Figure 7a) passing through KOH and Mn(OH)2 vertices, as well as the feed point of the precipitation process after adding KOH, the saturation curve for Mn(OH)2 can be obtained. Note that the other


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vertex of the cut is located at a point on the LiOH/H2O-Li2C2O4/H2C2O4 edge. The cuts passing through the feed points M1, M2, and M3 in Figure 7a correspond to the addition of 3.2, 3.3, and 3.4 mL of 1 M KOH (F2), respectively. Note that F1 in Figure 7a is the feed point before adding KOH, locating on the LiOH/H2O-Li2C2O4/H2C2O4-Mn(OH)2-MnC2O4 surface closest to the reader. The three cuts are superimposed in Figure 7b, along with its expanded version. L1, L2, and L3 in Figure 7b represent the composition of the mother liquor after Mn(OH)2 is precipitated for feed points M1, M2, and M3, respectively. Only 2.28% of Mn was precipitated when 3.2 mL of KOH was added (Table 3), thus L1 is very close to M1 in Figure 7b. As the percentage of Mn precipitated increases with KOH addition, the distance between the feed point (M) and the mother liquor (L) also increases. When NaOH (NH4OH) is used as chemical precipitant, phase diagrams similar to that of Figure 7 can be obtained. Compared to potassium oxalate, sodium (ammonium) oxalate has a lower solubility. Therefore, the compartment for precipitating Mn(OH)2 is relatively smaller when NaOH (NH4OH) is used. When similar cuts are made, they may not only pass through the compartment of Mn(OH)2 but also pass through the compartment of sodium (ammonium) oxalate. As a result, in addition to the single saturation curve for Mn(OH)2 as shown in Figure 7b, a single saturation region for sodium (ammonium) oxalate and a double saturation region for sodium (ammonium) oxalate and Mn(OH)2 may also appear in the cuts. Indeed, our experimental data show that sodium (ammonium) oxalate co-precipitated with Mn(OH)2 when NaOH (NH4OH) was used. Lithium recovery by chemical precipitation (step 3) After removing Mn in step 2, lithium is recovered by chemical precipitation, as illustrated in Figure 8. It is preferable that the precipitants used in step 2 and step 3 possess the same cation. If KOH (NaOH) is used for Mn removal, K2CO3 (Na2CO3) should be used to precipitate Li as Li2CO3.


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The chemical precipitation process has to be conducted at an elevated temperature (90 oC), as Li2CO3 has a lower solubility at a higher temperature.2 Considering all Mn was precipitated by KOH in step 2, the overall feed (F3) to the precipitation process contained Li+, K+, C2O42-, CO32-, and H2O. As the pH of the solution reached 11.36 after Mn removal by using KOH, water could be assumed to be non-dissociating. According to Eq. 2, the corresponding phase diagram has three independent coordinates, including one cationic coordinate, one anionic coordinate, and one solvent (H2O) coordinate. As LiOH and KOH are highly soluble and do not precipitate in the process, they are lumped with H2O in the solvent coordinate. This results in a simplified phase diagram, as shown in Figure 9, in which the solubility surfaces for Li2C2O4, Li2CO3, K2CO3, and K2C2O4 are presented. The compartment for precipitating Li2CO3 is larger than others, as it has a relatively lower solubility. Using K2CO3 as precipitant The stream after Mn removal (F3) lies on the plane of Li2C2O4- K2C2O4-H2O and is above the saturation boundary. When K2CO3 solution (F4) is added, its composition moves along the red line connecting F3 and F4. 2 mL, 4 mL, and 6 mL of 5M K2CO3 were added to a 25 mL synthetic feed of F3, giving an overall feed point to the precipitation process at M4, M5, and M6, respectively, in Figure 9. As the feed points are below the saturation surface of Li2CO3, Li2CO3 precipitates and the corresponding mother liquor compositions are indicated by L4, L5, and L6. Experimental data, as summarized in the second row of Table 4, show that only 44.8% of Li was recovered as solids when 2 mL of 5 M K2CO3 was added. It increased to 76% when 4 mL of precipitant was added, and further increased to 78.5% when 6 mL of precipitant was added. The XRD analysis showed that the recovered solid was Li2CO3 (Figure 10), with a purity of >99.5% as measured by ICP.


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Process alternative to increase product recovery To increase the recovery of Li2CO3, water can be removed before adding K2CO3. This can be illustrated by a cut passing through Li2CO3 and H2O vertices, as well as the feed point of the precipitation process after adding K2CO3 (M6), to yield Figure 11. The other vertex of the cut is located at a point on the K2C2O4/K2CO3 edge. Note that F3 and F4 do not lie on the cut and are projected from the prism surface to the cut. As potassium oxalate is much more soluble than lithium carbonate, the saturation region of Li2CO3 is much larger. When water is evaporated, the overall feed point moves from M6 to M6’ and the composition of the mother liquor moves from L6 to L6’. As the feed point (M6’) still lies in the single saturation region of Li2CO3, pure Li2CO3 can be precipitated and at a higher recovery because the feed point is farther away from the saturation boundary. If even more water is evaporated, the feed point moves to M6’’ and enters the double saturation region. Li2CO3 co-precipitates with K2C2O4 to yield a solid mixture (S6), instead of pure Li2CO3, and a mother liquor with a composition at the double saturation point (L6’’). Using Na2CO3 as precipitant When Na2CO3 is used as the chemical precipitant, similar phase behavior can be obtained. As sodium oxalate has a lower solubility than potassium oxalate, the saturation surfaces move towards the H2O vertex. Therefore, the overall feed point to the precipitation process is more likely to be located inside the double saturation region and co-precipitates sodium oxalate with Li2CO3. Experimental results showed that 45.2% of Li was recovered when 7 mL of 3 M Na2CO3 was added (Third row of Table 4). The recovery slightly increased to 48.1% when 15 mL of Na2CO3 was added. The solids composition was characterized by ICP and it contained 25 wt% Li salt and


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75 wt% sodium salt, indicating that the feed point was indeed located inside the double saturation region. As pure Li2CO3 could not be recovered, Na2CO3 is not recommended for use in this step. Possibility of using (NH4)2CO3 as precipitant Similarly, as the solubility of ammonium oxalate is close to that of sodium oxalate, it is expected that the feed point is located inside the double saturation region of (NH4)2C2O4 and Li2CO3 if ammonium carbonate is used as the precipitant. However, after co-precipitating (NH4)2C2O4 and Li2CO3, (NH4)2C2O4 can be decomposed by heat treatment to yield pure Li2CO3.45, 46 Lithium recovery by anti-solvent precipitation (step 3) Using a mixture of K2CO3 and ethanol Let us focus on the case where KOH was used to remove Mn in step 2. K2CO3 and ethanol were added to the 25 mL synthetic feed solution of F3 as chemical precipitant and anti-solvent, respectively, to recover lithium (Figure 12a), and the results are summarized in the second row of Table 5. By adding 5 mL of 5 M K2CO3 as precipitant and 30 mL of ethanol (an organic-to-aqueous (O/A) volumetric ratio of 1) as anti-solvent, 78.9% of Li was recovered as solids. The recovery increased to 90.1% when 60 mL of ethanol (an O/A volumetric ratio of 2) was added. The precipitated solid was K2C2O4·H2O, as confirmed by the XRD pattern in Figure 13. However, the ICP results show that Li2CO3 co-precipitated with K2C2O4·H2O, at a mole ratio of 6.5 between K2C2O4·H2O and Li2CO3. A dissolution step was then used to dissolve K2C2O4 at room temperature, while keeping Li2CO3 as solids. As Li2CO3 was slightly soluble in water at room temperature, the minimum amount of water to be added was determined to reduce the loss of Li. Experimental results in Table 6 show that 96.5% of potassium oxalate was dissolved when 4 mL


16

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of water was added to 2 g of solids. However, 48.8% of Li2CO3 was also dissolved. Additional Li2CO3 was dissolved when more water was added. Therefore, the low recovery of this process limits the applicability of this process alternative. Using NH4OH in step 2 and ethanol in step 3 If NH4OH were used as the precipitant to remove Mn in step 2, (NH4)2C2O4 also co-precipitates with Li2C2O4. Ethanol was now added as an anti-solvent, as illustrated in Figure 12b. Experimental results are summarized in Table 5. 84.1% of Li was recovered as solids when 37.5 mL of ethanol (an O/A volumetric ratio of 1.5) was added as anti-solvent. The recovery increased to 90.0% when 50 mL of ethanol (an O/A volumetric ratio of 2) was added. The precipitated solids contained (NH4)2C2O4·H2O and Li2C2O4. They were then heated at 700 oC to decompose (NH4)2C2O4·H2O and to convert Li2C2O4 to Li2CO3, as confirmed by the XRD pattern in Figure 14. Similar to previous cases, let us illustrate this promising precipitation process by phase diagrams. The overall feed to the anti-solvent precipitation process contained Li+, NH4+, C2O42-, H2O, and ethanol. Again, the H+ concentration in the solution was negligible as the solution pH reached 10.69 after Mn removal by ammonium hydroxide, and water was assumed to be non-dissociating. According to Eq. 2, a three dimensional phase diagram completely represents the phase behavior, and the phase diagram is shown in Figure 15. As (NH4)2C2O4 and Li2C2O4 are almost insoluble in ethanol, the saturation curves are close to the ethanol vertex. There is an additional compartment for (NH4)2C2O4·H2O, as illustrated by the light green saturation surface in Figure 15. Note that (NH4)2C2O4·H2O is located along the (NH4)2C2O4-H2O edge. The feed point to the anti-solvent precipitation process (F3’) locates at the base of the phase diagram. Adding ethanol moves the composition towards the ethanol vertex and moves inside the crystallization compartments of


17


Page 18 of 48

(NH4)2C2O4·H2O and Li2C2O4. M7, M8, and M9 in Figure 15 show the addition of different amounts of ethanol, equivalent to an O/A volumetric ratio of 1.5, 2, and 2.5, respectively. The tie lines in Figure 15 represent the corresponding solid (S) and mother liquor (L) compositions. Figure 16 illustrates a cut passing through the point representing (NH4)2C2O4·H2O and Li2C2O4 vertices, as well as the overall feed point to the precipitation process M9. The cut clearly illustrates that the overall feed point is located inside the double saturation region, leading to the co-precipitation of (NH4)2C2O4·H2O and Li2C2O4, which was indeed observed in our experiments. Similar cuts can be made for the other feed points M7 and M8. It is desirable to recover Li2C2O4 only in anti-solvent precipitation to reduce the energy required in the heat treatment step. By studying the phase behavior, this might be possible by evaporating the feed before ethanol addition. During evaporation, H2O and NH4OH are removed from the system, moving the feed point (F3’) away from the H2O vertex and possibly into the single saturation region of Li2C2O4. Application of oxalic acid selective lithium dissolution for other cathode materials This selective lithium recovery process was applied to other cathode materials, including NMC of different metal ratios, LiMnO2, and LiCo0.95Mn0.05O2. Selective lithium dissolution by using oxalic acid (step 1) was conducted under the conditions determined for NMC-532, but with an S/L ratio of 10 g/L to reduce the reaction time to 12 hours. The wt% of metals dissolved from cathode materials in F1 are listed in Table 7. More than 90% of lithium could be selectively dissolved for all cathode materials. As only Mn was leached out in these cathode materials, the process developed in this study can be applied.


18

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Conclusions Metal recycling from spent LIBs has received increasing attention as the disposal of these batteries is becoming an urgent environmental problem. A 3-step selective lithium recovery process has been synthesized. In step 1, the conditions under which the maximum amount of lithium dissolved from the cathode material were identified. In step 2, metals that were dissolved along with lithium were removed by chemical precipitation. In step 3, lithium was recovered as lithium carbonate by chemical precipitation or anti-solvent precipitation. Guided by insights from well-established crystallization techniques, the saturation boundaries and compartments for recovering various components were identified by conducting experiments. Process optimization alternatives were identified based on the same insights. Preliminary results indicate that this general separation process work for NMC 111, NMC 811, LiCo0.95Mn0.05O2, and LiMn2O4 as well. Considering the objective of the current paper is to recover lithium in a high yield, and high purity, the process alternative using ammonium hydroxide in step 2, ethanol in step 3 was selected as the best flowsheet. However, in an industrial application, the cost and energy consumption should be considered, and we might come up with another “best” flowsheet. This study is being extended in multiple directions. Cost analysis of the identified process alternatives is essential to identify the best alternative. Additional experiments should be conducted to track the minute amounts of Ni and Co neglected in our experimental work. Though they are not expected to affect the outcome of step 3 due to the co-precipitation of Ni and Co with Mn in step 2, the amount of chemical precipitant consumed would still vary. In addition, life cycle analysis should be performed to evaluate the environment impact. The separation of Ni, Mn, and Co in the solid residue from step 1 should be pursued. It is of commercial interest to study the applicability of this process for non-NMC cathode materials such as LiNi0.8Co0.15Al0.05O2.


19


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Acknowledgments The cathode materials were provided by Zoomwe Zhengyuan Material Co., Dalong Economic Development Zone, Tong Ren City, Guizhou, China.


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References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

Pillot, C. Battery Market for Hybrid, Plug-in & Electric Vehicles, 25th Edition - Avicenne Energy; 2016. Swain, B., Cost effective recovery of lithium from lithium ion battery by reverse osmosis and precipitation: a perspective. J. Chem. Technol. Biotechnol 2018, 93, (2). Jaskula, B. W., Mineral commodity summaries 2018: U.S. Geological Survey. In Survey, U. S. G., Ed. 2018; p 200. Graedel, T., Birat, J-P, BK, Sibley, SF, Sonnemann, G, Buchert, M, Hageluken, C. UNEP (2011) Recycleing rates of metals - A Status Report, A Report of the Working Group on the Global Metal Flows to the International Resource Panel. ; 2011. Cheret, D.; Santen, S. Battery recycling. US7169206 B2 2007. Zhang, P.; Yokoyama, T.; Itabashi, O.; Suzuki, T. M.; Inoue, K., Hydrometallurgical process for recovery of metal values from spent lithium-ion secondary batteries. Hydrometallurgy 1998, 47, (2), 259-271. Castillo, S.; Ansart, F.; Laberty-Robert, C.; Portal, J., Advances in the recovering of spent lithium battery compounds. J. Power Sources 2002, 112, (1), 247-254. Joulie, M.; Laucournet, R.; Billy, E., Hydrometallurgical process for the recovery of high value metals from spent lithium nickel cobalt aluminum oxide based lithium-ion batteries. J. Power Sources 2014, 247, 551-555. Lee, C. K.; Rhee, K.-I., Reductive leaching of cathodic active materials from lithium ion battery wastes. Hydrometallurgy 2003, 68, (1), 5-10. Lee, C. K.; Rhee, K.-I., Preparation of LiCoO 2 from spent lithium-ion batteries. J. Power Sources 2002, 109, (1), 17-21. Zou, H.; Gratz, E.; Apelian, D.; Wang, Y., A novel method to recycle mixed cathode materials for lithium ion batteries. Green Chemistry 2013, 15, (5), 1183-1191. Ferreira, D. A.; Prados, L. M. Z.; Majuste, D.; Mansur, M. B., Hydrometallurgical separation of aluminium, cobalt, copper and lithium from spent Li-ion batteries. J. Power Sources 2009, 187, (1), 238-246. Meshram, P.; Pandey, B. D.; Mankhand, T. R., Recovery of valuable metals from cathodic active material of spent lithium ion batteries: Leaching and kinetic aspects. Waste Manage. 2015, 45, 306-313. Nan, J.; Han, D.; Zuo, X., Recovery of metal values from spent lithium-ion batteries with chemical deposition and solvent extraction. J. Power Sources 2005, 152, 278-284. Georgi-Maschler, T.; Friedrich, B.; Weyhe, R.; Heegn, H.; Rutz, M., Development of a recycling process for Li-ion batteries. J. Power Sources 2012, 207, 173-182. Tedjar, F.; Foudraz, J. C. Method for the mixed recycling of lithium-based anode batteries and cells. US20070196725 A1, 2007. Gupta, R.; Manthiram, A., Chemical extraction of Lithium from layered LiCoO2. J. Solid State Chem. 1996, 121, (2), 483-491. Li, L.; Ge, J.; Wu, F.; Chen, R.; Chen, S.; Wu, B., Recovery of cobalt and lithium from spent lithium ion batteries using organic citric acid as leachant. J. Hazard. Mater. 2010, 176, (1), 288-293. Li, L.; Qu, W.; Zhang, X.; Lu, J.; Chen, R.; Wu, F.; Amine, K., Succinic acid-based leaching system: a sustainable process for recovery of valuable metals from spent Li-ion batteries. J. Power Sources 2015, 282, 544-551.


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20. Li, L.; Fan, E.; Guan, Y.; Zhang, X.; Xue, Q.; Wei, L.; Wu, F.; Chen, R., Sustainable Recovery of Cathode Materials from Spent Lithium-Ion Batteries Using Lactic Acid Leaching System. ACS Sustainable Chem. Eng. 2017, 5, (6), 5224-5233. 21. Nayaka, G. P.; Pai, K. V.; Santhosh, G.; Manjanna, J., Dissolution of cathode active material of spent Li-ion batteries using tartaric acid and ascorbic acid mixture to recover Co. Hydrometallurgy 2016, 161, 54-57. 22. He, L.-P.; Sun, S.-Y.; Mu, Y.-Y.; Song, X.-F.; Yu, J.-G., Recovery of Lithium, Nickel, Cobalt, and Manganese from Spent Lithium-Ion Batteries Using l-Tartaric Acid as a Leachant. ACS Sustainable Chemistry & Engineering 2017, 5, (1), 714-721. 23. Chen, X.; Kang, D.; Cao, L.; Li, J.; Zhou, T.; Ma, H., Separation and recovery of valuable metals from spent lithium ion batteries: Simultaneous recovery of Li and Co in a single step. Separation and Purification Technology 2019, 210, 690-697. 24. Sun, L.; Qiu, K., Organic oxalate as leachant and precipitant for the recovery of valuable metals from spent lithium-ion batteries. Waste Manage. 2012, 32, (8), 1575-1582. 25. Gao, W.; Zhang, X.; Zheng, X.; Lin, X.; Cao, H.; Zhang, Y.; Sun, Z., Lithium Carbonate Recovery from Cathode Scrap of Spent Lithium-Ion Battery: A Closed-Loop Process. Environ. Sci. Technol. 2017, 51, (3), 1662-1669. 26. Berry David, A.; Ng Ka, M., Separation of quaternary conjugate salt systems by fractional crystallization. AIChE J. 1996, 42, (8), 2162-2174. 27. Dye Susan, R.; Ng Ka, M., Fractional crystallization: Design alternatives and tradeoffs. AIChE J. 1995, 41, (11), 2427-2438. 28. Dye Susan, R.; Ng Ka, M., Bypassing eutectics with extractive crystallization: Design alternatives and tradeoffs. AIChE J. 1995, 41, (6), 1456-1470. 29. Berry David, A.; Ng Ka, M., Synthesis of reactive crystallization processes. AIChE J. 1997, 43, (7), 1737-1750. 30. Berry, D. A.; Dye, S. R.; Ng, K. M., Synthesis of drowning-out crystallization-based separations. AIChE J. 1997, 43, (1), 91-103. 31. Wibowo, C.; Ng, K. M., Unified approach for synthesizing crystallization-based separation processes. AIChE J. 2000, 46, (7), 1400-1421. 32. Cai, G.; Fung, K. Y.; Ng, K. M.; Wibowo, C., Process Development for the Recycle of Spent Lithium Ion Batteries by Chemical Precipitation. Ind. Eng. Chem. Res. 2014, 53, (47), 18245-18259. 33. Samant, K. D.; Ng, K. M., Representation of high-dimensional solid-liquid phase diagrams of ionic systems. AIChE J. 2001, 47, (4), 861-879. 34. Wibowo, C.; Ng, K. M., Visualization of high-dimensional phase diagrams of molecular and ionic mixtures. AIChE J. 2002, 48, (5), 991-1000. 35. Zhang, Y.-Z.; Zhao, J.; Xia, J.; Wang, L.; Lai, W.-Y.; Pang, H.; Huang, W., Room temperature synthesis of cobalt-manganese-nickel oxalates micropolyhedrons for highperformance flexible electrochemical energy storage device. Scientific reports 2015, 5, 8536. 36. Edwards, H. G. M.; Hardman, P. H., A vibrational spectroscopic study of cobalt (II) oxalate dihydrate and the dipotassium bisoxalatocobalt (II) complex. J. Mol. Struct. 1992, 273, 7384. 37. Bickley, R. I.; Edwards, H. G. M.; Rose, S. J., A Raman spectroscopic study of nickel(II) oxalate dihydrate, NiC2O4 2H2O and dipotassium bisoxalatonickel(II) hexahydrate, K2Ni(C2O4)26H2O. J. Mol. Struct. 1991, 243, (3), 341-350.


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38. Wang, D.; Belharouak, I.; Zhou, G.; Amine, K., Synthesis of lithium and manganese-rich cathode materials via an oxalate co-precipitation method. J. Electrochem. Soc. 2013, 160, (5), A3108-A3112. 39. Mancilla, N.; Caliva, V.; D'Antonio, M. C.; Gonzalez-Baro, A. C.; Baran, E. J., Vibrational spectroscopic investigation of the hydrates of manganese (II) oxalate. J. Raman Spectrosc. 2009, 40, (8), 915-920. 40. Donkova, B.; Mehandjiev, D., Mechanism of decomposition of manganese(II) oxalate dihydrate and manganese(II) oxalate trihydrate. Thermochim. Acta 2004, 421, (1), 141-149. 41. Gao, X.; Dollimore, D., The thermal decomposition of oxalates: Part 26. A kinetic study of the thermal decomposition of manganese(II) oxalate dihydrate. Thermochim. Acta 1993, 215, 47-63. 42. Allen, J. A.; Scaife, D. E., The Thermal Decomposition of Nickel Oxalate. J. Phys. Chem. 1954, 58, (8), 667-671. 43. Mansour, S. A. A., Spectrothermal studies on the decomposition course of cobalt oxysalts Part III. Cobalt oxalate dihydrate. Mater. Chem. Phys. 1994, 36, (3), 324-331. 44. Puigdomenech, I. Chemical equilibrium diagrams. https://www.kth.se/che/medusa/downloads-1.386254 (22 Feb. 2018), 45. Papazian, H. A.; Pizzolato, P. J.; Patrick, J. A., Thermal Decomposition of Oxalates of Ammonium and Potassium. J. Am. Ceram. Soc. 1971, 54, (5), 250-254. 46. Dollimore, D.; Tinsley, D., The thermal decomposition of oxalates. Part XII. The thermal decomposition of lithium oxalate. J. Chem. Soc. A 1971, 3043-3047.


23

Industrial & Engineering Chemistry Research 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47

STEP 1 Oxalic Acid

Cathode materials

Selective Dissolution

Solid residue (Ni, Co, Mn)

Li, Mn

(F1)

Page 24 of 48

STEP 2

STEP 3

Precipitant (F2)

Precipitant (F4)

Impurity Removal

Li

(F3/F5)

Mn(OH)2

Lithium Recovery

Mother Liquor

Li2CO3

Figure 1. Generic flowsheet of selective lithium recovery from cathode material.


24

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Figure 2. SEM photos of cathode materials (a, b) low and high magnification of cathode powder, (c, d) low and high magnification of solid residue after selective dissolution.


25


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Figure 3. XRD patterns of the solid residue in the base case.


26

491

827

1361 1315 1622

3376

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Transmission

Page 27 of 48

4000 3500 3000 2500 2000 1500 1000 500 -1

Wavenumber (cm ) Figure 4. FTIR result of the solid residue.


27


100 90

Weight %

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 28 of 48

20 % weight loss

80 70 60 50 40

0

100 200 300 400 500 600 700 800 900

Temperature (C) Figure 5. TGA profile of the solid residue.


28

Page 29 of 48

100 95

Lithium recovery in the leachate (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


90 35g/L 10 g/L

85 80 75 70 65 60 55 50

0

5

10

15

20

25

30

35

Dissolution time (h) Figure 6. Lithium recovery in the solution for different reaction times.


29


H2O

H2C2O4

LiOH

Li2C2O4

2[

M

n 2+ ]

(a) [K + ] [H + ]+

F1 L2

L1

[L

i +]

+

L3

[K + ]+

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47

Page 30 of 48

M1

F2

KOH [K

M3

M2

K2C2O4

+

]+

[L

i +]

2[

M n2 + [H +] + ]+ 2[

M

n 2+ ]

Mn(OH)2

2[C2O42-]

MnC2O4

2[C2O42-] + [OH-]


30

Page 31 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47


(b)

H 2O H 2C2O 4 LiOH Li2C2O4

0.810

L1 M1

0.2

0.2

0.814

L2 M2

0.818

0.822

L3

0.4

0.4 0.6

0.6

M3 0.8

KOH

0.8

0.8

0.6

0.4

0.2

Mn(OH)2

Figure 7. (a) Three-dimensional phase diagram of the precipitation process for Mn removal and the associated cuts, (b) a cut passing through Mn(OH)2, KOH, and the overall feed point after adding KOH.


31


Page 32 of 48

K2CO3 (F4)

Liquid solution after precipitating Mn by KOH (F3)

Mother Liquor

Precipitation

Li2CO3 Figure 8. Process flowsheet of lithium recovery by chemical precipitation.


32

Page 33 of 48

H2O KOH LiOH Saturation surface of Li2CO3 F3

L4 L5 L6 M4

M5 M6

F4

K2C2O4

K2CO3

[K + ] [K + ]+ [Li + ]

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Li2C2O4

[CO32-]

Li2CO3

[C2O42-] + [CO32-]

Figure 9. Three-dimensional conceptual phase diagram for lithium recovery by chemical precipitation using K2CO3.


33


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Figure 10. XRD patterns of solids recovered from chemical precipitation using K2CO3.


34

Page 35 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


H2O KOH LiOH

F3 L6 L6 '

K2C2O4 +Liq

L6''

M6 M6 '

Li2CO3 + Liq

M6''

F4

K2C2O4 K2CO3

Solid mixture +Liq S6''

Li2CO3

Figure 11. A cut, passing through Li2CO3, H2O, and the overall feed point after adding K2CO3, of the 3D conceptual phase diagram in Figure 9.


35


Page 36 of 48

EtOH, K2CO3 (a) Liquid solution after precipitating Mn by KOH (F3)

Mother Liquor H 2O

Precipitation

EtOH

Li2CO3 K2C2O4·H2O

Mother Liquor

Dissolution

(b) Liquid solution after precipitating Mn by NH4OH (F3’)

Li2CO3 Mother Liquor

Precipitation

Li2C2O4 (NH4)2C2O4 ·H2O

Heat Treatment

Li2CO3

Figure 12. Lithium recovery by anti-solvent precipitation.


36

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Figure 13. XRD patterns of (a) solids recovered from anti-solvent precipitation using K2CO3 as precipitant and (b) after the dissolution process.


37


Page 38 of 48

Figure 14. XRD patterns of solids recovered from (a) anti-solvent precipitation using NH4OH as precipitant and (b) thermal decomposition product


38

Page 39 of 48 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


EtOH Saturation surface of (NH4)2C2O4 Saturation surface of (NH4)2C2O4•H2O

L9 M9 M8 M7

(NH4)2C2O4•H2O (NH4)2C2O4 S9 S8 S7

Saturation surface of Li2C2O4

L8 L7

F3'

H2O

Li2C2O4 Figure 15. Three-dimensional conceptual phase diagram for lithium recovery by anti-solvent precipitation.


39


Page 40 of 48

EtOH, H2O

L9 (NH4)2C2O4•H2O + Liq

Li2C2O4 + Liq

Solid mixture + Liq

M9

(NH4)2C2O4•H2O

S9

Li2C2O4

Figure 16. A cut, passing through the (NH4)2C2O4·H2O, Li2C2O4 and the overall feed point after adding ethanol M9, of the phase diagram in Figure 15.


40

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Table 1. Experimental results of the selective lithium dissolution from cathode materials (step 1) at different conditions and a dissolution time of 12 hours.

Experimental conditions

% of metals being dissolved from the cathode materials Li

Ni

Mn

Co

(wt%)

(wt%)

(wt%)

(wt%)

95.4

0.1

24.0

0.52

0.75

84.0

N.D.

20.1

0.24

1.5

93.9

0.19

26.7

0.70

2

94.6

0.32

32.9

0.90

80

81.7

0.28

13.2

0.71

85

87.0

0.30

14.9

0.82

90

94.3

0.19

17.3

0.70

5

97.2

0.17

37.8

1.5

15

86.6

0.13

14.0

0.32

30

79.8

N.D.

4.8

0.17

35

79.0

N.D.

2.2

N.D.

Base case C = 1 M, T = 95 oC, S/L = 10 g/L

C (M)

T (oC)

S/L (g/L)


41


Page 42 of 48

Table 2. Experimental results of the selective lithium dissolution from cathode materials (step 1) at different solid-to-liquid ratios (C=1 M, T=95 oC) and a dissolution time of 35 hours % of lithium being S/L of liquid after dissolved (g/L) from the lithium cathode dissolution materials pH

Concentration of metals in the solution

Ni (ppm)

Mn (ppm)

Co (ppm)

35

1.63

96.3

N.D.

120.7

3.96

40

3.09

95.3

33.4

215.1

33.9

45

5.43

95.7

55.1

533.2

78.4

50

8.56

96.4

417.1

729.3

239.0


42

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Table 3. Experimental results on Mn removal in step 2 using different precipitants

Precipitant

KOH (1 M)

NaOH (1 M)

NH4OH (5 M)

Volume of precipitant (mL)

pH of the Concentration solution after of Mn in the precipitation solution (ppm)

% of Mn precipitated as solids (%)

3.2 (M1)

5.49

71.92

2.28

3.3 (M2)

8.84

53.22

26.8

3.4 (M3)

11.36

N.D.

100

2.7

5.58

76.59

2.28

2.8

10.61

N.D.

100

2.9

11.67

N.D.

100

1

9.76

73.53

26.9

2

10.38

15.60

81.9

4

10.69

N.D.

100


43


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Table 4. Experimental results on lithium recovery in step 3 by chemical precipitation (25 mL feed solution)

Precipitant

K2CO3 (5 M)

Na2CO3 (3 M)

Volume of precipitant (mL)

Amount of Li in the solution (mg)

% of Li precipitated as solids (%)

2 (M4)

38.3

44.8

4 (M5)

16.7

76.0

6 (M6)

14.9

78.5

7

38.0

45.2

10

37.5

46.0

15

36.0

48.1


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Table 5. Experimental results on lithium recovery in step 3 by anti-solvent precipitation (25mL feed solution)

Volume of ethanol (mL)

organic to aqueous volumetric ratio (O/A)

Amount of Li in the solution (mg)

% of Li precipitated as solids (%)

K2CO3,

30

1

14.6

78.9

(5 mL, 5 M)

45

1.5

9.6

86.2

Ethanol

60

2

6.7

90.1

37.5 (M7)

1.5

11.0

84.1

50 (M8)

2

7.0

90.0

62.5 (M9)

2.5

6.6

90.5

Precipitant and anti-solvent

None, Ethanol


45


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Table 6 Experimental results on the dissolution process after anti-solvent precipitation by K2CO3 H2O (mL)

% of Li2CO3 dissolved

% of K2C2O4 dissolved

4

48.8

96.5

6

65.0

99.8

8

100

100


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Table 7. Experimental results of selective lithium dissolution from different cathode materials (step 1) % of metals being dissolved from the cathode materials Cathode material

Li

Ni

Co

Mn

(wt%)

(wt%)

(wt%)

(wt%)

NMC (111)

102.3

0.26

1.12

22.8

NMC (532)

95.4

0.12

0.52

24.0

NMC (811)a

89.6

0.13

0.17

20.6

NMC (811)b

96.2

0.14

0.31

24.5

LiCo0.95Mn0.05O2

93.6

N.D.

1.8

30.0

LiMn2O4

100.5

N.D.

N.D.

26.2

Under determined condition (oxalic acid concentration = 1 M; temperature = 95 C; solid-toliquid feed ratio = 10) a

Under modified condition (oxalic acid concentration = 2 M; temperature = 95 C; solid-toliquid feed ratio = 10) b


47


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Table of Contents LiNixCoyMn1-x-yO2 H2C2O4

Li+ Mn2+

MC2O4·2H2O (M=Ni/Co/Mn)

Precipitant

Precipitant Li+

Li+, Mn2+ Li+

Mn(OH)2

Li2CO3


48

Process synthesis: Selective recovery of lithium from lithium ion battery

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