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Industrial & Engineering Chemistry Research pubs.acs.org/IECR Article Figure 7. Lithium concentration vs time without any divalent dissolved ions (I = 4.20 M, dashed line with square symbols) and with addition of (i) 0.17 M CaCl2 (dotted line with rhombus symbols), (ii) 0.25 M MgCl2 (dashed lines with cross-symbols), (iii) 0.17 M SrCl2 (dot-dashed lines with triangle symbols), and (iv) 0.17 M CaCl2 + 0.25 M MgCl2 + 0.17 M SrCl2 (dashed lines with circle symbols). Stirring speed = 300 rpm, CO32−/Li+ ratio = 1, and Na2CO3 solution flow rate = 10 mL/min. T = 50 °C. to the complete consumption of carbonates ions by precipitation of the added divalent cation salts. Li+ recovery and purity values in the presence of divalent cations are shown in Figure 8. Figure 8. Recovery and purity for Li2CO3 precipitation experiments in the presence of divalent cations in high-ionic strength solutions. No recovery was calculated in the simultaneous presence of Ca2+, Sr2+, and Mg2+ since no precipitation occurred. As already commented in Figure 6, Li+ recovery can reach a value around 70% for high-ionic strength solutions without any divalent ions. Here, the presence of divalent ions causes a Li+ recovery decrease to ∼60 and ∼40% in the case of Ca2+ or Sr2+ and Mg2+ ions, respectively. Li+ recovery is totally inhibited in the simultaneous presence of all three divalent salts (no recovery). The negative impact of the presence of divalent ions can be also observed on the low Li2CO3(s) purity, never exceeding 28% due to the co-precipitation of other carbonate compounds. Due to the considerable impact of divalent ion presence on the Li2CO3 precipitation process, the influence of Mg2+ concentration was further investigated considering only Mg2+ traces, which are likely to be present in the Li-MFCDI eluates of the actual SEArcularMINE treatment chain. In this case, precipitation was carried out at 80 °C (again, to focus on the expected condition in the actual treatment chain) by varying the Mg2+ concentration from ∼0.003 to ∼0.044 M. For the sake of brevity, only Li recovery and purity are reported in Figure 9 as functions of the initial Mg concentration. In this case, Li+ recovery values are close to ∼70% for all Mg2+ concentrations, thanks to the higher employed temper- Figure 9. Recovery and purity as a function of initial magnesium concentration. LiCl solutions of 0.70 M with added salts: 2.0 M NaCl and 1.5 M KCl. T = 80 °C. Stirring speed = 300 rpm, CO32−/Li+ ratio = 1, and Na2CO3 solution flow rate = 10 mL/min. ature; although, also in this case, they result in a lower recovery than that obtained with monovalent salts solutions (78%). A non-monotonic Li+ purity trend is observed with increasing Mg2+ concentration. Specifically, the purity increases from ∼80 to ∼90% up to a Mg2+ concentration of 0.01 M, which further decreases at higher Mg2+ concentrations. Purity decreases to values around 60% even at a low Mg concentration of 0.044 M, indicating that the presence of Mg2+ ions represents a crucial issue in Li2CO3 recovery processes from Mg2+-containing sources (a better combined strategy to by-pass this issue will be presented in Section 3.2.3). After the purification step with ethanol, purity values increase, leading to an almost monotonical decreasing trend, when increasing Mg2+ concen- tration. However, for higher Mg2+ concentrations, the washing step was unable to reach the 100% purity observed in the previous tests, thus again indicating the dramatic influence of Mg salts co-precipitation on the product purity. Also in this case, a loss of product is observed, resulting in an equivalent reduction of Li recovery from 70 to 57%. 3.1.4. Influence of Sulfates and Bromides on Li2CO3(s) Precipitation. The influence of sulfate and bromide anions on the Li2CO3(s) precipitation was studied by preparing six different solutions containing 0.70 M LiCl plus • 1.4 M Na2SO4 (I = 4.90 M) • 1.0MKCland1.4MNa2SO4 (I=5.90M) • 1.0MNaBr(I=1.70M) • 1.1MKCland1.0MNaBr(I=2.80M). Note that all salt concentrations refer to solutions before Na2CO3 solution addition. All precipitation tests were carried out at 50 °C with a stirring velocity of 300 rpm and a double excess of a 2.0 M Na2CO3 solution (CO32−/Li+ ratio of 1), fed at a flow rate of 10 mL/min. The Li+ concentration trends during the precipitation time in the presence of sulfate and bromide ions are shown in Figures 10 and 11, respectively. From Figure 10, it can be seen that the Li2CO3 precipitation rate considerably decreases in the presence of sulfate, in accordance with the reported delaying effect of sulfate ions on Li2CO3(s) nucleation.32 The delaying effect is reduced in high- ionic strength solutions, although no precipitation occurs within the experiments time; thus, no recovery and purity were calculated. It is worth noting that the dissolution of Na2SO4 salts also causes a salting-in effect that, in turn, leads to a Li2CO3 solubility increase, affecting the overall precipitation process. Figure 11 shows the Li+ concentration trend in the presence of Br−. It can be observed that Br− ions do not significantly affect the Li precipitation since similar concentration trends as those for pure LiCl solutions, see Figure 5a, are obtained. Furthermore, in the presence of KCl salt (I = 2.80 M), a final https://doi.org/10.1021/acs.iecr.2c01397 13595 Ind. Eng. Chem. Res. 2022, 61, 13589−13602PDF Image | Recovery of Lithium Carbonate from Dilute Li Rich Brine
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