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lithium from solutions with Li/Na of 1/100, 1/1000 and 1/10 000 was studied. In addition, a solution containing only NaCl was used as a control sample. The battery was a three-electrode flooded cell, with the electrodes immersed in a 500 ml flask. In the first step of the operation of the cell, Li+ ions were captured by applying a current density equal to 0.5 mA cm2 to the FePO4 cathode, according to the following equation: FePO4 + Ag + LiCl / LiFePO4 + AgCl In the second step, the two electrodes (LFP and Ag) were moved into a second cell with reduced volume (about 350 ml) containing the recovery solution. The recovery solution was made of a 50 mM KCl aqueous solution that guarantees an adequately low ohmic resistance while providing an internal standard (K+) for the subsequent water analysis. In the third step, the battery electrodes released the Li+ and Cl ions (backward reaction) into the recovery solution by applying a current density equal to 0.5 mA cm2 to the LiFePO4 cathode. Fig. 2(a) shows the typical voltage profile of the battery during the capture of ions (first step) at different Li/Na ratios in the brine obtained by applying 0.5 mA cm2 for two hours. The concen- tration of lithium in solution highly affects the shape of the curves. By decreasing the Li/Na ratio, the plateau shifts toward more negative voltages, which means that more energy is required for the ion capture, as expected from thermodynamic and kinetic considerations. From the Nernst equation, by decreasing the concentration of lithium, the voltage at which Li+ is captured is also shifted towards more negative voltages. At low Li+ concentrations, the concentrated Na+ intercalates into the FePO4 at a higher potential than the dilute Li+, so the electrode selectivity decreases. This reaction is slow, as Na+ has a lower diffusivity than Li in LFP.22 For the lower concentrations, namely 1/10 000 and 0/1, the results of the experi- ments are more subject to statistics due to the fact that the applied current density is much higher than the limiting diffusion current (see ESI†). Also, the kinetics of the insertion are strongly influenced by lithium concentration. A decrease in lithium concentration results in a decrease of the limiting diffusion current and therefore an increase of the concentration overvoltage, which further shifts the potential profile towards more negative voltages. The control sample without Li+ (0/1) in solution shows a clear plateau in the lower voltage region, which is attributed to sodium intercalation.23 Fig. 2(b) shows the corresponding third step of the battery, during which a current density of 0.5 mA cm2 is applied for 2 hours to the system after the transfer of the electrodes to the 50 mM KCl recovery solution. To avoid any side reactions such as hydrolysis of the elec- trolyte, a voltage limit of 0.5 V was imposed to the battery. From the voltage profile it is evident that the behavior of the cell during this step is almost independent of the conditions of the previous ion capture, although some differences can be seen. In particular, during the first 20 minutes the voltage profiles are considerably different. This is due to the release of some sodium ions, also captured by LFP. The lower the concentration of lithium in the brine, the greater the amount of sodium that is captured. At lower voltage the sodium is released first, which can be observed by the slope in the range of voltages between 0.2 and 0 V. The plateau above 0 V is due to lithium release.23 Since during the release the concentration of lithium changes significantly, the voltage necessary for the extraction also changes, according to the Nernst equation. Table 1 Mean ionic concentration of the solution and the total number of ions in solution in the brine cell and after the charging cycle in the recovery cell Brine cell [Li+] Li+/ /mM mmol 50 15 5 1.5 524 181 916 322 0.5 0.15 2711 94 1031 360.1 0.2 0.06 245 82 1116 392 Recovery cell [Li+]/ mM Li+/ mmol [Na+]/ Na+/ mM mmol View Online 1/100 1/1000 1/10 000 0/1 115 8 40 3 21 10 73 After the release of lithium ions into the recovery solution, a 200 ml sample was analyzed by inductively coupled plasma mass spec- trometry (ICP-MS) (Table 1) for quantitative analysis of the trans- ferred ions. The K+ added to the recovery solution was used as the internal standard, its concentration being 50 mM (not shown in the table). It should be noted that the final ratio Li/Na is up to 500 times the initial one. As observed from the voltage profiles of the capture and release (Fig. 2), the amount of sodium transferred increases by decreasing the ratio of Li/Na in the brine. It must be emphasized that the observed lithium transfer from the 0/1 brine is due to the lithium impurity in the NaCl source (50 ppm). The impurity generates a 0.25 mM Li+ concentration in the control solution and also influences the real concentration of Li+ in solution of the 1/10 000 brine. From the measured transferred ions, the volume of the cell for the recovery solution and the total charge flown during the capture and release steps, a charge efficiency of 63% was predicted for all the ions. The high ratio between the Li/Na in the 1/100 sample is indicative of the good selectivity of lithium iron phosphate for Li+.22,23 We want to stress that Mg2+ ions could eventually intercalate and decrease the selectivity of the LFP electrode. However, there are several methods to separate it from the lithium in a second step by using high surface electrodes (potential swing process24) or precipitation of Mg(OH)2 (lime soda process5). The energy consumption was calculated according to eqn (1) from the integral area of the plot DE vs. q. The resulting value was normalized by the amount of lithium recovered. An overview of the cycles is presented in Fig. 3 and the results are shown in Table 2. The theoretical energy was calculated based on the variation of the Gibbs free energy according to the experimental concentration in the brine and in the recovery solution. Due to the partial transfer of sodium, the resulting value is negative. This means that for this concentration range, the process should be thermodynamically favored. In contrast, the experimental results show that it is necessary to expend energy in order to transfer the lithium to the recovery solution. This discrepancy is related to the energetic loss in the transfer process due to ohmic drop, concentration overvoltage, water splitting and oxygen reduction (for more details see ESI†). For comparison, the energy consumption for the method proposed by Kanoh et al. is circa 33 W h mol1, calculated considering a charge efficiency of 100% and the oxygen evolution and the hydrogen evolution being the reactions at the counter electrode during the lithium capturing and release, respectively. Therefore the process presented here is much more energy-efficient than those of prior studies. The proposed battery for the recovery of lithium from brines shows low energy consumption when compared the previously This journal is a The Royal Society of Chemistry 2012 Energy Environ. Sci., 2012, 5, 9487–9491 | 9489 Downloaded by Stanford University on 24 October 2012 Published on 17 September 2012 on http://pubs.rsc.org | doi:10.1039/C2EE22977CPDF Image | Batteries for lithium recovery from brines
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