Waste Management 51 (2016) 234–238

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Waste Management

journal homepage: www.elsevier .com/ locate/wasman

Use of mild organic acid reagents to recover the Co and Li from spentLi-ion batteries

http://dx.doi.org/10.1016/j.wasman.2015.12.0080956-053X/� 2015 Elsevier Ltd. All rights reserved.

⇑ Corresponding author.E-mail address: vasantapai@gmail.com (K.V. Pai).

Girish Praveen Nayaka a, Karkala Vasantakumar Pai a,⇑, Jayappa Manjanna b, Sangita J. Keny c

aDept. of Industrial Chemistry, Kuvempu University, Shankaraghatta 577 451, IndiabDept. of Chemistry, Rani Channamma University, Belagavi 591 156, IndiacChemistry Group, Bhabha Atomic Research Centre, Mumbai 400 085, India

a r t i c l e i n f o

Article history:Received 24 August 2015Revised 11 December 2015Accepted 12 December 2015Available online 18 December 2015

Keywords:Spent lithium-ion batteryRecovery of cobaltChemical dissolutionIminodiacetic acidMaleic acidOxalic acid

a b s t r a c t

New organic acid mixtures have been investigated to recover the valuable metal ions from the cathodematerial of spent Li-ion batteries. The cathodic active material (LiCoO2) collected from spent Li-ion bat-teries (LIBs) is dissolved in mild organic acids, iminodiacetic acid (IDA) and maleic acid (MA), to recoverthe metals. Almost complete dissolution occurred in slightly excess (than the stoichiometric require-ment) of IDA or MA at 80 �C for 6 h, based on the Co and Li released. The reducing agent, ascorbic acid(AA), converts the dissolved Co(III)- to Co(II)-L (L = IDA or MA) thereby selective recovery of Co as Co(II)-oxalate is possible. The formation of Co(III)- and Co(II)-L is evident from the UV–Vis spectra of thedissolved solution as a function of dissolution time. Thus, the reductive-complexing dissolution mecha-nism is proposed here. These mild organic acids are environmentally benign unlike the mineral acids.

� 2015 Elsevier Ltd. All rights reserved.

1. Introduction

Li-ion batteries (LIBs) have been the power source for portableelectronic devices viz., mobile phones, personal computers, cam-eras and recently in electric vehicles due to their favorable charac-teristics such as high energy density, high voltage, long storage life,low self discharge rate and wide temperature range of use(Chagnes and Pospiech, 2013; Thyabat et al., 2011; Gonclaveset al., 2015). The production of LIBs will be increased further inthe upcoming years as the large numbers of electric vehicles areentering the market (Scrosati and Garche, 2010). Such an widespread use of these portable energy storage devices by billions ofpeople around the world, large number of spent LIBs are accumu-lated day-by-day (Jha et al., 2013a,b). The active cathode materialin most of the LIB is LiCoO2 due to its good performance, althoughnew materials such as LiMn2O4 and LiFePO4 are being developed.Such lithium cobalt oxide must be treated properly consideringthe environmental toxicity of Co and at the same time the scarcityof Li. Furthermore, Co is a rare and precious metal, and is a rela-tively expensive (Li et al., 2010b; Hayashi et al., 2009; Yang et al.,2011). Lithium is also vitally important for many industrial appli-cations. Hence, recovery of these valuable metals by a suitable

method would greatly benefit the society and environment. Thus,recycling of these spent LIBs is appropriate for at least two signif-icant reasons: it often pays to recover valuable materials, espe-cially if their supply is limited; and now it is necessary to recycleLIBs according to government regulations for the sake of sustain-ability and safety hazards associated with the disposal of spentLIBs (Chagnes and Pospiech, 2013). Therefore, spent LIB must beproperly recycled through pyrometallurgical or hydrometallurgicalprocesses (Maschler et al., 2012; Ziemann et al., 2012; Li et al.,2010a; Chen et al., 2011). The technologies existing for LIBs recy-cling can be categorized into physical (electrostatic, magnetic,gravity separations) and chemical (electrolysis, solvent extraction,bioleaching, leaching, precipitation) separations. Among these,hydrometallurgical process is advantageous from environmentalconservation viewpoint (Freitas et al., 2010).

There are several studies on dissolution of LiCoO2 from spentLIBs by strong acids like H2SO4 (Chen et al., 2011; Daniel et al.,2009; Kang et al., 2010), HCl (Wang et al., 2009), and HNO3

(Freitas et al., 2010; Ivano et al., 2009) with the addition H2O2 asreducing agents (Dorella and Mansur, 2007; Swain et al., 2007)with more than 95% recovery of Co and Li. However these strongacids are high cost, difficult to handle in large scale and are notenvironment friendly as they emit toxic gases. So it is importantand essential to develop simple, cost-effective and environmen-tally benign recycling processes to recover these valuable metalsfrom spent LIBs. Recently few studies have reported on mild

Fig. 1. SEM/EDXA images of calcined LiCoO powder after collecting from spent LIB.

G.P. Nayaka et al. /Waste Management 51 (2016) 234–238 235

organic acid with H2O2 as reducing agent viz., citric acid (Li et al.,2010c, 2013), malic acid, aspartic acid (Li et al., 2013), oxalic acid(Sun and Qiu, 2012) and ascorbic acid (Li et al., 2012). In our pre-vious study (Nayaka et al., 2015), we have used citric acid andascorbic acid as reducing agent because it was proved to dissolveeven the sintered metal oxides like Cr-substituted hematites(Manjanna et al., 2001). All of these reagents have not shownalmost complete dissolution of Co and Li. In order to investigatethe efficient dissolution media among various mild organic acids,this study is focused on the iminodiacetic acid (IDA) and maleicacid (MA) in presence of ascorbic acid. There are no reports onthese reagents, which are soluble in aqueous medium and environ-mentally benign.

2

0

20

40

60

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solu

tion,

%

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IDA % (Co) % (Li )

a

0 50 100 150 200 250 300 350 400

0 50 100 150 200 250 300 350 400

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MA % (Co) % (Li)

2. Materials and methods

All the reagents used in this study are of analytical grade, and allthe solutions were prepared in distilled water. The spent LIBs(BL-5CA, Nokia) used here were collected from the local marketand dismantled to obtain the cathode material as follows.

To prevent self-ignition and short-circuiting, the spent LIBswere discharged completely and dismantled to separate the cath-ode and anode materials coated on curled Al- and Cu-foil, respec-tively. The Al-foil was uncurled to cut in to small pieces andimmersed in n-methyl pyrrolidone (NMP) solvent and subjectedfor ultrasonication. Once all the active cathode material in the formof powder was detached from Al-foil, it was filtered and heated at700 �C for 2 h to burn off the organics such as carbon andpolyvinylidene fluoride. The powder was then ground to fine pow-der to increase the surface area for higher leaching efficiency. Thepowder sample was characterized by X-ray diffraction (XRD) usingCu Ka radiation (k = 1.542 Å) with a Ni-filter (Bruker D8 AdvanceX-ray diffractometer). Morphology (SEM, HITACHI S-3000H) andthe elemental composition of the sample were examined byrecording energy dispersive X-ray (EDXA) spectra. The samplewas found to be highly crystalline and coarse sized LiCoO2.

About 20 mM metal ion worth of above cathode material(LiCoO2, 0.2 g) was subjected for chemical dissolution in 100 mlaqueous mixture of (i) 100 mM IDA and 20 mM AA (I–A) (ii)100 mM MA and 20 mM AA (M–A) at 80 �C for about 6 h on tem-perature controlled magnetic stirrer. The periodically collectedsamples with syringe filters (0.2 lm) were analyzed for Co and Liusing atomic absorption spectrometer (Model: AA-7000F, ROMversion: 1.01, S/N: A30664801195). The UV–Vis spectra of the sam-ples were recorded using Ocean optics, DH-2000 BAL. At the end ofdissolution, the insoluble reside was collected after filtration andwashings. It was quantitatively estimated to account for the % dis-solution obtained. Some of the selected experiments were repeatedtwice or trice and the results here carry <5% error.

Fig. 2. Dissolution profiles of Co and Li in I–A and M–A mixtures at 80 �C.

1 For interpretation of color in Figs. 3 and 4, the reader is referred to the webversion of this article.

3. Results and discussion

Fig. 1 shows the SEM images and EDXA of the active cathodematerial (LiCoO2) obtained from spent LIB. The irregular andagglomerated particles (<1 lm) can be clearly seen in SEM images.The elemental composition by EDXA shows the presence of Cowhereas the low Z, Li cannot be detected and no other metal impu-rities. The XRD pattern of the sample confirmed the rocksalt struc-tured LiCoO2 (Nayaka et al., 2015).

Fig. 2 shows the percentage dissolution of Co and Li ions in I–Aand M–A (100–20 mM) mixtures at 80 �C as a function of time. Theamount of dissolution was estimated with respected to 20 mM ofeach metal ions expected on complete dissolution i.e., 0.2 g ofLiCoO2 in 100 ml. Accordingly, in both I–A and M–A mixtures,almost complete dissolution was obtained. In both the cases, the

dissolution showed a fast initial stage followed by a slow secondstage. For instance, 90–95% dissolution occurred in about 1 h. Sucha fast dissolution is ascribed to complexation ability of IDA and MAwhen compared to other organic acids (Armstrong and Bruce,1996; Lee and Reeder, 2006). Table 1 shows the dissolution ofLiCoO2 from spent LIB in different organic acids viz., oxalic acid,citric acid, ascorbic acid, malic acid, succinic acid, maleic acidand iminodiacetic acid in presence or absence of reducing agentslike H2O2. In this study, as the dissolution takes place, the solutionturned from colorless to pink1 due to complexation of cobalt ions

Table 1Dissolution of LiCoO2 (obtained from spent LIB) in different organic acids.

Reagents Conditions % Dissolution Reference

1 M OA + 15 vol.% H2O2 80 �C, 120 min 98% Li & 68% Co Sun and Qiu (2012)1.25 M AA 70 �C, 20 min 98% Li & 95% Co Li et al. (2012)1.25 M CA + 1 vol.% H2O2 90 �C, 30 min 100% Li & 90% Co Li et al. (2013)1.5 M malic acid + 2 vol.% H2O2 90 �C, 40 min 100% Li & 90% Co1.5 M aspartic acid + 4 vol.% H2O2 90 �C, 120 min 60% Li & 60% Co0.1 M CA + 0.02 M AA 80 �C, 360 min 100% Li & 80% Co Nayaka et al. (2015)1.5 M SA + 4 vol.% H2O2 70 �C, 40 min 96% Li & 100% Co Li et al. (2015)1 M IDA + 0.02 M AA 80 �C, 360 min 99% Li & 91% Co This study1 M MA + 0.02 M AA 80 �C, 360 min 100% Li & 97% Co

OA: oxalic acid; AA: ascorbic acid; CA: citric acid; MA: maleic acid; SA: succinic acid; IDA: iminodiacetic acid.

236 G.P. Nayaka et al. /Waste Management 51 (2016) 234–238

with IDA or MA. The intensity of pink color of the dissolving mediumincreased with dissolution time. The dissolution is relatively morepronounced in M–A when compared to I–A mixture. This could berelated to their stability constants, logK of Co with IDA and MA is6.95, and 1.83, respectively. The possible complexation of both Liand Co with L (=IDA or MA) is shown in Fig. 3. The main reasonfor using these mild chelating agents is for the selective recoveryof Co and Li through precipitation in the subsequent steps. If weuse strong chelating agents like EDTA, it is not possible to dissociatethe corresponding complexes. We have used ascorbic acid (reducingagent) along with the IDA or MA which are chelating and bufferingagents. In fact, ascorbic acid is not necessary for the dissolutionbecause we have used stoichiometrically excess chelating agents.However, ascorbic acid (AA) helps to reduce the Co(III)-L formed dur-ing dissolution to relatively weaker complex, Co(II)-L, which can beeasily recovered as Co(II)-oxalate precipitate. Thus, we have judi-ciously selected AA in the dissolution mixture to enable selectiverecovery of Co. Furthermore, AA is known to initiate the dissolutionof sintered oxides like Cr-substituted hematites (Manjanna et al.,2001). So, our aim here is to develop suitable chemical formula orreagents which can easily dissolve the active cathode materials fromspent LIBs. In some cases, the oxide lattice may be very stable, espe-cially when it is calcined to remove the organics. So, the presence ofAA takes care of stable oxide lattices and reduce the dissolved Co(III)- to Co(II)-L. Also, AA can provide buffering (i.e., H+ requirementduring dissolution to take care of oxide ions) and complexation withstoichiometrically excess metal ions, thereby increase the dissolu-tion efficiency. The pH of the solution before and after dissolutionwas 2.25 and 2.78 for I–A mixture and it was 2.27 and 3.15 forM–A mixture, respectively. Overall, it is necessary and advantageousto have AA in the dissolving medium. In the earlier reports, H2O2 (Liet al., 2013, 2012; Sun and Qiu, 2012) is commonly used as reducingagent with organic acids to increase leaching efficiency of Co.

O Li Co

OR

→

MA, C4H4O4

→

+

IDA, C4H7O4N

LiCoO2

Fig. 3. The possible complexation of Li and C

However, AA has many advantages over H2O2 in terms of its environ-mental compatibility and chelating action. Based on the preliminarystudies, we have optimized the concentration of reagents, 100 mM(=IDA or MA) and 20 mM of AA when the solid/liquid ratio is 2 g/Li.e., 0.2 g of LiCoO2 in 100 ml. As shown in Fig. 2, the amount of Lireleased is slightly higher than Co, although equal amounts areexpected if the LiCoO2 used here is in stoichiometric composition.However, during the life time charging–discharging of LIB, it isknown that the slight amount of Li is entrapped in graphite anodein the form of lithium carbide. Thus, the spent LiCoO2 becomesLi-depleted structure, although negligible. Accordingly, when we cal-culated the % dissolution here assuming nominal composition i.e.,20 mM each of Co and Li, we get relatively higher dissolution for Liwhen compared to Co. Nevertheless, we have not done completeanalysis of the elemental composition in the starting material. Basedon the dissolution behavior here, we estimate that the LiCoO2 usedhere is about 5% Li-depleted oxide.

Fig. 4 shows the UV–Vis spectra of dissolved solution collectedat different interval of time. The initially colorless solution wasturned to pink color and the intensity increased with dissolutiontime. This increase in intensity is an indication of Co(III) releasefrom oxide lattice (in LiCoO2, cobalt is present as Co3+) throughcomplexation with IDA or MA. However, it is clear from the spectrathat Co(III)-L is reduced to Co(II)-L due to the presence of reducingagent, AA in dissolution mixture. For comparison, the inset figuresshow the UV–Vis spectra of the dissolved solution in the absence ofAA. In the case of IDA, both Co(III)-L (kmax � 560 nm) and Co(II)-L(kmax 380 nm) are present, whereas only Co(III)-L (kmax � 512 nm)is present in MA. It shows that IDA acts as both chelating andreducing agent to reduce Co(III)- to Co(II)-L, unlike MA. In orderto obtain all the cobalt in the form of Co(II)-L, the presence of AAis essential in the dissolution mixture (more so in the case ofMA). Accordingly, in presence of AA all the cobalt released was

+

C H N

C4H2O4Li2 C4H2O4Co

+

C4H5O4NLi2 C4H5O4NCo

o with IDA and MA during dissolution.

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Fig. 4. UV–Vis spectra of the dissolved solution at different interval of time. Theinset figures show the spectra when there was no ascorbic acid in the dissolutionmixture.

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)(6

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(224

)

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)(0

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) (400

)

IDA

MA

Inte

nsity

(CPS

)

Scattering Angle, 2θ (Cu Kα radiation)

(202

) JCPDS file No. 25-0250

Fig. 5. XRD pattern of CoC2O4�2H2O precipitate recovered from the dissolvedsolution. Inset photo is given to show its actual color. (For interpretation of thereferences to color in this figure legend, the reader is referred to the web version ofthis article.)

G.P. Nayaka et al. /Waste Management 51 (2016) 234–238 237

present as Co(II)-IDA (kmax 330 nm) and Co(II)-MA (kmax 280 nm) atthe end of dissolution. Therefore, AA helps in the subsequent selec-tive recovery of cobalt as solid Co(II)-oxalate because only suchweak Co(II)-L (=IDA or MA with logK = 6.95, and 1.83, respectively)can be dissociated to form ca. Co(II)-oxalate. We have observed asimilar behavior in our pervious study (Nayaka et al., 2015). Toour knowledge, there are no in literature on the UV–Vis spectraof Co(III)- and Co(II)-L. It is important to note that the dissolvedsolution also contains Li-L, although it is not a colored complex like3d metal ion, it is expected to show absorption in UV region. Thus,the kmax values here cannot be ascribed to pure cobalt complexesbecause the influence from Li-L must be considered. For pure Co(III)-L, cobalt oxide like Co2O3 must be used. Nevertheless, thespectral features are useful from the coordination chemistry view-point for these important complexes.

At the end of dissolution here, the dissolved solution wasallowed to attain the room temperature and any undissolved resi-due (which was negligibly small) was removed by filtration. Then,100 mM worth of solid oxalic acid was added to obtain the Co(II)-oxalate precipitate, CoC2O4�2H2O. It was separated again by filtra-tion, washed thoroughly and dried in oven at 90 �C for 24 h beforerecording XRD (Fig. 5). The XRD pattern confirms the well crystal-lized orthorhombic structure with space group of Cccm (JCPDS file:25-0250), which is in agreement with literature (Sun and Qiu,2012). This Co(II)-oxalate can serve as suitable Co precursor to

obtaining the starting material. After such a selective recovery ofCo, it was ensured that there was no Co remained in the filtrateusing AAS. In order to recover Li from the filtrate solution, we usedNa2CO3 and CO2 purging, but, there was no Li2CO3 formation. Infact, such Li2CO3 formation is reported when the dissolution wascarried out with mineral acids. In the present study, Li is presentas Li-L (=IDA or MA) and hence it was not possible to precipitateas its carbonate. For academic interest, we have recovered all theLi as LiF by using NH4F, though it is not economical and environ-mental benign approach. Our efforts are on to recover the Li bysuitable method based on adsorption with some nanomaterials,which will be published separately.

4. Conclusions

The mild organic acids, iminodiacetic acid (IDA) and maleic acid(MA) in presence of ascorbic acid (AA), used here to dissolved theactive cathode material (LiCoO2) from spent LIB showed completedissolution under stoichiometric condition at 80 �C. In both thecases, the dissolution showed a fast initial stage with >90–95% dis-solution in about 1 h, followed by a slow second stage. The AA wasfound to play an important role during the dissolution process. Forinstance, the formation of Co(III)-L (=IDA with kmax � 560 nm orMA with kmax � 612 nm) was evident from the UV–Vis spectra ofthe dissolution mixture in the absence of AA. However, in the pres-ence of AA, it was reduced to Co(II)-IDA (kmax 330 nm) and Co(II)-MA (kmax 280 nm). Using oxalic acid, such weak Co(II)-L was easilydissociated to form CoC2O4�2H2O precipitate. Although the Li wasrecovered here as LiF, our efforts are to recover it by an economicaland environmental benign approach. Overall, in this study we havedeveloped a new organic acid mixture formulation to recover thevaluable metal ions from the cathode material of spent Li-ionbatteries.

Acknowledgment

One of the authors (G.P. Nayaka) gratefully acknowledges thefinancial support from the Kuvempu University and this workforms part of his Ph.D. thesis.

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