Hydrometallurgy 161 (2016) 54–57

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Hydrometallurgy

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

Dissolution of cathode active material of spent Li-ion batteries using

tartaric acid and ascorbic acid mixture to recover Co
G.P. Nayaka a, K.V. Pai a,⁎, G. Santhosh b, J. Manjanna c
a
Dept. of Industrial Chemistry, Kuvempu University, Shankaraghatta 577 451, India
b
Dept. of Mechanical Engineering, NMAM Institute of Technology, Nitte 574 110, India
c
Dept. of Chemistry, Rani Channamma University, Belagavi 591 156, India

a r t i c l e i n f o a b s t r a c t

Article history: Environmentally benign hydrometallurgical dissolution process is investigated for the recovery of cobalt from the
Received 3 August 2015 cathode active materials of spent lithium-ion batteries (LIBs). A mixture of tartaric acid and ascorbic acid is used
Received in revised form 15 December 2015 to dissolve the LiCoO2 collected from spent LIBs. The reductive-complexing mechanism led to N 95% dissolution
Accepted 21 January 2016
with 0.4 M tartaric acid and 0.02 M ascorbic acid in about 5 h at 80 °C. The dissolved Co was separated as cobalt
Available online 22 January 2016
oxalate from the mixture.
Keywords:
© 2016 Elsevier B.V. All rights reserved.
Spent lithium-ion batteries
Dissolution
Organic acids
Recovery of Co

1. Introduction process, hydrometallurgical process or the combination of both

(Bernardes et al., 2004; Al-Thyabat et al., 2013). Pyrometallurgical
Li-ion batteries (LIBs) are widely used as rechargeable power sup- process involves more energy consumption, high cost, low efficiency,
pliers because of their high capacity, high energy density, light weight loss of materials and emission of hazardous gases (Jha et al., 2013;
and small size (Chagnes and Pospiech, 2013; Bertuol et al., 2015; Shin Sun and Qiu, 2012; Chen et al., 2011; Joulie et al., 2014). Hence, recent
et al., 2015). Consequently, large quantities of spent LIBs are disposed studies have shown that the hydrometallurgical process is advanta-
as solid waste. Since they contain heavy metals like Co and other toxic geous to recover valuable metals from spent LIBs (Swain et al., 2007;
organic matters, they are hazardous to environment (Shapek, 1995). Lee and Rhee, 2003). Many studies have reported on the leaching of
Hence, it is imperative to recycle the valuable metals present in spent active cathode material (LiCoO2) using mineral acids such as HCl,
LIBs (Cai et al., 2014; Dunn et al., 2012; Wang et al., 2014). Thus, the HNO3, H2SO4 etc. (Espinosa et al., 2004; Fouad et al., 2007; Paulino
recovery of Co and Li from cathode active materials and Al and Cu et al., 2008; Granata et al., 2012). However, these processes introduce
from metallic components is important considering their uses in significant amount of secondary pollutants such as emission of Cl2,
commercial, industrial and military devices (Zhao et al., 2011; Swain SO3 and NOx. Also, the recovery processes become more complex due
et al., 2007; Nan et al., 2005). On the other hand, it is highly desirable to series of separation and purification steps for Li, Co and other compo-
to prevent environmental pollution from these elements, Co in particu- nents (Sun and Qiu, 2012).
lar. In view of the ever increasing consumption of metallic resources, re- To reduce such secondary pollution, without compromising leaching
sources recycling techniques/process have gained much importance. In efficiency, there are reports on using mild organic acids such as citric
addition to many basic R & D works, there are few companies viz., AEA acid, ascorbic acid, malic acid, oxalic acid, aspartic acid, succinic acid
Technology (U.K.), SNAM (France), Toxco (Canada), Umicore (Belgium) etc. (Li et al., 2010, 2012, 2013, 2015; Sun and Qiu, 2012) on adding
etc. which have developed the processes of recycling valuable metals H2O2 as reducing agent. Overall these mild organic acids are as effective
from the cathode of spent LIBs (Cai et al., 2014). as mineral acids. Recently, we have reported (Nayaka et al., 2015) using
In order to recover the metal ions from spent LIBs, the cathode mate- mixture of citric acid (chelating agent) and ascorbic acid (reducing
rial is usually dissolved in leaching agents followed by separation and re- agent). In this process, although Co was separated successfully as Co-
covery of metal salts from the solution in different steps. The conventional oxalate, the Li separation as LiF was not advantageous. Hence, we have
techniques for recycling spent LIBs mainly involve pyrometallurgical investigated on alternative mineral acids such as mixture of tartaric
acid (chelating cum buffering agent) and ascorbic acid (reductant).
⁎ Corresponding author. The details on dissolution behavior and separation of Co from dissolved
E-mail address: vasantapai@gmail.com (K.V. Pai). mixture of Co and Li ions are reported here.

http://dx.doi.org/10.1016/j.hydromet.2016.01.026
0304-386X/© 2016 Elsevier B.V. All rights reserved.
 G.P. Nayaka et al. / Hydrometallurgy 161 (2016) 54–57 55

2. Experimental used here is much higher in concentration. TA acts as a chelating

agent while the ascorbic acid (AA, 0.02 M) acts a reducing agent during
The spent LIBs (BL-5CA Nokia series) collected from the local market the dissolution process. As shown in Fig. 2, rapid dissolution to the
was discharged completely (to prevent self-ignition and short- extent of 80% occurred in about 30 min. Thereafter, marginal increase
circuiting), and manually dismantled to separate the cathode and in dissolution occurred over a period of 3–5 h as it reached to complete
anode parts. The cathode material coated on Al-foil was uncurled and dissolution. As shown in Fig. 3, on increasing the concentration of TA
cut to small parts. Upon ultrasonication in N-methyl-2-pyrrolidone from 0.1 to 0.5 M, overall there was about 15% increase of Co and
(NMP), all the oxide deposit was separated from Al–foil. After removing about 5% increase in Li release. It is clear that 0.4 M TA is enough to
Al metal part, the oxide powder was collected by filtration. It was heated bring complete dissolution and further increase in TA is not advanta-
to 700 °C for 2 h to burn off the organics such as carbon and geous. The chelating agent is responsible to leach the Co and Li ions
polyvinylidene fluoride. It was ground to fine powder for higher surface from the lattice through complexation process. As the lattice gets
area to increase the leaching efficiency. Fig. 1 shows the flow sheet of disturbed, both Co and Li ions are released from the LiCoO2. The metal
the process followed in this work. As reported in our previous study ions thus released are stabilized in the solution by complexation with
(Nayaka et al., 2015), the oxide thus obtained was found to be LiCoO2 tartarate (Sun and Qiu, 2012; Ferreira et al., 2009). Since Co is present
based on XRD, SEM/EDXA and FT-IR. as Co3 + in the oxide lattice, Co(III)-tartarate is expected to build-up
The rocksalt structured LiCoO2 was found to contain residual carbon, with dissolution time. However, the AA can reduce the Co(III)- to
5–10 wt.%, due to organic burn off. The origin for such carbon might be Co(II)-tartarate, and hence most of the Co is present as Co(II)-tartarate
from the acetylene black used to ensure the electronic conductivity in as the end of dissolution. The dissolution process can be shown as
the cathodes.
The above cathode material (LiCoO2, 0.2 g) was subjected for chem- 4 LiCoO2 + 12 C4H4O6 → 4 LiC4H3O6 + 4 Co(C4H3O6)2 + 6 H2O.
ical dissolution in 100 ml aqueous mixture of tartaric acid (TA) and
ascorbic acid (AA). TA concentration was varied from 0.1–0.5 M by Li et al. (2010, 2012, 2013, 2015) have clearly shown that leaching
keeping fixed concentration of AA (0.02 M). The dissolution was efficiency is enhanced by using reducing agent. However, in the present
monitored for about 5 h at 80 °C by collecting the samples periodically. study, the reducing agent AA is not found to increase the dissolution,
During sampling, the insoluble residue was separated by filtration nevertheless, reduces the dissolved Co(III)- to Co(II)-L. This is useful
through 0.2 μm syringe filter. The concentration of Li and Co metal for subsequent recovery of Co as Co(II)-oxalate. Manjanna and
ions in the dissolved solution was determined by using atomic absorp- Venkateswaran (2001, 2002) have shown that AA can reduce the lattice
tion spectrophotometer (Model: AA-7000F). The results obtained here metal ions such as Fe3+ in the dissolution of Cr-substituted hematite
are found to carry a maximum of about 5% error. The UV–vis spectrum and it was essential to initiate the dissolution. However, such a
of the dissolved solution was also recorded using UV–visible spectro-
photometer (Ocean optics, DH-2000 BAL).

3. Results and discussion

The amount of LiCoO2 used here (0.2 g) is worth of 0.02 M Co and

0.02 M Li. Thus, stoichiometrically, the tartaric acid (TA, 0.1–0.5 M)

Fig. 1. Flow sheet for the recovery of Co from active cathode material (LiCoO2) in spent Fig. 2. Dissolution of Co and Li as a function of time in aqueous mixture of tartaric acid
LIBs. (0.1–0.5 M) and ascorbic acid (0.02 M) at 80 °C.
56 G.P. Nayaka et al. / Hydrometallurgy 161 (2016) 54–57

Fig. 4. UV–vis spectra of the dissolved solution at different intervals of time showing the
Fig. 3. Effect of tartaric acid concentration on the dissolution of Co and Li from LiCoO2 at Co(II)–tartarate complex.
80 °C.

reduction of lattice Co3+ does not arise because the LiCoO2 used here is In order to recover the Co from dissolved solution, after complete
not sintered at high temperature to impose any lattice stability. Thus, dissolution, it was cooled to room temperature and stoichiometric
even in the absence of AA, complete dissolution occurred here with amount of oxalic acid was added. Since the Co(II)-tartarate is weak
tartaric acid. complex (low stability constant), it readily dissociated to form Co(II)-
As shown in Fig. 3, unlike Li, the Co released was about 95%. Such a oxalate precipitate. In fact, the main reason for choosing tartaric acid
discrepancy could be due to Li-depletion in the original sample. This is here is because of its weak complexation with Co ions (Gasser, 2014).
because, during its life-time charging/discharging process, fraction of Such a selective precipitation of Co(II)-oxalate from the mixture of Co
Li ions is irreversibly interacted with the anode (graphite). On the and Li ions in the dissolution mixture is highly advantageous for its
other hand, we have assumed nominal composition of LiCoO2 instead further processing as Co resource. The solid Co(II)-oxalate was separated
of Li1 − xCoO2. Also, there was about 5 wt.% carbon residue along with easily by filtration and washed with excess deionized water, dried in hot
the sample (due to organic burn-off) which is not taken into account air oven. Fig. 5 shows the XRD pattern and FTIR spectra (given as inset) of
for knowing the actual amount of metal ions released here. Further- this solid powder sample. The XRD patterns matches with JCPDS 25-0250
more, based on AAS analysis of dissolved samples, the results here confirming the well crystallized orthorhombic structure (Chen et al.,
carry about 5% errors. Thus, the % of dissolution obtained here is on 2011). The FTIR band ~3340 cm−1 has been assigned to O–H stretching
conservative side. In the previous studies, a similar leaching efficiency vibration and the band ~1613 cm−1 has been assigned to CO stretching.
is reported. For instance, N90% Co and 100% Li was obtained with The two peaks around 1317 cm−1 and 1200 cm−1 are due to the pres-
1.25 M citric acid (Li et al., 2010), 1.5 M malic acid (Li et al., 2013) and ence of carbonyl group (Sun and Qiu, 2011, 2012). A complete recovery
1.0 M oxalic acid (Sun and Qiu, 2011) using H2O2 as reducing agent of Co was ensured here.
over a period of 90 to 120 min. Although we have not studied the
variation of temperature, it is well established that the dissolution is 4. Conclusion
enhanced at with temperature (Sun and Qiu, 2011, 2012; Chen et al.,
2011). In the case of mineral acids, dissolution N 80 °C can volatilize An environmentally benign hydrometallurgical route for the dissolu-
the acid and induce accelerated corrosion of containers etc. The initial tion of active cathode material from spent Li-ion batteries (LIBs) is inves-
studies here have shown that there dissolution is negligible at room- tigated here. A mixture of tartaric acid and ascorbic acid was found to
temperature and the rate of dissolution increased almost linearly with completely dissolve sample obtained from the spent LIBs (BL-5CA Nokia
temperature. Hence, we have kept the maximum temperature of 80 °C series) at 80 °C in about 3–4 h. The UV–vis spectra of dissolved solution
(further increase will lead to loss of solvent during sampling through
boil off). Due to rapid dissolution at this temperature, the dissolution
kinetics is not determined here; unlike in our previous study (Nayaka
et al., 2015).
We have observed that the TA alone also resulted in complete dissolu-
tion. Fig. 4 shows the UV–vis spectra of dissolving mixture as a function of
time in presence and absence (inset) of AA. The absorbance of the solu-
tion increased as the dissolution time was increased. In the presence of
AA, all the Co(III)-tartarate was reduced to Co(II)-tartarate. The λmax
observed at 300 nm is attributed to Co(II)-tartarate in presence of AA. In
the absence of AA, the absorption band for Co(III)-tartarate is seen around
512 nm and Co(II)-tartarate is seen around 240 nm. The reduction of
Co(III)- to Co(II)-tartarate in the absence of AA is attributed to its oxidiz-
ing ability of water to stabilize as Co(II)-tartarate, a pale pink colored com-
plex. The purpose of using AA here is to ensure the reduction of all the
Co(III)- to Co(II)-tartarate for subsequent recovery of Co as Co(II)-oxalate.
Also, the AA helps to maintain de-aerated condition during dissolution
process due to its oxygen scavenging property. The concentration of AA
used here (0.02 M; H2A → 2 H+ +2e) is just stoichiometrically enough,
and hence the concentration was not varied. Fig. 5. XRD of CoC2O4 2H2O precipitate (inset is the FTIR spectra).
 G.P. Nayaka et al. / Hydrometallurgy 161 (2016) 54–57 57

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Acknowledgment 176, 288–293.
Li, L., Lu, J., Ren, Y., Zhang, X.X., Chen, R.J., Wu, F., Amine, K., 2012. Ascorbic-acid-assisted
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One of the authors (G.P. Nayaka) gratefully acknowledges the
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financial support from the Kuvempu University (Grant No.: 52/914). Li, L., Dunn, J.B., Zhang, X.X., Gaines, L., Chen, R.J., Wu, F., Amine, K., 2013. Recovery of
metals from spent lithium-ion batteries with organic acids as leaching reagents and
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