Application of Spent Li-Ion Batteries Cathode in Methylene Blue Dye Discoloration

This paper aims to present the mechanism study of methylene blue (MB) discoloration using spent Li-ion battery cathode tape and hydrogen peroxide.The recycled cathode used in this work is composed of 72% of LiCoO2, 18% of carbon, and 10% of Al.The value found for surface area is 8.9m/g and the ZCP value occurs in pH = 2.95. Different fromwhat is proposed in the literature, the most likely mechanism of methylene blue discoloration is the oxidation/delitiation of LiCoO2 and the reduction of H2O2 forming OH . Thus, in this paper, an important and promising alternative for discoloration of textile industry dyes using spent Li-ion battery cathode is presented.


Introduction
1.1.Li-Ion Battery Recycling: A Brief Overview.The fast popularization of portable devices such as cell phones is due mostly to the introduction of Li-ion batteries (LIBs) by Sony in 1991 [1][2][3].The introduction of LIBs to the market was due to characteristics such as high energy density, long lifespan, and lightweight [3,4].In general, LIBs have no Cd, unlike Ni-Cd batteries, and have much less Ni than Ni-MH batteries.The cathode material widely used in LIBs is lithium and cobalt oxide, LiCoO 2 [1][2][3]5].Due to their widespread use, LIBs' disposal is a topic of great relevance.The waste generated by LIBs will soon be a dramatic problem in countries such as China, where the quantity of discarded LIBs in 2020 can reach 25 billion units [6].Among many applications, it is reported that the spent LiCoO 2 catalyzes methylene blue degradation in the presence of H 2 O 2 [3].Thus, a very audacious project concerns the application of recycled materials from LIBs on waters contaminated with potentially toxic organic molecules.This process is very interesting concerning the environment due to the possibility of promoting the decontamination of organic pollutants using a residue coming from Li-ion battery waste.The substances traditionally used for organic decontamination are O 3 and H 2 O 2 [7].However, the cost of these compounds is relatively high (O 3 is $2.2 per kg and H 2 O 2 is $1.5 per kg) [7].The use of LiCoO 2 increases the velocity constant of methylene blue (MB) degradation 200 times if compared to pure H 2 O 2 with minimum additional cost since this material comes from spent LIB [3].Thus, the use of recycled LIB cathode is an excellent alternative for degradation of potentially toxic organic molecules.The literature reports hydrogen peroxide oxidation, promoted by LiCoO 2 reduction, with intermediate step correspondent to O 2 −• radical formation [3].However, we have reasons to believe that the methylene blue discoloration mechanism, using the H 2 O 2 and LIB spent cathode, is different from what is proposed in the literature [3].The main controversial point is that the reduction potential of LiCoO 2 in aqueous solution (see (1)) is around −0.95 V [4] and oxidation of H 2 O 2 (see (2)) is −0.89V [3].Thus, the global process (see (3)) is thermodynamically unfavorable (−1.84 V).Thus, the pathway to understand the reaction mechanism between LiCoO 2 and H 2 O 2 involves predominantly the study of the chemical behavior of LiCoO 2 in aqueous solution.
1.2.The Chemical Behavior of LiCoO 2 in Aqueous Solution.The LIB operating system basically consists of reversible extraction/insertion of Li + ions between two electrodes (see ( 4)) [1,8,9].For LIB with LiCoO 2 cathode, in the charge process, the Li + intercalated in the LiCoO 2 structure is forced out by an external applied potential in direction of a carbonaceous matrix (see ( 4)) [1].In the discharge (reverse process), the Li + ions return to the Li (1−) CoO 2 matrix due to a concentration gradient [1].Due to the large LIB operation potential, the electrolyte is a nonaqueous solution composed of lithium salts, such as 1 M LiPF 6 , in alkyl organic carbonates (ethylene, propylene, and dimethyl carbonates) [9].
Thus, to force the deintercalation of Li + ions from LiCoO 2 lamellar structure without an external power source, it is necessary to use an oxidant agent with reduction potential higher than 1.02 V.The reduction potential of H 2 O 2 /H 2 O [12] and O 2 /H 2 O [12] (considering  O2 = 0.21 atm) versus pH is shown in Figure 1.The dotted line represents the limit potential (1.02 V) for LiCoO 2 oxidation in aqueous solution (see (5)).The redox reaction represented by LiCoO 2 oxidation and the reduction of H 2 O 2 (see (6)) are spontaneous along the pH range 1 to 12.The potential for (6) considering pH =7 is 0.3 V (details in Figure 1).The interference of parallel O 2 reduction is observed only in pH <1 (Figure 1).
On the other hand, for acid pH values, the cointercalation of H + ions in the charged Li (1−) CoO 2 (see ( 7)) [13] normally promotes the obstruction of Li + diffusion pathways [11].
Thus, this paper aims to present the mechanism study of methylene blue (MB) discoloration using LIB spent cathode tape and hydrogen peroxide.The spent cathode tape (Eq. ( 6)) was characterized by X-ray diffraction, atomic absorption spectrophotometry (AAS), zeta potential, SEM, and the MB adsorption technique to determine the surface area.

The Characterizations of LIB Spent
Cathode.The LIB used in this paper was a spent cell phone battery.The LIB was dismantled using the manual procedure and the components were separated (Figure 2).The spent cathode in tape form was

Determination of Surface Area.
A successful method to determine the active area was described in the literature in recent years.It is based on the adsorption of MB [14].The measure of the surface area with MB adsorption has the main advantage of measuring the surface accessible to electrolyte.
To determine the surface area, the LIB spent cathode was cut into rectangular pieces of 4 cm 2 (∼60 mg) and put into a MB (Sigma-Aldrich, dye content ≥95%) solution with different concentrations (1, 2, 3, 5, 7, and 10 ppm) in different pH values (between 3 and 11).After 24 h, the MB concentration was monitored using a UV-spectrophotometer FEMTO Cirrus 80 PR at 659 nm.The MB dye adsorption capacity (  ) was determined by (8) [15], in which  and   are the initial and equilibrium MB concentration, respectively.For the determination of the surface area, the Langmuir isotherm (see ( 9)) was considered, in which  max is the monolayer adsorption capacity (mg/g) and   is a constant related to free energy of adsorption [15].To convert  max into surface area (m 2 /g), each MB molecule with 1,70 × 10 −18 m 2 [14] was considered.
The active radical was identified by the addition of 2 M isopropanol (i-PrOH) as a scavenger for OH • [16].The cyclic experiment was performed using the same spent cathode in 5 cycles of MB discoloration (600 min).All measurements were made in the absence of visible light to avoid the possible electron excitation of cobalt lithium oxide conduction band to the valence band generating the electron (BC) and hole (BV) pair.After MB discoloration, the resultant solution was analyzed with AAS aiming to determine the possible lixiviation of Co and Li.diffraction and AAS were used to elucidate the chemical composition of LIB spent cathode.The spent cathode powder was dissolved (Materials and Methods) and the resulting solution was analyzed by AAS.The presence of Fe, Mn, and Ni was not detected (detection limit: 0.04 ppm).The Li and Co concentrations were 55 ppm and 418 ppm, respectively.Considering the atomic mass of Co and Li (58.933 and 6.941 g/mol [12]), these concentrations are compatible with a stoichiometric proportion of approximately 1 : 1.The X-ray of LIB spent cathode tape was analyzed with Rietveld method using the free software FullProf© [17] (Figure 3).The X-ray diffraction showed the proportion of phases present in the cathode.The recycled cathode is composed of 72% of LiCoO 2 , 18% of carbon, and 10% of Al.Carbon is added to increase the electrical conductivity of the cathodic material [1].The Al presence is due to the current collector composition and its detection in X-ray diffraction is possible due to the porosity of the LIB cathode [1].It is worth pointing out that the presence of other Li  CoO 2 phases (with  varying from 1 to 0.5) is also possible, but their detection through X-ray diffraction is complex when 0.5 <  < 1 [11].

Results and Discussion
Figure 4 shows the Scanning Electron Microscopy (SEM) of the spent cathode from Li-ion battery.The values found for the porosity and the average size of grain were 42% and 31 m, respectively [18].
To clarify the adsorption process, the zeta potential was assessed (Figure 5).The zeta potential remains practically constant (∼−40 mV) in the pH ranging from 11 to 6. Considering the values of pH ranging from 6 to 3, the zeta potential varies considerably until it reaches the zero charge point (ZCP).The spent cathode has relatively acidic ZCP (close to pH = 2.95).
To determine the surface area, Figure 6(a) shows the   /  versus   plot for MB adsorption onto the spent LIB cathode.The adjustment of experimental points in Langmuir isotherm (see ( 9)) showed good concordance ( 2 = 0.99).Similar to the zeta potential measurement, the surface area measured by MB adsorption undergoes a variation with pH (Figure 6(b)).For values of pH ranging from 11 to 6, the surface area changes minimally.The value found for surface area based on MB adsorption was 8.9 m 2 /g in pH = 7.With the pH ranging from 6 to 3, the MB adsorption decreases quite probably due to the H + adsorption competition.

Methylene Blue Discoloration Mechanism.
To obtain the reaction order and the reaction determinant step (RDS), MB discoloration was performed in different conditions as shown in Figure 7. Table 1 shows  50 in different situations.The reaction orders obtained in relation to H 2 O 2 , SC (spent cathode), and MB are, respectively,  H 2 O 2 = 0.04 ≈ 0,  SC = 1,008 ≈ 1, and  MB = 1.102 ≈ 1.Thus, the reaction determinant step can be written as To prove that the reaction mechanism can be represented by ( 6), the pH and Co and Li concentration were measured before and after the discoloration reaction.The experimental design is given by Table 1 in condition 3.In the previous solution, the concentrations of Co and Li were below the detection limit (0.04 ppm).After MB discoloration, the Co concentration remained below 0.04 ppm and the Li concentration reached 1.2 ppm.Moreover, after the discoloration reaction, the pH value ranged from 7.01 to 7.25.It is important to note that, for charged Li (1−) CoO 2 phase in aqueous solution, the Li + intercalation (which corresponds to the LIB discharge process) only must be considered favorable in special conditions as reported in the literature (e.g., in 5 M LiNO 3 ) [10].Thus, the only spontaneous process for LiCoO 2 under the experimental conditions used in this work is the oxidation process shown by (6).With the decrease of Li content in the cathode (LiCoO 2 ), there is also a decrease in the kinetics of MB discoloration.This is easily seen in the cyclic test performed with the same LIB spent cathode tape (Figure 8).A slight drop of discoloration capacity is noticed, which can be evidence that the spent cathode is a reagent and not a catalyst as reported previously in the literature [3].Thus, the charged phase Li (1−) CoO 2 reacts in the oxidation direction by The X-ray diffraction was performed on the spent cathode after degradation of methylene blue.The main objective was to verify the presence of Co This reaction takes place at around 190 ∘ C [19].In this work, the most probable process involving LiCoO 2 is delitiation (LIB charge); in addition, the temperature of the reaction system did not exceed 25 ∘ C. Thus, after the methylene blue discoloration reaction, Co 3 O 4 presence is not expected.Figure 9 shows the positions of the main peaks of Co 3 O 4 phase (JCPDS 78-1970 [13]).In Figure 9, we observe only the presence of LiCoO 2 crystallographic planes.Thus, the presence of Co 3 O 4 phase in the spent cathode after degradation of methylene blue is not confirmed by X-ray diffraction.It is noted that, in the spent cathode after MB discoloration reaction, the interplanar distance undergoes a subtle decrease, showing that the H + cointercalation also is not evidenced [11].
To confirm the mechanism, the active radicals were identified by the addition of isopropanol (i-PrOH) as a scavenger for OH • [16].Figure 10 shows the discoloration of 12 ppm MB with 0.1 M H 2 O 2 and spent cathode with and without 2 M i-PrOH.The measure shows that OH • is the predominant radical formed by H 2 O 2 reduction (see (13)) promoted by LiCoO 2 oxidation.

Conclusion
In this paper, we presented an important and promissory alternative for discoloration of textile industry using LIB recycled cathode.The model molecule used was methylene blue.The recycled cathode is composed of 72% of LiCoO 2 , 18% of carbon, and 10% of Al.The value found for surface

3. 1 .Figure 3 :
Figure 3: The X-ray diffractogram of LIB spent cathode tape before MB discoloration reaction.The proportion of phase was analyzed with Rietveld method using the free software FullProf©.

Figure 5 :
Figure 5: Zeta potential of LIB spent cathode powder.

Figure 6 : 6 Figure 7 :
Figure 6: (a) The   /  versus   plot for MB adsorption onto spent LIB cathode.(b) The surface area measured by MB adsorption with pH variation.

Figure 8 :
Figure 8: The result of five cyclic tests of MB discoloration.The experimental conditions used were MB 12 ppm, spent cathode 4 cm 2 , and 0.1 M H 2 O 2 .

Figure 10 :
Figure 10: The discoloration of 12 ppm MB with 0.10 M H 2 O 2 and spent cathode with and without 2 M i-PrOH.

Table 1 :
The experimental conditions used in the reaction order study of MB discoloration.The variables were spent cathode (SC) area and H 2 O 2 and MB concentration.

)
[3])Methylene Blue Discoloration Mechanism.To study the kinetics of MB discoloration, a 100 mL glass cell containing 50 mL of MB solution, H 2 O 2 , and LIB spent cathode tape (cut into rectangular pieces) was stirred magnetically (10 rpm) in a bath with controlled temperature at 25 ∘ C. Before the H 2 O 2 addition, the MB solution was put in contact with LIB spent cathode tape for 24 hours to establish the adsorption equilibrium.The pH of the solution was adjusted at around 7 (exactly 7.02).The discoloration process was monitored by analyzing aliquots at 659 nm.The reaction order was evaluated by variation of spent cathode area and H 2 O 2 and MB concentration.The variations of the parameters are shown in Table1.Considering  50 as the interval of time to promote 50% of discoloration, the reaction order (  ) in relation to generic reactant "" can be obtained by(10)[3]in which  50 and   50 are the interval of time to promote 50% of discoloration in two different concentrations represented by [] and []  , respectively.