Synergy of Lithium, Cobalt and Oxygen Vacancies in Lithium Cobalt

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Synergy of Lithium, Cobalt and Oxygen Vacancies in Lithium Cobalt Oxide for Airborne Benzene Oxidation: A Concept of Reusing Electronic Wastes for Air Pollutant Removal Tianchi Dai, Hao Zhou, Yang Liu, Ranran Cao, Jingjing Zhan, Lifen Liu, and Ben W. L. Jang ACS Sustainable Chem. Eng., Just Accepted Manuscript • DOI: 10.1021/ acssuschemeng.8b05894 • Publication Date (Web): 09 Feb 2019 Downloaded from http://pubs.acs.org on February 9, 2019

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ACS Sustainable Chemistry & Engineering 1

Synergy of Lithium, Cobalt and Oxygen Vacancies in Lithium Cobalt Oxide for Airborne Benzene Oxidation: A Concept of Reusing Electronic Wastes for Air Pollutant Removal Tianchi Dai a, Hao Zhou a, Yang Liu a*, Ranran Cao c, Jingjing Zhan a, Lifen Liu a, Ben W.-L. Jang b* a. Key Laboratory of Industrial Ecology and Environmental Engineering (Ministry of Education), School of Food and Environment, Dalian University of Technology, Panjin 124221, China b. Department of Chemistry, Texas A&M University-Commerce, PO Box 3011, Commerce, TX 75429, USA c. School of Environment, Tsinghua University, Beijing 100084, China --------------------* Corresponding authors. Tel.: +86-427-2631799 (Y. Liu), Fax: +1-903-468-6020 (B.W.-L. Jang). E-mail addresses: [email protected] (Y. Liu), [email protected] (B.W.-L. Jang).

Abstract In this study, a recovery strategy of turning cathode materials of waste lithium-ion batteries, lithium cobalt oxide (LiCoO2), into high-performance catalysts for the oxidation of airborne benzene is investigated. Part of the Co3+ in LiCoO2 was leached out from the [CoO6] octahedra by HNO3 solution via the disproportion reaction with the deintercalation of interlayer Li+ ions, resulting in the formations of Li, Co and oxygen vacancies. The presence of Li and oxygen vacancies facilitated gaseous benzene adsorption and subsequent activation of adsorbed benzene, and the Co vacancies together with the oxygen vacancies induced the generation of plenty of

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active surface oxygen species, accordingly the catalytic activity of the acid-modified LiCoO2 was greatly improved with good resistance to the alternative “heating-cooling” operation and stability from 0.8 to 2.3 vol.% of humidity. More importantly, the acid modification process is GREEN since the HNO3 solution could be reused over times to produce effective catalysts and the final solution containing high concentrations of Li+ and Co2+ ions could be potentially recycled for cathode material manufacturing. The acid treatment method also worked effectively with the commercial LiCoO2. This study would inspire the development of novel and sustainable catalytic systems for environmental applications.

Keywords: Electronic waste; Lithium cobalt oxide; Catalytic oxidation; Benzene; VOCs


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Introduction "e-waste" commonly refers to the electrical or electronic equipment that is discarded or is no longer used.1 The type of e-waste that we focus on here is the complex metal oxide nanomaterials that compose lithium-ion (Li-ion) battery cathodes, widely used in electronic products such as computer, mobile phone and electric vehicle, etc. However, at the end of their life, improper disposal of the complex metal oxides can cause structural transformations such as dissolution and metal leaching, resulting in a significant exposure risk to the surrounding environment and health.2 On the other hand, it will be a much better choice if we can turn waste into “treasure” and recycle metal ions at the same time. Benzene is one of the commonly found volatile organic compounds (VOCs) in polluted air, recognized as hazardous contaminant. Due to its “teratogenesis, carcinogenesis and mutagenesis”,3-5 benzene is posing a serious threat to human health. Moreover, the residual benzene is detected in animals, plants, soil and water bodies, and this potential harm will last for generations since benzene is refractory and difficult to decompose.5 Among the several classes of chemical and physical reactions, catalytic oxidation at elevated temperatures is considered to be an effective and economical technique for benzene removal because of its low energy consumption, simple operation and high purification efficiency.6 So far, transitional metal oxides have been extensively studied for benzene oxidation as a result of their relatively low prices and a broad range of redox capacity.7-12 As reported in literature,13 Co oxide-based materials were effective for the abatement of VOCs due



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to their weaker Co-O bond and higher oxygen binding rate, thus attracting attention from many researchers. For example, Chen et al.14 synthesized a CoMnAl solid solution via co-precipitation method and found that the 550 oC-treated sample exhibited ∼90% of removal efficiency for 100 ppm of benzene under a space velocity of 60 L·g-1·h-1 at 208 °C, which is attributed to the large surface area (102 m2·g-1 vs. 85 and 34 m2·g-1 of single CoAlO and MnAlO samples, respectively) and abundant oxygen vacancies. They also found that the abundant surface Co3+ species could provide sufficient surface adsorbed oxygen species and acid sites, and promote low-temperature reducibility, which together accounted for the excellent benzene oxidation performance of CoOx-based catalysts.15 In this study, we chose the cathode material of Li-ion batteries, to say lithium cobalt oxide (LiCoO2), as the potential benzene oxidation catalyst, and presented the impacts of acid modification on its physiochemical and catalytic properties. It is the first time to apply LiCoO2 to the field of air pollutant remediation. Various characterizations were performed, demonstrating that the leaching effect of HNO3 solution resulted in the formations of Li, Co and oxygen vacancies, which synergistically improved the catalytic performance of the LiCoO2, and the effect of cation vacancy on catalytic oxidation of benzene has never been reported before. More importantly, the acid modification process was GREEN since the HNO3 solution could be recycled over times without lowering the catalyst activity and the final solution containing high concentrations of Li+ and Co2+ ions could be potentially recycled for cathode material manufacturing.


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Experimental The pristine LiCoO2 was prepared by thermally reacting the mixture of Li2CO3 and CoCO3 powders at 500 oC for 6 h in a furnace. The pristine LiCoO2 powder was further treated in HNO3 aqueous solution for modification, and the acid-treated samples were denoted as H-LiCo-t, where t indicates the period of acid treatment. For synthetic details, please refer to the supporting information (SI). For details of the instruments, procedures and parameters used for catalyst characterizations, catalytic benzene oxidation and temperature programmed desorption of adsorbed benzene (C6H6-TPD), see the SI. Results and discussion Catalytic activity evaluation 100

Benzene Conversion (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Sustainable Chemistry & Engineering

80

pristine LiCoO2

60

H-LiCo-2 H-LiCo-4 H-LiCo-6 H-LiCo-8 Li2O

40

CoOx LiCoO2-800

20 0 150

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350

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Temperature (oC)

Fig. 1. Temperature dependences of benzene conversion over the pristine LiCoO2 and H-LiCo catalysts. Benzene oxidation conditions: 200 mL·min-1 of synthetic air containing 370-430 ppm of benzene. First, benzene oxidation over the pristine LiCoO2 and H-LiCo catalysts was comparatively studied, and Fig. 1 depicts the conversion for benzene oxidation as a function of reaction temperature. Acid treatment of LiCoO2 could significantly enhance its benzene oxidation activity regardless of the treatment period. The pristine



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LiCoO2 was almost inactive toward benzene oxidation when reaction temperature was lower than 200 oC, and then benzene conversion was gradually increased with reaction temperature. However, even at a high temperature of 400 oC, only ~30% of benzene could be degraded. The H-LiCo-2, which was obtained by treating the pristine LiCoO2 powder in HNO3 solution for 2 h, exhibited ~13% and ~28% of removal efficiencies for 377 ppm of benzene at 175 oC and 200 oC, respectively, and ~91% of benzene could be destructed at 300 oC. When the acid treatment period was prolonged to 4 h, 6 h and 8 h, benzene conversion was further increased. However, there was no obvious difference in the low-temperature (≤ 200 oC) activities among H-LiCo-4, H-LiCo-6 and H-LiCo-8 samples whereas the H-LiCo-6 catalyst possessed relatively higher benzene conversions at 250 oC and 300 oC than the other three H-LiCo samples by 5%-9%. Although significant progress has been made in benzene oxidation over acid-modified LiCoO2 catalysts, the activity of the H-LiCo-6 catalyst still needed improving, especially when compared with our previously reported Cu2+-exchanged birnessite MnO2,3 which exhibited superior activity over the samples of similar type. Considering how challenging it is to break the refractory chemical bonds of benzene molecule, the H-LiCo-6 sample in this study was still promising for benzene oxidation in the gas phase. This observation highlights the possibility of using the acid modification method to transform waste Li-ion batteries into useful catalytic materials for removing air pollutants. In view of the outstanding catalytic performance of the H-LiCo-6, it was used for further studies thereafter.


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Water vapor is usually found detrimental to catalytic benzene oxidation16,17 due to the competitive adsorption of water with benzene as well as oxygen molecules.18 In this study, the H-LiCo-6 catalyst could stably work from 0.8 vol.% to 2.3 vol.% of moisture and exhibited good resistance to temperature fluctuation (Fig. S1A). In addition, benzene oxidation over the H-LiCo samples obtained with different washing times (Fig. S1B), recycling used acid solutions (Fig. S1C) and the effect of space velocity (Fig. S1D) were also discussed in detail in the SI. Catalyst characterizations Crystal structure, textual properties and elemental distribution







pristine LiCoO2

H-LiCo-6

B

pristine LiCoO2

H-LiCo-6

590

Intensity (a.u.)

{104}



{113}



{107}



 LiCoO2 (44-0145) {015}

{101}

{003}

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Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


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Raman Shift (cm-1)

Fig. 2. (A) XRD and (B) Raman spectra of the pristine LiCoO2 and H-LiCo-6 catalysts. The XRD patterns of the pristine LiCoO2 and H-LiCo-6 catalysts are shown in Fig. 2A. Legible diffraction peaks of the standard LiCoO2 (JCPDS No. 44-0145) were observed over the pristine LiCoO2 without detectable impurities, indicating the successful synthesis of LiCoO2 as the reference material. The facet indexes were also marked on the corresponding diffraction peaks. However, the two diffraction bands around 39o and 66o were a little different from the standard LiCoO2 pattern, in which



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each of these two peaks is divided into two distinct diffraction peaks. The possible reason is that 500 oC was not high enough to fully crystalize LiCoO2, testified by the well-indexed peaks of the 800 oC-treated LiCoO2 (LiCoO2-800) to the standard pattern. However, the well-crystallized LiCoO2-800 shows an inferior benzene oxidation activity to the 500 oC-treated sample (Fig. 1). Acid modification did not change the diffraction pattern of LiCoO2, but the relative intensities of the XRD peaks over the H-LiCo-6 were different from those over the pristine sample, which is ascribed to the partial destabilization of the layered structure via the deintercalation processes of Li+ ions by acid treatment19 (see later discussion). Raman spectra of the pristine LiCoO2 and H-LiCo-6 catalysts are exhibited in Fig. 2B, providing complementary information to the XRD results. Consistent with the Raman spectra of LiCoO2 reported in literatures,20,21 vibrational bands of the pristine LiCoO2 at 590 and 486 cm-1 are assigned to the symmetric Co-O stretching of the [CoO6] octahedra and the bridging Co-O-Co vibration, respectively. Acid treatment led to remarkable alteration of the Raman pattern of LiCoO2, no Raman peaks being detected over the H-LiCo-6 sample, but the XRD pattern of the H-LiCo-6 catalyst was identical to that of the pristine sample (Fig. 2A). This seemingly contradictory phenomenon is resulted from the fact that overall the H-LiCo-6 catalyst exhibited the characteristics of LiCoO2 phase and this long range order could be reflected by XRD; in the meanwhile, the spatial structure was damaged by the presence of crystal defects (such as cation vacancies and oxygen vacancies, later discussion) induced by acid modification, resulting in the significant reductions of the


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Raman peaks over the H-LiCo-6 sample. Song et al. investigated thermal reduction in hydrogen on the properties of ZrO2 and they ascribed the broader and weaker Raman peaks after hydrogen treatment to the formation of huge oxygen vacancies.22 The changes of particle size and morphology before and after acid modifications were investigated by SEM, with the SEM images displayed in Fig. S2. From the macroscopic view, the pristine sample was mainly composed of spherical particles (Fig. S2A) in agglomerated state (Fig. S2B) while chunks (Fig. S2D) consisted of smaller particles (Fig. S2E) were formed after acid treatment. Further enlargement of the SEM images (Fig. S2C and F) allowed us to find that the particles were in polygon structure regardless of whether the catalyst had been treated by acid or not, and the statistics of the particle size distribution could also be made accordingly. As displayed in Fig. S2C, particles of tens of nanometers, over 100 nm and even larger size (circled by green dashes) were clearly seen over the pristine LiCoO2. In stark contrast to the pristine LiCoO2, particles were much more uniformly distributed over the H-LiCo-6 sample (Fig. S2F), mainly in the range of 25-55 nm as seen from the inset histogram. These smaller particles rendered the H-LiCo-6 catalyst with a relatively larger BET surface area (28 m2·g-1) with respect to the pristine sample (15 m2·g-1), thus increasing the contact probability of the gaseous reactants with the acid-treated surface. Moreover, the acid treatment time barely influenced the final BET areas of the acid-treated samples, which were 24, 23, 28, 26 m2·g-1 for the H-LiCo-2, H-LiCo-4, H-LiCo-6 and H-LiCo-8 catalysts, respectively. The LiCoO2-800 sample possessed much poorer BET area of only ~7 m2·g-1, probably due



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to the sintering of particles at high temperature; therefore, the LiCoO2-800 sample exhibited much lower catalytic activity (Fig. 1). The BET-normalized reaction rates, which were more intrinsic to evaluate the catalytic efficiencies of different samples,23 were then calculated. Take the values obtained at 150 oC and 175 oC for example, as shown in Fig. S1E, the specific benzene reaction rates of the H-LiCo-6 catalyst were significantly more than doubled those of the pristine sample. Thus, it could be concluded that the benzene removal efficiency was not directly dependent on the specific surface area of catalyst in this study.


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Fig. 3. (HR)TEM images of the pristine LiCoO2 (A-C) and H-LiCo-6 (D-F) catalysts. The particle morphologies and exposed facets were observed with (HR)TEM, with the images depicted in Fig. 3. From Fig. 3A and D, the H-LiCo-6 catalyst (Fig. 3D) presented much more dispersed particles than the pristine LiCoO2 (Fig. 3A), unanimous with the SEM observations (Fig. S2C and F). The particle profiles are much more clearly shown in Fig. 3B and E, the pristine LiCoO2 displaying a polygon structure (Fig. 3B). Although it appears like that the particles of the H-LiCo-6 material were irregularly shaped (Fig. 3E), the polygon structure was also retained after acid modification as indicated by the corresponding inset image. From Fig. 3F, the surface of the H-LiCo-6 nanocrystals became disordered, forming the disorder layer of about 2 nm in thickness (white dashed lines). Nevertheless, there was no obvious disorder layer on the surface of the pristine sample (Fig. 3C), so the acid modification could be termed as acid-etching, and the thus-generated disorder layer might be resulted from the acid-induced crystal defects. In terms of exposed facets, edge-sharpened lattice fringes of 0.46 nm were observed on the rim of the pristine particle (Fig. 3C), corresponding to the LiCoO2



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{003} plane, which, however, gradually turned to be blurred and then intersected with the fringes in another direction (red dashed lines) as pointed by the left green arrow; finally, these two distinguishable intersected fringes were integrated into dot-shaped fringes (yellow dashed square) as indicated by the right green arrow, and the lattice space calculated from the Fast Fourier transform (FFT) image is also assigned to the LiCoO2 {003} plane (inset). This phenomenon might be due to the overlapping of two different crystals with the same exposed facet. In addition to the {003} facet, the LiCoO2 {104} plane of 0.20 nm was also detected but on another crystal separated by the purple curve, which seemingly grew below the former {003} surface, further testifying the co-growth of crystals over the pristine LiCoO2, and this would lead to the formation of large particles with relatively small surface area. There hardly existed any edge-sharpened lattice fringes over the acid-treated surface probably due to the formation of crystal defects, thus making the ordered lattices rearranged and disordered. Take Fig. 3F for example, the LiCoO2 {003} plane was also exposed over the H-LiCo-6 sample as calculated from the FFT image (inset), whereas the corresponding lattice space (= 0.48 nm) was increased with respect to that of the pristine sample (= 0.46 nm), which presented the same trend as their catalytic activities (Fig. 1). Since the {003} plane is parallel to the sheets of [CoO6] octahedra of LiCoO2,24 the expansion of the {003} lattice space is likely owing to the increasing interplanar distance from the severe deintercalation of interlayer Li+ ions,25 which will be discussed in the next section. Also shown in the inset FFT image of Fig. 3F, a new lattice spacing of 0.29 nm was formed over the H-LiCo-6 sample, which, however,


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does not belong to any of the crystal facets of LiCoO2. Further investigation reveals that CoOx possesses the crystal plane with a spacing of ~0.29 nm, e.g., Co3O4 {220} plane (0.286 nm, JCPDS No. 42-1467) and Co2O3 {002} plane (0.287 nm, JCPDS No. 02-0770). Moreover, based on the following XPS results, Co+3+ζ was generated over the H-LiCo-6 catalyst. Therefore, the lattice fringe of 0.29 nm can be ascribed to one of the crystal planes of the Co+3+ζOx. No integrated fringes by two or more overlapping crystals were observed over the H-LiCo-6, thereby resulting in the higher dispersion of particles with relatively large surface area. Deintercalations of Li+ and Co3+ ions by acid treatment induced the formations of lithium, cobalt and oxygen vacancies 4

A

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HNO3

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Co(NO3)2

Co2+

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H-LiCo-6 LiNO3

1.0

Absorbance

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


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pristine LiCoO2

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H Chemical Shift (ppm)

Fig. 4. (A) UV-vis spectrum of the supernatant after acid treatment for 6 h and the photo of the supernatant was inserted in the inset; (B) 1H-NMR spectra of the pristine LiCoO2 and H-LiCo-6 catalysts. After HNO3 treatment, the suspension was centrifuged. The supernatant after centrifuging was in pink color (seen from the photo in Fig. 4A), indicating that metal ions were leached out from the catalyst by acid. As shown in Fig. 4A, the UV-vis spectrum of the supernatant after centrifuging displays two adsorption bands centered at 300 nm and 510 nm, respectively. The 300 nm band is likely due to the NO3- ion



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since the individual solution of HNO3, Co(NO3)2 and LiNO3 all contributed to this band. The LiNO3 solution was colorless whereas the Co(NO3)2 solution was in pink color. What is more, as seen from the enlarged view in the inset, in contrast to the individual solution of HNO3 and LiNO3 that exhibited no adsorption around 510 nm, there was one legible adsorption band centered at 510 nm with the Co(NO3)2 solution. Conclusion could be thereby drawn from the above discussion that Co2+ was washed out from the LiCoO2 catalyst during acid treatment through the disproportion reaction of Co3+. The resulting higher valence of Co element over the H-LiCo-6 catalyst was confirmed by the following XPS measurements. According to the ICP-OES results, the weight contents of Li and Co of the pristine LiCoO2 and H-LiCo-6 materials were 6.1 wt.% & 60.8 wt.% and 2.5 wt.% & 47.4 wt.%, respectively. The atomic ratio of Li/Co was thus calculated to be ~0.86 over the pristine LiCoO2 and ~0.45 for the H-LiCo-6 sample. About 59% of Li+ ions were washed out from the pristine LiCoO2, while only ~22% for Co3+ ions. According to literatures, the layered structure of LiCoO2 was represented by different models.26,27 In this study, since Li+ ions were much more easily washed out than Co3+ ions, it is thereby inferred that Li+ ions existed in the interlayer positions while Co3+ ions were bonded with oxygen in the [CoO6] unit. Wang et al.28 also reported that the interaction between interlayer K+ ions and oxygen was not so strong as the Mn-O bond of [MnO6] octahedra over birnessite MnO2. The structural technique of NMR was also used to explore the chemical composition evolution during acid treatment, and the 1H-NMR spectrum investigation


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was used to analyze the proton-containing impurities at 80 oC, with the results displayed in Fig. 4B. The 1H-NMR curve of the pristine sample could be fitted into three bands, demonstrating the existence of hydrogen atoms in three different chemical environments. Surface adsorbed H2O ([H2O]ads) and hydroxyl group ([OH]ads) are the typically found hydrogen-containing species over metal oxides, and in our previous study,19 the hydrogen element in [H2O]ads appeared at higher chemical shift than that in [OH]ads. Besides, according to literature,29 there existed two kinds of hydroxyl groups over MnO2, namely Mn-OH and those ordinary ones directly bound to catalyst surface. Hence, the fitted band marked with P1 in Fig. 4B is assigned to the hydrogen element in [H2O]ads, and the P2 and P3 peaks are attributed to the hydrogen atoms in cation-bound [OH]ads and ordinary [OH]ads, respectively. The area ratio of NMR peaks is equal to the atomic ratio of the hydrogen atoms in corresponding chemical environments. After acid treatment (H-LiCo-6), the P1 peak increased while the P3 peak was barely influenced, thus excluding the H+-Li+ exchange, otherwise the intensities of both the P1 and P3 peaks should have been increased over the H-LiCo-6 catalyst due to the formation of [H3O]+ and [OH2]+. And the Li+ removal process might be dominated by Co oxidation process, i.e., disproportion of Co3+. A more straightforward evidence showing this mechanism is given by the above UV-vis measurements (Fig. 4A): while the original HNO3 solution showed no existence of dissolved Co2+, the solution used for 6-h acid treatment indeed showed significant increase in Co2+ concentration. Based on the assumption that the average oxidation



state of Co over the H-LiCo-6 was +3+ζ (ζ > 0) and the fact that the atomic ratio of Li/Co was 0.45 over the H-LiCo-6 (ICP-OES), the disproportion reaction of Co3+ could be written as follows: (4+4ζ)LiCoO2 (s) + (2.2+12ζ)H+ (l) → 4Li0.45CoOx (s) + 4ζCo2+ (l) + (2.2+4ζ)Li+ (l) + (1.1+6ζ)H2O (l) The acid-induced disproportion reaction was also observed over α-MnO2.30 The disproportion reaction of Co3+ resulted in both the formation of [H2O]ads and losses of Li+ and Co3+, therefore, resulting in the increase of the P1 peak and decrease of the P2 peak. A Li 1s

B

pristine LiCoO2

Co 2p

Intensity (a.u.)

H-LiCo-6

Intensity (a.u.)

B.E. = 1.4 eV

H-LiCo-6 Co 2p3/2 Co 2p1/2

satellite

satellite

pristine LiCoO2

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Binding Energy (eV)

C O 1s O

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Intensity (a.u.)

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H-LiCo-6

O

540

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531

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Binding Energy (eV)

Fig. 5. XPS spectra of the pristine LiCoO2 and H-LiCo-6 catalysts. The elemental composition and chemical states of the catalyst surface were


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characterized by XPS, with the results shown in Fig. 5. One peak appeared at 53.9 eV in the Li 1s spectrum of the pristine LiCoO2 (Fig. 5A), conforming to the reported binding energy for LiCoO2.31 After acid modification, not only the peak position downshifted to 52.5 eV, suggesting the increase in electron density around Li atoms, but also the peak intensity was greatly reduced over the H-LiCo-6 catalyst due to the above-mentioned leaching of Li+ ions. The decreased Co content in the H-LiCo-6 (ICP-OES results) might also cause its Li 1s peak to shift to lower binding energy based on the report of Schulz et al.31 that the Li2O appeared at lower binding energy than LiCoO2 in the Li 1s spectrum. In the Co 2p spectra (Fig. 5B), the starting compound LiCoO2 shows a Co 2p3/2 main line at 779.6 eV with a satellite peak at 789.4 eV and a Co 2p1/2 main line at 794.7 eV with a satellite peak at 804.7 eV. This result indicates that Co element was in trivalent state,32 which is in good agreement with the chemical formula of LiICoIIIO2. Besides, the Co 2p peaks were not symmetrical, and they also displayed shoulders toward higher binding energies as seen from the fitted peaks marked with dashed curves and corresponding insets. This feature could not be explained by a small amount of Co2+ ions because Co2+ ions in oxygen environment are usually characterized by the strong broadening of the main lines and the appearance of very intense satellite peaks at 786 eV (Co 2p3/2) and 803 eV (Co 2p1/2),33 which were not observed here. The deintercalation of Li+ ions did not change the Co 2p spectrum too much, and the overall profiles of the Co 2p core peaks were retained over the H-LiCo-6 sample but with two noticeable modifications in comparison with the



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pristine LiCoO2: (i) decrease of the peak intensities, including the main peaks and satellite peaks, and (ii) broadening of the main peaks toward higher binding energies as seen from the insets. On one hand, the reduction of the peak intensities of Co 2p is owing to the deintercalation of Co3+ ions by acid treatment (ICP-OES results); on the other hand, according to Dahéron et al.,33 the decrease of the satellite area together with the broadening of the main peaks could be explained by the partial oxidation of Co3+, which is also supported by the shift of Co 2p main peaks to lower binding energies over the H-LiCo-6 as shown in the insets.20 By comparing the Co 2p spectra with the Li 1s spectra, the deintercalation of Li+ ions was much severer than that of Co3+ ions, consistent with the ICP-OES results. The O 1s spectra are presented in Fig. 5C, and two characteristic peaks were detected over both samples. For the pristine LiCoO2, the Oβ and Oα peaks, evolved at 529.3 eV and 531.4 eV, respectively, are characteristic of the oxygen species from lattice oxygen and surface adsorbed oxygen, respectively.34,35 Three modifications in the O 1s spectrum were noticed over the H-LiCo-6 with respect to the pristine LiCoO2: (i) downshifts to lower binding energies, (ii) peak broadening and (iii) increase of peak area ratio of Oα to Oβ from 0.38 to 0.69. This peak broadening could be interpreted by a partial oxidation process of O2- ions,33 together with the Co 2p results confirming the hypothesis of literatures that upon deintercalation of Li+ ions, the valence charges of the lost Li+ ions were transferred to both cobalt and oxygen,36,37 resulting in the formation of Co+3+ζ and O-2+δ (ζ, δ > 0). For metal oxides, oxygen molecules are usually adsorbed at their oxygen vacancy sites;38 therefore, the


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relatively higher concentration of surface adsorbed oxygen (Oα/Oβ) over the H-LiCo-6 represents the formation of larger amount of oxygen vacancies over the acid-treated surface. Finally, in terms of the C 1s spectra (Fig. S3), the main difference lies in the disappearance of the CO32- peak over the H-LiCo-6 catalyst due to the reaction between CO32- and H+ during HNO3 treatment. Summarizing the ICP-OES and XPS results, the losses of Li and Co caused by acid modification would generate Li and Co vacancies, and more oxygen vacancies were formed over the acid-treated surface. Since much more Li+ ions were leached out than Co3+ ions, more Li vacancies were generated over the H-LiCo-6 sample. Both of the cation vacancies and oxygen vacancies were the catalytically active sites in oxidation reactions due to their ability to activate oxygen species, and this behavior was further studied by the following H2-TPR and O2-TPD experiments.

Fig. 6. Conceptual schematic of recycling waste Li-ion batteries for the synthesis of acid-modified benzene oxidation materials. In this study, the reference LiCoO2 material was prepared by thermally treating



the mixture of Li2CO3 and CoCO3. For practical application, the LiCoO2 could be obtained from the abandoned electronics. Based on the modification method proposed in this study, the conceptual schematic of recycling waste Li-ion batteries for the synthesis of acid-modified benzene oxidation materials is illustrated in Fig. 6. Activations of surface oxygen species and benzene molecules for gaseous benzene oxidation A

H-LiCo-6

245

120 0

40

80 120 160 200 240 280

Temperature (oC)

T = 170 oC

0

160

320

712

445

480

640

Temperature (oC)

800

T = 30

H-LiCo-6

-oxygen

817

300 204

756

407

pristine LiCoO2

-oxygen

TCD Signal (a.u.)

pristine LiCoO2

B

TCD Signal (a.u.)

TCD Signal (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 20 of 43

447

0

150

300

450

600

750

Temperature (oC)

900

Fig. 7. (A) H2-TPR and (B) O2-TPD curves of the pristine LiCoO2 and H-LiCo-6 catalysts. For H2-TPR: ~0.05 g of catalyst, 50 mL·min-1 of 5 vol.% H2/Ar, 10 o C·min-1. For O2-TPD: ~0.05 g of catalyst, 50 mL·min-1 of 5 vol.% O2/He for 30-min adsorption, 40 oC; 50 mL·min-1 of helium for desorption, 10 oC·min-1. The activities of catalyst surface to activate oxygen species were investigated by H2-TPR and O2-TPD, with the results depicted in Fig. 7A and B, respectively. For the H2-TPR of pristine LiCoO2, reduction mainly took place in the range of 300 oC to 800 o

C with two discernible H2 consumption peaks. Based on the deconvolution results,

four peaks were resolved at 407, 445, 712 and 756 oC, respectively. Besides, two other small peaks were evolved at 120 and 245 oC as seen from the enlarged view in the inset. Three types of surface adsorbed oxygen species (α-oxygen) existed over the pristine LiCoO2, which desorbed at 204, 300 and 447 oC, respectively, in inert gas, as


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indicated by the gray background of the O2-TPD curve (Fig. 7B). Therefore, the H2 consumption peaks at 120, 245 and 407 oC over the pristine LiCoO2 are likely due to the reductions of surface adsorbed oxygen. The alkali metal oxides possessed very potent metal-oxygen bonds and require high temperature for reduction;39 in addition, the peak area ratio of the 445 oC and 712 oC peaks was calculated to be 1:2, which is in line with the sequential reduction steps below and the complete reduction 1H2 (g) 2H2 (g) assumption: Co2O3 (s)  2CoO (s)  2Co (s) . Therefore, these two

peaks are attributed to the consecutive reduction of Co2O3 to CoO to Co, and the 756 o

C peak is reasonably ascribed to the reduction of Li+ to Li metal. Interestingly, in the profile of the H-LiCo-6 sample, its onset H2 consumption

temperature was ~170 oC below the temperature at which significant reduction occurred over the pristine LiCoO2, indicating that the surface oxygen associated with the acid-induced Co and oxygen vacancies in the [CoO6] octahedra became more reducible and reactive. Also in the O2-TPD curves (Fig. 7B), the main desorption peak of α-oxygen over the H-LiCo-6 was 30 oC lower than that of the pristine LiCoO2, further testifying the promoting effect of acid treatment on the generation of active surface oxygen species. Moreover, the overall reduction profile of the H-LiCo-6 was like an arch with no prominent reduction peaks. The total H2 consumption of the H-LiCo-6 (~41 mmol·g-1) was nearly three times that of the pristine sample (~14 mmol·g-1), suggesting that more active oxygen species were generated by acid treatment. This result was also confirmed by the significantly larger oxygen desorption peaks of the H-LiCo-6 catalyst (Fig. 7B), including surface adsorbed



oxygen (α-oxygen) and lattice oxygen (β-oxygen). 350

H-LiCo-6 Li2O

395 oC

pristine LiCoO2

300 100

400

CoOx

300

360 oC 175 oC

470 oC

75 50

200 100

25 0

Temperature (oC)

500

C6H6 desorption in N2 290 oC

325

Outlet Benzene (ppm)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 43

0 0

10

20

30

40

50

Time (min)

Fig. 8. C6H6-TPD curves of the pristine LiCoO2 and H-LiCo-6 samples. Conditions: ~0.1 g of catalyst, 200 mL·min-1 of N2 with ~320 ppm of benzene for 1 h adsorption, room temperature; 200 mL·min-1 of pure N2 for desorption, 10 o C·min-1. Benzene adsorption and then desorption in N2 was carried out to study the interaction of gaseous benzene with catalyst surface. Under the adsorption conditions, i.e., 200 mL·min-1 of N2 with ~320 ppm of benzene flowing through ~0.1 g of catalyst, there was no significant difference in the benzene adsorption behaviors between the pristine and acid-treated samples (Fig. S4). Benzene adsorption experiments were also held over CoOx and Li2O, which can clarify the changes of benzene adsorption sites after acid modification via subsequent benzene desorption experiments. The CoOx and Li2O were respectively obtained by calcining CoCO3 and Li2CO3 at 500 oC for 6 h. Compared with the pristine LiCoO2, benzene desorption behavior changed greatly over the H-LiCo-6 catalyst. As shown in Fig. 8, two benzene desorption peaks were presented at 290 oC and 395 oC, respectively, over the pristine LiCoO2. These two peaks very probably corresponded to the desorption of benzene adsorbed on the lithium sites since there was no benzene desorption above 200 oC over CoOx with


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only one weak peak at 175 oC. Li2O exhibited three benzene evolution peaks, centered at 175 oC, 360 oC and 470 oC, respectively, which, however, were different from those of the pristine LiCoO2 (290 oC and 395 oC). The plausible explanation is that the presence of cobalt oxide changed the chemical environment of lithium sites, thus influencing benzene adsorption on them. After acid treatment, benzene desorption over the H-LiCo-6 was shifted to higher temperatures of 360 oC and 470 oC, which agree well with those of Li2O. That is because the chemical environment of lithium sites of the H-LiCo-6 catalyst was more like that of Li2O on the basis of the XPS data that the Li 1s peak of the H-LiCo-6 (= 52.5 eV) was closer to that of Li2O (= ~53.0 eV40) in comparison with the pristine LiCoO2 (= 53.9 eV). According to literatures, Kim et al.41 reported that oxygen vacancies increased the binding strength of benzene molecules with adsorbent surface; cation vacancies also played an important role in the adsorption of gaseous pollutants, as found by Kim et al.41 and Wang et al.28 Therefore, it is reasonable to infer that Li vacancies and oxygen vacancies around Li atoms acted as the main adsorption sites for gaseous benzene over the LiCoO2 catalysts, but we could not exclude the possible interaction of benzene molecules with interlayer Li+. Last but not least, much less benzene was evolved from the H-LiCo-6 than from the pristine LiCoO2, but the stronger interaction of benzene molecules with the acid-treated surface, as indicated from the higher desorption temperatures, would promote benzene activation to contribute in a major way to benzene oxidation. Discussion



Page 24 of 43 24

Catalytic oxidation of VOCs over metal oxides typically proceeds through Mars-van-Krevelen mechanism,42 in which gaseous pollutants extract lattice oxygen from catalyst surface to be oxidized with the catalyst being reduced and the thus-formed oxygen vacancies are later replenished by gas phase oxygen for the next reaction cycle. Therefore, the activations of benzene and sufficient oxygen from gas phase are two important factors influencing the oxidation activity of catalysts. Besides, the activated surface adsorbed oxygen species by crystal defects were also reported to be crucial to achieve high VOCs removal efficiency.43 According to the UV-vis and 1H-NMR results, during acid treatment of LiCoO2, parts of the Co3+ were washed out from the [CoO6] octahedra as Co2+ ions via the disproportion reaction with the deintercalation of interlayer Li+ ions, and the resulting formations of Li, Co and oxygen vacancies were testified by the ICP-OES, Raman, HRTEM and XPS measurements. The C6H6-TPD experiments suggest that Li vacancies and oxygen vacancies around Li atoms together with the interlayer Li+ acted as the main adsorption sites for gaseous benzene over the LiCoO2 catalysts, and the acid-treated surface possessed much stronger bonds with the adsorbed benzene molecules. Mehar and co-authors44 found that the stronger binding of catalyst surface with CO molecules was responsible for the higher CO oxidation activity due to the lower energy barrier for the activation of CO molecules. Besides, Xu et al.45 reported that the CO adsorption capacity also positively influenced the catalyst activity for CO oxidation. However, the adsorption capacity of benzene was not the decisive factor for catalyst activity in this study since the pristine LiCoO2, which possessed larger


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benzene adsorption capacity, exhibited greatly inferior activity to the H-LiCo-6 sample. Therefore, it is concluded here that the stronger binding of benzene, and hence, the much lower activation energy for reaction, was responsible for the higher benzene oxidation activity of the H-LiCo-6 catalyst. In terms of oxygen activation, thanks to the acid modification, both the mobility and abundance of the oxygen sites increased over the H-LiCo-6 catalyst. Specifically, from O2-TPD, compared with the pristine LiCoO2, much more surface oxygen species, including surface adsorbed oxygen and lattice oxygen, were evolved from the acid-treated sample, and the oxygen species became significantly more reactive and facile over the acid-modified surface on the basis of the H2-TPR results. The large amount of oxygen species was very likely to originate from the acid-induced Co and oxygen vacancies, which might provide plenty of adsorption sites for gaseous oxygen. In order to prove this deduction, catalytic benzene oxidation over CoOx and Li2O was also performed, with the results shown in Fig. 1. The CoOx and Li2O were respectively obtained by calcining CoCO3 and Li2CO3 at 500 oC for 6 h. As shown in Fig. 1, Li2O was inactive toward benzene oxidation even at a high temperature of 300 o

C, performing much worse than the pristine LiCoO2, while CoOx exhibited much

higher activity, which demonstrates that the crystal defects of CoOx were the main active sites to activate oxygen species over LiCoO2 catalysts. However, because of the high activity of the acid-induced Co and oxygen vacancies in the [CoO6] octahedra to activate oxygen species, the H-LiCo-6 possessed higher benzene oxidation activity than CoOx, especially in the range of 150 to 200 oC. As reported in literatures, oxygen



Page 26 of 43 26

vacancies played two roles in the activation of oxygen species. First, active oxygen vacancies could increase the mobility and reactivity of lattice oxygen over metal oxides;46 secondly, oxygen vacancies also acted as the reservoir for gaseous oxygen, and the formed surface adsorbed oxygen was then activated to be the oxidants for pollutant removal.43 On the other hand, cation vacancy is another important defect. The presence of cation vacancies in metal oxides has been reported to influence the physiochemical and catalytic properties. Wang et al.28 discovered that creating Mn vacancies in birnessite MnO2 was beneficial to the generation of reactive surface oxygen, which in turn resulted in high formaldehyde oxidation activity. Finally, the commercial LiCoO2 sample, which was obtained from a LiCoO2 supplier (Aladdin Company) of domestic lithium battery manufacturers, was also treated with HNO3 aqueous solution and tested for benzene oxidation, with the results shown in Fig. S5. The acid treatment method also worked effectively with the commercial LiCoO2. The authors declare no competing financial interest. Acknowledgements Funding: This work is supported by National Natural Science Foundation of China (No. 21806016), Dalian University of Technology Fundamental Research Fund (No. DUT18RC(3)006) and Welch Foundation (#T-0014). All the authors of this article would like to deliver their best wishes to the coming 70th founding anniversary of Dalian University of Technology since the year of 1949.


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Supporting Information Instruments,

procedures

and

parameters

used

for

catalyst

preparation,

characterizations, catalytic benzene oxidation and temperature programmed desorption of adsorbed benzene (C6H6-TPD); Effects of alternative humidities and “heating-cooling-heating” process on benzene oxidation over the H-LiCo-6 catalyst; Benzene oxidation over the H-LiCo samples obtained with different washing times; Comparison of benzene oxidation over the samples prepared by the recycled acid solutions; Effects of space velocity on benzene oxidation over the H-LiCo-6 catalyst; Comparison of the BET-normalized reaction rates between the pristine LiCoO2 and H-LiCo-6 catalysts; SEM images; XPS spectra of C 1s; Benzene adsorption curves; Benzene oxidation over the commercial LiCoO2 samples before and after acid treatment. References [1] Poma, G.; Liu, Y.; Cuykx, M.; Tang, B.; Luo, X. J.; Covaci, A. Occurrence of Organophosphorus Flame Retardants and Plasticizers in Wild Insects from a Former E-Waste Recycling Site in the Guangdong

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Figure captions Fig. 1. Temperature dependences of benzene conversion over the pristine LiCoO2 and H-LiCo catalysts. Benzene oxidation conditions: 200 mL·min-1 of synthetic air containing 370-430 ppm of benzene. Fig. 2. (A) XRD and (B) Raman spectra of the pristine LiCoO2 and H-LiCo-6 catalysts. Fig. 3. (HR)TEM images of the pristine LiCoO2 (A-C) and H-LiCo-6 (D-F) catalysts. Fig. 4. (A) UV-vis spectrum of the supernatant after acid treatment for 6 h and the photo of the supernatant was inserted in the inset; (B) 1H-NMR spectra of the pristine LiCoO2 and H-LiCo-6 catalysts. Fig. 5. XPS spectra of the pristine LiCoO2 and H-LiCo-6 catalysts. Fig. 6. Conceptual schematic of recycling waste Li-ion batteries for the synthesis of acid-modified benzene oxidation materials. Fig. 7. (A) H2-TPR and (B) O2-TPD curves of the pristine LiCoO2 and H-LiCo-6 catalysts. For H2-TPR: ~0.05 g of catalyst, 50 mL·min-1 of 5 vol.% H2/Ar, 10 o

C·min-1. For O2-TPD: ~0.05 g of catalyst, 50 mL·min-1 of 5 vol.% O2/He for

30-min adsorption, 40 oC; 50 mL·min-1 of helium for desorption, 10 oC·min-1. Fig. 8. C6H6-TPD curves of the pristine LiCoO2 and H-LiCo-6 samples. Conditions: ~0.1 g of catalyst, 200 mL·min-1 of N2 with ~320 ppm of benzene for 1 h adsorption, room temperature; 200 mL·min-1 of pure N2 for desorption, 10 o

C·min-1.



Fig. 1 100

Benzene Conversion (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 34 of 43

80

pristine LiCoO2

60

H-LiCo-2 H-LiCo-4 H-LiCo-6 H-LiCo-8 Li2O

40

CoOx LiCoO2-800

20 0 150

200

250

300

350 o

Temperature ( C)


400

Page 35 of 43

35

Fig. 2

Intensity (a.u.)

{104}



{113}





{107}



 LiCoO2 (44-0145) {015}

{101}

{003}

A





pristine LiCoO2

H-LiCo-6 LiCoO2-800

0

10

20

30

40

50

2 ( )

60

70

80

90

o

B

pristine LiCoO2

300

H-LiCo-6

590

Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


486

400

500

600

700 -1

Raman Shift (cm )


800


Page 36 of 43 36

Fig. 3


Page 37 of 43

37

Fig. 4

4

A

H-LiCo-6 LiNO3

1.0

Co(NO3)2

Co2+

3

HNO3

Absorbance

Absorbance

0.8

2

0.0 400

450

500

550

600

Wavelength (nm)

1

0 300

400

500

600

700

800

Wavelength (nm)

B

P1

Abundance

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


P2 P3

H-LiCo-6

pristine LiCoO2

20

15 1

10

5

0

-5

H Chemical Shift (ppm)


-10


Fig. 5 A Li 1s

pristine LiCoO2

Intensity (a.u.)

H-LiCo-6

B.E. = 1.4 eV

56

55

54

53

52

51

50

Binding Energy (eV)

B

Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 38 of 43

Co 2p

H-LiCo-6 Co 2p3/2 satellite

Co 2p1/2

satellite

pristine LiCoO2

810 805 800 795 790 785 780 775 770

Binding Energy (eV)


Page 39 of 43

39

C O 1s

Intensity (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


540

O

pristine LiCoO2 H-LiCo-6

O

537

534

531

528

Binding Energy (eV)


525


Page 40 of 43 40

Fig. 6


Page 41 of 43

41

Fig. 7

pristine LiCoO2 H-LiCo-6

TCD Signal (a.u.)

TCD Signal (a.u.)

A

245

120 0

40

80 120 160 200 240 280

Temperature (oC)

0

160

712

445

T = 170 oC

320

480

756

407

640

800

o

Temperature ( C)

B

pristine LiCoO2

-oxygen

TCD Signal (a.u.)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


T = 30

H-LiCo-6

-oxygen

817

300 204 447

0

150

300

450

600

o

750

Temperature ( C)


900


Fig. 8 350

H-LiCo-6 Li2O

395 oC

pristine LiCoO2

300 100

400

CoOx

300

360 oC 175 oC

470 oC

75 50

200 100

25 0

0 0

10

20

30

40

Time (min)


50

Temperature (oC)

500

C6H6 desorption in N2 290 oC

325

Outlet Benzene (ppm)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 42 of 43

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This study provides a new concept of reusing e-wastes for air pollutant removal.


Synergy of Lithium, Cobalt and Oxygen Vacancies in Lithium Cobalt

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