Spectroscopic Investigation for Oxygen Reduction and Evolution

Subscriber access provided by UNIV OF CALIFORNIA SAN DIEGO LIBRARIES

Article

Spectroscopic Investigation for Oxygen Reduction and Evolution Reactions on Carbon Electrodes in Li-O Battery 2

Yu Qiao, and Shen Ye J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.6b01784 • Publication Date (Web): 23 Mar 2016 Downloaded from http://pubs.acs.org on March 31, 2016

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

The Journal of Physical Chemistry C is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 55

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

The Journal of Physical Chemistry

Spectroscopic Investigation for Oxygen Reduction and Evolution Reactions on Carbon Electrodes in Li-O2 Battery *

Yu Qiao and Shen Ye

Institute for Catalysis, Hokkaido University, Sapporo 001-0021, Japan

Abstract To develop a lithium-oxygen (Li-O2) battery with a high specific energy, the electrochemical oxygen reduction reaction (ORR) and oxygen evolution reaction (OER) on two typical carbon electrodes, a glassy carbon electrode and a porous carbon electrode made of Ketjenblack, have been systematically evaluated in a DMSO-based electrolyte solution. The galvanostatic cycling was employed to characterize the electrochemical properties on these carbon electrodes during the ORR and OER processes. The ex situ far/mid-infrared spectroscopy and in situ UV-vis absorption spectroscopy were employed to quantitatively analyze the reaction products adsorbed on the electrode surface and dissolved in solution, respectively. Our results demonstrated that lithium peroxide (Li2O2) is a major ORR product on the porous carbon electrode, while the lithium superoxide (LiO2) becomes a major ORR product on a glassy carbon electrode. On the other hand, a large amount of byproducts, such as lithium carbonate (Li2CO3), was also observed only on the porous carbon electrode surfaces at the end of the ORR. Li2CO3 gradually decomposed with the increasing potential, but remained on the electrode surface depending on the upper potential limit of the OER. The accumulation of the byproducts during the ORR/OER results in electrode passivation and capacity fading. A better understanding of the ORR/OER mechanism and possible reason for poor reversibility on these electrodes has been achieved in this study. 1

ACS Paragon Plus Environment


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

Introduction The lithium-oxygen (Li-O2) battery is recently attracting a great deal of research attention due to its high theoretical specific energy compared with the conventional Li-ion batteries and is regarded as one of the most promising energy storage devices.1-6 However, its development is fraught with many practical obstacles such as high charge overpotential and poor cycle life.7-13 The cathode reaction in the Li-O2 cell has been previously proposed as a 2-e– reaction of the oxygen molecule, 2Li+ + O2 + 2e– ⇌ Li2O2,1 with an equilibrium potential of 2.96 V (vs. Li/Li+).5 While in a Li-O2 cell using a porous carbon cathode, the potential for the oxygen evolution reaction (OER) can be as high as 4.5 V. The high overpotential for the OER causes a significant decrease in the energy conversion efficiency and oxidative decomposition of the carbon-based electrode and organic solvents.8, 11-12, 14 The chemical stability of the carbon electrode and electrolyte solutions towards the major reaction products, lithium peroxide (Li2O2) and lithium superoxide (LiO2) also limits its further applications.15-19 In order to overcome these drawbacks, many challenges have been devoted to using a porous carbon cathode due to its remarkable superiority including high conductivity, low cost and ease of fabrication to form the three-dimensional structure and three-phase interface.20-25 Typically, doping the carbon electrode with various metal oxides and noble metal nanoparticles has been employed to reduce the OER overpotential.26-28 The redox mediators have also been proposed to reduce the OER overpotential.29-31 However, the OER potential increased again after a few cycles of the oxygen reduction reaction (ORR) and OER on the porous carbon electrode.30, 32-33 In most cases, the reaction mechanism and reaction products for the ORR/OER on the porous carbon electrodes are fully understood on a molecular level. By help of a

13C-based

electrode, Bruce and co-workers worked to

distinguish the byproducts from the carbon electrode and organic solvents.34 2


Page 2 of 55

Page 3 of 55

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


Their pioneer work proved that the lithium carbonate (Li2CO3) was mainly generated from the carbon cathode as well as the solvents, while the lithium carboxylates were mainly formed from the solvents during the ORR. These products were simultaneously decomposed and produced during the subsequent OER. However, the relative amounts of all the reaction products have not yet been systemically discussed. FTIR is a powerful technique to analyze the reaction products on the electrode surface. Bruce and co-workers quantitatively investigated the molar ratio of reaction product (Li2O2) and byproducts (Li2CO3 and HCOOLi) on a nano-porous gold (NPG) cathode.35 They reported that the ratio of the byproducts was less than 1% even after 100 ORR/OER cycles. Xu et al. carried out comprehensive IR observations on a porous carbon electrode which is a practical material for the Li-O2 battery.19 They found that the growth of byproducts (such as Li2CO3, HCOOLi and LiAc) on a Super P porous carbon electrode became very serious after 40 cycles. By using a honeycomb-like palladium-modified hollow spherical carbon cathode, the formation of the byproducts was reduced but still could be observed after 100 cycles. However, most of the possible soluble products in solution have been ignored due to lack of characterizations.36-38 In this case, finding a practical way,

clarifying

the

details

and

making

the

mechanism

more

clear/comprehensive are essential in order to develop a practical carbon cathode for the Li-O2 batteries. Based on a systematic investigation for the ORR/OER on a gold electrode using in situ surface enhanced Raman spectroscopy (SERS) and UV-vis absorption spectroscopy in a DMSO-based electrolyte solution, we previously demonstrated that not only Li2O2 was formed on the electrode surface, but also a large amount of LiO2 was generated in the solution during the ORR.36 The reduction electron number per oxygen molecule on the gold electrode is expected to be much lower than 2e–/O2 (but higher than 1e–/O2), which is 3



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

unexpected regarding most of the differential electrochemical mass spectrometry (DEMS) observations on the porous carbon electrode that reported a 2e–/O2 discharge mechanism.7-8,

39-40

We expected that the

difference should be mainly attributed to the three-dimensional structure on the porous carbon electrode immersed only with a very small amount of the electrolyte solutions, which significantly promotes the Li2O2 formation. However, detailed results were not provided in the previous paper. In the present study, detailed experimental results obtained on a Ketjenblack-based porous carbon electrode have been provided and compared to that obtained on a glassy carbon electrode. By combining mid-IR and far-IR observations, we were able to determine the relative amount of the products and byproducts on these carbon electrodes during the ORR/OER cycle. The in situ UV-vis absorption spectroscopy was further used to analyze the soluble products under an electrochemical environment similar to that obtained in a coin-type cell. A number of facts have been experimentally confirmed on the carbon electrodes in the DMSO-based electrolyte solution: (1) Li2O2 is the dominant ORR product on the porous carbon electrode during the ORR along with many byproducts (Li2CO3, LiAc, Li2SO4 etc.); (2) A small amount of LiO2 is trapped in the porous carbon electrode during the ORR, independent of the depth of the ORR; (3). The reaction yield for the Li2O2 and LiO2 species are significantly dependent on the surface structure and morphology. (4) A typical byproduct of Li2CO3 partially decomposed with the increasing potential but remained on the porous carbon electrode surface depending on the upper potential limit of the OER. The accumulation of byproducts on the electrode surface results in the electrode passivation and capacity fading. A better understanding of the ORR/OER mechanism and origins for the degradation of the Li-O2 battery performance has been achieved. We expect that these findings will greatly benefit the development of the practical Li/O2 battery. 4


Page 4 of 55

Page 5 of 55

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


Experimental Section Materials: All chemicals were purchased from Kanto Chemical Co., Inc. (Tokyo, Japan) except these specially mentioned. The electrochemical grade DMSO was dried in freshly activated molecular sieves (type 3Å and 4Å) for several days before used. The DMSO-based electrolyte solution was chosen because it has been shown to exhibit a significant stability towards KO2 and Li2O2. Battery grade lithium perchlorate (LiClO4) from Stella Chemifa Co., Inc. (Osaka, Japan) was used for preparing the electrolytes. The typical concentration of the electrolyte was 0.5 M in the present study. The water content in the DMSO electrolyte solution was determined as 6-8 ppm by Karl Fischer titration method. After purging by O2 for 15 min, the water content in the electrolyte slightly increased (11-15 ppm). Ketjenblack EC600JD (Lion Co., Inc.), which has a BET area of 1400 m2/g, was dried at 120 °C for 12h under vacuum before the electrode preparation procedures. The as received Li2CO3, LiAc and Li2SO4 were dried at 100 °C for 12 h under vacuum. In the present study, the fresh lithium metal foil was kept in 0.5 M LiClO4-propylene carbonate (PC) electrolyte for 5 days before being used as the reference and counter electrodes in DMSO electrolyte. Bruce and co-workers reported that this procedure was useful to stabilize lithium in the DMSO.35 Lithium was rinsed by DMSO to remove any remained PC before use. Preparation of cathodes: The porous carbon cathodes were fabricated on a carbon paper (Toray) using a slurry of Ketjenblack. The slurry of Ketjenblack was prepared with a polyvinylidene fluoride (PVDF) binder and N-2-methyl pyrrolidone (NMP) dispersing agent in the ratio of 20∶1∶308 (m/m/m). The slurry was stirred for 3 min at 1000 rpm and then uniformly coated onto a carbon paper current collector. The electrode was then vacuum dried at 110 °C for 15 h and transferred into an argon-filled glove box (Labmaster MB10-C, MBraun) after vacuum drying to avoid exposure to the lab air. 5



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

Approximately 1 mg of Kentjenblack powder was coated on 1 cm2 of the carbon paper. Thus, the capacity of 1 mAh/cm2 used in this study, for example, corresponds to a capacity of 1000 mAh/g in weight. A glass carbon electrode (6.0 mm diameter) was purchased from Tokai Carbon Co., Inc. (Tokyo, Japan), then cut into a cylindrical electrode (2.0 mm height). The pre-treatment procedures and characterizations of the samples used in the study are briefly described in Scheme 1. SwagelokTM cell assembly: A SwagelokTM type cell (Figure S1a) was used to evaluate the electrochemical reactions in the Li-O2 battery. At the bottom of the anodic base, a piece of lithium foil (diameter, 16 mm; thickness, 0.8 mm) was fixed as the counter electrode. A porous glass microfiber filter (18 mm diameter, No. 1820-021, Whatman) was employed as a separator and placed on the lithium foil. Approximately 0.2 mL of electrolyte solution was dropped on the separator to make it totally wetted. The porous carbon cathode, as mentioned above, was then directly placed on the separator. The cell was sealed by screws and continuously purged by O2 (Hokkaido Air & Water) for 30 min (at the flow rate of nearly 5 mL/s) before cycling. Electrochemical characterizations: The electrochemical measurements using the SwagelokTM type cell were carried out under potential control using the battery tester system SPEC90476 (Kikusui Electronics Corp.) at room temperature. For the in situ UV-vis measurement, the potential was controlled by a potentiostat (Potentiostat/Galvanostat 2020, TOHO Technical Research) and a function generator (Function Generator 2230, TOHO Technical

6


Page 6 of 55

Page 7 of 55

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


Research) at room temperature. The current and potential outputs from the potentiostat were recorded by a multifunction data acquisition module (USB-6211, National Instruments) controlled by LabVIEW. All the potentials in the study are referenced to Li/Li+. Typically, the characterization of the cell was carried out under galvanostatic control at a current density of 0.1 mA/cm2 from the open circuit potential (OCP) unless otherwise noted. The OCP was approximately 3.3 V in most cases in this study. Potential limits for the galvanostatic cycling were selected between 2.0 V and 4.5 V in the study. After each OER characterization, the cell was kept on open circuit for 20 min before the next ORR characterization under galvanostatic control. Infrared Characterizations: A combined mid-IR/far-IR spectrometer, Perkin-Elmer Frontier, was also employed for the product characterization on the electrode surface, in both of the far-IR frequency region (30-700 cm–1) and mid-IR frequency region (500-4000 cm–1). Especially, the far-IR spectrometer is equipped with an auto-exchangeable grid-type beam splitter and a far-IR DTGS detector. The FTIR measurements were carried out using a GladiATR accessary with a diamond crystal (Pike Technology) purged with dry air. Typically, 64 interferograms were accumulated for one spectrum with a resolution of 0.5 cm–1 (far-IR region) or 1.0 cm–1 (mid-IR region). For the pre-treatment operation of the samples for the IR measurements, the cycled electrodes were first rinsed three times by DMSO to remove the electrolyte salt (LiClO4) and LiO2 on the electrode surface, and then continuously rinsed by dimethyl carbonate (DMC) to remove DMSO away in the glove box, then dried in a vacuum drying oven for 4 h at 55oC to make the DMC totally volatilized. UV-Vis Measurements: The in situ UV-vis bulk electrolysis cell (Figure S1b) was modified from a commercial UV-vis cell (10 mm optical path length) with a three-electrode configuration which does not block the light path. Lithium foil was used as the counter and reference electrodes. Typically, 3.0 7



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 8 of 55

mL of electrolyte solution was added to the UV-vis cell, and the geometric area of the porous carbon electrode and glass carbon electrode was 1.0 cm2 and 0.27 cm2, respectively. The electrolyte solution was purged in highly-purified O2 for 30 min before adding it to the cell. The details for the in situ UV-vis on a flat gold electrode have been described elsewhere.36 In the present study, the O2 dissolved in the DMSO-based electrolyte solution (ca. 2.0 mM/cm3) is not high enough to provide the ORR on the porous carbon electrode which has a very high specific area (nearly 1.0 m2/g) in comparison to that of the gold or glassy carbon electrodes. In order to supply a sufficient O2 to the porous carbon surface, the carbon electrode was sandwiched by two pieces of PTFE films (100 nm pore size, Toyo Roshi Kaisha, Ltd.) as the gas diffusion layer (GDL). O2 was continuously supplied towards the back side of electrode surface (at a flow rate of nearly 10 mL/min, Figure S1b). The cell was assembled in the glove box. The

UV-vis

absorption

spectra

were

recorded

using

a

UV-vis

spectrophotometer (Lambda 650, Perkin-Elmer) equipped with a double monochromator and a photomultiplier detector. As described in our previous paper,36 the UV-vis measurements were carried out to detect the superoxide product in the DMSO solution based on its 1πu → 1πg transition around 252 nm.41-43 Based on HASB theory and our previous results, the Li-ion is well solvated by DMSO and becomes a relatively soft acid.36, 44 The solvated-Li+ [Li+(DMSO)n] will combine with the superoxide anion (relative soft base) to form a stable ion-pair in the DMSO solution. In this case, the LiO2 mentioned in this study is actually a kind of free superoxide ion-pair in DMSO solution. The electrolyte solution with the same composition was used as a reference for the measurement. In order to get the LiO2 trapped in the porous carbon after cycling, the cycled porous carbon cathode directly underwent ultrasonic irradiation (15 s) in the same cell after the counter and reference electrodes were removed from the cell. After centrifuging the turbid solution in the cell 8


Page 9 of 55

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


(8000 rpm, 3 min), the clarified solution was evaluated by the UV-vis measurement.

Results and Discussion 1. Electrochemical behaviors of the carbon electrodes. Figure 1a shows a galvanostatic ORR/OER cycle on a Ketjenblack porous carbon electrode (black trace) in the SwagelokTM cell at a current density of 0.1 mA/cm2 in 0.5 M LiClO4-DMSO saturated by O2. Figure 1b shows a galvanostatic ORR/OER cycle on a glassy carbon electrode in a typical electrochemical cell under similar conditions. Generally, the ORR process on the carbon electrodes are dominated by a plateau around 2.6 V. The potential quickly drops to “death” as the ORR charge is close to its maximum capacity. Even the polarization curve of the ORR is similar in shape, the capacity obtained on the porous carbon electrode (2.3 mAh/cm2) is much higher than that of the glassy carbon (16 µAh/cm2). Figure 1a also shows the polarization curve of the glassy carbon (blue trace) and it is easy to see its low capacity in comparison to that of the porous carbon electrode (black trace). On the other hand, the OER on the porous carbon electrode exhibits several different regions,3, 39 similar to previous reports about the porous carbon electrodes.15, 20, 35, 45

The OER potential on the glassy carbon electrode monotonically increases

till the upper limit (4.5 V) with an efficiency of only 61%.

9



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

2. Spectroscopic Characterization during the ORR 2.1 Far-IR and Mid-IR Observations To fully understand the reaction mechanisms on the carbon electrode surface, one needs to quantitatively know the species formed both on the electrode surface and in the electrolyte solution. Figure 2 shows typical ex situ IR spectra on a Ketjenblack-based porous carbon electrode after the ORR with a different capacity in a DMSO-based electrolyte solution in (a) far-IR (200– 700 cm–1) and mid-IR region (500–1750 cm–1), respectively. The IR spectra for several standard materials, which are expected to be possible products or byproducts during the galvanostatic ORR process, such as Li2O2, LiAc, Li2CO3 and Li2SO4, are also given as references. A number of IR peaks appear and gradually increase with the ORR process. Typically, five far-IR peaks are observed at the end of the ORR on the porous carbon (purple trace, 2.25 mAh/cm2, Figure 2c). In comparison to the IR spectra of the standard materials, the three peaks at lower wavenumber (311,

10


Page 10 of 55

Page 11 of 55

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


420 and 488 cm–1) should be mainly attributed to those of Li2O2. Since the IR absorption bands for Li2O2 and Li2CO3 overlap each other in the spectral region, a quantitative analysis for the two species need more careful analyses (see below). A pair of IR peaks observed at 621 and 658 cm–1 (Figure 2c) should be specifically attributed to those of lithium acetate (LiAc) on the porous carbon electrode surface. As a complementary observation in the mid-IR region, the formation of similar species on the carbon surface during the ORR was also confirmed (Figure 2d). A broad IR absorption around 590 cm–1 in the mid-IR spectra should be attributed to the Li-O stretching in Li2O2. A group of peaks were observed in the frequency region between 1300 and 1650 cm–1, which should be attributed to the COO stretching modes from Li2CO3 and LiAc. The out-of-plane deformation mode of CO3 in Li2CO3 produces the peak at 862 cm–1. The peak at 1104 cm–1 can be assigned to the SO4 anti-symmetric stretch in Li2SO4, indicating the decomposition of the DMSO-based electrolyte solution during the ORR process, which also coincide with the DEMS observations.34 The IR observations clearly demonstrate that many byproducts, such as Li2CO3 and LiAc, were formed on the porous carbon electrode surface during the first ORR process, in addition to the main active ORR product of Li2O2. These byproducts are undesirable for the ideal ORR reaction and may significantly affect the round-trip efficiency and reversibility of the Li-O2 battery.5 It should be mentioned that the present fact has been partially ignored since most previous research interests have focused on the oxidation and decomposition reactions during the subsequent OER process. Based on our recent study, even the OER overpotential was significantly reduced using a redox mediator of trathiafulvalene (TTF); these byproducts were also generated on the porous carbon electrode surface during the ORR and accumulated on the electrode surface.33

11



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 12 of 55

On the other hand, the ORR/OER products and byproducts have been previously evaluated by many techniques, such as Raman, XRD, XANES and XPS,17, 46-48 but effective information for the relative amount of the byproducts towards Li2O2 during cycling is still lacking. This is important to understand the reaction mechanism and to improve the cycle reversibility and ORR/OER efficiency for the Li-O2 battery. 2.2 Mathematic Model for the IR Analysis Based on the general theory for the IR absorption spectroscopy, we are able to determine the ratio of the components in a mixture. The calculation becomes slightly complicated as the IR absorption peaks for the target species overlap each other. A general mathematic model for the IR spectral analysis for a multicomponent system will be briefly described below. The ATR-IR absorption spectrum of a species, i, in a mixed system can be expressed as, = × × ,

where εi(x) is the concentration extinction coefficient at wavelength x and Ci is concentration, and di,e is an effective penetration depth for the IR field, , = /2

where n21 is the ratio of the refractive indices of the ATR crystal and sample, E0 is the light field at the surface, θ is the incident angle and dp is the distance at which the electric field strength falls to 1/e.49-50 In the following analysis, di,e is assumed to be a constant for the experiment. If we can measure a ATR-IR spectrum for the pure species, fi,0(x), the equation above can be rewritten as =

× , ,

The IR absorption at an arbitrary wavelength (x) can be expressed as a linear combination of the IR absorption of all components in the system,

= × , + ,

,

× , + ! × ", + ⋯ !,

12


(1)

Page 13 of 55

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


Generally, if we can measure the IR absorbance of fmix(x) as well as fi,0(x) at different wavelengths at z points, we will have simultaneous equations of z-degrees, then the unknown variables, Ci/Ci,0, can be obtained by solving the equations. For example, as a model system containing three components (Scheme 2), we need at least IR absorbance at three wavelengths. By solving a ternary simultaneous equation, we can get a group of solutions for these species. Since the Ci,0 can be regarded as a constant, we are able to discuss the concentration changes in the same species, i. The solution is a ratio but not an absolute value, thus one is able to discuss the change in its own relative amount, but impossible to directly compare its amount with other species in the mixture without further conditions. However, we are able to determine the relative ratio for different components in the mixtures under some special cases if we can quantitatively measure the IR spectra for each pure component contained in the mixture (Figure S2b). We now consider a typical case for a two-component system (i=2) and their IR absorption bands overlap each other. Assuming we have the ATR-IR spectra for the system with a certain mixing ratio (r), the IR absorbance for the mixture at two wavelengths, x1 and x2, can be described as, , $ =

,% ,% × , + × , , ,

13



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 14 of 55

,% ,% × , + × , , ,

By dividing the two equations, we have ,% ( ∴ )*+ ()*+ ,%

=

,,- , , × ×( .( , , ,- ,, , ,,- , , × ×( .( , , ,- ,, ,

(2)

This equation shows that the ratio of the IR intensity at two wavelengths (left side) can be regarded as an inverse function of the mixing ratio (r), C1,r/C2,r. If we can prepare the mixed samples with different mixing ratios, a hyperbola relation is expected between the IR intensity ratio and mixing ratio (Eq. 2). By fitting the curve, we are able to determine the constants in Eq. 2. When we have a real IR spectrum for the same two-component system, we can estimate its mixing ratio by measuring the IR intensity ratio from Eq. 2 or reading from the hyperbola curve at the specified IR intensity ratio. The treatment can be further simplified if two IR absorption peaks are completely separated. Assuming components 1 and 2 only have an absorption at wavelength x1 and x2, respectively (as shown in Figure S2c), then. , $ =

,% × , ,

, $ =

,% × , ,

Similarly, we have

∴

()*+ ,% ()*+ ,%

=

, ,-

×

, ,

×

(, ( ,

(3)

Eq. 3 indicates that the IR intensity ratio is directly proportional to the ratio of the two components. By making a calibration curve at various mixing ratios, we can determine the parameters by a linear function. Based on the relationship, we can determine the ratio for the mixtures with an unknown ratio. 2.3 Quantitative Analysis Model for the IR Observations As shown Figure 2, several species are expected to form on the porous carbon surface after the ORR. Based on the mathematic models described in 14


Page 15 of 55

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


the last section, quantitative analyses were carried out to estimate the relative amounts of the ORR species in the mid-IR (Figure 3) and far-IR region (Figure 4). First, we worked to determine the relative amount of the two major byproducts, Li2CO3 and LiAc, from the mid-IR measurement (Figures 2b and 2d). As shown in Figure 2d, since the IR mode at 862 cm–1 for Li2CO3 and that at 1609 cm–1 for LiAc have no overlap, the relative amount of the two species can be quantitatively estimated based on Eq. 3. The IR measurements were first carried out on the mixtures of standard materials of Li2CO3 and LiAc mixed at known ratios. Figure 3a shows a good linear relationship between the intensity ratio for the two IR modes and the molar ratios of the two species (Figure 3a). Since the experimentally observed ratio of two peaks in the ORR sample (Figure 2d) was 0.94, the molar ratio of Li2CO3 and LiAc on the surface after the ORR was estimated to be 4.23 : 1 from the calibration curve (blue diamond, Figure 3a). Similarly, we can also estimate the molar ratio of Li2CO3 and Li2SO4 using the IR peaks at 862 cm–1 (Li2CO3) and 1104 cm–1 (Li2SO4) resulting in 13.5 : 1 on the porous carbon surface after the ORR (see the blue diamond in Figure 3b). Based on the analyses in the mid-IR region, we are able to determine that the ratio between Li2CO3, LiAc and Li2SO4 on the porous carbon electrode surface after the ORR is 1 : 0.24 : 0.07,

15



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

which is expected to be a constant for the following further analysis. A mixture using the standard materials based on the ratio was prepared as a standard for the partial ORR products (denoted as “M-powder” below). As shown above, we are able to determine the relative amount of the main byproducts for Li2CO3 and Li2O2 as well as Li2SO4, but it is still difficult to estimate the amount of the major product of Li2O2 at 593 cm–1 in the mid-IR region which only shows a broad background band (Figure 2d). To solve this problem, we further extended the present quantitative analysis to the far-IR region. Based on the ratio between Li2CO3, LiAc and Li2SO4 on the porous carbon electrode surface after the ORR determined above (1 : 0.24 : 0.07, i.e., M-powder), the mixtures of the Li2O2 and “M-powder” at different ratios were prepared and characterized in the far-IR region. Figure 4a shows the IR spectrum for a mixture at a ratio of 6 : 1 (pink trace) in addition to that of the “M-powder” (red trace), Li2O2 (black trace) and the porous carbon electrode after the ORR (blue trace). Since the spectral overlaps occur between these species in the far-IR region, the quantitative analyses were based on Eq. 2. We choose two peaks at 418 and 488 cm–1, and measured the intensity ratio of the two peaks as a function of the mixing ratio of Li2O2 and the “M-powder” (Figure 4b). The IR intensity ratio for the far-IR peak pairs (418/488 cm–1) shows a hyperbola-like relation to the molar ratio of Li2O2 and the “M-powder”, confirming that the

16


Page 16 of 55

Page 17 of 55

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


expectation by Eq. 2 is correct. Based on the IR intensity ratio determined from the experimentally observed IR spectrum of the ORR (black trace, Figure 4a), the molar ratio of the Li2O2 and other decomposition products was estimated as 2.9 : 1. To confirm that the present estimation is reliable, similar calculations were also carried out at the two frequencies of 538 and 621 cm–1 (Figure 4c). Although Figures 4b and 4c show different dependences, the relationship can be also fitted by a reciprocal function (dotted line, Figures 4b and 4c) and the ratio was determined to be 3.3 : 1. The discrepancy between the two calculations is reasonably small, then we believe that the present determination is reasonable and the molar ratio between the major product (Li2O2) and byproducts (the total amount of other decomposition products) on the Ketjenblack-based porous carbon electrode after the first ORR is approximately 3.1 (±0.2) : 1. In other words, the ratio of the Li2O2, Li2CO3, LiAc, and Li2SO4 after the ORR was estimated as 4.06 : 1 : 0.24 : 0.07, based on the composition of “M-powder” determined above. This relative amount of the ORR byproducts (Li2CO3, LiAc, and Li2SO4) is unexpectedly high, suggesting that approximately a quarter of the reaction products (Li2O2) are the byproduct on the porous carbon electrode surface only after one ORR polarization (see further discussion below). The present calculation was also applied to the IR results obtained on the glassy carbon electrode (Figure S3a). The IR peaks at 593 and 862 cm–1, attributed to Li2O2 and Li2CO3, respectively, were observed on the glassy carbon surface but the relative amounts seem to be quite different from that observed on the porous carbon electrode. By a similar method, the relative amount of the two species was estimated; the molar ratio of Li2O2 and Li2CO3 on the glassy carbon surface after the ORR was estimated to be 1 : 0.018. The ratio of the byproduct after the ORR seems to be much lower than that observed at the porous carbon electrode surface, i.e., less than 2% of byproducts were formed assuming that only the two species were formed on 17



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

the glassy carbon electrode after the ORR (in fact, more LiO2 was formed in DMSO, see below). Similar results have also been obtained on the gold electrode surface, indicating that the formation of the byproduct on the surface is significantly dependent on the electrode morphology and structure. 2.3 Soluble Species Determined by UV-Vis Measurements In the present section, the soluble products in the DMSO-based electrolyte solution were further evaluated by the UV-vis measurement. Our previous study showed that a large amount of LiO2 was observed in the DMSO solution after the ORR on the gold electrode surface, which can be detected by in situ UV-vis spectroscopy.36 Similar observations were also carried out for the ORR on a glassy carbon electrode. After a galvanostatic ORR polarization on the glassy carbon electrode with a capacity of 16

µAh/cm2 (Figure 1b), the UV-vis observation was used to measure the concentration of the LiO2 in DMSO as reported before (see Figure S4).36 The determination shows that a charge of 15 µAh/cm2 used to produce the LiO2 species on the glassy carbon electrode. This accounts that 93.4% of the ORR charges. Assuming that the ORR products only contains the three species, i.e., LiO2, Li2O2 and Li2CO3, the yields for Li2O2 and Li2CO3 are estimated as 6.3% and 0.1%, respectively. The results are generally agreement with that observed on the gold electrode.36 It should be mentioned that the absolute values for the yields are dependent on the polarization method (galvanostatic or potentiostatic) and ORR rate. As the similar measurement was carried out on the porous carbon electrode, we met some difficulties during the experiment (Figure 5). The black trace in Figure 5a shows the galvanostatic ORR process on a porous carbon electrode in the in situ UV-vis cell under the same conditions as that for the glassy carbon electrode (Figure S4) and gold electrode.36 No clear potential plateau was observed during the ORR and the ORR capacity was far 18


Page 18 of 55

Page 19 of 55

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


more less than that of the Swagelok-type cell (green trace, Figure 5a). As mentioned in the Experimental section, when the porous carbon electrode was sandwiched by two pieces of PTFE forming GDL structure (Figure S1b), the ORR capacity increased with the appearance of a potential plateau around 2.6V (red trace, Figure 5a). As O2 was directly bubbled towards the backside of the GDL covered porous carbon electrode, a similar ORR behavior to that of the SwagelokTM cell was observed (pink trace, Figure 5a). We believe that the amount of oxygen molecules dissolved in the DMSO-based electrolyte solution (2.1 mM/mL)44 is not high enough to provide for the ORR on the porous carbon electrode with a large active surface. By making the GDL structure on the porous carbon electrode surface, the mass transfer of O2 was significantly improved under the continuous O2 supply. It is interesting to note that even though only the backside of the porous carbon electrode was covered by the GDL structure (i.e., the surface of the porous carbon electrode is exposed to electrolyte solution), a similar ORR behavior was observed in the in situ UV-vis cell (blue trace, Figure 5a). The little difference in capacity between the one-side and two-side GDL structures will be discussed later based on the in situ UV-vis measurements. In order to determine the LiO2 species trapped inside the porous carbon electrode sandwiched by the PTFE films, after the ORR, the PTFE films and Li

19



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

foils were carefully removed from the cell, and then the carbon electrode was dispersed in the same cell by ultrasonic irradiation. After a centrifuge procedure, the LiO2 was evaluated by the UV-vis measurement (Figure 5b). After the ORR, a UV-vis absorption peak was clearly observed at 252 nm while nothing was observed at the OCP. As already extensively investigated in our previous study, this UV-vis peak is attributed to the 1πu → 1πg transition for the LiO2.36 Based on the absorption coefficient of LiO2 at this wavelength, the concentration of LiO2 in µmol/cm2 was calculated from the absorbance (right y-axis in Figure 6a; left y-axis in Figure 6b). Furthermore, this peak was not observed in the bulk solution (red block, Figure 6a top), indicating that the PTFE prevents diffusion of the LiO2 species trapped in the porous carbon electrode. It is interesting to note that the amount of the LiO2 trapped in the PTFE-sandwiched porous carbon electrode is almost constant (ca. 2.0 µmol/cm2, Figure 6a), independent of the ORR capacity (0.34 ~ 2.08 mAh/cm2, Figure 5b; also see blue blocks in top panel of Figure 6a). As an ORR product, the same amount of LiO2 was trapped in the porous carbon electrode during the ORR process in the DMSO-based solution. On the other hand, if only the surface of the porous carbon electrode was covered by a PTFE film, the amount of LiO2 in solution gradually increased during the ORR (red blocks, bottom panel of Figure 6a). The amount of the LiO2 can be much higher than that trapped in the porous carbon electrode (blue blocks, bottom panel of Figure 6a). In other words, since the amounts of the LiO2 trapped in the porous carbon electrode do not change with the ORR capacity, more LiO2 is expected to be further electrochemically reduced to Li2O2 in the PTFE-sandwiched porous carbon electrode where the PTFE films efficiently obstruct the diffusion of the LiO2 into the bulk solution. In comparison to the cell with the double-side PTFE films, the higher ORR capacity observed in the cell with the single-side PTFE (Figure 5a) may be

20


Page 20 of 55

Page 21 of 55

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


partially attributed to the formation of the LiO2, which easily diffuses into the bulk solution with the single-side PTFE geometry. Figure 6b shows the dynamic process of the diffusion process before/after the top PTFE film was removed from the PTFE-sandwiched porous carbon electrode after the ORR with different capacities at the OCP. No LiO2 was observed in the bulk solution within 20 min at the OCP after the ORR with the double-side PTFE geometry (point A and B, Figure 6b) although a certain amount of LiO2 was trapped inside the porous carbon electrode (see the black dotted line, Figure 6b). However, as soon as the top PTFE was removed, the absorbance at 252 nm (i.e., the amount of the LiO2 in solution) started to increase. After 20 min (point B, Figure 6b), more than one third of the LiO2 trapped in the porous carbon electrode diffused out, depending on the capacity. It is interesting to note that even though nearly the same amount of LiO2 was trapped inside the porous carbon electrode (ca. 2.0 µmol/cm2, Figure 6a), the diffusion rate of LiO2 is different. The lower the ORR capacity, the higher the diffusion rate (i.e., slope of the temporal change of the UV-vis absorbance at 252 nm), and the more the diffusion amount of LiO2 into the bulk solution (see Figure 6b). This suggests that with the increasing ORR capacity, the diffusion of the

21



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

LiO2 during the ORR can be significantly hindered. Although the detailed structures on the porous carbon electrode surface during the process need to be quantitatively characterized, different amounts of Li2O2 were deposited on the electrode surface during the ORR process and can significantly change the pore structures in the porous carbon electrode, as a result, the diffusion of the LiO2 trapped in the porous carbon electrode is modified. The higher the amount of the deposited Li2O2, the slower the diffusion rate for the LiO2.

2.4 4 Reaction Yield of the ORR Products. Figure 7 summarizes the reaction yield of the ORR products and byproducts after an ORR on (a) a glassy carbon electrode and (b) a porous carbon electrode. The diameter of the circle is proportional to the total amount of the ORR products and byproducts. Much more ORR species are formed on the porous carbon electrode in comparison to that of the glassy carbon electrode (due to the huge difference in capacity). The LiO2 is the major product on the glassy carbon electrode after the ORR, accounting for 93.6 % (red sector, Figure 7a) of the total ORR capacity, while the Li2O2 deposited on the glassy carbon surface only accounts for 6.3 % (blue sector, Figure 7a). The byproduct of Li2CO3 formed during the ORR on the glassy carbon electrode can be almost ignored (0.1%) in comparison to those observed at the porous carbon electrode surface (see below). Assuming that the Li2O2 thin-film is deposited on the perfectly smooth glassy carbon surface (as single crystal), the thickness of the Li2O2 was estimated to be 3.5 nm after the ORR, and this value is similar to that obtained on the flat gold surface (3.7 nm).36 On the other hand, the LiO2 only accouns for 2.4% on the porous carbon electrode after the ORR with a full capacity of 2.25 mAh/cm2, while the Li2O2 became the major product (73.8 %) with several byproducts. In addition to the LiO2 and Li2O2, the amount of byproducts accounts for nearly 24%. The amounts of the undesirable byproducts were surprisingly high. They not only 22


Page 22 of 55

Page 23 of 55

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


waste the ORR effective capacity but also irreversibly block the active sites of the carbon electrode surface. However, this issue has been ignored while most of interests were focused on the oxidation and decomposition reactions during the OER. Controlling the formation of the byproducts during the ORR is an essential issue to improve the stability and cycling of the Li-O2 battery. As mentioned above, the yields obtained on these electrodes are significantly dependent on the observation conditions, such as galvanostatic / potentiostatic control, ORR rate as well as the depth of the ORR. One has to compare these results under the same conditions. The main difference in the reaction yield of the products between the glassy carbon electrode and porous carbon electrode should be attributed to the difference in the morphology and surface structrue of the carbon electrodes as well as the experimental conditions. The porous carbon electrode has an extremely high surface area (~1000 m2/g) and the electrode with a surface area of approximately 1 m2 per cm2 geometric area can be typically prepared. A rough estimation shows that the surface area of the flat glass carbon electrode is only 1000th of the porous carbon electrode. The significant difference in the specific surface area directly leads to the difference in the specific capacity per unit area for the two electrodes. On the other hand, it is known that the eletrochemical reaction

23



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

activities of the gaseous molecules are significantly dependent on the interface between the electrocatalyst, electrolyte solution and gas phase (i.e., three-phase interface) where the diffusion and transport of gaseous molecules as well as their electron transfer reactions mainly occur．The pore strcucture of the porous carbon electrode is expected to provide many sites with three-phase phase structures for the ORR/OER processes. Furthermore, the normal coin-type Li-O2 cell only uses a very small amount of electrolyte solution (ca. ~200 µL) under the oxygen purging. In such a case, the solution layer above the electrode surface with a high surface area is extremely very thin in comparison to the bulk cell used for the glassy carbon electrode and the gold electrode.7, 36 The local concentration of the 1e–reduction product of O2, LiO2, in the pores with three-phase interface structures containing very limited electrolyte solution significantly increases with the ORR. Due to the special interfacial structure, the LiO2 is expected to be immediately reduced to Li2O2 under the galvanostatic or potentiostatic ORR conditions.36 As discussed in our previous paper, we believe that the electrochemical reduction of LiO2 to Li2O2 will play a main role in comparison to that of the disproportion reaction of LiO2.36 The Li2O2 is deposited on the electrode surface in the DMSO-based electrolyte solution and becomes a major product on the porous carbon electrode. The high local concentration of LiO2 near the carbon electrode can also significantly promote the oxidation of the carbon surface, resulting in a high coverage of the byproducts up to 25% just after one ORR. The decomposition of the solvent may be also partially included under the condition, especially when the ORR capacity is higher. For example, it was found on the porous carbon electrode, the yield for the Li2CO3 decreased to 9.6% with the lower ORR capacity of 200 µAh/cm2.33 In contrast, the glass carbon electrode used in the study has a very low surface area. Almost no three-phase interface area is available on the glass carbon electrode, thus only a small amount of oxygen dissolved in the 24


Page 24 of 55

Page 25 of 55

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


solution (ca. 2.1 mM) can react with the electrode surface during the ORR in comparison to the gas-diffusion condition for the porous carbon electrode. Furthremore, the LiO2 species are easy to diffuse into the DMSO bulk solution as a free ion-pair before they are further electrochemcailly reduced to Li2O2. As a result, LiO2 becomes the major ORR product on the flat glass carbon electrode (and the gold electrode). 3. Spectroscopic Characterization during the OER Furthermore, the OER behaviors have been systemically evaluated by ex situ far-IR and mid-IR measurements on the carbon electrode. Since the OER potential quickly increases higher than 4.5 V on the glassy carbon electrode surface (Figure 1b), it is hard to reach a 100% round-trip efficiency under the upper potential limit, thus our characterization of the OER mainly focused on the porous carbon electrode. As discussed in our previous paper, the quick increase of OER potential on the flat electrode should be mainly attributed to the low diffusion rate of the LiO2 from bulk solution, which is impossible to meet with the galvanostatic OER rate (100 µA/cm2).33 Figure 8 shows ex situ IR spectra for the OER subsequent to the initial ORR process (Figure 2) in the (a) far-IR and (b) mid-IR regions. Several standard materials (Li2O2, Li2CO3 and LiAc etc) expected as the ORR/OER products and byproducts are also shown on the bottom of the spectra. Similar peak assignments as that for the ORR (Figure 2) were also used for the OER

25



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

process. Generally, the IR peaks become weaker with the OER but do not completely disappear at the end of the OER in both regions. Figure 9 summarizes the intensities for the peaks for the two major products of Li2CO3 (862 cm–1, red symbols) and LiAc (1609 cm–1, blue symbols) at different OER capacities (indicated by the arrows from right to left), together with those observed in the ORR process (indicated by the arrows from left to right). The potential profiles (right-y axis) observed during the same galvanostatic ORR/OER processes are also shown (black trace). The intensities of the two IR peaks gradually increase during the ORR process in a similar way but show different capacity dependence during the OER process. The LiAc keeps dropping and almost disappears at the end of the OER. Li2CO3 decreases until 3.7 V and turns to increase but suffers to a quick drop from 4.35 V. Li2CO3 continues to decrease in the OER process when the potential is lower than 3.7 V (1.6 mAh/cm2), then turns to increase to a maximum at 4.3 V (0.8 mAh/cm2) and finally decreases until the end of the OER at 4.5 V. Since the IR peak for Li2CO3 (862 cm–1) and LiAc (1609 cm–1) are well separated, the observed capacity dependences should reflect the changes in the two species during the galvanostatic ORR/OER cycle. However, the capacity dependence of the major product, Li2O2, is difficult to directly obtain

26


Page 26 of 55

Page 27 of 55

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


from either the mid-IR region or far-IR region due to the baseline problems. To estimate the relative amount of the Li2O2 species during the ORR/OER cycle, we used the general mathematic model (Eq. 1) to analyze the far-IR spectra to obtain the capacity dependence of the surface species on the carbon electrodes. As discussed in the previous sections, the ORR products mainly contain Li2O2, Li2CO3, LiAc and Li2SO4. To simplify the calculation, only three main species (Li2O2, Li2CO3 and LiAc) were used. Li2SO4 was ignored in the calculation since it relative amount is small in comparison to the other products on the electrode surface (1.3 %, at the end of the first ORR, see Figure 7). The analysis and calculation were carried out based on Eq. 1 in which three peak positions (418, 488 and 538 cm–1) were used. Three sets of IR results are necessary to solve the simultaneous equations. As further confirmation, the calculations were also carried out in combination with the IR absorbance at 621 cm–1. Based on the calculations using four different combinations, the error distributions for each factor were also estimated which is reasonably low under the present conditions, confirming that the calculation is correct. Figure 10 summarizes the calculation results for the capacity dependence of (a) Li2O2, (b) Li2CO3 and (c) LiAc for ORR (blue circles) and the OER processes (red circles). As shown in the Figure 10a, the Li2O2 is quickly generated at the beginning of the ORR, while the increasing rate gradually slows with the further reaction. The Li2O2 rapidly decreased at the beginning of the OER and a very low amount of Li2O2 remained on the electrode surface at the end of the OER process. It is interesting to note that the capacity

27



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

dependences of Li2CO3 (Figure 10b) and LiAc (Figure 10c) are very similar to that obtained from the isolated IR bands in the mid-IR region (Figure 9). This confirms that the present analysis method (Eq. 1) is very reliable for a quantitative discussion. As mentioned above, although we are able to know the relative change in the same species (Ci/Ci,0), we are unable to get its absolute value since the constant of Ci,0 is unknown. For the same reason, it is impossible to directly compare it for other species. In addition for the decomposition of the ORR products of Li2O2 and LiO2, one expects that the electrochemical oxidation of the carbon electrode and the organic solvents occurs with the increasing OER potential, generating Li2CO3 and other byproducts on the surface. At the same time, the electrochemical decomposition of the Li2CO3 and other byproducts may also simultaneously occur. All of these reactions finally determine the capacity dependence of the species on the electrode surface. As shown in Figures 9 and 10, Li2CO3 shows a quite complicated profile during the OER (increases first and then decreases to a certain level) while LiAc simply decreases to zero at the end of the OER. This implies that the decomposition of Li2CO3 becomes more drastic with the increasing OER potential. Similar results have also be observed by DEMS, in which the CO2, as a final decomposition product from Li2CO3 (or/and other lithium carboxylates), significantly increased at a high potential.7-8, 34, 40 The present observations demonstrate that the many undecomposed Li2CO3 still remained on the electrode surface even at the end of the OER which was impossible to be detected by DEMS. In summary, the present IR analyses quantitatively determined the relative amount of the products and byproducts during the ORR/OER cycle. The analysis is highly sensitive to the chemical state of the species than other methods previously used, such as XRD and SEM, and it is expected to provide a comprehensive way to investigate the electrode reaction process in the study of the Li-O2 battery. 28


Page 28 of 55

Page 29 of 55

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


4. Cyclability of ORR/OER on the Porous Carbon Electrode. In order to understand the influence of the accumulated byproducts (such as Li2CO3 and LiAc) towards the cyclability, the continuous galvanostatic ORR/OER cycles were further evaluated (Figure 11). Each galvanostatic ORR process was recorded until the potential dropped to 2.0 V, then the OER process was characterized until a 100% round-trip efficiency. After a 20 min waiting period, the subsequent ORR/OER was evaluated in a similar way. As shown in Figure 11a, the second ORR capacity significantly dropped from 2.25 mAh/cm2 to 1.32 mAh/cm2, thus approximately 40% of the original capacity disappeared only after one cycle. The ORR capacity further decreased with the cycles and reached a value of 0.78 mAh/cm2 in the fifth cycle. Furthermore, the ORR plateau around 2.6 V became narrower with cycling and almost disappeared after the third cycle. The OER processes show a similar shape at the similar potential (4.5 V), but with much less capacity

29



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

with the cycling. These results indicate that the surface states of the porous carbon electrode significantly changed with the cycling. In addition to the electrochemical characterization, ex situ IR observations were carried out at the end of each ORR process (Figures 11b and 11c). In the far-IR region (Figure 11b), the spectral features for Li2CO3 significantly increase with cycling while that for LiAc do not change much. The mid-IR spectra show similar features and demonstrate the accumulation of byproducts with the cycling. Figure 11d summarizes the relative amount of Li2O2 (black bar), Li2CO3 (red bar), LiAc (blue bar) and Li2SO4 (green bar) on the Ketjenblack-based porous carbon electrode surface. The full length of the bars in each cycle (from bottom (first cycle) to top (fifth cycle) corresponds to the amount of species on the porous carbon electrode surfaces. The bottom bar for the yield in the first ORR is generally same as that in Figure 7b except a small amount of dissolved LiO2 in solution. Furthermore, by using the similar procedures used for Figure 10, we are able to determine the yield for each product in respective cycle. Based on the determined ratio for the first cycle (bottom bar, Figure 11d), we are able to calculate the relative amount of each species on the electrode during the cycles. The amount of the Li2O2 suffers from a significant drop with the cycling (black bar, Figure 11d). 74.5% of the adsorbates was Li2O2 after the ORR. The ratio decreased to 19.3% in the adsorbates after the 5th ORR (in other words, 6.7% of the adsorbates after the first ORR capacity).

The coverage for the active ORR product of Li2O2

significantly decreased with cycling, resulting in the capacity fading. On the contrary, Li2CO3, a major byproduct during the ORR/OER, significantly increased with the cycling. The ratio of Li2CO3 increased from 19.5% in the first cycle to 75.1% in the fifth cycle. The absolute amount of Li2CO3 on the porous carbon electrode surface increased nearly 2.2 times after 5 cycles. The amounts of the other two decomposition byproducts (LiAc and Li2SO4) also slightly increased with the cycling. Instead of Li2O2, the “byproduct” of 30


Page 30 of 55

Page 31 of 55

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


Li2CO3 gradually becomes the “major product” with the cycling. These byproducts, especially Li2CO3, irreversibly occupy the active sites on the carbon surface and seriously hinder/block the growth of Li2O2, resulting in electrode passivation and a serious decline in the capacity. The degradation rate of the cycling can be directly associated with the deposition and growth of these byproducts on the electrode surface. This is crucial in order to understand the reason for the capacity fading of the ORR/OER process on the porous carbon electrode surfaces and can provide a general indicator to solve the difficult problems in the development of Li-O2 battery. It should be noted here that the depth for the ORR process is selected at the lower potential limit of 2.0 V, which is very close to the “death point” for the cathodic reaction of oxygen. Of course, if we stop the ORR at a higher potential (i.e., lower capacity), for example, 2.6 V (on the ORR plateau), a “better” cycling performance can be obtained.33 However, a fundamental difference between the “deep discharge” and “shallow discharge” does not exist based on an assessment towards a specific carbon-based cathode material. The ORR with a lower depth means that some active sites remain empty and the cathode capacity is not fully used. However, the adsorption of irreversible byproducts will finally occupy the electrode surface and degrade the ORR activity. In fact, our previous studies using a shallow ORR depth, demonstrated that the ratio of Li2CO3 accounts more than 10% among all of the ORR products at the very beginning of the ORR process (0.2 mAh/cm2) and remains on the electrode surface after the OER process,33 blocking the active sites for the growth of the “active species (Li2O2). Therefore, reducing the formation of Li2CO3 on the porous carbon electrode is considered as an essential problem that needs to be solved.

31



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

5. Influence of the Upper Potential during the OER. To reach a 100% round-trip efficiency for the ORR/OER cycle discussed in the previous sections, the upper potential limit for the OER was not regulated. The potential increased to 4.5 V at the end of the OER (Figure 1). On the other hand, Li2O2 and Li2CO3 show totally a different dependence of the potential (capacity) during the OER (Figure 10). For example, Li2CO3 reached a maximum at 4.35 V, while Li2O2 almost disappeared at this potential. The upper limit potential may significantly affect the subsequent ORR/OER cycle. Figure 12a shows the second ORR curves at two different upper potentials, 4.50 V (blue trace) and 4.35 V (red trace) in comparison with the first ORR/OER curve. As mentioned in the last section, the second ORR capacity decreased by approximately 40% of the original capacity at the upper potential limit of 4.5V (1.32 mAh/cm2, blue trace, Figure 12a). However, as the upper potential limit was reduced to 4.3 V, the second ORR capacity decreased by approximately 68% of the original capacity (0.71 mAh/cm2, red trace, Figure 12a). Almost no plateau was observed in this ORR curve. The second ORR capacity with the upper potential limit of 4.3 V is lower than that after five ORR/OER cycles with the upper potential limit of 4.5 V. Figure 12b shows the mid-IR characterization results observed at the end of the first ORR (black trace), first OER and second ORR at the upper potential limit of 4.5 V (blue trace) and 4.3 V (red trace). A more intense IR peak for

32


Page 32 of 55

Page 33 of 55

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


Li2CO3 (862 cm–1) was observed at the end of the OER with 4.3 V (red dotted trace) than that with 4.5V (blue dotted trace). As mentioned in the last section, this should be attributed to the partial decomposition of Li2CO3 in the higher potential region, which reached a maximum at 4.35 V, i.e., more Li2CO3 was formed at 4.3 V than that at 4.5 V. The amount of Li2CO3 at 4.5 V increased in the second ORR (blue solid trace), but at 4.3 V, it was almost identical (red solid trace). The present IR observations demonstrated that the accumulation of Li2CO3 on the porous carbon electrode surface is significantly dependent on the upper potential limit of the OER process and thus can significantly affect the efficiency of the subsequent ORR/OER cycle. As discussed in the last section, active sites on the porous carbon electrode surface are blocked by the Li2CO3, resulting in a serious drop in the ORR capacity and increase in the overpotential. The OER with the upper potential of 4.3 V has more Li2CO3 on the electrode surface, thus a worse performance was observed under this condition compared to that at 4.5 V. Furthermore, the porous carbon electrode was characterized by cyclic voltammetry (CV) with a different upper limit potential for five cycles (Figure S6). As the potential was swept to 4.5 V, the cathodic peak at 2.4 V, corresponding to the formation of the Li2O2, gradually decreased with the cycling. However, if the positive-going sweep was swept back at 4.3 V, the cathodic peak in the second cycle already becomes very small. These results confirmed that the formation of the Li2O2 was strongly suppressed in the second cycle when the upper potential limit is not high enough to electrochemically decompose the Li2CO3 formed which can block the reaction sites for the subsequent ORR process. Electrochemical removal of the Li2CO3 needs a high overpotential, at which the carbon electrode and electrolyte solution will also suffer from electrochemical decomposition. However, the

33



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

latter one seems to be not the main reason for the worse electrochemical behavior observed on the carbon-based cathodes. In our recent study, a redox mediator, TTF, was dissolved in a DMSO-based electrolyte solution to reduce the OER overpotential.33 The OER overpotential was definitely reduced by the TTF, resulting in an apparent decrease of the byproducts (mainly Li2CO3 species) generated during the OER process. However, the Li2CO3 generated during the ORR was not affected by the redox mediator and still accumulated on the surface of the porous carbon electrode, hindering the active site for the growth of Li2O2. Therefore, not only for the OER process, one should suppress the formation of Li2CO3 and other byproducts in the ORR process to promote the performance of the Li-O2 battery.

Conclusion In summary, ex situ far/mid-IR and in situ UV-vis observations have been employed to systematically investigate the galvanostatic ORR/OER cycling on the glassy carbon electrode and porous carbon electrodes in a DMSO-based electrolyte solution. The relative amounts of products and byproducts formed during the ORR and OER have been quantitatively evaluated. The role of the three-phase interface and surface morphology on the ORR/OER mechanism was discussed. The incomplete decomposition of the byproducts (mainly Li2CO3), seriously block the reaction sites of the porous carbon electrode which is the most important issue to be solved in order to make a practical cathode catalysis for the Li-O2 battery. The development of novel carbon materials, such as graphene, reduced oxidized graphene and carbon nanofibers modified by different metal oxides, is significantly expected to solve these problems in the near future. If we are unable to solve the problems, we may unable to see the real application of the Li-O2 battery as a high energy power source. 34


Page 34 of 55

Page 35 of 55

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


Associated Content Supporting Information: Structures for the cells used in this study; electrochemical behaviors, in situ UV−vis and mid-IR characterizations on the glassy carbon electrode; additional information of mathematic model for the IR calibration; Cyclic Voltammetry (CV) behaviors for the Ketjenblack cathode between different potential regions.

Author Information Corresponding Author: *Tel: +81-117069126. Email: [email protected]. Notes: The authors declare no competing financial interest.

Acknowledgments This work was supported by the Advanced Low Carbon Technology Research and Development Program (ALCA), specially promoted research for innovative next generation batteries (SPRING) from the Japan Science and Technology Agency (JST). The authors want to give a special thanks to Prof. Kubo (NIMS) and Mr. Kimura for the preparation of porous carbon electrode. Y.Q. acknowledges a scholarship from the China Scholarship Council (CSC). The authors thank Prof. Kohei Uosaki for stimulating discussions and comments. The authors thank Dr. Can Liu and Miss Yingying Zhou for their assistance with some experiments.

35



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

Reference: 1. Abraham, K. M.; Jiang, Z., A Polymer Electrolyte-Based Rechargeable Lithium/Oxygen Battery. J. Electrochem. Soc. 1996, 143, 1-5. 2. Bruce, P. G.; Freunberger, S. A.; Hardwick, L. J.; Tarascon, J.-M., Li-O-2 and Li-S Batteries with High Energy Storage. Nat. Mater. 2012, 11, 19-29. 3. Adams, B. D.; Radtke, C.; Black, R.; Trudeau, M. L.; Zaghib, K.; Nazar, L. F., Current Density Dependence of Peroxide Formation in the Li-O-2 Battery and Its Effect on Charge. Energy Environ. Sci. 2013, 6, 1772-1778. 4. Lu, J.; Li, L.; Park, J. B.; Sun, Y. K.; Wu, F.; Amine, K., Aprotic and Aqueous Li-O-2 Batteries. Chem. Rev. 2014, 114, 5611-5640. 5. Luntz, A. C.; McCloskey, B. D., Nonaqueous Li-Air Batteries: A Status Report. Chem. Rev. 2014, 114, 11721-11750. 6. Chang, Z. W.; Xu, J. J.; Liu, Q. C.; Li, L.; Zhang, X. B., Recent Progress on Stability Enhancement for Cathode in Rechargeable Non-Aqueous Lithium-Oxygen Battery. Adv. Energy Mater. 2015, 5, 1500633. 7. McCloskey, B. D.; Scheffler, R.; Speidel, A.; Girishkumar, G.; Luntz, A. C., On the Mechanism of Nonaqueous Li-O-2 Electrochemistry on C and Its Kinetic Overpotentials: Some Implications for Li-Air Batteries. J. Phys. Chem. C 2012, 116, 23897-23905. 8. McCloskey, B. D.; Speidel, A.; Scheffler, R.; Miller, D. C.; Viswanathan, V.; Hummelshoj, J. S.; Norskov, J. K.; Luntz, A. C., Twin Problems of Interfacial Carbonate Formation in Nonaqueous Li-O-2 Batteries. J. Phys. Chem. Lett. 2012, 3, 997-1001. 9. Abraham, K. M., Electrolyte-Directed Reactions of the Oxygen Electrode in Lithium-Air Batteries. J. Electrochem. Soc. 2015, 162, A3021-A3031. 10.

Black, R.; Oh, S. H.; Lee, J. H.; Yim, T.; Adams, B.; Nazar, L. F., Screening for Superoxide

Reactivity in Li-O-2 Batteries: Effect on Li2o2/Lioh Crystallization. J. Am. Chem. Soc. 2012, 134, 2902-2905. 11.

Hummelshoj, J. S.; Luntz, A. C.; Norskov, J. K., Theoretical Evidence for Low Kinetic

Overpotentials in Li-O-2 Electrochemistry. J. Chem. Phys. 2013, 138. 12.

Viswanathan, V.; Norskov, J. K.; Speidel, A.; Scheffler, R.; Gowda, S.; Luntz, A. C., Li-O-2

Kinetic Overpotentials: Tafel Plots from Experiment and First-Principles Theory. J. Phys. Chem. Lett. 2013, 4, 556-560. 13.

Meini, S.; Solchenbach, S.; Piana, M.; Gasteiger, H. A., The Role of Electrolyte Solvent

Stability and Electrolyte Impurities in the Electrooxidation of Li2o2 in Li-O-2 Batteries. J. Electrochem. Soc. 2014, 161, A1306-A1314. 14.

Luntz, A. C.; Viswanathan, V.; Voss, J.; Varley, J. B.; Norskov, J. K.; Scheffler, R.; Speidel, A.,

Tunneling and Polaron Charge Transport through Li2o2 in Li-O-2 Batteries. J. Phys. Chem. Lett. 2013, 4, 3494-3499. 15.

Laoire, C. O.; Mukerjee, S.; Plichta, E. J.; Hendrickson, M. A.; Abraham, K. M., Rechargeable

Lithium/Tegdme-Lipf6/O2 Battery. J. Electrochem. Soc. 2011, 158, A302-A308. 16.

McCloskey, B. D.; Scheffler, R.; Speidel, A.; Bethune, D. S.; Shelby, R. M.; Luntz, A. C., On the

Efficacy of Electrocatalysis in Nonaqueous Li-O2 Batteries. J. Am. Chem. Soc. 2011, 133, 18038-18041. 17.

Marchini, F.; Herrera, S.; Torres, W.; Tesio, A. Y.; Williams, F. J.; Calvo, E. J., Surface Study of

Lithium–Air Battery Oxygen Cathodes in Different Solvent–Electrolyte Pairs. Langmuir 2015, 31, 9236-9245.

36


Page 36 of 55

Page 37 of 55

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


18.

Zhang, Y.; Zhang, X.; Wang, J.; McKee, W. C.; Xu, Y.; Peng, Z., Potential-Dependent

Generation of O2- and Lio2 and Their Critical Roles in O2 Reduction to Li2o2 in Aprotic Li-O2 Batteries. J. Phys. Chem. C 2016. 19.

Xu, J. J.; Wang, Z. L.; Xu, D.; Zhang, L. L.; Zhang, X. B., Tailoring Deposition and Morphology

of Discharge Products Towards High-Rate and Long-Life Lithium-Oxygen Batteries. Nat. Commun. 2013, 4, 2438-2448. 20.

Grande, L.; Paillard, E.; Hassoun, J.; Park, J. B.; Lee, Y. J.; Sun, Y. K.; Passerini, S.; Scrosati, B.,

The Lithium/Air Battery: Still an Emerging System or a Practical Reality? Adv. Mater. 2015, 27, 784-800. 21.

Jun, W.; Zhong, H. X.; Qin, Y. L.; Zhang, X. B., An Efficient Three-Dimensional Oxygen

Evolution Electrode. Angew. Chem.-Int. Edit. 2013, 52, 5248-5253. 22.

Xu, J. J.; Xu, D.; Wang, Z. L.; Wang, H. G.; Zhang, L. L.; Zhang, X. B., Synthesis of

Perovskite-Based Porous La0.75 Sr0.25 Mno3 Nanotubes as a Highly Efficient Electrocatalyst for Rechargeable Lithium-Oxygen Batteries. Angew. Chem.-Int. Edit. 2013, 52, 3887-3890. 23.

Liu, Q. C., et al., Flexible and Foldable Li-O2 Battery Based on Paper‐Ink Cathode. Adv.

Mater. 2015, 27, 8095-8110. 24.

Liu, Q. C.; Xu, J. J.; Xu, D.; Zhang, X. B., Flexible Lithium-Oxygen Battery Based on a

Recoverable Cathode. Nat. Commun. 2015, 6, 7892-7900. 25.

Liu, Q. C., et al., Artificial Protection Film on Lithium Metal Anode toward Long-Cycle-Life

Lithium-Oxygen Batteries. Adv. Mater. 2015, 27, 5241–5247. 26.

Jian, Z. L.; Liu, P.; Li, F. J.; He, P.; Guo, X. W.; Chen, M. W.; Zhou, H. S., Core-Shell-Structured

Cnt@Ruo2 Composite as a High-Performance Cathode Catalyst for Rechargeable Li-O-2 Batteries. Angew. Chem.-Int. Edit. 2014, 53, 442-446. 27.

Li, F. J.; Chen, Y.; Tang, D. M.; Jian, Z. L.; Liu, C.; Golberg, D.; Yamada, A.; Zhou, H. S.,

Performance-Improved Li-O-2 Battery with Ru Nanoparticles Supported on Binder-Free Multi-Walled Carbon Nanotube Paper as Cathode. Energy Environ. Sci. 2014, 7, 1648-1652. 28.

Zhang, T.; Liao, K.; He, P.; Zhou, H., A Self-Defense Redox Mediator for Efficient Lithium-O2

Batteries. Energy Environ. Sci. 2016. 29.

Chen, Y. H.; Freunberger, S. A.; Peng, Z. Q.; Fontaine, O.; Bruce, P. G., Charging a Li-O-2

Battery Using a Redox Mediator. Nat. Chem. 2013, 5, 489-494. 30.

Bergner, B. J.; Schurmann, A.; Peppler, K.; Garsuch, A.; Janek, J., Tempo: A Mobile Catalyst

for Rechargeable Li-O-2 Batteries. J. Am. Chem. Soc. 2014, 136, 15054-15064. 31.

Sun, D., et al., A Solution-Phase Bifunctional Catalyst for Lithium-Oxygen Batteries. J. Am.

Chem. Soc. 2014, 136, 8941-8946. 32.

Kwak, W. J., et al., Understanding the Behavior of Li-Oxygen Cells Containing Lii. J. Mater.

Chem. A 2015, 3, 8855-8864. 33.

Qiao, Y.; Ye, S., Spectroscopic Investigation for Oxygen Reduction and Evolution Reactions

with Ttf as a Redox Mediator in Li-O2 Battery. J. Phys. Chem. C 2016, 10.1021/acs.jpcc.5b11692. 34.

Thotiyl, M. M. O.; Freunberger, S. A.; Peng, Z. Q.; Bruce, P. G., The Carbon Electrode in

Nonaqueous Li-O-2 Cells. J. Am. Chem. Soc. 2013, 135, 494-500. 35.

Peng, Z. Q.; Freunberger, S. A.; Chen, Y. H.; Bruce, P. G., A Reversible and Higher-Rate Li-O-2

Battery. Science 2012, 337, 563-566.

37



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

36.

Qiao, Y.; Ye, S., In Situ Study of Oxygen Reduction in Dimethyl Sulfoxide (Dmso) Solution: A

Fundamental Study for Development of the Lithium-Oxygen Battery. J. Phys. Chem. C 2015, 119, 12236-12250. 37.

Sharon, D.; Afri, M.; Noked, M.; Garsuch, A.; Frimer, A. A.; Aurbach, D., Oxidation of

Dimethyl Sulfoxide Solutions by Electrochemical Reduction of Oxygen. J. Phys. Chem. Lett. 2013, 4, 3115-3119. 38.

Mozhzhukhina, N.; De Leo, L. P. M.; Calvo, E. J., Infrared Spectroscopy Studies on Stability of

Dimethyl Sulfoxide for Application in a Li-Air Battery. J. Phys. Chem. C 2013, 117, 18375-18380. 39.

McCloskey, B. D.; Bethune, D. S.; Shelby, R. M.; Girishkumar, G.; Luntz, A. C., Solvents'

Critical Role in Nonaqueous Lithium-Oxygen Battery Electrochemistry. J. Phys. Chem. Lett. 2011, 2, 1161-1166. 40.

McCloskey, B. D.; Bethune, D. S.; Shelby, R. M.; Mori, T.; Scheffler, R.; Speidel, A.; Sherwood,

M.; Luntz, A. C., Limitations in Rechargeability of Li-O-2 Batteries and Possible Origins. J. Phys. Chem. Lett. 2012, 3, 3043-3047. 41.

Kim, S.; Dicosimo, R.; Filippo, J. S., Spectrometric and Chemical Characterization of

Superoxide. Anal. Chem. 1979, 51, 679-681. 42.

Gampp, H.; Lippard, S. J., Reinvestigation of 18-Crown-6 Ether Potassium Superoxide

Solutions in Me2so. Inorg. Chem. 1983, 22, 357-358. 43.

Hunter-Saphir, S. A.; Creighton, J. A., Resonance Raman Scattering from the Superoxide Ion.

Journal of Raman Spectroscopy 1998, 29, 417-419. 44.

Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.; Hendrickson, M. A., Influence of

Nonaqueous Solvents on the Electrochemistry of Oxygen in the Rechargeable Lithium-Air Battery. J. Phys. Chem. C 2010, 114, 9178-9186. 45.

Johnson, L.; Li, C. M.; Liu, Z.; Chen, Y. H.; Freunberger, S. A.; Ashok, P. C.; Praveen, B. B.;

Dholakia, K.; Tarascon, J. M.; Bruce, P. G., The Role of Lio2 Solubility in O-2 Reduction in Aprotic Solvents and Its Consequences for Li-O-2 Batteries. Nat. Chem. 2015, 7, 1091-1099. 46.

Gallant, B. M.; Kwabi, D. G.; Mitchell, R. R.; Zhou, J. G.; Thompson, C. V.; Shao-Horn, Y.,

Influence of Li2o2 Morphology on Oxygen Reduction and Evolution Kinetics in Li-O-2 Batteries. Energy Environ. Sci. 2013, 6, 2518-2528. 47.

Zhai, D. Y., et al., Raman Evidence for Late Stage Disproportionation in a Li-O-2 Battery. J.

Phys. Chem. Lett. 2014, 5, 2705-2710. 48.

Meini, S.; Tsiouvaras, N.; Schwenke, K. U.; Piana, M.; Beyer, H.; Lange, L.; Gasteiger, H. A.,

Rechargeability of Li-Air Cathodes Pre-Filled with Discharge Products Using an Ether-Based Electrolyte Solution: Implications for Cycle-Life of Li-Air Cells. Phys. Chem. Chem. Phys. 2013, 15, 11478-11493. 49.

Harrick, N. J., Study of Physics and Chemistry of Surfaces from Frustrated Total Internal

Reflections. Phys. Rev. Lett. 1960, 4, 224-226. 50.

Harrick, N. J., Surface Chemistry from Spectral Analysis of Totally Internally Reflected

Radiation. J. Phys. Chem. C 1960, 64, 1110-1114.

38


Page 38 of 55

Page 39 of 55

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


Table of Contents (TOC)

39



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

Spectroscopic Investigation for Oxygen Reduction and Evolution Reactions on Carbon Electrodes in Li-O2 Battery Yu Qiao and Shen Ye*


Page 40 of 55

Page 41 of 55

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


Characterization for the products deposited on electrode Qualitative analysis Mid/Far-IR Calibration curves

Electrochemical performance SwagelokTM type cell

in situ UV-vis cell

Pre-treatment

Characterization for the products dissolved in electrolyte Pre-treatment

DMC washing cycled electrode Vacuum drying (55°C 4h)

Ultrasonication (30 s) Centrifuge (3000rpm 5min)

Quantitative analysis (components contents)

LiO2 species clarified solution

Scheme 1. The pre-treatment procedures and characterizations for the samples used in the study.


UV-vis


(a)

(b) 4.5

4.5

4.0

4.0

Potential (V vs. Li/Li+)


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

Page 42 of 55

Ketjenblack (porous) Glassy Carbon (flat)

3.5 3.0 2.5

(glassy carbon) 16 µAh/cm2

2.0 0.0

0.5

1.0

(Ketjenblack) 2.25 mAh/cm2

1.5

OER-Capacityb (4.5 V) 9.6 µAh/cm2

3.5 Glassy Carbon (flat)

3.0 2.5 ORR-Capacity (2.0 V) 16 µAh/cm2

2.0 2.0

0

2

4

6

8

10

12

14

16

2

2

Capacity (µAh/cm )

Capacity (mAh/cm )

Figure 1. (a) Galvanostatic discharge-charge curves observed on a Ketjenblack cathode (black trace, in a SwagelokTM type cell) and a glassy carbon cathode (blue trace, in a bulk electrolysis cell) at 0.1 mA/cm2 in 0.5 M LiClO4-DMSO electrolyte under O2. (b) Magnified ORR/OER curves observed on the glassy carbon electrode.


(b)


Discharge

2

1.6 mAh/cm

1.2 mAh/cm2 0.8 mAh/cm2 0.4 mAh/cm2 OCP

Li2CO3

200

300

400

Li2O2

500

LiAc

End of ORR 2.25 mAh/cm2 2 mAh/cm2 1.6 mAh/cm2 1.2 mAh/cm2 0.8 mAh/cm2 0.4 mAh/cm2 OCP

600

700

800

-1

Wavenumber (cm )

(d)

-1

313 cm

-1

420 cm

-1

488 cm

-1

621 cm-1 658 cm

LiAc Standard Substance

1000 1200 1400 1600

Wavenumber (cm-1) -1

593 cm

-1

1609 cm st

End of 1 ORR (2.25 mAh/cm2)

Li2O2 LiAc

Mid-IR Absorbance (a.u.)

(c)

Li2SO4

Li2O2 Li2CO3

Standard Substance

600

Discharge


End of ORR 2.25 mAh/cm2 2 mAh/cm2

Far-IR Absorbance (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

(a)


Page 43 of 55

-1

862 cm

-1

1104 cm

End of 1st ORR (2.25 mAh/cm2) Li2O2 LiAc Li2CO3 Li2SO4

Li2CO3 200

300

400

500

Wavenumber (cm-1)

600

700

600

800

1000

1200

1400

1600

-1

Wavenumber (cm )

Figure 2. The ex situ (a) far-IR and (b) mid-IR spectra on a Ketjenblack electrode surface are shown at different ORR ACS Paragon Plus Environment stages marked by different colors and the corresponding capacities. The spectra are offset for clarity. Figures 2(c) and 2(d) shows (c) far-IR and (d) mid-IR spectra observed recorded at the end of 1st ORR process during the galvanostatic


x=x2 x=x1

IR 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

fmix(x)

x=x3

Mixture of 1, 2 and 3

f3,0 (x)

Component-3

f2,0 (x)

Component-2

f1,0 (x)

Component-1 x1 x2 x3

Wavelength

Scheme 2. Schematic models for mathematic analysis of IR spectra. The IR spectrum (top) for a mixture, fmix(x), can be regarded as a linear combination of three different components, fi,0(x) (i=1,2,3). See text for details.


Page 44 of 55

Page 45 of 55

(b)

5

LiAc 1609 cm-1

ILi2CO3@862cm-1 / ILi2SO4@1104cm-1

(a) ILi2CO3@862cm-1 / ILiAc@1609cm-1

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


Li2CO3

4

862 cm-1

3 750

1000

1250

1500

Wavenumber (cm-1)

2 1 0 0

5

10

15

10

Li2SO4 Li2CO3

1104 cm-1

862 cm-1

8

6 750

1000

1250

1500

Wavenumber (cm-1)

4

2

0

20

Li2CO3: LiAc in moles

0

10

20

30

40

50

Li2CO3:Li2SO4 in moles

Figure 3. The calibration curves obtained by mixing known amounts of standard materials with different molar ratios. (a) Li2CO3 (862 cm-1) and LiAc (1609 cm-1); (b) Li2CO3 (862 cm-1) and Li2SO4 (1104 cm-1). Insets show the mid-IR spectra of the standard materials. The blue diamond highlights the value obtained from the mid-IR spectra observed on a Ketjenblack electrode after 1st ORR process (Fig. 2b) and the corresponding molar ratios are also listed in the blank. See text for details.



-1

538 cm

End of 1stORR (2.25 mAh/cm2) Mixture: Li2O2 and M powder (6 :1)

M powder

Li2O2 200

300

400

500

600 -1

Wavenumber (cm )

700

(c)

418cm-1 / 488cm-1

1.2

Far-IR I@538cm-1 / I@621cm-1

418 cm

(b)

621 cm-1

-1

Far-IR I@418cm-1 / I@488cm-1

488 cm-1

Abs = 0.1


(a)

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

Page 46 of 55

inverse proportional function fit

1.1

1.0

0.9

0.8

538cm-1 / 621cm-1

1.6

inverse proportional function fit

1.4

1.2

1.0

0.8 0

2

4

6

8

Li2O2 : M powder in moles

10

0

2

4

6

8

10

Li2O2 : M powder in moles

Figure 4. (a) The ex situ far-IR spectra observed on a Ketjenblack electrode surface recorded at the end of the first ORR (blue trace). The spectra of standard Li2O2 (black trace) and the “M powder” (red trace, see text) and a mixture of Li2O2 and “M powder” in the molar ratio of 6:1 (pink trace) are shown in the same figure. The far-IR calibration curves obtained by mixing known amounts of Li2O2 and “M powder” with different molar ratios were made by using a pair of peaks at (b) 418 and 488 cm-1; and (c) 538 and 621 cm-1. The ratio of peak intensity in each pair is observed as y axis in the calibration curve. The blue diamond highlights the value obtained from the far-IR spectra observed on the Ketjenblack electrode after the first ORR. The detail of the calibration method is shown in the Supporting Information (Fig. S2).


Page 47 of 55

3.6

(b)

in-situ bulk cell modification

3.4

1.8

Blowing O2 / PTFE film No / No No / double sides Yes / single side Yes / double sides Swagelok-type coin cell (for comparison)

3.2 3.0 2.8

UV-vis Absorbance

(a) Potential (V vs. Li/Li+)

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


2.6 2.4 Galvanostatic discharge (0.1 mA/cm2)

2.2

252 nm

1.5

0.9 0.6 0.3

2.0

OCP state

0.0 0.0

0.5

1.0

1.5

0.34 mAh/cm2 0.58 0.82 1.06 1.3 1.58 1.84 2.08

1.2

2.0

2.5

250

2

Discharge Capacity (mAh/cm )

260

270

280

290

300

Wavelength (nm)

Figure 5. (a) Galvanostatic ORR on a Ketjenblack carbon cathode at a current density of 0.1 mA/cm2 in O2-saturated 0.5 M LiClO4-DMSO electrolyte in a UV-vis bulk electrolysis cells with different kinds of geometries. (b) UV-vis spectra observed at different ORR depths in the double-sides PTFE covered condition (corresponds to pink trace in Fig. 5a).


3.0

4.5 3.0

2.5

1.5 0.0 7.5

2.0

single-side PTFE

3.5

6.0 3.0

4.5 3.0

2.5

1.5 2.0

0.0

Page 48 of 55

(b) 2.25 2.10 1.95 1.80

Superoxide (µmol/cm2)

Superoxide (µmol/cm2) Superoxide (µmol/cm2)

6.0

3.5

double-side PTFE


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

7.5


(a)


2.0 µmol/cm2 Molar amunt of LiO2 inside porous carbon

0.8

0.4 mAh/cm2 Point B remove the top PTFE film

0.6 0.4

0.5

1.0

1.5

2.0

2

1.2 mAh/cm2

Point A double PTFE

1.6 mAh/cm2

single PTFE

0.2

2 mAh/cm2 free-stay in UV cell

2.25 mAh/cm2 (end of 1st ORR)

0.0 0

0.0

0.8 mAh/cm2

2.5

Discharge Capacity (mAh/cm )

5

10

15

20

25

30

35

40

Time (min)

Figure 6 (a) Molar amount of superoxide species trapped in the porous carbon electrode (blue symbols) and in the bulk electrolyte solution (red symbols) at different ORR capacities for the porous carbon electrode sandwiched by PTFE films (top) and covered by one PTFE film on surface (bottom). The amount of superoxide is estimated from the UV-vis peak intensity at 252 nm. (b) Temporal profiles for UV absorbance at 252 nm recorded in the bulk solution before/after removal of the PTFE film on the surface of the porous carbon electrode at different ORR depths. See text for details.


Page 49 of 55

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


(a)

Glass Carbon (ORR capacity:16 μAh/cm2)

x100

(b)

Li2O2 (6.3 %)

LiO2 (93.6 %)

Ketjenblack Carbon (ORR Capacity: 2.25 mAh/cm2) Li2CO3 (18.2 %) Li2O2 (73.8 %)

LiAc (4.3 %) Li2SO4 (1.3 %) LiO2 (2.4 %)

Figure 7. The reaction yields at the end of the first ORR on (a) a glassy carbon electrode and (b) a porous carbon electrode. The area of the pie is proportional to the ORR capacity. Note that the area of the pie for the glass carbon (right) is magnified 100 times for clarity. ACS Paragon Plus Environment


(b)

OER end of OER 0.4 mAh/cm2 2

0.8 mAh/cm

2

1.2 mAh/cm

1.6 mAh/cm2 2

2 mAh/cm

2.25 mAh/cm2

Li2CO3

Li2O2

LiAc


(a) Far-IR Absorbance (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

charge end of OER 0.4 mAh/cm2 0.8 mAh/cm2 1.2 mAh/cm2 1.6 mAh/cm2 2 mAh/cm2 Li2O2

300

400

500

600

700

2.25 mAh/cm2 Li2CO3

Standard Materials 200

Page 50 of 55

600

-1

800

Li2SO4

Standard LiAc Materials

1000 1200 1400 1600

Wavenumber (cm-1)

Wavenumber (cm )

Figure 8. The ex situ IR spectra observed at different OER states in (a) far-IR and (b) mid-IR region. The ORR/OER states are marked by different colors and the corresponding capacity. The spectra are offset for clarity. IR spectra of standard materials (Li2O2, LiAc, Li2CO3 and Li2SO4) are also shown at the bottom of each figure for comparison.


Page 51 of 55

0.04

Li2CO3 @ 862 cm-1

Mid-IR Absorbacne

LiAc @ 1609 cm-1 0.03

4.5 4.0 3.5

0.02

3.0 0.01 2.5 0.00


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


2.0 0.0

0.5

1.0

1.5

2.0

2.5

2

Capacity (mAh/cm ) Figure 9. Capacity dependence of the mid-IR peak intensities at 862 cm-1 (red triangles, Li2CO3) and 1609 cm-1 (blue triangles, LiAc) observed on the Ketjenblack electrode surface during the first ORR/OER in Figure 8b. The galvanostatic ORR/OER curve is shown in the same figure for comparison. ACS Paragon Plus Environment


(b) 4.5 4.0

ORR OER 0.4

3.5 3.0

0.2

2.5 0.0

2.0 0.0

0.5

1.0

1.5 2

Capacity (mAh/cm )

2.0

ORR

Relative amount of Li2CO3

Li2O2 FIR Fitting Factor

0.6

(c) 0.020

0.04 OER

Relative amount of LiAc

(a)


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

Page 52 of 55

0.03

0.02

0.01

0.00

0.015

MIR absorbance @ 1609 cm-1 for LiAc ORR

0.010

OER

0.005

0.000 0.0

0.5

1.0

1.5

2.0

2

Capacity (mAh/cm )

0.0

0.5

1.0

1.5

2.0

2

Capacity (mAh/cm )

Figure 10. Relative amounts of (a) Li2O2, (b) Li2CO3 and (c) LiAc during the ORR/OER estimated based on Eq. 1 from the far-IR spectra. See text for details.



4.0 3.5 3.0 2.5 2.0

0.78 0.85 0.95 0.0

0.5

5th 4th 3rd 2nd 1st Li2O2 Li2CO3 LiAc

2.25

1.32

1.0

1-5th after ORR process (Far-IR)


1st 2nd 3nd 4th 5th

4.5

1.5

200

2.0

(c)

(d) 1-5th after ORR process (Mid-IR)

600

800

1000

Li2SO4

1200

500

600

5th

5th 4th 3rd 2nd 1st Li2CO3

400

LiAc

1400

700

Wavenumber (cm ) Different Component after ORR process (on the electrode surface)

Capacity (mAh/cm )

Li2O2

300

-1

2


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

(b)


1600

Cycle Number

Page 53 of 55

Li2O2 Li2CO3 LiAc Li2SO4

4th 3rd 2nd 1st 0

20

-1

40

60

80

100

Components Content

Wavenumber (cm ) ACS Paragon Plus Environment

Figure 11. (a) Galvanostatic ORR/OER curves during 1-5th cycles on a Ketjenblack cathode at a current density of 0.1 mA/cm2 in 0.5 M LiClO4-DMSO electrolyte under O2 atmosphere. Figures 11b and 11c show ex situ IR at the end of


(b)

4.5V 4.35V

4.5 4.0 3.5

2nd ORR

3.0

1st cycle

2.5 2.0



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

Page 54 of 55

2nd ORR (4.5V) OER (@4.5V)

2nd ORR (@4.35V) OER (@4.35V)

1st ORR Li2O2 Li2CO3 LiAc Li2SO4

2 0.71 mAh/cm2 1.32 mAh/cm

0.0

0.4

0.8

1.2

1.6

2.0

600

800

1000

1200

1400

1600

Wavenumber (cm-1) Capacity (mAh/cm2) Figure 12. (a) Galvanostatic ORR/OER curves of the first cycle (black trace) and second ORR at OER upper potential limits of 4.5V (blue trace) and 4.35 V (red trace). (b) Mid-IR spectra at different ORR/OER states. IR spectra of standard materials are show at bottom for comparison. See text for details.


Page 55 of 55

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


ORR Yields on the Different Electrodes Li2O2

Glassy Carbon (Capacity: 16 μAh/cm2)

Porous Carbon (Capacity: 2.25 mAh/cm2)

(on surface) LiO2 (in solution)

Li2O2

x100


Li2CO3 LiAc Li2SO4 LiO2

(in solution)

Spectroscopic Investigation for Oxygen Reduction and Evolution

Recommend Documents