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Reactivity of Amide Based Solutions in Lithium-Oxygen Cells Daniel Sharon, Daniel Hirsberg, Michal Afri, Arnd Garsuch, Aryeh A. Frimer, and Doron Aurbach J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/jp506230v • Publication Date (Web): 24 Jun 2014 Downloaded from http://pubs.acs.org on July 1, 2014

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The Journal of Physical Chemistry

Reactivity of Amide Based Solutions in Lithium-Oxygen Cells Daniel Sharon†, Daniel Hirsberg†, Michal Afri†, Arnd Garsuch‡, Aryeh A. Frimer† and Doron Aurbach†* †

Department of Chemistry, Bar Ilan University, Ramat-Gan, 52900, Israel ‡

BASF SE, GCI/E - M311, Ludwigshafen, 67056, Germany *E. Mail address: [email protected]

1 Environment ACS Paragon Plus


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Abstract The stability of electrolyte solutions during lithium-oxygen cells operation is of great importance and interest. This is because oxides formed during reduction are strong nucleophiles which can initiate solvent decomposition. The highly polar amide based solvents have come to the fore as possible candidates for Li-O2 applications. They show typical cycling behavior as compared to other solvents; however, their stability toward lithium oxides is shrouded in doubt. The present study has focused on Li-O2 cells containing electrolyte solutions based on DMA/LiNO3. We have used various analytical tools, to explore the discharge-charge processes and related side reactions. The data obtained from FTIR, NMR, XPS and EQCM all support a rational decomposition mechanism. The formation of various side products during the course the first discharge, leads to the conclusion that amide based solvents are not suitable for Li-O2 applications; however, electrolyte solution decomposition reduces the OER overpotential by forming oxidation mediators.

Keywords: Oxygen electrochemistry, Li-air batteries, EQCM, carbon electrodes, lithium oxides, redox mediator.


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1. Introduction The reality of non-aqueous lithium-air cells remains elusive. The high working potential of non-aqueous cells, and the infinite supply of oxygen as a redox couple, make this energy source ever so appealing. Yet, for this battery to be viable, some intrinsic problems must be resolved. The major challenge, according to many publications1,2,3,4, is the highly reactive nature of the lithium oxides (before precipitation as solid products), formed during discharge, toward both the electrolyte solution and to the carbon-based cathode5,6. The effect that the choice of electrolyte solution has on the cell has been the focus of many studies. Inter alia, there have been scores of studies on the role played by carbonates,7 polyethers,8 sulfoxides,9 silanes10 and ionic liquids11. New solvent candidates with possible stability toward oxygen radicals are the amide based solvents like DMA, DMF and NMP.12,13 One significant disadvantage of the amide based solvents is their very facile reactivity toward lithium metal. The reaction between the two is extremely fast and does not allow reversible lithium metal cycling that is necessary for metal air battery operation. To circumvent this problem, Uddin et al. proposed the use of LiNO3 as a passivation agent for lithium metal in these systems14. The goal of this paper is to determine whether amide based solutions are indeed suitable for oxygen reduction applications. Based on the analysis of the Li-O2 cells after cycling, and identification of the different side products by FTIR, NMR, EQCM, XPS XRD, we will propose a new rational decomposition mechanism of the amide solvents during ORR.



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2. Experimental Materials. Lithium nitrate (99%), and dimethylacetamide were purchase from Sigma Aldrich. Activated microfiber carbon fiber electrodes purchase from Cellguard, glassy separators were purchase from Whatman. Products analysis. An Inspect FEI microscope was employed for the high resolution scanning electron microscopy (HR-SEM) studies of pristine and cycled electrodes. The presence of lithium oxides on ACM electrodes was analyzed using a Bruker D8 Advance X-ray Diffractometer (2θ = 20-80°) working with Cu-Kα radiation (λ = 0.15418 nm). Surface species formed on discharged and charged ACM electrodes were characterized using a Nicolet 6700 FTIR spectrometer placed in a moisture and CO2-free glove box (Ravona Inc., Israel). The FTIR spectra of these electrodes (in the range of 4000-400 cm-1) were recorded with a diffuse reflectance accessory from Pike Technologies. Nuclear magnetic resonance (NMR) spectra were recorded using a Bruker 300 NMR spectrometer. X-ray photoelectron spectroscopy (XPS) analysis of cycled electrodes was carried out using a Kratos Axis HS spectrometer (England) equipped with an Al-Kα X-ray radiation source (photon energy 1486.6 eV). A homemade transfer system equipped with a gate valve and a magnetic manipulator were used for the transfer of the highly sensitive samples from the highly pure argon atmosphere of the glovebox to the XPS system EQCM measurements. Electrochemical quartz crystal microbalance (EQCM) measurements were carried out with a GAMRY (Warminster, PA, USA) EQCM system. The working electrodes for these studies were gold discs deposited on 5 MHz AT-cut quartz crystals (1.38 cm2) (SRS, Sunnyvale, CA, USA). A 3 ml electrochemical cell made of glass containing lithium counter and reference electrodes


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was used for EQCM measurements. Before measurements, the electrolyte solutions were bubbled with high-purity oxygen for 30 min., and the cells were purged with O2 (99.995%) throughout the duration of the experiment. Li-Oxygen cells operation. Cells were prepared from monolithic carbon paper cathodes in DMA/LiNO3 (1M) with a lithium metal anode. A glassy separator inbetween two polyethylene separators were incorporated in each cell. The cell contain a teflon body squeezed between two stainless steel plates and two valves for insertion of pure oxygen. Each cell was flushed with pure oxygen for one minute and than held in a 1 atm oxygen for the rest of the experiment. 3. Results and Discussion In previous studies, the EQCM method assisted our investigation of the ORR mechanism in various solvents, and allowed us to confirm their reversible behavior during cycling.15,16,17 The frequency changes occurring on the QCM during cycling voltammetry are then converted to mass by applying the basic Surbaey equation, which correlates between changes in frequency ( changes in mass (

) and

).

Accurate calculation of mass per electron (mpe) in the lithium air cell is very challenging for two reasons. First, some ORR products are soluble and do not absorb on the crystal. Second, products like Li2O2 can be formed by disproportionation of LiO2. These reactions do not allow us to assume that every transfer of an electron leads to deposition of ORR, and that every change in mass is accompanied by an electron transfer.



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Figure 1. (a) Cyclic voltammetry and (b) EQCM response of gold disks deposited on thin quartz crystals in DMA/LiNO3 (1M). Even with these limitations, EQCM can still serve as an effective qualitative method. As shown in Figure 1, the voltammetric and mass response of the cell is measured by the counter and reference lithium electrodes in LiNO3/DMA (1M). The voltage response in an argon atmosphere shows that the electrolyte solution is inert during the cathodic sweep. However, throughout the anodic sweep, starting at ca. 3.7 V, one observes reactions corresponding to the oxidation of the electrolyte solution. Indeed, the oxidation of LiNO3 is a well-known process14. The nitrate is first reduced by the lithium metal anode to the nitrite anion (NO2-), which is oxidized in turn to nitrogen dioxide (NO2) at higher potentials.

During the oxidation of the electrolyte, no

prominent changes in the mass response are observed - indicating the formation of soluble species like NO2, during oxidation.


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The CV of the cell operated under oxygen atmosphere corresponds to a multistep process during reduction. The peak at 2.7 V can be assigned to the formation of the superoxide radical. The stabilization of the superoxide radical by the highly polar DMA is not surprising. Like DMSO, amide solvents have a high Guttmann donor number9 that helps to stabilize the positively charged lithium cation. Without stabilization, lithium cation will react immediately to form LiO2 that disproportionates to Li2O2. DMA can delay these reactions at low overpotentials. Quantitatively, the EQCM and the electrochemical response cannot be the same in these systems, because some of the processes that change the electrode’s mass loading may relate to non-electrochemical precipitation and dissolution/desorption processes. The behavior of the mass and current changes shows a good fitting between them. During the cathodic scan, mass accumulates on the electrode and, as we sweep the potential back anodically, we observe mass removal from the electrode surface. It is noticeable that the mass curve has two distinctive slopes during ORR: one between 2.8 and 2.6 V, and a second between 2.6 and 2.0 V. The slope that corresponds to the first ORR peak does not change much from open circuit voltage conditions. Even with oxygen reduction taking place, there are no pronounced mass changes. Such behavior suggests the formation of a soluble oxide. For the second slope, we observe mass accumulation, suggesting the deposition of insoluble oxides. Our proposed ORR mechanism for the DMA/LiNO3 system is outlined in equations 1-3. Thus, soluble superoxide radical is first formed stabilized by DMA (eq. 1); while subsequent sweeping to lower potentials, deposits LiO2 and Li2O2 on the quartz crystal surface (eq. 2 and 3).



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During OER at 3.7 V, an oxidation peak is observed. The QCM response supports removal of most of the mass (conversion of Li2O2 to O2). Interestingly, the oxidation potential of the OER starts at the same potential as the LiNO3 electrolyte decomposition to NO2- and NO2, as discussed above. This oxidation of the electrolyte solution during cycling is clearly a significant disadvantage. Nevertheless, we believe that there is a silver lining, since the NO2/NO2- redox couple can behave as a redox mediator. [The latter is a soluble electroactive molecule that can interact with the oxides on the cathode surface, and oxidize them at lower potentials.18,19,20] Thus, the nonconductive Li2O2 deposits covering the cathode at the end of the discharge process, require high overpotentials to be oxidized. We assume that NO2- is first oxidized on the cathode, and then during oxidation the strong oxidizer NO2 is formed that can enhance Li2O2 oxidation by undergoing reduction back to NO2-. Since in other systems15,16 as well, the oxidation of the Li2O2 occurs almost simultaneously with electrolyte solution decomposition, we believe that this redox mediator behavior exists in many other electrolyte solutions, as well. Voltage profiles of the carbon paper electrode in DMA/LiNO3 (1M) are presented in Figure 2. The cells were operated in the voltage ranges of 2 to 3.9 V at a current density of 0.05 mA cm-2; the average discharge voltage of the cells was around 2.75 V. The charging profile is comprised of several steps that include the oxidation of both the Li2O2 and the electrolyte solution. As mentioned, we suspect that the relatively low potential used for Li2O2 oxidation is due to the redox behavior of the soluble electrolyte fragments.


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2

Figure 2. Galvanostatic discharge-charge profile of a carbon electrode at 0.05 mA/cm in oxygen atmosphere in DMA/LiNO3 (1M).

XRD patterns (Figure. 3) of the cathode polarized to 2.0 V exhibit crystalline peaks characteristic of Li2O2. The cathode pattern charged to 3.9 V was identical to that of pristine carbon cathode, supporting the oxidation of Li2O2.

Figure 3. XRD patterns of carbon cathodes discharged to 2 V and then charged to 3.9 V in DMA/LiNO3 (1M).

High resolution scanning electron microscopic (HR-SEM) images of pristine, discharged, and charged carbon electrodes in DMA/LiNO3 (1M) electrolyte solution



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under dry oxygen (1 atm.) are presented in Figure 4. These binder-free electrodes consist of carbon microfibers shown in Figure 4a. The carbon electrodes, discharged to 2.0 V in DMA/LiNO3 (1 M) solution, were uniformly covered with sub-micronic sized toroid shaped particles (Figure 4b and 4c), characterized in the literature as Li2O2 deposits2122. One can also observe that the toroid shaped particles are surrounded by dense forest of nano sized tubes that can point to the growth mechanism of the Li2O2 particles. Interestingly, electrodes charged to 3.9 V (Figure 3d) were free of any visible deposits.

Figure 4.

HR-SEM images of: (a) pristine carbon electrodes (b,c) electrode

discharged to 2 V and (d) electrode charged to 3.9 V in DMA/LiNO3 (1 M) . The composition of the surface film has a decisive influence on the performance of the Li-O2 cell. Thus, it is necessary to understand how amide-based electrolyte solutions affect the nature of the surface film. For this purpose, cycled cathodes in DMA/LiNO3 electrolyte solutions were characterized using ex-situ XPS and FTIR techniques. The cycled electrodes were thoroughly washed and dried under ultrahigh vacuum before the spectral measurements. It is, therefore, assumed that there was no residual electrolyte solution remaining on the electrodes during these measurements.


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The general elemental composition of the surface films obtained from XPS measurements can be calculated from Figure. 5. The pristine electrode contains 95% carbon and 5% oxygen. By discharging the cell to 2 V, we receive a carbon cathode that is covered by 64% oxygen, supporting the formation of oxygen rich species on the surface. Upon charging to 3.9 V, most of the oxygen was removed from the surface and we are left with 10% oxygen. The increase in oxygen content as compared to the initial state can be related to the oxidation of the carbon surface or to deposits of irreversible side products.

Figure 5. O 1s and C 1s XPS spectra of (a) pristine carbon cathodes (b) discharged to 2 V and (c) charged to 3.9 V in DMA/LiNO3 (1M).

The high-resolution XPS spectra in Figure 6 show the C, O, N and Li of the carbon cathode. The C 1s spectrum (c) of the pristine cathode is a broad peak, which together with the corresponding broad O 1s peak suggests that the carbon surface is composed of more than one kind of oxygen group. The discharged cathode C 1s spectrum (b) contains two additional broad high binding energy peaks in the 288-290 eV region,



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that readily correspond to the acetate and amide carbonate group located at 289.5 eV and 288.7 eV peaks. Turning now to the O 1s and Li 1s peaks of the discharged cathode spectra (b), we can confirm that oxygen containing species on the surface are mainly composed of Li2CO3 at 531.8 eV and to some extent Li2O2 at 531.2 eV. The absence of carbonate peaks in the XRD pattern and low concentration of Li2O2 on the carbon surface suggest that Li2CO3 is a surface phenomenon that deposits on the bulk Li2O2 during the discharge process. The low potentials that are applied during discharge support the assumption that this Li2CO3 is related to solvent decomposition and not to the reaction between the peroxide and the carbon electrode23. The N 1s spectrum (b) of the discharged cathode reveals two kinds of nitrogen groups: one over 400.5 eV associated with the amide group from the solvent decomposition, and second at high binding energy of 407.8 eV attributable to NO3 from the electrolyte. The charged cathode spectra of Li, C and O 1s confirm that Li2O2, Li2CO3 and LiO2CCH3 are oxidized at a voltage of 3.9 V. However, the N 1s spectrum (c) of the charged cathode exhibit irreversible behavior with respect to the nitrogen content. The peaks of the amide group around 400.5 eV and the nitrate moiety at 408 eV are still present in the charged cathode. An additional peak is detected at 406 eV and is associated with the nitrite (NO2) group. The presence of NO2 on the carbon surface is related to the oxidation mechanism of the electrolyte; as mentioned above, at high potentials, the NO2 formed in situ can oxidize Li2O2 and apparently the carbon cathode as well. The ability to remove Li2CO3 at 3.9V is very impressive since the potentials reported for the oxidation of carbonates are even higher than peroxides24.


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Figure 6. C, O, N, Li 1s spectra of carbon cathodes (a) charged to 3.9 V (top), (b) discharged to 2 V (middle), and (c) pristine (bottom), using DMA/LiNO3 (1M) as electrolyte solution. The FTIR spectrum of the pristine carbon paper is presented in Figure 7a. The peak at 1630 cm-1 is attributable to C=C bonds, while the broad peak between1270-1400 cm-1 can be assigned to the C-O and C-H bonds of the carbon. Figure 7b is an FTIR spectrum of the cathode discharged to 2 V, where the high intensity and multiple peaks indicate a complex interface formed during discharged. The peaks located between 450 and 650 cm-1 correspond to Li-O bonds of lithium peroxides and acetates. The two peaks at ca. 850 and 1470 cm-1 are the signature of Li2CO3, while the peak located at 1415 cm-1 corresponds to the amide group. The broad peak located at 1645 cm-1 corresponds to the C=O bond of the amide carbonyl. There are additional peaks located between 1260 to 1400 cm-1 and 960 to 1075 cm-1 which presumably belong to more complex lithium alkoxides and carbonates fragments formed during the ORR. In Figure 7c, the FTIR spectra of the charged cathode is similar to that of the pristine cathode, while the additional peak around 1660 cm-1 may correspond to the amide bond. This data support the XPS results and suggests that irreversible decomposition of the amide solvent is taking place on the carbon cathode surface.



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Figure 7. FT-IR spectrum of (a) pristine carbon electrodes (b,) electrode discharged to 2 V and (c) electrode charged to 3.9 V in DMA/LiNO3 (1 M). The residual electrolyte solution from DMA/LiNO3 cells was analyzed by 13C-NMR. NMR analysis of both the discharged and charged states did not show the formation of any soluble products. However, in Figure 8b we can see that, after prolonged cycling of the cell (Figure S1), the

13

C-NMR spectrum exhibits a peak at 39 ppm25

that corresponds to dimethylamine dissolved in the electrolyte solution.

Figure 8.13C-NMR spectrum of (a) pristine DMA/LiNO3 (b) and cycled to 2V DMA/LiNO3 electrolyte solution.


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The above results suggest a very troubling picture for those who have held out hopes for the future of Li-O2 batteries. The various species observed, including dimethylamine, acetate, carbonate, various N-O species and amide, indicate that the decomposition of both solvent and electrolyte is taking place during the first cycle. The original assumptions of Li-ion battery research had been that the side reactions involving electrolyte solution decomposition are negligible and will stop after a few cycles. However, these assumptions have proven incorrect in the context of Li-O2 systems. The surface film formed during Li-ion battery operation can transport Li cations. Actually, a passivation layer of this kind is detrimental for Li-O2 cells. In metal air batteries the most important location is the electrode-solution interface where the air, the electrolyte solution, and the conductive substrate make contact. Irreversible precipitation of solution decomposition products blocks active cathode area and any decrease in the active surface of the carbon is undesirable. Based on the above data, we suggest that the electrolyte solvent DMA (1) undergoes oxidative cleavage and decomposition mediated by the various superoxide and peroxide species present in solution, as outlined in Scheme 1.

Base Catalyzed Autoxidation

Path c

N CH3

MO 2

C CH 3 1

CH3 N M 3

CH3 N CH3

CH3 N

OO M

9

a

M C O

MO 2

C CH2

O

CH3

O

O

MO 2/O2

O

CH3 N

CH3

7

a,b

C

CH3 b CH3 O O M 2 b

Path a

M

M CH3 N 3 CH 3

CH3 N

CH3 MO2

CO3-2 8

CH3

M C O

O

+ O O C CH3 [H]

Path b

M

CH3 N MO 2

7

C O CH 3

CH3

MO2 = LiO 2 or Li2O2


O O C CH3

O

O + CO2

See below

6

4

5


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Scheme 1. Proposed Oxidation Mechanism of Dimethylacetamide The above analysis is of course qualitative. We cannot estimate yet the quantitative aspects and proportion between the main oxygen reduction/evolution reactions and the side reactions. Once LiO2 and then Li2O2 are generated, they can attack the carbonyl of DMA (1) yielding α-peroxyaminal 2. The latter can decompose (via path a) by eliminating dimethylamine anion 3 and peracetate 4 and unltimately acetate 5. 26 On the other hand, Baeyer-Villiger rearrangement (path b) yields methyl N,Ndimethylcarbamate 6.27 The latter presumably undergoes saponification to 7 and decarboxylation to the anion of dimethylamine 3 and carbonate 8. Alternatively, base catalyzed autoxidation (path c)28,29,30 of the acidic acetyl carbon of acetamide would lead to α-hydroperoxyketone 9 which would readily cleave yielding N,Ndimethylcarbamate anion 7. Decarboxylation would yield the anion of dimethylamine 3 and carbonate 8, as before. 4. Conclusion Amide based solvents were found to be unsuitable for oxygen reduction applications. Both solvents and the electrolyte salt are decomposed either by lithium oxides or high oxidation potential. The mechanism offered suggests, once again, that reduction of oxygen in lithium containing solutions leads to the formation of highly reactive oxides that govern the cell performance. However, it is important to remember that until now no other solvent has shown inert behavior towards ORR; hence, amide solvents may still be considered a default option. The effect of the soluble fragments from the decomposed electrolyte solutions on the OER needs more investigation. The NO2/NO2- is just one example that is easy to visualize; however, we believe that almost every system can produce these in-situ


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mediators - either from decomposition products or impurities, and thereby reduce the overpotential. AUTHOR INFORMATION Corresponding Author *E-mail: (([email protected])) ACKNOWLEDGMENT The research was supported by the BASF scientific network of electrochemistry and batteries. AAF thanks the Ethel and David Resnick Chair in Active Oxygen Chemistry as well as the Israel Science Foundation (Grant no. 1469/13) for their kind and generous support .

ASSOCIATED CONTENT Supporting Information. Electrochemical behavior of carbon electrode and EQCM theory and calculations. This material is available free of charge via the Internet at http://pubs.acs.org

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