Technical CommentsChemistry

Response to Comment on “Cycling Li-O2 batteries via LiOH formation and decomposition”

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Science  06 May 2016:
Vol. 352, Issue 6286, pp. 667
DOI: 10.1126/science.aad8843

Abstract

Lithium-oxygen (Li-O2) batteries cycle reversibly with lithium iodide (LiI) additives in dimethoxyethane (DME) to form lithium hydroxide (LiOH). Viswanathan et al. argue that because the standard redox potential of the four-electron (e) reaction, 4OH ↔ 2H2O + O2 + 4e, is at 3.34 V versus Li+/Li, LiOH cannot be removed by the triiodide ion (I3). However, under nonaqueous conditions, this reaction will occur at a different potential. LiOH also reacts chemically with I3 to form IO3, further studies being required to determine the relative rates of the two reactions on electrochemical charge.

We recently reported (1) on a highly reversible lithium-oxygen battery composed of a reduced graphene oxide (rGO) electrode and a dimethoxyethane (DME)/LiTFSI electrolyte. The additives (H2O and LiI) were key in controlling the nature of the battery reactions. LiOH, instead of the commonly reported Li2O2 phase, was the predominant product during discharge, the protons primarily coming from H2O in the cell. On charge, LiOH was removed [as seen by x-ray diffraction and proton nuclear magnetic resonance (1H NMR)], and we proposed that this occurs via a reaction with I3, O2 being observed by mass spectroscopy. In their Comment, Viswanathan et al. (2) argue that a charge mechanism involving reaction with I3 to produce O2 is not feasible, because the redox potential for reaction 1Embedded Image(1)is 3.34 V versus Li+/Li under standard conditions, whereas the redox potential for 6I ↔ 2I3 + 4e, is ~3.0 V in the DME-based electrolyte used here. We discuss this issue and possible mechanisms for LiOH removal.

LiOH formation in the presence of LiI has been previously reported in lithium-oxygen batteries (3), and the reversible formation/decomposition of LiOH—e.g., using ruthenium catalysts and tetraglyme (4)/ dimethyl sulfoxide (DMSO) (5)–based electrolytes with added water—was observed at 3.1 to 3.2 V. It is well established that the electrolyte affects the potential of a redox couple. Indeed, even the I3/I couple drops from 3.53 V in water under standard conditions (6, 7) to 3.35, 3.1, and 3.0 V in acetonitrile (6), tetraglyme, and DME (1), respectively. The O2/O22– couple varies from ~3.0 V in DMSO electrolyte to ~3.5 V in acetonitrile electrolyte (8), and the recent “water in salt” work (9) showed that the redox potentials for water reduction/oxidation shift considerably due to the chemical potential changes of water and Li+ in the electrolyte.

Changes in the redox potential of reaction 1 under nonstandard conditions arise from at least two factors. First, the concentrations of the species in the electrolyte deviate noticeably from standard conditions. Second, the different coordination environments of the species/ions differ dramatically between solvents. The standard Gibbs free-energy change, ΔGro, of the overall cell reactionEmbedded Image(2)is 1282 kJ/mol (10) at standard conditions resulting in Eo = –ΔGro /nF = 3.32 V. Under nonstandard conditions, the equilibrium voltage of cell reaction 2 can be expressed as a function of the activity (α) of the reactants and products.Embedded Image (3)where R is the gas constant, F is the Faraday constant, n is the number of electrons involved in the reaction, and T is the temperature.

Because Li and LiOH are solids, and gaseous O2 is at close to ~1 bar during the cell reaction—i.e., they are in their respective standard states—the above equation can be simplified.

Embedded Image (4)

The activity, or the chemical potential, of water in the batteries is given byEmbedded Image (5)where μ is the chemical potential, and ΔH and ΔS are the enthalpy and entropy difference of water in LiI/LiTFSI/DME electrolyte compared with those in a LiI/LiTFSI aqueous electrolyte at standard conditionsEmbedded Image(6)and “H2Ol” in reaction 2 should be replaced by “H2O(Li,LiTFSI,DME)”. It is then relevant to ask how large a value of ΔG for reaction 6 is required to shift the potential of reaction 2 so that it drops below that of the I/I3 couple in the same electrolyte. Although this value is currently not known for DME, a relatively small value of –60.0 kJ/mol is required to reduce the couple of reaction 2 down to 3.0 V. ΔS makes only small changes to voltage (of –20 mV to –77 mV, where the entropy is approximated by the concentration of water in the system). The ΔH for reaction 6 is controlled by the loss of water-water interactions and the difference between water–Li salt and DME–Li salt interactions. Evidence for strong water–Li+ interactions in DME comes from the hygroscopic nature of LiTFSI and LiI salts and their higher solubility in water than in ethers (7, 9). Theoretical calculations (9) also suggest a very high binding energy of water with Li+ cations in an aqueous LiTFSI electrolyte. Hydration enthalpies for LiI and LiTFSI in water/nonaqueous solvents are an order of magnitude larger [e.g., –828 kJ/mol for LiI in water and –756 kJ/mol for LiTFSI in acetonitrile (11, 12)] than the –60 kJ/mol assumed above, all suggesting that this value for ΔGvi is certainly plausible, but we stress that further measurements/calculations are required.

The chemistry is, however, more complicated because I3 can both react with LiOH forming I and liberating oxygenEmbedded Image(7)and also react to form metastable IO, which then disproportionates forming IO3 and I (13, 14).

Embedded Image(8)

The low concentrations of water in our electrolytes help drive both reaction equilibria to the right-hand side. Our Raman and ultraviolet (UV)–visible studies of the chemical reaction of LiOH and I3 show that reaction 8 dominates under aqueous conditions, the rate of the reaction slowing down noticeably, but still occurring in DME solutions containing 3 to 6 weight % water (Fig. 1). Interestingly, the addition of LiTFSI speeds up the reaction. However, electrochemical studies of the oxidation of I3 in aqueous solutions with carbon electrodes often observe a combination of O2 evolution and IO3 formation (15). Thus, it is clear that further studies are required to determine how much O2 is evolved on charge versus being tied up as IO3. Our 17O/1H NMR studies of electrodes cycled in and beyond the first cycle observe reversible LiOH formation, even when using the thick electrodes (>0.150 mg) required for the NMR experiments. Finally, Viswanathan et al. suggested that parasitic reactions (3) between I/I3 and the ether-based electrolyte represented the source of proton for LiOH and that such irreversible processes cannot form the basis of a practical Li-O2 battery. We agree that irreversible processes must be avoided. In our work, DME electrolyte decomposition is not, however, the dominant source of the protons in LiOH, and super P carbon, rGO, and titanium carbide electrodes operate for hundreds of cycles (1). In our current Li-O2 battery, the equilibria among water, oxygen, and iodide; the surface functionality of rGO; and the detailed reaction mechanism during discharge-charge all require further investigation. Nonetheless, the evidence for reversible LiOH formation, in the presence of O2 (and LiI), is compelling.

Fig. 1 Reaction kinetics of a LiOH + LiI3 mixture in DME/water solutions followed by UV-visible spectroscopy (with LiOH in a 103 excess).

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