Technical Comments

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.aad8689


Based on a simple thermodynamic analysis, we show that iodide-mediated electrochemical decomposition of lithium hydroxide (LiOH) likely occurs through a different mechanism than that proposed by Liu et al. (Research Article, 30 October 2015, p. 530). The mismatch in thermodynamic potentials for iodide/triiodide (I/I3) redox and O2 evolution from LiOH implies a different active iodine/oxygen electrochemistry on battery charge. It is therefore possible that the system described in Liu et al. may not form the basis for a rechargeable lithium-oxygen (Li-O2) battery.

Liu et al. (1) proposed that a highly rechargeable and electrically efficient Li-O2 battery could be produced by using LiI and small amounts of H2O in a nonaqueous electrolyte and a reduced graphene oxide (rGO) electrode. They suggested that LiI catalyzed oxygen reduction to LiOH when H2O impurities were present in the otherwise nominally nonaqueous electrolyte. They conjectured that I3 formed during charge at 3 V versus Li/Li+ (all potentials herein are referenced to Li/Li+) mediates the oxidation of LiOH to liberate O2 without substantial parasitic reactions.

Although no thermodynamic barriers exist for the LiI-mediated formation of LiOH during discharge of water-contaminated Li-O2 cells, the charging process warrants further analysis. We present a simple thermodynamic analysis of the triiodide-mediated LiOH oxidation mechanism proposed in their work (1). We consider their overall proposed charging reaction.Embedded Image(1)This reaction has an equilibrium potential of 3.34 V under standard conditions (2). The activity of water is not at standard conditions in these experiments, but even at parts per million (ppm) quantities of water, there are negligible changes to the equilibrium potential. For example, in an experiment with ~100 to 45,000 ppm, this corresponds to an equilibrium potential shift of ~0.02 to 0.009 V.

We now consider the iodide/triiodide redox reaction, shown in figure 4 of (1), given by

Embedded Image(2)

The redox potential for this reaction depends on the solvent (3). As reported in (1), we use the equilibrium potential—i.e., ~3.0 V—as shown in figure 1 of (1). Liu et al. point out in the caption that the crossing points of the charge/discharge curves indicate the positions of the redox potential of I/I3 in the specific electrode/electrolyte system.

The chemical reaction, shown in figure 4 of (1), to complete the charging cycle, is given by Embedded Image(3)This chemical reaction is uphill in free energy by ΔGc2 = 4 × (3.347 − 3.00) = 1.39 eV. This corresponds to an equilibrium constant, given by K = exp[−(ΔGc2/kBT)] ~ 10−24. Even if every cathode site were electrochemically active, the fraction that would transform to the product, O2, would be exceptionally small and therefore cannot account for a reversible process regenerating O2.

The overall free-energy diagram for the charging mechanism proposed by Liu et al. at U = 3 V is shown in Fig. 1. A simpler way to understand the inconsistency is that the minimum potential required to oxidize LiOH is 3.34 V. The iodide/triiodide redox couple has an oxidation potential of only ~3 V and hence it is not feasible to oxidize LiOH(s) with such a scheme.

Fig. 1 Free-energy diagram for the proposed charging mechanism involving iodide/triiodide–mediated LiOH oxidation at U = 3 V.

The chemical step associated with the oxidation of LiOH with triiodide is thermodynamically uphill by 1.39 eV.

In a rigorously anhydrous Li-O2 battery that produces Li2O2 as the discharge product (4, 5), the free-energy diagram involving an iodide/triiodide redox couple is shown in Fig. 2. In this case, the minimum potential required to oxidize Li2O2(s) is 2.96 V, and the iodide/triiodide redox couple possesses the required driving force to carry out this reaction, as shown in several experiments (6).

Fig. 2 Free-energy diagram for the iodide/triiodide–mediated Li2O2 oxidation in an anhydrous Li-O2 battery at U = 3 V.

The chemical step associated with the oxidation of Li2O2 with triiodide is thermodynamically downhill by 0.08 eV.

To summarize, our thermodynamic analysis shows that it is not possible to run an iodide-mediated reversible electrochemical cycle, as indicated in figure 4 of (1). However, given the clear evidence of LiOH disappearance during the charge process, some other electrochemical reaction or reactions involving LiOH are active. Whether these other reactions can form the basis for a rechargeable battery remains an open question. Previous reports have shown empirical evidence of parasitic reactions associated with iodide/triiodide redox couples in a Li-O2 battery with ethereal solvents (6, 7). Such irreversible routes obviously need to be avoided to allow long-term rechargeability in a Li-O2 battery.


Acknowledgments: Authors acknowledge helpful discussions with G. Chase.

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