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Reversible Mn2+/Mn4+ double redox in lithium-excess cathode materials

Abstract

There is an urgent need for low-cost, resource-friendly, high-energy-density cathode materials for lithium-ion batteries to satisfy the rapidly increasing need for electrical energy storage. To replace the nickel and cobalt, which are limited resources and are associated with safety problems, in current lithium-ion batteries, high-capacity cathodes based on manganese would be particularly desirable owing to the low cost and high abundance of the metal, and the intrinsic stability of the Mn4+ oxidation state. Here we present a strategy of combining high-valent cations and the partial substitution of fluorine for oxygen in a disordered-rocksalt structure to incorporate the reversible Mn2+/Mn4+ double redox couple into lithium-excess cathode materials. The lithium-rich cathodes thus produced have high capacity and energy density. The use of the Mn2+/Mn4+ redox reduces oxygen redox activity, thereby stabilizing the materials, and opens up new opportunities for the design of high-performance manganese-rich cathodes for advanced lithium-ion batteries.

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Fig. 1: Design and structural characterization of Li2Mn2/3Nb1/3O2F.
Fig. 2: Electrochemical performance of Li2Mn2/3Nb1/3O2F.
Fig. 3: Reaction mechanism of Li2Mn2/3Nb1/3O2F.
Fig. 4: Ab initio calculations of the redox mechanism of Li2Mn2/3Nb1/3O2F.
Fig. 5: Structural characterization and electrochemical performance of Li2Mn1/2Ti1/2O2F.

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Acknowledgements

Work by J.L., D.A.K., D.-H.K., Z.L., R.J.C. and G.C. was supported by Robert Bosch LLC, Umicore Specialty Oxides and Chemicals, and the Assistant Secretary for Energy Efficiency and Renewable Energy, Vehicle Technologies Office, of the U.S. Department of Energy under Contract No. DE-AC02-05CH11231, under the Advanced Battery Materials Research (BMR) Program. This research, in part, used resources of the Advanced Photon Source, a US Department of Energy (DOE) Office of Science User Facility operated for the DOE Office of Science by Argonne National Laboratory under contract no. DE-AC02-06CH11357. Work at the Advanced Light Source is supported by DOE Office of Science User Facility under contract no. DE-AC02-05CH11231. Work at the Molecular Foundry was supported by the Office of Science, Office of Basic Energy Sciences, of the US DOE under contract no. DE-AC02-05CH11231. The computational work relied on resources provided by the Extreme Science and Engineering Discovery Environment (XSEDE), which is supported by National Science Foundation grant no. ACI-1548562. J.K.P. acknowledges NSF Graduate Research Fellowship (grant no. DGE-1106400). B.D.M. acknowledges support from the Assistant Secretary for Energy Efficiency and Renewable Energy, Vehicle Technologies Office, of the US DOE under contract no. DEAC02-05CH11231, under the Advanced Battery Materials Research (BMR) Program. The authors thank S.-H. Hsieh for assistance in the soft XAS experiments and the California NanoSystems Institute (CNSI) at the University of California Santa Barbara (UCSB) for experimental time on the 500 MHz NMR spectrometer. The NMR experimental work reported here made use of the shared facilities of the UCSB MRSEC (NSF DMR 1720256), a member of the Material Research Facilities Network.

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J.L. and G.C. planned the project. G.C. supervised all aspects of the research. J.L. designed, synthesized, characterized (XRD) and electrochemically tested the proposed compounds. D.A.K. performed density functional theory calculations and analysed the data with J.L. D.-H.K. acquired and analysed TEM data. C.-W.L. and M.B. acquired and analysed hard XAS data. J.K.P. acquired and analysed DEMS data with input from B.D.M. Y.-S.L. and J.G. performed soft XAS measurements and analysed the data with J.L. Z.L. performed supportive electrochemical measurements. R.J.C. acquired and analysed the NMR data. T.S. performed SEM. The manuscript was written by J.L. and G.C. and was revised by D.A.K. and R.J.C. with the help of the other authors. All authors contributed to discussions.

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Correspondence to Jinhyuk Lee or Gerbrand Ceder.

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Extended data figures and tables

Extended Data Fig. 1 Solid-state NMR spectroscopy results.

a, b, 7Li spin echo NMR spectra acquired on as-synthesized Li2Mn2/3Nb1/3O2F (a) and Li2Mn1/2Ti1/2O2F (b) powders at 50 kHz MAS at a field B 0 = 11.7 T. The data have been fitted with a minimal number of Li sites: Li1, Li2 and Li3. Spinning sidebands of the three Li signals are indicated with asterisks. c19F spin echo sum spectra acquired on as-synthesized Li2Mn2/3Nb1/3O2F and Li2Mn1/2Ti1/2O2F powders at 50 kHz MAS at a field B 0 = 11.7 T. The spectra are compared to the spin echo spectrum collected on LiF under similar conditions. Spinning sidebands of the sharp LiF-like signals are indicated with asterisks. Detailed explanations of the results are given in Methods section ‘Supplementary Note 1’.

Extended Data Fig. 2 Structural characterization of Li2Mn2/3Nb1/3O2F.

a, TEM image of as-synthesized Li2Mn2/3Nb1/3O2F particles. Scale bar, 50 nm. b, A high-magnification TEM image of the area enclosed in a square in a. Scale bar, 10 nm. The yellow circle indicates the boundary of one of the many grains in the polycrystalline Li2Mn2/3Nb1/3O2F particle. c, An electron diffraction pattern of the Li2Mn2/3Nb1/3O2F particle. Scale bar, 5 nm−1. d, Fast Fourier-transformed (FFT) images of the dotted squared areas in b. e, The high magnification image across the squared areas 1, 2 and 3 in b. Scale bar, 5 nm. We can clearly observe lattice fringes and FFT peaks throughout the particle, indicating that our particles are made of small crystalline grains instead of amorphous phases.

Extended Data Fig. 3 Additional electrochemical data from Li2Mn2/3Nb1/3O2F.

a, b, Voltage profiles of the 60:30:10 electrode (that is, 60 wt% Li2Mn2/3Nb1/3O2F: 30 wt% carbon black: 10 wt% PTFE) when cycled between 2.0 V and 4.8 V (a), and 2.3 V and 4.6 V (b) at 20 mA g−1. c, d, Voltage profiles of the 70:20:10 (c) and the 80:15:5 (d) electrodes, when cycled between 1.5 V and 5.0 V at 20 mA g−1. e, Voltage profiles of the 80:15:5 electrode when cycled between 2.0 V and 4.8 V at 20 mA g−1. f, A comparison of the first discharge profiles of the 60:30:10, 70:20:10 and 80:15:5 Li2Mn2/3Nb1/3O2F electrodes (1.5–5.0 V, 20 mA g−1). The specific capacity was calculated on the amount of the Li2Mn2/3Nb1/3O2F powder in the cathode film.

Extended Data Fig. 4 Discharge capacity retention.

The 60:30:10 Li2Mn2/3Nb1/3O2F: carbon black:PTFE electrode was cycled between 1.5 V and 5.0 V at room temperature at 10, 20, 40, 100, 200, 400 and 1,000 mA g−1. A faster rate leads to less capacity fading during the initial 25 cycles. This is likely to be due to electrolyte decomposition per cycle occurring more (less) at a high voltage in a slower (faster) cycling test, which increases the impedance of a cell by creating a resistive surface layer and decreasing the ionic conductivity of the electrolyte.

Extended Data Fig. 5 Gas evolution measurements.

a, b, Initial voltage profiles (black solid line) of Li2Mn2/3Nb1/3O2F (a) and Li2Mn1/2Ti1/2O2F (b), when charged to 5.0 V at a rate of 20 mA g−1. DEMS results for O2 (red circles) and CO2 (blue triangles) evolution are also shown. c, Cumulative CO2 evolution from shaker-mixed Li2Mn1/2Ti1/2O2F and carbon black powder mixture, as a function of time during an acid titration test using 1 M H2SO4. Detailed explanations of the results are given in Methods section ‘Supplementary Note 2’. 1st c, first charge.

Extended Data Fig. 6 Evolution of the charge and discharge voltages.

Average charge voltage (triangles), discharge voltage (stars), and half of the charge–discharge voltage (circles) are shown when Li2Mn2/3Nb1/3O2F is cycled between 1.5 V and 4.6 V, 1.5 V and 4.8 V, and 1.5 V and 5.0 V, at 20 mA g−1. Detailed explanations of the results are given in Methods section ‘Supplementary Note 3’. c, charge; dc, discharge.

Extended Data Fig. 7 XANES of Li2Mn2/3Nb1/3O2F.

a, b, Manganese K-edge XANES spectra of Li2Mn2/3Nb1/3O2F: before cycle, after first charging to 135 mAh g−1, 270 mAh g−1 and 360 mAh g−1, and after first charging to 375 mAh g−1 then discharging to 320 mAh g−1. ce, First derivatives of normalized absorbance at the pre-edge region of Mn K-edge spectra of Li2Mn2/3Nb1/3O2F: c, before cycle and after first charging to 375 mAh g−1 then discharging to 320 mAh g−1; d, after first charging to 135 mAh g−1; and e, to 270 mAh g−1 and 360 mAh g−1. Data from MnO, Mn2O3 and MnO2 are presented for comparison. Detailed explanations of the results are given in Methods section ‘Supplementary Note 4’.

Extended Data Fig. 8 Niobium K-edge XANES spectra of Li2Mn2/3Nb1/3O2F obtained by hard XAS.

Results are shown before cycle, after charging to 135 mAh g−1, 270 mAh g−1 and 360 mAh g−1, and after charging to 375 mAh g−1 then discharging to 320 mAh g−1. The Nb K-edge XANES spectra of the Li2Mn2/3Nb1/3O2F samples are similar to that of Nb2O5 (Nb5+ reference), indicating that Nb in the compound stays as Nb5+ during cycling. The observable small shape changes are likely to be related to changes in local disorder and distortion60.

Extended Data Fig. 9 Structural characterization of Li2Mn1/2Ti1/2O2F.

a, TEM image of as-synthesized Li2Mn1/2Ti1/2O2F particles. Scale bar, 50 nm. b, A high-magnification TEM image of the area enclosed in a square in a. Scale bar, 10 nm. The yellow circle indicates the boundary of one of the many grains in the polycrystalline Li2Mn1/2Ti1/2O2F particle. c, An electron diffraction pattern of the Li2Mn1/2Ti1/2O2F particle. Scale bar, 5 nm−1. d, FFT images of the dotted squared areas in b. e, The high magnification image across the squared areas 1, 2 and 3 in b. Scale bar, 5 nm. We can clearly observe lattice fringes and FFT peaks throughout the particle, indicating that our particles are made of small crystalline grains instead of amorphous phases.

Extended Data Fig. 10 Electrochemical properties of Li2Mn1/2Ti1/2O2F.

ac, Voltage profiles and capacity retention of the 60:30:10 Li2Mn1/2Ti1/2O2F:carbon black:PTFE electrode when cycled at 20 mA g−1 at room temperature between 1.6 V and 5.0 V (a), 2.0 V and 4.8 V (b), and 2.3 V and 4.6 V (c). d, The initial charge–discharge profile of the 60:30:10 electrode when cycled between 1.6 V and 5.0 V at room temperature at 20, 40, 100, 200, 400 and 1,000 mA g−1. e,The discharge capacities during initial 25 cycles. f, Voltage profiles and capacity retention of the 80:15:5 electrode when cycled at 20 mA g−1 at room temperature between 2.0 V and 4.8 V. The specific capacity was calculated on the amount of the Li2Mn1/2Ti1/2O2F powder in the cathode film. Detailed explanations of the results are given in Methods section ‘Supplementary Note 5’.

Extended Data Fig. 11 XANES of Li2Mn1/2Ti1/2O2F.

a, b, Manganese K-edge XANES spectra of Li2Mn1/2Ti1/2O2F: before cycle (black), 120 mAh g−1 charged (navy), 240 mAh g−1 charged (wine), 400 mAh g−1 charged (grey), 330 mAh g−1 discharged after a 400 mAh g−1 charge (dark yellow). ce, First derivatives of normalized absorbance at the pre-edge region of Mn K-edge spectra of Li2Mn1/2Ti1/2O2F: c, before cycle and after first charging to 400 mAh g−1 then discharging to 330 mAh g−1; d, after first charging to 120 mAh g−1; and e, to 240 mAh g−1 and 400 mAh g−1. f, Titanium K-edge XANES spectra of Li2Mn1/2Ti1/2O2F during the initial cycle. Data from MnO, Mn2O3, MnO2, Ti2O3 and TiO2 are presented for comparison. Detailed explanations of the results are given in Methods section ‘Supplementary Note 6’.

Extended Data Table 1 Structural parameters from the Rietveld refinements
Extended Data Table 2 Target versus measured atomic ratio of Li2Mn2/3Nb1/3O2F and Li2Mn1/2Ti1/2O2F compounds

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Lee, J., Kitchaev, D.A., Kwon, DH. et al. Reversible Mn2+/Mn4+ double redox in lithium-excess cathode materials. Nature 556, 185–190 (2018). https://doi.org/10.1038/s41586-018-0015-4

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