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The use of non-precious catalysts (e.g., transition metal oxides) in electrochemical energy technologies in acid (e.g., proton exchange membrane fuel cells and electrolyzers) has been significantly hampered by the instability of such catalysts at low pHs. While first-row late transition metal oxides based on Fe, Co, and/or Ni have been reported with comparable or higher catalytic activity for the alkaline oxygen reduction and oxygen evolution reactions than their precious-metal-based counterparts (e.g., Pt and IrO2),[1] these oxide catalysts are not stable in acid.[2] Moreover, Mn-based oxides are more acid-stable than those based on late transition metals[1] and can be potentially stabilized by reaching dissolution-deposition equilibrium at moderately acidic pHs,[3] but they still corrode and deactivate in strong acid and after long operation.[4] To tackle this challenge, extensive efforts over the past decades have been focused on combining active, yet unstable metal oxides with corrosion-resistant oxides.[5-7] Nevertheless, having a higher fraction of corrosion-resistant elements results in more acid-stable, yet less active catalysts, indicating an activity-stability trade-off.[2,5,6] To design transition metal oxide catalysts (such as Mn-based oxides) with an optimal trade-off or bypass this limitation, it is critical to establish their stability descriptors in acid. These descriptors can offer a fundamental understanding of oxide dissolution and provide guiding principles to enhance their intrinsic stability. In this work, we employed a library of Mn-based oxides with diverse structures and Mn oxidation states to identify oxide stability descriptors in acid based on intrinsic electronic structure and energetic parameters. Using time-dependent dissolution experiments and density functional theory calculations, greater amounts and faster kinetics of oxide dissolution in acid were correlated with decreased Mn oxidation states, which are accompanied by lowered Mn-O covalency, weakened Mn-O bonds, and reduced barriers for key reaction steps (such as protonation, vacancy formation, and metal ion solvation). Such design principles were shown broadly with a computational screening across a vast chemical space of ~1,000 Mn-based oxides, where an Mn oxidation state of 4+ was found to give rise to the lowest energetic driving force for the dissolution of Mn-based oxides in acid. Moreover, limiting the percentage of ionic metal substituents in oxides and using more acidic substituents were shown to stabilize Mn-based oxides in acid. These findings can provide novel insights into designing acid-stable Mn-based oxides for renewable energy storage and conversion. References: [1] C. Wei, R. R. Rao, J. Peng, B. Huang, I. E. L. Stephens, M. Risch, Z. J. Xu, Y. Shao-Horn, Adv. Mater. 2019, 31, 1806296. [2] K. K. Rao, Y. Lai, L. Zhou, J. A. Haber, M. Bajdich, J. M. Gregoire, Chem. Mater. 2022, 34, 899. [3] M. Huynh, D. K. Bediako, D. G. Nocera, J. Am. Chem. Soc. 2014, 136, 6002. [4] A. Li, H. Ooka, N. Bonnet, T. Hayashi, Y. Sun, Q. Jiang, C. Li, H. Han, R. Nakamura, Angew. Chem. Int. Ed. 2019, 58, 5054. [5] R. Frydendal, E. A. Paoli, I. Chorkendorff, J. Rossmeisl, I. E. L. Stephens, Adv. Energy Mater. 2015, 5, 1500991. [6] L. Zhou, A. Shinde, J. H. Montoya, A. Singh, S. Gul, J. Yano, Y. Ye, E. J. Crumlin, M. H. Richter, J. K. Cooper, H. S. Stein, J. A. Haber, K. A. Persson, J. M. Gregoire, ACS Catal. 2018, 8, 10938. [7] L. Zhou, H. Li, Y. Lai, M. Richter, K. Kan, J. A. Haber, S. Kelly, Z. Wang, Y. Lu, R. S. Kim, X. Li, J. Yano, J. K. Nørskov, J. M. Gregoire, ACS Energy Lett. 2022, 7, 993. Figure 1

Fundamental understanding of the reactivity between electrode and electrolyte is key to design the safety and life of Li-ion batteries. Herein X-ray photoelectron spectroscopy was used to examine the electrode/electrolyte interface (EEI) on carbon-free, binder-free LiCoO 2 powder and thin-film electrodes in LP57 electrolyte as function of potential. Upon charging of LiCoO 2 a marked growth of oxygenated and carbonated species was observed on the surface, consistent with electrolyte oxidation at high potentials. We also demonstrated that LiCoO 2 oxide surface was prone to decompose the salt starting at 4.1 V Li , as evidenced by the increase of LiF and Li x PF y O z species upon charging. By DFT calculations we proposed a correlation between the interface composition and the thermodynamic tendency of the EC solvent for dissociative adsorption on the Li x CoO 2 surface, through the generation of reactive acidic OH groups on the oxide surface, which can have a role in the observed salt decomposition. This is consistent with the evidence of HF and PF 2 O 2− species at 4.6 V Li observed by solution 19 F-NMR measurements. Finally we compared EEI composition between composite and model electrodes and discussed the changes and mechanisms induced by the electrode composition or the use of electrolyte additives. We showed that the addition of diphenyl carbonate (DPC) in the electrolyte has a strong impact on the formation of solvent and salt decomposition products at the EEI layer.

Late first-row transition metal oxides, based on cobalt, nickel and iron [1] are reported to be the most active non-precious catalysts for oxygen evolution reaction (OER) in basic solution. Experimental and computational studies in the past decade have been focusing on elucidating OER mechanisms [2-3] and identifying activity and stability descriptors [3-6], where perovskites family (ABO3-δ) with immense structural, chemical and electronic flexibility associated with vast selections of A-site and B-site metal ions and oxygen deficiency [7] has been used to develop design principles of OER activity and stability. Recent works [6-8] have shown that lowering charge-transfer gap or increasing metal-oxygen covalency in perovskites can enhance the OER kinetics, by reducing the energetic barriers associated with electron transfer on the surface of metal oxides including the most active catalysts. However, reducing the charge-transfer gap also lowers the Fermi level on the absolute energy scale, making it below the OER redox potential in the basic solution for the most active catalysts and rendering weaker surface hydroxide affinity [6]. Consequently, the OER kinetics on highly covalent/active metal oxides are limited by proton transfer. Moreover, increasing the covalency typically moves the O 2p band closer to the Fermi level, rendering less energy penalty for the creation of oxygen vacancies and thus leading to surface or bulk instability at OER potentials [5]. Here we explored the substitution of A-site ions with high electronegativity or Lewis acidity in the cobalt perovskites to maintain high Co-O covalency by the inductive effect [9], and tune the surface acid-base chemistry by introducing highly Lewis acidic A-site ions to facilitate OER kinetics. Bismuth-substituted strontium cobalt perovskite, Bi0.2Sr0.8CoO3-δ, was shown to exhibit record OER specific activity in alkaline solution, exceeding those of other Co-based perovskite oxides reported to date, including SrCoO3-δ [8], at high current densities (> 1 mA cm-2 oxide). In addition, neither structural or chemical changes have been found for Bi0.2Sr0.8CoO3-δ, indicative of greater structural stability than other highly covalent oxide catalysts, e.g. Ba0.5Sr0.5Co0.8Fe0.2O3-δ [4]. The enhanced OER kinetics and high surface stability can be attributed to the stronger affinity towards hydroxide ions to facilitate surface deprotonation due to the presence of strong Lewis acidic surface Bi3+ ions, and the lowered O 2p-band center relative to the Fermi level upon bismuth substitution into the perovskite structure, respectively. This work exemplifies a novel strategy to facilitate the OER kinetics of highly active oxide catalysts by leveraging the inductive effect associated with rational metal substitution to maintain high metal-oxygen covalency and strengthen hydroxide affinity without the expense of surface stability. References: [1] W. T. Hong, M. Risch, K. A. Stoerzinger, A. Grimaud, J. Suntivich, Y. Shao-Horn, Energy Environ. Sci. 2015, 8, 1404. [2] H. Dau, C. Limberg, T. Reier, M. Risch, S. Roggan, P. Strasser, ChemCatChem 2010, 2, 724. [3] I. C. Man, H. Y. Su, F. Calle-Vallejo, H. A. Hansen, J. I. Martínez, N. G. Inoglu, J. Kitchin, T. F. Jaramillo, J. K. Nørskov, J. Rossmeisl, ChemCatChem 2011, 3, 1159. [4] J. Suntivich, K. J. May, H. A. Gasteiger, J. B. Goodenough, Y. Shao-Horn, Science 2011, 334, 1383. [5] A. Grimaud, K. J. May, C. E. Carlton, Y.-L. Lee, M. Risch, W. T. Hong, J. Zhou, Y. Shao-Horn, Nat. Commun. 2013, 4, 2439. [6] W. Hong, K. A. Stoerzinger, Y.-L. Lee, L. Giordano, A. J. L. Grimaud, A. M. Johnson, J. Hwang, E. Crumlin, W. Yang, Y. Shao-Horn, Energy Environ. Sci. 2017, 10, 2190. [7] J. Hwang, R. R. Rao, L. Giordano, Y. Katayama, Y. Yu, Y. Shao-Horn, Science 2017, 358, 751. [8] A. Grimaud, O. Diaz-Morales, B. Han, W. T. Hong, Y. L. Lee, L. Giordano, K. A. Stoerzinger, M. T. M. Koper, Y. Shao-Horn, Nat. Chem. 2017, 9, 457. [9] D. A. Kuznetsov, B. Han, Y. Yu, R. R. Rao, J. Hwang, Y. Román-Leshkov, Y. Shao-Horn, Joule 2018, 2, 225.