to cheerful mood. Thanks to the other lab members also. With such an international and collaborative working environment, science and research becomes very smooth and enjoyable. I would also like to thank Prof Silke Biermann and Dr Valérie Pralong for their time to review this thesis, and also to Prof Christel Laberty-Robert, Prof Philippe Barboux, and Prof Artem Abakumov for being part of the jury.

Finally, on more personal note, I would like to thank Frédéric Lavaud, Sumit Sarkar, and all my friends here, and also to Maison de L'Inde. Thanks also to my old teachers, supervisors and friends. Finally and most importantly, I would like to dedicate this work and effort to my Family for everything they have done for me. Without their silent love, encouragement and support, I could not go forward with science and research.

The global energy demand is rising every year and projected to increase by ~56% between 2010 and 2040 (Figure 1a). 1 To meet this energy demand, consumption of all fuels is accelerated, with fossil fuels meeting ~70% of the growth. As a result, global energy-related CO2 emissions is soaring at an alarming rate (by 1.7% in 2018). In efforts to limit the climate change, 'clean' electricity is being projected as the "fuel" of the future, and utilization of renewable energies is growing fast to meet the increase in demand for electricity (Figure 1b). Energy storage systems (ESS) are the basis for this future transformation, as they ensure energy availability when the wind is not blowing and the sun is not shining. Various ESS, such as, batteries, flywheel, pumped water storage, compressed air energy storage, etc., have been developed in the past. 2 Amongst them, batteries, specifically rechargeable Li-ion batteries (LIBs) outperform other systems, thanks to its high specific energy (250 -300 Wh•kg -1 ), while also providing a high specific power. 3 Historically, portable electronics have been the primary driving force for the LIB market, however, today the demand comes from electric vehicles (EVs) also. Furthermore, to enable large-scale energy storage for grids, etc., not only increasing the energy density, but also the safety (more important for EVs) of the LIBs, have become important. This thesis aims to contribute towards the cathode materials of LIBs, and all-solid-state Li-batteries which will be outlined below.

○ Thesis Outline

This work is divided into six Chapters, they consist, which are detailed below:

Chapter I briefly introduces the current Li-ion battery (LIB) and future 'all-solid-state battery' (ASSB) technology. The new phenomena 'anionic redox' will be explained and its associated 'limitations' will be identified. It is followed by a discussion on requirements for ASSBs, which highlights the importance to act on solid electrolytes (SEs) and the 'interface' of SE-cathode material in ASSBs.

Chapter II focuses on designing new oxide-based SEs, more specifically by stabilization of disordered high-temperature phases (HTPs) at room temperature (RT) within the Li-Zn-P-O phase diagram. The phases Li6Zn(P2O7)2, Li4Zn(PO4)2 and Na4Zn(PO4)2 were chosen for this study. Their structural transition accompanied by the increase of ionic conductivity has been characterized. Then the stabilization of their HTPs by means of chemical substitution and mechanochemical treatment (ballmilling) has been shown.

Chapter III deals with designing sulfide-based cathode materials that involves anionic sulfurredox to increase energy density of the LIBs. First we revisited TiS3 for its structural aspects and practicability as an anionic redox compound. Next, we designed a new family of Li-rich layered sulfides Li1.33-2y/3Ti 4+ 0.67-y/3Fe 2+ yS2 (i.e., Fe 2+ substituted Li2Ti 4+ S3) and unveiled the benefits of sulfur redox. Furthermore, we extended our studies to other Li/Na-rich sulfides.

Chapter IV presents practical assessment of anionic redox in Li-rich sulfides. We systematically benchmarked the performance of the Li-rich sulfides against Li-rich NMC oxide via detailed electrochemical analysis.

Chapter V shows application of the Li-rich sulfide Li1.13Ti0.57Fe0.3S2 (y = 0.3 composition) as cathode material in ASSB.

To conclude, Chapter VI summarizes the findings of this thesis together with its impact with respect to application.

Furthermore, an Appendix section can be found at the end, that includes the experimental methods and additional figures and tables.

Chapter I Brief introduction to the battery technologies I.1 Brief overview of Li-battery technology

Development of portable devices and electronics has been accelerated in last few decades. In addition, growing need for green mobility has led to emergence of electric vehicles (EVs). EVs, in combination with renewable energies, are important to tackle the global mobility needs. Batteries are a key enabler to limit the CO2-emission, and recent market developments of batteries strongly reflect this. The competitive market of electronics and EVs present a strong motivation to work on the development of 'better' batteries with increased energy and power density, with more safety, at a lower cost. Among a wide range of energy storage systems, rechargeable lithium-ion batteries (LIBs), are currently the dominant technology for powering portable devices, mostly due to their capability of delivering electrical energy with high power and energy densities. [START_REF] Larcher | Towards greener and more sustainable batteries for electrical energy storage[END_REF]5 A battery is a system consisting in a positive and a negative electrode, separated by an ionically conducting electrolyte. The external circuit allows for the flow of electrons, while the internal electrolyte solution allows for the flow of ions. Simultaneous electrochemical reactions at the two electrodes lead to energy storage or release (charge or discharge), depending on the direction of flow of electrons. In the next section we will introduce the Li-battery technology followed by its transition to 'next-generation' all-solid-state batteries.

I.1.1 Early Li-metal batteries with sulfide cathodes

Sulfide based electrodes played a crucial role to pave the way for today's Li-battery technology. The oil crisis in early 1970s had prompted serious investigations on renewable solar, wind, and domestic fossil fuels as alternative power source. [START_REF] Painter | Oil and Geopolitics: The Oil Crises of the 1970s and the Cold War[END_REF] However, the intermittency of these power sources, together with the emergence of portable electronic products, sparked worldwide research for energy storage system, especially for rechargeable batteries. [START_REF] Kurzweil | Lithium Battery Energy Storage[END_REF][START_REF] Goodenough | How we made the Li-ion rechargeable battery[END_REF] In 1975, transition-metal (TM) chalcogenide TiS2 was first reported by Whittingham, as Li insertion host (Figure I.1a) for Li-metal batteries. [START_REF] Whittingham | Electrical Energy Storage and Intercalation Chemistry[END_REF] The Li/TiS2 system shows excellent reversibility of Li-intercalation for LixTiS2 with 0 < x < 1, according to the reaction scheme, (at cathode) TiS2 + Li + + e -⇆ LiTiS2 (at anode) Li ⇆ Li + + e - Furthermore, other Li-free 3d/4d TM-chalcogenides, such as MoS2 VSe2, NbSe3 etc., were identified for reversible Li-intercalation. [START_REF] Whittingham | Electrical Energy Storage and Intercalation Chemistry[END_REF][START_REF] Whittingham | Chemistry of intercalation compounds: Metal guests in chalcogenide hosts[END_REF][START_REF] Whittingham | The Role of Ternary Phases in Cathode Reactions[END_REF] Among such Li-metal batteries, Li/TiS2 (at EXXON) and Li/MoS2 (at Moli Energy) system was initially considered very promising. [START_REF] Goodenough | How we made the Li-ion rechargeable battery[END_REF][START_REF] Whittingham | Lithium batteries and cathode materials[END_REF] However, due to the safety risks originated from the dendritic lithium growth at the metallic-Li anode, the EXXON's batteries could not be commercialized and Moli's batteries had to be recalled from the market after fire accidents. [START_REF] Whittingham | Lithium batteries and cathode materials[END_REF][START_REF] Blomgren | The Development and Future of Lithium Ion Batteries[END_REF]

I.1.2 Li-ion batteries with oxide cathodes for higher voltage

Even though early Li-metal batteries were rapidly terminated, research for safer and viable Libattery system was intense. In 1979, Goodenough proposed LiCoO2 as a Li-insertion host which demonstrated successive Li deintercalation-intercalation for Li1-xCoO2 with 0 < x < 0.5. 14,[START_REF] Mizushima | LixCoO2 (0<x⩽1): A new cathode material for batteries of high energy density[END_REF] Furthermore, as LiCoO2 contains Li, the metallic Li metal anode was not needed anymore and use of graphite was proposed as Li-insertion anode. [START_REF] Yazami | A reversible graphite-lithium negative electrode for electrochemical generators[END_REF] This led Akira Yoshino (at the Asahi Kasei Corporation) to use carbonaceous materials (with a certain crystalline structure that does not decompose in organic propylene carbonate (PC) electrolyte, e.g., vapor phase grown carbon fiber [START_REF] Oberlin | Filamentous growth of carbon through benzene decomposition[END_REF] ) as Li-insertion anode and demonstrated the first Li-ion rechargeable battery in 1986 (see Figure I.1b), according to the reaction scheme shown below. [START_REF] Yoshino | Secondary battery[END_REF][START_REF] Yoshino | Development of the Lithium-Ion Battery and Recent Technological Trends[END_REF][START_REF] Yoshino | The Birth of the Lithium-Ion Battery[END_REF] This configuration of LIB was commercialized by Sony in 1991. [START_REF] Goodenough | How we made the Li-ion rechargeable battery[END_REF][START_REF] Yoshino | Development of the Lithium-Ion Battery and Recent Technological Trends[END_REF][START_REF] Yoshino | The Birth of the Lithium-Ion Battery[END_REF] (at cathode) LiCoO2 ⇆ xLi + + Li1-xCoO2 + xe - (at anode) Cn + xLi + + xe -⇆ CnLix Followed by the successful demonstration by layered LiCoO2, electrochemical properties were discovered in other family of compounds, such as, 'spinel' LiMn2O4 and 'polyanionic' LiFePO4, etc., and also in other layered oxides. [START_REF] Larcher | Towards greener and more sustainable batteries for electrical energy storage[END_REF][START_REF] Manthiram | An Outlook on Lithium Ion Battery Technology[END_REF] Amongst the compounds, majority of today's batteries still use LiCoO2, and its derivatives as positive electrode. Both TiS2 and LiCoO2 has similar layered crystal structure (Figure I.1) and Li-(de)intercalation in them is charge compensated by cationic redox, i.e., the redox of the transition metals (TM n+/n+1 ). The voltage profiles of the sulfide and oxide electrodes are illustrated by their schematic band positions in Figure I.2. In sulfides the cationic redox level (Ti 3+/4+ ) has a higher energy (i.e., closer to the Li/Li + redox level) than the cationic redox level (Co 2+/3+ ) in oxides. [START_REF] Manthiram | An Outlook on Lithium Ion Battery Technology[END_REF]22 Consequently, low voltage can only be obtained from TiS2 sulfide. The voltage difference is reflected into their specific energies, as TiS2 offers much lower (~450 Wh•kg -1 ) than the commercialized LiCoO2 (~600 Wh•kg -1 ). [START_REF] Luo | Charge-compensation in 3d-transition-metal-oxide intercalation cathodes through the generation of localized electron holes on oxygen[END_REF][START_REF] Goodenough | Locating redox couples in the layered sulfides with application to Cu[Cr2]S4[END_REF] . The voltage profiles (1st cycle) of TiS2, LiCoO2 (LCO), and LiNi0.33Mn0.33Co0.33O2 (NMC) are shown in c-d. The purple circles show the starting points. TiS2 was discharged first (starting with x = 0) and LCO and NMC were charged first (starting with x = 1).

Despite low operating potential, the increased covalency in sulfides than in oxides plays a crucial role. In LiCoO2, only ~0. 5 Li can be reversibly (de)intercalated (i.e., capacity up to ~140 mAh•g -1 ), as more Li extraction leads to gradual collapse of the O3 lattice structure. [START_REF] Yano | LiCoO 2 Degradation Behavior in the High-Voltage Phase Transition Region and Improved Reversibility with Surface Coating[END_REF][START_REF] Zhang | Trace doping of multiple elements enables stable battery cycling of LiCoO2 at 4.6 V[END_REF] LiNiO2 with similar crystal structure could deliver a capacity >150 mAh•g -1 , however its thermal instability posed a major safety drawback. [START_REF] Rozier | Review-Li-Rich Layered Oxide Cathodes for Next-Generation Li-Ion Batteries: Chances and Challenges[END_REF] Contrarily, in TiS2, fully ~1 Li can be intercalated reversibly (i.e., capacity up to ~220 mAh•g -1 ), because of the robustness of the crystal structure due to greater covalency and bigger size of the S atoms.

The limited capacity could be pushed further by designing compounds such as Li[Ni1-y-zMnyCoz]O2 (NMC), and Li[Ni1-y-zCoyAlz]O2 (NCA) etc. This was made possible by partially replacing Co with Ni/Mn or by Ni/Al. The crystal structures of the LCO and NMC are compared in Figure I.3, which show that the Co is partially mixed with Ni/Mn in the metallic layer. "NMC 111" (LiNi1/3Mn1/3Co1/3O2, with equal parts of TM in metal layer) was reported in 2001, which showed reversible capacity of >150 mAh•g -1 . [START_REF] Ohzuku | Layered Lithium Insertion Material of LiNi 1/2 Mn 1/2 O 2 : A Possible Alternative to LiCoO 2 for Advanced Lithium-Ion Batteries[END_REF] In subsequent years the capacity in NMC could be increased by increasing the relative amount of Ni, e.g. NMC 622 (LiNi0.6Mn0.2Co0.2O2, ~190 mAh•g -1 ), that is mostly being used in most commercial LIBs. [START_REF] Lu | Layered Cathode Materials Li[END_REF] Thus capacity greater than 200 mAh•g -1 could be obtained for NMC 811 (LiNi0.8Mn0.1Co0.1O2) which is being considered to be the most promising for EV's in near future. [START_REF] Noh | Comparison of the structural and electrochemical properties of layered Li[END_REF] Similarly, in NCA compositions, up to ~200 mAh•g -1 could be obtained for LiNi0.8Co0.15Al0.05O2. [START_REF] Robert | Enhancement of the high potential specific charge in layered electrode materials for lithium-ion batteries[END_REF] Figure I.3. Improving the specific capacity of layered oxides cathode materials. LiCoO2 is currently being replaced in Li-ion batteries by the "NMC" materials in which Co is partially substituted by Ni and Mn. Increasing the amount of Li leads to Li-rich NMC (LrNMC) phases with Li in the metallic layers and capacities reaching ~270 mAh/g. Adapted from Rozier et al. [START_REF] Rozier | Review-Li-Rich Layered Oxide Cathodes for Next-Generation Li-Ion Batteries: Chances and Challenges[END_REF] Next legitimate question arises, is this the limit of capacity in cathode materials ? Indeed, the capacity could be pushed by putting additional Li in the metal layer (Figure I.3, right panel). This leads to a new paradigm "anionic redox" (Figure I.4) which will be briefly described next.

I.1.3 Anionic redox to push capacity limits

To comply with the ever-growing demands of energy storage for portable applications, increasing the energy density of LIBs has become a formidable challenge. [START_REF] Blomgren | The Development and Future of Lithium Ion Batteries[END_REF][START_REF] Schmuch | Performance and cost of materials for lithium-based rechargeable automotive batteries[END_REF] For many years, the cathode materials in LIBs relied solely on the transition metal (cationic) redox (Figure I.4a,c), until the recently discovered anionic redox, i.e., electrochemical participation of the oxygen ligands, became a new approach for designing higher energy cathode materials (Figure I.4b,d). Li-rich layered oxides, for example, Li1.2Ni0.2Mn0.6O2 and Li1.2Ni0.13Mn0.54Co0.13O2 (Li-rich NMC) hold the highest promises in this regard, as these cathodes can deliver a high specific energy approaching ~1000 Wh•kg -1 (at the materiallevel). These materials can potentially replace the currently used NMCs (e.g. LiNi0.6Mn0.2Co0.2O2, ~700 Wh•kg -1 ).

Looking back, Li-rich phase Li[Li0.33-2y/3Mn0.67-y/3Niy]O2 (y = 0.5 -0.17), with capacity >230 mAh•g -1 , was reported first in 2002 by Dahn and co-workers. [START_REF] Lu | Understanding the Anomalous Capacity of Li/Li[END_REF]34 Further breakthrough came in 2006 when the same group and Thackeray et al reported Li-rich phases with further Co substitution of Mn leading to unprecedented capacity of >270 mAh•g -1 for Li1.2Ni0.13Mn0.54Co0.13O2 , which is known as todays Li-rich NMC (Figure I.4b). [START_REF] Jiang | Insignificant impact of designed oxygen release from high capacity Li[(Ni1/2Mn1/2)xCoy(Li1/3Mn2/3)1/3]O2 (x+y=2/3) positive electrodes during the cycling of Li-ion cells[END_REF][START_REF] Thackeray | Li2MnO3-stabilized LiMO2 (M = Mn, Ni, Co) electrodes for lithiumion batteries[END_REF] Note that in this Li-rich phases Mn 4+ remains redox-inactive throughout charge and discharge, while the Ni 2+/3+/4+ and Co 2+/3+ redox can only account for part of the capacity (denoted by blue dotted circle in Figure I.4b,d). Thanks to a decade of intense research, it is now well established that the anomalous extra capacity of Li-rich cathodes arises from the redox of O 2-anions (denoted by red dotted circle in Figure I.4b,d), more specifically the 'non-bonding ' O 2p orbitals that point towards excess-Li in the metal layers. [START_REF] Luo | Charge-compensation in 3d-transition-metal-oxide intercalation cathodes through the generation of localized electron holes on oxygen[END_REF][START_REF] Koga | Reversible Oxygen Participation to the Redox Processes Revealed for Li 1[END_REF][START_REF] Yabuuchi | Solid-state Redox Reaction of Oxide Ions for Rechargeable Batteries[END_REF][START_REF] Seo | The structural and chemical origin of the oxygen redox activity in layered and cation-disordered Li-excess cathode materials[END_REF][START_REF] Saubanère | The intriguing question of anionic redox in high-energy density cathodes for Li-ion batteries[END_REF][START_REF] Okubo | Molecular Orbital Principles of Oxygen-Redox Battery Electrodes[END_REF] This new redox has been termed as 'anionic redox' which has emerged as a new paradigm to design higher energy Li-rich cathode materials. Invigorated by this fundamental understanding, solidstate chemists have extended the concept of oxygen redox to cation-disordered Li1+yM1-yO2 (M = Nb, Mn, Ti, Fe, V …) as well as Na-based layered oxides. [START_REF] Yabuuchi | Origin of stabilization and destabilization in solid-state redox reaction of oxide ions for lithium-ion batteries[END_REF][START_REF] Maitra | Oxygen redox chemistry without excess alkali-metal ions in Na 2/3 [Mg 0.28 Mn 0.72 ]O 2[END_REF][START_REF] Mortemard De Boisse | Highly Reversible Oxygen-Redox Chemistry at 4.1 V in Na 4/7x[END_REF] Despite this rich materials design-space, certain practical issues, such as voltage fade, poor kinetics, voltage hysteresis, and irreversible O2 loss, have delayed the commercialization of Li-rich NMCs. [START_REF] Assat | Fundamental understanding and practical challenges of anionic redox activity in Li-ion batteries[END_REF][START_REF] Li | High-voltage positive electrode materials for lithium-ion batteries[END_REF][START_REF] Li | Anionic Redox in Rechargeable Lithium Batteries[END_REF] As shown in the voltage profile of Li-rich NMC in Figure I.5a, the staircase charging profile converts to S-shaped discharge profile in the 1st cycle (known as 'activation cycle') with a large irreversibility of >0. 25 Li, along with huge voltage hysteresis. Furthermore, the voltage fades continuly in long cycling. In 2013, Sathiya et al. reported the electrochemical behavior of the Li2Ru1-ySnyO3 solid solution that possesses a similar voltage profile (Figure I.5b). [START_REF] Sathiya | Reversible anionic redox chemistry in high-capacity layered-oxide electrodes[END_REF] But, due to increased covalency of Ru-O bonds, a reversible capacity >200 mAh•g -1 could be obtained with good capacity retention and very little voltage decay. Although the crucial role of oxygen redox towards these issues was clearly highlighted by detailed investigations on a 'practical ' Li-rich NMC and also on a 'model' Li1.33Ru0.5Sn0.17O2 electrode, much remains to be understood for further fundamental insights that will ultimately lead to implementable solutions. [START_REF] Assat | Fundamental understanding and practical challenges of anionic redox activity in Li-ion batteries[END_REF][START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF][START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF] Moving further down to 5d Ir, the oxygen redox could be stabilized in α-Li2IrO3. [START_REF] Mccalla | Visualization of O-O peroxo-like dimers in high-capacity layered oxides for Li-ion batteries[END_REF] Pearce et al, reported the electrochemical properties of β-Li2IrO3 which has a layered 3D structure and it showed negligible voltage hysteresis due to the increased covalency of Ir-O bonds and robustness of the crystal structure. [START_REF] Pearce | Evidence for anionic redox activity in a tridimensional-ordered Li-rich positive electrode β-Li2IrO3[END_REF][START_REF] Pearce | Anionic and Cationic Redox Processes in β-Li2IrO3 and Their Structural Implications on Electrochemical Cycling in Li-Ion Cell[END_REF] Another approach was increasing the Li-rich character to access higher capacity in Li3IrO4. [START_REF] Perez | Approaching the limits of cationic and anionic electrochemical activity with the Li-rich layered rocksalt Li3IrO4[END_REF] More recently, a third approach with mild success has emerged, that is to tune the ligand by increasing electronegativity of the anionic sublattice via substituting fluorine for oxygen, but so far limited only to disordered rocksalt structures. [START_REF] House | Lithium manganese oxyfluoride as a new cathode material exhibiting oxygen redox[END_REF][START_REF] Lee | Reversible Mn2+/Mn4+ double redox in lithium-excess cathode materials[END_REF] Figure I.5. Voltage profiles of some Li-rich oxides. Figures are reproduced from references [START_REF] Luo | Charge-compensation in 3d-transition-metal-oxide intercalation cathodes through the generation of localized electron holes on oxygen[END_REF][START_REF] Sathiya | Reversible anionic redox chemistry in high-capacity layered-oxide electrodes[END_REF] . Now the question arises, that under what circumstances the 'non-bonding ' O 2p orbitals take part in 'anionic' redox ? Over the past five years, experimental and theoretical studies have correlated the origin of anionic redox as arising from oxygen lone-pair states to the compound's electronic structure. In classical layered oxides (LiMO2), the redox is limited to the oxidation/reduction of antibonding M-O* levels (often called as M d band for simplicity sake), which remains mostly cationic by nature (Figure I.6a). In compounds having higher O/M ratios, e.g. Li-rich Li2MO3 (i. (1) U << Δ, electrons are extracted from the filled LHB, as observed in case of classical oneband cationic redox (Figure I.6c). Widely applies to oxides and fluorides having highly ionic (large Δ) M-L bonds (L being the ligand).

(2) U/2 ≈ Δ, resulting in overlapping LHB and O 2p non-bonding bands, that are simultaneously available for electrochemical activity leading to reversible anionic redox (Figure I.6d).

(3) U >> Δ, when electrons are being removed directly from the non-bonding O 2p band situated above the filled LHB (Figure I.6d), where the highly reactive O n-species formed upon charging may not be stabilized and oxygen loss occurs, as observed in Li2MnO3. Hence for sustainable anionic redox, the second situation is favorable. The exemplary cases are evidenced in Li1.33Ru0.5Sn0.17O2 or in Li2IrO3, as shown before. This is a result of increased covalency in the 4d/5d metals, leading to lower energy difference between M d and O p orbitals (i.e., lower Δ). Now, instead of adjusting the metal d level, another direction to lower the Δ consists in raising the ligand p level, by replacing the oxygen with ligands X = S, Se, Te, etc., that has lower electronegativity compared to O. Taking the case of sulfides (X = S), the S 3p band, having a higher energy than the O 2p band (i.e., S 3p band is closer to the Li/Li + redox level), can help for reversible anionic redox, however with a lower voltage. Hence, when designing a Li-rich layered sulfides alike the oxides, the S 3d band will largely overlap with the M d bands, see Figure I.7a,b. In fact, sulfide like LiCoS2 cannot form, since the high-energy sulfur 3p bands will pour electrons into the low-lying Co 3d bands, which can also be interpreted as the oxidation of S 2-by Co, resulting into the (S-S) 2-dimer. [START_REF] Nowack | Charge densities in CoS2 and NiS2 (pyrite structure)[END_REF] Further intercalation of lithium ions into CoS2 will generate metallic Co and Li2S, via conversion reaction. [START_REF] Poizot | Rationalization of the Low-Potential Reactivity of 3d-Metal-Based Inorganic Compounds toward Li[END_REF][START_REF] Yan | A study of novel anode material CoS2 for lithium ion battery[END_REF][START_REF] Grayfer | Anionic Redox Chemistry in Polysulfide Electrode Materials for Rechargeable Batteries[END_REF] Hence, for Li intercalation, only few light transition metals (left side of the periodic table), e.g., Ti, and V, etc., can combine with sulfur, as the 3d redox level of these metals are situated above the S 3p band (Figure I.7b). [START_REF] Goodenough | Locating redox couples in the layered sulfides with application to Cu[Cr2]S4[END_REF] Works on anionic redox in sulfides will be carried out as a part of this thesis. substituted Li2MnO3 (a), and Li1.33Ti0.67S2 or Li2TiS3 (b). The relative band positions are estimated based on references [START_REF] Luo | Charge-compensation in 3d-transition-metal-oxide intercalation cathodes through the generation of localized electron holes on oxygen[END_REF][START_REF] Goodenough | Locating redox couples in the layered sulfides with application to Cu[Cr2]S4[END_REF] . The label nb stands for non-bonding.

Based on this findings, anionic redox chemistry was extended to Na-cathode materials. Na-rich layered oxide Na2MnO3 had drawn tremendous attraction, as shown by theoretical studies. [START_REF] Zheng | Na2MnO3 as cathode materials for Na ion batteries: From first-principles investigations[END_REF][START_REF] Gao | Electronic Properties, Phase Transformation, and Anionic Redox of Monoclinic Na 2 MnO 3 Cathode Material for Sodium-Ion Batteries: First-Principle Calculations[END_REF] Interestingly, for Na, despite having larger ionic radius than Li, similar migration barrier (~0.7 eV) was predicted due to elongated Na-O bond. Furthermore, possibility of better performance was predicted for Na2MnO3, because of enhanced electronic conductivity, for smaller band gap that arises from weaker octahedral crystal field. [START_REF] Zheng | Na2MnO3 as cathode materials for Na ion batteries: From first-principles investigations[END_REF] However, experimentally the Na-rich Na2MnO3 could not be synthesized yet. [START_REF] Kataoka | Spinel manganese oxide: A high capacity positive electrode material for the sodium ion battery[END_REF] Rather, phases with Li or vacancy (□) in place of Na in the metal layer have been reported in literature, such as Na[Li0.2Mn0.8O2], Na[Li0.2Ni0.2Mn0.6O2], Na0. 0.86O2], etc., that shows only cationic redox (with O3-type crystal structure, space group C2/m). [START_REF] Mortemard De Boisse | Highly Reversible Oxygen-Redox Chemistry at 4.1 V in Na 4/7x[END_REF][START_REF] Van Nghia | Synthesis and Electrochemical Properties of Sodium Manganese-based Oxide Cathode Material for Sodium-ion Batteries[END_REF][START_REF] Nghia | Synthesis and electrochemical performances of layered NaLi0.2Ni0.2Mn0.6O2 cathode for sodium-ion batteries[END_REF] Nevertheless, few phases, like Na0. 67[Mg0.28Mn0.72O2], Na0. 67[Mn1-yZnyO2], and Na0. 67[Ni0.33Mn0.67O2], etc., were reported which possess anionic redox capability (with P2-type crystal structure, space group P63/mmc)). [START_REF] Maitra | Oxygen redox chemistry without excess alkali-metal ions in Na 2/3 [Mg 0.28 Mn 0.72 ]O 2[END_REF][START_REF] Yabuuchi | A new electrode material for rechargeable sodium batteries: P2-type Na 2/3[END_REF][START_REF] Bai | Anionic Redox Activity in a Newly Zn-Doped Sodium Layered Oxide P2-Na 2/3 Mn 1-y Zn y O 2 (0 < y < 0.23)[END_REF][START_REF] Yabuuchi | New O2/P2-type Li-Excess Layered Manganese Oxides as Promising Multi-Functional Electrode Materials for Rechargeable Li/Na Batteries[END_REF][START_REF] Cheng | Elucidation of anionic and cationic redox reactions in a prototype sodium layered oxide cathode[END_REF] Along this line, in sulfides, NaCrS2, NaCr1-yTiyS2, etc., were reported showing anionic activity. [START_REF] Shadike | Antisite occupation induced single anionic redox chemistry and structural stabilization of layered sodium chromium sulfide[END_REF][START_REF] Wang | Anionic redox reaction in layered NaCr2/3Ti1/3S2 through electron holes formation and dimerization of S-S[END_REF] Interestingly, these studies reveal that excess Na is not required to activate oxygen redox, that demands further studies. Stoichiometric Na-rich layered Na2MO3 could be successfully synthesized, only by

O 2p

Mn 3+/4+ moving to 4d or 5d TMs, such as Na2RuO3, Na2Ru1-ySnyO3, Na2IrO3, Na1.2Mn0.4Ir0.4O2 etc. [START_REF] Tamaru | Layered Na2RuO3 as a cathode material for Na-ion batteries[END_REF][START_REF] Mortemard De Boisse | Intermediate honeycomb ordering to trigger oxygen redox chemistry in layered battery electrode[END_REF][START_REF] Perez | Strong Oxygen Participation in the Redox Governing the Structural and Electrochemical Properties of Na-Rich Layered Oxide Na 2 IrO 3[END_REF][START_REF] Pearce | β-Na 1.7 IrO 3 : A Tridimensional Na-Ion Insertion Material with a Redox R-vi Active Oxygen Network[END_REF][START_REF] Rozier | Anionic redox chemistry in Na-rich Na 2 Ru 1-y Sn y O 3 positive electrode material for Na-ion batteries[END_REF][START_REF] Zhang | Manganese-Based Na-Rich Materials Boost Anionic Redox in High-Performance Layered Cathodes for Sodium-Ion Batteries[END_REF] Contrarily, sulfides show the possibility to be synthesized as Na-rich Na2TMS3 with 3d metals, as recently electrochemical properties of Na2TiS3 (O3-phase) were reported, [START_REF] Pele | Nouveaux materiaux d'electrodes pour microbatteries au sodium[END_REF][START_REF] Nasu | Amorphous Na 2 TiS 3 as an Active Material for All-solid-state Sodium Batteries[END_REF] that will be discussed later as a part of this thesis.

I.2 From

LIBs to all-solid-state batteries (ASSBs)

I.2.1 Why all-solid-state batteries

In addition to the ongoing efforts to further increase the energy density of Li-ion batteries (LIBs), their safety has raised more serious safety concerns. Steps to increase energy density by raising the upper cutoff voltages (that causes poorer stability of electrode-electrolyte interfaces), [START_REF] Xu | Nonaqueous Liquid Electrolytes for Lithium-Based Rechargeable Batteries[END_REF] using ultrathin (< 10 µm) polymeric separators (despite reinforcements using ceramic materials), [START_REF] Lagadec | Characterization and performance evaluation of lithiumion battery separators[END_REF] etc. have exposed more vulnerability toward internal short circuits and thermal runaway. Even though today's LIBs and battery management systems (BMS) are well-advanced to prevent the flammability of the liquid electrolytes, several news of fires due to LIBs have been reported in the last decade (Figure I.8). 82- 84 Furthermore, the push for large-scale applications of LIBs, such as batteries for electric vehicles, and grid energy storages, etc. imposed more safety requirements for LIBs. [START_REF]Dell's Exploding Computer and Other Image Problems -The New York Times[END_REF][START_REF] Blaze On | 787 Dreamliner at Heathrow 'triggered by trapped battery wires[END_REF][START_REF]Tesla Model S catches on fire in France during test drive[END_REF][START_REF]Why Are Samsung's Galaxy Note 7 Phones Exploding?[END_REF] .

Therefore, replacing the conventional flammable organic liquid electrolytes with safer solid electrolytes (SEs) is considered to be an ideal solution, and bulk-type all-solid-state Li-ion batteries have been reinvestigated intensively in recent years. 5,[START_REF] Gallagher | Boeing's Dreamliner Batteries "Inherently Unsafe"-And Yours May be Too[END_REF][START_REF] Meier | Feds review third Tesla fire as shares fall again[END_REF] Several other advantages of solid electrolytes can be beneficial in this regards, 1) Possibility of electrochemical stability over higher voltages. [START_REF] Oh | Bulk-Type All Solid-State Batteries with 5 V Class LiNi 0.5 Mn 1.5 O 4 Cathode and Li 10 GeP 2 S 12 Solid Electrolyte[END_REF][START_REF] Han | Negating interfacial impedance in garnet-based solid-state Li metal batteries[END_REF] 2) Longer cycle life, with low self-discharge. [START_REF] Li | Solid electrolyte: The key for high-voltage lithium batteries[END_REF] 3) Higher energy density, due to possibility of using Li-metal anode, and with other promising battery systems (e.g., Li-air, Li-S, etc.). [START_REF] Manthiram | Lithium battery chemistries enabled by solid-state electrolytes[END_REF] 4) Capability of increasing cell-level energy density by designing ASSBs with stacked bipolar electrodes that allows minimization of inactive encasing materials. [START_REF]Research Progress : Next Generation Secondary Batteries[END_REF][START_REF] Zeier | A solid future for battery development[END_REF] Next, we will describe how SEs could be incorporated in actual batteries.

I.2.2 Battery technologies using solid electrolytes (SEs)

Battery technologies using SEs can be traced long back. In 1967, researchers at Ford discovered in sodium β-alumina (with conductivity ~10 -3 S•cm -1 ) which successfully led to commercialization of high-temperature Na-S batteries operating at ~300 °C (Figure I.9a). [START_REF] Yao | Ion exchange properties of and rates of ionic diffusion in beta-alumina[END_REF][START_REF] Lu | Sodium-beta alumina batteries: Status and challenges[END_REF][START_REF] Dunn | Electrical Energy Storage for the Grid: A Battery of Choices[END_REF] However, oxide materials like β-alumina are synthesized and sintered at very high temperatures (>1200 °C), to eliminate grain boundary resistance, and thus impose difficulties in the fabrication of batteries. In 1992, researchers at the Oak Ridge National Laboratory developed amorphous glassy SE LiPON (e.g., Li3.3PO3.8N0.24) that was commercially utilized by various companies, e.g., Sakti3, Excellatron, etc., in thin film Limicrobatteries (Figure I.9b) with capacity of ~10 -50 mAh. 100 The LiPON has a low Li-conductivity (2×10 -6 S•cm -1 at RT), but has good electrochemical stability and negligible interfacial resistance with both metallic Li and the transition metal cathodes. 101,102 However, because of the high cost (synthesized by RF Magnetron deposition from Li3PO4 target in N2) and low Li-conductivity, the applications of LiPON were limited to microbatteries. In contrast, in the last decade, the field have been revisited intensely in order to realize bulk-type ASSBs that can replace the traditional high-energy LIBs (Figure I.9c), which is quite challenging as will be described next.

I.2.3 Requirements for ASSBs

In bulk-type ASSB, the use of the SEs in powder form allows the adoption of battery configuration that was developed for LIBs. To comply with minimum requirements for a functional battery, the SEs must possess, 1) High ionic conductivity (>10 -4 S•cm -1 at RT) 2) A wide electrochemical stability window (preferably 0 -5 V), and, 3) Excellent chemical stability with anode (Li-metal) and cathode materials.

4) Facile processing capability for ASSB fabrication

In this regard, as Li-ion conductivities for SEs (will be discussed in Section I.3) is comparable to liquid electrolyte (>10 -3 S•cm -1 ), ASSBs show potential to compete with conventional LIBs. To estimate the merit, the required thickness of the SEs can be calculated from Li-conductivity according to the equation,

𝑅 = 1 𝜎 • 𝑙 𝐴
where R is the resistance, σ is the conductivity, l is the thickness, and A is the area of the electrolyte. Therefore the required thickness range of the SE can be calculated to meet an estimated range of resistance comparable to liquid electrolytes. As shown in Figure I.10, thin (< 40 μm) layer of SEs are required. Therefore this task is challenging, and will require materials processing techniques such as thin film deposition, SE casting etc., in future. [104][105][106] Nevertheless, the thickness can be increased further for SEs, because of its high Li-ion transference number t+, according to the relation D + ∝ σ × t+, where D + is the diffusion coefficient of Li + ions. In SEs t+ ≈ 1 (i.e., only Li + moves, anions are immobile), whereas for liquid electrolytes typically t+ in the range of 0.3 -0.4. 81 But, the advantage of single-ion conducting character of SEs is often offset by difficulties in forming and maintaining intimate ionic contacts with active materials in the composite electrodes. Hence the need of sintering SEs prior to cycling and high pressure during the cycling of ASSBs. Furthermore the kinetic issues, such as, the interfacial impedance of the SE/electrode interface, space charge layer, etc., and chemical compatibility, e.g., inter-diffusion and chemical reactions interface must be considered. Overall, to achieve bulk-type ASSBs that show comparable performance with conventional LIBs, several critical challenges must be addressed, thus highlighting the need of fundamental studies and exploration of various directions. As the SEs are the most important parameter, they will be discussed next, starting with its historical account till now.

I.3 Evolution of solid electrolytes (SEs)

I.3.1 Journey towards conductivity surpassing liquid electrolytes

The journey of 'electrolytes' is considered to be started way back in the 1830s, when Michael Faraday discovered ion conduction in solid materials. In 1834, the recorded remarkable property of conduction in heated solid Ag2S and PbF2, 107 as he described following, "I formerly described a substance, sulfuret of silver, whose conducting power was increased by heat; and I have since then met with another as strongly affected in the same way: this is fluoride of lead. When a piece of that substance, which had been fused and cooled, was introduced into the circuit of a voltaic battery, it stopped the current. Being heated, it acquired conducting powers before it was visibly red hot in daylight; and even sparks could be taken against it whilst still solid." 108 This was probably the first observations of transition (Figure I.11) from a poor conducting state to superionic conducting states in a materials which is now called 'solid electrolytes'. 108 During this period, he discovered the motion of mobile ions in both liquid and solid electrolytes, which subsequently took two different paths of 'Electrochemistry' and 'Solid State Ionics'. 108 However, the 1960s are generally considered the turning point for solid electrolytes (SEs) when they were studied extensively. Subsequent development of SEs are outlined in Figure I.12. Though started with report of high Li-conducting Li3N, discovery of fast 2D Na + -transport phenomenon in sodium β-alumina (Na2O•11Al2O3) marked the era, as it was readily used for high-temperature Na-S batteries, as discussed earlier. [START_REF] Yao | Ion exchange properties of and rates of ionic diffusion in beta-alumina[END_REF][START_REF] Lu | Sodium-beta alumina batteries: Status and challenges[END_REF][START_REF] Dunn | Electrical Energy Storage for the Grid: A Battery of Choices[END_REF]109 Following this, Goodenough et al. in 1976, discovered the fast Na-conduction in NASICON-type compounds (Na1+xZr2SixP3-xO12, 0 ≤ x ≤ 3). 110 Inspired by this, Li-ion conductivity in NASICON-like structures was investigated, as Subramanian et al. reported a NaSICONtype LiTi2(PO4)3 in 1986 (with an conductivity of 2•10 -6 S•cm -1 at RT). 111 The ionic conductivity of LiTi2(PO4)3 could be further improved by partially substitution as the compounds Li1.3Al0.3Ti1.7(PO4)3 (LATP) exhibits a conductivity of 7•10 -4 S•cm -1 at RT. 112 During this time sulfide-based solid electrolytes, especially glassy sulfides, were also investigated intensely. Binary Li2S-P2S5 and Li2S-SiS2 glasses doped with LiI, were reported to possess ionic conductivity higher than 10 -4 S•cm -1 . [113][114][115][116][117] The sulfide glasses, and also oxide glasses (e.g., Li3.6Si0.6P0.4O4, ~10 -6 S•cm -1 at RT) were mostly used for thin film batteries. 117,118 Perovskite type Li3xLa2/3-xTiO3 was reported also with high conductivity. 119 Thus, though this fruitful period has seen the discovery of many new Li-ion conducting compounds, the conductivities of these solid electrolytes were far behind practical utilization and this research fall into oblivion. 101,110,120,121 Figure I.12. Development of solid electrolytes towards superionic conductivity. The ionic conductivities (in S•cm -1 ) are indicated in the parentheses. Note that ionic conductivity comparable to liquid electrolytes was achieved in 2011. Based on references [START_REF] Yao | Ion exchange properties of and rates of ionic diffusion in beta-alumina[END_REF][START_REF] Lu | Sodium-beta alumina batteries: Status and challenges[END_REF][START_REF] Dunn | Electrical Energy Storage for the Grid: A Battery of Choices[END_REF]101,[109][110][111][112][113][114][115][116][117][119][120][121][122][123][124][125][126][127][128] .

Recent safety issues of LIB's in the early 20th century and foreseen aspiration of industries has led researchers to re-visit the field intensely in the past decade and several new interesting compounds were found or their conductivities were improved. 123,124 The conductivity of Li1.3Al0.3Ti1.7(PO4)3 (LATP) could be much improved to of 1.12•10 -3 S•cm -1 at RT by using spark plasma sintering to decrease the grain boundary resistance. 122 Garnet-type Li7La3Zr2O12 (LLZO) with a cubic structure was revisited in 2007 and a high ionic conductivity of 3.7•10 -4 S•cm -1 at RT was demonstrated, 123 which could be further improved to >10 -3 S•cm -1 at RT in 2013, by partial hypervalent substitution of Zr with Al, Ga, Ta, Nb, etc. 129,130 However the ionic conductivities achieved (~10 -3 -10 -6 S•cm -1 ) were still trailing behind the liquid electrolytes 125 . In 2000, based on LISICON-type oxides, by replacing O with S, Kanno et al. discovered a new crystalline material family Li4-xGe1-xPxS4 (thio-LISICON) with conductivity of 2•10 -3 S•cm -1 at RT. 126 Along this family, it is only recently, the ionic conductivities comparable to the liquids (~10 -2 -10 -3 S•cm -1 ) have been achieved with the discovery of thio-phosphates Li11-xM2-xP1+xS12 (M = Si, Ge, Sn). The Li10GeP2S12 (LGPS), reported by Kanno and coworkers, exhibits a very high lithium ionic conductivity of 1.2•10 -2 S•cm -1 at RT. 127 Seino et al. (2014) reported a heat-treated glass-ceramic conductor Li7P3S11 with an extremely high ionic conductivity of 1.7•10 -2 S•cm -1 at RT. 131 Thus, improvement in ionic conductivities has been stellar, however, achieved only with few special crystal structure, which will be summarized in next section.

I.3.2 State of the art

Inorganic solid electrolytes (oxide and sulfide-based), can be crystalline, glassy, and glassceramics, and they are summarized here briefly in Table I.1 and Table I.2 Cons-

Low conductivity

Na ion conductors. In parallel, the pursuit of Na-ion conductors also has been intense in the last decade. As solid state electrolytes, though the research started as early as in 1970s, only few inorganic materials have been found, e.g., β-alumina, NaSICON (Na1+xZr2P3-xSixO12), and Na3PS4, etc., which offer Na-ion conductivity close to the liquid electrolytes. 110,155 However, oxide materials like βalumina and NaSICON are synthesized and sintered at very high temperatures (>1200 °C), to eliminate grain boundary resistance and furthermore show interfacial incompatibility with metallic Na anode. 156 Cubic Na3PS4 has been reported for its potential application in a RT solid-state Na-ion battery, but the conductivity was still low (~10 -4 S•cm -1 ). 157 Further efforts to design new Na-ion conductor is concentrated on Na3PS4-based structure, for example, substitution of P to synthesize Na11Sn2PS12, Na3SbS4 etc., and even substitution of S leading to Na3-xPS4-xClx, Na3PSe4, etc., and slight increase in conductivity up to ~10 -3 S•cm -1 was obtained (except in recently reported Na2.88Sb0.88W0.12S4, 32 mS•cm -1 ). [158][159][160][161][162] However, due very small stability window and interfacial incompatibility, these sulfidebased electrolytes could not be implemented in practical ASSBs. 156,163 From theoretical calculation, a small electrochemical stability window has been predicted for Na3PS4 (1. 55 -2.25 V) andNa3PSe4 (1.8 -2 [165][166][167][168][169][170] In addition to the long-range connectivity of Li + sites, Li-ion diffusion probability depends on the relative energies of the lattice sites involved in hopping and the energy barriers that separate them. Hence, percolation of low energy pathways throughout the crystal lattice is also important to achieve good long range diffusivity. 171 Thus, a number of factors regarding crystalline structure and materials chemistry affect the overall ionic conductivity, σ, which can be expressed by,

𝜎 = 𝑞 2 𝑘 B 𝑇 𝑛 𝑐 𝐷 0 exp (- 𝐸 a 𝑘 B 𝑇 )
where, nc is the concentration of mobile-ion carriers (e.g., vacancies or interstitials), D0 is the diffusion coefficient, Ea is the activation energy for ion transport, T is temperature, and kB is the Boltzmann constant. Following the equation, the prerequisites to improve ionic conductivities are, 1) An open crystal structural framework with 3D conduction pathways with optimum free volume for ion migration.

2) A high nc, with low Ea.

3) Highly polarizable ions to lower the migration energy barriers. 172 As a consequence, high ionic conductivity is generally obtained in few SE families that have 'special' crystal structures including channel size, free volume, polarizability, etc. Furthermore, within a SE family, the maximum ionic conductivity is obtained upon introducing configurational disorder (partial site occupancies or site disorder) resulting from either vacancies or interstitials. 171 Therefore, aliovalent substitution is generally adopted to create Li vacancies or interstitials, to increase disorder in the Li/Na-sublattice. 110,120,136 In aliovalent solid solution, ions are substituted by other ions of different charge, which consequently causes additional changes involving creation of vacancies or interstitials (ionic compensation) to conserve the electro-neutrality and thereby increasing the value of nc. To illustrate, let us take the example of Li4-3xAlxSiO4 solid solutions (Figure I.15). 136 In this case, it is possible to fill a set of interstitial sites completely or to empty completely a particular set of lattice sites. At the halfstage, the ionic conductivity should pass through a maximum, as the product of the concentration of mobile species, nc and sites to which they may migrate (1nc) is maximum when the concentrations of filled and empty sites are equal. Indeed the same has been observed for the solid solution. 173 But where to find disorder ? Often, SEs possess polymorphism and the high temperature polymorph (HTP) are generally highly disordered with higher crystal symmetry. For example, the garnet-type Li7La3Zr2O12 (LLZO, Figure I.16), crystallizes in a thermodynamically more stable tetragonal phase at RT with a σRT of ~10 -6 S•cm -1 . 174 At a temperature >650 °C, it transforms to a highconducting cubic polymorph, which exhibits an extrapolated σRT of ~10 -4 S•cm -1 . 174,175 It is evident that in both polymorphs, the crystal framework formed by dodecahedral LaO8 and octahedral ZrO6 is influenced very little but Li-distribution is greatly disordered in the HTP. To stabilize the cubic phase at RT, the most efficient way is to introduce Li vacancy into the Li-sublattice by hypervalent substitutions. 129,130 Highest conductivity of ~1.4 × 10 -3 S•cm -1 at RT was obtained for Ga-substituted LLZO (Li6.25Ga0.25La3Zr2O12) 176 . Small amount of Ga-substitution creates Li-vacancy (though the framework remains invariant) which makes the crystal structure inherently disorder, hence the stabilization. Li, La, and ZrO6 are shown in green, cyan, and purple, respectively. Note that the cubic phase can be stabilized at RT via hypervalent substitutions, e.g., by Al 3+ , Ga 3+ etc. (in Li site, i.e., Li7-3xGaxLa3-xZr2O12), or by Ta 5+ , Nb 5+ , Te 6+ , W 6+ etc. (in Zr site, i.e., Li7-xLa3Zr2-xTaxO12), etc. 129,130 Similarly, the high-conducting thio-phosphate Li11-xM2-xP1+xS12 (M = Si, Ge, Sn) phases were designed by aliovalent substitution in Li3PS4. 121 Ionic conductivity of Li3PS4 at RT was reported back in 1984 by Tachez et al,121 (Figure I.17) to be ~3•10 -7 S•cm -1 . However, with increasing temperature, a jump in the conductivity was observed due to an order-disorder phase transition at ~200 °C. 143 Therefore, the Li-disorder leads to 4 times higher conductivity and a smaller activation energy (Ea) of 0.46 eV, compared to 0.49 eV for the low temperature polymorph (LTP). 121,143,177 Figure I.17. Ionic conductivity of Li3PS4. Figure taken from Reference 121 . Kanno et al (2003) demonstrated partial substitution of P 5+ by Ge 4+ in Li3PS4. 126 Thus prepared Li4-xGe1-xPxS4 (or Li3+yGeyP1-yS4) solid solution showed gradual increase of ionic conductivity with increasing substitution and highest conductivity of 2•10 -3 S•cm -1 was reported at RT for Li3.25Ge0.25P0.75S4, where x = 0.75 (or y = 0.25). Nevertheless, in quest for higher conductivity, Kanno and coworkers (2011) discovered further optimized Li10GeP2S12 (or Li3.33Ge0.33P0.67S4), as known as LGPS composition. Furthermore, the HTP can be stabilized by changing synthesis method also, as Liu et al reported stabilization of the β-Li3PS4 by nano-structuring (~80 -100 nm) via wet-chemical synthesis. 128 Stabilization of HTPs will be carried out as a part of this thesis.

What are the other factors ? The anionic sublattice has a significant effect on Li + diffusivity in SEs. Larger and more polarizable framework is preferable for ionic conduction, as it decreases the interaction of Li-X bond (X = anions), thus making the Li's more mobile. Hence, replacement of larger and more polarizable S for O improves the conductivity by magnitude of orders (Figure I.18a). 126,178 Furthermore, it was exemplified by Ong et al, with the LGPX (Li10GeP2X12, X= O, S, Se) structure. Li + diffusivity is significantly slower in Li10GeP2O12 with higher activation energy than in Li10GeP2S12 (Figure I.18b). 172 In contrary, cation substitution within the SE family has much lesser impact. 172 However, only disordered crystal structure is not enough for superionic conductivity. From, first principles computational studies, crystal structural frameworks with body-center cubic (bcc) packing of anions (e.g., S 2-) are found to be superior for Li-ion diffusion, than in other close-packed frameworks. In bcc anion packing, Li ions occupy and migrate between face-sharing tetrahedral sites, leading to a low migration barrier of ~0.2 eV (Figure I.19a), that is typical in Li sulfides, e.g., LGPS,Li7P3S11,etc. 179 In addition to crystal structural features, from AIMD (Ab Initio Molecular Dynamics) simulations, it was shown that Li superionic conductors in sulfides, and specially in oxides (despite having bcc packing, e.g., LLZO), exhibit a unique diffusion mechanism distinctive from typical solid materials. 180 Here, multiple Li ions migrate simultaneously through a concerted mechanism instead of an isolated ion hopping in typical solids. During the concerted migration of multiple ions, when ions located at the high-energy sites migrate downhill, they cancels out a part of the energy barrier felt by other uphill-climbing ions. Hence, the concerted migration with a decreased migration barrier is kinetically favorable (Figure I.19b). Therefore, to activate super-ionic conduction, one needs to insert mobile ions into high-energy sites that activates the concerted ion migration.

I.3.4 Choosing the electrolyte: Oxides or Sulfides ?

High ionic conductivities has been obtained in both oxide and sulfide based SEs, as shown above. So which is better for practical applications ? To answer this, few other properties, such as, electrochemical and chemical stability and compatibility, and, processing capability for fabrication of ASSB, etc., must be considered. In this regard, oxide and sulfide based solid electrolytes possess various advantages over one other, as compared overall in metal. SEI and MEIC denote solid electrolyte interface, mixed electronic-ionic conductor, respectively. Figures adapted from reference 181,182 .

Because of the low electrochemical stability, the solid electrolytes decompose at low and high voltage. Upon Li contact, few SEs, e.g., LiPON, Li3PS4, etc., decomposes to Li binary compounds like as Li2O, Li3P, Li3N, Li2S, etc., that are stable against the high chemical potential (μLi) of Li metal, which passivate the solid electrolyte and inhibit further decomposition (Figure I.21b). 183,184 The passivation mechanism relies on the electronic insulating properties of the decomposition interphase layers and is not protective if the layer is electronically conductive. In LGPS, the decomposition layer consists Li3P, Li2S and in addition electronically conductive Li-Ge alloys (Li15Ge4) or Ge metal that helps to propagate the decomposition and increase of interfacial impedance, as shown experimentally (Figure I.21c). [185][186][187] Similarly, the absence of passivation mechanism explains the lithiation and reduction of LLTO 188,189 ,LATP 190 ,LAGP 191 ,and LISICON 192 , etc., that were observed experimentally. 154,193 .

In short, even though oxide based SEs are desired over sulfides because of their higher voltage stability window and ease of handling, sulfide SEs provide advantages in terms of ionic conductivity, and more importantly for device integration (Figure I.20). Sulfides are suitable for room-temperature processing, as they possess the unique mechanical properties of 'RT pressure sintering', see Figure I.22. Furthermore, due to its moderate elastic modulus, sulfide SEs can sustain the volume changes of electrode during cycle. Whereas, oxide SEs, being mechanically rigid, need high temperature and pressure sintering and prone to fragmentations and contact loss during cycling. Therefore, sulfide-based SEs, e.g., LGPS, β-Li3PS4 etc., have been mostly used for ASSBs, and some examples will be shown next.

I.4 Progress in full ASSBs

The SEs, especially sulfide SEs, have a narrow electrochemical stability window. Thus the cyclability of the ASSBs depend on the SEI formation which effectively extends the electrochemical window of the SEs. 193,194,198 .

High resistance at SE/cathode interface, due to resistive decomposition products, was confirmed by the electrochemical impedance spectroscopy (EIS, Figure I.23c), that further grows upon chargedischarge. 150,189,198,199 The oxidative degradation at the cathode side represents a major factor contributing to the capacity fading of ASSBs. In particular, for the application of high-capacity nickel-rich layered oxides or high-voltage cathode materials, the electrochemical stability of thiophosphate electrolytes toward high potentials is insufficient and requires surface-coatings by ZrO2, LiNbO3, LiNb0.5Ta0.5O3 (LNTO) or [200][201][202] Indeed, the LNTO-coated LiCoO2 performs much better in ASSB than uncoated LiCoO2 (Figure I.23d). 193 The interfacial resistance was decreased largely (Figure I.23e), but still remains due to the formation of Li-depleted layer at the interface originated from μLi differences at S-O interface. [203][204][205] In short, though sulfide-based SEs are preferred for its higher conductivity and facile integration for ASSB fabrication, tackling the interfaces has become the next big challenge. Works on compatibility of sulfide-based electrolyte/electrode will be carried out as a part of this thesis.

I.5 Chapter Conclusions

This chapter initially discussed how to increase the energy density of cathode materials, via the utilization of materials that accumulate both cationic and anionic redox process, provided we are able to solve the bottlenecks associated with the anionic redox. This is our motivation to work on anionic redox in sulfides, which will be discussed in Chapter III. Secondly, we recalled that the safety of LIBs need to be improved. This could be achieved via fabricating safe all-solid-state-batteries (ASSBs), where solid electrolytes (SEs) are used in place of flammable liquid electrolytes. To realize functional ASSBs, SEs must possess good electrochemical/chemical stability with capability of easy handling. Oxide based SEs with improved ionic conductivity are desired in this regard. Hence, we decided search new ionically conducting oxides by selecting and optimizing disordered polymorphs, what exactly will be addressed in the next chapter.

Chapter II

Searching Ionic Conductors in the Li(Na)-Zn-P-O phase diagram # _______________________ # This chapter includes the following publications that I co-authored: Chem. Mater. 2018, 30, 1379-1390; Inorg. Chem. 2019, 58, 1774-1781; and Saha et al., manuscript under preparation.

II.1 Chapter Introduction.

This chapter deals with searching new oxide-based SEs for application in ASSBs. In oxides, despite sufficiently high conductivity (up to ~10 -3 Scm -1 ) in few materials, conductivities are still lower than the liquid electrolytes, which furthermore is affected by high interfacial resistance. 125,206 Owing to greater stability, ease of handling and larger electrochemical window, oxides can be better candidates than the sulfide SEs, provided that we improve the conductivity, hence our efforts to develop new materials.

Usually superionic conductivity is observed, when the material is highly disordered. As disorder increases in high-temperature polymorphs, numerous materials show temperature driven phase transition with the higher temperature polymorphs (HTPs) showing the best ionic conductivity. Thus, intense research is devoted to the stabilization of the HTPs at RT, by various methods, such as quenching, nano-structuring, varying synthesis condition like using certain precursors, and by aliovalent substitution, etc., as has been discussed in Chapter I. et al, 2009. 207 In this family, several compounds are shown to exhibit polymorphism extensively, 207 and we have spotted Li/Na6Zn(P2O7)2, and Li/Na4Zn(PO4)2 as potential candidates (Figure II.1). We furthermore decided to explore their Na-counterpart also, provided they can be synthesized. We have studied the relation between the polymorphism and disorder in these 'model compounds' along with the variation of conductivities with structural modifications. Finally, we have explored various ways to stabilize the high temperature polymorphs (HTP) at RT, with the conduction mechanisms rationalized by bond valence energy landscape (BVEL) maps, which confirms the high sensitivity of ionic conductivity on disorder. Our studies on Li-ion conductors Li6Zn(PO4)2 and Li4Zn(PO4)2 will be presented in Section II.2 and II.3, followed by its extension to Na-ion conductor Na4Zn(PO4)2 (Section II.4). 6 Zn(P 2 O 7 ) 2 The low temperature phase (LTP) of Li6Zn(P2O7)2, that was reported by Rao et al, 207 was showed to transforms to a HTP following a phase transition at 365 °C. XRD patterns for both the phases were reported, indexed with a possible monoclinic unit cells, but no structural model was presented. Herein we choose this compound to study the influence of polymorphism and disorder upon ion conduction.

II.2 Li

II.2.1 Probing the polymorphism with temperature

Li6Zn(P2O7)2 sample was prepared by classical solid state reaction of Li2CO3, ZnO, and NH4H2PO4 at 700 °C, detailed method is described in the Appendix (Section A1.1). The X-ray powder diffraction (XRD) pattern of the sample, when prepared by quenching from the synthesis temperature, is in agreement with the list of peak positions reported by Rao et al 207 , but our pattern indicates that the powder is composed of a mixture of different polymorphs of Li6Zn(P2O7)2. Hence, the quenched samples was further annealed at 300 °C (i.e. below the phase transition temperature) to get pure α-Li6Zn(P2O7)2. However, by annealing at temperatures higher than the phase transition temperature and successive quenching we could not succeed in stabilizing the other high-temperature polymorphs in pure form. For such a reason, hereafter if not otherwise specified, the pure α-Li6Zn(P2O7)2 sample was used as the starting material and its polymorphism was studied both via thermo-gravimetric analysis (TGA) and in situ X-ray diffraction as a function of the temperature. One should first mention that all patterns present similarities and only differ in peak's positions and in the presence, or not, of small additional peaks. This indicates that all six polymorphs are likely structurally related. The first transition was observed at ~335 °C. Till this temperature the compound crystalizes in α-Li6Zn(P2O7)2 form, but increasing the temperature to 380°C leads to a disappearance of its low angle peaks to the expense of new one at 2θ = 3.2°, hence leading to another polymorph denoted as the β-phase. Then upon further heating we observe the subsequent growth of the γ-polymorph (380°C to 470°C) and -polymorph (470°C to 569°C) which presents characteristic peaks at 2θ = 5.6° and 2θ = 4.1° and 5.18°, respectively. Afterwards, there is appearance of the -polymorph, whose existence is limited to a narrow temperature range (569-593°C), and XRD exhibits a pattern very close to the one with however the disappearance of the doublet around 2θ = 4.1°. Lastly, for temperature above 593°C, there is the appearance of the ζ-polymorph which shows XRD pattern with fewer peaks, indicative of a higher symmetry and/or smaller unit cell of this polymorph.

The synchrotron-XRD patterns collected during cooling (Figure II.3) indicate, in contrast to heating, two phase transitions are observed which start at ~487 °C and ~369 °C, in agreement with the DSC data (Figure II.2a). Moreover, upon cooling to RT, the sample does not transform back to pure α-Li6Zn(P2O7)2 polymorph, since its RT XRD consists in a mixture of the α-polymorph with other high temperature phases. However, we believe such a difference to be rooted in the slow kinetics of phase transitions upon cooling, as supported by the feasibility to transform the multiphase cooled sample into the pure αpolymorph by a short reannealing step at 300 °C.

II.2.2 Structural solution of ordered α-phase

Next, in the absence of any single crystal, we directly used the powder XRD patterns to solve the crystal structure of the two main phases. The pattern recorded at RT for the α-phase was indexed using the DICVOL program 208 with a monoclinic unit cell with lattice parameters a = 25.66942 (10) Å, b = 10.25856(4) Å, c = 9.10213(3) Å, and β = 104.7074(3)°. This leads to a unit cell volume of V = 2313.0(1) Å 3 which can accommodate eight formulas per unit cell (Z = 8). Systematic extinctions were found to be consistent with C 2/c space group (S.G.). Then a first set of Zn, P and O atomic positions were found using the 'FOX' program, the PO4 groups being treated as rigid bodies. 209 All Zn, P and O atoms are placed on general position 8f (except two oxygen atoms, denoted as O22 and O44), which

II.2.3 Structural solution of disordered ζ -phase

Similarly we proceed to solve the structure of the high temperature ζ-phase from SXRD and NPD patterns recorded at 600°C. The best fit (χ 2 = 3.51, RB= 9.38%, for NPD) could be found in a monoclinic unit cell with space group C 2/c, but with much smaller unit cell than the α-phase. The lattice parameters were found to be a = 13.24574 (6) Å, b = 5.19764(2) Å, c = 8.92388(4) Å, and  = 105.1842(4)°, leading to a unit cell volume V = 592.8(2) Å 3 that is one fourth of the volume of the RT monoclinic unit cell and therefore can accommodate two formulas per unit cell (Z = 2). The high temperature ζunit cell vectors are linked to the RT α-Li6Zn(P2O7)2 ones by the following relation:

( 𝐚 𝜁 𝐛 𝜁 𝐜 𝜁 ) = ( ½ 0 0 0 ½ 0 0 0 1 ) ( 𝐚 𝛼 𝐛 𝛼 𝐜 𝛼 )
The structural model for the ζ-phase was obtained from atomic positions of the α-phase at RT to which the transformation matrix was applied. None of these phase present XRD patterns that can be fully indexed with the αor the ζ-unit cells. However, the resemblance in the main peaks indicates that the P2O7 4-backbone remains identical,

a b α-Li 6 Zn(P 2 O 7 ) 2 ζ-Li 6 Zn(P 2 O 7 ) 2 (a) (b) α-Li 6 Zn(P 2 O 7 ) 2 ζ-Li 6 Zn(P 2 O 7 ) 2 (c) (d) c a
whereas gradual disappearance of low angle peaks in successive polymorphs indicates increasing disorder. We expect that further differences between structures arise from complex and partial Li/Zn orderings that may lead to pretty large unit cells, if not incommensurate. Nevertheless, looking at the subtle differences in the SXRD patterns of the βwith the α-polymorph, it could be inferred that the polymorphs have similar structure where both of the polymorphs possess the low angle peaks (2 < 3.5°). Similarly and phases have similar structure with characteristic mid-low angle peaks at 2 = 3.5° -5.8° and and ζ-are similar where these two sets of peaks are absent. Further studies, including a thorough NPD exploration versus temperature, will be needed to solve these intermediate structures.

Also, the phases observed on cooling remain largely unexplored since the kinetics of phase transformation seems to be fast on heating but slow on cooling, as three days of annealing the cooled phase at 300 °C are needed to recover the α phase.

II.2.4 Influence of disorder on ionic conductivity

The ionic conductivity of the compound was measured from the AC impedance spectra recorded on a sintered pellet of α-Li6Zn( P2O7 -12 F and ~10 -11 F associated with the Q of the deconvoluted semicircles at higher and lower frequency, respectively. Such values are respectively reflecting the grains (σg) and the grain boundaries (σgb) contributions, to the overall conductivity (see Appendix Section A1.1 for details).

The ionic conductivity of the α-Li6Zn(P2O7)2 sample is summarized in Figure II.8. The plot exhibits a linear variation of ln(σT) with inverse temperature indicating a thermally activated Arrheniustype behavior, σ(T)=A/T•exp(-Ea/kBT), where kB is the Boltzmann constant, A is a pre-exponential factor and T is the temperature. The bulk ionic conductivity (σg) at 200°C was found to be ~6 × 10 -8 S•cm -1 with an Ea of 0.9 eV and σg of ~2 × 10 -12 S•cm -1 at RT could be deduced by extrapolation of the Arrhenius plot. Variation of the total ionic conductivity (σtot) is also plotted (line a in Figure II.8b), which indicates a much lower σtot value of ~1 × 10 -14 S•cm -1 and Ea = ~1.1 eV at RT. This indicates that the conductivity in this material is limited by the grain boundary conductivity (σgb), which is ~1.4 × 10 -14 S•cm -1 at RT (Ea = ~1.1 eV).

T (°C) (a) (b)
Further increase was observed at ~590 °C, adhering to the phase transition to the ζ-polymorph (line c in Figure II.8b) and gives an extrapolated σtot value of ~1.5 × 10 -5 S•cm -1 at RT, with Ea value of 0.35 eV. For these polymorphs, the behavior of σg and σgb could not be obtained, because of the shift of the grains and grain boundary responses toward higher frequencies with temperature, which finally moves beyond the measurable frequency window of the instrument leaving only the Warburg tail observable (see Figure II.9). At this point, it's worth noting that while SXRD revealed multitude phase transitions, only two of them have a strong effect on the ionic conductivity.

II.2.5 Mechanism of Li-conduction from BVEL calculations

As an attempt to correlate the observed increase in ionic conductivity with the structural transition, we have performed bond valence energy landscape (BVEL) calculations of the αand ζ- In short, we thoroughly studied the polymorphism in Li6Zn(P2O7)2 and explored its impact on ionic conductivity. It showed two jumps in ionic conductivity with increasing temperature, with a high activation barrier for Li + conduction for the α than the -polymorph (0.9 eV vs. 0.35 eV). Altogether, these results suggest that not only disorder, but also cation-mixing can lead to new conduction pathways, which can increase the ionic conductivity. However, despite our efforts we could not stabilize the polymorph at RT by hypervalent synthesis. Next, we move to similar studies on Li4Zn(PO4)2 to verify whether the HTPs can be stabilized.

II.3 Li 4 Zn(PO 4 ) 2

The Li4Zn(PO4)2 was reported as early as in 1976, it was never explored in details for its transport properties despite its rich chemistry. 210 This compound received further attention in 1986 by West et al., who explored its thermal stability by differential thermal analysis of Li3PO4-Zn3(PO4)2 solid solutions. 211 Interestingly, they reported three polymorphs, low temperature the α-Li4Zn(PO4)2 and the high temperature βand γ-Li4Zn(PO4)2, and moreover they speculated that the β-phase derives from the α-phase by cation disordering. One year later, the same group, prepared the pure compound by solid state synthesis and explored its ionic conductivity. 212 They found that Li4Zn(PO4)2 is a poor Li + conductor, but noticed an increase of conductivity for the β-phase. Fifteen years later, Jensen et al., have prepared the compound by hydrothermal method and described an accurate structural model for the αphase. 213 However, nobody has described the crystal structures of the high temperature phases till date.

β-Li4Zn(PO4)2 has been speculated to be isostructural to γ-Li3PO4, 211,213 which is well-known since long as a good Li + conductor. 138 More specifically, this compound can be viewed as Li3PO4 with two Li + substituted by one Zn 2+ , and since there is no twofold positions in Li3PO4 to accommodate the extra Zn 2+ , the Zn 2+ and vacancies are probably disordered in fourfold Li sites. This could explain the higher conductivity reported for the β phase, bearing in mind that in solid electrolytes, disordered polymorphs possess greater conductivity than their ordered counterparts. 130,177,214 In this study, we studied its conductivity governed by its disorder in the structure and explored various ways to stabilize the β-polymorph at RT.

II.3.1 Crystal structure of Li4Zn(PO4)2

Li4Zn(PO4)2 sample was prepared by following classical solid state reaction of Li2CO3, ZnO, and NH4H2PO4 at 700 °C, see the Appendix Section A1.1 for details. The XRD pattern of the as-prepared pristine sample, α-Li4Zn(PO4)2 (RT form), can be perfectly fitted with a monoclinic unit cell having a P 21/c space group and lattice parameters a = 8.11990( 16

II.3.2 Order-disorder transition with temperature

The α-Li4Zn(PO4)2 phase is known to undergo phase transition at high temperature to β-Li4Zn(PO4)2, 213,212 but the crystal structure of β-Li4Zn(PO4)2 was never reported, to our best knowledge. Hence we have carried out systematic investigations of the phase transition and finally have been able to deduce a structural model for the β-Li4Zn( PO4

II.3.3 Structural solution of the β-polymorph

Note that our efforts to quench the high temperature β phase at RT were unsuccessful; therefore we directly used the pattern recorded at 600°C to solve the structure for the β phase. The β phase presents a powder pattern similar to the α-phase, except for some peaks that have disappeared, especially in the low angle region. This relative similarity in diffraction pattern hints for a structural relationship between the and the structures. The XRD pattern of β-Li4Zn(PO4)2 at 600°C could be indexed using the Dicvol program 208 with an orthorhombic unit cell with lattice parameters a = 10.30571(4) Å, b = 6.50156(3) Å and c = 5.00496(2) Å, leading to a unit cell volume V = 335.348(2) Å 3 that is half the volume of the RT monoclinic unit cell and therefore can accommodate two formulas per unit cell (Z = 2). The orthorhombic unit cell vectors are linked to the monoclinic α-Li4Zn(PO4)2 ones by the following relation:

( 𝐚 ortho 𝐛 ortho 𝐜 ortho ) = ( 0 1 0 ½ 0 -½ ½ 0 ½ ) ( 𝐚 mono 𝐛 mono 𝐜 mono )
Systematic extinctions were found to be consistent with Pnma space group. The orthorhombic unit cell is plotted on top of the α structure in Figure II.16a, and where, looking at two adjacent orthorhombic unit cells, it clearly appears that the high temperature β phase should involve a disorder between Zn and Li2 atoms to comply invariance through lattice translation. Therefore, the XRD and NPD of the β-phase recorded at 600°C were then simultaneously refined by the Rietveld method, introducing mixing between Zn and Li cations. Moreover, it should be noted that this volume increase results from anisotropic changes in the b, c and β unit cell parameters.

II.3.4 Stabilization of β-phase by ball-milling

Having found that the β-Li4Zn(PO4)2 is a disordered counterpart of the α-Li4Zn(PO4)2 polymorph, it became interesting to stabilize this phase at RT. Since ball-milling is known to increase disorder by repeated particles fracturing and welding, as was demonstrated recently for sulphates 215 , we tried to stabilize the β phase by ball-milling. More specifically, the as-prepared α-Li4Zn(PO4)2 was ball milled for 2 hours. The XRD pattern of the obtained powder shows broad peaks that can be nicely fitted with the structural model of the β-Li4Zn(PO4)2 polymorph deduced earlier (Figure II.17). From the width of the Bragg peaks, we could further deduce that the ball-milled powder is made of particles with crystallites size of 10(±2) nm.

II.3.5 Stabilization of β-phase by aliovalent substitution

Chemical substitution was then explored as an alternative to ball milling to introduce disorder. More precisely, we attempted the partial substitution of Zn 2+ with M = Al 3+ , Ga 3+ and In 3+ in Li4Zn(PO4)2 with the hope to create disorder in Li-sublattice by introducing Li + deficiency for charge compensation and hence to promote the β-phase stabilization. All samples of nominal compositions of Li4-xZn1-xMx(PO4)2 were prepared by quenching the pellet from the synthesis temperature (800°C). We found that Al 3+ and In 3+ are unable to form solid solutions with Li4Zn(PO4)2, leading to impurity phases ( Appendix Figure A1.3 and Appendix Figure A1.4).

This contrasts with the incorporation of Ga 3+ that was successful for x ≤ 0.

II.3.6 Structural solution of the Ga-substituted phases

Note that for x = 0.3 the collected XRD/SXRD pattern appears similar, at first sight, to the one of the high temperature β-Li4Zn(PO4)2 phase, hence suggesting that Ga-substitution is efficient to preserve some Li/Zn disorder even at RT. Taking a closer look at the Li3.7Zn0.7Ga0.3(PO4)2 pattern, it appears that all peaks can be indexed with the same orthorhombic cell as β-Li4Zn(PO4)2 with slightly smaller lattice parameters, a = 4.9391( 2 A1.5) will be denoted as β'hereafter, only differs from the βstructure by the fact that the 8d site in Pnma splits into two crystallographic sites of multiplicity 4, indicating that introducing Ga segregates Zn and Ga on two distinct positions in P212121, both of them being mixed with Li. Note that at high temperature (above 300 °C, Appendix At this stage, it is worth reminding that all Ga-substituted samples mentioned previously were prepared by quenching from their synthesis temperature (800°C). Samples prepared by slow-cooling (sc) to RT (1 °C/min) show slight but notable discrepancy in the XRD patterns compared to those prepared by quenching, for Ga contents of x ≥ 0.3. Two compounds, sc-Li3.7Zn0.7Ga0.3(PO4)2 and sc-Li3.5Zn0.5Ga0.5(PO4), were examined with special care. At RT, both samples present a XRD pattern characteristic of the β'-phase (P212121) along with many additional small superstructure peaks. The temperature-dependent XRD study (Appendix Figure A1.5) of the compositions mentioned above shows that the XRD patterns of the sample transforms into the β structure (Pnma space group, see

Appendix Figure A1.6) above 380°C.
Finally, to confirm the phase transitions with temperature in the quenched Ga-substituted samples, we measured the DSC curve of all quenched samples. The data have been summarized in Figure II.23a and are in perfect agreement with high-temperature XRD experiments. We see that the phase-transition temperatures to the high temperature β phase slightly decrease with increased Ga content, while the enthalpy of the phase transition reduces drastically and finally disappears for xGa= 0.45. This confirms that the substitution creates Li + deficiency and an inherently disordered Lisublattice, which in turn reduces the free energy associated with the ordering of Li-sublattice, eventually leading to stabilization of the disordered phase. Hence stabilization of the β-phase at RT can become

II.3.7 Transition in Ionic Conductivity

We measured the ionic conductivity of all compounds by ac impedance spectroscopy over a broad frequency and temperature range. The impedance spectra were recorded on sintered pellets between ionically blocking platinum electrodes (see the Appendix section A1.1 for details). Figure II.24a shows a typical Nyquist impedance plot of the sintered pellets of the pristine Li4Zn(PO4)2 sample at 300 °C. Each recorded spectra shows a depressed semicircle at high frequencies followed by a Warburg tail at lower frequency region. The low frequency spike represents the impedance response due to the blocking of mobile Li ions at the electrode interface with Pt electrode and it confirms that The total ionic conductivity (σtot) was obtained by using the value of intercept of the Warburg tail with the semicircle. We have explored the impedance spectra further by examining the depressed semicircle at high frequency. The depression of the semicircle cannot be accounted by the inhomogeneities at the sample-electrode interface, hence the need of deconvolution of the semicircle to separate the different contributions. The deconvolution to different contributions, as obtained from the fitting, has been shown in the same curve and also zoomed in the inset. Contributions from the grains and grain boundary, has been indicated by green and red semicircles, respectively, and Warburg tail by black line. Some typical frequencies has also been indicated, the highest two of them are the resonance frequencies corresponding to grains and grain boundaries, respectively. Inset shows the equivalent circuit used to fit the spectra. Arrhenius plot of σtot of Li4Zn(PO4)2 (b) and ball-milled β-Li4Zn(PO4)2 (c); in (b) the blue and red background refers to αand β-Li4Zn(PO4)2 polymorph, respectively.

Since the impedance response of polycrystalline materials is expected to be composed of grain (bulk) and grain boundaries (surface) contributions, we decouple the semicircle to two contributions by fitting the whole impedance spectra to an equivalent circuit model (Rg//Qg)(Rgb//Qgb)Wo (inset of Figure II.24a); where R and Q represent the resistance and constant phase element associated with the grains (g) or grain boundaries (gb) of the sample, respectively, and Wo is the Warburg resistance. The fitting produces an capacitances of ~10 -12 F and ~10 -11 F associated with the constant phase element, for the high and low frequency deconvoluted semicircles, respectively and hence can be assigned to grains (σg) and grain boundaries (σgb) contributions, respectively 216 . A typical fitting is shown in Figure II.24a. This assignment was further confirmed by calculating the characteristic resonance frequencies (fr) of the semicircles which generally appear at the apexes of the deconvoluted semicircles; as indicated in Figure II.24a. However, with increasing temperature, both grains and grain boundary responses gradually shift towards higher frequencies and finally move beyond the measurable frequency window of the instrument leaving only the Warburg tail observable. Hence deconvolution of grain conductivity (σg) and grain boundary conductivity (σgb) was not possible above ~300-400 °C for the samples and only σtot was obtained. plot to the RT, with a similar Ea of ~1.3 eV for all the conduction processes. Altogether, this implies that this compound is not a good conducting material at low temperature, at least in its α-form.

We measured the conductivity of the pristine Li4Zn(PO4)2 sample over a broader temperature range of 200 -600 °C and reported the measured σtot values in Figure II.24b. A sudden change in the slope of the Arrhenius plot occurs at ~425 °C; temperature at which the α→β structural transition was shown to occur. Note that the high temperature β-Li4Zn(PO4)2 polymorph shows a much higher conductivity with an Ea of 1.1 eV and a RT σtot of 1.8x10 -15 S•cm -1 . For sake of completion we also measured the transport properties of our new RT β-Li4Zn(PO4)2 phase obtained by ball-milling of the α-Li4Zn(PO4)2 phase (Figure II.24c). Interestingly, the ball-milled βsample shows an extrapolated RT σtot value of 2.4•10 -15 S•cm -1 , nearly similar to that obtained from the high-temperature β-phase of Li4Zn(PO4)2 with however a slightly lower activation energy (Ea = 1.0 eV). We exploited this data further by decoupling the grains and grain boundaries contributions and found activations energies of 1.1 and 1.0 eV, respectively, with corresponding RT conductivities of 1.7•10 -14 S•cm -1 (σg) and 3•10 -15 S•cm -1 (σgb). Overall, this further confirms the better transport properties of the β-polymorph even when prepared at RT.

II.3.8 Higher conductivity in RT-stabilized phases

Encouraged by these findings, we also explored the transport properties of the quenched Li4-xZn1-xGax(PO4)2 samples. The results, shown in Figure II.25, reveal a steady increase of the ionic conductivities together with a progressive smearing of the jump in conductivity at the α→β'/β transition with increasing the Ga 3+ substitution (xGa). More specifically, the change of slope in conductivity at the structural transition gradually decreases till a Ga 3+ content of 0.15 pfu. and then disappears completely with further increase of Ga content. This is consistent with the results obtained from XRD experiments which show that the β'-phase is formed beyond xGa = 0.25 and also from DSC measurements which indicate a shift/smearing of the phase transition with increasing xGa content. Further insight onto these transport properties of members of the Li4-xZn1-xGax(PO4)2 series can be gained by further analyzing the impedance data with the key values reported in Table 3. Firstly, there is an increase by several orders of magnitude of the σg RT values with increasing the xGa content so that σg of ~1-2 x10 -12 S•cm -1 can be obtained for xGa = 0.05, and 0.10 as compared to σg = 5.55•10 -18 S•cm -1 for xGa=0, while structure-wise the samples are still pertaining to the α-phase in the composition range. Such an effect is most likely nested in the increasing amount of disorder that we can spot by XRD from the broadening of Bragg reflections characterizing the α-polymorph (Figure II.19b). Moreover, we notice a drastic fall in the Ea to 0.8 and 0.9 eV, respectively for the above-mentioned compositions. Interestingly, we could observe that σgb tracks σg and increases to ~4 -8 x10 -14 S•cm -1 at RT, with σtot reaching ~4 -8 x10 -14 S•cm -1 at RT. Upon further increasing the xGa content beyond 0.05, the RT σg values slightly increase prior to saturate at values of ~6 x10 -12 S•cm -1 (β-phase, Table 3) while structurewise the compounds evolve from the α _ phase to the β'-phase. Similarly, gb follows the same variation over this composition range while the activation energy remains nearly constant. It is only for samples having higher Ga contents (xGa > 0.40) and crystalizing in the β-structure that we observe a noticeable decease in σg to ~2 x10 -12 S•cm -1 and an increase in Ea from 0.9 to 1.0 eV. Altogether, these findings fail to establish a direct correlation between structure and conductivity and further analysis is required to interpret this trend. Nevertheless, the ac conductivity is ionic in nature, and this was further confirmed by direct current (DC) polarization experiments, which indicates that the electronic contribution is negligible, see

II.3.9 Conduction mechanism

To retrieve qualitative information concerning the connection existing between structure and conductivity, BVEL was employed. The BVEL maps representing the Li + conduction pathways in the percolation paths, involving Li2, Li3 and Li4 sites and not the sterically hindered Li1 site, appear along the b-c plane at 1.05 eV above the minimum energy. This is fully consistent with its layered crystal structure and Li + percolation occurs along the layers, whereas hopping of Li + between layers is not possible because of the Zn hindrance between the Li layers. This explains the high Ea and very low σg of the sample. Although BVEL could account for the greater conductivity for the β'/β-polymorph as opposed to the α-one, the feasibility to obtain ionic conductivities for the α-polymorph as high as the β-phase at low Ga substitution level is not well understood. First let's recall that ionic conductivity is the result of a delicate balance between vacancies, disorder, and defects among others 125 . Numerous ionic conductors show order-disorder phase transition with the high temperature phase having usually a higher ionic conductivity, owing to its higher structural symmetry which flattens the potential landscape hence leading to reduced activation barrier for Li + jump as well as to its enhanced disorder that increases the number of energy equivalent Li-sites having partial occupancies. 217 The Ga-substitutions, besides promoting structural disorder in the samples, introduces Li vacancies. However, the role of the vacancies remains questionable for the above observation as the ionic conductivity is nearly similar for α-Li3.9Zn0.9Ga0.1(PO4)2 and β'-Li3.9Zn0.9Ga0.1(PO4)2. To understand the respective contributions of vacancies and disorder, we plotted the variation of the log(σg) as a function of the concentration of Li vacancies and of the broadening of the Bragg peaks characteristic to the α-polymorph as an estimation Note that the ionic conductivity drastically increases with increasing the amount of vacancies (x = 0.05) with afterwards another increase from x = 0.15 which seems to be related to an increase in structural disorder. We thus believe that the disorder, as evidenced from the peak broadening which indicates the presence of antiphase boundaries and/or small disordered β'/β domains in the structure, as was also observed in rocksalt-type oxides, 218 could trigger β'/β-like conduction even at low Ga-substitution level. Additionally, the conductivity of a sample depends upon the surface contact of a grain with its neighbors, hence stressing the importance of grain boundaries vs. grain. A simultaneous increase of (σb) and (σgb) with increasing the doping content is frequently reported 219,220 . Our samples, being synthesized and sintered identically, are expected to have similar grain size and surface contacts. Thus for compounds having the highest conducting grains, the number of charge carriers crossing the grain boundaries will be greater as well, hence providing a possible explanation to the systematic tracking of grain and grain boundary conductivities observed in our samples. Furthermore, we have collected FESEM images of the cross-sections of pellets of Li4Zn(PO4)2, and Li3.95Zn0.95Ga0.05(PO4)2, and they are shown in Figure II.30. As evident, Ga-substitution improves the grain boundaries also, which could be another reason for the improvement of the grain boundary conductivity. However, further studies, such as, TEM studies on grain boundaries for space charge region, 221 or for amorphous second phase, 222 etc., are needed to understand this behaviour. Next we extended our studies to Na-ion conductors, to study the implications in structural and conductivity aspects, when Li is replaced by Na.

II.4 Na 4 Zn(PO 4 ) 2

Conductivity in the Na-counterparts of these Li-phases, i.e., Na6Zn(P2O7)2, and Na4Zn(PO4)2, were not explored before in the literature. Our effort to synthesize the Na6Zn(P2O7)2 was not successful (shown in the Appendix Figure A1.7). In contrast, the Na4Zn(PO4)2 could be successfully synthesized, as indicated by Irvine and West (1989), 223 which can be visualized as Zn-substituted Na3PO4. Indeed Na3PO4 based systems are quite interesting, because they comprise dynamic rotational disorders in the anionic sublattice, which can provide an extra parameter for ionic conductivity, alongside the usual cationic disorder. The rotational disorders can enhance ionic conductivity, by the "paddle wheel" mechanism, where the cation hopping is accompanied by a simultaneous rotation of the anion leading to lowering of the activation barrier for Na-hopping. [224][225][226] Indeed research in 1980s on Na3PO4 showed quite promising Na-conductivity, as it shows high ionic conductivity, ~5•10 -3 S•cm -1 at 300 °C. 224,227,228 At RT, if quenched, it crystallizes into tetragonal polymorph (space group 𝑃4 ̅ 2 1 𝑐), which progressively transforms to hydrated phases with time by absorbing moisture. [228][229][230] The tetragonal polymorph has been stabilized at RT by making solid solution with AlPO4, Na2SO4, and Na4SiO4, but without any significant improvement in ionic conductivity yet. 227,231 Therefore, the Na3PO4 based systems not only provide an excellent playground to solid state chemists, but also significantly interesting to design new Na-ionic conductors.

II.4.1 Polymorphism in Na4Zn(PO4)2

Na4Zn(PO4)2 was synthesized by solid state reactions of Na2CO3, ZnO, and NH4H2PO4 in stoichiometric amounts, at 700 °C (details in the Appendix Section A1.1). As-synthesized Na4Zn(PO4)2 samples present different X-ray diffraction (XRD) patterns when the pellet is recovered by quenching from 700 °C (denoted hereafter as β-polymorph), or by cooling slowly (1 °C/min) to RT (denoted hereafter as α-polymorph), see the patterns at 50 °C in Figure II.31a,c. The origin of the differences in the XRD patterns were investigated by probing the temperature-dependent structural evolution of the as-synthesized Na4Zn(PO4)2 samples, as described next.

II.4.2 Structural solutions

The newly synthesized Na4Zn(PO4)2 not only shows polymorphic behaviour under the influence of temperature, but also presents interesting kinetic behaviour owing to the presence of Na. Hence, to explain the phenomena, we first explored the crystal structures of the polymorphs in detail. The structures were attempted to solve by studying their selected area electron diffraction (SAED) patterns, SXRD, and neutron powder diffraction (NPD) patterns and are described below.

a) Structure of γ-Na4Zn(PO4)2

The structure of γ-Na4Zn(PO4)2 was solved by treating its SXRD pattern (collected at 500 °C), which can be indexed in the 𝐹𝑚3 ̅ 𝑚 space group with lattice parameter a = 7.46901(4) Å, i.e., a structure similar to the one adopted by γ-Na3PO4, with a slightly smaller (3%) unit cell. 232 Several structural models were constructed from Na3PO4, replacing one Na by one Zn and a vacancy, either on the 8c or on the 4b Wyckoff sites. The best Rietveld refinement was obtained by placing Na and Zn on the 8c site, while the 4b site is only 50% occupied by Na. O was placed on the 96k site, instead of 24e in γ-Na3PO4, so that this structure exhibits a fully rotational disorder on PO4 with are seen as balls. The final refinement is shown in Figure II.36 and the structural parameters are gathered in Appendix Table A1.8.

b) Structure of β-Na4Zn(PO4)2

The structure of β-Na4Zn(PO4)2 was solved using SXRD and NPD patterns recorded on a quenched sample. As mentioned before, the XRD pattern of the β-phase shows one extra peak compared to the γ pattern. This indicates a close relation between both structures. Indexation of these peaks requires a doubling of the cubic unit cell parameters (see Figure S2). The pattern can be indexed with the cubic F23 space group and lattice parameter a = 14.786037(10) Å. Structural models were constructed from the high temperature γ-structure, and refined against both SXRD and NPD patterns. Thus, the obtained model led to a good Rietveld refinement, as shown in Figure II.33. The structural model is summarized in Appendix Table A1.9 and the resulting structure is shown in Figure 6. In β-Na4Zn(PO4)2, the distribution of Na, Zn and vacancies is similar as in γ-Na4Zn(PO4)2. The main difference, and reason for the doubling of the lattice parameter, is that 75% of the PO4 groups get ordered (namely P2 linked to O21 and O22) and are no longer described with oxygen being placed on a Wyckoff position with high multiplicity, while 25% of PO4 remain rotationally disordered (P1 and P3), see Appendix Table A1.9. This partial-crystallization of the PO4 tetrahedra, i.e., breaking of the rotational disorder for those tetrahedra, drives the γ to β transition on cooling.

c) Structure of α-Na4Zn(PO4)2

The XRD pattern of α-Na4Zn(PO4)2 presents many additional peaks compared to the patterns of γand β-Na4Zn(PO4)2, indicating again a structural filiation between phases, α-Na4Zn(PO4)2 having a lower symmetry. The SXRD pattern could be indexed with a with unit cell parameters a = 8.99824(4) Å, b = 21.05794 (10) Å, c = 9.08501(4) Å, and β = 110.51598(28)°. The lattice vectors of the α-phase can be deduced from those of the β one using the following relation:

(𝐚 𝛼 , 𝐛 𝛼 , 𝐜 𝛼 ) = ( 1/4 1 1/4 1/4 -1 1/4 1/2 0 -1/2 ) (𝐚 𝛽 ′ , 𝐛 𝛽 ′ , 𝐜 𝛽 ′ ) + ( 1 4 , 0, 1 4 )
The reflections conditions in the selected area electron diffraction (SAED) patterns of the α-Na4Zn(PO4)2 suggest that the probable space group is monoclinic P21/n. Solving the structure from powder diffraction with such a large unit cell is a challenge; however, one can suggest that the α-phase presents a complete PO4 ordering without any rotational disorder, and Zn/Na/Vacancy ordering. Further investigations using single crystals will be necessary to elucidate this unambiguously.

II.4.3 Stabilization of the β-polymorph.

Since the high-temperature γ-Na4Zn(PO4)2 presents a disordered crystal structure, we tried to stabilize the polymorph at RT by deliberately creating disorder in the system. For this purpose we have partially substituted Zn 2+ with hypervalent M 3+ (M = In, Ga, and Al) to create Li + vacancies, so that the inherently disordered phase is stabilized at RT. We found that M 3+ = In 3+ does not form solid solution with Na4Zn(PO4)2, see Appendix Figure A1.8.

However, M 3+ = Al 3+ , and Ga 3+ , was found to be successfully incorporated in the crystal structure. We prepared a series of compounds with nominal composition of Na4

II.4.4 Ionic Conductivity transition in pristine Na4Zn(PO4)2

The variations of ionic conductivities of the Na4Zn(PO4)2 samples synthesized by quenching or cooling slowly are summarized in Figure II.40. The linear variation for the ln(T) with inverse temperature confirms the thermally activated Arrhenius type behavior σ(T) = A/T•exp(-Ea/kBT), where kB is the Boltzmann constant, A is a pre-exponential factor and T is temperature. The thermodynamically stable α-Na4Zn(PO4)2 presents a poor ionic conductor at RT (extrapolated σtot ≈ 1 × 10 -14 S•cm -1 at 25 °C with Ea ≈ 0.92 eV). However, the quenched β-Na4Zn(PO4)2 sample, at low temperature, shows higher conductivity (σtot ≈ 3 × 10 -8 S•cm -1 at 150 °C) than the slow-cooled α-Na4Zn(PO4)2 sample (σtot ≈ 9 × 10 -10 S•cm -1 at 150 °C). Since at RT the quenched sample is in β-form, which is disordered, it shows higher conductivity than the slow-cooled α-form. Interestingly, at higher temperature, when T ≥ 250 °C, the conductivity of β-Na4Zn(PO4)2 decreases and traces the conductivity of the α-Na4Zn(PO4)2. But this behaviour is consistent with the structural analysis as portrayed before (Figure II.31). At slight higher temperature, the quenched sample transforms again to the α-form (Figure II.31c), hence showing same conductivity.

Upon increasing the temperature further, when T ≥ 375 °C, both the samples undergo a transition in ionic conductivity at ~375 °C, driven by the order-disorder transition (α → γ) in the compound, showing much higher ionic conductivity. Since the disordered phase favors fast ionic conduction, it becomes interesting to explore the transport properties of the RT-stabilized γ/β phases. 4. Note that Al-substitution over Ga exhibits similar transport properties and is shown in Appendix. For all the Ga-substituted samples, high ionic conductivities, comparable to the γ-Na4Zn(PO4)2, have been obtained even at RT. However, the Gasubstituted samples (xGa = 0.05 -0. This goes in line with the common observation in ionic conductors that the conductivity increases with increasing Na/Li vacancy leading to disorder. To recall, this also observed for the Li4Zn(PO4)2 (Section II.3) also where, despite pertaining to the α-phase, the α-Li4-xZn1-xGax(PO4)2 compositions (i.e., x = 0.05 -0.2) show higher conductivity with higher Ga-content, due to the increased disorder for Li-vacancy.

To summarize this section, a new compound Na4Zn(PO4)2 has been designed and synthesized, which shows a phase transition to disordered γ-phase at ~390 °C. Ga 3+ , and Al 3+ substitution for Zn 2+ successfully stabilizes the γ-phase at RT. As a consequence, Li-conductivity similar to the HT-stable γ-Na4Zn(PO4)2 was possible to obtain in the β-Na4-xZn1-xMx(PO4)2 phases (M = Ga, Al) at RT.

II.5 Chapter Conclusions.

We thoroughly studied polymorphism in Li6Zn(P2O7)2, and Li(/Na)4Zn(PO4)2, and explored its impact on ionic conductivity. The Li6Zn(P2O7)2 undergoes multiple phase transitions with temperature, elucidating the existence of six polymorphs. Full structural resolution for the αand ζpolymorphs has been deduced enabling us to account the variation of the ionic conductivity in the compound. The ζpolymorph was found to crystallize in a monoclinic unit cell (C2/c S.G.) as the α-polymorph but with a four-time smaller unit cell having a highly disordered and cation-mixed Li/Zn sublattice. We performed BVEL calculations on such newly resolved structures and found that the Li + ionic conduction in the polymorph is two dimensional in contrast to three-dimensional for the -polymorph owing to the Li/Zn mixed sites which promote interlayer ion hoping. Such predictions were confirmed by AC-impedance measurements which besides showing two jumps in ionic conductivity with increasing temperature, has confirmed an high activation barrier for Li + conduction for the α than the -polymorph (0.9 eV vs. 0.35 eV). Altogether, these results suggest that not only disorder, but also cation-mixing can lead to new conduction pathways, which can increase the ionic conductivity.

In Li4Zn(PO4)2 and Na4Zn(PO4)2, we have explored the rich crystal chemistry in great details and described how we could improve its ionic conductivity. We showed that both compounds undergo an order-disorder phase transition at HT, accompanied by a jump in its ionic conductivity. We solved the crystal structure of the high temperature β/β'-polymorphs, which appears to be derived from the low temperature α-structure by preserving the PO4 3-backbone, but presenting a disordered Li(Na)/Zn sublattice. The γ-polymorph was found to have a higher conductivity than the α-phase. We have demonstrated the feasibility to stabilize the γ-polymorph at RT either by ball-milling the α phase, or by partial substitution of Zn 2+ by Ga 3+ in Li4Zn(PO4)2 (also by Al 3+ in case of Na4Zn(PO4)2). With 0.5 mol pfu Ga, we could successfully stabilize the Li3.5Zn0.5Ga0.5(PO4)2 as γ-polymorph at RT, with a crystal structure pertaining to the β-Li4Zn(PO4)2 phase, but with the Zn and Ga ions fully mixed in a disordered Li-sublattice. The β-Li4-xZn1-xGax(PO4)2 samples exhibit ionic conductivities up to ~2 -5 x10 -13 S•cm -1 at RT, by several orders of magnitudes higher than α-Li4Zn(PO4)2. This could be explained by the higher crystal symmetry of the disordered β-Li4Zn(PO4)2 that enables to switch the conduction mechanism from 2-D to 3-D. Similarly, the β-Na4-xZn1-xMx(PO4)2, M = Ga and Al, samples exhibit ionic conductivities up to ~2.7 x10 -9 S•cm -1 at RT, by several orders of magnitudes higher than α-Na4Zn(PO4)2. Despite such a rich chemistry of Li6Zn(P2O7)2, Li4Zn(PO4)2/Na4Zn(PO4)2 and its Ga/Al substitutes, the ionic conductivity is still far from the targeted value (10 -3 S•cm -1 ). However, through this work we illustrate the role of disorder to ionic conduction and demonstrates how to use chemical substitutions to trigger RT stabilization of disordered high-temperature polymorphs (Figure II.42). Next, to fabricate practical all-solid-state-batteries (ASSB), we decided to implement sulfide-based solid electrolytes, e.g., β-Li3PS4, Li2S-P2S5 (LPS) glass, etc., which showed high ionic conductivity and easiness to be implemented in ASSB, as discussed in Chapter I. 154,177,233 In order to study interfacial compatibility in all-sulfide ASSBs 234 , we have revisited some sulfide cathode materials, and designed highly reversible Li-rich Li1.33-2y/3Ti0.67-y/3FeyS2 electrode for LIBs, that offer high energy density capitalizing the benefit of anionic redox, [START_REF] Assat | Fundamental understanding and practical challenges of anionic redox activity in Li-ion batteries[END_REF] and this will be the topic of the next chapter.

Chapter III

Sulfide-based cathode materials: Anionic redox in Li-rich sulfides #

______________________

III.1 Chapter Introduction

Layered sulfide TiS2, was one of the very first intercalation host that was considered to develop rechargeable Li-metal batteries (Figure III.1a). 235 However, with the emergence of Li-ion batteries (LIBs), layered oxide LiCoO2 (Figure III.1b) was proposed as positive electrode because of its high energy density, and today's commercial batteries use LiCoO2 (LCO), and its derivatives such as LiNixMnyCo1-x-yO2 (NMC), etc., as positive electrode, as discussed in Section I.1.3. [START_REF] Mizushima | LixCoO2 (0<x⩽1): A new cathode material for batteries of high energy density[END_REF]236 For LiCoO2, only ~0. 5 Li can be deintercalated reversibly. Even though there are ongoing efforts to push the limit by doping, [START_REF] Zhang | Trace doping of multiple elements enables stable battery cycling of LiCoO2 at 4.6 V[END_REF]237 deintercalation of >0.5 Li leads to a harmful phase transformation. [238][239][240] Nevertheless, Tarascon and colleagues (1999) showed the possibility of charge transfer from O 2p in the fullydelithiated "CoO2", which gave hints about anionic activity (Figure III.1b). 241,242 Then Li-rich NMC's (Li[LizNixMnyCo1-x-y-zO2]) was designed by partially substituting the transition metals with Li, where >1 Li can be extracted, thus leading to high capacity (Figure III.1c), by capitalizing the benefit of anionic redox (i.e., electrochemical participation of O anions). [START_REF] Lu | Layered Cathode Materials Li[END_REF][START_REF] Lu | Understanding the Anomalous Capacity of Li/Li[END_REF]34,[START_REF] Thackeray | Li2MnO3-stabilized LiMO2 (M = Mn, Ni, Co) electrodes for lithiumion batteries[END_REF] However, this comes with several limitations, as discussed before. The Li (de)intercalation in TiS2 is charge-compensated mostly by cationic Ti 3+/4+ redox process, with possibly very little contribution of S due to its high covalency, as pointed out by few reports. [243][244][245] Furthermore, possibility of oxidation of S upon Li-intercalation in LiTi0.5Fe0.5S2 was discussed by Tarascon et al (1983). 246 Interestingly, ligand to metal charge-transfer in layered disulfides, can be triggered upon lowering the metal d level, by tuning the transition metals (TM). Lowering the d-level energy on moving to the right-hand side of the periodic table can lead to holes on S 3s states by electron transfer to TM d-states, e.g., in FeS2. 242,[247][248][249] Therefore, unlike oxides, stable ligand-hole chemistry was well-known in sulfides, e.g., TiS3 (Ti 4+ S 2-(S2) 2-), Fe 2+ (S2) 2-, and V 4+ (S2 2-)2 etc., since the pioneering works by Rouxel et al. 247,248,250 Quasi-layered dichalcogenide FeS2, and trichalcogenides MQ3 (M = Ti, Zr, Hf, Nb, Ta; Q = S, Se), especially TiS3, TaS3 and NbSe3, have been subject of immense interest owing to the possibility of multielectron redox reaction, in early works, which presents the classical examples of anionic redox chemistry. [START_REF] Whittingham | Chemistry of intercalation compounds: Metal guests in chalcogenide hosts[END_REF]242,[251][252][253][254] In these electrode materials, S (or Se) fully or partially exists as dimerized S --S -pairs, and undergo breaking of S-S bonds to form standard S 2-states in discharge. Thus, the study of sulfide materials is quite important while discussing anionic redox and we decided to explore the practicability anionic redox in sulfides. In this chapter, we will present the investigation of anionic redox, in 'pseudo-layered' TiS3 first, and then, in layered Li-rich sulfides.

III.2 Pseudo-layered TiS 3

We first explore anionic S2 2-/S 2-redox and its structural consequence in pseudo-layered TiS3. The anionic redox in TiS3 was known since 1980s, even much before the oxygen redox. [START_REF] Whittingham | Chemistry of intercalation compounds: Metal guests in chalcogenide hosts[END_REF]242,255,256 We decided to revisit the anionic redox in TiS3 in order to reinvestigate in detail its structural aspects and assess the reversibility of the anionic redox together with the practical issues, such as voltage fade, kinetics, and voltage hysteresis, etc. Interestingly, TiS3 contains two kind of S atoms; bridging (S1) and disulfide pair (S2 and S3), see Figure III.2c, with the presence of the S-S dimer (S2 2-) with a bond length of 2.04 Å (S2-S3, marked in red). TiS3 remains the origin of great interest for various applications dealing with energy storage, photocatalytic, electronic, optoelectronic, etc. [START_REF] Whittingham | Chemistry of intercalation compounds: Metal guests in chalcogenide hosts[END_REF][259][260][261][262][263][264][265][266][267][268] However, despite immense attention, the crystal structure was taken for granted, until Guilmeau et al revisited the structure in 2014 and proposed slight different arrangement of the S atoms (Figure III.2d). 258 Thus, owing to these contradictory findings, we decided to revisit the structure and electrochemical properties of TiS3.

III.2.1 Synthesis of TiS3

TiS3 was synthesized as powder, by solid-state synthesis route. Required amount of TiS2 or Ti metal powder was reacted with S powder (20 wt% extra) in a vacuum-sealed quartz tube at 550 °C for one week (conditions were optimized). Thus shiny black colored TiS3 lumps were obtained (Figure III.3a) which was hand-grinded to powder for further use. For precaution, the TiS3 was always handled inside an Ar-filled glovebox. In parallel, we also tried to synthesize single crystals of TiS3 by vapor transport method. Ti powder and S powder (50 wt% excess) was mixed and vacuum sealed in quartz ampule (10 -13 cm long). The ampule were heated in furnaces at 600 °C, with a temperature gradient of ~50 °C to the cooler end and was kept for 2 weeks. Occasionally, up to 50 wt% of iodine was also added as additional vapor transporter, but no noticeable differences were observed. Ribbon shaped TiS3 crystals (up to ~1 cm long, with up to 1 mm of breadth and submicron thickness) could be obtained. A picture of the crystals is shown in Figure III.3b. However, closer look at the crystals by optical microscopy reveals that the crystals always consist a bunch of single crystals stick together, or with local exfoliations.

III.2.2 Structural characterizations: a different variant of TiS3

The purity of the as-synthesized TiS3 powder was confirmed by XRD. Regardless the precursor, whether TiS2 or Ti powder, was used, the XRD patterns of the synthesized TiS3 were found to be always similar. We collected scanning electron microscopy (SEM) images of the TiS3 powder (Figure III.4). Interestingly we found that the powder is comprised of micrometer sized ribbons.

III.2.3 Electrochemistry of TiS3

Then the TiS3 powder was characterized for electrochemical performance. They were mixed with 20 wt% conductive Csp and hand-grinded for 5 mins prior to cycling (SEM images were shown in Figure III.4). The powders were galvanostatically cycled in Li-half cells with LP30 electrolyte between 1.7 V and 3 V at a rate of C/20 and the voltage profiles are summarized in Figure III.9.

The TiS3 was discharged first (red curve in Figure III.9a) and up to 2 Li could be inserted till a cut off discharge potential of 1.75 V. At the beginning of first discharge we always observe a deep voltage drop (peak at ~2 V in dQ/dV curve in Figure III.9b), which is usually associated to nucleation of a second phase. [START_REF] Whittingham | Chemistry of intercalation compounds: Metal guests in chalcogenide hosts[END_REF]255,[269][270][271] The discharge plateau observed afterwards corresponds to the intense peak at 1.97 V in the dQ/dV curve. During the subsequent charge, a small nucleation overpotential followed by a charging plateau is obtained (peaking at 2.57 V in dQ/dV curve). However, the material does not go back to the initial composition (i.e., Li0TiS3), since ~0.2 Li remains in the charged phase (with an approximate composition of Li0.2TiS3). In the 2nd discharge, the nucleation overpotential is not observed, and the intercalation takes place at slightly higher potential than the 1st discharge. Overall a high polarization of >0.5 V is obtained.

III.2.4 Evolution of crystal structure during discharge-charge

To understand the structural change upon Li-(de)intercalation, we performed operando XRD during charge-discharge of the TiS3 powder and the results are summarized in Figure III.10a. During the first discharge, it goes through a biphasic process, as indicated by the appearance of a 2nd phase. At mid-discharge, the Li1TiS3 composition (shown in blue color) is a mixture of phases, with second phase solely remaining at the end of discharge (Li2TiS3, red pattern). Upon charging back, the evolution of the XRD patterns follows the reverse pathway. Note that, the full charged phase Li~0.2TiS3 shows a XRD pattern (shown in green) similar to the TiS3, but with much broader peaks. This indicates that, upon cycling, TiS3 loses crystallinity largely.

For more accurate characterization, we have also collected ex situ SXRD patterns (at 11-BM) of the samples at various SoCs (state of charges). The patterns are gathered in Figure III.10b. The midcharged LiTiS3 is clearly a mixture of the TiS3 and Li2TiS3, in agreement with the previously described the operando XRD. Mention must be made, that the SXRD pattern of the fully-discharged Li2TiS3 is different from the layered Li-rich Li2TiS3 (will be discussed in Section III.3). 272 This indicates that the crystal structure of the Li2TiS3 prepared electrochemically (by Li intercalating in TiS3) is different from the layered LiMS2 (M= transition metal, space group C2/m) family.

III.2.5 Tracking oxidation states during Li (de)intercalation

Next, we examined the charge compensation mechanism during the Li (de)intercalation in TiS3, which will also be used as a reference later. The ionic charge of TiS3 is expected to be formulated as Ti 4+ (S2) 2-S 2-. The oxidation changes in S and in Ti were probed at various SoCs. Figure III.15a shows the S 2p core XPS spectra collected on the cycled samples. C The pristine TiS3 shows S 2p core XPS peak containing both S 2-and S2 2-(dimer) at a ratio 32:68 %, consistent with the crystal structure drawn in Figure III.2, where the isolated S1 atom (bridging) is characterized by a larger negative charge than the S atoms in S2-S3 pair. 255,268,273,274 At mid-discharge (Li1TiS3), the intensity of the S2 2-peak decreases (56:44 %). The fully-discharged Li2TiS3 shows only the peak for S 2-(and similar to TiS2), as the S2 2- gets reduced to S 2-upon Li-intercalation. On charging back, in Li0.2TiS3, the spectra is almost restored back as the pristine. Ti 2p core XPS spectra were also collected on the same samples and they are shown in Figure III.15b. The spectra for the pristine TiS3 and mid-discharged LiTiS3 are positioned at same binding energies, indicating their formal oxidation state to be +4. 255 However, drastic change happens for the full-discharged Li2TiS3, where the peak moves to higher binding energy. Following the existing literature, this shift could be explained by the increase of ionicity of Ti-environment that arises from reduction of all S2 2-dimers to S 2-, while Ti stays as Ti 4+ . 255,256,273,275,276 However, we believe that the possibility of reduction of Ti 4+ cannot be discarded and we will continue working on this further. Furthermore, we are exploring the possibility of change in the Ti coordination from trigonal prismatic site to the more stable octahedron site. On charging back, the positon of the Ti XPS peak is restored back, despite the irreversibility of electrochemistry (Li~0.2TiS3).

To further probe the changes in the oxidation states, we performed Electron energy loss spectroscopy (EELS) in combination with X-ray absorption spectroscopy (XAS) studies. D The S L-edge EELS spectra of the pristine TiS3 and fully-discharged Li2TiS3 powders (Figure III.16a) comprise a weak pre-edge and intense broad edge feature. The spectra correspond to a series of transitions from the S 2p core level to unoccupied S 3s σ*, S 3d, etc. states, which are further split because of transitions from the S 2p3/2 (L3 edge) and S 2p1/2 (L2 edge) states due to spin-orbit coupling. 277 The S-L edge of the TiS3 and Li2TiS3 shows similar spectra, except the spectra for TiS3 contains more broad peaks. Worth noting that for the TiS3, the broad peaks may be associated to possibly presence of different coordinations of S atoms. spectra of TiS3 268,276 and has been associated to the Ti 4+ (in TiS8 coordination, as portrayed in the structural characterizations). Upon Li-intercalation (Li2TiS3), the spectra shifts towards higher energy by ~0.98 eV (similar to that of TiS2, 243 suggesting an octahedrally-coordinated of Ti 4+ ). Note that the observation in XPS (Figure III.15b) supports this shift.

To get further insights of S states, X-ray absorption spectroscopy (XAS) studies were performed at the 4-ID-C beamline (Argonne National Laboratory). E XANES spectra were collected in both total electron yield (TEY, probe depth is ~ 10 nm) and total florescent yield (TFY, probe depth is ~ 100 nm) modes. Both TEY and TFY spectra produces similar spectra and hence only TFY spectra (more bulk sensitive) has solely reported here, with the TEY spectra been described in the Appendix Figure A2.1. See the Appendix for details of the methods of XANES acquisition. The TFY-XANES spectra for pristine TiS3 consists a broad edge centered at ~2472 eV (denoted as B and an intense pre-edge at ~2471.5 eV (denoted as A). 243,268 The B edge (at ~2479.8 eV) represents E XAS studies were performed by Haifeng Li and Prof Jordi Cabana, University of Illinois at Chicago (USA). multiple photoelectron transitions enlisting hybridized sulfur 1s  3pσ* transition with the latter state being hybridized with metal 4s and 4p states. The pre-edge A, in contrast, depends on the degree of hybridization of the S unoccupied antibonding states with the empty transition metal 3d states. 250 The shape and position of the spectra are consistent with reported spectra of the TiS3. 256,268 Remarkable variations are observed upon Li intercalation, specifically decrease in the intensity of pre-edge A, and shift of edge B towards lower energy. In LiTiS3, the edge shifts to ~2477.3 eV, and the pre-edge becomes weak, whereas in Li2TiS3, the pre-edge further weakens. This is consistent with the reduction of S2 2-to S 2-, as similar shift in the edge B is observed during lithiation of VS4. 250 The S2 2-dimers having a lower energy 1s orbital compared to S 2-, hence leading to a higher energy edge peak. 250,268 Upon charging back, in Li0.2TiS3, the spectra is almost restored back. Overall, these studies unambiguously point towards the S 2-/S2 2-redox.

III.2.6 Practicability of the anionic redox

Having explored the strong involvement of anionic redox in TiS3 leading to an energy density as high as >700 Wh•kg -1 , the next legitimate question pertains to its practicability. To evaluate, first we cycled TiS3 powder over long cycling (Figure III.18), which shows that the capacity falls very rapidly. In summary, the rich crystallochemistry and electrochemistry of TiS3 has been explored in details. The structural change upon S-S bond breaking and reformation accompanied by Li (de)insertion is yet to be understood completely. Turning back to electrochemistry, utilization of the anionic S2 2-/S 2- redox lead to attractive capacity of ~370 mAh•g -1 , but to the expense of large hysteresis, and rapid capacity fading. Hence, to circumvent the issue, we decided to explore the feasibility of designing Lirich layered sulfides, alike the Li-rich oxides, which do not trigger complete S-S dimerization.

III.3 Li-rich layered sulfides

In layered family, Li-rich NMC oxides were derived from the layered Li1.33Mn0.67O2 (commonly written as Li2MnO3) with their anionic redox activity being a function of the competition between U (dd coulomb interaction) and Δ (charge transfer) terms. 279 Preparing a similar Li-rich Mn-based layered sulfide is not possible, simply because the S 3p band is situated much closer to the Li/Li + reference than the O 2p band, leaving the Mn 3+/4+ redox band too low-lying, as already discussed in Chapter 1 (recall Figure I.7a,b). Hence, to design an analogous layered sulfide Li1.33M0.67S2, (i.e., Li2MS3) an appropriate transition metal M needs to be chosen first. Amongst 3d metals, M = Ti 4+ presents the best choice for sulfides, because the Ti 3+/4+ redox band is located above the S 3p band. The Li2Ti 4+ (S 2-)3 offers a high theoretical capacity of 339 mAh•g -1 (considering removal of all Li's). Therefore, we decided to study anionic redox in Li2TiS3-based Li-rich layered sulfides, which will be described next.

III.3.1 Li2TiS3

III.3.1.1 Structural Characterizations

Li2TiS3 was prepared by reacting Li2S, and TiS2 in stoichiometric amounts in vacuum-sealed quartz tubes at 750 °C (see the Appendix for details). The SXRD pattern of the as-synthesized Li2TiS3

III.3.1.2 Electrochemical activity

The electrochemical performance of the Li2TiS3 sample was tested in Li-half cell between 1. The Li2TiS3 shows very poor electrochemical activity, since only ~0.1 Li could be extracted reversibly. This is in accordance with some recent literature also. 280,281 Attempts to enhance the activity by either ball milling the sample or by adding larger amounts of carbon additive were unsuccessful. Note that Li2TiS3 has a sufficiently high electronic conductivity (~2.7•10 -5 S•cm -1 , Appendix Figure A2.2), and ionic conductivity (~5.5•10 -4 S•cm -1 , Appendix Figure A2.3). This leads us to conclude that such a non-activity is intrinsic to the phase. Interestingly we noticed, that the voltage profile of Li2TiS3 changes in the 2nd cycle (Figure III.21a) and interestingly the 2nd charge occurs at a lower potential than the 1st one (Figure III.21b). Hence, the 1st cycle act as an 'activation cycle', alike the Li-rich NMC. This phenomenon is observed for all Li2TiS3-based compounds, as we will see in the subsequent sections.

Nevertheless, the pristine Li2TiS3 did not show any substantial electrochemical activity, which most likely nested in the fact that the Ti 3+/4+ redox band is empty (3d 0 ) and is situated far above the S 3p band, hence unable to stabilize oxidized sulfur (Figure I.7b). Hence, we move now for suitable partial substitution for Ti 4+ . Here onwards, the Li2TiS3 will be written as Li1.33Ti0.67S2, to be consistent with the layered LiMO2/LiMS2 formalism.

III.3.2 Fe 2+ -substituted Li2TiS3

In order to initiate electrochemical activity in Li1.33Ti0.67S2, suitable transition metal substituent M needs to be chosen. Since Ti 4+ has 3d 0 electronic configuration, Li1.33Ti0.66S2 (i.e., Li2TiS3) is apparently electrochemically inactive. This situation reminds that of the Li1.33Mn0.67O2 phase (poor electrochemical performance without nano-sizing) that required partial substitution with Ni 2+ to instigate electrochemical activity (Figure I.7a). [START_REF] Lu | Layered Cathode Materials Li[END_REF][START_REF] Lu | Understanding the Anomalous Capacity of Li/Li[END_REF]34 A first hint to address this problem in sulfides consists, as discussed by Li et al.,in using Co 2+ as a substituent to initiate reversible anionic redox. Pursuing their idea, the authors succeeded in preparing Li1.2Ti0.6Co0.2S2 showing anionic redox activity. 281 Following similar substitution strategy, another approach to overcome this unfavorable band situation in Li2TiS3 is to prepare Ti 3+ -doped Li1.33-y/3Ti 4+ 0.67-2y/3Ti 3+ yS2 so that the partially reduced Ti 3+ (3d 1 ) leads to a halffilled Ti 3+/4+ redox couple, as shown by Flamary-Mespoulie et al (2016). 272 Other successful strategies to electrochemically activate the Li2TiS3 are either preparing Li1.33Ti 4+ 0.67S2 having disordered rock-salt structures, or triggering anionic redox activity via antisite occupation as shown for NaCr 3+ S2 (leading to ~0.5 Na removal despite Cr 3+/4+ redox level is buried inside the S 3p band). [START_REF] Shadike | Antisite occupation induced single anionic redox chemistry and structural stabilization of layered sodium chromium sulfide[END_REF]282,283 We herein chosen Fe as a substituent to demonstrate the feasibility to activate the anionic redox activity in Li-rich layered Li1.33-2y/3Ti 4+ 0.67-y/3Fe 2+ yS2. This situation is favourable for reversible sulfur redox, since the Fe 2+/3+ redox couple with available electrons (3d [START_REF] Painter | Oil and Geopolitics: The Oil Crises of the 1970s and the Cold War[END_REF] ) is expected to be pinned at the top of the S 3p band (Figure III.22). 24 We prepared a series of materials of nominal composition Li1.33-2y/3Ti0.67-y/3FeyS2, with y = 0 -0.5, whose structural and electrochemical properties will be described below.

III.3.2.1 Structural Characterizations

The Li1.33-2y/3Ti0.67-y/3FeyS2 compositions were prepared by reacting Li2S, TiS2 and FeS in stoichiometric amounts in vacuum-sealed quartz tubes at 750 °C (see the Appendix Section A2.1 for

O 2p

Mn 3+/4+ Li +

S 3p

Ti 3+/4+ Therefore, the Fe-containing phases could be fitted in a hexagonal R3 ̅ m description analogous to the well-known Li-rich layered Li1+yM1-yO2 phases, with the obtained lattice parameters shown in Figure III.23c. With increasing Fe content, we observe a monotonic increase in the c parameter that is accompanied by a decrease in the a parameter, such that the overall effect is a monotonic decrease in unit cell volume (V).

Since among the Li1.33-2y/3Ti0.67-y/3FeyS2 series, the compound with y = 0.3 will be the center of interest in this study, we have characterized its structure in detail. The Li1.13Ti0.57Fe0.3S2 (y = 0.3) sample consists of submicron sized (5 -30 µm)

III.3.2.2 Electrochemical Properties

The electrochemical performances of the Li1.33-2y/3Ti0.67-y/3FeyS2 samples (mixed and handground with 20 wt% Csp) were tested in Li-half cells (in Swagelok type cells) between 1.8 V and 3 V at a rate of C/20 and the voltage profiles are summarized in Figure III.26. The Fe 2+ -containing phases are electrochemically active, as shown in Figure III.26, which furthermore shows that, with increasing Fe-content, more amount of Li's can be extracted from the compositions up to y = 0.3. With y > 0.3, the extractable amount of Li's does not increase anymore, because of the available amount of Li (i.e., 1.33-2y/3) decreases in the compositions. This improvement is most likely nested in the gradual amelioration of the band positioning with introduction of Fe content that triggers reversible anionic redox activity. However, despite the gradual electrochemical activation in the series, complete delithiation was not observed, as ~0.1 Li remains after full-charge of the compositions with y ≥ 0.3.

The voltage profiles shows a specificity that is nested in the second cycle's charge-discharge which occurs at a lower potential than the first cycle. Note also the appearance of a short sloppy voltage profile (marked by 'S' in Figure III.26) at the early stage of charge, and most likely related to Fe redox activity. This follows by a long plateau-like (marked by 'P') activity on further oxidation, and presumably from the S 2-/S2 n-redox.

III.3.2.3 Cycling behaviour over larger voltage range

From the electrochemical behaviour described above (within 1.8 -3 V), it is clear that ~0.1 Li remained electrochemically inaccessible, irrespective of the Fe-content. Hence, the legitimate question arises, is it possible to delithiate further ?

To verify this we decided to cycle the Li1.13Ti0.57Fe0.3S2 with an increased the upper voltage cutoff (> 4 V). The resulting voltage profile is shown in Figure III.29a, which shows that it does not lead to success, rather cell failure was observed above 3.8 V. Furthermore, it was confirmed by operando UV spectra that electrolyte decompositions occur at high voltage (as indicated by appearance of new absorption peak at ~250-300 nm wavelength) and furthermore by operando XRD that shows sample decomposition to unidentified phases (details in the Appendix Figure A2.5 and Appendix Figure A2.6), which is responsible for the short circuit like behaviour. Worth recalling, that this situation is alike the Li-rich oxides, where even pushing at higher voltage does not enable fully delithiation prior decomposition, and oxygen release. [START_REF] Perez | Approaching the limits of cationic and anionic electrochemical activity with the Li-rich layered rocksalt Li3IrO4[END_REF]284 Lastly, it remains to verify the electrochemical behaviour of the Li-rich sulfide at low voltage. Therefore, we have cycled the Li1.13Ti0.57Fe0.3S2 vs Li, up to 0.2 V (Figure III.29b). The sample indeed shows a staggering capacity increase, due to uptake of ~3 extra Li in discharge and consequently increase of the energy density up to ~800 mAh•g -1 . This can simply be explained by the conversion reaction type mechanism, 285 where the metal ions (Fe 2+ and Ti 4+ ) are reduced according to the following reaction, Li1.13Ti0.57Fe0.3S2 + 0.6 Li + + 2.28 Li + → 2 Li2S + 0.3 Fe + 0.57 Ti However, a large polarization and rapid capacity loss upon cycling is observed which again highlight the drawbacks of the conversion type redox. [START_REF] Poizot | Rationalization of the Low-Potential Reactivity of 3d-Metal-Based Inorganic Compounds toward Li[END_REF] Lastly, to assess the practical application of this cathode material, we tried to cycle the Li1.13Ti0.57Fe0.3S2 sample at 55 °C, which could not succeed (details can be found in Appendix Figure A2.7).

In short, the Li-rich sulfide shows excellent reversibility of Li-(de)intercalation process only within voltage range of 1.8 -3 V. Hence we will concentrate hereafter on the cycling performance of this compound within this range, starting with exploration of the structural aspects.

III.3.2.4 Structural integrity during charge-discharge

We first explore the effect Li uptake-removal process on the crystal structure of the Li1.13Ti0.57Fe0.3S2, by operando XRD measurements. XRD patterns were collected for every change in lithium stoichiometry of ~0.01 and summarized in Figure III.30. Upon charge, we observed a slight initial shift of the main peaks, indicating solid-solution behaviour. Over this regime, the unit cell volume decreases as a consequence of an anisotropic variation of the a and c lattice parameters (Figure III.30). Then, as the voltage reaches the first plateau, there is a gradual change in the intensity of the peaks with some peaks disappearing at the expense of new ones, which sharpen to give a well-defined XRD powder pattern at the end of the full charge, hence suggesting a two-phase de-intercalation process. This new phase with approximate composition ~Li0.1Ti0.57Fe0.3S2, whose structure remains as layered one as described in detail in the next paragraph, has a lower unit cell volume (by ~12.2%) than the pristine one.

On discharge, the compound undergoes first a solid-solution process for which both the a and c lattice parameters strongly increase (Figure III.30). Then, a biphasic process is observed with the growth of a phase with increased volume, followed again by solid-solution behaviour. We therefore have, for the first cycle, charge and discharge processes that proceed through different structural paths, even though the pattern returns close to the one for the pristine phase. This path-difference can be clearly observed from the evolution of the unit cell volume and of the a and c lattice parameters (Figure III.30). We further continued the operando XRD in the 2nd cycle (Figure III.31). Here also, the structural evolution goes through a biphasic process, as indicated by the evolution of the XRD patterns. However, the lattice parameters evolves via more symmetric, although not perfect, pathways on charge and discharge, compared to the 1st cycle. Overall, throughout cycling the phase remains crystalline with well-preserved long-range layered crystal structure.

III.3.2.5 Crystal structures of charged-discharged phases

Next, we performed an in-depth exploration of the crystal structures at different states of charge (pristine, fully-charged and fully-discharged after first cycle) using SXRD. In agreement with the lab XRD data, the patterns can be indexed in the R3 ̅ m space group. The pristine Li1.13Ti0.57Fe0.3S2 (V = 196.384(2) Å 3 ) presents an average Ti-S bond-length of 2.4792(3) Å, with average S-S distances of ~3.505 Å (structural model in Appendix Table A2.3). The Rietveld refinement of the SXRD pattern for the fully-charged Li0.13Ti0.57Fe0.3S2 phase is shown in Figure 4b. The corresponding crystal structure indicates a much smaller unit cell (V = 172.338(6) Å 3 ) (see the Appendix Table A2.4 for the structural model). This corresponds to ~12% contraction of the unit cell upon delithiation. Moreover, the average Ti-S bond length was found to shrink to 2.3635(1) Å, leading to decreased average S-S distances of ~3.344 Å. In parallel, we conducted transmission electron microscopy (TEM) studies of the fully charged phase to verify whether cationic migration 286 occurs. The HAADF-STEM image shows only the [Li0.1Ti0.57Fe0.3]S2 layers and thus discards the possibility of transition-metal migration to the interlayer sites (Figure III.32b). The HAADF intensity profile (inset, Figure III.32) clearly demonstrates no scattering density between the (Ti/Fe)S2 layers, hence no Fe 3+ cation migration. The selected-area electron diffraction (SAED) of the fully-charged phase was also taken (Figure III.32c) which could again be successfully indexed with an R3 ̅ m unit cell, in accordance with the SXRD pattern.

Finally, to assess the reversibility of Li-(de)intercalation, SXRD pattern of the fully-discharged sample was refined also. The Rietveld refinement is shown in Figure III.32d, and the structural parameters are gathered in the Appendix Table A2.5. The fully discharged phase is indeed very similar to the pristine phase (V= 196.896(2) Å 3 ) with an average Ti-S bond length of 2.4825(3) Å and average S-S distance of 3.51(2) Å, confirming the reversibility.

At this point it is interesting to verify whether S-dimerization occurs accompanying the anionic S 2-/S2 n-redox. We plotted the variation of bond lengths upon Li-(de)intercalation in Figure III.33. As we observe, despite the decrease of S-S bond lengths in fully-charged phase, such S-S distances are far larger than the (S-S) 2-bond length reported for TiS3 (2.04 Å) and FeS2 (2.16 Å), hence implying the onset of solely a partial S-S dimerization upon removal of Li from Li1.13Ti0.57Fe0.3S2. Although there is no doubt about the oxidation of S (which will be proved further via local-structure sensitive techniques in the later sections), it is evident that its extent is not enough to form S -that would condense as (S2) 2-. A simple charge balance (assuming one electron transfer in the Fe 2+/3+ process) suggests that the value of n in our compound is ~S1.65-, hence much smaller than S -. This may explain that the oxidized sulfur in our Li-rich layered sulfide does not condense as in TiS3 or pyrites FeS2. Another important distinction of these layered materials compared to pyrites is rooted in the nature of S-S bond, which are reported between two neighboring FeS6 octahedra in FeS2. 287 Similar S-S bonds or intra-octahedral S-S bond in Li1.13Ti0.57Fe0.3S2 would need significant rearrangement of the crystal structure of the compounds studied here. This is not overserved experimentally as the layered structure persists.

III.3.2.6 Charge compensation mechanism III.3.2.6.1 Cationic redox process

To identify the underlying redox mechanism which accounts for the high capacity obtained from the Li-rich sulfides, we performed operando Mössbauer spectroscopy F using an in-house designed electrochemical cell 288 , to probe the cationic Fe 2+/3+ redox process. Spectra were collected while charging and discharging a Li1.13Ti0.57Fe0.3S2/Li cell at a rate of C/70 (Figure III.34a). The Mössbauer spectra for pristine Li1.13Ti0.57Fe0.3S2, can neatly be fitted with four doublets (shown in the Appendix Table A2.6 L2,3-edge XAS data upon oxidation are consistent with previous literature of Fe 2+ and Fe 3+ sulfide species. 293 After discharge, the Fe L2,3 edge converts back to its initial phase.

Next we turn to the Ti atoms. To probe its redox, Ti L2,3-edge XANES spectra were also recorded for the aforementioned samples (Figure III.37c and Appendix Figure A2.12). The edge consists of four peaks attributed to the transitions 2p3/2 → 3d(t2g), 2p3/2 → 3d(eg) (the L3 edge) and 2p1/2 → 3d(t2g), 2p1/2 → 3d(eg) (the L2 edge). As expected, the spectrum for pristine Li1.13Ti0.57Fe0.3S2 is identical to that of Ti 4+ S2 (see Appendix Figure A2.12), confirming the formal oxidation state of Ti as 4+. 243 Moreover, no changes in the position of L2 and L3 peaks could be observed, irrespective of the sample state of charges (SoCs). Ti L2,3-edge electron energy loss spectroscopy (EELS) spectra were also recorded (shown in Appendix Figure A2.11) for the samples (further including a partially-charged sample) and they exhibit spectra similar to that of XANES. This indicates that the Ti 4+ remains invariant throughout the charge/discharge cycle.

III.3.2.6.2 Anionic redox process

After probing the electrochemical participations of the cations, we move to assess the participations of the S ligand. For this, S K-edge TFY-XANES spectra were collected at various SoCs and they are reported in Figure III.37d. The XANES spectra collected in TEY mode are gathered in the Appendix, which are similar to the TFY spectra. The spectrum presents two main regions. The pre-edge feature, below 2471 eV, arises generally from unoccupied S 3p/TM 3d hybridized states. Their position and intensity depend on their occupancy, the relative contribution of S and TM, and their position with respect to the core level, thus being a general measure of covalence of the TM-S bond. 294 Above this energy, the signals correspond to transitions to higher states, such as S 3p/TM 4s, p or those involving S 4p, and the photoionization of S atoms, involving complete ejection of the core electron to the continuum. Therefore, the position of the absorption threshold is strongly dependent on the effective nuclear charge (Zeff) on S, 294 being a measure of redox transitions at the ligands. Note, the S K-edge XANES spectra of all as-synthesized samples of Li1.33-2y/3Ti0.67-y/3FeyS2, y = 0 -0.5, are similar, confirming identical S-states in all compositions, see Appendix Figure A2.13. -164.5 eV, denoted as S m-) was also observed, which probably originated from some surface deposited species.

Coming back to the Li1.13Ti0.57Fe0.3S2 phase, the TFY-XANES spectra at different state of charges (SoCs) are shown in Figure III.38 (with the TEY spectra in Appendix Figure A2.14). The spectra for pristine Li1.13Ti0.57Fe0.3S2 also consists in a weak pre-edge at ~2465 eV (denoted as A) and a broad edge (denoted as B) centered at ~2472 eV similar to what is observed in amorphous rocksalt Li2TiS3 and hence characteristic of S 2-. 282 For the fully charged state (~Li0.1Ti0.57Fe0.3S2), the spectrum (red curve in Figure III.37d) undergoes remarkable changes with namely an increase in amplitude of the pre-edge A and a ~1.5 eV shift of the B edge towards higher energy, and beers resemblance to that of TiS3 (dashed purple curve). 268 Similarly, the opposite shift in the edge B is observed during re-lithiation, as was reported for TiS3, and VS4. 250,256 The 1.2 eV shift in B is therefore a clear indication of the oxidation of S 2-into S n-, n < 2. Similarly, the increase in intensity of pre-edge A is indicative of the increase in density of unoccupied states just above the Fermi level with an S character, in agreement with the expectation that the redox change is compensated by S 3p/Fe 3d electrons. The shift of the main absorption edge also suggests that this compensation, and the resulting states, have a significant S character. This unambiguously confirms the participation of S in the overall electrochemical activity of Li1.13Ti0.57Fe0.3S2 together with its reversibility since the S K-edge XANES spectra of the fully-discharged and pristine samples nearly superimpose (Figure III.37d).

To confirm the above observations S 2p core X-ray photoelectron spectroscopy (XPS) spectra were also taken and are shown in Figure III.37e. Clear variations of the spectra are observed upon charging-discharging the samples. The S 2p3/2-1/2 XPS core spectrum can be fitted with a single doublet (160.7 -161.9 eV) attributable to S 2-for the pristine material, as in the TiS2 reference. 295 As the sample gets oxidized, the spectrum markedly changes, with namely the appearance of one extra doublet at higher binding energies (161.8 -163.0 eV), which is indicative of the presence of oxidized S n-(n < 2), in light of early previous studies on TiS3 that is well-known to contain partially oxidized sulfur. 255,295 On discharge, the XPS spectrum is almost restored to the pristine-like state, confirming the reversibility of the anionic redox process.

At this point, though the anionic participation is confirmed, question remains whether the S atoms fully dimerize upon oxidation. To check this point, ex situ electron energy loss spectroscopy (EELS) spectra at the S L2,3 edge were collected and shown in Figure III.39. The spectra consist of a weak pre-edge and an intense broad edge feature that correspond to a series of transitions from the S 2p core levels to unoccupied states. 277 The S L2,3 edge for the pristine sample is similar to that of FeS indicating the predominance of S . 277 It remains nearly alike for the partially charged sample (at 2.66 V) with the exception of minor alterations that could be ascribed to a probable decrease of the Ti/Fe-S bond covalency. In contrast, for the fully-charged sample (at 3 V), the S L2,3 rising-edge shifts by ~1.6 eV towards higher energies and this is indicative of a partial oxidation of S 2- . 291 However the S L2,3 pre-edge does not show the typical peak splitting seen in pyrite-FeS2 or elemental S 0 (Figure S7b). 292 Lastly, it is worth mentioning that the initial shape as well as the energy of the risingedge of the S-L2,3 edge are fully recovered towards the end of the discharge, further indicative of the reversibility of the sulfur redox process. To complement the EELS data that were collected locally on individual electrode particles, we further probed the S L-edge by soft X-ray absorption spectroscopy (sXAS) performed in fluorescence mode, which provides bulk-sensitive information about sulfur redox activity (see Appendix for details). G Direct measurements of the electrodes on the S L-edge through conventional sXAS turned out to be challenging, because of interfering background signals from the carbon present in the electrodes which contributes an overwhelming background through 2nd order harmonic around 140 eV, right below the S-L signals around 160 eV (Appendix Figure A2.15a,b). We could successfully distinguish the S signals from the strong C background, by employing high-efficiency mapping of resonant inelastic X-ray scattering (mRIXS), see the Appendix Section A2.1 and Appendix Figure A2.15c for details. This advanced technique, which has emerged as a seminal technique for detecting oxygen redox, 296-302 further resolves the emitted photon energy, called emission energy, after each sXAS excitation process. 297 Figure III.39 displays these S-L mRIXS signals collected from the pristine, fully-charged, and fullydischarged electrodes. The integration of all the signals within the emission energy range (142 -151 eV) provides the partial fluorescence yield (PFY) signals of the clean S-L sXAS, as summarized in the bottom panel. The pristine Li1.13Ti0.57Fe0.3S2 without carbon additive displays the strongest S-L features. The relatively sharp features in mRIXS at 163 and 164.1 eV excitation energies correspond well with the pre-edge peaks observed in EELS. The sharp features suggest these are dominated by the localized TM 3d character that is hybridized with S orbitals. 277 At higher excitation energy, the continuous bandlike feature corresponds to the intrinsic sulfur states hybridized with TM 4s/4p orbitals. It is clear that the leading edge of these sulfur band-like feature shifts towards high energy for over 1.2 eV in the full charged phase, in complete agreement with EELS and indicating the oxidation of sulfur states. Furthermore, the leading edge position completely recovers after discharge, suggesting a reversible redox reaction of the sulfur.

Overall, the characterizations at the S L-edge confirms the redox-participation of S ligands without fully dimerization of S atoms, in accordance with the crystal structure that was deduced for the fully-charged phase (Figure III.33).

III.3.3 Discussions

We have shown the feasibility of triggering electrochemical activity in Li2TiS3 (i.e., Li1.33Ti0.67S3) by partially substituting Ti 4+ belonging to the metal layer with suitable metal substituent TM. Alike the Li-rich layered oxides having d 0 metals (Li2TiO3, etc…), we showed the feasibility to trigger Li electrochemical activity in Li-rich layered sulfides having d 0 metals by the injection of metal substituents. This does not come as a total surprise as such metal substitution modifies the U over  competition, so that it falls within the domain (lower ) to trigger reversible anionic redox activity as G sXAS and mRIXS were performed by Dr Wanli Yang and Dr Yang Ha (Lawrence Berkeley National Laboratory).

established from theoretical calculations (recall Section I.1.3). 279 Nevertheless, the unsubstituted Li2TiS3 demands further scrutiny, as it possess sufficiently high conductivity and low optical bandgap (black in color), and furthermore due to the fact that its disordered rocksalt (space group 𝐹𝑚3 ̅ 𝑚) counterpart (i.e., Li2TiS3 prepared by ball-milling), is electrochemically active. 282,283,303 We studied in detail the effect of substitution of TM = Fe 2+ , that can initiate the electrochemical activity in Li2TiS3. The Fe 2+ substituted phase, showing electrochemical activity due to cumulated cationic and anionic (S 2-/S n-, n ˂ 2) redox processes, were chosen to study as model materials for anionic redox. The anionic redox activity upon oxidation was spectroscopically confirmed in Li1.13Ti0.57Fe0.3S2 via clear energy shifts in the S L2,3-edge EELS and XANES spectra as well as the onset of a doublet signal in the S 2p3/2-1/2 XPS core spectra. Structure-wise on the other hand, our XRD and EELS evidence did not suggest the local formation of very short S-S dimers (like in pyrites). For comparing with oxides, let's recall that the XPS fingerprint of anionic redox activity in oxides was also the appearance of a new component at slightly higher binding energy (531.5 eV) in the O 1s XPS core spectra that we assigned to O n-(n ˂ 2). [START_REF] Sathiya | Reversible anionic redox chemistry in high-capacity layered-oxide electrodes[END_REF][START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF][START_REF] Pearce | Evidence for anionic redox activity in a tridimensional-ordered Li-rich positive electrode β-Li2IrO3[END_REF] Interestingly, the binding energy of this component was independent of the structure and composition as seen in various Li-rich oxides that show anionic redox activity, with or without evidences of O-O shortening. This observation indicates that although XPS features have been widely used as the spectroscopic signature of the anionic redox activity in various materials, caution has to be exercised when interpreting XPS spectra. Compared to Li-stoichiometric TiS2, the Li-rich sulfide phases offer higher voltage with availability of extracting more Li, as shown here for the Li1.13Ti0.57Fe0.3S2 (Figure III.41). Besides, we found that Li1.13Ti0.57Fe0.3S2 can deliver capacities as high as 245 mAh•g -1 with a near-zero irreversible capacity during the first cycle as compared to ~0.2 Li for Li-rich 3d metal-based oxides. This is consistent not only with the absence of cationic migration in the Li-rich sulfides, but also with less severe changes observed between the first charge and discharge profiles in Li1.13Ti0.57Fe0.3S2, as opposed to Li-rich NMCs that show a staircase charge curve drastically changing to a S-shaped discharge curve. Lastly, part of the initial irreversibility in oxides is associated to a small amount of oxygen release from the surface, either directly as O2 gas or indirectly by reacting with the electrolyte. [START_REF] Luo | Charge-compensation in 3d-transition-metal-oxide intercalation cathodes through the generation of localized electron holes on oxygen[END_REF]304 This is quite unlikely to occur with S, which is a softer element as compared to O, therefore showing less reactivity and greater stability for the same degree of electrochemical oxidation.

III.4 Expanding the S-redox chemistry

So far we have shown the reversible electrochemical activity in Li-rich layered Li2TiS3 by partially substituting Ti 4+ with suitable metal substituent TM = Fe 2+ . Its favourable electrochemical properties that positively contrasts with Li-rich oxides, raises the question about the electrochemical benefits in other Li/Na-rich sulfides, under similar testing protocol. To assess this point, we have extended our study to Li2TiS3 substituted by TM = Ti 3+ , Co 2+ , and also to Na2TiS3, that were reported elsewhere [START_REF] Nasu | Amorphous Na 2 TiS 3 as an Active Material for All-solid-state Sodium Batteries[END_REF]272,281 and briefly discussed below.

III.4.1 Ti 3+ -substituted Li2TiS3

Ti 3+ -substitution as a mean to electrochemically activate Li2TiS3, was explored by Flamary-Mespoulie et al (Thesis, 2018). 272 . They prepared nominal compositions of Li1.33-y/3Ti 4+ 0.67-2y/3Ti 3+ yS2 (or can be regarded as the solid solution of Li2Ti 4+ S3 and LiTi 3+ S2) and described its crystal structure and electrochemical properties in detail.

Hence, to study the practicability of anionic redox upon Ti 3+ -substitution, we have prepared the composition Li1.2Ti0.8S2 (i.e., Li1.2Ti 4+ 0.4Ti 3+ 0.4S2, with y = 0.4). Details of the synthesis and characterization can be found in the Appendix Section A2.1 and Appendix Figure A2.16. The electrochemical performance, the composition is shown in Figure III.42, that shows a capacity of ~235 mAh•g -1 , leading to a specific energy of ~560 Wh•kg -1 . A staircase like charging profile was observed, which becomes S-shaped in discharge, alike the Li-rich NMC. The broad peak at ~2.3 V in the dQ/dV is from the Ti 3+/4+ cationic redox, whereas, the peak at higher voltage corresponds to the S 2-/(S2) n-anionic redox. During discharge, the voltage profile changes its shape and an S-shaped curve is obtained comprising both cationic and anionic redox, which corresponds to the broad peak at 2.5 V in the dQ/dV curve.

III.4.2 Co 2+ -substituted Li2TiS3

The electrochemical property of the Li1.2Ti0.6Co0.2S2 composition was reported by Li et al (2018). 281 We prepared the compositions of Li1.33-2y/3Ti 4+ 0.67-y/3Co 2+ yS2, with y = 0.2, and 0.3 for our studies. Detailed structural characterizations can be found in the Appendix Figure A2.17 In short, electrochemical activity in Li2TiS3 can be activated via both Co 2+ , and Ti 3+ substitution, alike the Fe 2+ substitution. Therefore, these phases are also chosen for our study relate to the practicability of S-redox. large irreversibility >0.2 Na, and also a larger polarization (~700 mV) than the O3 phase. For both O3and P2-Na2TiS3, only one dQ/dV peak is observed during charge-discharge, indicating a single redox process, which can be assigned to the S 2-/S2 n-(n < 2) anionic redox, as Ti 4+ (d 0 ) is supposed to be redoxinactive.

To summarize this section, the Na2TiS3 also presents an interesting case of anionic redox in sulfides, which we will use next for our comparative studies.

III.5 Chapter Conclusions

Aiming to re-examine sulfide based cathode materials, we have studied various Listoichiometric and Li(Na)-rich sulfides. Li-stoichiometric TiS2 possess lower energy density (~450 mAh•g -1 ), however with excellent reversibility and rate capability, while relying solely on usual cationic redox of the transition metal. To push the energy density higher, we scrutinized the practicability of anionic S 2-/S n-(n < 2) redox, encouraged by the recent paradigm of anionic O 2-/(O2) n-redox in Li-rich oxides. The S 2-/S n-redox, based on intercalation chemistry was known either in 'pseudo-layered' TiS3, or in layered Li2TiS3. In fact, the S-redox in LiTi0.5Fe0.5S2 was predicted by Tarascon et al in 1983. 246 We have revisited TiS3, to understand the structural and electrochemical implications of the anionic redox. However in TiS3, the structural irreversibility and sluggish kinetics of the S 2-/S2 2-redox leads to rapid capacity fading and high voltage hysteresis. On the other hand, Li2TiS3 is nearly electrochemically inactive, probably because of unfavourable band positioning (recall Figure III.22). Therefore we designed a new class of Li-rich layered sulfides, namely Fe 2+ -substituted Li2TiS3, and studied for their electrochemical behaviour as cathode materials. Within this series, the phase Li1.13Ti0.57Fe0.3S2 offers the largest reversible capacity (245 mAh•g -1 ) and we have shown that this capacity mainly originates from anionic S 2-/S n-(n < 2) redox, besides the cationic Fe 2+/3+ redox. Energy-wise, these Li-rich layered sulfide positive electrodes display a specific energy of ~600 Wh•kg -1 at the material-level while being composed of earth-abundant elements (3d metals only). Furthermore, we extended our studies to Ti 3+ , and Co 2+ substituted Li2TiS3, and Na2TiS3, which are capable of undergoing reversible anionic redox.

The Li/Na-rich sulfide phases can serve as excellent 'model' electrodes to study the general properties of anionic redox chemistry, this is exactly what will be discussed in the next chapter, where we benchmark their performances against Li/Na-rich oxides. To further assess the hysteresis within the series, we performed GITT experiment on the LiTi0.5Fe0.5S2 (y = 0.5) composition. The GITT was performed in a two-electrode Swagelok cell with the sample employed as powder and the voltage profile is gathered in Figure IV.3. To recall, the LiTi0.5Fe0.5S2 composition has lesser amount of anionic S 2-/(S2) n-redox, since it contains higher amount of Fe than in the Li1.13Ti0.57Fe0.3S2. For LiTi0.5Fe0.5S2, only a small thermodynamic hysteresis of ~25 mV remains, with the voltage profile almost stabilized after 30 mins of relaxation. To verify whether the low hysteresis is true for other Li-rich sulfides, we have studied the Li1.2Ti0.8S2 phase (i.e., Li1.2Ti 4+ 0.4Ti 3+ 0.4S2) also. In Li1.2Ti0.8S2, under normal cycling at C/5 rate, there remains voltage gaps of ~80 mV in the sloppy voltage region and ~160 mV in the high-voltage plateau region (Figure IV.4a). The GITT of the Li1.2Ti0.8S2 is shown in Figure IV.4a. During the sloppy voltage region no hysteresis is observed after the relaxation (30 mins), whereas there remains a small thermodynamic hysteresis of ~60 mV in the high-voltage plateau region. Furthermore, the OCV stabilizes within few mins in the sloppy voltage region, but longer relaxation is needed for the 2nd plateau. This indicates that the small voltage hysteresis is due to the anionic S 2-/S n-(n < 2) redox. In short, small hysteresis is observed here, but much smaller compared to Li-rich oxides. [START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF][START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF] Lastly, we turn to the Li1.2Ti0.6Co0.2S2 composition. GITT was not performed in this case, but we again observed a small overpotential of ~165 mV during cycling under normal condition (Figure IV.4b).

Chapter

Overall, this observation clearly highlights that voltage hysteresis can be effectively mitigated by tuning the ligand, and not just by choosing appropriate transition metals (i.e., 4d / 5d based).

Next we asked the question whether the anionic redox has a path dependence, as in Li-rich oxides, for which the hysteresis was shown to be associated with the path difference during charge and discharge [START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF][START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF]278 . This aspect was further studied by progressive opening of voltage-windows during charge in subsequent cycles and the voltage profiles are summarized in Figure IV.5. First, the charging voltage window is opened gradually in subsequent cycles (Figure IV.5a), however no increase in the hysteresis was noticed. Thus irrespective of the charge window, the discharge voltage is never lowered noticeably. This is reflected in the dQ/dV curves (Figure IV.5b), as all the oxidative capacity from charging even in the high voltage (denoted by red arrows), could be recovered by ~2.48 V peak upon reduction (see the blue arrow). Next, the voltage window was opened gradually during discharge from 100% SoC (3 V) This is drastically different from Li-rich NMC, where charging beyond 4.1 V leads to a noticeable hysteresis that lowers the discharge potential around mid-SoCs (Figure IV.6a). The corresponding dQ/dV curves show that the oxidative capacity above 4.1V is only fully recovered on reduction up to significantly low potential (below 3.6 V), see Figure IV.6b. [START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF] Similar scenario was observed in 4d metal-based Li2Ru0.75Sn0.25O3 (LRSO). [START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF] Thus, in Li1.2Ti0.8S2 no hysteresis loop is observed, letting us to conclude that anionic redox in Li1.2Ti0.8S2 is not path dependent. Overall, since the anionic redox in Li-rich sulfides have the same dQ/dV peak in either chargedischarge (i.e., symmetrically between charge and discharge). This indicates the absence of any path dependence or hysteresis loop, and explains the very small voltage hysteresis observed.

IV.1.2 Electrochemical Kinetics

The low voltage hysteresis observed above could be related to fast interfacial and / or bulk Li diffusion kinetics. Concerning the kinetics, it was shown previously how oxygen redox displays sluggish kinetics in the Li-rich oxides. [START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF][START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF] Hence, to check whether the kinetics in Li-rich sulfides is fast or sluggish, we first deduced cell resistance from the voltage drop during the first 10 s of the relaxation steps of the GITT experiment. As showed for Li1.13Ti0.57Fe0.3S2, the electrochemical resistance remains quite low throughout the whole cycle (Figure IV.8), in contrast to the Li-rich NMC, for which the resistance is significantly larger at all SoCs with additionally an increase at low and high SoCs, the regions involving oxygen redox. 50 The kinetics can be better assessed from the evolution of charge-transfer resistance (Rct, i.e., the resistance associated to the Li + transfer process at electrolyte / electrode interface) at different SoCs. For Li-rich oxides, the Rct is shown to increase drastically in the anionic redox regime, compared to the cationic redox regime (Appendix Figure A3.2). [START_REF] Assat | Fundamental understanding and practical challenges of anionic redox activity in Li-ion batteries[END_REF][START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF] Therefore, to access the Rct in our Li-rich sulfides, we performed electrochemical impedance spectra (EIS) in Swagelok type cell in a two-electrode configuration. A typical Nyquist plot of the EIS is shown in Figure IV.9. The spectra consists a depressed semicircle followed by a rising tail could be fitted with an equivalent circuit R0(RHF//QHF)(RMF//QMF)WLF, where R, Q, and W represent the resistances, constant phase elements, and Warburg element. The resistance R0 corresponds to the resistance from electrolyte and electrical contacts of the cell configuration. The high-frequency (HF) semicircle (at ~300 Hz) comprises the contributions for interphase contacts, surface film, etc. The mid-frequency (MF) arc, at the onset of the rising tail (at ~1 Hz), enlists the charge-transfer resistance (Rct), and followed by low-frequency (LF) Warburg tail that comprises the resistance of Li-diffusion into the electrode. Overall, anionic redox in Li-rich sulfides is fast and no impedance build-up occurs upon cycling, which is positively reflected into the voltage hysteresis. This is quite contrary to Li-rich oxides where the resistance builds up drastically with deeper oxidation of oxygen (shown in Appendix Figure A3.2). [START_REF] Assat | Fundamental understanding and practical challenges of anionic redox activity in Li-ion batteries[END_REF][START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF] This clearly highlights the positive attribute of Li-rich sulfides concerning kinetics.

IV.1.3 Voltage fade

Lastly, regarding the critical issue of voltage fade, we found that it still afflicts Li-rich sulfides (Figure IV.11). However, the voltage fade exists in much lower extent, as shown in Figure IV.12. The voltage fade for Li1.13Ti0.57Fe0.3S2 can be divided in two regimes, starting first with a well-pronounced decrease from 2 nd cycle to the 7 th (a drop of ~35 mV) followed by a stabilization afterwards to reach an overall drop of ~50 mV after 64 cycles (the maximum we have cycled), see In each cycle, the average voltage is defined as the mean of the average charge and discharge voltages that was calculated by dividing the energy with the capacity obtained. Samples were cycled in Swagelok type cells, with samples as powders.

In short, anionic redox promotes voltage fade Li-rich sulfides, but not to the extent that was observed for Li-rich oxides.

IV.1.4 Na2TiS3

At this point we move to the Na-rich sulfide Na2TiS3 (i.e., Na1.33Ti0.67S2), in order to verify whether the positive attribute of anionic redox in sulfides is maintained for Na-(de)intercalation also.

For the O3-Na2TiS3, stable cycling occurs with a polarization of ~300 mV (Figure IV.14). In contrast, the P2-Na2TiS3 shows a large polarization of ~700 mV. Hence, we choose the O3-Na2TiS3 phase for comparison with the Na0. 67[Mg0.28Mn0.72O2] oxide and shown in Figure IV.14. It is clear that the hysteresis in Na2TiS3 (~0.3 V) is much smaller than the same observed for Na0.67Mg0.28Mn0.72O2 (~1.2 V), or recently reported P2-Na0.67Zn0.22Mn0.78O2. [START_REF] Bai | Anionic Redox Activity in a Newly Zn-Doped Sodium Layered Oxide P2-Na 2/3 Mn 1-y Zn y O 2 (0 < y < 0.23)[END_REF] Therefore, the S 2-/S2 n-anionic redox possess much lesser hysteresis than the anionic O 2-/O2 n-anionic redox in Na-cathode materials. Note that this value of hysteresis is comparable to that of cationic redox, e.g., in NaNi0.5Mn0.5O2 (~0.15 V), 306 etc., and approaches the state-of-art hysteresis-free situation observed in Na3V2(PO4)2F3 (NVPF). 307 Overall, amongst O3-and P2-Na2TiS3, the O3-phase offers better electrochemical performance. This is probably because the O3-phase, being thermodynamically stable, is more robust against Na-(de)insertion, whereas the kinetically stabilized P2-phase is less sustainable to structural evolution upon Na-(de)insertion. Furthermore, the O3-Na2TiS3 phase demonstrate excellent reversibility of anionic S 2-/(S2) n-redox in Na-cathode materials, with much lesser voltage hysteresis than the same in oxides, with small voltage and capacity fading. Further improvement here can be performed by fabricating Nafull cell (to avoid side reactions with Na metal anode) and also probably by partial metal-substitution in Na2TiS3, as shown for Li2TiS3 in this thesis.

IV.2 Chapter Conclusions

In summary, we have assessed the practical aspects of anionic redox in Li-rich layered sulfides, namely Fe 2+ /Co 2+ /Ti 3+ substituted Li2TiS3, and Na2TiS3, via detailed electrochemical analysis. When benchmarked against Li-rich NMCs, these phases present several positive attributes such as (i) very small overall voltage fade (as low as ~50 mV for Li1.13Ti0.57Fe0.3S2 even after 60 cycles) and (ii) low voltage hysteresis (~30 mV for Li1.13Ti0.57Fe0.3S2), along with fast kinetics, as compared to Li-rich NMC showing a voltage fade of 150 mV and a hysteresis of 300 mV besides sluggish kinetics. Similar positive attribution was observed for Na2TiS3, that shows low overpotential and decent energy density, compared to the oxides.

Overall, moving from oxygen to sulfur as the ligand turns out to be a correct strategy to partially mitigate the practical bottlenecks of anionic redox. However, we need to keep in mind that this comes at the expense of the overall energy density compared to Li-rich oxides, because of lower potential and higher molecular weight of sulfur. The energy density of Li1.13Ti0.57Fe0.3S2 is compared with Listoichiometric and Li-rich NMC in Figure IV.17. The Li1.13Ti0.57Fe0.3S2 offers an energy density at par with LiCoO2, which may not be attractive as this is much lower than the ~1000 Wh•kg -1 obtained for Li-rich NMC and >700 mAh•g -1 for todays commercialized NMC cathodes. Nevertheless, we believe that they could serve as excellent 'model' electrodes to study the general properties of anionic redox chemistry, especially via the realization of S-based solid-state batteries that can enable playing with temperature as an extra dimension. Both low hysteresis and structural integrity upon Li (de)intercalation are promising assets to design next-generation cathodes provided we can overcome the energy density penalty pertaining to the use of sulfur as a ligand. A compromise could consist in combining the energy advantage of oxygen redox with the practicability of sulfur redox by preparing 3d-metal-based oxysulfides, if they are made possible to synthesize ever. Interestingly, these sulfide based cathode materials can be good candidates for the all-solid-state batteries (ASSBs) with sulfide solid-state electrolytes, due to interfacial compatibility between sulfidebase electrodes and electrolyte materials. Hence, we decided to use the Li1.13Ti0.57Fe0.3S2 as positive electrode in ASSB as will be presented in the next chapter.

V.1 Chapter Introduction

Tremendous attention has been recently focused on all-solid-state-batteries (ASSB), to increase the safety and energy density of Li-ion batteries. This has been unveiled by the discovery of thiophosphate based solid electrolyte (SEs) Li10GeP2S12 by Kanno and colleagues,127 that offer an ionic conductivity (12 mS•cm -1 ) comparable to the liquid electrolytes. In spite of this advance, the practical development of ASSBs remains an issue, mostly because of the difficulties to master the interfaces in ASSBs. Current Li-ion batteries relies on layered oxides such as LiCoO2 (LCO), and LiNi1-x-yMnyCoxO2 (NMCs), etc., as cathode materials because of their high performance. [START_REF] Rozier | Review-Li-Rich Layered Oxide Cathodes for Next-Generation Li-Ion Batteries: Chances and Challenges[END_REF][START_REF] Noh | Comparison of the structural and electrochemical properties of layered Li[END_REF] For sulfide electrolyte based ASSBs using layered oxides as positive electrodes, the oxide-sulfide interface remains problematic, because of the formation of resistive decomposition layer and / or Li-depleted layer (i.e., the spacecharge layer) at the interface (Figure V.1a). 193,198,233,309,310 This leads to highly resistive cathode electrolyte interface (CEI) when oxide cathode materials are used with sulfide electrolytes and is detrimental for the ASSB performance. 193,198,233,309,311,312 This interfacial resistance could be partially mitigated by coating the cathode materials with a buffer redox-inactive (electronically insulating, but ionically conducting) layer of ZrO2, LiNbO3, LiNb0.5Ta0.5O3 or Li2O-ZrO2 etc., which effectively shields the chemical reactivity and high electrode potential of the oxides (Figure V.1b). 193,[200][201][202] Herein, we wanted to explore the performance of sulfide-based cathode materials in ASSB to alleviate the interfacial incompatibility with sulfide SEs. Indeed sulfide based positive electrodes, such as, TiS2, TiS3, etc., have been shown for small interfacial resistances and good performances in ASSB with sulfide SEs. 283,[314][315][316] Furthermore, as sulfide SEs are electrochemically instable at high voltage, sulfide cathode materials with moderate voltage (2 -3 V) can be helpful by limiting the cell potential. Therefore, we decided to utilize high-capacity Li-rich sulfide cathode material Li1.13Ti0.57Fe0.3S2 (LTFS) that has been reported in Chapter 3.

V.2 Designing a cell for ASSB assembly

In order to fabricate an ASSB, we first started building a suitable testing setup in lab. This was the beginning of a long journey in collaboration with Daniel Alves Dalla Corte. This task is quite challenging, and expensive as well, as a setup for ASSB must fulfill the following basic criteria:

(1) An insulating body to contain the ASSB.

(2) Strong body that can withstand with high pressure needed to fabricate (typically up to ~380 MPa) and cycle (at ~70 MPa) the ASSB. 127,157,193,198,[317][318][319] (3) Cycling in air-tight cell (i.e., in inert atmosphere). Hardened steel plungers were used to apply pressure, having a diameter of 8 mm. Moreover, the setup includes the capability of applying temperature (up to 200 °C) and pressure (up to 4 tons•cm -2 ) that can be controlled independently. Lastly, to cycle in inert atmosphere, the whole ASSB assembling and cycling was done inside an Ar-filled glovebox.

V.3 Protocol to fabricate ASSB in the setup

Using the cell previously described the ASSB was assembled by cold-pressing the cathode, solid electrolyte and anode materials, layer by layer, using a hydraulic press. A schematic of the ASSB stack (8 mm in diameter) is shown in Figure V.3. First the solid electrolyte powder (~30 -40 mg, corresponding to a thickness of ~40 -60 μm) was pressed inside the ASSB setup for 5 mins with a uniaxial pressure of 4 tons•cm -2 . Then the cathode composite powder (~5 -10 mg) was pressed on the solid electrolyte layer, with a uniaxial pressure of 4 tons•cm -2 for 1 min. The cathode composite was prepared by mixing cathode material and solid electrolyte powders in 70 : 30 weight-ratio and handgrinding them for 5 mins in an agate mortar. Note that, we avoided conductive carbon-additive Csp in the cathode composite, because of its detrimental effect on the electrolyte. 309,[320][321][322] Then a piece of indium (In) foil was placed on the other side of the electrolyte and pressed with 1 tons•cm -2 for 1 min, that acts as an anode. The In/InLix alloy-anode has a constant potential of 0.6 V vs. Li/Li + . At this point Li-metal anode was not used, to avoid dendrites issue and side reactions. 309,323,324 The process was performed inside an Ar-filled glovebox. Thus the three-layer ASSB was obtained and was cycled with a uniaxial pressure of 1 tons•cm -2 using the pressure frame of the ASSB setup.

V.4 Solid electrolytes used for ASSB

A prerequisite prior to start working on ASSB, is to select suitable solid electrolytes. Toward this end, we have chosen to work with sulfide-based solid electrolytes. We did not consider the oxidebased solid electrolytes that were explored in the Chapter 2, because of their low ionic conductivity. Nevertheless, oxide solid electrolytes in general are challenging to use in ASSBs, because of their processing difficulties as discussed in the Introduction Section. Generally, oxide solid electrolytes require sintering at high pressure and temperature to form intimate inter-particle contacts. In contrary, the sulfide solid electrolytes possess higher ionic conductivity than oxides while being easy to use in ASSBs with RT processing. 154 In this regards, we decided to work with Li2S-P2S5 glass (denoted as g-LPS) and crystalline β-Li3PS4 (denoted as LPS) according to the literature reports. 151,154,177,325,326 The Li2S-P2S5 glass (g-LPS) electrolyte used here was synthesized by mechanochemical synthesis (low-energy ball milling) following literature reported procedure. 151,325 Among g-LPS compositions, the 75•Li2S-25•P2S5 (in molar ratio) presents the highest Li-conductivity, hence this composition was chosen in this thesis. Required amount of Li2S (99.9%, Alfa Aesar) and P2S5 (98%, ACROS Organics) powder were first mixed homogenously by hand grinding for 10 mins in agate mortar. The mixture was transferred to an air-tight zirconia jar (500 ml) with zirconia balls (10 mm diameter, 12 balls) and closed. Then it was ball milled in a planetary ball mill (Frisch Pulverisette 7) for 30 h at 510 rpm. A rest of 10 mins was given after each 15 mins of milling. The resulting g-LPS was collected and hand-grinded again prior to use in ASSB. The as-synthesized g-LPS shows an ionic conductivity of ~0.1 mS•cm -1 at RT (Figure V.4a). Whole process was done inside an Ar-filled glovebox, except the ball-milling, to avoid air exposure. The β-Li3PS4 solid electrolyte was prepared (courtesy to Dr. Florencia Marchini) via a facile solvent-based route reported elsewhere. 177,326 Required amount of Li2S and P2S5 were thoroughly mixed using mortar and pestle, dispersed in anhydrous tetrahydrofuran (THF, 99.9%, Sigma Aldrich), and followed by stirring for 2 days. The white precipitate was recovered, washed with THF, and then dried in vacuum. Finally, the as-prepared powder was heated at 140 °C under vacuum. Whole synthesis was carried out in an Ar-filled glovebox. The synthesized nanoporous β-Li3PS4 shows an ionic conductivity Furthermore, in absence conductive additives, the electronic percolation in cathode composites are provided only through interconnected cathode particles. 309,320 Hence, it is crucial to have a good electronic conductivity of the LTFS cathode material and indeed we observed a high electronic conductivity of 3•10 -2 S•cm -1 In LTFS (deduced from DC polarization measurement, details in the Appendix Figure A4.1). We also measured the electronic conductivity of the cathode composite (LTFS + β-Li3PS4) and found an electronic conductivity of 5•10 -3 S•cm -1 (Figure V.7), which is sufficiently high to be used in ASSB. 233

V.5 Cycling results of ASSBs

Oxide cathode materials, such as, LiCoO2, LiNi0.8Mn0.1Co0.1O2, etc., could be cycled, as reported in literature, but show rapid capacity fading due to reactivity (at the surface) with the sulfide solid electrolytes (β-Li3PS4, LGPS, etc.) when they are not surface-coated, 193,198,309 hence our attempt to cycle sulfide-based cathode materials in ASSB. The Li1.13Ti0.57Fe0.3S2 (LTFS), the best composition amongst the Li-rich sulfides (see Chapter 3) was chosen for the studies. The performance of the LTFS with g-LPS and β-Li3PS4 as SEs in ASSB are summarized below.

V.5.1 Cycling with glassy LPS SE

The ASSB using Li1.13Ti0.57Fe0.3S2 (LTFS) as cathode material in ASSB using Li2S-P2S5 glass (with 75 : 25 molar-ratio) as solid electrolyte and indium as anode was cycled within the potential range of 1. A discharge capacity of ~125 mAh•g -1 was obtained. However, the 1st discharge delivers less capacity than the 1st charge, implying some Li was consumed in the 1st cycle, probably by interfacial reactions. 193,194 This irreversibility was not observed after the 1st cycle (Figure V.9). An overpotential of ~400 mV could be observed in the ASSB (Figure V.9).

V.5.2 ASSB with β-Li3PS4 SE

Next we have replaced g-LPS by β-Li3PS4 (LPS) as solid electrolyte and used a cathode composite comprised of LTFS and LPS (in 70:30 wt%). The ASSB was then cycled between 1.8 -3 V (courtesy of Florencia Marchini) and the voltage profile is shown in Figure V.10. Cycling of the ASSB with higher voltage cut off (up to 4.2 V vs Li) can be found in the Appendix Figure A4.2. Using β-Li3PS4, more capacity could be obtained (discharge capacity ~150 mAh•g -1 ). Interestingly, a polarization ~200 mV (in 2nd cycle) was observed, which is smaller than the value observed in the g-LPS ASSB (Figure V.9). This improvement became possible probably due to better interfacial compatibility or SEI formation with the β-Li3PS4 as solid electrolyte. Furthermore, we observed improvement of the overpotential at higher temperature, as a consequence of the increased ionic conductivity of the β-Li3PS4 SE at higher temperature.

V.5.3 EIS study of the ASSB

To further grasp insights into the interfacial compatibility in the ASSB, we collected in situ electrochemical impedance spectra (EIS) of the Li1.13Ti0.57Fe0.3S2│g-LPS│In cell at full-charged and full-discharged states. The EIS was recorded in two-electrode cell configuration, with 25 mV bias in the frequency range of 0.5 MHz -1 Hz (Figure V.11a). All spectra consists a semicircle at high frequencies (HF, >14 kHz) and a very small, almost non-visible semicircle at mid frequency region (MF, 14 kHz -100 Hz), followed by Warburg tail at low frequencies (LF, <100 Hz). Interestingly, slight increase was observed only until the 1st charge, whereas no significant increase of the impedance was observed 1st charge onwards. Importantly, the MF semicircle (R2), which comprises the interfaces between solid electrolyte and cathode (RSE-cathode), is negligible (barely visible). Note that a well-separated semicircle (typically >200 Ω) was often reported when oxide cathode materials (LiCoO2, LiNi0.8Mn0.1Co0.1O2, etc.) were cycled with sulfide electrolytes. 193,198 This finding suggests that sulfide based cathode Li1.13Ti0.57Fe0.3S2 forms stable interface with the Li2S-P2S5 sulfide glass. This imply that no coating is necessary for the Li1.13Ti0.57Fe0.3S2, whereas expensive LiNb0.5Ta0.5O3 or Li2O-ZrO2 coating is usually required for oxide cathode materials (e.g., LiCoO2, LiNi0.8Co0.15Al0.05O2, etc.) to decrease the interfacial resistance (to >40 Ω). 193,200

V.6 Chapter Conclusions

In summary, the sulfide-based Li1.13Ti0.57Fe0.3S2 (Li-rich sulfide) was successfully cycled in allsolid-state batteries and we found it is chemically/electrochemically compatible with the sulfide-based solid electrolytes. Moreover, its excellent electronic conductivity contributes positively to the cycling performance. However, we observed a large irreversibility in the 1st cycle which is probably originated from the Li-loss in SEI formation at the In-anode. To alleviate this Li-loss, further investigations are being pursued to cycle with Li anode. Lastly, through this study we showed sulfides could regain interest in ASSBs that use Li metal as negative electrode. Hence we are going back to the early time of Li-metal batteries, when sulfides were stellar. Whatever, whether sulfides can replace oxides as positive electrode in ASSBs remains to be proved.

Chapter VI

General Conclusions

In this thesis, two main axes were simultaneously pursued. The first one is related to the allsolid-state battery (ASSB) that are supposed to improves the safety and performance of Li-ion batteries. Over the last few years, ASSBs has emerged as a topic of intense research, and furthermore, with immense push from industries, especially EV makers. However, this task is indeed very challenging, and on the practical level, it still remains to be demonstrated. The second one has dealt with the anionic redox, which improves the energy density of cathode materials. Although anionic redox have equally been explored, from both application and fundamental point of view, numerous questions regarding its practicability remains to be addressed. This is what we have done in this thesis work was carried out and the findings are summarized below.

In the framework of the ASSBs, we first focused to develop new oxide-based solid electrolytes (SEs), and explored new phases in the Li(/Na)2O-ZnO-P2O7 phase diagram. We investigated the Li6Zn(P2O7)2 and Li4Zn(PO4)2/ Na4Zn(PO4)2 phases, that undergo an order-disorder phase transition(s) with temperature. Much higher Li/Na conductivity (and lower activation energy) could be obtained for the high-temperature phases (HTPs) due to disorder. To correlate this increment, we solved the crystal structures by combination of SXRD and NPD studies, we found for the HTPs, the structural backbones consisting of P2O7/PO4 groups remains invariant, whereas the mobile Li sublattice becomes disordered and cation-mixed (with Zn). Furthermore, the BVEL calculations for Li6Zn(P2O7)2 / Li4Zn(PO4)2 indicate, that the conduction mechanism, which is 2-D in the ordered low-temperature phase (LTP) becomes 3-D in the HTP. Having characterized the phase and determined the conductivity transitions, a challenge was to stabilize the HTPs at RT. We perused this question via hypervalent substitution. We have showed that the HTPs can be stabilized by partially substituting the Zn 2+ by M 3+ , thus creating 'inherently' disordered phases that are stable at RT. Such an effort was failed for Li6Zn(P2O7)2, however it was successful for Li4Zn(PO4)2/ Na4Zn(PO4)2. The newly designed β-Li4-xZn 2+ 1-xGa 3+

x(PO4)2 and β-Na4-xZn 2+ 1-xM 3+

x(PO4)2 (M = Ga, Al) phases are stable at RT while crystallizing with the HTP structure. Furthermore we showed that, by conventional ball-milling the disordered β-Li4Zn(PO4)2 phase could be stabilized.

Such study can guide as a 'model' to design new efficient SEs, especially by extending to its sulfide-counterparts. 328 However, despite such exploration of rich crystal chemistry and materials engineering, the ionic conductivities obtained (< 10 -10 S•cm -1 ) for these new phases at RT were too low to be implemented in functional all-solid-state batteries (ASSBs). Hence, we decided to use previously reported sulfide-based SEs, namely amorphous Li2S-P2S5 glass and crystalline β-Li3PS4, for their sufficiently high Li-conductivities (~10 -4 S•cm -1 ), and capability of RT processing, for practical ASSB fabrication. At this point, we took a different strategy to tackle the interfacial incompatibility of sulfide SEs with oxide cathode materials. Specifically, we decided to design Li-rich sulfide cathode materials with higher capacity, capitalizing the benefit of anionic redox.

We started with revisiting the well-known S2 0.0831 (7) 0.5831 (7) 0.3730 (7) 0.3835 (7) 0.2049 (6) 0.5179 (10) 0.4567 (7) 0.1672 (7) 0.3226(4) 0.1016 (7) 0.2117 (7) 0.4037 (8) 0.0818(4) 0.1437 (7) 0.1768 (7) 0.4620 (7) 0.0128 (9) 0.1790 (16) 0.3376 (16) 0.1760 (18) 0.5606 (17) 0.3229 (17) 0.0607 (17) 0.25 0.3883 (8) 0.2580 (9) 0.3533 (5) 0.5410 (8) 0.4367 (8) 0.2727 (7) 0.3449( 5 2.039 (36) 1.866 (39) 5.092 (61) 1.786 (36) 1.932 (31) 2.207 (43) 4.926 (56) 2.101 (41) 2.017 (42) 1.933 (35) 1.961 (21) 1.009 (29) 0.907 (23) 1.065 (28) 1.062 (28) 0.987 (26) 0.988 (24) a The relatively large Biso for Zn (2.05 Å 2 ), while the Biso for the Li atoms (refined with the constraint that they are all equal) are slightly negative (-0.15 Å 2 ). As Li + and Zn 2+ are similar in size, this could indicate a possible mixing between those atoms. Preliminary refinements using Li/Zn mixing indicate that this possible mixing would involve less than 2% of Zn on the Li sites. However, due to the complexity of the structure with six independent Li sites, we did not push further this quantification. A1.2 0.2706 (7) 0.1628 (7) 0.0620 (7) 0.3614( 6 0.535 (17) 0.535 (17) 0.535 (17) 0.535 (17) 0.535 (17) 0.535 (17) 0.535 (17) 0.535 (17) 0.39 (10) 0.39 (10) 0.39 (10) 0.39 (10) Occ.

Appendix Table

Figure 1 .

 1 Figure 1. Worldwide energy consumption from various sources. (a) The projection by year, and (b) their share (in percentage) of energy sources. Figure adapted from International Energy Outlook 2019. 1

Figure I. 1 .

 1 Figure I.1. Schematic configuration of Li-metal and Li-ion batteries. (a) Li-metal battery with Lifree insertion cathode (e.g., TiS2) and Li metal anode, which shows dendritic Li growth that can cause short-circuit. (b) Li-ion battery having a Li-containing insertion cathode (e.g., LiCoO2) with an insertion anode (i.e., graphitic-carbon) material.

Figure I. 2 .

 2 Figure I.2. Comparison of band positions and voltage profiles in LiMS2 sulfides and LiMO2 oxides. Simple schematics of the band positions of LiTi 3+ S2 (a) and LiCoO2 (b) are shown. The relative band positions are estimated based on references 23,24 . The voltage profiles (1st cycle) of TiS2, LiCoO2 (LCO), and LiNi0.33Mn0.33Co0.33O2 (NMC) are shown in c-d. The purple circles show the starting points. TiS2 was discharged first (starting with x = 0) and LCO and NMC were charged first (starting with x = 1).

Figure I. 4 .

 4 Figure I.4. Illustrating cationic and anionic redox. In Li-stoichiometric layered oxides, such as LiCoO2 (a), only transition metals (TM) are redox active (indicated by blue dotted circle). The Li-rich layered oxides, such as Li1.2Ni0.13Mn0.54Co0.13O2 (Li-rich NMC) (b) that contains extra Li within the metal layers, both TMs and the ligand O atoms (indicated in pink color) are redox active. The corresponding voltage profiles (c and d) indicate a near-doubling of capacity and specific energy for the Li-rich NMC due to cumulative cationic and anionic redox processes. Figure reproduced from reference 37 .

  e., Li[Li1/3M2/3]O2), one of the O 2p orbitals (the one pointing towards Li in the [Li1/3M2/3]O2 layer) is weakly bonded because of its relatively small overlap with the Li 2s orbital. Hence, it behaves like an O non-bonding state and is located above the stabilized (M-O) bonding band (Figure I.6b). This nonbonding O 2p orbital can donate electron once the (M-O)* is emptied, without destabilizing the structure. But the reversibility of the anionic redox depend on relative band positioning of (M-O)* antibonding and O 2p non-bonding bands. To describe, we need to introduce the d-d Coulomb interaction term U, that splits the partially filled (M-O)* band due to on-site electron repulsion within the d orbitals, as known as Mott-Hubbard splitting, leading to empty upper-and filled lower-Hubbard bands (UHB and LHB, respectively). Depending on the relative values of U vs. Δ (the energy difference between (M-O) and (M-O)*), three different scenarios can occur.

Figure I. 6 .

 6 Figure I.6. Schematic band diagrams that governs the anionic redox. In LiMO2 (a), no non-bonding O 2p exists, whereas in Li-rich Li[Li1/3M2/3]O2 (b), O atoms are coordinated by only two metals, so that part of the O 2p orbitals that does not have the good symmetry to participate in M-O bonds and remain as pure O non-bonding states. (c-e) Different relative positioning of the band that control the redox. Figure taken from reference 37 .

Figure I. 7 .

 7 Figure I.7. Moving from Li-rich layered oxides to sulfides. Schematic band structure of Ni 2+ substituted Li2MnO3 (a), and Li1.33Ti0.67S2 or Li2TiS3 (b). The relative band positions are estimated based on references[START_REF] Luo | Charge-compensation in 3d-transition-metal-oxide intercalation cathodes through the generation of localized electron holes on oxygen[END_REF][START_REF] Goodenough | Locating redox couples in the layered sulfides with application to Cu[Cr2]S4[END_REF] . The label nb stands for non-bonding.

Figure I. 8 .

 8 Figure I.8. Reported pictures of LIBs causing fire in electronic devices and vehicles. Figures taken from references85-88 .

Figure I. 9 .

 9 Figure I.9. Battery technologies using solid electrolyte. Na-S battery with molten-Na │ β-alumina │ molten-S configuration. Based on similar concept Na-NiCl2 battery (so called ZEBRA battery) was also commercialized using molten-NiCl2 cathode that operates at ~250 ° C. 99,103 (b) LiPON microbattery with LiCoO2 or LiMn2O4 cathode and vapor-deposited Li or Sn3N4 anode with Si3N4 insulation casing (Vacuum deposition for all layers), and (c) all-solid-state battery for future commercialization. Figures in a-b are adapted from references 99,100 .

Figure I. 10 .

 10 Figure I.10. Minimum thickness required for a SE to match the resistance of liquid electrolytes. LiPF6 in EC/DMC with an area of 1 cm 2 was used for comparison. Note that the effect of Li-ion lithiumion transference number t+ was not accounted in the calculation. Figure adapted from reference 105 .

Figure I. 11 .

 11 Figure I.11. Heat content (left y-axis) and ionic conductivity (right y-axis) of PbF2. Figure adapted from reference 108 .

a

  . Table I.1. Different classes of oxide-based solid electrolytes (crystalline). a Solid Electrolyte σRT (S•cm -1 ) Performance Garnet-type Li7La3Zr2O12 (LLZO) 123,129,130 e.g., Li7-3xM'xLa3-xZr2O12 (M'= Al 3+ , Ga 3+ ) or, Li7-xLa3M"2-xTaxO12 (M"= Nb 5+ , Te 6+ ) HT annealing for densification Does not wet Li metal NaSICON-type LiM2(PO4)3, M= Zr, Ti 112,132 e.g., Li1.3Al0.3Ti1.7(PO4)3 (LATP), 133 Li1.5Al0.5Ge1.5(PO4)3 (LAGP) Note that amorphous glassy-type SEs, such as Li3PO4/Li3BO3/Li2SO4 glass, LiPON (composition varies from Li3.3PO3.8N0.24 to Li3.6PO3.3N0.69), Li-La-Zr-O, etc. (with conductivities <10 -6 S•cm -1 at RT) were also explored that were prepared by mechanochemical synthesis, RF magnetron sputtering, etc. 139-142 Table I.2. Different classes of sulfide-based solid electrolytes. -type Li3PS4 121,143e.g., Li10MP2S12 (M= Si, Ge, Sn),127,144,145

3

 3 Design principle of superionic conductorsConductivities in different families being discussed, question arises about what is behind the conductivity in these solids ? Inorganic solids usually exhibit low ionic conductivity than traditional liquid electrolytes. In liquids, ion-conduction involves the following process-solvation and dissociation of electrolyte salt by/in polar solvent molecules, and migration of solvated ions through solvent media(Figure I.13a). Contrarily, in solids, the ionic conduction is mediated by empty Li/Na sites and due to their strong bonding inside the lattice, a large energy barrier ΔHmig must be overcome for the ion migration to occur(Figure I.13b).

Figure I. 13 .

 13 Figure I.13. Simplified schematics illustration of ionic conduction process in liquid (a) and solid (b) electrolytes.125,164

Figure I. 14 .

 14 Figure I.14. Illustration of Li-conduction in Li10GeP2S12 (LGPS). The LGPS shows an average Liconductivity σRT of 12 mS•cm -1 (Ea= 0.27 eV). 127 Li, P, and Ge atoms are shown in yellow, green, and blue. The anionic backbone consists of immobile PS4/GeS4 tetrahedra. The Li-ions forms the mobile sublattice. The Li-conduction pathways are highlighted in yellow. The Li1 and Li3 conducts along b direction (with σ[001] = ~27 mS•cm -1 ) and Li4 connect the 1D conduction channel (with σ[110] = ~7 mS•cm -1 ), while Li2 remains almost immobile.[165][166][167][168][169][170]

Figure I. 15 .

 15 Figure I.15. Ionic conductivity of Li4-3xAlxSiO4. The fraction of occupied mobile ion sites, nc varies from 1 (at x = 0) to 0 (at x = 0.5) and a broad conductivity maximum occurs around x = 0.25. Figure taken from reference 173 .

Figure I. 16 .

 16 Figure I.16. Crystal structures of LLZO. O atoms are not shown.Li, La, and ZrO6 are shown in green, cyan, and purple, respectively. Note that the cubic phase can be stabilized at RT via hypervalent substitutions, e.g., by Al 3+ , Ga 3+ etc. (in Li site, i.e., Li7-3xGaxLa3-xZr2O12), or by Ta 5+ , Nb 5+ , Te 6+ , W 6+ etc. (in Zr site, i.e., Li7-xLa3Zr2-xTaxO12), etc.129,130

Figure I. 18 .

 18 Figure I.18. Enhancement of ionic conductivity by moving from oxides to sulfides. (a) Optimizing ionic conductivity by introducing defects (Li vacancies or interstitials). 126 Note that the LGPS was designed by optimizing Li-concentration in Li3+yGeyP1-yS4 with y = 0.33. 127 (b) Variation in diffusivities (D) and activation energies (Ea) upon anion tuning in Li10GeP2X12 (X = O, S, and Se), calculated from MD simulations. 172 Figures adapted from references 126,172 .

Figure I. 19 .

 19 Figure I.19. Structural and configurational features of superionic conductors. (a) Calculated energy path in bcc and fcc sulfur-lattices (for single Li + migration). (b) Migration energy barrier LLZO for concerted migration of multiple Li + (in left) and single Li + (in right) along the diffusion channel. The Li + /path are shown as green spheres and O/S atoms as yellow spheres. T and O represent Li tetrahedra or octahedra. Note that the bcc lattice of concerted migration decreases the energy barrier of migration. Figures taken from references179,180 .

  Figure I.20.

Figure I. 20 .

 20 Figure I.20. Comparison of advantages/disadvantages between oxide and sulfide solid electrolytes. Figure taken from reference 94 .

Figure I. 21 .

 21 Figure I.21. Electrochemical stability and decomposition of solid electrolytes. (a) Electrochemical stability window (shown in solid color bar) and the oxidation potential to fully delithiate the material (marked in dashed line). (b-c) Schematic illustration of interphase layer of Li3PS4 and LGPS against Li

Figure I. 22 .

 22 Figure I.22. RT sinterability of sulfide SEs. Ionic conductivity of Li2S-P2S5 glass pellets pressed at RT. Inset shows the SEM image of cross sections of the pellet that was pressed at 360 MPa. Note that good ionic conductivity could be obtained only pelletizing at moderate pressure. 154 (b) SEM image of cross sections of the pellet of RT pressed LLZO oxide, showing inefficient densification. 154 (c) SEM image of cross sections of the RT-pressed pellet of stacked LGPS and LGPS-LiCoO2 showing good densification. 193 Figures adapted from references154,193 .

  Figure I.23a shows performance of LiCoO2 with LGPS-type SEs (indicated on the figure) and Li metal anode. 194 However because of interfacial reactions, full capacity from LiCoO2 could not be obtained. Moreover, dendrite issues were also shown persist with sulfide SEs and this is the reason why In-Li alloy is generally used. 195-197 NMC-811 cathode materials were used in ASSB with β-Li3PS4 SEs and In anode (Figure I.23b). 198 However, much less capacity could be obtained than the capacity obtained in LIBs with liquid electrolyte.

Figure I. 23 .

 23 Figure I.23. Performances of ASSBs with different sulfide SEs.Voltage profiles of LiCoO2 in ASSB using the SEs indicated on the figure. (b) Voltage profiles of NMC-811 (LiNi0.8Co0.1Mn0.1O2) in ASSB with β-Li3PS4 SE and in LIB. The EIS spectra of the ASSB is shown in (c). (d) Voltage profiles of surface-coated (with LiNb0.5Ta0.5O3, LNTO) and uncoated LCO (LiCoO2) in ASSB and LIB. The EIS spectra of the ASSBs are shown in (e). Figures adapted from references193,194,198 .

  Following this line, we have concentrated our search towards new Li-conducting oxides in the Li(/Na)-Zn-P-O phase diagram (Figure II.1), based on the facts that: 1) Zn-based phosphates are known to possess disordered polymorphs, 2072) Zn 2+ (d10) is electrochemically redox-inactive, and finally, that, 3) Zn-based phosphates can be synthesized at relatively lower temperature (< 800 °C), whereas most oxide SEs require high temperature (> 1000 °C) synthesis/sintering.

Figure II. 1 .

 1 Figure II.1. Phase diagram of the Li(Na)-Zn-P-O System. Diagram adapted from Rao et al, 2009. 207

Figure

  Figure II.2a shows the DSC profile of Li6Zn(P2O7)2 heated to 600°C and cooled to RT. A Several endothermic peaks of difference amplitudes were observed, as opposed to the single peak reported earlier 207 indicating the existence of multiple phase transitions upon heating. During the cooling, existence of fewer exothermic peaks suggests that the compound undergo different structural transitions. To grasp deeper insights in the phase transitions and the associated structural transformations, we have performed temperature dependent in situ synchrotron-XRD (SXRD) experiments over a broad temperature range up to 650 °C and the diffractograms are gathered in Figure II.2b,c. Diffraction patterns were collected at 5 °C intervals. They are plotted as 2D contour diagram in Figure II.2b and as a function

Figure II. 2 .

 2 Figure II.2. Probing temperature-dependent structural modifications by DSC and SXRD. (a) DSC curve of the Li6Zn(P2O7)2 sample. (b) 2D contour plot showing the temperature dependence of the SXRD patterns indicating different phases with temperature. (c) Evolution of the SXRD patterns during heating of the Li6Zn(P2O7)2 sample observed in the temperature dependent synchrotron-XRD experiment. Successive phase transitions have been denoted on the plot and corresponding temperatures has been indicated at the right. In (b) the horizontal lines correspond to the phase transition temperatures indicated in (c).

Figure II. 3 .

 3 Figure II.3. Evolution of the SXRD patterns during cooling of the Li6Zn(P2O7)2 sample.

  to form the pyrophosphate groups. Neutron powder diffraction (NPD) pattern recorded at RT was then used to determine the Li positions using Fourier difference maps. Finally, all atoms were refined using combined Rietveld refinement of both SXRD and the NPD. The refinements present good reliability parameters (χ 2 = 3.5, RB= 5.08% for NPD) and are shown in Figure II.4, with the structural parameters gathered in the Appendix TableA1.1.

Figure II. 4 .

 4 Figure II.4. Combined Rietveld refinement of the XRD (a) and NPD patterns (b) of α-Li6Zn(P2O7)2. The red circles, black continuous line, blue line and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively. The deduced crystal structure is shown in Figure II.6a,c. Four crystallographically distinct P atoms and fifteen crystallographically distinct O atoms form the pyrophosphate P2O7 4-groups which are all oriented in a parallel way. Zn and Li are placed in between those pyrophosphate groups and only coordinated by oxygen atoms from P2O7 4-. The Zn atoms (on a single Wyckoff site 8f) stay in a distorted trigonal bipyramidal coordination with three equatorial Zn─O bond lengths of ~2 Å and two slightly longer axial bonds of ~2.3 Å. The Li atoms are distributed upon six crystallographic sites and all sit in a distorted tetrahedral positions with average Li─O bond length of ~1.96 Å.

  It's a and b unit cell parameters being twice smaller than for the α-phase, we must have a mixing of Li and Zn atoms to comply with the invariance of the molecular formulae. This was supported by combined SXRD/NPD refinement (Figure II.5), which indicates that Li and Zn are statistically distributed over two Wyckoff sites (general position 8f). The preexisting Zn trigonal-bipyramidal positions in the P2O7 4-layer are converted to tetrahedral Li2/Zn2 positions with 0.9/0.1 occupancy and the preexisting Li positions in the Li layer lead to a trigonalbipyramidal mixed Li1/Zn1 position with 0.85/0.15 occupancy in the ζ-phase. The structural parameters are gathered in the Appendix Table A1.2 and the deduced crystal structure is shown in Figure II.6b,d.

Figure II. 5 .

 5 Figure II.5. Rietveld refinement of the XRD and NPD patterns of ζ-Li6Zn(P2O7)2, recorded at 600 °C. For legends refer to Figure II.4.

Figure II. 6 .

 6 Figure II.6. Crystal structure of the polymorphs of Li6Zn(P2O7)2. (a-d) Along [010] direction (a,b) and along [001] direction (c,d). Li is shown in yellow; ZnO4 and PO4 tetrahedra are colored in blue and orange, respectively. The structural determinations of the two main phases at RT and 600 °C shed light on the four β, , , and  intermediate phases, even if it is out of the goal of the present study to solve all structures.

  )2 (see Appendix Section A1.1 for details). Representative Nyquist impedance plots are shown in Figure II.7a. All recorded spectra consist of a depressed semicircle at the high frequency region followed by a Warburg tail at lower frequency region, which confirms the ionic nature of the ac conductivity. The slightly flattened semicircle sums up the impedance responses of the sample made of grains (g) and grain boundaries (gb). All the contributions were decoupled by fitting the whole impedance spectra to an equivalent circuit model (Rg//Qg)(Rgb//Qgb)Wo (inset of Figure II.7b); where R and Q represent the resistance and constant phase element associated with the grains (g) or grain boundaries (gb) of the sample, respectively, and Wo is the Warburg resistance.

Figure II. 7 .

 7 Figure II.7. AC impedance spectra of the Li6Zn(P2O7)2. (a) Representative AC impedance spectra of α-Li6Zn(P2O7)2 recorded at 200, 225 and 250 °C. Some typical frequencies have also been indicated.

  The fitting of the impedance spectra recorded at 200 °C. Inset shows the equivalent circuit used to fit the spectra. A typical fitting is shown in Figure II.7b. The fitting decouples the flattened semicircle into two deconvoluted semicircles and lead to capacitance values of ~10

Figure II. 8 .

 8 Figure II.8. Arrhenius plot of Ionic Conductivity of the Li6Zn(P2O7)2 sample. (a) Grains conductivity (σg) and grain boundary conductivity (σgb), (b) total conductivity (σtot). The blue, green and red background refers to the three regions of ionic conductivity. In (b) a, b, and c denote the solid lines in blue, green, and red, respectively. However, since this compound undergoes several order-disorder phase transitions at higher temperatures, we next measured the conductivity of the sample over a broader temperature range up to 650 °C. The variation of the measured σtot values are shown in Figure II.8b. Three jumps in conductivity can be observed in the Arrhenius plot, as denoted by lines a, b and c in the plot. First jump was observed at ~390 °C, as shown by line b, which probably is induced by the transition to the -polymorph. Extrapolation of the line b to RT indicates a σtot value of ~3.3 × 10 -10 S•cm -1 , with Ea of 0.73 eV.

Figure II. 9 .

 9 Figure II.9. AC impedance spectra of α-Li6Zn(P2O7)2 recorded at high temperatures. Some typical frequencies have also been indicated.

  polymorphs. The details of the calculations are described in the Appendix Section A1.1. The BVEL maps representing the Li + conduction pathways in the Li6Zn(P2O7)2 polymorphs are shown in Figure II.10. The calculations predict for α-Li6Zn(P2O7)2 an infinite 2-D Li + percolation path along the b-c plane at 1.05 eV above the minimum energy (Figure II.10a) which most likely results from the layered structure of the polymorph.This contrast with the calculated BVEL map for the ζpolymorph which predicts infinite 3-D Li + percolation paths at 0.65 eV above the minimum energy(Figure II.10b). From a comparison of the Li + conduction channels in both polymorphs we can deduce that the onset of the Li/Zn inter-sites mixing associated to the ζpolymorph creates new migration channels. Thus interlayer Li + hopping become possible through mixed Li + /Zn 2+ sites in agreement with the highest ionic conductivity measured for the ζas compared to the αpolymorph.

  Figure II.10. BVEL map of (a) α-Li6Zn(P2O7)2 and (b) ζ-Li6Zn(P2O7)2 at the percolation energies. The yellow domains indicates the migration paths for Li + in the structure, obtained using an iso-surface value of 1.05, and 0.65 eV above the minimum energy, respectively. For the color codes, refer to Figure II.6.

  ) Å, b = 10.24192(13) Å, c = 8.10549(15) Å,  = 104.8589(9)°, V = 651.538(2) Å 3 , with Z = 4, using the structural model reported earlier 213 .

Figure II. 11 .

 11 Figure II.11. Rietveld refinement of the XRD and NPD patterns of α-Li4Zn(PO4)2, measured at RT. The red circles, black continuous line, blue line and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

  )2 phase. Differential scanning calorimetry (DSC), shown in Figure II.13a, of the as-prepared Li4Zn(PO4)2 sample clearly depicts the α → β transitions at ca. 443 °C on heating, with an activation enthalpy of 11.7 kJ/mol (endothermic). On cooling, the DSC trace shows the reversible process with a thermal hysteresis of ∼15 °C indicating that the transition is of first order. This phase transition was monitored by temperature dependent XRD experiments and the diffractograms are gathered in Figure II.13b. Upon heating, XRD patterns of α-Li4Zn(PO4)2 show a small shift towards lower angles due to thermal expansion, until the temperature reaches 430°C. Above this temperature, the XRD pattern changes and looks similar to the one previously reported for the β form. On cooling, this process is reversible and the α-phase is recovered below 420°C, see Figure II.14.

Figure II. 12 .

 12 Figure II.12. Crystal structure of α-Li4Zn(PO4)2. (a) along [𝟏 ̅ 𝟎𝟏], (b) along [001]. Li is shown in yellow; ZnO4 and PO4 tetrahedra are colored in blue and green, respectively.

Figure II. 13 .

 13 Figure II.13. Probing temperature-dependent structural modifications of Li4Zn(PO4)2 by DSC and XRD. (a) DSC curve of Li4Zn(PO4)2, measured in air with a heating and cooling rate of 10 °C/min. (b) Evolution of the XRD patterns of Li4Zn(PO4)2 between 100 and 600 °C on heating. Values of temperatures corresponding to the patterns are indicated at the right. The patterns colored in blue and red corresponds to the and polymorphs, respectively, and the peaks that disappear for the phase has been indicated by blue backgrounds. (c) Evolution of lattice parameters with respect to temperature. The lattice parameters have been converted to monoclinic unit cell setting pertaining to the αcrystal structure for comparison. The blue and red regions correspond to the and phases, respectively.

Figure II. 14 .

 14 Figure II.14. Evolution of XRD patterns of Li4Zn(PO4)2 with temperature during cooling. Values of temperatures corresponding to the patterns are indicated at the right. The patterns colored in blue and red corresponds to the and polymorphs, respectively, and the peaks that reappear for the phase has been indicated by blue backgrounds.

Figure II. 15 .

 15 Figure II.15. Rietveld refinement of the XRD and NPD patterns of β-Li4Zn(PO4)2, measured at 600°C. The red circles, black continuous line, blue line and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively. The corresponding Rietveld refinements in Figure II.15 and the deduced structural model is shown in Figure II.16b, with the structural parameters are gathered in the Appendix Table A1.4. Lastly, all temperature-dependent XRD patterns were Rietveld refined and the evolution of lattice parameters of the Li4Zn(PO4)2 with temperature is summarized in Figure II.2c using the monoclinic description for the β phase for an easier comparison. It can be noticed that the unit cell volume monotonously increases with T and exhibits a slight discontinuity at the transition, as expected from a first order phase transition.

  .4. Lastly, all temperature-dependent XRD patterns were Rietveld refined and the evolution of lattice parameters of the Li4Zn(PO4)2 with temperature is summarized in Figure II.2c using the monoclinic description for the β phase for an easier comparison. It can be noticed that the unit cell volume monotonously increases with T and exhibits a slight discontinuity at the transition, as expected from a first order phase transition.

Figure II. 16 .

 16 Figure II.16. Crystal structure of (a) α-Li4Zn(PO4)2, (b) β-Li4Zn(PO4)2 and (c) β'-Li3.7Zn0.7Ga0.3(PO4)2. In (a), the monoclinic α unit cell is plotted in black and the orthorhombic β cell is shown in red. For sake of clarity, only half of the unit cell content (0 ≤ y ≤ ½ in the monoclinic setting) is shown. Li, Zn, P and Ga are shown in yellow, blue, green and pink, respectively and PO4 tetrahedra are colored in green.

Figure II. 17 .

 17 Figure II.17. Rietveld refinement of the XRD pattern of the ball-milled -Li4Zn(PO4)2 sample, measured at RT. The red circles, black continuous line, blue line and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

  5 mol pfu., beyond which secondary phases were found to form. The RT SXRD patterns of as-prepared Li4-xZn1-xGax(PO4)2 samples with 0 ≤ x ≤ 0.5 are shown in Figure II.18 (and the XRD patterns in Figure II.19a). The Bragg peaks gradually shift towards higher 2 angles, indicating that the unit cell volume decreases with insertion of Ga 3+ , as expected from the smaller ionic radii of Ga 3+ versus Zn 2+ (Figure II.19b and Figure II.18b). Moreover, low angle peaks gradually broaden and their intensity decreases with increasing the Ga content, prior to finally disappear completely for 0.3 mol Ga pfu.

Figure II. 18 .

 18 Figure II.18. SXRD patterns of the Li4-xZn1-xGax(PO4)2 samples. (a) SXRD patterns for the compositions with x values indicated in the figure. (b) Zoom of the SXRD patterns of selected 2θ range, marked in (a) with a rectangle. In (b), the peak marked by * is from some minor unknown impurity.

Figure II. 19 .

 19 Figure II.19. XRD patterns of the Li4-xZn1-xGax(PO4)2 samples. (a) XRD patterns of the samples, measured at RT. In (a), and (b), the values of x corresponding to the patterns are indicated at the right. The patterns colored in blue, green and red corresponds to the -, '-, and polymorphs, respectively, and the peaks that disappear from the and phases have been indicated by blue and green backgrounds, respectively. (b) Evolution of lattice parameters with respect to Ga content. The lattice parameters have been converted to monoclinic unit cell setting for comparison. The blue, green, and red regions correspond to the -, '-, and phases, respectively.

  ) Å, b = 6.4803(3) Å, c = 10.1468(5) Å, and V = 324.77(3) Å 3 . However the presence of peaks at 2 values of 16.2° and 29.8°, rules out the Pnma space group, and suggests its P212121 subgroup (011 and 013 peaks, respectively). A structural model was therefore constructed and Rietveld refined against both NPD and XRD patterns, simultaneously. The best structural model (shown in Figure II.16c), as deduced by combined NPD and XRD refinements (Figure II.20), was obtained by mixing Li and Zn on one site, and Li and Ga on another site. This structure (structural parameters are gathered in the Appendix Table

  Figure II.20. Rietveld refinement of (a) XRD, and (b) NPD patterns of β'-Li3.7Zn0.7Ga0.3(PO4)2, measured at RT. The red circles, black continuous line, blue line and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively. Further substitution of Zn with Ga leads to disappearance of the (011) and (013) peaks in the XRD patterns, as evident from Figure II.19b. This confirms the stabilization of the β-polymorph at RT with Ga concentration ≥0.35 mol pfu. The Li3.5Zn0.5Ga0.5(PO4)2 compound produces a XRD pattern fully consistent with the Pnma space group involving a full Li/Zn/Ga disorder, as the high temperature pure Zn compound. In order to deduce a structural model for Ga-substituted β-phase, we performed a combined Rietveld refinement of the NPD and XRD patterns of Li3.5Zn0.5Ga0.5(PO4)2 recorded at RT, as shown in Figure II.21. The patterns could be fitted perfectly in Pnma space group with lattice parameters of a = 10.08917(3) Å, b = 6.51079(2) Å, c = 4.935843(16) Å, and V = 324.16(5) Å 3 . The crystal structure is shown in Figure II.22 and the structural parameters are summarized in the Appendix Table A1.7.

Figure II. 21 .

 21 Figure II.21. Rietveld refinement of the RT XRD (a), and NPD (b) patterns of β-Li3.5Zn0.5Ga0.5(PO4)2. The red circles, black continuous line, blue line and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

Figure II. 22 .

 22 Figure II.22. Crystal structure of β-Li3.5Zn0.5Ga0.5(PO4)2 at RT. The unit cell is shown in black parallelogram. For sake of clarity, only half of the unit cell content (0 ≤ y ≤ ½ in the monoclinic setting) is shown. Li, Zn and Ga are shown in blue, yellow and pink, respectively and PO4 tetrahedra are colored in green.

  β-Li 3.5 Zn 0.5 Ga 0.5 (PO 4 ) 2 P n m a c b possible by quenching. The phase transitions in the Ga-substituted samples are summarized in the temperature-composition phase diagram in Figure II.23b, c.

Figure II. 23 .

 23 Figure II.23. DSC curves and simplified temperature-composition phase diagram of Li4-xZn1-xGax(PO4)2 samples. 'q' and 'sc' denotes the quenched and slow-cooled samples, respectively. (a) DSC curves of the samples, measured in air with a heating and cooling rate of 10 °C/min. The values of x corresponding to the patterns are indicated at the right. (b,c) The phase diagram, deduced from temperature-dependent XRD and DSC experiments (red stars). In (c), k1 and k2 are indicative of superstructure peaks. The black shaded region indicates the presence of secondary phases for Ga contents above 0.5 mol pfu.

Figure II. 24 .

 24 Figure II.24. Ionic conductivities of the Li4Zn(PO4)2 sample. (a) Representative AC impedance spectra of Li4Zn(PO4)2 sample recorded at 300 °C, the fitting has been shown by light green line on the curve. The deconvolution to different contributions, as obtained from the fitting, has been shown in the same curve and also zoomed in the inset. Contributions from the grains and grain boundary, has been indicated by green and red semicircles, respectively, and Warburg tail by black line. Some typical frequencies has also been indicated, the highest two of them are the resonance frequencies corresponding to grains and grain boundaries, respectively. Inset shows the equivalent circuit used to fit the spectra. Arrhenius plot of σtot of Li4Zn(PO4)2 (b) and ball-milled β-Li4Zn(PO4)2 (c); in (b) the blue and red background refers to αand β-Li4Zn(PO4)2 polymorph, respectively.

  The variations of the ionic conductivity for the pristine Li4Zn(PO4)2 over a temperature range are shown in Figure II.24b. There is a linear variation for the ln(T) with temperature, indicative of a thermally activated Arrhenius type behavior σ(T)=A/T•exp(-Ea/kBT), where kB is the Boltzmann constant, A is a pre-exponential factor and T is temperature). The RT b, gb, and σtot values of ~5.5•10 -18 , ~3.5•10 -19 , and ~1.8•10 -19 S•cm -1 , respectively, were deduced by linear extrapolation of the Arrhenius

Figure II. 25 .

 25 Figure II.25. Arrhenius plot ionic conductivity of Li4-xZn1-xGax(PO4)2 samples. The σtot (a), σg (b) and σgb (c) are shown. Values of x are indicated in (a).

  1

  Figure II.26.

Figure

  Figure II.26. DC polarization of β'-Li3.7Zn0.7Ga0.3(PO4)2 sample. (a) DC polarization plots at different potentials and temperatures, and, (b) Arrhenius plot of DC conductivity of the sample.

  Li4-xZn1-xGax(PO4)2 polymorphs are shown in Figure II.27. For -Li4Zn(PO4)2, infinite 2-D Li +

Figure II. 27 .

 27 Figure II.27. BVEL of (a) -Li4Zn(PO4)2, (b) β-Li4Zn(PO4)2 and (c) '-Li3.7Zn0.7Ga0.3(PO4)2 at the percolation energies.The yellow domains indicates the migration paths for Li + in the structure, obtained using an iso-surface value of 1.05, 0.7, and 0.9 eV above the minimum energy, respectively.

  Figure II.28. BVEL of β-Li3.5Zn0.5Ga0.5(PO4)2, at the percolation energies. The yellow domains indicate the migration paths for Li + in the structure, obtained using an iso-surface value of 1.1 eV above the minimum energy.For the high temperature β-Li4Zn(PO4)2 and '-Li3.7Zn0.7Ga0.3(PO4)2 phases, BVEL maps indicate a 3-D Li + percolation at 0.7 and 0.9 eV above the minimum energy, respectively. This comes in no surprise, because, in βand β'-phases, the mixing between Li and Zn/Ga atoms on the same crystallographic positions breaks the layered nature of the α-polymorph, for which Li atoms are confined in the b-c planes. Thus in the βand β'-polymorphs, Li ions can hop from sites to sites threedimensionally. This is consistent with the lower Ea in the βand β'sample compared to the α-phase. Moreover, in β'/βstructures that present the same (PO4) 3-framework, Li3 is the only Li site not being mixed with Zn or Ga which accommodates the vacancies resulting from the charge compensation,(Figure II.27), and BVEL maps indicate that ionic conduction takes place preferentially through this site. Therefore, this explains the little changes in σg values when the structure changes from β' to β-Li3.5Zn0.5Ga0.5(PO4)2.

  Li 3.5 Zn 0.5 Ga 0.5 (PO 4 ) 2 P n m a of disorder in the samples, see Figure II.29.

Figure II. 29 .

 29 Figure II.29. Variation of ionic conductivity and broadening of XRD peak with concentration of Li vacancy. The (100) peak of the XRD pattern at 2θ= 16.3° was chosen to estimate the FWHM. The light blue, green and red background represent α-, β'-, and βpolymorphs, respectively.

Figure II. 30 .

 30 Figure II.30. FESEM images of the cross-section of the sintered pellets.

Figure

  Figure II.31a,b summarizes the evolution of the XRD patterns with temperature of α-Na4Zn(PO4)2. During heating (1 °C/min), a phase transition at 380 -400 °C is observed, characterized by disappearance of many diffraction peaks to keep only a few and transforms to the HT γ-polymorph (Figure II.31a). On cooling, when cooled slowly (1 °C/min), two phase transitions are observed, at ~320 °C and ~260 °C, respectively, and the sample transforms back to the α-polymorph, as observed from the evolution of the XRD patterns (Figure II.31b). Hence on slow-cooling an intermediate βphase is observed, whose pattern only differs from the γphase by the presence of a peak at 2 = 35.8°. This indicates that the γ and β structure are very close, as will be shown later. If the sample is quenched during synthesis, the β phase can be stabilized at RT (Figure S1). However the β-Na4Zn(PO4)2 is a metastable phase, as the phase converts to the thermodynamically stable α-phase upon mild heating above  260 °C, as evidenced in Figure II.31c.

Figure II. 31 .

 31 Figure II.31. Evolution of as-synthesized Na4Zn(PO4)2 with temperature. (a-b) Evolution of the XRD patterns of the α-Na4Zn(PO4)2 sample (prepared by slow-cooling) on heating (a), and following

Figure II. 32 .

 32 Figure II.32. Rietveld refinement of the SXRD pattern of γ-Na4Zn(PO4)2 recorded at 500 °C. For symbol code, refer to the Figure II.31.

Figure II. 33 .

 33 Figure II.33. Combined Rietveld refinement of the SXRD (a), and NPD (b) patterns of β-Na4Zn(PO4)2 recorded at RT. For symbol code, refer to the Figure II.31.

Figure II. 34 .

 34 Figure II.34. Crystal structures of the γ-Na4Zn(PO4)2 (a), and β-Na4Zn(PO4)2 (b). The γunit cell (red line, a) is superposed on the βunit cell (black line, in b) as dashed red line.

  Figure II.35 shows the Le Bail refinement of both SXRD and NPD pattern of α-Na4Zn(PO4)2 with the above mentioned unit cell.

Figure II. 35 .

 35 Figure II.35. Le Bail refinement of the SXRD pattern (a) and the NPD pattern (b) of α-Na4Zn(PO4)2. The red circles, black, and blue lines, and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

  -xZn1-xMx(PO4)2, with M = Ga, and Al. When prepared by quenching the Ga/Al substituted compositions crystalize in the βform (Figure II.36), alike the pristine β-Na4Zn(PO4)2. But interestingly, even when they are prepared by cooling slowly to RT (1 °C/min), the compositions are also stabilized as the β-polymorph, see Figure II.37.

Figure II. 36 .

 36 Figure II.36. XRD patterns of the Na4-xZn1-xGax(PO4)2 synthesized by quenching. The values of x are indicated at the right of the patterns. All the compositions show XRD patterns like the pristine β-Na4Zn(PO4)2 (x = 0) phase.

Figure II. 37 .

 37 Figure II.37. XRD patterns of the Na4-xZn1-xMx(PO4)2 synthesized by slow cooling for M = Ga (a) and Al (b). The values of x are indicated at the right of the patterns. For comparison, the pattern corresponding to the pristine α-Na4Zn(PO4)2 (slow-cooled) is also shown (blue pattern).

Figure II. 38 .

 38 Figure II.38. Le Bail refinement of the SXRD pattern of slow-cooled β-Na3.85Zn0.85Ga0.15(PO4)2 sample. The red circles, black, and blue lines, and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

Figure II. 39 .

 39 Figure II.39. Evolution of XRD patterns of as-synthesized Na4-xZn1-xGax(PO4)2 samples (prepared by slow-cooling) with temperature upon heating (at 1 °C/min), for xGa = 0.15 (a), and xGa = 0.25 (b). The corresponding temperatures are indicated at the right side of the XRD patterns.

Figure II. 40 .

 40 Figure II.40. Arrhenius plot of total ionic conductivity of the Na4Zn(PO4)2 samples.

  Figure II.41 summarizes the total ionic conductivities (σtot) of all Ga-substituted samples and the conductivity values are gathered in Table4. Note that Al-substitution over Ga exhibits similar transport properties and is shown in Appendix. For all the Ga-substituted samples, high ionic conductivities, comparable to the γ-Na4Zn(PO4)2, have been obtained even at RT. However, the Gasubstituted samples (xGa = 0.05 -0.2 mol pfu) show interesting behaviour after slight heating, that the conductivity decreases (blue background in Figure II.41), but at T ≥ 390 °C it again increases. Furthermore, the conductivity transition temperature increases with increasing Ga content, but was not seen for the xGa = 0.25 sample. This behaviour is similar to what was observed for the quenched β-Na4Zn(PO4)2 sample. Hence this behaviour pertains to the γ → α transition, as was confirmed by the temperature dependent XRD measurement for the representative sample with xGa = 0.15 mol pfu (Figure II.39a). From the evolution of the XRD patterns, it is evident that the samples transforms to the α-phase at 280°C, before transforming back to the β-phase. Hence, the decreased conductivity can be attributed to the α-phase. Furthermore, it is interesting to note, that the transition temperature of the γ → α conductivity-transition moves to higher temperature (see blue background in Figure II.41) with increasing Ga-content for the Na4-xZn1-xGax(PO4)2 samples with x = 0.05 -0.2. For the xGa = 0.25 sample, the structural γ → α transition was not observed (Figure II.39b), hence the transition was not observed in ionic conductivity as well. Thus, total ionic conductivity of ~1.1•10 -10 at RT was obtained for the xGa

Figure II. 41 .

 41 Figure II.41. Arrhenius plot of total ionic conductivity of the Na4-xZn1-xGax(PO4)2 samples synthesized by slow cooling. The values of x are indicated in the legends.

Table 4 .

 4 Ionic conductivities and activation energies of the Na4-xZn1-xMx(PO4)2 samples (M = Ga, Al). Conductivities at RT were obtained by extrapolation of the Arrhenius plot. The red, and blue background denotes the Ga, and Al substitution, respectively. observed that, for the Ga-substituted compositions the conductivities in α-form (see blue background in Figure II.41) is higher than the pristine α-Na4Zn(PO4)2 (xGa = 0 in Figure II.41).

Figure II. 42 .

 42 Figure II.42. Strategy to design new ionic conductors. The conductivity plots are of Na4Zn(PO4)2 (in black) and of Na3.75Zn0.75Ga0.25(PO4)2, and shown as representatives.

Figure III. 1 .

 1 Figure III.1. Voltage profiles of TiS2, LiCoO2, and Li1.2Ni0.13Mn0.54Co0.13O2 (Li-rich NMC).

  The crystal structure of TiS3, determined on single crystal, was reported byFuruseth et al, way back in 1975. 257 The reported structures are shown in Figure III.2a,b. It crystallizes in monoclinic crystal structure (space group P21/m), with a = 4.958 Å, b = 3.4006 Å, c = 8.778 Å, and β = 97.32°. The structure consists TiS8 polyhedra linked via shared faces into an infinite chain parallel to the b axis with the neighboring chains offset by b/2 (Figure III.2a,c). The Ti bilayered chains bonded by van der Waals forces along c and connected via longer Ti-S bonds (2.86 Å) along a axis, which makes the structure quasi-layered (Figure III.2b,c), with Ti residing in an irregular eightfold coordination with S atoms (Figure III.2c).

Figure III. 2 .

 2 Figure III.2. Reported crystal structure of TiS3. (a-c) Structural model reported by Furuseth et al in 1975. 257 The structure is drawn along [100] (a), and, [010] (b) direction. Typical bond distances and Ti coordination is shown in (c). Ti and S atoms are shown in blue and yellow, respectively. Ti-S bonds

Figure III. 3 .

 3 Figure III.3. Picture of the as-synthesized TiS3. (a) Synthesized powder and (b) TiS3 crystals. Inset in (b) shows the ampule after the synthesis.

Figure III. 4 .

 4 Figure III.4. SEM images of the pristine TiS3 powders. Images were taken after the powders mixed and hand-grinded with 20 wt% Csp. The amorphous whitish Csp can be seen also in the SEM images.

Figure III. 5 .

 5 Figure III.5. Rietveld refinement of the SXRD pattern of the pristine TiS3 powder. The red circles, black continuous line, blue line, and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

Figure III. 6 .

 6 Figure III.6. Rietveld refinement of the NPD pattern of the pristine TiS3 powder collected at D1b. The red circles, black continuous line, blue line, and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

Figure III. 7 .

 7 Figure III.7. The crystal structure of as-synthesized TiS3. The structure deduced from Rietveld refinement of the NPD pattern (a) is compared with the reported model by Furuseth et al, 1975 257 (b).

Figure III. 8 .

 8 Figure III.8. HAADF-STEM image of TiS3 single crystal showing the atomic arrangement. Twinned nano domains with mirrored orientation of the c-axis are traced with the kinked white line. The inserts show the enlarged images with overlaid structural model obtained from the NPD Rietveld refinement (a) and from Furuseth et al 257 (b). Note that the reported model (b) does not match the HAADF-STEM image. The Ti and S atoms are shown as blue and yellow spheres, respectively.

Figure III. 9 .

 9 Figure III.9. Charge-discharge profile of TiS3 in the 1st cycle. The first Li-intercalation (discharge) and, subsequent second charge and discharge are shown in red, green and blue colors, respectively. The voltage profile are shown in (a) and the corresponding dQ/dV curves are in (b).

Figure III. 10 .

 10 Figure III.10. Operando XRD of TiS3. The corresponding charge (at C/20 rate) and discharge (at C/10 rate) profile is shown in the left panel. Some important compositions are indicated at the right of the patterns. (b) Ex situ SXRD patterns of some composition indicated on the figure. In contrast to the layered Li2TiS3 phase, 272 the crystal structure of this electrochemically prepared Li2TiS3 has never been reported in literature, hence our attempt to solve it by combined TEM and SXRD studies. Firstly, TEM studies were performed on Li2TiS3 crystals that were recovered from discharged TiS3/Li cell (i.e., few TiS3 crystals were cycled, without Csp additive, at a rate of C/50). The HRTEM images are shown in Figure III.11. We noticed significant exfoliation of the TiS3 crystals upon lithiation (Figure III.11a). HAADF-STEM image of Li2TiS3 shows that the crystal structure is retained, but the gap between the Ti-S slabs disappears (along a direction, compare Figure III.11b). However, we could reconstruct a reciprocal space by rotating the crystals around [010], and then it could be indexed with a primitive monoclinic unit cell with a = 6.35Å, b = 3.51Å, c = 17.87Å, β = 91.2°.

Figure III. 11 .

 11 Figure III.11. HRTEM images of Li2TiS3 single crystals. (a) Low magnification HAADF-STEM image, and (b) HAADF-STEM image of the Li2TiS3.

Figure III. 12 .

 12 Figure III.12. SEM images of the ex situ Li2TiS3 (discharged TiS3) powder. Images were taken after washing and drying of the cycled sample. The amorphous whitish region corresponds to the Csp that was added prior to cycling.

Figure III. 13 .

 13 Figure III.13. Le Bail refinement of ex situ Li2TiS3 powder. The red circles, black continuous line, blue line, and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively.

Figure III. 14 .

 14 Figure III.14. HAADF-STEM image of Li~0.2TiS3 after 1st cycle.

Figure III. 15 .

 15 Figure III.15. S 2p and Ti 2p core XPS spectra of the cycled TiS3 powder samples. The corresponding voltage profile during cycling is shown in the top panel. In (a), the spectra are deconvoluted by fitting to the doublets as indicated on the figure. In (b), the spectra shows traces of Ti-O, which is probably due to the reactivity with electrolyte.

Figure III. 16 .

 16 Figure III.16. EELS spectra of the S L2,3 and Ti L2,3 edge of the cycled TiS3 samples. The corresponding voltage profile during cycling is shown in Figure III.15.Turning to the Ti L-edge EELS of TiS3 powder (Figure III.16b), we found that the spectra is in good agreement with the previously reported Ti L-edge XANES (X-ray absorption near edge structure)

Figure III. 17

 17 Figure III.17. S K-edge TFY-XANES spectra collected for TiS3 and LiTiS3, (mid-discharged), Li2TiS3 (full-discharged), and Li~0.2TiS3 (fully-charged) samples.

Figure III. 18 .

 18 Figure III.18. Charge-discharge profile of TiS3 over first 25 cycles. The TiS3 (powder) was cycled in Swagelok cell as powder (hand-ground with 20 wt% Csp) at a rate of C/20.

Figure III. 19 .

 19 Figure III.19. Voltage profile of TiS3 recorded with a GITT protocol. The TiS3 was cycled in a twoelectrode cell at a rate of C/20 with relaxation time indicated in the figure. (a-b) Voltage profile in first 3 cycles with varying relaxation time. (c-f) Voltage profile during the relaxation indicated by boxes with dashed in a,b. Furthermore, a huge polarization of ~0.6 mV in the charge-discharge profile reveals that the Li uptake/removal in TiS3 lead to poor energy efficiency. To understand whether the polarization is nested in kinetic limitation, we performed galvanostatic intermittent titration technique (GITT) of TiS3. The voltage profile of the first 3 cycles recorded with GITT protocol (3 h pulse at C/20 followed by relaxation) is shown in Figure III.19. After a rest of 1 h a large hysteresis still remains as shown in the 1st cycle (Figure III.19a). The relaxation time was increased gradually in the 2nd and 3rd cycle (Figure III.19b), but even after rests of 4 h (2nd cycle) or 10 h (3rd cycle) the hysteresis does not improve. Typical voltage profile during the rest period is shown in Figure III.19c-e, which shows a never-ending voltage relaxation even for rest periods of 10 h. Therefore, the anionic redox process in TiS3 has a large time constant and is inherently limited thermodynamically. 278

  could be refined in the in C2/m space group alike for Li2MnO3. The Rietveld refinement of its synchrotron XRD (SXRD) pattern is shown in Figure III.20a, with the obtained parameters summarized in the Appendix Table A2.2. This crystal structure of Li2TiS3 is shown in Figure III.20b, that indicates the similarity to honeycomb-ordered structure reported for Li-rich layered oxides family. However with an expectedly expansion of unit cell in Li-rich sulfides than in oxides, to accommodate the bulkier S atoms.

Figure III. 20 .

 20 Figure III.20. Crystal structure of Li2TiS3 from Rietveld refinement of the SXRD pattern. (a) The Rietveld refinement. The red circles, black continuous line, blue line, and green tick bars represent the

  8 V and 3 V at a rate of C/20 and the voltage profiles are summarized in Figure III.21.

Figure III. 21 .

 21 Figure III.21. The electrochemical performance of Li2TiS3. (a) The voltage profile, and, (b) corresponding dQ/dV curves, are shown here. The curves in lighter color denote the 2nd cycle.

Figure III. 22 .

 22 Figure III.22. Fe-substitution in Li2TiS3. Schematic band structure of Li1.33Ti0.67S2 (i.e., Li2TiS3, in left), and its Fe 2+ substituted derivative (in right). The relative band positions are estimated based on Ref 24 . The label nb stands for non-bonding.

  details). The lab X-ray diffraction (XRD) patterns are gathered in Figure III.23a for all the compositions. For more accurate characterization, SXRD patterns were also collected at the 11-BM in transmission mode with λ = 0.41 Å and the patterns are gathered in Figure III.23b.

Figure III. 23 .

 23 Figure III.23. Structural behavior of the Li1.33-2y/3Ti0.67-y/3FeyS2 series. (a) XRD patterns, and (b) SXRD patterns of the as-synthesized Li1.33-2y/3Ti0.67-y/3FeyS2 samples. (c) Variation of lattice parameters (obtained from the Rietveld refinement of the SXRD patterns). For the y = 0 composition, the lattice parameters have been converted to the R𝟑 ̅ m description. Upon increasing the Fe content y, there is a progressive shift of the Bragg peaks (Figure III.24a) indicating the existence of a solid solution. Moreover, Fe-containing compositions (y = 0.1 -0.5) do not show the superstructure peaks, most likely because Fe 2+ disrupts the honeycomb Li + /Ti 4+ ordering.

  particles (shown in the Appendix Figure A2.4). We performed the Rietveld refinement of its SXRD pattern (Figure III.24b), which confirms that Fe, Ti and Li occupy same site in the metal layer (see the structural model in the Appendix). This structure was further confirmed by Rietveld refinement of its neutron powder diffraction (NPD) pattern also (see Figure III.24c).

Figure III. 24 .

 24 Figure III.24. Structural characterization of the Li1.13Ti0.57Fe0.3S2 (y = 0.3) sample. (a) Progressive shift of the peaks in SXRD patterns of the Li1.33-2y/3Ti0.67-y/3FeyS2 samples. (b) Rietveld refinement of the SXRD pattern of the Li1.13Ti0.57Fe0.3S2 sample. (c) Rietveld refinement of the NPD pattern of Li1.13Ti0.57Fe0.3S2 using the model deduced from the SXRD pattern. The refinement returns the cell parameters of a = 3.53563(4) Å, b = 18.08630(38) Å, and V = 195.801(5) Å 3 with GoFs of χ 2 = 1.6, Rwp = 0.355%, and RB = 1.15%. Few 2θ regions have been excluded to avoid peaks from the vanadiumcontainer. The red circles, black continuous line, blue line, and green tick bars represent the observed, calculated and difference patterns, and Bragg positions, respectively. The structures of the Li1.13Ti0.57Fe0.3S2 is compared with the Li-rich NMC (LrNMC) in Figure III.25, which indicates the similarity of the structures of Li-rich sulfides and oxides. In both cases the excess Li's are mixed and disordered with the transition metals in the metal layer (MX2 layer, X = S,

Figure III. 25 .

 25 Figure III.25. Comparing the crystal structures of Li1.13Ti0.57Fe0.3S2 (a) and Li-rich NMC (b).

Figure III. 26 .

 26 Figure III.26. Electrochemical behavior of Li1.33-2y/3Ti0.67-y/3FeyS2. Voltage profiles of the compositions obtained over cycling vs Li at C/20 for the first two cycles. The curves in lighter colors denote the second cycles. The arrows marked by 'S' and 'P' denote the slope and the plateau, respectively (d) Performance of the Li1.13Ti0.57Fe0.3S2 (y = 0.3) composition over cycling at a rate of C/2.5 (except the first formation cycle at C/20). The capacity obtained for all the compositions are summarized in Figure III.27a, which shows increase in capacity up to y = 0.3, and then decreases simply due to the lower amount of available Li. The maximum capacity of ~245 mAh•g -1 was obtained for the composition Li1.13Ti0.57Fe0.3S2 (y = 0.3). The maximum capacities available from the cationic Fe 2+/3+ redox for the compositions are shown in Figure III.27a (purple line), which constitute parts of the measured capacities. Even by assuming full utilization of the multi-electron oxidation of Fe 2+ to Fe 4+ (blue dotted line in Figure III.27a), which is quite unlikely to occur in sulfide frameworks, we cannot account for all the measured capacity, hence implying the activity of the anionic network. Lastly, the corresponding dQ/dV profiles are shown in

Figure III. 27 .

 27 Figure III.27. Capacity and dQ/dV of Li1.33-2y/3Ti0.67-y/3FeyS2. (a) Theoretical capacity (considering total Li-removal) and the actual discharge capacity obtained for the compounds over cycling at C/20. The capacity expected from cationic Fe 2+/3+ and hypothetical Fe 2+/4+ redox is also shown. (b) dQ/dV curves obtained in the first cycle for the various compositions.

Figure III. 28 .

 28 Figure III.28. More details of the electrochemical properties of Li1.13Ti0.57Fe0.3S2 (y = 0.3). (a) Evolution of the charge-discharge profile in first 7 cycles, at rates as indicated in the figure. The inset shows the zoom of the 2nd cycle, illustrating the overall voltage gap. (b) Cycling performance in Lihalfcell (at a rate indicated). The materials were cycled in Swagelok type cells, with samples as powders (mixed with 20 wt% Csp) only, without any further optimization. Inset shows zoom on the dQ/dV curve in the 1st cycle, illustrating the slope (S) and the plateau (P) features. The voltage profile of the Li1.13Ti0.57Fe0.3S2 is shown in Figure III.28a. Note that the voltage profile changes after the 1st cycle and then stable cycling has been observed. 2nd cycle onward, the polarization reduces to ~100 mV (inset, Figure III.28a), instead of ~220 mV during the first cycle. This indicates that the first cycle acts as an 'activation' cycle alike the Li-rich NMC oxides. The dQ/dV profile

Figure III. 29 .

 29 Figure III.29. Cycling Li1.13Ti0.57Fe0.3S2 in a broader voltage range. (a) Voltage profiles of the Li1.13Ti0.57Fe0.3S2 /Li half-cell cycled with a cutoff potential of 4.5 V (at a rate of C/20). Note that, even the cutoff potential was never reached because of sample decomposition, or parasitic reactions etc. (b)

Figure III. 30 .

 30 Figure III.30. Probing the Li (de)intercalation in operando. (a) Evolution of the in situ XRD patterns in the first cycle of a Li1.13Ti0.57Fe0.3S2 /Li half-cell at a rate of C/20. The left panel shows the corresponding voltage profile and the right panel shows the evolution of the unit cell volume (V), as obtained from the Rietveld refinements of the XRD patterns. Note that the dotted lines represent extrapolations, where the XRD patterns could not be refined because of poor intensity of the peaks. (bc) Rietveld refinement of the SXRD patterns (b) and the HAADF-STEM image (c) of the full-charged phase. The peaks denoted by * in (b) show minor Li2TiO3 impurity. (d) The bond-lengths observed in the pristine, full-charged and full-discharged phase (as obtained from the Rietveld refinement of the corresponding SXRD patterns).

Figure III. 31 .

 31 Figure III.31. 2nd cycle of the in situ cycling of Li/Li1.13Ti0.57Fe0.3S2 half-cell. The evolution of the XRD patterns during the in situ cycling of Li/Li1.13Ti0.57Fe0.3S2 half-cell at a rate of C/20 with the evolution of the unit cell parameters a, c, and V are also shown.

Figure III. 32 .

 32 Figure III.32. Structural characterization of Li1.13Ti0.57Fe0.3S2 after Li (de)intercalation. (a-c) Rietveld refinement of the SXRD pattern (a), the HAADF-STEM image (b), and, SAED patterns of the fully-charged sample along the [010] axis (c) of the full-charged phase. The inset in (c) shows the HAADF intensity profile. (d) Rietveld refinement of the SXRD pattern of the discharge phase of the Li1.13Ti0.57Fe0.3S2. The red circles, black continuous line, blue line, and green tick bars represent the observed, calculated and difference patterns, and Bragg positions, respectively. The peaks denoted by * are from minor Li2TiO3 impurity.

Figure III. 33 .

 33 Figure III.33. The bond-lengths observed in the pristine, full-charged and full-discharged phase (as obtained from the Rietveld refinement of the corresponding SXRD patterns). Some typical bond-lengths in TiS3 and FeS2 are also given for comparison (adapted from Ref 257 , and Ref 287 ).

  ), and thus highlights the presence of a distribution of high-spin (HS) Fe 2+ , as already encountered for FeS and various other Fe(II) -based sulfides. 289,290 The necessity of four doublets is simply due to different local arrangements of disordered Ti/Li/Fe cations around a given Fe site (shown in Figure III.35).

Figure III. 34 .

 34 Figure III.34. Operando Mössbauer spectroscopy to identify the Fe 2+/3+ redox process in Li1.13Ti0.57Fe0.3S2. (a) Typical Mössbauer spectra and their deconvolution for the pristine (comp. 1), comp. 2, comp. 3 (fully charged) and the fully-discharged product. (b) Contour plot of evolution of the Mössbauer spectra collected during in situ cycling of a Li1.13Ti0.57Fe0.3S2/Li half-cell at a rate of C/70. For reference, the voltage profile is shown in the top panel. The lowest panel shows the evolution of the reconstructed components during the cycling.

Figure III. 37 .

 37 Figure III.37. Probing the cationic redox processes in Li1.13Ti0.57Fe0.3S2 by ex situ XAS. TFY-XANES spectra at the Fe L2,3-edge (b), and at the Ti L2,3-edge (c). A representative is shown in (a), to indicate the SoCs of the ex situ samples.

Figure III. 38 .

 38 Figure III.38. Probing the anionic redox processes in Li1.13Ti0.57Fe0.3S2. S K-edge TFY-XANES spectra (b), and S 2p core XPS spectra (c) of the ex situ samples. A representative is shown in (a), to indicate the SoCs of the ex situ samples. In (c), another extra doublet at a higher binding energy (163.3

Figure III. 39 .

 39 Figure III.39. Probing S L-edge to identify the S-redox processes. The EELS spectra of the S L-edge (middle panel), mRIXS of the S L-edge with integrated PFY spectra on the bottom (right panel). The voltage profile is shown in the left panel to indicate the ex situ samples.

Figure

  Figure III.40. S-L2,3 edge EELS of FeS (a), and FeS2 and elemental S (b). The data have been adopted from References 277,292 .

Figure III. 41 .

 41 Figure III.41. Comparison of the voltage profiles of TiS2 and Li1.13Ti0.57Fe0.3S2 (green curve).

Figure III. 42 .

 42 Figure III.42. Electrochemical behaviour of the Li1.2Ti0.8S2 sample. (a) The voltage profile, and (b) the corresponding dQ/dV curves, of the composition. Cationic and anionic plateaus are indicated by C and A, on the figure.

  .The electrochemical behaviour for both compositions are shown in Figure III.43. Note that, much lesser capacity (up to 190 mAh•g -1 , with specific energy of ~420 Wh•kg -1 , for y = 0.2) could be obtained than the Fe-substituted samples. Furthermore, irrespective of Co-content, ~0.4 Li always remains inaccessible in the fully-oxidized samples. Interestingly, the voltage of dQ/dV peaks remains invariant amongst Co-substituted samples, in contrast to the Fe-substitution.

Figure III. 43 .

 43 Figure III.43. Electrochemical activity of the Li1.33-2y/3Ti 0.67-y/3CoyS2 samples. (a,c) Voltage profiles of the compositions obtained over cycling vs Li at C/20 for the first two cycles. The curves in lighter colors denote the second cycles. (b,d) the corresponding dQ/dV curves. Note that the y-axis break was provided to enlarge the peaks at 2nd cycle.

  Figure IV.1. GITT Voltage profile of Li1.13Ti0.57Fe0.3S2 and Li-rich NMC in a three-electrode cell. (a) The voltage profile was recorded in the 8th cycle (once the voltage profile stabilizes after 7 cycles), at C/5 rate (i.e., current density of 46 mA•g -1 ) with 30 min rests for equilibration. Inset shows the OCV profile during the relaxation period after the indicated step. Sample was used as a self-standing electrode in three-electrode Swagelok type cell. (b) GITT Voltage profile of Li-rich NMC recorded in threeelectrode Swagelok cell (in 4th cycle, with 40 mA•g -1 pulses and 4 h rests for equilibration).

Figure IV. 2 .

 2 Figure IV.2. GITT Voltage profile of Li2Ru0.75Sn0.25O3 (LRSO). GITT was recorded with a C/25 pulse for 90 min, and 10 h rest (in the 2nd cycle). Note that even after 10 h rest, ~100 mV hysteresis still remains. Figure adapted from reference 49 .

Figure IV. 3 .

 3 Figure IV.3. GITT Voltage profile of LiTi0.5Fe0.5S2 (y = 0.5) in a two-electrode cell. The voltage profile was recorded once the voltage profile stabilizes after 7 cycles, at C/5 rate with 30 min rests for equilibration. Sample was used as powder in two-electrode Swagelok type cell.

Figure IV. 4 .

 4 Figure IV.4. Voltage hysteresis in Ti 3+ and Co 2+ substituted Li2TiS3. (a) Voltage profile of Li1.2Ti 4+ 0.4Ti 3+ 0.4S2 recorded with a GITT protocol. Experiment was performed with the sample as powder in a two-electrode Swagelok type cell in the 8th cycle, at C/5 rate with 30 min rests for equilibration. Inset shows the OCVs during the relaxations. (b) Voltage profile of Li1.2Ti 4+ 0.6Co 2+ 0.2S2

  and the results are summarized in Figure IV.5c. It is clear that the charging voltage is always maintained irrespective of the discharge window. The corresponding dQ/dV curves (Figure IV.5d) also reveal that all the capacity obtained from discharge (blue arrows) are recovered by the 2.55 V peak during charge (red arrow) without creating any new peaks.

Figure IV. 5 .

 5 Figure IV.5. Probing the hysteresis of Li1.13Ti0.57Fe0.3S2 by voltage window opening experiment. (a) Voltage profiles with corresponding dQ/dV curves (b) during charge opening. (c) Voltage profiles with corresponding dQ/dV curves (d) during discharge opening.

Figure IV. 6 .

 6 Figure IV.6. Anionic redox in Li-rich NMC. (a) Voltage profiles and (b) corresponding dQ/dV curves of Li1.2Ni 2+ 0.13Mn 4+ 0.54Co 3+ 0.13O2 (Li-rich NMC), obtained as the charge window is opened stepwise from 2.0 to 4.8 V. In (b), the arrows indicate the splitting of the capacities obtained by charging above ~4.1 V. 50 Again we performed similar voltage window opening experiments for the Li1.2Ti 4+ 0.4Ti 3+ 0.4S2. During the charging window opening (Figure IV.7a and Appendix Figure A3.4), hysteresis was not increased in the sloppy voltage region. However, in the high-voltage plateau, the polarization increases slightly compared to 1st plateau, but deeper oxidation does not lead to a noticeable increase in the hysteresis. Thus irrespective of the charge window, the discharge voltage is never lowered noticeably.

Figure

  Figure IV.7. Probing the hysteresis of Li1.2Ti 4+ 0.4Ti 3+ 0.4S2 by voltage window opening experiment. (a) Voltage profiles and (b) corresponding dQ/dV curves during charge opening and (c) Voltage profiles and (d) corresponding dQ/dV curves during discharge opening.

Figure IV. 8 .

 8 Figure IV.8. Cell's electrochemical resistance as a function of SoC. The resistance (during charging) was estimated by Ohm's law from the voltage drop in first 10 s of rest from GITT experiment in a twoelectrode cell (shown in the Appendix).

Figure IV. 9 .

 9 Figure IV.9. Typical Nyquist plot of EIS of Li1.13Ti0.57Fe0.3S2 in Li-half cell. THE EIS was collected at full-discharged state after first cycle. The sample was used as a self-standing electrode. Observed data points are shown by black circles and fitting of the data are indicated by red line. The equivalent circuit used is shown above. The deconvoluted MF semicircle (corresponding to the Rct) and LF Warburg tail is illustrated by red semicircle and blue dotted line, respectively.

Figure IV. 10 .

 10 Figure IV.10. Nyquist plots of EIS recorded after each relaxation in GITT. EIS was recorded for Li1.13Ti0.57Fe0.3S2 (a), and LiTi0.5Fe0.5S2 (b) during charge and discharge. Experiment was done in a twoelectrode Swagelok cell and the sample was used as powder (hand-ground with 20 wt% Csp).

  Figure IV.12. This again positively contrasts with Li-rich NMC that shows a nearly continuous voltage fade upon cycling with an accumulated drop of ~150 mV after 60 cycles (Figure IV.12).

Figure IV. 11 .

 11 Figure IV.11. Discharge profiles of Li1.13Ti0.57Fe0.3S2 over long cycling to show the voltage fade. The sample was cycled as powder in Swagelok type cells with sample as powder, at C/2.5 rate. Note that the voltage profile almost stabilizes after the 7th cycle.

Figure IV. 12 .

 12 Figure IV.12. Average voltage during long cycling of Li1.13Ti0.57Fe0.3S2 and Li-rich NMC in Li-half cells.In each cycle, the average voltage is defined as the mean of the average charge and discharge voltages that and was calculated by dividing the energy with the capacity obtained. Data for Li-rich NMC is taken from reference 305 .

Figure IV. 13 .

 13 Figure IV.13. Average voltage of Li-rich sulfides in Li-half cells.In each cycle, the average voltage is defined as the mean of the average charge and discharge voltages that was calculated by dividing the energy with the capacity obtained. Samples were cycled in Swagelok type cells, with samples as powders.

Figure IV. 14 .

 14 Figure IV.14. The voltage profiles in 2nd cycle of (a) Na0.67[Mg0.28Mn0.72O2], and (b) O3-Na2TiS3. In (a) cationic Mn 3+/4+ redox is indicated by C and anionic O 2-/O2 n-by A, and in (b) by anionic S 2-/S2 n- is indicated by A. Cycling rate of 10 mAh•g -1 was used. Data for (a) is taken from Reference 44 .Next, to check the capacity fade, we plotted the capacity retention of the O3 and P2 phases in Figure IV.15. The O3-Na2TiS3, in the 1st cycle, shows a discharge capacity of >240 mAh•g -1 , leading to a decent discharge energy density of ~450 Wh•kg -1 (at material level). Furthermore, the phase shows a decent capacity retention, as shown in the cycling of first few cycles, even though the cycling condition was not optimized. Initially it shows slight capacity decay, but the capacity stabilizes at ~180 mAh•g -1 after first 10 cycles. The P2-Na2TiS3, despite having similar capacity in 1st cycle alike the O3-phase, exhibits rapid capacity fading. Nevertheless, the energy density (at material level) obtained for O3-Na2TiS3 is comparable to NVPF (~500 mAh•g -1 ).308

Figure IV. 15 .

 15 Figure IV.15. Capacity retentions of the O3-and P2-Na2TiS3. The samples (hand-ground with 20 wt% Csp) were cycled as powder in Swagelok cells without further optimization.

Figure IV. 16 .

 16 Figure IV.16. Average discharge voltages obtained in O3-and P2-Na2TiS3 over first few cycles.Average voltage here is defined by the ratio of energy density and the capacity obtained in each cycle.

Figure IV. 17 .

 17 Figure IV.17. Ragone plots of Li1.13Ti0.57Fe0.3S2. The sample was cycled as a self-standing electrode to minimize the Li-diffusion limitation inside the active material. Li-rich NMC, Li-stoichiometric NMC 811 (LiNi0.8Mn0.1Co0.1O2), and LiCoO2 are also included for comparison. Only values at material-level are considered and values for the NMC-811 and the Li-rich NMC are adapted from Reference 37 .

  of this chapter includes the following publication: Marchini and Saha et al., manuscript under preparation.

Figure V. 1 .

 1 Figure V.1. Schematic diagram of cathode/sulfide electrolyte interfaces with oxide or sulfide cathode. Figure adapted from reference313 .The 𝝁 𝐋𝐢 + is the chemical potential of Li ion, and eϕ is its electrostatic energy (e is elementary charge and ϕ is local electrostatic potential). The 𝝁 𝐋𝐢 + in the bulk of sulfide electrolytes is relatively high, because of Li ions weakly bonded to the anionic framework. High ϕ at the surface of oxide cathode materials lowers the 𝝁 𝐋𝐢 + at the interface, leading to decrease of Li ion concentration.313

( 4 )

 4 Additional capabilities such as controlling/monitoring temperature and pressure. A simple testing setup was designed and fabricated by other people and engineers in our lab-Florencia Marchini, Marc Duchardt, Daniel Alves Dalla Corte, and Thomas Dargon (Figure V.2). It consists of a core made of PEEK (polyether ether ketone) polymer, that is very tough, and rigid polymer and that can bear high mechanical and tensile strength. Furthermore, it is stable up to high temperature (260 °C) while showing no water absorption.

Figure V. 2 .

 2 Figure V.2. A picture of the ASSB setup fully designed and assembled at lab, that was used in this study. Further details can be found at https://sphere-energy.eu.

Figure V. 3 .

 3 Figure V.3. Schematics of the three-layer ASSB.

Figure V. 4 .

 4 Figure V.4. EIS spectra of the glassy 75•Li2S-25•P2S5 (g-LPS) and β-Li3PS4 (LPS) solid electrolyte. See the Appendix for details of the measurement protocol. The EIS was collected on sample-pellets, with 25 mV bias in the frequency range of 0.5 MHz -1 Hz.

  of ~0.2 mS•cm -1 at RT (Figure V.4b). The purity of the β-Li3PS4 was further confirmed from the Rietveld refinement of its synchrotron XRD pattern (Figure V.5).

Figure V. 5 .

 5 Figure V.5. Rietveld refinement of the SXRD pattern of the as-synthesized β-Li3PS4 SE. The red circles, black continuous line, blue line, and green tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively. The SXRD pattern was refined in Pnma space group with unit cell parameters a = 12.97721(27) Å, b = 8.03258(20) Å, c = 6.12629(13) Å, and V = 638.608(25) Å 3 , in agreement with the reported structural model. 143 Since spontaneous chemical reactivity of sulfide-based SEs remains an issue in ASSBs, 327 we first checked the chemical compatibility of β-Li3PS4 with the Li-rich sulfide Li1.13Ti0.57Fe0.3S2 (LTFS). For this purpose, we have collected SXRD pattern after annealing the cathode composite at 200 °C for 2 days. Even after the heat treatment, the SXRD pattern did not show any trace of impurity phase (Figure V.6), thus ruling out any chemical reaction between them.

Figure V. 6 .

 6 Figure V.6. Rietveld refinement of the SXRD pattern of the cathode composite. The red circles, black continuous line, blue line, and tick bars represent the observed, calculated, and difference patterns, and Bragg positions, respectively. The red and green tick bars represent the Bragg positions of LTFS

Figure V. 7 .

 7 Figure V.7. DC polarization plots of LTFS-LPS cathode composite. (a) The steady state current (I, in left axis) obtained by applying DC voltages (E, in right axis) to the sample pellet (thickness of 1.5 mm) placed between two steel electrodes. (b) Variation of I with E, which was used to deduce the resistance (by Ohm's law). The electronic conductivity (σel) was calculated from the resistance, according to the formulae σ = (d/A)•(1/R) and plotted in right axis.

  8 -3.6 V vs Li + /Li at a rate of C/30. The voltage profile is shown in Figure V.8.

Figure V. 8 .

 8 Figure V.8. Voltage profile of Li1.13Ti0.57Fe0.3S2 (LTFS) cycled in ASSB at RT. 75•Li2S-25•P2S5 glass (g-LPS) was used as solid electrolyte and In metal was used as anode. Cycling rate of C/30 was used.

Figure V. 9 .

 9 Figure V.9. Polarization in charge-discharge of Li1.13Ti0.57Fe0.3S2│g-LPS│In ASSB. The ASSB was cycled at C/30 rate. Note that in the first charge, the voltage drop at x = ~0.5 is due to a relaxation (at OCV) of 20 mins.

Figure V. 10 .

 10 Figure V.10. Voltage profile of Li1.13Ti0.57Fe0.3S2 cycled in ASSB at RT using β-Li3PS4 as solid electrolyte. The temperatures of cycling are indicated on the plot. The temperature were increased in for the same cell, in subsequent cycles. In metal was used as anode. Cycling rate of C/30 was used. ASSB: courtesy of Florencia Marchini.

Figure V. 11 .

 11 Figure V.11. EIS characterization of the Li1.13Ti0.57Fe0.3S2│g-LPS│In ASSB. (a) EIS spectra (enlarged) of the ASSB collected at various SoCs during the first two cycle. The abbreviations Pr, Ch, and Dis stands for pristine, full-charged, and full-discharged, respectively. Inset shows the whole spectra. (b) Fitting of the EIS for the pristine state. Inset shows the whole spectra. The equivalent circuit used for fitting is also shown in top left corner.

  . Structural Parameters for ζ-Li6Zn(P2O7)2, deduced from the Rietveld refinement of the NPD recorded at 600°C. Isotropic temperature factors (Biso) were constraint to be equal for a same chemical species. ζ-Li6Zn(P2O7)2, at 600 °C space group: C 2/c a = 13.24574(6) Å, b = 5.19764(2) Å, c = 8.92388(4) Å,  = 105.1842(4

  α-Li4Zn(PO4)2, RT Space group P 21/c a = 8.11990(16) Å, b = 10.24192(13) Å, c = 8.10549(15) Å,  = 104.8589(9

6 . 7 .

 67 Figure I.1. Schematic configuration of Li-metal and Li-ion batteries. .....................................................................

  Figure I.1. Schematic configuration of Li-metal and Li-ion batteries. .....................................................................

Figure I. 2 .

 2 Figure I.2. Comparison of band positions and voltage profiles in LiMS2 sulfides and LiMO2 oxides. ...................

Figure I. 3 .

 3 Figure I.3. Improving the specific capacity of layered oxides cathode materials. ...................................................

Figure I. 4 .

 4 Figure I.4. Illustrating cationic and anionic redox. ................................................................................................

Figure I. 5 .

 5 Figure I.5. Voltage profiles of some Li-rich oxides. ..............................................................................................

Figure I. 6 .

 6 Figure I.6. Schematic band diagrams that governs the anionic redox. ...................................................................

Figure I. 7 .

 7 Figure I.7. Moving from Li-rich layered oxides to sulfides. ..................................................................................

Figure I. 8 .

 8 Figure I.8. Reported pictures of LIBs causing fire in electronic devices and vehicles. .........................................

Figure I. 9 .

 9 Figure I.9. Battery technologies using solid electrolyte. ........................................................................................

Figure I. 10 .

 10 Figure I.10. Minimum thickness required for a SE to match the resistance of liquid electrolytes. ........................

Figure I. 11 .

 11 Figure I.11. Heat content (left y-axis) and ionic conductivity (right y-axis) of PbF2. ............................................

Figure I. 12 .

 12 Figure I.12. Development of solid electrolytes towards superionic conductivity. .................................................

Figure I. 13 .

 13 Figure I.13. Simplified schematics illustration of ionic conduction process in liquid (a) and solid (b) electrolytes. 126,165 ....................................................................................................................................................

Figure I. 14 .

 14 Figure I.14. Illustration of Li-conduction in Li10GeP2S12 (LGPS). ........................................................................

Figure I. 15 .

 15 Figure I.15. Ionic conductivity of Li4-3xAlxSiO4. ...................................................................................................

Figure I. 16 .

 16 Figure I.16. Crystal structures of LLZO. O atoms are not shown. .........................................................................

Figure I. 17 .

 17 Figure I.17. Ionic conductivity of Li3PS4. ..............................................................................................................

Figure I. 18 .

 18 Figure I.18. Enhancement of ionic conductivity by moving from oxides to sulfides. ...........................................

Figure I. 19 .

 19 Figure I.19. Structural and configurational features of superionic conductors. .....................................................

Figure I. 20 .

 20 Figure I.20. Comparison of advantages/disadvantages between oxide and sulfide solid electrolytes. ..................

Figure I. 21 .

 21 Figure I.21. Electrochemical stability and decomposition of solid electrolytes. ....................................................

Figure I. 22 .

 22 Figure I.22. RT sinterability of sulfide SEs. ..........................................................................................................

Figure I. 23 .

 23 Figure I.23. Performances of ASSBs with different sulfide SEs. ..........................................................................

Figure II. 1 .

 1 Figure II.1. Phase diagram of the Li(Na)-Zn-P-O System. ....................................................................................

Figure II. 2 .

 2 Figure II.2. Probing temperature-dependent structural modifications by DSC and SXRD. ..................................

Figure II. 3 .

 3 Figure II.3. Evolution of the SXRD patterns during cooling of the Li6Zn(P2O7)2 sample. ....................................

Figure II. 4 .

 4 Figure II.4. Combined Rietveld refinement of the XRD (a) and NPD patterns (b) of α-Li6Zn(P2O7)2. .................

Figure II. 5 .

 5 Figure II.5. Rietveld refinement of the XRD and NPD patterns of ζ-Li6Zn(P2O7)2, recorded at 600 °C. ..............

Figure II. 6 .

 6 Figure II.6. Crystal structure of the polymorphs of Li6Zn(P2O7)2. ........................................................................

Figure II. 7 .

 7 Figure II.7. AC impedance spectra of the Li6Zn(P2O7)2. .......................................................................................

Figure II. 8 .

 8 Figure II.8. Arrhenius plot of Ionic Conductivity of the Li6Zn(P2O7)2 sample. .....................................................

Figure II. 9 .

 9 Figure II.9. AC impedance spectra of α-Li6Zn(P2O7)2 recorded at high temperatures. .........................................

Figure II. 10 .

 10 Figure II.10. BVEL map of (a) α-Li6Zn(P2O7)2 and (b) ζ-Li6Zn(P2O7)2 at the percolation energies. ....................

Figure II. 11 .

 11 Figure II.11. Rietveld refinement of the XRD and NPD patterns of α-Li4Zn(PO4)2, measured at RT. .................

Figure II. 12 .

 12 Figure II.12. Crystal structure of α-Li4Zn(PO4)2. (a) along 𝟏𝟎𝟏, (b) along [001]. ................................................

Figure II. 13 .

 13 Figure II.13. Probing temperature-dependent structural modifications of Li4Zn(PO4)2 by DSC and XRD...........

Figure II. 14 .

 14 Figure II.14. Evolution of XRD patterns of Li4Zn(PO4)2 with temperature during cooling. .................................

Figure II. 15 .

 15 Figure II.15. Rietveld refinement of the XRD and NPD patterns of β-Li4Zn(PO4)2, measured at 600°C. ............

Figure II. 16 .

 16 Figure II.16. Crystal structure of (a) α-Li4Zn(PO4)2, (b) β-Li4Zn(PO4)2 and (c) β'-Li3.7Zn0.7Ga0.3(PO4)2. ............

Figure II. 17 .

 17 Figure II.17. Rietveld refinement of the XRD pattern of the ball-milled -Li4Zn(PO4)2 sample, measured at RT. ...............................................................................................................................................................................

Figure II. 18 .

 18 Figure II.18. SXRD patterns of the Li4-xZn1-xGax(PO4)2 samples. .......................................................................

Figure II. 19 .

 19 Figure II.19. XRD patterns of the Li4-xZn1-xGax(PO4)2 samples. ...........................................................................

Figure II. 20 .

 20 Figure II.20. Rietveld refinement of (a) XRD, and (b) NPD patterns of β'-Li3.7Zn0.7Ga0.3(PO4)2, measured at RT. ...............................................................................................................................................................................

Figure II. 21 .

 21 Figure II.21. Rietveld refinement of the RT XRD (a), and NPD (b) patterns of β-Li3.5Zn0.5Ga0.5(PO4)2. .............

Figure II. 22 .

 22 Figure II.22. Crystal structure of β-Li3.5Zn0.5Ga0.5(PO4)2 at RT. ............................................................................

Figure II. 23 .

 23 Figure II.23. DSC curves and simplified temperature-composition phase diagram of Li4-xZn1-xGax(PO4)2 samples. ...............................................................................................................................................................................

Figure II. 24 .

 24 Figure II.24. Ionic conductivities of the Li4Zn(PO4)2 sample. ...............................................................................

Figure II. 25 .

 25 Figure II.25. Arrhenius plot ionic conductivity of Li4-xZn1-xGax(PO4)2 samples. .................................................

Figure II. 26 .

 26 Figure II.26. DC polarization of β'-Li3.7Zn0.7Ga0.3(PO4)2 sample. .........................................................................

Figure II. 27 .

 27 Figure II.27. BVEL of (a) -Li4Zn(PO4)2, (b) β-Li4Zn(PO4)2 and (c) '-Li3.7Zn0.7Ga0.3(PO4)2 at the percolation energies. .................................................................................................................................................................

Figure II. 28 .

 28 Figure II.28. BVEL of β-Li3.5Zn0.5Ga0.5(PO4)2, at the percolation energies. ..........................................................

Figure II. 29 .

 29 Figure II.29. Variation of ionic conductivity and broadening of XRD peak with concentration of Li vacancy. ...

Figure II. 30 .

 30 Figure II.30. FESEM images of the cross-section of the sintered pellets. .............................................................

Figure II. 31 .

 31 Figure II.31. Evolution of as-synthesized Na4Zn(PO4)2 with temperature. ...........................................................

Figure II. 32 .

 32 Figure II.32. Rietveld refinement of the SXRD pattern of γ-Na4Zn(PO4)2 recorded at 500 °C. ............................

Figure II. 33 .

 33 Figure II.33. Combined Rietveld refinement of the SXRD (a), and NPD (b) patterns of β-Na4Zn(PO4)2 recorded at RT. .........................................................................................................................................................................

Figure II. 34 .

 34 Figure II.34. Crystal structures of the γ-Na4Zn(PO4)2 (a), and β-Na4Zn(PO4)2 (b). ...............................................

Figure II. 35 .

 35 Figure II.35. Le Bail refinement of the SXRD pattern (a) and the NPD pattern (b) of α-Na4Zn(PO4)2. ................

Figure II. 36 .

 36 Figure II.36. XRD patterns of the Na4-xZn1-xGax(PO4)2 synthesized by quenching. .............................................

Figure II. 37 .

 37 Figure II.37. XRD patterns of the Na4-xZn1-xMx(PO4)2 synthesized by slow cooling for M = Ga (a) and Al (b).

Figure II. 38 .

 38 Figure II.38. Le Bail refinement of the SXRD pattern of slow-cooled β-Na3.85Zn0.85Ga0.15(PO4)2 sample. ...........

Figure II. 39 .

 39 Figure II.39. Evolution of XRD patterns of as-synthesized Na4-xZn1-xGax(PO4)2 samples (prepared by slowcooling) with temperature upon heating (at 1 °C/min), for xGa = 0.15 (a), and xGa = 0.25 (b). ..............................

Figure II. 40 .

 40 Figure II.40. Arrhenius plot of total ionic conductivity of the Na4Zn(PO4)2 samples. ...........................................

Figure II. 41 .

 41 Figure II.41. Arrhenius plot of total ionic conductivity of the Na4-xZn1-xGax(PO4)2 samples synthesized by slow cooling. ..................................................................................................................................................................

Figure II. 42 .

 42 Figure II.42. Strategy to design new ionic conductors. ..........................................................................................

Figure

  Figure III.1. Voltage profiles of TiS2, LiCoO2, and Li1.2Ni0.13Mn0.54Co0.13O2 (Li-rich NMC). ..............................

Figure III. 2 .

 2 Figure III.2. Reported crystal structure of TiS3. .....................................................................................................

Figure III. 3 .

 3 Figure III.3. Picture of the as-synthesized TiS3. ....................................................................................................

Figure III. 4 .

 4 Figure III.4. SEM images of the pristine TiS3 powders. ........................................................................................

Figure III. 5 .

 5 Figure III.5. Rietveld refinement of the SXRD pattern of the pristine TiS3 powder. .............................................

Figure III. 6 .

 6 Figure III.6. Rietveld refinement of the NPD pattern of the pristine TiS3 powder collected at D1b. ....................

Figure III. 7 .

 7 Figure III.7. The crystal structure of as-synthesized TiS3. .....................................................................................

Figure III. 8 .

 8 Figure III.8. HAADF-STEM image of TiS3 single crystal showing the atomic arrangement. ..............................

Figure III. 9 .

 9 Figure III.9. Charge-discharge profile of TiS3 in the 1st cycle. .............................................................................

Figure III. 10 .

 10 Figure III.10. Operando XRD of TiS3. ...................................................................................................................

Figure III. 11 .

 11 Figure III.11. HRTEM images of Li2TiS3 single crystals. .....................................................................................

Figure III. 12 .

 12 Figure III.12. SEM images of the ex situ Li2TiS3 (discharged TiS3) powder. ........................................................

Figure III. 13 .

 13 Figure III.13. Le Bail refinement of ex situ Li2TiS3 powder. .................................................................................

Figure III. 14 .

 14 Figure III.14. HAADF-STEM image of Li~0.2TiS3 after 1st cycle. ........................................................................

Figure III. 15 .

 15 Figure III.15. S 2p and Ti 2p core XPS spectra of the cycled TiS3 samples. .........................................................

Figure III. 16 .

 16 Figure III.16. EELS spectra of the S L2,3 and Ti L2,3 edge of the cycled TiS3 samples. .........................................

Figure III. 17

 17 Figure III.17. S K-edge TFY-XANES spectra collected for TiS3 and LiTiS3, (mid-discharged), Li2TiS3 (fulldischarged), and Li~0.2TiS3 (fully-charged) samples. .............................................................................................

Figure III. 18 .

 18 Figure III.18. Charge-discharge profile of TiS3 over first 25 cycles. .....................................................................

Figure III. 19 .

 19 Figure III.19. Voltage profile of TiS3 recorded with a GITT protocol. ..................................................................

Figure III. 20 .

 20 Figure III.20. Crystal structure of Li2TiS3 from Rietveld refinement of the SXRD pattern. .................................

Figure III. 21 .

 21 Figure III.21. The electrochemical performance of Li2TiS3. .................................................................................

Figure III. 22 .

 22 Figure III.22. Fe-substitution in Li2TiS3. Schematic band structure of Li1.33Ti0.67S2 (i.e., Li2TiS3, in left), and its Fe 2+ substituted derivative (in right). .....................................................................................................................

Figure III. 23 .

 23 Figure III.23. Structural behavior of the Li1.33-2y/3Ti0.67-y/3FeyS2 series. .................................................................

Figure III. 24 .

 24 Figure III.24. Structural characterization of the Li1.13Ti0.57Fe0.3S2 (y = 0.3) sample. .............................................

Figure III. 25 .

 25 Figure III.25. Comparing the crystal structures of Li1.13Ti0.57Fe0.3S2 (a) and Li-rich NMC (b). ............................

Figure III. 26 .

 26 Figure III.26. Electrochemical behavior of Li1.33-2y/3Ti0.67-y/3FeyS2. .......................................................................

Figure III. 27 .

 27 Figure III.27. Capacity and dQ/dV of Li1.33-2y/3Ti0.67-y/3FeyS2. ...............................................................................

Figure III. 28 .

 28 Figure III.28. More details of the electrochemical properties of Li1.13Ti0.57Fe0.3S2 (y = 0.3). ................................

Figure III. 29 .

 29 Figure III.29. Cycling Li1.13Ti0.57Fe0.3S2 in a broader voltage range. ...................................................................

Figure III. 30 .

 30 Figure III.30. Probing the Li (de)intercalation in operando. ...............................................................................

Figure III. 31 .

 31 Figure III.31. 2nd cycle of the in situ cycling of Li/Li1.13Ti0.57Fe0.3S2 half-cell. ..................................................

Figure III. 32 .

 32 Figure III.32. Structural characterization of Li1.13Ti0.57Fe0.3S2 after Li (de)intercalation. ....................................

Figure III. 33 .

 33 Figure III.33. The bond-lengths observed in the pristine, full-charged and full-discharged phase (as obtained from the Rietveld refinement of the corresponding SXRD patterns). ..........................................................................

Figure III. 34 .

 34 FigureIII.34. Operando Mössbauer spectroscopy to identify the Fe 2+/3+ redox process in Li1.13Ti0.57Fe0.3S2. ....

Figure III. 35 .

 35 Figure III.35. Coordination of Fe atoms in Li1.13Ti0.57Fe0.3S2. .............................................................................

Figure III. 36 .

 36 Figure III.36. Operando Mössbauer spectroscopy for the LiTi0.5Fe0.5S2 (y = 0.5) sample. .................................

Figure III. 37 .

 37 Figure III.37. Probing the cationic redox processes in Li1.13Ti0.57Fe0.3S2 by ex situ XAS....................................

Figure III. 38 .

 38 Figure III.38. Probing the anionic redox processes in Li1.13Ti0.57Fe0.3S2. .............................................................

Figure III. 39 .

 39 Figure III.39. Probing S L-edge to identify the S-redox processes. .....................................................................

Figure III. 40

 40 Figure III.40. S-L2,3 edge EELS of FeS (a), and FeS2 and elemental S (b). .........................................................

Figure III. 41 .

 41 Figure III.41. Comparison of the voltage profiles of TiS2 and Li1.13Ti0.57Fe0.3S2 (green curve). .........................

Figure III. 42 .

 42 Figure III.42. Electrochemical behaviour of the Li1.2Ti0.8S2 sample. ...................................................................

Figure III. 43 .

 43 Figure III.43. Electrochemical activity of the Li1.33-2y/3Ti 0.67-y/3CoyS2 samples. ..................................................

Figure III. 44 .

 44 Figure III.44. Electrochemical behaviour of the O3-and P2-Na2TiS3. ...............................................................

Figure

  Figure IV.1. GITT Voltage profile of Li1.13Ti0.57Fe0.3S2 and Li-rich NMC in a three-electrode cell. ..................

Figure IV. 2 .

 2 Figure IV.2. GITT Voltage profile of Li2Ru0.75Sn0.25O3 (LRSO). .......................................................................

Figure IV. 3 .

 3 Figure IV.3. GITT Voltage profile of LiTi0.5Fe0.5S2 (y = 0.5) in a two-electrode cell. ........................................

Figure IV. 4 .

 4 Figure IV.4. Voltage hysteresis in Ti 3+ and Co 2+ substituted Li2TiS3. .................................................................

Figure IV. 5 .

 5 Figure IV.5. Probing the hysteresis of Li1.13Ti0.57Fe0.3S2 by voltage window opening experiment. ....................

Figure IV. 6 .

 6 Figure IV.6. Anionic redox in Li-rich NMC........................................................................................................

Figure IV. 7 .Figure V. 3 .Figure V. 5 .Figure V. 6 .Figure V. 8 .Figure V. 9 .

 735689 Figure IV.7. Probing the hysteresis of Li1.2Ti 4+ 0.4Ti 3+ 0.4S2 by voltage window opening experiment. .................. 125 Figure IV.8. Cell's electrochemical resistance as a function of SoC. .................................................................. 126 Figure IV.9. Typical Nyquist plot of EIS of Li1.13Ti0.57Fe0.3S2 in Li-half cell. ..................................................... 127 Figure IV.10. Nyquist plots of EIS recorded after each relaxation in GITT. ....................................................... 128 Figure IV.11. Discharge profiles of Li1.13Ti0.57Fe0.3S2 over long cycling to show the voltage fade. .................... 129 Figure IV.12. Average voltage during long cycling of Li1.13Ti0.57Fe0.3S2 and Li-rich NMC in Li-half cells. ...... 129 Figure IV.13. Average voltage of Li-rich sulfides in Li-half cells. ...................................................................... 130 Figure IV.14. The voltage profiles in 2nd cycle of (a) Na0.67[Mg0.28Mn0.72O2], and (b) O3-Na2TiS3. ................. 131 Figure IV.15. Capacity retentions of the O3-and P2-Na2TiS3. ........................................................................... 131 Figure IV.16. Average discharge voltages obtained in O3-and P2-Na2TiS3 over first few cycles. .................... 132 Figure IV.17. Ragone plots of Li1.13Ti0.57Fe0.3S2. ................................................................................................. 133 Figure V.1. Schematic diagram of cathode/sulfide electrolyte interfaces with oxide or sulfide cathode. ............ 136 Figure V.2. A picture of the ASSB setup fully designed and assembled at lab, that was used in this study. ...... 137 Figure V.3. Schematics of the three-layer ASSB. ................................................................................................ 138 Figure V.4. EIS spectra of the glassy 75•Li2S-25•P2S5 (g-LPS) and β-Li3PS4 (LPS) solid electrolyte. ............... 139 Figure V.5. Rietveld refinement of the SXRD pattern of the as-synthesized β-Li3PS4 SE. ................................. 140 Figure V.6. Rietveld refinement of the SXRD pattern of the cathode composite. ............................................... 140 Figure V.7. DC polarization plots of LTFS-LPS cathode composite................................................................... 141 Figure V.8. Voltage profile of Li1.13Ti0.57Fe0.3S2 (LTFS) cycled in ASSB at RT. ................................................ 142 Figure V.9. Polarization in charge-discharge of Li1.13Ti0.57Fe0.3S2│g-LPS│In ASSB. ......................................... 142 Figure V.10. Voltage profile of Li1.13Ti0.57Fe0.3S2 cycled in ASSB at RT using β-Li3PS4 as solid electrolyte. ... 143 Figure V.11. EIS characterization of the Li1.13Ti0.57Fe0.3S2│g-LPS│In ASSB...................................................... 144

Table 3 . Ionic conductivities of the Li4-xZn1-xGax(PO4)2 samples at RT obtained by extrapolation of the Arrhenius plot and at 200 °C and their activation energies (Ea).

 3

			σg			σgb			σtot
	xGa		(S•cm -1 ) RT	200 °C	Ea (eV)	(S•cm -1 ) RT	200 °C	Ea (eV)	(S•cm -1 ) RT	200 °C	Ea (eV)
		0	5.6•10 -18 2.1•10 -9 1.32 3.5•10 -19 9.5•10 -11 1.30 3.6•10 -19 9.1•10 -11 1.3
		0.05 1.4•10 -12 5.7•10 -7 0.87 4.2•10 -14 7.4•10 -9	0.82 4.1•10 -14 7.3•10 -9	0.82
	α-	0.1	1.9•10 -12 1.3•10 -6 0.92 8.3•10 -14 2.4•10 -8	0.85 8.1•10 -14 2.3•10 -8	0.85
		0.15 6.5•10 -12 3.6•10 -6 0.89 2.7•10 -13 3.1•10 -8	0.79 1.8•10 -13 3.1•10 -8	0.82
		0.2	5.2•10 -12 1.9•10 -6 0.95 3.5•10 -13 6.9•10 -8	0.82 3.4•10 -13 6.7•10 -8	0.83
	β'-	0.25 6•10 -12 0.3 5.7•10 -12 4.3•10 -6 0.91 3.9•10 -13 1.4•10 -7 5.9•10 -6 0.93 4.9•10 -13 1.1•10 -7	0.84 4.5•10 -13 1.1•10 -7 0.86 3.6•10 -13 1.4•10 -7	0.84 0.87
	β-	0.4 0.5	5.6•10 -12 9.2•10 -6 0.95 2.1•10 -13 8.7•10 -8 1.6•10 -12 5.9•10 -6 1.02 1.4•10 -13 1.3•10 -7	0.88 2.7•10 -13 8.6•10 -8 0.93 1.3•10 -13 1.3•10 -7	0.87 0.93

  2-/S 2-redox in 'pseudo-layered' TiS3. Interestingly, we identified a subtle discrepancy in the reported structural model (byFuruseth et al, 1975 257 ) of TiS3. Our structural model supports its structure alike the 'ZrSe3-type variant', where S atoms form 'isosceles

	atom	site	x	y	z	Biso (Å 2 )	occ. BVS
	Zn	8f	0.18337(9)	0.4266(2)	0.4387(3)	2.05(5)	1	1.842(19)
	P1	8f	0.19763(18)	0.3036(4)	0.1283(5)	0.27(3)	1	5.159(61)
	O11	8f	0.2545(3)	0.2542(6)	0.2282(8)	0.72(3)	1	2.194(32)
	O12	8f	0.2116(3)	0.4081(7)	0.0320(8)	0.72(3)	1	2.114(37)
	O13	8f	0.1695(3)	0.1884(7)	0.0445(8)	0.72(3)	1	2.073(37)
	O14	8f	0.1631(3)	0.3638(7)	0.2255(9)	0.72(3)	1	1.946(37)
	P2	8f	0.44386(16)	0.0784(4)	0.1524(5)	0.27(3)	1	5.193(60)
	O21	8f	0.4518(3)	0.1461(6)	0.0168(8)	0.72(3)	1	2.031(40)

Table A1 .3. Structural Parameters for α-Li4Zn(PO4)2, deduced from the Rietveld refinement of the NPD at RT

 A1 . Isotropic temperature factors (Biso) were constraint to be equal for a same chemical species.

	Appendix							A-15
	O3	8f	0.4115(5)	0.8063(13)	0.0136(8)	4.44(10)	1	2.064(35)
	O4	8f	0.3381(6)	0.7934(13)	0.2446(9)	4.44(10)	1	2.290(41)
	Li1/Zn1	8f	0.0969(6)	0.3587(18)	0.0311(11)	9.3(4)	0.852(3)/	0.964(17)/
							0.148(3)	1.834(33)
	Li2/Zn2	8f	0.2277(12)	0.652(4)	0.3082(16)	17.2(8)	0.898(3)/	0.843(27)/
							0.102(3)	1.706(51)
	Appendix
								5.241(55)
								1.833(30)
								2.155(14)

# This chapter includes the following publication: Saha et al., Nature Energy 2019.

B TEM studies were performed by Prof Artem M. Abakumov, Skolkovo Institute of Science and Technology (Russia).

F Mössbauer studies were performed by Dr Moulay Tahar Sougrati, Institut Charles Gerhardt (Montpellier).

Acknowledgements i

Acknowledgements

Time flies. Three years have passed, since I started my PhD at College de France. Now, I take immense pleasure for the chance to thank the people whom I am indebted for their help and enriching my experience during this time.

First and foremost, I would like to thank Prof. Jean-Marie Tarascon. He not only introduced me

The evolution of the spectra obtained during in situ cycling is shown as a contour plot in Figure III.34b, with the spectra analyzed using principal component analysis (PCA), as described in Appendix Section A2.1 and the fitting parameters are summarized in Appendix Table A2.6. All measured spectra could be adequately fitted as linear combinations of three reconstructed spectral components. As expected, the first component (Comp. 1) is identical to the pristine material's spectrum. The Comp. 2 can be fitted using at least two doublets (Figure III.34a). The major one (90%) has an isomer shift of 0.48 mm/s which stands between what is expected for HS Fe 2+ and HS Fe 3+ , alike what has been seen in Fe3S4 and FeV2S4. 289 The second doublet with a 10 % contribution to the overall spectrum is indicative of the HS-Fe 3+ signature, as seen in NaFeS2. 289 Overall, the average oxidation state of Comp. 2 is higher than Fe 2+ but not fully reaching Fe 3+ . Note that Comp. 2 reaches its maximum around xLi = ~0.93, just before the voltage plateau (P). This indicates an exchange of ~0.2 Li, that takes place in the sloppy profile (S), before the plateau and in agreement with the voltage window opening experiment (see the Appendix Figure A2.8). During the plateau, the Comp. 2 converts progressively to the Comp. 3 (fullycharged, 3 V), which can be fitted as low-spin Fe 3+ state alike in the previous report 246 . On discharge, the evolution of the components is reversed, except for a much lower contribution from the Comp. 2 (Figure III.34b). This is fully consistent with the path dependence observed above with XRD. Overall, these data indicate the progressive oxidation of Fe 2+ to Fe 3+ on charge and its full reduction back to Fe 2+ on discharge.

Similarly, operando Mössbauer spectroscopy was performed for the LiTi0.5Fe0.5S2 (y = 0.5) composition that contains more Fe, and the results are summarized in Figure III.36 with the fitting results in Appendix Table A2.7. The scenario is similar as before except the intermediate phase (comp 2, with mixed valence) that contains more HS-Fe 2+ , since it is well fitted using 27% of HS-Fe 2+ and 73% of Fe 2+ /Fe 3+ environments and reaches its maximum around xLi = ~0.6. Furthermore, the spectra were collected over the 2nd cycle, which shows evolution of the spectra similar to the 1st cycle (shown in Appendix Figure A2.9), thus confirming the trend of the Fe 2+/3+ redox among the compositions. Furthermore, X-ray absorption near edge structure (XANES) spectroscopy was performed to confirm the above observation in Li1.13Ti0.57Fe0.3S2. See the Appendix Section A2.1 for details of the methods of XANES acquisition. The results of the Fe L2,3-edge TFY-XANES spectra are shown in Figure III.37b, with the spectra in TEY mode gathered in Appendix Figure A2.10. The spectra for the pristine material is similar to that of FeS confirming the presence of Fe 2+ . 291,292 It enlists two main peaks corresponding to 2p3/2 → 3d (L3) and 2p1/2 → 3d (L2) transitions. These peaks shift to higher energy upon charge indicating gradual conversion to Fe 3+ . The position and shifts in energy observed in our Fe

III.4.3 Na2TiS3

In contrary to the Li2TiS3, its Na-counterpart is shown to be redox-active, without any substitution. [START_REF] Pele | Nouveaux materiaux d'electrodes pour microbatteries au sodium[END_REF][START_REF] Nasu | Amorphous Na 2 TiS 3 as an Active Material for All-solid-state Sodium Batteries[END_REF] The Na2TiS3 (also written as Na1.33Ti0.67S2) has a low temperature O3-polymorph and high temperature P2-polymorph. [START_REF] Pele | Nouveaux materiaux d'electrodes pour microbatteries au sodium[END_REF][START_REF] Nasu | Amorphous Na 2 TiS 3 as an Active Material for All-solid-state Sodium Batteries[END_REF] Amongst the two phases, only the voltage profile for O3-Na2TiS3 was reported, by Pele et al, 2016 (in Na-half cell) [START_REF] Pele | Nouveaux materiaux d'electrodes pour microbatteries au sodium[END_REF] , and by Nasu et al, 2019 (in ASSB) [START_REF] Nasu | Amorphous Na 2 TiS 3 as an Active Material for All-solid-state Sodium Batteries[END_REF] . Thus, to grasp further into the anionic redox in sulfides, we decided to extend our studies to Na2TiS3.

The O3 and P2 phases were synthesized slow-cooling and quenching, respectively (see methods in Appendix Section A2.1). Their structural characterizations can be found in Appendix Figure A2. 19, Appendix Figure A2.20 and Appendix Figure A2.21. Their electrochemical performances were tested in Na-half cell, using 1 M NaPF6 in PC as electrolyte. For the O3 phase, ~1. 16 Na could be deintercalated in 1st charge, and there remains an irreversibility only ~0.1 Na in the 1st discharge and stable cycling occurs afterwards with a polarization of ~300 mV (Figure III.44a,b). Note that we obtained much higher capacity (~225 mAh•g -1 ) than the reported (~100 mAh•g -1 ) [START_REF] Pele | Nouveaux materiaux d'electrodes pour microbatteries au sodium[END_REF] , which is most likely due to the electrolyte used (1 M NaPF6 in TEGDME), see the Appendix Figure A2.22 for details. For the P2 phase, the Na-(de)intercalation occurs at a lower voltage than the O3 phase (Figure III.44c,d). Up to ~1.2 Na could be extracted in the 1st cycle of the P2 phase, but in discharge it shows a

IV.1 Chapter Introduction

In the previous chapter, we showed that redox activity in Li-rich layered sulfides can be triggered by suitable TM-substituents. More specifically, the Li1.33-2y/3Ti 4+ 0.67-y/3Fe 2+ yS2 family (with TM = Fe 2+ ) displays excellent reversibility of Li-(de)intercalation while involving anionic S-redox. Furthermore, we spotted that TM = Ti 3+ , Co 2+ substitutition can equally trigger reversible redox activity in Li2TiS3, whereas the Na2TiS3 is active without any substitution. At this stage, having evidenced the strong participation of anionic redox in the charge compensation mechanism in these Li-rich layered sulfides, the next legitimate question remains about the practicability of anionic redox in sulfides. Despite the promise of large capacity in Li-rich NMC oxide, certain practical issues, such as voltage fade, sluggish kinetics, large voltage hysteresis, and irreversible oxygen loss, have delayed the commercialization of Li-rich NMCs. Although the crucial role of oxygen redox towards these issues was clearly highlighted by detailed investigations on 'practical ' Li-rich NMC and also for 'model' Li1.33Ru0.5Sn0.17O2 electrode, much remains to be understood prior to provide implementable solutions. [START_REF] Assat | Fundamental understanding and practical challenges of anionic redox activity in Li-ion batteries[END_REF][START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF]278 Using these Li-rich sulfides as model compounds, we investigated whether these practical issues, still persist when oxygen is replaced by sulfur. Therefore in this chapter we will systematically assess the practical figures of merit in Li-rich sulfides starting with the voltage hysteresis issue.

IV.1.1 Voltage hysteresis

Let's recall first that the Li-rich Li1.13Ti0.57Fe0.3S2 sulfide shows an overpotential of only ~100 mV during cycling at C/5 rate (with current density of 46 mA•g -1 ). To get further insight into the voltage hysteresis, we performed a galvanostatic intermittent titration technique (GITT) experiment (See Appendix for the method). For this, a rest of 30 mins was applied after each 30 mins of charge/discharge of the Li-half cell and the voltage was recorded. The GITT voltage profile of Li1.13Ti0.57Fe0.3S2 is shown in When compared with Li-rich oxides, this is a significantly better scenario than in Li-rich NMC (Li1.2Ni0.13Mn0.54Co0.13O2), where a severe OCV hysteresis up to 270 mV has been reported and further shown to be associated with oxygen redox (Figure IV.1b). [START_REF] Assat | Fundamental interplay between anionic/cationic redox governing the kinetics and thermodynamics of lithium-rich cathodes[END_REF] It was shown before that by moving to Li2Ru0.75Sn0.25O3 (4d metal-based), the voltage hysteresis was considerably mitigated but it could not be fully eliminated (Figure IV.2). [START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF] The performance of Li1.13Ti0.57Fe0.3S2 is also better than the ~100 mV OCV hysteresis in Li2Ru0.75Sn0.25O3 and approaches the favorable hysteresis-free situation observed for 5d metal-based β-Li2IrO3. [START_REF] Assat | Choice-Practical Assessment of Anionic Redox in Li-Rich Layered Oxide Cathodes: A Mixed Blessing for High Energy Li-Ion Batteries[END_REF][START_REF] Pearce | Evidence for anionic redox activity in a tridimensional-ordered Li-rich positive electrode β-Li2IrO3[END_REF] triangles' around the Ti atoms, as was shown by Guilmeau et al, 2014. 258 Whatever, TiS3 shows a rapid capacity fading upon Li-(de)intercalation with large hysteresis, due to unsustainable structural rearrangement during S-S bond breaking/reformation. Towards this end, we decided to design Li-rich layered sulfides that undergo sulfur redox without involving complete S-S dimerization. We have chosen Li2TiS3 for this purpose; however, was found to be nearly electrochemically inactive. To initiate the reversible redox activity we demonstrated the benefit of partial substitution with proper transition metal (TM). We showed that with TM = Fe 2+ , we can initiate reversible redox activity. Along this line, we have synthesized the Li-rich layered Li1.33-2y/3Ti 4+ 0.67-y/3Fe 2+ yS2 family that shows reversible redox activity. We have isolated the composition Li1.13Ti0.57Fe0.3S2 (y = 0.3) that was shown to have a specific energy of 600 Wh•kg -1 (which compares favorably with LiCoO2). Detailed charge-compensation mechanism was investigated indicating the presence of Fe 3+ and S n-(n < 2) in the fully-charged phase, thus demonstrating simultaneous cationic and anionic redox.

Next, inspired by this reversibility of the anionic S 2-/S n-(n < 1) redox, we have benchmarked the Li-rich sulfides against the Li-rich NMC oxides. The anionic O 2-/O n-(n < 2) redox has emerged as a new paradigm to design high-energy Li-rich oxide cathode materials (e.g., Li-rich NMC), however with several practical drawbacks, such as large voltage hysteresis, voltage fade, sluggish kinetics, etc.

In contrarily, our Li-rich sulfide Li1.13Ti0.57Fe0.3S2 shows (i) a voltage hysteresis as low as ~30 mV under OCV condition, (ii) small voltage fade (~50 mV in 60 cycles), along with fast electrochemical kinetics. Furthermore, we extended our study to Ti 3+ /Co 2+ substituted Li2TiS3 and Na2TiS3 that also highlight the positive attributions of S-redox. However, we need to keep in mind that this comes with a price of energy density, because of lower voltage in sulfides. The best compromise would consists in designing oxysulfides-based Li-rich cathode materials, which could be an attractive direction for future. [329][330][331][332][333][334] Having designed the Li-rich sulfides, we proceed to assemble full ASSB using Li1.13Ti0.57Fe0.3S2 as positive electrode. We assembled the Li1.13Ti0.57Fe0.3S2│SE│In ASSBs, with glassy Li2S-P2S5 and β-Li3PS4 as SEs, that demonstrate excellent cyclability. Such performance could be attributed to the high electronic conductivity (>10 -2 S•cm -1 ) of Li1.13Ti0.57Fe0.3S2 and its small interfacial resistance. Thus, use of Li-rich sulfide based cathode materials could be beneficial to tackle that interfacial challenge and low electrochemical stability window of sulfide SEs in ASSBs. Will sulfide replace the oxides remains an open question. Interfaces are obviously more friendly with sulfides, but their use could be hindered because of their moisture-sensitivity and unwanted H2S gas release.

Appendix

Appendix A-2 A1 Supporting Information for Chapter II A1.1 Methods Material Synthesis. Li6Zn(P2O7)2 sample was prepared by following classical solid state synthesis. Li2CO3 (Sigma Aldrich, 99%), ZnO (Sigma Aldrich, 99%), and NH4H2PO4 (Alfa Aesar, 98%) were used as precursors. Required amounts of precursors were mixed homogeneously by ball-milling for 20 mins with stainless-steel balls in a stainless-steel vial using SPEX-8000M mixer-miller and pressed into a pellet subsequently. The pellet was heated on an alumina boat for 36 hours at 700 °C in air after a ramp of 3 °C/min, followed by quenching in air. The pellet was grinded again and pelletized, followed by subsequent annealing at 300 °C for 12 hours in air to obtain the pure -Li6Zn(P2O7)2 of the compound.

To prepare samples of Li4Zn(PO4)2 and Li4Zn1-xGax(PO4)2 stoichiometric amounts of Li2CO3, ZnO, NH4H2PO4 and Ga2O3 (Alfa Aesar, 99.99%) were used as precursors.

Na2CO3 (Sigma Aldrich, 99.95%), ZnO, NH4H2PO4, and Al2O3 (Sigma Aldrich, 99%)/ Ga2O3/ In2O3 (Sigma Aldrich, 99.99%), were used as precursors to prepare samples of Na4Zn(PO4)2 and Na4Zn1-xMx(PO4)2 (M= Al, Ga, and In).

In case of Li4Zn1-xGax(PO4)2, different pellets were heated in alumina boats for 36 hours at 800 °C after a ramp of 1 °C/min, followed by quenching in air. For Na4Zn(PO4)2 and Na4Zn1-xMx(PO4)2 (M= Al, Ga, and In), pellets were heated in alumina boats for 24 hours at 700 °C after a ramp of 1 °C/min, followed by quenching or cooling slowly (1 °C/min) to room-temperature (RT), as indicated in the chapter.

For the high-energy ball-milling synthesis of the β-Li4Zn(PO4)2, the as-prepared α-phase of the samples was continuously milled for 2 hrs using balls-to-powder weight ratio of 40.

Structural Characterizations. Powder X-ray diffraction (XRD) patterns were recorded in Bragg-Brentano geometry using a Bruker D8 Advance diffractometer equipped with a Cu K X-ray source (λ1 = 1.54056 Å, λ2 = 1.54439 Å) and a LynxEye detector. Evolution of XRD patterns of the samples with temperature were monitored in situ in the same diffractometer equipped with an Anton Paar HTK 1200N furnace to control the temperature. A ramping rate of 2 °C/min was applied.

Complementary synchrotron X-ray powder diffraction (SXRD) and neutron powder diffraction (NPD) data were also obtained for more accurate characterization. Temperature-driven structural transitions were monitored in situ in the MSPD beamline at the ALBA synchrotron beamline (Barcelona, Spain) equipped with a hot-air blower to control the temperature. All SXRD data were collected in transmission mode with λ = 0.4948 Å, placing the powder in a sealed quartz capillary with a diameter of 0.5 mm and a ramping rate of 5 °C/min was applied for the temperature-driven experiment.

A-3

In addition, some SXRD patterns at RT were also collected at the 11-BM beamline of the Advanced Photon Source (APS), Argonne National Laboratory, in transmission mode with λ = 0.4127 Å, with the powder sealed in a kapton capillary of 0.7 mm diameter. Complementary high-resolution NPD patterns were collected with the HRPT instrument at SINQ-PSI (Villigen, Switzerland). The sample powder was placed in a vanadium container of 8 mm diameter to record the NPD patterns in transmission mode with a neutron wavelength of 1.494 Å.

All patterns were refined using the Rietveld method, 208 or by Le Bail fitting 335 , as implemented in the FullProf program 336 .

All crystal structures in this thesis are drawn using the program VESTA 337,338 . DSC measurements. Differential scanning calorimetry (DSC) measurements were done in air using a STA 449C Netzsch apparatus at a rate of 10 °C/min or at a rate mentioned otherwise. TEM studies. The samples for selected area electron diffraction (SAED) were prepared by depositing the sample onto holey carbon grids by drop-casting a sample-ethanol suspension. Electron diffraction patterns were collected with a FEI Tecnai G2 electron microscope operated at 200 kV.

Bond Valence Energy Landscape maps (BVEL) Analysis. Bond Valence Energy

Landscape maps (BVEL) were generated according to the method developed by Adams, 339 using the program BondSTR as implemented in the FullProf Suite 336 .

The 'bond valence modeling' is a simple theoretical tool, which gives a necessary first approximation for the mechanism of lithium diffusion pathways. 340 Generally, ion transport from one Li equilibrium site to the other one follows a pathway, along which the valence mismatch ΔV = |V -Vnom| between the bond valence V and the nominal valence Vnom remains as small as possible. Following Pauling's rule (1929), the total bond valence sum V of cation 𝑀 + can be expressed as,

Here the individual bond valences 𝑆 𝑀 + -𝑋 are calculated using the bond lengths 𝑅 𝑀 + -𝑋 between cation 𝑀 + and the adjacent anions X:

The bond valence parameters 𝑅 0 and b are deduced from a large number of crystalline compounds containing the same 𝑀 + -𝑋 pairs, but having different coordination number and bond length distributions. In all these different compounds, a monovalent cation 𝑀 + should have a bond Appendix A-4 valence sum (BVS) close to Videal = 1. Hence for crystalline compounds, the process of the ion 𝑀 + transport from one equilibrium site to another should follow the channel which requires the lowest valence mismatch ΔV = |V -Vnom|, which in turn corresponds to the energetically most favored pathway. Thus in principle, the activation energy Ea should be directly related to ΔV. However, this model based on valence V still does not include all interactions other than the 𝑀 + -𝑋 interaction. Hence, in 2011, Adams et al proposed BV-based force-field approach by introducing the interaction energies. In this method, the Li ions are assumed to move within a framework composed of immobile anions and other cations. Thus, the interaction between a dummy Li ion in the space and the ions surrounding it are empirically evaluated by the summation of the attraction between Li + and anions described by Morsetype potential and the repulsion between Li + and cations expressed by Coulombic potential follows,

Where D0, α, and Rmin are Morse potentials determine from large amount of stable compounds, q and R refer to the charge and Li-A distance respectively. The summation of above two terms on grids of points with a small resolution (e.g., a unit cell divided into 100s smaller subcell) builds the maps of the total potential energy E (Li) of the unit cell. The region enclosed by the isosurfaces of constant E (Li) are considered as the space where Li + can move through and the threshold value of E (Li) to form a continuous pathway is estimated as the Li + migration energy barrier.

However the E (Li) does not necessarily give the real Ea of a compound, as in real systems lot of other complex interactions exist which this model cannot take into account. Hence, in this thesis, the percolation energy above the minimum, E (Li) obtained is only used qualitatively and the isosurfaces is used to predict possible Li conduction channel. For the BVEL maps, calculations of Li/Na conduction paths were done after the removal of pre-existing Li/Na in the unit cell and also after removing Zn and Ga ions located at mixed Li sites. Anionic neighbors up to 8 Å were considered.

Electrochemical Impedance Spectroscopy (EIS). EIS was used to determine the Li-ion

conductivity of the samples. The AC impedance measurements were made using a Bio-Logic MTZ-35 impedance analyzer equipped with a HTF-1100 furnace. To perform the measurements, pellets of ~8 mm diameter and ~1.5 mm thickness were obtained by cold-pressing using a uniaxial pressure of ~5 bars in a hydraulic press and were densified by sintering at 600 °C for 6 hours following a heating rate of 1 °C/min and then quenched to RT and reannealed at 300°C to obtain the pure α-Li6Zn(P2O7)2 phase. In case of Li4Zn1-xGax(PO4)2 and Na4Zn1-xMx(PO4)2, pellets of ~8 mm diameter and ~1.5 mm thickness were obtained by cold-pressing using a uniaxial pressure of ~5 bars in a hydraulic press and were densified by sintering at 700 °C for 6 hours while heated and cooled at a rate of 1 °C/min. In exceptional case of low-temperature sintering, that was applied to the pellet of ball-milled β-Li4Zn(PO4)2 or γ-Na4Zn(PO4)2, pellets were sintered at 200 °C for 4 days. The sintered pellet was found to have a density A-5 of 75 -80 % with respect to the crystallographic densities; hence, our reported conductivity values were not corrected by the pellet porosities.

The AC impedance measurements were performed using platinum (Pt) blocking electrodes, by applying a sinusoidal signal of 500 mV amplitude over a frequency range of 10 MHz to 1 Hz. A thin foil of carbon was placed between the Pt electrode and the pellet in order to improve the electrode contact, as shown as in Appendix Figure A1.1.

Appendix Figure A1.1. Schematics of the AC impedance measurement procedure.

The measurements were carried out in the temperature range of 50 -650 °C (or mentioned otherwise) with an accuracy of ±1 °C and were soaked for 15 mins at each measurement temperature prior triggering each sweep of frequency dependent impedance measurement at the temperature.

Analysis of EIS.

The Appendix Figure A1.2 shows a typical Nyquist impedance plot of a sintered pellet. Each spectra recorded shows a depressed semicircle at high frequencies followed by a Warburg spike at lower frequency region. The low frequency spike represents the impedance response due to the blocking of mobile Li ions at the electrode interface with Pt blocking electrode and it confirms that the ac conductivity is ionic in nature. The total ionic conductivity (σtot) can be obtained from the value of intercept of the extrapolated low frequency spike with the real axis. However, we have explored the impedance spectra further by examining the semicircle at high frequency, which is too depressed to be accounted by the inhomogeneities of the sample-electrode interfaces and hence needs to deconvolute to different contributions. The impedance response of polycrystalline materials are expected to be composed of contributions from grains and grain boundaries (see Appendix Figure A1.2). Hence, we decoupled the spectra by fitting the complex impedance plot to an equivalent circuit model (RgCPEg)( RgbCPEgb)WO, where R and CPE represent the resistance and constant phase element (CPE) associated with the grains (g) or grain boundaries (gb) of the sample respectively and WO is the Warburg resistance. The contributions were assigned to either bulk/grains (σg) or grain boundaries (σgb) based on their associated capacitances (C), calculated according to the following formulae 341 ,

. Typical Nyquist impedance plot of a sintered pellet and its interpretation.

For visualization, the contributions from bulk (grain), grain boundary and Warburg resistances are shown by dashed semicircles or line. Inset shows the equivalent circuit that was used to fit the Nyquist plot.

Here Q is the pre-factor of the CPE, and n its exponent. Typically, the capacitances were calculated to be ~10 -12 -10 -11 F for grains or ~10 -11 -10 -10 F for grain boundaries. 216,342 This assignment was confirmed by calculating the relaxation frequency (fr) which generally appears at the apex of the deconvoluted arcs and is always higher for grains response than that of the grain boundaries. 216,342 The conductivities were calculated from the resistances according to the formulae, σ = (d/A)•(1/R), where d and A are the thickness and the area of the pellet measured, respectively. A-7

A1.2 Supporting Figures

A1.3 Supporting Tables

Appendix Table A1.1. Structural Parameters for α-Li6Zn(P2O7)2, deduced from the Rietveld refinement of the NPD. Isotropic temperature factors (Biso) were constraint to be equal for a same chemical species. 0.6966 (5) 0.7319 (7) 0.8766 (6) 0.8123(4) 0.900 (7) Biso 1.16 (7) 1.89 (5) 1. 89(5) 1.89 (5) 1.37 (8) 1.37( 8 a Isotropic temperature factors (Biso) were constrained to be equal for the chemical species in same sites.

A2 Supporting information for Chapter III

A2.1 Methods Synthesis. Li1.33-2y/3Ti0.67-y/3FeyS2 samples were prepared by solid-state reaction of stoichiometric amounts of Li2S (Alfa Aesar, 99.9%), TiS2 (Sigma Aldrich, 99.9%) and FeS (Alfa Aesar, 99%). Homogeneously mixed and hand-grinded precursor powders were filled in quartz tubes in an Arfilled glovebox followed by sealing the tubes under vacuum (~10 -5 mbar). The sealed tubes were subsequently annealed at 750 °C for 36 h followed by quenching in water. TiS3, was prepared similarly, by reacting TiS2 and elemental S (15 wt% extra), in a vacuum sealed quartz tube at 550 °C, followed by cooling slowly.

The Li1.2Ti0.8S2 (Li1.33-y/3Ti 4+ 0.67-2y/3Ti 3+ yS2 with y = 0.4) was synthesized similarly, by controlling the stoichiometry and by reacting stoichiometric amount of Li2S, and TiS2 in a vacuum-sealed quartz tube at 750 °C.

Nominal compositions of Li1.33-2y/3Ti 4+ 0.67-y/3Co 2+ yS2, with y = 0.2, and 0.3, were prepared by reacting Li2S, TiS2 and CoS (Alfa Aesar, 99.5%).

The Na2TiS3 was prepared by prepared via solid-state synthesis, by reacting Na2S (Sigma Aldrich), and TiS2 in stoichiometric amounts in vacuum-sealed quartz tubes at 750 °C for 36 h, followed by either cooling to RT at 1 °C/min (for O3 phase) or quenching in water (for P2 phase).

The as-prepared samples were collected inside a glovebox and hand-grinded prior to further use. In the whole process, air contact was avoided and subsequent processing was done in an Ar-filled glovebox.

Structural Characterizations. Synchrotron X-ray powder diffraction (SXRD) patterns were collected at the 11-BM beamline of the Advanced Photon Source (APS), Argonne National Laboratory. All SXRD data were collected in transmission mode with λ= 0.4127 Å, with the powder sealed in a quartz capillary of 0.7 mm diameter. Operando X-ray powder diffraction (XRD) was performed in an airtight electrochemical cell equipped with a Be window. XRD patterns were recorded in reflection mode in Bragg-Brentano geometry using a Bruker D8 Advance diffractometer equipped with a Cu-K X-ray source (λ1 = 1.54056 Å, λ2 = 1.54439 Å) and a LynxEye detector.

High-resolution neutron diffraction was performed for the TiS3 sample on the D1B powder diffractometer (Institut Laue-Langevin) with a wavelength λ = 1.286 Å. The neutron diffraction of the Li1.13Ti0.57Fe0.3S2 sample was performed on the HRPT instrument at SINQ-PSI (Villigen, Switzerland) with a wavelength λ = 1.494 Å. In both instances, the sample powders were placed in a vanadium container of 8 mm diameter to record the NPD patterns in transmission mode.

The refinements of the XRD or NPD patterns were done using the Rietveld method 208 as implemented in the FullProf program 336 . For some cases, Le Bail refinement was also performed 335 in the FullProf program. All crystal structures are drawn using the program VESTA 337,338 .

Electrochemical characterizations. All samples were cycled in galvanostatic mode in Li half-cells assembled in Swagelok-type cells. The cathode materials were mixed with 20 wt% conductive carbon Super-P by hand-grinding for 5 min in an agate mortar and pestle prior to cycling. LP30 (1M LiPF6 in ethylene carbonate/dimethyl carbonate in 1:1 weight ratio) was used as the electrolyte and was soaked in a Whatman GF/D borosilicate glass fiber membrane that was used as separator. Typical loadings of 10 mg•cm -2 of active materials were used and metallic Li was used as the negative electrode. The cells were assembled in an Ar-filled glovebox and were cycled at a C/20 rate between 1.8 V and 3 V if not specified otherwise. To check whether Li-vacancy exists in the pristine samples, fresh cells or all compositions were cycled starting with discharge. The cells showed negligible Li-intercalation in those discharges, thus confirming the absence of Li-vacancies.

The electrochemical performances of the Na2TiS3 (mixed and hand-ground with 20 wt% Csp) were tested in Na-half cells, where 1 M NaPF6 in PC was used as electrolyte.

Charged/discharged samples from the Swagelok cells were recovered for ex situ characterizations by disassembling the cells inside glovebox, rinsed thoroughly with anhydrous DMC and dried under vacuum.

Mössbauer spectroscopy. Room-temperature [START_REF] Nowack | Charge densities in CoS2 and NiS2 (pyrite structure)[END_REF] Fe Mössbauer spectra were recorded in transmission geometry in the constant acceleration mode and with a 57 Co(Rh) source with a nominal activity of 370 MBq. The velocity scale (±4 mms -1 ) was calibrated at RT with α-Fe foil. The in situ cell was prepared with 32 mg.cm -2 of active material mixed with 8 mg of carbon black. The hyperfine parameters IS (isomer shift) and QS (quadrupole splitting) were determined by fitting Lorentzian lines to the experimental data. The isomer shifts values are calculated with respect to that of α-Fe standard at RT.

The whole Mössbauer dataset was analyzed using a chemometric approach including Principal component Analysis (PCA) followed by Multivariate Curve Resolution -Alternate Least Squares A-21 (MCR-ALS) approach. [343][344][345] PCA is a tool, which extracts the maximum amount of useful information from a large data set by considering each spectrum as a vector in a multidimensional space. The "principal components" are the independent (orthogonal) vectors necessary to describe the whole set of spectra, but these components are only orthogonal mathematic functions, not real Mössbauer spectra. The MCR-ALS method is then needed to reconstruct the real components. The fitting was then performed on these reconstructed spectral components with least square conventional Mössbauer methods. Details on this fitting approach can be found in the Reference 346 .

TEM studies. The sample was prepared in an Ar-filled glove box by crushing the crystals in a mortar in DMC and depositing drops of the suspension onto holey carbon grids. The samples were transported to the transmission electron microscope (TEM) column while completely avoiding contact with air. High angle annular dark field scanning transmission electron microscopy (HAADF-STEM) images and Energy loss spectra (EELS) spectra were obtained with a FEI Titan G3 electron microscope operated at 120 kV and equipped with a monochromator and a Gatan Enfinium ER spectrometer. Energy resolution measured by full width at half maximum of the zero loss peak is 0.15 eV.

X-ray absorption spectroscopy (XAS).

S K-edge, Fe L-edge and Ti L-edge X-ray absorption near edge spectroscopy (XANES) measurements were performed at 4-ID-C beamline of APS at Argonne National Laboratory. Spectra were recorded simultaneously under both the total electron yield (TEY) mode from the sample photocurrent at ∼10 -9 Torr and total fluorescence yield (TFY) mode using a silicon drift diode detector at a spectral resolution of ∼0.2 eV, with a 2 s dwell time. The energy scales of the spectra were calibrated with the references of Mo metal, Fe metal and SrTiO3 measured simultaneously, for S, Fe and Ti edges, respectively. sXAS and mRIXS. Soft X-ray absorption spectroscopy (sXAS) and mapping of resonant inelastic X-ray scattering (mRIXS) experiments were performed in the iRIXS endstation of beamline 8.0.1 at the Advanced Light Source (ALS) of Lawrence Berkeley National Laboratory. 347 The pristine powder and cycled electrodes were mounted in high purity Ar glove box and transferred into the experimental vacuum chamber through a home-made kit to avoid any air exposure. The experimental energy resolution of sXAS is better than 0.1 eV without considering the intrinsic core hole broadening (~0.2 eV). The energy resolution along the emission energy in mRIXS is about 0.2 eV, 347 which is sufficient for separating the intrinsic S-L signals from the strong C-K (2 nd order) background that are different for about 10 eV. The X-ray exposure area on the samples were kept moving throughout the mRIXS experiments to avoid irradiation damage.

X-ray photoemission spectroscopy (XPS).

XPS spectra were collected on a sample (analyzed area = 300×700 μm 2 ) with a Kratos Axis Ultra spectrometer, using focused monochromatic A-22 Al Kα radiation (hν = 1.4866 keV). The pressure in the analysis chamber was around 5 × 10 -9 mbar. The binding energy scale was calibrated using the C 1s peak at 285.0 eV from the invariably present hydrocarbon contamination (for the pristine sample), and using the S 2-position of the S 2p3/2 at 160.7 eV for a better accuracy (for the other cycled samples). Peaks were recorded with constant pass energy of 20 eV. Core peaks were analyzed using a nonlinear Shirley-type background. 348 The peak positions and areas were optimized by a weighted least-squares fitting method using 70 % Gaussian, 30 % Lorentzian line shapes. Quantification was performed on the basis of Scofield's relative sensitivity factors. 349 For the UV-Vis spectra, the starting spectra before cycling was taken as the blank spectra. It is clear that when the cell was cycled up to 3 V (green curve in (a)), the spectra do not show any trace of dissolved species. In contrast, the spectra drastically change when charged to >3 V (red curve in (a)) indicating the presence of copious amount of dissolved species. However, the position of the absorption peaks is rather reminiscent of alkyl carbonates than polysulfides. 307 Hence, electrolyte decomposition might occur in presence of aggressive S2 n-species within the electrode material. See reference 307 for the experimental details. Note that, when we cycled the Li1.13Ti0.57Fe0.3S2 in all-solid-state-battery, we found a stability up to > 4.2 V, which will be shown in the Appendix Thus, the cationic contribution, which was assumed to be manifested in the slope (S), involving the Fe 2+/3+ redox, can be decoupled from the presumed anionic plateau (P). For this purpose, the upper cut-off voltage was gradually increased in subsequent cycle by 0.5 V during charging. The dQ/dV curves reveals that initially charging leads to sloppy dQ/dV curves (green curves). Only after charging above ~2.67 V, the intense dQ/dV peaks appear (in charge and discharge, blue curves), which can be regarded as the initiation of the anionic process. The corresponding voltage profiles indicate only a capacity of ~40 mAh•g -1 is obtained from solely cationic redox (i.e., before the blue lines appear) which corresponds to ~0.16 e -transfer. This gives a hint that the Fe 2+/3+ redox (capable of 0.3 e -transfer) does not complete prior to the anionic redox. edge mRIXS results of the pristine powder, fully-charged electrode, and full-discharged electrode and full-discharged electrode (but with a different color scale to emphasize the weak-signals). The strongest line feature in all the maps is from the elastic peak, i.e., the same energy of the outgoing fluorescence photons (emission energy, horizontal axis) and the incident photons (excitation energy, vertical axis). The mRIXS is collected at each excitation energy that is used for X-ray absorption spectroscopy, 350 but provides the information of the emission energy of the fluorescence photons. 297 2019) [START_REF] Nasu | Amorphous Na 2 TiS 3 as an Active Material for All-solid-state Sodium Batteries[END_REF] . They cycled the phase vs Na15Sn4 anode in ASSB, using Na3PS4 glassceramics as solid electrolyte. Figures are adapted from references [START_REF] Pele | Nouveaux materiaux d'electrodes pour microbatteries au sodium[END_REF][START_REF] Nasu | Amorphous Na 2 TiS 3 as an Active Material for All-solid-state Sodium Batteries[END_REF] . Note that only a capacity of ~110 mAh•g -1 was observed which is much less than what we found in Na-half cell, probably because of the electrolyte used.

Appendix

A2.3 Supporting Tables

Appendix Table A2. 1.27( 4)

1.19 (8) 1.5 (2) 1. 75(9) 1. 75(9) A3 Supporting Information for Chapter IV

A3.1 Methods

All electrochemical tests were performed similar to the way described in previous chapter. In case the samples were used as self-standing electrodes, they were prepared by mixing the sample with polytetrafluoroethylene (PTFE) polymer and conductive carbon Csp (in 75 : 5 : 20 wt%), with a sample loading of ~4 mg•cm -2 .

GITT experiments were performed in two or three electrode (as specified in the text) Swagelok type cells, after 7 cycles (the first cycle at C/20 and the next six at C/5) so that the voltage profiles are stabilized. All electrochemical cycling and EIS measurements (in 10 mHz -200 kHz frequency range applying a 10 mV sinusoidal wave) were performed with BioLogic potentiostats. List of Tables Table I.1

A3.2 Supporting Figures

List of Publications

This thesis includes the following articles:

1)

Sujoy Saha, Gwenaëlle Rousse, François Fauth, Vladimir Pomjakushin, and Jean-Marie Tarascon; Influence of Temperature-Driven Polymorphism and Disorder on Ionic Conductivity in Li6Zn(P2O7)2. Inorg. Chem. 58, 1774-1781 (2019). http://pubs.acs.org/doi/10.1021/acs.inorgchem.8b01800

2)

Sujoy Saha, Gwenaëlle Rousse, Ignacio Blazquez Alcover, Matthieu Courty, Daniel Alves Dalla Corte, and Jean-Marie Tarascon; Polymorphism in Li4Zn(PO4)2 and Stabilization of its Structural Disorder to Improve Ionic Conductivity. Chem. Mater. 30, 1379-1390 (2018). http://pubs.acs.org/doi/10.1021/acs.chemmater.7b05139

3)

Sujoy Saha, Gaurav Assat, Moulay Tahar Sougrati, Dominique Foix, Haifeng Li, Jean Vergnet, Soma Turi, Yang Ha, Wanli Yang, Jordi Cabana, Gwenaëlle Rousse, Artem M. Abakumov & Jean-Marie Tarascon; Exploring the bottlenecks of anionic redox in Li-rich layered sulfides. Nat. Energy 4, 977-987 (2019). http://dx.doi.org/10.1038/s41560-019-0493-0

4)

Sujoy Saha et al., Improvement of Ionic Conductivity in Na4-xZn1-xMx(PO4)2 (M= Ga, Al) driven by a structural transition, Manuscript in preparation.

5)

Florencia Marchini and Sujoy Saha et al., A Li-rich layered sulfide as cathode active material in all-solid-state Li-metal batteries, Manuscript submitted.