General introduction

The constant growth of world population along with the enhancement of living standards leads to a continuous increase of the energy demand, forecasted to increase by 30% from current 1.5•10 14 kWh/year in the two following decades [1][2]. Around 80% of the currently produced energy is coming from the combustion of fossil fuels, e.g. coal, petroleum and natural gas.

However, due to its limited feedstock, its extraction is likely to decline after 2025 [3]. Apart from its limited availability, fossil fuels burning is the main source of carbon dioxide emission into the atmosphere. Counting for 72 % of all the greenhouse gases emitted [4], CO2 is the main contributor to the climate change with all the threatening consequences of global warming such as sea level elevation, ice melting, more frequent and stronger hurricanes, extreme heat and wildfires, flooding, etc.

Based on numerous scientific reports, the United Nations (UN) framework convention on climate change has been signed in Paris in 2015 by 174 states and the European Union as a political answer to the uprising issues [5]. The agreement stated the aim to keep the maximum of the global temperature rise well below 2°C by tackling the problem of the greenhouse gas emissions. Following shortly after, at the 21 st Conference of the Parties to the UN Framework Convention on Climate Change (COP21) the emission-cutting pledges were made by all the 170 countries attending [6]. The adopted strategy of decarbonization is primarily based on the replacement of the fossil fuels by the alternative sustainable and renewable energy sources.

Nowadays, among all the alternative energy sources employed, renewables and nuclear have a 34% share of the total electricity generation [1]. As for the nuclear, even though the CO2 release in the production cycle is very low (90-140 g of CO2 per kWh) compared to conventional energy sources (1000-1100 g of CO2 per kWh for coal), it can hardly be called sustainable because of the finite and toxic nature of the Uranium-235 and the harmful radioactive wastes after nuclear chain reactions. Except for geothermal (1.3•10 10 kWh), most of the renewable energy sources originate directly or indirectly from solar activity. Wind power (4.9•10 14 kWh), bio-energy (1.1•10 11 kWh), hydropower (1.1•10 12 kWh), and marine power (3.6•10 13 kWh) are secondary products of the solar irradiation of the Earth due to the transformation of the solar radiation into thermal heat, photosynthesis, evaporation of the water, and waves respectively (potential energy generation capacities numbers of installed facilities from [7][8]).

The amount of energy produced at the Sun surface is tremendous, ca. 9.5•10 22 kWh, as it originates from the hydrogen to the helium gas transformation through the nuclear fusion reaction. Even taking into account losses due to the back reflections, scattering and absorption in the atmosphere, the energy available for the harvesting at the Earth surface (1.1•10 18 kWh/year assuming 1500 h/year of sunshine) can still easily satisfy the actual yearly energy demand of 1.5•10 14 kWh. More than that, the amount of solar energy reaching the earth in one hour (7.3•10 14 kWh) corresponds to the annual energy consumption of the world population.

There are two main reasons explaining why solar is still not the primary energy source despite its credibility. The first is that the light-to-electricity conversion used to be not cost competitive compared to the nuclear or fossil fuels [9], though it has seriously changed over the last years owing to the continuous drop in the cost of silicon-based photovoltaic cells ( -86% from 2009 to 2017) in parallel to fossil fuels' price rising (Fig. 1.1) [10]. Nowadays, 50 USD for a megawatt-hour (MWh) of solar energy becomes competitive with the cost of electricity produced from the natural gas (60 USD/MWh) and twice cheaper than that produced from coal (102 USD/MWh). Therefore, even though the introduction of the solar energy still requires substantial financing of new infrastructures and electronic management systems for smart grid implementation, the technology becomes worth the investment in the long-term. The second issue is more difficult to solve as it concerns the intermittent nature of the solar energy on long time-scale range from second up to day extent (Fig. 1.2). The time-dependent solar energy output and its unpredictability constrains the direct implementation of the photovoltaics as it cannot be produced in a steady and reliable manner to supply society's fluctuating electricity demands at any time. The current application of PV technologies in connection with a grid is possible due to the backup of energy supply from another source (often a stable non-renewable one). It allows to level on time the energy production to the energy demand from customers [11]. [12].

For off-grid application, the costly backup energy generation systems, often based on fossil fuels, are required to manage the PV power output disruption [13]. However, constant energy output can be also assured by storage of the excess of energy produced during peak generation or lower demand time for later on-demand supply. The generated energy can be stored chemically (i.e. batteries, capacitors or hydrogen), mechanically (pumped hydroelectric power, compressed air or flywheels), or in thermal systems [14]. Each of these technologies has its application place depending on the employment conditions. However, the most appealing among them is electrochemical storage systems such as batteries due to easy integration, flexibility of storage time, and high versatility from µWh to MWh energy density scale.

However, the incorporation of the electrochemical storage part into the PV generation system at least doubles the initial costs, rising up to 28 000 -33 000 USD for a 5.6 kW PV plant associated with a 3 kW / 6 kWh lithium-ion battery system including installation costs, DC/DC converter/inverter and battery management system [15]. Thus, despite the mature technology, the cost return time of the installation remains too long (20-30 years) compared to the service lifetime of silicon PV and Li-ion batteries.

With the aim to lower domestic cost of PV installation as well as to prompt new innovation and science streams, the development of a single device combining both energy conversion and storage is highly desirable but also very challenging from a scientific point of view when dealing with the functions gathered at the molecular level. One shorter-term alternative is based on a hybridization of the two technologies. It consists in developing a single device comprising two compartments, one corresponding to PV, such as dye-sensitized solar cell as often proposed, and a lithium-ion battery part sharing a same electrode in common [16]. This approach is potentially less expensive and easier for integration compared to the conventional external coupling. The second approach, inscribed to longer term perspectives, consists in tailoring molecular assembly to combine both sunlight conversion of energy and its chemical storage. This approach was first demonstrated on inorganic materials by Cahen et al. [17] and then Tributsch et al. [18], who highlighted an interfacial light-driven ion transfer in the mixed electronic/ionic semiconductor by making use of the efficient charge separation process within the depletion layer. A direct consequence is that such a mechanism is resticted to the surface, thus representing one important bottleneck of such an approach. To overcome this issue, ca. 5 nm crystalline anatase TiO2 nanocrystals were synthesized in our group affording to drastically enhance the surface-to-volume ratio [19][20]. As a result, a quantitative photorecharge reaction of the electrode was obtained in ca. 2 hours when exposed to incident standard A.M. 1.5G illumination [21]. This proof-of-concept showed for the first time that the band gap excitation can induce a complete electrode recharge, and paved the way to new scientific and technological direction towards bifunctional materials combining conversion and storage of sunlight power. Nevertheless, further developments call for deepening our current understanding on these photoinduced charge transfer processes in order to maximize the performance in terms of stability and recharge time and developing an appropriate counterelectrode affording a regenerative full photorechargeable Li-ion battery device by collecting free-electron carriers from the photoanode.

One objective of this thesis is to explore what are the optoelectronic properties of most common class of insertion materials which are in discharged or charged states (i.e. lithiated or not) and investigating how these properties are modified during the battery operation (i.e. upon lithium insertion / deinsertion). It has been already discussed in the literature that the potential of intercalated ions can contribute to the electrode free energy causing a shift of the energy band edges, band gap, and thus affecting the Debye length and interfacial electric field of the spacecharge layer [22][23]. With this aim, the second chapter of this thesis is dedicated to describing the results on the in-depth study of the evolution of the optoelectronic properties and dynamics of the ultra-fast charge transfer processes at electrode/electrolyte interface of the anatase TiO2 upon lithium insertion on the basis of in situ / in operando experiments using UV-Visible absorption spectroscopy, steady-state and time-resolved fluorescence spectroscopy and Mott-Schottky electrochemical capacitance measurements.

One important obstacle in the development of fully photorechargeable battery system is finding a suitable counter electrode for which optical absorption is complementary to the photoanode, while it has the opposite-type of major carriers and suitable position of conduction and valence bands with respect to the other electrode in order to collect the major carriers. For this reason, the third chapter will describe our work related to the creation of a material database gathering the optoelectronic properties of the most widely used battery electrode materials in order to reveal which materials present potentialities for this application but can also present potentialities for other type of applications such as in the fields of electrochromism or photocatalysis. Three groups of materials are investigated: transition metal oxides (Li4Ti5O12, LiCoO2, MoO3, WO3, CuO, Bi2O3, LiMn1/3Ni1/3Co1/3O2), transition metal sulfides (TiS2, MoS2, WS2), and lithiated polyanion strtuctures (LiFePO4, LiMnPO4, LiVPO4F).

Chapter 1. State of the art on photoelectrochemical energy storage

Solar energy harvesting, conversion, and storage

History of photovoltaics

The story of photovoltaics starts back in 1839 when French scientist Edmond Becquerel presented the photoelectric effect to the Academie des Sciences in Paris. In his experiment, generation of a current induced by light action was noticed when illuminating a halogenated metallic electrode connected to a platinum electrode placed in acidic solution [1]. It is only a century later that the first practical photovoltaic cell with worth power conversion efficiency (PCE) of 6 % was developed by Chapin, Fuller, and Pearson based on silicon p-n junction [2].

The discovery of p-n junction has fostered the development of the off-grid energy production devices in the contest of Cold War's spacial programs. Since then, the vast majority of fabricated solar cells are based on either single-crystalline or large grain poly-crystalline silicon.

Still nowadays, the PV technology is dominated by p-n silicon hetero-junction technologies with a market share of ca. 90 % owing to the performances, cost and elemental availability as Si is the second most abundant element in Earth. The highest reported PCE value is 26.1% for Si single crystal technology [3] and 26.6% for Si heterostructure [4]. The values are approaching the 32% Shockley-Queisser limit determined for single junction solar cells having a band gap of 1.10 eV (Si) under ASTM G173-03 conditions [5][6].

After the two petrol crisis in the 70's, development of a second generation of solar cells was explored to democratize solar energy to large public using a strategy to reduce modules manufacturing cost with new technologies. This second generation explores thin-film inorganic technologies based on lower cost amorphous or nanocrystalline silicon PV (14.0%) or on polycrystalline chalcogenide III-V and II-VI heterojunction cells, amongst CdTe (22.1%), CuInGaSe (22.9%), and GaAs with a record of 28.9% achieved by Alta Devices (Fig. 1.3) [7][8]. Multi-junctions to panchromatically capture the light through complemental junctions allowed to reach the world record efficiencies of 38.8% using a 5-junction GaInAs/GaInP/GaAs/AlGaInAs/AlGaInP cell under standard A.M.1.5G illumination conditions (100 mW/cm 2 , 1 sun) [9] and as high as 46.0% with a 4-junction solar cell based on GaInP/GaAs//GaInAsP/GaInAs under concentrated sunlight conditions (50.8 W/cm 2 , 500 sun) [10]. Nevertheless, the toxicity of some of these elements, their scarcity and the extremely high production cost of multiple junction modules restrain such technology to laboratories and to spatial applications niche [11]. [8]).

At the end of the 80's, the development of conducting polymers for organic electronics paved the way towards a third solar cell technology generation which promises low-cost and new market opportunities thanks to roll-to-roll printing processing yielding flexible solar cells.

Organic materials also offer an added value with respect to inorganic counterparts owing to their versatility, in particular band gap engineering. This has led to organic photovoltaics (OPV) [12], quantum dots solar cells, which remains inorganic but with the same aim of band gap tunability [13], hybrid organic / inorganic devices such as dye-sensitized solar cells (DSSCs) [14], and more recently halide perovskite solar cells [15]. Among all the abovementioned technologies, even though not the best performing (PCE = 14.3% [39]), DSSCs stand out as the most mature technology in terms of stability. They also gather advantages in converting very efficiently low diffuse light reaching as high as PCE = 29.3% [40]. In addition, the DSSC technology offers power conversion efficiency independent of external temperature, tunable coloration, bifacial conversion, and semi-to full transparency possibilities.

Working principles of conventional solar cells and dye-sensitized solar cells

Conventional solar cells are based on a p-n junction between either differently doped semiconductor of the same nature (homojunction) or two different types of semiconductor (heterojunction). When the quantum of light with an energy equal to or greater than the semiconductor's band gap is absorbed, an electron is promoted from the valence band into the conduction band leaving a positively charged hole in the valence band. The formed excitons are diffusing towards the junction. The interfacial electric field is sufficient to dissociate electron-hole pairs, and trigger the transportation of electrons towards the n-type semiconductor and the holes in opposite direction towards the p-type semiconductor [18] (Fig. 1.4). Thus, the p-n junction allows to spatially and energetically separate the excitons while avoiding their recombination, leading to the establishment of a photovoltage which under short-circuit condition drives the carriers towards the external load. The working principle of dye-sensitized solar cells has opened a new paradigm in PV field because of (i) the utilization of nanocrystalline particles instead of defect-free single crystals, (ii) a principle relying on electrochemical processes, (iii) no electric field separating the excitons, (iv) separation of charge carrier generation and charge transport processes. The technology is composed of two electrodes, a photoanode and a counter-electrode (CE), separated by a liquid electrolyte containing a redox couple to regenerate the dye (e.g. I 3 -/I -) (Fig. 1.5) [14]. In the photoanode, light absorption and charge separation take place in the dye molecules, typically based on polypyridyl ruthenium complexes, which are anchored on the wide band gap anatase TiO2 nanoparticles (Eg=3.2 eV). The photogenerated electrons located in the lowest unoccupied molecular orbitals of the dye's (LUMO) are injected from the triplet states of Ru(+II) into the conduction band of TiO2. The oxidized dye molecule (S + ) is regenerated by a two-hole transfer process leading to the oxidation of I -to I 3 -. The free electrons from TiO2 are collected and transferred through the back contact towards the counterelectrode to ensure closing the electrochemical loop by reducing I 3 -back to I - [14,19].

Principles of energy storage in electrochemical batteries and capacitors

Electrochemical batteries and capacitors are widely used for nomade applications as a power storage feedstock [20]. The Li-ion battery technology is leading the field owing to higher cell voltage (3.2 -4.2 V), high energy density Wh/kg), low self-discharge (0.03% to 5%), and high stability (>1000 cycles lifespan) [20][21][22]. The working principle of a battery is based on the conversion of electrical energy stored in a form of a chemical energy by means of redox reactions on transition metal-based electrode materials where the charges are compensated by ion insertion into or deinsertion from the structure. A typical configuration consists of two electrodes, one at high potential (cathode) and one at low potential (anode), separated by a lithium-ion conducting medium (solid or liquid) (Fig. 1.6 (a)). The discharge of the battery is ensured by a release of lithium ions from the negative electrode into the electrolyte with a concomitant lithium insertion into the positive electrode. This leads to a simultaneous transfer of electrons in the same direction via the external circuit. This process can be inversed by imposing an external current such as from the grid to provide the battery recharge.

In supercapacitors, the electrical energy is stored in the form of the surface accumulation of ions [23]. A typical device consists of two conductive high surface-area electrodes separated by an electrolyte solution (Fig. 1.6 (b)). The energy is stored via the Helmholtz double-layer formation at both electrode/electrolyte interfaces by the charge carriers (electrons or holes) from the electrode side and the ions (cations or anions respectively) from electrolyte. This technology stands out for the high capacitance, fast and reversible charge storage/release function and very high lifespan (>10 5 cycles). However, in contrast to electrochemical batteries, supercapacitors offer a high power but a low energy density as showed in the Ragone plot (Fig. 1.7) [24][25]. Ragone plot for various energy storage devices [25].

Photorechargeable energy storage devices

As an alternative to the coupling of photovoltaic and energy storage technologies, a range of studies have been carried out aiming at developing a two-in-one solar energy conversion and storage device by means of hybridization of technologies.

Most of such devices reported can be classified into one of the two following sub-groups depending on the engineering of the cell: (i) devices with a multifunctional electrode and (ii) separate planar sub-devices. In the first case, the charge generation material and charge storage material are put in junction forming a multifunctional electrode. The assembly is quite similar to that of a dye-sensitized solar cell (Fig. 1.8 (a)) [26]. Here, charging of the device takes place at the multifunctional electrode at open-circuit condition, and the system is discharged at closed-circuit. In the second case, charge storage and charge separation functions happen at two spatially separated electrodes. Most of these devices are assembled in a three-electrode configuration (Fig. 1.8 (b)) [27]. The photorecharge takes place when the illuminated photoanode is connected to the charge storage electrode. The photogenerated electrons are transferred from the photoanode to the charge storage electrode via the external circuit. The discharge takes place between the charge storage electrode and the internal electrode while leaving the photoanode disconnected.

Devices based on multifunctional charge generation and storage electrode

The first photorechargeable device based on multifunctional electrode was proposed by Miyazaki et al. back in 2000 using TiO2 and carbon fiber (CF) as energy generation and energy storage materials, respectively [26]. The photorecharge was ensured at the multifunctional electrode at open-circuit condition by means of a charge generation in the anatase TiO2 particles and its subsequent storage in the carbon fibers leading to the insertion of lithium ions into CFs.

The discharge was performed through an external load of 1kΩ at short-circuit condition. The authors reported a rather small discharge capacity of ca. 250 µC/cm 2 which has been attributed

to the low rate of lithiation of the carbon fibers under illumination [28].

The majority of the following devices were assembled accordingly to a dye-sensitized solar cell configuration or at least using a dye-sensitized TiO2 photoanode to improve light harvesting of the electrode. In 2002, Orel et al. proposed a device including a photoanode composed of two layers: an external layer of nanocrystalline TiO2 particles sensitized with N3-dye that has a role of carrier generation under illumination and an internal layer of WO3 nanoparticles that are able to store the electrons throughout lithium insertion in WO3 (Fig. 1.9) [29]. The intimate contact between these two redox active layers is responsible for self-discharge as mentioned by the authors. This leads to the chemical oxidation of LixWO3 by TiO2. The device reached an opencircuit voltage of 0.6 V and a discharge capacity of 500 µAh/cm 2 after 1 hour of photocharging under standard A.M. 1.5G illumination conditions (100 mW/cm 2 ). Later on, Nomiyama et al.

proposed a mixture of TiO2 particles to generate carriers and doped polyaniline (PANI) polymer for a storage function, which is ensured by anions (ClO4 -or SO4 2-) transfer from PANI to the electrolyte [30][31]. The authors consider TiO2/PANI as a promising bifunctional composite because of its good specific capacity between 60 and 120 mAh/g and an energy conversion efficiency over 12% [32]. However, experimentally the overall electricity-to-storage conversion efficiency of the device was reported to be less than 1% owing to the low charge transfer rate between TiO2 and PANI upon photorecharge. Indeed, the halftime of charge transfer is 10 -1 second which is ten times longer than the lifetime of the excited electrons in TiO2. Finally, Zaghib et al. recently presented results on N719 dye grafted onto LiFePO4 electrode [35]. The photorecharge of the electrode was ensured under standard A.M. 1.5G conditions leading to a photorecharge time of ca. 10 hours under O2 atmosphere and above 100 hours under Ar. The device exhibited 340 mAh/g discharge capacity stable over 15 cycles by combining discharge load and light illumination. The authors suggested that the twice higher capacity than theoretically expected is the result of the two concommitant mechanisms of galvanostatic discharge and photorecharge. Nevertheless, the band alignment and the photorecharge mechanism proposed remain unclear and can be largely debated. This feature will be discussed in more details in Chapter 3.

Separate planar sub-devices

The approach of separate planar sub-devices lies in the presence of two compartments, one photovoltaic (DSSC or OPV) and one for electrochemical storage (battery or capacitor). In most cases, the photoelectrochemical cell is composed of three electrodes. The charge storage electrode alternately plays the role of counter electrode for the photovoltaic compartment and working electrode for the battery part. Below we present these devices dividing them into two sub-groups based on the charge storage type, i.e. capacitor storage or battery storage. The carbon nanotubes (CNT) are known to have a good specific capacitance (54 F/g) and an excellent electrochemical stability [39]. When incorporated into a photorechargeable electrochemical device, it led to the highest solar-to-storage conversion efficiency, corresponding to the overall device efficiency achieved up to date. They reported a value of 5.12 %, where a capacitance of 48 F/g was obtained under 100 mW/cm 2 irradiation [40]. The discharge capacity of 1.36 µAh/cm 2 was maintained over 100 cycles. The incorporation of PANI into CNT electrode increased the electrode's capacitance to 208 F/g. However, it caused a decrease in the storage capacity from 84% to 70% resulting in a drop of overall device efficiency down to 4.29%. After photorecharge, a discharge capacity of 56 µAh/cm 2 was reached, for which a capacity retention of 80% was obtained after 100 cycles. The combination of poly(3,4-ethylenedioxythiophene) (PEDOT) with the multi-walled carbon nanotubes has also increased the initial capacitance of CNT electrode to 95 F/g, which corresponds to an energy density of 3.15 mWh/g [41]. The authors assign this improvement to originate from a lowering in internal resistance of high surface area and high conductivity PEDOT. When employed alone, PEDOT reached 0.52 F/cm 2 storage capacity under standard irradiation [42].

Capacitor storage

Most

However, note that authors have used different units (F/cm 2 vs F/g) which in absence of some information related to the size and loading of the electrode impede a direct comparison between the numbers. Both CNTs and PANI have also been incorporated in combination with dyesensitized TiO2 into so-called "energy-fibers" for smart textiles. The device exhibits rather limited performances of up to 3.32 mF/cm 2 under 100 mW/cm 2 irradiation [43][44] and 20 mF/cm 2 under 65 mW/cm 2 irradiation [45]. Bae et al. presented a single fiber-based device including ZnO nanowires and graphene as principal energy conversion and storage materials.

The electrode reaches a storage capacitance of 0.4 mF/cm 2 after photorecharge under standard A.M. 1.5G conditions [46]. The nanocomposite of PVDF and ZnO nanowires as a capacitive energy storage material in combination with N719-sensitized TiO2 exhibited 59.4 mAh/g storage capacity achieving 3.70% overall device efficiency under standard conditions [47].

Finally, single-walled carbon nanotubes were also incorporated into an OPV device based on P3HT:PCBM ( poly(3-hexylthiophene):1-(3-methoxycarbonyl)propyl-1-phenyl [6,6]C61 ).

This assembly yielded a specific capacitance of 28 F/g after 70 seconds of illumination under 100 mW/cm 2 and 3.39% overall device conversion efficiency [48]. However, only 47% of initial discharge capacity of 4.9 mAh/g was maintained after 10 cycles. Using two OPV in series allowed reaching 79.8 F/g storage capacity while exhibiting a higher power density of 10 W/g and energy density of 5 mWh/g.

Battery storage

A range of other inorganic and organic materials has been used for developing photoelectrochemical devices including a battery. A surface-oxidized tungsten oxide in combination with a classical DSSC part reached 141 µAh/cm 2 discharge capacity with 0.4 V discharge output voltage after 10 hours of photorecharge under standard illumination conditions (A.M. 1.5G) [49]. Zhao et al. presented a similar device based on nanoparticles of WO3 and Z907-sensitized TiO2 which reached 40.6 mAh/g (24.4 µAh/cm 2 ) discharge capacity after a photorecharge under 7.5 mW/cm 2 irradiation during 10 minutes [50]. In the following cycle, the discharge capacity dropped to 34 mAh/g and was further maintained during 4 cycles only.

The authors ascribed this capacity loss to the possible reduction of WO3 to W and side reactions with the electrolyte leading to a SEI formation. Gao et al. presented a photorechargeable battery based on WO3-coated carbon nanotubes with a TiN nanotube arrays on Ti mesh as an electrocatalytic electrode for I3 -to I -reduction (Fig. 1.12 (a)) [51]. After 30 minutes of photorecharge under 100 mW/cm 2 , the open circuit voltage of the complete device reached 0.7 V and yielded a discharge capacity of 139 µAh/cm 2 . In the following cycle, the discharge capacity dropped to 124 µAh/cm 2 (11% loss) and remained stable over 10 cycles. In contrast to the above described three-electrode systems, Jiang et al. presented a photorechargeable device based on tungsten oxide hydrate in a two-electrode assembly [52]. The nano-WO3•H2O/CNTs/PVDF composite electrode was combined with N719-sensitized TiO2 photoanode providing a discharge capacity of 383 µAh/cm 2 after 10 minutes of photorecharge under 100 mW/cm 2 irradiation. The overall device efficiency was 2.12 %. nanotube array electrode, one side of which is a photoanode of a dye-sensitized solar cell for charge generation and the other is a charge storage electrode combined with the LiCoO2 positive electrode in the battery compartment (Fig. 1.13 (a)) [53]. In order to provide high enough photovoltage for the battery recharge, the authors integrated three series-connected tandem N719-N749 DSSCs to afford above 2 V generation under illumination (Fig. 1.13 (b)).

Photorecharging power pack under standard illumination for 440 seconds resulted in the voltage rise up to 3 V and a consequent discharge capacity of 33.9 μAh. The entire conversion efficiency of the device was limited to ∼0.82%, which authors assign to the growing inner resistance of the cell over the operational time. A set of organic materials have also been proposed in photorechargeable systems for ensuring the energy storage function. Segawa et al. were first to present the polypyrrole (PPy) as a charge storage electrode in a three-electrode cell [27,54]. Here, the storage of photogenerated energy is performed by deinsertion of ClO4 -from the PPy structure. A maximum discharge capacity of 0.5 μAh/cm 2 was achieved after 30 minutes of photorecharge under standard A.M. 1.5G conditions. This capacity, although low compared to the inorganic counterparts, remained stable over 10 cycles. An overall device energy conversion efficiency of 0.5 to 1% was reported. A few studies presented photoelectrochemical devices based on different light harvesting materials and redox couples. Combination of WO3-decorated BiVO4 photoanode and aqueous I3 -/I -and Br3 -/Br -redox couples as catholyte and anolyte respectively resulted in 250 µAh discharge capacity after first photorecharge and 0.4 V output voltage stable over 20 cycles [60].

The overall device efficiency reached 1.25%, however without any precision on the type of light and the power used, thus making the direct comparison with the other device architectures impossible. Utilization of other combinations between absorbers and redox couples have also been presented, i.e. CdS associated to V 3+ /V 2+ [61] and α-Fe2O3 associated to Fe(CN)6 Jin et al. presented a photorechargeable redox flow battery composed of regenerative silicon solar cell in association with 9,10-anthraquinone-2,7-disulfonic acid (AQDS) and 1,2benzoquinone-3,5-disulfonic acid (BQDS) based redox couples as catholyte and anolyte, respectively [64]. After a photorecharge under 100 mW/cm 2 during 7.3 hours, the device reached an excellent discharge capacity of 3.5 mAh/mL leading to a solar-to-storage efficiency of 1.7%. Silicon absorber and AQDS-based catholyte have also been studied in combination with Br3 -/Br -redox couple based anolyte [65]. After a photorecharge under A.M. 1.5G illumination conditions, the battery reached a discharge capacity values of up to 730 µAh/mL and 0.78 V stable output voltage with a 3.2% solar-to-storage conversion efficiency.

The main characteristics and performances of all above-discussed photoelectrochemical storage devices are gathered in Table 1.1.

Charge generation and storage at the molecular level

One longer-term approach to overcome the solar intermittency issue is the development of materials capable of both energy storage and energy conversion at a molecular level. Both organic and inorganic materials have been considered for this application. The second approach addressed bi-functional inorganic materials and started back in the 80's by Gerischer et al., who extended the fundamental knowledge of the semiconductor/electrolyte interface and the theory on photoinduced ion transfer reactions through the interface [67][68][69].

In the same period of time, Cahen et al. were also working in similar direction and reported the first photoelectrochemical system capable of conversion and storage of solar energy based on CdSe electrode, where the partial reversible substitution of Se with S from electrolyte takes place under illumination [70]. A few years later Tributsch et al. experimentally proved the possibility of the light-driven intercalation reactions in p-type layered compounds (ZrS2 and ZrSe2) and discussed the theoretical prediction of the photoinduced ion deinsertion from the ntype Hf dichalcogenides for simultaneous energy conversion and storage [71], which was experimentally demonstrated a couple of years later [72][73]. However, the main drawback of transition metal dichalcogenides was the too low quantum efficiencies and a quasi-metallic behavior leading to a loss of the photo-effect. Tributsch et al. considered the possibility of using a typical for batteries Li + ion for photo-insertion but abandoned the idea because of lithium's low ionization energy and too high electropositivity for the practical use [74]. Therefore, they focused on the Cu + insertion systems that are compatible with a copper metal counter-electrode.

The most promising performances were achieved on Cu6-xPS5I [75] and Cu3PS4 [74] p-type semiconductors, where Cu 2+ ion insertion into structure was triggered by homogeneous electron transfer to the transition metal. With similar approach, Tributsch et al. reported a successful proton deinsertion from the n-type HxTiO2(B) by the action photogenerated holes oxidizing Ti 3+

to Ti 4+ [76]. However, focusing almost exclusively on the photoinduced phenomenon, the storage capability of these systems has been relatively undertreated.

The principle of the photoinduced insertion/deinsertion reaction lies in the effective charge separation in the depletion layer of the semiconductor [68][69]. When a semiconductor electrode is immersed in the electrolyte, the equalization of the Fermi levels between the semiconductor and the electrolyte (chemical potential) takes place (Fig. 1.18). It results in the space charge depletion of the minority charge carriers and bending of semiconductor band edges at the surface. When the light is absorbed in the depletion layer, the photogenerated charges are effectively separated due to the migration of the main charge carriers into the bulk of semiconductor and the minor charge carriers towards the surface. In absence of a suitable redox mediator in the electrolyte, the minor charge carriers accumulate at the semiconductor surface, trigger redox reaction with the transition metal and interfacial ion transfer for charge compensation. Unfortunately, so far, the interfacial light-driven ion transfer in mixed electronic/ionic semiconductor is only an extreme surface reaction limited to the Debye length (space charge layer width). This was confirmed by Hirakawa et al. on the CuFeTe2-based electrode, in which the storage capability was only 0.562 mWh/g, thus representing less than 1% of the theoretical energy density (60 mWh/g) [77].

Conclusions

In this chapter, we presented the evolution of the solar cells development since the discovery of the photo-effect until today. We introduced the principles of the electrochemical energy storage devices, i.e. batteries and capacitors, commonly used to store energy. The concept of combining the two functions to solve solar intermittency or to prolong battery capacity for nomad applications has been reviewed. We presented them into two groups depending on the approach, namely hybridization of two technologies in one device or development of bifunctional materials capable of combining energy conversion and its storage at molecular level.

The main advantage of the solar-to-chemical conversion and storage systems based on hybridization lies in the use of the third generation PV benefiting from the ease of its integration with the storage counterpart and the potential overall cost reduction. Nevertheless, the devices developed so far are highly limited in terms of performances (maximum 1.0 V output voltage and the overall device efficiency of 5.12%) and stability. One of the reasons is that up to now focus has been on the development of the prototypes and hardly any on the optimization of the system, which often resulted in poorly matched conversion and storage parts leading to inefficient electron transfer. In addition, the issue of battery self-discharge when the PV part is exposed to lower light illumination has never been tackled. One complication in analyzing the publications on photorechargeable devices arises from the different sources of illumination and the different units used to report the performances.

The longer-term and more fundamental second approach to overcome the solar intermittency is based on combining charge generation and charge storage functions as a result of light-induced charge separation within the depletion layer and ion-transfer reaction in a mixed ionic/electronic conductor. Such materials open up new scientific and technological directions that can potentially close the gap between PV and batteries in terms of functions as well as cut the cost of energy conversion and storage device. Currently, the main limitation of such device is the low performances nested either in poor charge separation or in photoelectrochemical reactions confined to the extreme surface of the material.

Introduction

The proof-of-concept of an interfacial light-driven ion transfer at the interface between a mixed electronic/ionic semiconductor and electrolyte has been presented by Tributsch et

𝑊 = ( 2∆𝛹𝜀𝜀 0 𝑞𝑁 ) 1 2 ⁄ (2.2)
where ∆𝛹 is the electrical interfacial energy drop in the space charge layer (directly related to the flat band potential Vfb), 𝜀 0 the permittivity of the vacuum (8.85•10 -12 F/m), 𝜀 𝑟 the relative permittivity (dielectric constant) of the material to that of the vacuum, 𝑞 the elementary charge (1.6•10 -19 C), and 𝑁 the main charge carrier concentration (donor density). One additional approach led at LRCS under the post-doctoral work of Dr. Christian Andriamiadamanana prior to my Ph.D. thesis, was to downsize the particles towards nano-scale to promote the surfaceto-volume ratio of the active material and thus to increase the yield of the photoelectrochemical reaction per mass unit.

The anatase polymorph of titanium dioxide appears as an excellent candidate owing to its good semiconducting properties, a high dielectric constant of 31-48 [3][4] and low charge carrier density of 10 17 -10 19 cm -3 [5][6][7]. It displays a reversible lithium insertion / deinsertion process at a redox potential of ca. 1.75 V vs. Li + /Li making this material attractive as a negative electrode for lithium-ion batteries [8][9]. Prior to this Ph.D. work, we elaborated a synthetic procedure under room-temperature control yielding to nanocrystalline anatase TiO2 crystals of ca. 4-5 nm in size [10].

This second chapter starts with a summary of this preliminary work that laid the background for the current thesis, including the details of the nanoparticle TiO2 synthesis, powder and electrode characterization, and the photoelectrochemical performances. In the following parts, the results gathered during my thesis will be presented and discussed aiming at better understanding of the mechanisms behind the photorecharge as well as the evolution of optoelectronic properties of TiO2 upon lithium insertion, such as the band edge positions, charge carrier concentration and photoluminescence properties. Finally, this chapter will end with a study led by time-correlated single photon counting experiments to determine precisely the dynamics of the charge transfer processes in TiO2 and Li0.6TiO2 nanocrystals and incorporated in a photo-battery.

Preliminary work on the photorechargeable anatase TiO2 nanoparticles

The data and concepts presented below make a part of the research activities led by Dr.

Snehangshu Patra and Dr. Christian Andriamiadamanana (Post-doctoral fellows between 2012 and 2015 at LRCS) and Gaspard Bouteau (Ph.D. thesis defended in February 2019).

Synthesis and characterization of the anatase TiO2 nanoparticles

In order to bypass the surface limitation of the long-lived light-induced charge separation in semiconductor [1], two-step room-temperature synthesis procedure was proposed in our group to shape anatase-type TiO2 nanoparticles of ca. 4-5 nm in size [10]. In my work, we slightly adapted this procedure as following: 40 mL of Ti(iOPr)4 (titanium (IV) isopropoxide) was added into 400 mL of distilled water to achieve complete hydrolysis of the alkoxyde precursor.

The solution was vigorously stirred for 4 hours. The white precipitate was retrieved by multiple centrifugations with water, ethanol, and then dried in an oven at 60°C overnight. The resulted white powder is x-ray amorphous corresponding to hydrated titanium dioxide TiO2•1.6H2O as deduced by TGA analysis. The dehydration of powder was performed by adding concentrated sulfuric acid (H2SO4) until pH = 2. Thermolysis of the solution was carried out for one week at 60 °C under these conditions. The resulting white powder was retrieved by three subsequent centrifugations using distilled water. The high resolution transmission electron micrograph (Fig. 2.1 (a)) showed mostly spherical nanoparticles with an average diameter of around 5 nm.

The x-ray diffraction (XRD) pattern shows successful crystallization of the resulted power after controlled thermolysis leading to an anatase-type of TiO2 structure in good agreement with the selected area electron diffraction (SAED) (Fig. 2.1(b)). Note that there is a minor content of the brookite polymorph, visible with the small peak at 2θ = 30.8°. The presence of this polymorph as a structural impurity is explained by a possible increase of pH during the thermolysis [11].

Using Debye-Scherrer equation, we determined a crystallite size of 4.3 nm based on (004) orientation at 2θ = 37.8° [12]. N2 adsorption-desorption isotherm of these nanoparticles features a type IV isotherm corresponding to a mesoporous network (Fig. 2.2) [13]. Brunauer-Emmett-Teller (B.E.T.) surface area is 248.6 m 2 /g (± 3.2 m 2 /g, CBET = 327) [14]. The distribution of pores is monomodal centered at 3.5 nm as deduced by Barrett-Joyner-Halenda (BJH) method [15].

Formulation of the screen-printable paste and electrode printing procedure

In order to obtain mesoporous electrodes with reproducible thicknesses in the range of 2-15 µm, size and porosity, we privileged in this work the screen-printing methodology. For this, a printable paste containing the synthesized TiO2 nanoparticles was prepared. We adapted a protocol reported in the literature for the preparation of dye-sensitized solar cell photoanode [16][17]. For this, 6 g of dry TiO2 powder was dispersed within 175 mL of distilled water prior acidification by adding 1 mL of HNO3 (68%, Aldrich)). The solution was stirred at 78°C for 90 minutes with a 45-minute ramp to reach the temperature to allow particle peptization. This can be visualized by the solution color change from white to bluish. The colloidal solution was further dispersed using first Ultra-Turrax homogenizer (for 5 minutes at 25000 rpm) and then an ultrasonic horn with a titanium probe for one minute, alternating every 0.5 seconds ultrasound (at 50% power of 200 W at 20 kHz) and rest. After repeating this procedure three times between homogenizer and high power ultra-sound, the solution was concentrated in rotary evaporator under pressure of 10 mbar at 45°C until an aqueous suspension containing 20 weight percent (wt%) of TiO2 was obtained. The water and remaining HNO3 were removed by centrifugation three times using ethanol. Then, 20 wt% of ethylene cellulose 30-50 mPa.s (1.32 g), 34 wt% of ethylene cellulose 5-15 mPa.s (2.28 g) with respect to the initial amount of dry TiO2, and 24.3 g of terpineol were added to adjust the viscosity of the paste and final porosity. The mixture was diluted by ethanol to a volume of ca. 130 mL and homogenized by stirring for few hours at 800 rpm. The solution was redispersed following the same procedure as described above (Ultra-Turrax homogenizer and ultrasonic horn) and then concentrated again in the rotary evaporator to reach optimal viscosity for screen-printing. The obtained paste was then homogenized in a three-roll mill system, and heat-treated at 80°C for 4 hours in order to enhance its storage longevity. The TiO2 electrodes were screen-printed upon iron-free NSG10 Nippon Sheet Glass FTO. The glass surface was pretreated in 40 mmol/L solution of TiCl4 at 70°C for 30 minutes. This leads to a few nanometers thick deposit of TiO2 to enhance the film adhesion of the screen-printed TiO2 layer onto the FTO substrate. The screen utilized was composed of a network of 90 threads per cm (90T mesh) leading to 16 squares of 16 mm 2 . Three layers of TiO2 were deposited on the substrate leading to a final thickness of around 10 μm. The printing procedure is the following: printing, mechanical relaxation for 3 minutes and 7 minutes of heat treatment at 110°C on a hot plate to dry the electrodes. The removal of the remaining terpineol and ethyl cellulose, which plays the role of porogen, was carried out in four steps at 325, 375, 450 and 500°C under air according to ramps and step durations schematized in Figure 2.3. These steps follow precisely the thermal degradation stages of the ethyl cellulose. Prior to any use in the photoelectrochemical cell, the electrodes were always dried at 70°C under vacuum overnight.

We found reproducibly that working with dry electrodes is important for the photoelectrochemical performances. After the heat-treatment procedure, N2 adsorption-desorption measurements on the films still show a type IV isotherm (Fig. 2.4 (a)) with a H3 type hysteresis which typically met for porous materials formed from aggregates despite the different dispersion steps led during the paste preparation [13]. This is also visible based on the white color of the printed-electrodes, as discussed below, which shows Mie scattering from small aggregates in the film. The B.E.T. surface area decreases to 135.6 m²/g implying that the heating procedure induces also grain coarsening. The B.J.H. pore distribution in the film indicates a pore size centered at 4 nm in the film with a kind of bimodal distribution leading to an average pore diameter of 7.5 nm (Fig. For all photoelectrochemical experiments, a screen-printed TiO2 electrode was incorporated into an air-tight double-wall thermostated (at 20°C) photoelectrochemical cell made out of borosilicate (Fig. 2.5). This glass reaches ca. 90% transmittance in the UV and visible range with a cut-off transmittance below 300 nm. The airtight cap was mounted with three platinum wires to one of which the TiO2 photoanode (WE) was attached by soldering with indiumgallium alloy, and the other two were equipped with crocodile clips. Prior to the cell assembly, the cup with the working electrode and crocodile clips was dried at 70°C under vacuum for at least 6 hours. Meanwhile, the cell body was dried at 100°C in an oven. All the cell components were assembled in an Ar-filled glove box. Two lithium strips were serving as counter (CE) and reference (RE) electrodes. The cell was filled with an electrolyte containing 1 mol/L lithium hexafluorophosphate (LiPF6) in 1/1 ratio ethylene carbonate (EC) and dimethyl carbonate (DMC) solution from Solvionic. The cup was sealed to the cell by high-vacuum Glisseal grease.

The light excitation was ensured by using an Oriel LCS-100 sun simulator equipped with class ABB xenon lamp filtered at AM 1.5G (100 mW/cm 2 ). This light power condition was regularly checked with a silicon reference diode certified at NREL. The photoelectrochemical characterization was carried out using a galvanostat/potentiostat/impedancemeter Modulab ® XM connected to an optical shutter enabling automatized on/off light exposure.

Proof-of-concept of quantitative photorechargeable battery electrode based on nanocrystalline anatase TiO2 particles

The downsizing of TiO2 nanoparticles is known to modify the thermodynamic stability of the polymorphism chemistry in TiO2 [18][19]. At ambient conditions, rutile is thermodynamically the most stable polymorph followed by brookite and anatase. However, an interesting cross over point in this phase stability is observed when the particle is downsized to nanoscale. Banfield et al. concluded in his work that after crossing the threshold of 14 nm anatase becomes the most stable polymorph [20]. A similar observation was reported by Navrotsky et al. based on calorimetry experiments for particles having a B.E.T. surface area greater than 50 m 2 /g, corresponding to a size limit of around 30 nm assuming spheres [21]. This greater phase stability of the anatase is the result of its lower surface energy which becomes preponderant at the nanoscale, thus ruling the total free Gibbs energy of the system [21][22]. In accordance with literature, our particles of 5 nm demonstrated an increased thermal stability of the anatase polymorph which irreversibly transforms into rutile structure at 1050°C instead of 750°C as deduced from in situ x-ray diffraction and Raman spectroscopy [23]. In addition, we demonstrated that the greater structural stability of the anatase TiO2 is also experienced electrochemically with the occurrence of a complete solid solution domain between TiO2 and LiTiO2, thus bypassing the well-known tetragonal to orthorhombic to quadratic rock-salt phase transitions upon lithium insertion [24]. This led to enhanced electrochemical performances of these nanoparticles in terms of capacity and reversibility [23].

The preliminary work demonstrating the complete photorecharge proof-of-concept of the anatase TiO2 nanoparticles was carried out in our group by C. Andriamiadamanana (Post-(b) (a) doctoral fellow 2013-2015) and is herein presented for the sake of clarity to the context of my work [25][26].

The cyclic voltammetry reported in Figure 2.6 compares the electrochemical behavior of TiO2 nanoparticles under dark and illumination conditions in 1 mol/L LiPF6 EC/DMC electrolyte at 2 mV/s scan rate. As one can see, under illumination the energy separation between the anodic and cathodic peaks is almost twice lower (469 mV in dark vs. 246 mV under illumination) and is accompanied by the increase of the peak intensity as well as its narrowing. For the first time, it was evidenced that the electrochemical properties of TiO2 electrode can be modified by light.

In this case, light absorption affords to improve the kinetics of the redox processes as a result of the photo-generation of free charge carriers leading to an increase in conductivity [27]. In addition, a slight thermodynamic changes in the system were observed, represented by an increase of equilibrium redox potential from 1.834 V (±0.009) (vs. Li + /Li) in dark to 1.881 V (±0.004) (vs. Li + /Li) under illumination. In order to evaluate the photorecharge capacity of the electrode, it was first galvanostatically discharged to 1.5 V (vs. Li + /Li) at j = -100 µA/cm 2 (corresponding to ~ C/15, i.e. one lithium inserted in 15 hours) (Fig. 2.7 (a)). The plateau at 1.8 V vs. Li + /Li corresponds to the lithium ion insertion by two-phase reaction between the quadratic anatase LiαTiO2 (0 < α < 0.25) and the orthorhombic Li0.5+δTiO2 (0 < δ < 0.15) [9,23,28]. Upon discharge, the electrode's color changes from white to dark blue. It is induced by the lithium insertion into the structure forming Ti 3+ centers. The quantity of lithium inserted in the structure was estimated to be 0.6 using inductively coupled plasma -atomic emission spectroscopy (ICP-AES), thus corresponding to a composition of Li0.6TiO2. At this discharged state, the potential relaxation behavior of the electrode at open circuit conditions has been compared for one hour between dark and under light illumination. In dark, the electrode's potential reached a steady-state value of 1.92 V vs.

Li + /Li. This corresponds to the potential of the electrode close to equilibrium. The electrode remained dark blue in this condition. The subsequent discharge down to 1.5 V vs. Li + /Li under darkness revealed a very limited capacity of 46 µAh/cm 2 recuperated. In contrast, when the discharged electrode was left relaxing under illumination under the same other conditions, a faster rise of the electrode potential was observed reaching ca. 3.0 V vs. Li + /Li. The electrode color changed from dark blue to white and almost full electrode capacity of 608 µAh/cm 2 was retrieved after subsequent discharge (Fig. 2 Based on the above presented results, the first proposed mechanism to account for the photoinduced recharge of the anatase TiO2 nanoparticles is as following. The illumination of the LixTiO2-composition leads to the creation of photogenerated electron-hole pairs in the particle (eq. 2.

3). Considering a model including a depletion layer, which can be questionable given the size of the particles, the holes will drift towards particles' surface where they will onset the oxidation reaction with the transition metal Ti 3+ to Ti 4+ and subsequent release of lithium ions out from the crystal structure for charge compensation (eq. 2.4). On the other side, the electrons in the center of the particle are ideally collected towards the external circuit. However, in our case, the charge collection is not taking place because of the open-circuit condition and the lithium metal foil used at the counter-electrode which is too reductive with respect to the energy of the photogenerated electrons located in the conduction band of TiO2 (ca. -3.25 eV derived from Mott-Schottky measurements, section 2.3.2). In that configuration, they are scavenged by fluoride-based radicals formed in the electrolyte [29].

𝐿𝑖

When discharged at j = -100 µA/cm 2 under illumination, the device provides a stable output potential of around 3 V (vs. Li + /Li), with a remaining white electrode coloration, by virtue of the two antagonist processes: the lithium insertion into TiO2 triggered by the galvanostatic discharge current and the photo-induced lithium deinsertion from LixTiO2 (Fig. 2.8 (a)). The extracted capacity of the system is in theory infinite if the cell is completely regenerative without any side reactions with the electrolyte. Experimentally, we reached as high as 20 mAh/cm 2 until the electrode detached from the FTO. The electrolyte acquired an orange coloration indicating some side reactions and modification of solvation of Li + and PF6 -driven by continuous light exposure as recently demonstrated in our group [29]. One additional aspect of the two functions of the electrode, namely sun energy conversion and storage, has been demonstrated through the employment of the device as a photovoltaic cell during the day (simultaneous galvanostatic load under illumination) and as a simple battery in discharge during the night (Fig. 2.8 (b)) [26]. Thus, a constant current can be provided regardless of external light power fluctuation and day/night intermittency over 3 days.

Optical properties of TiO2 and Li0.6TiO2 by in situ UV-visible absorption spectroscopy

In order to evaluate the evolution of the nanoparticle anatase TiO2 light absorption in general and the band gap value in particular as a function of Li + content in the structure, in situ spectroelectrochemical experiments were carried out by UV-visible absorption spectroscopy on TiO2 at different states of discharge.

Absorption spectroscopy and methodology of band gap determination

Light absorption is an intrinsic property of semiconductors. It relates to the energy of interband electronic transitions when the semiconductor is illuminated with an incident photon exceeding the band gap value. As a result, the electrons from the valence band (VB) are promoted to the conduction band (CB) while leaving positive charges in valence band. The correlation between the band gap value and the absorbance onset is typically described by the linear relationship proposed by Tauc, Davis, and Mott [30][31] (eq. 2.5).

(ℎ𝑣𝛼)

1 𝑛 ⁄ = 𝐴(ℎ𝑣 -𝐸 𝑔 ) (2.5)
where ℎ is the Planck's constant (J•s), 𝑣 the frequency of vibration (s -1 ), 𝛼 the absorption coefficient, 𝐸 𝑔 the band gap (eV), 𝐴 is proportionality constant and 𝑛 a tabulated constant.

Capacity (mAh/cm

Depending on 𝑛 value, different types of transitions can be derived: 𝑛 equals to 1/2 for direct allowed transition, 3/2 for direct forbidden transition, 2 for indirect allowed transition and to 3

for indirect forbidden transition.

In order to determine the band gap value of a non-transmitting material, i.e. powder pressed into a pellet, we used an integration sphere and measured the diffuse reflectance spectrum of the pellet. In this case, the diffuse reflectance is inversely related to the absorption property.

The band gap value is then derived from its spectrum using the Kubelka-Munk function (F(R))

(eq. 2.6) [32][33][34].

𝐹(𝑅) = (1 -𝑅) 2 2𝑅 (2.6)
Substitution of 𝛼 by 𝐹(𝑅) gives the Tauc equation [35][36], from which Eg value can be determined from the intercept with the x axis in the (ℎ𝑣 ; (ℎ𝑣𝐹(𝑅))

1 𝑛 ⁄ ) coordinates (eq. 2.7).

(ℎ𝑣𝐹(𝑅))

1 𝑛 ⁄ = 𝐵(ℎ𝑣 -𝐸 𝑔 ) (2.7)

Evolution of TiO2 absorption upon lithium insertion

Prior to studying the evolution of TiO2 absorption as a function of the lithium content (depth of discharge), the band gap of the pristine 5 nm anatase TiO2 powder has been determined by UV-Visible absorption spectroscopy in diffuse reflectance mode. For this, 1 wt% of TiO2 was mixed into KBr before to be pressed to form a pellet. Figure 2.9 shows the reflectance spectrum recorded between 250 and 800 nm using a matt Teflon sample as a nominal 100% reflectance reference. The high reflectance above 400 nm is in a good agreement with the visual white color of the powder. The anatase TiO2 is reported to have an indirect band gap involving the lowest 𝛸 1,2 → 𝛤 1 transition [37]. Assuming an indirect allowed transition, the Tauc plot leads to a band gap value of 3.24 eV corresponding to 𝛤 3 → 𝛸 1𝑏 transition (Fig. 2.10). This value is in a good agreement with both experimental [38][39][40][41][42] and theoretical [37] values reported in literature. The result suggest the absence of any quantum confinement effect at a size of 5 nm particles. This is in good agreement with the literature reporting a threshold value of ca. 2 nm [39]. To provide more insight on the origin and absorption consequences of the electrode's color change from white to deep blue upon lithium insertion into TiO2, we developed in my work an air-tight three electrodes quartz cuvette allowing spectroelectrochemical experiments on films in transmittance configuration. However, for this study, we replaced the aggregated 5 nm particles by perfectly dispersed 20 nm particles for which screen-printed electrodes of 10 µm thickness are transparent to visible light in agreement with Mie scattering theory. The TiO2 working electrode was assembled with two lithium strips as counter and reference electrodes.

The electrolyte is composed of 1 mol/L LiPF6 in 1/1 EC/DMC solvent mixture (LP30, Solvionic). The lithium insertion was performed by applying j = -100 µA/cm 2 discharge current using two-channel biologic VSP potentiostat/galvanostat. The transmittance spectra were recorded in situ at different states of discharge (Fig. 2.11). The sequence of galvanostatic curves providing these different states of discharge is presented in Figure 2.11 (a). The cumulative discharge curve is composed of a solid-solution domain until ca. 1.75 V (vs. Li + /Li) before a well-defined plateau at 1.75 V vs. Li + /Li corresponding to a two-phase reaction between lithium poor and lithium richer TiO2 [9,28]. The evolution of transmittance spectrum at different TiO2 states of discharge is presented in Figure 2.11 (b). Before lithium insertion (OCV condition), (hv

F(R)) 0.5 hv (eV)
there is very little absorption between 400 to 1500 nm in agreement with the white color of TiO2. The LP30 electrolyte has an absorption threshold below ca. 250 nm, a value well inferior to the absorption wavelength of TiO2 / LixTiO2 compositions. The electrolyte exhibits also two absorption bands at 890 and 1022 nm, and two much stronger at 1162 nm with a shoulder at 1116 nm and a larger one containing at least 4 components at around 1400 nm. FTO in electrolyte absorbs all wavelengths below 360 nm, a few nanometers before the absorption of TiO2 nanoparticles. Upon lithium insertion, the main outcome comes from a significant increase of absorption in the visible and infra-red part of light. This suggests that the color change from white to dark blue upon lithium insertion into TiO2 stems from the increased absorption by free electrons in the conduction band. In addition, as the powder remains blue in darkness, this coloration also in part originates from the localization of electrons in the Ti 3+ centers occupying the 3d metal orbitals. It may be accompanied by an additional contribution from the antibonding O 2p -Ti 4s hybridized orbitals owing to the nanosize of the particles in agreement with Vayssieres et al. [43]. Our result is in agreement with literature, reporting similar evolution of absorption spectra and color change as a function of lithium insertion for Li4/3Ti5/3O4 [44] and WO3 [45].

The band gap of TiO2 and Li0.6TiO2 electrode inside the electrolyte was derived from the transmittance spectrum when the battery assembled at open circuit coltage (OCV) condition and discharged down to 1.5 V vs. Li + /Li (Fig. 2.13). The absorption coefficient 𝛼 was calculated using a derivation of Beer-Lambert law (eq. 2.8), where %𝑇 corresponds to the transmittance value and 𝑑 is the film thickness.

𝛼 = (ln ( 100 %𝑇

))/𝑑 (2.8)

The band gap value was derived from the absorbance onset in the Tauc plot in (ℎ𝑣 ; (ℎ𝑣𝛼)

1 𝑛 ⁄ )
coordinates considering an indirect transition in agreement with literature [38][39][40][41][42] (Fig. 2.13).

For TiO2, two distinct onsets are observed corresponding to lowest indirect 𝛸 1 → 𝛤 1 transition at 2.85 eV and 𝛤 3 → 𝛸 1𝑏 transition at 3.21 eV in agreement with literature [37,42]. Compared to the band gap measured on powder samples, i.e. 3.24 eV, this suggests a noticeable band gap narrowing when TiO2 is in contact with electrolyte. This can be attributed on the one hand to a

Stark effect induced by the specific adsorption of lithium cation on the surface of the particles as demonstrated in the Ph.D. work of Gaspard Bouteau [29], and on the other hand by the coulombic interactions created at the interface between TiO2 and the high dielectric constant and polarity of the EC/DMC solvent mixture.

This bandgap is maintained until 1.77 V vs. Li + /Li. At this potential range, the amount of lithium inserted into the anatase structure remains relatively low (< 0.2 Li + per formula unit).

The electrochemical phase transition taking place in the plateau at around 1.76 V vs. Li + /Li contributes to a drastic change in the Tauc plot with a band gap increases. Indeed, for 1.76 V, 1.73 V and 1.5 V vs. Li + /Li, LixTiO2 exhibits exclusively a transition of 3.21 eV (Fig. 2.12).

This higher band gap value upon lithium insertion is attributed to a Burstein-Moss effect. This arises because the electron density increases nearby the conduction band edge of TiO2 as a result from the continuous formation of Ti 3+ (4s 0 3d 1 ) upon lithium insertion [46][47].

1.5 V 1.73 V 1.76 V 1.77 V 2.0 V 2.2 V 3.0 V TiO 2 Li x TiO 2 (h) 0.5 Energy (eV)
To verify this hypothesis of donor density enhancement upon lithium insertion, the transmittance spectra were fitted by the Drude model. For this, we used an effective mass of electron of 0.4 as reported in literature by Deák et al. [48][49] (eq. 2.9-2.10) (Fig. 2.13).

𝑒1 = 1 - 𝜔 𝑝 2 𝜔 2 + 𝑖𝜔 (2.9) 𝑒2 = 𝜔 𝑝 2 𝛾 𝜔(𝜔 2 + 𝑖𝛾) (2.10)
where 𝑒1 and 𝑒2 correspond to real and imaginary part of the complex dielectric function (𝜀 = 𝑒1 + 𝑖𝑒2), 𝜔 is the resonant frequency (in order of 10 14 Hz), 𝜔 𝑝 is the plasmon frequency and 𝛾 is the damping constant. Comparison of the TiO2 data with LP30 and FTO + LP30 reference spectra suggests that the hills and valleys in the transmittance spectra of complete cell stem from the electrolyte absorption (Fig. 2.11 (b)). Therefore, corresponding parts of the spectra were excluded from the Drude model fitting for the sake of simplification. The simulation of the OCV, 2.6 V, and 2.4 V spectra to determine the optical carrier concentration was not successful. Nevertheless, excellent simulation was obtained for the 2.2 V spectrum. This should correspond closely to the pristine TiO2 structure, as no (or almost no) lithium is normally inserted above this potential.

The evolution of the charge carrier concentration as a function of the depth of discharge is reported in Figure 2.14. The results highlight that lithium insertion into TiO2 increases by 20

times the charge carrier concentration, from 1.3•10 19 cm -3 to 2.9•10 20 cm -3 .

Figure 2.13. Drude model fitting of the transmittance spectra of 20 nm nanocrystalline TiO2 in 1M

LiPF6 EC/DMC at different depth of discharge: 2.20 V, 2.00 V, 1.77 V and 1.50 V vs. Li + /Li.

Figure 2.14. Evolution of charge carrier concentration upon lithium insertion in 20 nm anatase TiO2

nanoparticles determined from the Drude model fitting of the transmittance spectra.

Band edge and charge carrier concentration evolution upon Li + insertion

The determination of the energy location of the valence and conduction band in TiO2 and

LixTiO2 can be carried out by Mott-Schottky experiments in combination with the steady-state UV-Visible absorption spectroscopy. This approach has been undertaken all along my Ph.D. thesis.

Mott-Schottky electrochemical impedance spectroscopy theory

Mott-Schottky experiments afford through an experimental approach to determine not only the type of main carriers and their concentration but also the flat-band position on the basis of the Schottky barrier establishment at the semiconductor/blocking electrolyte interface [50][51]. The measurement demands a blocking interface. It thus requires the absence of any possible interfacial redox reactions under applied bias voltage which may affect the electrode's built-in capacitance. Under such conditions, the equilibration of the Fermi level in the semiconductor and the chemical potential of the electrolyte leads to the occurrence of band bending (Fig.

(a)

). It consists in the depletion of the main charge carriers (electrons) underneath the electrode surface, thus establishing a Schottky barrier at the semiconductor/electrolyte interface. From the electrolyte side, the Helmholtz layer is formed due to the repulsion of the positively charged ions and attraction of the negatively charged ions in the electrolyte to the electron-depleted surface. This is the case for the n-type semiconductor, and the opposite takes place in the p-type semiconductor.

Mott-Schottky theory assumes the semiconductor to be ideal such as a single-crystal without any intermediate surface states and a negligibly small Helmholtz layer potential drop at the interface. In this case, the interfacial capacitance can be simplified to a pure double-layer capacitance.

TiO2 flat band potential and charge carrier concentration in different media: aqueous and non-aqueous aprotic electrolytes

Anatase TiO2 was first studied. For this, 3 µm thick films of 5 nm size particles of anatase TiO2

were prepared by doctor-blading. The electrolyte was composed of 1 mol/L NaOH in deionized water. The Mott-Schottky capacitance measurements were performed in a three-electrode cell using a saturated calomel electrode and a platinum grid as a reference and counter-electrode, respectively. The most reproducible conditions were obtained in the range of 100 Hz to 1 kHz.

No clear frequency dependence was noticed on the electrode's thickness when varying this latter between 3 to 15 µm. The results are reported in Figure 2.16 for a given electrode thickness and for different size of particles, namely 5 nm, 20 nm and 200 nm. To reach reproducible results and lower frequency dependence on the results especially at higher perturbation frequencies (i.e. > 200 Hz), we found crucial to encapsulate FTO to avoid direct contact with the electrolyte. In all experiments, FTO was systematically covered with a layer of silicone resin. Direct contact of FTO to the electrolyte leads also to an increase by almost two orders of magnitude of the carrier density as a result of the predomination of charge carriers from the highly doped SnO2 layer, which typically has ca. 10 21 cm -3 carrier concentration [52].

Our results show that there is no evident trend for the flat band potential (Vfb) and the carrier density (Nd) dependence on the particle size (Fig. 2.16). Taking into account such a dependency for each set of measurements, all throughout this study, one representative Mott-Schottky plot was selected based on the most reproducible Vfb value.

The measurement frequency of the chosen data is systematically reported.

Analogous electrochemical experiments were carried out in aprotic non-aqueous electrolyte in order to approach the real properties in photo-batteries. These experiments were performed in argon-filled glove box. The electrolyte is composed of 1 mol/L TBAPF6 in EC/DMC. The measurements were performed in a two-electrode cell in which a lithium foil was used as both counter and quasi-reference electrode. This lithium quasi-reference electrode's potential was determined from two cyclic voltamperometry measurements on fresh lithium electrodes in Fc + /Fc electrolyte (Fig. V fb (V vs SCE) Similarly to the measurements in aqueous electrolytes, frequency dependence phenomenon was also experienced in 1 mol/L TBAPF6 EC/DMC electrolyte, i.e. flat-band potential decreases and negligible donor density fluctuation while increasing the frequency (Fig. 2.19). It is important to mention that while carrying out Mott-Schottky measurements of electrodes in a non-aqueous aprotic electrolyte in glovebox, a weakly defined semi-circle was observed in impedance coordinates which may result from the greater length of cab les between the cell in glovebox and the impedancemeter. However, due to its low resistive contribution with respect to the total cell impedance and its small capacitance contribution, we assumed its effect into the measurements to be negligible. The values of flat band potential and donor density were derived using the Mott-Schottky equation (eq. 2.11) (Table 2.1): )

Frequency (kHz) (a) (b) 0.1 1 8.0x10
0.1 1 2.8 2.9 3.0 3.1 S1 S2 V fb (V vs Li QRE) Frequency (kHz) 0.1 1 2.0x10
Voltage (V vs Li QRE) 1 C 2 = ( 2 qεε 0 NA )(V -V 0 - kT q ) (2.11)
where C is the differential capacitance of the space-charge layer (C), ε0 is the permittivity (dielectric constant) of vacuum, εr is the relative dielectric constant of material, N is the charge carrier density (cm -3 ) (Nd, i.e. donor density for n-type semi-conductor and Na acceptor density for p-type semi-conductor), A is the sample's area (cm 2 ), V is the electrode potential (V), Vfb is the flat band potential (V), k is the Boltzmann's constant (1.38•10 -23 J/K), q is the elementary charge (1.6•10 -19 C), and T is the absolute temperature (K). reported much higher flat band potentials of TiO2 in DMF (-0.9 V vs. NHE) and THF (-1.0 V vs. NHE) comparable to those measured in water at pH = 14 [70]. To the best of our knowledge, there is no previous literature on the determination of the flat-band potential of TiO2 in TBAPF6 EC/DMC electrolyte whereas it represents from a fundamental point of view an interesting approach and data collection to better understand equilibrium of carriers between battery materials with its carbonate-based electrolyte.

Making direct comparison of our results with literature values in other non-aqueous aprotic electrolytes is complicated due to the strong dependence of flat band potential on the nature and concentration of the salt in the electrolyte [58]. Following the conclusion of Fitzmaurice et al.

who discussed that anions in the electrolyte have negligible influence on the Vfb value, we examined the results obtained in the electrolytes containing the same cation (TBA + ) and tried to rationalize a trend based on the electrolyte solvent nature. The -0.20 V vs. NHE determined in 1 mol/L TBAPF6 EC/DMC is much more positive than those reported for 0.2 mol/L TBAP (tetrabutylammonium perchlorate) in different types of solvents, i.e. DMF (-0.9 and -1.8 V vs.

NHE), THF (-1.0 and -2.1 V vs. NHE) and ACN (-2.0 V vs. NHE) [58]. Table 2.2 gathers all flat-band potential values reported in the literature for TiO2 in different aprotic electrolytes including their dielectric constants. This comparison highlights that there is no clear relationship between the flat-band potential and the dielectric constant of the solvent. This is in agreement with the conclusions made by Fox et al. who suggested that Vfb in aprotic solvent is constant for a given electrolyte but may depend on the solvent nature in absence of any added protic solute.

Besides the flat-band potential value, the charge carrier concentration can also vary to a large extent. Our results show more than one order of magnitude higher values in EC/DMC-based electrolyte, i.e. 2.3•10 19 cm -3 compared to 9.8•10 17 cm -3 in aqueous solution. This means that depending on the electrolyte nature, this latter can dope the nanocrystals during the Fermi level equalization process.

Evolution of the flat band potential and charge carrier concentration upon Li + insertion into TiO2

The evolution of TiO2 band edge positions and donor density upon lithium insertion was studied using a similar approach and methodology as aforementioned. For this experiment, a mesoporous film of 10 µm thickness composed of 20 nm nanoparticles of TiO2 was used to preserve the consistency with the above presented in situ UV-visible absorption spectroscopy measurements. We remind that our previous results show no size dependence of the particles on the flat band potential and carrier concentration. The Mott-Schottky capacitance was measured in a two-electrode transparent electrochemical cell using a lithium strip as both counter and quasi-reference electrode (stable at -3.402 V vs. Fc + /Fc) in 1 mol/L TBAPF6 EC/DMC. Lithium insertion was carried out in 1 mol/L LiPF6 EC/DMC by galvanostatic discharge at j = -100 µA/cm 2 until 1.5 V (vs. Li + /Li) to reach final Li0.6TiO2 composition. The lithiated electrode was then thoroughly rinsed inside the glovebox with DMC to remove all possible traces of lithium. The electrode was then assembled back into the cell including 1 mol/L TBAPF6 EC/DMC electrolyte for Mott-Schottky measurements.

As a result of lithium insertion into the structure, there is a noticeable shift of the flat band potential towards positive potentials (Fig. propylene carbonate [6]. Even though the reliability of the Mott-Schottky measurements in an ion-conducting electrolyte can be questioned, the authors claimed that the capacitance measured can be entirely attributed to the space charge capacitance based on electrochemical impedance spectroscopy. Even though the charge carrier concentration reported by Van der Krol et al. was calculated using a dielectric constant of 50 for both pristine and lithiated TiO2, the authors also report an increase in the dielectric constant of lithiated anatase TiO2 (Li0.5TiO2) to as high as 500 and 900 at a surface region in the film extending to at least 7 nm in depth [6]. A comparable increase in the dielectric constant up to 880 has also been reported upon the proton insertion into the anatase TiO2 [79]. At the same time, Boschloo et al. attributed the changes in the (C-V) characteristics, from which the dielectric constant is derived, rather to the increase in donor density upon H + doping [80]. We support this hypothesis since reaching such a very high dielectric constant is questionable in light of the rather low differences of crystallography between the two phases. In addition, if we consider a dielectric constant of 500 for Li0.6TiO2, this would result to a donor density of 2.4•10 18 cm -3 ; a value lower than TiO2 which comes in contradiction to results discussed in paragraph 2.2.2 based on Drude model. Consequently, on the basis of these results and knowing that crystal structure is not drastically modified by lithium insertion, we consider that dielectric constant should not vary at all or only to a very low extent.

Thus, in the following calculations a dielectric constant of 31 as for TiO2 was kept constant [3].

One can notice that the charge carrier concentration values derived from the Mott-Schottky experiment are at least one order of magnitude lower than those determined by the Drude model.

This discrepancy can be explained by the presence of surface traps and grain boundaries in the nanocrystalline TiO2 which can be optically excited and not electrochemically because of their shallow position in the band gap that prevents any excitation in a low frequency range of 100 Hz to 1 kHz.

The positive shift of Vfb and increase of Nd as above discussed is representative for all the Mott-Schottky experiments in a frequency range from 100 Hz to 1 kHz (Fig. 2.22). Li0.6TiO2 follows a similar frequency dependence as discussed above for TiO2, i.e. decrease of both flat band potential and carrier concentration values while increasing the small perturbation frequency.

However, it seems more pronounced in the determination of the donor density for Li0.6TiO2. This experimental observation gives credit to the DFT calculations reported by Wagemaker et al. who predicted such a movement upon lithium insertion [81]. As a consequence, this means that the photogenerated holes in Li0.6TiO2 have a stronger oxidizing character than in TiO2 counterpart. In this scheme, we also present HOMO and LUMO positions of the electrolyte deduced from both electrochemical cyclic voltamperometry experiments and the optical band gap measurements. The HOMO value is in agreement with literature data even though a broad range of values have been reported for the oxidation potential of PF6 -between 4.94 V (vs.

Li + /Li) [82] and 6.8 V (vs. Li + /Li) [83]. Upon illumination of Li0.6TiO2, the photogenerated holes oxidize the Ti 3+ transition metal centers leading to a release of Li + out from the structure (i.e. electrode photorecharge). However, in parallel, the oxidizing strength of these holes is sufficient to also oxidize PF6 -.

Fluorescence spectroscopy

The anatase TiO2 is known to be an excellent UV-absorber (320 nm < λ < 410 nm) emitting in the visible range (430 nm < λ < 730 nm) [84]. From a practical point of view, better understanding of the material's photoluminescence in the different media can provide further information regarding the excited-state dynamics of the photo-anode depending on its oxidation state and the media.

Photoluminescence spectroscopy basics

Photoluminescence is a physical phenomenon probing photon emission in response to light excitation. It is typically explained by the Jablonski diagram presented in Figure 2.24 [85].

Upon photon absorption, the electron is excited from the ground singlet state (S0) into the excited singlet states (S1, S2, ...). A non-radiative relaxation from higher energy level to lower lying energy levels such as for instance from S2 to S1 or inside S1 or S2 excited singlet states is called internal conversion. This is most often an ultra-fast relaxation process taking place in the femtosecond time-scale that can be monitored by transient absorption spectroscopy. The radiative decay of electrons back to the valence band is called direct band-to-band recombination. However, when dealing with nanocrystalline particles instead of large single crystals, longer photoluminescence processes can also take place through either surface shallow traps or deeper traps in the band gap (ns time-scale). For molecules typically containing heavy metals, non-radiative singlet-to-triplet transitions involving so-called intersystem crossing (ISC) mechanism can also take place. This leads at end to generate very long lifetimes that can reach up to phosphorescence (10 -6 -10 2 s). This latter can also take place when very deep trapping of carriers occurs at the defect or impurity sites in the crystal (T1) [86].

Steady-state fluorescence of anatase TiO2 nanocrystals

The steady-state excitation/emission map of the 5 nm nanocrystalline anatase TiO2 powder was recorded on FLS980 spectrometer (Edinburgh Instruments) equipped with double excitation/emission monochromators to reduce stray light rejection to less than 1:10 -10 . Figure 2.25 (a) shows excitation/emission map of the nanocrystals in the range of exc. 300-700 nm / em. 312-712 nm. The maximum emission is reached at 380 nm excitation wavelength leading to two distinct emission bands at 416 nm and 438 nm (Fig. 2.

(b)). This corresponds to a

Stokes shifts of 36 nm (2277 cm -1 ) and 58 nm (3485 cm -1 ) corresponding to the lattice relaxation energy based on the Franck-Condon theory [87]. Our values are close to the 60 nm reported by Niederberger et al. [88], but significantly smaller than the range from 141 to 174 nm as reported in other studies [89][90][91][92]. The emission spectrum of TiO2 can be deconvoluted using a Gaussian function into four bands located at about 413, 435, 450, and 472 nm (Fig. 2.25 (с)). This implies that the radiative recombination takes place through different relaxation pathways. The first band at 413 nm corresponds to the fluorescence of the self-trapped excitons localized on TiO6 octahedra in nanocrystalline anatase TiO2 [93][94]. The broad luminescence over the range of 420-600 nm, comprising the bands 2-4, is attributed to recombination pathways through different distributed surface states [95] and shallow traps in the band gap often associated with oxygen vacancies located 0.51-0.82 eV below the conduction band [42,94,96]. The absolute fluorescence quantum yield (QY) of the nanocrystals was determined using both direct and indirect measurement methodologies using an integration sphere (Fig. 2.26 (a)) [97].

As a result, the integration of the fluorescence spectra leads to a value of quantum yield from scratch equal to -0.4 %. The negative sign of the value stems from the very low emission signal close to the noise and the high character of light scattering from the samples. This result highlights the low quantum yield that falls below the detection limit, i.e. <1%. In order to bypass this intrinsic issue, we have attempted to compare the absolute intensities of emission from TiO2 to that of the scattering barium sulfate reference recorded in a steady-state fluorescence mode at 380 nm excitation wavelength (Fig. at room temperature is reported from 0.25% [98] to 1% [88] by comparison of the wavelengthintegrated emission intensity to that of a standard reference molecule. Looking more closely at the emission spectrum at 277 nm excitation wavelength, one can see two emission bands: a main band at around 335 nm and a second broader and much weaker in intensities at around 500 nm (Fig. 2.28). This latter is hardly visible on the maps due to its low intensity. The maximum of this broad contribution tends to shift towards red from 498 to 515 nm when discharging the electrode in contrast to the first band for which only the intensity is dependent on the electrode's potential (i.e. no effects on the band energy). The first band is attributed to the fluorescence of the solvated Li + and PF6 -in EC/DMC electrolyte (Fig. 2.29 (a)) [99]. The origin of the second band is attributed to deeper trap-mediated fluorescence of TiO2 nanocrystals in equilibrium with the electrolyte that contains potential-determining cations (i.e. Li + ). In addition, another contribution seems to appear at around 700 nm upon the discharge. It is however ascribed to the measurement noise, the amplitude of which above 630 nm becomes comparable or even higher than the intensity of TiO2 fluorescence.

Interestingly, both TiO2 and electrolyte fluorescence yields decrease upon lithium insertion in TiO2 nanocrystals. The quenching of electrolyte luminescence, up to five times in intensity, depending on the bias voltage applied to TiO2 is surprising at first sight. This result suggests that a part of the radiative energy from the electrolyte is transferred to the electrode, and, that this energy transfer process in benefit to TiO2 is more efficient when the electrode is lithiated.

At same time, the fluorescence QY of the electrode is decreasing when lithiated despite the energy transfer from the electrolyte. This means that the radiative recombination process in LixTiO2 is quenched compared to TiO2 as one could expect as a result from the hole transfer from Ti 3+ to Ti 4+ which is responsible for the electrode's photorecharge. Intensity (counts)

Wavelength (nm) OCV 2.8 V 2.6 V 2.4 V 2.2 V 2.0 V 1.77 V 1.73 V 1.6 V 1.5 V 334 nm 498 nm 515 nm
The FRET mechanism is clearly feasible for three reasons. The QY of luminescence of the electrolyte is much greater than TiO2 (QYLP30 = 7.7%), the high molecular proximity between solvated Li + and PF6 -responsible for the electrolyte fluorescence and TiO2 nanocrystals, and thirdly, there is an energy overlap between the emission band of the electrolyte and the excitation band of TiO2 and Li0.6TiO2 nanocrystals. The observed quenching of fluorescence of both TiO2 and electrolyte is consistent with the fluorescence measurements obtained through a colloidal solution, i.e. without any electrical bias on electrode, containing either 5 nm TiO2 or Li0.6TiO2 in LP30 electrolyte (Fig. 2.29 [29]).

In this case, 40 mg of powder was introduced in 3 mL of electrolyte. TiO2 nanocrystals quench the luminescence of the electrolyte by energy transfer of likely Förster type (FRET) from 3.5•10 6 counts to 1.0•10 5 counts (Fig. 2.29 (a-b)). This process is clearly more efficient in the case of lithiated nanocrystals as observed on biased electrodes. In the case of colloidal solution, we found a nearly total bleaching of the electrolyte (2.5•10 4 counts) which is prone for an enhanced FRET mechanism in contact with the lithiated particles (Fig. 2.29 (c)). This is consistent with the greater light absorption of the particles when lithiated in agreement with the UV-visible absorption spectra as a function of the depth of discharge reported in Fig. 2.11. This quenching process is thus not originating from any ions drift in the electrolyte or different solvation mechanism induced by the bias voltage. The yield of luminescence of Li0.6TiO2 is lower as a result from the photorecharge process in agreement with the films and may also in part be due to traps passivation by lithium insertion [100].

Determination of the kinetics of the charge transfer processes in TiO2 by TCSPC

The The numerical reconvolution of each decay was performed by considering systematically the instrument response function (IRF) determined on 100% refractive BaSO4 powder. The IRF full width at half maximum (FWHM) was ca. 80 ps. The lifetime values were determined using a multi-exponential mathematical function (eq. 2.12):

𝑅(𝑡) = ∑ 𝐵 𝑖 4 𝑖=1 𝑒 - 𝑡 𝜏 𝑖 (2.12)
This algorithm used in FAST software allows to perform an exponential fit of the decay using up to four fixed lifetime components (𝜏 𝑖 ) and corresponding pre-exponential factors (𝐵 𝑖 ), while automatically excluding the constant background level of the curve. The mathematical reconvolution was carried out resulting in the reduced chi-square (χ 2 ) of the fits of less than 1.1.

In our case, the most robust reconvolution of the decays was obtained by introducing two singleexponential components. The results show that for both TiO2 and Li0.6TiO2 a very similar dynamics of excited-states in obtained very close to the IRF value (Table 2.3). The faster contribution is clearly the major, reaching 29 ps after reconvolution. It corresponds to the direct recombination from conduction band down to valence band. This value is at the real limit of the spectrometer capability, but is in good agreement with Bowman et al. who reported an average lifetime of 23 ps for 2 nm TiO2 nanoclusters based on femtosecond transient absorption spectroscopy [104]. The longer component of ca. 2 ns is physically assigned to the trapmediated fluorescence process [105]. It shows higher amplitude in the case of lithiated phase, i.e. 18.1% for Li0.6TiO2 vs. 8.5% for TiO2. This difference originates from the redistribution of the traps in the band gap subsequently to lithium insertion. However, this result tends to contradict the traps passivation induced by lithium insertion discussed by Tsui et al. [100], but is in agreement with the work of Islam et al. who reported the formation of additional energy levels and trapping sites below the bottom of the conduction band induced by the defect clusters of Li + interstitials and Ti 3+ electronic species based on atomistic simulation study [106]. In addition to this slight but existing and well-reproducible difference in the photoluminescence decays, we also experienced a tenfold decrease in fluorescence yield for Li0.6TiO2 in agreement with the steady-state experiments (section 2.4.3). When the photoluminescence of Li0.6TiO2 nanocrystals is measured in contact with EC/DMC and EMITFSI-based electrolytes, we observed significantly longer decays compared to Argon.

This stems from the contact of the nanocrystals with a liquid interface of higher dielectric constant which tends to trap the electrons on the predominant surface of the particles. Indeed, the interfacial contact with the electrolyte hinders the band-to-band recombination to the benefit of slower processes through traps. This may also be the result of the Li + and/or TBA + cations forming "quasi-particles" with trapped electrons in the surface of the particles, thus prolonging the excited-state lifetime [107][108][109]. In both media, Li suggesting the occurrence of an additional fast charge transfer process which we attribute to the hole transfer to Ti 3+ that is responsible for the electrode's photorecharge (Table 2.5). This hole transfer is fast, requiring 145 ps in EC/DMC-based electrolyte, accounting for 7.8% of the total decay, and within 97 ps in EMITFSI ionic liquid representing 12.5% of the total decay. The difference in the amplitude tends to suggest that the photogenerated hole transfer process is more favorable in more inert EMITFSI ionic liquid media than in the EC/DMC-based electrolyte. The type of solvent influences the trap-mediated fluorescence mostly in terms of amplitude in the decay and less in terms of the lifetimes. This is likely due to a redistribution of energy and density of states of the intra-band gap traps. The longer lifetime is obtained in EMITFSI ionic

Conclusions

This chapter firstly described the preliminary work leading to the first proof-of-concept of a complete photorechargeable electrode based on TiO2/Li0.6TiO2. To achieve this, we developed a low-temperature synthesis yielding to nanocrystals of anatase TiO2 of 5 nm size. Such a downsizing of the particles affords to maximize the surface reactivity with respect to the volume which is one key to ensure a quantitative photorecharge process. This part of the thesis memoire focused on the study of the processes behind the effective charge separation in Li0.6TiO2 nanoparticles triggering photo-induced lithium deinsertion out from the structure. For this, we specifically developed a photoelectrochemical cell which allowed us to study in situ the evolution of the optoelectronic properties of TiO2 upon lithium insertion. This increased absorbance is assigned to the optical absorption of free electron carriers in the conduction band of TiO2 and the electrons localization into the Ti 3+ centers occupying the 3d metal orbitals.

In order to determine the band edge positions, which is an important parameter to consider in photochemistry, and their evolution upon lithium insertion, Mott-Schottky measurements on both TiO2 and Li0.6TiO2 were performed in TBAPF6-based ionic blocking electrolyte. The results highlight a positive 0.5 V shift in potential of the conduction band. This is accompanied by one order of magnitude increase in charge carrier concentration. In addition, our results stress a strong influence of the flat band potential and carrier concentration depending on the protic or aprotic nature of the electrolyte.

To grasp insights into the dynamics of the excited states of TiO2 and Li0.6TiO2 in different types of electrolyte, we used steady-state fluorescence spectroscopy and time-correlated single photon counting spectroscopy. In this chapter, we identified that in addition to the direct bandto-band recombination which takes place in the range of 29 ps for both TiO2 and Li0.6TiO2, a significant part of the luminescence stems from the recombination through shallow traps associated with oxygen vacancies and surface states when included inside the electrolyte (τ > 1 ns). Upon lithium insertion into the structure, we highlighted a quenching of electrolyte luminescence attributed to FRET mechanism from the electrolyte excited states to TiO2 and even more efficient to Li0.6TiO2 nanoparticles owing to its greater absorption properties.

Lithium insertion into TiO2 also substantially quenches the luminescence of the particles.

Finally, on the basis of time-correlated single photon counting experiment comparing the dynamics of excited state of Li0.6TiO2 in contact with either a lithium ion-blocking electrolyte or lithium-based electrolyte, we were able to evaluate that the hole transfer to Ti 3+ is in a range of 100 ps, namely 145 ps in EC/DMC-based electrolyte and 97 ps in EMI.TFSI based ionic liquid and the existence of Stark-Lo Surdo effect as a result from lithium specific adsorption on the surface of the anatase TiO2 which modifies the band gap value, the vertical excitation E0-0, excitation/emission spectra and the dynamics of the excited states relaxation.

Introduction

Most of the information related to the optoelectronic properties of inorganic materials reported in literature is connected to either water-splitting, solar fuel generation or photocatalysis applications [1][2][3][4][5]. Only few materials known for their lithium insertion properties are presented, namely TiO2, WO3, CuO, V2O5, Bi2O3, and BiVO4. For photoelectrochemical applications, the position of both valence band and conduction band is critical as it controls the reductive and oxidizing strength of the photogenerated carriers and defines the possibility for electrons or holes to be transferred to the transition metal and to the counter-electrode [6][7][8][9]. In contrast, when employed for batteries application [10], the redox potential and crystal structure to evaluate ionic conduction pathways and number of available sites for lithium to be accommodated are the most important criteria. Consequently, for the majority of insertion materials developed for batteries, no or little information is reported on the band gap value, band edge positions, or even charge carrier nature and its concentration. Moreover, with the exception of TiO2, which is a reference material for a broad range of applications, no studies are reported aiming at measuring these characteristics in non-aqueous aprotic solvents such as those typically used as electrolytes in batteries. This latter is important as the position of band edges is known to vary significantly depending on doping [11][12][13], surface states [14], and the nature of electrolyte in contact [15].

This chapter focuses on the creation of a database gathering the optoelectronic properties of the most widely used battery electrode materials. This database enables to evaluate the potentiality of a material for photobattery application and it represents an interesting contribution to the field of photocatalysis. Some first results related to photobatteries are discussed at the end of the chapter.

Three groups of materials were chosen for this study (Table 3.1):

transition metal oxides (Li4Ti5O12, LiCoO2, MoO3, WO3, CuO, Bi2O3, LiMn2O4, LiMn1/3Ni1/3Co1/3O2),

transition metal sulfides (TiS2, MoS2, WS2),

lithiated polyanions (LiFePO4, LiMnPO4, LiVPO4F).

Study of transition metal oxides

Li4Ti5O12

The spinel lithium titanate Li4Ti5O12 attracted much attention as an anode material for Li-ion batteries owing to its theoretical gravimetric capacity of 175 mAh/g, bi-phasic plateau at 1.5 V (vs. Li + /Li), high power, and excellent cycling stability owing to almost zero volume change between Li4Ti5O12 and Li7Ti5O12 [16][17]. One major disadvantage of Li4Ti5O12 is its poor electronic conductivity (σ < 10 -13 S/cm) [18], which stems from the empty Ti 4+ 3d orbital, and low lithium-ion diffusion coefficient (~10 -12 cm 2 /s) [19]. Its crystal structure, crystallizing in cubic phase with Fd3 ̅ m space group, is often presented in a form of Li8a[Ti5/3Li1/3]16dO4, denoting that 75 % of lithium ions are located at tetrahedral (8a) sites, while the rest of the lithium ions and titanium occupy octahedral (16d) sites (Fig. 3.1). The oxygen ions are located at 32e sites.

The XRD pattern of Li4Ti5O12 powder was successfully refined and indexed into the spinel structure with the Fd3 ̅ m space group (Fig. 3.2 (a)). The lattice cell parameter determined is a = 8.3588(3) Å (Vcell = 584 Å 3 ), in good agreement with the value reported by Dolotko et al. [20]. The sample is free of any impurities based on x-ray diffraction. The conclusion of a direct vs. indirect nature of the bandgap is often experimentally not so straightforward. However, some correlations were found between the band gap transition nature and the shape of the plot of the absorption coefficient vs. wavelength [30][31][32][33]. -0.8 -0.7 -0.6 -0.5 6.0x10 The flat band potential and donor density value of Li4Ti5O12 is summarized in table 3.2. In 1 mol/L NaOH aqueous solution (pH = 14), the flat band potential is situated at -0.99 V vs. SCE, which translates into -0.76 V vs. NHE. This value is 0.37 V lower than the one reported by Ge et al. (-1.13 V vs. NHE at pH = 14) in 0.2 mol/L Na2SO4 aqueous solution. In a nonaqueous aprotic electrolyte, based on 1 mol/L TBAPF6 in EC/DMC, we obtained a flat band potential of 1.86 V vs. Li QRE, which represents -1.34 V vs. NHE. To the best of our knowledge, this is the first report on the flat band potential determination of Li4Ti5O12 in a nonaqueous aprotic solvent. Considering the n-type conductivity and assuming the close proximity of the fermi level to the conduction band, the flat band potential corresponds to the lower edge of the conduction band.

Taking into account the bandgap of 3.79 eV, the valence band edge was calculated as 3.03 V in aqueous electrolyte and 2.45 V in non-aqueous aprotic electrolyte denoting a considerable influence of the nature of the electrolyte on the band edge positions in Li4Ti5O12 (Fig. 3.5).

The charge carrier concentration was determined by considering a dielectric constant of 20 as the average of the values reported by Zhang et al. [35] and Liu et al. [36] based on the microwave dielectric properties measurements (Table 3.2). In 1 mol/L NaOH aqueous electrolyte, donor density was determined as 2.7•10 18 cm -3 . In contrast, a significantly higher value of 6.9•10 19 cm -3 is obtained when measured in 1 mol/L TBAPF6 EC/DMC. This indicates that the donor density of Li4Ti5O12 measured in non-aqueous aprotic electrolyte is more than one order of magnitude higher than in the aqueous electrolyte. Such a trend was also observed in chapter II for TiO2 (see section 2.3.2). The steady-state photoluminescence map of Li4Ti5O12 powder was recorded between 270-800 nm in excitation and 278-808 nm in emission using a front-face configuration (0.5 nm excitation / 0.5 nm emission slits) (Fig. 3.6 (a)). The maximum of emission is reached for an excitation at 333 nm leading to a broad asymmetric band at 728 nm (Fig. 3.6 (b)). Our results are in relatively good agreement with the fluorescence data reported by Wen et al. who obtained a strong emission at 716 nm when excited at 325 nm [37]. The very large Stokes shift of 395 nm (16294 cm -1 ) is assigned to the presence of deep traps in the bandgap through which most of the energy relaxation is taking place in a non-radiative way (i.e. phonon dispersion and heat dissipation, etc.). The absolute fluorescence quantum yield of Li4Ti5O12 powder was measured using an integrating sphere combining both direct and indirect measurements (Fig. 3.7). The quantum yield of -0.57% derived from the measurement suggests that the value lies in the error bar range due to the highly reflective sample nature. Thus, we conclude that the quantum yield must be below the detection limit of the set-up, i.e. < 1%. The decay was the best fitted using a stretch exponential function (eq. 2.10) including three components leading to a reduced chi-square (χ 2 ) value of 1.025. The results after reconvolution of the IRF decay are presented in Table 3.

LiCoO2

LiCoO2 is the most common positive electrode used in lithium-ion batteries owing to its excellent cycling stability and high intercalation / de-intercalation potential of around 4 V (vs.

Li + /Li) [38][39]. Its structure consists of edge sharing CoO6 octahedra forming layers separated by lithium cations ensuring a good bi-dimensional ionic conduction (Fig. 3.9). One of the main drawbacks of this material is the irreversible structural change taking place upon deintercalation above 0.5 Li + extracted inducing a phase transition to a hexagonal close-packed structure of CoO2 above 4.5 V, which limits the practical capacity to around 140 mAh/g [54]. The band gap of LiCoO2 was evaluated as previously described assuming an indirect allowed transition (Fig. 3.11 (a)). A very low optical band gap of 0.79 eV was determined in agreement with black colour of the particles (Fig. 3.11 (b)). In this case, there is no uncertainties on the origin of the transition since a direct allowed transition would lead to a value of 5.17 eV which should provide a white powder (Fig. 3.11 (c)). In literature, a direct band gap value of LiCoO2 has been reported between 2.3 to 2.5 eV based on optical absorption spectra measurements [41] and photocurrent spectra onset determination [42]. However, the sluggish increase of reflectance suggests the indirect nature of the optical transition. electrolyte, which agrees well with our results in non-aqueous aprotic electrolyte [42]. LiCoO2 has also been reported by Menetrier et al. to be p-type based on thermopower measurements [44]. It is worth to mention that we can exclude the possibility of lithium / proton exchange in aqueous electrolyte, as we repeated the Mott-Schottky measurements in aqueous electrolytes at different pH from 1 to 14 without any visible change in the slope. Note that such a verification has been made for all other sensitive materials which showed a conductivity-type changeover depending on the electrolyte nature (i.e. MoS2 and WS2 presented in the following). This means that the donor or acceptor character depends on the electrolyte nature during the Fermi level equalization for such weak degenerated semiconductor, i.e. when the concentration of electrons is relatively comparable to the concentration of holes. When measured in a non-aqueous aprotic 1 mol/L TBAPF6 EC/DMC solution, the flat band potential is at 4.44 V vs. Li QRE, corresponding to 1.24 V vs. NHE. In contrast to non-aqueous aprotic electrolyte, this value is attributed to the position of the higher edge energy of the valence band because of the conductivity-type changeover. Considering the band gap of 0.79 eV, the lower conduction band edge is positioned at 0.37 V vs. NHE, which is 1.09 V lower compared to 1 mol/L TBAPF6 EC/DMC (Fig. 3.13). The charge carrier concentration of LiCoO2 for each experiment was determined by considering a dielectric constant of 2. This value was proposed by Khatun et al. [45] and Nageswara Rao et al. [46] (Table 3.4). In 0.5 mol/L TBANO3 aqueous solution we determined a concentration of n-type charge carriers of 4.9•10 19 cm -3 compared to 2.6•10 20 cm -3 of holes in 1 mol/L TBAPF6 EC/DMC electrolyte. This latter value is almost one order of magnitude lower than the value of 10 21 cm -3 reported by Mizutani et al. based on termopower measurements [47]. As one can see, the concentration of electrons and holes is not considerably different which supports our explanation driving the n-type vs. p-type changeover. Unfortunately, the photophysical properties of LiCoO2 could not be measured because of the too low bandgap value. This would correspond at best to an excitation band at around 1570 nm, which is not appropriate with a Xe lamp source for excitation and far from the detection limit of a PMT-type detector (ca. 850 nm). Note that this will be the case for all the materials having a band gap value lower than 1.5 eV presented in this study.

MoO3

MoO3 was one of the first materials studied for lithium-ion battery owing to its layered This value is in good agreement with the value of 2.9 eV reported by Anwar et The charge carrier concentration of MoO3 was determined using a dielectric constant of 6 as proposed by Lajaunie et al. based on calculations [65]. In 0.5 mol/L TBANO3 aqueous solution, a charge carrier concentration of 4.1•10 18 cm -3 was determined. This is in a good agreement with the value of 1.5•10 18 cm -3 reported by Nadkarni et al. based on the alternating current measurements [66][67]. In 1 mol/L TBAPF6 EC/DMC electrolyte, we found a significantly higher donor concentration of 1.2•10 20 cm -3 , which is almost two orders of magnitude higher compared to aqueous electrolyte. Interestingly, we found that MoO3 gives no fluoresence in the range of 250 to 800 nm. This means that all energy relaxation from the excitated states proceeds according to non-radiative mechanisms or that the low luminescence is quenched by the particles agglomeration. Such a behavior is not intrinsic to MoO3 but has also been encountered in other materials presented in the following, i.e. WO3, Bi2O3, LiFePO4, LiMnPO4, and LiVPO4F.

WO3

Tungsten trioxide is a well-known wide band gap semiconductor, which has attracted attention for its electrochemical and electrochromic properties. This material found extensive applications in smart windows [68], gas sensors [69], electronic displays [70], and photocatalysis [71][72]. The WO6 octahedra are forming a corner-sharing structure allowing open diffusion channels for alkali cation insertion (Fig. 3.19). Its main advantage is the high reversible electrode capacity reported to maintain ca. 650 mAh/g over 100 cycles for lithiumion batteries [73]. WO3 band gap was determined to be 2.49 eV assuming an indirect allowed transition (Fig. 3. 21 (a-b)). This result lies close to the lower edge of the 2.51 -2.87 eV indirect band gap values reported in literature based on optical measurements [75][76][77][78][79]. In the case when we assume a direct transition, a band gap of 3.52 eV is obtained, which is not consistent with the yellow/green coloration of the particles (Fig. 3.21 (c)). 3.6) [80]. A value of 1.4•10 18 cm -3 was extracted in 0.5 mol/L TBANO3 aqueous electrolyte and 2.8•10 19 cm -3 in 1 mol/L TBAPF6 EC/DMC. This indicates that more than one order of magnitude higher WO3 donor density is obtained in non-aqueous aprotic electrolyte. Our value determined in aqueous electrolyte in one to two orders of magnitude lower than those reported in literature in Na2SO4 and H2SO4 aqueous electrolytes, i.e. from 10 19 to 10 20 cm -3 [80][81][82].

CuO

Cupric oxide is an interesting negative electrode material for Lithium-ion batteries because of its high gravimetric capacity of ca. 600 mAh/g obtained through a conversion reaction involving Cu +II / Cu 0 redox couple according to a plateau at 1.3 V vs. Li + /Li [84], [85]. CuO has a tenoritetype structure with a square planar coordination of copper by oxygen in a monoclinic crystal These values are in good agreement with the values reported by Chaudhary et al. [86]. The band gap value of CuO was determined from the diffuse reflectance spectrum considering both indirect allowed and direct allowed transitions (Fig. 3.26). Assuming that CuO has an indirect allowed transition, we determined a band gap value of 1.19 eV (Fig. 3.26 (b)). This result is in good agreement with the 1.2 eV calculated from the high temperature intrinsic conductivity measurements reported by Jeong et al. [87]. The value also lies in the bottom of the wide range of values reported in literature between 1.20 eV and 1.79 eV based on the absorption spectroscopy measurements assuming an indirect allowed transition [88][89][90][91][92].

Assuming a direct allowed transition, the band gap is equal to 1.66 eV (Fig. 3.26 (c)). This value is considerably lower than the values reported in literature for a direct allowed transition ranging from 2.4 to 3.2 eV [88,92,93]. However, the slow change in the reflectance spectrum suggests the indirect nature of the transition and the 1.19 eV band gap comes in a good agreement with the black color of the powder.

Bi2O3

Bismuth oxide is a wide band gap semiconductor that can be used as anode material for lithiumion batteries. It delivers a gravimetric capacity of ca. 357 mAh/g obtained through a conversion reaction involving Bi +III / Bi 0 redox processes [95]. Among all the existing polymorphic forms of Bi2O3, monoclinic αand tetragonal β-Bi2O3 are the most important. In both cases, the crystal structure is constituted by the equivalently distant layers, parallel to the (011) plane. The tetragonal β-Bi2O3 comes of particular interest for lithium-insertion because of the large tunnels The band gap of Bi2O3 was also determined from diffuse reflectance experiments (Fig. 3.31 (a)). An absorption edge is observed at around 650 nm. The reflectance maximum is in the NIR region attaining 35% under this experimental condition. We determined a band gap value of 0.98 eV assuming an indirect allowed transition (Fig. 3.31 (b)). This is significantly lower than the values of 2.08 eV reported by Chitrada et al. [98] and 2.15 eV reported by Schlesinger et al. [99] both based on the optical absorption measurements.

Assuming a direct allowed transition, a band gap of 2.12 eV is obtained (Fig. 3.31 (c)). This value is closer to the broad range of values reported in literature from 2.26 to 2.88 eV based on the optical absorption spectroscopy measurements [98][99][100][101][102][103][104]. The charge carrier concentration in Bi2O3 was calculated assuming a dielectric constant of 32 as reported in literature [107] (Table 3.8). A donor density of 9.9•10 17 cm -3 has been determined in aqueous electrolyte, in agreement with literature [98,108], and 3.3•10 19 cm -3 in non-aqueous aprotic electrolyte. Once again, we found a more than one order of magnitude increase of the carrier concentration in non-aqueous electrolyte. This is the first report determining the flat band potential, carrier type and concentration of Bi2O3 in such media.

LiMn2O4

Lithium manganese oxide LiMn2O4 is one of the most studied cathode material for lithium-ion batteries because of its high redox potential of ca. 4 V vs. Li + /Li, natural abundance of Mn, and easy synthetic procedure, despite its quite low gravimetric capacity of 120 mAh/g [109][110].

LiMn2O4 crystallizes into a cubic spinel structure, which consists of cubic close-packed arrangements of oxygen atoms at the 32e sites, lithium occupying half of the 8a sites, and Mn 3+ and Mn 4+ ions occupying 1/8 of the octahedral 16d sites (Fig. The diffuse reflectance spectrum of LiMn2O4 is presented in Figure 3.36. The band structure seems to be either of a not semiconducting nature or to have a smaller bang gap value than 0.7 eV ( > 1800 nm). This contrasts with literature, which reports a direct band gap value between 1.16 to 1.20 eV from experimental optical absorption [112][113]. DFT calculations predicts from the calculation of the DOS a band gap lower than 1.0 eV [114][115]. Similarly to LiMn2O4, we did not achieve to determine the optical band gap of LiMn1/3Co1/3Ni1/3O2 from the diffuse reflectance spectrum owing to the absence of a clear absorption transition below 1800 nm ( > 0.7 eV) (Fig. 3.39). To the best of our knowledge,

Transition metal sulfides

The layered transition metal sulfides (MS2) are among the first materials studied for their ability to reversibly insert / deinsert lithium cations [119][120]. Their structure is typically constituted of layers, each of which is composed of transition-metal ions in edge-shared interstices between the pairs of close-packed anion planes (Fig. 3.40). These hexagonally close packed MS2 layers are held together by Van der Waals forces enabling lithium intercalation between the layers.

Among the family of transition metal sulfides, TiS2 is particularly attractive owing to its low molecular weight and good electrical conductivity stemming from high charge carrier concentration of ~10 20 cm -3 [49,[121][122]. [123]. The scanning electron micrograph reveals a nano-platelet morphology ranging from 5 to 15 µm in size (Fig. 3.41 (b)). The HR-TEM micrograph confirms a good crystallinity of the particles until the edges (Fig. 3.41 (c)). It also reveals the presence of pores in the particles as showed in the micrograph by semi-transparent "bubbles" in the core of the particle. The selected area electron diffraction pattern was indexed into TiS2 (Fig. 3.41 (e)). The electron diffraction pattern recorded over a particle consists of differently oriented sheets of TiS2 according to the same zone axis as showed by the absence of some reflections in the SAED pattern. absorption edge below 1800 nm suggests that the band structure is not of a semiconducting nature. Indeed, even though TiS2 is theoretically a semiconductor, it has always been experimentally found to have a semi-metallic nature. This is supported by most of DFT calculations on the electronic structure of TiS2 which suggest TiS2 to be an indirect band overlap semimetal [124][125][126][127][128][129][130]. It was also supported by the experimental studies reporting that TiS2 is a metal or semimetal with an indirect overlap of the bottom of the conduction band and the top of the valence band ranging from 0.2 to 1.5 eV [131][132][133][134][135]. This semi-metallic behavior is assigned to the access of the titanium metal electrons filling the d-type conduction band of the material [136][137]. Table 3.9 gathers the flat band potential and donor density values derived from the Mott-Schottky equation (eq. 2.7). A flat band potential of -0.24 V vs. SCE was determined in 0.5 mol/L TBANO3 aqueous solution (pH = 3.45). The value translates into -0.62 V vs. NHE at pH = 14 [59][60]. In the non-aqueous aprotic electrolyte, we obtained a flat band potential of 2.01 V vs. Li QRE, which corresponds to -1.18 V vs. NHE. No data is reported in literature regarding the flat band potential of TiS2. To calculate the charge carrier concentration, a dielectric constant of 15 was assumed by us, taking as a first hypothesis of a similar ratio between the dielectric constant of TiS2 and TiO2 with respect to WS2 (ε=8) and WO3 (ε=20), i.e. roughly 2:1. A donor density of 4.2•10 18 cm -3 was obtained in 0.5 mol/L TBANO3 aqueous electrolyte and 6.9•10 19 cm -3 in 1 mol/L TBAPF6 EC/DMC (Table 3.9). As previously observed, the charge carrier concentration in non-aqueous aprotic electrolyte is one order of magnitude higher compared to the aqueous counterpart. At the same time, these values are one to two orders of magnitude lower compared to the values from 1.1•10 20 to 2.8•10 20 cm -3 reported in literature based on calculations from Hall coefficient measurements [122,123,134,138,139]. This tends to suggest that the dielectric constant of TiS2 might be far from assumed value of 15. doping [148][149][150][151][152] and the experimental conditions [153][154]. However, this is the first time that a crossover in the main charge carrier in MoS2 crystals depending on the electrolyte nature is reported. These results suggest that MoS2 is a weak degenerated semiconductor for which the Fermi level is close to the middle of the band gap, resulting in a comparable concentration of electron and hole carriers as previously observed in LiCoO2. The charge carrier concentration of MoS2 in both media was calculated with a dielectric constant of 11 based on the calculations reported in literature (Table 3.10) [143,155,156]. In 0.5 mol/L TBANO3 aqueous electrolyte, we determined an electron concentration of 3.0•10 18 cm -3 . This value is one order of magnitude lower than that reported by Anand et al.,

MoS2

i.e. 3.5•10 19 cm -3 [141]. In 1 mol/L TBAPF6 EC/DMC, the hole carrier concentration is 4.0•10 19 cm -3 . This number lies in the middle of the 10 17 -10 19 cm -3 range reported for p-type MoS2 based on thermopower measurements [147]. ). This value is lower than those predicted by DFT calculations [146,159,160] and experimentally determined values between 1.12 eV and 1.30 eV based on different techniques such as absorption spectroscopy [161],

photoelectrochemical study [162], and measurements of the gate-voltage dependence of the source-drain current in an ambipolar ionic liquid-gated transistor [163]. At the same time it is 0.97 eV higher than 1.79 eV indirect band gap reported by Kam et al. based on the photocurrent spectroscopy measurements [164]. Assuming a direct transition would lead to a value of Mott-Schottky experiments were carried out on a 4 µm thick doctor-bladed film (Figure 3.50).

Interestingly, as it was the case for MoS2 and LiCoO2, WS2 shows a changeover in the conductivity type depending on the electrolyte nature. A n-type conductivity is observed in

aqueous 1 mol/L NaOH electrolyte and a p-type conduction appears in non-aqueous aprotic 1mol/L TBAPF6 in EC/DMC. This changeover suggests a weak extrinsic semi-conduction in WS2 in which the process of Fermi level equalization control either a n-type conduction (hole transfer) or a p-type conduction (electron transfer) depending on the electrolyte nature. Single crystals WS2 is reported in the literature to be n-type [165]. However, both n-and p-type conductivity are reported for polycrystalline WS2 samples [161,162,[166][167][168][169]. This changeover is mainly controlled by the synthetic method leading to different levels of punctual defects driven mainly by the sulfur stoichiometry. The literature also reports that the nature of semi-conduction in WS2 can be easily changed by a careful control in aliovalent doping [170][171][172] or by changing the biais voltage during thermopower measurements as demonstrated by Kawai et al. on nanotubes [173]. The flat band potential and donor density values are reported in Table 3.11. In 1 mol/L NaOH aqueous electrolyte (pH = 14), the flat band potential of WS2 is equal to -0.72 V vs. SCE, which corresponds to -0.48 V vs. NHE. This value is comparable to -0.45 V vs. NHE at pH = 14 reported in literature for the n-type WS2 by Morrish et al. [166] and is ca. 0.35 V more positive than that reported by Bassaid et al. [167] based on Mott-Schottky measurements in Na2SO4 aqueous electrolytes. This value of flat band potential corresponds to the lower edge of the conduction band. When measured in a non-aqueous aprotic electrolyte, the flat band potential is at 3.34 vs. Li QRE, corresponding to a higher edge of the valence band at 0.15 V vs. NHE.

In literature, all measurements on p-type WS2 were carried out in aqueous electrolyte, giving p-type semi-conduction and a large range of values for the valence band position from 0.00 to 0.87 V vs. NHE at pH = 14 based either on Mott-Schottky experiments [161,167,168] or photoelectrochemical measurements [170,174]. Taking into account the bandgap of 0.82 eV, the lower conduction band edge is calculated as -0.67 V vs. NHE. This denotes that there is solvation effect on the flat band potential in WS2 of 0.19 eV higher in 1 mol/L TBAPF6 EC/DMC electrolyte (Fig. 3.51). Charge carrier concentration of WS2 for each experiment was determined assuming a dielectric constant of 8 [167] (Table 3.11). An electron density of 6.7•10 18 cm -3 is obtained in 1 mol/L NaOH aqueous electrolyte in agreement with the broad range of values reported in literature [166][167]. A strong p-type doping is experienced in 1 mol/L TBAPF6 EC/DMC with a hole density of 1.8•10 20 cm -3 . This value is very comparable to the 1.2•10 20 cm -3 hole concentration determined by Wang et al. based on termopower measurements [175].

Lithiated polyanions

LiFePO4 and LiMnPO4 are isostructural and belong to the family of olivine (LiMPO4, M = Fe 2+ , Mn 2+ , Co 2+ , Ni 2+ ). These two materials are very attractive as high rate cathodes for lithium-ion batteries operating at 3.42 V and 4.13 V (vs. Li + /Li). They have lower environmental impact and lower cost by substituting cobalt by iron or manganese [176][177].

The olivine structure consists of a corner-sharing MO6 octahedra and edge shared LiO6 octahedra localized along the b-axis linked together by PO4 tetrahedra forming a threedimentional architecture (Fig. 3.52).

LiFePO4

The full-pattern matching refinement of LiFePO4 is presented in Figure 3 gradient approximation (GGA) report a band gap between 0.1 and 1 eV [181][182][183] which is not really consistent with the colour of the material (sand to green depending on the particles size).

In contrast, more recently, Zhou et al. reported a large band gap value of 3.7 eV calculated using GGA method adding Hubbard U term (GGA+U) [184]. These authors support their calculations by experimental data based on the diffuse reflectance measurements leading to a band gap value curve indicates n-type conductivity in agreement with Xu et al. [186]. However, these authors underline that a p-type conductivity may arise in the case of extrinsic doping owing to the higher mobility of the holes. N-type conductivity has also been observed by other studies reporting a negative sign in Seebeck coefficient [187][188]. All these results tends to disapprove the assumptions made on p-type conductivity mechanism in LiFePO4 [189][190]. The charge carrier concentration of LiFePO4 for each experiment was calculated with a dielectric constant of 3 based on the complex permittivity calculations from the impedance data by Bharathi et al. [188] (Table 3.12). In 0.5 mol/L TBANO3 aqueous electrolyte, we obtained 1.9•10 18 cm -3 donor charge carrier concentration. This is two orders of magnitude lower than the donor density derived from the measurements in 1 mol/L TBAPF6 EC/DMC, i.e.

3.9•10 20 cm -3 . To the best of our knowledge, this is the first report of the charge carrier concentration of LiFePO4.

LiMnPO4

Similar to LiFePO4, the XRD pattern of LiMnPO4 was refined using same space group leading to the lattice cell parameters of a = 10.460(9) Å, b = 6.100(6) Å and c = 4.744(5) Å (Vcell = 303 Å) in agreement with Koleva et al. [192]. The cell volume is ca. 3% greater than LiFePO4 owing to the larger ionic radii of Mn 2+ (r = 97 pm) than Fe 2+ (r = 92 pm) in high spin configuration. The SEM study revealed the nanoflake morphology of LiMnPO4 particles, which is rather difficult to determine precisely as we can extrapolate an onset at ca. 2.22 eV followed by a more marked transition above 5 eV (Fig. 3.59 (b)). For a direct electronic transition, the extrapolation of the linear part of (hνF(R)) 2 curve results in a 5.19 eV intercept. However, zooming in shows that there is also a small transition onset at around 3.93 eV, which has a low absorption coefficient compared to that of the 5.19 eV transition. We argue that LiMnPO4 has a direct band gap nature based on the Tauc plots. Our value is in agreement with 3.8 -4.0 eV reported by Xu et al. based on GGA+U calculations [184]. It is however significantly higher than 1.7 to 1.9 eV calculated by other groups [181,182,193]. It is recognized in literature that LiMnPO4 has a stronger insulating character than LiFePO4, thus consistent with the greater band gap value [187]. Using the Mott-Schottky equation, the flat band potential and donor density of LiFePO4 were determined (Table 3.13). The flat band potential, corresponding to the position of the bottom of the conduction band, was measured at 2.07 V vs. Li QRE, which translates into -1.12 V vs.

NHE. Taking into consideration a 3.93 eV band gap, the top of the valence band should be at 2.81 V vs. NHE. This shows that compared to LiFePO4 both conduction band and valence band of LiMnPO4 are destabilized in energy by 0.54 eV and 0.24 eV, respectively. (Fig. 3.61). A charge carrier concentration of 2.2•10 20 cm -3 was determined using a dielectric constant similar to LiFePO4 as an approximation [188]. This value is two times lower than in LiFePO4

(3.9•10 20 cm -3 ). To the best of our knowledge, this is the first report on the flat-band potential and donor density in LiMnPO4.

LiVPO4F

The XRD diffractogram of LiVPO4F is presented in Among the transition metal oxides herein studied, the photogenerated hole carriers (h + ) of the n-type Li4Ti5O12, MoO3, WO3, and Bi2O3 have sufficient oxidizing strength to be transferred to the transition metal to onset the lithium deinsertion reaction from their lithiated composition (Table 3 .14). This is the same case for the two polyanions LiFePO4 and LiMnPO4. For the ptype LiCoO2, the electrons are also sufficiently reductive to conduct the lithium insertion reaction from its delithiated composition Li1-xCoO2, then making this material promising for a photorecharge application. In contrast, some materials can be discarded for thermodynamic reasons, such as the p-type CuO, MoS2, and WS2, for which the electrons are not reductive enough to reduce Cu 2+ , Mo 4+ and W 4+ .

First results on the photoelectrochemistry of chosen insertion materials

All photoelectrochemical studies developed in this part were carried out using the airtight threeelectrode photoelectrochemical cell presented in chapter 2. LiCoO2 leading to a current density of up to 106 µA/cm 2 . In reduction, only a very low cathodic current of around -6 µA/cm 2 is observed, denoting the irreversible character of the lithium insertion into carbon-free electrode. This leads to a half wave redox potential of 3.84 V (vs.

LiCoO2

Li + /Li) in a good agreement with literature [39]. No further capacity is observed in the subsequent cycle. Interestingly, there is an important modification under illumination. An anodic current is starting to rise from 3.3 V (vs. Li + /Li) leading to a current density almost six times higher than in dark. We estimate that this faradaic current is resulting from a combination of the lithium extraction from LiCoO2 and a photo-activation of the electrolyte decomposition in contact with LiCoO2 excited states. This role played by LiCoO2 as a photoanode towards electrolyte degradation is somewhat surprising given its p-type semi-conduction. However, as discussed in the first part of this chapter, LiCoO2 has a weak degenerated p-type semiconduction which may either cross to n-type depending on the applied potential or lithium presence in the electrolyte, or, have a leak of electrons across the space charge layer. On top of this, it is also known that electrolyte degradation through LiPF6 can be prominent under illumination [211]. As it will be discussed in the following, this reaction is here significantly catalyzed by the excited state of LiCoO2. Upon reverse scan, one can decipher a peak at around 3.8 V (vs. Li + /Li) which still maintains a positive current. This is a result of two antagonist processes, namely lithium reinsertion into the structure (reduction) and the continuation of the electrolyte degradation (oxidation) at this potential; this latter being the predominant reaction.

Nevertheless, a better reversibility observed under illumination can be a result of (i) the higher conductivity of the electrode, by photo-action (i.e. photoconductivity) and/or (ii) the possibility of photon-induced lithium insertion, which will lead to a higher reductive current. Potential (V vs Li + /Li) Illumination

LiFePO4

In order to evaluate the electrochemical behavior of LiFePO4 in dark and under illumination, we carried out a cyclic voltammetry study on a 50 µm thick doctor-bladed film at different scan rates (ν) from 0.5 to 20 mV/s (Fig. Comparison of the current density of anodic (Jpa) and cathodic (Jpc) peaks as a function of scan rate in dark and under illumination is presented in Figure 3.68 (a). Both in dark and under illumination, the non-linear dependence of the peak current density as a function of scan rate (ν) and scan rate square root (ν 0.5 ) indicates a kinetic behavior that is intermediate between capacitive and faradaic. This is ascribed to the low electro-activity of the carbon-free films, which undergo only surface reactions in agreement with the film capacity that is lower than 1 µAh/cm 2 (Fig. 3.68 (b)). We can notice a close to pure faradaic behavior only in reduction regardless in dark or under illumination conditions. There is no drastic differences under illumination in terms of film capacities (Fig. 3.68 (b)). We only found slightly higher peak current under illumination in a very reproducible manner. To clarify in some points if the photorecharge arises from the dye excitation or from the nanosize of LiFePO4, we collaborated with the group of Dr. Kulka (AGH University of Science and Technology, Cracow, Poland) who provided us nanoparticles of carbon-free LiFePO4. The synthetic procedure of this powder is reported in the following reference [212]. The full-pattern matching refinement of the x-ray diffractogram revealed minor impurities of Li3PO4 corresponding to ~4 wt% of the sample as reported in their publication (Fig. 3.72 (a)) [212].

All diffraction peaks of nano-LiFePO4 are corresponding to the orthorhombic lattice adopting The very low film activity is also reflected in the in situ UV-visible absorption spectroscopy experiments showing the evolution of the transmittance of the film as a function of the applied potential (Fig. 3.74). For this, the electrode was prepared by spin coating LiFePO4 nanoparticles onto the FTO glass resulting in a ca. 300 nm thin film. The lithium insertion was performed by applying j = 0.5 µA/cm 2 current density and the transmission was recorded at different states of charge. No noticeable change in the transmission spectrum is observed depending on the applied potential. The main absorption bands are related to the electrolyte as discussed in chapter 2. This also suggest the very poor film electro-activity even using nanocrystalline particles and thinner films made by spin coating. We sensitized these nanoparticles with N719 dye. There is an additional 20 mV when the film is sensitized, which is attributed to a photovoltage built-in from N719 excited states. This low compared to a dye-sensitized solar cells value indicates a very limited charge separation process induced by the low electronic coupling between the dye and LiFePO4. The triplet states of the N719 dye (LUMO ≈ -1.5 V vs. NHE) can inject electrons into the conduction band of LiFePO4 measured at -0.58 V vs. NHE, as determined in this thesis, but the dye cation formed after injection cannot be regenerated to its fundamental initial state (no redox mediator in electrolyte). It thus gives a clear clarification to the wrong charge transfer mechanism between LiFePO4 and N719 presented by Paolella et al. [191]. One limitation of LiFePO4 / FePO4 system can also stem from the optoelectronic properties of the delithiated phase. For this, the iron phosphate was obtained by chemical oxidation of the lithium iron phosphate powder by means of a strong oxidizing agent [213]. The band gap of FePO4 was determined by UV-visible spectroscopy in diffuse reflectance mode (Fig. 3.77). The powder is highly reflective in the NIR part of the light, twice more than LiFePO4. Assuming an indirect allowed transition, we found a band gap value of 2.26 eV which is close to the range of values between 1.9 and 2.1 eV reported in literature based on DFT calculations [184,216]. However, this value is not consistent with the light pink colour of the powder. Assuming a direct allowed transition as for LiFePO4, the onset of absorption is determined at 3.43 eV with a smaller transition at an energy of 2.73 eV that corresponds to the band gap value of FePO4. Lithium extraction leads then to a band gap narrowing of 0.58 eV. With same procedure as previously described, we determined the band edge positions of FePO4 by Mott-Schottky measurements in aqueous and non-aqueous aprotic electrolytes (Fig. 3.78).

In both cases, the positive slope of the Mott-Schottky plot indicates that the material has a ntype conductivity as for LiFePO4 in agreement with Zhu et al. [217]. Table 3.15 presents the

Li4Ti5O12

The kinetics of lithium insertion/deinsertion in the spinel Li4Ti5O12 was investigated on 14 µm thick electrode by cycling voltammetry at different scan rates from 0.05 to 0.5 mV/s (Fig. 3.80).

These experiments were performed under darkness and illumination. Higher scan rates lead to stronger kinetic limitations. Regardless of the scan rate, the degenerative influence of the light on the electrochemical cycling is observed. The small cathodic peak at around 2.3 V vs. Li + /Li during the first cycle upon illumination is attributed to the filling of the surface trap states. To better understand this phenomenon, the TEM investigations were carried out and showed the formation of a SEI layer in both cases. Its thickness reaches up to 40 nm in dark and more than 50 nm under illumination with few isolated SEI parts, which can attain up to 100 nm (Fig. 3.83 (a-b)). This implies that the illumination of the electrode upon cycling promotes the formation of SEI. This was also observed for the samples cycled by cyclic voltamperometry.

EELS analysis combined with electron diffraction study showed that the SEI layer formed under illumination is exclusively composed of LiF without any byproducts from carbonates decomposition as also deduced from Raman analysis of cycled electrodes. to study the dynamics of the excited state of Li7Ti5O12, which could have provided hints to explain if the absence of the photorecharge could be attributed to a too fast recombination dynamics. The quenching of the lithiated phase can stems either from the lithium insertion into Li4Ti5O12 structure, too fast dynamics for the spectrometer, or from the formation of the LiF layer on the surface of the electrode, which completely quenches the luminescence. We argue that the two last hypothesis are the most likely because: (i) the electrode recharge to 2.6 V vs. Li + /Li does not allow retrieving the original luminescence signal of Li4Ti5O12, even though there is a small irreversible capacity for the first recharge, which will still results in an electrode containing mostly Li4Ti5O12 and a small part of Li7Ti5O12; (ii) fluorescence of the chemically reduced Li7Ti5O12 powder, which should not have a LiF layer, also shows a quenching of the luminescence (as well as that of the electrolyte). Note that the emission band of the electrolyte is partially but not totally retrieved during the electrode oxidation. This tends to suggest that the FRET mechanism might be more efficient in contact with Li7Ti5O12. The UV-visible absorption spectroscopy study of Li4Ti5O12 in 1 mol/L LiPF6 EC/DMC was carried out aiming at understanding the optical changes in Li4Ti5O12 upon lithium insertion. For this, a semi-transparent 250 nm thick film of Li4Ti5O12 was prepared by spin-coating on FTO glass and then discharged at -10 µA/cm 2 . The transmittance spectra were recorded in situ at different states of discharge (Fig. 3.100 (a)). The changes in the transmission spectrum take place at 1.7 V vs. Li + /Li and stay almost unaffected for all the potentials below. There is a significant decrease in the transmittance observed in a range from 300 nm to 1500 nm. The FTO glass absorbs all the light below 300 nm and all absorption bands are corresponding to the electrolyte fingerprint as detailed in chapter 2. The optical modifications are accompanied by a color change of the electrode turning from white to a dark blue. This is in agreement with the increase of absorbance in the visible and infra-red region as reported by Kanamura et al. [START_REF] Kanamura | [END_REF]. we argue that the band movement is not related to the lithium insertion. The important interfacial built-in capacitance is mainly controlled by the formation of LiF below 1.5 V vs.

V 1.6 V 1.5 V 1.4 V 1.3 V 1.2 V 3.0 V 2.5 V 2.3 V Transmittance (%) Wavelength (nm) (a) (b)
Li + /Li. Note that a similar experiment has been carried out on pure FTO and the SEI formed on FTO. Even though the nature of the SEI is probably different on FTO, we found consistency with the results above, that the flat band potential of the FTO shifts up in potential by more than 1 V (Fig. 3.90 (a)). The evolution of the charge carrier concentration was calculated using a dielectric constant of 6

[65] (Table 3.17). The initial 2.9•10 19 cm -3 charge carrier concentration increases upon lithium insertion to only 3.4•10 19 cm -3 whereas the literature reports a 6 orders of magnitude increase in conductivity between Li4Ti5O12 and Li7Ti5O12 [221]. This is an additional evidence that the interfacial capacitance become governed by LiF and not the insertion material anymore (ε ≈ 10 In the second part of this chapter, we aimed to investigate the photo-electrochemical properties of some materials for the time being of this thesis. The first results clearly highlight the difficulties of this research field. Many materials, which were expected to have a photorecharge property did not express it, i.e. LiCoO2, Li4Ti5O12, LiFePO4, Bi2O3 and MoO3 (results for two latter are not reported in this manuscript). Besides the thermodynamic considerations, we think these difficulties are related to (i) the poor electronic conductivity of the electrode which leads to an excessive interfacial impedance value, (ii) the size of the particles which are often far from a threshold of few nanometers to lead to the Debye length comparable to particles radius, and (iii) faster excited state lifetimes compared to carrier transfer. Further research is therefore needed to point out which of these challenges is the most crucial to tackle. In addition, as showed in the case of LiCoO2 and Li4Ti5O12, the electrolyte's stability under electrodes excited states strongly limits the applicability of photorechargeable electrodes. We demonstrated for instance in the case of LiCoO2 at excited states that this latter can be strongly photocatalytic towards the electrolyte's oxidation above 3.3 V vs. Li + /Li or that LiF-SEI type formed at lower potential than 1.5 V vs. Li + /Li totally quenches the excited state of the electrode (Li4Ti5O12).

Even though the photorecharge was not observed in LiFePO4 likely due to a too insulating nature of the electrode, we investigated further the possible reasons of its absence and potential solutions. Two strategies were attempted to overcome the possible obstacles: (i) decreasing the particle size to increase the electronic conductivity and (ii) grafting N719 dye on the particle's surface to increase light harvesting. None of these approaches led to noticeable improvements. Even though we showed some kinetic improvements of the electrochemical processes of lithium insertion/deinsertion process in Li4Ti5O12/Li7Ti5O12 electrode, and faster reaching of equilibrium state under illumination, this spinel material did not express any sign of photorecharge neither. The in-depth study showed that on the one hand the Li7Ti5O12 excited state may be ultrafast and faster compared to the carrier transfer, and on the other hand that SEI formation leading exclusively to insulating LiF on the electrode surface has an important influence in agreement with previous study on TiO2 [START_REF] Bouteau | Etude des propriétés de l'interface semi-conducteur / électrolyte sous illumination pour batteries à ion-lithium photo-rechargeables[END_REF]. We showed by time-resolved fluorescence study that there is a strong quenching of electrolyte luminescence when including Li4Ti5O12. This is explained by energy transfer via Förster resonance effect (FRET). This FRET mechanism is even more efficient for Li7Ti5O12 owing to its band gap narrowing by 0. 500 µAh/cm 2 ). A patent for the molecule included into the electrolyte is currently under filling process. We hypothesize that a function of this molecule is to avoid the formation of LiF upon discharge and under illumination, and it may also play a role in promoting the carriers transfer towards the transition metal.

study of Li4Ti5O12 showed that an important limitation of the system is related to the formation of a SEI layer exclusively made of LiF under dark but also under illumination with a thicker layer which can attain up to 100 nm. To present level of understanding, we estimate that LiF formation annihilates the photorecharge process of the electrode. On the other hand, we also speculate that excited state lifetime of Li7Ti5O12 is too fast with respect to a possible charge transfer towards the transition metal. The formation of LiF on the surface prevented us from determination of how the band structure is modified in Li7Ti5O12. However, in this part of the study, we revealed a FRET mechanism taking place in the UV from the electrolyte to the electrode. This FRET process, which induces the quenching of the electrolyte fluorescence, seems even more efficient in contact with Li7Ti5O12 owing to its narrower bandgap by 0.23 eV and improved absorption capability in the whole visible range.

This study reveals mainly three important challenges that need to be addressed in the field: (i) electrode nanostructuralization to favor surface reactions and to bring the Debye length comparable to the particle size and carbon-free electrodes with high conductivity, (ii)

developing electrolytes specific to photorechargeable systems with enhanced stability in both reduction and oxidation with respect to electrode's photo-excited states, and (iii) choosing the right photo-anode with the right photo-cathode that allows an adequate band energy alignments to collect carriers and thus yielding to a complete photorechargeable lithium-ion batteries.

Preliminary results on the second point shows that the addition of a specific scavenger in the electrolyte can likely prevent the formation of LiF over the electrode, thus offering not only improved electrochemical performances of carbon-free electrodes but also a photorechargeable process in Li7Ti5O12 under white light illumination in ca. 15 hours. The third point also remains very challenging, as the database herein developed shows that none of the materials can be associated to form a fully regenerative bifacial p-n photobattery.

This pioneering work in a complex subject of energy storage and energy conversion at the molecular level requires still a lot of profound investigations from fundamental standpoints, particularly in the field of the time-resolved spectroscopy. In our laboratory, further timeresolved spectrometers are under setting up, which will allow to provide more details on the different charge carrier transfer taking place in the photo-electrode (i.e. carrier transfer to the transition metal vs. to the electrolyte). This study will be able to be carried out by a picosecondpump / three probes transient absorption spectrometer including a high dynamic range streak camera detection that is delivering today its first results and two other time-resolved spectrometers in fluorescence associating a fs tunable Ti:Sa laser with a streak camera in synchroscan mode for TR-PL (ca. 4 ps time resolution) and a up-conversion setup (IRF ca.

200 fs) which will be able to monitor faster processes than ca. 30 ps. Establishing the time-scale of the different carrier transfer in the photo-electrode will allow understanding all limiting processes taking place in the electrode under illumination. It will be also a basis to improve the materials and the electrode mesostructuralization towards optimizing the electrode's performance.

As mentioned in the manuscript, this work has often faced a lack of data/knowledge from the literature to guide us more efficiently into this thesis development. However, the results of this

work not only open up important lines of research that will be pursued in the laboratory to continue development of these systems, but can also establish a background for the pursuit of this field in the community. In addition, the created database of the optoelectronic properties of common insertion materials can be beneficial for a broad range of other applications like photocatalysis, electrochromic devices, and even to battery itself.

Figure 1 . 1 .

 11 Figure 1.1. Evolution of levelized cost of energy (LCOE) from various sources over time adapted from [10]. Solar mean LCOE values primarily relate to North American alternative energy landscape but reflect broader/global cost declines.

Figure 1 . 2 .

 12 Figure 1.2. Evolution of real power output from a 4.6 MW solar photovoltaic array in Springerville AZ on (a) tenths of seconds over two hours and (b) minutes over a week timescale [12].

1. 1 .

 1 Solar energy harvesting, conversion, and storage ............................................................. 1.1.1. History of photovoltaics ............................................................................................ 1.1.2. Working principles of conventional solar cells and dye-sensitized solar cells ......... 1.1.3. Principles of energy storage in electrochemical batteries and capacitors.................. 1.2. Photorechargeable energy storage devices ........................................................................ 1.2.1. Devices based on multifunctional charge generation and storage electrode ............. 1.2.2. Separate planar sub-devices ....................................................................................... 1.2.2.1. Capacitor storage ............................................................................................... 1.2.2.2. Battery storage ................................................................................................... 1.2.3. Charge generation and storage at the molecular level ............................................... Conclusions .............................................................................................................................. References ................................................................................................................................

Figure 1 . 3 .

 13 Figure 1.3. Evolution of certified record power conversion efficiencies for existing technologies since 1976 (Data from National Renewable Energy Laboratory from 17.07.2018[8]).

Figure 1 . 4 .

 14 Figure 1.4. Schematic representation of the photoinduced charge separation in the p-n junction.

Figure 1 . 5 .

 15 Figure 1.5. Schematic representation of a dye-sensitized solar cell.

Figure 1 . 6 .

 16 Figure 1.6. Schematic representation of (a) Li-ion battery and (b) an electrochemical supercapacitor.

Figure 1 . 7 .

 17 Figure 1.7. Ragone plot for various energy storage devices [25].

Figure 1 . 8 .

 18 Figure 1.8. Schematic representation of the two configurations proposed for photorechargeable devices based on the use of (a) a multifunctional electrode or (b) spatially separated photoanode and storage electrode.

Figure 1 . 9 .

 19 Figure 1.9. Schematic representation of the "photovoltaically self-charging battery" proposed by Orel et al. [29].

Figure 1 . 10 .

 110 Figure 1.10. Schematic representation of the "organic photorechargeable air battery" proposed by Nishide et al. based on BODIPY-sensitized (3) anthraquinone polymer layers (2)[33].

  photocapacitor. However, the performance was limited by the formation of a Schottky barrier at TiO2/carbon interface which penalizes fast electron transfer from TiO2 to carbon upon discharge. To overcome this issue, the authors proposed a three-electrode configuration by introducing an internal Pt electrode. By minimizing the inner resistance, they reached 0.8 V open circuit photovoltage and a discharge capacity as high as 130 µAh/cm 2[37][38]. No cyclability data are reported.

Figure 1 . 11 .

 111 Figure 1.11. Schematic representation of the "carbon-incorporated photocapacitor" proposed by Miyasaka et al. in a two-and three-electrode configuration [38].

Figure 1 . 12 .

 112 Figure 1.12. Schematic representation of the "solar storable rechargeable air battery" proposed by Gao et al. [51].

Figure 1 . 13 .

 113 Figure 1.13. Schematic representation of (a) two-compartment photorechargeable device based on double-sided TiO2 nanotube arrays proposed by Wang et al. and its modification with the three tandem DSSCs integrated in the photovoltaic compartment [53].

  Gao et al. improved this concept by introducing PEDOT polymer into the DSSC compartment to ensure the dye regeneration by hole trapping and forming PEDOT + ClO4 -with the anions from the electrolyte (Fig. 1.14) [55]. This photoelectrochemical cell reached 0.76 V opencircuit voltage under standard light conditions. The subsequent discharge capacity was 8.3 mAh/g and a negligible capacity fade of less than 4 % was experienced over 10 cycles. The authors ascribed the limited performances of the device, i.e. solar-to-storage conversion (a) (b) efficiency of ∼0.1%, to the low power conversion efficiency (PCE = 0.5%) of the TiO2/PEDOT anode.

Figure 1 . 14 .

 114 Figure 1.14. Schematic representation of the "photorechargeable battery based on polymeric charge storage electrodes" proposed by Gao et al. [55].

Figure 1 . 15 .

 115 Figure 1.15. Schematic representation of a "solar rechargeable redox flow battery" proposed by Gao et al. [57].

Finally, more advanced

  photoanode and GaN/Si photocathode (Fig.1.16), providing the highest up-to-date device photovoltage of 1.5 V[63]. This high voltage allows to operate a 1.2 V alkaline anthraquinone/ferrocyanide redox battery with a 3.0% solar-to-storage conversion efficiency under standard light conditions (A.M. 1.5G). The discharge capacity reached 200 µAh/mL, even though this latter is limited by the catholyte degradation.

Figure 1 . 16 .

 116 Figure 1.16. Schematic representation of a "solar rechargeable redox battery" proposed by Wang et al. [63].

  The only functional energy harvesting and storage device based on the organic molecule was presented by Samuel et al. inspired by a homojunction OPV cell (Fig.1.17)[66]. The bifunctional cyanine dendrimer is sandwiched between a hole transporting polymer based on PEDOT/PSS and an aluminum plate to collect the free carriers. Upon light excitation, the electron is trapped in the central indoline unit triggering the movement of I -or PF6 -across the cyanine to the second indoline. This process is reversible. Even though the performances are modest (0.003 µAh/cm 2 ), likely because of the low light capture and short lifetime of the charge separation, this work stands for the first proof-of-concept of a photorechargeable device based on molecular energy conversion and storage in an organic molecule.

Figure 1 . 17 .

 117 Figure 1.17. Cell configuration and bifunctional organic cyanine dendrimer performing light conversion and charge storage at molecular level proposed by Samuel et al. [66].

Figure 1 . 18 .

 118 Figure 1.18. Schematic representation of the band bending at the semiconductor/electrolyte interface for an electrolyte containing a redox mediator.

Figure 2 . 1 .

 21 Figure 2.1. (a) High Resolution Transmission Electron Micrograph including SAED pattern and (b) xray diffraction pattern of the synthesized anatase TiO2 powder.

Figure 2 . 2 .

 22 Figure 2.2. (a) N2 adsorption/desorption isotherm and (b) BJH pore size distribution determined in desorption of the synthesized anatase TiO2 powder.

  5nm

Figure 2 . 3 .

 23 Figure 2.3. Schematic representation of the 4-step heat-treatment procedure to remove ethyl cellulose and terpineol from the screen-printed TiO2 electrodes.

2

 2

Figure 2

 2 Figure 2.4. (a) N2 adsorption-desorption isotherm of the screen-printed TiO2 electrodes and (b) pore size distribution determined by the B.J.H. method.

Figure 2 . 5 .

 25 Figure 2.5. (a) Photograph of a mesoporous screen-printed TiO2 electrode deposited upon the conductive FTO glass substrate and (b) air-tight thermostated borosilicate of the photoelectrochemical cell used in this work.

Figure 2 . 6 .

 26 Figure 2.6. Comparison of cyclic voltammetry of the TiO2 electrode in 1 mol/L LiPF6 EC/DMC in dark (blue curve) and under illumination (red curve) at 2 mV/s scan rate

  .7 (b)). The complete lithium extraction was also confirmed by ICP-AES. These results demonstrate a quantitative nature of the photo-induced Li + deinsertion from the nanoparticle anatase TiO2. It gives credit to the downsizing of the particles that provides an answer to the surface limitation issue of the photoinduced ion transfer reaction [2]. This has been further demonstrated in the Ph.D. thesis work of Gaspard Bouteau comparing different particles size of anatase TiO2 from 4 to 400 nm [29].

Figure 2 . 7 .

 27 Figure 2.7. Electrochemical properties of anatase TiO2 nanoparticles in a half-cell configuration: (a) galvanostatic discharge in dark until 1.5 V (vs. Li + /Li) at C/15 (j = -100 µA/cm 2 ) followed by either 1 hour open-circuit condition in dark (black curve) or under illumination (red curve), and (b) galvanostatic discharge profile of the electrode after one hour open-circuit condition in dark (black curve) or under illumination (red curve).

Figure 2 . 8 .

 28 Figure 2.8. (a) A comparison of galvanostatic discharge at j = -100 μA/cm 2 of nanocrystalline anatase TiO2 electrode in dark (black curve), under illumination (red curve) and under illumination in lithiumfree TBAPF6 EC/DMC electrolyte (blue curve). (b) Galvanostatic discharge at j = -100 μA/cm 2 under reproduced intermittent light intensity conditions with a sequence of 16 hours under illumination and 8 hours in dark conditions.

Figure 2 . 9 .

 29 Figure 2.9. Evolution of diffuse reflectance as a function of wavelength for the 5 nm anatase TiO2 nanoparticles.

Figure 2 . 10 .

 210 Figure 2.10. Tauc plot for 5 nm anatase TiO2 particles for an indirect allowed transition (n=2).

Figure 2 . 11 .

 211 Figure 2.11. In situ UV-visible absorption experiment of 20 nm nanocrystalline anatase TiO2 upon lithium insertion: (a) succession of galvanostatic discharge curves and (b) corresponding evolution of transmittance after each discharge step (spectrum of electrolyte and FTO in electrolyte is provided as reference).

Figure 2 . 12 .

 212 Figure 2.12. Tauc plot of 20 nm anatase TiO2 electrode at different states of discharge in 1mol/L LiPF6 EC/DMC electrolyte. The fitting of the linear part of TiO2 and LixTiO2 curves and the resulting band gap values are provided for information.

Figure 2 . 15 .

 215 Figure 2.15. Schematic representation of the band bending at n-type semiconductor/electrolyte interface.

Figure 2 . 16 .

 216 Figure 2.16. Evolution of (a) flat-band potential and (b) donor density as a function of perturbation frequency for different size of anatase TiO2 particles in 1 mol/L NaOH-based aqueous electrolyte.

Figure 2 . 17 .

 217 Figure 2.17. Frequency dependence and reproducibility over five fresh samples of (a) the flat band potential and (b) the main charge carrier concentration of the 5 nm anatase TiO2 nanoparticles measured between 100 Hz and 1 kHz in 1 mol/L NaOH aqueous electrolyte.

  54]. Taking this as a reference, the values measured vs. Li QRE can be translated into potentials versus other reference electrodes following the chart presented below (Fig. 2.18 (b)).

Figure 2 .

 2 Figure 2.18. (a) Cyclic voltammetry of the Li quasi-reference electrode (QRE) used for Mott-Schottky measurements in non-aqueous aprotic electrolyte containing Fc + /Fc redox couple in EC/DMC-based electrolyte and (b) a chart representing its energetic position versus conventional reference electrodes.

Figure 2 . 19 .

 219 Figure 2. 19. Frequency dependence and reproducibility over two fresh samples (S1 and S2) of (a) the flat band potential and (b) the charge carrier concentration of 5 nm nanocrystalline anatase TiO2 recorded between 100 Hz and 1 kHz in 1M TBAPF6 EC/DMC electrolyte.

Figure 2 . 20 .

 220 Figure 2.20. Mott-Schottky plot of 5 nm nanocrystalline anatase TiO2 measured in (a) 1 mol/L NaOH aqueous electrolyte at 750 Hz and (b) 1 mol/L TBAPF6 EC/DMC electrolyte at 100 Hz.

Figure 2 .

 2 Figure 2.21. Mott-Schottky plot comparing TiO2 and Li0.6TiO2 in 1mol/L TBAPF6 EC/DMC electrolyte at 316 Hz perturbation frequency.

Figure 2 . 22 .

 222 Figure 2.22. Frequency dependence of (a) the flat band potential and (b) charge carrier density of TiO2 and Li0.6TiO2 between 100 Hz and 1 kHz in 1 mol/L TBAPF6 EC/DMC electrolyte.

Figure 2 . 23 .

 223 Figure 2.23. Schematic energy diagram of the valence band and conduction band positions of TiO2 and its discharged form Li0.6TiO2 determined in 1 mol/L TBAPF6 EC/DMC electrolyte. It is compared to HOMO and LUMO positions of the electrolyte (1 mol/L LiPF6 EC/DMC).

Figure 2 . 24 .

 224 Figure 2.24. Jablonski diagram adapted from reference [85].

Figure 2 .

 2 Figure 2.25. Steady-state photoluminescence of 5 nm nano-crystals of anatase TiO2 particles: (a) excitation / emission map, (b) normalized excitation / emission spectrum at maximum intensity, and (c) deconvolution of the emission spectrum measured at 380 nm excitation wavelength.

  2.26 (b)). Being 100% reflective, the reference emission intensity, i.e. 4•10 5 counts, equals to the full incident light intensity received by TiO2 upon excitation. The absolute value of TiO2 emission intensity varying from 1.5•10 2 to a maximum of 1•10 3 counts which in turn represents a fluorescence QY value of 0.25%, thus confirming the very low photoluminescence yield of TiO2 when the integrating sphere was used. This value is in good agreement with literature, where the quantum yield of anatase TiO2

Figure 2 . 26 . 2 . 4 . 3 .

 226243 Figure 2.26. Fluorescence of 5 nm nanocrystalline anatase TiO2 particles: (a) quantum yield spectra measured in direct and indirect mode using an integrating sphere and (b) steady-state emission spectra excited at 380 nm compared to the reference BaSO4 sample.

Figure 2 .

 2 Figure 2.27. Steady-state photoluminescence map of 5 nm nanocrystalline TiO2 electrode in 1 mol/L LiPF6 EC/DMC at different depth of discharge. Note that intensity scale is same for all maps.

Figure 2 .

 2 Figure 2.28. Emission spectra of 5 nm nanocrystalline TiO2 in 1 mol/L LiPF6 EC/DMC at different depth of discharge (Li + insertion) using 277 nm excitation wavelength.

Figure 2 .

 2 Figure 2.29. Steady-state photoluminescence map of (a) LiPF6 EC/DMC electrolyte, (b) colloidal solution of TiO2 in LiPF6 EC/DMC electrolyte and (c) colloidal solution of Li0.6TiO2 in LiPF6 EC/DMC electrolyte. Note that the intensity scale has been reduced 35 times between (a) and (b-c).

2 . 5 . 1 .

 251 photoluminescence (PL) decay.

Figure 2 . 31 .

 231 Figure 2.31. Evolution of the photoluminescence decay of TiO2 and Li0.6TiO2 electrodes in an argon filled cuvette cell. The corresponding instrument response function decays IRF1 and IRF2 are presented for information.

Figure 2 . 30 .

 230 Figure 2.30. Schematic representation of the two-input stopwatch working principle of TCSPC (from Edinburg Instrumets ® TCSPC technical note)

  4). The first component is still ascribed to the direct band-to-band recombination while the second to the trap-mediated fluorescence. The lifetime values are very comparable to those under argon except that more traps are involved in the fluorescence process as a result from the built-in depletion layer at the semiconductor/ electrolyte interface. Note that a difference between 29 ps and 25 ps cannot be discussed because of the IRF value of 80 ps. These results suggest the absence of any interfacial charge transfer process between the semiconductor and the electrolyte, or at least down to ca. 20-30 ps.

Figure 2 . 32 .

 232 Figure 2.32. Photoluminescence decay of the anatase TiO2 nanocrystals deposited upon FTO glass in contact with an electrolyte composed of 1 mol/L TBAPF6 in EC/DMC. The photoluminescence decay of TiO2 under argon and IRF is provided for comparison.

2. 5 . 2 .

 52 Kinetics of charge transfer processes in Li0.6TiO2 in contact with an electrolyteSimilar experiments were performed on the discharged composition Li0.6TiO2. We herein investigated two types of electrolyte, one based on the benchmark organic EC/DMC solvent mixture typically used in batteries (Fig.2.33 (a)), and a second, more inert, based on EMITFSI ionic liquid (Fig.2.33 (b)). The electrolytes were tailored by the nature of the added salt to be either ionic-blocking to hamper the electrode photorecharge (TBAPF6) or ionic-conducting to enable the electrodes' photorecharge (LiPF6). For the measurements in lithium-free electrolytes, Li0.6TiO2 was obtained by galvanostatic discharge inside an Ar-filled glove-box at j = -100 µA/cm 2 in LP30 electrolyte until 1.5 V vs. Li + /Li. After lithiation, the electrode was thoroughly rinsed with DMC as for the Mott-Shottky experiments. The electrode was then transferred into a TCSPC cuvette filled with the electrolyte and sealed inside the glove-box to ensure its air-tightness. For the lithium-ion conducting electrolytes, the time-resolved electroluminescence experiments were performed in a three-electrode air-tight cuvette with TiO2 as a working electrode, and two lithium strips as counter and reference electrode. By combining the galvanostat with the TCSPC spectrometer, the galvanostatic discharge at j = -100 µA/cm 2 down to 1.5 V vs. Li + /Li was directly performed inside the spectrometer chamber. Then, the electrode was potentiostatically stabilized at 1.5 V vs. Li + /Li until steady-state current close to 0 was obtained before any fluorescence decay was measured. The evolution of the photoluminescence decays of Li0.6TiO2 electrode in lithium ion-blocking and lithium ion-conducting electrolytes based on either EC/DMC solvent mixture or EMITFSI ionic liquid is presented in Figure2.33.

Figure 2 . 33 .

 233 Figure 2.33. Evolution of the photoluminescence decay of Li0.6TiO2 electrode in (a) EC/DMC mixture or (b) EMITFSI ionic liquid based electrolytes containing or not lithium ions. The decay of Li0.6TiO2 in argon and IRF are provided for comparison.

  + -containing electrolyte's decays are faster than those determined in lithium-free electrolytes. The reconvolution of the PL decays in lithium-based electrolytes requires systematically the means of a third fast contribution

  liquid which reaches up to 14.6 ns with an amplitude of 69.4%. Along this line, we have evidenced experimentally a Stark-Lo Surdo effect on our nanocrystalline anatase TiO2 particles induced by the specific adsorption of lithium cation upon the surface of the particles which is visible on the excitation/emission spectrum (Fig.2.34), band gap value from 2.90 eV to 3.60 eV for TiO2 in pure DMC and in highly concentrated 1.5 M of LiPF6. Stark-Lo Surdo effect has also important consequences in the photoluminescence dynamics which we think is a beneficial mechanism for the electrode's photorecharge.

Figure 2 .

 2 Figure 2.34. (a) Steady-state excitation/emission spectrum, (b) Tauc plot, and (c) time-resolved emission spectroscopy of 5 nm nanocrystalline anatase TiO2 film in contact with different types of electrolytes.

  By means of UV-visible absorption spectroscopy, we established that the color change from white to dark blue upon lithium insertion into TiO2 is associated on the one hand with an increase of the bandgap value, which can be ascribed to Burstein-Moss effect, and on the other hand with the rise of absorbance in both visible and near-infrared part of the light spectrum.

Figure 3 . 1 .

 31 Figure 3.1. Crystal structure of the spinel Li4Ti5O12.

Figure 3 .

 3 Figure 3.2. (a) Full pattern matching refinement of x-ray diffractogram of Li4Ti5O12, (b) scanning electron microscopy image, (c) transmission electron micrograph, (d) high-resolution transmission electron micrograph, and (e) selected area electron diffraction pattern of the particles.

Figure 3

 3 Figure 3.3 (a) presents the diffuse reflectance spectrum of Li4Ti5O12 recorded over the range of 1800 nm down to 200 nm. For this, 1 wt% of Li4Ti5O12 was mixed within KBr before to be pressed into a pellet. Note that this ratio was kept for all materials herein studied unless mentioned otherwise. An absorption tail is observed at around 400 nm suggesting a wide band gap in agreement with the white color of the particles. The band gap value was determined from the Tauc plot using Kubelka-Monk function. Assuming an indirect allowed transition, the band gap energy of Li4Ti5O12 is 3.79 eV, which is in agreement with the white color of the particles (Fig.3.3 (b)). It is 4.05 eV if we assume a direct transition (Fig.3.3 (c)). Our results are significantly higher than 1.7 -2.3 eV reported by early calculations[21][22][23][24], but close to 3.87 eV reported recently by Verde et al. based on DOS calculations[25]. The obtained values are also in agreement with the 3.8 eV indirect band gap reported by Ge et al. based on diffuse reflectance measurements[26] and with the range of 3.5 -3.8 eV band gap with a direct allowed transition reported by a number of other authors[27][28][29].

Figure 3 . 3 .

 33 Figure 3.3. (a) Evolution of diffuse reflectance of Li4Ti5O12 as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed band gap transition (n=2) and (c) direct allowed band gap transition (n=0.5).

  For a direct transition, the absorbance increases slowly by contrast to an indirect transition for which the absorbance change becomes steep. Taking into account the inverse relation between absorbance and reflectance in opaque materials, thus, the steep reflectance onset of Li4Ti5O12 suggests an indirect band gap. The flat band potential as well as the charge carrier type and concentration in Li4Ti5O12 was determined as previously from the Mott-Schottky plot in two different kinds of blocking electrolytes, i.e. aqueous (protic) and non-aqueous aprotic. For this, a film with 12 µm thickness was prepared by doctor-blading. In aqueous electrolyte, we carried out a set of measurements in different electrolytes, namely in NaOH and in TBANO3 at different pH values adjusted by adding a small quantity of HNO3. We then chose one representative measurement based on quality of the Mott-Schottky curve and the highest reproducibility of the flat band potential (Vfb) and the charge carrier concentration (N) values. The frequency dependence of the Mott-Schottky curves for Li4Ti5O12 and for the other materials described in this chapter follows the same trend as observed for TiO2, i.e. a decrease of Vfb and negligible changes in N as a function of frequency in the range from 100 Hz to 1 kHz. This is experienced for measurements in both aqueous and non-aqueous aprotic electrolytes. Figure 3.4 presents the Mott-Schottky plots of Li4Ti5O12 in aqueous and in non-aqueous aprotic EC/DMC-based electrolytes. In both media, the positive slope of the Mott-Schottky curve confirms the n-type nature of conductivity in agreement with literature [26, 34].

Figure 3 . 4 .

 34 Figure 3.4. Mott-Schottky plot of Li4Ti5O12 measured in (a) 1 mol/L NaOH aqueous electrolyte at 250 Hz and (b) 1 mol/L TBAPF6 EC/DMC electrolyte at 200 Hz.

Figure 3 . 5 .

 35 Figure 3.5. Representation of the band edge positions of Li4Ti5O12 film in aqueous electrolyte at pH = 14 (in black) and in non-aqueous aprotic EC/DMC-based electrolyte (in red).

Figure 3 . 6 .

 36 Figure 3.6. Steady-state (a) excitation/emission map and (b) normalized excitation/emission spectra of nanocrystalline Li4Ti5O12 powder.

Figure 3 . 7 .Figure 3 . 8 .

 3738 Figure 3.7. Quantum yield fluorescence spectra of Li4Ti5O12 particles measured in direct and indirect method using an integrating with BaSO4 as a reference sample.

3 .

 3 The fast component at 0.10 ns cannot be related to a direct band-to-band recombination due to the large Stokes shift. Therefore, all the three components, i.e. 0.10, 3.00 and 27.87 ns, are assigned to a very deep trap-mediated fluorescence. A low amplitude of the first component reaching only 6.9% indicates that a very small portion of recombination takes place through the shallow trap states. The 61.0% and 32.1% of the total fluorescence amplitude belonging to the second and third components supports the radiative recombination through the deep traps. This can also explain the origin of the very low quantum yield of Li4Ti5O12.

Figure 3 . 9 .

 39 Figure 3.9. Crystal structure representation of LiCoO2 in (011) plane.

Figure 3 .Figure 3 .

 33 Figure 3.10 (a) shows the x-ray diffraction pattern of LiCoO2 powder used in this study. All diffraction peaks are indexed into the rhombohedral R-3m space group with lattice cell parameters of a = b = 2.8149(3) Å, c = 14.050(51) Å (Vcell = 96 Å 3 ). These values are in good agreement with those reported for instance by Shinova et al. [40]. The diffraction peaks are intense and narrow indicating excellent crystallinity with large crystallite size. The SEM micrographs show well-separated particles having no specific morphology and the size ranging from 2 to 20 µm. There are also spherical nanosized particles of ca. 50 nm (Fig. 3.10 (b-c)).

Figure 3 .

 3 Figure 3.11. (a) Evolution of diffuse reflectance of LiCoO2 powder as a function of wavelength and (b) corresponding Tauc plot assuming an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

Figure 3 . 12 .

 312 Figure 3.12. Mott-Schottky plot of LiCoO2 measured at 100 Hz in (a) 0.5 mol/L TBANO3 aqueous electrolyte and (b) 1 mol/L TBAPF6 EC/DMC electrolyte.

  6 EC/DMC 0.5M TBANO 3 aq.

Figure 3 . 13 .

 313 Figure 3.13. Representation of the band edge positions of LiCoO2 film in aqueous electrolyte at pH = 14 (black) and in non-aqueous aprotic EC/DMC-based electrolyte (red).

  structure. Each layer is formed by two sheets of MoO6 octahedra sharing corners along [001] and [100] directions, and connected together by edge sharing along [001] direction (Fig. 3.14).The layers are alternatively stacked along the c-axis and maintained by electrostatic van der Waals' forces. This two-dimensional structure allows a reversible intercalation of ca. 1.5 lithium per molybdenum transition metal[48][49], resulting in a lithiated phase (LixMoO3) with a good electronic conductivity (10 -2 S/cm) and Li + diffusion coefficient (10 -10 cm 2 /s)[50][51].

Figure 3 . 14 .

 314 Figure 3.14. Crystal structure representation of MoO3 according to (101) plane.

Figure 3 .

 3 Figure 3.15. (a) Full pattern matching refinement of x-ray diffraction pattern of MoO3 with its corresponding (b,c) transmission and high-resolution transmission electron micrographs and (d) selected area electron diffraction pattern.

Figure 3 .

 3 Figure 3.16 (a) shows the evolution of the reflectance as a function of wavelength for MoO3 between 200 and 1800 nm. Assuming an indirect allowed transition, which we argue to be the case based on the steepness of the transition, we found a band gap of 2.93 eV (Fig. 3.16 (b)).

Figure 3 .Figure 3 . 17 .Figure 3 . 17 .Figure 3 . 18 .

 3317317318 Figure 3.16. (a) Evolution of diffuse reflectance of MoO3 powder as a function of wavelength using an integrating sphere and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

Figure 3 .Figure 3 .

 33 Figure 3.19. Crystal structure representation of WO3 according to (011) plane.

Figure 3 .Figure 3 . 22 .

 3322 Figure 3.21. (a) Evolution of diffuse reflectance of WO3 nanoparticles as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

Figure 3 . 23 .

 323 Figure 3.23. Representation of the band edge positions of WO3 film measured in aqueous electrolyte and recalculated to pH = 14 (black) and in non-aqueous aprotic EC/DMC-based electrolyte (red).

  The charge carrier concentration of WO3 was determined with the dielectric constant of 20 as proposed byWang et al. (Table

  structure, forming a zigzag Cu-O chains oriented along the [101 ̅ ] and [101] directions (Fig. 3.24). The refinement of x-ray diffraction pattern of CuO shows single phase that crystallizes into a monoclinic-type structure with C2/c space group (Fig. 3.25 (a)). The refined lattice cell parameters are a = 4.690(8) Å, b = 3.420(16) Å, c = 5.131(3) Å, and β = 99.54° (V = 81 Å 3 ).

Figure 3 . 24 .

 324 Figure 3.24. Crystal structure representation of CuO.

Figure 3 .

 3 Figure 3.25. (a) Full pattern matching refinement of x-ray diffraction profile of CuO powder with corresponding (b) transmission electron micrograph and (c) high-resolution transmission electron micrograph with (d) selected area electron diffraction pattern.

Figure 3 .

 3 Figure 3.26. (a) Evolution of diffuse reflectance of CuO powder as a function of wavelength. The corresponding Tauc plot is reported assuming (b) an indirect allowed transition (n=2) and (c) a direct (n=0.5) allowed transition.

  position of conduction and valence band, as well as the type of conductivity, Mott-Schottky impedance was measured on 3 µm thick films deposited on FTO by doctor blading. When measured in a range of aqueous protic electrolytes including NaOH, KOH, Na2HPO4, and TBANO3 at different pH values, the obtained Mott-Schottky plots were not reproducible. Apart from the significant shift in the Mott-Schottky onset, which can vary to an extent of 0.4, in certain cases, the slope of the curve was changing from positive at lower potentials to negative at higher potentials. This traduces either some kind of side reactions taking place in the electrode depending on the applied potentials, or it can be associated with a crossover change of n to p-type conductivity depending on the applied voltage owing to relative comparable density of holes and electrons in CuO. In addition, for half of the electrodes, we faced issues of bad adhesion onto the FTO substrate leading to the film pilling off at the end of the experiment. In contrast, in 1 mol/L TBAPF6 EC/DMC electrolyte, a reproducible Mott-Schottky plot was successfully obtained (Fig.3.27). The negative slope of the linear part of the plot indicates a p-type conductivity in agreement with literature[87][88][89]. This p-type conductivity stems from the existence of copper vacancies that facilitate hole conduction through Cu-O planar rectangles.

Figure 3 .

 3 Figure 3.27. Mott-Schottky plot of CuO measured in 1 mol/L TBAPF6 EC/DMC electrolyte at 300 Hz.

  vs Li QRE) 1M TBAPF 6 EC/DMC Taking into account the 1.19 eV band gap of the material, the position of the conduction band edge is -0.59 V vs. NHE (Fig. 3.28).

A

  hole concentration of 2.7•10 19 cm -3 has been determined based on the dielectric constant of 10 as reported by Nakaoka et al. [89]. This number is in good agreement with the values reported in literature between 4.0•10 18 to 4.0•10 20 cm -3 [87-89].

Figure 3 . 28 .

 328 Figure 3.28. Band edge positions of CuO film in non-aqueous aprotic EC/DMC-based electrolyte.

  along the a-and c-direction as a result of the different coordination polyhedra sharing only corners in the three-dimensional array (Fig.3.29)[96].

Figure 3 . 29 .Figure 3 .

 3293 Figure 3.29. Crystal structure representation of β-Bi2O3.

Figure 3 .

 3 Figure 3.31. (a) Evolution of diffuse reflectance of Bi2O3 powder as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

Figure 3 .Figure 3 . 32 .

 3332 Figure 3.32 presents the Mott-Schottky plots of 5 µm thick doctor-bladed Bi2O3 films in either aqueous electrolyte (a) or non-aqueous aprotic EC/DMC-based electrolyte (b). The positive slope of the curves indicates a n-type conductivity in agreement with previous literature (Fig. 3.32) [94, 98, 105].

Figure 3 . 33 .

 333 Figure 3.33. Representation of the band edge positions of Bi2O3 film measured in aqueous electrolyte at pH = 14 (black) and in non-aqueous aprotic EC/DMC-based electrolyte (red).

  3.34). The empty sites in the [Mn2O4] -framework represent a diamond type network of tetrahedral 8a and octahedral 16c sites, which share common faces and edges and form 3D open channels for Li + ion diffusion. The full pattern matching refinement of LiMn2O4 diffractogram is presented in Figure 3.35 (a).All diffraction peaks can be indexed into Fd-3m space group. The refined lattice cell parameter is a = 8.238(3) Å (Vcell = 559 Å 3 ) in agreement with the value of a = 8.2404(2) reported by Mukai et al.[111]. The SEM micrograph shows particles with octahedral morphology and very inhomogeneous particle size ranging from ca. 500 nm to 5 µm (Fig.3.35 (b)). Significant level of particles agglomeration is observed.

Figure 3 . 34 .

 334 Figure 3.34. Crystal structure representation of the spinel LiMn2O4.

Figure 3 .

 3 Figure 3.35. (a) Full pattern matching of x-ray diffraction pattern of LiMn2O4 powder and corresponding (b) scanning electron micrograph and (c,d) high-resolution transmission electron micrograph and selected area electron diffraction pattern.

Figure 3 . 36 . 3 . 1 . 8 .

 336318 Figure 3.36. Evolution of diffuse reflectance of LiMn2O4 powder as a function of wavelength.

Figure 3 . 37 .O

 337 Figure 3.37. Schematic representation of LiMn1/3 Co1/3 Ni1/3O2 crystal structure.

Figure 3 .

 3 Figure 3.38. (a) Full pattern matching refinement of the x-ray diffraction pattern of LiMn1/3Co1/3Ni1/3O2 with the corresponding (b) scanning electron micrograph, (c) transmission electron micrograph and (d,e) high-resolution transmission electron micrograph including the selected area electron diffraction pattern.

  there are no reports related to the determination of the LiMn1/3Co1/3Ni1/3O2 band gap value in literature.

Figure 3 . 39 .

 339 Figure 3.39. Evolution of diffuse reflectance of LiMn1/3Ni1/3Co1/3O2 powder as a function of wavelength.

Figure 3 . 40 .

 340 Figure 3.40. Crystal structure representation of layered MS2

Figure 3 .

 3 Figure 3.41. (a) Full pattern matching refinement of x-ray diffraction pattern of TiS2 and corresponding (b) scanning electron micrograph, (c) transmission electron micrograph and (d,e) high-resolution transmission electron micrograph with selected area electron diffraction pattern.

Figure 3 .

 3 Figure 3.42 presents the diffuse reflectance spectrum of TiS2 powder. The absence of any

Figure 3 . 42 .Figure 3 .

 3423 Figure 3.42. Evolution of diffuse reflectance of TiS2 powder as a function of wavelength.

Figure 3 .

 3 Figure 3.44 (a) exhibits the full pattern matching refinement of the XRD pattern of MoS2. All diffraction peaks were successfully indexed into hexagonal crystal structure of P63/mmc space group. The obtained cell parameters are a = b = 3.160(7) Å and c = 12.294(9) Å (Vcell = 106 Å 3 ). They are in good agreement with the results reported by Bronsema et al. [140]. The electron micrographs show that the particles display a nanoflake-like morphology with flakes diameter ranging from 1 to 10 µm without amorphous domain at the particles surface (Fig. 3.44 (b-d)). The SAED pattern is consistent with MoS2 crystal structure (Fig. 3.44 (e)).

Figure 3 .

 3 Figure 3.44. (a) Full pattern matching refinement of x-ray diffraction pattern of MoS2 and (b) scanning electron micrograph, (c) transmission electron micrograph and (d,e) high-resolution transmission electron micrograph including selected area electron diffraction pattern.

  photocurrent study[142]. DFT calculations predict an indirect bandgap nature between 1.04 eV to 1.29 eV based on the DOS calculation under vacuum[143][144][145][146][147]. The indirect bandgap is consistent with the slow increase of reflectance near the absorption threshold.

Figure 3 .

 3 Figure 3.45. (a) Evolution of diffuse reflectance of MoS2 powder as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2).

Figure 3 . 46 .

 346 Figure 3.46. Mott-Schottky plot of MoS2 measured at 100 Hz in (a) 1 mol/L NaOH aqueous electrolyte and (b) 1 mol/L TBAPF6 EC/DMC electrolyte.

Figure 3 . 47 .

 347 Figure 3.47. Representation of the band edge positions of MoS2 film measured in aqueous electrolyte at pH = 14 (black) and in non-aqueous aprotic EC/DMC-based electrolyte (red).

  corresponding Tauc plot (Fig.3.49 (b)

1 .

 1 31 eV (Fig. 3.29 (c)). Nevertheless, the very broad reflectance onset clearly suggests an indirect transition.

Figure 3 .

 3 Figure 3.49. (a) Evolution of diffuse reflectance of WS2 powder as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (b) a direct allowed transition (n=0.5).

Figure 3 .

 3 Figure 3.50. Mott-Schottky plot of WS2 measured at 100 Hz in (a) 1 mol/L NaOH aqueous electrolyte and (b) 1 mol/L TBAPF6 EC/DMC electrolyte.

Figure 3 . 51 .

 351 Figure 3.51. Representation of the band edge positions of WS2 film measured in aqueous electrolyte at pH = 14 (black) and in non-aqueous aprotic EC/DMC-based electrolyte (red).

Figure 3 . 52 .

 352 Figure 3.52. Crystal structure representation of LiMPO4 materials down to [001] and [010] directions.

Figure 3 .

 3 Figure 3.53. Crystal structure representation of LiVPO4F down to [100] direction.

  .54 (a). All diffraction peaks are corresponding to the orthorhombic lattice of LiFePO4 adopting Pnma space group. The lattice cell parameters refined are a = 10.332(9) Å, b = 6.019(3) Å and c = 4.704(5) Å (Vcell = 293 Å) in agreement with the literature [180]. The SEM study revealed particles without well-defined morphology and broad size distribution ranging from few micrometers to ca. 80 µm (Fig. 3.54 (b)). The HR-TEM micrograph and SAED pattern confirm the XRD indexation and the absence of amorphous regions in the particles (Fig. 3.54 (d-e)).

Figure 3 .

 3 Figure 3.54. (a) Full pattern matching refinement of x-ray diffraction profile of LiFePO4, (b) scanning electron micrograph, (c) transmission electron micrograph and (d,e) high-resolution transmission electron micrograph including selected area electron diffraction pattern.

of 3 .

 3 8 eV determined as the absorption onset and 4.0 eV using the Kubelka-Monk function. Shi et al. reported a band gap of 3.2 eV based on same DFT method (GGA+U) [185]. Our 3.11 eV direct band gap value comes in agreement with these calculations. The light sand-like colour of the particles is consistent with such a bandgap value.

Figure 3 .

 3 Figure 3.55. (a) Evolution of diffuse reflectance of LiFePO4 powder as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

Figure 3 .Figure 3

 33 Figure 3.56. Mott-Schottky plot of LiFePO4 measured in (a) 0.5 mol/L TBANO3 aqueous electrolyte at 250 Hz and (b) 1 mol/L TBAPF6 EC/DMC electrolyte at 100 Hz.

Figure 3 . 57 .

 357 Figure 3.57. Representation of the band edge positions of LiFePO4 measured in aqueous electrolyte and recalculated to pH = 14 (black) and in non-aqueous aprotic EC/DMC-based electrolyte (red).

  are forming aggregates of up to 10 µm in diameter (Fig. 3.58 (b)). The size of the primary particles is ranging from 20 to 200 nm (Fig. 3.58 (c)). The high-resolution TEM micrograph and SAED pattern confirms that particles are single crystals of LiMnPO4 without the presence of amorphous regions (Fig. 3.58 (d-e)).

Figure 3 .

 3 Figure 3.58. (a) Full pattern matching refinement of x-ray diffraction pattern of LiMnPO4 with its corresponding (b) scanning electron micrograph, (c) transmission electron micrograph and (d,e) high-resolution transmission electron micrograph including selected area electron diffraction pattern.

Figure 3 .

 3 Figure 3.59(a) shows that LiMnPO4 has good reflecting properties in all visible and NIR part of light. Using Kubelka-Monk function, the optical band gap for an indirect allowed transition

Figure 3 .

 3 Figure 3.59. (a) Evolution of diffuse reflectance of LiMnPO4 powder as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

Figure 3 .

 3 Figure 3.60. Mott-Schottky plot of LiMnPO4 measured in 1 mol/L TBAPF6 EC/DMC electrolyte at 300 Hz.

Figure 3 . 61 .

 361 Figure 3.61. Comparison of the band edge positions of LiFePO4 and LiMnPO4 films in non-aqueous aprotic 1 mol/L TBAPF6 EC/DMC-based electrolyte.

Figure 3 .

 3 62 (a) including its full-pattern matching refinement into a triclinic unit cell of P-1 space group. The sample is free of any impurities. The lattice cell parameters refined are a = 5.168(4) Å, b = 5.309(8) Å, c = 7.263(6) Å, α = 107.599°, β = 108.02°, γ = 98.336° (Vcell = 174 Å 3 ). These values are in a good agreement with the crystallographic data reported by Ateba Mba et al. [179]. The SEM and TEM micrographs show the nanometer size of the particles forming rock-like agglomerates of few tenth of micrometers (Fig. 3.62 (b-c)). HR-TEM investigations confirmed the absence of amorphous regions while showing good crystallinity of the material to the edges of the particles (Fig. 3.62 (d-e)).

Figure 3 .

 3 Figure 3.62. (a) Full pattern matching refinement of x-ray diffraction pattern of LiVPO4F powder with its corresponding (b) scanning electron micrograph, (c) transmission electron micrograph and (d,e) high-resolution transmission electron micrograph including selected area electron diffraction pattern.

Figure 3 .Figure 3 . 65 .

 3365 Figure 3.63. (a) Evolution of diffuse reflectance of LiVPO4F powder as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

Figure 3 .

 3 Figure 3.66 shows a comparison in dark and under illumination of the cyclic voltammogram of 16 µm thick film of LiCoO2 at a scan rate of 0.1 mV/s. In dark, the first cycle shows an irreversible anodic peak between 4 and 4.2 V corresponding to the lithium extraction from

Figure 3 . 66 .

 366 Figure 3.66. Comparison of the first two cycles of cyclic voltammetry recorded on 16 µm thick LiCoO2 electrode at 0.1 mV/s scan rate in 1 mol/L LiPF6 EC/DMC, in dark (black curve) and under A.M. 1.5G illumination (red curve).

  3.67). The electrochemical behavior shows the typical redox activity of LiFePO4 with a redox potential determined at half-wave at ca. 3.42 V vs. Li + /Li. In contrast to the case of LiCoO2, we do not see here any predominant electrode degradation up to 4 V under illumination. This means that the excited states of LiFePO4 have no photo-catalytic properties with respect to the electrolyte oxidation.

Figure 3 . 67 .

 367 Figure 3.67. Evolution of the cyclic voltammograms of 50 µm thick LiFePO4 film measured in 1 mol/L LiPF6 EC/DMC as a function of the scan rate (0.5 mV/s < ν < 20 mV/s) (a) in dark and (b) under standard A.M. 1.5G illumination conditions.

Figure 3 . 69 .

 369 Figure 3.69. Evolution of the (a) potential of the faradaic peaks separation and (b) resistance of anodic and cathodic processes as a function of scan rate in dark and under illumination for 50 µm LiFePO4 film in 1 mol/L LiPF6 EC/DMC.

Figure 3 . 70 .

 370 Figure 3.70. Evolution of the open circuit potential of LiFePO4 film under standard A.M. 1.5G conditions.

Figure 3 . 71 .

 371 Figure 3.71. Open circuit voltage and discharge curves of LiFePO4 on FTO glass reported byPaolella et al.[191]. OCV charge (red lines) were performed under illumination at100 mW/cm 2 and galvanostatic discharge (blue lines) at C/24.

  Pnma space group. The refined lattice cell parameters are a = 10.301(6) (9) Å, b = 5.993(4) Å and c = 4.694(5) Å (Vcell = 290 Å). The result is close to the values obtained for the bulk LiFePO4 (a = 10.332(9) Å, b = 6.019(3) Å and c = 4.704(5) Å (Vcell = 293 Å)), even though a small decrease of the unit cell volume by 3 Å is observed. These minor differences in the cell structure can be assigned to the higher number of the punctual defects present in the nanopowder. The TEM micrographs indicate that the particles are single crystals of LiFePO4 with a platelet morphology of about 100-150 nm in diameter (Fig. 3.72 (b-d)).

Figure 3 .

 3 Figure 3.72. (a) Full pattern matching refinement of x-ray diffraction pattern of LiFePO4 nanoparticles with the corresponding (b-d) transmission electron micrographs and selected area electron diffraction pattern.

Figure 3 . 73 .

 373 Figure 3.73. Evolution of open circuit potential in dark or under illumination under A.M. 1.5G conditions (a) from LiFePO4 (b) from "FePO4" obtained after galvanostatic charge to 4.0 V vs Li + /Li at 10 µA/cm 2 .

Figure 3 . 74 .

 374 Figure 3.74. Evolution of the 300 nm thick nanoparticle LiFePO4 film transmittance at different states of charge from 3.0 to 4.0 V vs. Li + /Li. The transmittance spectra of the electrolyte (LP30) and transparent conducting glass (TCO) substrate of electrode in LP30 are provided for comparison.

Figure 3 .

 3 75 shows a comparison of the open circuit potential evolution of LiFePO4 electrode in dark, under illumination, and under illumination when sensitized. Regardless of the electrode, there is no photorecharge process.

Figure 3 . 75 .

 375 Figure 3.75. Evolution of the open circuit voltage of nanoparticles of LiFePO4 measured in dark (black) and under standard A.M. 1.5G illumination conditions with (blue) or without (red) N719 dye.

First, 2

 2 g of LiFePO4 was added into a solution containing 3.36 g of NO2BF4 in 20 mL of acetonitrile. The mixture was stirred for four days at room temperature in an argon-filled glovebox to allow complete chemical oxidation of LiFePO4 (800 to 1500 ppm of O2, 3.5 ppm of H2O). The resulting powder was retrieved after three times centrifugation with acetonitrile. The XRD pattern of the resulting powder is presented in Figure 3.76. All diffraction peaks were indexed into Pnma space group in agreement with the heterosite FePO4 phase without the presence of unoxidized LiFePO4. The lattice cell parameters are a = 10.306(14) Å, b = 6.006(4) Å and c = 4.704(7) Å (V = 291 Å), in agreement with literature [214-215].

Figure 3 . 76 .

 376 Figure 3.76. Full pattern matching refinement of the chemically obtained FePO4 powder.

Figure 3 .

 3 Figure 3.77. (a) Evolution of diffuse reflectance of the chemically obtained FePO4 powder as a function of wavelength and corresponding Tauc plot assuming (b) an indirect allowed transition (n=2) and (c) a direct allowed transition (n=0.5).

  and charge carrier concentration values. In an aqueous 0.5 mol/L TBANO3 electrolyte, a flat band potential of -0.33 V vs. SCE was determined at pH = 6.1, which corresponds to -0.56 V vs. NHE at pH = 14. In a non-aqueous aprotic 1 mol/L TBAPF6 EC/DMC electrolyte, the flat band potential value is equal to 2.36 V vs. Li QRE, corresponding to -0.83 V vs. NHE.

Figure 3 .

 3 Figure 3.78. Mott-Schottky plot of FePO4 measured in (a) 0.5 mol/L TBANO3 aqueous electrolyte at 250 Hz and (b) 1 mol/L TBAPF6 EC/DMC electrolyte at 200 Hz.

Figure 3 . 79 .

 379 Figure 3.79. Energy diagram of the band edge positions of LiFePO4 and FePO4 measured in 0.5mol/L TBANO3 aqueous electrolyte and recalculated for pH = 14 (on the left), and in non-aqueous aprotic 1 mol/L TBAPF6 EC/DMC-based electrolyte (on the right). The information on the HOMO and LUMO levels of N719-dye is provided for comparison[218].

Figure 3 . 80 . 3 .

 3803 Figure 3.80. Comparison of the cyclic voltammograms of 14 µm thick Li4Ti5O12 electrode measured in 1 mol/L LiPF6 EC/DMC as a function of the scan rate from 0.05 to 0.5 mV/s (left to right) and cycle number (top down).

Figure 3 . 81 .

 381 Figure 3.81. Evolution of (a) peak current density and (b) electrode capacity as a function of scan rate for a 14 µm thick Li4Ti5O12 film in dark or under illumination.

Figure 3 .

 3 Figure 3.82 (a) depicts the galvanostatic cycling of Li4Ti5O12 electrode in dark and underillumination. Interestingly, we can observe that light illumination leads to a lower overpotential for the redox process at an applied current of +/-100 µA/cm 2 . This latter is more visible for the oxidation in agreement with the n-type character of Li4Ti5O12. There is no drastic influence of light on the electrode's capacity in agreement with the presented above cyclic voltamperometry experiment. Nevertheless, as we also experienced in cyclic voltamperometry and for TiO2 study, the electrode cyclability tends to fade faster when cycling under illumination. This is clearly showed by reporting the coulombic efficiency as a function of number of cycles (Figure3.82 (b)). For the first cycle, the coulombic efficiency of the electrode

Figure 3 .

 3 Figure 3.82. (a) Galvanostatic cycling of 13.5 µm thick Li4Ti5O12 electrode versus lithium at 100 µA/cm 2 in dark or under illumination and (b) corresponding evolution of the coulombic efficiency as a function of cycle number.

Figure 3 . 83 .Figure 3 . 84 .

 383384 Figure 3.83. Transmission electron micrographs of Li4Ti5O12 electrode after 10 cycles of galvanostatic cycling at 100 µA/cm 2 (a) in dark and (b) under illumination.

Figure 3 . 85 .

 385 Figure 3.85. Comparison of the photoluminescence decay of Li4Ti5O12 in contact with argon (red) and in contact with 1 mol/L LiPF6 EC/DMC electrolyte (blue) recorded at 728 nm emission wavelength.

  86 (a). It shows the presence of two bands. The high intensity band at 333 nm excitation and 735 nm emission wavelengths corresponds to the fluorescence of Li4Ti5O12 (Fig.3.86 (b)). However, when in contact with the electrolyte, the intensity of Li4Ti5O12 emission drops tenfold due to the fluorescence quenching. The low intensity band at around 470 nm excitation and 520 nm emission wavelengths can be assigned to the tail of the electrolyte photoluminescence stemming from the solvated Li + and PF6 -in EC/DMC (Fig.3.86 (c))[START_REF] Bouteau | Etude des propriétés de l'interface semi-conducteur / électrolyte sous illumination pour batteries à ion-lithium photo-rechargeables[END_REF]. Interestingly, the highest intensity band of 1 mol/L LiPF6 EC/DMC electrolyte (LP30) located at 280 nm excitation and 310 nm emission wavelengths disappears completely when in contact with Li4Ti5O12. This can be explained by a Förster resonance energy transfer (FRET) from the electrolyte to the electrode due to a partial overlapping of the electrolyte emission (centered at ca. 310 nm) with Li4Ti5O12 excitation (excitation band at 333 nm).

Figure 3 .

 3 Figure 3.86. (a) Steady-state photoluminescence mapping of Li4Ti5O12 powder in 1 mol/L LiPF6 in EC/DMC electrolyte. The photoluminescence maps of (b) 1 mol/L LiPF6 in EC/DMC electrolyte (LP30) and (c) Li4Ti5O12 powder are provided for comparison.

Figure 3 .

 3 Figure 3.87 presents the evolution of the emission spectra of Li4Ti5O12 electrode in 1 mol/L LiPF6 EC/DMC upon lithium insertion into the structure recorded in situ at 333 nm excitation wavelength. The emission spectrum of the fresh Li4Ti5O12 electrode in 1 mol/L LiPF6 EC/DMC atOCV has two emission bands: one at around 390 nm assigned to the electrolyte, and a second at 728 nm stemming from Li4Ti5O12. Upon discharge down to 1.2 V vs. Li + /Li, leading to Li7Ti5O12, we observed a complete bleaching of the 728 nm emission band coming with a fluorescence decrease of the electrolyte at 390 nm. Because of this complete quenching, we have not been able

Figure 3 . 87 .

 387 Figure 3.87. Evolution of the emission spectrum of Li4Ti5O12 electrode in LiPF6 EC/DMC recorded at 333 nm excitation wavelength in the initial state (OCV 3.4 V vs. Li + /Li), discharged (1.2 V vs. Li + /Li) and after first recharge (2.6 V vs. Li + /Li).

4

 4 Ti 5 O 12 in LP30 (at OCV 3.4V) Li 7 Ti 5 O 12 in LP30 (discharged to 1.2 V) Li 4 Ti 5 O 12 in LP30 (recharged to 2.6V) Counts Emission wavelength (nm)This phenomenon is assigned to the absorption by the free carriers populating the conduction band. The derived Tauc plot assuming an indirect allowed transition shows that Li4Ti5O12 has a band gap moving down to 3.48 eV when in contact with the electrolyte (Fig.3.88 (b)). This is ca. 0.3 eV lower than the 3.79 eV determined for the powder in air. For the discharged to Li7Ti5O12 composition electrode, the band gap decreases to 3.25 eV. It is assigned to the electron filling of the lower energy levels close to the conduction band (Ti 3+ : (4s 0 3d 1 ) electronic configuration)[25].

Figure 3 .

 3 Figure 3.88. (a) Evolution of the transmittance spectrum of 250 nm thick film of Li4Ti5O12 in 1 mol/L LiPF6 in EC/DMC electrolyte as a function of lithium insertion at -10 µA/cm 2 and (b) Tauc plot of Li4Ti5O12 (3.0 V spectrum) and its lithiated form Li7Ti5O12 (1.2 V spectrum) assuming an indirect allowed transition.

Figure 3 . 89 .

 389 Figure 3.89. Evolution of Mott-Schottky analysis of Li4Ti5O12 (3.3 V) and Li7Ti5O12 (1.2 V) electrodes recorded at 200 kHz in 1 mol/L TBAPF6 EC/DMC electrolyte. A recharged Li4Ti5O12 (2.6 V) electrode is presented for additional information.

[ 222 ]

 222 ). After the recharge, the charge carrier concentration decreases down to 2.4•10 19 cm -3 .The resulting donor density value is probably a combination of the charge carrier concentration of Li7Ti5O12, which increases upon lithium insertion, and LiF. This would be also consistent with the carrier concentration measured on the SEI formed upon FTO, which is in same order of magnitude (Fig.3.90 (b)).

Figure 3 . 90 .

 390 Figure 3.90. Frequency dependence of the (a) flat band potential and (b) the charge carrier concentration on pristine FTO glass (black line) and FTO glass discharged to 1.2 V vs. Li + /Li for which SEI is formed on the surface (red line). The measurements were performed on three different samples between 100 Hz and 1 kHz in 1 mol/L TBAPF6 EC/DMC electrolyte.

  band edge positions in EC/DMC-based electrolyte and the redox energy of the transition metal of the material, we deciphered materials which can or those which cannot trigger the Li + insertion / deinsertion from the structure by photo-induced minor carrier transfer.As a result, among all the materials herein studied, n-type Li4Ti5O12, MoO3, WO3, Bi2O3, LiFePO4, LiMnPO4 and p-type LiCoO2 are potential candidates for photorechargeable electrode application based on thermodynamic considerations. The p-type LiCoO2 though should be discarded for application in a full device since p-type cathode material would lead to a photodischarge process (i.e. lithium insertion under excited states).

However, this study

  photoluminescence study are underway in the group to determine the rate constant of charge injection from the N719* into LiFePO4 and FePO4 conduction bands.

  photo-recharged in about 15 hours under standard A.M. 1.5G conditions (capacity of ca.

  Cela pénalise le retour sur investissement financier (20-30 ans) par rapport à la durée de vie des panneaux photovoltaïque et des batteries Li-ion.Dans le but de réduire les coûts d'installation et de ferme germer de nouvelles innovations et de nouvelles filières scientifiques, la mise au point d'un seul système combinant à la fois la conversion et le stockage de l'énergie est hautement souhaitable. Toutefois, la réalisation d'un tel dispositif combinant conversion et stockage de l'énergie solaire à l'échelle du matériau et de la molécule est particulièrement complexe scientifiquement et requière donc de nombreuses recherches scientifiques. Une alternative à plus court terme est basée sur une hybridation des deux technologies qui consiste à développer un dispositif unique comprenant deux compartiments : un système de conversion PV, tel qu'une cellule solaire à colorant comme souvent proposé, et l'autre une batterie lithium-ion partageant ou non une même électrode. Cette approche est potentiellement moins coûteuse et plus facile à intégrer par rapport au couplage externe classique. La deuxième approche, avec des perspectives à plus long termes, consiste à adapter un assemblage moléculaire permettant de combiner la conversion de l'énergie solaire et son stockage chimique. C'est justement le cadre de ce travail de doctorat où l'approche consiste à créer des transferts d'ions à l'interface semi-conducteur mixte et électrolyte induit par l'excitation du niveau de la bande interdite du semi-conducteur. Ce transfert ionique résulte d'un transfert du porteur de charge minoritaire en phase homogène vers l'élément de transition. Un point bloquant de ce mécanisme provient de la nature purement surfacique de cette réaction qui est limitée par la longueur de Debye correspondant à la zone de déplétion du semi-conducteur. Afin de contourner cette limitation importante du mécanisme de photo-recharge, notre équipe a mis au point des nanocristaux d'anatase TiO2 de l'ordre de 4 à 5 nm de diamètre via une synthèse et thermolyse à température ambiante ou à basse température. Ces nano-cristaux ont permis de promouvoir considérablement le rapport surface / volume des particules. Une conséquence directe, ces particules permettent de maximiser la réactivité de surface vis-à-vis du volume, conduisant à un processus de photorecharge quantitatif. Une électrode mesoporeuse constituée par des nanoparticules de TiO2 peut être photorechargée intégralement (≈ 600 µAh/cm 2 ) en seulement 2 heures sous des conditions d'éclairage réaliste et standardisée (A.M. 1.5G, 100 mW/cm 2 ). Cette preuve de concept a démontré pour la première fois que l'excitation de la bande interdite peut induire une recharge complète d'une électrode de batterie. Cette recherche a permis d'ouvrir de nouvelles perspectives scientifiques et technologiques sur le développement de matériaux bifonctionnels combinant la conversion et le stockage de l'énergie solaire à l'échelle moléculaire.Des développements ultérieurs ont appelé à approfondir notre compréhension actuelle de ces processus de transfert de charge photoinduits afin de maximiser les performances en termes de stabilité et de temps de recharge. En plus, afin de passer d'une demi-cellule photorechargeable à une cellule complète, il est nécessaire de développer une contre-électrode appropriée permettant la collection des porteurs de charge majoritaires de l'électrode photo-active.Il a déjà été discuté dans la littérature que le potentiel des ions intercalés peut contribuer à l'énergie libre des électrodes provoquant un décalage des bords de la bande d'énergie, de la bande interdite, et affectant la longueur de Debye et le champ électrique de la zone de déplétion. Dans ce but, nous nous sommes concentrés dans cette thèse sur l'étude approfondie de l'évolution des propriétés optoélectroniques et de la dynamique des processus de transfert de charges ultra-rapides à l'interface électrode/électrolyte de l'anatase TiO2 lors de l'insertion du lithium. Pour cela, nous avons développé une cuvette en quartz à trois électrodes étanche permettant de coupler des expériences d'électrochimie et de spectroscopies optiques in situ / in operando du type absorption UV-Visible mais aussi de fluorescence à l'état stationnaire et résolue en temps. Nos travaux décrivent l'évolution des spectres de transmittance des films de TiO2 in operando à différents états de décharge. Nous avons établi que le changement de couleur passant du blanc au bleu foncé lors de l'insertion du lithium est associé d'une part à une augmentation de la valeur de la bande interdite par effet Burnstein-Moss, passant de 2,85 eV à 3,21 eV, et d'autre part liée à l'augmentation de l'absorbance dans la partie visible et dans le proche-infrarouge. Cette absorbance accrue est attribuée à l'absorption optique par des électrons libres dans la bande de conduction de TiO2 et ainsi qu'à la localisation des électrons dans les centres Ti 3+ occupant les orbitales métalliques 3d en raison de la taille nanométrique des particules. A ce titre, l'ajustement des spectres de transmittance à différentes profondeurs de décharge en utilisant le modèle de Drude montre que l'insertion de lithium dans TiO2 augmente par un facteur 20 la concentration en porteurs de charges, passant de 1,3•10 19 cm -3 à 2,9•10 20 cm -3 . Afin de déterminer les positions des bords de bandes et leurs évolutions lors de l'insertion du lithium, des mesures en trois électrodes par la méthode Mott-Schottky ont été entreprises en combinaison avec la spectroscopie d'absorption UV-visible. Il est établi dans ce travail que l'insertion de lithium dans la structure provoque un décalage important du potentiel de la bande plate vers des valeurs positives passant de 2,84 V à 3,36 V vs. Li + /Li. Cela se traduit par un décalage négatif de 0,5 eV de la bande de conduction. Ce processus est accompagné par une augmentation d'un ordre de grandeur de la concentration d'électrons passant de 2,1•10 19 cm -3 à 3,9•10 19 cm -3 sans aucun changement dans la nature des porteurs principaux. De plus, nos résultats soulignent systématiquement une forte influence de la nature protique ou aprotique de l'électrolyte sur la valeur de potentiel de bande plate et la concentration des porteurs. Enfin, en confrontant les mesures Mott-Schottky avec celles de spectroscopie d'absorption UV-Visible à l'état solide, nos résultats montrent une diminution de l'énergie de la bande de valence de 800 mV, passant de 5,7 V à 6,5 V et de la bande de conduction de 500 mV, passant de 2,8 V à 3,3 V (vs. Li + /Li) entre TiO2 et sa forme réduite Li0,6TiO2. Pour comprendre la dynamique des états excités de TiO2 et de Li0,6TiO2 dans différents types d'électrolytes, la spectroscopie de fluorescence à l'état stationnaire et en temps de vie à comptage simple photon a été un outil important et incontournable durant ce travail de doctorat. Les cartes d'excitation/émission de fluorescence sur un film de TiO2 à différentes profondeurs de décharge ont systématiquement montré la présence d'une bande principale pour laquelle le maximum se situe à 277 nm et 335 nm, ainsi qu'une seconde contribution moins intense et plus large vers 500 nm. Cette dernière se déplace de 498 à 515 nm lors de l'insertion de lithium. La première bande est attribuée aux phénomènes de fluorescence des ions solvatés Li + et PF6 -dans l'EC/DMC. L'origine de la deuxième bande moins intense est attribuée à la fluorescence des nanocristaux de TiO2 en équilibre avec l'électrolyte. Nos travaux mettent en évidence que les rendements de fluorescence de TiO2 et de l'électrolyte diminuent drastiquement lors de l'insertion du lithium dans des nanocristaux de TiO2. Ce phénomène suggère très fortement qu'une partie de l'énergie radiative de l'électrolyte est transférée à l'électrode par résonance de Förster (FRET) et que ce processus de transfert d'énergie au bénéfice de l'électrode est plus efficace lorsque le TiO2 est lithié en raison de ses meilleures propriétés d'absorption. En même temps, le rendement de fluorescence de l'électrode diminue à son état réduit. Ceci signifie que le processus radiatif de recombinaison dans le Li0.6TiO2 est désactivé comparé à TiO2. Ce phénomène est attendu en raison du transfert des trous d'électrons vers le Ti 3+ , mécanisme responsable de la photorecharge de l'électrode. Un mécanisme de type FRET est clairement réalisable pour trois raisons : (i) le rendement quantique (QY) de luminescence de l'électrolyte est largement supérieur à celui du TiO2 (QYLP30 = 7,7%, positionnement des bandes. Les résultats ont montré que parmi tous les matériaux, Li4Ti5O12, MoO3, WO3, Bi2O3, LiFePO4, LiMnPO4 de types n et LiCoO2 de type p sont des candidats potentiellement prometteurs. Toutefois, les résultats préliminaires de photo-rechargeabilité de ces matériaux ont toutefois montré que la recherche d'un matériau photorechargeable n'était pas aussi simple qu'attendu. En effet, de nombreux matériaux censés avoir une propriété de photorecharge n'ont pas montré les résultats escomptés pour des raisons encore peu claires. D'une part, la faible performance électrochimique peut être attribuée en partie à la faible conductivité électronique des matériaux purs. D'autre part, les particules utilisées ne sont pas systématiquement aussi divisées que celles du TiO2 et enfin il est aussi possible pour certains matériaux que les temps de vie des états excités sont trop courts pour permettre un quelconque transfert de charge sous illumination. Nous nous sommes attardés sur le cas de LiFePO4 qui ne montre pas d'effet de photo-recharge. Tout d'abord, la première approche a été de réduire la taille des particules et deuxièmement de greffer un colorant (N719) à la surface des particules afin d'accroître l'absorption de lumière. Nos résultats montrent qu'aucune amélioration significative des propriétés photoélectrochimiques n'a été observée dans les nanoparticules de LiFePO4. La présence du colorant a augmenté le potentiel de tension ouverte de l'électrode d'un apport mineur de 20 mV ce qui est loin d'être suffisant pour déclencher une réaction de photorécharge. Nous avons également rencontré des difficultés pour désinsérer le lithium de la structure en absence de carbone en raison de la nature isolante du matériau. Afin de mieux comprendre les mécanismes entravant la photorécharge de LiFePO4, nous avons déterminé les positions des bords de la bande de LiFePO4 et de sa forme chargée FePO4 obtenue par oxydation chimique. Notre étude montre que la valeur de la bande interdite diminue en fonction de la désinsertion du lithium, passant de 3,31 eV pour LiFePO4 à 2,73 eV pour la phase délithiée. La conductivité de type n exprimée par FePO4 est identique à celle de la phase lithiée. Le potentiel de bande plate et les valeurs de densité de donneur ont été déterminés et comparés à ceux de LiFePO4. Nous avons établis que les bords des bandes de valence et de conduction se déplacent vers le haut dans l'échelle d'énergie lors de la désinsertion du lithium. Le bord de la bande de conduction de FePO4 est supérieur de 0,06 eV à celui de LiFePO4 dans un électrolyte aqueux et de 0,25 eV dans un électrolyte aprotique non aqueux. Les décalages de la bande de valence et conduction sont respectivement de + 0,64 eV et +0,83 eV. Les positions des bords des bandes de Les résultats préliminaires qui ne sont pas rapportés dans ce manuscrit montrent que l'inclusion d'un piégeur non sélectif de radicaux dans l'électrolyte montrent une considerable amélioration des performances de l'électrode en capacité et en réversibilité sans moyen de carbone. D'autre part, l'ajout de cette molecule permet de déverrouiller la barrière de photorécharge. En conséquence, une électrode en Li7Ti5O12 peut être complètement photo-rechargée en 15 heures environ sous conditions d'éclairage standardisée (A.M. 1.5G, 100 mW/cm 2 ) avec une capacité d'environ 500 µAh/cm 2 . Un brevet pour la molécule incluse dans l'électrolyte est en cours de remplissage. Nous émettons l'hypothèse qu'une fonction de cette molécule est de piéger les radicaux fluorés avant que ceux-ci réagissent pour former du LiF à la surface des particules, même s'il est nécessaire d'approfondir des mécanismes liant cette molecule à la photo-recharge et la dégradation de l'électrolyte. Ce travail est pionnier dans un sujet complexe où l'on tente de combiner la fonction de conversion de l'énergie lumineuse avec celle du stockage électrochimique de cette énergie, ceci à l'échelle de l'électrode et de la molécule. Ce travaille apporte une brique complémentaire aux choix possibles des matériaux. Toutefois elle laisse aussi place à un certain nombre de zones d'ombres qui nécessitent de nouvelles recherches plus approfondies, en particulier dans le domaine de la chimie douce pour concevoir des matériaux nanostructurés d'une part et d'autre part sur la spectroscopie résolue en temps en absorption et en émission. Dans notre laboratoire, l'installation de la spectroscopie d'absorption transitoire en régime picoseconde avec caméra Streak est finalisée, ce qui doit permettre d'étudier plus en profondeur la cinétique du transfert de charge photogénéré dans les électrodes seules et au contact de l'électrolyte. D'autres outils sont également en cours d'installation comme la photo-luminescence en associant une caméra streak avec un laser Ti:Sa qui permettra de descendre à une résolution temporelle de l'ordre de 5 ps et à la spectroscopie de fluorescence par up conversion permettant d'obtenir une résolution temporelle de l'ordre de 200 fs. L'établissement de la chronologie des processus dans le système permettra de reconnaître les processus limitants et offrira la possibilité d'améliorer les performances et la stabilité des cellules en guidant les expérimentalistes vers des orientations précises. Comme il est mentionné dans le manuscrit, ce travail a souvent été confronté à un manque de données et d'observables de la littérature. Cependant, les résultats de ces travaux ouvrent à de nombreuses perspectives qui seront poursuivis au laboratoire pour continuer au développement de ces systèmes. Ce travail de thèse constitue une première démarche raisonnée dans la recherche 203 d'autres matériaux et d'une contre-électrode appropriée pour une photoanode de TiO2 qui permettra de collecter les électrons photo-induits. Ceci ouvrira la voie vers une photo-batterie complète et régénérative. Enfin, il est également important de souligner que ce travail de thèse est aussi profitable pour d'autres applications telles que la photocatalyse, le stockage d'informations optiques, les dispositifs électrochromiques, etc…

Table 1 .

 1 1. Overview of photorechargeable electric energy storage systems including the overall devicedischarge performance characteristics and solar-to-storage efficiencies when available.

							Solar-to-
		Conversion		Storage	Discharge performance	Light power	storage	Ref.
							efficiency
		Dye-sensitized		Carbon fibers	Ohmic load discharge	n/a	n/a	[28-28]
	Multifunctional electrode	TiO2 BODIPY-dye SrTiO3	Faradaic storage	WO3 PANI AQ 2-/AQ LaNi3.75Al1.24Hn	500 µAh/cm 2 n/a 143 mAh/g; 0.63 V 950 mAh/g	100 mW/cm n/a 100 mW/cm 100 mW/cm	n/a <1% n/a n/a	[29] [31] [33] [34]
		N719-dye		LiFePO4	340 mAh/g	100 mW/cm 0.06-0.08%	[35]

Table 2 .

 2 1. Flat band potential and charge carrier concentration of 5 nm nanocrystalline anatase TiO2 in 1 mol/L NaOH aqueous electrolyte and 1 mol/L TBAPF6 EC/DMC electrolyte.It also falls in the wide range of values reported in literature from 10 17 to 10 19 cm -3[5,[59][60][61][65][66][67][68][69]. Such a variation of values can stem from the fact that TiO2 is often not a single crystal, thus containing intraband gap surface states and can be easily doped by impurities with important consequences on carrier density as a result from its 3d 0 electronic configuration. flat-band potential values for measurements in water (-1.03 V vs. NHE at pH = 14) and in non-aqueous protic solvents (-1.2 V vs. NHE for MeOH and -0.9 V vs. NHE for EtOH), but significantly lower values in non-aqueous aprotic solvents (-1.8 V vs. NHE for DMF, -2.0 V vs. NHE for ACN, and -2.1 V vs. NHE for THF)[58]. Interestingly, Fox et al.

		Flat band potential	Donor
	Electrolyte	as measured	recalculated, V vs. NHE	density, cm -3
	1 mol/L NaOH aqueous	-0.87 V vs. SCE	-0.63 (pH = 14)	9.8•10 17
	1 mol/L TBAPF6 EC/DMC	2.99 V vs. Li QRE	-0.20	2.3•10 19

In a 1 mol/L NaOH aqueous solution (pH = 14), the flat band potential value was determined at -0.87 V (vs. SCE). It translates into -0.63 V vs. NHE at pH = 14. This result is in agreement with literature even though a quite large range of values is reported based on various techniques, such as Mott-Schottky measurements (-1.01 to 0.56 V vs. NHE at pH = 14)

[7,[58][59][60][61]

, photocurrent measurements (-0.07 V vs. NHE at pH = 14)

[62]

, electron affinity measurements (0.6 V vs. NHE)

[63]

, and calculations (-0.70 V vs. NHE at pH = 14)

[64]

. The calculation of the donor density considering a dielectric constant of 31

[3][4]

leads to a value of 9.8•10 17 cm -3 . because of the establishment of proton specific adsorption-desorption equilibrium in protic solvents and the absence of such in aprotic counterpart. Such an important dependence on the nature of the electrolyte is in agreement with the work reported by Fitzmaurice et al. who presented a comparable

Table 2 .

 2 2. Flat-band potential derived from Mott-Schottky measurements in aprotic electrolytes

	Solvent	ε solvent	Vfb, V vs. NHE	Salt	Ref.
	THF	7-8	-2.1	0.2 M TBAP	[58]
			-1.0	unknown	[70]
	DMF	37	-1.8	0.2 M TBAP	[58]
			-0.9	unknown	[70]
	ACN	37	-1.3	unknown	[70]
			-1.9	0.2 M TBAP	[58]
			-2.1	TBAP	[71]
			-2.0	TBAP	[72]
			-1.8	0.2 M TBAP	[73]
	ACN/IPA (4/1)	18-37	-1.7	0.2 M TBAP	[73]
	Me2SO	49	-0.7	unknown	[70]
	EC/DMC (1/1)	34-41	-0.2	1 M TBAPF6	This work

  2.21). This latter moves from 2.84 to 3.36 V vs.

	Li + /Li. To the best of our knowledge, there is no other experimental studies reported in the
	literature aiming at determining band movements in lithiated phase Li0.6TiO2 and in aprotic
	carbonate-based electrolyte. However, in good trend with our results, Fitzmaurice et al. reported
	a flat band potential increase by 1.14-1.23 V assigned to the lithium insertion into TiO2 based
	on absorbance measurements. This was performed in 0.2 mol/L TBAP ACN-based electrolyte
	containing 10 -3 mol/L of LiClO4 supporting electrolyte [58]. It has to be noted that lithium
	2.0x10
	2.4	2.6	2.8	3.0	3.2	3.4	3.6
	Lithium insertion also leads to an increase of carriers' density from 2.1•10	19 to 3.9•10 19 cm -3
	without any change in the nature of the main carriers. This trend is in good agreement with Van

cation in electrolyte is a potential determining ion due to its specific adsorption mechanism, as it is extensively reported in literature for protons

[74][75]

, DMPIm +

[76]

, and Li + ions

[58, 76- 78]

. One consequence of this surface adsorption mechanism is the shift of potential towards positive values when increasing the concentration of lithium in agreement with a Nernst law.

der Krol et al. who reported by Mott-Schottky experiments higher donor densities for lithiated phase, i.e. 8.7•10 17 cm -3 , compared to pristine TiO2, i.e. 1.4•10 17 cm -3 in 1.0 mol/L LiClO4 in

Table 2 .

 2 3. Lifetime and amplitude of the fluorescence decay after reconvolution from the IRF of TiO2 and Li0.6TiO2 in contact with argon.In a lithium-free electrolyte based on 1 mol/L TBAPF6 in EC/DMC electrolyte, corresponding to an ionic blocking interface that will prevent any photorecharge mechanism, the photoluminescence decay of TiO2 was again best fitted with two components leading to a lifetime values of 25 ps (82.1%) and 2.39 ns (17.9%) (Fig.2.32, table 2.

		TiO2		Li0.6TiO2
		Amplitude, %	Lifetime, ns	Amplitude, %	Lifetime, ns
	τ1	91.5	0.029	81.9	0.029
	τ2	8.5	2.19	18.1	1.96

Table 2 .

 2 4. Comparison of the lifetime values determined after reconvolution of the photoluminescence decay with the IRF for anatase TiO2 nanocrystals either in contact with argon or with the electrolyte composed of 1 mol/L TBAPF6 EC/DMC at open-circuit condition.

		TiO2 in contact with argon	TiO2 in 1 mol/L TBAPF6 EC/DMC
		Amplitude, %	Lifetime, ns	Amplitude, %	Lifetime, ns
	τ1	91.5	0.029	82.1	0.025
	τ2	8.5	2.19	17.9	2.39

Table 2 .

 2 5. Reconvolution results of the time-correlated single photon counting decays recorded onLi0.6TiO2 in EC/DMC and EMI.TFSI based electrolytes with or without Li + -containing salt.

		1 mol/L TBAPF6	1 mol/L LiPF6	1 mol/L TBAPF6	1 mol/L LiPF6
		EC/DMC	EC/DMC	EMITFSI	EMITFSI
		Amplitude,	Lifetime,	Amplitude,	Lifetime,	Amplitude,	Lifetime,	Amplitude,	Lifetime,
		%	ns	%	ns	%	ns	%	ns
	τ1	-	-	7.8	0.145	-	-	12.5	0.097
	τ2	77.9	1.34	38.5	1.33	30.6	1.49	58.4	1.21
	τ3	22.1	5.88	53.7	5.44	69.4	14.6	29.1	4.85

Table 3 .

 3 1. Visual aspects of the different classes of materials studied in this work.

		Transition metal oxides		Transition metal sulfides	Lithiated polyanions
	Li4Ti5O12	LiCoO2	MoO3	TiS2	LiFePO4
	WO3	CuO	Bi2O3	MoS2	LiMnPO4
	LiMn2O4	LiMn1/3Co1/3Ni1/3O2	WS2	LiVPO4F

Table 3 .

 3

	2. Flat band potential and charge carrier concentration of Li4Ti5O12 in 1 mol/L NaOH
	aqueous and 1 mol/L TBAPF6 EC/DMC electrolytes.
		Flat band potential
	Electrolyte	as measured	recalculated, V vs. NHE	Charge carrier concentration, cm -3
	1 mol/L NaOH in H2O	-0.99 V vs. SCE	-0.76 (pH = 14)	2.7•10 18
	1 mol/L TBAPF6 in EC/DMC	1.86 V vs. Li QRE	-1.34	6.9•10 19

Table 3 .

 3

	3. Results after reconvolution of the time-correlated single photon counting decay recorded on
	Li4Ti5O12 powder using stretched exponential function (375 nm excitation, 728 nm emission).
	Component	Pre-exponential factor 𝐵 𝑖	Amplitude, %	Lifetime, ns
	τ1	0.114	6.9	0.10
	τ2	0.035	61.0	3.00
	τ3	0.004	32.9	27.87

Table 3 .

 3

	4. Flat band potential and charge carrier concentration of LiCoO2 in 0.5 mol/L TBANO3
	aqueous electrolyte and 1 mol/L TBAPF6 EC/DMC electrolyte.
			Flat band potential
	Electrolyte		Recalculated	Charge carrier concentration,
		as measured	V vs. NHE	V vs. NHE	cm -3
				at pH=14
	0.5 mol/L TBANO3 in H2O -0.23 V vs. SCE 0.01 (at pH=1.65)	-0.72	4.9•10 19
	1 mol/L TBAPF6 in EC/DMC	4.44 V vs. Li QRE	1.24	-	2.6•10 20

Table 3 .

 3

			Flat band potential
			recalculated	Charge carrier concentration,
	Electrolyte	as measured	V vs. NHE	V vs. NHE at	cm -3
				pH=14
	0.5 mol/L TBANO3 in H2O -0.01 V vs. SCE 0.23 (pH=8.84)	-0.08	4.1•10 18
	1 mol/L TBAPF6 in EC/DMC	2.94 V vs. Li QRE	-0.26	-	1.2•10 20

5. Flat band potential and charge carrier concentration of MoO3 in 0.5 mol/L TBANO3 aqueous electrolyte and in 1 mol/L TBAPF6 EC/DMC electrolyte.

Table 3 .

 3

	6. Flat band potential and charge carrier concentration of WO3 in 0.5 mol/L TBANO3 aqueous
	electrolyte and 1 mol/L TBAPF6 EC/DMC electrolyte.
			Flat band potential		Charge carrier
					concentration,
	Electrolyte		recalculated	cm -3
		as measured	V vs. NHE	V vs. NHE
				at pH=14
	0.5 mol/L TBANO3 in H2O	0.03 V vs. SCE 0.27 (pH=1.65)	-0.46	1.4•10 18
	1 mol/L TBAPF6 in EC/DMC	3.22 V vs. Li QRE	0.03	-	2.8•10 19

Table 3 .

 3

	7. Flat band potential and charge carrier concentration of CuO in 1 mol/L TBAPF6 EC/DMC
		electrolyte.
		Flat band potential	Charge carrier
	Electrolyte			concentration,
		as measured	Recalculated	cm -3
	1 mol/L TBAPF6 in EC/DMC	3.80 V vs. Li QRE	0.60 V vs. NHE	2.7•10 19

Table 3 .

 3

	8. Flat band potential and charge carrier concentration of Bi2O3 in 1 mol/L NaOH aqueous
	electrolyte and 1 mol/L TBAPF6 EC/DMC electrolyte.
		Flat band potential
	Electrolyte	as measured	recalculated	Charge carrier concentration, cm -3
			V vs. NHE
	1 mol/L NaOH in H2O	-0.90 V vs. SCE	-0.66 (pH=14)	9.9•10 17
	1 mol/L TBAPF6 in EC/DMC	2.42 V vs. Li QRE	-0.77	3.3•10 19

Table 3 .

 3

	9. Flat band potential and charge carrier concentration of TiS2 in 0.5 mol/L TBANO3 aqueous
	electrolyte and 1 mol/L TBAPF6 EC/DMC electrolyte.
		Flat band potential
					Charge
	Electrolyte	as measured	Recalculated V vs. NHE V vs. NHE	carrier concentration, cm -3
				at pH=14
	0.5 mol/L TBANO3 in H2O	-0.24 V vs. SCE	0.00 (pH=3.45)	-0.62	4.2•10 18
	1 mol/L TBAPF6 in EC/DMC	2.01 V vs. Li QRE	-1.18	-	6.9•10 19

Table 3 .

 3

		Flat band potential
				Charge carrier
	Electrolyte		recalculated	concentration,
		as measured		cm -3
			V vs. NHE
	1 mol/L NaOH in H2O	-0.75 V vs. SCE	-0.51 (pH=14)	3.0•10 18
	1 mol/L TBAPF6 in EC/DMC	4.20 V vs. Li QRE	1.01	4.0•10 19

10

. Flat band potential and charge carrier concentration of MoS2 in 1 mol/L NaOH aqueous electrolyte and 1 mol/L TBAPF6 EC/DMC electrolyte.

Table 3 .

 3 11. Flat band potential and charge carrier concentration of WS2 in 1 mol/L NaOH aqueous and 1 mol/L TBAPF6 EC/DMC-based electrolytes.

		Flat band potential	Charge carrier
	Electrolyte	as measured	recalculated V vs. NHE	concentration, cm -3
	1 mol/L NaOH in H2O	-0.72 V vs. SCE	-0.48 (pH=14)	4.2•10 18
	1 mol/L TBAPF6 in EC/DMC	3.34 V vs. Li QRE	0.15	1.8•10 20

Table 3 .

 3 12. Flat band potential and charge carrier concentration of LiFePO4 in 0.5 mol/L TBANO3 aqueous solution and 1 mol/L TBAPF6 EC/DMC electrolyte.

			Flat band potential
	Electrolyte		Recalculated	Charge carrier concentration,
		as measured	V vs. NHE	V vs. NHE at	cm -3
				pH=14
	0.5 mol/L TBANO3 in H2O -0.04 V vs. SCE 0.20 (pH=1.65)	-0.53	1.9•10 19
	1 mol/L TBAPF6 in EC/DMC	2.62 V vs. Li QRE	-0.58	-	3.9•10 20

Table 3 .

 3 13. Flat band potential and charge carrier concentration of LiMnPO4 in 1 mol/L TBAPF6EC/DMC electrolyte.

		Flat band potential	Charge carrier
	Electrolyte			concentration,
		as measured	recalculated	cm -3
	1 mol/L TBAPF6 in EC/DMC	2.07 V vs. Li QRE	-1.12 V vs. NHE	2.2•10 20

Table 3 .

 3 14. Potential of the photogenerated minor charge carriers (h + for n-type and e -for p-type SC)compared to the redox potential corresponding to lithium insertion/deinsertion (V vs. Li + /Li) for the transition metal oxides, transition metal sulfideas, and lithiated polyanions in study.

	Material	Potential of minority charge carrier, V vs. Li + /Li	Li + insertion / deinsertion potential, V vs. Li + /Li	Overportential for charge transfer (V)	[Reference]
	n-Li 4 Ti 5 O 12	5.55	1.55	0.2-0.25	[202]
	p-LiCoO 2	4.25	3.8	0.1	[39, 203]
	n-MoO 3	5.75	2.2-2.8	0.5	[204]
	n-WO 3	5.55	1.4-2.3	0.8	[205]
	p-CuO	2.45	1.3	1.2	[206]
	n-Bi 2 O 3	4.45	0.64-0.75	0.6	[207]
	p-MoS2	3.55	1.1	1.0	[208]
	p-WS2	2.35	1.4	0.9	[209]
	n-LiFePO4	5.75	3.4	0.2	[176]
	n-LiMnPO4	5.55	4.0	0.2	[210]

Table 3 .

 3 15. Flat band potential and charge carrier concentration of FePO4 in 0.5 mol/L TBANO3 aqueous electrolyte and 1 mol/L TBAPF6 EC/DMC electrolyte.

		Flat band potential
	Electrolyte		Recalculated	Charge carrier concentration,
		as measured	V vs. NHE	V vs. NHE	cm -3
				at pH=14
	0.5 mol/L TBANO3 in H2O -0.33 V vs. SCE	-0.09 (pH=6.1)	-0.56	1.2•10 18
	1 mol/L TBAPF6 in EC/DMC	2.36 V vs. Li QRE	-0.83	-	1.3•10 19

Table 3 .

 3 2.63 ns, are comparable with the first two components measured in contact with argon, i.e. 0.10 ns and 3.00 ns. The third very long component of 27.8 ns disappeared in contact with the electrolyte. The main difference comes from a redistribution of the lifetime values towards the fast component when in contact with the electrolyte (73.3 % in electrolyte vs. 6.9 % in argon).16. Reconvolution results of the time-correlated single photon counting decays recorded onLi4Ti5O12 powder and Li4Ti5O12 electrode in 1 mol/L LiPF6 EC/DMC electrolyte.

	This confirms the quenching of Li4Ti5O12 luminescence in the presence of LiPF6 EC/DMC
	electrolyte.
		Li4Ti5O12 powder		Li4Ti5O12 electrode in 1 mol/L LiPF6 EC/DMC
	Pre-exponential	Amplitude,	Lifetime,	Pre-exponential	Amplitude,	Lifetime,
		factor 𝐵 𝑖	%	ns	factor 𝐵 𝑖	%	ns
	τ1	0.572	6.9	0.10	0.395	73.3	0.09
	τ2	0.578	61.0	3.00	0.621	26.7	2.63
	τ3	2.752	32.1	27.87	-	-	-
	The steady-state excitation/emission map of Li4Ti5O12 electrode in 1 mol/L LiPF6 EC/DMC
	electrolyte recorded over 250-700 nm excitation and 265-800 nm emission wavelengths is
	reported in figure 3.

Table 3 .

 3 17 gathers the information on the conductivity type, flat band potential and charge carrier concentration values. Regardless of the discharge state, the electrode remains n-type.

	For Li4Ti5O12, the flat band potential is 1.91 V vs. Li QRE corresponding to -1.28 V vs. NHE.
	Upon lithium insertion, the latter increases significantly to 3.60 V vs. Li QRE (+ 0.41 V vs.
	NHE). This means that upon lithium insertion, forming Li7Ti5O12, the conduction band moves
	down by + 1.69 V in energy, which is hardly comprehensible based on the low structural
	modifications made by the lithium insertion. After electrode recharge, the flat band potential
	remains very close to Li7Ti5O12, at 3.42 V vs. Li QRE (+ 0.23 V vs. NHE). This value cannot
	be solely explained by the first cycle irreversibility (coulombic efficiency ca. 65%). Most part
	of this irreversibility stems from the formation of LiF SEI layer. Based on these experiments,

Table 3 .

 3 17. Conductivity type, flat band potential, and charge carrier concentration of Li4Ti5O12-based electrode measured at 200 kHz in 1 mol/L TBAPF6 EC/DMC electrolyte under argon at initial state (OCV 3.4 V vs. Li + /Li), discharged state down to 1.2 V vs. Li + /Li and charge state after one cycle (2.6 V vs. Li + /Li).

		Li4Ti5O12	Li7Ti5O12	Li4Ti5O12 1 cycle
	Conductivity type	n	n	n
	Li QRE	1.91	3.60	3.42
	Flat band potential, V vs.
	NHE	-1.28	0.41	0.23
	Charge carrier concentration, cm -3	2.9•10 19	3.4•10 19	2.4•10 19

) (a) (b)

1M TBAPF 6 EC/DMC 0.5M TBANO 3 aq.

WS2

The full-pattern matching refinement of WS2 x-ray diffractogram shows an important preferential orientation along (00l) planes. All observed diffraction peaks can be indexed into the hexagonal lattice of P63/mmc space group (Fig. 3.48 (a)). The refined lattice cell parameters are a = b = 3.154 (5) Å and c = 12.362 (12) Å (Vcell = 106 Å 3 ). These values are very close to the isostructural MoS2 previously investigated, and are in good agreement with the values reported in literature [157], [158]. The SEM micrograph shows nanoplates morphology with a particle size ranging from 200 nm to 6 micrometers (Fig. 3. 48 (b-c)). This morphology is responsible for the preferential orientation observed. The HR-TEM study confirms that there are no amorphous regions on the sample, especially at the surface of the particles (Fig. 3.48 (d)). The SAED pattern shows single crystal electron diffraction with an indexation in agreement with XRD (Fig. 3.48 (e)). WS2 particles show a very broad reflectance within the NIR light portion (Fig. 3.49 (a)).

Assuming an indirect allowed transition, a band gap value of 0.82 eV is determined from the

3.4.Fundamental evaluation of the possibility of Li + photo-insertion / photo-deinsertion reaction.

In the late 1970s, Gerisher et al. have established that during the photo-induced charge separation in the depletion layer of a mixed semiconductor, the main charge carriers migrate into the bulk of the material while the minor carriers migrate towards the surface [195][196].

The latter can subsequently trigger an interfacial ion-transfer at the electrode / electrolyte interface induced by a redox reaction between the surface carriers and the transition metal center in the electrode [197][198][199][200][201]. Based on this pioneering work, we evaluated the possibility of the materials to be employed as a photorechargeable electrode relying on the information gathered in the database. To be used as a photorechargeable electrode, material needs to have an energy of the minor carriers adequate to the redox potential of the transition metal to allow its transfer.

The figure 3.65 gathers the positions of the valence and conduction bands for different materials investigated in this thesis with respect to the type of semi-conduction and the redox energy of the transition metal. This small improvement can be originating from a minor enhancement of electrode's conductivity under illumination. This can also explain why there is a ca. 40 mV smaller anodic/cathodic peak separation energy (ΔEP) for scan rates greater than 1 mV/s, indicating lower kinetic issues for lithium insertion/deinsertion process (Fig. 3.69 (a)). This is also consistent with the global cell resistance that we determined from the linear portion of the current rise at low overpotentials, which shows lower but not drastic ohmic drop under illumination (Fig. 3.69 (b)).

Conclusions

In this chapter, we focused on the development of a database gathering the optoelectronic properties of most common insertion materials used in lithium-ion batteries. The aim of this study was to decipher the materials, which can have perspective application in photorechargeable electrochemical device. Three groups of material were studied: transition metal oxides (Li4Ti5O12, LiCoO2, MoO3, WO3, CuO, Bi2O3, LiMn2O4, LiMn1/3Ni1/3Co1/3O2), transition metal sulfides (TiS2, MoS2, WS2), and lithiated polyanions (LiFePO4, LiMnPO4, LiVPO4F).

All materials were characterized before studying the optoelectronic properties by using x-ray diffraction, scanning electron microscopy, and transmission electron microscopy techniques to ensure that materials are phase pure, lattice cell parameters to detect punctual defects, and evaluate size and morphology of the particles.

Using UV-visible absorption spectroscopy measurements in diffuse reflectance mode, the band gap values were determined. For a set of materials herein studied, we did not find optical bandgap likely owing to a too narrow band gap value below 0.7 eV (LiMn2O4 and LiMn1/3Co1/3Ni1/3O2) or to the semi-metallic nature of the material (TiS2). We have identified that for all the materials, except Li4Ti5O12, that the fluorescence cannot be measured in the range from 250 to 800 nm. This can be explained either by a band gap which is not in the visible range (Eg < 1.6 eV), an emission in the NIR region, or simply by materials' nature, which leads to exclusively non-radiative excited state relaxation.

By means of Mott-Schottky electrochemical impedance spectroscopy, we measured the type of conductivity, flat-band potential, and charge carrier concentration. In this chapter, three interesting phenomena are highlighted. First, for certain materials, i.e. LiCoO2, WS2, and MoS2, a change of semi-conduction type found depending on the nature of the electrolyte. More specifically, these materials are behaving as n-type semiconductor in aqueous electrolyte and turn p-type in non-aqueous aprotic EC/DMC-based electrolyte. This change is assigned to the weak degenerated nature of these semiconductors. Second, except the case of LiFePO4, all the materials have band edge positions that also depend on the nature of the solvent. Third, the charge carrier concentration is systematically higher in non-aqueous aprotic solution by one to two orders of magnitude, which means that materials can be "doped" depending on the electrolyte nature in contact through equilibrium process of Fermi level equalization.

General conclusions

This work is dedicated to the study of mixed electronic/ionic semi-conducting materials that can be capable of combining sunlight energy conversion reaction and its energy storage at the electrode and molecular levels. The objective of this thesis work is the exploration of other materials than the anatase TiO2. Our strategy, against simply blindly screening materials, was focused on deepening the current knowledge on the light-induced charge transfer through the semiconductor/electrolyte junction following up the example of the nanoparticles of anatase TiO2 and to develop a database that can be used to decipher which materials can have potentialities for the application in photo-rechargeable electrode and in full devices that require a perfect band alignments between the two electrodes.

In the first part of this thesis, we developed an airtight photoelectrochemical cell allowing to monitor the change of the optoelectronic properties of TiO2 upon lithium insertion. We have established that few factors stand behind the color changeover from white in TiO2 to dark blue in Li0.6TiO2. On the one hand, there is a 0.3 eV increase in the bandgap value upon lithium insertion following Burnstein-Moss effect. On the other hand, the blue coloration is connected to the enhancement in absorption within the visible and near-infrared region of the solar spectrum as a result of the enrichment of free carriers in the conduction band of LixTiO2 and electrons localized in Ti 3+ centers occupying the 3d metal orbitals. The in situ Mott-Schottky measurements showed a positive shift in potential of 0.5 V of the conduction band. Our study also revealed that the charge carrier concentration increases by circa one order of magnitude upon lithium insertion into the structure.

We have established the ultrafast nature of the direct recombination, i.e. 29 ps, for both TiO2 and Li0.6TiO2. We demonstrated that even though the majority of the photoluminescence stems from the direct recombination, a significant part of it belongs to the recombination through carrier traps associated with oxygen vacancies and surface states. This trap-mediated fluorescence was found to increase when electrode was put in contact with electrolytes.

Depending on the electrolyte composition, the band gap, excitation/emission spectrum and photoluminescence dynamics can vary to a large extent. This phenomenon was assigned to an experimental demonstration of a Stark effect taking place in the nanocrystals as a result from the specific adsorption of lithium cation from the electrolyte. Today, Stark effect in nanocrystals is considered as being a beneficial effect for the electrode's photorecharge. We also demonstrated elsewhere that Stark effect is in part or in whole responsible for two phase vs.

single phase transition during lithium insertion by downsizing the particles towards nanoscale.

Comparing the fluorescence decays in lithium ion-blocking vs. ion-conducting interface, we support that the half-time of the hole transfer to Ti 3+ is in the scale of 100 ps, thus entering in direct completion with the deactivation of the excited states in crystals.

The second important part of this work was dedicated to the creation of a database describing the optoelectronic properties of common insertion materials. Three groups were studied: transition metal oxides (Li4Ti5O12, LiCoO2, MoO3, WO3, CuO, Bi2O3, LiMn2O4, LiMn1/3Ni1/3Co1/3O2), transition metal sulfides (TiS2, MoS2, WS2), and lithiated polyanions (LiFePO4, LiMnPO4, LiVPO4F). This work reports for the first time that for some materials, i.e.

LiCoO2, MoS2, and WS2, the type of conductivity can change of sign depending on the nature of the electrolyte in contact, i.e. from n-type in an aqueous (protic) electrolyte to a p-type in a non-aqueous aprotic EC/DMC-based electrolyte. We also show that the flat band potential can be drastically modified by the electrolyte up to 1.4 eV except for LiFePO4, and that the main carrier density is always increased by at least one order of magnitude in EC/DMC electrolyte.

Comparing the minor charge carrier energy to the redox energy of the transition metal for the lithium insertion or deinsertion reaction, we established from this database that p-CuO, p-MoS2 and p-WS2 are materials which should not work for thermodynamic reasons. The case of LiCoO2 is somewhat different. Although this material can be of interest from a fundamental point of view, this latter is not interesting for the application since its high redox potential associated to its p-type semi-conduction will lead to a photo-dischargeable electrode. By contrast, the photoelectrochemical recharge of the electrode is thermodynamically realizable in n-Li4Ti5O12, n-MoO3, n-WO3, n-Bi2O3, n-LiFePO4, and n-LiFePO4.

Unfortunately, the photoelectrochemical study of some specifically chosen materials underlines that reaching the photo-rechargeable mechanism is not as straightforward as anticipated. None We focused to extend this study in the case of two materials: LiFePO4 and Li4Ti5O12 Les valeurs de durée de vie déterminées pour Li4Ti5O12 en contact avec l'électrolyte, i.e. 0,091 ns et 2,63 ns, sont comparables aux deux premières composantes de durée de vie du déclin de fluorescence de la poudre de Li4Ti5O12, soit 0,102 ns et 3,00 ns. La troisième composante de durée de vie observée sur la poudre a disparu lorsque le matériau a été mesuré dans l'électrolyte. La redistribution des valeurs de temps de vie vers la composante rapide (73,3% pour l'électrode dans l'électrolyte contre 6,9% pour la poudre) a confirmé l'extinction de la luminescence de Li4Ti5O12 en présence d'électrolyte. Ceci résulte soit du transfert de charges photo-induites vers l'électrolyte conduisant à sa dégradation et/ou par une fraction plus importante d'un mode de désactivation de l'état excité de manière non radiative. 208

Appendix

Résumé

Le problème de la nature intermittente de l'énergie solaire est souvent résolu par le couplage traditionnel des unités PV et batterie. Notre approche plus fondamentale vise le développement de matériaux capables de combiner la conversion et le stockage de l'énergie solaire au niveau moléculaire. Les nanocristaux de TiO2 d'anatase à 5 nm ont été synthétisés dans notre groupe, ce qui a permis une réaction quantitative de photorecharge par la seule contribution de l'illumination. Nous présentons ici une étude de l'évolution des propriétés optoélectroniques et de la dynamique du transfert de charge dans une électrode de TiO2 à l'aide d'expériences in situ / in operando effectuées pendant le fonctionnement de la batterie (spectroscopies de UV-visible, Mott-Schottky, fluorescence). L'augmentation de la valeur de la bande interdite et de l'absorbance a été observée lors de l'insertion du lithium dans TiO2. Un décalage négatif en énergie de la bande de conduction indique un potentiel plus oxydant des trous photogénérés dans le Li0.6TiO2 par rapport au TiO2 initial. En analysant les processus de recombinaison dans Li0.6TiO2, nous avons établi une compétition entre les processus ultra-rapides (gamme ps) de recombinaison directe et de transfert de charge vers Ti 3+ dans Li0.6TiO2, ce qui limite potentiellement le rendement de la réaction de photorécharge. Cette étude a été étendue à d'autres matériaux d'insertion généralement utilisés dans les batteries lithium-ion (Li4Ti5O12, LiCoO2, LiFePO4, MoO3, etc.). Les positions de bord de bande, la bande interdite, le type de porteurs de charge et leur concentration ont été mesurées et rassemblées dans une base de données. Basé sur ces résultats, la possibilité de photorécharge induite par la lumière a été évaluée et les premiers résultats discutés.

Mots-clés: photo-électrode, battery photo-rechargeable, photoelectrochimie, matériaux d'insertion, propriétés optoélectroniques, Mott-Schottky, photoluminescence, bande interdite

Abstract

The problem of intermittent nature of solar energy is often addressed by the traditional coupling of the PV and battery units. Our more fundamental approach targets the development of materials able to combine solar energy conversion and storage at the molecular level. The 5 nm anatase TiO2 nanocrystals were synthesized in our group affording a quantitative photorecharge reaction by a sole contribution of illumination. Here, we present a study of the evolution of the optoelectronic properties and dynamics of charge transfer in TiO2 electrode using in situ / in operando experiments performed during the battery functioning (UV-visible, Mott-Schottky, fluorescence spectroscopy). The increase of the bandgap value and the rise of absorbance are observed upon lithium insertion into TiO2. A negative shift of the conduction band indicates a more oxidizing potential of the photogenerated holes in Li0.6TiO2 compared to TiO2. By analysis of the recombination processes in TiO2 upon lithium insertion, we established a competition of the ultra-fast (ps range) processes of direct recombination and charge transfer towards Ti 3+ in Li0.6TiO2, potentially limiting the yield of the photorecharge reaction. This study was extended to other insertion materials typically used in lithium-ion batteries (Li4Ti5O12, LiCoO2, LiFePO4, MoO3, etc.). The measured band edge positions, band gap, charge carrier type and concentration were gathered into a database, based on which the fundamental evaluation of the possibility of the light-induced photorecharge was conducted. The first results of the photoelectrochemical study of chosen materials are also discussed. Key-words: photoelectrode, photorechargeable battery, photoelectrochemistry, insertion materials, optoelectronic properties, Mott-Schottky, photoluminescence, band gap