Low temperature synthesis of γ-LixMnO2 powders in ethylene glycol

International Journal of Inorganic Materials 2 (2000) 389–396

Low temperature synthesis of g-Li x MnO 2 powders in ethylene glycol ´ D. Larcher*, B. Gerand, J.-M. Tarascon ´ ´ et Chimie des Solides, UMR 6007, Universite´ de Picardie Jules Verne, 80039 Amiens Cedex, France Laboratoire de Reactivite Accepted 27 June 2000

Abstract The reaction of Electrolytic Manganese Dioxide (EMD g-MnO 2 ) with anhydrous LiOH in ethylene glycol (EG) was studied. It was found that this reaction proceeds in three steps while the reaction temperature increases: (1) chemical insertion of lithium in g-MnO 2 structure leading to g-Li x MnO 2 phase, (2) progressive dissolution of this reduced phase and (3) precipitation of a Li-free organo-metallic compound Mn(OCH 2 CH 2 O) with lamellar structure (Mn-EG). If the reaction is performed at moderate temperature (T ,1208C), we can limit the dissolution reaction and completely avoid the precipitation of the organo-metallic phase. The x value in the resulting Li x MnO 2 phases is found to be mainly linked to the temperature of the reaction medium. Although a maximum value of x50.35 can be obtained at 308C, Li 0.84 MnO 2 can be prepared at 1208C. We showed that the Lix MnO 2 formation mechanism consists of a direct lithium insertion in the structure of MnO 2 concomitant with Mn reduction by the organic medium. By changing the medium cationic content of the medium, we were able to prepare various reduced MnO 2 compositions, such as H x MnO 2 , (Li,H)x MnO 2 , (Na,H)x MnO 2 , (K,H)x MnO 2 . These Lix MnO 2 phases, when tested in half Li-cells, show poor cyclability but large first charge capacity (260 mAh / g for x50.84) that takes place at significantly higher voltage than observed for electrochemically inserted EMD. This finding was related to the presence of chemically inserted protons besides lithium ions.  2000 Elsevier Science Ltd. All rights reserved. Keywords: Lithium manganese oxides; Ethylene glycol; Chemical insertion; Polyol; Cathode materials

1. Introduction Typically, the electrode materials used in a lithium battery are metallic lithium (anode) and an intercalation material (cathode) in which the lithium is inserted (discharge) and extracted (charge). For safety purposes, the metallic lithium is commonly replaced by another intercalation phase such as graphite, leading to the so-called ‘rocking-chair’ technology. In this technology, the positive electrode is the only source of lithium. Actually, LiNiO 2 , LiCoO 2 and LiMn 2 O 4 are the most promising candidates for cathode, and LiCoO 2 / C cells are the only ones up to now to be commercialized. Since one Li can be extracted per 3d metal in LiNiO 2 and LiCoO 2 , a theoretical capacity of 280 mAh / g is expected for these two oxides. In fact, for stability reasons of the partially Li x NiO 2 and Li x CoO 2 delithiated phases, only 0.5 Li per metal can be practically extracted from them. In contrast, LiMn 2 O 4 can be fully delithiated (0.5Li / Mn) without any drastic effect on its *Corresponding author. E-mail address: [email protected] (D. Larcher).

cycling efficiency. Finally, these three phases have similar practical reversible capacities of about 140 mAh / g but in terms of toxicity, cost and recyclability, the Mn based oxides have greater advantages compared to the Ni- and Co-based ones. Then, in order to improve the intrinsic capacity of the batteries (e.g. the initial total lithium content) and / or to reduce their cost, numerous studies have been focused on the synthesis of new Mn-based cathode materials or on new synthetic processes involving much lower temperatures than the classical solid state reactions (8008C) used for the spinel LiMn 2 O 4 . Among the several possible alternatives, the low temperature synthesis of LiMnO 2 phases showing high reversible capacity and good performances upon cycling is of great interest. Ohzuku et al. proposed a new Low Temperature (LT) synthesis (T , 4508C) of orthorhombic LiMnO 2 consisting of the heating under argon a mixture of g-MnOOH and LiOH?H 2 O [1]. Reimers et al. synthesized LiMnO 2 by Li / H cationic exchange from MnOOH [2]. Recently, Gummow et al. used higher temperatures (T|6008C) to prepare the orthorhombic form of LiMnO 2 from g-MnO 2 and LiOH with

1466-6049 / 00 / $ – see front matter  2000 Elsevier Science Ltd. All rights reserved. PII: S1466-6049( 00 )00091-X

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additional carbon acting as reducing agent [3], and Rossouw et al. [4] and Thackeray et al. [5] prepared Li x MnO 2 phases by reacting ramsdellite-MnO 2 with reducing LiI in acetonitrile at 808C. Also well documented is the stabilization / lithiation of g-MnO 2 by reaction at 300–4008C with lithium salt (LiOH, LiNO 3 ) leading to g-Li x MnO 2 (referred to as CDMO) in which x ranges from 0.2 to 0.3 [6–8]. Whatever the process used, the resulting LiMnO 2 materials do not show good cycling efficiency but LTmade phases exhibit higher capacities than the HT-made ones. It’s interesting to note that very few [2,4,5] of the methods proposed to prepared LiMnO 2 can be considered as genuine low temperature or even ‘Chimie Douce’ alternatives. Based on these observations, we decided to search for a different and simple preparation way for electrochemically active g-Li x MnO 2 phases. The investigated route involves low temperature (,1208C) and conditions / reactants not prohibitive on a manufacturing scale. More specifically, we will herein present and study the use of an organic liquid medium as solvent / reducing agent (ethylene glycol) in presence of lithium hydroxide (LiOH) for the preparation of Li x MnO 2 from g-MnO 2 . The use of reducing organic medium has been largely used to insert protons in the open structure of g-MnO 2 [9]. To our knowledge, no attempt has been made to take advantage of this reducing effect in view of a controlled insertion of other cationic species.

2. Experimental section All syntheses were made in a glass vessel equipped with both a refrigerating unit and a mechanical stirring unit to ensure constant reflux and constant mechanical agitation, respectively. Anhydrous lithium hydroxide (Merck.96%) and g-MnO 2 (IBA[15, Japan Metals & Chemical Co, Ltd) were successively added to cold ethylene glycol (Prolabo, Reagent grade) and the heating rate was tuned to 18 / mn. At room temperature, LiOH is soluble in EG. The reaction medium temperature was monitored by means of a temperature controller and the final temperature was selected between 308C and 1968C. This temperature is maintained from several hours to some days and the cooling rate was set at 18 / mn. The solid and liquid phases were separated by vacuum filtration through a paper filter sheet, and the recovered solid phase was washed with acetone, and dried at 558C under air for 1 h. Phase identification was carried out by means of X-rays diffraction (Philips PW 1729) using CuK a radiation ( l 5 ˚ and a diffracted beam monochromator. The 1.5418 A) counting time is 10 s. Infrared spectra of solid phases were obtained at room temperature with a Nicolet 510 FT-IR spectrophotometer. For that, a few milligrams of powders were mixed with 200 mg of KBr dried at 1258C and pelletized at 8t / cm 2 . The lithium and manganese contents in the solid and organic liquid phases were determined by

Atomic Absorption (AA) using a Perkin Elmer 3030 Spectrophotometer. Solid materials were dissolved in a H 2 O / HNO 3 / H 2 O 2 (90 / 5 / 5) solution. Scanning electron microscopy was performed with a Field Emission Gun microscope Philips FEG XL30 equipped with an EDX analyzer (Link Isis-Oxford Instruments). EDS was used to evaluate the Na / Mn and K / Mn atomic ratios. The manganese mean oxidation state (Mn o.s.) was obtained by ferrous sulfate method that consists of the reduction of the Mn 21y to Mn 21 by ferrous ammonium sulfate and back titration by KMnO 4 of the non-consumed Fe 21 . The calculated amount of electron consumed and the Mn weight percentage in the solid phase determined by AA are used to evaluate the mean oxidation state of the manganese. This method leads to an accuracy evaluated to 60.02 [10]. Electrochemical tests were made in SwagelokE cells. The Lix MnO 2 powders were dispersed into a plastic according to Bellcore PLiONE technology [11]. The positive electrode plastic layer was made by mixing 56 parts of lithium metal oxide powder and 6 parts in weight of carbon black (SP), and by intimately dispersing this mixture in a binder matrix of 15 parts of an 88:12 vinylidene fluoride: hexafluoropropylene (PVDF:HFP) copolymer (2801) plasticized with 23 parts of dibutylphtalate (DBP), to which the required amount of acetone was added to obtain a viscous paste. The resulting paste was spread on a glass substrate, the acetone evaporated and the resulting dried plastic layer cut into 1 cm 2 discs. Prior to being used as the positive electrode in the Swagelok cell, DBP was extracted by three successive treatments in dried ether. Two sheets of Wattman GF / D borosilicate glass fiber, soaked with 1 M LiPF 6 in 2EC / DMC electrolyte, electrically separated DBP extracted cathode discs from the lithium metal disc negative electrode. Cells were tested on a MacPile Macintosh controlled cycler (Bio-Logic Co., Claix, France) in a galvanostatic mode with current charge / discharge stability better than 1%. Cycling data were collected at 258C, between 2.0 and 4.0–4.8 V and at a C / 5 rate. Specific surface areas were measured with a Micromeritics (Norcross, GA, USA) Gemini 2375 system according to the BET multi-point method [12] by nitrogen physisorption at 77 K. For typical measurements, prior to measuring the free space by means of helium gas, the samples were dried under argon flow at 1108C during 1 h.

3. Results and discussion

3.1. Characterization of the starting EMD powder A huge number of works has been devoted to the crystal structure and XRD patterns study of the different MnO 2 polymorphs, and especially of g-MnO 2 . The structure of

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the latter is interpreted as intergrowth of rutile-type pyrolusite (b-MnO 2 ) in a ramsdellite-type matrix, which are two of the various known forms of manganese dioxide. The structure of both pyrolusite and ramsdellite are described as infinite chains of edge-sharing MnO 6 octaedra connected by corners to neighbor chains. The difference between the two structures resides in the single chains (pyrolusite) or double chains (ramsdellite) delimiting 131 or 231 tunnels, respectively, along the c-axis. Therefore, intergrowth is not surprising since these two polymorph phases exhibit important similarities in their oxygen frameworks packing. Recently, the differences observed in the cristallinity and in the XRD patterns of various synthetic g -MnO 2 have been explained by the combined effect of two types of disorder [13]. The first one (de Wolff disorder) is related to the random distribution of single (pyrolusite) and double (ramsdellite) chains in the structure. The second one originates from the presence of microtwinning in the Ramsdellite structure on the (021) / (061) planes. Based on the presence of both perturbations, a calculation model was proposed to evaluate the cell parameters and their effects on displacements and broadening of the XRD Bragg peaks. Despite an important broadening of the Bragg peaks, we evaluated the cell parameters of our starting IBA[15 material to be a59.45 ˚ b52.82 A, ˚ c54.44 A ˚ in an orthorhombic system using A, (101), (210), (002), (211), (401), (212) and (402) set of lines. The pyrolusite content (Pr ) was evaluated to 37%. The BET surface area of IBA[15 is found to be around 20 m 2 / g, consistent with SEM observations showing very small particles highly agglomerated. The mean Mn oxidation state is determined equal to 3.9260.02.

3.2. Reaction of EMD with LiOH in ethylene glycol Descriptions of the experimental conditions and chemi-

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cal / structural characteristics of the samples recovered along this study are listed in Table 1. The first experiment (A) was performed with an initial LiOH / MnO 2 ratio of 0.1 M / 0.01 M in 200 ml of EG. By atomic absorption, we followed both the lithium ([Li]) and manganese ([Mn]) contents in the organic medium. In Fig. 1 are plotted the evolution of the [Li] and [Mn] together with the temperature. From this graph, it can be seen that above 808C, Mn is dissolved, as clearly seen from the appearance of an orange coloration of the liquid. Upon further heating at 1608C, there is no solid MnO 2 left in solution and [Mn] reaches a value of 0.01 M. At higher temperature (1908C), the amount of Mn in the solution decreases. This corresponds to the precipitation of an organo-metallic Mn-containing phase that was described and characterized in a previous work [14]. The main characteristic of this new phase is a very strong XRD reflection located at 10.68 2u. A second observation is that the lithium concentration is quite constant along the experiment, with the exception of a slight decrease during the first heating stage. If this tiny decrease in [Li] is due to a reaction with MnO 2 , one would expect to enhance the amplitude of this feature by performing a reaction with higher MnO 2 content. The above experiment was therefore repeated (B) with a lower initial LiOH / MnO 2 ratio (0.1 M / 0.1 M). The metallic contents evolution with temperature is plotted in Fig. 2. At room temperature, [Li] is lower (15%) than initially dissolved in the reaction medium. No solubility limit can be invoked to explain this observation since the initial LiOH quantities, and therefore the pH, are the same in both tests. Only the amount of MnO 2 changed, implying a room temperature reaction between MnO 2 and Li in EG. XRD analyses of reacted MnO 2 show an important shift towards low angles of the Bragg peaks without the appearance of any new phase. These modifications in the XRD patterns are consistent with those reported in the literature in the case

Table 1 Experimental parameters and composition of the recovered powders Sampleexperiment

Medium (200 ml)

MnO 2 (M)

LiOH (M)

NaOH (M)

KOH (M)

T (8C)

Reaction Time

Final product

Li / Mn Solid

Mn o.s.

A B C D E F G H I J K L M N O

EG EG EG EG EG EG EG EG EG H2O H2O EG EG EG EG

0.01 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1

0.1 0.1 0.1 0.2 0.4 1 0.4 0.4 0.4 0.4 / / / / /

/ / / / / / / / / / / / / 0.4 /

/ / / / / / / / / / / / / / 0.4

30–196 30–196 30 30 30 30 80 100 120 30 30 30 80 30 30

/ /

/ / Li 0.20 H 0.12 MnO 2 Li 0.24 H 0.09 MnO 2 Li 0.33 MnO 2 Li 0.35 MnO 2 Li 0.44 MnO 2 Li 0.80 MnO 2 1Li 2 CO 3 Li 0.84 MnO 2 1Li 2 CO 3 MnO 2 MnO 2 H 0.32 MnO 2 H 0.67 MnO 2 Na 0.05 H 0.27 MnO 2 K 0.04 H 0.29 MnO 2

/ / 0.20 0.24 0.33 0.35 0.44 0.96 1.13 0 / / / / /

/ / 3.60 3.59 3.61 3.58 3.47 3.12 3.08 3.92 3.92 3.60 3.25 3.60 3.59

4 3 4 2 21 20 15 4 24 2 24 3 2

days days days days h h h days h days h days days

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Fig. 1. Evolution of the concentrations of dissolved lithium and manganese in liquid ethylene glycol as a function of reaction time and temperature. The initial LiOH / MnO 2 molar ratio is 0.1 / 0.01.

of lithium insertion in the g-MnO 2 structure (Mn 41 → Mn 31 ) that proceeds with conservation of the Mn–O skeleton [5]. Furthermore, the Li / Mn ratio in the recovered solid phase is measured to be 0.16, in agreement with the [Li] left in solution. While the temperature increases, the [Li] further decreases until 1208C and the

dissolution of Mn begins (i.e. [Mn] in solution increases, visible orange coloration) at 808C. From 258C to 1208C, the Li / Mn ratio measured on each withdrawn solid phase continuously increases to reach a maximum value of around 0.30. Beyond 1208C, the dissolution of the Li x MnO 2 phase (i.e. simultaneous increase of both [Li]

Fig. 2. Evolution of the concentrations of dissolved lithium and manganese in liquid ethylene glycol as a function of reaction time and temperature. The initial LiOH / MnO 2 molar ratio is 0.1 / 0.1.

D. Larcher et al. / International Journal of Inorganic Materials 2 (2000) 389 – 396

and [Mn] in solution, red coloration) competes with the insertion of lithium in the g-MnO 2 structure. A large drop in [Mn] (i.e. precipitation of organo-metallic Mn-rich phase) occurs at around 1508C and XRD shows that no new phase is precipitated between 80 and 1508C. Due to the precipitation of small amounts of Li 2 CO 3 at high temperature, the final Li content at 1908C is slightly lower than initially introduced in the medium. Finally, 25–1508C seems to be the temperature range within which pure g-Li x MnO 2 phase can be prepared by this route. To know what the maximum x value is and what the experimental factors mainly involved in the insertion reaction are, we performed experiments at constant temperature, and followed the liquid and solid phases behaviors while the reaction proceeded. The first row of experiments was done at 308C with 0.1 M of MnO 2 in 200 ml of EG, and the LiOH concentration was altered from 0.1 M to 1 M (experiments C–F). For a starting LiOH / MnO 2 ratio50.1 / 0.1 (sample C), we observed that the x value reaches a maximum after ca. 48 h of reaction and that no significant Mn dissolution could be detected (max: 0.45%). After 4 days of reaction, we measured x50.20 in Li x MnO 2 . When varying the initial LiOH concentrations (0.2, 0.4 and 1 M), x reaches a maximum value of 0.33–0.35 for 0.4–1 M LiOH, and an intermediate value of 0.24 was obtained for 0.2 M LiOH. The Mn dissolution never exceeded 1%. XRD patterns of the EMD and the reduced phase are plotted in Fig. 3 that shows a higher shift towards low angles for Li 0.35 MnO 2 than for Li 0.20 MnO 2 . The Mn o.s. of all the phases

Fig. 3. X-rays diffraction patterns of the samples obtained by reaction EMD in ethylene glycol at 308C for several days and with different initial LiOH / MnO 2 molar ratios: 0.1 / 0.1 (sample C) and 1 / 0.1 (sample F). The X-rays diffraction pattern of non-reacted initial EMD is shown for reference on the bottom panel.

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prepared at 308C is evaluated around 3.60 (see Table 1). The difference in oxidation state (0.34) between starting EMD (3.92) and sample Li 0.35 MnO 2 (3.58) is consistent with the amount of lithium inserted per Mn (0.35). In contrast, these differences in Mn o.s. for Li 0.24 MnO 2 and Li 0.20 MnO 2 are not consistent with the measured lithium contents, suggesting that other species, namely H 1 , are inserted as will be further discussed. To investigate the temperature effect on the Li amount uptake, the following experiments were carried out. We fixed the initial LiOH / MnO 2 ratio to 0.4 M / 0.1 M and changed the temperature of the reaction medium. It must be kept in mind that increasing the temperature of the ethylene glycol medium induces a progressive dissolution of the Li x MnO 2 material. In the next syntheses, the reacting time will thus be limited to about 20 h in order to avoid too much dissolution. The X-rays diffraction patterns of the solid phase obtained after 20 h of reaction at 808C, 1008C and 1208C (experiments G, H, I) are plotted in Fig. 4. They show a progressive shift of the g-Li x MnO 2 reflections while the reaction temperature increases and the

Fig. 4. X-rays diffraction patterns of the samples obtained by reaction EMD and LiOH (initial molar ratio: 0.4 / 0.1) in ethylene glycol at different temperatures: 308C (sample E), 808C (sample G), 1008C (sample H) and 1208C (sample I). The X-rays diffraction pattern of non-reacted initial EMD is shown for reference on the bottom panel. For samples H and I, arrows indicate Li 2 CO 3 impurities.

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appearance of some new reflections for the sample made at 100 and 1208C, suggesting a structural modification. This was previously explained [5] as a result of increasing electrostatic interactions between the inserted species and the manganese ions. The XRD pattern collected on sample I is similar to that of Li 0.9 MnO 2 prepared in Ref.s [5,15]. For samples H and I, the presence of lithium carbonate impurities (confirmed by infrared spectroscopy) resulted in an over evaluation of the Li / Mn atomic ratios determined by AA. This over evaluation is obvious when one compares the measured Li contents with the changes in Mn o.s. However, we should stress that the determination of the Mn oxidation state is transparent to the presence of Li 2 CO 3 . Thus, the Li stoechiometry within these phases was evaluated from the changes in Mn o.s, assuming that only Li is inserted.

3.3. Reaction mechanism None of the reduced MnO 2 samples prepared in this work showed significant changes in the particles morphology or in their BET surface area, ruling out a dissolution / crystallization process. To better understand the mechanism of this organicdriven insertion reaction, we first performed various syntheses at 308C showing 1) that no insertion occurs when EG is replaced by water (samples J and K), and 2) noticeable modifications (shift towards low angles) in the X-rays pattern of EMD treated for 48 h in EG without LiOH (sample L) (Fig. 5). The Mn o.s. of the as-reduced g-MnO 2 is measured to 3.60 compared to 3.92 for the starting

Fig. 5. X-rays diffraction patterns of the samples recovered after treatment of 0.1 mole of EMD in LiOH-free EG at 308C (sample L) and at 808C (sample M). The X-rays diffraction pattern of non-reacted initial EMD is shown for reference on the bottom panel.

EMD. This clearly indicates that the ethylene glycol acts even at room temperature as a reducing agent for Mn. At 808C (sample M), the reaction of EMD in EG leads to Mn o.s. of 3.25 and drastic changes in XRD pattern (Fig. 5) as previously reported in the case of treatment of EMD by organic reducing agent [9]. The reduction of Mn 41 to Mn 31 is then compensated by the insertion of protons (produced by the oxidation of EG) in the structure leading to H x MnO 2 products. Indeed, the XRD pattern of the sample M, very different from that of Li x MnO 2 phases, matches well that of a-MnOOH (groutite). By reacting EMD in ethanol at 308C for 24 h, we also observe a significant shift of the Bragg peaks and a decrease in the Mn o.s. down to 3.85. Thus, ethanol acts also as a reducing agent at room temperature, and has therefore not been used as solvent to wash the powders. Two reaction paths can be proposed for the transformation of EMD into Li x MnO 2 phases. This process can be regarded either as a direct insertion of Li or as a two-step mechanism involving first the initial insertion of protons provided by the EG medium and secondly a subsequent

Fig. 6. X-rays diffraction patterns of the samples recovered after reacting 0.1 mole of EMD at 308C: in LiOH-free EG (sample L), with 0.4 mole LiOH in EG (sample E), with 0.4 mole NaOH in EG (sample N), with 0.4 mole KOH in EG (sample O). All these samples have a mean Mn oxidation state around 3.60. The X-rays diffraction pattern of non-reacted initial EMD is shown for reference on the bottom panel.

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exchange for lithium. To distinguish between these two possibilities, we performed a two-step preparation. First, the reaction of 0.1 M of MnO 2 in EG at 308C was carried out for a 48 h period to produce reduced H 0.32 MnO 2 . Next, we added 0.4 M of powdered LiOH to the resulting suspension, and observed the evolution of Li in the reaction medium at 308C during 2 days. Neither noticeable [Li] evolution nor XRD modification were detected. This led to the conclusion that the formation of Li x MnO 2 is likely a direct insertion of lithium in the MnO 2 rather than H 1 insertion followed by a proton for lithium exchange reaction. From the previous experiments, we can easily conclude that the nature of the intercalates depends on the cationic species present in the solution, and of their ability / competition to insert into the structure. To further dwell on this point we investigated the possibility of forming Na or K intercalated MnO 2 . We then reacted EMD in NaOH or KOH loaded EG and obtained H 0.27 Na 0.05 Mn 3.60 O 2 (sample N) and H 0.29 K 0.04 Mn 3.60 O 2 (sample O) materials, respectively., The difference in the composition of the final products, compared to the proton free Li 0.35 MnO 2 phase obtained in the same experimental conditions with LiOH, can be related to the larger ionic radii of Na 1 and K 1 vs. that of Li 1 . We cannot therefore totally rule out that the presence of such small amounts of Na and K could also be due to adsorbed species. Thus, caution has to be exercised regarding the localization of the Na 1 / K 1 species in these materials. Finally, the competition between cationic insertion can also be further exemplified by looking at the (H,Li) x MnO 2 phases (samples C and D) obtained with low LiOH concentrations.

3.4. Stability Highly reduced metal oxides (e.g. Li 2 Mn 2 O 4 , gLiMnO 2 ) are generally unstable when exposed to moisture [16–18]. Surprisingly, sample I (Li 0.84 MnO 2 ) was found to remain unchanged, at least as deduced from XRD results, after being exposed to ambient air for several weeks. However, longer exposure periods resulted in modifications of the X-rays pattern and in a progressive shift of Bragg peaks consistent with a transformation towards lower lithium stoichiometry. Simultaneously, peaks due to lithium carbonate clearly grow in intensity. After a resting period of 3 years, X-rays pattern of sample I matches that of Li |0.3 MnO 2 phases. Despite an unambiguously long term instability, the relatively slow transformation of our reduced phases could be ascribed to the protective effect of surface carbonates and / or adsorbed EG, both detected in small amounts by IR. For instance, it is well known that a thin EG coverage onto the surfaces of Ni particles can considerably limit their oxidation upon heating under air. Consistently, EG / carbonate free highly lithiated g-Li x MnO 2 prepared according to Ref. [5] undergoes a very fast degradation upon exposure to air.

395

Interestingly, it has been shown that moisture reactive lithiated manganese oxides are stabilized in aqueous LiOH solution with a pH higher than 10 (.10 4 M LiOH) [17]. In light of these results, a combination of much higher LiOH concentrations (0.01 M 1 M) in a polar reducing solvent (EG) could well account for the in situ formation and stability of the present highly lithiated g MnO 2 phases.

3.5. Electrochemical tests Fig. 7 depicts the electrochemical curve collected at C / 5 for sample I. The capacity amplitude (e.g. length) of the first charge is in good agreement with the amount of inserted species as deduced from the change in Mn o.s. between the starting EMD and the chemically reacted sample (0.84). However, this oxidation is occurring at significantly higher voltage than that observed for the subsequent ones. Similar important differences between the first and the following charges, related to irreversible structural transitions, have also been observed for LiMnO 2 materials prepared at low temperatures [2]. Based on previous works [8,19] which have shown that either electrochemical or chemical lithium inserted g-Li |0.8 MnO 2 phases are isostructural, one would expect our material to have the same oxidation signature. Surprisingly, the first charges for EG lithiated EMD and for electrochemically Li inserted EMD g-Li |0.8 MnO 2 (sample I) are drastically different (Fig. 8). Therefore, drastic irreversible structural transformation cannot be retained to explain this point. In Fig. 8 is also displayed the first charge for Hx MnO 2 (sample L) prepared in Li free EG. The removal of protons from the MnO 2 structure obviously results in a shift to higher voltage when compared to that of Li extraction suggesting the presence of protons within our chemically made materials. Although indirectly proved, we believe that the co-insertion of protons is inherent to the present

Fig. 7. Electrochemical curve vs. lithium metal for sample I at a constant rate of C / 5. Note the high voltage of the first charge compared to the subsequent ones. After one cycle, the cutoff voltage was lowered down to 4.0 V in order to limit electrolyte degradation.

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Fig. 8. Comparison of the first charge (C / 5) for sample I (a), electrochemically inserted EMD (Li 0.8 MnO 2 ) (b) and proton inserted EMD in Li free ethylene glycol (sample L) (c).

EG driven preparation method, and can never be totally avoided. Such high voltage first charge has also been reported for electrochemically lithiated phases prepared in aqueous medium from g-Li 0.36 MnO 2 [20]. Finally, the absence of high first charge voltage for untreated EMD demonstrates that this feature is intrinsic to EG prepared materials. Beyond the first cycle, our materials basically behave like pristine EMD vs. Li. No noticeable evolution is seen in the shape of the successive discharge curves showing two distinct plateaus located at 3.0 and 3.4 V, whereas a unique plateau at 3.2 V is observed upon charge. These plateaus have been carefully studied and linked to biphasic insertion / extraction mechanisms [8]. Finally, the overall capacity for the EG prepared g-Li x MnO 2 rapidly drops over the first 10 cycles followed by a slower rate decay. Nevertheless, the capacity retention is very poor since 50% of the initial capacity is lost after only 50 cycles.

4. Conclusion The chemical insertion of cations into the structure of g-MnO 2 has been studied in organic reducing medium (ethylene glycol). We showed that the extent of the reduction / insertion is mainly linked to the reaction temperature, and that the nature / proportion of the inserted species can be tuned by carefully controlling the com-

position of the liquid medium. Although it appears that insertion of protons can never be totally avoided, high lithium contents (Li |0.9 MnO 2 ) can be obtained in lithium rich EG and at temperature of 1208C. Despite an important extent of reduction, these materials show improved air / moisture stability likely due to carbonates and EG coverage. Electrochemical tests performed vs. lithium on highly chemically Li inserted EMD phases show, beyond the first charge, similar electrochemical performances as obtained for commercial EMD. We believe that this soft and cheap preparation method is promising, and could find multiple applications for the making / stabilization of tailor made reduced oxides (e.g. lithium metal oxides). As an example, preliminary results indicate that cationic species can easily be inserted in oxides with open structure such as b-MnO 2 , V2 O 5 or RuO 2 .

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