Electrochemical lithium intercalation in disordered manganese oxides

Solid State Ionics 138 (2001) 203–212 www.elsevier.com / locate / ssi

Electrochemical lithium intercalation in disordered manganese oxides A. Ibarra Palos, M. Anne, P. Strobel* Centre National de la Recherchere Scientifique, Laboratoire de Cristallographie, BP 166, 38042 Grenoble Cedex 9, France Received 31 May 2000; received in revised form 10 August 2000; accepted 9 October 2000

Abstract Four highly disordered manganese oxides were prepared by reduction of sodium permanganate by chloride, iodide, hydrogen peroxide or oxalate in aqueous medium containing a large excess of Li 1 ions, yielding hydrated oxides with Mn valence in the range 3.80–3.92. Thermogravimetric studies showed that the iodide-reduced oxide can be dehydrated to 92% at 2408C, while the other three ones retain water at temperatures up to ca. 4008C, where crystallization is significant. The electrochemical behaviour was studied potentiostatically and galvanostatically in lithium cells on samples dried at 2408C. All samples give a unique, broad, reversible oxidation–reduction peak in the range of 2.0–3.6 V. Cycling capacities vary in decreasing order iodide . hydrogen peroxide . chloride . oxalate. The best compound, Li 0.60 Na 0.16 MnO 2.33 I ¯0.05 , has an initial capacity of . 170 Ah / kg, slowly decreasing on cycling to reach 140 Ah / kg after 70 cycles. These performances, which are far superior to those of manganese spinels, are compared to those of recently reported amorphous manganese oxi-iodides prepared in anhydrous conditions.  2001 Elsevier Science B.V. All rights reserved. Keywords: Manganese oxide; Amorphous oxide; Lithium intercalation; Lithium batteries

1. Introduction Manganese oxides are among the most attractive cathode materials for lithium batteries. Compared to the cobalt oxide used in presently commercialized batteries, they present significant advantages in terms of cost and environmental impact. The most widely studied manganese oxide in this respect is LiMn 2 O 4 (and substituted variants) with the spinel structure. Its use, however, is beset by a constant decrease in capacity with cycling or with storage at temperatures *Corresponding author. Tel.: 133-476-887-940; fax: 133-476881-038. E-mail address: [email protected] (P. Strobel).

in excess of ca. 508C [1]. Other crystalline manganese oxides such as b-MnO 2 (with the rutile structure) or g-MnO 2 have early been shown to possess very limited reversible lithium intercalation capacity. Finally, layered manganates of the birnessite or LiMnO 2 type are attractive, but the structures are unstable on delithiation and these materials convert progressively to spinel on cycling [2–4]. Much fewer studies have been devoted to noncrystalline manganese oxides. Since 1997, several amorphous phases were reported as cathode materials. Most of them were prepared from permanganates, using various reducing agents such as oxalic acid [4], fumaric acid [5] or potassium borohydride [6]. The most promising of these materials is an

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oxi-iodide proposed by Kim and Manthiram [7,8], which showed capacities in excess of 200 mAh / g at the 40th cycle (in the potential window 1.5–4.5 V) and no significant tendency to convert into spinel. Peculiarities of Kim and Manthiram’s route include (i) the use of lithium iodide as a reducing agent, (ii) a synthesis process completely excluding water. The latter aspect is a serious drawback in view of practical applications. In addition, a more recent report on this oxi-iodide showed capacities of only ca. 120 mAh / g within limits 2–4 V [9]. These data prompted us to investigate amorphous manganese oxides, using a slightly different synthesis procedure. In this paper, we present new syntheses of nearly amorphous lithium / sodium manganates from the reaction of permanganate with various reducing agents in an aqueous lithium hydroxide medium. Thermogravimetric studies will show a particular behaviour for the material obtained with lithium iodide, which can be almost totally dehydrated at 2408C. Its electrochemical behaviour in lithium cells at room temperature will be presented. Galvanostatic and step-potential cyclings show a single-phase behaviour centered around 2.85 V and a remarkable cycling stability.

2. Experimental The starting material was a 0.5 M aqueous solution of sodium permanganate (Aldrich). The reducing agents used were lithium chloride, lithium iodide, oxygen peroxide and lithium oxalate (see Table 1), and the samples obtained will be abbreviated to

Mn–Cl, Mn–I, Mn–HO and Mn–ox, respectively, throughout this paper. All reactions were carried out in a 5–6-fold excess of Li 1 with respect to the NaMnO 4 concentration at room temperature under vigorous stirring for 15 h. Products were washed with distilled water, filtered and dried at 808C in air. The samples were characterized by X-ray diffraction using a Siemens D-5000 diffractometer (Cu Ka radiation) and scanning electron microscopy (JEOL 840). Chemical compositions were established using atomic absorption spectrophotometry for Li, Na and Mn contents and standard oxalate / permanganate volumetric titration for manganese oxidation state determination. For samples made from LiI reduction, iodine / manganese ratios were obtained from EDX measurements coupled to SEM. The EDX acquisition conditions (in particular the sample surface) were not optimized, in view of the rather small iodine contents initially found. Water losses were measured by thermogravimetry in air, using a 18C / min heating / cooling rate in a Perkin-Elmer TG-7 thermobalance. Electrochemical tests were carried out in liquid electrolyte at room temperature using Swagelok  or 2430 coin cells. The electrolyte was a 1 M solution of LiPF 6 in EC-DME 1:2. Positive electrodes were cut from a paste made of an intimate mixture of oxide, carbon black (Y50A grade, SNNA, Berre, France) and PTFE in weight ratio 70:20:10. Typical quantities used in cells were 6–12 mg / cm 2 of active material. Negative electrodes were 200-mm thick lithium foil (Metall Ges., Germany). Cells were assembled in a glove box under argon with # 1 ppm H 2 O. Electrochemical studies were carried out using

Table 1 Preparation conditions and analytical data Sample abbrev.

Mn–Cl Mn–I Mn–HO Mn–ox Ref. [7] a

Reaction medium

Product analysis

Reducing agent

Li / Mn ratio in solution

Na / Mn

Li / Mn

Mn valence

LiCl LiI H 2 O 2 –1.25 M LiOH H 2 C 2 O 4 –1.25 M LiOH LiI

6.0 6.0 5.0 5.0 1.5 a

0.15 0.16 0.17 0.20 1.51

0.315 0.60 0.305 0.27 0.51

3.90 3.91 3.915 3.805 3.80

Sample with best electrochemical performances.

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a MacPile Controller (Bio-Logic, Claix, France), in either galvanostatic mode or step-potential electrochemical spectroscopy (SPES) [10], using typically 10 mV/ h steps. X-ray powder diagrams after cycling were measured directly on the pellets after disassembling the cells; the diameter of 2430 coin cells is especially suited for our diffractometer. However, such diagrams contained extra lines due to carbon, lithium salts and collector materials.

3. Results

3.1. Synthesis and chemical composition

Fig. 1. X-ray powder diffraction diagrams of samples as synthesized (dried at 808C). A typical X-ray diagram of crystallized sodium phyllomanganate [2] is included for comparison (top).

The permanganate decomposition is especially fast in the presence of hydrogen peroxide. All reactions but that with lithium iodide resulted in rather low yields ( , 30% as MnO 2 ), probably due to spontaneous decomposition of the reducing agent (in the hydrogen peroxide and oxalic acid case) or to partial redissolution of manganese. A blue–green color was observed for the first filtrate in the H 2 O 2 case, which seems to indicate the presence of some MnO 432 anionic species in the alkaline solution. In contrast, the lithium iodide reaction resulted in a quantitative yield based on manganese. Subsequent chemical analyses (see Table 1) showed that all solids obtained have manganese oxidation states in the range 3.80–3.92. Samples Mn–Cl and Mn–HO are especially close in composition, with formula Li 0.31 Na 0.16 MnO ¯2.2 ? nH 2 O within experimental errors. Sample Mn–ox has the lowest Mn oxidation state and the highest Na / Li ratio. The iodide-derived sample Mn–I is atypical regarding its lithium content, which is twice as large as that in other samples. EDX analysis showed that it contains minor amounts of iodine (0.03–0.06 per Mn atom). We note, however, that the composition obtained here is markedly different from that obtained in anhydrous medium [7,8], where much higher sodium contents were reported (see Table 1, last line).

centered near 128 2u (Cu Ka). This feature reminds us of phyllomanganates with the birnessite structure, ˚ which have a typical interlayer distance of ca. 7.0 A [11–13]. The trend towards structural organization decreases in the order sample Mn–HO . Mn–Cl . Mn–I . Mn–ox. But only sample Mn–HO shows the second typical line at 258 corresponding to the ˚ layer next possible reflection connected with a 7 A stacking (see asterisk in Fig. 1). In addition, the first ˚ a rather peak of sample Mn–Cl yields d 5 6.58 A, odd value for such layered manganates stabilized by a layer of water molecules, giving a fairly constant ˚ The other two broad periodicity of 7.160.1 A. reflections at 37 and 668, which are common to all our samples, are typical of Mn 41 –O spacings in octahedral coordination and can be indexed as 100 and 110 in a basic hexagonal subcell with a 5 2.80 ˚ equal to the edge of such a Mn 41 –O octahedron. A, The large width of all these lines shows that the compounds as prepared exhibit considerable disorder around these average periodicities. SEM micrographs (Fig. 2) show that these samples are highly porous and consist of an agglomeration of particles with irregular, sharp edges and grain size in the 1–5 mm range.

3.2. Structure and morphology

3.3. Thermal behaviour

Samples as prepared and dried at 808C show a few very broad X-ray lines (Fig. 1), especially one

Thermogravimetric and DTG curves are given in Fig. 3. All samples lose mass, starting from room

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Fig. 2. Scanning electron micrographs of samples Mn–Cl (a), Mn–I (b) and Mn–HO (c).

temperature. The mass losses up to 4008C are 9– 11% for samples Mn–Cl and Mn–ox and 18–19% for Mn–I and Mn–HO. Only a small fraction of this

mass is recovered on cooling. The irreversible part of the mass loss is attributed to adsorbed water. It is completed below 3508C in sample Mn–I, and near

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Fig. 3. Thermogravimetric analysis of the four samples in air (heating rate 18C / mn). (a) Mass loss, (b) DTG as a function of temperature.

380–4008C in other samples (see exact figures in Table 2). The water contents calculated from the irreversible mass loss are in the range of 0.5–1.2 per formula unit (see Table 2). DTG curves (Fig. 3b) confirm the similarity between samples Mn–Cl and Mn–ox, which are

characterized by the presence of two DTG peaks around 150 and 3608C. Sample Mn–HO shows an almost featureless mass loss variation up to 3808C. The most remarkable sample is Mn–I, which exhibits two well-defined mass loss steps, corresponding to sharp peaks in DTG (Fig. 3b). These

Table 2 Features of thermogravimetric analysis on amorphous Mn oxides Sample

Mass loss at 2408C (A, %)

Mass loss at 4008C (%)

Irreversible mass loss (B, %)

T irr (8C)

H 2 O / Mn a

A /B b

DTG maxima (8C)

Mn–Cl Mn–I Mn–HO Mn–ox

5.03 14.9 11.7 7.00

8.85 17.8 18.8 11.25

9.145 16.1 18.3 11.5

413 325 375 415

0.53 1.07 1.20 0.69

0.55 0.925 0.64 0.61

1501360 951210 1401275 1501360

a b

From irreversible mass loss B. Fraction of water lost at 2408C.

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peaks are not only sharper than those observed in the other samples, but occur at much lower temperatures. In particular, the steep mass loss at 2108C stops abruptly and is followed by a near plateau vs. temperature. As a result, this sample looses 92.5% of its total irreversible mass loss below 2408C, compared to 55–64% for the other samples. X-ray diffraction carried out after annealing shows a displacement of the first broad line from 128 to 188 in the 2u range (see Fig. 4). The higher the annealing temperature, the better the definition of diffraction lines in these samples. This set of reflections corresponds roughly to the positions of the main spinel or Li 2 MnO 3 reflections, indicating a tendency towards ordering into three-dimensional structures based on octahedral coordination of manganese by oxygen. Similar crystallization trends were observed on other samples annealed at 400 or 5008C. The plateau above 2108C in sample Mn–I provides an interesting way of removing most or all adsorbed water in a temperature range where the sample remains nearly amorphous and oxygen losses do not yet take place. This is an important advantage over other samples, which exhibit no clear limit in mass loss below ca. 3808C. As a result, most electrochemical studies were carried out after drying under vacuum at 2408C.

Fig. 4. X-ray powder diffraction diagrams after annealing. From top to bottom: sample Mn–Cl at 4008C, sample Mn–I dried at various temperatures. The 2408C diagram was measured in an air-tight sample holder to avoid re-hydratation in ambiant air.

3.4. Electrochemical behaviour 3.4.1. SPES Step-potential cyclings carried out in similar conditions (10 mV/ h) on samples Mn–Cl, Mn–I and Mn–HO dried at 2408C are compared in Fig. 5, which also includes LiMn 2 O 4 as a typical crystallized phase reference. All samples give a unique, broad current peak centered around 2.85 V in reduction and 3.05 V in oxidation. The peak widths (especially large for sample Mn–Cl) and the significant overlap between reduction and oxidation peaks differ markedly from the LiMn 2 O 4 case. For the latter, sharp peaks with a common initial potential are typical features of a two-phase intercalation process [14]. As shown in Fig. 6, the current peak shape for

Fig. 5. Compared SPES voltamograms of samples Mn–Cl, Mn–I and Mn–HO measured in similar conditions (active mass: 5–15 mg / cm 2 , potential scanning rate: 10 mV/ h). Also shown is a LiMn 2 O 4 spinel, as a typical case of a crystallized compound giving rise to a two-phase intercalation process (dashed line).

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Fig. 6. Second and eighth cycle SPES voltamograms of sample Mn–I at 10 mV/ h. The higher intensity at the beginning of eighth discharge (3.0–3.4 V) is due to an acceleration condition (reduction of the potential step duration in regions with small initial intensity).

sample Mn–I is very stable on cycling and does not show any significant evolution towards spinel behaviour, as was frequently observed for layered manganates and other poorly crystallized manganese oxides [2,4,15]. Voltammograms measured on sample Mn–I at three different scan rates are displayed in Fig. 7. The i(U ) curve shape and the potential scale are very similar for all speeds, indicating rather fast kinetics and a probable single-phase behaviour. This is confirmed by a detailed examination of the current evolution during incremental steps in the reduction peak (Fig. 7, inset), which shows that this phenomenon is diffusion-controlled. However, we note a constant shift between reduction and oxidation peaks. These curves can be regarded as the superposition of a very broad peak (comparable to the one given by sample Mn–Cl in Fig. 5) and a sharper, central part, for which an interpolation of charge and discharge branches allow to define a characteristic potential near 2.94 V. This feature is especially prominent on oxidation, where a second peak is almost resolved at 10 mV/ h.

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Fig. 7. Influence of scanning speed on the SPES voltamogram of sample Mn–I. Inset: incremental current evolution through the reduction peak at 10 mV/ h steps.

Voltammograms were measured on sample Mn–I in potential ranges as wide as 1.8 V in reduction and 4.5 V in oxidation. No additional feature was observed; in particular, no significant capacity was detected around 4 V, in agreement with the initial manganese oxidation state close to 14 in the samples studied.

3.4.2. Galvanostatic cyclings Fig. 8 shows the first cycle behaviour (galvanostatic and potentiostatic) of the four samples dried at 2408C. In agreement with the broad current peaks observed in SPES mode, the charge–discharge curves are S-shaped, with fairly low polarization (ca. 200 mV) between reduction and oxidation. Capacities are in the order Mn–I.Mn–HO.Mn–Cl.Mn– ox. The same trend applies on cycling (see compared data for cycles 1 and 33 in Table 3). Sample Mn–I again stands out for its remarkable performances. Measurements under three sets of conditions on this sample (Fig. 9) show that SPES cycling gives a

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Fig. 8. First galvanostatic charge–discharge curves of samples dried at 2408C at C / 12 rate.

quantitative lithium intercalation down to the Mn 31 oxidation state (x Li 50.97), while the initial C / 12 capacity is also outstanding (x Li 50.76, Q s 5175 Ah / kg). These figures are rather high for manganese oxides, especially considering that the reference formula contains one Mn per formula unit, against two in Li 11x Mn 2 O 4 (the theoretical capacity at 3 V for the latter system, which is never reached, is 148 Ah / kg). Fig. 9 also shows the detrimental effect of water on both initial capacity and rechargeability: the water-containing sample discharge capacity is ca. 25% lower than that of samples dried at 2408C. Moreover, several cells containing sample Mn–I dried at 1208C only could not be charged above 3.25 V, and showed instead a parasitic oxidation reaction,

Fig. 9. Charge–discharge curves of sample Mn–I dried at 1208C (one cycle shown) and 2408C (20 cycles) under galvanostatic control at C / 12 rate. Dashed line: first cycle under potentiostatic control at 10 mV/ h.

very likely water oxidation (giving rise to artificial negative x values on charge). In contrast, cells containing sample Mn–I dried at 2408C were very stable when cycled in the 1.8–3.8 V range, as shown by the first 20 cycles displayed in Fig. 9. The capacity of sample Mn–I dried at 2408C decreases slowly on cycling (see Fig. 10), but remains high for a manganese oxide (140 mAh / g after 65 cycles). The X-ray analysis of batteries disassembled after various numbers of charge–discharge cycles (Fig. 11) shows no clear evidence of crystallization: the bumps present in the material dried at 2408C are hardly enhanced and remain very broad, even after 71 cycles.

Table 3 Electrochemical capacities of amorphous Mn oxides Sample no.

Mn–Cl Mn–I Mn–HO Mn–ox

Potentiostatic cycling (10 mV/ h)

Galvanostatic cycling (¯C / 12)

First discharge

First discharge

33th discharge

x Li / Mn

Q s (Ah / kg)

x Li / Mn

Q s (Ah / kg)

x Li / Mn

Q s (Ah / kg)

0.52 0.97 0.66 0.37

121 222 154 86

0.31 0.76 0.50 ,0.25

72 175 116

0.34 0.67 0.43 n/a

79 152 99

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Fig. 10. Evolution of the capacity with cycle number for sample Mn–I at C / 12 regime (potential window: 1.8–3.8 V).

Fig. 11. Evolution of the X-ray diffraction diagram of sample Mn–I with cycling.

4. Discussion The samples obtained in this study present analogies with previously studied amorphous manganese oxides [4–8]: a highly divided morphology and / or porosity, and a manganese oxidation state in the range 3.7–3.9. However, they differ widely in alkali metal content, A / Mn, which varies from 0.06 for the fumarate reduction [5] to .1.5 for the LiI reduction in anhydrous medium [7,8]. The compositions of the samples obtained in this study are fairly close to those of Leroux and Nazar [4] (A / Mn in the range of 0.33–0.47); interestingly, the exception is the LiI-

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reduced sample, which is richer in lithium, confirming the trend observed by Kim and Manthiram [7,8] (see Table 1). The dehydration behaviour of the LiI-reduced sample is also peculiar, as it is the only sample in our series which loses more than 90% of its water below 2508C. Regarding X-ray diffraction, it must be noted that permanganate reduction is a classical route to phyllomanganates [11,12]; the presence of broad lines corresponding to birnessite interlayer spacings is thus not surprising. Leroux and Nazar also noted the presence of such broad diffraction peaks in samples prepared by oxalate reduction [4]. The electrochemical properties resemble those of phyllomanganates [2,12], characterized by a main broad single reduction / oxidation peak centered around 2.8–3.0 V. However, deintercalated layered manganates tend to collapse on deintercalation [2,15], unless they are stabilized by larger cations and / or structural water [16]. The absence of crystallization in sample Mn–I after 71 cycles is remarkable in this respect. The presence of sodium seems to have not any detrimental effect on cycling, as shown by the good performances of the products obtained in this study and in Ref. [7] by iodide reduction of NaMnO 4 ; thus the sodium–lithium ion exchange is not a necessary step to obtain interesting cathode materials. It is especially worth pointing out that Kim and Manthiram [7,8] investigated a wide range of Na / Mn and Li / Mn compositions, and found larger capacities for samples with Na / Mn.0.5 than for sodium-poor samples. For phyllomanganates and Leroux’s samples, there is also ample evidence of better cycling performances for Na or K-containing compounds [2,4,15]. The presence of heavier alkali cations (Na, K) induces an increase in molar mass, hence a lower specific capacity (in Ah / kg) at constant intercalation level. On the other hand, the presence of such cations, which have too large ionic radii to enter the spinel structure, may play an important role in preventing the progressive emergence of the unwanted Li–Mn–O spinel phase, a key feature for long-time cycling capacity. The performances of samples obtained by oxalate, hydrogen peroxide or chloride reduction of permanganate are not outstanding. This may be due in part to the presence of an important fraction of

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residual water. No data about the capacity evolution with cycling were given in Refs. [4,5]. For fumaric acid reduction, Xu et al. [5] stated a capacity decrease of 2–3% per cycle, compared to ,0.3% for our sample Mn–I. The reasons for the superior performances of amorphous oxi-iodides are not clear at the moment. Because of its high atomic mass, iodine increases the ‘dead weight’ in a lithium-intercalation oxide, unavoidably reducing the specific capacity Q s at a given intercalation level, x Li . However, this effect can be more than compensated by its possible redox and structural role: iodine may indeed influence the redox processes, in spite of its low content in our samples. More important, iodine can help in preventing the crystallization of rigid framework structures which would be detrimental to lithium mobility, just as sodium. Further work is necessary (i) to determine the actual influence of iodine, (ii) to improve the high intercalation capacity in this material.

5. Conclusion The recent announcement of an amorphous manganese oxide obtained by iodide reduction of permanganate [7] prompted us to investigate this source of lithium battery positive materials, using aqueous reduction reactions. Four highly disordered and hydrated manganese oxides with Mn oxidation state close to 14 were obtained. We showed that the sample obtained by iodide reduction, Li 0.60 Na 0.16 MnO 2.33 I ¯0.05 , emerges as a remarkable material in this series: it is the most disordered one, the one that is the easiest to dry at moderate temperature and the one exhibiting by far the best electrochemical behaviour. Its capacity on cycling in the 3 V range (140 Ah / kg at the 70th cycle) are far superior to that of crystallized manganese oxides such as Li–Mn–O spinels. The presence of residual sodium is believed to play a key role in opposing the

trend towards crystallization on cycling. Further insights in the structural mechanisms are planned via in-situ X-ray absorption measurements.

Acknowledgements A. Ibarra Palos’s work is supported by a PhD grant from the Mexican–French cooperation agreement CONACYT / SFERE. The authors wish to thank Y. Chabre for helpful discussion of SPES results.

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