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Hydrothermal Synthesis and Characterization of KxMnO2‚yH2O Rongji Chen, Peter Zavalij, and M. Stanley Whittingham* Chemistry Department and Materials Research Center, State University of New York at Binghamton, Binghamton, New York 13902 Received November 17, 1995. Revised Manuscript Received February 15, 1996X
We report here the direct synthesis of a hexagonal form of manganese dioxide using mild hydrothermal methods. The reaction of potassium or sodium permanganate in water at 170 °C leads directly to potassium or sodium manganese dioxide, Alk≈0.25MnO2‚0.6H2O, with a R3 h m rhombohedral structure like Lix(H2O)TiS2 and a 7 Å repeat distance indicative of a monolayer of water between the manganese dioxide layers. This manganese oxide reacts readily and reversibly with lithium ions.
Introduction In recent years, there has been much interest in generating a rational approach to the synthesis of inorganic compounds, so that new structures with the desired optimum chemical and physical properties can be formed on demand. Several low-temperature techniques may yield such rational synthesis approaches. Such techniques include hydrothermal reactions, solgel process, intercalation, and ion exchange. Among these techniques, we have used the hydrothermal method1,2 in a search for new transition-metal oxides that might have open crystalline structures that would be electrochemically active. Hydrothermal reactions occur naturally in the crust of the earth aiding in the formation of minerals. The goal in using mild hydrothermal reactions is to precipitate new solids from the reaction nutrient which may have open crystalline structures different from the thermodynamically stable phases formed under normal high-temperature synthesis conditions. Such “metastable” phases may then be able to undergo ion ex* To whom enquiries should be sent. Phone: (607) 777-4623. Fax: (607) 777-4623. E-mail:
[email protected]. X Abstract published in Advance ACS Abstracts, May 1, 1996. (1) Reis, K. P.; Ramanan, A.; Whittingham, M. S. Chem. Mater. 1990, 2, 219. (2) Whittingham, M. S.; Guo, J.-D.; Chen, R.; Chirayil, T.; Janauer, G.; Zavalij, P. Solid State Ionics 1995, 75, 257. (3) Barrer, R. M. Hydrothermal Chemistry of Zeolites; Academic Press: London, 1982. (4) Beck, J. S.; Vartuli, J. C.; Roth, W. J.; Leonowicz, M. E.; Kresge, C. T.; Schmitt, K. D.; Chu, C. T.-W.; Olson, D. H.; Sheppard, E. W.; McCullen, S. B.; Higgins, J. B.; Schlenker, J. L. J. Am. Chem. Soc. 1992, 114, 10834. (5) Girnus, I.; Pohl, M.-M.; Richter-Mendau, J.; Schneider, M.; Noack, M.; Venzke, D.; Caro, J. Adv. Mater. 1995, 7, 711. (6) Johnson, J. W.; Jacobson, A. J.; Brody, J. F.; Rich, S. M. Inorg. Chem. 1982, 21, 3820. (7) Haushalter, R. C.; Mundi, L. A. Chem. Mater. 1992, 4, 31. (8) Whittingham, M. S.; Li, J.; Guo, J.-D.; Zavalij, P. Soft Chemistry Routes to New Materials; Rouxel, J., Tournoux, M., Brec, R., Eds.; Trans Tech Publications Ltd.: Nantes, France, 1993; Vols. 152 and 153, p 99. (9) Komarneni, S.; Li, Q. H.; Roy, R. J. Mater. Chem. 1994, 4, 1903. (10) Stein, A.; Fendorf, M.; Jarvie, T.; Mueller, K. T.; Benesi, A. J.; Mallouk, T. E. Chem. Mater. 1995, 7, 304. (11) Harrison, W. T. A.; Dussack, L. L.; Jacobson, A. J. J. Solid State Chem. 1995, 116, 95. (12) Feng, Q.; Kanoh, H.; Miyai, Y.; Ooi, K. Chem. Mater. 1995, 7, 1226. (13) Whittingham, M. S. Curr. Opinion Solid State Mater. Sci. 1996, 1, 227.
S0897-4756(95)00550-3 CCC: $12.00
change, intercalation, and other reactions with solution species to form additional novel compounds. Mild hydrothermal synthesis has been shown to be a powerful technique in the synthesis of novel compounds. This technique has been previously used in the synthesis of zeolites,3-5 of vanadium phosphates,6,7 and more recently of new metal oxides.1,2,8-13 Manganese oxides, such as LiMn2O4, are of particular interest14-16 because they readily intercalate lithium into their structures and are therefore potentially useful as the cathode of lithium batteries. However, only 0.5 lithium can be cycled per manganese atom so that their energy densities are not sufficiently high. In contrast the SONY lithium ion cell uses the very expensive LiCoO2 cathode. Hence, extensive research is currently underway to find promising candidates for cathode materials in lithium secondary batteries to replace the expensive cobalt used in the present commercial cells and a manganese oxide that behaved like the layered LiCoO2 would be a prime candidate for this application because of its high free energy of reaction with lithium and relatively low cost. The effectiveness of using potassium permanganate as the starting reagent to prepare layered manganese oxides, such as the ranciete form, has been reported;17,18 the hydrogen form was first synthesized, and the alkali forms were prepared by ion exchange. Bach et al.19 formed an amorphous bismuth-doped birnessite phase by the slow reduction of a potassium permanganate solution by concentrated nitric acid containing bismuth (14) Tarascon, J. M.; Wang, E.; Shokoohi, F. K.; McKinnon, W. R.; Colson, S. J. Electrochem. Soc. 1991, 138, 2859. (15) Guyomard, D.; Tarascon, J. M. J. Electrochem. Soc. 1992, 139, 937. (16) Thackeray, M. M.; Roussow, M. H.; Kock, A. d.; Harpe, A. P. d. l.; Gummow, R. J.; Pearce, K.; Liles, D. C. J. Power Sources 1993, 43-44, 289. (17) Tsuji, M.; Komarneni, S.; Tamaura, Y.; Abe, M. Mater. Res. Bull. 1992, 27, 741. (18) Leroux, F.; Guyomard, D.; Piffard, Y. Solid State Ionics 1995, 80, 299. (19) Bach, S.; Pereiraramos, J. P.; Cachet, C.; Bode, M.; Yu, L. T. Electrochim. Acta 1995, 40, 785. (20) Bach, S.; Henry, M.; Baffier, N.; Livage, J. J. Solid State Chem. 1990, 88, 325. (21) LeGoff, P.; Baffier, N.; Bach, S.; Pereiara-Ramos, J. P.; Messina, R. Solid State Ionics 1993, 61, 309. (22) Ching, S.; Landrigan, J. A.; Jorgenson, M. L.; Duan, N.; Suib, S. L.; O’Young, C.-L. Chem. Mater. 1995, 7, 1604.
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nitrates; this phase had compositions such as Bi0.1MnO2.1‚0.95H2O. Feng et al.12 reported the formation of sodium birnessite, Na0.40Mn2.15O2‚0.6H2O with an interlayer spacing of 7.19 Å, by the hydrothermal treatment of γ-MnO2 with NaOH; others20-22 reported that manganese oxides can be formed by sol-gel reactions at 400-600 °C, where the gel was formed by the reduction of permanganate by organics such as fumaric acid and sugars. We decided to try and form a layered lithium manganese oxide directly under mild hydrothermal conditions by the reduction of potassium or sodium permanganate in the presence of lithium ions, for example, by adding lithium chloride to the reaction medium. This led to the formation of a layered phase. We then explored the use of a range of alkali halides or just permanganate, and here report our findings. The most common manganese mineral in soils is birnessite, with the general formula Na0.1Ca0.04MnO2‚0.4H2O.23 Interestingly, birnessite was reported24 to be unstable to hydrothermal treatment at 155 °C converting to buserite, 10 Å spacing, and todorokite; these are essentially the same conditions under which we report here the formation of birnessite. We also explored the relationship between these manganese phases and the wellknown layered dichalcogenide phases.25 Experimental Section The initial experimental studies involved reacting potassium permanganate with hydrochloric acid18 and then hydrothermally treating the reaction mixture at around 170 °C; this resulted in a poorly crystalline material. The use of readily oxidized ions produced precipitates; for example, KMnO4 reacted immediately with KI giving an amorphous compound, which gave a crystalline phase after heating to 600 °C for 4 days in air (no hydrothermal treatment). NMe4Cl behaved similarly. Therefore, manganese oxides were synthesized by acidifying 0.01 mol of KMnO4 from Fisher Chemicals in 20 mL of water with 2 drops of 4 M HNO3 giving a pH of around 3.5 and a Mn:H+ ratio of 100:1. The resulting solution was transferred to a 125-mL Teflon-lined autoclave (Parr bomb), sealed, and reacted hydrothermally for 4 days at 170 °C. The resulting black crystals were filtered, and dried at 45 °C in air. The pH of the solution after reaction was alkaline, pH 12-13. X-ray powder diffraction was performed using Cu KR radiation on a Scintag θ-θ diffractometer. The data were collected from 10° 2θ to 70° 2θ with 0.03° 2θ steps and 15 s/step. The TGA data were obtained on a Perkin-Elmer Model TGA 7, the electron microprobe on a JEOL 8900, and the chemical analysis was obtained on an ARL Spectrospan-7 DCP atomic emission spectrometer. The degree of reduction of the manganese oxide by lithium was determined by reaction with n-butyllithium from Aldrich Chemicals, following standard procedures.26,27 Electrochemical studies were conducted in lithium cells using lithium perchlorate in dioxolane as the electrolyte; a MacPile potentiostat was used to cycle the cells. The manganate was hot-pressed, at ≈200 °C, with 10 wt % Teflon powder and 10 wt % carbon (23) Sposito, G. The Chemistry of Soils; Oxford University Press: New York, 1989. (24) Golden, D. C.; Chen, C. C.; Dixon, J. B. Clays Clay Miner. 1987, 35, 271. (25) Whittingham, M. S. Prog. Solid State Chem. 1978, 12, 41. (26) Dines, M. B. Mater. Res. Bull. 1975, 10, 287. (27) Whittingham, M. S.; Dines, M. B. J. Electrochem. Soc. 1977, 124, 1387. (28) Chen, R.; Chirayil, T.; Zavalij, P.; Whittingham, M. S. Proceedings of the 10th International Symposium on Solid State Ionics, Singapore, Dec 1995; Solid State Ionics, in press.
Chen et al. black into a stainless steel Exmet grid, and discharged at 0.1 ma/cm2.
Results and Discussion Our initial experimental studies involved reacting potassium permanganate with hydrochloric acid and then hydrothermally treating the reaction mixture. This gave poorly crystalline material. Similarly, the use of readily oxidized ions such as I- and NMe4+ led to amorphous materials that had to be heated to around 600 °C to promote crystallization. We therefore decided to place emphasis on reactants that did not show significant reaction at room temperature and would crystallize out only under the hydrothermal conditions. Optimal crystallinity was found when using potassium permanganate alone; it was dissolved in water and the pH was adjusted to around 3.5 by the addition of 1 or 2 drops of nitric acid. The ratio of H+ to Mn is around 1:100, so that there is insufficient acid for bulk reaction with the permanganate. A black crystalline solid was precipitated out after hydrothermal treatment. The nitric acid was found to accelerate the reaction, and in the case of NaMnO4 the reaction did not go to completion in the absence of nitric acid.28 The yield of product based on manganese approached 100%. Consistent with this yield analysis, the basic liquid product contained potassium species (as potassium carbonate due to reaction with CO2 in the air) and no manganese species; there was no indication of the formation of K2MnO4 as noted in the dry thermal decomposition of KMnO4.29 The overall reaction is consistent with the equation
KMnO4 + (1 - x + 2y)/2H2O f KxMnO2‚yH2O + (1 - x)KOH + (3 + x)/4O2 Thus, we can conclude that potassium permanganate decomposes in aqueous solution at 170 °C to give a black crystalline solid. DCP analysis of the potassium sample showed 0.25 K/Mn. Unlike most of our earlier hydrothermal studies, the SEM images showed no particular evidence for a layered structure, consisting of small spheres of diameter 2-5 µm. The SEM analysis showed the presence of manganese, oxygen, and potassium. Thermogravimetric analysis of the potassium manganese compound under oxygen (Figure 1) shows a 10% weight decrease to 150 °C, indicating the loss of 0.55 H2O/100 g of material, i.e., K0.25MnO2‚0.6H2O. The second weight loss between 300 and 400 °C may possibly be associated with loss of oxygen from the lattice associated with a partial reduction of Mn(IV) to Mn(III). Delmas et al.30 reported that a monoclinically distorted P′3 phase K0.5MnO2 prepared at 450-700 °C takes up 0.12 oxygen on oxidation at 400 °C converting to the hexagonal P3 phase K0.47Mn0.94O2 (manganese vacancies based on density measurements) with a ) 2.88 Å and c ) 19.00 Å. Herbstein et al.29 also reported a weight loss of 3.7 wt % when “K4Mn7O16” is heated at around 500 °C, which they associated with loss of oxygen. After thermal analysis to 600 °C a black compound was found, and its X-ray diffraction pattern showed no difference from the starting material, indicating that (29) Herbstein, F. H.; Ron, G.; Weissman, A. J. Chem. Soc. A 1971, 1821. (30) Delmas, C.; Fouassier, C. Z. Anorg. Allg. Chem. 1976, 420, 184.
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Figure 1. Thermal gravimetric analysis of the manganese oxide phase in oxygen at a heating rate of 3 °C/min: (a) first heating cycle; (b) second heating cycle after exposure of the sample to the atmosphere at room temperature; (c) second heating cycle after sample kept in dry oxygen between the two heating cycles.
the structural water was regained in the time taken to transfer the sample from the TGA to the X-ray diffractometer. This was still the case, even when heated to over 900 °C at 3 °C/min. However, prolonged heating (28 h) of the sample at 600 °C resulted in the formation of the oxide tunnel framework found in hollandite, cryptomelane, and manjiroite. Gently heating the sample to around 160 °C and maintaining it in a dry environment, thus removing the structural water, led to a contraction of the c lattice repeat distance as shown by Figure 2a,b. The water lost is rapidly reabsorbed on exposure to air, so that the sample had to be X-rayed under a mylar film. The X-ray pattern after subsequent water uptake is shown in Figure 2c and is essentially identical with Figure 2a showing the reversibility of the hydration reaction. If the sample was not exposed to any moisture-containing atmosphere after the first heating cycle, then no weight loss was observed on the second heating cycle, as indicated in Figure 1c. Powder X-ray diffractograms (Figure 2) of the asprepared potassium manganese compound indicated the presence of just one phase, a layered compound of repeat distance 7.0 Å. The dehydrated sample of Figure 2b shows a contraction from the 7.0 to 6.44 Å. The pattern of the hydrated sample could be indexed on a simple three-block cell with hexagonal lattice parameters a ) 2.849(8) Å and c ) 21.536(7) Å;31 the data are listed in Table 1 and with all reflections fitting the criteria -h + k + l ) 3n the structure is rhombohedral with space group R3 h m. The structure of this compound appears to be the same as that of Lix(H2O)TiS2, which has lattice parameters of a ) 3.430 Å and c ) 25.98 Å (3 × 8.66). Just as in that case,32 the mixed reflections 0 1 3n + 2 are stronger than the 1 0 3n + 1 reflections. This indicates that the manganese are octahedrally coordinated and that the oxygens in the layers provide trigonal prismatic sites for the waters in the interlayer region; thus the stacking arrangement is AbC(water)CaB(water)BcA(water), where ABC represents the positions of the oxygen atoms and abc the positions of the (31) The lattice parameters were found to vary from sample to sample; for example another sample had a ) 2.853(6) Å and c ) 21.01(1) Å (3 × 7.0); this is probably associated with a variation of the relative humidity. (32) Whittingham, M. S. Mater. Res. Bull. 1974, 9, 1681.
Figure 2. X-ray diffraction pattern of the manganese oxide phase, using Cu KR radiation: (a) as-prepared material; (b) after heating to 160 °C (under mylar); (c) after reabsorption of water from the air. Table 1. X-ray Data for KxMnO2‚nH2Oa h
k
l
d(obs)
d(calc)
I(obs)
I(calc)
0 0 1 0 1 0 1 0 0 1 0 1 1
0 0 1 1 0 1 0 1 0 0 1 1 1
3 6 1 2 4 5 7 8 12 10 11 0 3
7.1668 3.5872 2.4524 2.4043 2.2429 2.1413 1.9228 1.8204 1.7964 1.6220 1.5343 1.4240 1.3982
7.1798 3.5896 2.4514 2.4051 2.2431 2.1411 1.9249 1.8190 1.7948 1.6226 1.5337 1.4246 1.3973
1000 338 77 109 1 219 30 113 28 34 40 64 91
999 335 79 112 13 222 32 119 11 25 68 44 84
a
a ) 2.849(8), c ) 21.536(7) Å.
manganese cations. The potassium ions and water molecules share the interlayer space as shown in Figure 3b, just as in the layered disulfides;32 the composition is consistent with each potassium being surrounded by six water molecules in a plane as in (NH4)0.33‚(H2O)0.67TaS2.32 A recent report on sol-gel formed birnessite pictures two water layers22,33 even though the lattice repeat distance is around 7 Å, comparable to that reported here. (33) Guzman, R. N. D.; Awaluddin, A.; Shen, Y.-F.; Tian, Z. R.; Suib, S. L.; Ching, S.; O’Young, C.-L. Chem. Mater. 1995, 7, 1286.
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Chen et al. Table 2. Crystallographic Parameters for K0.27MnO2‚0.54H2O formula space group a c cell vol calc density diffractometer radiation 2θmax no. of reflns 3 Mn in B(Mn) 6 O in z(O) B(O) 0.81K + 1.62H2O in B(K, H2O)
Figure 3. Schematic of layered manganates, showing (a) anhydrous material, (b) one water layer (birnessite), and (c) two water layers (buserite). (gray-shaded circles) ) H2O; (b) ) M ion.
To confirm this structure, an integrated intensity refinement was carried out using the CSD crystallographic package. The powder pattern of the potassium compound was indexed in the trigonal system, space group R3 h m, with hexagonal axes a ) 2.849(8), c ) 21.536(7) Å. The starting model of the MnO2 layer was taken from KTiS2 (3 Mn in 3a 0 0 0 and 6 O in 6c 0 0 0.039(3)) and refined by using integrated intensities. On the differential Fourier map an additional peak was found in position 9d 1/2 0 1/2 between MnO2 layers. Its occupancy was refined as K and H2O mixing with a fixed ratio of 1:2. This atom is placed not in the center of trigonal prism but on its face. Refinement with the atom in the center of the prism leads to zero occupancy. The resulting refined composition is K0.27MnO2‚0.54H2O which agrees well with the chemical analysis. Table 1 compares the observed and calculated parameters, and Table 2 the crystallographic parameters for K0.27MnO2‚0.54H2O. The water-free compound, shown in Figure 3a, has been prepared at high temperatures. Thus, KxMnO2, for 0.45 e x e 0.67 has probably30,34 the same structure as proposed here for K0.25(H2O)0.55MnO2 but with the potassium ions in trigonal prismatic sites. However, K0.25MnO2 has been prepared at high temperatures and then has the hollandite structure,35 just as found here after heating of the layer phase for 28 h. Na0.2MnO2 also has the hollandite structure, but as the sodium content increases lamellar phases are formed, and in the Na0.70MnO2+y phase the sodium ions are in trigonal (34) Delmas, C.; Fouassier, C.; Hagenmuller, P. Physica 1980, 99B, 81. (35) Parant, J.-P.; Olazcuaga, R.; Delalette, M.; Fouassier, C.; Hagenmuller, P. J. Solid State Chem. 1971, 3, 1.
K0.27MnO2‚0.54H2O R3h m 2.8490(8) Å 21.536(7) Å 151.2 Å3 3.53 g/cm3 Scintag XDS 2000 Cu KR 70° 29 3a (0 0 0) 0.7(9) Å2 6c (0 0 z) 0.039(3) 0.9(9) Å2 9d (1/2 0 1/2) 5 Å2
prismatic sites.36 Similar phases have been reported for KxCrO2, 0.50 e x e 0.60, where a ) 2.918 Å and c ) 18.44 Å37 and for KxCoO2, x ) 0.5, where a ) 2.829 Å and c ) 18.46 Å (3 × 6.15).38 The latter was reported to absorb moisture with an expansion of the lattice, though no data was reported. Most fully “intercalated” oxides, such as LiCoO2, where a ) 2.816 Å and c ) 14.08 Å (3 × 4.69), have the alkali metal in octahedral sites but otherwise have the same Figure 3a structure. Just as reported for the transition-metal disulfides,32,39,40 the maximum alkali content of these oxides in aqueous environment may be constrained by the evolution of hydrogen which will occur at the higher thermodynamic potentials. Indeed LeGoff et al.21 showed that Na0.7MnO2.14 lost sodium when treated with water to give Na0.45MnO2.14‚0.76H2O. This synthesis approach was extended to the sodium compound, and as noted above nitric acid was required to drive the reaction to completion. X-ray analysis of this compound showed a repeat distance of 7.28 Å which contracted to 5.6 Å on dehydration. This compares with the value of 5.5 Å reported for Na0.7MnO2.14.21,35 Further details of the sodium compound will be discussed elsewhere.28 Structure Comparison. There is much confusion and disagreement in the literature as to the structures of these layered manganates. Delmas et al.30 were the first to report that there are two forms of the layered manganates that differ from one another in the manganese to oxygen ratio. The more oxygen-rich phase, K0.47Mn0.94O2, stable at lower temperatures, has a hexagonal structure, whereas the phase formed at 450700 °C has a monoclinically distorted structure; at 400 °C the latter takes up 0.12 oxygen atoms converting to the former. Natural manganates, known as birnessite, have recently been reported to be hexagonal while synthetic samples are monoclinic.41 Sodium manganate, birnessite, had been assigned42 a large orthorhombic structure (a ) 8.54 Å, b ) 15.39 (36) Mendiboure, A.; Delmas, C.; Hagenmuller, P. J. Solid State Chem. 1985, 57, 323. (37) Delmas, C.; Devalette, M.; Fouassier, C.; Hagenmuller, P. Mater. Res. Bull. 1975, 10, 393. (38) Delmas, C.; Fouassier, C.; Hagenmuller, P. J. Solid State Chem. 1975, 13, 165. (39) Scho¨llhorn, R.; Weiss, A. Z. Naturfo¨ rsch. 1973, 28b, 711. (40) Lerf, A.; Scho¨llhorn, R. Inorg. Chem. 1977, 16, 2950. (41) Manceau, A.; Gorshihkov, A. I.; Drits, V. A. Am. Miner. 1992, 77, 1144. (42) Giovanoli, R.; Sta¨hli, E.; Feitknecht, W. Helv. Chim. Acta 1970, 53, 209.
Hydrothermal Synthesis and Characterization of KxMnO2‚yH2O
Å, c ) 14.26 Å JCPDS 23-1046) modeled from electron diffraction studies43 on that of chalcophanite,44 which is still often referenced.22,45-47 This model suggests45 that one-seventh of all the manganese is in the interlayer region forming a Frenkel defect, that is, it takes the place of the zinc in chalcophanite, ZnMn3O7 or Zn2/7Mn6/7O2 ≡ MnIII1/7MnIV6/7O2. There is no conclusive evidence at present for these defects, and the ready expansion of the lattice by long-chain alkylamines45 probably speaks against the presence of pinning ions such as manganese. Wadsley48 noted that some birnessites such as the sodium form can incorporate a second water layer expanding to a 10 Å repeat distance; it is this expanded lattice that can be further expanded by alkylamines. An inspection of the literature, as pointed out by Delmas et al.,30 indicates that many of the reported X-ray patterns can be indexed to a simple hexagonal structure. De Carvello49 reported a potassium manganate which is probably a hydrate,30 all of whose reflections can be indexed to a hexagonal cell with a ) 2.85 Å and c ) 20.6 Å.50 The seven reported lines of the potassium manganate prepared by sucrose reduction of permanganate22 can also be indexed to a hexagonal cell, a ) 2.84 ( 0.02 Å and c ) 20.9 ( 0.3 Å. Herbstein et al.29 reported the formation of the compound K4Mn7O16 as an intermediate in the thermal decomposition of potassium permanganate. Although their published X-ray figure appears similar to that in Figure 2, their tabulated data have many more lines and fits well to a tetragonal unit cell as they propose. They noted that it and its hydrate, “K4Mn7O16‚5H2O” have different (but unknown) crystal structures; the hydrate loses water rapidly and reversibly on heating above 100 °C. It is possible that this phase, which can be formulated as K0.5Mn0.875O2‚0.625H2O, is similar to or the same as the phases reported here and by Delmas30 with a contaminating second phase. Post and Veblen51 revisited the structure of synthetic manganates and found a monoclinic lattice with a ) 5.149 Å, b ) 2.843 Å, c ) 7.176 Å, and β ) 100.76° with essentially zero vacancies on the manganese lattice; this was recently confirmed in two other studies.41,52 These studies also reported that the ions sodium, potassium, and magnesium in the interlayer region ordered, resulting in a superstructure. However, these electron diffraction studies caused the loss of water and the collapse of the lattice,51 so it cannot be concluded that such ordering occurs when water is present. One would expect that the hydrated alkali ions are rather mobile (43) Burns, R. G.; Burns, V. M. Philos. Trans. R. Soc. London A 1977, 286, 283. (44) Wadsley, A. D. Acta Crystallogr. 1955, 8, 165. (45) Golden, D. C.; Dixon, J. B.; Chen, C. C. Clays Clay Miner. 1986, 5, 511. (46) Chen, C.-C.; Golden, D. C., Dixon, J. B. Clays Clay Miner. 1986, 5, 565. (47) Pereira-Ramos, J. P.; Baddour, R.; Bach, S.; Baffier, N. Solid State Ionics 1992, 53-56, 701. (48) Wadsley, A. D. J. Am. Chem. Soc. 1950, 72, 1782. (49) DeCarvalho, A. J. C. J. Appl. Chem. 1957, 7, 145. (50) The compounds in Table 2 (a ) 2.84 ( 0.01 Å and c ) 20.49 ( 0.07 Å) and Table 7 (a ) 2.86 ( 0.01 Å and c ) 20.8 ( 0.2 Å) of ref 49 are probably the same; these hexagonal parameters give a better fit than the proposed tetragonal unit cell, a ) 7.05 ( 0.05 Å and c ) 7.0 ( 0.1 Å. (51) Post, J. E.; Veblen, D. R. Am. Miner. 1990, 75, 477. (52) Kuma, K.; Usui, A.; Paplawsky, W.; Gedulin, B.; Arrhenius, G. Miner. Mag. 1994, 58, 425.
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Figure 4. Electrochemical reduction of the manganese oxide in a lithium cell; the lithium metal anode was also used as the reference electrode.
and therefore would not order on the time scale of the diffraction experiment. As noted by Leroux et al.18 the ranciete form of manganese dioxide has a larger spacing of 7.4 Å. The precise differences between the phases formed here (birnessite-type) and ranciete is not clear. However, it may be possible that the ranciete formed by acid hydrolysis contains manganese between the oxide layers and therefore has a larger deficit of manganese in the layers leading to a different charge on the layers. These interlayer manganese are then exchanged for alkali ions such as potassium. An extended study of the defect structure of these manganates will be required to resolve the precise structure of the hexagonal and rhombohedral forms and of the difference between birnessite and ranciete. Redox Behavior of Layered Manganates. Reaction of dehydrated potassium manganese oxide with n-butyllithium showed the uptake of 1.0 Li/Mn. There have been no prior reports of butyl lithiation of the potassium compound; however, the 7.3 Å hydrate compound MnO1.85‚nH2O reacts with lithium with a gradual decrease in the c lattice repeat distance; only lithium contents up to 0.45/Mn were studied.53 The electrochemical intercalation of lithium into the dehydrated potassium manganese oxide lattice is shown in Figure 4. This shows a continuous curve indicative of single-phase behavior, LixK0.25MnO2, where 0 e x < 0.6, similar to that observed in the layered disulfides.54 The lower value of x ) 0.6 observed here than the x ) 1 in the n-butyllithium reaction is related to the 2 V lower limit used for the electrochemical reduction, compared to the effective about 1 V reduction potential for butyllithium.27 The curve of the electrochemical intercalation of lithium into LixK0.25MnO2 also suggests that the oxidation state of the manganese is around 3.5, as the cell open circuit voltage is 3.5 V, close to the value reported for lithium insertion into LiMn2O4.14 The cycling behavior of this hydrothermally formed manganate is now underway. There have been a few studies of the use of manganates in lithium cells, but most have contained water, which might be expected to degrade the cyclability of the cell. Lecas et al.55 have reported on the cycling (53) LeGoff, P.; Baffier, N.; Bach, S.; Pereira-Ramos, J.-P. J. Mater. Chem. 1994, 4, 133. (54) Whittingham, M. S. Science 1976, 192, 1126. (55) Lecas, F.; Rohs, S.; Anne, M.; Strobel, P. J. Power Sources 1995, 54, 319.
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behavior of phyllomanganates in lithium batteries and found a gradual degradation caused by conversion of the active material to cubic spinel. Leroux et al.56 studied the cycling behavior of the ranciete-type H, Li, and Na manganic salts and found that the interlayer water played a key role, with dehydrated materials performing the best. Le Cras et al.57 have found rather poor reversibility for the related lithium and sodium phyllomanganates in lithium cells. Conclusion This paper reports the direct synthesis, under mild hydrothermal conditions, of a layered manganese dioxide from aqueous potassium permanganate. Its structure has MnO2 layers between which lie potassium ions and water molecules; on heating, the water is reversibly removed and lithium ions are readily incorporated into (56) Leroux, F.; Guyomard, D.; Piffard, Y. Solid State Ionics 1995, 80, 307. (57) Cras, F. L.; Rohs, S.; Anne, M.; Strobel, P. J. Power Sources 1995, 54, 319.
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the structure both chemically and electrochemically. This compound may thus be of interest as the cathode in secondary lithium batteries. The structure and chemistry of the MxMnO2‚yH2O phases is similar to the corresponding well-characterized MxTiS2‚yH2O phases. Mild hydrothermal reactions appear to be a general useful method for the synthesis of transition-metal oxides. In many cases, new structures are formed as recently noted in the case of tungsten, molybdenum, and vanadium.1,2,8,10 Extensive research is now underway to explore the possibilities of this approach to new materials, and has recently been reviewed.13 Acknowledgment. We thank the Department of Energy through Lawrence Berkeley Laboratory and the National Science Foundation through Grant DMR9422667 for partial support of this work. We also thank Professor Richard Naslund for the use of the DCP atomic emission spectrometer, Bill Blackburn for electron microprobe assistance, and Sean Kelly for some initial experiments in this area. CM950550+
