摘要:

J. MATER. CHEM., 1994. 4( 1 ), 113-1 18 Vanadium Phosphate Glasses. Effect of Composition on their Structure and Performance as Cathodes in High-temperature Lithium Polymer-electrolyte Cells Andrew Davies: Richard J. Hobson: Michael J. Hudson: William J. Macklinb and Robin J. Neat!' a Department of Chemistry, University of Reading, Whiteknights, Reading, UK RG6 2AD Applied Electrochemistry Department, A€A Industrial Technology, Harwell Laboratory, Oxfordshire, UK OX77 ORA Five vanadium phosphate (v@,-P@,) glasses, containing between 58 and 88 mol% V205, have been prepared using the melt-quenching technique. IR spectroscopy, in conjunction with powder X-ray diffraction (XRD) of the devitrified materials, has been used to show that the glasses containing d70molo/o V205 have a microstructure which is similar to that of B-VPO, whereas the remainder show both orthorhombic-V,O, and p-VPO, structural features. Molar-volume data suggest that there is a fairly abrupt change in microstructure at ca.75 mol% V,05. The glasses have been employed as the active component of the cathode in lithium polymer-electrolyte cells operating at 120 "C and the cycling performance has been investigated as a function of glass composition. After an initial capacity decline the cells showed good reversibility, although this was not achieved at capacities as high as that of V,O,, in a similar cell. Glasses containing less than 78 mol% V205 cycled at considerably lower capacities than those with higher vanadium content and this has been related to the above change in microstructure from a layered V,O,-type to a network p-VPO, type.The electrochemical insertion of lithium into vanadium (v) oxide containing glasses has been studied by a number of researchers. Many of these studies involved ternary materials with a glass-forming oxide also present, such as the systems V205-P205,1p4 V205-B203,5 and V205-Te0,.6 The V205-P205 system, which may be prepared using a moder- ately fast quenching technique ( 102-103 "C s-'), has shown some interesting features. Lithium-ion insertion into a glass of this type containing 60 mol% V205 was first demonstrated by Pagnier et al. in 1983.' Unfortunately, when tested in a lithium polymer-electrolyte cell operating at 85 "C the glass exhibited poor reversibility which was attributed to the elec- trolyte becoming oxidized near the upper voltage limit.However, a cathode consisting of a single solid piece of glass rather than an intimately mixed composite cathode was employed, which would in any case have led to a poor cell performance regardless of the cathode material because inti- mate contact between the cathode active material and the electrolyte is critical in determining cell perf~rmance.~ Reversible insertion of Li' ions in V205-P@5 glasses was first reported in ambient-temperature liquid-electrolyte cells by Sakurai and Yamaki? who later demonstrated that the glassy materials were superior to crystalline V205 in terms of long-term re~ersibility.~ Reversibility over several hundred cycles was reported for all glass compositions.A relationship between microstructure and composition has been shown for V205-P205 gla~ses.~.~Bhargava and Condrate; on the basis of IR spectral evidence, suggested that the short-range order of the glasses, particularly those with a high P205content (up to 50mol0/0), resembles that of c(-VP05. Sakurai and Yamaki,4 however, proposed that glasses containing <75 mol% V205have a structure which resembles that of the p-VP05 network, whereas above this boundary a change to a 'V205-like' structure occurs. These authors found a significant difference in electrochemical properties either side of this boundary. The good long-term reversibility reported for vanadium phosphate glasses in room- temperature liquid-electrol yte cells24 has led us to reinvestigate the rechargeability of these materials in lithium polymer-electrolyte cells operating at 120"C.The percentage theoretical capacities and reversibility of the glasses are compared with those of V,OI3 for the first time; to warrant further study a material would have to show a significant improvement in reversible capacity and/or energy density over V6OI3. In addition, we have re-examined the IR spectra of the glasses and compared these with the spectra of their crystalline devitrification products, for which we were able to obtain definitive structural information. Experimental Synthesis The vanadium phosphate glasses were prepared from in timate mixtures of powdered V205 and P,O,.Mixtures containing initially 55, 65, 75, 80, and 85 mol% V,O, were employed. These were prepared in a nitrogen-filled dry box and then heated in air to 750°C in a platinum crucible for 1 h. The melt was quenched by pouring it onto a pre-cooled (5°C) stainless-steel plate. Devitrified samples were prepared by heating the glasses at 100"C above their crystallization tem- peratures for 2 h. Techniques Scanning electron micrographs were recorded using a JEOL JXA 840 scanning electron microscope. The vanadium :phos-phorus molar ratios of the glasses were established using electron microprobe X-ray analysis on the same instrument. Reduced vanadium [V"] and total vanadium were deter- mined by potassium permanganate titration and atomic absorption spectroscopy (AAS), respectively.Densities were obtained by pycnometry using toluene or cyclohexane. Electronic conductivities were measured on solid pieces of glass, cut using a diamond stylus, of area =1 cm2 and thickness 2.5 1.0mm (measured using a micrometer). Conducting adhesive copper tape was attached to cover both faces and the resistance through the thickness of the sample was meas- ured at 25°C. Powder XRD patterns were recorded using graphite monochromated Cu-Kcc radiation on a Philips PW1710 diffractometer controlled by a Citrons Cray 112 system run- ning Sietronics 112 software. Data were collected for the 28 range 4-64" at a rate of 2" min-' in steps of 0.04". Differential thermal analyses were carried out using a Stanton Redcroft STAlOOO instrument equipped with data manipulation software using a heating rate of 10"C min-'.IR spectra were measured between 1400 and 400 cm-' as KBr discs using a Perkin-Elmer 1720-X FT-IR spectrometer. Cell Fabrication and Cycling Composite cathodes containing ground vanadium phosphate glass (45 vol.%, particle size <50 pm), ketjenblack carbon (5 vol.%), poly(ethy1ene oxide) (PEO) (Union Carbide MW 4 000 000) and LiC104 (Aldrich) (50 vol.% PEO-LiC104, CEO units]," Li] = 12) were prepared via doctor blade casting from the appropriate solvent slurry onto a nickel current collector. Sheets of the electrolyte PEO-LiC104 (CEO units]/[Li] =12) were cast from acetonitrile solution onto silicone release paper.Solid-state cells, Fig. 1, incorporating a lithium foil anode (Lithco 150 pm) with an active area of 40 cm2 were constructed in a dry room (T=20 "C, dew-point temperature -30 "C) using a combination of heat and pressure. The cells had a capacity, C, of ca. 30mA h. This was calculated using the theoretical capacity of V6013, 417 mA h 8-l for Li,V,O,,, to enable a direct comparison between the performance of the glasses and V601,. Cell cycling of the packaged cells was performed galvanostatically under computer control between limits of 3.5 and 1.7 V at a rate of C/10 and a temperature of 120°C. Prior to cycling the cells were allowed to equilibrate at 120 "C for 2 h when their open- circuit voltages were measured.For each cathode material two cells were cycled, typically 30 times, in order to establish the reproducibility of the results. Results and Discussion The glasses prepared were hard and smooth with no visible crystallites. Their amorphous nature was confirmed using powder XRD, which showed no Bragg peaks, and scanning electron microscopy (SEM). The electron micrographs of fractured samples showed no crystallites (Fig. 2) at any magni- fication up to 2800, at which crystallites as small as 5 x m should be clearly visible. Micrographs of the powdered samples employed in the fabrication of cells showed particles of random size (<50 pm) and shape with no regular faces. The amount of V205 found in the glasses was in all cases slightly larger than the initial content of the mixtures, Table 1, owing to the sublimation of a small amount of P,O, prior to 45 vol% V2O5-P2O5glass Li / PEOl2:LiC1O4/ 5 vol% carbon / Ni 50 vol% PEOl2:LiCIO4 Fig. 1 Lithium polymer-electrolyte cell configuration J.MATER. CHEM., 1994, VOL. 4 Fig. micrograph Of a powdered Of 82 molo/o V,O, glass Table 1 Analytical data of glasses V,O, (mol%) initial found V'"/V,,, density/g cmP3 TJC lJT 55 58 0.19 2.78 390 572 65 70 0.17 2.79 295 410 75 78 0.12 2.84 266 403 80 82 0.09 2.90 241 326 85 88 0.07 3.OO 235 275 melting. The figures given are based on the vanadium :phos-phorus molar ratios determined using electron microprobe analysis and were confirmed by the AAS analyses for total vanadium.All of the glasses were dark blue and contained some vanadium(Iv), see Table 1, produced during the prep- arations with concomitant loss of oxygen. (No correction for this has been applied to the mol% of V205.) The proportion and amount of V" decreases with increasing vanadium con- tent. The measured densities increase with vanadium content as would be expected. Differential thermal analyses of all of the glasses showed an endothermic discontinuity in the baseline indicating the glass-transition temperature,' q,followed by an exotherm due to crystallization of the glass, the onset of which' gives the crystallization temperature, T,. The data are summarized in Table 1. As expected, both Tp and T, increase with the amount of the better glass-forming component, i.e.P,Os. The measured electronic conductivities, which increased with increasing V205 content, ranged from 7.9 x to 3.2 x lop5S cm-l, of the same order as those determined previously" for similar materials at ambient temperature. The conductivities of the glasses would be expected" to be higher at the cell operating temperature of 120 "C. Microstructures of Glasses The microstructures of the glasses would be expected to have a profound influence on their performance as cathode active materials in lithium batteries. Amorphous materials contain- ing solely vanadium(v) oxide have microstructures which are related" to that of orthorhombic V2OS and therefore have sites which are suitable for lithium insertion and V-0-V linkages which, together with the presence of some VIV, are necessary for electronic conduction.The vanadium phosphate glasses, however, are likely to have microstructures based on VPO,, which is dimorphic. Previous investigator^'^^^^ have demonstrated the presence of VO, and PO4 moieties in the glasses and on the basis of IR spectral studies, the glasses J. MATER. CHEM., 1994, VOL. 4 have been assigned microstructures related to r-VPO, by one group' and p-VPO, by a second,, (see earlier). VPO, crystallizes in the tetrag~nal'~ (a) or orthorh~mbic'~ (p)systems. Both modifications contain highly distorted VO, octahedra which are linked into chains via corner sharing along (001)(a)or (100) (p).However, the V-0, bonds linking the oc!ahedra are unequal (a, 1.580 and 2.857 A; p, 1.566 and 2.591 A) and a better description of the structures can be obtained in terms of chains of weakly linked VO, square-based pyramids, the four basal oxygen atoms of which are corner-shared with four PO, tetrahedra. The PO4 tetrahedra in a-VPO, link four chains of VO, units giving rise to layers lying parallel to (001) with only weak links between them.Thus there are sites which may be suitable for lithium-ion insertion, similar to those16 in orthorhombic V205. However, the PO4 tetrahedra in p-VPO, link three chains of VO, units and bridge two adjacent VO, moieties in the same chain, giving rise to a network. This bridging is likely to hinder the free movement of lithium ions.Both structures lack symmetric V-0-V bridges so a glass with a microstructure based on either form would be expected to have a poorer electronic conductivity than one with a V,O,-like microstructure. Molar Volume Drake17 et al. studied a series of vanadium phosphate glasses containing between 47 and 74 mol% V205.They showed that the molar volume per gram atom of oxygen, V2;, showed a 'monotonic and quasi-linear' change with composition, indi- cating that there was no phase separation or structural change in their composition range. Sakurai and Yamaki4 examined glasses containing between 58 and 95 mol% V205 and sug- gested that an inflection in the Vg us. composition plot at ca. 75 mol% V205 indicated a change in microstructure in this region.We calculated Vg using the expression where x =molar fraction, M =relative molar or atomic mass, r =molar ratio VIV:Vtot,and p =density. Similar values to those of Sakurai and Yamaki4 were obtained. However, rather than a point of inflection in the V2; us. composition plot, a maximum was observed between 70 and 78 mol% V205, see Fig. 3, confirming that a change in microstructure or amorph- ous phase separation occurs in this composition range. The latter interpretation is more consistent with our IR evidence given later. The differences between our results and those of Sakurai and Yamaki can be attributed largely to significant 7' 13.4 lP4t A T W W 2.8 b10.8. Fig. 3 Dependence of (0)V;S (see text) and (H)density on composi-tion.The 100mol% data are for crystalline V,O,. The smoothed lines were drawn by fitting a third-degree polynomial piece-wise to the data. differences in the measured densities, see Fig. 3, to which VT, is particularly sensitive. The inclusion of the data for crystalline V205 shows that it fits well our observed trends. Powder X-Ray DifSraction of the Deuitrifed Glasses The diffraction patterns of the devitrified glasses, examples of which are shown in Fig. 4, demonstrate clearly that the products of devitrification are p-VPO, and orthorhombic- V205. The devitrified 58 mol% V,O, glass gave a diffraction pattern which is an excellent match with the calc~lated'~~'' diffraction pattern of p-VPO,, with the exception of two weak peaks at 20=23.22 and 28.34' which do not correspond with either a-VPO, or orthorhombic V205.This match is in contrast with the result of Sakurai and Yamaki4 who obtained an unidentified crystalline material upon devitrification of a similar glass. The devitrified 70 mol% glass gave a very similar diffraction pattern to the 58 mol% material. In both of these some vanadium atoms must presumably occupy some of the tetrahedral phosphorus sites. Such a suggestion' with respect to the glasses has been made previously on the basis of a low angle X-ray scattering study. The other devitrified glasses all gave diffraction patterns which showed a mixture of V20, and p-VPO,, Fig. 4. We found no evidence for the existence of solid solutions in the crystalline V205-P,O, system :is has been reported previ~usly.~ IR Spectra of Glasses and Devitrifed Glasses The IR spectra of the glasses and their devitrification products, for which definitive structural information had been obtained, were studied in order to resolve the differences in the interpret- ation4q8 of their microstructure.The spectra of the glasses showed bands which were substantially broader than those of the crystalline materials, as shown in Fig. 5. This is consist- ent with structures lacking long-range order that have a range of bond lengths and strengths. However, the general features 10 20 30 40 50 60 28ldegrees Fig. 4 Powder XRD patterns of (a)orthorhombic V20,, (b)delritrified 82 mol% V20, glass, (c)devitrified 58 mol% V,O, glass and (d) the pattern calculated for p-VPO, using Lazy Pulverix.18 1200 800 4 wavenumberkm-' Fig.5 IR spectra, (a) orthorhombic V20,; (b)88 mol% V205 devitri-fied material; (c) 88 mol% V205 glass; (d) 58 mol% V20, devitrified material; (e)58 mol% V20, glass of the IR spectra of the vitreous and corresponding non-vitreous materials are similar. Between 1250 and 1050 cm-' the spectra of the glasses show a strong, broad feature and those of the devitrified materials two strong bands which are characteristic' of p-VPO, (1052 and 1149 cm-I). a-VPO, has no bands in this region but instead shows a weak band' at 1211 cm-I absent in all our spectra. The V=O stretching vibration in crystalline V205, p-VPO, and a-VPO, is at 1029, 1000, and 990 cm-', respectively,' and is also evident in the vitreous (1008-1021 cm-') and non-vitreous (1000-1010 cm-') materials.In the region 950-800 cm-' both V205 and p-VPO, show a single, strong absorption (828 and 942 cm- ', respectively) whereas a-VPO, has only very weak features.' The three devitrified materials with higher vanadium content display both bands, whereas the other two show only the 940cm-' band. The maxima of the broad absorption envelopes are at 840 cm-' for the three glasses containing 378 mol% V205 and 920 and 928 cm-' for the 70 and 58 mol% glasses, respectively. This is consistent' with the loss of V-0-V linkages as the proportion of P205increases. Between 800 and 400cm-' the three devitrified materials with higher vanadium content show the two intense bands of orthorhombic V205 (602 and 478 cm-') whereas the 58 and 70mol% materials display the five bands' of p-VPO,.The glasses with >,78 mol% V205show a strong band at 640 cm- and a broad indistinct band around 400 cm-' whereas the 70 and 58 mol% glasses both show a medium intensity shoulder at 780cm-' and a medium or low intensity band around 630 cm-'. The broad absorption envelopes of these glasses, with the exception of the shoulder at 780cm-', encompass the bands of the devitrified materials, see Fig. 5. The superficial similarity between the spectra of the low vanadium glasses and a-VPO, in this region (a-VPO, has two bands: 780 crn-', very weak; 602 cm-', sharp, medium) does not, we believe, show that they have an a-VP0,-like microstructure.' These bands are difficult to assign specifically to an isolated group vibration' and because of the lack of long-range periodicity and differences in local symmetry the vibrational modes of J.MATER. CHEM., 1994, VOL. 4 the glasses are likely to differ from those of the crystalline materials. The considerable similarities between the TR spectra of the glasses and the equivalent non-vitreous materials strongly suggest that the components present in the latter reflect the microstructural domains present in the glasses. Thus the glasses containing <70 mol% V205 have microstructures which are predominantly p-VPO,-like, whereas the glasses containing 378 mol% V205 have domains which are V20,- like and domains which are p-VPO,-like.Without further evidence it is not possible to state unequivocally that this indicates amorphous phase separation although this has been reported previously by several groups of u'orkers.8*'s21 The consequence of this structural change is a profound depen- dence on composition in the performance of the glasses as cathode-active materials. Performance in High-temperature Lithium Polymer-electrolyte Cells Cells containing four of the glasses (70, 78, 82, and 88 mol%) were constructed and cycled as described. The open-circuit voltages of the cells, ca. 3.5 V, were similar to the values forrep~rted~.~ vanadium phosphate glasses in room-temperature liquid-electrolyte cells and or thorhombic V20, in an identical polymer-electrolyte cell but higher than that of cells incorporating V6OI3 (2.8 V). The first discharge curves of the cells are shown in Fig.6 in comparison with those of crystalline V6OI3 and V205 in identical cells. The smooth curves given by the glasses show that, unlike crystalline vanadium oxides, there are no specific sites for lithium inser- tion, which is characteristic of the amorphous state. The 70 mol% glass discharges to a much lower capacity on the first cycle (73% theoretical) than the glasses with higher V205 content, which all gave similar first discharge capacities (> 90% theoretical). This large difference cannot be attributed solely to the vanadium content of the different glasses but rather seems to be associated with the change in microstruc- ture which occurs at ca.75mol% V205 (see earlier). A network structure similar to that of 0-VPO, would restrict the diffusion of Li+ ions and give rise to a lower electronic conductivity (as observed). Our molar volume data, which showed a maximum around 75 mol% V20,, are not consistent with a previous suggestion4 that lithium diffusion is being restricted by a more densely packed structure in the 70 mol% glass. The later discharges for all of the glass-containing cells were also monotonous, as shown in Fig. 6, demonstrating that cycling at elevated temperatures does not induce crys- tallization of the vanadium phosphate glasses.The percentage theoretical capacity obtained is plotted against cycle number in Fig. 7. All of the cells showed a large capacity decline over the first four or five cycles which was primarily associated with a steepening of the voltage-capacity curve between 3.5 and 2.7 V, see Fig. 6,indicating that lithium was being retained in the lower energy sites. This was followed by a region of slow recovery so that by the 25th cycle the cells containing the three glasses with 378 mol% V205 cycled reversibly at around 60% theoretical capacity, with the highest utilization occurring for the highest vanadium content, as expected. This recovery, we believe, is due to the formation of a more open structure with continued cycling. Cathodes containing crystal- line V6013 show very similar behaviour to those containing the glasses, see Fig.7. This oxide has been shown" to become amorphous in high-temperature polymer-electrolyte cells after a few cycles, so the similarities seen here are perhaps not surprising. None of glasses we examined showed a utilization approaching that of V6013 (80% theoretical capacity). J. MATER. CHEM., 1994, VOL. 4 (b) 1 l.I.I I. 0 20 40 60 80 100 (c 1 3.0-1.51 I 1.51 II I I I I 1 1 I0 20 40 60 80 100 0 20 40 60 80 100 theoretical capacity (%) Fig.6 First and later discharge curves (as numbered) for cells containing V205-P205glasses as the cathode-active component and the first discharge curves for similar cells containing V6OI3 or Vz05.(a) 82 mol% V205 (b) 70 molo/o V205; (c) V6OI3;(d)crystalline V205 1001 = I 01 II I I 1 0 6 12 18 24 30 cycle number Fig. 7 Cycling performance of cells containing V205-P205 glasses as the cathode-active component in comparison to a cell containing V6OI3.(a) VOOl3;(h) 82 mol% V205;(c) 70 mol% Vz05 Cells containing the 70 mol% glass recovered after 30 cycles to a much lower capacity than those containing the other glasses (cci. 30% theoretical) which cannot be attributed solely to the lower vanadium content but, like the lower capacity on the first discharge, must be associated with the change in microstructure discussed above. However, our results show a significant improvement in reversibility compared with those of Pagnier et a/.' who experienced difficulty in recharging their cells.The disappointing reversibility they observed may have been due to the low electronic conductivity of their cathodes, which were prepared without the addition of either ketjenblack carbon, for increased electronic conductivity, or polymer electrolyte, for improved interfacial contact between cathode material and electrolyte. Our results, we believe, demonstrate the significant contribution of the composite cathode to cell performance. Fig. 8 compares the observed energy densities of two of the glasses with that of V6013in a similar cell. Although cells incorporating the glasses gave higher open-circuit voltages than V6OI3, and exhibited comparable average cell voltages on cycling (2.1-2.4 V), their specific capacities were slightly lower than that of V6OI3(ca.330 mA h g-') at the capacity plateau. This results in lower observed energy densities 1OOOr -800 iY0) .c 6oo. c).-v)c % 400-P i im -*-------'\A-+---+--------0 5 10 15 20 25 cycle number Fig.8 Variation of observed energy density with cycle number for (*) 70 and (A)82 mol% V205glasses compared with that of V6OI3(m)in a similar cell (500-600 W h kg-' for glasses containing 278 mol% V205) after 25 cycles than those observed for V6013 (ca. 770 W h kg-'). Conclusions The microstructures of vanadium phosphate glasses contain- ing < ca. 75 mol% V205 are predominantly fl-VPO,-like whereas those containing 3 ca. 75 mol% V205 have micro- structures with both V,O,-like and P-VP0,-like domains. All of the glasses we examined can be cycled reversibly when employed as cathodes in secondary polymer-electrolyte cells operating at 120"C, with those containing greater amounts of V205 giving the largest specific energy densities.This difference in performance is not solely related to the vanadium content but is also profoundly influenced by the microstruc- ture of the glasses. The better materials gave lower specific 118 J. MATER. CHEM., 1994, VOL. 4 energy densities over the first 30 cycles than obtained for V6OI3in a similar cell. 9 10 11 M. E. Brown, Introduction to Thermal Analysis, Chapman and Hall, London, 1988. M. Sayer and A. Masingh, Phys. Rev. B, 1972.6,4629.P. Aldebert, H. W. Haesslin, N. Baffier and J. Livage, J. Coll. We thank SERC and AEA Industrial Technology, Harwell Laboratory, for a CASE award to A.D. 12 Interface Sci., 1984,94,484, and references therein. Bh. V. Janakirama-Rao, J. Am. Ceram. Soc., 1966, 49, 605; Bh. V. Janakirama-Rao, J.Am. Ceram. Soc., 1965,48, 31 1. 13 F. R. Landsberger and P. J. Bray, J. Chem. Pkys., 1970,53,2757. References 14 15 B. Jordan and C. Calvo, Can. J, Chem., 1973,51,2621. R. Gopal and C. Calvo, J. Solid State Chem., I972,5,432. 1 T. Pagnier, M. Fouletier and J. L. Souquet, Solid State Ionics, 1983.9 & 10,649. 2 Y. Sakurai and J. Yamaki, J. Electrochem. Soc., 1985, 132,512. 3 Y. Sakurai, S. Okada, J. Yamaki and T. Okada, J. Power Sources, 1987,20, 173. 4 Y. Sakurai and J. Yamaki, J. Electrochem. SOC.,1988,135,791. 5 A. C. Leech, J. R. Owen and B. C. H. Steele, Solid State Ionics, 16 17 18 19 20 D. W. Murphy, P. A. Christian. F. J. DiSalvo and J. V. Waszczak, Inorg. Chem., 1979,18,2800. C. F. Drake, J. A. Stephan and B. Yates, J. Non-Cryst. Solids, 1978,28,61. K. Yvon, W. Jeitschko and E. Parthe, J. Appl. Cryst., 1977, 10, 73. G. W. Anderson and F. U. Luehrs, J. Appl. Phys., 1968,39, 1634. G. W. Anderson and D. W. Compton, J. Chem. Phys., 1970, 52, 6166. 1983,9 & 10,645. 21 E. J. Freibele, L. K. Wilson and D. L. Kinser. J.Am. Ceram. Soc., 6 M. Levy, P. Rousseau and M. J. Duclot, Solid State lonics, 1988, 1972,55, 164. 28-30,736. 22 W. J. Macklin, R. J. Neat and S. S. Sandhu. Electrochim. Acta, 7 A. Hooper and J. M. North, Solid State lonics, 1983,9 & 10, 1161. 1992,37,1715. 8 R. N. Bhargava and R. A. Condrate Sr., Appl. Spectrosc., 1977, 31,230. Paper 3/04197B; Receired 19th July, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400113

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. MATER. CHEM., 1994, 4( l), 119-124 Chemomechanical Polishing of Lithium Niobate using Alkaline Silica Sol and Alkaline Silica Sol modified with Hydrogendifluoride Anion Margaret Beveridge," Laurence McGhee," Scott G. McMeekin,bt Max I. Robertson,b Alexander Rossb and John M. Winfield*" a Department of Chemistry, University of Glasgow, Glasgow, UK G12 8QQ Logitech Ltd, Old Kilpatrick, Dunbartonshire, UK G60 5EU Lithium niobate undergoes a slow reaction with anhydrous HF at room temperature to give LiNbOF,, which is readily hydrolysed to give Nb205 as one product. This sequence of reactions is believed to be responsible for the superior performance of aqueous alkaline silica sol modified by the addition of [HFJ-for the chemomechanical polishing of LiNbO, wafers, compared with silica sol alone.Lithium niobate is an important material for optoelectronic applications;' for example the fabrication of planar optical waveguides by hydrogen-lithium exchange processes involv- ing LiNbO, single-crystal wafers has received considerable attention.24 For many applications the device performance is determined by the quality of the surface finish. The trend towards smaller devices requires wafer surfaces whose rough- ness, the peak-valley variation, should approach 1 nm as measured by a stylus instrument. Careful control of surface polishing is therefore necessary. Polishing of optical materials has a long history, many of the guiding principles were formulated by Lord Ra~leigh,~ but in many respects it is still as much art as science.Studies of LiNbO, polishing have either focussed on practical aspects or have emphasised mechanical interactions between abrasive particles and the LiNbO, surface.8p'' It is increasingly appar- ent, however, from the large number of polishing studies involving silica and other oxide glasses, that chemical reactions at the surface may be just as important as mechanical action in determining the quality of the surface fini~h,'~.'~ hence the term 'chemomechanical polishing'. We have examined the effect on LiNbO, polishing of the hydrogendifluoride anion by comparing the results obtained using an alkaline silica sol'' (Syton) with those obtained under identical polishing conditions but with [HF,] -added to the reagent.The latter reagent results in a faster process without loss in surface quality. The chemical events occurring at the LiNb03 surface have been investigated using model experiments and the results used to suggest a mechanism for the polishing reaction. Experimental Materials Congruent lithium niobate wafers, X-, Y-, or Z-cut, (Pilkington Electrooptics) were cut to the required dimensions using a diamond saw and were edge lapped to prevent mechanical fracture during subsequent face-polishing operations. Congruent LiNb0, powder (Johnson Matthey) was used as received. The aqueous alkaline silica sol was a commercial material (Syton, Monsanto; nominal particle size 0.04 pm) Potassium hydrogendifluoride was prepared using the method of Clark et ~1.'~Anhydrous potassium fluoride (General Purpose Reagent BDH; 174g 3.0mol) was added to distilled water (300 cm3).The solid dissolved on stirring, then glacial acetic acid (Pronalys A.R, May and Baker, 300 g, t Present address: Department of Electronic and Electrical Engmeering, University of Glasgow. 5 mol) was added slowly with stirring. The resulting solution was allowed to stand at room temperature overnight, during which time a colourless crystalline solid formed. This was filtered and dried at 120°C. The total solid product obtained after 72 h was 35g, 30% based on the KF taken. (Found: HF, 23.6; KHF, requires: HF, 25.85%). Potassium hydrogen- difluoride (1.0 g) was dissolved in alkaline colloidal silica sol (125 cm3), stirring the mixture at room temperature until all the solid had dissolved.The pH was in the range 7-8. The reagent appeared to be stable, i.e. a colloidal solution, in use over 4 h. After longer periods however, coagulation occurred but the reagent was easily restored to its former state by vigorous shaking at room temperature. Retention of HF by silica gels has been noted previously." SEM examination of the solid obtained by evaporation of the colloidal solution showed it to be very similar to that obtained by evaporation of the silica sol alone, the only difference being a greater proportion of larger particles. However much of the fluorine may have been lost during the evaporation process." General Instrumentation IR spectra were recorded using PE 983 and Nicolet 5DXC spectrometers using a SpectraTech collector for diffuse reflec- tance spectra.Powder XRD measurements were made using a Philips PW1050-35 diffractometer with a vertical goniometer and Co-Kcr radiation, samples being mounted on adhesive tape. BET areas were determined using N, as adsorbate. Lithium in polishing solutions was determined by atomic absorption spectroscopy (AAS) using a Perkin-Elmer 1 100 spectrometer with a lithium hollow cathode lamp. Standard solutions were prepared from Spectrosol lithium nitrate solu- tion (1000 ppm) to give lithium concentrations in the range 0.25-2.00 ppm. Three integrated absorbance measurements were made for each solution and the relationship between absorbance and Li' concentration was linear over the range studied.Absorbances from the analyte solutions were deter- mined similarly and the lithium concentrations measured from the experimentally derived concentration us. absorbance relationship. Analyte solutions were prepared by taking an aliquot (25 cm3) of material collected from the polishing process and digesting it with aqueous HF (4070,10 cm3) in a PTFE beaker. The resultant solution was heated to dryness, digested with HNO,, evaporated to dryness and finally made up as a standard solution using aqueous HNO,. Duplicate analyses were carried out throughout. Polishing Procedure A set of 12 wafers (2 cm x 1cm) was cut from a large LiNb03 wafer, edge lapped and wax bonded in a cruciform arrange- ment to a glass puck, normally 83 mm diameter and 6 mm thick.This was attached to a Logitech PP5 polishing jig by means of a vacuum chuck. The polishing jig ensured that the samples were held parallel to the lapping and polishing plates. The samples were lapped on a Logitech PM2 precision polishing instrument which was fitted with a grooved, cast- iron plate, at 30-40 rpm using 3 pm a-alumina in water as the abrasive. The jig loading was ca. 3.2 kg. The lapping time was dependent on the initial height and condition of the samples, lapping being continued until all the samples were in a single plane. Surface-roughness measurements of lapped wafers were in the range 100-140 nm. Samples were polished using the same instrument with an expanded polyurethane plate (grade LP87) at 10 rpm, with 1.5 kg loading and alkaline silica sol with or without added KHF,, as the polishing fluid.Feed rates were ca. 100 cm3 h-l. The progress of a polishing reaction was followed by determination of stock removal, surface roughness and Li analyses of polishing fluids at regular intervals. Stock removal was calculated using a contact gauge to measure the change in thickness at the centre point of each wafer. The precision of these measurements was limited by the resolution of the gauge (0.1 pm). Surface roughness determinations were made on a Rank Taylor Hobson Talystep instrument, data being presented as the amplitude parameter, R, (defined as the arithmetic mean of the departures of the roughness profile from the mean line).R, values were calculated from data taken over a 0.5 mm trace at magnifications of 1 x 105-2 x lo6 depending on the degree of lapping damage that remained. The standard stylus was used for all traces, except after the final polishing time interval when the fine stylus was used. The quantity of polishing reagent used over a specified time interval was collected quantitatively and to this volume was added the washings from the wafers and polishing plate. Analytical data are expressed as total Li, i.e. the product of the total solution volume and the Li concentration as deter- mined by AAS. Replicate experiments indicated that there was no marked dependence of total Li determination on the volume of solution collected, thus the solubility of the Li-containing product was not a determining factor.The quality of surface finish obtained and the time required for polishing with alkaline silica sol alone proved to be very dependent on the procedure used for conditioning the expanded polyurethane pad. A grooved pad required frequent reconditioning to prevent a film of silica being formed on the surface, inhibiting transport of the reagent. A satisfactory, reproducible procedure was to scroll the pad on a lathe, using a 45" diamond tool to form a 0.5 mm deep spiral groove. This procedure was less important when using the silica sol-KHF, reagent but was usually adopted in order to achieve an objective comparison. Fluorine-18 Measurements and the Reaction of LiNbO, with Anhydrous HF Fluorine-18 (rl,, = 110min, p' emitter) was prepared by neu- tron irradiation of Li,CO, at the Scottish Universities Research and Reactor Centre, East Kilbride. The irradiated material was converted to H18F and then to solid Cd8F using inactive CsF as a carrier and the solid was transported to Glasgow.16 Aqueous H18F was prepared by addition of a small quantity (ca.5 mg) of solid Cs18F to 40% aqueous HF (20 cm3) contained in a PTFE beaker. The mixture was warmed to equilibrate the solution, cooled and diluted with 40% aqueous HF until an aliquot (0.5 cm3) counted in an FEP (perfluoroethylene propylene) tube gave a count of lo4 in 30s. Anhydrous HF was labelled with [''F] by exchange with solid Cs18F at 200°C in a Monel metal pressure vessel J. MATER.CHEM., 1994, VOL. 4 connected to a calibrated Monel vacuum line. The specific count rate (count minmmol-l) of each batch of H18F pre- pared was determined by condensing a measured quantity of gaseous Hi8F onto a sample of calcined ;'-alumina contained in an evacuable FEP tube equipped with a Monel or PTFE valve. The interaction of H18F vapour with LiNbO, wafers, cu. 5 x 20 mm, and with congruent LiNbO, powder (0.5 g) was studied by expanding a measured quantity of H18F into an evacuable FEP counting tube. In some cases the sample was pumped for 24 h before Hl'F admission. Count rates were determined at regular intervals using a Scaler ratemeter and an NaI well scintillation counter (NE and Ecko).The counting geometry used was such that ["F] activity determined arose mainly from the activity deposited on the solid sample. Calibration experiments using counting tubes containing H18F vapour only, confirmed this and thus ["F] activity originating from the gas phase could be discounted without serious error. The variation of count rate with time for a sample defined the ['*F] growth curve for the solid. When apparent equilib- rium was reached, volatile material was removed by pumping at room temperature for ca. 10min. The solid count rate corresponded to ["F] activity permanently retained by the wafer or powder. This count rate, combined with the specific count rate of H18F used, enabled the uptake of HF to be determined.For some samples the effect of addition of further quantities of H18F or the addition of inactive HF was deter- mined using a similar procedure. Congruent LiNbO, powder (4g) was added to an FEP tube and connected to a Monel vacuum line by Swagelok fittings. Anhydrous HF (3 cm3) was added by distillation and the mixture was allowed to react for 20 h. Volatile material was removed by distillation and the solid pumped to remove remaining traces of HF. The solid material was washed with distilled water through a plastic frit, the colourless washings evaporated, and the resulting solid dried at 120"C. Analysis of the solid gave: F, 31.8; Li, 3.55%; Li: F = 1:3.2. A second analysis carried out later gave: F, 25.0; Li, 3.6%; Li :F =1:2.5. XRD and vibrational spectroscopic data obtained are dis- cussed below. Results Polishing Experiments Several series of polishing experiments carried out under identical conditions, demonstrated the beneficial effect of adding KHF, to alkaline silica sol for the polishing of Y-cut LiNbO, wafers. The criteria used were stock removal measure- ments, surface-finish determinations as measured by R, data (an average value determined from two traces taken on five of the twelve wafers polished in any one experiment) and the lithium contents of the polishing reagents collected over 0.5 h time periods during polishing.Stock removal data are com- pared in Fig. 1, data averaged from two separate experiments for each reagent.After 120 min polishing ca. 8 pm of material had been removed using silica sol-KHF, whereas only cu. 3.5 pm was removed using silica sol alone over the same time period. Subnanometre surface roughness was achieved in under 1 h using silica sol-KHF, whereas ca. 2 h was required using silica sol alone, see Fig. 2 (data averaged from two experimental runs). Using either reagent, final R, values were in the range 0.5-1.0 nm. Lithium content, determined as the product of the Li concentration in the reagent collected over a 0.5 h period and the reagent volume, was always greater using the silica sol-KHF, reagent. However, data obtained from polishing with silica sol alone depended markedly on the degree of J. MATER. CHEM., 1994, VOL.4 I I I I I I I 0 20 40 60 80 100 120 140polishing time/min Fig. 1 Stock removal from LiNbO, during polishing with alkaline silica sol (a)and alkaline silica sol-KHF, (0) 60 I 50 .i 40 E5 30cr" 20 10 0 0 20 40 60 80 100 120 140 poIis h in g tim e/m in Fig. 2 Surface roughness measurements on LiNb0, made during polishing with alkaline silica sol (a)and alkaline silica sol-KHF, (0) polishing pad conditioning. Data in Fig. 3 are typical of those obtained, however Li contents from silica sol polishing experi- ments could be increased by up to a factor of two, particularly in the first hour of polishing, by careful attention to pad conditioning. The effect of anisotropy in LiNbO, on polishing was investigated by Li determinations in polishing fluids. Similar- sized pieces of X-, Y- and Z-cut LiNbO, were lapped and polished using identical conditions after 1.5 h and Li contents 0.7 0.6 0.5 0.4 h -I-a c0 0.3 0.2 10.1 0.0 1 I 2 5 6 7 8 time intervaV0.5 h Fig.3 Total Li determined from polishing fluids during successive 0.5 h periods during polishing of LiNbO, with alkaline silica sol (dark columns) or alkaline silica sol-KHF, (light columns) determined (Table 1). The orientation of the LiNbO, wafer had an effect on the quantity of soluble Li determined, particularly when silica sol alone was used. Addition of KHF, produced the greatest enhancement of the polishing reaction when Y-cut material was polished, but appeared to have no effect for X-cut. In view of the mechanical effects described above, any correlation of the small differences in the lithium removal with the different atomic arrangements of the three surfaces is not justified.Interaction between Anhydrous HF and Silica The concentration of [HF,] -in the fluoride-modified silica- sol reagent was 0.1 mol dmP3, this being the upper limit for avoidance of rapid coagulation of the reagent. In order to model the silica sol-[ HF,] -interaction, silica powder obtained by evaporation and vacuum drying of alkaline silica sol was exposed to fluorine-18 labelled anhydrous hydrogen fluoride vapour whose [18F] specific count rate had been determined at room temperature. Exposure resulted in an immediate [18F] count rate which increased slowly thereafter.The interaction between H1'F and the solid appeared to be essentially complete within a few minutes. On the removal of excess H"F by pumping, the [l'F] count rate of the solid corresponded to an HF uptake of 0.30f0.01 mmol g-' This was indicative of a significant interaction, although SEM examination showed no difference in particle size as a result of HF vapour treatment. Table 1 Total Li determined from polishing fluids after 90 min polishing of LiNb0, Li determined/mg after polishing with wafer orientation silica sol silica sol-KHF2 X-cut 0.26 0.25 Y-cut 0.05 0.3 1 Z-cut 0.18 0.42 J. MATER. CHEM., 1994, VOL. 4 Interaction between LiNbO, Wafers and Water A large Y-cut LiNbO, wafer was cut into twelve 2 cm x 1 cm wafers; these were polished with alkaline silica sol and their transmission IR spectra recorded.The spectra contained a complex band envelope in the v(0H) region, v,,, =3485 cm-' with shoulders at 3465 and 3502cm-'. These were very similar to previous observations which have been taken to indicate the presence of three types of hydroxy group in the material." Band integrations were uniform from sample to sample, indicating there was no gross difference in chemical action over the total face of the larger wafer as a result of the polishing process. Irradiation of an LiNbO, wafer-water mixture in a commer- cial microwave oven led to the incorporation of additional hydrogen. The IR spectrum of a polished wafer irradiated for 30 min in liquid r2H]-H2O contained a complex absorption between 2650 and 2560cm-' due to v(02H) vibrations in addition to the v(0'H) absorption.The latter was essentially unaffected by the irradiation process, indicating that the new bands arose from the incorporation of additional hydrogen rather than being the result of 2H/'H isotopic exchange. After 16 days under ambient laboratory atmosphere the v(02H) absorption could not be observed. In a separate experiment the v(02H)absorption was observable after 20 min irradiation and the band area increased significantly after a further 20 min irradiation period. Attempts to follow the decrease in the v(02H)band area quantitatively with time were not entirely successful; however, it had decreased significantly after 11 days.Irradiation of an LiNbO, wafer in the presence of H20 for 35 min resulted in the v(0H) band area increasing by a factor of two. After a total of 1.5 h irradiation the band intensity had increased by a factor of three. In all cases it appeared that additional hydroxy groups were formed during the irradiation process and that the 'extra' hydrogen incorpor- ated was lost from a wafer under ambient conditions. Prism- coupling experiments on an LiNbO, wafer irradiated for 1.5 h in H20 indicated that the sample behaved as a monomode waveguide, i.e. the extraordinary refractive index of the wave- guiding layer was greater than that of bulk LiNbO,. By analogy with previous work in which multimode planar waveguides have been fabricated by reactions of LiNbO, wafers with protonic acids,2" it is likely that lithium is lost from the wafer during this process.The observed behaviour was unexpected from previous studies in which hydrogen ('H or 2H) diffusion into LiNbO, was observed only at high temperatures.18 It is, however, related to the facile 'H/2H isotopic exchange processes that have been observed between water vapour and partially exchanged Li, -,H,NbO, materials, and implies that the reversible reaction (eqn. 1)19 is possible 02-(surface)+H20 (liquid)+20H- (surface) (1) Interaction between LiNbO, and Anhydrous HF This was examined by two methods, by exposing LiNbO, to H18F vapour at room temperature and by physico-chemical examination of the products from the reaction between LiNbO, and liquid anhydrous HF.Admission of HI8F vapour to LiNbO, powder at room temperature resulted in significant uptake of ["F] by the solid (Fig. 4). The bulk of the radio- activity (ca. 80%) appeared to be retained on pumping at room temperature and the uptake of HF, determined by using H"F of measured specific count rate, was in the range (0.19-0.35)kO.Ol mmol g-'. The [18F] growth curve indi- cated that the reaction was relatively slow. An apparent equilibrium was reached after 100 min; however, removing excess H18F at this stage then adding a further aliquot of 1000 3 900 F I .-E E 800 1 81 Q,c s! c1 8 700 LL 7 600 II I Ill 0 20 40 60 80 100 120 140 t im e/m in Fig.4 Growth in ['*F] count rate in LiNbO, on exposure to HI8F vapour at room temperature H18F, led to a second rapid uptake.The [18F] activity was reduced by exposure of the labelled solid to inactive HF vapour at room temperature, suggesting that some fluoride was labile and all ["F] appeared to be removed completely by washing with H,O. Aqueous H18F was placed on an LiNbO, wafer and left for 0.5 h; after rinsing with H20 all the activity was removed. The BET area of LiNbO, powder was 0.3 m2 g-'. Assuming the van der Waals radius of the HF molecule is 2.55 x lo-'' m then the uptake of HF required to form a monolayer on the surface of LiNbO, is 0.002mmol g-'. The observed uptakes were far higher, suggesting that the reaction involved bulk material.The IR transmission spectrum of the solid recorded after ["F] activity had decayed, was similar to that of untreated LiNbO,. In particular, there was no evidence for new bands in the region expected for Nb-F stretching vibrations. Diffuse reflectance IR Fourier transform spectroscopy (DRIFTS) of LiNbO, powder and LiNb0, powder treated with anhydrous HF produced good quality spectra. The main features in the LiNbO, spectrum (4000-650 cm-l) were envelopes in the hydroxy group stretching and deformation regions and a strong band at 960 cm-'. Treatment of LiNbO, powder with HF vapour followed by subsequent exposure to moist air, resulted in a significant increase in the intensity of the hydroxy group stretching mode envelope.Before HF treatment this occurred at 3700-2750 cm-' with vmax= 3468 cm-'; the corre- sponding data after treatment were 3725-3000 cm-' and v,,,=3528 cm-'. In addition, the profile of the 960 cm-' band was changed. There was no evidence in the spectra for the presence of (HF), oligomers nor for bands attributable to Nb-F stretching modes. The XRD powder pattern of LiNbO, that has been treated with liquid anhydrous HF overnight [Fig. 5(a)] showed the presence of LiNbOj2' and LiNbOF4.21 This product was partially soluble in water and the solid isolated from solution after evaporation and drying contained LiNbOF42' and Nb20520 as shown from its XRD powder pattern [Fig.5(b)]. The main features in its Raman spectra were strong bands at J. MATER. CHEM., 1994, VOL. 4 76 68 60 52 44 36 28 20 12 4 20ldegrees Fig. 5 XRD powder results from (a)the solid product from LiNb03 and anhydrous HF (liquid); (h)the water soluble material from this reaction; ( W) LiNbOF,, (A)LiNb03, (a)Nb,O, 810 and 607cm-’. Its fluorine content decreased with time presumably owing to hydrolysis. The behaviour of LiNb0, towards HF may be compared with the reaction between LiNbO, and liquid BrF,. At room temperature a mixture of LiNbOF, and LiNbF, is formed; conversion to LiNbF, is not complete even at 1260C.21Under more extreme conditions, for example using molten hydrogen- difluoride salts, LiNbO, can be fluorinated to give LiF and M,NbOF6 (M=NH, or K).,, Discussion Both processes studied in this work can be described as chemomechanical, in which chemical reactions are induced at the lithium niobate surface via mechanical energy generated from forces between the polishing pad and the wafer surface.The results obtained are inconsistent with a simple grinding process involving silica particles for two reasons. First, a water soluble lithium species is produced and, secondly, addition of the hydrogendifluoride anion has a beneficial effect. The latter would not be expected for a purely mechan- ical process because of the destabilizing effect of [HF,]- on the negatively charged silica particles (cf. ref. 15). The lithium- containing product is unlikely to be simply LiNbO, since the latter is insoluble in water and can be dissolved only with difficulty in aqueous HF.The reaction between LiNbO, and [HF,] -under polishing conditions does not lead to retention of fluorine-containing species on the surface. Removal of LiNbO, is enhanced however (Fig. 1 and 3) but without loss of surface finish (Fig. 2). Simple etching of LiNbO, in aqueous HF, in which the major species is the ‘tight’ ion pair H30+F-,23 destroys the surface finish. The most obvious way of moderating an etching reaction for polishing purposes is via the formation of an insoluble product (a passivating layer) which accumu- lates in the valleys on the surface (cf. ref. 12, 13). Continued reaction on the peaks eventually produces a plane surface from which the passivating layer is removed by mechanical wiping.The observations made in the polishing of LiNbO, by KHF,-silica sol and the results of the model reaction between LiNbO, and anhydrous HF, particularly the XRD data (Fig. 5), lead us to propose the following model. Hydrogendifluoride anion reacts with the LiNbO, surface to give water-soluble LiNbOF,, which undergoes hydrolvsis to give hydrated Nb205 and LiF. Hydrated Nb205 accumulates on the surface and acts as a passivating layer. Further reactions involving Nb205 must be considered, since it is known to form the [NbOF5I2- anion in dilute aqueous HF,,, and to react with MHF, (M=K or NH,) giving M,NbOF,.25 However [NbOF,I2- and [NbOF4( H20)] -exist in equilib- rium with each other in water,,, hydrolysis of [NbOF,I2- in alkaline media is reported to give hydrated Nb205” and [NH,], [NbOF,] gives [NH, J[NbOF,] on thermal decomposition.22 In reality therefore, LiNbOF, and hydrated Nb205 can be considered to be two of the key member.; (eqn.2 and 3) in a complex series of reactions occurring on the LiNbO, surface. LiNbO3+4[HF,]--+LiNbOF,+4F-+2H20 (2) 2LiNbOF, +6H,O+Nb,O, +2LiF +6HF (3) Attempts to identify the presence of NbV 0x0 or oxofluoro species by vibrational spectroscopy were unsuccessful, because of the dominant effect of LiNb03.27 There were, howeher, no bands that could be obviously attributed to [NbF,]-or [NbF7I2-. The lack of evidence for LiF in the XRD powder data from the model reaction (Fig.5) can be attributed to the dominant effect of the heavy NbV atoms. The activation of the surface observed when an LiNbO, wafer is exposed to microwave radiation in the presence of water (eqn. l), leads us to suggest that the chemical reaction between LiNbO, and alkaline silica sol under polishing con- ditions may involve the attack of OH- or -Si-0- maieties on the surface. The products expected would be LiOH and Nb,O,. The absence of additional reactions (eqn. 2 and 3) results in a slower polishing process in agreement with the observations made. The authors thank staff at the Scottish Universities Research and Reactor Centre, East Kilbride for assistance with the neutron irradiations. Financial assistance from the Department of Trade and Industry and Logitech Ltd, under the auspices of the LINK Nanotechnology programme, is gratefully acknowledged.Prism-coupling experiments were carried out by Dr. A. Loni of the Department of Electronic and Electrical Engineering, University of Glasgow. References 1 A. Rauber, Current Topics in Materials Science, ed. E. Kaldis, North-Holland, 1978, vol. 1, ch. 7, p. 481. 2 J. L. Jackel. C. E. Rice and J. J. Veselka, Jr., Appl. Phys. Lett., 1982, 41, 607; C. E. Rice, J. L. Jackel and W. L. Brown, J. Appl. Phys., 1985,57,4437. 3 A. Loni, R. M. De La Rue and J. M. Winfield, J. Appl. Phys , 1987, 61, 64; A. Loni, G. Hay, R. M. De La Rue and J. M. Winfield, J. Lightwave Technol., 1989,7,911; A. Loni, R. W. Keys, R. M.De La Rue, M. A. Foad and J. M. Winfield, IEE Proc. Part J, 1989, 136,297;A. Loni, R. M. De La Rue, J. McCaig and J. M. W infield, J. Appl. Phys., 1990,67, 3968. 4 J. T. Cargo, A. J. Filo, M. C. Hughes, V. C. Kannan, F. A. Stevie, J. A. Taylor and R. J. Holmes, J. Appl. Phys., 1990,67,627. 124 J. MATER. CHEM., 1994, VOL. 4 I. T. Savatinova, M. Kuneva, B. Jordanov and D. Kolev, J. Mol. 18 R. Gonzalez, Y. Chen, K. L. Tsang and G. P. Summers, Appl. Struct., 1990,219, 165. Phys. Lett., 1982, 41, 739; N. Schmidt, K. Betzler, M. Grabs, V. A. Ganshin and Yu. N. Korkishko, Opt. Commun., 1991, 86, S. Kapphan and F. Klose, J. Appl. Phys., 1989.65, 1253. 523. Lord Rayleigh, Proc. Optical Convention, 1st Convention, London, 19 20 W. Bollmann, Phys.Status Solidi A, 1987,104. 643. Joint Committee for Diffraction Standards, International Centre 1905, p. 73; see also, F. Twyman, Prism and Lens Making, 2nd edn, for Diffraction Data, LiNbO,, 20-631; Nb205.18-91 1. 8 9 10 11 12 13 14 15 16 17 Institute of Physics, 1988, ch. 3, p. 49; G. Fynn and W. J. A. Powell, Cutting and Polishing Optical and Electronic Materials, IOP, Bristol, 1979;W. J. Rupp, Optica Acta, 1971,18, 1. J. Noda and I. Ida, Rev. Electr. Commun. Lab., 1972,20, 152. B. Furch, E. Bratengeyer and H. Rauch, J. Opt. Commun., 1983, 4, 47. E. Neumann and H. Schulz, Cryst. Res. Technol., 1985,20, K115. S. D. Poulsen, FerroeIectrics, 1987,75, 79. N. J. Brown, Precision Engineering, 1987,9, 129. L. M. Cook, J. Non. Cryst. Solids, 1990, 120, 152.J. H. Clark, J. Emsley, D. J. Jones and R. E. Overill, J. Chem. Soc., Dalton Trans., 1981, 1219. E. M. Rabinovich and D. L. Wood, Muter. Rex Soc., Symp. Proc., 1986.73, 251; E. M. Rabinovich, D. M. Krol, N. A. Kopylov and P. K. Gallagher, J.Am. Ceram. Soc, 1989,72, 1229. K. W. Dixon and J. M. Winfield, J. Chem. Soc., Dalton Trans., 1989,937. J. R. Herrington, B. Dischler, A. Rauber and J. Schneider, Solid 21 22 23 24 25 26 27 E. G. Rakov, M. V. Melkumyants and V. F. Sukhoverkov, Russ. J. Inorg. Chem., 1990, 35, 632. A. I. Agulyanskii, Yu. I. Balabanov, \I. A. Bessonova, A. G. Babkin and P. T. Stangrit, Iza. Akud. Nauk SSSR, Neorg. Muter., 1985,21,98; Chem. Abstr., 1985, 102, 124587r. P. A. Giguere and S. Turrell, J. Am. Chem. Soc., 1980, 102, 5473; D. Mootz, U. Ohms and W. Poll, 2. Anorg. Allg. Chern., 1981, 479,75; J. Khorami, R. Beaudoin and M. Menard, Can. J. Chem.. 1987,65,817. J. A. S. Howell and K. C. Moss, J. Chem. SOL‘.,A, 1971,2481. A. I. Agulyanskii, S. S. Pochivalov and V. M Mel’nikova, Russ. J. Inorg. Chem., 1989,34, 1567. Yu. A. Buslaev, E. G. Ilk, V. D. Kopanev and V. P. Tarasov. J. Struct. Chem., USSR, 1972, 13, 865. J.-M. Jehng and I. E. Wachs, Chem. Muter, 1991,3, 100; J. H. von Barner, E. Christensen, N. J. Bjerrum and B. Gilbert, Inorg. Chem.. 1991,30, 561. State Commun., 1973, 12, 351; L. Kovacs, V. Szalay and R. Capelletti, Solid State Commun., 1984,52, 1029. Paper 3/0348 1J; Receiccii 17th June, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400119

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. MATER. CHEM., 1994, 4(1), 125-131 Microcalorimetric Study of the Acidity of Tungstic Heteropolyanions Frederic Lefebvre, Feng Xian Liu-Cai and Aline Auroux lnstitut de Recherches sur la Catalyse, Laboratoire Propre du CNRS conventionne a I’Universite Claude-Bernard Lyon I, 2 avenue Albert Einstein, 69626 Villeurbanne Cedex, France The number and strength of the acid centres of tungstic heteropolyacids have been determined by absorption calorimetry of ammonia. The initial heats are in the order H3PW,20,0 >H,SiWl,O,O >H6P2W21071(H20)3>H,p2W1&62, varying from 200 to 155 kJ mol-’. An increase in the number of protons in Keggin heteropolyanions decreases the acidic strength. Moreover, the influence of the activation temperature on the acidity has been studied and confirms that there is a drastic modification of the solid at ca.350°C.The heat capacities of the heteropolyanions and of the corresponding ammonium salts, the thermokinetic parameters of ammonia absorption and the heat capacities of the acids and ammonium salts have been related to the porosity of the various samples. Heteropolyanions are complexes used in acid and/or redox catalysis whose general formula may be represented by (X,M,0,)4-(x<rn) where M is usually molybdenum or tungsten in their highest oxidation states, X is a heteroatom (P, Si, . . .) and q is the anion charge. When the counter-cations are protons, they are called heteropolyacids. In the formula, these exchangeable protons are placed before the heteroatoms as in simple oxoanions.Owing to their high number and diversity, classification, nomenclature and formulation of these complexes cause problems. For this reason, a classification based on structural relationships and acidic properties should be of great interest. They are used industrially mainly in the reaction of hydration of alkenes (acid catalysis) and in the synthesis of methacrolein or isobutyric acid (redox catalysis). The first experiments concerning the activity of heteropolyanions in catalysis were performed in Japan’,2 but recently more and more studies devoted entirely to this field developed owing to the industrial challenge of finding alternatives for the usual acid catalysts (H2S04,HF) for pollution reasons.3p5 The purpose of this work is to have a better understanding of the acidic properties of heteropolyacids. Indeed, because of the considerable number of available polyanion structures, it is possible to vary their acidic properties as well as their redox ones.Molybdic polyacids, which can be easily reduced, display mainly redox properties, whereas tungstic acids are often stronger. This study concerns tungstic compounds whose acidic properties are not simultaneously linked to redox character. The aim is to measure their acidity using ammonia absorption microcalorimetry and also to determine the specific heats of the initial acids and the corresponding ammonium salts. Four tungstic compounds were studied (see Table 1). Two of them, H3PW,2040 and H4SiW12040, display a Keggin Table 1 Physical characteristics of the heteropolyanions formula molecular weight/g density/ g cmP3 a BET surface area/m2g - 2882 6.4 5 2933 3.3 130 2880 2948 3.7 - 5 3.3 5104 6.5 3.5 5206 4.3 - 4372 5.6 4 4474 5.8 - structure6 [Fig.l(a)] with, respectively, 3 and 4 protons for the same structure, allowing the study of the charge effect. These compounds display a tetrahedral symmetry based on a central X04 tetrahedron, surrounded by twelve M06 octahedra arranged in four groups M3013 of three edge- bridged octahedra. The two other compounds under study, H6P2W18062 and &P,W2,0,, (H20)3, are phosphotungstic heteropolyacids but they present structures that are different from the previous ones.[P2w&62]6-, with a ‘Dawson’ heteropolyanion structure6 is represented in Fig. l(b) and can be described in terms of two PW,03, groups (obtained by removing three W06 octahedra from a Keggin anion) connected by six terminal oxygen atoms. [P2W21071(H20)3]6-[Fig. 1(c)] is coniposed of two PW903, units linked by three WO,(H,O) octahedra. All the compounds can be prepared by acidification of aqueous solutions of the simple oxoanions and the required heteroatoms, for example, according to the reaction pattern: 12(WO,)’-+(HPO,)’-+23H+-+(PW12040)3-+ 12H,O The equilibrium constants and the formation rates are suffi- ciently high to allow the crystallization of the salts of the polyanions starting from stoichiometric mixtures of the compounds at room temperature.6 Table 1 shows the chemical formulae of the compounds and of the corresponding ammonium salts, the molecular weight of the anhydrous compounds, their density relative to water (measured in n-hexane) and their BET surface area.Their acidity was determined using ammonia absorption microcalor- imetry, since this is one of the most reliable methods for measuring the number and strength of acid sites of a ~atalyst;~ we discuss these results in terms of the macroscopic and microscopic properties of the samples. In particular, we shall consider their porosity. For this purpose, the samples are separated in two groups, those with a low BET surface area (< 10m2g-’) and those with a high BET surface area (> 100 m2 g-’).The former will be referred to as ‘non-porous’ even if they have some meso- or macro-pores, but which represent only a small part of the total surface area. The latter, referred to as ‘microporous’ compounds, have micropores (diameter < 1.5nm) which represent most of the surface area. These micropores are arguably zeolitic channels (see later). Some of the calorimetric data have been published in preliminary form’ but were not then related to porosity and particle size. Experimental Measured by using a pycnometer; bmeasured by N2 absorption after The solids were purchased by the Laboratoire de Chimie des evacuation of the sample during 2 h at 150 “C. Metaux de Transition (University of Paris VI). Thev were J.MATER. CHEM., 1994, VOL. 4 Fig. 1 Structure of various tungstic heteropolyanions: (a) structure of the Keggin heteropolyanion [PW,,0,,]3-; (h) structure of the Dawson heteropolyanion [P2W18062]6-;(c) structure of [P2W21071( H20)3]6- characterized by infrared and UV-VIS spectroscopies, polar- ography on mercury drop, X-ray diffraction and chemical analysis. The amount of water of hydration was determined by thermogravimetry. All compounds were used without further purification. The microcalorimetric study has been performed using a heat-flow microcalorimeter (HT from Setaram). The heteropolyacid sample was outgassed under vacuum (150"C, 1.33mPa) prior to any absorption of ammonia in the calorimetric cell maintained at 150°C. A 150°C temperature has been chosen in order to avoid too much physisorption and thus to favour the accessibility of the probe to the active acid sites of the solid.The experiment consists in admission of successive doses of reactive gas (NH,) on to the catalyst and waiting for the thermal equilibrium after each pulse. The pressure evolution is measured with a differential Barocel gauge (Datametrics) until a residual equi- librium pressure of ca. 100 Pa is reached. The evaluation of the absorbed volume and of the evolved heat for each dose allows us to plot the differential heat curve as well as the integral heat and the differential entropy as a function of the number of sites. Ammonia from Air Liquide (purity >99.9%) has been purified by successive freeze-thaw cycles after drying on sodium chips. The specific heats of the acids and the corresponding ammonium salts were determined using a differential scanning calorimeter (Setaram TG-DSC 11l),by a discontinuous scan- ning temperature programme, the temperature varying by 5°C steps and the heat consumed being measured for each increment.This method allows the sample to return to thermal equilibrium after each increment but needs two runs with and without a sample so as to measure the specific heat with accuracy. Results and Discussion Fig. 2 represents the differential heats of NH, absorption as a function of the sorbed volume for the four samples under study. It appears from these results that the heteropolyacids can be separated into two groups: those compounds showing a curve with a constant differential heat (H3PW1204, and H6P2W18062)and those for which it is continuously decreas- ing [H&W210,1(H20), and H,SiW,,O,,].However, in all cases, the initial heats of absorption are much higher than the values reported in the literature for acidic catalysts such as zeolites or simple oxides.' The total volume of sorbed ammonia can be deduced from these results, being 1004, 899, 1257 and 996 pmol g- (hydrated catalyst) for H,PWl2O4,, H4SiW,,040, H6P2W18062 and H6P2W2 I 071(H20)3, re-spectively. If these values are calculated as the number of NH, molecules for one heteropolyacid, the following values are obtained: 2.89 (H,PW,,O,,), 2.36 (H4SiW120,,), 5.0 (H6P2wl8o6,) and 4.65 [H6P2W210,1(H20),].These data show that most of the protons of the solids have reacted. So the reaction: H,(HPA)+xNH,+(NH,),(HPA) can be considered as complete in a first approximation and, as a consequence, it appears more reliable to speak of ammonia absorption instead of adsorption for these compounds. Most of our results are in disagreement with previous studies on the acidity of heteropolyacids. Indeed, studies using Hammett indicators showed that there is a distribution of the acidic strength for heteropolyacids with the Keggin struc- ture.l0,l' Our results disagree also with those of a recent paper where the reaction of H,PW120,, with ammonia was studied by microcalorimetry,12 but the activation conditions were not the same.In order to understand why our results are different we have undertaken a more detailed study of the absorption of NH, by H3PW1204,. This heteropolyacid was chosen for many reasons: (i) it is, to our knowledge, the only compound for which other microcalorimetric experiments were per-formed; (ii) many data on its acidic properties have been reported; (iii) we have previously studied its thermal behaviour and so we know its temperature range of ~tability;'~ (iv) the structure of its ammonium salt is isomorphous to that of H3PW12040-6H,0'4~'5and it was found that it has a porous structure like zeolites.16 The curves showing the variation of the differential heat of absorption of ammonia as a function of the pretreatment temperature of H3PW1204, are shown in Fig.3. When the solid is treated under vacuum at 150 or 250°C, the curves are quite similar and do not show a very important modification of the acidity of the proton. However, when the catalyst is evacuated at 400'C, the curve is very different with an initial heat of absorption of 140 kJ mol-' (200-210kJ mol-' when the pretreatment is performed at 150 "C) and the heat of absorption is continuously decreasing with the sorbed amount of ammonia. This difference can be explained easily if we take into account the fact that at temperatures below 300 "C the H3PWI20,, polyacid is in its J. MATER. CHEM., 1994, VOL. 4 -6 200 400 600 800 1000 1200 amount NHJ~O-~mol g-‘ Fig.2 Differential heat of NH, absorption uersus the ammonia uptake for samples: 0 H3PW120,,; 0 H4SiW12040;* H6P2W1,062; H6P2W21071(H20)3 0 .I I 1 I . I u m .I 0 200 400 600 800 1000 1200 amount NH~IO-~ mol g-’ Fig. 3 Influence of the outgassing temperature on the acidity of H,PW,,O,,: 0,150; 0,250; I?, 400 ”C anhydrous form while at 400°C its transformation into the anhydride form is ~omplete,’~this latter phase being meta-stable and leading slowly to the constitutive oxides. This thermal evolution of the polyacid can be described by the following equations: H3PW1204,,, 6H20-+6H20+H3PWI2O4,,(anhydrous acid) nH3PWl2O4,+3nH20 +[PW12038.5]n (anhydride phase) [PW12038.5]n+$nP205+ 12nW03(oxides) The anhydride phase can be described as Keggin units linked by W-0-W bridges. Its formation has been deduced from thermogravimetry and BET surface area measurements. By reaction with water, it transforms into the polyacid.Thus, we can conclude that, in the anhydrous form of the H3PWI2O4,,polyacid, all protons display the same acidity, which is very high and corresponds to super-acidic centres, as has been demonstrated by cataly~is,””~while the anhydride (or the melt of the constitutive oxides) phase is less acidic with a heterogeneous distribution of sites and strength. The results of Kapustin et all2 can be interpreted in terms of an intermediate structure between the anhydrous and the anhy-dride phases as their pretreatment temperature is 300’C but with an evacuation time of 50 h (as the transformation of the anhydrous acid into the anhydride phase is a kinetic phenom-enon, the longer the time of evacuation of the sample the more complete will be the dehydration).It seems more difficult to explain the differences between the present results and the J.MATER. CHEM., 1994, VOL. 4 studies of the acidity by the method of the indicators'0*'' which showed that there is a heterogeneous distribution of the acidic strength. However, from a chemical point of view, our results seem more realistic, as it is difficult to obtain a distribution of proton strengths and so of chemical environ- ments in a crystalline phase where the atoms are located in defined positions.The only case should be a delocalization of protons in the solid, but it is not reasonable to think that such a distribution could induce chemical environments so different that weak and strong acidic centres could be created. Moreover, it seems from literature data that using coloured indicators in aromatic hydrocarbons solutions is not the most reliable method for the characterization of the acidity of solid catalysts. For example, in the present case, while the base used for the titration of the protons goes inside the polyacid grain, the indicator stays at its surface and the observed result is probably correlated to the acidity of the surface, not of the bulk. It can also be pointed out that before these studies with Hammett indicators, it was generally admitted that all acid sites of the solid polyacids displayed the same strength.We can now analyse the results obtained with the other heteropolyacids. In a first step, we can compare the initial heats of absorption of ammonia (in kJ mol-' of NH,), which give a measure of the acidity of the compounds. The values are: 196 kJ mol-' for H,PW,2040; 185 kJ mol-' for H4SiWl2O4,; 164 kJ mol-' for H,P2W2,O7,(H2O),; 156 kJ mol-' for H6P2W1gO,,. These values show that the order of acid strength is: H3PW12040>H4SiW12040> H6P2W21071(H20),>H6P2W18062, in agreement with pre- vious results which prove that Keggin compounds are much more acidic than Dawson ones and that increasing the number of protons in the Keggin heteropolyacid decreases the acid ~trength.'~*".'~ These values are also in agreement with the results of Tzumi et al.who found, by ammonia TPD, that silica-supported tungstic heteropolyacids show the acidity order: H3PW,2040>H4SiW,20,0.20 They agree also with recent results on the acidity of several solids which showed that H3PW12040is more acidic than the H-ZSM-5 zeolite but less than SO~-/Zr02.21 Let us now discuss the shape of the curves in Fig. 2. The difference between H,PW12040 and H6PzW18062 on the one hand and H4SiW12040 and H6P,W21071(H20), on the other could be related to a different stability of the heteropolyacids, a partial degradation occurring for the last two, as observed above for H3PW1204, as a function of the temperature of prior activation (Fig.3). However, differential thermal analysis showed that the stabilities of these compounds decrease in the order: H,PW12040 =H6P2W2,07,(H20), > H4SiW12040>H6P2W18062, and, thus, such an explanation is not appropriate. In order to explain these results, it is neces- sary to study the decreasing slopes of the calorimetric peaks for each pulse of ammonia. Indeed, the decrease of the heat can be approximated by an exponential function AH =AHo exp(-t/z) where t is the time and z the thermokinetic parameter. The value of z gives an indication of the speed of the sorption reaction and also in the case of heteropolyacids it allows a study of the diffusion of ammonia through the crystal. Fig.4 shows the evolution of the thermokinetic parameter for the four experiments presented in Fig. 2 as a function of the sorbed volume. For H3PW12040. the thermokinetic parameter does not increase with the amount of ammonia. This result is in agreement with the studies of Moffat et al. who showed that the ammonium salt of the phosphotungstic acid has a porous str~cture.'~*'~-~~ Thus, it is easy to understand that ammonia can diffuse very easily through the solid: it reacts first with the protons near the surface forming (NH4),PW1204,, which has a porous structure allowing NH, to react more easily with the acidic centres inside the solid. For H,SiW,,O,,, the same authors reported a similar stru~ture,'"'~ but the thermokinetic parameter shows that there is a drastic increase with the amount of sorbed ammonia, the diffusion through the crystal occurring more and more slowly.In order to elucidate this point, we undertook the study of the ammonium salt of H4SiW12040. We prepared it by the same way as Moffat (precipitation from an aqueous solution of H,SiW120,, by addition of ammonium carbonate, solid 1) and measured its BET surface area. This compound was then dissolved in hot water and solid 2 was obtained after precipi- tation by ammonium chloride. The BET surface area of solid 1 was ca. 120 m' g-' in agreement with the results of Moffat et al. and decreased to ca. 50 m2 g-' for solid 2. In addition, keeping solid 1 in a desiccator under ammonia pressure (by contacting it with NH3 vapour arising from some drops of a concentrated ammonia solution) led to solid 3 which had a BET surface area of 3.3 m' g-'.The infrared spectra of solids 1-3 are very similar to those reported for H4SiW12040 and its salts." Indeed, if a degradation of the anion occurred amount NH3/1o4 mol g-' Fig. 4 Thermokinetic parameter (in seconds) as a function of the ammonia uptake for: 0,H,PWl,O,O; a, H,SiW,,O,,,; Ir, H,P,W,,O,,; 0, H6P,W210,1(H,O), J. MATER. CHEM., 1994, VOL. 4 during the treatment by ammonium carbonate and/or ammonia, the first degradation product should be SiWl10;;, whose infrared spectrum is very different from that of SiW120:,.26 For example, the v(Si-0) band, found at 1020cm-' for SiW,,O& is shifted to lo00 cm-' in SiW,,O&.As a consequence, we can assume that the polyanionic struc- ture remains intact in compounds 1-3. Fig. 5 shows the X-ray diffraction spectra of solids 1 and 3, of K,SiWl,O,, and (NH4)3PW12040. lead to ident- Solid 1 and (NH4)3PW12040 ical XRD patterns showing that their structures should be quite similar, in agreement with their high BET surface areas. On the other hand, solid 3 and the potassium salt of the silicotungstic acid give patterns very different from the pre- vious ones. It must also be pointed out that infrared experi- ments showed that the heteropolyanion structure had been retained in these two compounds. Unfortunately, it was not possible to record the X-ray diffraction pattern of the sample of H4SiW,2040 which has been submitted to react with ammonia in the calorimetric cell but we can assume that its spectrum should be quite similar to that of solid 3. The XRD spectra of solid 1 and (NH4)3PW,2040 are also similar to those of H3PW12040.6H20 [Fig.5(e)]whose struc- ture has been refined by Brown et Only small variations of position arise from the smaller width of the peaks of the hexaaquo acid and from the replacement of (H502)+ by NH;. Intensity variations can be due to the different cations and/or to preferential orientations. These authors pointed out also that many alkaline salts of heteropolyanions showed similar XRD patterns and they gave the examples of the caesium salts Cs3PW1204,, Cs,HSiW,,O,, and CS,H,BW,~O,~. It is well known that CS~PW,,O~~ has a 2oool I1000 1 1 50011 I1 2000~ 1 1 20 40 60 80 2tYdegrees Fig.5 X-Ray diffraction patterns of: (a) solid 1 [(NH,),HS~W120,,~xH,0(x=5-6)]; (b) solid 3 [(NH4),SiW1204,-5 H201; (4 WLM"12040.8H20; (4 K4SiW12040*6H20;(4H3PW,,040~6H20 porous structure as (NH4)3PW,,0,,,16~22-24and so the follow- ing interpretation can be proposed. The alkaline salts of heteropolyanions having an XRD pattern similar to that of H3PW1,O4,-6H2O have a microporous structure and corre- spond to salts with only three alkaline cations [such as (NH4)3HSiWl,0,0]. The XRD patterns of solid 3 and of K4SiW12040 correspond to compounds with four alkaline cations and with no microporosity, as it can be shown by the BET surface-area measurements (K4SiW12040 has a BET surface area lower than 2m2g-').It is surprising to obtain (NH4)4SiW,,0,0 by the reaction of ammonia at room tem- perature, only the triammonium salt should be obtained by reaction at 150"C during the microcalorimetric experiment. However, the experimental conditions are completely dii€erent, time and ammonia pressure being some orders of magnitude higher in the first case. In addition, the presence of water in the reaction leading to solid 3 probably has a non-negligible role. Indeed, if it is well known that the structure of the porous water-insoluble heteropoly salts is not dependent on their water content (as water is only physisorbed), very probably there is initially reaction of water with the protons of the solid, leading to the formation of the highly mobile (H30)+ species which can then react with ammonia at the surface of the grain.Water can then be considered as a catalyst of the reaction. However, two questions remain. First, H3PW120,0~6H,0 should be microporous, but experimentally its BET surface area is found to be very small. This could be due to slightly different positions of the cations in the structure or to a blocking effect of residual water. This problem is, however, under study.27 Secondly, McMonagle and Moffat report16 the preparation of a non-porous ammonium 12-phosphotungstate salt by reaction of the acid with ammonia at high temperature. However, this result was not reported in their subsequent papers on this subject.As the preparation method was not well described, in particular the pretreatment of the polyacid and the quantity of ammonia introduced in the cell, it is difficult to discuss this point. Indeed, the thermokinetic param- eter increases with the activation temperature of 1Zphosphotungstic acid showing that there is a blocking effect due to the surface degradation of the poljanion. Moreover, if a high quantity of ammonia was introduced, it is also possible that this base reacts not only with the protons but also with the polyanionic species, leading also to a blocking effect due to a modification of the surface of the grains. In these two cases, the X-ray diffraction patterns should not be modified (in agreement with the experience) but the BET should not show microporosity.The shape of the microcalorimetric curves for the silico- tungstic acid can now be explained easily. Indeed, there is initially formation of the triammonium salt which is micro- porous and during the course of the experiment this salt reacts with ammonia leading to the formation of the tetra- ammonium salt. But this latter compound is not porous and ammonia diffuses through it very slowly leading to (i) an increase in the thermokinetic parameter and (ii) an apparent non-stoichiometric reaction, as it proceeds more and more slowly. Thus, if one waits sufficiently long, one will obtain an almost stoichiometric reaction, leading to nearly complete neutralization, with essentially no decrease of the initial heat of absorption with an increasing quantity of ammonia admit- ted.However, the decrease of the reaction rate with time makes the evaluation of the heats difficult, because it is not easy to integrate quantitatively broad signals. This explains the apparent decrease of the absorption heats (Fig. 2). It has been proposed by some authors that the crystals of alkaline salts of phosphotungstic acid are non-porous by themselves, the microporosity being in fact the space between J. MATER. CHEM., 1994, VOL. 4 mY' 100 120 140 160 180 200 7°C Fig. 6 Heat capacities C, as a function of temperature for acidic heteropolyanions (continuous line) and for their corresponding ammonium salts (dashed line): 0,H,PW,,O,,; 0,H,SiW,20,,; *,H6PZW18062;0,H6P2WZ1071(H20)3 ultrafine non-porous particles which were detected by electron microsc~py.~,~~This idea is also consistent with the structure determined by Brown et a2.I4 However, it seems to us that such an explanation is not reasonable for the following reasons.(i) If we assume that the microporosity is due to the interparticle space, such a phenomenon should also be observed for highly dispersed oxides such as Aerosil but it is not; (ii) the ultrafine particles have a diameter corresponding to less than 10 unit cells and so the X-ray diffraction pattern should be very broad, while in some cases we obtained very nice spectra allowing a structure determination by Rietveld refinement; (iii) finally we studied the p-xylene adsorption on the potassium salt of phosphotungstic acid and we found the same curve than for the ZSM-5 zeolite, in agreement with an intrinsic microporosity.As a consequence, it is necessary to assume that the alkaline salts of phosphotungstic acid are bidispersed systems with ultrafine particles and big crystals. This duality, which was previously observed for supported heteropolyacid~,~~*~~is responsible of the broadening of the foot of the X-ray diffraction peaks.27 As no data are available on the ammonium salts of H6P2W&6, and H,P,W,,O,,( one can only propose that a similar behaviour as above occurs for these compounds, the shapes of the curves being related to the porosity of the material. However, it is worth noticing the presence of two maxima in Fig.4 for H,P,W,,O,,(H,O)3 which could be explained by the formation of a mixed intermediate salt. Fig. 6 shows the evolution of the specific heats as a function of the temperature for the acidic and ammonium compounds of the four polyanions which were studied here. The determi- nation has been performed between 100 and 2OO0C, the temperature increasing by increments of 5 "C.Differences are more or less important between the acid and the correspond- ing ammonium salt, depending on the sample. For example, H3PW,,0,, presents a decrease of C, when going from the acid to the ammonium salt. This result is in good agreement with the fact that its density decreases after absorption of ammonia (Table 1).Indeed, the ammonium salt structure is more microporous and organized than the correspond- ing acid. In the same way, a decrease of C, for [P,W,,07,(H,0)3]6-is observed when going from the acid to the ammonium salt, corresponding to a decreasing density. On the contrary, the two other polyanions, whose density is only slightly modified by ammonia absorption, lead to a significant increase in C, when going from the acid to the salt, in agreement with the absence of microporosity and the lower organization of this latter. Conclusions The calorimetric study of the acidity of tungstic heteropoly- acids by ammonia absorption has shown that these are very acidic compounds displaying high absorption heats, which accounts for their use as substitution catalysts for sulfuric acid.The acidity varies between compounds, which so giving inorganic solid compounds of different acidities but with sites of the same strength. Diffusion phenomena are important for some of the compounds and are related to the structure of the heteropolyacids and their ammonium salts. F.X. L.-C. thanks the CNRS for financial support. References Y. Onoe, Kagaku Kogyo, 1975,26,355. Y. Oda and K. Uchida, Jpn. Petrol. Inst., 1977, 20, 1054. I. V. Kozhevnikov, Russ. Chem. Rev., 1987,56.811. M. Misono, Catal. Rev.-Sci. Eng., 1987,29, 269. I. V. Kozhevnikov and K. I. Matveev, Appl. Catal., 1983,5135. M. T. Pope, Heteropoly and Isopolyoxometalates (Inorganic Chemistry Concepts, Vol. 8), Springer Verlag, New-York, 1983. A.Auroux, J. C. Vedrine and P. C. Gravelle, Stud. Surf. Sci. Catal., 1982, 10, 305. A. Auroux, F. X. Cai and F. Lefebvre, Calorim. Anal. Therm., 1990,20-21,103. J. MATER. CHEM., 1994, VOL. 4 9 A. Auroux and J. C. Vedrine, Stud. Surf. Sci. Catal., 1985,20,311. 20 Y. Izumi, R. Hasebe and K. Urabe, J. Catal., 1983,84,402. 10 A. K. Ghosh and J. B. Moffat, J. Catal., 1986,101,238. 21 T. Okuhara, T. Nishimura, H. Watanabe and M. Misono. J. Mol. 11 B. Viswanathan, M. J. Omana and T. K. Varadajan, Znd. J. Chem., Catal., 1992,74,247. 1988,27A, 674. 22 J. B. Moffat, J. Mol. Catal., 1989,52, 162. 12 G. I. Kapustin, T. R. Brueva, A. L. Klyachko, M. N. Timofeeva, S. M. Kulikov and I. V. Kozhevnikov, Kinet. Catal., 1990,31,896. 23 D. B. Taylor, J. B. McMonagle and J. B. Moffat, J. Colloid Interface Sci.,1985, 108, 278. 13 14 15 A. Oulmekki, D. Olivier, G. Herve and F. Lefebvre, C. R. Acad. Sci. Paris,Sbie ZI, 1990,311, 619. G. M. Brown, M. -R. Noe-Spirlet, W. R. Busing and H. A. Levy, Acta Crystallogr. B, 1977,33, 1038. P. Pascal, Nouveau Traitk de Chimie Minkrule, Volume XIV, 24 25 26 J. B. Moffat, J. B. McMonagle and D. B. Taylor, Solid Statr Ionics, 1988,26, 101. C. Rocchiccioli-Deltcheff, R. Thouvenot and R. Franck, Spectrochim. Acta Part A, 1976,32, 587. C. Rocchicioli-Deltcheff and R. Thouvenot, J. Chein. Rex 16 17 18 Masson, Paris, 1959, p. 964. J. B. McMonagle and J. B. Moffat, J. Colloid Interface Sci., 1984, 101,479. K. Nowinska, J. Chem. Soc., Chem. Commun., 1990,44. A. Oulmekki and F. Lefebvre, React. Kinet. Catal. Lett., 1992, 27 28 29 30 Synopses, 1977,46. F. Lefebvre and B. F. Mentzen, unpublished results, 1992. N. Mizuno and M. Misono, Chem. Lett., 1987,967. F. Lefebvre, J. Chem. Soc., Chem. Commun., 1992,756. A. Oulmekki and F. Lefebvre, React. Kinet. Catal. Lett., 1992, 48, 601. 48,607. 19 B. Viswanathan, M. J. Omana and T. K. Varadajan, Catal. Lett., 1989,3, 217. Paper 3/03391K; Received 14th June, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400125

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. MATER. CHEM., 1994, 4( l), 133-137 Chemical Lithium Insertion into Sol-Gel Lamellar Manganese Dioxide MnO,.,,mH,O Philippe Le Goff," Noel Baffier," Stephane Bach,*b and Jean-Pierre Pereira-Ramosb a Laboratoire de Chimie Appliquee de /'€tat Solide, C.N.R.S. URA 302 ENSCP 7 7, rue Pierre et Marie Curie, 75237 Paris Cedex 05, France Laboratoire d'Electrochimie, Catalyse et Synthese Organique, C. N. R.S. UM 28 2, rue Henri Dunant, 94320 Thiais, France Chemical lithium insertion in the sol-gel manganese oxide MnO,.,,.nH,O has been performed using n-butyllithium as a reducing agent. Open-circuit voltage (OCV) experiments, IR, XRD and thermal analysis on lithiated samples Li,MnO,~,,~nH,O 0 <x 60.45show that lithium insertion into the host lattice induces a slight structural change: from x =0.05, a relative gliding of MnO, layers occurs, leading to the transformation of the initial hexagonal Structure into a closely related monoclinic phase.The higher the lithium content, the stronger the Li-hydrogen bond interactions which explains the contraction of the lattice observed along the c direction. The low magnitude of the structural changes is consistent with the high reversible behaviour found for this compound used as rechargeable cathodic material for lithium batteries. During the last several years, extensive research has been developed on manganese dioxide in order to obtain rechargeable Li/MnO, cells. From the wide variety of manga- nese dioxides, only the birnessite group exhibits a layered structure' which would theoretically provide attractive proper- ties as a rechargeable cathodic material for secondary Li batteries.Reversible insertion of ca. 1 Li ion per mole of oxide would be effectively possible without destruction of the host lattice thus leading to a theoretical Faradaic capacity of ca. 300 A h kg-'. The birnessite is usually synthesized by dehydration of buserite.2 This latter compound is prepared from an oxidation of an aqueous Mn(OH), suspension by oxygen or chlor-ine leading to sodium birnessite with the formula Na,,2sMnOl .,,-0.64H20. After removal of sodium ions from the structure by an acid treatment performed at room tempera- ture,, the Mn0,.84.0.71H,0 oxide is obtained. Regarding the corresponding sol-gel compound, a layered manganese di- oxide is obtained from mixed oxides Na,.,MnO,, KMnO, or K,.,,MnO, synthesized via a sol-gel pro~ess.~ These com- pounds are then transformed into a lamellar sol-gel oxide MnO,~,,.nH,O (0.60G n G0.76) by an acid treatment at room ternperat~re.~Contrary to what happens for the classical compound Mn01~,,-0.7 1H20,the highly preferred orientation of these latter is maintained during acid treatment.The layered structure of the sol-gel MnO,,,,~nH,O com-bined with these textural properties have been proved to induce remarkable electrochemical behaviour. Indeed, ca. 0.85 Li ions enter reversibly into the host lattice with high kinetics for Li transport, which corresponds to a specific capacity of ca.250 A h kg-' available in the potential range 4.2/2 V. Its good cycling behaviour has been previously ascribed to texture and structural properties. However, its electrochemical behav- iour has only been correlated to preliminary X-ray diffraction experiments. In particular, no data about the role of structural water have been provided. In order to achieve further insight into the Li insertion process in sol-gel MnO,.,,.nH,O and to specify the structural changes occurring during Li insertion-extraction, we have undertaken the structural investigation of chemically lithiated samples. In this paper, we report additional IR, X-ray, DSC and TG experiments as well as new potentiometric measure- ments which are discussed in relation to the electrochemical properties of the cathodic material.Experimental X-Ray diffraction experiments were performed with a Philips diffractometer using Cu-Kcc radiation. Thermal analysis measurements were performed in air at heating rates of 10 "C min-' using a Netszch STA 409 analyser with the simultaneous recording of weight losses (TG) and temperature variations (DSC). The mean oxidation state 'Z' of manganese was determined by a chemical titration using iron@) sulfate6 with an accuracy of f0.02. Chemical analysis of the compounds was made by atomic absorption measurements with a Varian 2150 appar- atus. Two successive sulfuric acid treatments of sol-gel Nao.,MnO, prepared as mentioned in ref. 5 ensure a complete removal of sodium ions from the compound.This was checked by atomic absorption experiments. Infrared spectra were recorded on a Perkin-Elmer 783 spectrophotometer by grinding the powder into KBr pellets. The electrochemical measurements were performed in pro- pylene carbonate (PC), twice distilled, obtained from 'Fluka' and used as received. Anhydrous lithium perchlorate was dried under vacuum at 200°C for 12 h. The electrolytes were prepared under a purified argon atmosphere. Results and Discussion According to the literature,' the structural model for the manganese(1v) oxides with layered structure is chalcophanite ZnMn,O7-3H,O. This structure consists of single sheets of water molecules between layers of edge-sharing MnOb octahedra, with Zn atoms located between the water layer and oxygens of the MnO, layer.In the case of classical birnessite, it is necessary to substitute Mn2+ ions for Zn2+ ions. The stacking sequence along the c-axis would thus be: O-Mn'V-O-Mn"-H,O-Mn"-O-Mnlv-O-and the perpendicultr distance between two consecutive MnO, layers is ca. 7.24 A. In this type of structure, vacancies exist in the layer of linked MnO, octahedra and the total water content is variable. Chemical Synthesis of Lithiated Compounds From thermodynamic data, butyllithium is well known to be particularly suitable to perform the Mn4+ reduction into Mn3+ ions., Since the lithiation reaction occurs at room temperature the chemically reduced compound should be similar to the electrochemical one.The reduction reaction, which has already been described in the literature,' was carried out as follows. The sol-gel compound (1g) was mixed with a hexane solution in a test-tube first placed in an inert atmosphere. An appropriate ratio of n-butyllithium in hexane was added and the mixture was stirred for a period of 1 or 2 days at room temperature. The reaction took place according to: MnO,.,,.nH,O +xC,H,Li~Li,MnO,,,,.nH,O +x/2C,H,, (1) The final product was obtained after filtering, washing with hexane and drying at room temperature. The amount of lithium in the compound was determined by atomic absorp- tion spectroscopy. The following x values were obtained: 0.05; 0.1; 0.15; 0.20; 0.25; 0.30; 0.35; 0.40; 0.45. Electrochemical Data Electrochemical Li Insertion Li insertion into sol-gel MnOl,8s.nH,0 has been shown to occur reversibly in two steps in the potential range 4.25/2 V us.Li/Lif with a maximum Faradaic yield equal to 0.85 F per mol of oxide [Fig. l(a)]. From these results it was suggested that only Mn" ions were involved in the charge transfer so that the most convenient formula for the sol-gel compound is Mn'V,~,5Mn",~,s01,8,.nH,0.9 Previous thermodynamic and kinetic results as well as the XRD experiments performed on electrochemically lithiated electrodes indicated a two-step process for the Li insertion reaction: for the Li content 0 <x <0.4, Li insertion occurs in the potential range 4.25-2.85 V with high kinetics for Li transport, while for a higher Li concentration range 0.4 <x <0.9 Li insertion is more difficult and a voltage quasi- plateau at ca.2.8 V is observed. In terms of cycling capacity, the specific capacity slightly decreases from an initial value of 200 A h kg-' (0.7 F mol-') to ca. 150 A h kg-' by the 50th cycle. In a few words, the major advantage of the sol-gel lamellar manganese dioxide consists of an unusually high depth of discharge which makes this compound a very promis- ing rechargeable MnO, in comparison with various MnO, forms studied for secondary Li batteries.'@I2 However, some unclear points remain; in particular, the unusual presence of water molecules in the structure of the material raises some questions about its effect on the Li insertion process.Chemical Li Insertion: OCVMeasurements For each composition, ca. 20 mg of pure chemically lithiated compound was pressed on a stainless-steel grid with a geo- 4400 40001. Fig. 1 OCV curves for (a) electrochemically and (b)chemically lithi- ated Li,MnO,,,,.nH,O samples J. MATER. CHEM., 1994, VOL. 4 metric area of 1 cm2. Under these conditions, equilibrium is considered to be reached when the open-circuit voltage remained stable (21 mV) for 20 h. Conversely, the composition change is ensured by coulo- metric titration for electrochemical experiments. As shown in Fig. 1, very similar OCV curves are found for the samples reduced according the two lithiation methods. This indicates the suitability of using data drawn from the present study on chemically lithiated samples to investigate the electrochemical behaviour of the sol-gel Mn0,,85.nH,0.X-Ray Diffraction Analysis of the Lithiated Compounds Fig. 2 shows the evolution of X-ray diffraction patterns for different amounts of Li in the compounds. For x =0, starting sol-gel lamellar manganese dioxide, two well defined diffrac- tion peaks, 001 and 002, correspond to the d-spacing between two consecutive Mn0, layers of the lamellar structure. This parameter is weakly dependent oq the water content, n: d= 7.24 A for n =0.64 H20up to 7.32 A for n =0.76 H20. Accord- ing to the position of the other diffraction peaks, th,e structure is of hexagonal type, with parameters a= 2.84 A and c= 14.64A. This sLructure is consistegt with the a-Na,.,MnO, form (c = 11.12 A, &spacing= 5.56 A), but with a higher inter- layer distance., Increasing of the interlayer distance is due to the extraction of Na' ions from the crystal framework, i.e.to an increasing MnIV content, during the acid treatment. From x =0.05, new diffraction peaks appear, the intensity of which grows with x,whereas intensities of peaks corre-sponding to the hexagonal structure decrease. From x =0.25, only the new diffraction peaks are present on the diagram. They correspond to .a monoclir$c phase with the following parameters: a =5.15 A; b =2.86 A; c = 14.29 A; p= 102.6". For higher Li contents, only the monoclinic phase exists, but ma- terials become less and less well crystallized. The hexagonal- monoclinic transformation is a well known phenomenon in Mn0,-based compound^.'^ The relationships between the parameters are (Fig.3): aMono ahexJ3 bMono ahex CMvlono sin P ZZ chex Such a structural change can be explained by a glide of the layers in the ab plane, one MnO, layer relative to another. The evolution of the spacing distance (c/2) uersus x is shown in Fig. 4. Two sets of data are achieved in the composition ranges 0 <x <0.02 and 0.05 <x <0.45 for which the interlayer spacing of the hexagonal and monoclinic phases decreases slightly (ca. 2%). A biphasic region is evidenced for 0.05<x<0.20 while a and b parameters of the monoclinic phase do not change significantly all along the lithium inser- tion process.Hence, Li insertion originates a hexagonal-monoclinic distortion corresponding to a maximum contrac- tion of the interlayer space (of the order of 5%) for 0 <x <0.9, with a volume change < 1%. In other respects, the lower interlayer space of the monoclinic phase combined with the higher lithium content makes the kinetics of lithium transport slower from ~~0.25.~ Nevertheless as reported,' no loss of reversibility is noted whatever the depth of discharge. IR Analysis of the Lithiated Compounds The evolution of infrared spectra uersus Li content x is shown in Fig. 5. The OH stretching region in the high-frequency spectrum between 4000 and 2500 cm-' corresponds, for x = 0 to a broad asymmetric absorption band centred near 3350 cm-', constituted of two components (3500 and J.MATER. CHEM., 1994, VOL. 4 * 0 I-. 6 10 15 20 25 6 10 15 20 25 * 0 I. 6 10 15 20 25 , (f) 6 10 15 20 25 Hdegrees Fig. 2 Evolution of X-ray diffraction patterns for x=O (a); 0.05 (b); 0.1 (c); 0.15 (d); 0.25 (e); 0.45 (f) in chemically lithiated LiXMnO,,85.tIH,Osamples. For x =0, the diffraction peaks are charac-teristic of a hexagonal phase. From x=O.O5, new diffraction peaks (*) appear, corresponding to a monoclinic phase. For x=O.25, all the diffraction peaks are characteristic of the monoclinic phase only 3300cm-l) due to two kinds of hydroxyl group. As lithium accommodation proceeds, the stretching vibration band .assigned to hydroxyl groups situated near 3500 cm-' is pro-gressively shifted towards lower frequencies, leading for x= 0.45 to a more symmetric absorption band centred near 3400 cm-'.Such an evolution indicates the strengthening of amon I I .z I I 1sI I 13 I I II I igI I I II I I I I Fig. 3 The hexagonal-monoclinic transformation is a well known phenomenon in Mn0,-based compounds from ref. 13;(-) monoclinic; (---) hexagonal 6.80 0 0.10 0.20 0.30 0.40 0.50 lithium ratio Fig.4 Evolution of interlayer spacing with lithium content x for chemically lithiated Li,Mn0,,85.tIH,0 samples. The interlayer spacing is c/2 and c sinb/2 for the hexagonal (W) and monoclinic (0)phases, respectively HO-Li interactions during lithium accommodation.In fact, IR measurements show that the higher the degree of reduction, the stronger the interaction of Li' ions with hydrogen bonds between the water layer and oxygen atoms of the MnO, layers. Hence, it can be assumed that such a phenomenon explains the notable decrease of the interlayer spacing pointed out in XRD, especially in the Li composition range 0 <.Y d 0.2. However, since the samples have not been preserved from moisture, the basic character of interlayer water due to the presence of lithium ions results in a notable dissolution of atmospheric CO, to give Li,C03 as shown by the emergence of intense specific absorption bands at 1380 and 88Ocm-' for x>0.2. Thermal Analysis of the Lithiated Compounds Thermal analysis of lithiated compounds has been performed in order to study the dependence of the water content on the temperature. In the case of the starting sol-gel material MnO,,,,.nH,O, the simultaneous recorded thermal analyses (TG and DSC) were characterized by a total weight loss of ca.20% (Fig. 6).From room temperature to 250 "C,only one endothermic peak appeared at 120 "C corresponding to a weight loss of 11.5%, due to the departure of the interfoliar water (weakly bonded water). The weak weight loss (ca. 3%) observed up to 450 "C is probably due to the removal of more strongly bonded water or hydroxy groups. The last endo-thermic peaks, located at 540 and 950°C correspond to the transformation into a-Mn,03 and Mn304,respectivel\-.. J.MATER. CHEM., 1994, VOL. 4 n I I I I I I,,I I I I I 10 3200 2400 1600 1200 800 400 wavenumbedcm-' Fig. 5 IR spectra of Li,MnOl,,,~nHzO samples where x=(a) 0; (h)0.05;(c) 0.10; (d)0.15; (e)0.20; (f) 0.25; (g)0.30 (h)0.40 (i) 0.45 I-251 , , , 1-25 (u -30 0 200 400 600 800 1000 TI"C Fig.6 Simultaneous thermal analysis (TG and DSC) of the sol-gel lamellar manganese oxide for x =0 (20<T/"C<900). Thermal analy- sis measurements were performed in air at heating rates of 10 "Cmin-l The evolution of TG and DSC curves between 20 and 500 "C versus Li content x is shown in Fig. 7. From x =0.05, the main endothermic peak at 128°C splits into two well shaped peaks. The first one is located at ca. 1OO"C, while the other, more important, appears at 170 "C.Thus, insertion of Li' ions induces a new localisation of water leading to the formation of more strongly bonded water. In addition, the emergence of a broad exothermic peak located at ca. 220°C which increases with the lithium content, corresponding to the strengthening of the HO-Li interactions, is consistent with the IR analysis. Beyond 350 "C, X-ray diffraction analysis reveals the pres- 10--0---10--I I I -20 40 --20 -jj90 i 0--2010 100 .-200 300 400 500 1°C Fig. 7 Simultaneous thermal analysis (TG and DSC) of chemically lithiated sol-gel lamellar manganese oxide for x=O (a); 0.05 (b); 0.10 (c); 0.45 (d) in Li,Mn0,.,,~nH,0 samples (20d T/"C<500). Thermal analysis measurements were performed in air at heating rates of 10°Cmin-l ence of the Li,,,Mn,O, spinel structure which could be correlated to the small additional exothermic peak observed in thermal analysis (Fig.7). TG data show that the quantity of water which is given off before 200 "C remains approximately constant (ca. 0.6 H,O/Mn) indicating that lithium insertion does not affect the concentration of interlayered water. However, beyond J. MATER. CHEM., 1994, VOL. 4 x=0.25, a slight and regular decrease of interlayered water is observed. This departure of water is correlated with the formation of Li,CO, evidenced on IR spectra. Conclusions The present work has clarified the lithium insertion process into the lamellar compound Mno1,8,*nH2O synthesized via the sol-gel process.IR, XRD and thermal analysis of chemi-cally reduced samples Li,MnOl,8,.nH,0 (0<x <0.45) have shown that lithium ions enter the interlayer space leading to strong HO-Li interactions. Interlayer hydrogen bonds are then significantly affected which could explain the contraction of the host lattice in the c Girection: the interlayer distance diminishes from 7.29 to 6.9 A for O<x <0.35. A similar trend was found for this parameter in the case of electrochemically li thia ted samples.’ The lithium insertion process into the hexagonal structure of the sol-gel MnO,,,,.nH,O results from x=O.O5 in a glide of MnO, layers, leading to the formation of a monoclinic phase from x=O.25 with a biphasic region for 0.05dxd0.2.However, the monoclinic phase is very close to that of the initial hexagonal structure lattice. From x =0.25, the shorter monoclinic interlayer distance induces a decrease in the kinetics of lithium transport in the lamellar compound Mn0,,8,.nH20. Finally, the structural rearrangements are then minimized as lithium insertion proceeds, which is in good agreement with the high electrochemical reversibility encountered for this MnO,,,,.nH,O material. The authors are grateful to Dr. J. P Labbe for valuable discussions and to H. Lepesant for IR spectra recording. Financial support by the Direction des Recherches, Etudes et Techniques (DRET) is gratefully acknowledged. References 1 R. G. Burns and V. M. Burns, in Proceedings of the Manganese Dioxide Symposium, I.C.Sample Office, Cleveland, OH 1975, p. 305. 2 R. Giovanoli, E. Stahll and W. Feitknecht, Helu. Chim. Acta, 1954, 31,2322. 3 R. Giovanoli, E. Stahll and W. Feitknecht, Helu. Chim. act^, 1970, 53,453. 4 S. Bach, M. Henry, N. Baffier and J. Livage, J. Solid State Chem., 1990,88,325. 5 N. Baffier, S. Bach and J. P. Pereira-Ramos, in Solid Statt lonics, ed. M. Balkanski, T. Takahashi and H. L. Tuller, Ihevier, Amsterdam, 1992, p. 55. 6 M. J. Katz, R. C. Clarke and W. F. Nye, Anal. Chem, 1956, 28, 507. 7 M. B. Dines, Muter. Res. Bull., 1975,10,287. 8 J. Rouxel, J. Chim. Phys., 1986,83, 850. 9 S. Bach, J. P. Pereira-Ramos, N. Baffier and R. hlessina, Electrochim. Acta, 1991,36, 1595. 10 T. Nohma, Y. Yamamoto, K. Nishio, I. Nakane and N. Furukawa, J. Power Sources, 1990,32,373. 11 J. M. Tarascon, E. Wang, F. K. Shokoohi, W. R. McKinnon and S. Colson, J. Electrochem. SOC.,1991, 138,2859. 12 J. M. Tarascon and D. Guyomard, J. Electrochem. Soc., 1991, 138,2864. 13 J. P. Parant, R. Olazcuaga, M. Devalette, C. Fouassier and P. Hagenmuller, J. Solid State Chem., 1971,3, 1. Paper 3/03381C; Received 1lth June, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400133

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. MATER. CHEM., 1994,4( l), 139-144 Electroluminescence of Ho3+ Ions in Semiconducting Polycrystalline Zinc Oxide Electrodes in Contact with Aqueous Electrolyte Jean-Claude Ronfard-Haret, Kazufumi Azuma, Sai'd Bachir, Djakaridia Kouyate and Jean Kossanyi Laborafoire de Photochirnie Solaire, CNRS, 2-8, rue Henri Dunant, 94320 Thiais, France Sintered semiconducting zinc oxide has been doped with varying amounts of Ho3+ ions and its electroluminescence has been studied under various polarization conditions. Under cathodic polarization and in the presence of persulfate ions, the emission of the ZnO matrix only was observed. Conversely, under anodic under anodic polarization, character- istic emission bands of the Ho3+ ions were generated. The intensity of these bands was studied as a function of both the polarization potential and the rare earth doping level.In the 9.8-10.2 V vs. SCE range, the log of the emission intensity varied inversely with the square root of the applied potential, according to the Alfrey-Taylor relationship. It showed a maximum for a Ho3+ doping level of ca. 0.5 atom%. Analysis of the samples' polycrystalline structure using scanning electron microscopy (SEM) and energy dispersive spectroscopy (EDS) techniques indicates that the rare- earth ions create islands on the surface of ZnO which induce dislocations in the polycrystalline structure. The electroluminescence of rare earth (RE3 +)-doped semi- conductors is attractive because of its numerous applications.' Surprisingly, of the RE3 + -doped 11-VI semiconductors, most literature deals with sulfides, selenides or tellurides, and only do our recent articles report the electroluminescence of the RE3+ ions inserted into semiconducting zinc o~ide.~-~ Rare earth-doped zinc oxide electrodes show two types of electroluminescence depending upon the polarization of the electrode in the electrochemical cell. Under negative bias the light originates from the recombination of electron-hole pairs (e-/h+).The spectrum shows two features: a sharp band centered at 390 nm corresponding to the direct band to band radiative recombination, and a broad band centred around 550 nm which corresponds to recombinations at self-activated centres intrinsic to ZnO. This electroluminescence observed under negative bias is very similar to the photoluminescence of zinc oxide obtained by 390 nm ex~itation.~,~The only difference lies (i) in the way in which the holes are created in the valence band: by injection from a redox couple in solution for the electroluminescence process and by the promotion of an electron from the valence band to the conduction band in the case of photoluminescence; (ii) in the intensity of the reabsorption, by the rare earth incorporated into the matrix, of the light emitted by zinc o~ide.~?~?~ Under positive bias, the characteristic emission of the RE3+ ions is observed in addition to the luminescence originating from zinc oxide.8 Under such conditions, the luminescent centres of zinc oxide are excited by two different processes: direct electron impact excitation or impact ionization of the medium, while excitation of the RE3+ ions occurs only by a direct electron impact-excitation process.The above fundamental results which differ with the polariz- ation mode have been interpreted in terms of the absence of coupling between the energy levels of the semiconducting matrix and those of the lanthanide. This is corroborated subsequently by SEM which shows that the RE3 + ions induce dislocations inside the zinc oxide microcrystalline structure5 as it does inside zinc sulfide.' This result lead us to question the method of rare earth insertion into zinc oxide. In a previous paper,, we reported some preliminary results concerning the electroluminescence of Ho3+ and Sm3 ions+ when inserted in semiconducting zinc oxide.In the present article we report a systematic study of the electroluminescence of the Ho3+ ions inserted in ZnO, and the results obtained by SEM and EDS techniques carried out upon polycrystal- line samples. Experimental Sintered RE3 +-doped polycrystalline ZnO pellets were pre- pared by mixing intimately 0.04-5.0 atom% of the lanthanide (99.9% pure from Rhbne-Poulenc) to zinc oxide (99.99% pure from Koch-Light) in the presence of a small amount of ethanol. The mixture of the two oxides was pressed (4 tons cmP2) in a Specac press, then the pellets were heated for 5 h at 1100"C, under atmospheric pressure in an Adamel Lhomargil furnace. The electrodes were prepared by covering one face of the pellet with the In-Ga eutectic (both In and Ga metals from Prolabo). A copper wire was fixed to this ohmic contact by silver resin (Eco-Bond, Emerson and Cummings).The resulting preparation was then glued to a glass support by a chemically inert epoxy resin (Araldite). The electrodes were polished and etched with 1mol dm-3 HCI. The electrochemical cell used was a classical three-electrode system with a saturated calomel electrode (SCE) as reference and a platinum counter-electrode. Potential polarizations were generated by a P.A.R. model 173 potentiostat monitored by a P.A.R. model 175 programmer system. For the electrofumi- nescence measurements, the electrochemical cell replaced the usual cell holder of a Perkin-Elmer MPF-44B spectrofluori-meter. The photocurrent spectra and the photoresponse of the electrodes were measured using an irradiation set-up built in the laboratory with a 450Watts XBO Xenon lamp (Cunow) and an H25 monochromator (Jobin-Yvon). In order to pre-vent rapid ageing of the electrode, pulsed (instead of continu- ous) polarizations, between 0 V and the working potential, were applied for recording the electroluminescence spectra.Characterization of the electroluminescence under cathodic polarization was performed with an NaCl (1 mol dmP3) + Na,S208 (0.5 mol dmP3) aqueous solution; for the anodic polarization measurements, the electrolyte was a 1mol dmP3 NaCl aqueous solution sometimes containing 0.3 rnol dm -3 Na,S,03.For time-resolved experiments, the duration of the anodic bias pulse and the relaxation time were adjusted separately to achieve a 95% signal decay. The synchronized electrical signals one originating from the potentiostat and the other from the photomultiplier, were recorded on a Tektronix 7834 storage oscilloscope. The charge resistance was adapted to the resolution needed. The energy dispersive detector (Tracor Northern system) was mounted on a scanning electron microscope (JEOL JSM 840 A). The major components of the sample were identified by superimposing labeled cursors to each element over a recorded spectrum. The chemical analysis of the spots of micrographs was determined by looking at the energy disper- sive X-ray spectra obtained under electronic irradiation. X-ray fluorescence spectrometry is a valuable technique for element analysis at the submicrometre level.The electron beam focussed onto the sample generates X-rays which correspond to the electronic relaxation between the initial and final excited states of the atom and is then analysed by the spectrometer. Results and Discussion Semiconducting properties The semiconducting properties of the Ho3 +-doped ZnO elec- trodes were verified by photocurrent measurements in a neutral medium (pH 7, phosphate buffer) in the presence of 0.4 mol dmP3 Na,SO,. Despite the absorption of light by the Ho3+ ions around 450 and 550 nm6,7, the photocurrent action spectra display only, as in the case of Er3+-doped ZnO electrode^,^ the characteristic response of the band to band absorption of zinc oxide below 390 nm.The current-voltage response of the electrodes is typical of an n-type semicon- ductor. As observed here and already reported for the Sm3+- doped electrode^,^ the cyclic voltammograms indicate that the semiconducting properties of the RE3+-doped electrodes are less marked than those of the undoped ones. Electroluminescence measurements Cathodic Polarization In the absence of persulfate ions and for applied potentials ranging from 0 to -4 V us. SCE, no light emission is observed which would originate from the electrode. When persulfate ions are added to the solution a white-greenish light emission is observed the spectrum of which is given in Fig.1. As reported previously for other RE3+ -doped ZnO electrodes, this spectrum exhibits two patterns: a sharp band centred at 390 nm which corresponds to band-to-band e-/h+ recombina- tions, and a broad band centred around 550 nm corresponding to e-/h+ recombinations at ZnO intrinsic self-activated centres. The mechanisrn6,l0,l' put forward to explain the electroluminescence phenomenon involves the injection of holes into the valence band of the semiconductor. These holes are produced by the second reaction of the two-step reduction of the persulfate ions: I I I 400 500 600 700 wavelengthhm Fig. 1 Electroluminescence spectrum under cathodic polarization of a 0.6 atom% Ho3+-doped ZnO electrode pulsed between 0.0 and -4.0 V in 1 mol dm-3 aqueous Na,S,O, electrolytic solution J.MATER. CHEM., 1994, VOL. 4 where the subscripts CB and VB refer to conduction and valence bands, respectively. The redox potential of the second reaction lies below the valence band edge." Therefore, although electron/hole pairs are created in the semiconductor, neither emission from the rare-earth ion, nor any emission originating from a new emitting centre could be detected. The only difference with the electroluminescence of pure zinc oxide comes from the weak reabsorption around 560nm by the Ho3+ ions of the light emitted by ZnO, as observed previously under the same conditions when Er3+ (ref. 3) and Nd3+ (ref. 4) ions are the dopant. Anodic Polarization In 1 mol dm-3 NaC1, the etched Ho3+-doped polycrystalline semiconducting zinc oxide electrodes show a weak green light emission when submitted to a polarization potential higher than +5.5 V us.SCE (Fig. 2). An anodic current is also detected. Both reach a plateau between 7 and 9.7 V us. SCE. Increasing the applied potential above 9.7 V us SCE induces a strong increase of both the emitted light and the anodic current. The variation of the anodic current is similar to the one already reported by Kiess12 for pure zinc oxide monocrys- talline electrodes. The only difference is the value of the second potential threshold found for the anodic current which occurs after the plateau: 9.7 V us. SCE as compared to the ca. 50 V reported by Kiess12 for monocrystalline ZnO.The spectrum obtained is characteristic of the Ho3+ ions (Fig. 3). By analogy with published spectra13p15 of the Ho3+/ZnS electro-luminescent system, we have attributed the series of bands centred around 500, 550, 650 and 760nm to the 5F3+518, 5s2+518, 5F3-+517, transitions of and 5s2-+s17 the Ho3+ion, respectively. In the 5.5-9.7 V t's. SCE potential range, only the most intense emission at 550 nm, correspond- ing to the 5S2-+518 transition, could be detected. If the electrode has not been etched or if it has been purposefully aged, the electroluminescence and the anodic current can only be detected when the applied potential reaches ca. +9.7 V us. SCE. However, in a pure solution of NaC1, the light emission is unstable, and its intensity decreases with time.A partial dissolution of the electrode and an oxygen evolution are observed. As already disc~ssed,~-~ the addition of thiosulfate ions to the electrolytic solution stabilizes the emission of light and prevents both the dissolution of the electrode and the evolution of gas. But once the electrode has been aged, the addition of thiosulfate ions to the electrolytic solution does not allow recovery from the loss of both current and luminescence observed previously at low voltage; only 0.51 1 I, I "'I' I, ~ E 0.30.41 0.21 I I lloo 0 24 6 8 10 12 VN vs. SCE Fig.2 Dependence of the anodic current i (solid line) and of the emitted light B (points) observed at 550nm i:ersus the applied potential V vs.SCE for an etched 0.5 atom% Ho3+-doped ZnO electrode. Electrolyte: 1 mol dm-3 NaCl aqueous solution. J. MATER. CHEM., 1994, VOL. 4 55, -51, 5s2-51, 5F3-518 A A 500 600 700 800 wavelengthlnm Fig. 3 Electroluminescence spectra under anodic polarization of a 0.6 atom% Ho3+-doped ZnO electrode pulsed between 0.0 and 9.88 V (us. SCE) in a 1 mol dmP3 NaCl and 0.3 mol dm-, Na,S,O, aqueous electrolytic solution fresh etching creates the conditions necessary for observation of the current and the luminescence below +9.7 V. This behaviour suggests that the oxidation of an intrinsic ZnO surface state, evidenced by Kiess,12 does occur. The variations of both the emitted light and the anodic current around 5V us.SCE are too weak to be analysed correctly, but in the 9.7-10.0 V us. SCE potential range, the analysis of the variation of the emitted light intensity B as a function of the applied potential V shows that the studied electrodes follow an Alfrey-Taylor-type relationship (Fig. 4)16: B=B, exp(b/V1/2) which indicates clearly that the light emitted by the electrodes is controlled by a Mott-Schottky type potential barrier, the width of which varies with the square root of the applied potential. In previous paper^,^.^' two different mechanisms for gener- ating hot electrons at the ZnO/electrolyte interface were proposed: (i) electrons originate from the oxidation of a surface state of ZnOs O:;f-+@2 +2e~~ (3) 10 ooor 1000?-100-10: t 11 I I I I I I 0.31 0.312 0.314 0.316 0.318 0.32v-'/2 Fig.4 Alfrey-Taylor plot of the electroluminescence intensity observed at 550 nm for an unetched 0.6 atom% Ho3+-doped ZnO electrode, in a 1mol dm-, NaCl and 0.3 M Na,S,O, aqueous electrolytic solution and depending on the relative position of the 0fuTf and of the electrolyte redox potentials, either the surface of the electrode or the electrolyte is decomposed: ZnOsurf+Zn2+ +402+2e- (4) H20surf+2H++302+2e-(5) Addition of thiosulfate ions to the solution results in the replacement of eqn. (4)and (5) by 2S20:-440g-+2e-(6) and consequently prevents the degradation of the electrode. Due to the proximity of the redox potential of the S20i-/S40g-couple (Eo=0.09 V us.normal hydrogen elec- trode)18 to that of ZnO surface states8,I2 (which act as an electron source in the absence of thiosulfate ions), these surface states catalyse the electron exchange at the electrode s~l-face.'~ (ii) Electrons are originating from the valence band by a band-to-band tunnelling process." This process generates holes in the valence band of ZnO at the surface of the electrode which must induce an electrochemical dissolution of ZnO and an oxygen evolution ZnO+2h+-+Zn2+ +to2 (7) H20+2h++2H+ +to2 (8) Holes are able to react not only with the electrode surface or with water but also with any redox system present in the electrolyte. Thus with thiosulfate ions, the reaction (9) equival- ent to (6) occurs rather than (7) and (8) 2S20i- +2h+-+S4O2- (9) Once injected into the conduction band, the electrons are accelerated by the high electric field until they gain enough energy to impact-excite the rare-earth ions.The concomitant variation of both the luminescence and the anodic current agrees with the electron impact-excitation mechanism pro- posed in our previous st~dies.~-~ That only the emission characteristic of the transitions between the 4f levels of the Ho3+ ions inserted into ZnO is observed, led us to disregard the hypothesis20,21 of new emissive donor levels created by the RE3+ ions in the gap of the semiconducting ZnO; instead, a direct impact-excitation process of the RE3+ ions, induced by the high energy electrons as a result of their acceleration by the applied potential, is proposed.If the excitation of the rare earth were the result of an energy transfer from the host semiconductor to the RE3' ions, then the same results should have been obtained what- ever the excitation mode (anodic or cathodic). Therefore the I I I0.10.001 0.01 0.1 1 10 N (atom%) Fig.5 Dependence of the relative values of Vphys upon the Ho3+ doping level for different electrodes observed at 550 nm in d 1 rnol dm-, NaCl and 0.3 mol dm-3 Na2S,03 aqueous electrolytic holution Fig.6 SEMs of the ZnO samples doped with (a) 0.4,(h) 1.0, (c) 3.0 atom% Ho3+.S and B refer to the EDS spectra of Fig. 7. Note that the magnification of (a) is 1.5 x greater than (b)and (c). emission observed in both polarization modes should be identical.This is not strictly the case here, since the only emission observed is characteristic of either ZnO or Ho3+ depending upon the polarization mode of the electrode. Hence, no coupling can exist between the energy levels of the RE3+ ions and those of the semiconductor. Furthermore, in the emission spectra obtained under both polarizations, no new structure originating from a new species could be observed and all the observed luminescence bands could be attributed to transitions originating from ZnO itself or from HO~+ions. This verifies the absence of any donor level created by the RE3+ions inside the gap of ZnO. The technical efficiency qtech of an electrode2' is given by where Pelecrefers to the electrical power consumed in the J.MATER. CHEM., 1994, VOL. 4 Zn Zn Fig. 7 EDS spectra (a) recorded over the entire Fig. 6(c) micrograph; (b)recorded over the area S, of Fig. 6(c) micrograph; (c) recorded on 6(a)micrograph: solid line area B,, dotted line area S, electrode and Plumto the power emitted by the same electrode. But instead of qtech the physical yield qphys. is often used for experimental ease and for conceptual simplicity: number of photons emitted out qphys = number of charges transferred and this physical yield has been decomposed into three parts23 rphys =qexc qrad vopt where qexc is the ratio of the number of luminous centres excited over the number of charges transferred; qrad is the number of luminous centres which deactivate radiatively over the number of excited luminous centres and rIopt is the number of photons which emerge from the electrode over the number of luminous centres which do deactivate radiatively.Using the same method as that described previously for evaluating the luminescence of Er3+ (ref. 3) and Sm3+ (ref. 5) in ZnO, relative values of qphys have been measured for several +electrodes of different RE3 doping levels ranging from 3.0 x to 3.0 atom%. The results shown in Fig. 5 show a maximum efficiency for an Ho3+-doping level around 0.5 atom%. Below this value qphys varies linearly with the Ho3+-doping level. As such a dependence is expected for qexc,23'24both qrad and qOptmust remain constant. As already de~cribed,~,'qopt, which depends upon the experimental con- ditions, was found to be constant for all the observations.The invariance of qrad was verified by measuring the luminescence decay of the 5S2+518transition observed at 550nm for two different Ho3+-doping levels, i.e. 0.04 and 0.6 atom%. In both J. MATER. CHEM., 1994, VOL. 4 cases the luminescence created by a short polarization pulse decays exponentially with a rate constant equal to 1.10kO.2 x lo4SKI.Thus, the dependence of qphys upon Ho3+ is limited to the variations of qexc. Scanning Electron Microscopy and Energy Dispersive Spectroscopy The SEMs of Ho3+-doped ZnO are given in Fig. 6. They show that an increase of the Ho3+ concentration induces a decrease of the average size of the grains and an increase of the number of holes at the grain boundaries.As reported previously for Sm3+ ions,5 the presence of the rare earth prevents the growth of the grains, limits their microcrystalline structure, and induces dislocations in the crystal lattice of ZnO. But in addition a new feature appears [Fig. 6(a), 6(b) and (c)] as clear irregular spots. The EDS analysis of the samples of Fig. 6 is shown on Fig. 7. The spectrum shown in Fig. 7(a)is recorded over the entire 3 atom% sample [Fig. 6(c)] whereas that shown in Fig. 7(h)is recorded over a selected area centred around the white spot, S,, on the right side of Fig. 6(c). The difference between these two spectra arises from the intensity of the signals of Ho atoms. This intensity, higher on the spectrum of Fig.7(h)than on that of Fig. 7(a),indicates a higher Ho concentration on the area centred around the spot. The EDS analysis of the SEM shown in Fig. 6(a)is reported in Fig. 7(c). Two spectra are presented. The first is recorded over an area B,, centred on the left side of Fig. 6(a) and excluding any spot. The second is recorded over an area S, centred around the white spot near the centre of Fig. 6(a). Clearly both spectra present the same features corresponding to Zn atoms whereas only the second one shows the patterns of Ho atoms. The same results are observed on Fig. 6(b).The patterns of Ho atoms are present only on the spectrum recorded on an area including a white spot. Hence the white spots on the SEMs are attributed to rare earth aggregates, the presence of which has been postulated in previous on the basis of luminescence kinetics measurements, and already characterized in zinc oxide varistors.26 Furthermore the number of these clear spots on a given constant area increases with the Ho3+-doping level.The EDS analysis of the surroundings of the holes at the grain boundaries shows also a higher Ho concentration in the area centred around such holes. Fig. 8 reports a micro- graph of an Ho3+-doped ZnO sample and its EDS analysis. The signals arising from Ho atoms are present only on the spectrum recorded over the area H centred around the hole near the centre of the micrograph. This last result can be interpreted as the fact that the rare earth oxide acts as an inhibitor of grain growth.Conclusion The large bandgap of semiconducting zinc oxide enables the observation of the electroluminescence of the inserted hol- mium ions, as in the case of zinc sulfide. The yield of this electroluminescence reaches a maximum value when the con- centration of the dopant is ca. 0.5 atom%. The totally different behaviour of the electrode when changing the polarization mode from anodic to cathodic indicates that the energy levels of the rare earth are not coupled with those of the host. Such lack of coupling is corroborated by the SEM results which show that the incorporated rare-earth does not substitute zinc atoms, but creates islands and dislocations in the ZnO crystal lattice.This results in a progressive decrease of the semicond- ucting properties of the host. The rare-earth-doped zinc oxide has to be taken as a composite material. The electrochemical Fig. 8 (a) SEM and (b) EDS spectra of a 0.4 atom% doped ZnO sample. Solid line: spectrum recorded over the area B, centred on a ZnO monocrystal. Dotted line: spectrum recorded over the area H, centred around the hole at the grain boundaries near the centre of the micrograph. technique used to induce the electroluminescence of the rare earth needs rather low polarization voltages compared with those usually needed for all solid devices. The presence in the electrolyte solution of an electron source external 10 the semiconductor, and of a judiciously chosen redox potzntial, suppresses considerably the electrode decomposition.from Rhone-Poulenc recherches (France) for SEM and EDS measurements, and for fruitful discussions. PA -. qhinnnva and H Knhavachi <i>rinoer1 Fl~rtml~~min~cr~nr~ -..---. J..Y...,-""I.."."r..rr.r""""..cr) .,-. ',""'b.,' Verlag, Berlin, Heidelberg, 1989. 2 J. C. Ronfard-Haret, D. Kouyate and J. Kossanyi, SoliJ State Comm., 1991,79,85. 3 D. Kouyate, J. C. Ronfard-Haret and J. Kossanyi, J. Electroanal. Chem., 1991,319, 145. 4 D. Kouyate, J. C. Ronfard-Haret and J. Kossanyi, J.Luininesc., 1991,50,205. 5 D. Kouyate, J. C. Ronfard-Haret and J. Kossanyi, J. Muter. Chem., 1992,2,727. 6 J. Kossanyi, D. Kouyate, J. Pouliquen, J-C. Ronfard -Haret, P. Valat, D. Oelkrug, U.Mammel, G. P. Kelly and F. Willcinson, J. Luminesc., 1990, 46, 17. 7 D. Kouyate, J.-C. Ronfard-Haret, P. Valat, J. Kossanyi, U. Mammel and D. Oelkrug, J. Luminesc., 1990, 46,329. 8 D. Fichou and J. Kossanyi, J. Electrochem. SOC.,1986, 133, 1607. 9 P. K. Patil, J. K. Nandgave and R. D. Lawangar-Pawar, Solid State Commun., 1990,76, 571. 10 B. Pettinger, H. R. Schoeppel and H. Gerischer, Ber.Bur:senges. Phys. Chem., 1976,80,849. 144 J. MATER. CHEM., 1994, VOL. 4 11 T. Yamase and H. Gerischer, Ber. Bunsenges. Phys. Chem., 1983, 20 S. Bhushan, B. R. Kaza and A. N. Pandey, Phys. Status. Solidi. A, 87, 349. 1978,49, K167. 12 H. Kiess, J. Phys. Chem. Solids, 1970,31, 2379. 21 S. Bhushan, A. N. Pandey and B. R. Kaza, J. Luminesc., 1979, 13 E. W. Chase, R. T. Hepplewhite, D. C. Krupka and D. Kanhg, 20,29. 14 J. Appl. Phys., 1969,40,2512. R. Mach and G. 0.Mueller, Phys. Status Solidi A, 1982,69, 11. 22 23 R. Mach and G. 0.Mueller, Phys. Status Solidi A, 1984,81,609. J. W. Allen, J. Luminesc., 1984,31132, 665. 15 16 17 18 G. Boulon, Rev. Phys. Appl., 1986,21,689. G. F. Alfrey and J. B. Taylor, Proc. Phys. SOC. London, 1955, 68B, 775. B. Pettinger, H. R. Schoeppel, T. Yokoyama and H. Gerischer, Ber. Bunsenges. Phys. Chem., 1974,78,1024. Handbook of Chemistry and Physics, ed. R. C. Weast, CRC Press, 24 25 26 A. Krier and F. J. Bryant, Phys. Status Solidi 4,1984,83, 315. D. Kouyate, J. C. Ronfard-Haret and J. Kossanyi, J. Luminesc., 1993,55209. P. Williams, 0. L. Krivanek, G. Thomas and M. Yodogawa, J. Appl. Phys., 1980,51 3930. 19 Boca Raton, 1985, 65th edn., Table D155-62. H. Gerischer, Faraday Discuss., 1980,70, 137. Paper 3/03322H; Receired 9th June, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400139

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. MATER. CHEM., 1994, 4( l), 145-151 New Fluoroionophores for Alkali-metal Cations based on Tetrameric Calixarenes Consuelo Perez-Jimenez, Stephen J. Harris and Dermot Diamond* School of Chemical Sciences, Dublin City University, Dublin 9, Ireland The synthesis and fluorescence behaviour of two novel fluorescent calix[4]arenes containing four anthracene moieties on the lower rim are described. The calixarenes exhibit an optically selective fluorescent response to complexation with group 1 metal cations. Interesting optical responses to Li', Na+ and K+ have been found with the tetraester derivative and high affinity for Na' ions with the tetraamide derivative. They may prove to be useful components for the fabrication of optical sensors for the determination of these ions.'H NMR studies enable the molecular basis for differences in the optical behaviour of the two ligands to be understood at least partially, and perhaps explain why the tetraamide calixarene has better potential as an Na' sensor material. There is currently considerable interest in the development of optical sensors for chemical analyses especially for the selective determination of clinically important species such as lithium, sodium and potassium. The inherently sensitive nature of fluorescence signalling makes it an attractive option, particu- larly when combined with a selective complexing process involving the target species.' Several groups have recently reported fluorescent alkali- metal sensors derived from macrocyclic compounds such as cyclic and non-cyclic pol yet her^,^,^ crown ethers4 and cryp- tand~.',~ Calixarenes, another major class of macrocyclic receptor molecule^,^'^ have recently evoked some interest in the area of fluorogenic sensors.Calixarene derivatives incorpo- rating ionophoric functional groups linked to the phenolic oxygen atoms exhibit excellent properties as neutral receptors for metal ions. In particular, they form stable complexes and show sharp size-related selectivity because of their well preor- ganized ionophoric group^.^-" More recently, Sat0 and Shinkai12 found very interesting luminiscence properties of modified calix [4] arenes. This led us to prepare fluorescent calixarenes13 by incorporation of aromatic fluorophores at sites adjacent to the polar ionophoric cavity.It was anticipated that the inclusion of a positively charged species in the cavity would result in sufficient perturbation of the fluorophoric groups to enable complexation to be monitored using fluor- escence, as demonstrated recently with pyrene based system^.'^ We now report the synthesis and spectroscopic study of new calix [41 arene derivatives (I and 11), containing four anthracene moieties on the lower rim, designed to combine the specific complexing ability of the tetramer and amide" series for alkali-metal ions with the photophysical behaviour of the anthracene ring.20 Experimental Instrumentation The structures of the final products and intermediates were elucidated by a variety of spectral methods.IR spectroscopy was carried out using a Perkin-Elmer 9836 infrared spectro- photometer and proton ('H) NMR spectroscopy was carried out using a Bruker AC 400 nuclear magnetic resonance spectrometer. Fluorescence spectra were recorded on a Perkin- Elmer LS-5 luminescence spectrometer equipped with a red- sensitive Hamamatsu R928 detector and a thermostatted cell compartment. Materials The calix[4]arenes I and I1 were used as reagents for Auor- escence measurements. Extra-pure-grade lithium, sodium and potassium perchlorates and thiocyanates were obtained from Aldrich. The solvents chloroform (Riedel De Haen), tetra- hydrofuran (THF)and methanol (Fluka) were used without further purification. Tetraester I and tetraamide I1 were prepared (see Scheme 1) from the known p-tert-butylcalix [4] arenetetraacetic acid.21 The tetraacid 3 was treated with thionyl chloride to afford the tetraacid chloride 4 which was treated with 9-anthra- cene methanol in THF containing pyridine to furnish the tetraester I (m.p.105.5-109.0"C) and with 1 4 I * kH2-C-OEt II I 0 2 ii OCH2-C-OH II0 3 CH2RH 0 IR=O 11 R = NCH3 Scheme 1 Reagents (i) BrCH2C02CH2CH, K2C03 in acetone; (ii) KOH, H20,EtOH, HC1; (iii) SOC1,; (iv) pyridine, THF 9-(methylaminomethy1)anthraceneto furnish the tetraamide I1 (m.p. 158.0-161.0 "C), respectively. The analyses of the structures of the products and intermediates by spectroscopic methods were found to be consistent with the predicted structures.The analytical data for the calix [41 arene deriva- tives follow. Tetraester Z Yield 65%; Found: C, 80.00; H, 6.42. Calc. for C,,,H1,,0,2: C, 81.92; H, 6.38; IR spectrum (KBr) vmaX/cm-': 1755 (C=O). 'H NMR (400 MHz; solvent CDCl,; standard TMS) dH 1.05 (36 H, s, CMe,), 2.85 (4 H HB, d, ArCH,Ar), 4.65 (4 H HA, d, ArCH2Ar), 4.70 (8H, s, OCH2C02), 6.10 (8 H, s, CH,anthr), 6.60 (8 H, s, ArH), 7.31-8.30 (36 H, m, anthracene). Tetraamide ZZ Yield 60%; Found: C, 79.31; H, 6.86; N, 3.00. Calc. for Cl16Hl,6N40,: C, 78.98; H, 6.69; N, 3.15. IR spectrum (KBr) vmaX/cm-': 1658 (C=O). NMR (400 MHz; solvent CDCl,; standard TMS) 6, 1.12 (36 H,s, CMe,), 2.60 (12 H, s, NCH,), 3.20 (4 H HB, d, ArCH2Ar), 5.05(8 H, s, OCH,CON), 5.18 (4 H HA, d, ArCH,Ar), 5.52 (8 H, s, CH,anthr), 6.83 (8 H, s, ArH), 7.20-8.40 (34 H, m, anthracene).Effect of Complexation on Fluorescence Emission Solutions of ligand I (5 x mol drn-,) and ligand I1 (5 x lop5mol drnp3) were made up in CHCI, and THF, respectively. 5 cm3 aliquots of these solutions were taken and incremental concentrations of alkali metals were added. During the work with ligand 11, access to the fluorescence instrument was restricted and samples had occasionally to be stored for several days. Routine stability checks showed that with CHC1,-based experiments, the emission spectra changed significantly suggesting that the complex had decomposed. However, this problem did not occur with THF-based experi- ments, and it was noted that there was no discernible change in the emission spectrum after storage of the complex for periods of up to one week.Furthermore, it was established that the trends in emission behaviour observed on com-plexation with both ligands were independent of whether THF or CHC1, was used as the ligand solvent, whether the source of the metal ions was SCN- or ClO,, or whether the salt was dissolved in methanol or water. Hence although different solvents were used for fluorescence experiments with ligands I and 11, valid comparisons can be made for the two sets of results. Methanolic solutions of lithium, sodium and potassium thiocyanates were used with ligand I, whereas aqueous solu- tions of metal perchlorates were added in the case of ligand 11.As mentioned above, comparative experiments with both ligands showed the same general trends whether aqueous or methanolic solutions of the metal salts were used, and as very small volumes of these salts were used as sources of the metal ions in the titration experiments, this is not unexpected. NMR experiments in all cases employed CDC1, as the solvent. The fluorescence-intensity changes of ligand solutions upon stepwise addition of alkali-metal salts were recorded from 300 to 700 nm. Quartz glass cells (1 x 1 cm2) were used for each measurement. Other measurement conditions are given in the figure legends. Effect of Complexation on 'H NMR Spectra The complexing ability of the new fluorescent calix [41 arenes synthesized was determined by 'H NMR titration experiments.The 'H NMR spectra were obtained from a 5.1 mmol dm-, solution of the ligands in CDC1, (Aldrich). The salts used for J. MATER. CHEM., 1994, VOL. 4 these experiments were thiocyanates in CD,OD (Aldrich). 'H NMR titrations were performed by adding incremental amounts of MSCN in CD,OD (M=Li+, Na', K') directly to a CDCl, solution of I or I1 in an NMR tube. Results and Discussion The fluorescence spectra of both calix [4] arene derivatives exhibit a monomer emission with a fluorescence maximum around 418.0 nm (excitation 388 nm) in the case of tetraester I, and at around 415.5 nm (excitation 263 nm) in the case of tetraamide 11. However, substantial differences appeared when we examined the effect of the optical response of the ligands to addition of alkali-metal ions.When Li' or Na+ thiocyanate (in MeOH) is added to the CHC1, solution containing I (5 x lop6 mol dm-3), the fluor- escence intensity of the entire spectrum decreases markedly with increasing salt concentration (Fig. 1A). The addition of KSCN has a different effect. As shown in Fig. lB, the maxi- mum emission (ca. 418 nm) decreases with increasing KSCN concentration and the emission at 443 nm increases, with an isoemissive point at 432 nm. In contrast, the addition of alkali-metal perchlorate (in H20) to the THF solution containing I1 (5 x lop5mol dmP3) produced, in all instances, an increase of the fluorescence intensity compared to that of the free ligand, although the addition of lithium or potassium had only a slight effect compared with sodium.As Fig. 2 shows, the fluorescence intensity increased markedly in the concentration range 10-6-10-4 mol dmP3 of NaClO,, whereas relatively minor changes were in fact observed on the addition of lithium or potassium ions. Some indication of an inner filter effect is suggested by the reduction in the fluorescence emission at the lower edge of the spectrum (380-400 nm) compared to Fig. 1, where an emission peak is obtained, but the remainder of the spectrum appears unaffected. This arises from the very large extinction coefficient for anthracene at the excitation wave- length 263 nm (log E =4.2).22 Consequently, given the concen- tration of ligand employed, the absorbance of the solution will be very high, and a large proportion of the ligand will be in the ground state at any instant in time.However, as the degree of overlap between the emission and absorption spectra is very slight, and coincides only with the lower edge of the emission spectrum (380-400nm) we are confident that the emission spectra above 400 nm are essentially unaffected. In fact, it is very probable that similar results to those described in this research could be obtained with much lower ligand concentrations. The inset in Fig. 2 shows a plot of I/Z, (I= emission intensity of the complex emission at 415.5 nm, I,= emission intensity of the free ligand at the same wavelength) us. log[Na+].It is clear that there is a sensitive response on complexation, centred at log "a+] =4.3-4.4. As the concen- tration of I1 is 5.0 x mol dm-3, this is convincing evidence in support of the formation of a 1 : 1 complex with Na' ions. The fluorescence study of the two calix [4] arenes reveals significant differences between the anthracenemethyl ester I and the anthracenemethyl amide 11, ie. the substitution of the group -COO-by -CONCH,- between the polar cavity of the calixarene and the fluorescent anthracene groups pro- vokes a drastic change of the optical response to alkali-metal cations. With the tetraester I, the difference in behaviour obtained with K+ compared to that observed with Li' and Na' is not easy to interpret. To obtain some insights into the confor- mational changes occurring in the polar cavity on com-plexation with the alkali-metal cations, we carried out 'H NMR titration experiments. In the absence of NaSCN, compound I possesses the cone conformation as evidenced by J.MATER. CHEM., 1994, VOL. 4 147 280] A c 200-c Y 2nm 2nm Fig. 1 Fluorescence spectra of I (5.00x mol dmP3) in chloroform at different concentrations of NaSCN (A) and KSCN (B). [SCN-] = (u) 0, (b)1.0x lop6,(c) 4.0 x (d) 6.0 x (e) 8.0 x loP6mol dm-3. The spectra were measured with excitation at 388 nm. Unm Fig. 2 Fluorescence spectra of I1 (5.00 x mol dmP3) in THF at different concentrations of NaC10, (a)0; (b)0.1 x lo-'; (c) 0.5 x (d) 1.0~loP5; (e) 2.0~ (f) 4.0 x lop5; (g) 6.0~lo-'; (h) 10 x lo-' mol drn-,.The spectra were measured with excitation at 263 nm. Inset shows plot of I/Z, us. [Na'], where Z=fluorescence intensity of ligand 11-Na' complex and I, =fluorescence intensity of free ligand I1 measured at 415.5nm. the splitting pattern of the ArCH,Ar methylene protons [Fig. 3(a)]. When NaSCN is added directly to the CDC1, solutions of I, the signals in the 'H NMR spectrum change greatly. With a salt :I ratio of less than 1 [Fig. 3(b)],signals for both complexed and uncomplexed ligand were present in the spectrum, indicating that on the NMR timescale, the exchange rate between the two species was slower at room temperature. Upon reaching a 1:1 stoichiometry [Fig. 3(c)] all the signals for the free ligand disappeared and an increase 8.0 6.0 4.0 210 6 Fig.3 Partial 'H NMR spectra of the fluorescent calix[4]arene I in CDCl, at 25 "C: (a)R =[NaSCN]/[I] =O; (b)R =0.5; (c) R = 1, where [I] =5.1 mmol drn-,. Aliquots from CD30D solution of 1 mcil dm-3 NaSCN were added to a CDC1, solution of I in a NMR tube. 0, 0=ArH; V,V =ArCH,Ar; El,-. =CH,anthr. in the sa1t:I ratio beyond unity produced no further spectral shifts. This finding indicates a 1: 1 stoichiometry for the NaSCN complex with I. In contrast, titration of ligand I with potassium thiocyanate (Fig.4) produced spectral changes up to the point of 1:l stoichiometry and did not show separate signals for complexed and uncomplexed ligand. As shown in Fig. 4(c), upon reaching K+ :I molar ratio of 1, the signals of the free ligand disappear and those corresponding to the K+ complex start to appear and continue to increase in intensity until a K+ :I molar ratio of 2 is reached [Fig.4(d)], whereas at the same stage with Na', all signals for the free ligand have already disappeared and an increase in the Na' :I ratio beyond unity produced no further spectral shift. Similar effects have been reported previously" and interpreted as arising from differences in the rate of formation/disassociation of the metal-ion-ligand com-plexes. In the case of potassium, the process of complex formation involves some rearrangement of the pendant polar groups in order to accommodate the larger-than-ideal size of the ion. The ILK+ complex would thus be less stable than the corresponding I-Na' complex and the net residence time of the potassium ion within the cavity shorter.Hence the I-K+ complex is not seen in the NMR spectrum until the potassium concentration is sufficiently high to enable the complex to exist for long enough to be picked up on the NMR timescale. Additionally, a comparison between the ID 1 I.,,,'I 2.0 I+6.0 4.0 6 Fig. 4 Partial 'H NMR spectra of the fluorescent calixC4larene I in CDCl, at 25 "C: (a) R= [KSCN]/[I] =O; (b) R=0.5; (c) R = 1; (d)R =2, where [I] =5.1 mmol drn-,. Symbols as Fig. 3. J. MATER. CHF.M., 1994, VOL. 4 NMR spectra of the I-metal salts reveals that the proton chemical shifts relative to the free ligand are much larger on Na' complexation than K+ (downfield shift aromatic proton: 0.47 with Na', 0.34 with K+; H, of AB methylene ArCH,Ar quartet: 0.47 with Na', 0.28 with K+ and upfield shift HA of AB methylene ArCH2Ar quartet: 0.39 with Na' and 0.25 with K').The 'H NMR data suggest that the formation of an I-Na' complex of 1:1 stoichiometry occurs. In contrast, with K', a 1:2 molar excess of ligand :metal ion is required before 100% complexation occurs. This suggests that the complexation process needs to be driven to completion by excess K+ ions. The observed quenching of fluorescence arising from com- plexation of ligand I with Lif and Na' may be the result of the I : Li+ and I :Na+ complexes being less sterically hindered than the free ligand.23 Similar perturbation differences on fluorescence have been noted by alkali-metal complexation of 2,3-naphtho-20-crown-6 (quenching) and its close relative 1,8-naphtho-21-crown-6 (enhancement).In this case, the differing fluorescence behaviour was accounted for by differ- ences in the geometrical orientation on co~nplexation.~~ As mentioned previously, the same general pattern of quenching is obtained when perchlorate ions are substituted for thiocyan- ate, or THF is used as the reaction medium in place of chloroform, or the metal salt added in aqueous instead of methanolic solution. Hence the observed quenching is not due to the anion used or the medium in which the reaction is performed, but is a function of the metal ion-ligand interaction.Enhancement of the longer wavelength emission spectrum of ligand complexed with K + suggests that complexation with potassium has a significantly different effect on the molecular conformation of the ligand than either lithium or sodium, which may simultaneously encourage certain modes of vibrational relaxation to occur (lower energy) at the expense of other higher energy modes. In the case of ligand 11, the high affinity shown for sodium ions is probably a size-related phenomenon. Thus, the cavity defined by the four amide carbonyl groups and four phenoxy oxygens best matches the size of Na+ cations. However, the molecule may allow the inclusion of the smaller Li' cation or the larger K+ cation by a flexing movement of the pendant ligating groups and/or a change in the tilt angle of the aromatic ring, but the resulting contraction or expansion of the cavity will be energetically expensive and will lead, in relative terms, to a destabilization of the complex.Support for this interpretation was obtained by the 'H NMR study of compound I1 both in the absence and presence of MSCN (M =Li', Na+, K'). The 'H NMR data established unequivocally from the characteristic Ar-CH2-Ar (AB quartet) resonances that the compound I1 and the complexes with the three cations possess the cone conformation (Fig. 5-7). The addition of incremental amounts of NaSCN in CD,OD to 5.1 mmol dm-3 solution of I1 in CDCl, affected resonances arising from the ArCH,Ar protons, the aromatic protons (0.40 downfield) and the anthracene protons (resonance at 7.20 is displaced to 7.42, the signal at 8.20 disappears and the intensity of the signal at 8.4 increases).Thus signals for both complexed and uncomplexed ligand were present in the spectrum with a sa1t:II ratio of less than 1 [Fig. 5(b)].Upon reaching a 1: 1 stoichiometry [Fig. 5(c)] the corresponding signals for the free ligand disappeared and an increase in the sa1t:II ratio beyond unity produced no further spectral changes. In contrast, titration of compound I1 with lithium or potassium thiocyanate produced new additional spectral changes apart from those observed with sodium. Thus, as Fig. 6 and 7 show the CH,anthr proton signal, which was not affected by Na+ complexation, experi- J.MATER. CHEM., 1994, VOL. 4 I, 8.0 6.0 4.0 6 Fig. 5 Partial 'H NMR spectra of the fluorescent calix [4] arene I1 in CDCl, at 25 'C: (a)R = [NaSCN]/[ 111=0 (b) R =0.5; (c) R = 1;where [II] =5.1 mmol drn-,. Symbols as Fig. 3. enced a downfield shift of 0.13 and 0.08 for K+ and Li', respectively. Similarly, the NCH, protons were shifted down- field by 0.07 and 0.13 by K+ and Li+, respectively but showed no change with Na'. In addition, in the case of Li' com-plexation alone, the OCH,CON protons were also displaced (0.25 downfield). This supports the view that the inclusion of K' and especially Li' cations inside the cavity is followed by a more extensive conformational reorganisation of the cavity region than that observed with Na' which hence has the optimum size for best fit with the polar cavity.In further support of this view is the fact that during incremental addition of both Lif and K+ salts, the presence of a 50% excess of metal salt was required for 100% formation of the new AB quartet of the calixarene-metal-ion complex (together with the disappearance of the AB quartet of the uncomplexed ligand 11). However, as a ligand I1 :ISf or Lif complex of 1:1.5 stoichiometry is highly unlikely, this prob- ably arises from an equilibrium phenomenon. Thus, excess metal ion is needed to drive complexation to completion. The same ratios are obtained from plots of I/I, us. metal-ion concentration [Fig. 2 (inset)]. From these plots, dissociation constants corresponding to the equilibrium ML+eM++L with M =Li+, Na+ and K+ and L =ligand 11, were determined + I I0 8.0 7.0 6.0 5.0 4.0 3.0 2.0 6 Fig.6 Partial 'H NMR spectra of I1 in CDCI, at 25 "C: (u) R= [KSCN]/[II]=0.5;(b) R= l;(c) R=1.5;where [II]=5.1 mmoi dm-,. V, V, 0, 0, 0,W, as Fig. 3; 0,+=NCH,. to be 2.2~ lop5 and 1.5~lop4, 4.4~ 10-4mol dmp3, respectively. The preceding NMR data suggest that, in the case of ligand 11, the observed enhancement in fluorescence upon com-plexation with all three metal ions arises from the increased rigidity of the calixarene-metal-ion complex over that of the free ligand.23 With less freedom to interact with the surround- ing solvent molecules or for intramolecular movement, there is a corresponding increasing probability of fluorescence occurring.For the three cations studied, the largest enhance- ment was obtained with Na+, in agreement again with the NMR data which suggest that Na+ complexation confers more order on the molecule (as evidenced by the much simpler 'H NMR spectrum obtained from the complex compared with the free ligand) and least perturbation to the cavity region. In a similar way, the much smaller enhancement observed on lithium complexation may be considered a conse- quence of the higher perturbation that the inclusion of Li' causes in the polar cavity. Our results reveal that the tetraamide I1 possesses improved fluorescence behaviour over that of the tetraester I. To verify the utility of I1 as a fluorescent Na+ sensor, we have examined the effect of this ion on the fluorescence intensity in the presence of other alkali-metal ions.As shown in Fig. 8, the addition of concentrations from to 10-4mol dmF3 NaClO, causes an increase in the emission intensity which is proportional to the salt concentration. Similar behaviour to that illustrated in Fig. 8 has been also found in the presence of KClO,. To obtain further insight into the ligand I1 selec-tivity we recorded the changes in fluorescence intensity caused by the addition of incremental amounts of NaC10, to the THF solution of 11, in the presence of 10-3mol dr~i-~ of MClO, (M =Li+, K+). As Fig. 9 shows, the addition of low Na' concentration, in relation to Li' or Kf Concentration, J.MATER. CHEM., 1994, VOL. 4 --_/---Nat 820 6.7 x mot dm-38601 + 8.0 7.0 6.0 5.0 4.0 3.06 Fig. 7 Partial 'H NMR spectra of I1 in CDCl, at 25 "C: (a) R= [LiSCN]/[ I13 =0.5; (b) R =1; (c) R = 1.5; where [111=5.1 mmol dm-,. V, V, 0, 0, 0,H, as Fig.3; 0, +=NCH,; A, A= OCH,CON. Nnm Fig.8 Na' titrations of the fluorescence emission spectra of I1 (5.00x mol drnp3) in THF in the presence of lo-, mol dm-, of LiClO,. [NaClO,] =(a) 0, (b)1, (c) 4, (d) 6, (e) 10 x lop5mol dm-,. Na+ [ 5.3x 10" mol dm3li 660-1 I IIIIIIIIIII20 60 100 140 180 220 tlS Fig. 9 Effect of Na' on the fluorescence intensity of I1 (5.00x lop5mol drn-,) in THF, in the presence of lop3mol dm-3 of MC10, (M=Li+, K'). Fluorescence intensity was monitored at 415nm.Aliquots from 10-2mol dm-, of NaClO, in water were added directly to the solution of I1 and MCIO, in a cuvette. provides a marked increase on the fluorescence intensity. It should be noted that small changes in the Na+ concentration are translated into a large increase in the fluorescence intensity. This sensitivity suggests that it should be possible to make precise and accurate determinations of Na at concentrations + lower than the limit of detection of potentiometric sensor^.^^,^' These results indicate that there is little observable inter- ference, with the fluorescence intensity of I1 with sodium ions in the range 1x 10-6-1 x mol dm-3, from much higher concentrations of Li+ and K'. In addition, Fig. 9 demon- strates clearly the fast dynamics of the emission response to complexation with sodium ions, with the final emission inten- sity being reached within a few seconds after addition of the metal ions.The limiting factor for this is the rate of dispersion of the sodium ions after injection rather than the kinetics of the complexation reaction. Conclusions Two ligands which exhibit a selective fluorescent response on complexation with sodium ions have been synthesized, with the amide derivative (ligand 11) exhibiting better selectivity than the ester (ligand I). 'H NMR experiments suggest that a 1: 1 complex is formed between I1 and Na+ ions but with K+ ions, a 50% molar excess of the metal ion is required for complete reaction of the ligand.The authors gratefully acknowledge financial help for C.P.J. from the Cornissib Interpartamental de Recerca e Innovacib Tecnolbgica (CIRIT), Generalitat de Catalunya (grant BE92-244), and for S.J.H. from the Irish Science and Technology Agency (EOLAS) (grant no. SC,192/319). References 1 R. A. Bisell, A. P. de Silva, H. Q. N. Gunaratne, P. L. M. Lynch, G. E. M. Maguire and K. R. A. S. Sandanayate, Chem. SOC.Rev., 1992,187. J. MATER. CHEM., 1994, VOL. 4 K. Haratini, J. Chem. SOC., Chem. Commun., 1987, 960; K. Haratini, Analyst, 1988,113, 1065. A. P. de Silva and T. Koyama, J. Chem. SOC., Chem. Commun., 1989,1183. S. Fery-Forgues, M.-T. le Bris, J.-P. Guette and B. Valeur, J. Chem. SOC., Chem. Commun., 1988,384. F. Fages, J.-P.Desvergne, H. Bouas-Laurent, P. Marsau, J.-M. Lehn, F. Kotzyba-Hibert, A.-M. Albrecht-Gary and M. Al-Joubbeh, J. Am. Chem. SOC., 1989,111,8672. 6 F. Fages, J.-P. Desvergne, H. Bouas-Laurent, J.-M. Lehn, J. P. Konopelski, P. Marsau and Y. Barrans, J. Chem. SOC., Chem. Commun., 1990,655. 7 C. D. Gutsche, Calixarenes, Royal Society of Chemistry, Cambridge, 1989. 8 J. Vicens and V. Bohmer, Calixarenes: A Versatile Class of Macrocyclic Compounds, Kluwer Academic Publishers, Dordrecht, 1991. 9 M. A. Mckervey, E. M. Seward, G. Ferguson, B. Ruhl and S. J. Harris, J. Chem. SOC., Chem. Commun., 1985, 388. 10 F. Arnaud-Neu, E. M. Collins, M. Deasy, G. Ferguson, S. J. Harris, B. Kaitner, A. J. Lough, M. A. Mckervey, M. J. Schwing-Weill and E.M. Seward, J. Am. Chem. SOC., 1989, 111,8681. 11 J.-D. van Loon, L. C. Groenen, S. S. Wijmenga, W. Verboom and D. N. Reinhoudt, J. Am. Chem. Soc., 1991,113,2378. 12 N. Sato and S. Shinkai, J. Chem. SOC.,Perkin Trans. 2, 1993,621. 13 C. Perez-Jimenez, S. J. Harris and D. Diamond, J. Chem. SOC., Chem. Commun., 1993,480. 14 I. Aoki, H. Kawabata, K. Nakashima and S. Shinkai, J. Chem. SOC., Chem. Commun., 1992,1771. 15 T. Jin, K. Ichikawa and T. Koyama, J. Chem. SOC.. Chem. Commun., 1992,499. 16 I. Aoki, T. Sakaki and S. Shinkai, J. Chem. Soc., Chem. Commun., 1992,730. 17 A. Cadogan, D. Diamond, M. R. Smyth, M. Deasy, M. A. McKervey and S. J. Harris, Analyst, 1989, 114, 1551. 18 R. J. Forster, A. Cadogan, M. Telting-Diaz, S. J. Harris and M. A. McKervey, Sens. Actuators B, 1991,4,325. 19 A. Arduini, E. Ghidini, A. Pochini, R. Ungaro, G. D. Andreetti, G. Calestani and F. Ugozzoli, J. Incl. Phenom. Mol. Recognit. Chem., 1988,6,119. 20 H. Bouas-Laurent, A. Castellan and J.-P. Desvergne, Pure Appl. Chem., 1980,52,2633. 21 R. Ungaro, A. Pochini and C. D. Andreetti, J. Incl. Phenom. Mol. Recognit. Chem., 1989,2, 199. 22 W. Kemp, Organic Spectroscopy, Macmillan, London, 1975, p. 168. 23 H. H. Willard, L. L. Merritt Jr., J. A. Dean and F. A. Settle Jr., Instrumental Methods of Analysis, Wadsworth, Belmont, CA, 1988,199. 24 L. R. Sousa and J. M. Larson, J. Am. Chem. SOC.,1977,99,307. Paper 3/02936K; Received 24th May, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400145

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. MATER. CHEM., 1994,4(1), 153-159 153 Nickel-Iron Hydroxide Carbonate Precursors in the Synthesis of High-dispersity Oxides Elly Uzunova,* Dimitar Klissurski* and Stoyan Kassabov Institute of General and lnorganic Chemistry, Bulgarian Academy of Sciences, Acad. G. Bontchev Str. bl.1 1, Sofia 1113, Bulgaria Nickel-iron hydroxide carbonates with layered reevesite-li ke structures have been synthesized by coprecipitation and their evolution to high-dispersity oxides has been studied. Ni"-Fe"l cation substitution in the octahedral positions of the cationic layers for a wide interval of nickel :iron ratios (0.33-3) is deduced from the XRD and DTA data and the Mossbauer spectra. No divalent iron is present in the dried samples and no admixtures are detected in the diffraction patterns and the IR spectra.Thermal decomposition of the precursor hydroxide carbon- ates occurs at temperatures <580 K and high-dispersity oxides are obtained. The average particle size of nickel ferrite, synthesized at 623 K is 7.5 nm and the specific surface area is 160 m2 g-'. A mixture of oxides is obtained from the precursors with iron :nickel ratios different from that of ferrite, in agreement with the phase diagram of that system. At 873 K a partly inverse nickel ferrite with the cation distribution Nio,l~s~Feo,~~s~[Fe,~lo]04is obtained. The surface composition of the oxide samples displays deviation from the bulk ratio in favour of nickel that increases with the iron : nickel ratio. For samples with Fe: Ni 20.5 the oxygen-to-metal surface ratio is higher than that expected on the basis of the formula of spinels and sesquioxides.Multicomponent oxide materials containing different trans- ition-metal cations are interesting for their electrical, magnetic and catalytic properties. Their application as electrode or magnetic materials, catalysts or catalyst precursors for many processes of industrial importance such as aldol condensation, epoxide polymerization, alkene metathesis and selective oxi- dation presupposes that they are obtained with high dispersity and without imp~rities.'-'~ Among the wide variety of precur- sors for the synthesis of binary and multicomponent mixed oxides, layered-type hydroxides with different intercalated cations proved to be reliable with respect to these Hydroxide carbonates deserve special attention, since they decompose at low temperatures, thus ensuring the formation of small-particle polycrystalline oxide phases with basic properties in general and good interdispersion of the active components.'0~'3-'5 A la rge number of hydrotalcite-like com- pounds have been studied as precursors in the synthesis of catalysts for CO hydr~genation.~,~,~After controlled thermal decomposition and reduction the catalysts consist of an oxide phase containing well dispersed metallic particles.Variations in the number and ratio of the elements of the hydrotalcite and pyroaurite structures is all~wed.'~.'~ The lattice is built from positively charged brucite-like cationic layers [M", -xM111x(OH)2]X+, intercalated with disordered anionic layers of the type [CO3'4H20l2 -.Different transition-metal cations may be introduced in the alumina matrix by using similar precursors and high stability of the oxide phase in catalytic reaction conditions is attained.7-'3 Generally, ratios M"':M" in the range 0.33-0.50 are ~btained:.~.'~ but for Ni-A1 hydrotalcite-like compounds a metal cation ratio of 1:1 can be reached.'* Introduction of cobalt in hydrotalcites is limited to Co :A1 =0.5,19 while in pyroaurite-like hydrox- idecarbonates, the cobalt :iron ratios range from 0.33 to 2.00.14 A number of MI1-Fe"' hydroxide carbonates with layered structure were synthesized by coprecipitation using Fe"' salts as the initial reagent.'6v20*21 N o substantial change of the unit-cell dimensions with composition was detected for magnesium-iron hydroxide carbonates prepared in this way.Fe203 was found to be present in the samples when higher crystallisation temperatures were used.21 Nickel-iron hydroxide carbonate exists as the mineral reevesite. Nickel(I1) hydroxide carbonate is an amorphous substance, as well as the iron(1rr) hydroxide carbonate. The OH :C03 ratio in these compounds depends strongly on the preparation conditions. In the present study, nickel-iron hydroxide carbonates have been synthesized by a coprecipi- tation method and the limits of existence of the reevesite-like structure with variation of the nickel :iron ratio have been studied.Experimental Synthesis The hydroxide carbonates were synthesized as follows: a solution of Ni" and Fe" salts, in the desired molar ratio with a total M" concentration 0.6 mol dmP3 was added to a sodium carbonate solution with a two-fold excess volume at pH 9. The initially formed dark-green precipitates were stirred continuously for 1 h at 300-310 K. After they had been filtered and washed carefully with cool water, they were dried at room temperature, during which Fe" was oxidized to Fe"' and yellow-brown substances with the reevesite-like structure were formed. Characterization Atomic absorption analysis of nickel and cobalt in the hydrox- ide carbonates was performed; in the oxide products the Ni :Fe ratio was controlled by X-ray fluorescence analysis; the OH :C03 ratio was determined by absorption of the C02 and H20 evolved in Natron asbestos. X-Ray diffraction analysis was performed on a diffractometer with Cu-Ka monochromatized radiation.Differential thermal analysis was carried out at heating rate of 10 K min-' using cr-Al,O, as the reference. The IR spectra were recorded on a Perkin-Elmer 9836 spectrometer as KBr pellets. Mossbauer spectra were recorded on a spectrometer operating in constant-acceleration mode at room temperature. The source used as a 57Co/Pd matrix. The isomer shifts, presented in Tables 4 and 6 (later) are referred to a-Fe. In estimating the ferrite cation distribution a correction factor of the recoil-free fraction ratio for octahedral (Bi) and tetrahedral (A) sites fb/.f, =0.94 was used.XP spectra were recorded on an ESCALAB MkII spectrometer with Al-Ka radiation (1486.6 eV). Transmission electron microscopy was used for the determination of mean J. MATER. CHEM., 1994, VOL. 4 particle sizes. Specific surface areas were measured according to the modified BET single-point method, using N2 physisorp- tion at 77 K. Results and Discussion Characterization of the Precipitates The XRD patterns (Fig. 1) of samples with 0.33 dNi :Fe d3.00 display reflections that can be indexed in a hexagonal unit cell (Tables 1 and 2) and are similar to those of pyroaurite (ASTM file 25 521) and reevesite (ASTM file 20 786). Sample NFHCll is highly crystalline, while beyond the range 0.33-3.00 there is a tendency for samples to form amorphous compounds.The continuous decrease in unit-cell volume in agreement with Vegard's law is an indication of Ni"-Fe"' substitution in the octahedral positions of the cationic layers (Fig. 2). As the total positive charge of the cationic layer increases when the trivalent cation is present at higher concen- tration, this is compensated by a higher number of carbonate groups per formula unit. The distance between the two nearest cationic layers is c/3 and the decrease exceeds by about seven times the change in the a parameter when the Ni :Fe ratio is varied from 0.33 to 3 (Table 1). This effect may be explained, bearing in mind that a divalent cation in the lattice is replaced by a trivalent cation with a smaller ionic radius, thus leading to a decrease in unit-cell volume and influencing mainly the hN II 66 58 50 42 34 26 18 10 219ldegrees Fig.1 X-Ray diffraction patterns of (a) NFHC31, (b) NFHCl 1 and (c) NFHC12 Table 1 Chemical composition and unit-cell parameters of iron-nickel hydroxide carbonates sample a/nm c/nm chemical formulae 0187-rn5 183--3 9--a, .-Y 179-t 175 1 I 1 I 0.0 1.o 2.0 3.0 Fe:Ni molar ratio Fig. 2 Dependence of unit-cell volume on the composition of iron--nickel hydroxide carbonates Table 2 Interplane distances and Miller indices of NFHCl1 dexplnm 0.750 0.374 0.263 0.258 0.250 0.239 0.229 --0.1931 0.1873 0.1821 0.1716 0.1621 0.1534 0.1503 -0.1447 dca,clnm 0.749 0.374 0.264 0.258 0.250 0.240 0.229 0.217 0.205 0.1930 0.1872 0.1819 0.1715 0.1619 0.1533 0.1 502 0.1498 0.1449 index (hexagonal) 1/10 003 100 006 37 101 8 012 40 009 6 104 2 015 12 106 - 107 - 018 10 0012 3 109 2 1010 6 Olii 5 110 20 113 18 0015 1OE 4 - a parameter.The higher positive charge, related to Ni2 + -Fe3 + substitution is compensated by the interlayer negative charge, resulting in a stronger electrostatic interaction between adjac- ent layers, estimated by the change in the c parameter. Compared with other similar systems of hydrotalcite-like or pyroaurite-like structures, the highest degree of M"-M"' substitution is detected for nickel-iron, followed by ~obalt-iron'~ divalent-trivalent cation combinations.This may be explained bearing in mind that nickel@), like cobalt(Ir), has a preference towards regular octahedral coordination and the difference between the high-spin octahedral ionic radii is smallest for Ni" (0.690 A) and Fe"' (0.645 In the IR spectra (Fig. 3) of the coprecipitated hydroxide carbonates, bands characteristic of carbonate groups with D3,, symmetry are observed (Table 3).24-26 The distance between two adjacent layers, 46, does not allow coordination of the carbonate groups or water molecules from the anionic layers with the metal cations; nevertheless, a small effect of splitting of the asymmetric stretching vibration v3(E'), characteristic for free carbonate groups is observed in the spectrum along with a weak band at ca.1060 cm-l, due to activation of the symmetric valence vibration v,(A;). This can be explained by NFHC31 0.3071 2.2960 Nio,7,Feo,2,(OH)2(C03)o,125~0.53H20NFHC21 0.3070 2.2800 Ni,,67Fe,,,,(OH)2(CO,)o,~6,.0.54H,O assuming an interaction between carbonate groups and water .0.48H20NFHCl 1 0.3066 2.2471 Nio~,oFeo,50(OH)2(C03)o,25 molecules in the interlayers and/or existence of hydrogen.0.43H20NFHC12 0.3050 2.2296 ~io,33~eo~67(~~)2(~~3)o,335 bonding. The deformation vibration in the plane of the.0.42H20NFHC13 0.3033 2.21 70 Nio~2,Feo.74(0H)2(C03)o,37 carbonate group is in most cases obscured by the stronger J.MATER. CHEM., 1994, VOL. 4 6(MOH) band. The same number of bands is observed in hydrotalcite-like comp~unds,~,~~ cobalt-iron hydroxide car- bonate~'~and nickel hydroxide carbonate.26 As in the spectra of cobalt-iron pyroaurite-like hydroxide carbonates, the band corresponding to the metal-oxygen valence vibrations is observed below 400 cm-', indicating the lower energy of that bond, than in the iron(u1) or nickel@) hydroxide carbonates. For these two compounds the position of the v(M0) band is found at 454 and 410 cm-', respectively. Iron(1Ir) and nickel(I1) have similar masses and the absence of substantial change with composition at that frequency means that in the cationic layers of the lattice they have identical coordination and similar MI1-0 force constants.The Mossbauer spectra (Fig.4) of all samples consist of central doublets with isomer shifts (8)corresponding to high- spin Fe"' in octahedral c~ordination,~~ no divalent iron being detected. Since they are broad and asymmetric, they may be regarded as being constituted by a number of doublet compo- nents with different quadrupole splittings. At least three components with relative areas > 10% can be distinguished for all samples, with quadrupole splittings (A), ranging from 0.48 to 1.05 mm s-' (Table 4). The first neighbours of all cations are the same, since they are statistically distributed 372 I,, 1700 1300 900 500 wavenumbedcm-' Fig. 3 IR spectra of iron-nickel hydroxidecarbonates: (a) NFHC31, (b) NFHCl1, (c) NFHC12, (d) NFHC13 among the octahedral sites of the brucite-like cationic layers.The metal cation environment may be denoted as the second coordination sphere and the average number of Fe"' neigh-bours is different for each sample (Table 4).The main doublet component may be assigned to the predominant Fe"' cation environment, but as the distribution of metal cations is statistical, the full spectrum represents a superposition of the components of FelIr with deviations from the average second- order environment. The increase in the quadrupole splitting of the full spectrum (Fig. 5) and in the linewidth of both the constituent lines (Table 4)and the full line, depending on the iron content are in agreement with the assumption of a continuous Ni"-Fe"' substitution.Thermal Decomposition of the Hydroxide Carbonate Precursors The DTA profiles show three endothermic peaks (Fig. 6); the first corresponds to the release of weakly bound water mol- ecules from the interlayers and the second is related to the evolution of C02. In the first stage of thermal decomposition the mass loss is 6-9%, while in the second stage 7-25% mass losses are found for the different samples. The third endo- thermic peak is assigned to the destruction of the cationic layers and formation of oxides. This diminishes with increasing Fe:Ni ratio and shifts to lower temperatures, so that for samples NFHC12 and NFHC13 it is indistinguishable from the second peak. Nickel(rr) and iron(@ hydroxide carbonates decompose at 565 and 408 K, respectively.The shift of all endothermic peaks towards lower temperatures reflects the effect of Ni"-Fe"' substitution in the cationic layers. For iron(u1) hydroxide carbonate, calcite-like hydroxide- carbonates and some carbonate complexes, a splitting of the doubly degenerate asymmetric valence vibration v3(E) of the carbonate group, due to coordination with the metal cations, is observed and a new band at 1500-1650 cm-' appear^.^' The lowering of the carbonate group symmetry from DJhto C2" or C, at intermediate stages of thermal decomposition can be interpreted as being due to the partial destruction of the anionic layers and direct coordination with metal cations. This corresponds to a degree of decomposition when only water molecules from the interlayers are lost (Fig.7). The assignment of the other bands in the spectra can be made according to Table 3. The position of the v(M0) band differs negligibly with composition both in the initial and the partly decomposed samples. After heating the samples at 400 K the diffraction patterns of the reevesite-like phase almost disappear and a shift of the 003 line towards smaller interlayer distance (0.619 nm) is observed. Unlike most hydrotal~ites,~.~.'~the dehydration process is irreversible and the diffraction patterns of the initial sample do not reappear after contact of the partially decom- Table 3 Assignment of IR bands in the spectra of iron-nickel hydroxidecarbonates ;/cm -vibration NFHC3 1 NFHC 1 1 NFHC 12 NFHC 13 1644 m 1625 w 1632 w CQ.1630 b ca. 1580 w -ca. 1600 sh 1378 s 1380 s 1379 s{ ;E:h 1060 s ca. 1055 vw 1056 vw 1032 w 820 sh 800 m, b 808 m, b 820 m, b 712 m, b ca. 705 sh ca. 700 sh 690 m _-ca. 660 sh ca. 670 sh ca. 680 sh 524 m 524 m -ca. 515 sh 364 s, sp 368 s, sp 368 s, sp 372 s, sp w, weak; v, very; m, medium; s, strong; sp, sharp; b, broad; sh, shoulder; sym, symmetric: d, degenerate. J. MATER. CHEM., 1994, VOL. 4 100 99 9ai \VA 98 -96 97 -94 96-92 -2 -1 0 1 2 -2 -1 0 1 2 100 98 96-94 -I -2 -1 0 1 2 -2 -1 0 1 2 h8 Y c .-0 .-v) E c ?! c -2 -1 0 1 2 velocity/mm s-’ Fig. 4 Mossbauer spectra of iron-nickel hydroxidecarbonates: (a) NFHC3 I, (b) NFHC21, (c) NFHCl 1, (d) NFHCl2, (e) NFHC 13 posed sample with an alkali-metal carbonate solution. Even at 440 K broad and diffuse patterns of a spinel phase are observed. Bulk and Surface Characterizationof Oxides, obtained by Thermal Decomposition of Hydroxide Carbonate Precursors In the samples with Fe :Ni <2, after thermal decomposition at 623 K, nickel oxide is present, while from precursors with Fe:Ni>2, the excess of iron forms a haematite phase.High specific surface areas are measured for all oxide products (Table 5). The particle shape is close to spherical and the average diameter is in the range 6-1 1 nm. The phase composi-tion of the oxide samples is in agreement with the phase diagram of the Ni-Fe-0 ~ystem.~~,~’ The IR spectra of nickel ferrite contain bands typical of a spinel structure (Fig.8).30,31There is a considerable line broadening and a fine structure of the major bands, leading to the suggestion that they correspond to a number of vibrations with close energies. The presence of a band at ca. 630 cm-’ in all ferrite samples is characteristic of spinel systems with ordered cation vacancies, located predominantly in octahedral position^.^^-^^ The peak maxima move to higher frequencies for samples treated at 673 and 873 K, but even for the ferrite synthesized at 873 K, the fine structure of the IR spectra is retained. The nickel ferrite, synthesized at 873 K, is a partially inverse spinel with a lattice parameter of 0.8031 nm for which the cation distribution, deduced from the Mossbauer spectrum may be written as Nio,l(5tFeo.8~,,[Fe, Iron.1(5~Ni0.8(51]04. cations located in octahedral and tetrahedral sites can be distinguished in the spectrum [Fig.~(u)].~’v~~The linewidth is larger and the spectrum can be fitted on the basis of the assumption2’ that the octahedral component is composed of a number of sextets, corresponding to octahedral iron(rr1)with a different number nearest tetrahedral nickel@)neighbours (0, J. MATER. CHEM., 1994, VOL. 4 Table 4 Mossbauer parameters of doublet components in the spectra of iron-nickel hydroxidecarbonates sample d/mm s-' A/mm s-l NFHC3 1 0.377 f0.005 0.95 f0.01 0.364 f0.006 0.68 f0.0 I 0.346 f0.002 0.4 19 & 0.004 NFHC52" 0.375 f0.006 1.05& 0.01 0.367 f0.004 0.799 f0.09 0.347 f0.002 0.467 & 0.004 NFHC21 0.374 f0.005 1.05f0.01 0.344f0.003 0.495 f0.006 0.369f0.003 0.803 f0.006 NFHCl 1 0.369 k0.003 1.041f0.009 0.347 f0.003 0.485 & 0.005 0.369 f0.003 0.795 f0.005 NFHC12 0.384f0.009 1.25f0.02 0.36 1 f0.005 0.469 f0.009 0.369 f0.006 0.94f0.01 0.37 1 f0.007 0.71 +O.Ol NFHC 13 0.425 f0.0 1 1.61 f0.02 0.382 f0.006 1.14f0.0 1 0.370 f0.003 0.537 fO.005 0.372 f0.004 0.838 & 0.007 "Sample NFHC52 has Fe :Ni =215.0.80 7 0.70 v) E5 0.60 0.50 C 10 1 .oo 2100 3100 Fe:Ni molar ratio Fig. 5 Quadrupole splitting of the full doublet of Fe"' as a function of composition Table5 Phase composition and specific surface area of oxides, obtained by thermal decomposition of iron-nickel hydroxidecarbon- ates at 623 K precursor phases S/m2 g-' NFHC3 1 NiO +spinel 150 NFHC2 1 NiO +spinel 135 NFHCl 1 NiO +spinel 140 NFHC12 spinel 1 60 NFHC 13 spinel +haematite 168 1 and 2).The internal magnetic fields for each component are presented in Table 6. Bearing in mind that nickel@) has the highest crystal-field stabilization energy for octahedral coordi- nation, compared with other divalent transition-metal cations (e.g. MnZ+, Fez+, Coz+ and CU~+),~~it can be concluded that a metastable cation distribution is obtained. At 1073 K an inverse spinel Fe[NiFe]O, is obtained and the octahedral iron cations have no nickel cations as nearest tetrahedrally situated neighbours. The two sublattices are significantly different in the hyperfine magnetic fields [Fig.9(b) and relative average number Fe"' FWHM/mm s-' area (YO) near neighbours ~~ 0.250 & 0.004 15 1.s 0.250f0.004 25 0.250 & 0.004 0.28 1 +0.004 60 17 1.7 0.28 1 f0.004 38 0.28 1 k0.004 45 0.294 f0.004 22 2 0.294 f0.004 27 0.294 & 0.004 51 0.301 f0.003 22 3 0.301 f0.003 26 0.301 f0.003 52 0.293 & 0.004 7 4 0.293 & 0.004 26 0.293 L-0.004 33 0.293 f0.004 34 0.321 f0.004 5 4.5 0.321 f0.004 14 0.321 & 0.004 39 0.321 f0.004 42 DTA t B 463 a, 373 573 773 T/K Fig.6 DTA profiles of iron-cobalt hydroxidecarbonates. (u)-(e) as in Fig.4 Table 6 Mossbauer parameters of nickel ferrite, synthesized by ther- mal decomposition of NFHC12 at 873 and 1073 K Fe"' site d/mm s-A/mm s-' HeffikOe NiFe,O,, 873 K A 0.273 f0.002 -0.002 f0.004 481.2 k0.2 B, 0.374&0.001 -0.064&0.003 513.2 fO.l B2 0.252k0.008 -0.073f0.016 459.4 k0.9 B3 0.31 kO.1 -0.220 f0.024 416.9f1.2 NiFe,O,, 1073 K A 0.255 kO.001 -0.007 fO.OO1 488.52 f0.08 B 0.363 &0.00 1 -0.033 f0.00 I 522.735f0.066 Table 6) and the Mossbauer parameters are close to those reported in previous studies.35 The surface composition of the oxides synthesized at 623 K display a deviation in favour of nickel, when the Fe :Ni ratio increases (Fig.10). The same effect was observed with iron- J. MATER. CHEM., 1994, VOL. 4 1595 30 1600 1200 800 400 wavenu m bedcm-' Fig. 7 IR spectra of iron-nickel hydroxide carbonates (partly decom- posed samples, 400 K). (a)-(d) as in Fig. 3 (Om zg s 1000 600 300 wavenumberkm-' Fig. 8 IR spectra of nickel ferrite: (a)623 K, (b) 673 K, (c) 873 K cobalt mixed and was explained by the lower content of iron on the surface of the precursor. The electron binding energies of nickel (855.1 kO.2eV) and iron (71 1.1 k0.2 eV) in the oxide products exhibit no substan- tial changes associated with the Fe:Ni ratio and are in 100 99 98 100 99 98 97 96 -9-6-3 0 3 6 9 velocity/rnrn s-l Fig. 9 Mossbauer spectra of nickel ferrite, synthesized at 873 K (a) (A denotes tetrahedral positions, while Bistands for octahedral iron with a different number of tetrahedral Ni" cation neighbours) and 1073 K (b) 3.2-1 /c? 2.41 / / , / / / / / 0 / / / 0 / 0.0Oa81-----0.0 1 .o 2.0 3.0 Fe:Ni bulk ratio Fig.10 1ron:nickel ratios on the surface and in the bulk of the oxides, obtained by thermal decomposition at 623 K; (---) nominal ratio agreement with the reference data.37-39 The 0 1s binding energy (530.2k0.2eV) is the same as that determined for the oxides of nickel and ir~n.~~,~ The presence of a-(or y-) Fe,O, on the surface cannot be excluded, having in mind the similarity of their binding energies to that of nickel ferrite.39 There are no indications of the presence of Ni3+ on the surface of the oxides.The oxygen :metal ratio on the surfaces of the oxides increases with the iron:nickel ratio, and for samples with Fe:Ni20.5 it exceeds both the values for the J. MATER. CHEM., 1994, VOL. 4 2.21 M304 1.oj, 0.0 1.o 2.0 3.0 Fe:Nibulk ratio Fig. 11 Oxygen-to-metal ratio for the oxides, synthesized at 623 K, compared with the nominal ratio for spinels and sesquioxides spinel oxides and the sesquioxides (Fig. 11). As the 0 1s peak is rather symmetric in shape, a single type of oxygen surface species should be present. Cationic vacancies, already deter- mined in the IR spectra, may be responsible for the excess of oxygen. Conclusion Iron-nickel hydroxidecarbonates with reevesite-like structures can be synthesized by coprecipitation of the corresponding divalent metal salts with sodium carbonate and a wide vari- ation of the nicke1:iron ratio is attained.The linear change in unit-cell volume, the shift of the DTA peaks and the dependence of the Mossbauer parameters on composition confirm the assumption of continuous divalent-trivalent cation substitution in the cationic layer. No additional phases have been detected. The nickel ferrite synthesized at 873 K from the corresponding precursor has a metastable, partially inverse cation distribution. For all oxide samples synthesized at 623 K enrichment of the surface with nickel(I1) and high oxygen-to-metal ratios are found. References M.Tarasevich and B. Effremev, in Electrodes of Conductive Metallic Oxides, ed. S. Trassati, Elsevier, Amsterdam, 1980, p. 221. G. Fornasari, S. Gusi, F. Trifiro and A. Vaccari, Ind. Eng. Chem. Res., 1987,29, 1500. 0. Clause, M. Gazzano, F. Trifiro, A. Vaccari and L. Zatorski, Appl. Catal., 1991,73,217. 0. Clause, B. Rebours, E. Merlen, F. Trifiro and A. Vaccari, J. Catal., 1991, 133,231. Ph. Courty, D. Durand, Ed. Freund and A. Sugier, J. Mol. Catal., 1982,17, 241. S. Gusi, F. Pizzoli, F. Trifiro, A. Vaccari and G. Del Piero, Studies in Surface Science and Catalysis, Preparation of Catalysts IV, ed. B. Delmon, P. Grange, P. Jacobs and G. Poncelet, Elsevier, Amsterdam, 1987, vol. 31, p 753. 7 F. Trifiro, A. Vaccari, 0.Clause and M. Gazzano, Proc.2nd Eur. Con$ on Advanced Materials and Processes ‘Euromat ’91’, ed. T. Clyne and P. Withers, The Institute of Materials, London, 1992, vol. 3, p. 304. 8 J. Nunan, P. Himelfarb, R. Herman, K. Klier, Ch. Bogdan and G. Simmons, Inorg. Chem., 1989,28,3868. 9 A. Vaccari, Chimica Industria, 1992,74, 174. 10 A. McKenzie, C. Fishel and R. Davis, J. Catal., 1992,138,547. 11 W. Cares and J. Hightower, J. Catal., 1971,23, 193. 12 H. H. Kung, Transition Metal Oxides: Surface Chemi.stry and Catalysis, in: Studies in Surface Science and Catalysis, ed. B. Delmon and J. Yates, Elsevier, Amsterdam, 1989, vol. 45, p. 171. 13 F. Cavani, F. Trifiro and A. Vaccari, Catal. Today, 1991, 11, I 73. 14 E. Uzunova, D. Klissurski, I. Mitov and P. Stefanov, Chem.Muter., 1993,5, 576. 15 D. Klissurski and E. Uzunova, Chem. Mater., 1991,3, 1ObO. 16 R. Allmann, Chimia, 1970,24,99. 17 M. Ulibari, M. Hernandez and J. Cornejo, J. Muter. Sci., 1991, 26, 1512. 18 G. Busca, V. Lorenzelli and V. Escribano, Chem. Mater., 1992, 4,595. 19 G. Busca, V. Lorenzelli and V. Bolis, Muter. Chem. Phjs., 1992, 31,221. 20 T. Sato, H. Fujita, T. Endo and M. Shimada, React. Solids,1988, 5,219. 21 W. Reichle, Solid State lonics, 1986,22, 135. 22 G. Sawatzky, F. van der Woude and A. Morrish, Phys. Rrv., 1969, 187,747. 23 R. Shannon, Acta Crystallogr., Sect A., 1976,32,751. 24 G.Busca and V. Lorenzelli, Muter. Chem., 1982,7,89. 25 K. Nakamoto, Infrared and Raman Spectroscopy of Inorganic and Coordination Compounds, Wiley, New York, 1978, pp.283; 380. 26 N. Minkova, M. Krusteva and G. Nikolov, J. Mol. Struct., 1984, 115,23. 27 N. Greenwood and T. Gibb, Mossbauer Spectroscopy, Chapman and Hall, London, 1971, p. 148. 28 F. Schneider and H. Schmalzried, 2. Phys. Chem. NF, 1990,. 166,l. 29 A. Pelton, H. Schmalzried and J. Sticher, J. Phys. Chem. Solids, 1979,40,1103. 30 P. Tarte, Spectrochim. Acta, 1962, 18,467. 31 V. Farmer, Infrared Spectra of Minerals, Mineralogical SOC., London, 1974, p. 191. 32 B. Gillot, F. Jemmali and A. Rousset, J. Solid State Chem., 1983, 50, 138. 33 W. White and B. De Angelis, Spectrochim. Acta, Part A, 1967, 23,985. 34 B. Gillot, F. Bouton, J. Ferriot, F. Chassagneux and A. Rousset, J. Solid State Chem., 1977,21,375. 35 J-P. Morel, J. Phys. Chem. Solids, 1967,28,629. 36 A. Wells, Structural Inorganic Chemistry, Clarendon Press, Oxford, 5th edn., 1986, ch. 13, p. 314. 37 G. Allen, S. Harris, J. Jutson and J. Dyke, Appl. Surf. Sci., 1989, 37, 11 1. 38 N. McIntyre and D. Zetaruk, Anal. Chem., 1977,49,1521. 39 C. Brundle, T. Chuang and K. Wandelt, Surf. Sci.,1977,68,459. 40 N. McIntyre and M. Cook, Anal. Chem., 1975,47,2208. Paper 31027581; Received 17th May, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400153

出版商:RSC    

年代:1994

数据来源: RSC

摘要:

J. MATER. CHEM., 1994, 4(1), 161-165 Intermolecular Magnetic Interactions in 2,4,6=Tri(tert=butyl)phenoxy and 2,6=Di(tert=butyl)=4=tritylphenoxyRadicals Karsten Jung," Paul M. Lahti,*" Peter Sheridan," Shawn Britt? Wan-ru Zhang,b and Christopher Landee*b a Department of Chemistry, University of Massachusetts, Amherst, MA 01003 USA Department of Physics, Clark University, Worcester, MA 01610 USA Variable temperature magnetic susceptibility and electron paramagnetic resonance studies of neat solid-state samples of 2,4,6-tri(tert-butyl)phenoxy radicals (3)and 2,6-di(tert-butyl)-4-tritylphenoxy radicals (4) showed antiferromagnetic coupling in both materials. Spin counts indicated appreciable dimerization of radical content in both species (up to ca. 70%) but the remaining spins in each sample were well described by a model of antiferromagnetically exchange- coupled Heisenberg S =1/2 dimers.Antiferromagnetic behaviour was more pronounced in 3 than in 4, with exchange constants for these being ca. -7 and -1 K, respectively. There is evidence for substantial numbers of isolated radical spin sites in 3 but not in 4, with the Weiss constants for these being -2.4 and -0.85 K, respectively. The changes in behaviour from 3 to 4 are attributable to changes in the average inter-radical distance, induced by the large 4-trityl group in 4. The study of intermolecular exchange effects in crystalline and powdered stable radicals is of substantial interest as part of ongoing efforts to design new types of magnetic material.While much is known about the relationship between molecu- lar structure and intramolecular exchange effects (ground- state spin multiplicity), much less is known about exchange coupling between molecules as a function of molecular pack- ing. In addition, it is much harder to control crystal packing of molecules in the solid state, even where plausible models for intermolecular exchange have been suggested. As a result, a substantial amount of work is needed to build up a database correlating bulk magnetic properties with the structures of different types of stable radicals. The galvinoxyl radical has been studied in the solid state by various workers.'-3 Its magnetic nature in the solid state is fascinatingly variable and depends on the technique used to precipitate the neat radical or whether the radical is admixed with its non-radical phenolic precursor.A variety of nitroxide- based radicals have been ~tudied,~with some examples showing ferromagnetic coupling between radicals in the solid state at extremely low temperatures. Studies of verdazyl-based radicals by Allemand et aL5 have elegantly shown that magnetic exchange effects can be 'fine-tuned' in some cases by a change in substituent. Whereas 1-(4-nitropheny1)-3,5-diphenylverdazyl,1, exhibits antiferro- magnetic coupling in the solid state, the very closely related derivative with an added phenyl ring, 2, exhibited one-dimensional chain ferromagnetic coupling with a Weiss con- stant of @ = + 1.6 K.5 The change in behaviour could be attributed to a change in relative intermolecular geometry in the crystal packing.This last work is notable, in being an attempt to affect in a rational and systematic manner the relationship between molecular structure in a radical and its solid-state bulk magnetic behaviour. A fairly modest change in molecular structure leads to a substantial change in bulk magnetic nature. However, to make use of this strategy, more data are needed correlating the structures of radicals to bulk magnetic nature. In this contribution, we describe our study of two differently substituted phenoxy radicals, 2,4,6-tri( tert-buty1)phenoxy (3) and 2,6-di( tert-butyl)-4-tritylphenoxy (4).We chose to investi- gate these systems because: (1) their syntheses were well described by Muller et al.in previous work,6 (2) they are examples of the simplest of reasonably stable conjugated 1 X=H 2 X=CeHs radicals; (3) they are known to have substantial retention of unpaired spins, even in the solid state. These radic;ils were synthesized as shown in Scheme 1, and their bulk magnetic behaviours were studied as functions of temperature. Experimental General All solvents were distilled and thoroughly purged with nitrogen or argon before use. 2,4,6-Tri( tert-buty1)phenol and 2,6-di( tert-buty1)phenol were obtained from Aldrich. 2,4,6-Tri( tert-buty1)phenol was recrystallized from ethanol and dried thoroughly under vacuum before use. All melting points were uncorrected.Infrared spectra were obtained on a Perkin-Elmer 1420 ratio-recording spectropho- tometer, and were referenced to the 1601 cm-' peak of polystyrene. Electron paramagnetic resonance (EPR 1 spectra were obtained on a Bruker Instruments ESP-300 X-band spectrometer. EPR peak positions were referenced to solid diphenylpicrylhydrazyl radical at g =2.0037. 2,4,6-Tri(tert-butyl) phenoxy (3) Method A 2,4,6-Tri( tert-buty1)phenol (0.25 g, 0.96 mmol) was dissolved in 25 ml of hexane in a Schlenk inert atmosphere reaction tube, and the resultant solution degassed thoroughly by bubbled nitrogen. Then, manganese dioxide (0.78 g, 9 1 mmol) was added, and the reaction stirred under nitrogen for 1.5 h to give a deep-blue solution, which was filtered through a 162 OH 0' MnO-Jhexane or I But ootassium hexacvanoferrate 1111)..H,O/NaOGpentane I But 3 OH OH (C6H5)3C4H But*But AcOWH,SO* A potassium hexacyanoferrate(Ill) H20/NaOWpentane I 0' Scheme 1 Schlenk fritte into a nitrogen-filled receiver. The solvent was then removed by a stream of nitrogen to give radical 3 as a blue-black solid in nearly quantitative yield, which could be stored under vacuum at -20 "C. Method B Potassium hexacyanoferrate(n1) (1.4 g, 4mmol) and 0.3 g of potassium hydroxide were dissolved in 20 ml of water, covered with 20ml of pentane, and degassed with nitrogen in a Schlenk reaction vessel equipped with stir bar. 2,4,6-Tri( tert-buty1)phenol (0.39 g, 1.5 mmol) was then added with brisk stirring, and the reaction further stirred under nitrogen for 2 h.The lower aqueous layer was removed by pipette with nitrogen flowing over the reaction surface, 6g of anhydrous potassium carbonate (ignited and cooled under vacuum before use) added, the solution stirred for a further 2 h, filtered under nitrogen, and the solvent removed with a steady flow of nitrogen. The resultant blue-black powder 3 (0.38 g, 96%, m.p. 87 "C) could be quickly sealed in ampoules under vacuum and stored at -20 "C. Infrared spectroscopy of radical 3 was consistent with that noted by Muller in earlier work,6 with radical formation indicated by disappearance of the sharp reactant phenolic -OH absorption at 3600 cm-l. EPR spectra of samples in uucuo up to a year old retained much of their spin intensity.Upon exposure to the air, the sample rapidly lightened to a greenish colour; even 3 min of exposure caused up to 7% oxygen uptake as determined by elemental analysis. EPR: 9.77 GHz, room temperature, neat solid, giso=2.0046 (slightly asymmetric gaussian lineshape); 9.57 GHz, 11 K, neat solid, J. MATER. CHEM., 1994, VOL. 4 2,6-Di(tert-butyl)-4-tritylphenoxy (4) This procedure was essentially identical to Method B used for radical 3 above. 1.93 g (3.5 mmol) of potassium hexacyano- ferrate(m), 0.3 g of potassium hydroxide, 0.5 g (1.17 mmol) of 2,6-di(tert-butyl)-4-tritylphenol, and 6 g of freshly ignited pot- assium carbonate were used. Removal of pentane solvent from the green-black solution yielded a greenish solid 4 (0.4 g, 85%), which could be stored at -20°C in ampoules under vacuum.Infrared spectroscopy showed radical formation by disappearance of the sharp -OH absorbance at ca. 3600 cm-l. EPR: at room temperature, 9.77 GHz, neat solid, giso=2.0052, gaussian lineshape; at 11 K, 9.58 GHz, neat solid, giso=2.0055, gaussian lineshape. IR (solid, KBr) v/cm-' 2940, 1660, 1550, 1470, 1440, 1350, 1250, 690. Magnetic Susceptibility Studies The magnetic moments of 3 and 4 were measured as a function of temperature in a PAR Model 155 vibrating sample magnet- ometer (VSM), which was calibrated against the magnetic moment of high-purity nickel. The temperatures of the samples were varied between 2 and 60K in a custom-built helium cryostat,* and were measured with a commercial Ga/Al/As diode (Lake Shore Cryogenics model TG- 120) which had been checked in situ against a calibrated diode.The tempera- ture scale was verified by measurements of HgCo(SCN),.9-'o The dependence of moment as a function of field at 4.2 K was linear up to a maximum applied field of 10 kG. The sample moments were therefore obtained in an applied field of 10 kG, and the magnetic susceptibility (xM) calculated as the ratio 250 200 7 I 150 Q)3 100 I& 50 0 10 20 30 40 0.20 I 1 ~ 0.15--E Y = 0.10-6 kr: 0.05 -simulated powder pattern a good fit to g, =2.006(2), g,,= 2.002(0). IR (solid, KBr) v/cm-l 2950, 1720, 1630, 1350, 0.00$1 1250, 690.0 10 20 30 40 2,6-Di(tert-butyl)-4-tritylphenol(5) temperature and (b)X,Tuersus absolute temperature for radical 3. In (a) squares are experimental data xi ',the solid line is a Curie-Weiss T/K Fig. 1 Plots of (a) inverse molar susceptibility (xi') uersus absolute This compound was synthesized from commercially available 2,6-di(tert-butyl)phenoland triphenylmethanol by the pro- fit to the data at T>10K, where C=0.180(5) emuKmol-' and 0= -2.4( 1) K. In (b),squares are experimental xhlTdata, the solid cedure of Cook and Gilmour7 to give light-yellow crystals of 5 with m.p. 179-182°C (lit.7 m.p. 179.5-180°C). line is an S= 1/2 Heisenberg dimer fit with J/k,= -6.5 K and C=0.178 emu K mol-' J. MATER. CHEM., 1994, VOL. 4 M/H.Corrections for diamagnetism were made using Pascal's constants. Variable Temperature Electron Paramagnetic Resonances Studies All variable temperature EPR work was carried out using an APD Cryogenics CS-202 Displex closed-cycle helium circulat- ing cryostat with APD Cryogenics DMX-15 ESR Suprasil X-band vacuum shroud and OHFC copper sample spindle. A sample of neat solid radical was quickly (<60 s) attached to the copper sample spindle with Apiezon N grease, the vacuum shroud emplaced, and the sample evacuated to ca. 0.005 mTorr. The sample was cooled to 10 K, the sample zone of the Displex manoeuvered into the EPR spectrometer cavity, and spectra obtained as a function of temperature. Temperature measurements were monitored using an Fe-doped gold-chrome1 thermocouple (Scientific Instruments, Inc.; Model CG07FC-4) placed at the top of the copper sample spindle.EPR signal intensities were obtained by double integration of the radical spin resonances over the same magnetic field range for all spectra, using standard Bruker software routines on the ESP-300. Results and Discussion Fig. 1 shows both l/xm us. Tand xmTus.Tplots for radical 3. The paramagnetic susceptibility of samples made by entirely different oxidation procedures showed varying amounts of free radical spin (34-47%) in the solid state, but exhibited the 1ool (bJ 80. o' fB60. k 40. 0 20. OD 0 20 40 60 T/K Fig. 2 Solid-state EPR spectrum of neat 3 at 10 K, over the range 3360-3460 G at v, =9.58 GHz (a),and a plot of EPR spectral intensity times absolute temperature (IT)versus absolute temperature (b).At each datum in (b),I =double-integrated peak intensities of the g =2 spectral region same thermal behaviour. The data presented in Fig. 1 for sample method A are typical. A Curie-Weiss fit to the inverse susceptibility gave excellent agreement with the data above 10 K, with a Curie constant of C =0.180(5) emu K mol -and a Weiss constant of 0 = -2.4(1) K, indicating that antiferro- magnetic interactions are present. The onset of antiferromagnetic behaviour in samples of 3 occurs at sufficiently high temperatures to be observed by means of the thermal dependence of the EPR spectral line intensity.Fig. 2(a) shows the solid-state EPR spectrum for the neat solid 3 at 10 K, and Fig. 2(b) shows the EPR line intensity times temperature us. temperature (ITus. T)behaviour. Thus, both EPR and magnetic susceptibility plots exhibit substantial antiferromagnetic curvature. There was some slight variation in the EPR lineshape and peak-to-peak linewidth of solid 3 as a function of temperature, but the effect was not substantial at temperatures up to 77 K. No half-field band was observed in the g=4 region of the EPR spectrum, which would have been indicative of well defined triplet radical pairs. Given the the antiferromagnetic sign of the qualitative interaction, a physically plausible model for exchange between the radicals would involve interaction between pairs of mol- ecules.Hence, the magnetic model of an antiferromagnetic S =1/2 Heisenberg dimer was selected for quantitative com- parison studies. The susceptibility data for 3 prepared by method A were compared to the predictions for the susceptibil- ity of a model spin system? in which the interaction Hamiltonian is given by eqn. (l), where J is the exchange constant and Si is the electron spin quantum number: H = -2J1Si.Si+ 1 The data and the results of the fitting procedure are presented in Fig. l(b) where the data are plotted as ~,Tus. T. Since the Curie law T-' dependence is removed by this method of plotting, the x,T product clearly reveals the presence of antiferromagnetic interactions by giving a decrease in this product as the temperature is reduced, while a simple param- agnet would have a ~~Tproduct which is temperature inde- pendent and equal to the Curie constant.Reasonable agreement was found over the entire temperature range for eqn. (1) with J(3)/kB=-2.7K and g=1.33 (C= 0.167 emu K mol-I), but systematic deviations were found that were largest at low temperatures. Given the possibility of the existence of isolated paramagnetic spins in the material, the model was expanded to include a term for the susceptibility of uncoupled spins. Excellent agreement was found for three- parameter fitting to the data of Fig. 1, using J and g with eqn (l),and a paramagnetic term involving a fraction of uncoupled spins F. However, the three parameters were found to be t The dimeric model is based upon work described in ref.11. While the dimeric model is physically logical, we found that a one-dimensional Heisenberg chain model also gives a good fit to the susceptibility data, using a procedure such as that of Bonner and Fisher.12 The fit lines for the susceptibilities of dimer and chain models are very similar, both predicting a maximum in the suscepti- bility at low relative temperatures [T,,, (dimer)= 1.25)Jl/kB, T,, (chain)= 1.281JlkBl. The exchange constants derived from the linear chain model were J(3)/kB= -3.05(5) K and J(4)/kB=0.61(2) K. As was the case for the dimer model, it was necessary to include a correction for uncoupled spins for 3, but not for 4. Below T,,, the susceptibility of a dimer should decrease exponentially to zero while that of the chain should reach a finite value as T approaches zero. While in principle the data for 3 were at sufficiently low temperatures to distinguish between the dimer and chain models, the presence of the large fraction of uncoupled spins obscured the predicted maximum and prevented a definitive determination of the correct model.The chain model, however, seems harder to explain, in terms of the physical arrangement of radicals leading to such behaviour within the material. highly correlated. A non-linear least-squares fit of the data yielded the parameter set J(3)/kB= -5.19 K, g= 1.373, and F = -0.33, where a negative F is physically meaningless by our model. A set of parameters with virtually the same statistical excellence of fit was obtained by simultaneously increasing the magnitude of the antiferromagnetic interaction and increasing the fraction of uncoupled spins, to yield J(3)/kB=-6.5 K, g= 1.38, (C =0.178 emu K mol-I), and F = 0.40. The Curie constant calculated for 1 mol of S= 1/2 with g=2.00 is 0.375 emu K mol-', so the experimental value is 47% of the theoretically available number of radical spins.When sample preparation method B was used for 3, the exchange constant was found by the same method just described to be quite similar to that found from method A, with J(3)/k,= -8.0 K, g=O.99 (C=O.128 emu K mol-'), and F=0.40. The value for the exchange constant of 3 can be estimated to be ca.J(3)/kB= -7 1 K for both preparation methods A and B, with the considerable uncertainty being due to the difficulty of determining the exchange constant in the presence of the substantial fraction of uncoupled spin sites in 3. The similarity of behaviour between the two different preparation methods indicates that no solvent is incorporated into the antiferromagnetically coupled portion of these samples. In addition, elemental analysis of these samples does not show solvent incorporation. Fig. 3 shows the l/xmand ~~TversusTplots for radical 4, while Fig. 4 shows the solid-state EPR spectrum. The product xmTalso decreases at low temperature but at a slower rate than is found for 3. A good fit to the Curie-Weiss law was found for the data in Fig.3(a) at T>15 K, with C= 0.12( 1) emu K mol-' [solid line, Fig. 3(b)], corresponding t I E 0 5 10 15 0.12 I I 0.00 i II 1 I 0 5 10 15 20 TIK Fig. 3 Plots of inverse molar susceptibility (x;') versus T/K (a) and of zrnversus T/K (b)for radical 4. In (a), squares are experimental data x&', the solid line is a Curie-Weiss fit to the data at T> 15 K, where C=0.12(1) emu K mol-' and @= -0.85(5) K. In (b),squares are experimental X,Tdata, the solid line is a S= 1/2 Heisenberg dimer fit with J/k,= -0.98(2) K and C=0.114(5)emu K mol-' J. MATER. CHEM., 1994, VOL. 4 Fig. 4 Solid-state EPR spectrum of neat 2 at 10 K, over the range of 3360-3460 G at v, =9.58 GHz to 31% of the theoretical number of free-radical spins.Interestingly, it was not found necessary to make a paramag- netic correction for the contribution of non-interacting spins in 4. Overall, the qualitative result of using the large 4-trityl substituent in 4 is to diminish substantially the magnitude of the antiferromagnetic interaction in comparison to the interaction in 1, approximately from J(3)/kB=-6 K to J(4)/kJj=-1 K. Owing to the oxygen instability and dimerization proclivity of these radicals, we were not able to obtain crystal structures to reveal the crystal packing geometries in 3-4. Hence, we can interpret the bulk behaviour of the radicals only in fairly qualitative terms. Presumably the very large 4-trityl group in 4 keeps these radicals further separated than is true in samples of 3, resulting in a smaller antiferromagnetic cooperation in 4.Despite apparently substantial amounts of spin-destroying dimerization in both 3 and 4, the spins which are present behave as antiferromagnetically coupled dimers, with allow- ance made in the samples of 3 for some isolated radical spins. We conclude that the spins are not distributed randomly within these samples but that regions of stable co-operative behaviour occur to a substantial degree. Conclusions 2,4,6-Tri( tert-buty1)phenoxy radical synthesized by the methods of Muller et al. yields a blue-black powder solid that rapidly degrades in the presence of oxygen, but which is much more stable under vacuum at -20 "C. Although the blue solid has only 34-47% of the possible theoretical spins based upon the observed Curie constant, a substantial number of the spin sites interact in an antiferromagnetic fashion at temperatures below 25 K, as shown by analysis of both magnetic susceptibility and EPR peak intensity as functions of temperature.2,6-Di(tert-butyl)-4-tritylphenoxy radicals synthesized under the same conditions show similar stability and ca. 31% of the theoretically available spins according to the Curie constant. The spins in the latter radical also interact antiferromagnetically, but much more weakly than in 2,4,6-tri( tert-buty1)phenoxy radicals. The results are consistent with substantial dimerization of both phenoxy radicals in the solid state, leaving regions of free radicals that interact to give antiferromagnetically coupled radical pairs, as well as some isolated paramagnetic regions of radical spins in the case of 3.While the oxygen instability of these radicals has prevented us from obtaining single crystals of them, the magnetic properties of the neat powdered materials is intriguing, show- ing that even these very simple radicals can exhibit substantial exchange interaction in the solid state. As a result, it should be possible to study a variety of related phenoxy-based J. MATER. CHEM., 1994, VOL. 4 165 systems with reasonable hope of observing co-operative mag- netic behaviour. In addition, the appreciable effect of substi-tution (4-trityl for 4-tert-butyl) on the exchange strength exemplifies the possibilities for alteration of exchange strength through molecular structural changes, even in cases where crystalhne structure is unknown and hence not usable as a guide to morphologically based property changes.K. Awaga, T. Sugano and M. Kinoshita, Solid State Commun., 1986,57,453; J. Chem. Phys., 1986,85,2211. L. Y. Chiang, R. B. Upasani, H. S. Sheu, D. P. Goshorn and C. H. Lee, J. Chem. Soc., Chem. Commun., 1992,959. e.g. M. Kinoshita, in Magnetic Molecular Materials, ed. D. Gatteschi, 0. Kahn, J. S. Miller and F. Palacio, Kluwer, Dordrecht, 1991, pp. 87ff; also Proc. Symp. Chem. Phys. Molec. Based Mater., Tokyo, Japan, 25-30 October 1992, Mol. Cryst. Liq. Cryst., 1993,232-233, 1. This work was supported by the University of Massachusetts Materials Research Laboratory and the National Science Foundation (DMR 9006470 and CHE 9204695). We acknowledge with pleasure the exchange programme between the University of Massachusetts (Amherst) and the Universitat 5 6 7 8 P. M. Allemand, G. Srdanov and F. Wudl, J. Am. Chem. Soc., 1990,112,9391. E. Muller, R. Mayer, B. Narr, A. Schick and K. Schemer, Just. Liebigs Ann. der Chem., 1961,645, 1. C. D. Cook and N. D. Gilmour, J. Org. Chem., 1960,25,1429. C. P. Landee, R. E. Greeney and A. C. Lamas, Rev. Sci. Instr., Freiburg, which enabled K.J. to do research at the former ins tit u tion. 9 1987,543,1957. B. N. Figgis and R. S. Nyholm, J. Chem SOC.,1958,4190. 10 H. St. Rade, J. Phys. Chem., 1973,77,424. 11 B. Bleaney and K. D. Bowers, Proc. Roy. SOC. London Ser. A., References 1952,214,451. 1 K. Mukai, H. Nishiguchi and Y. Deguchi, J. Phys. Soc. Jpn., 1967, 12 J. C. Bonner and M. E. Fisher, Phys. Rev. A, 1965,135,640. 23, 125; K. Mukai, Bull. Chem. SOC.Jpn., 1969, 42, 40 K. Mukai and J. Sakamoto, J. Chem. Phys., 1978,68,1432. Paper 3/01886E;Received 2nd April, 1993

ISSN:0959-9428    

DOI:10.1039/JM9940400161

出版商:RSC    

年代:1994

数据来源: RSC