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Chem. Mater. 1997, 9, 1036-1041
A Chimie Douce Route to Pure Iridium Oxide N. Bestaoui Institut des Mate´ riaux de Nantes (CNRS UMR110), 2 rue de la Houssinie` re, BP 32229, 44322 NANTES Cedex 03, France
E. Prouzet*,† Laboratoire des Membranes et Proce´ de´ s Membranaires (CNRS UMR 5635) E.N.S.C.M, 8 rue de l’Ecole Normale, F-34296 Montpellier Cedex 5, France Received December 10, 1996. Revised Manuscript Received February 3, 1997X
Pure iridium oxide has been synthesized through a chimie douce process. This synthesis proceeds by a slow basic hydrolysis of hydrated iridium(III) chloride, at room temperature. A highly disordered oxihydroxide IrO1.45(OH)1.10‚1.5H2O is first obtained. It has been characterized by thermal analysis, particle-induced X-ray emission (PIXE), X-ray diffraction, and X-ray absorption at the iridium LIII edge. Unlike commercial iridium oxide powders, this phase is entirely free of chlorine. Calcination at 900 °C for 12 h leads to a well-crystallized iridium oxide, with the expected rutile-type structure (P42/mnm symmetry, a ) b ) 4.4990(2) Å, c ) 3.1533(2) Å).
I. Introduction Use of new synthesis approaches in solid-state chemistry has been expanding for several years, and, among them, chimie douce processes have opened the field of new materials or new phases that could not have been obtained by high-temperature processes. However, while the classical “shake and bake” approach of solidstate chemistry still remains a good way to reach both purity and high degree of crystallization in many cases, this report shows that a chimie douce process, by its better control of the reaction kinetics, can sometimes succeed where classical processes do not manage to achieve these goals. Initially, we tried to solve problems exhibited by sputtered iridium oxide films. This material is being studied so that its electrochromic properties can be enhanced for the development of large display devices known also as “smart windows”. Such devices require transparent flat batteries with large surface areas, part of which can be reversibly colored by an oxidationreduction mechanism. Some metal oxides exhibit this behavior. Among them, the most studied are the tungsten (WO3) and iridium (IrO2) oxides. Sputtered iridium oxide films have been studied and most of the authors agree that the deep blue to transparent color change occurs through the reduction of Ir(IV) to Ir(III) or Ir(II), assisted by both hydrated and amorphous states.1-6 However, among the problems that still remain, we have established that the presence of chlorine in these thin layers, even a trace, prevents †
E-mail:
[email protected]. Abstract published in Advance ACS Abstracts, March 15, 1997. (1) Gottesfeld, S.; Srinivasan, S. J. Electroanal. Chem. 1978, 86, 89. (2) Rice, C. E.; Brindenbaugh, P. M. Appl. Phys. Lett.. 1981, 38, 59. (3) Hackwood, S.; Dautremont-Smith, W. C.; Beni, G.; Schiavone, L. M.; Shay, J. L. J. Electrochem. Soc. 1981, 128, 1212. (4) Hackwood, S.; Dayem, A. H.; Beni, G. Phys. Rev. B 1982, 26, 471. (5) Granqvist, C. G. Solid State Ionics 1993, 60, 213. (6) McIntyre, J. D. E. J. Electrochem. Soc. 1979, 126, 2171.
them from achieving a perfectly colorless state upon reduction.7 Chorine arises from the classical iridium synthesis process that begins with iridium chloride. Unfortunately, most of the commercial iridium or iridium oxide powders, which are used as starting materials for thin layer synthesis, contain residual chlorine that cannot be totally removed. Moreover, even for small amounts, the influence of the chlorine is strong. For example, it induces a crystalline distortion in the iridium oxide framework, which prevents any reliable crystallographic parameter refinement to be achieved.7 As the electronic properties are highly associated with the degree of purity, this previous work convinced us that the synthesis of a chloride-free iridium oxide or iridium metal is a prerequisite if one wishes to prepare iridium oxide films that can reach a perfectly colorless state upon reduction. Unfortunately, purification of commercial powders through high-temperature processes cannot be employed since chlorine remains trapped in the iridium oxide framework. Besides, the temperature involved (≈1000 °C), which is also both time and energy consuming, induces a partial reduction of iridium oxide to iridium. Since few attempts have been made to explore other synthesis processes,8,9 we decided to investigate new ways of synthesis through chimie douce processes,10 especially metal salt hydrolysis. This report shows how a controlled hydrolysis of the iridium(III) chloride salt, the most readily available and least expensive commercial salt, leads to a chlorine-free material.11 It describes the hydrolysis conditions that allow total chlorine removal and then the preparation
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S0897-4756(96)00628-X CCC: $14.00
(7) Bestaoui, N.; Prouzet, E.; Deniard, P.; Brec, R. Thin Solid Films 1993, 235, 35. (8) Harriman, A.; Thomas, J. M.; Millward, G. R. New J. Chem. 1987, 11, 757. (9) Osaka, A.; Takatsuna, T.; Miura, Y. J. Non Cryst. Solids 1994, 178, 313. (10) Rouxel, J., Tournoux, M., Brec, R., Soft Chemistry Routes to New Materials. Mater. Sci. Forum 1994, 152-153. (11) Bestaoui, N.; Prouzet, E. Patent FR 94/06773, 1994.
© 1997 American Chemical Society
A Chimie Douce Route to Pure Iridium Oxide
of a material that is 10 times purer than commercially available products. A chlorine-free iridium oxihydroxide is obtained, which readily crystallizes upon moderate heating to an iridium oxide that exhibits an undistorted rutile-type structure. The outline of the present article is as follows: First, we define the parameters for the hydrolysis of iridium(III) chloride, which allows a total substitution of the chlorine bound to iridium by water molecules. The second section describes the synthesis of a pure (i.e., a chlorine-free) iridium oxide. Structure and composition of the as-synthesized phase are described in the third section. Finally, the iridium oxide obtained after calcination is characterized in the fourth section. II. Experimental Section The starting material was IrCl3‚2.6H2O7 (Aldrich). We compared the purity of our material to IrO2 powders produced by Aldrich, STREM, and Jansen Chimica. Analyses were also carried out on the chloride β-IrCl3 (Aldrich). The compounds are referenced as follow: a-IrO2 for the as-synthesized phase obtained after hydrolysis and drying of the iridium(III) chloride, c-IrO2 for the phase obtained after calcination of a-IrO2, and r-IrO2 for a rutile-type oxide provided from Aldrich and used as a reference. pH and potential measurements were performed at 25 °C on a 702 SM TITRINO automatic titrimeter coupled with an Ag/AgCl electrode (Metrohm) as a potentiometric reference and a chloride selective crystal membrane electrode (Metrohm). The pH was increased by adding small amounts of a molar lithium hydroxide solution. The course of the reaction was characterized by the parameter r ) [OH-]/[IrCl3], which represents the ratio of hydroxyl moles added to the number of iridium chloride moles initially present in solution. Hence the amount of added hydroxyl ions is equal to the total amount of chlorine in solution when r ) 3. The extent of hydrolysis was characterized by the parameter q expressed by the ratio q ) [Cl-]sol/3[IrCl3], where [Cl-]sol is the concentration of free chloride, released by the iridium salt. Hydrolysis is then completed when q ) 1. A 0.025 M aqueous solution was prepared at room temperature by dissolving 0.864 g of iridium chloride (Aldrich) IrCl3‚2.6H2O7 in 100 cm3 of deionized water. After 1 day of stirring, an acidic green-brown solution (pH ) 1.8) was finally obtained. The samples for potentiometric measurements were prepared in a 1.7 M KNO3 solution in order to maintain a constant ionic strength. We tested two speeds of base addition arbitrary chosen in order to check if this parameter affects the hydrolysis yield. So-called fast hydrolysis was performed by adding equal amounts of lithium hydroxide (M) in order to reach the value r ) 3 after 1 h (0.1 mL every 48 s). Slow hydrolysis was performed by adding equal amounts of lithium hydroxide (M) in order to reach the value r ) 3 after 24 h (≈0.016 mL every 3 min). Lithium hydroxide was chosen as possible residual lithium can be easily extracted by further electrochemical deintercalation. X-ray absorption spectra were recorded at LURE, on the DCI ring using 1.85 GeV positrons with an average intensity of 250 mA. They were collected in transmission mode at the iridium LIII edge (≈11 210 eV), on the EXAFS I spectrometer. The monochromator was a Si(331) channel-cut and the flux before and after the sample was measured by partially argon filled ionization chambers. The EXAFS (extended X-ray absorption fine structure) spectra were recorded from 11 100 to 12 700 eV with a 3 eV step (2 s time counting). The EXAFS analysis was performed following the curvedwave single scattering theory12 once the absence of significant multiple scattering contributions had been checked.13 The (12) Teo, B. K. EXAFS: Basic principles and Analysis, SpringerVerlag: Berlin, 1986. (13) Prouzet, E. J. Phys: Condensed Matter 1995, 7, 8027.
Chem. Mater., Vol. 9, No. 4, 1997 1037 background absorption was calculated by using a theoretical expression14 and the single atomic absorption of the absorber was interpolated by a fifth degree polynomial between about 11 220 and 12 700 eV. Each spectrum was carefully extracted by varying both the degree and first point of the polynomial, and the best removal of low-frequency noise was checked by further Fourier transformation. The energy of the edge E0 was taken equal to 11 210 eV at the half-height of the absorption. The pseudo radial distribution function around the iridium (RDF) was obtained by a Fourier transform of the weighted ω(k)k2χ(k) spectra, where ω(k) is a window using a Kaiser function (τ ) 2.5) defined between 2 and 19 Å-1. All the further back-Fourier transforms included a removal of this window. The refinements were performed on EXAFS spectra obtained by a back-Fourier transforming between 1 and 2.5 Å. Structural parameters were fitted by using both simplex and least-squares calculations.15 These structural parameters are N, the number of neighbors of a defined species around the iridium, at the distance R, with both thermal and structural disorders described by σ, the Debye-Waller factor (DW). The energy shift ∆E0 between experimental spectra and theoretical phase shift files was also fitted. Theoretical curvedwave of oxygen backscattering amplitude and iridium-oxygen phase shift files were used.16 Γ ) k/λ(k), that is the electron mean free path calculated for an absorber/backscatterer couple, was deduced from references (β-IrCl3 and c-IrO2) and kept at this value for all further fits. The multielectron scaling factor S02 was set at 0.8.12 Finally, the fit residue F was calculated following the formula
∑k[kχ
exp(k)
- kχcalc(k)]2
k
(1)
F)
∑
k[kχexp(k)]2
k
Powder diffraction patterns were recorded on an INEL CPS 120 diffractometer using a curve position sensitive detector with a 0°-120°range in 2θ (Debye-Scherrer geometry) and a 15 min counting time (Cu KR1 radiation). TGA measurements were made with a SETARAM TGDTDA 92 in air between room temperature and 1200 °C with a 10 °C‚min-1 rate. Quantitative elemental analyses were performed at the Centre d’Etudes Nucleaires de Bordeaux-Gradignan (France) by particle induced X-ray emission (PIXE) using a 10 nA beam of 1 MeV protons produced by a Van der Graaff generator. This technique was used once we had checked that the usual analyses (chemical analysis or scanning electron microscopy (SEM) correlated energy-dispersive X-ray element analysis (EDAX)) do not provide reliable results due either to the high absorption of iridium or its difficult dissolution in nonhydrochloric acid medium.
III. Results and Discussion III.1. Hydrolysis of the Iridium(III) Chloride. Little is known about the hydrolysis of chloroiridates. The literature reports only of the kinetics studies of the hydrolysis of (IrCl6)3- species in Na2IrCl6 and Na3IrCl617-19 or of the aquairidium(III) [Ir(H2O)6]3+ 20 but no study of the hydrated iridium(III) chloride was reported. In the present report, the initial species have been first identified by XAS, and then the hydrolysis process of (14) Lengeler, B.; Eisenberger, P. Phys. Rev. B 1980, 21, 4507. (15) James, F.; Roos, M. CERNID Internal Report 75/20 1976. (16) McKale, A. G.; Veal, B. W.; Paulikas, A. P.; Chan, S. K.; Knapp, G. S. J. Am. Chem. Soc. 1988, 110, 3763. (17) Poulsen, I. A.; Garner, C. S. J. Am. Chem. Soc. 1962, 84, 2032. (18) Chang, J. C.; Garner, C. S. Inorg. Chem. 1965, 4, 209. (19) Fine, D. A. Inorg. Chem. 1969, 8, 1014. (20) Gamsja¨ger, H.; Beutler, P. J. Chem. Soc., Dalton Trans. 1979, 9, 1415.
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Figure 1. RDF’s of IrO2 (dashed), IrCl3‚2.6H2O (circles), and the initial solution (line). The solution RDF is mostly marked by a sharp peak at 2.1 Å, due to chlorine backscatters, plus a shoulder at 1.5 Å, due to oxygen neighbors. (Distances are uncorrected from atomic phase shift.) Table 1. Results of Fits for the EXAFS Analysis of the First Shell in β-IrCl3, C-IrO2, and the Initial Solutiona c-IrO2
β-IrCl3
initial solution
Oxygen N(O) R(O) (Å) σ(O) (Å) Γ(O) (Å-2) ∆E0(O) (eV) Chlorine N(Cl) R(Cl) (Å) σ(Cl) (Å) Γ(Cl) (Å-2) ∆E0(Cl) (eV) F (%)
6.0 1.98 0.05 0.5 15.9
1.5
2.0 2.09 0.00 0.5 14.5 6.0 2.35 0.07 0.7 12.4 0.5
4.0 ) 6-N(O) 2.38 0.00 0.7 13.8 2.9
a
N, number of neighbors in the shell; R, real distance between iridium and a backscattering atom; σ, Debye-Waller factor; Γ, parameter linked with the mean free path λ (Γ ) k/λ(k); ∆E0, energy shift; F, residue. Values in italics were not fitted.
iridium(III) chloride by lithium hydroxide was followed by combined pH and potential measurements. Figure 1 compares the RDF of this initial 24 h old solution with iridium oxide c-IrO2 and iridium chloride β-IrCl3. This RDF is characterized by the appearance of a single peak in the same position as for β-IrCl3, plus a small shoulder a shorter distance away, caused by oxygen neighbors. The peak at 1 Å is physically meaningless. Structural refinement confirms a mean iridium environment consisting of two oxygen and four chorine atoms (Table 1). The presence of chlorine around the iridium is rather surprising because MCln complexes are generally unstable in aqueous solution: dissolution in water leads quickly to aquo cations with a total release of chloride ions in solution. Unlike this general statement, our XAS experiment shows that chloroiridium(III) complexes are very inert, possibly due to their stable low spin d6 electron configuration. Their stability is also confirmed by free-chloride measurements (see below) that prove that the initial iridium chloride solution does not hydrolyze. Both the presence of water molecules in the close environment of iridium and the chloride/oxygen ratio prove that iridium chloride does not dissolve as a monomer. Indeed, as we start with IrCl3 and as iridium is in an octahedral site, a monomer would lead to a Cl/O atomic ratio equal to one. On the contrary, the lack of chlorine release and the presence of one-third of water molecules in the ligands imply the formation of oligomers. Existence of oligomers is in line with Harriman’s results who reported that radiolytically prepared iri-
Figure 2. Fast hydrolysis: evolutions of pH (line) and q ) [Cl]sol/3[IrCl3] (dashed), the rate of hydrolysis, in function of r ) [OH]/[IrCl3], the ratio of hydroxyl moles added per moles of iridium chloride. A basic pH is quickly reached whereas no chlorine is released in solution until a large excess of base has been added.
dium oxide clusters at pH ) 2.8 contain four or five iridium atoms.21 LiOH was progressively added to this solution and we recorded the variations of pH and free chloride concentration as a function of the addition of base. We started to record the evolution of free chloride concentration versus the amount of base once the pH ) 2 value had been reached, as no reliable measurement can be performed below this value. The initial pH is low (≈1.7) and its variation with r is marked by a jump. For the fast hydrolysis (Figure 2), the pH jump is observed at r ) 1.0 and it quickly reaches a plateau (pH ≈ 11) at r ) 1.4. This evolution resembles the titration of an acid solution (initial pH ≈ 1.8) with a base, without any additional reaction being involved. Indeed, no chloride ions are detected in the solution below the value r ≈ 4. This observation validates the previous XAS analysis and the hypothesis of the formation of stable oxochloroiridate oligomers in the pristine solution. After r ) 4, q rises rapidly to the value 1.0 (the higher theoretical value) at r ) 5.8. This rise does not stop at 1.0 because the potential/chloride concentration correlation is disrupted by some precipitation due to the oxidation of Ir(III) to Ir(IV) in a basic medium. However, this experiment shows that there is no correlation at all between the pH rising and the chlorine release. Hence, one cannot expect that pH is the only parameter for achieving full hydrolysis of iridium(III) chloride. Nevertheless, the basic pH remains a thermodynamic parameter of the hydrolysis as this occurs when a huge excess of base is added. When the slow hydrolysis is undertaken (Figure 3), the pH jump happens at a higher value of r (r ) 1.33) and this pH jump is lower (final pH at 7 instead of 11). Unlike fast hydrolysis, no plateau is observed but the pH rising slope remains almost similar before and after the jump. The measure of the free chloride rate is also totally different from what was observed for the fast hydrolysis. When we start the measurements (r ) 1.2), 10% chlorine is soon present in solution (q ) 0.1); then, there is a quick increase in chlorine release up to q ) 0.8 between r ) 1.5 and 2. After this value, little precipitation occurs, which prevents us from correctly interpreting the evolution of q. Nevertheless, we may assume that for r ) 3, i.e., the stoichiometric ratio, (21) Nahor, G. S.; Hapiot, P.; Neta, P.; Harriman, A. J. Phys. Chem. 1991, 95, 616.
A Chimie Douce Route to Pure Iridium Oxide
Figure 3. Slow hydrolysis: evolutions of pH (line) and q ) [Cl]sol/3[IrCl3] (dashed), the rate of hydrolysis, in function of r ) [OH]/[IrCl3], the ratio of hydroxyl moles added per moles of iridium chloride. Compared with fast hydrolysis, the pH rising is slower but hydrolysis occurs more rapidly and it is total for r ≈ 3.0.
almost all the chlorine atoms linked to iridium have been released (q ) 1). Comparison of fast and slow hydrolyses reveals how the speed of base addition alters the hydrolysis of iridium chloride. When the base is slowly added, no excess of base is required and total hydrolysis occurs for the stoichiometric ratio of base. The hydrolysis is completed at the value of r ) 3 whereas it has not yet started for the fast hydrolysis. Actually, there is a large difference between the two speeds of base addition since fast hydrolysis is 24 times faster than slow hydrolysis. Fast hydrolysis demonstrates that basic pH is a thermodynamic parameter, as expected, but due to the poor reactivity of initial oligomers, the iridium chloride hydrolysis is unexpectedly a very slow mechanism. Hence, kinetics is the main parameter to control if one wishes to achieve a total hydrolysis of iridium(III) chloride. For example, dynamic light scattering analyses (not shown in this report) confirm that hydrolysis and further particles formation appears, even without base addition, if one waits long enough (more than two months). The speed of base addition also affects the final material structure. TEM observations show that proceeding too quickly with the addition of the base leads to the precipitation of an oxochloroiridate that will trap chlorine in the solid framework. Chlorine that remains trapped in this precipitate cannot be removed by washing. On the contrary, slow hydrolysis progressively gives well-defined small particles with a mean size of 10 nm, which do not hold chlorine, as shown by further chemical analyses. III.2. Synthesis of an Iridium Oxide. The previous study has allowed us to establish a pathway for the synthesis of iridium oxide through a slow hydrolysis of iridium chloride. For this purpose, a 0.0025 M solution of iridium chloride is prepared by dissolving, by stirring for 24 h, 1.6 × 10-3 moles of hydrated iridium(III) chloride (IrCl3‚2.6H2O) in 670 cm3 of deionized water. A slightly brown acidic solution (pH ) 1.8) solution is obtained. A solution of lithium hydroxide (M) is then slowly added in order to get a basic pH (≈12) after 24 h: the addition of 0.01 mL of LiOH every 3 min is monitored by an automatic titroprocessor until the molar ratio [LiOH]/[IrCl3] ) 3 has been reached. Since both the basic pH conditions and the atmospheric oxygen induce the oxidation of Ir(III) to Ir(IV), a blue colloidal suspension is then obtained. This solution is let for 12 h in a 22 cm3 stoppered Teflon bottle, 80%
Chem. Mater., Vol. 9, No. 4, 1997 1039
filled, in a stainless steel autoclave heated at 180 °C, prior to extraction. This step allows the initial particles to aggregate. When the container is opened, a black powder has settled in a colorless LiCl solution. This powder is filtered on 0.2 µm pore size filters and washed with boiling water over a 30 min period, this procedure being repeated four times. The powder is finally dried for 2 h at 200 °C and a black fluffy powder named a-IrO2, is obtained. SEM observations show aggregates of small spherical particles, with a mean size of 80 nm. The c-IrO2 phase is obtained by calcination of this dried powder at 900 °C for 12 h. III.3. Composition and Structure of the AsSynthesized Phase a-IrO2. Table 2 presents the PIXE analysis of chlorine in the dried material a-IrO2 (the formula presented in this table is discussed below). It is compared with commercial iridium oxide powders. First, the analysis confirms that commercial powders contain chlorine. The amount of chlorine is low but even a small amount can modify the electrochromic behavior of IrO2 by inducing structural distortions, which also explains the slight brown color in the reduced state.7 Also, this analysis reveals that the material synthesized through our new process contains almost 10 times less chlorine than the purest commercial powder analyzed. Taking into account the raw result (200 ( 100 ppm), we can even estimate that the amount of chlorine is almost below the detection limit of the technique. Besides, PIXE analysis of a-IrO2, which is a multinuclei sensitive technique, does not reveal any other impurity. The TGA analysis of a-IrO2 exhibits three steps in the weight loss (Figure 4). The first occurs between room temperature and 440 °C, with a 11.7 wt % loss. The second extends up to 900 °C with a weight loss of 4.5%, and the last weight loss of 14.8 wt % begins after 900 °C, increases mostly after 1000 °C, and stops at 1100 °C. This latter step leads to metal iridium.7 As crystallization studies (not shown in this report) establish that the as-synthesized material (a-IrO2) crystallizes directly in the rutile-type symmetry (c-IrO2), it must be assumed that the calcination process leads, from the initial dried material a-IrO2, to iridium oxide c-IrO2 and then to iridium. Therefore, the composition of a-IrO2 can be deduced by considering first the last weight loss (3) as being due to the complete reduction of IrO2 to iridium. Starting with the final weight at 1200 °C, the theoretical weight loss due to the IrO2 reduction should be 14.3 wt %, which is a value close to the observed loss (3) (14.8%). Thus, the 900 °C phase is confirmed to be iridium dioxide IrO2. The purity of the as-synthesized material, checked by elemental analysis, proves that the second weight loss (2) between 440 and 900 °C cannot be due to a loss of impurities but to the loss of water during a transition from an oxihydroxide Ir(IV)O1.45(OH)1.10 to the iridium oxide IrO2. Finally, the first 11.7 wt % weight loss (1) is due to the dehydration of the initial material, with the removal of ≈1.5 mole of water per molar formula. The as-synthesized compound a-IrO2 is thus an hydrated oxihydroxide with the composition IrO1.45(OH)1.10‚1.5H2O. The modulus of the Fourier transform obtained from the EXAFS spectra of a-IrO2, is presented in Figure 5. It is compared with the RDFs of both r-IrO2 and c-IrO2. The oxihydroxide RDF exhibits only a single peak that corresponds to the first shell of oxygen neighbors around
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Table 2. Particle Induced X-ray Emission (PIXE) Analysis of the Amount of Chlorine in Three Commercial Iridium Oxide Powders and the Iridium Oxihydroxide a-IrO2 Obtained through Our Synthesis Cl (ppm) % mol (Cl/Ir) formula a
r-IrO2a
r-IrO2b
r-IrO2c
a-IrO2, this work
4250 ( 50 1.35(2) IrO1.986Cl0.014
2580 ( 70 0.80(3) IrO1.992Cl0.008
2240 ( 90 0.70(3) IrO1.993Cl0.007
200 ( 100 0.06(3) IrO1.45(OH)1.10-Cld ≈ 0.0006
Aldrich. b STREM. c Jansen Chimica.
d
Formula based on thermogravic analysis (see text).
Figure 4. Thermogravimetric analysis of the as-synthesized material a-IrO2 iridium oxihydroxide. From the weight losses (I) due to the dehydration of a-IrO2, the second one (II) to the oxidation of the oxihydroxide, and the last one (III) to the reduction of iridium oxide to iridium, one can deduce the formulation IrO1.45(OH)1.10‚1.5H2O for a-IrO2.
Figure 5. Modulus of the EXAFS spectra Fourier transforms at the Ir LIII edge (RDF) for a-IrO2 (IrO1.45(OH)1.10‚1.5H2O) (circles), r-IrO2 (dashed), and c-IrO2 (line) obtained by calcination of a-IrO2 at 900 °C for 12 h. The highly disordered character of the iridium oxihydroxide is proved by the lack of any outer shell above 2 Å. (The distances are uncorrected from atomic phase shift.) Table 3. Results of the Fits for the EXAFS Analysis of the Oxygen First Shell in a-IrO2, r-IrO2, and c-IrO2a N R (Å) σ (Å) ∆E0 (eV) F (%)
a-IrO2b
r-IrO2c
c-IrO2d
5.8 1.99 0.09 10.6 1.8
5.7 1.97 0.06 12.3 3.5
5.6 1.97 0.05 13.6 1.3
a N, number of neighbors; R, real distance between the iridium and the neighbors; σ, Debye-Waller factor; ∆E0, energy shift; F, fit residue. b IrO1.45(OH)1.10‚1.5H2O. c From Aldrich. d Obtained after calcination of the oxihydroxide.
the iridium atoms. The absence of further shells shows that this compound is disordered. Refinements for the oxygen shell are given in Table 3. The result confirms that iridium atoms in a-IrO2 are in octahedral environment as for iridium oxide: we find 5.8 oxygen neighbors (≈6) at 1.99 Å. The structural disorder appears in the Debye-Waller parameter σ, which is higher for this compound (0.09 Å) than for both the crystallized powders (0.06 and 0.05 Å for r-IrO2 and c-IrO2, respectively). The X-ray diffraction pattern of the oxihydroxide a-IrO2 is characteristic of a highly disordered material. Five broad peaks are observed at 2θ ) 7°, 12°, 35°, 60°,
Figure 6. X-ray pattern of the iridium oxihydroxide IrO1.45(OH)1.10‚1.5H2O dried at 200 °C for 2 h. It is characteristic of a highly disordered material and exhibits only five broad peaks at 7°, 12°, 35°, 60°, and 85° that are usually observed both in rutile or hollandite type structures.
and 85°, respectively (Figure 6). Due to the weakly ordered character of the compound, the Bragg relationship cannot be applied to correlate any of these peaks with reticular planes.22 Nevertheless, these broad peaks can be related to the most intense diffraction peaks either of iridium oxide rutile7 or hollandite structures.23 Thus, a-IrO2 is not entirely amorphous but constructed of IrO6 clusters linked together to give a highly disordered rutile or hollandite type structure. III.4. Structure of the Calcined Phase c-IrO2. Upon calcination, the initial phase a-IrO2 can be observed up to 450 °C. Above this temperature, it begins to crystallize to the rutile type structure. Heating a sample at 550 °C for 1 week leads to a crystallized material but the best crystallization is obtained after a short calcination at 900 °C for 12 h which gives crystals with a mean size of 10 µm. The X-ray pattern exhibits narrow symmetric peaks with a mean half-width of 0.27° and the pattern refinement in the P42/mnm symmetry agrees very well with the parameters of the rutile-type structure (a ) b ) 4.4990(2) Å, c ) 3.1533(2) Å, mean error of 0.007° in 2θ).7 EXAFS also confirms the high degree of crystallization state of c-IrO2 (Figure 6) since, compared with the commercial powder r-IrO2, the peaks intensities are higher for c-IrO2, especially for the farthest peaks which mainly correspond to the outer iridium shells. For r-IrO2, almost no signal can be detected above 4 Å whereas the c-IrO2 RDF exhibits peaks up to 8 Å. The better crystallization of c-IrO2 is also confirmed by the DW parameter of the oxygen shell fit (Table 3), which is smaller for c-IrO2 (0.05 Å) than for the commercial powder (0.06 Å). This study confirms that one can only obtain a well-crystallized iridium oxide from a chlorine-free starting material. IV. Conclusion We have shown in this study that pure iridium oxide can be obtained through a chimie douce process. The (22) Guinier, A. The´ orie et Technique de la Radiocristallographie; Dunot: Paris 1964. (23) Bestaoui, N.; Deniard, P.; Brec, R. J. Solid State Chem. 1995, 118, 372.
A Chimie Douce Route to Pure Iridium Oxide
first step is the hydrolysis of the hydrated irididium(III) chloride that leads to a hydrated iridium oxihydroxide, with the formula IrO1.45(OH)1.10‚1.5H2O. Unlike commercial materials, this compound does not exhibit any trace of chlorine. Preliminary tests have confirmed the electrochromic behavior of a-IrO2 and its ill-ordered state may be helpful to enhance electrochromic properties as disorder should increase the diffusion rate of intercalated species. Calcination of this material yields a rutile-type iridium oxide, with a perfectly
Chem. Mater., Vol. 9, No. 4, 1997 1041
ordered crystalline framework. This synthesis is without doubt a progress if compared with either the commercial products or our previous purification process that only gave a small amount of pure crystals after a 2 month calcination at 1050 °C under oxygen pressure.7 Acknowledgment. The authors thank N. Casey who kindly polished the English. CM9606282
