Mössbauer Characterization and in Situ Monitoring of Thermal

4280

J. Phys. Chem. B 2007, 111, 4280-4286

Mo1 ssbauer Characterization and in Situ Monitoring of Thermal Decomposition of Potassium Ferrate(VI), K2FeO4 in Static Air Conditions Libor Machala,† Radek Zboril,† Virender K. Sharma,*,‡ Jan Filip,† Oldrich Schneeweiss,†,§ and Zolta´ n Homonnay| Nanomaterial Research Centre, Palacky UniVersity in Olomouc, SVobody 26, 771 46 Olomouc, Czech Republic, Chemistry Department, Florida Institute of Technology, 150 West UniVersity BouleVard, Melbourne, Florida 32901, Institute of Physics of Materials AS CR, Zˇ izˇkoVa 22, 61662 Brno, Czech Republic, and Department of Nuclear Chemistry, Eo¨tVo¨s Lora´ nd UniVersity, H-1117 Budapest, Pa´ zma´ ny P. s. 1/A, Budapest, Hungary ReceiVed: December 1, 2006; In Final Form: February 16, 2007

Solid orthorhombic crystals of potassium ferrate(VI) (K2FeO4) of a high-chemical purity (>99.0%) were characterized by low-temperature (1.5-5 K), high-temperature (463-863 K), and in-field (1.5 K/3 T) Mo¨ssbauer spectroscopy. Potassium ferrate(VI) reveals a Ne´el magnetic transition temperature (TN) of ∼3.8 K and a saturation hyperfine magnetic field of 13.8 T at 1.5 K. Spectral line intensities recorded below TN in an external magnetic field of 3 T manifest a perfect antiferromagnetic ordering. For the in situ monitoring of the thermal behavior of K2FeO4, high-temperature Mo¨ssbauer data were combined with those obtained from thermogravimetry, differential scanning calorimetry, and variable-temperature X-ray diffraction measurements. Such in situ approach allowed the identification of the reaction products and intermediates and yielded the first experimental evidence for the participation of CO2 in the decomposition process. As the primary conversion products, KFeO2 and two potassium oxides in equivalent molar ratio, KO2 and K2O, were suggested. However, the KO2 phase is detectable with difficulty as it reacts very quickly with CO2 from air resulting in the formation of K2CO3. The presented decomposition model is consistent with thermogravimetric data giving the mass loss of 8.0%, which corresponds to the participation of 1/6 mol of CO2 and liberation of 3/4 mol of O2 per 1 mol of K2FeO4 (K2FeO4 + 1/6CO2 f KFeO2 + 1/3K2O + 1/6K2CO3 + 3/4O2). An explanation of the multistage reaction mechanism has an important practical impact for the optimization of the solid-state synthesis of potassium ferrate(VI).

Introduction Iron commonly exists in the +2 and +3 oxidation states; however, in a strong oxidizing environment, higher oxidation states of iron such as +4, +5, and +6 can be also obtained.1-4 In recent years, there has been increasing interest in the +6 oxidation state of iron, ferrate(VI) (FeVIO42-), because of its potential use in high-energy density rechargeable batteries, in cleaner (“greener”) technology for organic synthesis, and in treatment of contaminants and toxins in water and wastewater.5-11 Potassium ferrate(VI) (K2FeO4) is the most studied compound among the family of ferrate(VI) derivatives, and its powder remains stable for a long period of time in a moisture-free environment. Chemical, electrochemical, and thermal techniques have been used to prepare solid K2FeO4.12-17 The chemical technique requires large amounts of chemicals, and several steps are involved in the production of solid K2FeO4.12,13 Electrochemical synthesis applies electrolysis of iron (or iron salt) under concentrated hydroxide solution, and a separation step is needed to obtain the solid K2FeO4 product. The formation of passive iron oxide on the electrode reduces the electrochemical yield.14 A possible reduction of iron(VI) in water (2FeO42- + 5H2O f * To whom correspondence should be addressed. Phone: 321-674-7310. Fax: 321-674-8951. E-mail: [email protected]. † Palacky University. ‡ Florida Institute of Technology. § Institute of Physics of Materials AS CR. | Eo ¨ tvo¨s Lora´nd University.

2Fe3+ + 3/2O2 + 10OH-) lowers the product yield of K2FeO4 in wet techniques.18 A dry technique is thus attractive as it can address difficulties associated with the wet techniques. In such a thermally induced solid-state process, potassium- and ironcontaining salts are heated together to obtain K2FeO4.16 However, the decomposition of ferrate(VI) occurs at elevated temperatures used in a thermal synthesis technique, which results in a low yield of K2FeO4 product. An increase in the K2FeO4 yield may be achieved by optimizing temperature conditions under which the secondary decomposition is prevented. This would thus require understanding the mechanism of the thermal decomposition of K2FeO4. Surprisingly, only a few thermal studies of K2FeO4 have been performed and most of the work was carried out in the 1960s and 1970s.19-23 Scholder19 was the first one who studied the thermal behavior of K2FeO4 under a stream of oxygen. The microscopic images of samples heated between 350 and 550 °C showed a mixture of two crystalline phases of dark- and light-green particles; the latter ones were assumed to be KFeO2 with a high probability. The darker phase was calculated to be a solid solution of K2FeO4 and K3FeO4 in a 1:2 molar ratio. The overall mean oxidation number of iron species was measured to be +4.4 and eq 1 was suggested to explain the decomposition process:

5K2FeO4 f K2FeO4‚2K3FeO4 + 2KFeO2 + 2O2

(1)

Ichida20 applied Mo¨ssbauer spectroscopic and X-ray diffraction techniques to determine the decomposition products of K2FeO4

10.1021/jp068272x CCC: $37.00 © 2007 American Chemical Society Published on Web 04/04/2007

Thermal Decomposition of K2FeO4 in Static Air

J. Phys. Chem. B, Vol. 111, No. 16, 2007 4281

in air. The intermediate valence states, Fe(V) or Fe(IV), were not observed during the decomposition process, which was in disagreement with a hypothesis proposed by Scholder.19 Ichida20 noticed light-green residues at the temperatures between 250 and 1000 °C, and all of the X-ray diffraction peaks of heated K2FeO4 samples were identified to belong to those of potassium orthoferrite(III), KFeO2. The suggested chemical process (eq 2) indicates an uncertainty in the chemical composition of potassium oxide residue:

K2FeO4 f KFeO2 + KOx + (2 - x/2)O2

(2)

Later, Fa¨tu and Schiopescu21 used simultaneous thermogravimetry and differential thermal analysis to investigate the thermal behavior of K2FeO4 in air. In this study, K2FeO4 sample decomposed in one step between 50 and 320 °C accompanied with 14.3% loss of the initial mass. The observed mass loss was entirely ascribed to the release of 3/4 mol of O2 for each mole of decomposed K2FeO4 (eq 3):

2K2FeO4 f K2O‚Fe2O3 + K2O + 3/2O2

(3)

However, the theoretical mass loss calculated using eq 3 is 12.1% and thus significantly smaller than that observed in thermogravimetry (TG) experiments. Recently, Tsapin et al.22 studied the thermal decomposition of K2FeO4 under nitrogen atmosphere. The two main steps on the TG curve gave an overall 16.2% mass loss, which is again higher than that expected from eq 3. However, 7.2% from the initial mass loss was ascribed to the dehydration of ferrate(VI) at low temperatures. In a more recent work performed by Madarasz et al.,23 a mass loss of 7.6%, observed during the thermal decomposition of K2FeO4 in nitrogen, was interpreted by the formation of potassium orthoferrite(III) and the mixture of oxide, peroxide, and superoxide of potassium. To sum it up, the literature has principal discrepancies in the results regarding the proposed intermediate oxidation states of iron, Fe(V) and Fe(IV), identification and characterization of iron oxide phases, and mass loss in the thermal decomposition of K2FeO4. Moreover, a difference in the decomposition process in air and inert atmosphere is not clear. In the present paper, the thermal decomposition mechanism of K2FeO4 in static air is described, which is consistent with the results of a combination of in situ thermal analysis, hightemperature Mo¨ssbauer spectroscopy, and variable-temperature X-ray powder diffraction techniques. The in situ techniques offer a possibility of direct monitoring of phase composition during thermally induced process, including the identification of reaction intermediates such as Fe(V) and Fe(IV). Moreover, secondary chemical transformations of the decomposition products because of the interaction with air humidity could be prevented. To the best of our knowledge, application of in situ techniques to investigate the thermal decomposition of K2FeO4 have not been reported yet. Additionally, in-field Mo¨ssbauer spectroscopic analysis of K2FeO4 was performed for the first time, which gave complementary information on the local surrounding of the iron atoms in the FeO42- ion. Experimental Section Crystals of K2FeO4 were prepared by a method of Thompson et al.12 Briefly, oxidation of Fe(NO3)3·9H2O salt by hypochlorite in strong NaOH solution resulted in sodium ferrate(VI) (Na2FeO4), which was then precipitated as potassium ferrate(VI) (K2FeO4) by adding KOH. Crystals were then dried by ethanol and were stored in a vacuum desicator. The purity of

Figure 1. X-ray diffraction pattern of K2FeO4 measured at room temperature.

K2FeO4 crystals was higher than 99.0% as determined by X-ray fluorescence. The thermal behavior of K2FeO4 was studied in static air by DSC and TG using a THASS XP-10 thermal analyzer with a heating rate of 5 °C/min. The transmission 57Fe Mo¨ssbauer spectra were recorded using a Mo¨ssbauer spectrometer in a constant acceleration mode with a 57Co(Rh) source. The isomer shift values were related to metallic iron at room temperature. The measurements were carried out at various temperatures ranging from 1.5 to 863 K in a zero external magnetic field and at 1.5 K in an external magnetic field of 3 T, applied parallel to the direction of the gamma ray propagation. Low-temperature and in-field data were collected using a cryomagnetic system (Oxford Instruments). The in situ high-temperature measurements were performed in a furnace under static air conditions and in a vertical measuring geometry to keep the similar reaction conditions (full air access) as in the VT XRD and DSC/TG experiments. Variable-temperature XRD measurements were performed in static air with a PANalytical X’Pert PRO instrument (Co KR radiation), equipped with a primary monochromator, X’Celerator detector, and a conventional Anton Paar THC heating chamber. The samples were spread on a Pt sample holder and were scanned in the 2θ range of 10-90° at temperatures ranging from 25 to 590 °C. Results and Discussion Characterization of K2FeO4. X-ray fluorescence, Mo¨ssbauer spectroscopy, and X-ray powder diffraction analyses confirm the single-phase character of the sample. XRD pattern of the potassium ferrate(VI) sample, used in this study (see Figure 1), reveals the presence of the orthorhombic K2FeO4 as a single phase, without any indications of diffraction lines of other crystalline compounds. With the assumption of the Pnam space group,24-28 the refined lattice parameters a ) 7.702(2), b ) 10.346(1), c ) 5.862(1) Å, and cell volume 467.1(1) Å3 are close to those determined by other researchers.25,28 In the potassium ferrate(VI) structure, the iron atom of a ferrate ion, FeO42-, has a tetrahedral environment with the mean O-Fe-O bond angles of 109.5°, representing a nearly perfect tetrahedron.24 Such a high symmetry of the neighboring iron(VI) environment is reflected through a singlet Mo¨ssbauer spectrum with a zero quadrupole interaction above the ordering temperature. With respect to a second-order Doppler thermal shift, the value of the isomer shift parameter at 5 K (δ ) -0.79

4282 J. Phys. Chem. B, Vol. 111, No. 16, 2007

Machala et al.

Figure 3. Mo¨ssbauer spectrum of potassium ferrate(VI) measured at 1.5 K in an external field of 3 T (Bext || γ-ray).

Figure 2. Temperature evolution of Mo¨ssbauer spectra of potassium ferrate(VI).

TABLE 1: Hyperfine Parameters and Relative Spectrum Areas of the Mo1 ssbauer Spectra of the Potassium Ferrate(VI) Samplea T [K] Bext [T] component δ [mm/s] Q [mm/s] Bhf [T] RA [%] 5 4

0 0

3.5 3 1.5 1.5

0 0 0 3

singlet singlet sextet sextet sextet sextet sextet

-0.79 -0.79 -0.79 -0.80 -0.80 -0.80 -0.80

-0.04 -0.03 -0.04 -0.04 0

9.0 10.6 11.9 13.7 14.5b

100 64 36 100 100 100 100

a T, temperature of measurement; Bext, external magnetic field; δ, isomer shift; Q, quadrupole shift; Bhf, hyperfine magnetic field; RA, relative spectrum area. b Effective magnetic field Beff ) Bext + Bhf.

mm/s) corresponds well with the previously published roomtemperature data (δ ∼ -0.90 mm/s).22,29,30 The temperature evolution of the hyperfine magnetic field in Mo¨ssbauer spectra (see Table 1, Figure 2) indicates the magnetic ordering temperature of about 3.8 K, which is comparable with the values reported by Herber and Johnson31 (4.2 K) and Corson and Hoy32 (3.6 K). Below the Ne´el temperature, TN, the hyperfine magnetic field increases from 10.6 T at 3.5 K to the saturation value of about 13.7 T at 1.5 K (see Table 1, Figure 2). Interestingly, the magnetic ordering induces a slight decrease of a high symmetry of iron(VI) environment as illustrated by the nonzero quadrupole shift parameters of sextet spectra (see Table 1). To understand the magnetic ordering below TN, in-field Mo¨ssbauer spectroscopy was applied (see Figure 3). Mo¨ssbauer spectroscopy realized in an external magnetic field brings a deeper insight into the magnetic behavior of the studied samples. No such study for Fe(VI) samples has been published yet. It is known that in a random distribution of the magnetic moments within the particle, the intensity ratio of the six lines of the magnetically split Mo¨ssbauer spectra is 3:2:1:1:2:3. In most cases, when the studied magnetic sample is exposed to the

applied magnetic field, the magnetic moments of the atoms respond to it by a rearrangement of their orientations, which is then seen by a change of the line intensities in the Mo¨ssbauer spectrum. If the direction of the external magnetic field is parallel to that of propagation of gamma-rays (as in this study), the second and the fifth lines of the Mo¨ssbauer spectrum vanish for well-crystalline ferromagnetic and ferrimagnetic material, giving an intensity ratio of 3:0:1:1:0:3.33 In the case of in-field analysis of well-crystalline antiferromagnetic samples with a low magnetic anisotropy, a sharp spin reorientation occurs on reaching a critical field when the directions of all magnetic moments flip to a plane perpendicular to the direction of the propagation of γ-rays. This is manifested by an enhancement of intensities of the second and the fifth lines, resulting in an ideal intensity ratio of 3:4:1:1:4:3.34,35 In the case of the K2FeO4 sample, the spectral line intensities ratio of 3:3.9:1:1: 3.9:3 was determined from the spectrum recorded at 1.5 K in an external magnetic field of 3 T (see Figure 3). This ratio reflects almost a perfect antiferromagnetic ordering. Thermal Behavior of K2FeO4 in Static Air. Initially, TG and DSC analyses on K2FeO4 under static air were conducted (see Figure 4). As shown in Figure 4a, the K2FeO4 is thermally stable up to 230 °C. A slight weight loss of 0.3% below 230 °C can be ascribed to a release of adsorbed water, which was also observed through an endothermic effect with a minimum at 80 °C in the DSC curve (see Figure 4b). Between 230 and 280 °C, a weight loss of 8.0% was observed, which suggests a release of oxygen from the ferrate structure. This weight loss is significantly lower compared to the values previously reported in studies under both static air19-21 and inert atmospheric conditions.22 This can be connected with a very low content of hydrated water and a high-chemical purity of the sample used in this study. The main decomposition step is related to an endothermic effect having a minimum at 256 °C (see Figure 4b). In DSC curve, this endo-effect is immediately followed by a broadened exo-effect with a maximum at 270 °C. Importantly, the exothermic peak observed in the present study was not previously reported in an inert atmosphere.22,23 Between 280 and 750 °C, no significant change in the sample weight was recorded. Above 750 °C, the mass loss proceeded because of melting and evaporation of the decomposition products. Next, in situ high-temperature Mo¨ssbauer spectra were collected at four different temperatures (190, 300, 420, and 590 °C) to monitor the transformation process of Fe(VI) and to identify the iron-bearing conversion intermediates (see Figure 5). The in situ measurement started by the insertion of the K2-

Thermal Decomposition of K2FeO4 in Static Air

J. Phys. Chem. B, Vol. 111, No. 16, 2007 4283

Figure 4. TG (a) and DSC (b) curves of K2FeO4 measured in static air.

Figure 5. High-temperature Mo¨ssbauer spectra exhibiting KFeO2 as the only Fe-bearing decomposition product.

TABLE 2: Hyperfine Parameters of the High-Temperature Mo1 ssbauer Spectraa T [°C]

component

δ [mm/s]

Q [mm/s]

Bhf [T]

190

singlet sextet sextet sextet sextet

-1.01 0.06 0 -0.08 -0.19

0.06 0.06 0.05 0.03

45.8 42.9 38.2 28.0

300 420 590

a T, temperature of measurement; δ, isomer shift; Q, quadrupole shift; Bhf, hyperfine magnetic field.

FeO4 sample into the preheated furnace (190 °C) to minimize an interaction of the sample with air humidity. The temperature was kept constant for 5 h during Mo¨ssbauer data recording. The hyperfine parameters of the Mo¨ssbauer spectra are given in Table 2. The spectrum of the sample heated at 190 °C (see Figure 5) consists of two components including a singlet (δ ) -1.01 mm/s) and a sextet with δ ) 0.06 mm/s, Q ) 0.06 mm/ s, and Bhf ) 45.8 T. The singlet corresponds clearly to nontransformed K2FeO4 while the isomer shift value of the latter subspectrum is much lower than expected for an octahedral highspin iron(III) compound.36 Using hyperfine parameter values, the subspectrum was assigned to potassium iron(III) oxide, KFeO2, where Fe3+ ions are tetrahedrally coordinated.20,23 The

sextet of KFeO2 represents only one spectral component in the Mo¨ssbauer spectra measured at 300, 420, and 590 °C (see Figure 5, Table 2). A gradual decrease of its isomer shift with increasing temperature is in a full agreement with the second-order Doppler shift. The quadrupole shift parameter remains practically unchanged (Q ) 0.03-0.06 mm/s) while the hyperfine magnetic field is significantly reduced as the temperature of the measurement approached TN of KFeO2 at 710 °C.17 Overall, Mo¨ssbauer spectroscopy revealed potassium iron(III) oxide, KFeO2, to be the only iron-bearing phase formed during the thermal decomposition of K2FeO4 in air. Contrary to some earlier postulations,19-21 the formation of intermediates containing Fe(V) or Fe(IV) or other Fe(III) oxides (e.g., Fe2O3, FeO(OH)) during in situ thermal decomposition of K2FeO4 were not identified. Finally, VT XRD measurements on K2FeO4 sample were carried out (see Figure 6). The VT XRD spectra demonstrate that the potassium ferrate(VI) readily transforms to potassium iron oxide (KFeO2) upon heating at 190 °C, although some content of K2FeO4 was still observed. This process is consistent with the Mo¨ssbauer measurements (see Figure 5). The decomposition of ferrate(VI) into the KFeO2 phase was completed at 300 °C, where new additional phases including monoclinic potassium carbonate (K2CO3) and potassium oxide (K2O) clearly

4284 J. Phys. Chem. B, Vol. 111, No. 16, 2007

Machala et al.

Figure 6. Variable-temperature in situ XRD patterns of decomposition products of potassium ferrate(VI) in static air. Unmarked peaks at 190 °C pattern correspond to non-decomposed potassium ferrate(VI).

appeared in the XRD pattern (see Figure 6). However, potassium iron(III) oxide KFeO2 and hexagonal high-temperature K2CO3 were the only phases detected in the XRD patterns recorded at 420 and 590 °C without any indications of K2O. The hexagonal high-temperature K2CO3 structure appeared at the expense of the monoclinic K2CO3 and indicates the thermally induced polymorphous transformation of potassium carbonate. Decomposition Mechanism. In situ high-temperature Mo¨ssbauer and VT XRD techniques suggest that the main decomposition step results in the formation of a mixture of KFeO2, K2CO3, and K2O. The experimental weight loss of 8%, found from the TG curve above 300 °C, may be explained by considering a participation of 1/6 mol of CO2 and liberation of 3/ mol of O per 1 mol of K FeO (∆m 4 2 2 4 theor ) 8.2%) according to the reaction

1 1 1 3 K2FeO4 + CO2 f KFeO2 + K2O + K2CO3 + O2 6 3 6 4 (4) However, such a direct thermally induced reaction of carbon dioxide with potassium ferrate(VI) is highly improbable. One would rather expect the primary formation of the mixture of potassium oxide and superoxide together with KFeO2 followed by the rapid secondary reaction of carbon dioxide with KO2 (eqs 5 and 6):

1 1 1 K2FeO4 f KFeO2 + K2O + KO2 + O2 3 3 2

(5)

1 1 1 1 KO + CO f K CO + O 3 2 6 2 6 2 3 4 2

(6)

This assumption is in agreement with the well-known high affinity of KO2 to CO237 and, mainly, with the presence of slightly overlapping endo- and exo-effects appearing in DSC curve between 250 and 280 °C, which indicate a two-step formation mechanism of potassium carbonate (see Figure 4b). As mentioned above, the secondary exo-effect with a maximum at 270 °C was absent in the DSC curve recorded in an inert atmosphere23 (see Figure 7a). This reflects the principal difference in the decomposition mechanisms if K2FeO4 is heated in static air and inert atmospheres. Concerning a consumption of carbon dioxide during the decomposition of potassium ferrate(VI), it can be calculated from eq 4 that 57 mL of air with a natural content of CO2 is needed for the complete conversion of 10 mg of K2FeO4. Such a requirement was fulfilled in the case of all in situ techniques used in the present study. Equation 5 represents the primary decomposition of Fe(VI) to Fe(III) by three-electron-transfer process, which is unlikely to occur. The possible intermediates, Fe(V) and Fe(IV), can be primarily formed by one- and two-electron-transfer processes. Such high oxidation states of iron have been identified independently both in solid state and in solutions.2,3,19,38-42 However, these ferrate species are quite unstable and may be rapidly transformed to the Fe(III) phase. Thus, they were not detectable even by in situ techniques. However, the possible formation of Fe(IV) and Fe(V) species during the thermal decomposition of K2FeO4 was further explored indirectly by carrying out a detailed zoom of DSC curve of the “critical range” between 240 and 265 °C (Figure 7b). Interestingly, DSC revealed the change of the curve trend at 253 °C, probably because of the overlapped endo-effects, which indicate rather two-step conversion of Fe(VI) to Fe(III) phase. The primary

Thermal Decomposition of K2FeO4 in Static Air

J. Phys. Chem. B, Vol. 111, No. 16, 2007 4285

Figure 7. DSC curve of K2FeO4 measured in nitrogen (a). The detailed zoom of the endothermic effect observed in DSC curve during the main decomposition step (240-265 °C) in static air (b).

step may be one-electron transfer to Fe(V), which then converted to Fe(III) phase with a release of oxygen (eqs 7 and 8).

1 1 K2FeO4 f KFeO3 + K2O + KO2 3 3

(7)

1 KFeO3 f KFeO2 + O2 2

(8)

If conversion occurs through two-electron transfer, Fe(IV) would be formed (eq 9), which gives Fe(III) phase (eq 10). A release of oxygen would then be in the first step (eq 9).

1 K2FeO4 f K2FeO3 + O2 2

(9)

1 1 K2FeO3 f KFeO2 + K2O + KO2 3 3

(10)

Both mechanisms (eqs 7, 8 and 9, 10) give the same final phase composition and the same weight loss as shown in eq 5. Thus, it is very difficult to distinguish experimentally which primary step occurs preferentially. Additionally, there is another possibility that Fe(IV) formed in eq 9 goes through redox process to result in Fe(V) phase and the equimolar mixture of KO2 and K2O (eq 11).

1 1 1 K2FeO3 + O2 f KFeO3 + K2O + KO2 2 3 3

(11)

The KFeO3 formed can be subsequently reduced by a final twoelectron process to yield KFeO2 (eq 8). This three-step process (a combination of eqs 9, 11, and 8) also suggests the same phase composition and weight loss as has been found experimentally (eq 5). At higher temperatures above 300 °C, no other chemical reactions were observed in static air although the primary decomposition products undergo the various phase transitions as demonstrated by VT XRD data. Thus, potassium oxide, clearly identified at 300 °C, is absent in the XRD pattern, recorded at 420 °C, because of its melting. The melting point of K2O is ∼350 °C.43 Similarly, the potassium carbonate changed at the higher temperatures from monoclinic to hexagonal structure because of the thermally induced polymorphous transition as observed in VT XRD patterns.44 Above 750 °C, melting and evaporation of decomposition products occur as documented by a drastic decrease in the sample weight (see TG curve in Figure 4a).

In summary, in situ experimental approach was successful in revealing the principal influence of CO2 in the decomposition mechanism of K2FeO4. Moreover, it also allowed the monitoring of the secondary transformations of intermediates including the chemical transformation of KO2 at the lower heating temperatures, the melting of K2O, and the polymorphous transition of K2CO3 at the higher temperatures. Such steps could not be observed experimentally using a “standard” approach of analysis of the cooled products after the thermal decomposition. Acknowledgment. Financial supports from the Ministry of Education of the Czech Republic (MSM6198959218 and 1M6198959201) are gratefully acknowledged. The authors also thank Miroslav Mashlan and Jiri Tucek (Palacky University in Olomouc, Czech Republic) for low-temperature and in-field Mo¨ssbauer measurements. We wish to thank anonymous reviewer for useful comments. References and Notes (1) Rush, J. D.; Bielski, B. H. J. J. Am. Chem. Soc. 1986, 108, 523. (2) Jeannot, C.; Malaman, B.; Gerardin, R.; Oulladiaf, B. J. Solid State Chem. 2002, 165, 266. (3) Delattre, J. L.; Stacy, A. M.; Young, V. G.; Long, G. J.; Hermann, R.; Grandjean, F. Inorg. Chem. 2002, 41, 2834. (4) Kopelev, N. S.; Perfiliev, Yu. D.; Kiselev, Yu. M. J. Radioanal. Nucl. Chem. 1992, 162, 239. (5) Licht, S.; Wang, B.; Ghosh, S. Science 1999, 28, 1039. (6) Licht, S.; Alwis, C. D. J. Phys. Chem. B 2006, 110, 12394. (7) Waltz, K. A.; Suyama, A. N.; Suyama, W. E.; Sene, J. J.; Zeltner, W. A.; Armacanqui, E. M.; Roszkowski, A. J.; Anderson, M. A. J. Power Sources 2004, 134, 318. (8) Delaude, L.; Laszlo, P. J. Org. Chem. 1996, 61, 6360. (9) Sharma, V. K.; Mishra, S. K.; Nesnas, N. EnViron. Sci. Technol. 2006, 40, 7222. (10) Sharma, V. K; Kazama, F; Jiangyong, H; Ray, A. K. J. Water Health 2005, 3, 45. (11) Sharma, V. K. Water Sci. Technol. 2004, 49, 69. (12) Thompson, G. W.; Ockerman, G. W.; Schreyer, J. M. J. Am. Chem. Soc. 1951, 73, 1279. (13) Schreyer, J. M.; Thompson, G. W.; Ockerman, L. T. Inorg. Synth. 1953, 4164. (14) Bouzek, K.; Schmidt, M. J.; Wragg, A. A. Collect. Czech. Chem. Commun. 2000, 65, 133. (15) Lescuras-Darrou, V.; Lapicque, F.; Valetin, G. J. Appl. Electrochem. 2002, 32, 57-63. (16) Kiselev, Y. M.; Kopelev, N. S.; Zav’yalova, N. A.; Perfiliev, Y. D.; Kazin, P. E. Russ. J. Inorg. Chem. 1989, 34, 1250. (17) Scholder, R.; Bunsen, H. V.; Kindervater, F.; Zeiss, W. Z. Anorg. Allgm. Chem. 1955, 282, 268. (18) Goff, H.; Murmann, R. K. J. Am. Chem. Soc. 1971, 93, 6058. (19) Scholder, R. Anges. Chem. 1962, 1, 220. (20) Ichida, T. Bull. Chem. Soc. Jpn. 1973, 4, 79. (21) Fa¨tu, D.; Schiopescu, A. ReV. Roum. Chim. 1974, 19, 1297.

4286 J. Phys. Chem. B, Vol. 111, No. 16, 2007 (22) Tsapin, A. I.; Goldfeld, M. G.; McDonald, G. D.; Nealson, K. H.; Moskovitz, B.; Solheid, P.; Kemner, K. M.; Kelly, S. D.; Orlandini, K. A. Icarus 2000, 147, 68. (23) Madarasz, J.; Zboril, R.; Homonnay, Z.; Sharma, V. K.; Pokol, G. J. Solid State Chem. 2006, 179, 1426. (24) Audette, R. J.; Quail, J. W. Inorg. Chem. 1972, 11, 1904. (25) Audette, R. J.; Quail, J. W.; Black, W. H.; Robertson, B. E. J. Solid State Chem. 1973, 8, 43. (26) Licht, S.; Naschitz, V.; Halperin, L.; Halperin, N.; Lin, L.; Chen, S.; Ghosh, S.; Liu, B. J. Power Sources 2001, 101, 167. (27) Nowik, I.; Herber, R. H.; Koltypin, M.; Aurbach, D.; Licht, S. J. Phys. Chem. Solids 2005, 66, 1307. (28) Hoppe, M. L.; Schlemper, E. O.; Murmann, R. K. Acta Crystallogr. 1982, B38, 2237. (29) Shinjo, T.; Ichida, T.; Takada, T. J. Phys. Soc. Jpn. 1970, 29 (1), 111. (30) Dedushenko, S. K.; Perfiliev, Y. D.; Goldfeld, M. G.; Tsapin, A. I. Hyperfine Interact. 2001, 136/137, 373. (31) Herber, R. H.; Johnson, S. Inorg. Chem. 1979, 18, 2786. (32) Corson, M. R.; Hoy, G. R. Phys. ReV. B 1984, 29, 3982.

Machala et al. (33) Tucek, J.; Zboril, R.; Petridis, D. J. Nanosci. Nanotechnol. 2006, 6, 926. (34) Pankhurst, Q. A., Pollard, R. J. J. Phys.: Condens. Matter 1990, 2, 7329. (35) Pankhurst, Q. A. Hyperfine Interact. 1994, 90, 201. (36) Maddock, A. G. Mo¨ssbauer Spectroscopy: Principles and Applications of the Techniques; Howard Publishing Limited: Horwood, U.K., 1998; p 108. (37) Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M. AdVanced Inorganic Chemistry, 6th ed.; John Wiley & Sons, Inc: New York, 1999; p 461. (38) Ayers, K. E.; White, N. C. J. Electrochem. Soc. 2005, 152, A467. (39) Sharma, V. K.; O’Connor, D. B.; Cabelli, D. E. J. Phys. Chem. B 2001, 105, 11529. (40) Rush, J. D.; Bielski, B. H. J. Inorg. Chem. 1989, 28, 3947. (41) Hoppe, R.; Mader, K. Z. Anorg. Allg. Chem. 1990, 586, 115. (42) Rush, J. D.; Bielski, B. H. J. J. Am. Chem. Soc. 1988, 110, 4957. (43) Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Butterworth: U.K., 1997. (44) Becht, H. Y.; Struikmans, R. Acta Crystallogr. 1976, B32, 3344.