Synthesis, structure, and Mössbauer spectroscopic studies on the heat-induced solid-phase redox reactions of hexakis(urea-O)iron(III) peroxodisulfate

Anhydrous hexakis(urea-O)iron(III)]peroxydisulfate ([Fe(urea-O)6]2(S2O8)3 (compound 1), and its deuterated form were prepared and characterized with single-crystal X-ray diffraction and spectroscopic (IR, Raman, UV, and Mössbauer) methods. Six crystallographically different urea ligands coordinate via their oxygen in a propeller-like arrangement to iron(III) forming a distorted octahedral complex cation. The octahedral arrangement of the complex cation and its packing with two crystallographically different persulfate anions is stabilized by extended intramolecular (N–H⋯O = C) and intermolecular (N–H⋯O–S) hydrogen bonds. The two types of peroxydisulfate anions form different kinds and numbers of hydrogen bonds with the neighboring [hexakis(urea-O)6iron(III)]3+ cations. There are spectroscopically six kinds of urea and three kinds (2 + 1) of persulfate ions in compound 1, thus to distinguish the overlapping bands belonging to internal and external vibrational modes, deuteration of compound 1 and low-temperature Raman measurements were also carried out, and the bands belonging to the vibrational modes of urea and persulfate ions have been assigned. The thermal decomposition of compound 1 was followed by TG-MS and DSC methods in oxidative and inert atmospheres as well. The decomposition starts at 130 °C in inert atmosphere with oxidation of a small part of urea (~ 1 molecule), which supports the heat demand of the transformation of the remaining urea into ammonia and biuret/isocyanate. The next step of decomposition is the oxidation of ammonia into N2 along with the formation of SO2 (from sulfite). The main solid product proved to be (NH4)3Fe(SO4)3 in air. In inert atmosphere, some iron(II) compound also formed. The thermal decomposition of (NH4)3Fe(SO4)3 via NH4Fe(SO4)2 formation resulted in α-Fe2O3. The decomposition pathway of NH4Fe(SO4)2, however, depends on the experimental conditions. NH4Fe(SO4)2 transforms into Fe2(SO4)3, N2, H2O, and SO2 at 400 °C, thus the precursor of α-Fe2O3 is Fe2(SO4)3. Above 400 °C (at isotherm heating), however, the reduction of iron(III) centers was also observed. FeSO4 formed in 27 and 75% at 420 and 490 °C, respectively. FeSO4 also turns into α-Fe2O3 and SO2 on further heating.


Introduction
Heat-induced solid-phase quasi-intramolecular redox reactions of [ML n ](XO 4 ) m complexes (M = Cu, Zn, Cd, Fe, Co, Cr) containing reducing ligands (NH 3 , pyridine, or urea. n = 2-6) and oxidizing anions (X = Cl, Mn, Mo, Re, S; m = 0.5-3) ensure an easy and convenient way to prepare simple or mixed transition metal oxides with particle size in the nanometer region [1][2][3][4][5][6][7][8][9][10][11][12][13][14][15][16][17]. These nanosized oxides are important materials in various industrial processes, especially the transition metal oxides with perovskite (ABO 3 ) or spinel (AB 2 O 4 ) structures (A, B = Mn, Fe, Co, Ni, and Cu) that can replace the noble-metal-based redox catalyst [18]. For example, the cobalt-and copper-manganese mixed spinel oxides are widely studied and used in various processes like nitrate-to-ammonia conversion [19], oxidation of ammonia, propane, and CO [20][21][22][23] or reduction of NO [24,25]. Co-Mn or Cu-Mn mixed spinel oxides are active catalysts in the oxidation processes of ciprofloxacin or other organic pollutants in wastewater, in the activation of peroxymonosulfate, a promising technology for water treatment [26,27], or in the degradation of pharmaceuticals [28]. In the reduction reactions of oxygen, it is possible to use not only Co or Cu but also other AMn 2 O4 (A = Mn(II), Mg(II), Li(I), and Cd(II)) mixed oxides [29]. In our previous work, we dealt with preparation of stoichiometric AgMnO 2 [7] from [Ag(NH 3 ) 2 ]MnO 4 and from silver manganese oxides with various Ag-Mn ratio via mixed [Ag(NH 3 ) 2 ](ClO 4 ,MnO 4 ) solid solutions [7]. Besides Ag + , it is possible to dope the manganese oxides with Fe 3+ (or Zn 2+ , Cu 2+ ) and this can improve their photocatalytic efficiency [30]. These help us contextualize our previous research about [Fe(urea-O) 6 ] (MnO 4 ) 3 from which we prepared (Fe,Mn)O x [1]. These (Fe,Mn) O x phases contain oxygen vacancies which are important in a catalyst structure since it can help reduce the activation energy of the given process [31] and help the catalytic oxidation of NO with Fe-Mn bimetal oxides [32]. Fe-Mn mixed oxides are widely used (just like Co-Mn mixed spinel oxides) in important technological processes like reduction of nitroaromatic compounds [33] or CO 2 [1], gasification of biomass [34], catalytic oxidation of ethylene oxide [35] or formaldehyde [36], and transforming carbon dioxide into jet fuel [37]. Preparation and decomposition of [hexakis(urea-O)iron(III)] peroxydisulfate (compound 1) fit to this systematic work, because the sulfur content of the oxidative peroxydisulfate can be eliminated as SO 2 , and studies on this compound can serve as a strong basis to prepare various solid solutions of peroxydisulfate with metal containing oxygen bridged anions, e.g., dichromate that facilitates low-temperature synthesis of nanosized Fe,Cr-oxides with adjusted Fe/Cr ratio.

Results and discussion Preparation and properties
[Hexakis(urea-O)iron(III)] peroxydisulfate (compound 1) was first prepared by Barbieri in the metathesis reaction of an aq.
[hexakis(urea-O)iron(III)] chloride and a saturated aq. ammonium persulfate solution as a blueish-green crystalline mass, but neither the yield nor the properties of compound 1 were given [41]. A similar reaction of in situ prepared hexakis(urea-O) iron(III) nitrate (from iron(III) nitrate and six equivalents of urea) and sodium persulfate at room temperature resulted in light-blue blocks during a week with 65% yield. Compound 1 is anhydrous, easily soluble in water (57.5 g/100 mL at 25 °C), but insoluble in organic solvents such as benzene, toluene, chloroform, and carbon tetrachloride. Its saturated aqueous solution is yellow and strongly acidic (pH 2.02). Its pycnometric density is 1.94 g cm −3 at 25 °C. No polymorph phase transition of compound 1 was found between 77 K and its decomposition temperature (396 K) (ESI Fig. S1). Its powder XRD diffractogram is given in ESI Fig. S2.
The Fe(urea) 6 3+ complex cation has an octahedral arrangement of urea ligands, the rotation direction of the urea propellers is different at the two sides of the complex [ Fig. 1   Two different peroxydisulfate anions (labeled as anions B and C) could be found in the asymmetric unit of the crystal structure of compound 1 with different hydrogen bond patterns. The hydrogen bond network formed by anion "C" is symmetrical, that is, the two halves of the anion form the same number of hydrogen bonds with the same number of cations [ Fig. 3(d)]. In contrast, H−bond interactions of anion "B" are arranged asymmetrically [ Fig. 3 Fig. 3) form different number of hydrogen bonds Ligand 1 forms 5 hydrogen bonds with 3 "B" anions: N1B is connected to the terminal oxygens of two different anions with 1 and 2 hydrogen bonds, while N1A forms 1 hydrogen bond with a terminal oxygen atom and 1 hydrogen bond with a bridging peroxy oxygen atom of a different anion. Ligand 2 forms a total of 6 hydrogen bonds with 4 "B" and 1 "C" anions: N2A forms 2 hydrogen bonds with 2 different "B" ligands, while N2B forms 4 hydrogen bonds to terminal oxygens of 2 "B" and a terminal and a peroxy oxygen of 1 "C" anions. Ligand 3 forms the smallest number of hydrogen bonds: it makes only 4 hydrogen bonds with terminal oxygens of 2 "B" and 1 "C" anions. Ligand 4 and 5 both make 5-5 hydrogen bonds with 2 "B" and 1 "C" anions but in a different arrangement. In the case of ligand 4, N4A makes 4 hydrogen bonds with 2 "B" and 1 "C" anions, while N4B forms 1 hydrogen bond with a "B" anion. For ligand 5, N5A forms 1 hydrogen bond with a "C" anion, while N5B forms 4 hydrogen bonds with 2 "B" anions and the same "C" anion that interacted with N5A. Ligand 6 makes 6 hydrogen bonds with 1 "B" and 1 "C" anions: N6A forms 3 hydrogen bonds with terminal oxygens of 1 "B" and 1 "C" anions, while N6B makes 3 hydrogen bonds with terminal and peroxy oxygens of a single "C" anion. The packing of the molecules is presented in Fig. 4  The shortest distances between the iron ions in the structure are 6.433 Å and 6.532 Å connecting neighboring complexes along unit cell axis a. The iron atoms along the other two axes have greater separation by the urea molecules and the channels between them are filled by peroxydisulfate ions. Their distances are 9.9-11.8 Å (ESI Fig. S4).

Vibrational spectra of compound 1
There are 3 crystallographically different persulfate anions, 2 of which are at positions of trivial symmetry and there are 6 crystallographically different urea ligands. Altogether 74 atoms build the unit cell. Since there are 3 × 12 and 6 × 27 internal normal modes in 3 sets of persulfate and 7 sets of urea bands, respectively, are expected, together with external modes and overtone/combination bands, it is not worth doing correlation analysis. Instead of that, deuteration experiments were done on compound 1 to separate the heavily overlapping modes belonging to the urea NH 2 other groups. The IR and Raman spectra of compound 1 and compound 1-D were recorded at room temperature in the range of 4000-100 cm −1 ( Fig. 5 and ESI Fig. S5). Due to the deuteration, the modes that belong to the N-D bonds shifted to lower wavenumbers as compared to the modes of N-H bonds (υ(NH)/υ(ND)≈1.35, ESI Fig. S6). Substantial shifts in the positions of the persulfate normal modes were also observed due to differences in the strength of O…H and O…D interactions (Table 1, Fig. 5). Low-temperature (263, 193 and 93 K) Raman spectra were also recorded (ESI Fig. S7). The assignments of the vibrational modes of the anionic (Table 1) and the cationic (Table 1 and ESI Table S9) parts of compound 1 are given. Out of the 24 normal modes of the peroxydisulfate ion [44,45] modes [44][45][46][47]. The presence of two kinds of peroxydisulfate ion and the hydrogen bond induced asymmetry of the "B" peroxydisulfate ion results in more complicated experimental IR and Raman spectra than that of alkali peroxydisulfates [45] (ESI Fig. S8). The IR and Raman bands located around 400 cm −1 were assigned as the δ SOOS modes by Cleaver [45] and as the terminal δ SO3 mode by Skogareva [48]. Comparing the available IR and Raman data of the alkali, alkali metal and some [M(NH 3 ) 4 ] 2+ (M = Cu, Zn) complex peroxydisulfates [45,49], increasing the polarizing effect of cation (decreasing the ionic size or increasing the charge) increases the wavenumber of δ(SOOS) and υ(SOOS) modes. Due to the presence of the trivalent iron cation in compound 1, the bands at 409 and 391 cm −1 might be assigned to δ SOOS mode of the two different peroxydisulfate ions. The size of the complex cation, however, increased due to the presence of six urea ligands, resulting in a strong decrease in the polarizing effect of the complex cation. Comparison of the υ as (SOOS) IR wavenumbers found in the IR spectra of alkali, alkaline earth metal and divalent metal (Cu, Zn) ammine complex peroxydisulfates, and compound 1, it is obvious that the complex cations ([M(NH 3 ) 4 ] 2+ (M = Zn, Cu, (υ as (SOOS) = 1048 cm −1 ) and [Fe(urea) 6 ] 3+ (1045 an 1035 cm −1 ) have less polarizing effects than the less polarizing alkali (cesium) cation (υ as (SOOS) = 1052 cm −1 ). Based on this relationship, the bands at 391 and 409 cm −1 in the IR spectrum of compound 1 belong to the terminal S-O bonds and not to the deformation mode of SOOS linkage. The peroxy-bond is not IR active in peroxydisulfates [44]. However, due to the strong distortion (e.g., the antisymmetric hydrogen bond in the "B" peroxydisulfate anion) of symmetries, the bands appearing at 796 cm −1 and 830 cm −1 in the IR and at 810 and 823 cm −1 in the Raman spectra of compound 1 might be attributed both to the υ 2 (O-O) [44] and to υ s (SOOS) modes. The symmetric and antisymmetric stretching and deformation modes of the terminal SO 3 group were also assigned ( Table 1). There  was no pyrosulfate ion (S 2 O 7 2− ) as a decomposition product detected.
The vibrational modes assigned to the urea ligands and iron-oxygen interactions are given in ESI Table S9. The coordination of urea through its oxygen atom influences on the C=O and the C-N bond lengths. The C=O bond length elongates, whereas the C-N bond length shortens compared to the gaseous urea. Therefore, the corresponding ν CO and ν CN stretching modes shift to lower and higher wavenumbers, respectively (ESI Table S9). Due to the hydrogen bonds formed by the -NH 2 groups, the N-H bonds become longer; thus, the NH 2 stretching vibrations will shift toward lower wavenumbers as was observed in the spectra of solid urea [49] and hexakis(urea-O)iron(III) permanganate [1]. The IR spectrum of compound 1 contains two high-and two lowintensity bands in the range above 3000 cm −1 . The overlapping bands were separated via deuteration of compound 1, by comparing the IR spectra of the deuterated and non-deuterated samples. Furthermore, the combination/overtone bands could also be used to determine some original band positions as well. Following the method given in [50], together with the deuteration experiments, ν(C=O) components in the complex band systems were located at 1504 cm −1 and 1503 cm −1 in the IR and Raman spectra, respectively. A wide band appears at 1625 cm −1 in the IR spectrum of compound 1 and splits into two well-defined peaks (1693 cm −1 and 1620 cm −1 ) in the Raman spectra. The δ s (N-H) band in the Raman spectra could be assigned as sharp peaks at 1692 cm −1 and at 1693 cm −1 at room temperature and at 100 K, respectively. The δ s (N-H) position in the IR spectrum was calculated to be 1685 cm −1 . The δ as (N-H) mode is at 1625 cm −1 and at 1620 cm −1 in the IR and Raman spectra, respectively.

UV spectroscopy
Compound 1 has three well-defined absorption bands in its solid-phase UV-Vis spectrum (ESI Fig. S9). The UV spectra of compound 1, in principle, are built up from the increments of Fe III crystal field bands, the bands belonging to the urea ligand and to the persulfate anion.
Compound 1 contains a high-spin Fe III state with a 6S ground term. The 6S term is not split by any crystal field, and the electronic transitions are spin and Laporte forbidden. Thus, among the expected 4T 1g (G) ← 6S, 4T 2g (G) ← 6S, and the 4A g , 6 ] 3+ cation, the first does not appear, the second one appears at 590 nm as a very weak and wide band, and the third transition was found at 432 nm. The 4A g ,4 E g (G) ← 6S transition is degenerate in octahedral symmetry.
The urea n-π* and π − π* bands in the UV spectrum of compound 1 show up as a wide band in the UV range centered at 212 and 228 nm, due to various strength and bond lengths of N-H and C=O linkages, including possible Fe III -O=C LMCT bands. The uncoordinated urea has bands in the UV range at 211, 230-260, 299, 337, 361, and 380 nm due to various tautomer and hydrogen-bonded clusters present in solid urea [50]. The absorption belonging to the excitation of the peroxide bond in the persulfate anion [51] is located around 265 nm, near to the band found in the UV spectrum of sodium persulfate at 260 nm (ESI Fig. S10a, b).

Room-temperature and low-temperature (liq. N2) Mössbauer spectrum of compound 1
The Mössbauer measurements of compound 1 were done both at room temperature and at liquid nitrogen temperature. At both temperatures, the Mössbauer spectra show only broad singletlike envelopes attributable to only one iron(III) site with magnetic relaxation (Fig. 6). This line shape is quite common for this type of complexes [1,[52][53][54]. The spectra were evaluated by the Blume-Tjon two-state magnetic relaxation model (B-T model) provided by the Mosswinn code [55]. The Mössbauer parameters calculated by this model are as follows: the isomer shift (δ), the amplitude of the relaxing hyperfine magnetic field (H), the line width (Γ), the component V zz of the electric field gradient (EFG), the asymmetry parameter of the EFG (ETA) (these latter two are not calculated in a Lorentzian fit, only the quadrupole splitting), and, finally, the most specific Mössbauer parameter of this model is the jump up the rate of the relaxing magnetic field (forced to be equal to the jump down rate because of the two-state relaxation). This parameter is the base ten logarithm of the jump-up frequency (W), from which the spin relaxation time may be calculated as π /W. At the first attempt, we used B-T model for the evaluation of our experimental Mössbauer data without any fixed values. In this case, a very broad distribution of fitted parameters was the result of the evaluation. In some cases, even unphysical parameters were obtained. We had a similar experience in the case of hexakis(ureao)-iron(III)-permanganate complex that we have discussed in ref. [1]. At the second attempt, we used the B-T model again, but now the previously published ETA = 0.232 and V zz = -0.762 × 10 21 V m −2 parameters obtained by density functional theory calculations for the complex cation only (ESI Table S10) were used as fixed values. The Mössbauer parameters that we obtained in this way (ESI Table S11) met the Mössbauer parameters expected for other highspin iron(III) urea complexes [1,[52][53][54].
The experimental Mössbauer spectra were recorded at 295 K and liquid N 2 (80 K) temperatures. The isomer shifts obtained were 0.416 mm s −1 and 0.542 mm s −1 , respectively. This confirms that compound 1 is a high-spin iron(III) complex. The isomer shifts are similar to those we have found for hexakis(urea-o)-iron(III)-permanganate (0.412 mm s −1 and 0.501 mm s −1 ) [1] and are smaller than those for the previously published hexakis(urea-o)-iron(III)-chloride (0.58 mm s −1 and 0.60 mm s −1 ) [53]. Since single-crystal X-ray diffraction proved the presence of hydrogen bonds between [Fe(urea)6] 3+ cations and (S 2 O 8 ) 2− anions, the isomer shifts of compound 1 are in good agreement with our previously published finding [1]. We can conclude that the hydrogen bonds reduce the 3d electrons density on the iron(III) central ion and thereby the isomer shift decreases. The line widths are 0.608 mm s −1 and 0.966 mm s −1 at 295 K and 80 K, respectively, and it is similar to the [Fe(urea) 6 ] (MnO 4 ) 3 line widths (0.626 mm s −1 and 0.875 mm s −1 ) which is consistent with the spin relaxation effect [56]. The shortest distance between the Fe-Fe atoms is rather high, 6.433 Å, which supports this type of relaxation.

Solid-phase decomposition reactions of compound 1
The thermal decomposition of compound 1 was followed by TG-MS and DSC methods in inert (argon or N 2 ) and oxidative (air) atmospheres. The overall weight losses until 800 °C were found to be 75% and 85% in the inert and in the oxidative atmosphere, respectively (ESI Figs. S11-S14). The final product of the decomposition was α-Fe 2 O 3 in every case as was confirmed by powder XRD and Mössbauer spectroscopy as well (ESI Figs. S15, S16).
The thermal decomposition is a multistep process with wellseparated steps and starts at around 130 °C with exothermic character in both atmospheres. Thus, the aerial oxygen does not play a role in the initiation of the decomposition reaction (ESI Figs. S11-S14). The peak temperatures of the first decomposition steps are 151 °C and 159 °C(DTG) with ∆H r = 386.45 and 375.93 kJ/mol (DSC peak temperatures are 154 and 155 °C) in inert and oxidative atmospheres, respectively (ESI Figs. S13, S14). The thermal decomposition reaction of compound 1 has a complicated character. First, instead of an endothermic ligand loss, an exothermic process takes place as was observed in the case of [Fe(urea) 6 ](MnO 4 ) 3 [1], where the permanganate ions oxidized the urea in the first stage of the decomposition reaction. The TG-MS study of the decomposition of compound 1 confirmed the presence of such urea oxidation products as H 2 O (m/z = 18), CO 2 (m/z = 44), and N 2 (m/z = 28). NO (m/z = 30) and SO 2 (m/z = 64) as a reduction product of the persulfate ion could also be detected in small amounts (Fig. 7). There was no gaseous ammonia found (the m/z = 17 peak found belongs to the OH fragment of water based on the m/z = 18 and m/z = 17 peak intensity ratio [ Fig. 7(a)], which agrees well with the value for H 2 O → OH fragmentation [57].
In order to study the phase relations in the decomposition intermediates formed in the thermal decomposition step, isotherm heatings of compound 1 at some particular temperatures  ESI  Table S12). Thus, there was no redox reaction occurring between the iron(III) center and the urea ligand. Due to the lack of oxygen, the only oxidizing agent could only be the persulfate ion of compound 1. In accordance with this, the bands belonging to the persulfate ion completely disappeared from the IR and Raman spectra of I-160-N 2 [ Fig. 9(a) and ESI Figs. S17, S18]. The IR and Raman results of the intermediate I-160-N 2 confirmed the presence of sulfate ion (ν 3 (F 2 ) = 1104 cm −1 , ν 1 (A 1 ) = 986 cm −1 (strongest Raman peak of SO 4 2− ), ν 4 (F 2 ) = 597 cm −1 , and ν 2 (E) = 466 cm −1 [58]) as a persulfate reduction product. Surprisingly, bands belonging to ammonium ions appeared in the IR and Raman spectra of I-160-N 2 [ν 1 , ν 2 , ν 3 , and ν 4 = 3201, 3077, 2834, 1419 cm −1 Fig. 9(a)] (and in the IR spectra of I-160-O 2 too (ESI Fig. S18a)). The ammonia source during the decomposition of compound 1 is the direct decomposition or hydrolysis, shown by equations E1, E2, E3, (water forms as oxidation products from urea and persulfate detected by TG-MS) as was observed in the case of [Fe(urea) 6 ](MnO 4 ) 3 [1] (E1) 2 H 2 NCONH 2 = H 2 NCONHCONH 2 + NH 3 The presence of biuret and isocyanate could be detected with IR [ Fig. 9(a) at 160 °C], where the ν(NH) bands and mixed amide ν(CONH) bands do not change much on heating up to 300 °C. An IR band at 2213 cm −1 belongs to an intermediate with carbon-nitrogen multiple bond (e.g., coordinated or noncoordinated isocyanate) and appears as a weak band in the IR spectrum obtained for sample I-160-N 2 . The formation of isocyanate intermediates was detected in the oxidative degradation of hexakis(urea-O)iron(III) nitrate as well [59].
The main iron-containing phase of the decomposition intermediates formed at 160 °C either in air or in inert atmosphere was assigned to pyracmonite, (NH 4 ) 3 Fe(SO 4 ) 3 [60] [ Fig. 9(b) and ESI Fig. S19]. The Mössbauer spectrum obtained for sample I-160-N 2 shows that ca. 73% of the iron-containing products is (NH 4 ) 3 Fe(SO 4 ) 3 , whereas the remaining 27% of iron is in the trivalent state located in octahedral (71%) and tetrahedral (29%) coordination environment [ Fig. 8(b) Fig. 8(a) and ESI Table S13] in the same amount and ratio. Thus, external oxygen does not play any role in this process.
The formation of (NH 4 ) 3 Fe(SO 4 ) 3 in huge amounts shows that the persulfate ion is mainly reduced into sulfate ion, and the urea ligand is decomposed mainly into biuret and ammonia (without redox reaction). The persulfate-driven oxidative urea degradation with the formation of water, CO 2 , and SO 2 also occurs, as was detected in the first decomposition step by TG-MS [ Fig. 7(c), (d)]. This exothermic degradation reaction can serve as a source of activation energy for the decomposition of urea with NH 3 elimination; however, the majority of urea decomposes without a redox reaction; thus, the persulfate has to be consumed by another reducing agent. The large amount of (NH 4 ) 3 Fe(SO 4 ) 3 formed (71%) (ESI Tables S12, S13) shows that the main redox reaction does not involve oxygen transfer from the persulfate ion (no SO 2 formation), because the Fe-to-S ratio in (NH 4 ) 3 Fe(SO 4 ) 3 is the same as that in compound 1 (2:6); thus the huge SO 2 loss would prevent the formation of (NH 4 ) 3 Fe(SO 4 ) 3 in such a large amount (71%) as we found. The (NH 4 ) 3 Fe(SO 4 ) 3 formation cannot be 100%, of course, because a fraction of the urea ligands is oxidized by persulfate consuming actually oxygen from the persulfate and resulting in the formation of H 2 O, CO 2 , and SO 2 as detected by TG-MS [ Fig. 7(a), (c), (d)]. The relatively low mass loss found (9.2 and 14.6% in air and argon, respectively) confirms that the complete urea degradation may only be a minor side-reaction (theoretical mass loss related to one urea is 9.9%) and the reduction of persulfate results mainly in sulfate ions.
The formation of pyracmonite with NH 4 + to Fe III to SO 4 2− stoichiometry of 3:1:3 can be explained if we assume that the main persulfate-driven redox reaction is the oxidation of ammonia and not directly of urea.
Since two iron(III) belong to three persulfate ions in compound 1, the ammonium and sulfate stoichiometry for pyracmonite can be provided by the ammonium sulfate formed in reaction E4. If one assumes an oxygen transfer reaction, e.g., then the amount of sulfate formed decreases and could not fulfill the stoichiometry needs for the formation of two iron(III) since the Fe III -to-SO 4 2− ratio would only be 2:3 at maximum.
The main decomposition route of (NH 4 ) 3 Fe(SO 4 ) 3 is The primary thermal decomposition product of (NH 4 ) 3 Fe(SO 4 ) 3 was identified as NH 4 Fe(SO 4 ) 2 by XRD [ Fig. 9(b)] and confirmed by the Mössbauer spectrum as the only product in air [ Fig. 8(a)]. In the samples made in inert atmosphere, three different iron coordination environments could be found: the main component NH 4 Fe(SO 4 ) 2 (69%), and two iron environments belonging to transition states, one contains iron(III) (17%), whereas the other iron(II)-containing species has an environment similar to that in (NH 4 ) 2 Fe(SO 4 ) 2 (14%) [ Fig. 8(b) and ESI Table S12].
The amount of biuret-like condensation products decreases (ν(C=O = 1669 cm −1 ), and the most important change in the IR and Raman spectra is shifting the ν s (S-O) stretching mode from ν 1 = 986 cm −1 and 996 cm-1 (IR) to 1036 and 1024 cm −1 , respectively [ Fig. 9(a)]. It may be attributed to the change in the sulfate coordination environment (formation of a polymer-like sulfate-bridged structure) [47].
On further heating, NH 4 Fe(SO 4 ) 2 decomposes [69] in air in a redox reaction with the formation of water ad SO 2 around 400 °C according to equation E6 [66]: Both H 2 O an SO 2 could be detected by TG-MS [ Fig. 7(a), (d)], the reaction product at 565 °C is solely Fe 2 (SO 4 ) 3 (ESI Fig. S19). Further decomposition of iron(III) sulfate into iron(III) oxide with SO 2 evolution above 500 °C [ Fig. 7(d)] can also be detected [1,67,68]. By the end of the heat treatment in air, the final product at 800 °C is hematite based on the powder XRD and Mössbauer spectroscopy (ESI Fig. S15). In the case of I-640-O 2 , the sample contains 6% of wüstite. This phase could form during the decomposition of the Fe 2 (SO 4 ) 3 . The release of the SO 2 results in the formation of vacancies in the microenvironment of iron that can produce wüstite-like parameters for this metastable phase.
The sample made in an inert atmosphere at 420 °C consists of unreacted NH 4 Fe(SO 4 ) 2 and Fe 2 (SO 4 ) 3 as decomposition product, and Mössbauer spectroscopy shows the presence of an iron(II) compounds as well (58%) [Figs. 8(b), 9(b) and ESI Table S12]. The low-temperature Mössbauer measurement showed that there are two iron(II) signals with the same quadrupole splitting but different isomer shifts (ESI Fig. S20 and ESI Table S14). These two phases could be assigned to iron coordination environments found in (NH 4 ) 2 Fe(SO 4 ) 2 and FeSO 4 . At room temperature, the Mössbauer parameters of the (NH 4 ) 2 Fe(SO 4 ) 2 and FeSO 4 cannot be distinguished, but at 86 K, they can easily be resolved. The appearance of the second doublet only at 86 K is due to the different temperature dependence of the Mössbauer parameters of the different iron(II) microenvironments.
XRD of the sample I-420-N 2 contains the peaks belonging to NH 4 Fe(SO 4 ) 2 , Fe 2 (SO 4 ) 3 , and FeSO 4 , and heating until 490 °C resulted in decreasing peak intensities of NH 4 Fe(SO 4 ) 2 and increasing peak intensities of FeSO 4 [ Fig. 9(b)]. It means that there is a substantial reduction rate driven by the ammonium ion of NH 4 Fe(SO 4 ) 2 and the parent of FeSO 4 is the NH 4 Fe(SO 4 ) 2 itself. The transformation of NH 4 Fe(SO 4 ) 2 into Fe 2 (SO 4 ) 3 around 400 °C is a well-known process [63], but neither FeSO 4 nor (NH 4 ) 2 Fe(SO 4 ) 2 was found. Our experiments were done at 420 °C. Therefore, we checked the effect of temperature on the quasi-intramolecular redox reaction of NH 4 Fe(SO 4 ) 2 . The Mössbauer data showed the increase of the iron(II) phase content up to 75% [ Fig. 8(b) and ESI Table S12] on heating the sample I-420-N 2 until 490 °C in inert atmosphere. The iron(II) compound is FeSO 4 , according to the powder XRD results [ Fig. 9(b)]. Its formation is summarized in equation E7.
Two peaks in the TG-MS belonging to SO 2 (m/z = 64) appear around 540 °C and 590 °C, and this is consistent with the DTG peaks found at 526 °C and 579 °C [ Fig. 7(d)]. The first peak belongs to the redox reaction of NH 4 Fe(SO 4 ) 2 leading to FeSO 4 formation, whereas the second peak belongs to the SO 2 evolution from Fe 2 (SO 4 ) 3 and FeSO 4 with the formation of α-Fe 2 O 3 (E8) [65][66][67][68].
The iron(III) content consists of two phases. Regular hematite and defect hematite (supposedly with oxygen vacancies) phase formed through the decomposition of Fe 2 (SO 4 ) 3 [64][65][66][67]. Further increasing the temperature until 590 °C in inert atmosphere, the hematite content increased to 95% and the amount of the defect phase decreased to 5% [ Fig. 8(b) and ESI Table S12]. In the powder XRD of the sample, a shoulder can be seen on a few reflections that can indicate the defect phase [ Fig. 9(b)].
Both I-590-N2 and I-640-O 2 contain small amount of FeOOH, 5 and 12%, respectively ( Fig. 8 and ESI Tables S12, 13). This relatively small amount shows that this phase may form only at the surface of the samples during the removal from the furnace due to the humidity of air. The thermal studies of compound 1 showed that its thermal degradation reaction is initiated by oxidation of ~ 1 molecule of urea with persulfate ion, together with the releasing of ammonia from the remaining urea molecules while biuret and isocyanate form. The remaining persulfate ions are reduced into sulfate by the liberated ammonia through SO 2 and ammonium ion formation. The solid phase contains mainly (NH 4 ) 3 Fe(SO 4 ) 3 ,

Conclusion
which decomposes stepwise into NH 4 Fe(SO 4 ) 2 , Fe 2 (SO 4 ) 3 , and α-Fe 2 O 3 upon increasing the temperature. In inert atmosphere, some amount of iron(II) compound formed even at 220 °C. The thermal decomposition of the NH 4 Fe(SO 4 ) 2 intermediate also depends on the atmosphere and the temperature. Basically, Fe 2 (SO 4 ) 3 formed together with N 2 and SO 2 , but at 420 °C and 490 °C, in inert atmosphere, FeSO 4 also formed in an amount of 27 and 75%, respectively. Finally, both FeSO 4 and Fe 2 (SO 4 ) 3 decomposed into α-Fe 2 O 3 and SO 2 , but the temperature of α-Fe 2 O 3 formation was not the same, 590 and 800 °C, in inert atmosphere and in air, respectively.

Synthesis of compound 1
Iron(III) nitrate nonahydrate (8.08 g, 0.02 mol) and urea 7.20 g, (0.12 mol) were dissolved in 9 mL of water, and then, the orangecolored solution obtained was mixed with an aq. solution of Na 2 S 2 O 8 (7.86 g Na 2 S 2 O 8 in15 mL of water) under stirring. The resulted dark orange-colored solution was left to crystallize at room temperature for a week when light-blue block single crystals were formed. The solid crystalline mass was separated by filtration on a G3 glass filter, washed with a copious amount of cold (0 °C) water, and dried with abs. ethanol and then with diethyl ether. The yield was 65% (18.27 g).
X-ray powder diffraction patterns were recorded by using a Bragg-Brentano parafocusing goniometer manufactured by Philips (PW-1050) equipped with a Cu anode (40 kV, 35 mA tube power), a secondary beam graphite monochromator, and a proportional counter. Every scan was recorded in a step mode and the diffraction patterns were evaluated by full-profile fitting techniques.
A light-blue, block-like single crystal of 0.3 × 0.2 × 0.1 mm was selected for single -crystal X-ray diffraction measurement. The crystal was mounted on a loop and measured on an XtaLAB Synergy R diffractometer equipped with PhotonJet-R rotating anode source [Cu-K-alpha radiation (λ = 1.54184 Å)], confocal mirrors as a monochromator, and Hypix-6000HE detector. Data collection and data reduction were carried out using CrysAlisPro v.1.171.40.68a program [70]. Data were collected at 100 K using 104 ω scans. A total of 17,580 frames were collected (0.5° rotation, 0.05 s exposure time/frame). The structure was solved using direct methods as implemented in the SHELXS program [71]. Refinement was carried out using SHELXL [71] and Olex2 [72] with full matrix least-squares method on F 2 . All non-hydrogen atoms were refined anisotropically. Hydrogen atoms were generated based on geometric evidence and their positions were refined by the riding model. Hydrogen atoms were found in difference Fourier maps. Atomic displacement parameters of the S 2 O 8 2− ions were refined with rigid bonds restraints applied.
Olex2 [72] and Mercury [73] were used for molecular graphics and analyzing crystal packing. Crystal data and details of the structure determination and refinement are listed in ESI Table S1, and selected parameters of the structure are summarized in ESI Tables S2-S8. The X-ray crystallographic data have been deposited in the Cambridge Crystallographic Data Centre with CCDC ID 2,192,412 and can be obtained free of charge from the CCDC via www. ccdc. cam. ac. uk/ getst ructu res The analytical range FT-IR and far-IR spectra of compound 1 were recorded in an attenuated total reflection (ATR) mode at room temperature using a BioRad-Digilab FTS-30-FIR and a Bruker Alpha IR spectrometer for the 4000-400 and 400-100 cm −1 range, respectively.
The Raman measurements of compound 1 were carried out at 93 K, 193 K, 263 K, and 298 K on a Horiba Jobin-Yvon LabRAM microspectrometer. External (532 nm and 785 nm) Nd-laser sources operated at ~ 40 mW and an Olympus BX-40 optical with a temperature-controllable microscope stage (Linkam THMS600) were used in the 4000-100 cm -1 spectral range with 3 cm −1 resolution. The laser beam (50 × objective) was focused and a D0 intensity filter was applied. A confocal hole of 1000 µm and 1800 groove mm −1 grating monochromator was used for light dispersion resolution. The exposure time was 15 s.
The solid-phase UV-Vis (diffuse reflectance) spectrum of compound 1 was measured at room temperature (Jasco V-670 UV-Vis spectrophotometer, equipped with a NV-470 integrating sphere, BaSO 4 standard). 57 Fe Mössbauer spectroscopy measurements were performed at room temperature and at T = 80 K with a conventional WissEl Mössbauer spectrometer (Starnberg, Germany) operating in constant acceleration mode ( 57 Co source in Rh matrix). Low-temperature measurements were done using a SVT-400-MOSS cryostat (Janis, Woburn, MA, USA) filled with liquid nitrogen. The powdered samples were mixed with polyethylene powder to ensure a perfectly random orientation of the crystallites. The Mössbauer spectra recorded were evaluated by standard computer-based statistical analysis methods, including fitting the experimental data by a sum of Lorentzians or a magnetic relaxation model using a least-squares minimization procedure (MossWinn 4.0 program [55]). The isomer shifts are given relative to α-Fe standard at room temperature. The TG-MS measurements were carried out in both argon and air by a simultaneous thermal analyzer SDT Q600 from TA Instruments, online coupled with an HPR20/QIC mass spectrometer from Hiden Analytical. The sample holder was an alumina crucible, and an empty alumina crucible was used as reference. The sample mass was 1-3 mg. The carrier gas flow rate was 50 mL min −1 . The decomposition was followed from room temperature to 650 °C by a heating rate of 10 °C min −1 . During evolved gas analysis, ions selected between m/z = 1-98 were monitored in multiple ion detection modes (MIDs).
The non-isothermal DSC curve of compound 1 was recorded up to 400 °C with a Perkin Elmer DSC 7 apparatus. Another measurement was done between − 130 °C and room temperature to check polymorphic phase transitions. Sample masses varied between 3 and 6 mg. The heating rate was selected to be 5 °C/min in both cases under a continuous nitrogen flow (20 cm 3 min −1 ). The aluminum pans used as sample holders were unsealed.

Funding
Open access funding provided by ELKH Research Centre for Natural Sciences.

Conflict of interest
The authors declare no conflict of interest.

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