Thermodynamics of LiFePO4 Solid-Phase Synthesis Using Iron (II

Article pubs.acs.org/jced

Thermodynamics of LiFePO4 Solid-Phase Synthesis Using Iron(II) Oxalate and Ammonium Dihydrophosphate as Precursors Alexei V. Churikov,*,† Alexander V. Ivanishchev,† Arseni V. Ushakov,† Irina M. Gamayunova,† and Ilya A. Leenson‡ †

Institute of Chemistry, Saratov State University, 83 Astrakhanskaya Str., 410012 Saratov, Russian Federation Department of Chemistry, Moscow State University, Lenin’s Hills, 199992 Moscow, Russian Federation

‡

ABSTRACT: A detailed thermodynamic analysis and an experimental study of the thermolysis mechanism of Li2CO3 + NH4H2PO4 + FeC2O4 blends were performed from the viewpoint of the usage of these compounds for the LiFePO4 solid-phase synthesis. Thermodynamic calculations of the assumed chemical reactions have been done within a practically important (for LiFePO4 synthesis) temperature range of (25 to 900) °C. The thermodynamic parameters of the related compounds (Li2O, Li3PO4, (NH4PO3)4, NaFePO4, FePO4 2H2O, FePO4, and LiFePO4) were determined more precisely than earlier. The temperature dependences of changes in the standard Gibbs energy, standard enthalpy, standard entropy, and standard molar heat capacity for the chemical reactions in LiFePO4 synthesis were calculated. Various paths of oxalate decomposition may well proceed concurrently with the predomination of this or that path under slight changes in the experimental conditions. The formation of orthorhombic lithium orthophosphate Li3PO4 was detected just in a blend grinded at room temperature. Heating up to 360 °C results in full destruction of the reaction mixture; Li3PO4 in its activated state is a crystal component at this synthesis stage. Lithium orthophosphate structurally belonging to the same spatial Pnma group as the target product LiFePO4 is the basis of its synthesis.

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INTRODUCTION As is known, orthorhombic lithium iron phosphate LiFePO4 (triphylite, lithium iron phosphate, iron-based orthophosphate, some other names are also in use), structurally belonging to the olivine group, has become the recent-year leader among the cathodic materials used in rechargeable lithium-ion batteries (LIB). The variety of the olivine synthesis methods is so wide that it is hard to meet two identical synthetic techniques. Main efforts were directed either to achieving unique properties of the product or to adaptation of some technique as a commercial technology. In this connection, attempts are made to simplify the scheme of the process or to reduce the cost of the materials used. Besides practical questions concerning selection of optimal synthesis conditions, attention is paid to fundamental problems, namely, elucidation of the mechanism of chemical reactions, identification of the intermediate products, and revealing of some regularities of the processes.1−3 As our analysis of the literature shows, iron(II) oxalate FeC2O4 (usually as its dihydrate FeC2O4·2H2O) remains the most popular source of iron for LiFePO4 synthesis. The compound FeC2O4·2H2O is met in the nature as the humboldtine mineral.4 Undoubtedly, the iron(II) oxalate usage is due to the apparent easiness of its thermal decomposition and the supposed stoichiometric simplicity of the proceeding reactions, since stoichiometric calculations underlie the development of a precursor blend for the LiFePO4 synthesis. For this reason (the easiness and stoichiometry of © XXXX American Chemical Society

thermal decomposition), ammonium dihydrogen phosphate NH4H2PO4 is chosen as a phosphorus source and Li2CO3 (or LiOH) as a lithium source. Though the thermolysis processes of such mixtures are studied rather insufficiently, it is implicitly supposed that at the LiFePO4 solid-phase synthesis each precursor in the mixture undergoes changes similar to those when in its individual state. Unlike complex mixtures, the individual thermolysis of the precursors was studied repeatedly, with the most determining role played by the thermal decomposition of ferrous oxalate, which can form various products at even small variations of the reaction conditions. Table 1 lists the set of proposed chemical reactions by eqs 1 to 14. A Review of Thermal Decomposition of Ferrous Oxalate and Ammonium Dihydrophosphate. We have traced the long and controversial history of the oxalate thermolysis researches, since this analysis bears an immediate relation to the olivine synthesis.5−53 For the first time, the decomposition by eq 1 was described by Magnus in 1825 and, later, by Vogel and Phipson.5−8 However, Liebig first pointed to another direction (eq 2), though the concurrent reaction by eq 1 cannot be excluded.9 Later, Moissan obtained impurityfree FeO by cooling the decomposition product in a CO flow.10 FeC2O4·2H2O losses its crystallization water by eq 3 in vacuum Received: February 16, 2013 Accepted: May 2, 2013

A

dx.doi.org/10.1021/je400183k | J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

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Table 1. Chemical Reactions Which Can Proceed at the LiFePO4 Solid-Phase Synthesis Using FeC2O4, NH4H2PO4, and Li2CO3 as Precursors chemical reaction equation

FeC2O4 → Fe + 2CO2

(1)

5−8, 14−16, 22, 23, 26, 34

FeC2O4 → FeO + CO2 + CO

(2)

9, 10, 13, 14, 17−21, 24, 25, 28

FeC2O4 ·2H 2O → FeC2O4 + 2H 2O

(3)

2FeC2O4 → Fe + FeO + 3CO2 + CO FeO + CO → Fe + CO2

2CO → C + CO2

references

11, 12, 28, 43

(4)

16

(5)

16, 17

(6)

4FeO → Fe3O4 + Fe

16, 17

(7)

17, 25, 30, 38, 48, 49

2FeC2O4 ·2H 2O + 3/2O2 → Fe2O3 + 4CO2 + 4H 2O

3FeC2O4 → Fe3O4 + 2CO2 + 4CO

(8)

16, 36−40

(9)

2FeC2O4 + FeO → Fe3O4 + 3CO + CO2

43, 47

(10)

36, 37, 45

FeC2O4 ·2H 2O → 0.26Fe3O4 + 0.07Fe3C + 0.89CO + 1.03CO2 + 2H 2O

Fe3O4 + CO → 3FeO + CO2

FeC2O4 → FeCO3 + CO

FeCO3 → FeO + CO2

(12)

51

(14)

51

(15)

55

NH4H 2PO4 → 1/2(NH4)2 H 2P2O7 + 1/2H 2O

NH4H 2PO4 → NH4PO3 + H 2O 2FeC2O4 → Fe2O3 + CO2 + 3CO 3FeC2O4 → Fe3O4 + 4CO2 + 2C

(16)

57, 58

(18)

57, 58

(19)

20

(20)

13

(21)

9Fe2O3 + 2NH3 → 6Fe3O4 + N2 + 3H 2O

(22)

3Fe3O4 + 2NH3 → 9FeO + N2 + 3H 2O

(23)

3Fe3O4 + 8NH3 → 9Fe + 4N2 + 12H 2O

(24)

3FeO + 2NH3 → 3Fe + N2 + 3H 2O

2NH3 → N2 + 3H 2

57, 58

(17)

2NH4PO3 → 2NH3 + P2O5 + H 2O

3Fe2O3 + CO → 2Fe3O4 + CO2

43 43

(13)

NH4H 2PO4 → NH3 + H3PO4

(11)

(25)

(26)

3Li 2CO3 + 2NH4H 2PO4 → 2Li3PO4 + 2NH3 + 3CO2 + 3H 2O Li3PO4 + 2NH4PO3 + 3FeO → 3LiFePO4 + 2NH3 + H 2O

(27)

(28)

at 142 °C and in a dry nitrogen flow between (100 to 200) °C.11,12 Subsequently, many researchers also detected Fe, its oxides, and carbon in various quantitative ratios in the decomposition products of FeC2O4, namely: the mixed product FeO + Fe3O4 + Fe + C;13 the product corresponded to the formula Fe5O4, which could be interpreted as a FeO:Fe = 4:1 blend;14 pyrophoric Fe in a hydrogen flow,15 and other variants. A rather detailed study was made by Herschkowitsch, who stated decomposition by eq 4.16 The “correct” CO:CO2 = 1:1 ratio takes place rarely because of the secondary reactions by eqs 5 to 7. The author himself assumes that pure metal and CO2 are initially formed by eq 1, and a secondary reaction of Fe with CO2 by eq 5 then proceeds; iron reacts with CO2 incompletely, so a Fe + FeO mixture appears. However, the reverse course of reaction 5 is thermodynamically possible at a rather high temperature only (see the Discussion section). The decomposition by eq 2 goes at temperatures above 300 °C, and very pure FeO results at temperatures above 850 °C.17 This contradictory information has subsequently been reflected in numerous monographs, handbooks, and textbooks, unfortunately, with insufficient critical analysis.18−27

These reactions were verified and corrected many times in later years, and many new results can be cited.28−46 Macklen studied the decomposition of FeC2O4·2H2O in CO2 and N2 atmospheres and discovered disproportionation of the primary FeO product into Fe and Fe3O4 by eq 7 above 320 °C.30 According to Boldyrev, the decomposition half-period of iron oxalate is 1 min at 270 °C, irrespective of whether dihydrate FeC2O4·2H2O or a salt previously dehydrated in vacuum at 200 °C is heated.32,33 Ugay made a comprehensive thermographic study of decomposition of the oxalates of many bivalent metals (Mg, Ca, Sr, Cd, Fe, Cu, Zn, Ba, Co, Ni, Pb, Mn, Sn, and Hg) and established, in particular, that FeC2O4·2H2O dehydrates about 200 °C and decomposes above 370 °C by eq 1.34 A more comprehensive investigation was conducted by Kornienko, who analyzed the gaseous products for CO and CO2, and the solid ones for the metal:oxide ratio.35 Kornienko, analyzing his own and literature data, suggests that all of the reactions proceed uniformly: first, a metal oxide is formed, which then partially or fully reduces (it is just opposite to the aforementioned model by Herschkowitsch). For example, FeC2O4·2H2O decomposes at 378 °C to release 1.16 mol of CO2 and 0.74 mol of CO, the solid residue being a Fe3O4:Fe = 9:1 blend.35 The initial B



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since the detailed description of the experimental technique does not correspond to the stated atmosphere of the conversion gases.43 An original mechanism has been recently proposed for the thermolysis of oxalates and carbonates of 16 metals.51−53 Hannesen51 has found that dehydration by eq 3 is the first stage of the thermal decomposition of FeC2O4·2H2O (as well as that of the corresponding oxalates of Zn, Co, Mn, Ni, Cu), while at the second one the metal carbonate is formed which then decomposes down to oxides (eqs 13 and 14). In sum, the longknown eq 2 results, so the proposed scheme reveals the detailed mechanism of the oxalate decomposition to FeO. There are fewer contradictions as to the paths of the NH4H2PO4 thermolysis; however, several versions have also been proposed, including ammonia and water evolution.54−58 According to the reference data, ammonium dihydrophosphate melts at 190 °C without decomposition; 54 however, NH4H2PO4 decomposes at 210 °C by eq 15 to form orthophosphoric acid and ammonia.55 When temperature increases, H3PO4 turns into pyrophosphoric acid H4P2O7 and then turns into metaphosphonic one HPO3. On the contrary, there was no evolution of ammonia from NH4H2PO4 heated up to 580 °C.56 The thermolysis of NH4H2PO4 proceeds with water evolution only to form (NH4)2H2P2O7, then to form ammonium metaphosphate NH4PO3, and may finish with full destruction of the compound (the sequential reactions of eqs 16 to 18).57,58 The processes occurred in a complex way, which could not be separated into elementary reactions, with the main weight loss occurring below 600 °C.57,58 Thus, the presented literature analysis shows a variety of ferrous oxalate thermolysis versions and no unambiguous relation between the synthesis conditions and the final product composition. Not always the authors take account of the thermodynamic possibility or impossibility of some processes. The thermolysis of ammonium dihydrophosphate may yield various solid, liquid, and gaseous products as well. The aim of the present work is an experimental study and thermodynamic calculations of the thermolysis mechanism of FeC2O4 and NH4H2PO4 from the viewpoint of the usage of these compounds for the solid-state synthesis of LiFePO4 as the electrode material for LIBs. To synthesize the electrode material olivine, of special importance are the plurality and ambiguity of the thermolysis products, and the possible formation of various compounds of iron in various oxidation degrees under similar conditions, which hinders selection of optimal conditions of the process.

dehydration is quickly followed by oxidative decomposition to form the hematite α-Fe2O3 as the final product in an oxidative atmosphere (air or oxygen), which is described by the longstanding eq 8.16,36−40 Both FeO and Fe3O4 are assumed as primary intermediates of FeC2O4 oxidation in an oxygen or dry air atmosphere.36−39 Unlike this, the maghemite γ-Fe2O3 is formed in a wet oxidative atmosphere.36,37 Zboril et al. observed the simultaneous formation of α-Fe2O3 and γ-Fe2O3 in a superparamagnetic state at the early initial stage of oxalate decomposition.40 The formation of γ-Fe2O3 was observed as well.41,42 In an inert (N2, Ar) or reductive (H2) atmosphere, the thermal decomposition of FeC2O4·2H2O also proceeds in two stages (dehydration and decomposition), but these processes go within different temperature ranges.43 Depending on the atmosphere and thermal conditions, FeO, Fe3O4, α-Fe2O3, αFe, and Fe3C can be identified in the solid residue. FeO is thought as the primary thermolysis product, while the magnetite Fe3O4 and iron carbide Fe3C are transformed into α-Fe at the reductive stage of the reaction.44,45 If ammonia is present in the gaseous atmosphere along with hydrogen, oxide reduction to pure iron is complicated by the formation of its various nitrides.46 Recently, Małecka found that the thermal decomposition of iron(II) oxalate after its dehydration proceeds in two concurrent paths, namely, by eqs 1 and 9 (the latter reaction never met before).47 The kinetics of both reactions has been established to be described by the Avrami−Yerofeyev equation with a high accuracy, which gives evidence of the determining role of nucleation and the growth of these nuclei. However, some authors38,48,49 consider that FeO is formed first in an inert atmosphere by eq 2, which decomposes into Fe3O4 and αFe by eq 7 above 570 °C (below, the reaction by eq 7 will be shown to be thermodynamically forbidden above 570 °C). According to some other data, an oxalate sample isothermally heated at 440 °C in dry nitrogen converted into a mixture of FeO, Fe3O4, and α-Fe2O3, while gradual heating of the oxalate in an argon medium resulted in the formation of a α-Fe4C, Fe3O4, and Fe blend.50 An unusual thermolysis mechanism (eq 10) was proposed by Rane et al.: the previously formed oxide FeO enters into a reaction with the nonreacted oxalate in a dry nitrogen atmosphere to finally form the magnetite Fe3O4, so eq 9 results in total.37 The exotic nature of this scheme (in comparison with those considered above), nevertheless, is confirmed by some publications, where Fe3O4 has been detected as the only product in a wet nitrogen or argon atmosphere.36,44 A very detailed study was recently made by Hermanek et al.: solid phase transformation was systematically studied in a temperature range of (25 to 640) °C; the results obtained are of significant interest.43 Unlike the majority of other studies, here, the wüstite FeO has turned out as the final product of oxalate thermolysis rather than the primary one, along with Fe3O4 and α-Fe. As to the initial stage, the cementite Fe3C and magnetite Fe3O4 are formed almost simultaneously in the stoichiometry of eq 11. Equation 11 is very close to the stoichiometry proposed by Kornienko.35 The full transformation of iron(II) oxalate was described by the authors as the following sequence: crystalline water elimination within (170 to 230) °C by eq 3, conversion of anhydrous oxalate to magnetite by eq 9 at a temperature above 230 °C, and final reduction of magnetite to wüstite at a temperature above 535 °C by eq 12 (CO previously formed by eq 9 at a low temperature). However, the proposed mechanism is doubtful,

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EXPERIMENTAL SECTION The source substances lithium carbonate (Li2CO3, purity w > 0.99, JSC “Plant of rare metals”, Russia), ammonium dihydrophosphate (NH4H2PO4, purity w > 0.99, Reachim, Russia), and iron(II) oxalate dihydrate (FeC2O4·2H2O, purity w > 0.99, Aldrich, Germany) were mixed in the 1:2:2 (stoichiometric) molar ratio, which corresponded to the elementary composition of the target product Li:P:Fe = 1:1:1. The synthesis of lithium iron phosphate LiFePO4 included the stages of mechanochemical activation in an AGO-2 ball miller-activator (Noritz, Russia) at room temperature and 560 rpm for 20 min and thermal treatment in a tube furnace in flowing argon with a heating rate of 10 °C·min−1 with subsequent storage at some fixed temperatures (360, 600, 650, 800, 850, and 900) °C during (1 to 600) min. The lower temperature limit corresponds to the full decomposition of the C



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Table 2. Conditions under Which the Experiments Were Performed and the Synthesis Products XRD analysis resultsa

synthesis conditions treatment type

a

temperature conditions/°C

time/min

crystal phases Li3PO4 FeC2O4·2H2O NH4H2PO4 Li3PO4 NH4PO3 LiFePO4 Li3PO4 LiFePO4 Fe2O3 Li4P2O7 FePO4 LiFePO4 FeO LiFePO4 Fe2O3 LiFePO4 Li3PO4 Fe3O4 FePO4 LiFePO4 Li3PO4 FeO Fe3O4 Fe LiFePO4

mechanical activation

r.t.

20

heat treatment of the precursor mixture in an Ar atmosphere

r.t. to 360 at 360 r.t. to 400 at 400

34 1 38 5

r.t. to 500 at 500

48 5

r.t. to 600 at 600 r.t. to 650 at 650

58 1 63 1

r.t. to 650 at 650

63 600

r.t. to 650 at 650

63 600

r.t. to 800 at 800

78 180

Bold values indicate the basic phase according to XRD analysis results.

lattice parameters/Å a

b

c

10.518

6.119

4.922

10.403

6.024

4.710

10.321

6.009

4.688

10.323

6.003

4.789

10.312

5.977

4.663

10.328

6.009

4.698

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RESULTS AND DISCUSSION Chemical and Phase Transformations during the LiFePO4 Synthesis. Let us analyze here the mainstream of the solid-phase synthesis of LiFePO4 from iron oxalate FeC2O4·2H2O, ammonium dihydrophosphate NH4H2PO4, and lithium carbonate Li2CO3 (the source of lithium is no matter of principle in this case; LiOH can be used as well). The basic results of our experimental investigation are summarized in Table 2 and are also illustrated in Figure 1. According to them, Li2CO3 completely disappears from the reaction mixture as early as at the stage of mechanoactivation at room temperature. This is proven by both complete disappearance of the diffraction pattern of Li2CO3 (Figure 1a) and negative analytical reaction with concentrated hydrochloric acid for carbonates: carbonic gas evolves only from the initial reagent mixture, but not from the milled blend, which means decomposition of the carbonate with CO2 removal at mechanoactivation. In the blend milled at room temperature, the formation of orthorhombic lithium phosphate Li3PO4 (Figure 1a) is reliably fixed by X-ray analysis. Heating up to 360 °C (by this temperature, intense gas release is mainly over) results in the full destruction of all of the source components in the reaction mixture. According to XRD data (Figure 1b, Table 2), Li3PO4 and NH4PO3 are the only crystal components at this synthesis stage (ammonium metaphosphate is detected less reliably; some difference at 2θ can be explained by increased lattice parameters in the amorphized NH4PO3 phase). All iron compounds are part of the amorphous mass. Further heating up to (400 and 500) °C results in the appearance of the target product LiFePO4, which

precursors, while the upper one is determined by consolidation of the product particles (grains), which leads to an insufficient activity of the product as the electrode material in LIBs.2,3 For the structural characterization of the materials, X-ray diffraction (XRD) and laser diffraction analyses were used. Xray diffraction patterns were recorded on X-ray diffractometers ARLX’TRA (Thermo ARL, Switzerland) and DRON-4 (Burevestnik, Russia) with filtered Cu Kα radiation. The ICDD search indexing software and PDF-4 database (2007) were used for XRD analysis. Hygroscopic samples were prepared in a dry chamber; they were covered with an X-rayamorphous scotch during exposition for moisture protection. X-ray diffraction patterns were processed, and the cell parameters were calculated with the usage of the WinPlotr and CellRef software. The mean size of crystallites in the synthesized LiFePO4 materials was calculated by the Debye Scherrer formula with the aid of a corundum reference with this value ≫100 nm, with the diffraction peak shape to be simulated by the WinPlotr and FullProf software, on elimination of the instrumental peak broadening. The linear sizes of the particle agglomerates were estimated on a laser diffraction particle size analyzer SALD-2201 (Shimadzu) by monochromatic radiation diffraction. Powders were analyzed with the help of the SALD WEAL software in an aqueous suspension. For thermogravimetric and differential thermal analyses with gas release analysis, a TGAQ500 research-grade thermogravimetric analyzer (TA Instruments, USA) and a FTIR Thermo Scientific Nicolet 6700 spectrometer (Thermo scientific, USA) were used; heating of the samples in the range (25 to 1000) °C was performed at 10 °C/min. D



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Fe2O3, α-Fe, FePO4, Li4P2O7, and graphite are sometimes detected as well. Neither phase is accumulated in a considerable amount; however, their detection enables one to hypothesize a complex staged mechanism of the process. So, lithium orthophosphate Li3PO4 structurally belonging to the same spatial Pnma group as the target product LiFePO4 is the basis of its synthesis. The lattice parameters of both substances are similar. In the considered solid-phase synthesis, lithium orthophosphate Li3PO4 is formed in an activated state, with somewhat increased lattice parameters (a = 10.518 Å; b = 6.119 Å; c = 4.922 Å; ⟨D⟩ ≈ 9 nm). The phase composition of the material sampled at (800, 850, and 900) °C is the same, namely, pure lithium iron phosphate LiFePO4, with no side phases detected (Figure 1d). The most low-temperature LiFePO4 has a somewhat increased parameter a of the orthorhombic lattice (Table 2), but it decreases and stabilizes in the course of further heat treatment. The well-crystallized and phase-pure triphylite results with the size of crystallites ⟨D⟩ regularly changing with temperature, namely, increasing from 15 up to 60 nm as heat treatment progresses, reaching an optimal value for the electrode material. The average lattice parameters of LiFePO4 (a = 10.327 Å; b = 6.006 Å; c = 4.693 Å) are closest to the card no. 01-070-6684 of the PDF-4 database. The mean size ⟨D⟩ of LiFePO4 nanoparticles is close to 40 nm, and the macrostructure of the powdered material is formed by nanoparticle agglomerates. According to granulometric curves obtained by monochromatic radiation diffraction, these agglomerates form a disperse macrosystem with a distribution maximum near 2 μm. Thermodynamics of Reactions at LiFePO4 Synthesis. A thermodynamic assessment of the possibility of some processes

Figure 1. X-ray diffraction pattern of the Li2CO3 + NH4H2PO4 + FeC2O4·2H2O precursor blend after its mechanochemical activation at 25 °C (a), the semiproduct of LiFePO4 synthesis after its thermal treatment at 360 °C (b), and the final product of the LiFePO4 synthesis at 600 °C (c), 650 °C (d), and 800 °C (e); PDF cards of NH4PO3, Li3PO4, and LiFePO4 are indicated.

becomes the main component of the reaction mixture. Further heating up to 600 °C, even for a short time, is enough for almost full proceeding of the reactions of the final stage (as XRD shows, Figure 1c). Some intermediate crystal forms in small or trace quantities (w < 0.05), including FeO, and Fe3O4, Table 3. Standard Thermodynamic Properties of Substances ΔfH298 °

S298 °

ΔfG298 °

Cp,298 °

substance

kJ·mol−1

J·(mol·K)−1

kJ·mol−1

J·(mol·K)−1

FeC2O4 Na3PO4 NH4H2PO4 Li2CO3 Li3PO4 (NH4)2H2P2O7 (NH4PO3)4 NaFePO4 FePO4 2H2O FePO4 (trig) FePO4 (orth) FePO4 LiFePO4 LiFePO4 LiFePO4 α-Fe

−849.68 −1924.64 −1445.07 −1216 −2095.64 −2521.95 −4475.62 −1571.8 −1888 −1112.55 −1220.4 −1289.56 −1392.45 −1591.23 −1514.8 0

228.01 224.68 151.96 90.17 105

−783.87 −1811.31 −1210.56 −1132.67 −1967.17

211.33 153.57 142.5 96.2 145

465.01

−3772.96

461.20 123.26 180.5

FeO P4O10 P4O10 P (cr, white) Li2O

−264.85 −2984.03 −3009.9 0 −598.73

103 b

10−5 c′

T/K

reference

191 136.1 34.33 42.53 125.02

161.09 67 362.8 177.34 67

−24.62

(298 to 700) (298 to 1600)

117.15 107

1154.9 89.91

102.68

95.7

51a 65 61, 68, 70b 65, 70 54, 65, 68, 71 60 62, 63c 69 54 66 66 d 66, 72 64, 68e f 65

a

(298 to 623) 298

171.3

108.51 136.75 136.75 27.15 60.75 228.86 228.78 41.09 37.91

−1110.14 −1179.3 −1282.19 −1480.98 −1404.54 0 −244.3 −2697.6 −2723.3 0 −561.91

145.93 115.08 24.98 49.92 211.71

17.24 −159.8 50.8 93.3

24.77 181 8.61 407.19

23.82 54.1

23.82 66.92

17.27

−9.371

−14

332.36 −3.31

−17.7

(295 298 298 298 298 298 (290 298 (298 (700 (298 (298 298 (273 (265

to 470)

to 780) to to to to

700) 1000) 1650) 630)

to 317) to 1733)

65 65 67 65, 67 54, 70, 73g

a

Calculated by eq 34 with the use of ref 65 data for simple substances. bCoefficients of eq 29 for NH4H2PO4 are calculated from Figure 2a. Coefficients of eq 29 for (NH4PO3)4 are calculated from Figure 2b. dCalculated from the data for the crystal hydrate FePO4·2H2O. eCalculated from the data for the oxides Li2O, P2O5, and FeO. The coefficients in eq 29 for LiFePO4 are calculated from Figure 2c. fCalculated from the data for Li3PO4, NaFePO4, and Na3PO4. gThe coefficients in eq 29 for Li2O are calculated from Figure 2d. c

E



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Figure 2. Temperature dependences of the standard molar heat capacity under constant pressure Cp° for NH4H2PO4 (a), (NH4PO3)4 (b), LiFePO4 (c), Li2O (d); the Cp° values taken from refs 61, 62, 64, and 73, respectively.

can help their substantiation. Thermodynamic calculations of the assumed chemical reactions will be performed in a practically important (for LiFePO4 synthesis) temperature range, that is, (25 to 900) °C. It is evident that any of eqs 1 to 18 can proceed under these conditions in principle. Besides the main reactions, some side ones may also take place to form iron nitrides and phosphides;46,59 we will not consider here these possibilities which are realized in special conditions only. Besides, on the basis of experimental data (Figure 1, Table 2), some paths of oxalate thermolysis are also possible to form Fe2O3 and carbon and secondary reactions with ammonia involved (eqs 19 to 26). Just at a low temperature, the solidphase acid−base interaction by eq 27 can go in the reaction mixture, and the chemical process at the final (high-temperature) stage of olivine synthesis can be represented by eq 28. The accuracy of our calculations is restricted by the quality of the thermodynamic data available for the source reagents and reaction products, namely, the standard enthalpy of formation ΔfH298 ° (the enthalpy change at the formation of a compound from its constituent elements, at their standard states), the standard Gibbs free energy of formation ΔfG°298 (the Gibbs free energy change at the formation of a compound from its constituent elements, at their standard states), the standard entropy S298 ° , the standard molar heat capacity under constant pressure C°p,298, and the coefficients a, b, and c′ in the absolute temperature (T) dependence of heat capacity C p◦ = a + bT + c′/T 2

In view of eq 29, the expressions for the temperature dependences of changes of the standard molar heat capacity ΔCp°, the standard enthalpy of reaction ΔrHT°, the standard entropy of reaction ΔrST°, and the standard Gibbs energy of reaction ΔrG°T, have the form ΔC p◦ = Δa + ΔbT + Δc′/T 2

(30)

◦ Δr HT◦ = Δf H298 + Δa(T − 298.15) + Δb(T 2 − 298.152 )

/2 − Δc′(T −1 − 298.15−1)

(31)

◦ Δr ST◦ = ΔS298 + Δa ln(T /298.15) + Δb(T − 298.15)

− Δc′(T −2 − 298.15−2 )/2 Δr GT◦ = Δr HT◦ − T Δr ST◦

(32) (33)

where Δa, Δb, and Δc′ are the algebraic sums of the corresponding coefficients. For our work, taking thermodynamic data from a sole source was possible only for well-known substances (α-Fe, Li2CO3, etc) whose characteristics completely coincide in few reference books. Table 3 presents the standard thermodynamic values needed for our calculations and taken from various sources with a necessary indication of the references and comments on how the calculated thermodynamic values have been obtained by us, if absent in the source explicitly.51−73 In every case, preference

(29) F



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(hi(Li2O) = (−817.30 ± 12.97) kJ·mol−1; gi(Li2O) is not given). This enables us to evaluate, by this method, the standard enthalpy of formation only: ΔfH298 ° (LiFePO4) = (−1591.23 ± 13) kJ·mol−1, which is very near to the corresponding value for orthorhombic sodium iron phosphate: ΔfH298 ° (NaFePO4) = −1571.8 kJ·mol−1,69 but is in contradiction with the value ΔfH298 ° (LiFePO4) = −1392.45 kJ·mol−1 calculated from the data from ref 66. Taking various data sets and applying the quasi-additivity principle to compounds close by composition and structure, we have additionally found ΔfH298 ° (LiFePO4) = −1514.8 kJ·mol−1 (from the standard enthalpies of formation of Li3PO4, NaFePO4, and Na3PO4) and S298 ° (LiFePO4) = 136.75 J·K−1·mol−1 (from the standard entropies of the constituent oxides Li2O, P2O5, FeO). Table 3 presents the values of ΔfG298 ° (LiFePO4) corresponding to different calculation techniques. A justified choice between them can be made from easily implementable electrochemical measurements. For example, for the electrochemical circuit Li|LiFePO4 the experimental emf value of 3.425 V was measured (ref 74; close values can be found in many works), which entails ΔrG°298 = −330.46 kJ·mol−1 for the current-generating reaction

was given to the primary experimental papers containing fuller data, including the coefficients of eq 29. The thermodynamic information for the following compounds is most scarce: (NH4)2H2P2O7, NH4H2PO4, NH4PO3, FeC2O4, LiFePO4, and FePO4. Only the value of ΔfH°298 has been found for ammonium dihydropyrophosphate, which hinders using (NH4)2H2P2O7 in the complete calculation.60 For a number of compounds we have additionally calculated their thermodynamic quantities on the basis of segmental experimental data, and in separate cases, estimation techniques were used. So, Cp,298 ° and the coefficients a, b, and c′ in eq 29 for NH4H2PO4, (NH4PO3)4, and LiFePO4 have been estimated by means of processing of the heat capacity data,61−64 as shown in Figure 2. The values of ΔfG298 ° for Li3PO4, (NH4PO3)4, FePO4 2H2O, FePO4, and LiFePO4, Cp,298 ° for NaFePO4, the coefficients of eq 29 for Li2O have been calculated by us as well. Because of the lack of more reliable data, the same value of the linear coefficient b for Na3PO4 and Li3PO4 was accepted. Further, according to Hannesen, the temperature dependence of the standard Gibbs free energy of FeC2O4 formation within (298 to 700) K is described by the equation51 Δf GT◦(FeC2O4 ) = −878002 − 77.12T ln T + 773T − 0.06T 2 (J·mol−1 )

Li + FePO4 ↔ LiFePO4 (34)

Our closest calculated value is ΔrG298 ° = ΔfG298 ° (LiFePO4) − (ΔfG298 ° (FePO4) + ΔfG298 ° (Li)) = −1480.97 − (−1179.30 + 0) = −301.67 kJ·mol−1. The discrepancy of 29 kJ·mol−1 between our calculation and experiment is not small, though an exacter coincidence of thermodynamic calculations with experimental data is hardly possible at the existing level of our knowledge of the thermodynamic properties of these substances. Moreover, as shown below, a smaller discrepancy is hardly possible today because of the specificity of phosphorus-containing substances, namely, the discrepancy of the thermodynamic quantities for P, P2O5, and P4O10 recommended in different years. It should be taken into account that white phosphorus as well as red phosphorus were regarded as simple substances in different works.67 This causes a significant difficulty at comparing data taken from various sources and obtained in different years. According to recent Rard and Wolery’s detailed analysis,67 various reference databases give the value of ΔfH298 ° (P4O10) equal to (−3012.5 to −2984.0) kJ·mol−1; similar discrepancies are observed for the entropies, Gibbs energies of formation, and enthalpies of formation of other phosphorus-containing substances and ions. The basic cause of these discrepancies is a jump from older databases to more recent ones without adjusting for the enthalpy difference between red phosphorus and the α-form of white phosphorus, that is, between the standard state in earlier tabulations and the current standard state.67 As is noted, the failure to adjust for this difference results in an error of (35 to 40) kJ·mol−1.67 On the basis of the aforesaid, the most justified set of the thermodynamic properties of orthorhombic lithium iron phosphate should be: ΔfH298 ° (LiFePO4) = −1591.23 kJ·mol−1, ΔfG°298(LiFePO4) = −1480.97 kJ·mol−1, and S°298(LiFePO4) = 136.75 J·K−1·mol−1. We have used the data for P4O10 and P2O5, where they contain the coefficients of eq 29 (Table 3, ref 65) and the data for P4O10 of 1998 recommended in ref 67. The usage of other values for P, P2O5, and P4O10 would somewhat change the calculated values with no influence on the conclusions under discussion. It should be noted that the data in ref 66 have not been used in ref 74 correctly enough, where the enthalpy of formation of

From this equation, all of the thermodynamic data of FeC2O4 presented in Table 3 have been calculated by using reference data for the simple substances α-Fe, C, and O2.65 We have calculated the thermodynamic properties of FePO4 and LiFePO 4 by using a method of evaluating the thermodynamic parameters of phosphate compounds as the sum of oxide units at high and low temperatures, recently developed by La Iglesia.68 This method can calculate the values of ΔfH°298 and ΔfG°298 for mineral phosphates with a high accuracy, by using the sets of constituent hi and gi values for each basic oxide unit (P2O5, FeO, Fe2O3, H2O, etc.), respectively, for example, hi(H2Ocryst) = −299.22 kJ·mol−1; gi(H2Ocryst) = −239.10 kJ·mol−1.68 Hereof we have found ΔfH298 ° (FePO4) = −1289.56 kJ·mol−1, ΔfG298 ° (FePO4) = −1179.30 kJ·mol−1, and S298 ° (FePO4) = 108.51 J·K−1·mol−1, with the aid of the known thermodynamic data for crystal hydrate FePO4 ·2H2O (Table 3). This consistent data set for FePO4 should be treated as most justified for now. We note that a close value (S298 ° (FePO4) = 100.92 J·K−1·mol−1) was obtained at calculations by the additivity technique from the standard entropy values of the constituent oxides Fe2O3 and P4O10. Moreover, the enthalpies of formation ΔfH298 ° for LiFePO4 and FePO4 were calculated from the measured enthalpies of the reactions 1/2Li 2O + FeO + 1/4P4 O10 → LiFePO4

(35)

1/2Fe2O3 + 1/4P4 O10 → FePO4

(36)

(37)

which are −151.52 kJ·mol−1 (eq 35), −113.68 kJ·mol−1 (eq 36, o-FePO4), and −102.01 kJ·mol−1 (eq 36, t-FePO4) at 25 °C, respectively,66 with the usage of the ΔfH°298 of α-Fe2O3, FeO, Li2O, and P4O10 (the nesessary numeric data are given in Table 3). As a result of the application of a thermodynamic cycle, a close value ΔfH°298(FePO4) = −1220.4 kJ·mol−1 was calculated on the basis of these numeric data. Due to the poorer experimental database, La Iglesia68 describes the lithium−phosphate compounds with a lower accuracy than the soduim and potassium compounds: G



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Table 4. Standard Thermodynamic Parameters of Reactions with the Temperature Ranges of Applicability of the Coefficients in Equations 30, 31, 32, and 33 Indicated ΔrH°298 reaction number

−1

kJ·mol

ΔrS°298 J·(mol·K)

ΔrG°298 −1

kJ·mol

−1

ΔC°p,298 J·(mol·K)−1

Δa

103 Δb

10−5 Δc′

−85.48 −67.65 −153.13 −17.83 −194.87 4.18 −99.69 −276.73 −284.81 −217.16 −90.4 81.86 −113.93 46.28 101.57 24.96 77.68 −154.89 −276.45 −104.95 −303.76 256.67 129.47 −1463.89 −42.4 −573.52 50.12 335.81 61.64 −4.91 68.53 51.01 115.13 646.25

−118.24 −139.34 −257.58 21.1 177.33 5.61 199.25 355.48 −239.87 −100.53 −93.81 −178.15 −44.86 −94.48 −337.32 −63.37 −312.19 −228.71 −228.65 206.39 589.79 −563.83 −462.07 944 33.92 502.61 −36.91 −1013.41 −321.51 −23.34 −151.03 −915.31 −384.81 −853.5

7.54 12.31 19.85 −4.77 327.59 −16.16 13.24 345.6 54.94 42.63 15.41 −18.01 24.16 −11.85 −1.67 0.33 −3.01 26.43 22.62 30.59 120.34 −25.46 17.32 3008.56 14.26 1011.34 4.84 −27.97 −35.08 −1.84 −26.24 −7.65 −20.77 −1017.85

1 2 4 5

62.66 80.79 143.44 −18.13

226.46 243.95 470.42 −17.49

−4.87 8.05 3.18 −12.92

−112.13 −95.16 −207.3 −16.97

6 7

−172.45 −57.73

−175.7 −69.66

−120.07 −36.97

−12.63 −23.91

9 10 11 12 13 14 15 17 18 19 20 21 22 23 24

202.76 121.97 57.29 39.6 0.995 79.79 132.23 84.35 412.11 152.09 −142.15 −50.76 63.11 334.19 817.19

679.69 435.74 213.21 52.17 64.94 179.01 241.53 153.01 455.97 437.75 328.29 46.14 462.43 480.52 1295.14

0.11 −7.95 −6.67 24.05 −18.37 26.42 60.08 38.71 276.11 21.58 −240.03 −64.54 −74.83 190.94 431.03

−292.43 −197.27 −100.55 6.94 −98.93 3.77 −1.24 6.41 −20.82 −194.38 −317.69 −1.73 30.53 56.54 10.97

25

161

271.54

80.03

−15.19

26 27 28 35 38 41 42

91.88 349.01 20.63 −281.01 224.59 404.13 75.02

197.74 1228.05 464.52 −0.17 161.1 1410.95 516.99

32.96 −17.13 −117.92 −281.31 176.6 −16.61 −79.16

45.29 −1.12 −76.2 −13.85 −6.93 −232.95 −25.29

LiFePO4 from FePO4 and Li was calculated as −37.84 kJ·mol−1. From the emf value equal to 3.425 V for the Li|LiFePO4 electrochemical cell (micrometer-size LiFePO4 material), one can calculate ΔfG298 ° (LiFePO4) = −39.75 kJ·mol−1 and ° (LiFePO4) ≈ 6.4 J·K−1·mol−1. This seems impossible S298 because it corresponds to the emf value equal to 0.412 V rather than to the correct value of 3.425 V for eq 37. Comparison of Experimental Data with Thermodynamic Estimations. Of considerable interest would be to compare the results of thermodynamic calculations and experimental studies with the aim to elucidate the mechanism, stage consequence, and intermediates of the synthesis of LiFePO4. Table 4 presents the calculated coefficients of eqs 30 to 33 for the above reactions, and Figure 3 shows plots of the temperature dependences of ΔrHT°, −TΔrST°, and ΔrGT° for several cases. As is shown in this section, the analyzed reactions are preferably endothermic; their spontaneous proceeding becomes only possible at increasing temperatures above a certain level just due to the predominating contribution of the entropic term −TΔrS°T (Figure 3). Let us consider consistently the synthesis process of LiFePO4, beginning from the precursor mixture at room temperature. Our thermodynamic calculations confirm the

T/K (298 (298 (298 (298 (700 (298 (298 (700 (298 (298 (298 (298 (298 (298

to to to to to to to to to to to to to to

700) 700) 700) 700) 1000) 2500) 700) 1000) 700) 700) 700) 1000) 700) 855)

(298 (298 (298 (298 (298 (298 (298 (298 (700 (298 (700 (298 (298 (298 (298 (298 (298 (298 (700

to to to to to to to to to to to to to to to to to to to

700) 630) 700) 700) 1000) 1000) 1000) 700) 1000) 700) 1000) 1800) 700) 780) 780) 700) 700) 700) 1000)

experimental observation: the solid-phase acid−base interaction between Li2CO3 and NH4H2PO4 to form Li3PO4 and to remove the gaseous products by eq 27 becomes quite possible just at usual temperature. The process is additionally promoted by mechanical activation, which changes the energetic state of the solid reagents. As a result, the whole lithium carbonate and 1/3 of the ammonium dihydrophosphate introduced into olivine synthesis are spent by eq 27. Then, above 100 °C, dehydration of FeC2O4·2H2O with simultaneous oxalate decomposition down to carbon, α-Fe, FeO, Fe3O4, Fe3C, FeCO3, or other iron compounds becomes possible by eqs 1 to 4, 9 to 11, 13, 19, and 20. Here, thermodynamic calculations do not allow ascertaining the preferred direction: all of these reactions become potentially possible within (100 to 250) °C. The detection of insignificant amounts of these semiproducts in the reaction mixture with no explicit relation with the conditions of mechanical and heat treatments (Table 2) confirms that the thermolysis proceeds by several paths. Around 300 °C, the thermolysis of the remaining NH4H2PO4 by eq 17 becomes possible. While the thermolysis of the individual substance FeC2O4·2H2O more often produces iron and magnetite Fe3O4 (eqs 1, 9, and 10), which is confirmed by the expressed ferromagnetic properties of the product, FeO is H



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Figure 3. Temperature dependences of the standard enthalpy change of reaction ΔrHT°, the −TΔrST° value, and the standard Gibbs free energy change of reaction ΔrGT° by eq 2 (a), eq 5 (b), eq 17 (c), eq 27 (d), eq 28 (e), and eq 38 (f).

and that of enthalpy of 14 kJ·mol−1 at the melting point of NH4H2PO4, which is taken into account in our calculations and is reflected as a step on the corresponding dependences (Table 4 and Figure 3). Our thermodynamic estimation does not confirm the preference of the staged thermolysis of NH4H2PO4. Really, ° being equal to both eqs 16 and 17 are endothermic [ΔrH298 (63.19 and 84.35) kJ·mol−1, respectively]. Though the factor ΔrTS°T for eq 16 cannot be calculated for the lack of the

preferably formed in a reductive medium, apparently, by eqs 2, 13, and 14. The reaction mass at this synthesis stage (360 °C) shows no ferromagnetic properties. During the olivine synthesis, NH4H2PO4 melts (by different data, within (180 to 190) °C). With no exact calorimetric data, the heat effect of ionic salt melting can be estimated by the ° as (11 to 14) J·K−1 per 1 mol of ions.76 entropy increase ΔSm Assuming that (2 to 3) ions per formula unit are formed, this gives an additional entropy increment about 30 J·K−1·mol−1 I



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necessary data, it is approximately half as much as that for eq 17 due to half as much gas release. Our calculation based on these grounds shows that, at increasing temperature, the formation of ammonium metaphosphate NH4PO3 becomes possible around 300 °C, earlier than the formation of (NH4)2H2P2O7. Hence, NH4H2PO4 directly converts into NH4PO3 by eq 17, this process starting at ca. 300 °C. Wang et al. conducted a similar study with gas release analysis at heating of Li2CO3 + NH4H2PO4 + FeC2O4·2H2O blends and offered their own mechanism including the thermal decomposition of Li2CO3 to Li2O and CO2 within (600 to 700) °C77 Li 2CO3 → Li 2O + CO2

Li 2CO3 + 2NH4H 2PO4 + 2FeC2O4 → 2LiFePO4 + 3CO2 + 2NH3 + 2CO + 3H 2O

Any other decomposition products of iron oxalate (Fe, Fe3O4, etc.) require coparticipation of oxidative or reductive agents in the reaction, for example, Li3PO4 + 2NH4PO3 + 3Fe + 3CO2 → 3LiFePO4 + 2NH3 + 3CO + H 2O

(38)

(39)

but neither details of the formation ways of FeHPO4 are reported nor its presence in the reaction mixture is anyhow proven. We detected no presence of FeHPO4 in the reaction mixture. As for eq 38, this reaction is thermodynamically possible at a rather higher temperature only (Table 4 and Figure 3). This is confirmed by our thermogravimetric analysis and the absence of the thermal decomposition of lithium carbonate with apparent CO2 release up to 900 °C. Li2CO3 melts at ≈730 °C and begins to evaporate slightly (Figure 4). According to ref 78, the saturation pressure of CO2 is only 4 Torr near the melting point of lithium carbonate (≈1000 K).78

■

CONCLUSIONS In spite of plenty of the works concerning the thermolysis of iron(II) oxalate in various conditions, there are no grounds to conclude that the question of the reaction chemistry has been completely revealed and closed. The necessity of a new experimental study is justified by the practical importance of the decomposition products for various applications, including electrode materials, but the predecessors’ experience and the existing developments are not always taken into account. As shown in this paper, the proposed mechanisms of solid-state reactions are numerous but contradictive; the authors not always adequately estimate the thermodynamic probability of proceeding of these or those processes. So, our conclusion of the extreme ambiguity of the mechanism on the published data basis is not casual. This question gains a special significance when FeC2O4·2H2O is used as a precursor in the LiFePO4 synthesis for its application as an electrode material in LIBs. There are grounds to think that the known variability of the electrochemical and physicochemical properties of LiFePO4 with apparently similar synthesis techniques is caused by different mechanisms of the precursor decomposition reaction. The synthesis of the target electrode material with slight variations in the process conditions (especially, in the composition and renewal rate of the gaseous medium, the

Figure 4. Curves of thermogravimetric analysis (TG) and differential thermal analysis (DTA) of lithium carbonate Li2CO3 in air. The MP is indicated.

The uncertainty in the intermediate products of iron oxalate decomposition does not allow one to reliably formulate the reaction equation at the high-temperature stage. However, the solid-state process by eq 28 with no changes in the iron oxidation degree is only possible with FeO. In the whole, this provides grounds to write the following total equation of chemical reactions proceeding in parallel and in sequence. At the initial (low-temperature) stage of the olivine synthesis, the chemical processes described by the following total equation takes place: 3Li 2CO3 + 6NH4H 2PO4 + 6FeC2O4 → 2Li3PO4 + 4NH4PO3 + 6FeO + 9CO2 + 2NH3 + 6CO + 7H 2O

(42)

In principle, such reactions are quite possible as well, as follows from Table 4. By combining, summating, and subtracting the above equations, one can get plenty of more complex ones. Actually, eq 42 is a combination of eq 5 and 28; the situation with other redox processes is similar. The total eq 41 is the same in every case. The considered decomposition reaction with gas release is actually irreversible; the majority of them are thermodynamically possible just at low temperatures, and the experiment confirms their proceeding within (25 to 360) °C. Various paths of the oxalate decomposition may well proceed concurrently with the predomination of this or that path under slight changes in the experimental conditions. Just this explains such a considerable scattering of the published experimental data. The role of secondary reactions with participation of the formed semiproducts and gases is incorrectly overestimated. These secondary reactions would have equalized the initial differences in the composition and structure of the semiproducts by changing them toward the thermodynamically equilibrium ones. However, the experimental conditions at the solid-state synthesis of LiFePO4 as well as at studying thermolysis almost exclude any noticeable secondary participation of the conversion gases, provided that the considerable difference between the temperatures of their formation and assumed consumption is taken into account. This would be only possible in a closed synthesis device, but not in a protective gas atmosphere.

Then, according to ref 77, the target product is formed due to the interaction of lithium oxide Li2O with iron hydrophosphate FeHPO4 as Li 2O + FeHPO4 → LiFePO4 + H 2O

(41)

(40)

the final (high-temperature) stage is described by eq 28, and the total reaction of the olivine synthesis is J



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rate and program of heating, final temperature, the mode and degree of dispersing) leads to the formation of various intermediate products possessing different reactivities. “Deactivation” of even small part of FeC2O4· in Fe, FeO, Fe2O3, Fe3O4, Fe3C, or other iron compounds will inevitably affect the properties of the target product LiFePO4. This effect will combine the inevitably negative component (reduction of the theoretical electrochemical capacity of the product) as well as the possible positive one due to an improved conductivity of the composite and changes in its physicochemical properties. One should critically evaluate the literature and experimental data and allow for possible realization of many paths and mechanisms of this solid-phase synthesis rather than one scenario. From the viewpoint of the usage of this information for the solid-state synthesis of lithium iron phosphate for LIBs, several limiting combinations of the heating rate, blowing rate, and protective gas composition can be resolved: (1) Inert gas + low heating rate + high blowing rate: no secondary reactions with gas participation, the composition, and characteristics of the product are mainly determined by the conditions of the low-temperature thermolysis of the precursors. (2) Inert gas + high heating rate + low blowing rate: secondary reactions with gas participation may play a noticeable role, so the initial semiproducts undergo further conversion, approaching the calculated composition. (3) Reductive atmosphere of the conversion gases (CO + CO2+ NH3): complete proceeding of the thermodynamically possible reactions is possible.

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(9) Liebig, J. Simple preparation of ferrous, manganese, and stannous oxides. (Einfache Darstellung von Eisenoxydul, Manganoxydul und Zinnoxydul). Liebigs Ann. 1855, 95, 116−120. (10) Moissan, A. About metal oxides of iron family. (Sur les oxides metalliques de la famille du fer). Ann. Chim. 1880, 21, 199−255. (11) Birnie, S. Decomposition of ferrous oxalate by heat in nitrogen and hydrogen. (Sur la decomposition de l′oxalate ferreux par la chaleur, dans l′azote et dans l′hydrogene). Rec. Trav. Chim. 1883, 2, 273−275. (12) Baur, E.; Orthner, R. Zeitschrift für physikalische Chemie; Band 91, Heft 2. Naturwiss. 1916, 4, 75. (13) Hilpert, S.; Beyer, J. About iron (II) and iron (III) oxides (Ü ber Eisenoxyduloxyde und Eisenoxydul). Chem. Ber. 1911, 44, 1608−1619. (14) Mixter, W. G. The heat of formation of oxides and sulphides of iron, zinc, and cadmium. Ninth Communication on the bond energies of acid oxides with sodium oxide. (Die Bildungswärme der Oxyde und Sulfide von Eisen, Zink und Cadmium. Neunte Mitteilung über die Verbindungswärme von sauren Oxyden mit Natriumoxyd). Z. Anorg. Allg. Chem. 1913, 83, 97−112. (15) Finzel, T. G. Pyrophoric Iron. I. Preparation and properties. J. Am. Chem. Soc. 1930, 52, 142−149. (16) Herschkowitsch, M. Decomposition of oxalates. (Ü ber die Zersetzung der Oxalate). Z. Anorg. Allg. Chem. 1921, 115, 159−167. (17) Günter, P. L.; Rehaag, H. Thermal decomposition of oxalates. Communication II. Preparation of pure ferrous oxide. (Ü ber Die thermische Zersetzung von Oxalaten. II Mitteilung. Darstellung von reinem Ferrooxyd). Z. Anorg. Allg. Chem. 1939, 243, 60−68. (18) Smith, A. Introduction to inorganic chemistry; The Century Co.: New York, 1917. (19) Partington, J. R. A Text-Book of Inorganic Chemistry; Macmillan: London, 1953. (20) Latimer, W. M.; Hildebrand, J. H. Reference Book of Inorganic Chemistry; Macmillan: New York, 1951. (21) Remy, H. Lehrbuch der Anorganischen Chemie, Vol. 2; Akademische Verlagsgesellschaft m b H: Leipzig, 1942. (22) Nekrasov, B. V. Course in General Chemistry; Goskhimizdat: Moscow, 1962. (23) Jacobson, C. A. Encyclopedia of chemical reactions, Vol. 4; Reinhold: New York, 1951. (24) Cotton, F. A.; Wilkinson, G. Basic inorganic chemistry; Wiley: New York, 1976. (25) Ripan, R.; Ceteanu, I. Chimia metalelor, Vol. 2; Didactică şi Pedagogică: Bucureşti, 1969. (26) Karapetiants, M. Kh.; Drakin, S. I. The practical works on general and inorganic chemistry; Vysshaya Shkola: Moscow, 1969. (27) Brauer, G. Handbuch der präparativen anorganischen Chemie, Vol. 3; F. Enke: Stuttgart, 1981. (28) Robin, J. Bull. Soc. Chim. Fr. 1953, 20, 1078−1084. (29) Boullé, A.; Dorémieux, J.-L. Influence of nature of atmosphere on the thermal decomposition of iron, cobalt, and nickel oxalates. (Influence de la nature de l′atmosphére sur la décomposition thermique des oxalates de fer, cobalt et nickel). C. R. Hebdom. Séances Acad. Sci. 1959, 248 (part 2), 2211−2213. (30) Macklen, E. D. Influence of atmosphere on the thermal decomposition of ferrous oxalate dehydrate. J. Inorg. Nucl. Chem. 1967, 29, 1229−1234. (31) Balek, V. Application of inert radioactive gases in the study of solids − Part 2. Thermal decomposition of various iron salts and the preparation of α-Fe2O3 in its active state. J. Mater. Sci. Lett. 1970, 5, 166−170. (32) Boldyrev, V. V. Influence of external surface of the kinetics of thermal decomposition of oxalates. Proc. Tomsk State Univ. 1954, 126, 41−50. (33) Boldyrev, V. V. Temperature dependence of the rate of thermal decomposition of lead, mercury, copper, and iron oxalates. Proc. Tomsk State Univ. 1954, 126, 51−69. (34) Ugay, Ya. A. The TGA study of the decomposition of divalent metals’ oxalates. Russ. J. Gen. Chem. 1954, 24, 1315−1321.

AUTHOR INFORMATION

Corresponding Author

*Tel.: +7-8452-516413; fax: +7-8452-271491; e-mail address: [email protected] (A.V.C.). Funding

We thank the Russian Foundation for Basic Research (project nos. 12-03-31839 and 13-03-00492) for financial support. Notes

The authors declare no competing financial interest.

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