Dissolution of Iron Oxides Highly Loaded in Oxalic Acid Aqueous Solution for a Potential Application in Iron-Making

Abstract

Oxalic acid has been identified as a sustainable chemical enabling an efficient recovery of target metals from industrial minerals by dissolution. The dissolution process recently has attracted attention as a key reaction in a potential clean iron-making. In this application to efficiently produce a high-purity iron, the dissolution is required to occur in the absence of light, with no addition of other chemical reagents, and to produce high concentration iron oxalate aqueous solution as fast as possible. To reveal the chemistry of iron oxide dissolution for this application, in the present study, the dissolution experiments are carried out under various conditions with a particular focus on the iron oxide highly loaded in the oxalic acid aqueous solution. Highly acidic oxalic acid solution for dissolving the highly loaded iron oxide enabled the production of iron oxalates aqueous solution with the concentration of up to 0.56 mol-Fe/L. Different from conventional studies under diluted conditions with pH control, the dissolution followed a non-reductive mechanism, producing [Fe³⁺HC₂O₄]²⁺ as a dominant iron species, and highly correlated with a concentration of proton in the solution. The experimental results and proposed stoichiometries identified a minimum amount of oxalic acid required for the complete dissolution of iron oxide independently from the concentration and type of loaded iron oxide. Among iron oxides tested (α-Fe₂O₃, FeOOH and Fe₃O₄) as the feedstock, Fe₃O₄ had an advantage in the dissolution rate, but showed a relatively low iron recovery in the solution (80–90%) because of an unavoidable formation of FeC₂O₄·2H₂O precipitates.

1. Introduction

Dissolution of iron oxides such as hematite (α-Fe₂O₃), magnetite (Fe₃O₄), and goethite (α-FeOOH) is an important chemical process in industries to improve properties of industrial minerals (e.g., ore, clay, quartz, soil, and ceramic). During this process, the iron oxides, as impurity metals, are removed by acids.^1,2,3,4,5) Inorganic (HCl and HNO₃) and organic acids (oxalic, acetic, L-ascorbic, and citric acids) have been studied for the iron oxide dissolution.^{1,2,3,4,5,6,7,8)} Among them, oxalic acid has been suggested to be the most promising acid due to its high acid strength, complexing ability, and reducing property.^{8,9,10,11,12)} Moreover, since it can be easily decomposed by calcination, there is no risk of contaminating the treated materials.^5,11)

The dissolution of iron oxides in oxalic acid aqueous solution is generally explained by three different mechanisms: (1) adsorption of oxalate anions on the surface of iron oxide via protonation and complexation, (2) non-reductive dissolution, and (3) reductive dissolution.^13,14) The adsorption of oxalate anions is the initial step, followed by non-reductive dissolution and/or reductive dissolution, depending on operating conditions. The non-reductive dissolution is a simple desorption process for removing only active sites over iron oxides and is more likely to occur at a low pH and high temperatures, which increases the number of active sites. The reductive dissolution consists of induction and autocatalytic periods. The induction period is to generate ferrous ions (Fe²⁺). When the concentration of Fe²⁺ reaches a sufficient level, the autocatalytic period is initiated by the generated catalytic species, [Fe²⁺(C₂O₄)₂]²⁻, accelerating the iron dissolution. When an iron oxide has Fe(II) in the structure (e.g., magnetite), the rate of dissolution is fast because the catalytic anion is directly formed upon adsorption of oxalate anions. The reductive dissolution of iron oxides without Fe(II) (e.g., hematite and goethite) in the structure requires electron transfer from adsorbed oxalate anions to Fe(III) in the induction period, which is the rate-determining step in the dissolution.

The efficiency of dissolution using oxalic acid is affected by several factors.^{6,13,15,16,17,18,19,20)} The pH of aqueous solution remarkably influences the rate of iron dissolution, and studies have revealed the optimum pH in the range of 2.5–3.0.^10,15,16) Several types of thermodynamically stable iron oxalate complex ions under different pH have been reported.¹⁷⁾ In a high pH solution (pH > 3), [Fe²⁺(C₂O₄)₂]²⁻ and [Fe³⁺(C₂O₄)₃]³⁻ are thermodynamically stable, while [Fe³⁺(C₂O₄)₂]⁻ and [Fe³⁺C₂O₄]⁺ are stable at the pH of 1–2. Only [Fe³⁺HC₂O₄]²⁺ presents in a highly acidic solution (pH < 1). Uncomplexed Fe²⁺ is also found in the highly acidic aqueous solution, but uncomplexed Fe³⁺ is unlikely to form in oxalic acid solutions.

Apart from the pH effect, the rate of dissolution increases with the concentration of oxalic acid.²⁾ However, in the case of magnetite dissolution, the effect of oxalic acid concentration is diminished by the richness in autocatalytic species.²¹⁾ The temperature also contributes to the acceleration of dissolution, where reasonable dissolution rates are found above 80°C.^{2,6,13,16,18)} Besides, the ratio between leachate (oxalic acid solution) and loaded iron oxide-containing materials, so-called liquid-solid ratio (L/S ratio), is an important factor for the efficient dissolution.^19,20)

We have recently proposed a new and sustainable iron-making process that starts with the iron oxide dissolution from iron ore as presented in Fig. A1.²²⁾ Iron dissolved mainly as Fe(III) oxalate (Eq. (1) for Fe₂O₃) is reduced to metallic iron in two steps; photochemical reduction and pyrolytic reduction. Unique chemical properties of iron oxalates enable iron-making at low temperatures. To extract iron selectively and produce a high purity iron, the iron oxide dissolution process is operated in the absence of light. For achieving high iron productivity of the process, it is required to produce the solution with a high concentration of iron oxalates as fast as possible. However, previous studies on iron oxide dissolution have been carried out basically at diluted conditions with the addition of buffering agents to control the pH in the optimum range.^{12,14,15,16,23,24,25)} To the best of our knowledge, there has been no research that focused mainly on the dissolution of highly loaded iron oxides. Moreover, the amount of oxalic acid required to completely dissolve iron (molar ratio of oxalic acid to iron loaded in the solution: OxA/Fe) has not been reported regardless of its importance for minimizing the consumption of oxalic acid. The dissolution of iron oxides is unlikely to simply follow the Eq. (1), considering the presence of several types of iron oxalates in the solution as mentioned above.   

$$\mathrm{F}{\mathrm{e}}_2{\mathrm{O}}_3+3{\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4\to \mathrm{F}{\mathrm{e}}_2{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_3+3{\mathrm{H}}_2\mathrm{O}$$(1)

Therefore, in the present study, the dissolution of iron oxides highly loaded in the oxalic acid aqueous solution (up to 1.3 mol-Fe/L) in the absence of light and with no addition of other chemical reagents under atmospheric pressure has been investigated. Hematite (α-Fe₂O₃), goethite (FeOOH) and magnetite (Fe₃O₄) were used as iron precursors typically contained in iron ores. We believe that the findings serve as a guideline for achieving an efficient iron oxide dissolution in the potential iron-making process, and provide further insight into the mechanism of iron oxide dissolution by oxalic acid as well.

2. Materials and Methods

2.1. Materials

α-Fe₂O₃ and Fe₃O₄ were purchased from Sigma-Aldrich, and FeOOH was purchased from Nacalai Tesque. Anhydrous oxalic acid (Wako Pure Chemical) was used as a reagent for the iron dissolution. 1, 10-phenanthroline (Aldrich), ammonium acetate (Wako Pure Chemical), and hydroxylamine (Wako Pure Chemical) were used for quantification of dissolved iron.

2.2. Iron Dissolution

The iron oxide dissolution experiments were carried out in a dark room to protect the solution from light. The irradiation of light, particularly within the tropospheric solar UV-visible region (290–570 nm), causes reduction of Fe(III) oxalate to water-insoluble Fe(II) oxalate.²²⁾ The quantum efficiency of this photochemical reaction is close to unity.²⁶⁾ Upon the light irradiation, the amount of dissolved iron, therefore, readily decreases. This phenomenon is not preferred for the proposed iron-making process, because it decreases the yield of metallic iron on a feedstock iron basis.

A 50 mL of centrifuge tube with a silicone cap was used as a reactor. A needle for syringe was penetrated through the silicone cap for keeping atmospheric pressure and sampling of the solution during the iron dissolution. Various loading amounts of iron oxides (4.7–100 g-iron oxide/L: 0.05–1.3 mol-Fe/L) and oxalic acid concentrations (0.1–1.0 mol/L) were investigated. The iron oxide with particle sizes below 38 μm was added to a hot oxalic acid aqueous solution (47, 75 or 92°C), and then the slurry was stirred for 360 min using an oil bath with a magnetic stirrer. The dissolution temperatures are hereafter denoted by 45, 75, and 90°C, respectively, for the simplification. During the dissolution, 30 μL of the solution was sampled at predetermined times to quantify the dissolved iron by absorption spectroscopy. The volume of sampling amount was small enough to neglect the influence on the analysis of the dissolution. The pH of the solution was not controlled throughout the experiment.

2.3. Characterizations

The concentration of iron dissolved in the solution was determined by absorption spectroscopy with a UV-Vis spectrophotometer (PerkinElmer, Lambda 365). For investigating chemistry of the dissolution in detail, the dissolved iron was analytically categorized into two types, Fe²⁺ species and Fe³⁺ species. For determining the total concentration of Fe²⁺ and Fe³⁺ species, prescribed amounts of hydroxylamine, 1,10-phenanthroline, ammonium acetate, and deionized water as reducing agent, complexing agent, buffering agent, and diluting solution, respectively, were added to the solution immediately after the sampling. In this pretreatment, Fe³⁺ was reduced to Fe²⁺ by the reducing agent, and then all the dissolved iron species formed tris (1,10-phenanthroline) iron (II) ([Fe(phen)₃]²⁺) having a typical absorbance at 510 nm. The concentration of iron as [Fe(phen)₃]²⁺ was determined with spectroscopy using the standard sample. The concentration of Fe²⁺ species was determined by the analysis of the solution prepared in the absence of the reducing agent. The concentration of Fe³⁺ species was calculated by the difference between total Fe and Fe²⁺ concentrations. The concentrations of Fe²⁺ species, Fe³⁺species, and their total in the solution are hereafter denoted by Fe²⁺, Fe³⁺, and total Fe concentrations, respectively. The rate of iron dissolved in the solution from feedstock (Fe dissolution) was calculated from the dissolved iron and iron loaded in the solution. Crystalline structures of solid samples were analyzed by X-ray diffraction, XRD, on a Rigaku TTR-III X-ray diffractometer with Cu Kα radiation at 50 kV and 30 mA.

3. Results and Discussion

3.1. Effect of Iron Oxide Type

Figure 1 shows the experimental results of dissolution of α-Fe₂O₃, FeOOH, and Fe₃O₄ in 0.5 M oxalic acid aqueous solution at 75°C. The loaded iron oxides were 0.33 mol-Fe/L, corresponding to 26.3 g/L, 29.3 g/L, and 25.4 g/L for α-Fe₂O₃, FeOOH, and Fe₃O₄, respectively. As seen from the results, the rate of dissolution depended significantly on the type of iron oxide; Fe₃O₄ was the fastest, followed by α-Fe₂O₃ and then FeOOH. Fe dissolution of Fe₃O₄ reached 80% within 30 min and then remained almost unchanged in 6 h of the run. On the other hand, Fe dissolutions of α-Fe₂O₃ and FeOOH gradually increased and, at 6 h, reached 77% and 71%, respectively. It should be mentioned that, although the dissolution of Fe₃O₄ was fast, the dissolution was incomplete, and the product contained a solid residue that was a precipitated Fe(II) oxalates, the details of which are discussed later. The fast dissolution of Fe₃O₄ was due to the presence of Fe(II) in the structure. Fe²⁺ was dissolved into the solution at an early stage of the dissolution, forming the complex with oxalate anions to work as a type of catalyst for the iron dissolution.^13,14,24,27) These events nearly completed before the first sampling at 15 min, but the occurrence was apparent from the abundance of Fe²⁺ in the solution, compared to those using α-Fe₂O₃ and FeOOH as the feedstock.

Fig. 1.

Time-dependent change of (a) Fe dissolution and (b) Fe²⁺ or Fe³⁺ concentration. Conditions: 75°C, 0.33 mol-Fe/L loaded iron oxide, and 0.5 mol/L oxalic acid. (Online version in color.)

The distribution of dissolved Fe species (Fig. 1(b)) showed that the majority of dissolved iron in all solutions was trivalent iron (Fe³⁺) species. The result indicated that the dissolution mainly followed the non-reductive dissolution, where active sites over iron oxide surface were removed by active hydrogen ions in the solution.^{10,13,23,24,25,28)} This conflicted with a generally accepted mechanism of reductive dissolution for the iron dissolution using organic acids. On the other hand, the iron dissolution in the present study occurred under a highly acidic solution because more oxalic acid was required to dissolve highly loaded iron oxides, compared to general studies. In addition, the dissolution was carried out at a relatively high temperature (75°C) without the exposure to light that promotes the generation of Fe(II) oxalates. Therefore, under the conditions examined in this study, the non-reductive dissolution is suggested to be a dominant pathway, because those conditions contribute to the increase in the number of active sites over iron oxides.^24,29)

Fe(III) oxalate in the solution is supposed to be in the form of [Fe³⁺HC₂O₄]²⁺, having dissociation constant (K_d) = 2.95×10⁻¹⁰, because it is the most stable form in a highly acidic solution of pH < 1.¹⁷⁾ With this in consideration, reactions of α-Fe₂O₃ and FeOOH dissolution are described by Eqs. (2) and (3), respectively. Trace amounts of Fe²⁺ in the solution of α-Fe₂O₃ and FeOOH occurred possibly via the reductive dissolution pathway by electron transfer from oxalate anion,¹³⁾ resulting in the formation of uncomplexed Fe²⁺, as shown in Eqs. (4) and (5), where the symbol, “⟩”, is used to represent species at the surface of iron oxides.   

$$\begin{array}{l}\mathrm{F}{\mathrm{e}}_2{\mathrm{O}}_3+2\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4{}^{–}{}_{(\mathrm{a}\mathrm{q})}+6{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\\ \to 2{\left[\mathrm{F}{\mathrm{e}}^{3+}\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4\right]}^{2+}{}_{(\mathrm{a}\mathrm{q})}+3{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\end{array}$$(2)

  

$$\begin{array}{l}2\mathrm{F}\mathrm{e}\mathrm{O}\mathrm{O}\mathrm{H}+2\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4{}^{–}{}_{(\mathrm{a}\mathrm{q})}+6{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\\ \to 2{\left[\mathrm{F}{\mathrm{e}}^{3+}\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4\right]}^{2+}{}_{(\mathrm{a}\mathrm{q})}+3{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\end{array}$$(3)

  

$$\left[⟩\mathrm{F}{\mathrm{e}}^{\mathrm{I}\mathrm{I}\mathrm{I}}–{\mathrm{C}}_2{\mathrm{O}}_4{}^{2–}\right]\leftrightarrow \left[⟩\mathrm{F}{\mathrm{e}}^{\mathrm{I}\mathrm{I}}–{\mathrm{C}}_2{\mathrm{O}}_4{}^{–}\right]$$(4)

  

$$\begin{array}{l}2\left[⟩\mathrm{F}{\mathrm{e}}^{\mathrm{I}\mathrm{I}}–{\mathrm{C}}_2{\mathrm{O}}_4{}^{–}\right]+2{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\\ \leftrightarrow 2\mathrm{F}{\mathrm{e}}^{2+}{}_{(\mathrm{a}\mathrm{q})}+2\mathrm{C}{\mathrm{O}}_2+{\mathrm{C}}_2{\mathrm{O}}_4{}^{2–}{}_{(\mathrm{a}\mathrm{q})}+2⟩\mathrm{H}\end{array}$$(5)

The dissolution of Fe₃O₄ with oxalic acid has been reported as Eq. (6), where iron is dissolved as [Fe³⁺(C₂O₄)₃]³⁻ and [Fe²⁺(C₂O₄)₂]²⁻ at the ratio of 2.^24,27) However, OxA/Fe of 2.67, assumed in Eq. (6), is much higher than that applied in the present experiment (= 1.5). Considering the most stable chemical forms of dissolved Fe³⁺ and Fe²⁺ as [Fe³⁺HC₂O₄]²⁺and uncomplexed Fe²⁺, respectively, under conditions of this study, the dissolution of Fe₃O₄ is rewritten by Eq. (7). Furthermore, the presence of Fe(II) oxalate precipitate, which has little solubility in water (K_sp = 2 ×10⁻⁷ at 25°C),¹⁷⁾ in the solution indicated the occurrence of a reaction described by Eq. (8). In the concentration profiles of Fig. 1(b), the reaction system, consisting of Eqs. (7) and (8), reached the equilibrium in 30 min, and the concentrations of Fe³⁺ and Fe²⁺ were almost unchanged over the time period thereafter. Fe³⁺/Fe²⁺ ratio in the solution after reaching the equilibrium was roughly consistent with the stoichiometry of Eq. (7).   

$$\begin{array}{l}\mathrm{F}{\mathrm{e}}_3{\mathrm{O}}_4+8\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4{}^{–}{}_{(\mathrm{a}\mathrm{q})}\\ \to 2{\left[\mathrm{F}{\mathrm{e}}^{3+}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_3\right]}^{3–}{}_{(\mathrm{a}\mathrm{q})}+{\left[\mathrm{F}{\mathrm{e}}^{2+}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_2\right]}^{2–}{}_{(\mathrm{a}\mathrm{q})}+4{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\end{array}$$(6)

  

$$\begin{array}{l}\mathrm{F}{\mathrm{e}}_3{\mathrm{O}}_4+2\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4{}^{–}{}_{(\mathrm{a}\mathrm{q})}+8{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\\ \to 2{\left[\mathrm{F}{\mathrm{e}}^{3+}\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4\right]}^{2+}{}_{(\mathrm{a}\mathrm{q})}+\mathrm{F}{\mathrm{e}}^{2+}{}_{(\mathrm{a}\mathrm{q})}+4{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\end{array}$$(7)

  

$$\mathrm{F}{\mathrm{e}}^{2+}{}_{(\mathrm{a}\mathrm{q})}+{\mathrm{C}}_2{\mathrm{O}}_4{}^{2-}{}_{(\mathrm{a}\mathrm{q})}\leftrightarrow \mathrm{F}\mathrm{e}{\mathrm{C}}_2{\mathrm{O}}_{4(\mathrm{s})}$$(8)

3.2. Effect of OxA/Fe on the Dissolution of Iron Oxides

OxA/Fe is an important parameter for enabling a complete iron oxide dissolution while minimizing the consumption of oxalic acid. The experiments of Fig. 1 employed OxA/Fe = 1.5, which was determined from the stoichiometry of Eq. (1). However, the product solutions from all iron oxides contained undissolved iron oxides, indicating that OxA/Fe = 1.5 was insufficient for achieving the complete dissolution. To investigate the effect of OxA on the iron oxide dissolution and involved reactions, the experiments were carried out at various oxalic acid concentrations (0.1–1.0 mol/L) and iron oxides loadings (0.05–1.3 mol-Fe/L), providing the OxA/Fe from 0.4 to 10, at 90°C for 360 min in the absence of light. The dissolution temperature of 90°C was employed for achieving an equilibrium Fe dissolution in the limited dissolution time of 360 min because higher temperature resulted in a faster dissolution of iron oxides as confirmed by Fig. A2. Temperatures higher than 90°C were effective for further facilitating the dissolution, but 90°C was chosen to avoid boiling the solution in the experiment under atmospheric pressure. Figure A3 presents time-dependent changes in Fe dissolution and total Fe concentration. The result showed steady Fe dissolutions at the late stage of experiments, indicating that the conditions of 90°C and 360 min were sufficient for achieving Fe dissolution at the equilibrium.

The data set obtained from the experiments under 167 different conditions (oxalic acid concentration, iron oxide loadings and iron oxide types) are listed in Tables A1, A2, A3. It is to be noted that oxalic acid concentration was limited up to 1.0 mol/L due to its solubility in water at room temperature. It was confirmed from the results that the majority of iron dissolved in the solution presented as Fe³⁺ species. Total Fe concentrations from α-Fe₂O₃, FeOOH, and Fe₃O₄ reached 0.56, 0.55, and 0.50 mol/L, respectively. The data were further analyzed under the assumption that Fe dissolution reaches the equilibrium value in the experiment of 360 min, i.e., Fe dissolution would not change after longer dissolution.

Figure 2 plots Fe dissolution or the ratio between Fe²⁺ and Fe³⁺ species concentrations (Fe²⁺/Fe³⁺) against OxA/Fe. It can be seen from Fig. 2(a) that the Fe dissolution increases linearly with OxA/Fe until it reaches plateaus that are almost 100% for α-Fe₂O₃ and FeOOH, and 80–90% for Fe₃O₄. The linear correlation between Fe dissolution and OxA/Fe, obtained from the plots, is approximated by y = 54.8x with the regression coefficient R² = 0.95, where y and x are Fe dissolution and OxA/Fe, respectively. For α-Fe₂O₃ and FeOOH, Fe dissolution reaches 100% at OxA/Fe = 1.82 (= 100/54.8). This is the required minimum OxA/Fe to achieve the complete dissolution. In the case of Fe₃O₄, because a portion of dissolved iron precipitates as Fe(II) oxalate, the Fe dissolution does not reach 100%, as observed in Fig. 1.

Fig. 2.

Fe dissolution (a) and Fe²⁺/Fe³⁺ (b) plotted against OxA/Fe. The data were obtained from the dissolution experiments at 90°C for 360 min with 0.1–1.0 mol/L oxalic acid and 0.05–1.3 mol-Fe/L loaded iron oxide. (Online version in color.)

As shown in Fig. 2(b), Fe²⁺/Fe³⁺ decreased with OxA/Fe, irrespective of the type of iron oxides. This was because of the stability of Fe(III) oxalates, which, compared to Fe(II) oxalates, increased with acidity.¹⁷⁾ Under the conditions employed in this experiment, the initial pH of oxalic acid aqueous solution was generally below 1, excepting 0.1 mol/L oxalic acid (pH = 1.23). Therefore, the iron dissolution was believed to follow the non-reductive mechanism, as evidenced by the presence of Fe(III) oxalates as the main Fe species for all the solutions after dissolution. Fe²⁺/Fe³⁺ was thermodynamically determined by the chemical equilibrium, resulting in more abundance of Fe³⁺ with OxA/Fe.

According to the mechanism of non-reductive dissolution, the most important factor influencing the dissolution is active hydrogen ions, protons (H⁺), generated from the dissociation of oxalic acid because protons create active sites over iron oxides surface by protonation. In Eqs. (2), (3), and (7), H⁺ required for obtaining 1 mol of [Fe³⁺HC₂O₄]²⁺ is 3 mol or 4 mol. Oxalic acid dissociates in two steps (Eqs. (9) and (10)) and potentially provides 2 mol of H⁺ per mol-OxA. When oxalic acid is fully dissociated, the required minimum of OxA/Fe for the complete Fe dissolution is 1.5 as presented by Eq. (1). However, the complete Fe dissolution at 90°C, as shown in Fig. 2(a), needed more OxA/Fe, 1.82. This was attributed to the incomplete dissociation of oxalic acid. Using the dissociation constants, H⁺ concentrations in 0.1–1.0 mol/L oxalic acid aqueous solution at 25°C was calculated to be 0.07–0.24 mol/L. The dissociation of oxalic acid was temperature-dependent, resulting in higher H⁺ concentrations in the range of 0.09–0.27 mol/L at 90°C.   

$${\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_{4(\mathrm{a}\mathrm{q})}\leftrightarrow \mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4{}^{-}{}_{(\mathrm{a}\mathrm{q})}+{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\hspace{1em}(K_{\mathrm{a}1}=5.9\times {10}^{-2})$$(9)

  

$$\mathrm{H}{\mathrm{C}}_2{\mathrm{O}}_4{}^{–}{}_{(\mathrm{a}\mathrm{q})}\leftrightarrow {\mathrm{C}}_2{\mathrm{O}}_4{}^{2–}{}_{(\mathrm{a}\mathrm{q})}+{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\hspace{1em}(K_{\mathrm{a}2}=6.2\times {10}^{–5})$$(10)

To correlate H⁺ concentration with Fe dissolution, Eqs. (2) and (3) were employed, and Fe dissolution was plotted against 3H⁺/Fe in Fig. 3, where H⁺ was a molar concentration of proton generated by dissociation of oxalic acid at 90°C, and Fe was a molar concentration of Fe loaded as iron oxides. As with the case of plots against OxA/Fe, Fe dissolution increased in proportion to 3H⁺/Fe, irrespective of iron oxide types, before reaching the plateau values. The linear relationship had a slope of 54.3, which was close to that found in Fig. 2(a). Therefore, it was concluded that the reason for the required minimum OxA/Fe (= 1.82) to complete the Fe dissolution was the requirement for dissociated protons to create a sufficient number of active sites over the iron oxides surface by protonation. It is worth mentioning that this finding corresponds to previous reports stating that when the dissolution is assisted by proton (proton-assisted dissolution or protonation-induced dissolution), three protons are required to detach Fe(III) into the solution.^30,31,32,33) One proton is adsorbed on the surface to generate a positively charged site (Eq. (11)), followed by a weakening of the Fe-O bond with the additional two protons to dissolve Fe(III) into the solution (Eq. (12)). The present experimental results revealed that this pathway dominates the dissolution of highly loaded α-Fe₂O₃, FeOOH, and Fe₃O₄ under the conditions employed.   

$$⟩\mathrm{F}{\mathrm{e}}^{\mathrm{I}\mathrm{I}\mathrm{I}}\mathrm{O}\mathrm{O}{\mathrm{H}}_{(\mathrm{s})}+{\mathrm{H}}^{+}\to ⟩\mathrm{F}{\mathrm{e}}^{\mathrm{I}\mathrm{I}\mathrm{I}}{\left(\mathrm{O}\mathrm{H}\right)}_2{}^{+}{}_{(\mathrm{s})}$$(11)

  

$$⟩\mathrm{F}{\mathrm{e}}^{\mathrm{I}\mathrm{I}\mathrm{I}}{\left(\mathrm{O}\mathrm{H}\right)}_2{}^{+}{}_{(\mathrm{s})}+2{\mathrm{H}}^{+}\to \mathrm{F}{\mathrm{e}}^{3+}{}_{(\mathrm{a}\mathrm{q})}+2{\mathrm{H}}_2\mathrm{O}$$(12)

Fig. 3.

Fe dissolution plotted against H⁺ concentration per Fe loaded in the solution, 3H⁺/Fe, derived from Eqs. (2) and (3). The data were obtained from the dissolution experiments at 90°C for 360 min with 0.1–1.0 mol/L oxalic acid and 0.05–1.3 mol-Fe/L loaded iron oxide. (Online version in color.)

3.3. Fe Dissolution at a Constant Oxalic Acid Concentration

To study and clarify the influence of OxA/Fe in more detail, a part of the data set in Tables A1, A2, A3 was extracted and replotted in Fig. 4. Figure 4(a) presents the relationship between Fe dissolution and OxA/Fe at 0.5 mol/L of oxalic acid concentration. The loaded amount of iron oxides ranged from 10 to 100 g/L (0.1–1.3 mol-Fe/L). The plot confirmed that OxA/Fe = 1.82 was required for the complete dissolution of iron oxides at 90°C. Figure 4(b) presents the pH of the solution after dissolution. The pH of 0.5 mol/L fresh oxalic acid aqueous solution was 0.74. When the dissolution of iron oxides was completed (OxA/Fe ≥ 1.82), the pH of the solution was lower than 0.72. This was likely because the consumption of bioxalate anion to form [Fe³⁺HC₂O₄]²⁺ during the dissolution caused more dissociation of oxalic acid, contributing to the decrease in pH of the final solution. On the other hand, at the OxA/Fe below 1.82, the dissolution is incomplete, and the pH gradually increases in the range of 1–2 with the decrease of OxA/Fe because of the insufficient oxalic acid loaded in the solution. The difference in pH of the resulting solution between below and above OxA/Fe = 1.82 indicated the occurrence of other chemical events in addition to Eqs. (2), (3), (4), (5), (7), and (8). At the pH of 1–2, the stable Fe(III) oxalates are [Fe³⁺(C₂O₄)₂]⁻ (K_d = 6.31×10⁻¹⁷) and [Fe³⁺C₂O₄]⁺ (K_d = 3.98×10⁻¹⁰).¹⁷⁾ Considering higher stability of [Fe³⁺(C₂O₄)₂]^–, Eqs. (13), (14), (15) are supposed to be involved in the dissolution at the oxalic acid loadings insufficient for the complete dissolution of iron oxides. Fe(III) oxalate ([Fe³⁺HC₂O₄]²⁺) in Eqs. (2), (3), and (7) forms from an equivalent mole of oxalic acid, but [Fe³⁺(C₂O₄)₂]⁻ in Eqs. (13), (14), (15) requires two moles of oxalate, which reasonably explains the shortage of oxalates at OxA/Fe < 1.82, causing the high pH of the solution.   

$$\begin{array}{l}\mathrm{F}{\mathrm{e}}_2{\mathrm{O}}_3+4{\mathrm{C}}_2{\mathrm{O}}_4{}^{2–}{}_{(\mathrm{a}\mathrm{q})}+6{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\\ \to 2{\left[\mathrm{F}{\mathrm{e}}^{3+}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_2\right]}^{–}{}_{(\mathrm{a}\mathrm{q})}+3{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\end{array}$$(13)

  

$$\begin{array}{l}2\mathrm{F}\mathrm{e}\mathrm{O}\mathrm{O}\mathrm{H}+4{\mathrm{C}}_2{\mathrm{O}}_4{}^{2–}{}_{(\mathrm{a}\mathrm{q})}+6{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\\ \to 2{\left[\mathrm{F}{\mathrm{e}}^{3+}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_2\right]}^{–}{}_{(\mathrm{a}\mathrm{q})}+4{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\end{array}$$(14)

  

$$\begin{array}{l}\mathrm{F}{\mathrm{e}}_3{\mathrm{O}}_4+4{\mathrm{C}}_2{\mathrm{O}}_4{}^{2–}{}_{(\mathrm{a}\mathrm{q})}+8{\mathrm{H}}^{+}{}_{(\mathrm{a}\mathrm{q})}\\ \to 2{\left[\mathrm{F}{\mathrm{e}}^{3+}{\left({\mathrm{C}}_2{\mathrm{O}}_4\right)}_2\right]}^{–}{}_{(\mathrm{a}\mathrm{q})}+\mathrm{F}{\mathrm{e}}^{2+}{}_{(\mathrm{a}\mathrm{q})}+4{\mathrm{H}}_2{\mathrm{O}}_{(\mathrm{l})}\end{array}$$(15)

Fig. 4.

Fe dissolution and pH of the solution after dissolution plotted against OxA/Fe in the experiments with 0.5 mol/L oxalic acid and 0.05–1.3 mol-Fe/L loaded iron oxide at 90°C for 360 min. (Online version in color.)

Figure 5 plots the total Fe, Fe²⁺, and Fe³⁺ concentrations against the Fe loaded as iron oxides in 0.5 mol/L oxalic acid. As shown in Fig. 5(a), the total Fe concentration increased with the loaded Fe and then remained constant at around 0.3 mol-Fe/L, when the loaded Fe was over 0.3 mol-Fe/L. This result is explained by the required minimum OxA/Fe = 1.82, which corresponds to the maximum Fe concentration of 0.27 mol/L in 0.5 mol/L oxalic acid. Thus, when the loaded Fe was over 0.27 mol/L, the dissolution reached its maximum, and further increases in Fe loading did not result in the increase in total Fe concentration because of the shortage of oxalic acid. In other words, the excess loadings of Fe did not affect the chemistry of dissolution. It should be noted that the maximum total Fe concentration was not caused by the solubility of iron oxalates in water. In a preliminary experiment, the solubility of Fe(III) oxalate (Fe₂(C₂O₄)₃·6H₂O) was confirmed to be over 1.0 mol-Fe/L. Indeed, Tables A1, A2, A3 contained experimental data that showed total Fe concentrations over 0.27 mol/L. It was seen from Fig. 5(b) that Fe³⁺/Fe²⁺ ratio in all cases was around 2, and the ratio was unchanged by the loaded Fe at the full dissolution with 0.5 mol/L oxalic acid. This result agrees with Fig. 2(b), where Fe²⁺/Fe³⁺ approaches around 0.5 with a decrease in OxA/Fe.

Fig. 5.

Total Fe, Fe²⁺ and Fe³⁺ concentrations plotted against iron oxide loaded in the solution in the experiments with 0.5 mol/L oxalic acid and 0.05–1.3 mol-Fe/L loaded iron oxide at 90°C for 360 min. (Online version in color.)

3.4. Analysis of Precipitates in the Dissolution of Fe₃O₄

Three types of iron oxides, examined in this study, showed similar trends in most experimental results, but the faster dissolution rate and lower Fe dissolution of Fe₃O₄ was the exception. The difference was caused by the presence of Fe(II) in the structure of Fe₃O₄. A facile and quick release of catalytic Fe(II) oxalate results in the fast dissolution of overall Fe₃O₄ as reported in literature.²¹⁾ On the other hand, the formation of precipitates was assumed by Eq. (8) in this study. To confirm the occurrence of Eq. (8), precipitates recovered from the dissolution at different OxA/Fe at 90°C were analyzed by XRD. As shown in Fig. 6, all the precipitates had a monoclinic structure of Humboldtine, FeC₂O₄·2H₂O, although higher crystallinity was observed for the precipitates generated at OxA > 2.

Fig. 6.

XRD patterns of precipitates obtained from the dissolution of Fe₃O₄ at 90°C for 360 min with different OxA/Fe. (Online version in color.)

In our scheme of iron-making using oxalic acid,²²⁾ FeC₂O₄·2H₂O is prepared by photochemical reduction of Fe(III) oxalate in the solution and used as a direct feedstock for producing metallic iron by its pyrolytic reduction. The formation of FeC₂O₄·2H₂O without the need of a photochemical reduction step and as its fast dissolution rate seems like an advantage of Fe₃O₄ over α-Fe₂O₃ and FeOOH. However, when Fe₃O₄ contained in natural iron ore is used as feedstock, a difficulty is found in the separation of precipitated FeC₂O₄·2H₂O from the dissolution residue consisting mainly of metals other than iron. The use of FeC₂O₄·2H₂O mixed with other metals for pyrolytic reduction results in the production of low-quality iron with impurities. When the precipitates are not considered as the feedstock of metallic iron in consideration of this issue, 10–20% of the iron in Fe₃O₄ is lost in the dissolution. On the other hand, as experimentally evidenced in the present study, Fe₂O₃ and FeOOH are recovered, with little loss in the dissolution, as Fe(III) oxalates and Fe²⁺ in the aqueous solution, which are subjected to photochemical reduction to produce FeC₂O₄·2H₂O with a near-complete recovery.

4. Conclusions

The dissolution of iron oxides was studied to reveal the chemistry under various conditions with a particular purpose for its application to iron-making. The experimental results demonstrated that Fe concentration of over 0.5 mol/L could be achieved by the dissolution. Among iron oxides tested, Fe₃O₄ showed a distinguished fast dissolution, completing within 30 min even at 75°C, due to the facile and fast release of catalytic Fe(II) oxalate. A drawback of Fe₃O₄ as a feedstock of the iron-making was the formation of FeC₂O₄·2H₂O precipitates, which accounted for 10–20% of the iron in the feedstock, during the dissolution. Near-complete dissolution was confirmed for Fe₂O₃ and FeOOH. On the other hand, highly acidic oxalic acid for preparing the high concentration iron oxalate solution resulted in the predominant occurrence of non-reductive dissolution independently of the type of iron oxides. Fe dissolution under various conditions could be described by a single line when plotted against a factor represented by proton concentration, which was indicated by the occurrence of non-reductive dissolution. This correlation also revealed the required minimum OxA/Fe for achieving the maximum Fe dissolution. The detailed analysis of dissolved iron species and suggested stoichiometries evidenced that Fe³⁺ and Fe²⁺ presented mainly as [Fe³⁺HC₂O₄]²⁺ and Fe²⁺ in the solution, respectively, but other iron species such as [Fe³⁺(C₂O₄)₂]⁻ was possibly involved in the solution when oxalates were not sufficiently provided, i.e., at OxA/Fe lower than the required minimum.

Acknowledgements

This work was financially supported by New Energy and Industrial Technology Development Organization (NEDO), Japan, for the Feasibility Study Program on Uncharted Territory Challenge 2050. The authors are also grateful to the Cooperative Research Program of Network Joint Research Center for Materials and Devices that has been supported by the Ministry of Education, Culture, Sports, Science, and Technology (MEXT), Japan.

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Appendix:

Concept of the proposed sustainable iron-making process (Fig. A1), time-dependent change of Fe dissolution at different temperatures (Fig. A2), Time-dependent change of Fe dissolution and total Fe concentration at 90°C for 360 min (Fig. A3), and dissolution of α-Fe₂O₃ (Table A1), FeOOH (Table A2), or Fe₃O₄ (Table A3) with different feedstock loadings at 90°C for 360 min.

Fig. A1.

Concept of the proposed sustainable iron-making process. OA: oxalic acid. Reprinted with permission from ref. 22 (Santawaja et al. 2020). Copyright 2020 American Chemical Society. (Online version in color.)

Fig. A2.

Time-dependent change of Fe dissolution at different temperatures. Conditions: 0.33 mol-Fe/L loaded iron oxide, 0.5 mol/L oxalic acid, and 180 min. (Online version in color.)

Fig. A3.

Time-dependent change of Fe dissolution and total Fe concentration at 90°C for 360 min. Conditions: 0.33 mol-Fe/L loaded iron oxide and 0.67 mol/L oxalic acid. (Online version in color.)

Table A1. Dissolution of α-Fe₂O₃ with different feedstock loadings at 90°C for 360 min.

Feedstock			Fe dissa) (%)	Dissolved Fe (mol/L)			Feedstock			Fe dissa) (%)	Dissolved Fe (mol/L)
Fe (mol/L)	OxA (mol/L)	OxA/Fe (–)					Fe (mol/L)	OxA (mol/L)	OxA/Fe (–)
				Total	Fe3+	Fe2+					Total	Fe3+	Fe2+
0.07	0.10	1.5	74	0.05	0.04	0.01	0.42	0.50	1.2	68	0.28	0.20	0.08
0.10	0.20	2.0	95	0.09	0.08	0.02	0.42	0.75	1.8	88	0.37	0.29	0.07
0.10	0.30	3.0	99	0.10	0.09	0.01	0.50	0.50	1.0	61	0.30	0.21	0.10
0.10	0.40	4.0	100	0.10	0.09	0.01	0.50	0.70	1.4	81	0.41	0.29	0.11
0.10	0.50	5.0	100	0.10	0.09	0.01	0.50	0.75	1.5	84	0.42	0.31	0.11
0.10	0.60	6.0	99	0.10	0.10	0.00	0.50	1.00	2.0	96	0.48	0.37	0.11
0.10	0.80	8.0	100	0.10	0.10	0.00	0.53	0.80	1.5	83	0.44	0.32	0.12
0.10	1.00	10.0	99	0.10	0.10	0.00	0.58	0.75	1.3	70	0.41	0.30	0.11
0.13	0.50	3.9	96	0.12	0.11	0.01	0.60	0.90	1.5	78	0.47	0.37	0.10
0.13	0.75	5.9	96	0.12	0.11	0.01	0.63	0.50	0.8	49	0.30	0.21	0.10
0.13	1.00	8.0	95	0.12	0.11	0.01	0.63	0.75	1.2	69	0.44	0.32	0.12
0.14	0.20	1.5	82	0.11	0.08	0.03	0.63	1.00	1.6	79	0.49	0.39	0.11
0.17	0.50	3.0	99	0.17	0.14	0.02	0.67	1.00	1.5	78	0.52	0.40	0.13
0.20	0.30	1.5	83	0.17	0.13	0.04	0.75	0.50	0.7	41	0.31	0.22	0.09
0.20	0.50	2.5	100	0.20	0.17	0.03	0.75	0.75	1.0	59	0.44	0.32	0.12
0.22	0.50	2.2	100	0.22	0.17	0.05	0.75	1.00	1.3	72	0.54	0.40	0.15
0.25	0.50	2.0	89	0.22	0.19	0.04	0.88	0.50	0.6	36	0.31	0.22	0.10
0.25	0.75	3.0	95	0.24	0.21	0.03	0.88	0.75	0.9	51	0.45	0.32	0.13
0.25	1.00	4.0	94	0.23	0.21	0.03	0.88	1.00	1.1	63	0.55	0.39	0.16
0.27	0.40	1.5	86	0.23	0.16	0.07	1.00	0.75	0.7	45	0.45	0.32	0.13
0.33	0.50	1.5	87	0.29	0.20	0.09	1.00	1.00	1.0	55	0.56	0.39	0.16
0.33	0.67	2.0	100	0.32	0.25	0.08	1.01	0.50	0.5	31	0.32	0.21	0.10
0.33	0.75	2.2	99	0.33	0.24	0.08	1.13	0.50	0.4	27	0.31	0.21	0.10
0.34	0.67	2.0	98	0.33	0.26	0.07	1.13	0.75	0.7	39	0.44	0.31	0.13
0.38	1.00	2.7	90	0.34	0.29	0.05	1.13	1.00	0.9	48	0.54	0.39	0.15
0.38	0.50	1.3	74	0.28	0.21	0.07	1.25	0.75	0.6	36	0.45	0.32	0.13
0.38	0.75	2.0	88	0.33	0.27	0.06	1.25	1.00	0.8	43	0.54	0.39	0.15
0.40	0.60	1.5	83	0.33	0.24	0.10	1.26	0.50	0.4	25	0.31	0.21	0.11

^a)  Fe dissolution.

Table A2. Dissolution of FeOOH with different feedstock loadings at 90°C for 360 min.

Feedstock			Fe dissa) (%)	Dissolved Fe (mol/L)			Feedstock			Fe dissa) (%)	Dissolved Fe (mol/L)
Fe (mol/L)	OxA (mol/L)	OxA/Fe (–)					Fe (mol/L)	OxA (mol/L)	OxA/Fe (–)
				Total	Fe3+	Fe2+					Total	Fe3+	Fe2+
0.05	0.20	4.0	97	0.05	0.05	0.00	0.45	0.75	1.7	86	0.39	0.29	0.10
0.05	0.30	6.0	99	0.05	0.05	0.00	0.45	1.00	2.2	99	0.45	0.36	0.09
0.05	0.40	8.0	99	0.05	0.05	0.00	0.47	0.70	1.5	88	0.41	0.29	0.12
0.05	0.50	10.0	98	0.05	0.05	0.00	0.49	0.50	1.0	64	0.31	0.21	0.11
0.07	0.10	1.5	82	0.06	0.04	0.01	0.50	0.50	1.0	62	0.31	0.21	0.10
0.11	0.50	4.4	97	0.11	0.10	0.01	0.50	0.75	1.5	86	0.43	0.31	0.12
0.11	0.75	6.7	98	0.11	0.10	0.01	0.50	0.75	1.5	91	0.46	0.32	0.14
0.11	1.00	8.8	100	0.11	0.11	0.01	0.50	1.00	2.0	97	0.49	0.37	0.11
0.13	0.20	1.5	88	0.12	0.08	0.03	0.53	0.80	1.5	86	0.46	0.32	0.14
0.17	0.50	3.0	100	0.17	0.14	0.02	0.56	0.50	0.9	51	0.29	0.20	0.09
0.20	0.30	1.5	89	0.18	0.12	0.06	0.56	1.00	1.8	98	0.55	0.42	0.13
0.20	0.50	2.5	100	0.20	0.17	0.04	0.60	0.90	1.5	86	0.52	0.37	0.15
0.22	0.50	2.2	100	0.22	0.17	0.05	0.67	1.00	1.5	81	0.54	0.40	0.14
0.23	0.50	2.2	98	0.22	0.17	0.05	0.68	0.50	0.7	44	0.30	0.19	0.11
0.23	0.75	3.3	97	0.22	0.19	0.03	0.68	0.75	1.1	68	0.46	0.32	0.14
0.23	1.00	4.4	100	0.23	0.21	0.02	0.68	1.00	1.5	81	0.55	0.40	0.15
0.25	0.50	2.0	100	0.25	0.20	0.05	0.79	0.50	0.6	37	0.29	0.19	0.10
0.27	0.40	1.5	89	0.24	0.16	0.08	0.79	0.75	1.0	56	0.44	0.32	0.12
0.30	0.50	1.7	96	0.29	0.21	0.08	0.79	1.00	1.3	68	0.54	0.39	0.14
0.33	0.50	1.5	91	0.30	0.21	0.10	0.90	0.50	0.6	32	0.29	0.19	0.10
0.33	0.67	2.0	100	0.33	0.29	0.05	0.90	0.75	0.8	48	0.43	0.32	0.11
0.33	0.75	2.2	96	0.32	0.24	0.08	0.90	1.00	1.1	58	0.52	0.40	0.12
0.34	0.50	1.5	82	0.28	0.19	0.09	1.01	0.75	0.7	40	0.40	0.30	0.10
0.34	0.75	2.2	93	0.31	0.25	0.06	1.01	1.00	1.0	49	0.49	0.38	0.12
0.34	1.00	3.0	100	0.34	0.29	0.04	1.02	0.50	0.5	29	0.30	0.20	0.10
0.38	0.75	2.0	94	0.35	0.28	0.07	1.13	0.50	0.4	25	0.28	0.19	0.09
0.40	0.60	1.5	88	0.35	0.25	0.10	1.13	0.75	0.7	36	0.40	0.30	0.10
0.45	0.50	1.1	62	0.28	0.19	0.09	1.13	1.00	0.9	43	0.49	0.37	0.11

^a)  Fe dissolution.

Table A3. Dissolution of Fe₃O₄ with different feedstock loadings at 90°C for 360 min.

Feedstock			Fe dissa) (%)	Dissolved Fe (mol/L)			Feedstock			Fe dissa) (%)	Dissolved Fe (mol/L)
Fe (mol/L)	OxA (mol/L)	OxA/Fe (–)					Fe (mol/L)	OxA (mol/L)	OxA/Fe (–)
				Total	Fe3+	Fe2+					Total	Fe3+	Fe2+
0.07	0.10	1.5	94	0.06	0.04	0.02	0.50	0.50	1.0	55	0.27	0.19	0.09
0.13	0.20	1.5	94	0.13	0.09	0.04	0.50	0.75	1.5	68	0.34	0.25	0.09
0.13	0.50	3.8	85	0.11	0.09	0.02	0.50	1.00	2.0	71	0.36	0.29	0.07
0.13	0.75	5.8	90	0.12	0.11	0.01	0.52	0.50	1.0	58	0.30	0.20	0.10
0.13	1.00	7.6	91	0.12	0.11	0.01	0.52	0.75	1.4	70	0.36	0.26	0.10
0.15	0.20	1.3	81	0.12	0.08	0.04	0.52	1.00	1.9	74	0.38	0.32	0.06
0.15	0.25	1.7	95	0.14	0.10	0.04	0.53	0.80	1.5	72	0.39	0.29	0.09
0.15	0.30	2.0	88	0.13	0.11	0.03	0.60	0.90	1.5	70	0.42	0.32	0.10
0.15	0.40	2.7	86	0.13	0.11	0.02	0.65	0.50	0.8	45	0.29	0.19	0.10
0.15	0.50	3.3	87	0.13	0.12	0.01	0.65	0.75	1.2	61	0.40	0.28	0.12
0.15	0.60	4.0	88	0.13	0.12	0.01	0.65	1.00	1.5	70	0.45	0.35	0.10
0.15	0.80	5.3	88	0.13	0.12	0.01	0.67	1.00	1.5	68	0.46	0.35	0.11
0.15	1.00	6.7	87	0.13	0.13	0.00	0.78	0.50	0.6	38	0.30	0.20	0.10
0.17	0.50	3.0	83	0.14	0.12	0.02	0.78	0.75	1.0	51	0.40	0.28	0.12
0.20	0.30	1.5	82	0.16	0.11	0.05	0.78	1.00	1.3	64	0.50	0.34	0.15
0.20	0.50	2.5	82	0.16	0.14	0.03	0.91	0.50	0.6	32	0.29	0.19	0.10
0.25	0.50	2.0	83	0.21	0.17	0.04	0.91	0.75	0.8	41	0.37	0.26	0.12
0.26	0.50	1.9	80	0.21	0.16	0.04	0.91	1.00	1.1	55	0.50	0.34	0.16
0.26	0.75	2.9	86	0.23	0.20	0.03	1.04	0.50	0.5	28	0.29	0.19	0.10
0.26	1.00	3.8	85	0.22	0.20	0.02	1.04	0.75	0.7	37	0.38	0.25	0.13
0.27	0.40	1.5	83	0.22	0.16	0.06	1.04	1.00	1.0	49	0.50	0.34	0.17
0.33	0.50	1.5	81	0.27	0.20	0.07	1.17	0.50	0.4	25	0.29	0.19	0.10
0.33	0.67	2.0	78	0.26	0.21	0.05	1.17	0.75	0.6	33	0.38	0.26	0.13
0.39	0.50	1.3	66	0.26	0.17	0.08	1.17	1.00	0.9	43	0.50	0.33	0.17
0.39	0.75	1.9	79	0.31	0.25	0.06	1.30	0.50	0.4	23	0.29	0.19	0.10
0.39	1.00	2.6	80	0.31	0.27	0.04	1.30	0.75	0.6	29	0.37	0.25	0.12
0.40	0.60	1.5	80	0.32	0.24	0.08	1.30	1.00	0.8	35	0.46	0.31	0.15
0.47	0.70	1.5	74	0.35	0.26	0.09

^a)  Fe dissolution.