Dissolution of iron oxide using oxalic acid

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Abstract

Iron oxide is the main contaminant of clay and silicate minerals used during the production of high quality ceramics. Its content has to be removed to generally less than 0.1% for achieving the required whiteness of 90% ISO or higher for clay and silicate materials. Oxalate has been used to dissolve iron oxide from various sources. The dissolution is affected by oxalate concentration, solution pH and temperature. The mineral phase is also critical in determining the reaction rate. Hematite is slow to dissolve whereas iron hydroxide and hydroxyoxides such as goethite and lepidocrosite can be easily dissolved. As the dissolution requires a pH controlled in the region 2.5–3.0 for maximum reaction rate, it is essential to create a hydroxide-oxalate mixture for use in the leaching process. The characteristics of NaOH-, KOH- and NH4OH-oxalic acid mixtures were also determined in this study. Due to the precipitation of salts such as Na2C2O4(s) and NaHC2O4(s) the NaOH-oxalic acid could act as pH buffer for the leaching. Such precipitation also reduces the concentration of the free bioxalate, HC2O4 required for the dissolution of iron oxide. KOH behaves the same as NaOH whereas NH4OH precipitates the less stable salt NH4HC2O4(s) which easily re-dissolves forming soluble oxalate species. Ammonium hydroxide is therefore the most suitable reagent that can be used for pH control during the leaching of iron oxide using oxalate. Using STABCAL, several Eh–pH and stability diagrams were developed to explain the dissolution process.

Introduction

The production of high quality ceramics requires the iron oxide content of clay or silicates minerals be lowered to less than 0.1% to achieve an acceptable whiteness (higher than 90% ISO). Iron therefore has to be removed from these minerals by physical, physicochemical or chemical processing before being used. The use of different inorganic and organic acids for dissolving iron compounds has been evaluated in several studies in an attempt to replace the costly high-temperature chlorination technique. Sidhu et al. (1981) evaluated the dissolution of iron oxides and oxyhydroxides in hydrochloric and perchloric acids. Chiarizia and Horwitz (1999) studied the dissolution of goethite in several organic acids belonging to the families of the carboxylic and diphosphonic acids in the presence of reducing agents. Ambikadevi and Lalithambika, (2000) tested several organic acids (such as acetic, formic, citric, ascorbic acid, etc.) and concluded that oxalic acid is the most efficient that can be used to dissolve iron oxide from ceramic minerals.

Oxalic acid was found to be the most efficient acid that can be used for dissolving most iron oxides as it presents a lower risk of contamination of the treated materials after calcinations. It also has a good complexing characteristics and high reducing power, compared to other organic acids. Many researchers have studied the use of oxalic acid to dissolve iron oxide as a result (Veglio et al., 1998, Segal and Sellers, 1984, Jepson, 1988, Panias et al., 1996, Cornell and Schindler, 1987). Biological processes have also been evaluated, most recently by Mandal and Banerjee (2004) who presented their results of a study on the use of Aspergillus niger and their cultural filtrates for removing iron from a China clay. Using oxalic acid, the dissolved iron can be precipitated from the leach solution as ferrous oxalate, which can be further re-processed to form pure hematite by calcination (Taxiarchou et al., 1997a).

Ambikadevi and Lalithambika (2000) found oxalic acid (0.05–0.15 M) to be the best extractant for removing iron from a kaolinite material finely ground to 90% passing 2 μm. The dissolution efficiency was found to increase with acid concentration within the range 0.05–0.15 M studied. Both oxalate and hydrogen ion concentrations were increased in this case. The pH of the leach systems was not controlled nor any measurement given for the solution pH, making it harder to interpret their results. Using a 0.15 M oxalic acid approximately 70% of the iron could be extracted from a kaolinite slurry (20% w/v) containing 0.93% iron oxide (of goethite and hematite phases) at 100 °C within 90 min. The iron oxide concentration in the leachates is equivalent to 1.86 g/L Fe2O3. The high loading and fine grain size of the raw material could be the reasons explaining the faster iron dissolution found in their study compared to others reported by Taxiarchou et al., 1997a, Taxiarchou et al., 1997b.

The presence of Fe2+ was found to significantly enhance the leaching of iron extraction from silica sand at a temperature even as low as 25 °C (Taxiarchou et al., 1997a). Ferrous oxalate however is oxidized quickly by air during the dissolution and in general an induction period of a few hours was observed unless a strong acidic environment (< pH1) or an inert atmosphere is maintained. Maintaining the high level of ferrous oxalate in the leach liquor using an inert gas will enhance the reaction kinetics according to these authors.

Most studies reported that the dissolution of magnetite and goethite by oxalic acid reached a maximum rate at around pH2.7–3.0, outside which range the dissolution rate dropped dramatically (Cornell and Schindler, 1987, Panias et al., 1996). Dissolution of hematite was found to be slower than for magnetite (FeO · Fe2O3) and other hydrated iron oxide such as goethite (α-FeOOH) and lepidocrocite (γ-FeOOH) and iron hydroxide (Fe(OH)3).

The dissolution of iron oxide is believed to take place via a photo-electrochemical reduction process, involving a complicated mechanism of charge transfer between the predominant oxalate species, namely ferric oxalate, Fe(C2O4)33−, ferrous oxalate, Fe(C2O4)22− acting also as an auto-catalyst, and the oxalate ligand on the iron oxide surface (Taxiarchou et al., 1997b, Blesa et al., 1987). In the absence of light the reaction proceeds slowly which complicates the reaction further.

The solution pH governs the distribution of various oxalate ions in the leach system. Below pH 1.2, oxalic acid exists mainly as H2C2O4, whereas HC2O4 is the most predominant species (mole fraction > 0.92) at pH 2.5–3.0. Above pH 4, C2O42− is the predominant species. The speciation of Fe(III) oxalate and Fe(II) oxalate is also governed by pH and total oxalate concentration (Panias et al., 1996). For a solution having pH > 2.5 and an oxalate concentration higher than 0.1 M, the most predominant Fe(III)-oxalate species is Fe(C2O4)33−. At these conditions (pH > 2.5 and oxalate concentration higher than 0.1 M) the predominant Fe(II) complex species is Fe(C2O4)22−.

The dissolution process also has to be optimized with respect to oxalate concentration and pH to minimize the precipitation of ferrous oxalate. On Eh–pH diagrams (Sukhotin and Khentov, 1980) reproduced in Fig. 1, the predominance of FeC2O4(s) is clearly shown for the system containing 0.21 M oxalate (right-sided graph). Without oxalate, Fe2O3 and Fe3O4 will be dissolved in acid forming Fe2+, whereas in the presence of oxalic acid, solid FeC2O4(s) is the predominant species existing over a wide range of pH from acidic zone to pH > 7 in the potential range where reductive dissolution of iron oxides takes place for 0.21 M oxalate. This implies that solid FeC2O4(s) will be finally formed when the oxalate concentration is 0.21 M (as shown in this graph). Unfortunately there is no reference to the concentration of total Fe used for these diagrams, making it difficult for the interpretation of the process involved. As a result, these diagrams however could not be used to explain the fact that iron oxide could finally be dissolved by oxalate. There must be another reaction step involved which causes the solid ferrous oxalate to re-dissolve if formed, or there must be conditions which allow the dissolution to take place, indicating the shortfall of Sukhotin and Khentov's Eh–pH diagrams.

The iron dissolution process therefore takes place via an electrochemical process, summarised below:

Oxidation of oxalate to form carbonic acid or carbon dioxide,HC2O4 = H+ + 2CO2 + 2eReduction of hematite forming Fe(II) oxalate,2H+ + Fe2O3 + 4HC2O4 + 2e = 2Fe (C2O4)22− + 3H2OThe dissolution reaction is therefore:H+ + Fe2O3 + 5HC2O4 = 2Fe (C2O4)22− + 3H2O + 2CO2.

The overall reaction indicates that species involved in the leaching would be hydrogen ions, oxalate and iron oxide particles. At the optimum pH 2.5–3.0 temperature, concentration of oxalate, iron oxide mineralogy and its particle size will determine the reaction kinetics. The charge transfer mechanism could also be assisted by the presence of Fe(II) as experienced in previous studies.

This paper presents results of a leaching study on the use of oxalate to dissolve iron oxide. Industrial clay samples were tested and compared with model and pure iron oxides (hematite and iron rust materials containing iron hydroxide and hydroxyoxides). Equilibrium studies were also conducted to determine the best reagent (NaOH, KOH or NH4OH) that can be used for the control of the reaction pH. Attempts were also made to explain the major steps involved in the dissolution process using STABCAL (Huang, 2006) to produce various Eh–pH and stability diagrams relevant to the process.

Section snippets

Experimental

Leaching experiments were conducted using different iron oxides or iron-containing clay materials. The clay samples were obtained from Haeng Nam Chinaware Ltd, Korea who purchased their raw products from Jangsan Mine. Hematite and iron rust samples were ground and wet screened to obtain the 105–149 μm size range. For the raw clay the − 149 μm size fraction was used for the test, of which the iron content was 1.06% Fe2O3 and contained mainly iron hydroxyoxide (FeOOH) and iron aluminium silicate.

Leaching of various iron oxides and iron-containing clay

All leaching experiments in this study were conducted at 100 °C. The dissolution of the rust material and hematite at different concentrations of oxalic acid is shown in Fig. 1.

Non-hematite iron oxides in the iron rust were found to dissolve faster than hematite, confirming previous studies by other workers. At the stoichiometric ratio (oxalate/iron oxide) of 5:1 as per Eq. (3), a concentration of 0.048–0.05 M represents only 50–60% of the stoichiometric requirement for oxalate. Therefore for

Conclusions

Oxalic acid was used to dissolve iron oxide from hematite, synthetic iron rust (containing iron hydroxide and hydroxyoxide) and a clay material. The critical effect of solution pH was confirmed with optimum dissolution taking place at pH2.5–3.0. A precipitation of ferrous oxalate, either on the oxide surface or in the bulk also affects the level of iron dissolved from hematite where its dissolution was limited to less than 50% at 100 °C. Over 90% of the iron rust material was also dissolved.

Acknowledgement

The authors acknowledge the assistance of Prof. H. H. Huang (Montana Tech, USA) in the development of STABCAL thermodynamic diagrams presented in this study.

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