Mechanisms of dissolution of iron oxides in aqueous oxalic acid ...

hydrometallurgy ELSEVIER

Hydrometallurgy42 (1996) 257-265

Mechanisms of dissolution of iron oxides in aqueous oxalic acid solutions D. Panias, M. Taxiarchou, I. Paspaliaris, A. Kontopoulos * Laboratory of Metallurgy, National Technical University of Athens, P.O. Box 64056, GR-157 80 Zographos, Greece

Received 20 April 1995; accepted 30 October 1995

Abstract The dissolution of pure iron oxides by organic acids has been extensively reviewed. The mechanism of dissolution comprises three distinct steps: (1) adsorption of organic ligands on the iron oxide surface; (2) non-reductive dissolution; and (3) reductive dissolution. Reductive dissolution involves two stages: an induction period and an autocatalytic period. The overall dissolution process is affected by the pH of the initial solution, temperature, the exposure of solution to UV radiation and the addition of bivalent iron in the initial solution.

1. Introduction The dissolution of metal oxides is an important process in several fields, such as hydrometallurgy, the passivity of metals and cleaning of metal surfaces. Iron oxides are the most studied of the various metal oxides, due to their wide occurrence in natural systems. Many experimental studies have been reported in the literature concerning the dissolution of pure iron oxides by means of inorganic or organic acids. All studies aimed to identify: 1. the most promising acids (organic and inorganic) as solvent agents; 2. the conditions that intensify their action; 3. the probable mechanisms prevailing during the dissolution process. In hydrometallurgical processes, bacterially produced organic acids have been proposed as an alternative, and probably less expensive, leaching agent. Organic acids such as oxalic, citric, ascorbic, acetic, fumaric and tartaric acid, have been screened for their

* Corresponding author. E-mail: [email protected]. 0304-386X/96/$15.00 Copyright © 1996 Elsevier Science B.V. All rights reserved. SSDI 0304-386X(95)00104-2

258

D. Panias et al. / Hydrometallurgy 42 (1996) 257-265

ability to solubilise iron and other metal oxides. Of the above organic acids, oxalic [1-3], citric [4] and ascorbic [5] are the most used carboxylic acids, due to their effectiveness as solvent reagents. Iminodiacetic acid (IDA), N-hydroxyethyl-iminoacetic acid (HIDA) and ethylenediaminotetraacetic acid (EDTA) have also been studied [6,7] as solvent reagents in aqueous solutions with satisfactory results. The present review aims to collect literature data concerning iron oxide dissolution in organic solvents and is especially focused on the mechanisms prevailing in this process.

2. Mechanism of dissolution of iron oxides in organic acids Numerous kinetic studies have been performed on the dissolution of iron oxides with organic acids in order to reveal the reaction mechanism. Most of them have been conducted with pure iron oxides or synthetic ferrites [8-12] rather than with iron-bearing minerals. In such systems interference by other metallic compounds on the dissolution process is excluded and attention to the real chemical pathways is given. The dissolution mechanism involves three different processes taking place simultaneously: 1. adsorption of organic ligands from the solution on the system interface; 2. non-reductive dissolution [6,10,13]; 3. reductive dissolution [1-3,5-7,10,14]. 2.1. Adsorption o f organic ligands on the system interface

It is generally accepted that the first step of the dissolution reaction is the adsorption of carboxylic acid on the surface of iron oxides. When an iron oxide particle is suspended in an acidic solution an electrical double layer [15-20] is established on the system interface: H , L ~ nH+ + L "-

Acid ionisation

)Fem-O + H +~ )Fem-O... H +

(1)

Protonation of oxygen

(2)

The surface of the oxide can behave as a Lewis base (electron pair donor) according to Eq. (2) and, consequently, the system interface appears positively charged. At a given pH value of the initial solution, the higher the ionic strength of the solution, the higher the surface charge associated with adsorbed protons [19]. The above pathway plays an important role in oxide dissolution as the metal-oxygen bond is loosened due to the protonation process. The surface hydroxyl groups (-OH) become active sites for the subsequent adsorption of organic ligands [1-4,13,19,21-24], as described by: )Fern-OH + + L"- + H + ~ [)FellI-L] -(n-2) q- H20

Surface complexation

(3)

The above is a competitive reaction between active surface hydroxyl groups and anions from the acidic solution, such as oxalates, citrates and ascorbates. The greater the stability of [FenI-L] -(~ -2) complexes, the greater the yield of Eq. (3). Generally, as the chelating capacity of a ligand increases, the stability of [Fern-L] -~" -2) complexes also increases and the above reaction shifts to the right.

259

D. Panias et al. /Hydrometallurgy 42 (1996) 257-265 100"/o

100% [#1 0

0

2] 2-

6O%

"~

4O%

t~

40*/0

20%

1~

20%

o

eO*/,

o

0%

0%

0

4

2

6

2

3

4

pH

pH

Fig. 1. Speciation in bulk solution as a function of pH for (left) oxalic acid and (right) Fe 2+/oxalic acid [1].

2.1.1. Adsorption of oxalates on the system interface A pH-dependent equilibrium is established between the two different surface > Fem-oxalate complexes [2,3,24]: o

OH

IIC ~ C'-I 0 ife--

O~

~-'-- 0

e--

0 /

C---- 0

~Pe-- 0 /

+H + + HC:Oi

OH c--c__--o

II o

The speciation of oxalate solutions as a function of pH is shown in Fig. 1. This figure also shows the speciation of Fe2+/oxalic acid solution. 2.2. Non-reductive dissolution The non-reductive dissolution pathway is a simple desorption process. It involves the desorption of adsorbed surface ferric complex ions and their transfer to the acidic solution: [)FeIII_L] -(n-2) + H + ~ - [Fe 3+-LJ(,q) 1-(n-3) + )H -

(4)

The non-reductive dissolution mechanism removes only the more reactive sites of the oxide surface. The number of the latter increases with decreasing pH and increasing temperature [6,13]. As observed in the potential energy-reaction coordinate diagram (Fig. 2), the desorption process is characterised by a high activation energy. As a result, non-reductive dissolution at low temperature is not an operative pathway. As the temperature is raised, the above pathway may become increasingly important and, eventually, override the reductive dissolution process as the most important pathway [13].

260


activation energy of lesorption

=4

Fe"l(~

)rocees

[Fe'"L]~){,}

[Fe"lL]~-2~o ~__t/

Reaction Coordinate Fig. 2. Reaction coordinate diagram for non-reductive dissolution process in acidic solution.

2.3. Reductive dissolution Reductive dissolution can be divided in two different stages interacting with each other: 1. induction period (first stage); 2. autocatalytic dissolution period (second stage).

2.3.1. Induction period During the induction period the generation of ferrous ions in the solution takes place. In the presence of lattice Fe n , as in magnetite, ferrous ions are dissolved [2,3,5] and their concentration in the solution is slowly increased, as described by the following equation: [)Fen

T 1 - ( n) - 2 ) L ] - (n- 2) --o tr e~ e 2+ -L.](aq

(5)

This observation is consistent with theories concerning the dissolution of metal oxides [19]. In general, the thermodynamic solubility, the ionic character of the M - O bonds [2,25] and the rate of dissolution decrease as the charge of the cation increases and its radius decreases. Iron as a transition element appears in two oxidation states. Fe II can be transferred to the solution more readily than Fe III because of the greater lability (kinetic instability) of the F e l l - o bond as compared to the FetII-o bond. In the absence of lattice Fe II ions, as in hematite, the generation of Fe 2+ ions in the solution is a very slow process. It involves electron transfer [2,7,10,13] from the adsorbed complex to surface Fe llI ions: DFeIII

L ._]-(n-2)~_~

[)Feli

L(n_I)_]-(n-2)

Electron transfer from the ligand to F e I11 DFen_L(n - 1)- ]-{n-2)._, Fe~aq) + products of ligand oxidation 2+ + L n - ~ [Fe " 2+ -L](aq) _ ~-(n-2) Fe(aq)

Complexation

(6) Dissolution (7) (8)

261

D. Panias et a l . / Hydrometallurgy 42 (1996) 257-265

At this stage, the dissolution of iron takes place in the form of Fe 2+ rather than in the form of Fe 3+.

2.3.1.1. Induction period in the Fetn-oxalates system. When oxalic acid is used as solvent reagent, the above mechanism can be described by the following equations [10]: [)Fem-C20~ -] ~ [ ) F e n - C 2 0 ; ]

Electron transfer

2 [ ) F e ' I - c 2 0 4 ] + 2H +---> 2Fe(2aq)+ 2CO 2 + C 2 0 ~- + 2)H

(9) Dissolution as Fe 2 ÷ (10)

At this stage, the factors affecting the rate of dissolution are: Temperature: It is reported [2,7,10] that, at high temperature (150°C), complete dissolution is rapidly achieved; the solution contains only Fe 2+ species generated by the reduction of Fe nl ions with oxalate ions, through a partially or totally heterogeneous process. Visible light impacts: Photochemical electron transfer on surface > FeIn-oxalate complexes provides an additional pathway for the start of dissolution. Photolysis of [Fe(C204)3] 3- takes place over a wide range of wavelengths (visible and ultraviolet region) according to the following equation [1-3]: 2[ Fe3* ( C204 )3 ]~aq, "~ 2[ Fez * (C204) ] (aq) + 3C2 O2- + 2CO2

( 11 )

2.3.2. Autocatalytic dissolution period When a sufficient amount of ferrous oxalate ions has been formed, the secondary reductive dissolution step becomes operative and the whole process is accelerated. This pathway is described by the following equations: - -,-(n-2) ~ DFelU-L] -(n-2) "" [Fe2+-L] -(n-2) [ ) F e m - L ] - ( n - 2 ) + [We2+ -Ll(aq) Adsorption of complex to surface [)FeIll L ] - ( n - 2 ) . . .

(12)

[We2+_L]-(n-2)¢¢..~ [)Well L ] - ( n - l ) . . .

[We3+ L] -(n-3)

Electron transfer

(13)

DFe,,_L]-(~-1) . . . [Fe3 + _ L ] - ( ~ - 3 ) . [)FeII_L]-(.-,) + [Fe3 +_L]-tn-3) Desorption

(14) "

H +'-" tFe

2+ - -~-(n-2)

-t.l aq

+ )H

Dissolution

(15)

The above reaction scheme can be divided into three elementary steps: 1. adsorption of aqueous solution ferrous complexes on the surface ferric complexes; 2. a fast outer-sphere or inner-sphere electron transfer [18,20,26] and formation of Fe II on the system interface; electron microscopy has shown that Fe n generated on the surface by electron transfer is more reactive than normal lattice Fe n [3], as a result, its transfer to the solution by Eq. (15) is easier in relation to Eq. (5) and Eq. (6), (7) and Eq. (8); 3. desorption of ferric complexes and transfer of trivalent iron in the solution.

262

D. Panias et al. / HydrometaUurgy 42 (1996) 257-265

60 50. 40

E 30 20 10 ¸

0 0

0.5 [ Fe(ll). ]l[Fe(ll)~d I

I1 concentration,expressed as its ratio respective to the Fig. 3. Induction period, "r, as a function of Fe(aq) concentrationof FeII originatedby total dissolutionof the solid (pH 4.1, temperature30°C, magnetite/oxalic acid solution) [2].

2.3.2.1. Autocatalytic dissolution period in iron-oxalate system. In the case of dissolution in oxalic acid the above reductive dissolution pathway can be illustrated by the following equations [2,3]. [ ) F e " ' - o x ] + [Fe2+-oX]Caq)~ ) F e " I - o x . . . Fe2+-ox

(16)

) F e " I - o x . . . Fe 2+-ox ~ ) F e " - o x . . . Fe 3÷-ox

(17)

)FeI'-ox... Fe 3+-ox --, ) F e l [ - o x q- [Fe 3+ OX](aq)

(18)

) F e n - o x --~ [Fe2+-oX](aq)

(19)

where ox denotes any species derived from oxalic acid: [ox] = [H2C204] + [HC204] + [C2042- ]

2.3.3. Factors affecting reductive pathway The rate of reductive dissolution is affected by two factors: Addition of ferrous ions in the initial acidic solution: Addition of bivalent iron in the initial solution exerts a beneficial influence on the overall dissolution rate because it reduces the time-consuming induction period (Fig. 3). As can be understood by the previous reductive dissolution mechanism, the presence of Fe E+ ions in the initial solution favours the autocatalytic dissolution period according to Eqs. (6), (7) and (12). Illumination of iron oxides suspensions: The illumination of metal oxide suspensions with UV radiation gives exactly the same result [1,4,27]. As has been previously reported, UV radiation assists the reduction of surface ferric ions and accelerates the overall reductive pathway. 3. Factors affecting the dissolution m e c h a n i s m

The dissolution process is affected by several parameters. The most important are: pH of the initial aqueous solution: pH values in the range 2 - 3 give satisfactory


263

2.5

m

2

E

1.5

g

1

"-

0.5 0

I

I

I

2

3

4

pH

Fig. 4. Influenceof pH on the initial dissolution rate of magnetite, R0 ([Oxalic acid]= 0.1 M, temperature

30oc) [3]. results, with the optimum value being 2.6 [1,5,10,17,19]. The influence of pH on the dissolution rate of magnetite in the system oxalic acid-magnetite is plotted in Fig. 4. Temperature: It is generally accepted that high temperatures, above 90°C, enhance the dissolution of iron oxides [1,5,16,17,19]. Chelating capacity of polycarboxylic acids: Chelating capacity is the ability of polycarboxylic acids to form chelated compounds (chelates) with metal ions in the solution. It is well known that the dissolution process is assisted by chelating agents such as EDTA and oxalic acid [4-7,14,16,19,22,28,29]. Photochemical effects: The dissolution of iron oxides suspended in organic acid solutions is promoted by ultraviolet and visible light radiation [1,4,16,27,30]. This process involves photochemical reduction of trivalent iron to bivalent, exerting a beneficial influence on the dissolution process.

4. C o n c l u s i o n s The dissolution of iron oxides with organic acids comprises three distinct steps: 1. Adsorption of organic ligands on the iron oxide surface. This is a chemisorption rather than a physical sorption process. The Lewis acid-base properties of surface iron oxides must be taken into account for the better understanding of this step. 2. Non-reductive dissolution. This is a simple desorption process of surface iron-ligand complexes. It is characterised by high activation energy and, therefore, it is an important dissolution pathway only at high temperatures. 3. Reductive dissolution. This is the main mechanism of iron oxides dissolution with organic acids. It consists of two different stages: 3.1. induction period; 3.2. autocatalytic period. The overall dissolution process is affected by: 1. pH of the initial solution; 2. temperature; 3. illumination of the solution with UV radiation; 4. the addition of bivalent iron to the initial solution.

264

5. L i s t o f

D. Panh~s et al. / Hydrometallurgy 42 (1996) 257-265

symbols

L n- :

):

)FelIl: )FelI: [)Fe-L]: II, HI: n+,n-:

a n y organic ligand with oxidation n u m b e r n, such as oxalate (L n - = C2042- or H C 2 0 4 ) , citrate ( L n - = C 6 H 5 0 7 3 or C6H6072- or C6H70 7) particle surface trivalent lattice iron on the particle surface b i v a l e n t lattice iron on the particle surface surface c o m p l e x a d s o r b e d species on the particle surface oxidation n u m b e r of surface lattice iron v a l e n c e o f aqueous species

Acknowledgements The financial support of the E u r o p e a n C o m m i s s i o n w i t h i n the framework o f the B r i t e - E u r a m II P r o g r a m (Contract No. B R E 2 - C T 9 2 - 0 2 1 5 ) is gratefully acknowledged.

References [1] Comell, R.M. and Schindler, P.W., Photochemical dissolution of goethite in acid/oxalate solution. Clays Clay Miner., 35(5) (1987): 347-352. [2] Baumgartner, E.C., Blesa, M.A., Marinovich, H.A. and Maroto, A.J.G., Heterogeneous electron transfer as a pathway in the dissolution of magnetite in oxalic acid solutions. Inorg. Chem., 22 (1983): 2224-2226. [3] Blesa, M.A., Marinovich, H.A., Baumgartner, E.C. and Maroto, A.J.G., Mechanism of dissolution of magnetite by oxalic acid-ferrous ion solutions. Inorg. Chem., 26 (1987): 3713-3717. [4] Waite, T.D. and Morel, F.M.M., Photoreduction dissolution of colloidal iron oxide: effect of citrate. J. Colloid Interface Sci., 102(1) (1984): 121-137. [5] Afonso, M.S., Morando, P.J., Blesa, M.A., Banwart, S. and Stumm, W., The reductive dissolution of iron oxides by ascorbate. J. Colloid Interface Sci., 138(1) (1990): 74-82. [6] Tongs, R., Blesa, M.A. and Matijevic, E., Interactions of metal hydrous oxides with chelating agents. VII. Dissolution of hematite. J. Colloid Interface Sci., 131(2) (1989): 567-579. [7] Tones, R., Blesa, M.A. and Matijevic, E., Interactions of metal hydrous oxides with chelating agents. IX. Reductive dissolution of hematite and magnetite by aminocarboxylic acids. J. Colloid Interface Sci., 134(2) (1990): 475-485. [8] Valverde, N., Investigations on the rate of dissolution of ternary oxide systems in acidic solutions. Ber. Bunsenges. Phys. Chem., 81(4) (1977): 380-384. [9] Regazzoni, A.E. and Matijevic, E., Interactions of metal hydrous oxides with chelating agents. VI. Dissolution of a nickel ferrite in EDTA solutions. Natl. Assoc. Corrosion Eng., 40(5) (1984): 257-261. [10] Sellers, R.M. and Williams, W.J., High-temperature dissolution of nickel chromium ferrites by oxalic acid and nitrilotriacetic acid. Faraday Discuss. Chem. Soc., 77 (1984): 265-274. [11] Biesa, M.A, Maroto, A.J.G. and Morando, P.J., Dissolution of cobalt ferrites by thioglycolic acid. J. Chem. Soc. Faraday Trans. I, 82 (1986): 2345-2352. [12] Segal, M.G. and Sellers, R.M., Kinetics of metal oxide dissolution: reductive dissolution of nickel ferrite by tris(picolinato)vanadium(II). J. Chem. Soc. Faraday Trans. I, 78 (1982): 1149-1164. [13] Rucda, E.H. and Grassi, R.L., Adsorption and dissolution in the system goethite/aqueous EDTA. J. Colloid Interface Sci., 106(1) (1985): 243-246.


265

[14] Blesa, M.A, Borghi, E.B., Maroto, J.G. and Regazzoni, A.E., Adsorption of EDTA and iron-EDTA complexes of dissolution of magnetite by EDTA. J. Colloid Interface Sci., 98(2) (1984): 295-305. [15] Godchev, I.G. and Kipriyanov, N.A., Regular kinetic features of the dissolution of metal oxides in acidic media. Russ. Chem. Rev., 53(11)(1984): 1039-1061. [16] Valverde, N. and Wagner, C., Considerations on the kinetics and the mechanism of the dissolution of metal oxides in acidic solutions. Ber. Bunsenges. Phys. Chem., 80(4) (1976): 330-333. [17] Gorichev, I.G. and Kipriyanov, N.A., Kinetics of the dissolution of oxide phases in acids. Russ. J. Phys. Chem., 55(11) (1981): 1558-1568. [18] Vermilyea, D.A., The dissolution of ionic compounds in aqueous media. J. Electrochem. Soc., 113(10) (1966): 1067-1070. [19] Blesa, M.A. and Maroto, AJ.G., Dissolution of metal oxides. J. Chim. Phys., 83(11/12) (1986): 757-764. [20] Pankow, Aquatic Chemistry Concepts. Lewis, Michigan (1991), pp. 603-641. [21] Comell, R.M. and Schindler, P.W., Infrared study of the adsorption of hydroxycarboxylic acids on ct-FeOOH and amorphous Fe(lII) hydroxide. Colloid Polymer Sci., 258 (1980): 1171-1175. [22] Rubio, J. and Matijevic, E., Interactions of metal hydrous oxides with chelating agents. I. fl-FeOOHEDTA. J. Colloid Interface Sci., 68(3) (1979): 408-421. [23] Grauer, R. and Stumm, W., Die Koordinationschemie oxidischer Grenzflachen and ihre Auswirkung auf die Auflosungskinetik oxidischer Festphasen in wal3dgen Losungen. Colloid Polymer Sci., 260 (1982): 959-970. [24] Buckland, A.D., Rochester, C.H. and Topham, S.A., Infrared study of the adsorption of carboxylic acids on hematite and goethite immersed in carbon tetrachloride. J. Chem. Soc. Faraday Trans. I, 76 (1980): 302-313. [25] Valverde, N., Investigations on the rate of dissolution of metal oxides in acidic solutions with additions of redox couples and complexing agents. Ber. Bunsenges. Phys. Chem., 80(4) (1976): 333-340. [26] Marcus, R.A., Electron, proton and related transfers. Faraday Discuss. Chem. Soc., 74 (1982): 7-15. [27] Buxton, G.V., Rhodes, T. and Sellers, R.M., Radiation-induced dissolution of colloidal hematite. Nature, 295 (1982): 583-585. [28] Baumgartner, E.C., Blesa, M.A. and Maroto, A.J.G., Kinetics of the dissolution of magnetite in thioglycolic acid solutions. J. Chem. Soc. Dalton Trans., (1982): 1649-1654. [29] Hidalgo, M.d.V., Katz, N.E., Maroto, A.J.G. and Blesa, M.A., The dissolution of magnetite by nitrilotriacetatoferrate(II). J. Chem. Soc. Faraday Trans. I, 84(1) (1988): 9-18. [30] Litter, M.I. and Blesa, M.A., Photodissolution of iron oxides. I. Magnetite in EDTA solutions. J. Colloid Interface Sci., 125(2) (1988): 679-687.