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New formulations for iron oxides dissolution
Hydrometallurgy, 27 ( 1991 ) 339-360
339
Elsevier Science Publishers B.V., Amsterdam
New formulations for iron oxides dissolution R. Chiarizia and E.P. Horwitz Chemistry Division, Argonne National Laboratory, 9700 South Cass Avenue, Argonne, IL 60439, USA (Received May 15, 1990; revised version accepted March 12, 1991 )
ABSTRACT Chiarizia, R. and Horwitz, E.P., 1991. New formulations for iron oxides dissolution. Hydrometallurgy, 27: 339-360. The dissolution of 59Fe-labelled synthetic goethite, o~-FeOOH, has been studied in solutions of mineral (HCI, HNO3, H2S04) and organic acids, belonging to the families ofcarboxylic and diphosphonic acids, alone, or in the presence of reducing agents. The data have been interpreted with a firstorder rate law, although the shrinking core dissolution rate law has been shown to apply generally. Tile effect of several reducing agents on the rate of dissolution of o~-FeOOH by inorganic and organic acids has been measured. Very high accelerations of the dissolution process were exhibited by ascorbic acid, sodium dithionite and sodium formaldehydesulfoxylate. In particular, the mixture of 1-hydroxyethane-l,l-diphosphonic acid and sodium formaldehydesulfoxylate proved to be extremely effective for a-FeOOH dissolution even at room temperature.
INTRODUCTION
The dissolution of iron (III) oxides is a very important technological process with applications in several areas; such as: ( 1 ) leaching of oxides ores in hydrometallurgy; (2) removal of iron oxides from non-metallic minerals, such as kaolin and silica, of importance for the glass and ceramic manufacturing industry; ( 3 ) removal of the scale from a metal prior to rolling; (4) removal of deposits from thermal power engineering equipment. An understanding of the laws and mechanisms that rule the dissolution of iron oxides is also important for studies on corrosion and passivation of the metal, and in geochemical studies, because the dissolution of iron oxides plays an important role in the availability and transformations of iron in soils, water and sediments, and in biochemical reactions. Improved and more efficient ways to dissolve iron oxides could also have application in two other areas of relevance for the nuclear industry. The first one is the decontamination of steel equipment, where the removal of a surface layer of oxides could possibly facilitate the removal of radioactive contami0304-386X/91/$03.50
© 1991 Elsevier Science Publishers B.V. All rights reserved.
340
R. CHIARIZIA AND E.P. HORWITZ
nants, such as, the transuranium ( T R U ) elements. Such a decontamination procedure, if successful, would reduce the cost of long-term confinement of TRU-contaminated equipment. The second application is related to the retrieval and treatment of high-level nuclear waste, at present stored in single or double shell tanks at a number of Department of Energy sites. The nuclear waste solutions, generally acidic, have been neutralized in the tanks by addition of alkali, to reduce their corrosive power. If, because of environmental considerations, the waste has to be retrieved and further processed to remove the T R U component, the availability of effective leaching solutions for the metal oxides and hydroxides becomes a very important factor. Based on these considerations, we have performed an investigation of the dissolution rate of synthetic goethite ( a - F e O O H ) by several inorganic and organic acids, alone or in mixture with reducing agents, with the aim of finding new formulations capable of effectively dissolving iron oxides. Although iron is probably present in the waste predominantly as amorphous ferrihydrate (FeOOHam), goethite ( a - F e O O H ) has been chosen for the investigation as a worst case; because it is known that the dissolution of goethite is more difficult than that of hematite (a-Fe203) and especially of ferrihydrate 111,21.
Mechanism of iron oxide dissolution Two generally accepted mechanisms for the dissolution of metal oxides are reported in the literature. The first one is referred to as the "adsorption" mechanism [ 3-7 ], the second one as "electrochemical theory" [ 6-8 ]. It is also known that when two oxidation states of the metal forming the oxide are possible (as in the case of iron), the presence of redox couples in the leaching solution can lead to redox reactions which affect the dissolution rate [ 2,3,10-14 ]. An exhaustive discussion of the mechanism of the dissolution of metal oxides in the presence of redox couples can be found in [9]. Based on electrochemical considerations, the authors reached the conclusion that, in the case of iron oxide, the rate of dissolution can be expressed as: d a i.~0.5 ~0.5 ~-o.5 dt =,~.u+.ve2+.Fe3+
(1)
where: ~ = the fraction of oxide dissolved; a = activity, and the potential of the solution is determined by the reaction: Fe 3+ + e - ~ F e 2+
(2)
Equation ( 1 ) predicts that the presence of reducing agents capable of low(.'ring the Fe 3+/Fe 2+ ratio, will accelerate the iron oxide dissolution, as exper-
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
341
imentally observed. It also predicts that the order of reaction with respect to hydrogen ions is 0.5. This has been verified in a n u m b e r of investigations (see table 3 of [9] ), although the formal reaction order of H + can show different values when the dissolution is complicated by hydration, redox reactions, complex formation, and other factors [9 ], becoming as high as 1 in H F or even 2 in HC1 [ 15 ]. Another factor, besides acidity and solution potential, that strongly affects the dissolution of metal oxides, is the presence in solution of complexing agents. It has been recognized [ 15 ] that, in the case of iron oxides, for exampie, the rate constant for the dissolution increases linearly with the logarithm of the stability constant of the complex of Fe (III) with the anion of the electrolyte; that is in the order H C 1 0 4 < H N O a < H B r < H z S O 4 < H C I < H 3 P O 4 < HF. Equation ( 1 ) explains these findings by considering the effect that the presence of complexing agents has on the potential of reaction given in eq. (2). D e p e n d i n g on the nature of the ligand, the potential of eq. (2) is displaced into the cathodic region, which favors dissolution, or into the anodic region, which prevents dissolution. In general, the rate constant ofeq. ( l ) in the presence of complexing agents is linearly d e p e n d e n t on the difference between the stability constants of the complexes of Fe (III) and Fe ( II ) with the ligand. Based on the previous discussion, the ideal leaching agent for iron (III) oxides should be a strong acid of an anion which is a good ligand for F e ( I I I ) and possibly also a reducing agent of Fe 3+. It is very difficult to find all these characteristics in the same c o m p o u n d , because strong complexing agents are generally anions of weak acids. A c o m p o u n d that comes close to the above requirements is oxalic acid, a strong complexant of Fe 3+ (fl~ = 7.53, f12= 13.64, f13--= 18.49) with moderate acid strength (PKa~= 1.25) [16]. It is known to dissolve iron oxide quite effectively, especially at high temperature, and especially in the presence of L-ascorbic acid [ 17 ]. Ascorbic acid is also a somewhat effective dissolving m e d i u m [17], in spite of its low acid strength (pKaj = 4.17 ) [ 16 ], because of its reducing power toward Fe 3+ and its ability to form a complex with Fe 3+ in the intermediate state of the reaction [18]. The requirements of high acidity strength and complexing power toward Fe 3+ led us to give a special consideration, as possible dissolving agents for iron oxides, to the relatively little-studied d i p h o s p h o n i c acids. The commercially available 1-hydroxyethane- 1, l-diphosphonic acid ( H E D P A ) is, in fact, a quite strong acid (pKa~= 1.56, pK,2=2.20, PK~l+PK~2=3.76 as opposed to 4.59 for oxalic acid [ 19] ), and a very good complexant for Fe 3+ (fire3+ = 16.2, fife2+ = 3.0 [20] ). According to Gorichev et al. [21 ], who studied the dissolution of magnetite (Fe304) with HEDPA, the dissolved iron was present in solution in the divalent state, which would imply that H E D P A also has some reducing power toward Fe 3+. The potential of H E D P A for the removal of
342
R. CHIAR1ZIA A N D E.P. H O R W I T Z
magnetite deposits and scale from steel was fully recognized by the authors of [21 ].
Rate laws The kinetics of dissolution of oxides under chemical control, that is when the rate-limiting step is a surface reaction, can be expressed by the rate law:
1 - (1-o~ ) l/3=kCt=k' t pro
(3)
where: k and k' = kinetic constants; (7 = the concentration of the leaching agent (assumed constant ); r and p = radius and density, respectively, of the dissolving particles (assumed of spherical shape ); c~= the fraction of solid dissolved. Equation (3) is often referred to as the "shrinking core" model [ 22,23 ], or the "compressive sphere" model [24,25]. A plot of the left side of eq. (3) against t should result in a straight line having a slope k' (s- ~). A similar rate law would be observed if the rate-limiting step were diffusion through a limiting boundary layer of constant thickness, 3, at the receding interface [26]. It is not possible from the kinetics alone to determine if k' contains a rate constant for a surface reaction or includes diffusion through a limiting boundary layer of solution adjacent to the solid surface. To distinguish between the two, the effect of temperature must be measured. The activation energy for solution diffusion is usually of the order of 5 Kcal or less, while that for chemical reactions usually occurs between 10 and 25 Kcal. A particular case of the shrinking core model that is often met in practice is when a solid product, which coats the dissolving particles, is formed as a result of the dissolution reaction. This happens, for example, when elemental sulfur is formed in the leaching reaction [26] and is relevant for our investigation involving the use of sulfur-based reducing agents for the dissolution of goethite. The sulfur which forms can adhere to the mineral surface forming a tightly bonded diffusion layer and reduce the speed of the dissolution process. In this case various rate laws have been derived, which consider the diffusion of the reactant through a solid product layer. The most generally used is the Jander's equation:
[1--(1--a)l/3]2=k't
(4)
The limits of validity of eq. (4) as well as other better approximations are discussed at length in [22]. Equations (3) and (4) are necessarily approximate when the dissolving solution, as in our case, is not in very large excess
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
343
compared to the a m o u n t of dissolved oxide, but still useful for a qualitative interpretation of the kinetic data. Although our data, as will be shown in the next section, are generally quite well represented by the rate law given in eqs. (3) or (4), we found it more straightforward to interpret them according to a first-order rate equation, following the approach of Noyes and Whitney, summarized in [22 ]. By expressing the dissolution rate as:
~
= k ( C ~ s - C )
where: C = the concentration of dissolved metal at time t; Cois~--the concentration when the oxide has been completely dissolved; it tbllows upon integration that: In (Cdiss - -
C)
-
In
Cdiss =
--
kt
By expressing the concentration of the metal as dissolved fraction o~= C/Cdiss, we obtain: ln(1 -o~) = - k t
(5)
Equation (5) predicts that a semilogarithmic plot of the fraction of undissolved material, ( 1-o~) versus time, should give a straight line of slope k (s- l ), and that the kinetic constant is related to the 50% dissolution time, tl/2, by the familiar first-order rate expression: In 2 tl/2-- k
(6)
Equation (5) is a very simple tool for handling the dissolution rate data, but does not tell us whether a surface reaction or a diffusion process through a product layer is the rate-determining process. For this purpose the use of the shrinking core model (eqs. (3 and 4) ) is still necessary. EXPERIMENTAL
~-FeOOH Synthetic goethite was prepared following the method of Atkinson et al. [27]. The first step of the procedure consists in the preparation of a basic iron nitrate sol. During the preparation of the sol, containing 0.05 mol of Fe ( NO3 ) 3, 7.5 X 107 counts per minute (cpm) of 59Fe ( III ) (Amersham) were added to the solution. The sol was then aged for a few days. After aging, a gel was precipitated at 45 °C by addition of a slight excess of NaOH. The gel was heated for 4 d at 62°C, during which time crystal growth takes place [27].
344
R. CHIARIZIA AND E.P. HORWITZ
During the 4 day period the solid-liquid mixture was stirred manually twice daily. The slurry was filtered, the solid dried in air at 100°C, and ground in an agate mortar. The final product had a specific activity of 1.7× l04 c p m / mg dried solid; 5°mg of oxides dissolved in 2 ml of leaching solution gave a final 59Fe activity of 4.25 × 103 cpm/10¢tl and an iron concentration equal to 0.28 M. Another sample of synthetic goethite was prepared following the same procedure, but without the addition of 59Fe (III). Samples of inactive goethite were analyzed by Mtissbauer spectroscopy and X-ray powder diffraction. The results confirmed the presence of only orthorhombic o~-FeOOH.
Dissolution experiments The dissolution experiments were performed by contacting in screw-cap glass tubes a weighed amount of radioactive goethite (approximately 50 mg) with the dissolving m e d i u m at a volume to weight ratio of 40 (2 ml for 50 rag). The sizes of the particles of iron oxide were estimated to be in the range 60-120 mesh. The glass tubes, which contained small magnetic bars, were inserted into a large circulation-jacketed beaker containing glycerol. The temperature was maintained at the desired value by a circulating thermostat. A magnetic stirrer rotated the magnetic bars at such a speed (several hundred r p m ) that the iron oxide was suspended in the dissolving medium. The glass tubes containing the samples of a - F e O O H to be dissolved were thermally equilibrated in the thermostatically controlled bath for several minutes before adding the required volume of leaching solution. The m o m e n t when the solution was added to the glass tube was taken as time zero for the dissolution reaction. At various time intervals, the test tubes were removed from the bath, rapidly quenched by dipping in ice/sodium chloride-water mixture ( - 5 °C) and centrifuged. The dissolution reaction was considered terminated at the m o m e n t of immersion in the cold bath. After centrifugation, 10 ¢tl samples of the solutions were withdrawn, the tubes were reinserted into the hot bath and the agitation restarted. At this time the dissolution reaction was considered to begin again. The time required for the quenching-centrifugation-sampling procedure was less than 3 min. This time can be considered as negligible for the experiments characterized by a slow dissolution rate. The kinetic parameters obtained for experiments with a 50% dissolution rate of a few minutes can be affected by the relatively slow procedure described above. However, duplicate experiments involving fast kinetics showed that, by rigorously following the same sampling procedure, it was possible to obtain results reproducible within about 10%. All the results reported in the following as induction time, tj/2, kinetic constants and activation energies must be considered affected by at least a 10% uncertainty, when not otherwise specified. The activity of the samples withdrawn at time intervals was counted with a
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
345
Beckmann Biogamma well-type gamma counter. The results were converted in o~ (fraction dissolved) by measuring the final activity (all dissolved) or from the knowledge of the specific activity of the radioactive goethite. By plotting the data as o~ versus t, in most cases the usual S-shaped curves mentioned by many authors [9] were obtained. These curves are characterized by an initial period with a slow dissolution rate, an intermediate period with m a x i m u m dissolution rate, and a final period where again the dissolution rate is slow. The data were analyzed as log (1-o~) versus t. In most cases, straight lines were obtained, as predicted by eq. (5). In several cases, the data points were aligned on a straight line only after an induction time. The presence of such an induction or initiation time has been noticed by many authors (see, for example, [ 11,15, 25, 28-30 ] ). Although attempts have been made to correlate the induction time to the presence on the particle surface of potential nuclei capable of being activated under the reaction conditions [ 30 ], no rational explanations for these anomalies are available [ 11,28 ]. The presence of an induction time must also be taken into account when plotting the data according to the equations of the shrinking core model.
Reagents used in the dissolution experiments HC1, HNO3, H2504, oxalic, malonic, succinic, maleic, glycolic, a-hydroxyisobutyric, citric, tartaric, ascorbic and tetrahydrofurantetracarboxylic acid ( T H F T C A ) were analytical reagent grade products used without further purification. Furantetracarboxylic acid (FTCA) was provided by R.C. Gatrone, Chemistry Division, Argonne National Laboratory. 1-hydroxyethane-1, 1diphosphonic acid (HEDPA) was obtained from the Albright and Wilson Company and recrystallized from glacial acetic acid. Vinylidene-1,1-diphosphonic acid (VDPA) and 1,2-dihydroxyethane-l,l-diphosphonic acid (DHEDPA) were synthesized and purified as reported in [31 ]. VDPA and DHEDPA are new derivatives of HEDPA, with a higher acid strength than the parent c o m p o u n d ( P K a i +pga2= 3.47 for VDPA, 2.65 for DHEDPA, 3.76 for HEDPA) [ 19 ], but much more readily decomposable upon heating [ 31 ]. As reducing agents, analytical grade hydroxylamine hydrochloride (Ntt2OH-HC1), SnC12, hydroquinone, metallic zinc, Na2SO3, sodium dithionite (Na2S204) and sodium formaldehydesulfoxylate (HOCH2SO2Na) were used. The latter c o m p o u n d (hereafter called SFS) was obtained from Kodak under the commercial name of Rongalite. Although it has been used for a long time as a reducing agent in the textile industry and for reduction reactions of certain actinides [32,33 ], we are not aware of any previous use of SFS as an accelerator of metal oxide dissolution. The oxidation reactions and standard potentials of the above reducing agents can be found in [34].
346
R. CHIARIZIAAND E.P. HORWITZ
The chemical formulas of the organic acids and some reducing agents used in the experiments are reported in the Appendix. RESULTS AND DISCUSSION
Inorganic acids The a-FeOOH dissolution data obtained with mineral acids are plotted in Fig. 1 according to eq. (5). From the slope of the straight portion of the lines, the first-order kinetic constant k was calculated. The induction time, t~, when present, was obtained as the difference of the experimental 50% dissolution time, tl/2, and the theoretical half dissolution time calculated using eq. (6): In 2 ti=h/2-- k B
A
\
eL
HCI .01
0
50
1~'
100
150
Time (minutes)
.ol 0
50
100
50
100
HNO 3 150
Time (minutes)
H2SO4
150
Time (minutes)
Fig. 1. Dissolution of a-FeOOH by inorganic acids. (A) HCI. a = 2 M, 25°C; b = 2 M+0.1 M SFS, 25°C; c= 1 M, 80°C; d = 2 M+0.1 M Na2S204, 80°C; e = 2 M+0.1 MSFS, 8 0 ° C ; f = 2 M, 80°C. (B) HNO3. a = 1 M, 80°C; b = 1 M+0.1 MSFS, 25°C; c = 6 M, 80°C; d = 12 M, 80°C. (C) H2SO4. a = 3 M, 80°C; b = 3 M+0.1 M Na2SO3, 80°C; c = 3 M+0.1 M Na2S204, 80°C; d = 3 M+0.1 MSFS, 80°C.
347
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
TABLE 1 K i n e t i c p a r a m e t e r s for the d i s s o l u t i o n of a - F e O O H with i n o r g a n i c a c i ds D i s s o l v i n g so lution
2MHC1 2 M H C I + 0 . 1 M SFS 1 MHCI 2 MHCI 2 MHCI+0.1 MSFS 2 M HCI+0.1 MNazSzO 4 1 MHNO3 6 MHNO3 12 M H N O 3 3 M H2SO 4 3 M H2SO4 + 0.1 M Na2SO3 3 M H2SO4+ 1 M Na2S204 3 M H2SO4+0.1 M S F S
T
k a
t, a
t ~/ 2 a
(°C)
(s - I )
(min)
(min)
25 25 80 80 80 80 80 80 80 80 80 80 80
2)< 2.0)< 2.0)< 1.3X 1.1)< 6.7× 4.8 X 1.1)< 7.9)< 5.7)< 5.6 X 6.3)< 1.9)<
3.8 25 20 29 24 13 -
58 8.9 10 21 130 35 49 45 31 6.0
10 -619 10 -5 10 -4 10 -3 10 -3 10 - 4 10 -6 10 4 10 -4 10 4 10 -4 10 -4 10 -3
a-- 1 0 % bq- 1 0 0 %
This calculation was possible only when the 50% dissolution was actually achieved. The values of k ( s - I ) , ti (min) and tl/2 (rain) are listed in Table 1. The data obtained with HC1 confirmed the already known pronounced effect of acid concentration and temperature [ 3,5,15 ]. The calculation of the activation energy for 2 M HC1, using the Arrhenius equation [ 26 ], although performed from only two temperatures, provided a value for Ea equal to about 24 Kcal/mol in good agreement with the value of 20-24 Kcal/mol reported in [ 5 ] for 0.5 MHC1 and various crystal morphologies of natural goethite. The addition o f NazS204 or SFS did not bring about an acceleration of the a-FeOOH dissolution rate at 80°C. The presence of Na2S204 actually reduced the kinetic constant by a factor of 2. An increase of k of about 10 times was caused by SFS at 25 oC, but the dissolution process remained very slow. The different behavior of SFS and Na2S204 can be probably explained by the observation that the presence of the latter led to the formation of colloidal sulfur. Elemental sulfur formed during the dissolution of c~-FeOOH coats the dissolving particles, slowing the dissolution process. The formation of colloidal sulfur has been reported in [35 ] in a study of the reduction of acid solutions of UO 2+ by Na2S204, and is attributed to the rapid disproportionation of dithionite into thiosulfate and sulfite: 2 $2042- + H 2 0 ~ $ 2 0 2 - + 2 HSO~followed by the disproportionation of thiosulfate:
3,48
F,. CHIARIZIA AND E.P. HORWITZ
$202- + 2 H+--+H20+SO2+S In agreement with our observations, the formation of colloidal S was not noticed by the authors of [ 35 ] if SFS was used instead of Na28204. Although sulfoxylic acid is also known to be unstable [ 34 ], decomposing according to: 2 H2802--+$2 O2- + 2 H + + H 2 0
which should also lead to the formation of elemental sulfur, its decomposition may be retarded by the presence of the - C H 2 O H moiety in the SFS molecule. The formaldehydesulfoxylate ion, CH20-HSO2, is known to be rather stable and this property is utilized in polarographic procedures for the analysis of commercial dithionite preparations [36 ]. The formaldehyde-stabilized sulfoxylate anion can then be oxidized directly to sulfite, according to: H2SO2 + H20-+H2803 + 2e + 2 H + "['he failure of SFS to accelerate the dissolution of a-FeOOH in 2 M HC1 at 80 °C is, however, not completely understood. The oz-FeOOH dissolution data in HNO3 show that a very high concentraA
B
0.80
m
0.80
HCI
0.60
~
09 0.8 07 ~
~:- 0.40
i
0.20 0.00
20
Time 40
~
HNO 3
0.60 09
d_.c
O8
0.40
07
0.6
0.6
0.20
04 O2
60
02
0.00
80
(minutes)
20
0
cy
Tim(mi e nutes) 40
60
80
C
.2So4
0.80 0.60
09 08
b
0.40
07 06
0.20
04 0.2
0.00
0
20
Time 40
60
80
(minutes)
Fig. 2. Same data o f Fig. 1 plotted according to eq. ( 3 ) . For the meaning of the symbols see Fig. 1.
349
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
tion of the acid ( > 12 M) is required to achieve relatively fast dissolution of the oxide at 80°C, with k approaching the value of 10 -3 s-1. An induction time in the range 20-25 min was observed. The attempt to use SFS as an accelerator failed even at 25°C (curve b of Fig. 1B). In the presence of 1 M HNO3 + 0.1 M SFS after an initial acceleration, the dissolution process became again very slow, after about 2 h, indicating the destruction of SFS by HNO3. The data for H2SO 4 (Fig. 1C) show that, in this non-oxidizing medium, reducing agents are effective in accelerating the dissolution. While sulfite and dithionite have only a modest effect on the value of k, and primarily reduce the induction time, with SFS the dissolution rate is increased by a factor of about 6, while the induction time is reduced practically to zero. With dithionite, the formation of elemental S was again noticed. The better performance of SFS as compared to Na28204 may be explained as in the case of HCI data. If the data of Fig. 1 are plotted according to the shrinking core model--eq. ( 3 ) - - t h e plots of Fig. 2 are obtained. The data are well aligned on straight lines, when allowance is made for the induction time; indicating that the rate of the dissolution process is primarily controlled by a surface chemical reaction or by diffusion through a limiting boundary layer. With HCI the measured value of the activation energy excludes diffusion as a controlling factor. No activation energy measurements were performed with HNO3 and H2SO4.
Carboxylic acids Figures 3 and 4 report the data on the dissolution rate of a - F e O O H by oxalic acid alone or in the presence of reducing agents. The kinetic parameters are listed in Table 2, together with the results of the measurements performed with other carboxylic acids. The data in Fig. 3A and Table 2 show B
--q:>~ d
1.41 1.2 co
"
1.0
09
0.8
.1
0.6k t
O8 a
,,/,,c
d
07 06 04 02
01
100 200 Time (minutes)
o o~ 0
'-~+~ 50 1O0 150 Time (minutes)
2OO
Fig. 3. Dissolution of a - F e O O H by oxalic acid. (A) a = 1 M, 80°C; b = 0 . 5 M, 80:C; c = 1 M, 60~C; d = 1 M, 40°C; e = 1 M, 25°C. (B) Same data as in A (a,c,d, and e) plotted according to eq. (3).
350
R. C H I A R I Z I A A N D E.P. H O R W I T Z
A
z~
T-
.1
co
1.0
"~" --
0.6 0.4 0.2 0.6
,,/~/,~/~::
-.~
U 0
i 20
40
,
i
j,,~
c _ 0.9 ~.~ oZ ~' 0.2
/ ~
0.4
~
~-~
~ .01
B
f 0.4
/
~//
dG, d/
~
~
c o.~
0.2
o.8
0.0
60
o.~ 0
10
Time (minutes)
20
30
Time (minutes)
Fig. 4. Dissolution of c~-FeOOH by oxalic acid in the presence of reducing agents. ( A ) a = l M, 6 0 ° C ; b = 1 M + 0 . 0 1 MNazSO3, 60°C; c = 1 M + 0 . 0 1 MNa2S204, 6 0 ° C ; d = 1 M + 0 . 0 1 M S F S , 60°C; e = 1 M + 0 . 0 1 M ascorbic acid, 8 0 ° C ; f = 1 M + 0 . 0 1 M S F S , 80°C. ( B ) Same data as in A (c,d,e a n d f ) plotted according to eq. ( 3 ) ( u p p e r p a r t ) and to eq. ( 4 ) (lower part ).
'TABLE 2 Kinetic parameters for the dissolution of c~-FeOOH with carboxylic acids
Dissolvings o l u t i o n 0.5 M succinic or glycolic or o~-hydroxyisobutyric acid 0.5 M maleic or citric or tartaric acid 0.5 M malonic acid 1 M THFTCA or FTCA 1Mascorbicacid FTCA +0.1 M ascorbic acid 1 M H2C204 I M H2C204+0.01MSFS IMH2C204 1 M H2C204 1M H2C204+0.01MSFS 1 M H2C204+0.01 MNa2SO3 IMH2C204+0.01MNa2S204 0.5 MH2C204 1 M H2C204 I MH2C2O4+0.1 M H E D P A 1 MH2C204+0.01Mascorbic acid 1M H2C204+0.01MSFS a± 10% b+ 100%
T
ka
ti a
lj/2 a
(°C)
(s - l )
(min)
(min)
80 80 80 80 80 80 25 25 40 60 60 60 60 80 80 80 80 80
NO DISSOLUTION 1 × l 0-6 b
_
-
4×
10 -6b
-
-
IX10 -6b
-
9.3)<10 5 4.1)<10 -4 5.0× 10 -5 5.2X 10 -4 1.1×10 -4 5.2×10 -4 2.5)<10 -3 4.0)< 10 - 4
1.6×10 3 1.2×10 -3 2 . 1 × 1 0 -3 9.4× 10 -5 4.7× 10 3 6.3X 10 -3
5.8 39 49 25 15 16 _ 3.4 3.5 17 _ -
-
130 67 280 22 130 37 4.6 45 7.2 13 9.0 140 2.5 1.8
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
35 [
that the effect of temperature is more pronounced than the effect of the acid concentration. An activation energy of 16 Kcal/mol has been calculated from the data of Fig. 3A, well in the range of kinetic control by surface chemical reactions. The same data, plotted in Fig. 3B according to eq. (3), show that the shrinking core model also applies to this case. Attempts to accelerate the dissolution reaction by increasing the acidity of the solution by adding HC1 or HNO3 to 1 M H 2 C 2 0 4 caused a slight opposite result; probably by depressing the dissociation of oxalic acid, reducing the concentration of the active ligand, C2042- . A more evident reduction in k was caused by the addition of HEDPA to the 1 M oxalic acid solution. This may be an indication that HEDPA has some reducing power towards Fe (III), as claimed in [21 ]. Fe(II) oxalate is, in fact, only sparingly soluble and its formation on the surface of the dissolving particles would reduce the dissolution rate. Other explanations, based on the different adsorption of the oxalate and EDPA anions on the goethite surface, are, however, available. The precipitation of FeC204 or, rather, the transformation of some of the o~-FeOOH into FeC204, was obtained in the experiments where the dissolving m e d i u m was 1 M oxalic acid containing 0.1 M ascorbic acid or SFS at 60 ° C. The dissolution was initially faster, but after going through a maximum, it stopped at about 25% dissolution, probably corresponding to the solubility of Fe204 in the dissolving medium. All other dissolution experiments in the presence of reducing agents were performed at a much lower concentration of the reducing agent. The results are reported in Fig. 4 a n d / o r in Table 2. Na2SO3 is practically ineffective; Na2S204 doubles the first-order rate constant at 60 °C and generates elemental sulfur; SFS is the most effective, raising k by a factor of 3, 5 and 10, at 80, 60 and 25°C, respectively. An effect similar to SFS is shown by 10 -2 M ascorbic acid at 80°C. The acceleration of c~-Fe203 dissolution by an ascorbate-oxalate mixture has been recently reported [ 17 ]. Figure 4B shows the data obtained with some oxalic acid-reducing agents mixtures plotted according to eqs. ( 3 ) and (4). The fact that eq. (4) provides a better fit of the data indicates that in this case the rate-controlling process is the diffusion of the reagents through a solid surface layer, presumably FeC204, covering the dissolving particles. Regarding the other carboxylic acids used in this work, the data of Table 2 show that a significant dissolution of o~-FeOOH takes place only with ascorbic acid and, to a much lesser extent, with malonic acid. The former acid is a reducing agent, which may explain its effectiveness. Especially striking is the difference of the constant k for oxalic acid and malonic acid ( 1.2× 1 0 - 3 v e r s u s 4.1 × 1 0 - 6 S l, at 0.5 M a n d 80°C). Since malonic acid is also a very good complexing agent for Fe 3+ [ 16 ], the difference has to be ascribed either to its lower acidity or to the presence of a CH2 group in the molecule that makes the interaction of the anion of the acid with the surface of the oxide more difficult. With succinic acid, having another CH2 group and an even lower
352
R. C H I A R I Z I A A N D E.P. H O R W I T Z
acidity, there is no measurable dissolution of oz-FeOOH. Similar results have been reported in [37 ] in the study of the dissolution of 5-A1203 by oxalic, malonic and succinic acid, and have been attributed to the formation on the oxide surface of five-, six-, and seven-membered chelate rings, with the fivernembered rings more readily detachable from the surface. The tetracarboxylic acids also performed very poorly. However, the mixture of FTCA with ascorbic acid showed a remarkable enhancement of the dissolution rate compared to the FTCA alone, (k increased by a factor of 400) or to ascorbic acid alone (k larger by a factor of 5). No experiments have been run with tetracarboxylic acids and SFS or Na2S204, because of the limited availability of the acids. An acceleration effect similar to that of ascorbic acid would, however, be expected.
Diphosphonic acids Figure 5 shows a comparison of the oz-FeOOH dissolution data in the three acids investigated, HEDPA, VDPA and DHEDPA. The first two acids have practically the same (disappointingly low) kinetic constant, although VDPA shows a somewhat lower induction time. DHEDPA, on the other hand, is much slower. Data with DHEDPA have been obtained at 60°C instead of 80 ° C, because this acid spontaneously decomposes, even at room temperalure [31 ]. For all three acids the shrinking core model applies (figure not .shown for brevity). The data obtained with the three diphosphonic acids in the presence of reducing agents, at different concentrations and in the temperature range 2580°C are summarized in Table 3. Some data with HEDPA are reported in Fig. 6 ( A - D ) . A very high acceleration of the dissolution rate by 1 M H E D P A at 80 °C is exhibited by ascorbic acid, dithionite, and SFS, with a kinetic constant increase of about 100. A strong acceleration is also shown at 80°C by Na2SO3 and Zn, which affects the dissolution rate through the reducing power
\\ \ 01 0
2
4 6 Time (hours)
8
10
Fig. 5. Dissolution of ~ - F e O O H by diphosphonic acids, a = 0 . 5 M HEDPA. 80°C: b = 1 M HEDPA. 80~C; c = 1 M VDPA, 80°C; d = 1 M D H E D P A , 60°C.
353
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
TABLE 3 K i n e t i c p a r a m e t e r s for t h e d i s s o l u t i o n o f ~ - F e O O H w i t h d i p h o s p h o n i c a c i d s Dissolving solution
T (°C)
ka (s 1)
0.5 M H E D P A M HEDPA M VDPA M DHEDPA
80 80 80 60
2.7× 1.8 × 1.7 × 4×
25 25 25 25 25 25 25 25 25
1 × 1 0 6b 1 × 1 0 6b 2.3?< 10 -5 2 . 2 × 10 -4 5.0X10 4 1.9N10 3 2 . 3 × 10 -3 2 . 5 X 10 4 3 . 4 × 10 4
_ _ -47 23 _ 8.0 _ -
_ _ 500 100 46 6.1 13 46 34
M M M M M M M M M
HEDPA+0. l Mhydroquinone H E D P A + 0. l M N H 2 O H " H C I H E D P A + 0 . 1 M SnCI2 H E D P A + 0.1 M Na2SO3 H E D P A + (0.2 M ) Z n C HEDPA+ 1 MNazS204 HEDPA+0.1 MSFS V D P A + 0.1 M SFS DHEDPA+0.1 MSFS
10 . 5 10 -- 4 10 - 4 10 - 6 b
ti a (min)
tl/2 a (min)
52 120 82 -
480 190 [ 50 -
1 M H E D P A + 0.1 1 At H E D P A + 0.1 1 M H E D P A + 0.1 1 M HEDPA+0.1 1 MDHEDPA+0.01
M ascorbic acid M Na2SO3 M Na2S204 MSFS MSFS
40 40 40 40 40
1.0X 10 5 1.5)< 10 -3 5 . 7 X 10 - 3 7.1 × 10 - 3 1.2X10 3
_ 5.3 2.7 _
_ 13 1.7 4.3 9.6
1M 1M 1M 1M 1M
M ascorbic acid MNa2SO3 MNa:S204 MSFS MSFS
60 60 60 60 60
1.4×10 3 2 . 9 X 10 . 3 1.1X 10 - z I.IXIO 2 5 . 8 × 10 -3
30 1.2 _ -
38 5.2 1.0 1.0 2.0
80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80
2 . 6 X 10 - 4 2 . 4 × 10 - 4 1.4X 10 -3 8.1X10 3 1 . 6 × 10 . 2 9.1XI0 3 1 . 6 × 10 -2 1.6X 10 . 2 1 . 6 × 10 - 2 8.4×10 3 4 . 8 X 10 -3 1 . 4 × 10 -2 1.4×10 2 1.4X 10 -2 8.5 X 10 -3 4.2 X 10 -3 2.4×10 3 4 . 6 × 10 -3 2 . 9 × 10 - 3
14 62 4.7 1.1 4.8 44s d _ _ -
58 110 13 2.5 5.5 2.0 45 s o 45 s d 45 s a 1.4 2.4 49 s d 49s d
M M M M M M M M M M M M M M M M M M M
H E D P A + 0.1 HEDPA+0.1 HEDPA+0.1 HEDPA+0.1 DHEDPA+0.1
H E D P A + 0.1 M h y d r o q u i n o n e HEDPA+0.1 M NH2OH.HCI H E D P A + 0.1 M SnC12 H E D P A + (0.2 M ) Z n c HEDPA+0.1 Mascorbic acid H E D P A + 0 . 1 M Na2SO3 H E D P A + 0.15 M N a 2 S 2 0 4 H E D P A + 0.1 M N a 2 S z O 4 H E D P A + 0.05 M Na2S204 H E D P A + 0 . 0 1 M Na2S204 H E D P A + 0 . 0 0 5 M Na2S204 HEDPA+0.15 MSFS HEDPA+0.1 MSFS HEDPA+0.05 MSFS HEDPA+0.01 MSFS HEDPA+ 0.005 MSFS V D P A + 0.01 M a s c o r b i c a c i d VDPA+0.1 M ascorbic acid V D P A + 0.02 M SFS
" + 10% b_+ 100% CThe c o n c e n t r a t i o n refers to t h e Z n 2+ in s o l u t i o n a f t e r c o m p l e t e d i s s o l u t i o n o f the m e t a l . aS = s e c o n d s .
49
sa
1.4 2.7 4.8 2.5 4.0
R. CHIARIZIAAND E.P. HORWITZ
35;4 A
1
~ , ~-..
a
1TTx
b i
\
.01
20
%
J
\\,
h'~
80 C
10
0
\
~g
1 g ',, +,e
x\
.01
30
0
c
60 C 10
20
30
~,
25C
D
d
d
!~
-
.1:
1
g'\
x 40'C .01
"g
20 40 Time (minutes)
60
.01
20 40 Time (minutes)
60
Fig. 6. Dissolution of o~-FeOOH by 1 M HEDPA in the presence of reducing agents al ( A ) 80cC, ( B ) 60°C: ( C ) 4 0 ° C and ( D ) 25°C. Reducing agents: a = 0 . 1 M NH2OH-HCI; b = 0 . 1 .'~l hydroquinone; c = 0.1 M SnCI2; d = 0.1 M ascorbic acid; e = (0.2 M ) Zn; j - 0.1 M Na2SO3; g = 0 . 1 M S F S : h = 0 . 1 MNa2S204.
of native hydrogen generated during its dissolution by HEDPA. Hydroxylamine and hydroquinone, on the other hand, show a negligible effect, while Sn(II) occupies an intermediate position. The behavior of ascorbic acid is very characteristic. Its effectiveness as a dissolution accelerator is strongly influenced by the temperature. Even at 60°C it is much less effective than dithionite and SFS. The activation energy for the dissolution of o~-FeOOH by 1 M HEDPA in the presence of 0.1 M ascorbic acid has been calculated from the data of Table 3 and is equal to 34 Kcal/mol. This very high value of Ea indicates that a complex chemical interaction between Fe (III), HEDPA and ascorbic acid takes place and is rate-determining. The comparison of the data of Fig. 6A and D (80°C and 25°C, respectively) shows that the effect of Sn ( II ), Zn and Na2SOB on the dissolution rate of o~-FeOOH by 1 M HEDPA also depends strongly on the temperature, with activation energies in the range 10-15 Kcal/mol. With these reducing agents, therefore, chemical reactions seem to still be, at least in part, the rate-determining process for the iron oxide dissolution. A completely different behavior is exhibited by the dithionite and sulfoxylate species. They show a very pronounced accelerating effect even at 25 oC, where the dissolution of c~-FeOOH by the diphosphonic acids alone is
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
355
negligible. With both reducing agents, the application of Arrhenius' equation to the data of Fig. 6 allows calculation of an activation energy of 5.7 Kcal/ mol; which shows that, in this case, a diffusion process, not a chemical reaction, is the rate-limiting factor. The formation, at all temperatures, of colloidal sulfur was again observed when Na2S204 was used as an accelerator. This fact does not seem to have any effect on the goethite dissolution reaction at 80 and 60 °C. However, it does change the shape of the dissolution rate data at 40°C and especially at 25°C (Fig. 6C and D). At these temperatures the data for Na2S204 show a curvature toward slower dissolution rates. In other words, the initially extremely fast dissolution rates of oz-FeOOH become progressively slower, requiring a much longer time than with SFS to proceed to completion. This behavior of Na2S204 may be ascribed to the progressively slower diffusion, with decreasing temperature, of the reagents through a layer of sulfur coating the dissolving particles. Supporting evidence for sulfur coating is shown by Fig. 7, where the dissolution data at 25 °C with 1 M HEDPA in the presence of 0.1 M Na2S204 or SFS is plotted according to eqs. (3) and (4). With Na2S204, the data are better aligned when plotted as required by Jander's equation (eq. (4)). This indicates that, at low temperature, the rate-determining process is the diffusion of the reagents through a solid product layer. When VDPA or DHEDPA were used instead of HEDPA, acceleration effects similar to those described in the case of HEDPA were observed with ascorbic acid at 80 °C and with SFS throughout the whole temperature range. With VDPA, however, the presence of 0.1 M SFS seemed to promote a rapid alteration or decomposition of the diphosphonic acid, leading to the formation, after a few minutes at 80°C, and within 1 hour at 25°C, of a massive 0.8 -~
0.6
"~ 0.4 EL ~- 0.2 A
0.0
~[ ~: :,% 0.2 B
09 08 07 06 0.4 0.2 O0
h
0.9 h .~...//~g
O.C
,
0
~/)¢/ )~ .
10
~0.8~0.7 .
.
~ O.6 40.4
.
20
30
Time (minutes)
Fig. 7. Data for the dissolution of ~ - F e O O H by 1 M H EDPA in the presence of 0.1 M SFS (g) or 0.1 M Na2S204 (h) at 25 oC, plotted according to (A) eq. ( 3 ) and (B) eq. (4).
356
R. CHIARIZIA AND E.P. HORWITZ
whitish precipitate, presumably of FePO4. The same kind of precipitate was observed after a longer time in the solutions containing 1 M DHEDPA and 0.1 M SFS, after dissolution of a-FeOOH. The iron-SFS-promoted decomposition of VDPA and D H E D P A is at present under investigation in our laboratory. The data in Table 3 show that with HEDPA there is no correlation between the final concentration of iron in solution and the concentration of either Na28204 or SFS. Both Na28204 and SFS maintain their accelerating effect at a concentration as low as 5 X 10-3M. These results seem to indicate that SFS or Na2S204 are required only for the reduction and subsequent removal of a surface layer of Fe (III) atoms. The successive layers of oxide must be activated by this process and more easily dissolved by HEDPA, possibly through a redox mechanism, as reported in [21 ]. If FeOOH is dissolved in the presence of a substoichiometric amount of SFS, in fact, a colorless solution forms. On the other hand, when no SFS is present, iron is dissolved slowly by HEDPA as yellow Fe(III) contrary to the findings of [21 ]. The real nature of the interaction between Fe (III) and Fe (II) with HEDPA, however, requires further investigation. With MnO2, where the redox couples of the metal is characterized by a more positive Eo value than the Fe3+/Fe 2+ couple, HEDPA clearly shows its reducing power. MnO2 is dissolved by 1 M HEDPA slowly at room temperature and extremely rapidly at 80 ° C, generating a purple complex of Mn (III) stabilized by HEDPA. In the presence of SFS, the reduction of MnO2 proceeds further, and a colorless solution of M n ( I I ) is obtained, with the dissolution reaction becoming almost instantaneous even at 25°C. Preliminary experiments, performed in the same conditions as with o~FeOOH, showed that SFS also has a very pronounced acceleration effect on the dissolution of CeO2 by 2 M HC1 at 80 ° C. With an HEDPA-SFS mixture the dissolution of CeO2, after reaching 25% dissolution rapidly, practically stops due to the probable decomposition of HEDPA and precipitation of either CePO4 or the C e - H E D P A complex. In contrast, when no SFS is present, practically no CeO2 is dissolved by either 2 M HC1 or 1 M HEDPA over a period of 4 h. SUMMARY AND CONCLUSION
A pronounced acceleration effect by some reducing agents on the rate of dissolution of synthetic c~-FeOOH by sulfuric acid, oxalic acid and diphosphonic acids has been measured. The most effective reducing agents, among those tested, are ascorbic acid, sodium dithionite and sodium formaldehydesulfoxylate (SFS). The acceleration effect is relatively small in H2804 and in I-~2C204. In oxalic acid the concentration of the reducing agents must be kept very low to prevent precipitation of FeC204. The risk of FeC204 precipitation, together with the known capability of oxalic acid to rapidly dissolve iron
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
357
oxides without reducing agents (at least at high temperature), makes the acceleration effect in oxalic acid of limited value. With diphosphonic acids, showing by themselves a rather low dissolution rate of a-FeOOH, in spite of their high acidity and complexing power for Fe(III), the accelerating effect brought about by the reducing agents is remarkable. It is best shown by SFS, because ascorbic acid is effective only at high temperature, and Na2S204 rapidly decomposes in the acidic medium generating colloidal sulfur which can slow down the dissolution process at low temperature. Although the chemistry involved in the dissolution of o~-FeOOH by a HEDPA-SFS mixture is not completely understood, this novel system, investigated for the first time in this work, could have practical applications in all those fields where a rapid dissolution of iron oxides at room temperature is required. The capabilities of HEDPA-SFS mixtures to dissolve metal oxides can probably be extended to all other cases of oxides of metals having lower oxidation states. Preliminary experiments have shown that the presence of SFS in 1 M H E D P A also accelerates the dissolution of MnO2 and, to some extent, of Ce02 and PuO2. ACKNOWLEDGMENTS
This work was performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Sciences, U.S. Department of Energy, under contract number W-31-109-ENG-38. The authors wish to thank Dr. Ralph Gatrone and Dr. Kenneth Nash for providing samples of FTCA, VDPA and DHEDPA, and Dr. K.A. Martin, formerly at the Institute of Gas Technology, Chicago, and Dr. L. Soderholm, of Argonne National Laboratory, for the characterization of the a-FeOOH used in this work by M6ssbauer spectroscopy and X-ray powder diffraction, respectively. APPENDIX:
Chemical formulae of the organic acids and some reducing agents used oxaric malonic succinic maleic glycolic
HOOC-COOH HOOC-CH2-COOH HOOC-CH2-CH2-COOH HOOC-CH=CH-COOH HO-CH2-COOH
358
R. C H I A R I Z I A A N D E.P. H O R W I T Z
Q -hydroxybutyric
/COOH (0H3)20\O H COOH
I
citric
HOOC-CH~I-CHzCOOH OH
tartaric
COOH-CHOH-CHOH-COOH CH2OH
ascorbic H ~ O H
tetrahydrofurantetracarboxylic THFTCA
furantetracarboxylic FTCA
hyd roxyethane-l,l-diphosphonic
HEDPA
vinylidene-l,l-diphosphonic VDPA
2-dihydroxyethane-l,l-diphosphonic DHEDPA
CH~',,,_/~(OH)= Ho/C'~(OH)z / ~(OH) 2
H2C--C'~I~(OH)2
HOCH2\_/~(OH)2 Ho/C"~(O H)2
hydroquinone
sodium formaldehydesulfoxy[ate SFS
HOCH2\s_ O- Na +
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
359
REFERENCES 1 Cornell, R.M., Posner, A.M. and Quirk, J.P., J. Appl. Chem. Biotechnol., 25 ( 1975): 701706. 2 Zinder, B., Furrer, G. and Stumm, W., Geochim. Cosmochim. Acta, 50 (1986): 18611869. 3 Surana, V.S. and Warren, H.J., Trans. Inst. Min. Metal. C, 28 ( 1969): 133-139. 4 Warren, I.H. and Devuyst, W., Leaching of metal oxides. In D.I. Evans and R.S. Shoemaker (Editors), Int. Syrup. on Hydrometallurgy. A.I.M.E., Chicago, (1973) pp. 229-264. 5 Cornell, R.M., Posner, A.M. and Quirk, J.P., J. Inorg. Nucl. Chem., 38 ( 1976): 563-567. 6 Osseo-Asare, K., Interfacial phenomena in leaching systems. In R.G. Bautista (Editor), Hydrometallurgical Process Fundamentals. Plenum, New York, (1982) pp. 227-268. 7 Nicol, M.J., The non-oxidative leaching of oxides and sulfides: an electrochemical approach. In. K. Osseo-Asare and J.D. Miller (Editor), Hydrometallurgy Research, Development and Plant Practice. A.I.M.E., New York, (1982) pp. 177-195. 8 Diggle, J.W., Dissolution of oxide phases. In: J.W. Diggle (Editor), Oxides and Oxide Films, Vol. 2. Marcel Dekker, New York, (1973) pp. 281-386. 9 Gorichev, I.G. and Kipriyanow, N.A., Regular kinetic features of the dissolution of metal oxides in acidic media. Russian Chem. Rev., 53(11 ) ( 1984): 1039-1061, and references therein. 10 Valverde, N. and Wagner, C., Ber. Bunsenges, Phys. Chem., 80(4) ( 1976): 330-333. 11 Valverde, N., Ber. Bunsenges. Phys. Chem., 80(4) ( 1976): 333-340. 12 Lu, Z.Y. and Muir, D.M., Hydrometallurgy, 21 ( 1988): 9-21. 13 Kunda, W., Rudyk, B. and Mackiw, V.N., Bull. Can. Inst. Min. Metall., 61 (1968) 819835. 14 Warren, I.H., Removal of iron oxide from silicate minerals. In: M.E. Wadsworth and F.T. Davis (Editor), Unit Processes in Hydrometallury, Vol. 1. A.I.M.E., New York (1964) pp. 300-307. 15 Azuma, K. and Kametani, H., Trans. Metall. Soc. AIME, 230 ( 1964): 853-862. 16 Kotrly, S. and Sucha, L., Handbook of Chemical Equilibria in Analytical Chemistry. Ellis Horwood, Chichester, Sussex, UK ( 1985 ). 17 Banwart, S., Davies, S. and Stumm, W., Colloids and Surfaces, 39 ( 1989): 303-309. 18 Laurence, G.S. and Ellis, K.J., J. Chem. Soc. Dalton Trans., 1667 ( 1972): 1667-1670. 19 Nash, K.L. and Horwitz, E.P., lnorg. Chim. Acta, 169 ( 1990): 245-252. 20 Rizkalla, E,N., Metal chelates of phosphonate-conlaining ligands. Rev. Inorg. Chem., 5 ( 1983): 223-304. 21 Gorichev, I.G., Gorsheneva, V.F., Kipriyanov, N.A. and Klyuchnikov, N.G., Kinetika i Kataliz, 21 (6) ( 1980): 1422-1425. 22 Habashi, F., Principles of Extractive Metallargy. Gordon and Breach, New York, Vol. l ( 1969 ) Chap. 7. 23 Paspaliaris, Y. and Tsolakis, Y., Hydrometallurgy, 19 ( 1982): 259-266. 24 Gorsheneva, V.F., Gorichev, I.G. and Klyuchnikov, N.G., Russian, J. Phys. Chem., 53(9) (1979) 1296-1298. 25 Vainman, S.K., Gorichev, I.G., Klyuchnikov, N.G. and Makordei, F.V., Kinetika i Kataliz, 17(5) (1976): 1323-1325. 26 Wadsworth, M.E. and Miller, J.D., Hydrometallurgical Processes. In: Rate Processes of Extractivd Metallurgy, H.Y. Sohn and M.E. Wadsworth, Eds., Plenum, New York, 1979. 27 Atkinson, R.J., Posner, A.M. and Quirk, J.P., J. Inorg. Nucl. Chem., 30 (1968): 23712381. 28 Gorichev, I.G. and Kipriyanov, N.A., Russian J. Phys. Chem., 55 ( 11 ) ( 1981 ): 1558-1568.
360 29 3(I 31 3:!
33 34 35
36 3"7
R. CH1ARIZIA AND E.P. HORW1TZ
Gorichev, I.G., Klyuchnikov, N.G., Bibikova, Z.P. and Boltovskaya, I.G., Russian J. Phys. Chem., 50(12) (1976): 1853-1855. Vainman, S.K., Gorichev, I.G. and Klyuchnikov, N.G., Russian J. Phys. Chem., 50(5) (1976): 803-804. Gatrone, R.C., Horwitz, E.P., Rickert, P.G. and Nash, K.L., Sep. Sci. Technol., 25 ( 1315) (1990): 1607-1627. Marsh, S.F., Ortiz, M.R. and Rein, J.E., Coprecipitation of uranium and plutonium oxalates using sodiumformaldehydesulfoxylate reduction and diethyloxalate hydrolysis precipitation. Los Alamos Rep. LA-5876-MS, Los Alamos (1975). Horwitz, E.P., Diamond, H. and Martin, K.A., Solv. Extr. Ion Exch. 5 ( 1987): 447-470. Bard, A.J., Parsons, R. and Jordan, J. (Editor), Standard Potentials in Aqueous Solutions. IUPAC, Marcel Dekker, New York (1985), p. 104. Grinberg, A.A., Nikolskaya, L.E. Petrzhak, G.I., Ptitsyn, B.V. and Filinov, F.M., Preparation of slightly soluble compounds of quadrivalent uranium using rongalite. In: Soviet Research on the Lanthanide and Actinide Elements, 1949-1957. Part I. Basic Chemistry, Consultants Bureau, New York, (1959), p. 65-67. Kolthoft; I.M. and Lingane, J.J., Polarography, Vol. II. Interscience, New York (1952), Second Ed., p. 562. Furrer, G. and Stumm, W., Geochim. Cosmochim. Acta, 50 ( 1986): 1847-1860.
339
Elsevier Science Publishers B.V., Amsterdam
New formulations for iron oxides dissolution R. Chiarizia and E.P. Horwitz Chemistry Division, Argonne National Laboratory, 9700 South Cass Avenue, Argonne, IL 60439, USA (Received May 15, 1990; revised version accepted March 12, 1991 )
ABSTRACT Chiarizia, R. and Horwitz, E.P., 1991. New formulations for iron oxides dissolution. Hydrometallurgy, 27: 339-360. The dissolution of 59Fe-labelled synthetic goethite, o~-FeOOH, has been studied in solutions of mineral (HCI, HNO3, H2S04) and organic acids, belonging to the families ofcarboxylic and diphosphonic acids, alone, or in the presence of reducing agents. The data have been interpreted with a firstorder rate law, although the shrinking core dissolution rate law has been shown to apply generally. Tile effect of several reducing agents on the rate of dissolution of o~-FeOOH by inorganic and organic acids has been measured. Very high accelerations of the dissolution process were exhibited by ascorbic acid, sodium dithionite and sodium formaldehydesulfoxylate. In particular, the mixture of 1-hydroxyethane-l,l-diphosphonic acid and sodium formaldehydesulfoxylate proved to be extremely effective for a-FeOOH dissolution even at room temperature.
INTRODUCTION
The dissolution of iron (III) oxides is a very important technological process with applications in several areas; such as: ( 1 ) leaching of oxides ores in hydrometallurgy; (2) removal of iron oxides from non-metallic minerals, such as kaolin and silica, of importance for the glass and ceramic manufacturing industry; ( 3 ) removal of the scale from a metal prior to rolling; (4) removal of deposits from thermal power engineering equipment. An understanding of the laws and mechanisms that rule the dissolution of iron oxides is also important for studies on corrosion and passivation of the metal, and in geochemical studies, because the dissolution of iron oxides plays an important role in the availability and transformations of iron in soils, water and sediments, and in biochemical reactions. Improved and more efficient ways to dissolve iron oxides could also have application in two other areas of relevance for the nuclear industry. The first one is the decontamination of steel equipment, where the removal of a surface layer of oxides could possibly facilitate the removal of radioactive contami0304-386X/91/$03.50
© 1991 Elsevier Science Publishers B.V. All rights reserved.
340
R. CHIARIZIA AND E.P. HORWITZ
nants, such as, the transuranium ( T R U ) elements. Such a decontamination procedure, if successful, would reduce the cost of long-term confinement of TRU-contaminated equipment. The second application is related to the retrieval and treatment of high-level nuclear waste, at present stored in single or double shell tanks at a number of Department of Energy sites. The nuclear waste solutions, generally acidic, have been neutralized in the tanks by addition of alkali, to reduce their corrosive power. If, because of environmental considerations, the waste has to be retrieved and further processed to remove the T R U component, the availability of effective leaching solutions for the metal oxides and hydroxides becomes a very important factor. Based on these considerations, we have performed an investigation of the dissolution rate of synthetic goethite ( a - F e O O H ) by several inorganic and organic acids, alone or in mixture with reducing agents, with the aim of finding new formulations capable of effectively dissolving iron oxides. Although iron is probably present in the waste predominantly as amorphous ferrihydrate (FeOOHam), goethite ( a - F e O O H ) has been chosen for the investigation as a worst case; because it is known that the dissolution of goethite is more difficult than that of hematite (a-Fe203) and especially of ferrihydrate 111,21.
Mechanism of iron oxide dissolution Two generally accepted mechanisms for the dissolution of metal oxides are reported in the literature. The first one is referred to as the "adsorption" mechanism [ 3-7 ], the second one as "electrochemical theory" [ 6-8 ]. It is also known that when two oxidation states of the metal forming the oxide are possible (as in the case of iron), the presence of redox couples in the leaching solution can lead to redox reactions which affect the dissolution rate [ 2,3,10-14 ]. An exhaustive discussion of the mechanism of the dissolution of metal oxides in the presence of redox couples can be found in [9]. Based on electrochemical considerations, the authors reached the conclusion that, in the case of iron oxide, the rate of dissolution can be expressed as: d a i.~0.5 ~0.5 ~-o.5 dt =,~.u+.ve2+.Fe3+
(1)
where: ~ = the fraction of oxide dissolved; a = activity, and the potential of the solution is determined by the reaction: Fe 3+ + e - ~ F e 2+
(2)
Equation ( 1 ) predicts that the presence of reducing agents capable of low(.'ring the Fe 3+/Fe 2+ ratio, will accelerate the iron oxide dissolution, as exper-
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
341
imentally observed. It also predicts that the order of reaction with respect to hydrogen ions is 0.5. This has been verified in a n u m b e r of investigations (see table 3 of [9] ), although the formal reaction order of H + can show different values when the dissolution is complicated by hydration, redox reactions, complex formation, and other factors [9 ], becoming as high as 1 in H F or even 2 in HC1 [ 15 ]. Another factor, besides acidity and solution potential, that strongly affects the dissolution of metal oxides, is the presence in solution of complexing agents. It has been recognized [ 15 ] that, in the case of iron oxides, for exampie, the rate constant for the dissolution increases linearly with the logarithm of the stability constant of the complex of Fe (III) with the anion of the electrolyte; that is in the order H C 1 0 4 < H N O a < H B r < H z S O 4 < H C I < H 3 P O 4 < HF. Equation ( 1 ) explains these findings by considering the effect that the presence of complexing agents has on the potential of reaction given in eq. (2). D e p e n d i n g on the nature of the ligand, the potential of eq. (2) is displaced into the cathodic region, which favors dissolution, or into the anodic region, which prevents dissolution. In general, the rate constant ofeq. ( l ) in the presence of complexing agents is linearly d e p e n d e n t on the difference between the stability constants of the complexes of Fe (III) and Fe ( II ) with the ligand. Based on the previous discussion, the ideal leaching agent for iron (III) oxides should be a strong acid of an anion which is a good ligand for F e ( I I I ) and possibly also a reducing agent of Fe 3+. It is very difficult to find all these characteristics in the same c o m p o u n d , because strong complexing agents are generally anions of weak acids. A c o m p o u n d that comes close to the above requirements is oxalic acid, a strong complexant of Fe 3+ (fl~ = 7.53, f12= 13.64, f13--= 18.49) with moderate acid strength (PKa~= 1.25) [16]. It is known to dissolve iron oxide quite effectively, especially at high temperature, and especially in the presence of L-ascorbic acid [ 17 ]. Ascorbic acid is also a somewhat effective dissolving m e d i u m [17], in spite of its low acid strength (pKaj = 4.17 ) [ 16 ], because of its reducing power toward Fe 3+ and its ability to form a complex with Fe 3+ in the intermediate state of the reaction [18]. The requirements of high acidity strength and complexing power toward Fe 3+ led us to give a special consideration, as possible dissolving agents for iron oxides, to the relatively little-studied d i p h o s p h o n i c acids. The commercially available 1-hydroxyethane- 1, l-diphosphonic acid ( H E D P A ) is, in fact, a quite strong acid (pKa~= 1.56, pK,2=2.20, PK~l+PK~2=3.76 as opposed to 4.59 for oxalic acid [ 19] ), and a very good complexant for Fe 3+ (fire3+ = 16.2, fife2+ = 3.0 [20] ). According to Gorichev et al. [21 ], who studied the dissolution of magnetite (Fe304) with HEDPA, the dissolved iron was present in solution in the divalent state, which would imply that H E D P A also has some reducing power toward Fe 3+. The potential of H E D P A for the removal of
342
R. CHIAR1ZIA A N D E.P. H O R W I T Z
magnetite deposits and scale from steel was fully recognized by the authors of [21 ].
Rate laws The kinetics of dissolution of oxides under chemical control, that is when the rate-limiting step is a surface reaction, can be expressed by the rate law:
1 - (1-o~ ) l/3=kCt=k' t pro
(3)
where: k and k' = kinetic constants; (7 = the concentration of the leaching agent (assumed constant ); r and p = radius and density, respectively, of the dissolving particles (assumed of spherical shape ); c~= the fraction of solid dissolved. Equation (3) is often referred to as the "shrinking core" model [ 22,23 ], or the "compressive sphere" model [24,25]. A plot of the left side of eq. (3) against t should result in a straight line having a slope k' (s- ~). A similar rate law would be observed if the rate-limiting step were diffusion through a limiting boundary layer of constant thickness, 3, at the receding interface [26]. It is not possible from the kinetics alone to determine if k' contains a rate constant for a surface reaction or includes diffusion through a limiting boundary layer of solution adjacent to the solid surface. To distinguish between the two, the effect of temperature must be measured. The activation energy for solution diffusion is usually of the order of 5 Kcal or less, while that for chemical reactions usually occurs between 10 and 25 Kcal. A particular case of the shrinking core model that is often met in practice is when a solid product, which coats the dissolving particles, is formed as a result of the dissolution reaction. This happens, for example, when elemental sulfur is formed in the leaching reaction [26] and is relevant for our investigation involving the use of sulfur-based reducing agents for the dissolution of goethite. The sulfur which forms can adhere to the mineral surface forming a tightly bonded diffusion layer and reduce the speed of the dissolution process. In this case various rate laws have been derived, which consider the diffusion of the reactant through a solid product layer. The most generally used is the Jander's equation:
[1--(1--a)l/3]2=k't
(4)
The limits of validity of eq. (4) as well as other better approximations are discussed at length in [22]. Equations (3) and (4) are necessarily approximate when the dissolving solution, as in our case, is not in very large excess
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
343
compared to the a m o u n t of dissolved oxide, but still useful for a qualitative interpretation of the kinetic data. Although our data, as will be shown in the next section, are generally quite well represented by the rate law given in eqs. (3) or (4), we found it more straightforward to interpret them according to a first-order rate equation, following the approach of Noyes and Whitney, summarized in [22 ]. By expressing the dissolution rate as:
~
= k ( C ~ s - C )
where: C = the concentration of dissolved metal at time t; Cois~--the concentration when the oxide has been completely dissolved; it tbllows upon integration that: In (Cdiss - -
C)
-
In
Cdiss =
--
kt
By expressing the concentration of the metal as dissolved fraction o~= C/Cdiss, we obtain: ln(1 -o~) = - k t
(5)
Equation (5) predicts that a semilogarithmic plot of the fraction of undissolved material, ( 1-o~) versus time, should give a straight line of slope k (s- l ), and that the kinetic constant is related to the 50% dissolution time, tl/2, by the familiar first-order rate expression: In 2 tl/2-- k
(6)
Equation (5) is a very simple tool for handling the dissolution rate data, but does not tell us whether a surface reaction or a diffusion process through a product layer is the rate-determining process. For this purpose the use of the shrinking core model (eqs. (3 and 4) ) is still necessary. EXPERIMENTAL
~-FeOOH Synthetic goethite was prepared following the method of Atkinson et al. [27]. The first step of the procedure consists in the preparation of a basic iron nitrate sol. During the preparation of the sol, containing 0.05 mol of Fe ( NO3 ) 3, 7.5 X 107 counts per minute (cpm) of 59Fe ( III ) (Amersham) were added to the solution. The sol was then aged for a few days. After aging, a gel was precipitated at 45 °C by addition of a slight excess of NaOH. The gel was heated for 4 d at 62°C, during which time crystal growth takes place [27].
344
R. CHIARIZIA AND E.P. HORWITZ
During the 4 day period the solid-liquid mixture was stirred manually twice daily. The slurry was filtered, the solid dried in air at 100°C, and ground in an agate mortar. The final product had a specific activity of 1.7× l04 c p m / mg dried solid; 5°mg of oxides dissolved in 2 ml of leaching solution gave a final 59Fe activity of 4.25 × 103 cpm/10¢tl and an iron concentration equal to 0.28 M. Another sample of synthetic goethite was prepared following the same procedure, but without the addition of 59Fe (III). Samples of inactive goethite were analyzed by Mtissbauer spectroscopy and X-ray powder diffraction. The results confirmed the presence of only orthorhombic o~-FeOOH.
Dissolution experiments The dissolution experiments were performed by contacting in screw-cap glass tubes a weighed amount of radioactive goethite (approximately 50 mg) with the dissolving m e d i u m at a volume to weight ratio of 40 (2 ml for 50 rag). The sizes of the particles of iron oxide were estimated to be in the range 60-120 mesh. The glass tubes, which contained small magnetic bars, were inserted into a large circulation-jacketed beaker containing glycerol. The temperature was maintained at the desired value by a circulating thermostat. A magnetic stirrer rotated the magnetic bars at such a speed (several hundred r p m ) that the iron oxide was suspended in the dissolving medium. The glass tubes containing the samples of a - F e O O H to be dissolved were thermally equilibrated in the thermostatically controlled bath for several minutes before adding the required volume of leaching solution. The m o m e n t when the solution was added to the glass tube was taken as time zero for the dissolution reaction. At various time intervals, the test tubes were removed from the bath, rapidly quenched by dipping in ice/sodium chloride-water mixture ( - 5 °C) and centrifuged. The dissolution reaction was considered terminated at the m o m e n t of immersion in the cold bath. After centrifugation, 10 ¢tl samples of the solutions were withdrawn, the tubes were reinserted into the hot bath and the agitation restarted. At this time the dissolution reaction was considered to begin again. The time required for the quenching-centrifugation-sampling procedure was less than 3 min. This time can be considered as negligible for the experiments characterized by a slow dissolution rate. The kinetic parameters obtained for experiments with a 50% dissolution rate of a few minutes can be affected by the relatively slow procedure described above. However, duplicate experiments involving fast kinetics showed that, by rigorously following the same sampling procedure, it was possible to obtain results reproducible within about 10%. All the results reported in the following as induction time, tj/2, kinetic constants and activation energies must be considered affected by at least a 10% uncertainty, when not otherwise specified. The activity of the samples withdrawn at time intervals was counted with a
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
345
Beckmann Biogamma well-type gamma counter. The results were converted in o~ (fraction dissolved) by measuring the final activity (all dissolved) or from the knowledge of the specific activity of the radioactive goethite. By plotting the data as o~ versus t, in most cases the usual S-shaped curves mentioned by many authors [9] were obtained. These curves are characterized by an initial period with a slow dissolution rate, an intermediate period with m a x i m u m dissolution rate, and a final period where again the dissolution rate is slow. The data were analyzed as log (1-o~) versus t. In most cases, straight lines were obtained, as predicted by eq. (5). In several cases, the data points were aligned on a straight line only after an induction time. The presence of such an induction or initiation time has been noticed by many authors (see, for example, [ 11,15, 25, 28-30 ] ). Although attempts have been made to correlate the induction time to the presence on the particle surface of potential nuclei capable of being activated under the reaction conditions [ 30 ], no rational explanations for these anomalies are available [ 11,28 ]. The presence of an induction time must also be taken into account when plotting the data according to the equations of the shrinking core model.
Reagents used in the dissolution experiments HC1, HNO3, H2504, oxalic, malonic, succinic, maleic, glycolic, a-hydroxyisobutyric, citric, tartaric, ascorbic and tetrahydrofurantetracarboxylic acid ( T H F T C A ) were analytical reagent grade products used without further purification. Furantetracarboxylic acid (FTCA) was provided by R.C. Gatrone, Chemistry Division, Argonne National Laboratory. 1-hydroxyethane-1, 1diphosphonic acid (HEDPA) was obtained from the Albright and Wilson Company and recrystallized from glacial acetic acid. Vinylidene-1,1-diphosphonic acid (VDPA) and 1,2-dihydroxyethane-l,l-diphosphonic acid (DHEDPA) were synthesized and purified as reported in [31 ]. VDPA and DHEDPA are new derivatives of HEDPA, with a higher acid strength than the parent c o m p o u n d ( P K a i +pga2= 3.47 for VDPA, 2.65 for DHEDPA, 3.76 for HEDPA) [ 19 ], but much more readily decomposable upon heating [ 31 ]. As reducing agents, analytical grade hydroxylamine hydrochloride (Ntt2OH-HC1), SnC12, hydroquinone, metallic zinc, Na2SO3, sodium dithionite (Na2S204) and sodium formaldehydesulfoxylate (HOCH2SO2Na) were used. The latter c o m p o u n d (hereafter called SFS) was obtained from Kodak under the commercial name of Rongalite. Although it has been used for a long time as a reducing agent in the textile industry and for reduction reactions of certain actinides [32,33 ], we are not aware of any previous use of SFS as an accelerator of metal oxide dissolution. The oxidation reactions and standard potentials of the above reducing agents can be found in [34].
346
R. CHIARIZIAAND E.P. HORWITZ
The chemical formulas of the organic acids and some reducing agents used in the experiments are reported in the Appendix. RESULTS AND DISCUSSION
Inorganic acids The a-FeOOH dissolution data obtained with mineral acids are plotted in Fig. 1 according to eq. (5). From the slope of the straight portion of the lines, the first-order kinetic constant k was calculated. The induction time, t~, when present, was obtained as the difference of the experimental 50% dissolution time, tl/2, and the theoretical half dissolution time calculated using eq. (6): In 2 ti=h/2-- k B
A
\
eL
HCI .01
0
50
1~'
100
150
Time (minutes)
.ol 0
50
100
50
100
HNO 3 150
Time (minutes)
H2SO4
150
Time (minutes)
Fig. 1. Dissolution of a-FeOOH by inorganic acids. (A) HCI. a = 2 M, 25°C; b = 2 M+0.1 M SFS, 25°C; c= 1 M, 80°C; d = 2 M+0.1 M Na2S204, 80°C; e = 2 M+0.1 MSFS, 8 0 ° C ; f = 2 M, 80°C. (B) HNO3. a = 1 M, 80°C; b = 1 M+0.1 MSFS, 25°C; c = 6 M, 80°C; d = 12 M, 80°C. (C) H2SO4. a = 3 M, 80°C; b = 3 M+0.1 M Na2SO3, 80°C; c = 3 M+0.1 M Na2S204, 80°C; d = 3 M+0.1 MSFS, 80°C.
347
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
TABLE 1 K i n e t i c p a r a m e t e r s for the d i s s o l u t i o n of a - F e O O H with i n o r g a n i c a c i ds D i s s o l v i n g so lution
2MHC1 2 M H C I + 0 . 1 M SFS 1 MHCI 2 MHCI 2 MHCI+0.1 MSFS 2 M HCI+0.1 MNazSzO 4 1 MHNO3 6 MHNO3 12 M H N O 3 3 M H2SO 4 3 M H2SO4 + 0.1 M Na2SO3 3 M H2SO4+ 1 M Na2S204 3 M H2SO4+0.1 M S F S
T
k a
t, a
t ~/ 2 a
(°C)
(s - I )
(min)
(min)
25 25 80 80 80 80 80 80 80 80 80 80 80
2)< 2.0)< 2.0)< 1.3X 1.1)< 6.7× 4.8 X 1.1)< 7.9)< 5.7)< 5.6 X 6.3)< 1.9)<
3.8 25 20 29 24 13 -
58 8.9 10 21 130 35 49 45 31 6.0
10 -619 10 -5 10 -4 10 -3 10 -3 10 - 4 10 -6 10 4 10 -4 10 4 10 -4 10 -4 10 -3
a-- 1 0 % bq- 1 0 0 %
This calculation was possible only when the 50% dissolution was actually achieved. The values of k ( s - I ) , ti (min) and tl/2 (rain) are listed in Table 1. The data obtained with HC1 confirmed the already known pronounced effect of acid concentration and temperature [ 3,5,15 ]. The calculation of the activation energy for 2 M HC1, using the Arrhenius equation [ 26 ], although performed from only two temperatures, provided a value for Ea equal to about 24 Kcal/mol in good agreement with the value of 20-24 Kcal/mol reported in [ 5 ] for 0.5 MHC1 and various crystal morphologies of natural goethite. The addition o f NazS204 or SFS did not bring about an acceleration of the a-FeOOH dissolution rate at 80°C. The presence of Na2S204 actually reduced the kinetic constant by a factor of 2. An increase of k of about 10 times was caused by SFS at 25 oC, but the dissolution process remained very slow. The different behavior of SFS and Na2S204 can be probably explained by the observation that the presence of the latter led to the formation of colloidal sulfur. Elemental sulfur formed during the dissolution of c~-FeOOH coats the dissolving particles, slowing the dissolution process. The formation of colloidal sulfur has been reported in [35 ] in a study of the reduction of acid solutions of UO 2+ by Na2S204, and is attributed to the rapid disproportionation of dithionite into thiosulfate and sulfite: 2 $2042- + H 2 0 ~ $ 2 0 2 - + 2 HSO~followed by the disproportionation of thiosulfate:
3,48
F,. CHIARIZIA AND E.P. HORWITZ
$202- + 2 H+--+H20+SO2+S In agreement with our observations, the formation of colloidal S was not noticed by the authors of [ 35 ] if SFS was used instead of Na28204. Although sulfoxylic acid is also known to be unstable [ 34 ], decomposing according to: 2 H2802--+$2 O2- + 2 H + + H 2 0
which should also lead to the formation of elemental sulfur, its decomposition may be retarded by the presence of the - C H 2 O H moiety in the SFS molecule. The formaldehydesulfoxylate ion, CH20-HSO2, is known to be rather stable and this property is utilized in polarographic procedures for the analysis of commercial dithionite preparations [36 ]. The formaldehyde-stabilized sulfoxylate anion can then be oxidized directly to sulfite, according to: H2SO2 + H20-+H2803 + 2e + 2 H + "['he failure of SFS to accelerate the dissolution of a-FeOOH in 2 M HC1 at 80 °C is, however, not completely understood. The oz-FeOOH dissolution data in HNO3 show that a very high concentraA
B
0.80
m
0.80
HCI
0.60
~
09 0.8 07 ~
~:- 0.40
i
0.20 0.00
20
Time 40
~
HNO 3
0.60 09
d_.c
O8
0.40
07
0.6
0.6
0.20
04 O2
60
02
0.00
80
(minutes)
20
0
cy
Tim(mi e nutes) 40
60
80
C
.2So4
0.80 0.60
09 08
b
0.40
07 06
0.20
04 0.2
0.00
0
20
Time 40
60
80
(minutes)
Fig. 2. Same data o f Fig. 1 plotted according to eq. ( 3 ) . For the meaning of the symbols see Fig. 1.
349
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
tion of the acid ( > 12 M) is required to achieve relatively fast dissolution of the oxide at 80°C, with k approaching the value of 10 -3 s-1. An induction time in the range 20-25 min was observed. The attempt to use SFS as an accelerator failed even at 25°C (curve b of Fig. 1B). In the presence of 1 M HNO3 + 0.1 M SFS after an initial acceleration, the dissolution process became again very slow, after about 2 h, indicating the destruction of SFS by HNO3. The data for H2SO 4 (Fig. 1C) show that, in this non-oxidizing medium, reducing agents are effective in accelerating the dissolution. While sulfite and dithionite have only a modest effect on the value of k, and primarily reduce the induction time, with SFS the dissolution rate is increased by a factor of about 6, while the induction time is reduced practically to zero. With dithionite, the formation of elemental S was again noticed. The better performance of SFS as compared to Na28204 may be explained as in the case of HCI data. If the data of Fig. 1 are plotted according to the shrinking core model--eq. ( 3 ) - - t h e plots of Fig. 2 are obtained. The data are well aligned on straight lines, when allowance is made for the induction time; indicating that the rate of the dissolution process is primarily controlled by a surface chemical reaction or by diffusion through a limiting boundary layer. With HCI the measured value of the activation energy excludes diffusion as a controlling factor. No activation energy measurements were performed with HNO3 and H2SO4.
Carboxylic acids Figures 3 and 4 report the data on the dissolution rate of a - F e O O H by oxalic acid alone or in the presence of reducing agents. The kinetic parameters are listed in Table 2, together with the results of the measurements performed with other carboxylic acids. The data in Fig. 3A and Table 2 show B
--q:>~ d
1.41 1.2 co
"
1.0
09
0.8
.1
0.6k t
O8 a
,,/,,c
d
07 06 04 02
01
100 200 Time (minutes)
o o~ 0
'-~+~ 50 1O0 150 Time (minutes)
2OO
Fig. 3. Dissolution of a - F e O O H by oxalic acid. (A) a = 1 M, 80°C; b = 0 . 5 M, 80:C; c = 1 M, 60~C; d = 1 M, 40°C; e = 1 M, 25°C. (B) Same data as in A (a,c,d, and e) plotted according to eq. (3).
350
R. C H I A R I Z I A A N D E.P. H O R W I T Z
A
z~
T-
.1
co
1.0
"~" --
0.6 0.4 0.2 0.6
,,/~/,~/~::
-.~
U 0
i 20
40
,
i
j,,~
c _ 0.9 ~.~ oZ ~' 0.2
/ ~
0.4
~
~-~
~ .01
B
f 0.4
/
~//
dG, d/
~
~
c o.~
0.2
o.8
0.0
60
o.~ 0
10
Time (minutes)
20
30
Time (minutes)
Fig. 4. Dissolution of c~-FeOOH by oxalic acid in the presence of reducing agents. ( A ) a = l M, 6 0 ° C ; b = 1 M + 0 . 0 1 MNazSO3, 60°C; c = 1 M + 0 . 0 1 MNa2S204, 6 0 ° C ; d = 1 M + 0 . 0 1 M S F S , 60°C; e = 1 M + 0 . 0 1 M ascorbic acid, 8 0 ° C ; f = 1 M + 0 . 0 1 M S F S , 80°C. ( B ) Same data as in A (c,d,e a n d f ) plotted according to eq. ( 3 ) ( u p p e r p a r t ) and to eq. ( 4 ) (lower part ).
'TABLE 2 Kinetic parameters for the dissolution of c~-FeOOH with carboxylic acids
Dissolvings o l u t i o n 0.5 M succinic or glycolic or o~-hydroxyisobutyric acid 0.5 M maleic or citric or tartaric acid 0.5 M malonic acid 1 M THFTCA or FTCA 1Mascorbicacid FTCA +0.1 M ascorbic acid 1 M H2C204 I M H2C204+0.01MSFS IMH2C204 1 M H2C204 1M H2C204+0.01MSFS 1 M H2C204+0.01 MNa2SO3 IMH2C204+0.01MNa2S204 0.5 MH2C204 1 M H2C204 I MH2C2O4+0.1 M H E D P A 1 MH2C204+0.01Mascorbic acid 1M H2C204+0.01MSFS a± 10% b+ 100%
T
ka
ti a
lj/2 a
(°C)
(s - l )
(min)
(min)
80 80 80 80 80 80 25 25 40 60 60 60 60 80 80 80 80 80
NO DISSOLUTION 1 × l 0-6 b
_
-
4×
10 -6b
-
-
IX10 -6b
-
9.3)<10 5 4.1)<10 -4 5.0× 10 -5 5.2X 10 -4 1.1×10 -4 5.2×10 -4 2.5)<10 -3 4.0)< 10 - 4
1.6×10 3 1.2×10 -3 2 . 1 × 1 0 -3 9.4× 10 -5 4.7× 10 3 6.3X 10 -3
5.8 39 49 25 15 16 _ 3.4 3.5 17 _ -
-
130 67 280 22 130 37 4.6 45 7.2 13 9.0 140 2.5 1.8
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
35 [
that the effect of temperature is more pronounced than the effect of the acid concentration. An activation energy of 16 Kcal/mol has been calculated from the data of Fig. 3A, well in the range of kinetic control by surface chemical reactions. The same data, plotted in Fig. 3B according to eq. (3), show that the shrinking core model also applies to this case. Attempts to accelerate the dissolution reaction by increasing the acidity of the solution by adding HC1 or HNO3 to 1 M H 2 C 2 0 4 caused a slight opposite result; probably by depressing the dissociation of oxalic acid, reducing the concentration of the active ligand, C2042- . A more evident reduction in k was caused by the addition of HEDPA to the 1 M oxalic acid solution. This may be an indication that HEDPA has some reducing power towards Fe (III), as claimed in [21 ]. Fe(II) oxalate is, in fact, only sparingly soluble and its formation on the surface of the dissolving particles would reduce the dissolution rate. Other explanations, based on the different adsorption of the oxalate and EDPA anions on the goethite surface, are, however, available. The precipitation of FeC204 or, rather, the transformation of some of the o~-FeOOH into FeC204, was obtained in the experiments where the dissolving m e d i u m was 1 M oxalic acid containing 0.1 M ascorbic acid or SFS at 60 ° C. The dissolution was initially faster, but after going through a maximum, it stopped at about 25% dissolution, probably corresponding to the solubility of Fe204 in the dissolving medium. All other dissolution experiments in the presence of reducing agents were performed at a much lower concentration of the reducing agent. The results are reported in Fig. 4 a n d / o r in Table 2. Na2SO3 is practically ineffective; Na2S204 doubles the first-order rate constant at 60 °C and generates elemental sulfur; SFS is the most effective, raising k by a factor of 3, 5 and 10, at 80, 60 and 25°C, respectively. An effect similar to SFS is shown by 10 -2 M ascorbic acid at 80°C. The acceleration of c~-Fe203 dissolution by an ascorbate-oxalate mixture has been recently reported [ 17 ]. Figure 4B shows the data obtained with some oxalic acid-reducing agents mixtures plotted according to eqs. ( 3 ) and (4). The fact that eq. (4) provides a better fit of the data indicates that in this case the rate-controlling process is the diffusion of the reagents through a solid surface layer, presumably FeC204, covering the dissolving particles. Regarding the other carboxylic acids used in this work, the data of Table 2 show that a significant dissolution of o~-FeOOH takes place only with ascorbic acid and, to a much lesser extent, with malonic acid. The former acid is a reducing agent, which may explain its effectiveness. Especially striking is the difference of the constant k for oxalic acid and malonic acid ( 1.2× 1 0 - 3 v e r s u s 4.1 × 1 0 - 6 S l, at 0.5 M a n d 80°C). Since malonic acid is also a very good complexing agent for Fe 3+ [ 16 ], the difference has to be ascribed either to its lower acidity or to the presence of a CH2 group in the molecule that makes the interaction of the anion of the acid with the surface of the oxide more difficult. With succinic acid, having another CH2 group and an even lower
352
R. C H I A R I Z I A A N D E.P. H O R W I T Z
acidity, there is no measurable dissolution of oz-FeOOH. Similar results have been reported in [37 ] in the study of the dissolution of 5-A1203 by oxalic, malonic and succinic acid, and have been attributed to the formation on the oxide surface of five-, six-, and seven-membered chelate rings, with the fivernembered rings more readily detachable from the surface. The tetracarboxylic acids also performed very poorly. However, the mixture of FTCA with ascorbic acid showed a remarkable enhancement of the dissolution rate compared to the FTCA alone, (k increased by a factor of 400) or to ascorbic acid alone (k larger by a factor of 5). No experiments have been run with tetracarboxylic acids and SFS or Na2S204, because of the limited availability of the acids. An acceleration effect similar to that of ascorbic acid would, however, be expected.
Diphosphonic acids Figure 5 shows a comparison of the oz-FeOOH dissolution data in the three acids investigated, HEDPA, VDPA and DHEDPA. The first two acids have practically the same (disappointingly low) kinetic constant, although VDPA shows a somewhat lower induction time. DHEDPA, on the other hand, is much slower. Data with DHEDPA have been obtained at 60°C instead of 80 ° C, because this acid spontaneously decomposes, even at room temperalure [31 ]. For all three acids the shrinking core model applies (figure not .shown for brevity). The data obtained with the three diphosphonic acids in the presence of reducing agents, at different concentrations and in the temperature range 2580°C are summarized in Table 3. Some data with HEDPA are reported in Fig. 6 ( A - D ) . A very high acceleration of the dissolution rate by 1 M H E D P A at 80 °C is exhibited by ascorbic acid, dithionite, and SFS, with a kinetic constant increase of about 100. A strong acceleration is also shown at 80°C by Na2SO3 and Zn, which affects the dissolution rate through the reducing power
\\ \ 01 0
2
4 6 Time (hours)
8
10
Fig. 5. Dissolution of ~ - F e O O H by diphosphonic acids, a = 0 . 5 M HEDPA. 80°C: b = 1 M HEDPA. 80~C; c = 1 M VDPA, 80°C; d = 1 M D H E D P A , 60°C.
353
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
TABLE 3 K i n e t i c p a r a m e t e r s for t h e d i s s o l u t i o n o f ~ - F e O O H w i t h d i p h o s p h o n i c a c i d s Dissolving solution
T (°C)
ka (s 1)
0.5 M H E D P A M HEDPA M VDPA M DHEDPA
80 80 80 60
2.7× 1.8 × 1.7 × 4×
25 25 25 25 25 25 25 25 25
1 × 1 0 6b 1 × 1 0 6b 2.3?< 10 -5 2 . 2 × 10 -4 5.0X10 4 1.9N10 3 2 . 3 × 10 -3 2 . 5 X 10 4 3 . 4 × 10 4
_ _ -47 23 _ 8.0 _ -
_ _ 500 100 46 6.1 13 46 34
M M M M M M M M M
HEDPA+0. l Mhydroquinone H E D P A + 0. l M N H 2 O H " H C I H E D P A + 0 . 1 M SnCI2 H E D P A + 0.1 M Na2SO3 H E D P A + (0.2 M ) Z n C HEDPA+ 1 MNazS204 HEDPA+0.1 MSFS V D P A + 0.1 M SFS DHEDPA+0.1 MSFS
10 . 5 10 -- 4 10 - 4 10 - 6 b
ti a (min)
tl/2 a (min)
52 120 82 -
480 190 [ 50 -
1 M H E D P A + 0.1 1 At H E D P A + 0.1 1 M H E D P A + 0.1 1 M HEDPA+0.1 1 MDHEDPA+0.01
M ascorbic acid M Na2SO3 M Na2S204 MSFS MSFS
40 40 40 40 40
1.0X 10 5 1.5)< 10 -3 5 . 7 X 10 - 3 7.1 × 10 - 3 1.2X10 3
_ 5.3 2.7 _
_ 13 1.7 4.3 9.6
1M 1M 1M 1M 1M
M ascorbic acid MNa2SO3 MNa:S204 MSFS MSFS
60 60 60 60 60
1.4×10 3 2 . 9 X 10 . 3 1.1X 10 - z I.IXIO 2 5 . 8 × 10 -3
30 1.2 _ -
38 5.2 1.0 1.0 2.0
80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80
2 . 6 X 10 - 4 2 . 4 × 10 - 4 1.4X 10 -3 8.1X10 3 1 . 6 × 10 . 2 9.1XI0 3 1 . 6 × 10 -2 1.6X 10 . 2 1 . 6 × 10 - 2 8.4×10 3 4 . 8 X 10 -3 1 . 4 × 10 -2 1.4×10 2 1.4X 10 -2 8.5 X 10 -3 4.2 X 10 -3 2.4×10 3 4 . 6 × 10 -3 2 . 9 × 10 - 3
14 62 4.7 1.1 4.8 44s d _ _ -
58 110 13 2.5 5.5 2.0 45 s o 45 s d 45 s a 1.4 2.4 49 s d 49s d
M M M M M M M M M M M M M M M M M M M
H E D P A + 0.1 HEDPA+0.1 HEDPA+0.1 HEDPA+0.1 DHEDPA+0.1
H E D P A + 0.1 M h y d r o q u i n o n e HEDPA+0.1 M NH2OH.HCI H E D P A + 0.1 M SnC12 H E D P A + (0.2 M ) Z n c HEDPA+0.1 Mascorbic acid H E D P A + 0 . 1 M Na2SO3 H E D P A + 0.15 M N a 2 S 2 0 4 H E D P A + 0.1 M N a 2 S z O 4 H E D P A + 0.05 M Na2S204 H E D P A + 0 . 0 1 M Na2S204 H E D P A + 0 . 0 0 5 M Na2S204 HEDPA+0.15 MSFS HEDPA+0.1 MSFS HEDPA+0.05 MSFS HEDPA+0.01 MSFS HEDPA+ 0.005 MSFS V D P A + 0.01 M a s c o r b i c a c i d VDPA+0.1 M ascorbic acid V D P A + 0.02 M SFS
" + 10% b_+ 100% CThe c o n c e n t r a t i o n refers to t h e Z n 2+ in s o l u t i o n a f t e r c o m p l e t e d i s s o l u t i o n o f the m e t a l . aS = s e c o n d s .
49
sa
1.4 2.7 4.8 2.5 4.0
R. CHIARIZIAAND E.P. HORWITZ
35;4 A
1
~ , ~-..
a
1TTx
b i
\
.01
20
%
J
\\,
h'~
80 C
10
0
\
~g
1 g ',, +,e
x\
.01
30
0
c
60 C 10
20
30
~,
25C
D
d
d
!~
-
.1:
1
g'\
x 40'C .01
"g
20 40 Time (minutes)
60
.01
20 40 Time (minutes)
60
Fig. 6. Dissolution of o~-FeOOH by 1 M HEDPA in the presence of reducing agents al ( A ) 80cC, ( B ) 60°C: ( C ) 4 0 ° C and ( D ) 25°C. Reducing agents: a = 0 . 1 M NH2OH-HCI; b = 0 . 1 .'~l hydroquinone; c = 0.1 M SnCI2; d = 0.1 M ascorbic acid; e = (0.2 M ) Zn; j - 0.1 M Na2SO3; g = 0 . 1 M S F S : h = 0 . 1 MNa2S204.
of native hydrogen generated during its dissolution by HEDPA. Hydroxylamine and hydroquinone, on the other hand, show a negligible effect, while Sn(II) occupies an intermediate position. The behavior of ascorbic acid is very characteristic. Its effectiveness as a dissolution accelerator is strongly influenced by the temperature. Even at 60°C it is much less effective than dithionite and SFS. The activation energy for the dissolution of o~-FeOOH by 1 M HEDPA in the presence of 0.1 M ascorbic acid has been calculated from the data of Table 3 and is equal to 34 Kcal/mol. This very high value of Ea indicates that a complex chemical interaction between Fe (III), HEDPA and ascorbic acid takes place and is rate-determining. The comparison of the data of Fig. 6A and D (80°C and 25°C, respectively) shows that the effect of Sn ( II ), Zn and Na2SOB on the dissolution rate of o~-FeOOH by 1 M HEDPA also depends strongly on the temperature, with activation energies in the range 10-15 Kcal/mol. With these reducing agents, therefore, chemical reactions seem to still be, at least in part, the rate-determining process for the iron oxide dissolution. A completely different behavior is exhibited by the dithionite and sulfoxylate species. They show a very pronounced accelerating effect even at 25 oC, where the dissolution of c~-FeOOH by the diphosphonic acids alone is
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
355
negligible. With both reducing agents, the application of Arrhenius' equation to the data of Fig. 6 allows calculation of an activation energy of 5.7 Kcal/ mol; which shows that, in this case, a diffusion process, not a chemical reaction, is the rate-limiting factor. The formation, at all temperatures, of colloidal sulfur was again observed when Na2S204 was used as an accelerator. This fact does not seem to have any effect on the goethite dissolution reaction at 80 and 60 °C. However, it does change the shape of the dissolution rate data at 40°C and especially at 25°C (Fig. 6C and D). At these temperatures the data for Na2S204 show a curvature toward slower dissolution rates. In other words, the initially extremely fast dissolution rates of oz-FeOOH become progressively slower, requiring a much longer time than with SFS to proceed to completion. This behavior of Na2S204 may be ascribed to the progressively slower diffusion, with decreasing temperature, of the reagents through a layer of sulfur coating the dissolving particles. Supporting evidence for sulfur coating is shown by Fig. 7, where the dissolution data at 25 °C with 1 M HEDPA in the presence of 0.1 M Na2S204 or SFS is plotted according to eqs. (3) and (4). With Na2S204, the data are better aligned when plotted as required by Jander's equation (eq. (4)). This indicates that, at low temperature, the rate-determining process is the diffusion of the reagents through a solid product layer. When VDPA or DHEDPA were used instead of HEDPA, acceleration effects similar to those described in the case of HEDPA were observed with ascorbic acid at 80 °C and with SFS throughout the whole temperature range. With VDPA, however, the presence of 0.1 M SFS seemed to promote a rapid alteration or decomposition of the diphosphonic acid, leading to the formation, after a few minutes at 80°C, and within 1 hour at 25°C, of a massive 0.8 -~
0.6
"~ 0.4 EL ~- 0.2 A
0.0
~[ ~: :,% 0.2 B
09 08 07 06 0.4 0.2 O0
h
0.9 h .~...//~g
O.C
,
0
~/)¢/ )~ .
10
~0.8~0.7 .
.
~ O.6 40.4
.
20
30
Time (minutes)
Fig. 7. Data for the dissolution of ~ - F e O O H by 1 M H EDPA in the presence of 0.1 M SFS (g) or 0.1 M Na2S204 (h) at 25 oC, plotted according to (A) eq. ( 3 ) and (B) eq. (4).
356
R. CHIARIZIA AND E.P. HORWITZ
whitish precipitate, presumably of FePO4. The same kind of precipitate was observed after a longer time in the solutions containing 1 M DHEDPA and 0.1 M SFS, after dissolution of a-FeOOH. The iron-SFS-promoted decomposition of VDPA and D H E D P A is at present under investigation in our laboratory. The data in Table 3 show that with HEDPA there is no correlation between the final concentration of iron in solution and the concentration of either Na28204 or SFS. Both Na28204 and SFS maintain their accelerating effect at a concentration as low as 5 X 10-3M. These results seem to indicate that SFS or Na2S204 are required only for the reduction and subsequent removal of a surface layer of Fe (III) atoms. The successive layers of oxide must be activated by this process and more easily dissolved by HEDPA, possibly through a redox mechanism, as reported in [21 ]. If FeOOH is dissolved in the presence of a substoichiometric amount of SFS, in fact, a colorless solution forms. On the other hand, when no SFS is present, iron is dissolved slowly by HEDPA as yellow Fe(III) contrary to the findings of [21 ]. The real nature of the interaction between Fe (III) and Fe (II) with HEDPA, however, requires further investigation. With MnO2, where the redox couples of the metal is characterized by a more positive Eo value than the Fe3+/Fe 2+ couple, HEDPA clearly shows its reducing power. MnO2 is dissolved by 1 M HEDPA slowly at room temperature and extremely rapidly at 80 ° C, generating a purple complex of Mn (III) stabilized by HEDPA. In the presence of SFS, the reduction of MnO2 proceeds further, and a colorless solution of M n ( I I ) is obtained, with the dissolution reaction becoming almost instantaneous even at 25°C. Preliminary experiments, performed in the same conditions as with o~FeOOH, showed that SFS also has a very pronounced acceleration effect on the dissolution of CeO2 by 2 M HC1 at 80 ° C. With an HEDPA-SFS mixture the dissolution of CeO2, after reaching 25% dissolution rapidly, practically stops due to the probable decomposition of HEDPA and precipitation of either CePO4 or the C e - H E D P A complex. In contrast, when no SFS is present, practically no CeO2 is dissolved by either 2 M HC1 or 1 M HEDPA over a period of 4 h. SUMMARY AND CONCLUSION
A pronounced acceleration effect by some reducing agents on the rate of dissolution of synthetic c~-FeOOH by sulfuric acid, oxalic acid and diphosphonic acids has been measured. The most effective reducing agents, among those tested, are ascorbic acid, sodium dithionite and sodium formaldehydesulfoxylate (SFS). The acceleration effect is relatively small in H2804 and in I-~2C204. In oxalic acid the concentration of the reducing agents must be kept very low to prevent precipitation of FeC204. The risk of FeC204 precipitation, together with the known capability of oxalic acid to rapidly dissolve iron
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
357
oxides without reducing agents (at least at high temperature), makes the acceleration effect in oxalic acid of limited value. With diphosphonic acids, showing by themselves a rather low dissolution rate of a-FeOOH, in spite of their high acidity and complexing power for Fe(III), the accelerating effect brought about by the reducing agents is remarkable. It is best shown by SFS, because ascorbic acid is effective only at high temperature, and Na2S204 rapidly decomposes in the acidic medium generating colloidal sulfur which can slow down the dissolution process at low temperature. Although the chemistry involved in the dissolution of o~-FeOOH by a HEDPA-SFS mixture is not completely understood, this novel system, investigated for the first time in this work, could have practical applications in all those fields where a rapid dissolution of iron oxides at room temperature is required. The capabilities of HEDPA-SFS mixtures to dissolve metal oxides can probably be extended to all other cases of oxides of metals having lower oxidation states. Preliminary experiments have shown that the presence of SFS in 1 M H E D P A also accelerates the dissolution of MnO2 and, to some extent, of Ce02 and PuO2. ACKNOWLEDGMENTS
This work was performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Sciences, U.S. Department of Energy, under contract number W-31-109-ENG-38. The authors wish to thank Dr. Ralph Gatrone and Dr. Kenneth Nash for providing samples of FTCA, VDPA and DHEDPA, and Dr. K.A. Martin, formerly at the Institute of Gas Technology, Chicago, and Dr. L. Soderholm, of Argonne National Laboratory, for the characterization of the a-FeOOH used in this work by M6ssbauer spectroscopy and X-ray powder diffraction, respectively. APPENDIX:
Chemical formulae of the organic acids and some reducing agents used oxaric malonic succinic maleic glycolic
HOOC-COOH HOOC-CH2-COOH HOOC-CH2-CH2-COOH HOOC-CH=CH-COOH HO-CH2-COOH
358
R. C H I A R I Z I A A N D E.P. H O R W I T Z
Q -hydroxybutyric
/COOH (0H3)20\O H COOH
I
citric
HOOC-CH~I-CHzCOOH OH
tartaric
COOH-CHOH-CHOH-COOH CH2OH
ascorbic H ~ O H
tetrahydrofurantetracarboxylic THFTCA
furantetracarboxylic FTCA
hyd roxyethane-l,l-diphosphonic
HEDPA
vinylidene-l,l-diphosphonic VDPA
2-dihydroxyethane-l,l-diphosphonic DHEDPA
CH~',,,_/~(OH)= Ho/C'~(OH)z / ~(OH) 2
H2C--C'~I~(OH)2
HOCH2\_/~(OH)2 Ho/C"~(O H)2
hydroquinone
sodium formaldehydesulfoxy[ate SFS
HOCH2\s_ O- Na +
NEW FORMULATIONS FOR IRON OXIDES DISSOLUTION
359
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360 29 3(I 31 3:!
33 34 35
36 3"7
R. CH1ARIZIA AND E.P. HORW1TZ
Gorichev, I.G., Klyuchnikov, N.G., Bibikova, Z.P. and Boltovskaya, I.G., Russian J. Phys. Chem., 50(12) (1976): 1853-1855. Vainman, S.K., Gorichev, I.G. and Klyuchnikov, N.G., Russian J. Phys. Chem., 50(5) (1976): 803-804. Gatrone, R.C., Horwitz, E.P., Rickert, P.G. and Nash, K.L., Sep. Sci. Technol., 25 ( 1315) (1990): 1607-1627. Marsh, S.F., Ortiz, M.R. and Rein, J.E., Coprecipitation of uranium and plutonium oxalates using sodiumformaldehydesulfoxylate reduction and diethyloxalate hydrolysis precipitation. Los Alamos Rep. LA-5876-MS, Los Alamos (1975). Horwitz, E.P., Diamond, H. and Martin, K.A., Solv. Extr. Ion Exch. 5 ( 1987): 447-470. Bard, A.J., Parsons, R. and Jordan, J. (Editor), Standard Potentials in Aqueous Solutions. IUPAC, Marcel Dekker, New York (1985), p. 104. Grinberg, A.A., Nikolskaya, L.E. Petrzhak, G.I., Ptitsyn, B.V. and Filinov, F.M., Preparation of slightly soluble compounds of quadrivalent uranium using rongalite. In: Soviet Research on the Lanthanide and Actinide Elements, 1949-1957. Part I. Basic Chemistry, Consultants Bureau, New York, (1959), p. 65-67. Kolthoft; I.M. and Lingane, J.J., Polarography, Vol. II. Interscience, New York (1952), Second Ed., p. 562. Furrer, G. and Stumm, W., Geochim. Cosmochim. Acta, 50 ( 1986): 1847-1860.