The reductive dissolution of iron oxides by ascorbate

The Reductive Dissolution of Iron Oxides by Ascorbate The Role of Carboxylate Anions in Accelerating Reductive Dissolution M A R I A D O S S A N T O S A F O N S O , * P E D R O J. M O R A N D O , * M I G U E L A. BLESA, *'1 S T E V E N B A N W A R T , ] " A N D W E R N E R S T U M M t '1 *Departamento de Qufmica lnorgdnica, Analitica y Quimica Fisica, Facultad de Ciencias Exactas y Naturales, Universidad de Buenos Aires, Ciudad Universitaria, Pabell6n 2, N~gez, Buenos Aires, Argentina, and ~flnstitute for Water Resources and Water Pollution Control (EA WAG), Swiss Federal Institute of Technology (ETH), Zurich, Switzerland

Received January 19, 1989; accepted October 2, 1989 There are four general pathways of dissolution of reducible metal oxides in acidic aqueous solution: proton-assisted (acid), ligand-promoted acid, reductive, and ligand-promoted reductive dissolution. The presence and reactivity toward the surface of protons, chelating ligands, and reductants dictate the mechanism (s) controlling the dissolution. For the massive reductive dissolution of magnetite by ascorbic acid, the experimental rate law R = k [HA - ] 1/ 2[H +] suggeststhe involvement of surface ~FelHA - complexes. Adsorption isotherms of ascorbic acid onto hematite at pH 3 and 25°C yield a Langmuir-type surface complexation constant Ks = (9.57 × 108 M-~). Slow dissolution follows with an empirical rate law R = kobs(~--FenIA). It is concluded that the formation and kinetic reactivity of surface complexes determine the rate of dissolution. Dehydroascorbic acid also dissolves magnetite, but at slower rates. Oxalate accelerates the reductive dissolution of hematite by ascorbate even though it competes with ascorbate for surface sites; enhanced detachment of ~ F e n surface species by oxalate complexation may be involved. Autoacceleration of the reductive dissolution by dissolved Fe~Lcarboxylatecomplexes is observed in EDTA / ascorbic acid mixtures; the rate reaches a maximum at intermediate [EDTA ] values, where synergistic effects between EDTA and Fen-EDTA complexes are important. Autoacceleration may also operate in oxalate solutions. © 1990AcademicPress,Inc. INTRODUCTION

o f the surface m e t a l leaving group. P r o t o n s b o u n d to a d j a c e n t lattice o x y g e n ions a n d ligands b o u n d to m e t a l ions accelerate the diss o l u t i o n b y w e a k e n i n g the c o v a l e n t b o n d s bet w e e n the surface m e t a l a t o m a n d the lattice oxygen. In i r o n ( I I I ) oxides, acid a t t a c k is generally less i m p o r t a n t t h a n the a t t a c k b y red u c t a n t s (2, 9 - 2 1 ). T w o b r o a d classes o f red u c t a n t s are effective: r e d u c i n g m e t a l ions, usually in the f o r m o f c a r b o x y l a t e c o m p l e x e s , a n d r e d u c i n g o r g a n i c a n i o n s able to f o r m stable c o m p l e x e s with iron. In principle, all organic m a t t e r is t h e r m o d y n a m i c a l l y a p t to red u c e F e ( I I I ) , a n d d u e to the high affinity o f c a r b o x y l a t e g r o u p s for F e ( I I I ) in s o l u t i o n or o n t h e surface o f a h y d r o u s i r o n ( I I I ) oxide, this t y p e o f process is very likely to o c c u r in n a t u r a l water, soils, a n d sediments. T h e c o m -

T h e d i s s o l u t i o n o f i r o n oxides plays a n i m p o r t a n t role in several fields o f a p p l i c a t i o n o f c h e m i s t r y , such as the passivity o f i r o n alloys, the d e v e l o p m e n t o f f o r m u l a t i o n s for the chemical cleaning and decontamination of steel surfaces ( 1, 2 ) , the availability a n d transf o r m a t i o n s o f i r o n in soils, water, a n d sedim e n t s , a n d t h e b i o c h e m i s t r y o f i r o n ( 2 - 6 ) . In principle, two c h e m i c a l p a t h w a y s m a y lead to dissolution: a t t a c k b y acids ( u s u a l l y in the presence o f c o m p l e x i n g agents) or a t t a c k b y r e d u c t a n t s . In t h e case o f n o n r e d u c i b l e m e t a l oxides (7, 8), fast p r o t o n a n d ligand exchanges o n t h e surface p r e c e d e t h e slower d e t a c h m e n t

i To whom all correspondence should be addressed. 74

0021-9797/90 $3.00 Copyright © 1990 by Academic Press, Inc. All rights of reproduction in any form reserved.

Journal of Colloid and Interface Science, Vol. 138, No. 1, August 1990

DISSOLUTION

OF IRON OXIDES

bination of a reductant and a chelating ligand (dithionite-citrate, ascorbate-oxalate) is especially effective in promoting the dissolution of iron oxides (14). The adsorption of protons is also essential in reductive dissolution. In the present paper the systems magnetite / ascorbic acid and hematite/ascorbic acid are explored, in an attempt to unravel the mechanism of the reductive dissolution, and, in particular, the possible operation of synergistic effects of the reductant with strong complex forming agents such as oxalate and EDTA anions. The use of two oxides of different reactivities was useful in checking the validity of mechanistic postulates: shifts in kinetic control may occur in the complex mechanisms of dissolution under different experimental conditions, yielding apparently conflicting rate laws. The dissolving effect of ascorbic acid on iron oxides has already been used in some formulations of solvents for the chemical cleaning and decontamination of steel surfaces (22).

MATERIALS

AND

METHODS

Magnetite was prepared by reaction of K N O 3 with a slurry of iron(II) hydroxide in

the presence of hydrazine (23). It was characterized by its powder X-ray diffractogram, scanning electron microscopy, chemical analysis, and specific surface area. Most of the experiments were performed with a sample having a BET area of 3.3 _ 0.3 m 2 g-1 composed of polydispersed cubo-octahedral particles in the size range 0.1-1.0/~m. Magnetite dissolution experiments were performed in a magnetically stirred cylindrical beaker provided with a water jacket connected to a thermostat. Experiments were carried out at 64.8, 51.8, and 31.0°C under N2 atmosphere scrubbed through alkaline pyrogallol. Magnetite (25 mg) was suspended in doubly distilled water and dissolution was initiated by adding the appropriate a m o u n t of ascorbic acid (HzA). Periodical samples were taken with a syringe and poured into a large volume of cold water to quench further dissolution. This solution

75

was then filtered through a 0.45-#m Nuclepore membrane into a flask containing o-phenantroline and the absorbance was measured at 510 nm. In some experiments, dissolved iron was measured by AAS using a Varian AAR5 spectrophotometer. If needed, H2SO4 or N a O H was added to the reaction flask during the experiment to maintain constant pH. Redox potential was measured with a Keithley 610C electrometer and a Fluke 8050A digital multimeter. Partially reacted magnetite samples were examined occasionally by SEM using a Phillips apparatus. The hematite suspension was prepared by the Matijevic and Scheiner (24) method, as modified by Penners and Koopal (25). Electron microscopy showed smooth homodispersed spherical particles, with diameters of 5.0 X 10-2 #m. X-ray diffraction detected only hematite (o{-Fe203 ). The ion-exchange capacity measured by fluoride adsorption was 1.78 × 10-4 mol g-1 and the surface area from particle geometry was 17.5 m 2 g-1. The dissolution experiments were carried out under N2 in the dark at 25°C and at an ionic strength of 10 -2 mol dm -3 (NaNO3). The initial pH was adjusted to 3 with 0. I mol din-3 HNO3. Sampling was started after a conditioning time of one to several hours following the addition of the reagents. Samples (typically 12-15 cm 3) were taken periodically with a syringe to minimize contact with air, filtered through a 0.2# m pore-sized membrane filter under nitrogen, and the filtrate was collected in a glass bottle to which 0.10 cm 3 of 1.0 tool dm -3 HNO3 had been added to quench the autooxidation of Fe +2 and ascorbate. Iron was measured by the o-phenantroline method and ascorbate by UV absorption at 243 nm, after correction for the presence of iron. At pH 3 no measurable change in pH occurred during the whole experiment. The total amount of oxide dissolved during an experiment was typically 1.0 X 10 -5 to 3.0 × I0-5 mol dm -3. Adsorption experiments were carried out using the same procedure. A single sample was taken at t = 60 rain and analyzed for iron and ascorbate. Journal of Colloid and Interface Science, Vol. 138, No. 1, August 1990

76

DOS RESULTS

AND

SANTOS

AFONSO

ET AL.

1.0.

DISCUSSION

Acidity constants of HeA. The speciation of ascorbic acid in solution was required to interpret the kinetic experiments. The necessary acidity constants were derived from acid-base potentiometric titration experiments. The values thus obtained, shown in Table I, are in reasonable agreement with literature data for other temperatures ( 2 7 - 2 9 ) .

f

0.5

o

T = 64.8"C

0,0 50

Dissolution of Magnetite by Ascorbic Acid The shape off / t plots. Dissolution of magnetite could be followed up to high degrees of conversion because of the intrinsic reactivity of the solid, the high ascorbic acid concentrations, and higher temperatures used. Figure 1 shows the shape of the curves of dissolved iron t?action f ( f = (Wo - w)/wo, where w is the mass of magnetite and Wo is the initial mass) as a function of time for some selected experiments. The profiles are not monotonously deceleratory, and a contracting geometry kinetic law (e.g., 1 - (1 _f)1/3 = kt) is not totally adequate. However, affine transformation curves show that the basic topochemical characteristics of the reaction are unchanged when the solution variables and temperature were varied. Figure 2 shows the plot of f as a function of [ t~ tl/2 ], where tl/2 is the time for 50% dissolution; within experimental error, all the curves fall in a single trace, proving that tl/2 values m a y be used to derive rate laws and activation energies. The sigmoidal shape of the profiles is probably due to the influence of Fe(II) built up in solution; the effect is however not large (cf. below).

100

150 t (mini

200

250

FIG. 1. Dissolution fraction f [= (w0 - w)/wo] as a function of time, for magnetite in the presence of various concentrations of ascorbic acid: ©, 7.53 × 10-3, n, 4.55 × 10-2, V, 7.55 × 10-2 mol dm-3; pH 3.46; t = 64.8°C.

Demonstration of kinetic control by heterogeneous chemical reaction. The reaction rate was found to be proportional to the instantaneous surface area S in experiments carried out with various amounts of magnetite in the same volume of reaction: in all casesf/t curves overlap within experimental error. Such a dependence demonstrates a truly heterogeneous process (as opposed, for instance, to a control by chemical reactions of ascorbic acid in bulk solution). 1.0 o •

o°

o

oo el

t

o

¢,

•o ÷ °~o

0.5

oe

T = 64.8"C

{

TABLE I Acidity Constants of Ascorbic Acid K,X

T(°C)

1 = 1M

/

105

1 = 0.5 M

I = 0.1M

0 1~.0

31.0 51.8 64.8

8.4 11.5 13.7

8.5 9.6 11.7

7.6 8.9 10.3

Journal of Colloid and Interface Science, Vol. 138, No. 1, August 1990

21o

3'.0

4.0

" t / ' t 112

FIG. 2. D i s s o l u t i o n f r a c t i o n f as a f u n c t i o n o f ( t / t l / 2 ) f o r m a g n e t i t e in v a r i o u s e x p e r i m e n t s at 6 4 . 8 ° C .

DISSOLUTION OF IRON OXIDES 15

Experiments carried out at various stirring rates rule out diffusion control. The value of the apparent activation energy obtained from an Arrhenius plot in the temperature range 31-65°C (Ea = 68.0 kJ mo1-1) is also indicative of chemical control.

10. "Tc

=

E

Dependence of the rate on solution composition. The influence of ascorbic acid concentration and acidity on ti-/~2are shown in Figs. 3 and 4. The empirical relationship is R

77

e,l

Ill

=

where R is the instantaneous rate (assumed to be proportional to t i-/2), k is a proportionality constant, and [H2A]T is the analytical concentration of ascorbic acid (irrespective of protonation state). Log-log plots show that the best fitting value for the exponent on [HzA]T and on [ H +] are both indeed 0.50. The rate law [1] m a y be interpreted in two ways: (a) Ascorbic acid promotes dissolution irrespective of its protonation state, probably because of the operation of outer-sphere electron transfer, and the proton dependency represents a typical Freundlich-type adsorption of protons (with n = 0.5) to yield adequately reactive sites. (b) The electroactive species is the ascotbate anion H A - , the concentration of which varies with p H according to Eq. [ 2 ] (activity

0

! 1,0

0

[ 2.0

i 3.0

i 4.0

i 5.0

6.0

pH

FIG. 4. Dependence of t?/~zon p H for the dissolution of magnetiteat [HzA];r = 4.55 10 z mol dm -3, t = 64.8°C.

coefficients were ignored) and the proton influence on rate is the result of opposing effects on the protolysis equilibrium and on the density of reactive sites. The plot of log { R / [ H A - ] °5 } vs pH, shown in Fig. 5, demonstrates that [ 1] and [ 3 ] represent the experimental data equally well. [ H A ] = [HzA]T{I + K a l [ H + ] } -1 R = k[HA-]I/2[H+].

[2] [3]

2-.0 0

1,o. '<

o

To=B4.8~C o.o.

Co

~'.o

~:o

~io

51o

61o

-1.0-

7.'o

o.o

[H2A] . 102, mo[ - ctm"3

FIG. 3. Dependence of t~]2on total ascorbic acid concentration for the dissolution of magnetite at pH 3.46, t = 64.8oC.

8'c

;.o

Lo

5'.o 6.0 p~

FIG.

5. Dependence

on

pH

of

]og{tT)2/[HA-]l/2);

conditions as in Fig. 4. Journal of ColloM and lnteOCaceScience, Vol.

138, No. 1, August 1990

78

DOS SANTOS AFONSO ET AL.

Interpretation (b) favors a chemisorptive interaction with formation of ascorbate-iron (III) surface complexes that become reactive sites for dissolutions on adequate protonation of adjacent oxide groups. The Freundlich-type exponent on [H +] is larger than in case (a) because of a low coverage density of protons. The kinetic scheme [ 4 ] is further substantiated by the results with hematite (see below)

1.0-

,, 3.28-162rnol-dm3 A "/4.47 "10"2moi'drn3N2A / o / 3.55"102rno[drn3 A o/ 4.44"1d mol'dm3N2A

f

/ 0.5

5"10 mo[.dm I-t2A

/ / p u r e

ascorbate

HA-(KA)

~Fem_OH

.•

~_Fem__A 2- H+(K/)) ( ~ F e n I _ _ A 2 - ) H +

•

•

• puredehydroascorbate

FelIq (-- ( ~ F e n - - A = ) H +. [41 Ascorbate radical A ~ may either be disproportionate to dehydroascorbic acid (A) and HzA, or reduce further iron(III) centers. A itself is not unreactive toward magnetite. Figure 6 shows the influence of dehydroascorbic acid on the dissolution profiles; the effect of H2A and A are simply additive. A is less effective than H2A, and the contribution of A formed by reduction of HzA in experiments without added A is below 10%. From Fig. 6, the solution redox potential is seen to be irrelevant in determining the rate of reductive dissolution. Indeed, at least in the case of less strong reductants reacting with both iron (III) oxides and manganese(IV) oxides (1, 11), electron transfer is rate determining and the reverse process does not take place. The case of more reversible (metallic) redox couples has been discussed elsewhere ( 1 ). The possible involvement of dissolved Fe(II) in an autocatalytic reaction ( 1, 17, 30) was explored in additional experiments carried out with various concentrations of added Fe(II) and Fe(III). Although the values of t T)2 do not change appreciably in either case, the sigmoidal shape of the dissolution profiles is a strong indication of acceleration by Fe(II ) [ 2 ]. The lack of influence of added Fe(II) on tl/~2 can be attributed to a complex rate law (31 ). Because of the fast reduction of Fe(III) by ascorbic acid, ferric salt addition is cornJournal of Colloid and Interface Science, Vol. 138,No. 1, August1990

0.0 ~ 0

,

,

50

100

150

t, min FIG. 6. Dissolution profiles for magnetite in the presence ofdehydroascorbic acid (3.2 × 10 -2 tool d m -3, e ) , ascorbic acid (4.55 X 10 -2 tool d m -3, ©), and mixtures of both (at approximately the same concentrations, [] and []); pH 3.46, t = 64.8°C.

pletely equivalent to Fe(II) addition, and the Fe(III) / Fe(II) redox potential is irrelevant. From the temperature dependence of the rate of magnetite dissolution, an apparent activation energy of Ea = 68 kJ mo1-1 is calculated that embodies contributions from K~, Kv, KA, and kzT (see [2] and [4]).

The Interaction of Hematite with Ascorbic Acid When ascorbic acid is added to the hematite dispersion, in our experimental conditions, two processes take place in different time scales: (a) a fast decrease of the concentration ofascorbic acid from the solution, without appreciable dissolution of the metal oxide, and (b) a slower release of iron ions into the solution. The first stage is therefore the adsorption of ascorbic acid on the surface. Figure 7 shows the adsorption isotherm at pH 3 and

DISSOLUTION OF IRON OXIDES 25

79

F

Q5

% 20-

o . 4

b OZ

35 REDUCTiVE DISSOLUTION

°

"a, -6 E

15-

•

Q000

Q1 02 [H2A] 4 {10-6rr~i.dm"3 )

t °20

o

10E~

/o

,....---"-T-o

o

o TION,

0

0

o

5'o

40 [H2A]T

~'o

~o

~5o

III

1v

1'o zo

o

3o 4~o 50

TIME {HOURS)

(10-6 mot. din-3 )

FIG. 7. Adsorption isotherm ofascorbic acid on a-FezO3 at pH 3, t = 25°C. Insert: Double reciprocal plot.

25 °C. T h e degree o f c o v e r a g e r e a c h e d at the highest c o n c e n t r a t i o n is a b o u t 1.3 × 10 6 tool m -z, c o r r e s p o n d i n g to a surface area o f 130 A 2 p e r a d s o r b e d m o l e c u l e . T h e shape o f the a d s o r p t i o n i s o t h e r m c a n be expressed b y L a n g m u i r Eq. [5] w i t h Kf = 7.50 × 10 -4 a 7 ] mo1-1 d m 3 ( a l = [HA-]/[H2A]T = 0 . 0 7 8 4 at p H 3.0, I = 0.01 m o l d m -3) a n d Smax = 1.01 × 10 _8 m o l m - 2 : {~FemA - }

= KuSmax[HA-]/(1 + Kf[HA ]).

PROTONASSISTED DISSOLUTION /

[5]

T h e d a t a m a y b e d e s c r i b e d e q u a l l y well b y a F r e u n d l i c h i s o t h e r m w i t h n = 0.3 a n d K = 1.70 × 10 -5 m o l °7 d m °9 m -2. After several h o u r s o f c o n t a c t t i m e , i r o n is slowly released i n t o t h e s o l u t i o n . S o m e t y p i c a l k i n e t i c profiles are s h o w n i n Fig. 8, a n d Fig. 9 shows t h a t t h e o r d e r o n a s c o r b i c acid c o n c e n t r a t i o n is 0.3 w i t h k = 2.0 × 10 -6 m o l °7 d m °9 m -2 h -] (cf. Eq. [1]). T h e a l t e r n a t i v e L a n g m u i r - H i n s h e l w o o d e x p r e s s i o n m a y also be used. T h e i d e n t i t y o f t h e k i n e t i c r e a c t i o n o r d e r o n a s c o r b i c acid a n d t h e F r e u n d l i c h exp o n e n t for t h e a d s o r p t i o n p r e e q u i l i b r i u m d e m o n s t r a t e t h a t t h e rate is p r o p o r t i o n a l to

FIG. 8. Dissolution profile for oL-Fe20 3 in the presence of 0.0 to 5.0 X 10 -4 tool dm -~ ascorbic acid, 5 X 10 -5 tool dm-3 oxalic acid, and mixtures of both; II, proton assisted dissolution, [3, [ox] = 5.0 × 10 -5 mol dm -3 and no ascorbate; O, [H2A ] = 1.0 × 10-5 mol din-3; O, [H2A ] = 5.0 X 10 5 mol dm-3; O, [ H z A ] = 1.0 X 10 -4 tool dm 3; ~, [ H2A ] = 5.0 X 10 -4 tool dm-3; l~, [ H2A ] = 1.0 × 10 -4 mol dm -3, and [ox] = 5.0 × 10 5 mol dm-3; Cs = 0.613 g dm-3; I = 0.01 mol din-3; pH 3, t = 25°C; dark,

the c o n c e n t r a t i o n o f surface c o m p l e x e s ~ F e m A - (see Fig. 10). D i s s o l u t i o n is therefore m e d i a t e d b y a n e l e c t r o n t r a n s f e r w i t h i n a n i n n e r - s p h e r e a s c o r b a t o - F e ( I I I ) surface c o m p l e x i n a d e q u a t e l y p r o t o n a t e d sites. T h e q u e s t i o n o f w h e t h e r electron transfer or F e ( I I ) p h a s e t r a n s f e r c o n t r o l the rate c a n n o t be a n -

3 i o~ 'E 73 ,~" 2. E £

% ~o

, 0

i 1

,

t 2

.

[M2A]OT 3

i 3

,

i 4

,

i 5

i

i 6

,

(10-6 m o l . d r n - 3 ) 0 3

FIG. 9. Influence of [ HzA]T on the rate of dissolution of a-Fe203 at pH 3, t = 25°C. Journal of Colloid and Interface Science, Vol. 138, No. 1, August 1990

80

DOS SANTOS AFONSO ET AL. 3-

#~," 2, E "5 E t;-,

o

rate decreases again. In this system, ascorbic acid initiates an autocatalytic process in which FeY 2- (Y = E D T A tetraanion) is the actual dissolution agent. During the process, HzA further behaves as sacrificial reductant. The reaction scheme is

:H:3 C s = 0.613g.dm -3 = 0.01 rnol.dm - 3 dark

=25%

mm ~

J=

1.

Initiation: ~0P 0

' ,[r~VeIA -}

I~) (I0-7moi

'

20

~Fem--X

~-~A- F d I ( a q ) H'vyn-4 ~. FeIIy-2aq

. rn-2 )

FIG. 10. Influence of surface complexconcentration on the rate of dissolution of a-Fe203 at pH 3, t = 25°C.

[61 Acceleration: ~ - F e I n - - x + F e n y a 2 .--}

swered from our data. The long time elapsed between adsorption and dissolution may imply that the slow step is the phase transfer o f = F e H formed rapidly by internal electron transfer within the surface complex (note that electron transfer in homogeneous solution is fast (32, 33)). However, electron transfer may be more reversible in surface ~ F e l l I - - A 2- than in aqueous Fem-A 2- because -~Fe n centers can efficiently scavenge A = before these are oxidized by Fe hI. In fact, m a n y internal electron transfer reactions of F e m - L n- complexes become irreversible only when a scavenger traps the one-electron reduction product of L n- and prevents the back reaction. It seems likely that hematite and magnetite m a y represent two different cases, the breakage o f - O - F e n being slower, and therefore rate determining, in the former but not in the latter case.

~___FerI--X

+

[7]

FemY2q

dissolution

Feii y - 2

X represents any exposed ligand ( A - 2 0 H - I o r y4-) bound to an adequate surface site. When Fe +2 is added to E D T A / m a g n e t i t e or goethite suspensions, a m a x i m u m rate is obtained when the concentration of E D T A is high enough to complex all dissolved Fen(34); at higher concentrations, E D T A binds surface ~ F e m - - O H sites yielding ~ F e m Y-4 sites of reduced reactivity. The m a x i m u m rate depicted in Fig. 11 is very similar to that found previously in the ascorbic acid-free system: 5.7

3.0

Dissolution by Ascorbic Acid in the Presence of Chelating Carboxylic Acids

T = 31.0 *C p H : 3.46 = . - -

2.O

.

-

%

In the case of magnetite, at high H2A concentrations, there is a strong influence of increasing E D T A concentration on the rate of dissolution in 7.5 × 10 -3 tool d m -3 ascorbic acid at p H 3.46 and 31.0°C (Fig. 11). There is an obvious synergistic effect raising the rate when [ E D T A ] is increased up to 2.5 × 10 -3 mol dm-3; at higher E D T A concentrations the Journal of Colloid and Interface Science, Vol.138,No. 1,August1990

@

1.0

t

L

i

1.o [EDTA] "102 mol.drff 3

1

z.o

FIG. 11. Dependence on EDTA concentration of the rate of magnetite dissolution by ascorbic acid at [H2A]r = 7.5 × 10-3 tool dm -3, pH 3.46, and t = 31.0°C.

DISSOLUTION OF IRON OXIDES

X 10 -3 and 5.9 × 10 -3 s -l, respectively (34); furthermore, the rates obtained at very high EDTA concentration in both studies, corresponding to a surface fully covered by EDTA, are also similar (7.37 × l 0 -4 and 6.40 X 10 -4 s -~, respectively). In the case of hematite, at lower H2A concentrations, the rate of dissolution also increases upon addition of 5.0 N 10 _5 tool dm -3 o x a l a t e to 1.0 X 10 -4 t o o l dm -3 a s c o r b a t e sol u t i o n (Fig. 8 ) even though the surface density of ascorbate complexes measured in the adsorption experiments decreases from 1.26 to 0.41 X 10 -6 tool m -2. Khan and Martell (32, 33) showed that the oxidation ofascorbate by Fe +3 in solution is inhibited by chelating ligands and proposed that electron transfer within the inner-sphere iron(III)-ascorbate complex was the rate-determining step. If a similar effect were operative in the dissolution, the competition for surface sites between the carboxylate and ascorbate should also give rise to a decrease in dissolution rate. This evidence therefore also favors rate control in this case by phase transfer of Fe(II), the acceleration being due to the complexation of surface Fe(II) by oxalate. Again, the evidence is not conclusive, because phase transfer of ~ F e m - - o x and redox reactions (homogeneous or heterogeneous) between Fem and H A - may produce Fe n complexes capable of triggering a mechanism similar to that described above for magnetite. Additive transfer of oxalate and ascorbate complexes can be ruled out; the total rate is more than three times greater than the sum of the rates for the oxalate-promoted dissolution and the reductive dissolution, about 2.0 X 10 .8 mol m -2 h -~ for the conditions of Fig. 9. The effect of oxalate is synergistic rather than additive. ACKNOWL~EDGMENTS We thank Rudolph Giovanolli, University of Bern, for the X-ray diffraction analysis and electron microscopy. We also thank Barbara Sulzberger and Janet Hering, EAWAG, for their useful contribution to the discussion and interpretation of our results. M. A. Blesa and P. J.

81

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