Langmuir 1992,8, 1671-1675
1671
Reductive Dissolution of Iron(111) (Hydr)oxides by Hydrogen Sulfide Maria dos Santos Afonso*f+and Werner Stumm Swiss Federal Institute for Water Resources and Water Pollution Control (EA WAG), Swiss Federal Institute of Technology (ETH), Duebendorf, Switzerland Received November 11,1991.I n Final Form: March 6, 1992 The reductive dissolution of hematite by H2S was investigated under conditions of constant partial atm), pH = 3-7 and 25 "C, and under conditions where the solubility product pressure of HzS (lo4 to of iron sulfide is not exceeded. The reaction proceeds by an initial (relatively fast) formation of SFeSand =FeSH surface complexes; subsequently, electron transfer and detachment of Fe(I1) from the oxide surfaceoccur. The rate law for the dissolution can be interpreted by assuminga surface-controlledreaction depending on the surface concentration of the surface complexes =FeS- and EFeSH: Rt = k,(=FeS-) k,'(=FeSH) where (GFeS-) and {=FeSH) are surface concentrationsin mol.m-2. The experimentaldata are compatible with the constants k, = 30 h-l and k,' = 400 h-1 and surface complex formation constants for EFeS- (=FeOH + HzS(g) F? =FeS- + H+ + H20) and =FeSH (EFeOH + HzS(g) $ EFeSH + H2O) of log Ks = -2.70 and 2.82, respectively. Preliminary experiments with lepidocrocite and goethite give results on the reaction rates that show the same kind of dependence on solution variables as those with a-FezO3 and are compatible with the mechanism proposed here.
+
Introduction Sulfide can reductively dissolve iron(II1) (hydrloxides. This process is of importance in the corrosion of iron (destruction of passive films)and in the cycling of electrons in watersediment systems and in soil systems and is of interest in the storage and transfer of iron in biological cells. Iron(II1) on the surface of an oxide is reduced by an adsorbed reductant (reducing organic or inorganic ligand or reducing metal complex);iron(I1)is then released to the solution faster than iron(II1) because the bonds between reduced iron and 02-ions of the crystalline lattice are weakened (decrease of the Madelung energy). All these reactions have in common that they are surface controlled and the dissolution rate depends on the concentration of the dissolution-promoting species on the surface. The dissolution of iron(II1) oxides by different reductants has been extensively reviewed.'s6 Pyzik and Sommer1 have described reactions of sulfide with Fe(OH)3and FeOOH, proposing a model based on sorption of aqueous sulfide on the surface ferric iron and subsequent dissolution of surface Fe(I1). Rickard2p3proposed a two-stage mechanism (diffusion of H2S to the solid surface and reaction of H2S with dissolved Fe3+).Furthermore, Canfield and Bernel.4 studied the dissolution and pyritization of magnetite in anoxic marine sediments. Luther: on the basis of the analysis of the involved frontier orbitals, proposed, for low pH conditions, inner-sphere coordination of sulfide to Fe(II1) followed by electron transfer. In this paper we present the results of experiments carried out under conditions of constant partial pressure of H a a t a pH range of 3-7 in order to establish the reaction rate and in an effort to elucidate the mechanism of the reaction. Under our conditions the dissolution rates are t Present ad+
Depto de Qdmica Inorghica, Andtica y Q u h ica Fhica, Facultad de Ciencias Exactas y Naturales, Ciudad Universitaria Pabell611 11, 1428 Buenos Aires, Argentina. (1) Pyzik, A. J.; Sommer, S. E. Geochim. Cosmochim. Acta 1981,45, 687. (2) Rickard, D. T. Am. J. Sci. 1974,274,941. (3) Rickard, D. T. Am. J. Sci. 1975,275, 636. (4) Canfield, D. E.;Berner, R. A. Geochim. Cosmochim. Acta 1987,51, 645.
(5) Luther, G. W., 111, In Aquatic Chemical Kinetics; S t u " , W., Ed.; Wiley-Interscience: New York, 1990.
surface controlled and depend on the surface concentration of chemisorbed HS- and S2-. Background We assume a mechanism schematically given by the following reaction sequence: ki
EFeOH + HS- s =FeS-
+ H20
(1)
k-i
ket
&'eS- e E F eI1s' k-et
ki
=Fe"S'
+ H 2 0 s =Fe"OH2+ + So-
-
H+
=Fe"OH2+
(2)
k-2
new surface site + Fe2+,
ka
(3)
A similar reaction sequence can be given for the interaction of cr-FezO3 with H2S (see eq 9). Sulfide can bind to and reduce Fe(II1) through an innersphere coordination of sulfide to the Fe(II1) followed by d p Y )to ?r(dy,) electron transfers that is expected to be faster than an outer-sphere process alone. The So-formed in reaction 2 reacts with Fez03 in subsequent fast steps to form Sod2-
S'-
+ 7rFeOH
--
feat
SO:-
+ 7Fe2+aq
(4)
The rate law for Fe2+aqformation a t a given pH is (5)
Assuming, fast formation of =FelIIS- the steady-state approximation d(=FeIIOHz+)/dt = 0 is used to obtain the expression
Braces in eqs 5 and 6 refer to surface concentration (mol.m-2). Equation 6 gives the dissolution rate, R, of hematite by hydrogen sulfide in terms of a first-order rate
0743-7463/92/2408-1671$03.00/0 0 1992 American Chemical Society
1672 Langmuir, Vol. 8, No. 6, 1992
dos Santos Afonso and Stumm
law, dependent only on the surface concentration of sulfide d[Fe2+ I aq = k,{=FeS-} (7) A dt provided that the intermediate product [S'l is at steady state. A is the surface area concentration in m2.L-'. The surface species =FeS-must be assumed to undergo protonation, i.e. an acid-base equilibrium of the form R=
prevails. From surface complex formation equilibria it is known (Sigg and Stumm,G Stumm et al.,' and Dzombak and Morel8)that this pKaSvalue is several units lower than the pK, of dissolved HS-, because of the effect of the neighboring Fe(II1) surface center. Recently, R6nngren et aL9 determined the following acidity constant (25 "C) for =ZnSH+
~i GZnS
+ H+
Table I. Matrix for All the Equilibria Needed in the System. species EFeOH exp(-Fk/RT) HzS(gas) H+ logK EFeOH 1 0 0 0 0.00 =FeOHz+ 1 1 0 1 7.25 =FeO1 -1 0 -1 -9.75 EFeS1 -1 1 -1 -2.70 =FeSH 1 0 1 0 2.82 1 0 -0.98 HzS(aq) 0 0 HS0 0 1 -1 -8.00 S20 0 1 -2 -21.90' HzS(ga) 0 0 1 0 0.00 H+ 0 0 0 1 0.00 The equilibrium expression can be read from this matrix; e.g. for =FeOH2+, (=FeOH2+) = (=FeOH)[H+110726exp(-m/RT) or for SFeSH, (EFeSH) = (=FeOH)(pHZS)102.82. This equilibrium is characterized by an uncertain constant. The modelingof the surface concentrations and the dissolution rate is independent of this equilibrium. 10 -6
pK,S = 6.91
10 -7
In this study we achieved a good fit with the acidity constant of EFeSH (pKaS= 5.5). It is reasonable to assume that the pKas of =FeSH is somewhat lower than that of =ZnSH+. Thus, at lower pH values, =FeSH is formed; its equilibrium is given by
...
5
E
.-
10 -8
0
C
E
c
e
10 -9
Q)
0
E
0
k4
EFeOH + H2S(aq)e e F e S H + H 2 0
(9)
0
10 -10
h-4 10 -11
A similar reaction scheme as that given for the reaction involving S F e S - can be given for the reaction involving =FeSH
0
2
R, = k,{=FeS-j
+ k,'{=FeSH)
+ H+ GFeOH F? ZFeO- + H+
+ R') (11)
where {=FeS-) and {EFeSH)can be expressed in terms of surface complex formation constants Kls = kl/k-l (cf. eq 1) and Kz8 = kdk-4 (cf. eq 9). In order to calculate the equilibrium composition of the surface species, one needs to consider that the surfaces involved are charged; electrostatic corrections for Coulombic interaction, Kapp = Kin, exp(-F+IRT), using the diffuse double layer model (Stumm et allo) need to be made. The mass balance for hematite surface sites is given by [=FeOHl,,
= [=FeOH2+l
+ [=FeOHl+
[=FeO-l+ [=FeSHl + 1 ~ F e S - l(12) where the acidity constants of =FeOHZ+ are defined as (6) Sigg, L.; Stumm, W. Colloids Surf. 1980/1981, 2 , 101. (7) Stumm, W.; Kummert, R.; Sigg, L. Croat. Chem. Acta 1980, 53, 291.
(8)Dzombak, D. A.; Morel, F. M. M. In Surface Complexation Modeling-Hydrous Ferric Oxide; Wiley-Interscience: New York, 1990. (9) RBnngren, L.; SjBberg, S.; Sun, Z.; Forsling, W.; Schindler, P. W. J. Colloid Interface Sci. 1991, 145, 396. (10) Stumm, W.; Huang, C. P.; Jenkins, S. R. Croat. Chem. Acta 1970, 42, 223.
8
10
12
PH Figure 1. Surface speciation as a function of pH for lo4 atm of partial pressure of HzS and total concentration of [=Fe0Hlbt = 1.63 X mo1.L-l as calculated using MICROQL and the information given in Table I.
EFeOH; The overall dissolution rate is (Rt = R
6
4
~i=FeOH
Kels
(13)
K,; (14) Table I gives the matrix for all the equilibria needed in the system and the equilibrium constants applied. Figure 1shows the calculated surface speciation as a function of pH for a partial pressure of H2S of atm and [=FeOHlbt = 1.63 X lo-' mo1.L-l using a computer calculation based on the MICROQL program by Westall.11 Experimental Section Dissolution experiments were performed in a magnetically stirred cylindrical beaker provided with a water jacket. The temperature was held constant to 25 0.5 OC with a constant temperature circulation bath. The hematite suspension was prepared by the method of Matijevic and Scheiner,12as modified by Penners and K00pal.l~The hematite was analyzed by X-ray and transmission electron microscopy. Electron microscopy showed homodispersed spherical particles with diameters of 0.3-0.4 pm. All the kinetics experiments were made in an anoxic environment under Nz atmosphere scrubbed through a V(I1) solution (Jones' reagent) and ionic strength M of NaC104. The dissolution experiment was initiated by adding a volume of hematite suspension to a solution exposed to hydrogen sulfide.
*
(11) Westall J. MICROQL,A Chemical EquilibriumPrograminBasic. Internal Report, EAWAG, Duebendorf. (12) Matijevic,E.; Scheiner,P. J. Colloid Interface Sci. 1978,63,509. (13) Penners, N. H. G.; Koopal, L. K. Colloids Surf. 1986, 19, 337.
Langmuir, Vol. 8, No. 6, 1992 1673
Reductive Dissolution of Iron(III) Oxides
60
40
20
0
0
time (h)
Figure 2. Correlation between the H+ consumption and Fe2+ formed in a typical experiment. Table 11. Possible Dissolution of Fez03 Leading to Different 9-Oxidation Products stoichiometry 4Fez03 + 2HzS + 15H+ a S Z O ~+~12Hz0 + 8Fe2+, 4Fez03 + HzS + 14H+ e SO,z- + 8Hz0 + 8Fe2+., 3Fez03 + HzS + 10H+ a s03'- + 6HzO + 6FeZ+, 8FezO3 + 8HzS + 39H+ a Ss0 + 24&0 + 16FeZ+, ( x - 1)FezOs + xH&3 + (4x - 6)H+ a Sz2- + 3(x - 1)HzO + 2(x - l)FeZ+.,
-A[H+l/ A[Fe2+l 1.87 1.75 1.66 2.44 1.0 < y < 1.8 if 2 < x < 6
Gaseous hydrogen sulfide was supplied to the reaction mixture by continuous bubbling of special H2S/N2 mixtures. The pH was kept constant by adding small aliquots of 5 X or M of HCIOl using a Dosimat pH-stat system. Light was excluded from all experiments. The progress of the dissolution was followed by measuring the concentration of Fe(I1) as a function of time and H+ ion consumption (HC104 volume added by pH-stat). Figure 2 gives an example of the relationships between H+ consumed and Fe(11) produced. The stoichiometries for the dissolution reactions leading to different sulfur-oxidation products are given in Table 11. The values of the dissociation constants14of H2S at 25 OC and Z = 0 in water are pK1 = 7.02 and pK2 = 13.9, but more recent papers16J6report pKd = 19. The arguments presented are not dependent on t h k constant because S2-wasnot used as a variable. The experiments were carried out at different pH values (4 to 7) and partial pressures of H2S (lo4 to lo-* atm). During the experiments, the partial pressure of H2S, pH, and temperature were held constant and under conditions of undersaturation with regard to the solubility product15of amorphous FeS (FeS(,, + H+ F? Fe2+ + HS-; log *K, = 2.95). No formation of solid iron sulfide was observed. The specific surface area of hematite was calculated on the basis of the geometry of the homodisperse a-FezO3. Earlier investigations have shown a good agreement between the calculated area and the BET-measured area. The hematite used was characterized as follows: specific surface area = 4.5 m2.g-l; site density = 1nm-2; a-Fe203concentration range between0.017 and 0.036 gL-l ([=Fe0Hlbt = 1.27 X loe7to 2.69 X 10-7mol.L-1). Surface acidity constants, as determined for a-Fe203, equally synthesized as our own preparations, by Liang and Morgan1' (obtained for diffuse double layer model), pK.1, = 7.25 and pKds = 9.75, respectively, were used. Periodic samples were taken with a syringe and then filtered through a O.Z-pm Nuclepore membrane. All the samples was (14) Stumm, W.; Morgan, J. J. In Aquatic Chemistry, 2nd ed.; Wiley-Interscience: New York, 1981. (15) Davison, W. Aquat. Sci. 1991,53, 309. (16) Schoonen, M. A. A.; Barnes, H. L.; Geochim. Cosmochim. Acta 1988, 52,649. (17) Liang; Morgan, J. J. Aquat. Sci. 1990, 51, 32.
Figure 3. Progress of the dissolution reactions at different experimental conditions; H, pH = 7.0, pH2S = 10-4 atm; 0,pH atm; 0 , pH = 5.6, pH2S = atm; A, pH = 4.0, pH2S = = 4.0, pH2S = atm; A, pH 7 5.6, pH2S = lo4 atm; O , pH = 4.0, proton-promoted dissolution only (in absence of H2S). stored under an oxygen-free N2 atmosphere until analyzed. Filtration of the homodisperse spherical a-FesO3 particles with 0.2 pM membrane filters was quantitative as checked independently. Dissolved Fe(I1) was usually measured by atomic absorption spectroscopy. In some experiments dissolved iron(11)was measured using the o-phenanthroline method (absorption at 510 nm). No difference was found between both methods. Dissolved iron(I1) could be determined with a precision better than i 5 % . Errors in the determination of Fe(I1) increase with an increase in pH and decrease of Fe(I1) concentrations because Fe(I1) is adsorbed on particles, glass walls, and filtration apparatus. At pH 5 and versus [HS-l-l[H+l-l for pH 7 where FeS(s) was formed. Preliminary results on the reduction of iron(II1) (hydr)oxides other than hematite have been carried out. The results (Figure 6) obtained, compatible with the mechanism proposed, show that the reduction rate (eq 11) decreases with increasing the free energy of the following reductionreactions: FeOOH + 3H++ e- e Fe2+, + 2H20; Fez03 + 6H+ + 2e- 2Fe2+as+ 3H20; Fe3Or + 8H+ 2e3Fe2+8q+ 4Hz0. The free energy values for these reactions were taken from Heusler and Lorenz22and Fischer .23 The reductive dissolution of magnetite by hydrogen sulfide has a different stoichiometry than that of the dissolution of iron(II1) (hydr)oxides, i.e., more Fe(I1) is released per Fe(II1) reduced. Preliminary experiments reveal that the dissolution of Fe304 also proceeds through the formation of S(-11) surface complexes. The reaction mechanism proposed, in the broadest term, is in agreement with the ideas presented earlier by Kumai,24R i ~ k a r d , ~and s Pyzik and S0mmer.l The latter authors' empirical rate expression for the dissolution and sulfidation of goethite is in agreement with our analysis. The main difference is whether the electron transfer occurs at the iron oxide surface (inner-sphere complexes) or far away from the surface. It is well stablished (e.g. in biological systemsz5)that S(-11) has a strong affinity to the Fe(II1). Thus it is justified to postulate a mechanism of an initial formation of inner-sphere surface complexes of =FeS- and =FeSH that undergo subsequent electron transfer. Similar reaction sequences have been proposed
*
+
(22) Heusler, K. E.; Lorenz, W. J. In Standard Potentials in Aqueous Sohtion; Bard, A. J., Parsons, R., Jordan, J., Eds.; Marcel Dekker, Inc.: New York, 1985. (23) Fischer, W. R. In Iron in Soils and Clay Minerals; Stucki, J. W., Goodwan, B. A., Schwertmann, U. Eds.; Dordrecht, 1988. (24) Kumai T. Nippon Kagaku Zasshi 1960,81, 1690. (25) Frausto da Silva, J. J. R.; Williams, R. J. P. In The Biological Chemistry of the Elements: The Inorganic Chemistry of Life; Clarendon Press: Oxford, 1991.
Reductive Dissolution of Iron(III) Oxides
in other research on reductive dissolution of iron(II1) (hydr)oxides.18*21The rate law given in eq 11 permits estimation of the dissolution rates as a function of pH and solution variables.
Acknowledgment. M.d.S.A. has been a visiting scholar from Buenos Aires University; she is grateful to Buenos Aires University for granting her a leave of absence. We thank Rudolph Giovanoli, University of Bern, for the X-
Langmuir, Vol. 8, No. 6,1992 1675
ray diffraction and electron microscopy analyses and M. A. Blesa, University of Buenos Aires, for useful contributions to the discussion and interpretation of our results. We are indebted to George W. Luther, 111, and two anonymous reviewers for their constructive criticisms. Registry No. Fe203, 1309-37-1;FeaOl, 1317-61-9;H& 778306-4; iron hydroxide oxide, 20344-49-4.
