consisting of hydroxide sheets, [Fe4Fezn III(oH)12 ]2+, charged ... iron(II) hydroxide is initially precipitated from iron(II) sulphate and ... dissolved iron(II) sulphate is left in solution after ...... Cuttler A.H., Man V., Cranshaw T.E. & Longworth G.
Clay Minerals (1999) 34, 499-510
Chemical composition and Gibbs standard free energy of formation of Fe(II)-Fe(III) hydroxysulphate green rust and Fe(II) hydroxide PH. R E F A I T ,
C. B O N , L. S I M O N , G . B O U R R I I ~ * ' * * , F. T R O L A R D * J. B E S S I I ~ R E AND J . - M . R . G ] ~ N I N 1
Laboratoire de Chimie Physique pour l'Environnement, UMR 7564 CNRS-Universitk H. Poincard, Equipe sur la Rdactivit~ des Espbces du Fer and Ddpartement de Science des Mat&iaux, ESSTIN, 405, rue de Vandoeuvre, F 54600 Villers-les-Nancy, *INRA - UR de Science du Sol et de Bioclimatologie, 65 rue de Saint Brieuc, F 35042 Rennes Cedex, and **UPR 4661 CNRS-Univ. Rennes 1, Gkoscienees Rennes, Campus de Beaulieu, F 35042 Rennes Cedex, France (Received 20 January 1997, revised 15 October 1998)
A B S T RA C T: The redox potential and pH of aerated suspensions of iron(li) hydroxide in sulphatecontaining aqueous solutions are measured during the oxidation process. Plateaux corresponding to the equilibrium conditions between Fe(OH)2(s) and Fe(II)-Fe(III) hydroxysulphate GR2(SO2-)(s) on the one hand, and between GR2(SO]-)(s~ and FeOOH(s) on the other hand, are displayed. Potentiometry, voltammetry, pH-metry and M6ssbauer spectroscopy are applied to follow all reactions. The thermodynamic meaning of the measured potential of the first plateau which corresponds to the GR2(SO42 )(s)/Fe(OH)2(~)equilibrium is demonstrated. The chemical composition of GR2(SO42 )(s) is found to be Fe~Fe~II(oH)12SOg.nH20 all along the oxidation process implying that this compound must be considered as a pure phase with a well-defined composition. The Gibbs standard free energy of formation or ehemical potential g~ 2 )(s)] in 'anhydrous form' (n = 0) is determined at -3790_+10 kJ tool 1. A consistent value of g~ at -490___ 1 kJ mol I is obtained.
The Fe(II)-Fe(III) hydroxysulphate, with formula [FeIIFem(OH)lz]z+.[SO4.nH20]2 (Hansen et al., 1994; G6nin et al., 1996) is more commonly k n o w n as green rust two and d e s i g n a t e d GR2(SO ] )(s). As with other green rust (GR) compounds, it is a layered double hydroxysalt, consisting of hydroxide sheets, [Fe4Fezn III(oH)12 ]2+, charged positively due to the presence of ferric cations, which alternate regularly with negatively charged interlayers, [SO4.nH20] 2-, composed of
1 Corresponding author
sulphate anions and water molecules. The GRs have been observed as a corrosion product of steel (G6nin et al., 1991) and proposed in ochre sludge (Bender Koch & Morup, 1991). Their occurrence in anaerobic soils and sediments was suggested by several workers (e.g. Ponnamperuma et al., 1967), but it was only very recently that the definite presence of a GR compound as a mineral in hydromorphic soils was demonstrated (Trolard et al., 1996, 1997), even though its precise nature still remains uncertain, i.e. in relation to the type of anions involved (G6nin et al., 1998b). Of special interest is the reduction of nitrite and nitrate by green rust compounds, i.e. the possibility that GRs
9 1999 The Mineralogical Society
500
Ph. Refait et al.
could influence the denitrification of soils and sediments. Hansen et al. (1994, 1996) focused on the role that GR2(SOeZ-)(s) could play; they showed that it should be able to reduce nitrite and nitrate since their estimation of the Gibbs standard free energy (free enthalpy) of formation or chemical potential translated in terms of redox potential infers the fact that the reduction of nitrite is thermodynamically possible. Some other works were also devoted to the chemical and electrochemical properties of GR2(SO2-)(~) and various computations of the Gibbs standard free energy of formation have been proposed (Detournay et al., 1975; Olowe & G~nin, 1989; Refait & G~nin, 1994) but the chemical formula considered for GR2(SOZ-)(s) was erroneously referred to as [Fe~lFe~n(OH)1412+.[SO4-nH2O12 ; thus, these results cannot be used. In contrast, the composition [Fe~IFem(OH)12]2+.[SO4.nH20] 2- was recently ascertained and the standard free energy of formation re-evaluated by G~nin et al. (1996) at -3770_+4 kJmo1-1 while Hansen et al. (1994) obtained -3669_+4 kJ tool 1, taking the number n of water molecules incorporated in the interlayers to be equal to zero for easier comparison (add -237.18 kJ mol 1 per water molecule). The method we used, originally devised by Detournay et al. (1975) and also applied to other GRs (Refait & G6nin, 1993; Drissi et al., 1995), consists of an electrochemical survey of the oxidation of an aerated Fe(OH)2(s) suspension. If iron(II) hydroxide is initially precipitated from iron(II) sulphate and sodium hydroxide, GR2( SO] )(s) is obtained at the end of a first oxidation stage, provided that a sufficient excess of dissolved iron(II) sulphate is left in solution after precipitation of Fe(OH)2(s) (Olowe & G~nin, 1989, 1991). The pH and redox potential of the solution do not vary during the process and their values were assumed to verify equilibrium conditions between Fe(OH)2(~) and GR2(SO42 )(s~, allowing the computation of the standard equilibrium potential of the reaction and of the standard chemical potential I.t~ The measured redox potential was assumed to be close to that of the equilibrium, but the difference between them, which can g o v e r n the a c c u r a c y of the g~ 2 )(~)] value, was not established. Since some discrepancy exists between the value we found and that given by Hansen et al. (1994), it appears necessary to validate the procedure used in our own method. This determination implies a
thorough study of the electrochemical behaviour of GR2(SO ] )(s~, realized here by voltammetry. It allows us to propose a reliable and definite estimation of the standard chemical potential of the Fe(II)-Fe(III) hydroxysulphate. Moreover, the characterization by M6ssbauer spectroscopy of the solid phases which form during the oxidation process will allow us to follow the chemical composition of GR2(SOZ-)(s) with time, a key point in discovering whether true equilibrium is reached. METHODS The Fe(OH)2(s) suspensions were obtained by mixing 100 ml of an aqueous solution of melanterite FeSO4.7H20 (Normapur| provided by PROLABO | and containing a maximum impurity content of 1%, in particular GR
0.02
0.01
Eh (V)
-0.6__ 0
'
~J__.__
I
,
-I
-0.4 -0.3
-0.01
9
(a) GR --+ Fe(II)
-0.02 F~G. 3. Voltammetry of the Fe(OH)2(~)/GR2(SO42 )(s~ system9 (a) Current-potential curves obtained by cyclic voltammetry with an Fe(OH)2 and GR2(SO]-)(~) suspension (Eh - --0.5 V and pH - 8.0) at a Pt electrode at a sweep rate of 1 mV s 1. (b) Currentpotential curves of de-aerated 0.1 M Na2SO4 solutions at pH - 8.0.
thermodynamically possible. Nevertheless, such a process, if it takes place, would be very slow since we never observed in any of our previous studies (e.g. Refait & G6nin, 1996) any change with ageing in the TMS spectrum of Fe(OH)2(s> after as long as 72 h. It is clear, therefore, that the Fe(OH)2(s)/ GR2(SO ] )(s~ electrode is reversible and the measured redox potentials have a thermodynamic meaning9 Similar experiments were performed in plateau B. However, the overall features are less pronounced and the thermodynamic meaning of the measured potential at the corresponding equilibrium cannot be guaranteed9 Transmission M d s s b a u e r spectroscopy M6ssbauer spectra measured at 15 K are displayed in Fig. 4 and corresponding hyperfine parameters are gathered in Table 1. Seven examples from samples (a)-(g) which are reported, relate to
503
various reaction times which are indicated in the Eh and pH vs. time curves of Fig. 1. For experimental facility, spectrum (a) which would correspond to the very beginning of the oxidation process is in fact that of a Fe(OH)2(s) standard sample precipitated in a basic medium (pH ~ 11.5) where GRs do not form (cf. Olowe et al., 1991). Besides the analyses at the beginning, at the inflection point Tg and at the end of the overall reaction, i.e. spectra (a), (d) and (g), two product analyses were performed at each reaction stage, one at the beginning and one at the end. Therefore, spectra (a) and (d) are those of Fe(OH)2(s) and GR2(SO42-)(~), respectively, whereas spectrum (g) is that of a mixture of 7-FeOOH(s) and ~-FeOOH(s). In plateau A, i.e. spectra (b) and (c), both Fe(OH)2(s) and GR2(SOZ-)(s) are always present in different amounts. In contrast, plateau B is only found at the beginning of the second reaction stage and corresponds to spectrum (e) where only GR2(SO24-)(~) and 7-FeOOH(o are present. After plateau B, ~-FeOOH(s) is also present in spectrum (f) along with GR2(SO] )(s) and 7-FeOOH(s). in the spectra, GR2(SO42-)(s) is still paramagnetic and always displays two different quadrupole d o u b l e t s D 1 and D2 attributed to Fe(II) and Fe(III), respectively. At 15 K, the respective quadrupole effect values AEQ are -2.8 and 0.4 mm s 1 and isomer shifts 6 ~1.3 and 0.5 m m s 1 (Table 1). The abundance ratio D1/D2 varies little, ranging from 1.87 to 2.11, effectively equal to 2. This proves that the compound corresponds to a well-defined chemical composition and that departures from this stoichiometry do not occur in the experimental conditions considered here. 7-FeOOH(s) is antiferromagnetic (Johnson, 1969) and corresponds to sextet Sv with hyperfine field of - 4 5 0 k O e and AEQ of 0 at 1 5 K (G6nin et al., 1998a); some additional sextet S' r may be used to fit correctly the asymmetry of the spectrum, due to the poor crystallinity of lepidocrocite (Murad & Schwertmann, 1984)9 ~-FeOOH(s) is antiferromagnetic (D6zsi et al., 1967) corresponding to sextet S~ with a field o f - 5 0 0 kOe and AEQ of ~ - 0 . 2 5 mm s -1 (Van der Woude & Dekker, 1966; Forsyth et al., 1968; Murad, 1982). Finally, Fe(OH)2(s) is also antifen-omagnetic at 15 K under the N6el temperature of 34 K (Miyamoto et al., 1967) and gives rise to a magnetically split spectrum consisting of eight peaks, due to a mixture of quantum states with a hyperfine field, -200 kOe at 4 K (Miyamoto et al., 1967; G6nin et
504
Ph. Refait et al.
lOO "" o c t~ :~
100
.
98
,=_
96
tO I--
96
94
d) 9
E t~ t-
p
E
1 tt~
m
98
-5
~
.
-4
I
,
i
-3 -2
,
E
-1
,
i
0
,
i
,
1
i
2
,
i
3
,
i
4
,
i
5
94
F---
9
-1(
i
,
i
,
i
,
i
,
i
-8-6-4-2
,
i
0
,
i
2
,
4
i
,
6
[
Sy
,
8 10
V (mm/s)
98 o c
E
96
!:
94
(e)
9
-10-8-6-4-2
0
2
4
6
8
10
V (mm/s) 100 (11 o e-
E tO F--
Z
'~i!:...~D:"
99
oo_ ! FH~;
100
99
;
(b)
98
i
I
,
I
,
I
9
I
-10-8-6-4-2
o~._4 9 8 .
,
I~
0
,
I
2
,
I
4
,
I
6
,
I
8
,
I
10
o c t~
E t/) c E I--
FH
98 )
97
-10-8-6-4-2
0
2
4
6
8
10
100
o 96
E c
I--
o c
94
92 ~ -10-8-6-4-2
,, 0
(c) 2
V (mm/s)
4
6
8
10
99
E taq c
I---
) -10-8
-6 -4 -2
0
2
4
6
8
10
V (mm/s) FIG. 4. Transmission M6ssbauer spectra measured at 15 K of the solid phases sampled at various times of the oxidation of ferrous hydroxide. The times (a)-(g) of sampling are indicated on the Eh vs. time curve of Fig. 1. (a) Ferrous hydroxide prepared in basic medium; (b) and (c) in Plateau A during the first stage; (d) at Tg; (e) in Plateau B; (f) during the second stage; and (g) at Tf.
Fe(II)-Fe(iiI) hydroxysulphate green rust and Fe(II) hydroxide
505
TABLE 1. M6ssbauer hyperfine parameters at 15 K of the precipitates (a)-(g) sampled during the reaction of oxidation of Fe(OH)2(~) in sulphate containing aqueous medium. Sample (a)
FH
1.4
Ol D2
. -
Sample (b)
Sample (c)
AEQ
H
RA
6
AEQ
H
RA
~i
AEQ
3.1
175
100
1.4
3.1
175
74.6
1.4
-3.1
-
1.33 0.50
2.78 0.39
-
16.9 8.5 1.99
1.34 0.53
.
.
.
-
D~/D2
2.82 0.41
H
RA
175
11.1
-
60.3 28.6 2.11
Sample (d) GR2(SOZ-)(s) obtained at Tg AEQ RA D1
1.30 0.47
1)2 D1/D2 Sample (e) AEQ H
S~ S' v
0.51 .
S~
.
D~
D2
1.34 0.51
0 . .
440 . .
2.87 0.43
RA
6
28.2 .
.
. --
Da/O2
2.85 0.45
47.7 24.1 1.98
65.2 34.8 1.87
Sample (f) AEQ H
RA
~
Sample (g) AEQ H
RA
0.49 .
0 .
445
43.5
0.50 0.48
0 0
454 420
47 26
0.49
-0.25
505
19.6
0.49
-0.28
504
27
1.34 0.49
2.89 0.40
--
24.7 12.2 2.02
.
.
. --
. --
= isomer shift with respect to metallic et-iron at room temperature in mm s--l; AEQ = quadmpole splitting in mm s-l; H = hyperfine field in kOe; RA = relative area in %. FH = Fe(OH)2(s~; D1 and D2 = GR2(SO] )(s); S~ and S'~ = 7-FeOOH(s); S~ = et-FeOOH(s).
al., 1986) and 166 kOe at 20 K (Refait et al., 1998), which lie perpendicular to the axis of symmetry of the electric field gradient, and the quadrupole effect AEQ, which is - - 3 m m s 1 (Miyamoto et al., 1967; G6nin et al., 1986), whereas the isomer shift is 1.4 m m s -1.
DISCUSSION
Chemical f o r m u l a o f GR2(S02-)C~) Various authors concluded previously that the Fe(II)/Fe(III) ratio of GR2(SO]-)(s) was close or exactly equal to 2 (Detournay et al., 1975; Cuttler et al., 1990; Schwertmann & Fechter, 1994; G6nin et al., 1996). The TMS analyses confirm that this is the case but also demonstrate that this ratio is kept
constant all along the formation and oxidation process of the GR, from the beginning of the oxidation of Fe(OH)2(s~ to the end of the formation of ferric oxyhydroxides. Therefore, only pure phases are present, and no solid-solution between Fe(OH)2(s~ a n d GR2(SO42-)(s~ or b e t w e e n GR2(SO]-)(s) and FeOOH(s) is observed. Similar conclusions were drawn recently for Fe(II)-Fe(III) hydroxyoxalate GRl(C2042-)(s) (Refait et al., 1998). GR2(SO]-)(~) is generally assumed to be a p y r o a u r i t e - s j 6 g r e n i t e - l i k e hydroxide ( A l l m a n n , 1970; Taylor, 1973; Hansen et al., 1994) and consists of [Fe~IFe~n(OH)12] 2+ positively charged hydroxide sheets due to the Fe(III) cations, alternating with negatively charged interlayers composed of sulphate ions and water molecules. To m a i n t a i n the whole electroneutrality, the
Ph. Refait et al.
506
TABLE 2. Gibbs standard free energy of formation AG~ or go in kJ mo1-1 and thermo-dynamic constants used for calculations and results of this study. Species
AG~'f or g~
Fe(OH)2(s) FeI~FelII(OH)~2S04(s~ 7-FeOOH(s) Fe2+
-277.4
FeSO4aq
-823.49
H20 SO 2-
-237.18 -744.56
Complexes
log K
Fe 2+ + H20 = FeOH + + H + Fe z+ + SO2 = FeSO4aq Na + + SO,] = NaSO4-
o f the
This work This work Computed from Lindsay (1979) Yang, 1982 Bard et al., 1985; Kelsall & Williams, 1991 Wagman et al., 1982; Bard et aL, 1985; Kelsall & Williams, 1991 Bard et al., 1985; Kelsall & Williams, 1991 Bard et al., 1985; Kelsall & Williams, 1991 Bard et at., 1985; Kelsall & Williams, 1991
-490 -3790 -475.5 -91.5
FeOH +
composition
References
interlayers
References
-8.98 2.2 0.7
should
Computed from the AG~ values Id Jenkins & Monk (1950)
be
[SO4.rIH20] 2-, the presence of two Fe(III) ions being counterbalanced by one sulphate ion, leading to the overall formula Fe~tFe~H(OH)12SO4-nH20 already proposed (Hansen et al., 1994, 1996; G~nin et al., 1996). The results given by the measurement of Fe and S concentrations can be used to verify this hypothesis. With an initial concentration of Fe(OH)z(s) equal to 0.1 mol 1-1, the amount of Fe in the GR is thus equal to 0.1 + A[Fe], that is 0.1165 mol 1 1. The amount of sulphate in the GR is A[S], that is 0.025 mol 1-1. This gives an Fe/SO4 ratio of 0.1165/0.025 4.66. Since the Fe(II)/Fe(III) ratio is found to be 2 by TMS, the chemical formula given by c[Fe] and c[S] measurements would be: Fe~Fe~u(OH)11.48(SO4)l.26.nH20. The departure from the proposed formula is more likely to be due to slight experimental errors. Actually, such a composition would imply a release of OH ions during the formation of GR2(SO42-)(s~ from Fe(OH)2(s~, since the chemical balance would be: 5 Fe(OH)z(s)+Fe 2++ 1.26SO~ + 1 / 2 0 2 + H20 ~ Fe4Fe2 11 llI(OH)11.48(804)1.26 nt- 0.52 O H - (3) Since c[Fe(OH)2(~)] is equal to 0.1 tool 1 1, the release of 0.52 OH would correspond to ~0.01 mol 1-a of OH ions, which would have increased the
pH to -12. This is not observed: formula FelIFe~n(OH)12SO4.nH20 must be maintained (cf. Eqns. 1 & 2)
S t a n d a r d Gibbs energy o f f o r m a t i o n o f GR2(S02-)(.~) a n d Fe(OH)2(~) The standard chemical potential of GR2(SO] )(s) can be computed using the equilibrium conditions given by Nernst's law. Before undertaking this determination, it is necessary to dispose of a consistent and reliable set of thermodynamic constants for all species involved in the computations. In particular, the value admitted for ~t~ 2+) is of the utmost importance. On the one hand, M. Pourbaix (1966) and subsequently followed by us recently (Refait & G6nin, 1993; Drissi et al., 1995; G6nin et al., 1996) suggested a value o f - 8 4 . 9 kJ mol 1, initially determined by Randall & Frandsen (1932). On the other hand, the value suggested by NBS (Wagman et al., 1982), and used by Hansen et al. (1994) for their detelmination of ~t~ )(~], is larger, - 7 8 . 9 kJ tool - I . Finally, recent studies lead to a smaller value of - - 9 1 . 5 kJ mol a (Yang, 1982; Bard et al., 1985; Kelsall & Williams, 1991), which will be used in this work. Accordingly, the ~t~ values retained for the other Fe species, listed in Table 2, were also
Fe(II)-Fe(llI) hydroxysulphate green rust and Fe(II) hydroxide taken from the work of Kelsall & Williams (1991); this allows us to have a complete set of consistent values and equations. The activities of dissolved species were computed by means of the MINTEQ A2 computer program (CLAM, 1991) considering the various complexes likely to form in the considered experimental conditions, that is FeOH +, FeSOaaq and NaSO4. The corresponding log K values given in Table 2 were computed using the retained I.t~ values for the sake of consistency. Since Fe(OH)2(s) and GR2(SO 2 ){s) are the only solid phases present during the first reaction stage in plateau A (Table 1, Fig. 4 b - c ) , the standard Gibbs energy of formation of GR2(SO]-)(~) can be determined from the corresponding equilibrium conditions. One can estimate first the standard chemical potential of Fe(OH)2(s) by using the e q u i l i b r i u m c o n d i t i o n s b e t w e e n Fe 2+ and Fe(OH)2(s) and given by: Fe 2+ + 2 H20 ~- Fe(OH)2(s) + 2 H + (4) pK = 2 pH(A) + log a[Fe 2+] with (4') pK = {g~
- I-t~ 2+]
2 la~
where (In 10) x (RT ~ = 5.708 kJ mol 1. As the formation of GR2(SOe-)(s) proceeds, dissolved Fe species are consumed and the solution becomes undersaturated with respect to Fe(OH)2(~). The equilibrium conditions between Fe 2+ and Fe(OH)2(s) are encountered at the beginning of the reaction, corresponding to pH = 8.0• (Fig. 1), c[Fe] = ( 1 5 • -3 mol 1 1 and c[S] = (113• x10 -3 mol 1 1 (Fig. 2, point 1). From these concentrations, an activity a[Fe 2+] of 2.09x 10 3 is computed. This leads to pK = 13.32• and I.t~ = -490• k J m o l - 1 ; this value matches that of -492.0 kJ mo1-1 given by Bard et al., (1985) and by Kelsall & Williams (1991). Thus, even though Fe(OH)2(s) is slightly oxidized as soon as it precipitates, the resulting variations of its chemical potential are kept within the range of the experimental error. Therefore, the standard Gibbs energy of formation of GR2(SO 2 )(s) can be evaluated from the equilibrium conditions between Fe(OH)2(s) and GR2(SO 2 )(~) met in plateau A. These conditions are given by: 6 Fe(OH)2(s) + SO2- , ~ FelIFe~1(OH)12SO4(s) + 2e Eh(A) = E~ 0.029 log a[SO 2 ]
(5) (5')
507
with E~ -{Ia~ )(s)] - It~ SO2-] 6g~ }/(2 x 96.485) if chemical potentials are expressed in kJ mol -I and Eh in V. Taking Eh(A) 0.497 + 0.007 V, pH(A) - 8.0 • 0.1, c[Fe] = (13+l)x 10 -3 m o l 1 1 a n d c [ S ] = ( 1 1 0 • 2 1 5 10 -3 mol 1-I, the activity of sulphate is obtained at 2.25 x 10 2, which leads to E~ = -0.545_+0.008 V and, taking into account the value of I.t~ determined above with its corresponding error, I.t~ )(s)] becomes - 3 7 9 0 • 10 kJ mo1-1. This value confirms that GR2(SO 2-) is able to reduce species like NO2 or NO3 into N2, N20 or NH~ spontaneously with formation of 7-FeOOH, as shown by the conditional redox potentials given in Table 3.
Plateau B: verification of the 7-FeOOH/ GR2(SO24-)(sj equilibrium conditions Knowing the value of g~ 2 )(s)], it is possible to verify the meaning of parameters measured at plateau B, which would correspond to the equilibrium conditions between GR2(SO 2 )(s) and T-FeOOH(s> But neither NBS (Wagman et al., 1982) nor Kelsall & Williams (1991) gave any value for the chemical potential of lepidocrocite. However, it can be calculated from the log K value given for the reaction: 7-FeOOH(s) + 3 H ~ ~,~ Fe 3+ + 2 HeO since log K = {g~ bt~ 3+) - 2 ~t~
(6)
TABLE 3. Conditional redox potentials for N species and GR2(SO2-). The values given are computed at pH = 7.0 and all activities equal to 1. Couples NO2/N2 NO2/N20 NO3/N2 NO3/N20 NOz/NO 3 NO2/NHI NO3/NH~ 7-FeOOH/GR2(SO] ) GR2(SOZ-)/Fe(OH)2
E(V) 0.96 0.76 0.75 0.60 0.44 0.34 0.33 0.11 -0.55
508
Ph. Refait et al.
According to Lindsay (1979), log K = 1.39. With the g~ and g ~ ) considered here ( T a b l e 2), it g i v e s l a ~ at - 4 7 5 . 5 kJ mol 1. The equilibrium conditions between GR2(SO ] )(s) and 7-FeOOH(s) are given by: FeI41FenI(OH)12SO4 ~_ 6 y-FeOOH(~) + SOl- + 6 H + + 4 e (7) Eh(B) = E~ + 0.015 log a[SO]-] - 0.089 pH(B) (7') with E~ = {g~ + 6 g~ g~ 2 )(s)]}/(4 x 96.485). Using the previous procedure and taking pH(B) - 6 . 7 + 0 . 1 , c[Fe] = ( 1 2 + 3 ) x 1 0 -3 m o l l i,e[S] = (110 + 5)x 10 3 mol 1-1, the activity of sulphate is obtained at 2.26 x 10 -2. Taking into account g~ as d e t e r m i n e d above and g~ = 3790+ 10 kJ mol 1, EOh(B) - 0.50_+0.03 V is found which would lead to Eh(B) 0.13_+0.03 V according to (7'), a value close to the e x p e r i m e n t a l Eh(B) found at 0.075_+0.01 V. Comments
The actual f o r m u l a of G R 2 ( S O ] - ) ( s ) is Fe~IFe~n(OH)12SO4.nH20 but the value of n cannot be ascertained. The actual value of g~ )~yd.] is thus, g~ 2 )hya.] = g~ + n g~ However, by using as a formula, FeIIFenI(OH)12SO4, and its corresponding value for g~ there is no influence on the equilibrium conditions of the reactions assuming a[H20] to be close to 1. Let us write, for instance, eqn. (5) with the actual formula of GR2(SO]-)(s): 6 Fe(OH)2(s ) + SO42- + nH20 FenFem(OH)12SO4(s2.nH20 + 2 e (8) Eh = E~ -- 0.029 log a[SO~-] with (8) E%(8) -
{g~
2 6 g~
As
g~
)hyd.] -- g ~
--
- ng~ 2 )hya.]
=
x 96.485) g~
Discussion
of other data
The work by Hansen et al. (1994) led to the value of -3669 kJ mo1-1 for GR2(SOZ-)(s) in the 'anhydrous form', i.e. with formula Fe~n)Fe~UI)(OH)12SO4. This value appears to be 121 kJ mol 1 larger than that found here. Such a difference corresponds to a difference in the standard equilibrium potential of the GR2(SOZ-)(s)/Fe(OH)z(s) couple of 630 mV, which is completely out of the range of the experimental errors. It is possible to illustrate how important this difference is. Using the value given by Hansen et al. (1994), we compute the redox potentials at pH = 7 and [SOl-] = 1 of GR2(SOZ-)(s)/Fe(OH)2(s) with g~ -486.5 kJ mol -a as given by NBS (Wagman et al., 1982) and 7-FeOOH(s)/ G R 2 ( SO2 )(s) w i t h g ~ = -475.5 kJ mo1-1 as computed here from Lindsay (1979). Eh = --0.03 V and - 0 . 4 4 V is found, respectively. This means that the formation of 7-FeOOH(s) would take place directly from Fe(OH)2(s) and that GR2(SO 2 )(s) would not form. The discrepancies between our results and those given by Hansen et al. can be explained. First, the influence of the value considered for the standard free energy of formation of Fe z+ comprises about half of it; the NBS value (Wagman et al., 1982), agreed by Hansen et al., is -78.9 kJ mol 1 whereas a smaller value of 91.5 kJ mol - l is suggested here. Secondly, Hansen et al. (1994) used a completely different synthesis route to obtain the GR2(SO 2 )(s) product, involving an assumption that could be a source of error: the activity of Fe~q would be controlled by the solubility of ferrihydrite, implying the use of the solubility product of this compound (Fox, 1988), the composition of which is not ascertained. Moreover, it is not demonstrated that the GR compound prepared by Hansen et al. (1994) has the same composition, in particular the same Fe/S ratio, as here. In contrast, the results presented here, which confirm and refine the previous results by G6nin et al. (1996), are experimentally grounded measurements.
+
ng~ O) we have E~ = E~ = E~ and eqn. (8) = eqn. (5); the equilibrium conditions are exactly the same. This is of course true since this so called 'anhydrous form' is totally fictitious and designates GR2(SO]-)(s) without considering the intercalated water molecules.
CONCLUSIONS The discovery of the mineral 'foug&ite' as being a GR compound in hydromorphic soils (Trolard et al., 1996, 1997) highlights the influence of such Fe(II)Fe(III) compounds on soil reactivity and the interest of firmly established thermodynamic data (G~nin et
Fe(II)-Fe(III) hydroxysulphate green rust and Fe(II) hydroxide al., 1998b). The methodology for determining the Gibbs standard free energy of formation of G R 2 ( S O 2 )(s), A G ~ or g~ = --3790+ 10 kJ m o l i in the ' a n h y d r o u s f o r m ' , i.e. c o r r e s p o n d i n g to [Fe~IFeI~II(OH)12]2+.[SO4.nH20] 2- where n is assumed to be 0, yields a reliable value. It also demonstrated that GR2(SO 2 )(s~ is characterized by a well-defined composition, which does not vary all along the oxidation process. The Fe(II)/Fe(III) ratio stays at 1.99-1-0.12 and departures from the stoichiometry of 2 are not observed. REFERENCES Allmann R. (1970) Doppelschichtstrukturen mit brucithiinlichen Schichtionen [Me(II)l xMe(III)x(OH)2] x+. Chimia, 24, 99 108. Bard A.J., Parsons R. & Jordan J. (1985) Standard Potentials" in Aqueous Solution. Marcel Dekker Inc., New York. Bender Koch C. & Morup S. (1991) Identification of green rust in an ochre sludge. Clay Miner. 26, 577-582. Centre for Exposure Assessment and Modelling (CEAM) (1991) MINTEQ A2. USEPA Office of Research and Development, College Station Rd., Athens, GA, USA. Chariot G. (1959) Les Rdactions l~lectrochimiques. Masson, Paris. Cuttler A.H., Man V., Cranshaw T.E. & Longworth G. (1990) A MSssbauer study of green rust precipitates: I. Preparations from sulphate solutions. Clay Miner. 25, 289-301. Detournay J., De Miranda L., D6rie R. & Ghodsi M. (1975) The region of stability of GRII in the electrochemical potential-pH diagram of iron in sulphate medium. Corros. Sci. 15, 295 306. DNzsi I., Keszthelyi L., Kulgawczuk D., Molnar & Eissa N.A. (1967) M6ssbauer study of 13- and ~z-FeOOH and their disintegration products. Phys. Stat. Sol. 22, 617-629. Drissi S.H., Refait Ph., Abdelmoula M. & GNnin J.-M.R. (1995) Preparation and thermodynamic properties of Fe(iI)-Fe(III) hydroxide-carbonate (green rust one), Pourbaix diagram of iron in carbonate-containing aqueous media. Corros. Sci. 37, 2025 2041. Forsytb J.B., Hedley I.G. & Johnson C.E. (1968) The magnetic structure and hyperfine field of goethite (~-FeOOH). J. Phys. C, 1, 179-188. Fox L.E. (1988) The solubility of colloidal ferric hydroxide and its relevance to iron concentrations in river water. Geoehim. Cosmochim. Acta, 52, 771-777. G~nin J.-M.R., Bauer Ph., Olowe A.A. & R~zel D.
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