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In situ determination of the apparent solubility product of amorphous iron sulphide
Applied Geochemistry, Vol. 9, pp. 379-386, 1994 Copyright(~ 1994Elsevier Science Ltd Printed in Great Britain. All rightsreserved 0883-2927/94 $7.00+ 0.00
Pergamon
0883-2927(94)E001 I - S
In situ determination of the apparent solubility product of amorphous
iron sulphide
LARS EmK BAGANDER a n d ROLF CARMAN Department of Geology and Geochemistry, Section of Biogeochemistry, Stockholm University, S-106 91 Stockholm, Sweden (Received 12 November 1992; accepted in revised form 14 February 1994)
Abstract--The apparent (pK~) and pKs~Bs solubility products of amorphous iron sulphide (FeS) were NBS determined experimentally, in situ, in the Baltic Sea. The apparent and pKsp equilibrium constants were 2.6 and 3.15, respectively, at a mean experimental temperature of 15.8°C. The temperature dependence of the apparent equilibrium constant (within the temperature range 12.8-18.8°C) is described by the equation: pK~ = -0.068t + 3.70, where t denotes the temperature in centigrade. The composition of the amorphous precipitate was estimated to be Fe~86. The charge of the non-stoichiometric iron sulphide compound is most likely charge balanced by some major sea salt anion, e.g. chlorine and/or fluorine. Such charge balance has been observed for colloidal amorphous iron 0II) hydroxides, another common and kineticly fast-forming precipitate in natural waters.
INTRODUCTION BECAUSE of bacterially mediated sulphate reduction, thermodynamically metastable sulphide is formed in reduced environments. At low temperatures (<100°C) this is the only known mechanism of sulphate reduction (TRUDINGER et al., 1985; MACHEL, 1988). Precipitation of iron sulphides will normally follow the qualitative Ostwald step rule for sequential reactions (OsTVv'ALD,1897; MORSE and CASEY, 1988), i,e. the thermodynamically least stable iron sulphide will precipitate first. However, the formation of sedimentary iron sulphide minerals involves a complex set of reactions in which kinetics play an important role (BERNER, 1970; PYZIK and SUMMER, 1981). Different reaction pathways of iron sulphide formation have been suggested but a general conclusion is that amorphous iron sulphide (FeS) regulates the first precipitation of iron sulphide (BERNER, 1970; MORSE and CASEY, 1988) because the precipitation proceeds rapidly (PYzm and SUMMER, 1981). This phase is unstable and will crystallize to mackinawite and/or greigite (BERNER, 1967; RICKARD, 1969; HALLBERG,1972; SWEENEYand KAPLAN,1973). Pyrite is the most stable form of iron sulphide and usually the end product in natural waters. However, the mechanisms involved in the formation of pyrite are still debated. Reactions involving thiosulphate (VOLKOV and OSTROUMOV, 1957), polysulphides (HALLBERG, 1972; PANKOW and MORGAN, 1980; LUTHER, 1991), excess H2S or elemental sulphur reacting with precursor iron sulphides (RICKARD,
1969; FELD, 1977; BERNER et al., 1979) or slight oxidation of spherical greigite (SWEENEYand KAPLAN, 1973) have been proposed for pyrite formation in natural waters. Kinetics of pyrite formation in hypoxic-anaerobic environments may vary considerably, depending on sulphate availability, presence of organic matter and reactive iron, as well as influence by oxygen or oxygen-carrying species (BERNER, 1984). As an example, BOESEN and POSTMA (1988) found high contents of iron monosulphide in the Bornholm basin with an estimated turnover time to pyrite of several hundred years. However, pyrite may form rapidly in sulphide-rich environments (BOESEN and POSTMA, 1988; SKEI, 1988). Bacterial mediated reduction of iron(III) to iron(II) is one of the most important geochemical reactions in anaerobic aquatic environments. This is because of its broad influence on the overall chemistry of these environments. For example, excess sulphide in reduced marine environments results in scavenging of iron(II), largely inhibiting formation of iron carbonates and phosphates. Hence, the apparent solubility equilibrium constant of the amorphous iron sulphide within different aquatic environments is an important tool which can be used to define geochemical controls on sediment sink-source of, for instance, the fertilizer phosphorus as well as for many t r a c e and heavy metals because of their close connection with the iron and sulphur cycle. The aim of this work was to determine the solubility product of amorphous FeS under natural conditions. Experiments were carried out in situ during
379
380
L . E . Bhgander and R. Carman
the s u m m e r on a n e a r - s h o r e Baltic s o f t - b o t t o m sediment.
250 - - O2(I-I M ) 200
--
150
--
100
-
20
(*c)-~+
16 METHODS
Field methods
50-
The in situ experiments were conducted at 10 m water depth in closed sediment-water systems consisting of Plexiglass boxes (HALLaERGet al., 1972) with a salinity of 4.83%0 (calculated from concentration of major sea salt ion within the boxes; Table 1). The boxes, with the bottom open, were gently pressed down into the sediment and were closed at the start of the experiment. One transparent and two opaque boxes were used, the former to investigate the influence of light on oxygen production and H 2S consumption. The temperature varied with the ambient water temperature and was measured with thermometers inserted through rubber stoppers into the boxes. Oxygen concentrations in the box water were measured with a YSI model 57 electrode and the redox potential with a platinum electrode. Water samples were taken with 50 ml syringes through self-sealing rubber membranes at time intervals indicated in Fig. 1A. All work under water was carried out using SCUBA equipment.
A I
0
I
12~ [l=e~Tt t~M) V
t x 8 6-4
B
2-0 1500
I
I
I
I
I
i
I
10
15
20
25
H2S(tot)(t~M)
1000
500
I
Laboratory methods
12
I
5
Day
All water samples were stored, in closed syringes, in darkness at 5°C. Analysis was performed within 1 h of sampling. Sulphide and iron analyses were performed in 100 ml incubation flasks flushed with nitrogen gas to prevent oxidation. Water samples were filtered through 0.45/~m Millipore filters directly into the flasks through self-sealing rubber membranes and the reagents were added with syringes. All glassware was cleaned with 6 M HC1 and rinsed with double-distilled and deionized water. Sulphide was analysed spectrophotometrically at a wavelength of 670 nm using the methylene blue complex method (CLINE, 1969). Dissolved Fe(II) and total dissolved Fe were determinated spectrophotometricallyat 533 nm by using the method described by EICHELSDORFERand ROSOPULO(1968) with modifications according to LINDSTRGM(1980). NO significant difference was found between concentration of Fe(II) and Fe(total), and between duplicates of both the Fe and H2S analyses. The pH of water samples was measured on the National Institute for Standards and Technology pH scale (BATES, 1964). Chloride was measured with an Orion 94-17-06 solid-state ion specific electrode. Sulphate was analysed after precipitation as BaSO4 and stabilized in the solution by gum arabic and measured spectrophotometrically at 450 nm (modified from V O G E L , 1961). Sodium, magnesium, calcium and potassium were analysed with a
Table 1. Major sea salt ion concentration in the experimental solution
Ion CINa ÷ Mg 2+ SO]Ca 2+ K÷ HCO 3
Concentration (M) 1.1 6.9 1.1 5.7 1.6 1.6 1.6
x x x x x x x
10 -1 10 -2 10-2 10-3 10-3 10-3 10 -3
F]G. 1. Experimental variations of: (A) oxygen concentration and temperature for the three systems; (B) dissolved iron(II) concentration; (C) dissolved sulphide concentration. In (B) and (C) A = opaque system, 1, + = opaque system 2 and O = transparent system. The equilibrium constants have been calculated on data between days 7 and 22.
Varian Techtron model AA-6 atomic absorption spectrophotometer according to the manual supplied by the manufacturer. Alkalinity and total carbonate concentration were determined by potentiometric titration with HCI (DvRSSEN and SILL~N, 1967).
THEORY
T h e r e are three types of equilibrium constants used to describe activities of reactants and products, t h e r m o d y n a m i c , pKs0 (e.g. SILL~N and MARTELL, 1964), a p p a r e n t , pK" (e.g. GOLDHABER and KnPLAN, 1975) and p K ~ Bs ( D n v I s o n , 1980). For t h e r m o dynamic constants the activity scale is defined so that the activity coefficient (y), that is, the q u o t i e n t activity/concentration, is equal to 1 at infinite dilution. W h e n using an a p p a r e n t constant the q u o t i e n t is equal to 1 at the ionic strength (1) of the e x p e r i m e n t a l m e d i u m . C o n s t a n t s derived by Davison's m e t h o d use the e x p e r i m e n t a l data g a t h e r e d at some ionic strength, along with individual ion activity coefficients to correct to zero ionic strength. In the usual analytical d e t e r m i n a t i o n of a specific e l e m e n t all dissolved species are included, not only the free ion. The single ion activity coefficient
Solubility product of amorphous iron sulphide depends on the ionic strength and successive extent of ion association. The extent of ion-pair and complex formation is generally not well known over the ranges of temperature, pressure and concentration that occur in nature. This is the main difficulty when using thermodynamic equilibrium constants in concentrated multi-electrolyte solutions (PYTKOWICZ, 1969). Previous authors have determined the thermodynamic solubility product (Ks0) for precipitated amorphous FeS and FeS(solid), according to the reaction FeS ~-~ Fe z+ + S2- , either by laboratory experiments (e.g. BERNER, 1967) or by theoretical calculations (SILLI~Nand MARTELL,1964). To calculate the/(so for the reaction above, the value of the second dissociation constant for H2S is needed. This constant is very small and has not yet been determined accurately (MORSEet al., 1987; SCHOONENand BARNES, 1988, 1992). The value of/(so for FeS is, thus, uncertain. At the pH of most natural sea waters (7-8) the concentration of S2- is negligible compared to the concentration of H2S and H S - and can thus be disregarded. By studying the pH dependence of metal sulphide dissolution, POnE (1954) deduced that the rate-determining step involves the H S - ion rather than S2 . This agrees with ELLIS and GIGGENBACH (1971) who assert that a reaction such as H + + FeS ~ Fe 2+ + H S - is a better explanation of the equilibrium. The apparent solubility product involves the sum of the concentration of the free and complex species (GOLDHABERand KAPEAN, 1975). However, this precludes comparison between waters of different ionic strength. Therefore, DAVISON (1980) introduced a new operational equilibrium constant to overcome some of these comparison problems, which accounts for differences in ionic strength and ionic constituency.
CALCULATION OF THE EQUILIBRIUM C O N S T A N T S OF A M O R P H O U S I R O N S U L P H I D E
Influences of complexation in calculation of the equilibrium constants of amorphous iron sulphide The subscript zero in the thermodynamic equilibrium constant pKs0 denotes that the equilibrium of the solid ML (M = metal; L = ligand) only considers the uncomplexed species M and L (SILE~N and MARTELL, 1964). Thus, Ks0 for amorphous FeS is /(so = YFe2+[Fe2+]ys2-[S2-],
(1)
where y denotes the single ion activity coefficient and [] the concentration. The total dissolved concentration of sulphide can be expressed as [H2S]tot = [H2S] + [HS-] + [Sz-] + polysulphides + thio complexes.
(2)
381
Calculations of polysulphide concentrations in the experimental solution were made based on data from BOt~LECUE and MICHARO(1978). At pH 7 the polysulphide concentration can reach a maximum of 1% of the total H2S(aq) concentration and was thus disregarded. The complexation between sulphur and heavy metals to form thio complexes is also negligible at the experimental conditions (S~UMM and BRAUNER, 1975; DAVISON, 1980). Besides free Fe 2+ ions, the complex's F e O H ÷, FeCI- and FeSO4° may affect the total dissolved concentration of Fe 2÷ (DAvISON, 1979). Thus, [Fe2+]tot--- [Fe 2+] + [FeOH +]
+ [FeCl-] + [VeSO°].
(3)
We have used selected stability constant data for iron(II) complexation to calculate the iron(II) speciation. The maximum percentage complexation under experimental conditions was around 9% (calculated by using a modified version of the equilibrium program H A L T A F A L L (INGR! et al., 1977)) of the total dissolved iron content where FeSO °, FeC1- and F e O H + constitutes 4, 4 and 1%, respectively. However, since there is a general uncertainty regarding these complexation constants (DAvISON, 1979) we assume that the influence from complexation is negligible compared with the uncertainty in the colorimetric measurements of Fe(II) (HEANEV and DAVISON, 1977). In addition, a test of the analytical method with and without C1- and SO]- ions at experimental concentrations showed that only free Fe z+ ions reacted with the iron reagent. Hence, we have assumed that Fe 2+ = Fetot.
Calculation of pK s of amorphous iron sulphide The second dissociation constant, K2, is difficult to determine. ELLIS and GIGGENaACH(1971) suggested therefore that the solubility product of iron sulphide should be defined according to the following reaction: FeS + H + = H S - + Fe 2+.
(4)
The apparent solubility product, K~, is then K ~ - [Fe2+]t°t[HS-]t°t (H+) ,
(5)
and the first apparent dissociation constant of hydrogen sulphide is (GOLDHABERand KAPEAN, 1975): ,
K x-
(H+)*[HS -] [HzS(aq)]
(6)
in which (H+) * is an empirical quantity defined by pH = - l o g (H+) *, where pH is measured on the National Institute for Standards and Technology pH
382
L . E . Bfigander and R. Carman Table 2. Measured and calculated parameters for the experimental systems. The values are given for the opaque 1, 2 and transparent systems for each day, respectively
[Fe2+]tot
[H2S]tot
Day
Temp. (°C)
pH
Q~M)
(~M)
pKs0
pK~
7
12.8
8
15.0
9
14.3
11
18.1
12
16.8
13
15.8
14
16.7
15
18.0
16
18.3
22
18.8
7.01 6.99 7.01 7.01 6.99 7.01 7.00 6.98 7.01 7.00 6.98 7.00 7.00 6.98 7.00 7.00 6.97 7.00 6.99 6.97 7.00 7.00 6.98 7.00 7.00 6.98 7.00 7.04 7.01 7.01
5.55 5.01 4.66 4.30 3.58 3.94 3.94 3.58 3.58 1.79 1.79 1.79 1.79 1.97 2.15 1.43 1.43 1.43 1.07 1.25 1.25 1.07 1.07 1.07 0.716 0.895 1.43 0.716 0.716 0.716
77.4 38.7 50.6 140.0 140.0 89.4 167.0 176.0 149.0 363.0 432.0 372.0 524.0 429.0 444.0 545.0 448.0 506.0 575.0 539.0 569.0 685.0 748.0 563.0 795.0 727.0 563.0 1330.0 1310.0 903.0
17.53 17.91 17.89 17.30 17.41 17.55 17.30 17.35 17.38 17.15 17.11 17.14 17.04 17.12 17.04 17.17 17.26 17.20 17.25 17.24 17.17 17.10 17.10 17.20 17.20 17.17 17.05 16.90 16.95 17.11
2.76 3.14 3.02 2.60 2,71 2,83 2.58 2.63 2.66 2.55 2.51 2.54 2.40 2.48 2.39 2.49 2.59 2.52 2.60 2.59 2.52 2.50 2.49 2.58 2.61 2.58 2.46 2.32 2.39 2.53
scale (BATES, 1964). The total dissolved concentration of sulphide at the experimental p H can be expressed as [H2Sltot = [H2S ] + [ H S - ] .
p K ~ Bs : K~ YFe:+YHS •
(8)
This equilibrium constant is corrected for differences in ionic strength. H o w e v e r , it is not entirely independent of ionic strength since the effect of the liquid junction potential of p H measurements has not been accounted for.
[HeS]tot
(H+) ,
- - + 1
The expression below is from DAVlSON (1980):
(7)
Rearranging (6) and (7) gives
[HS-I =
Calculation o f p K NBs of amorphous iron sulphide
(11)
Ki
RESULTS AND DISCUSSION The negative logarithm of (8) combined with (5) gives pK~ = - l o - [Fez+ ]t°--t[H2S]t°t
(9)
g (H+)* + (H+) * K~ K[ was calculated for the experimental temperatures using the equation pK~ = 2.527 - 0.169C11/3 + 1359.96/T
(10)
from GOLDHABER and KAPLAN (1975), where C1 stands for chlorinity and T for temperature in Kelvin.
A t the beginning of the experiment the sediment was oxidized to a depth of 2-3 m m (Eh > 200 mV; MORXIMER, 1942). The variations of the experimental temperature are shown in Fig. la. During the first days there was a rapid p H decrease, around 1 p H unit, in all three boxes. As a result of oxidation of organic matter CO2(g ) produced in the box cannot leave the closed system. Subsequent to that, p H stabilizes around 7 (Table 2) in all boxes, indicating a sulphate-sulphide system control of p H (BENYAAKOV, 1973). Figure 1 A - C shows the variations in temperature and concentrations of oxygen, Fe(II) and H2S. Dissolved oxygen in the water was con-
Solubility product of amorphous iron sulphide sumed within 5 d of the beginning of each experiment in all three boxes (Fig. 1A). Redox potential in the three experimental systems was measured occasionally during the experiment and used only as an environmental parameter (BAGANDER and NIEMISTO, 1978). In nalural waters with high dissolved oxygen content and near-neutral or higher pH, the concentration of soluble iron(Ill) is very low. Most of the soluble and suspended iron constituents occur as solid iron(III) oxyhydroxides. The increase in total dissolved Fe(II) during the first 4 days was, hence, due to reduction of the Fe(III) constituents in the water phase. At oxygen concentrations below the detection limit (<3/~M) dissolution of solid iron(Ill) phases released dissolved iron(II) from the sediment to the overlaying water. Between days 6 and 22, Millipore filters (0.45 ktm) had a residue of black precipitated iron sulphide, which was X-ray amorphous. Hydrogen sulphide could be detected by odour from day 6 and detected analytically from day 7. The decrease in H2S in the transparent box from day 15 (Fig. 1C) was due to oxidation by photosynthetic sulphur bacteria. The experiment continued after day 22 but Fe(II) was then below the detection limit (0.36/~M). Different solubility products of amorphous FeS were indirectly calculated from solution chemistry between day 7 [Fe(II) decrease] and 22 [limit of Fe(II) detection] (Fig. 1B). At the experimental pH, around 7 (Table 2), the HCO~ concentration was almost equal to the total carbonate concentration (deduced from STUMMand MORGAN, 1981). The concentration of NH2, H2PO4, H P O ] - and H S - (maximum 0.2 x 10 -3 M measured in earlier experiment) was too small to contribute to the overall ionic strength. The ionic strength, 0.128, of uncomplexed media, was calculated from the values in Table 1. This value is too high since the formation of ion-pairs has not been considered. To determine the maximum influence of ion-pairing we have used a modified computer version of HALTAFALL (INGRI et al., 1977) which includes a large thermodynamic database. The program uses the Davies" equation to consider the ionic strength and calculate the activity coefficients (log y). The results from these calculations are expressed in Table 3, which give the ionic strength 0.122. The activity coefficients of Fe e+ and HS- were determined by Davies' equation (e.g. Sxurar~ and MORGAN, 1981). The activity coefficient of H2S(aq) differs only slightly from unity (1.03) for sea water (DOUABULand RILEY, 1979). We have thus calculated pK~0 of amorphous FeS with YH.~S,,,,,~equal to 1.
383
Table 3. Maximum ion-pairing concentration Ion compound
% of total concentration
CI Na + Mg2+ MgSO° Ca2+ CaSO~ K+ SO] HCO~
100 99 87 8 90 9 99 76 89
Concentration (M) 1.10 x 6.90 X 9.60 x 4.60 x 1.44 x 1.40 x 1.60 x 4.30 x 1.42 x
10-I 10 -2
10 3 10 4 10 3 I0 4 10-3 10 3 10-.3
3.2-
3.0-
2.8 pK"s 2.6
:4-
2.4
2.2
'
12.0
'
I
'
14.0 16.0 Temperature (°C)
I
180
'
I
20.0
FIG. 2. Temperature dependence of the apparent equilibrium constant of amorphous iron sulphide (r = 0.75. n = 30).
pK; = -0.068t + 3.70.
(12)
with r = 0.75 (n = 30) at 1% level of significance (t denotes the temperature in celsius). Table 4 shows measured apparent solubility products from different environments together with the equilibrium constant defined by DAVISON(1980). The apparent and pK~ Bs equilibrium constant of the present study is given for the average experimental temperature (15.8°C) together with equilibrium constants from other investigations, p K ~ Bs values above 4 are probably related to greigite or mackinawite equilibria, according to data from BERNER (1967). The other values agree rather well but differences in analytical methods can partly explain some discrepancies as well as different stoichiometric compositions depending on the ionic media.
Apparent solubility constant, pK~, and Davison's equilibrium constant, p K ~ Bs
Solubility diagram
Linear regression analysis on the apparent equilibrium constants (Fig. 2) gave the equation:
The concentration HS- in the water was calculated from analytical values of [H2S]tot using Eqns (8) and
AG 9: 4-C,
L. E. B~gander and R. Carman
384
Table 4. pK~ and pKNpBs solubility products from different investigations Site
pK~
pK NBs
Lake Nitinat Cariaco Trench Skejennungen Laboratory Laboratory Baltic Sea Black Sea Byfjorden Lake Greifensee Clear Lake Esthwaite Abereiddy Kiel Bight Framvaren Baltic Sea
3.79 3.34 1.95
4.43 4.04 2.06 2.97 3.68 4.34 3.53 3.74 2.99 3.23 2.59 2.55 2.47 3.67 3.15
3.15 3.7 2.85 3.1 2.8 3.12 2.48 1.94 1.8 3.0 2.6
pH
7.1 0.6 7.8 0.7 6.8 0.002 3.7 ---)0 4.0 0.1 -0.24 7.7 0.45 -0.62 7.05 0.007 6.9 0.002 7.13 0.002 7.17 0.7 7.6 0.42 7.0 0.40 7.0 0.12
(10). Values of p[HS-] were then plotted against p[FeZ+]tot (Fig. 3). Linear regression analysis gave the equation: p[Fe2+]to, = - 0 . 6 7 p [ H S - ] + 8.44.
(13)
The values determined prior to the [Fez+ ]tot decrease in the box water are excluded. The correlation coefficient is 0.95 and significant at the 1% level (n = 30). The abscissa (Fig. 3) also could be looked upon as a time-axis since [HS-] increased with time during the experiment. Ideally, at constant temperature, amorphous FeS would precipitate when the equilibrium solubility product is reached as [HS-] increases along a p[Fe 2+ ]tot line. The concentrations of [FeZ+]tot and [HS-] would then follow a line with the slope - 1, if the mole fraction of amorphous FeS is 1. Thus, Eqn (13) indicates precipitation of a solid amorphous iron sulphide phase deficient in sulphur i.e. FeS0.67. However, a "true" equilibrium was probably not reached immediately in this in situ experiment due to a number of "uncontrolled" factors. From the data presented in Fig. 3 it appears that there
5.2
5.6 p F e 2+ 5.8
:
~, ~.
6.0-
6.2
I 5.2
4.8
'
I 4.4
I
I
4.0
3.6
pHS -
'
I 3.2
3. Solubility diagram of amorphous iron sulphide (r = 0.95, n = 30). Days 11-22 = pHS- values below 4.
FIG.
[I]
t (°C) 10 17 4 25 25 5 9 5 5 10 10 10 14 9 16
Reference RICHARDSet al. (1965) RICHARDS(1965) KJENSMO(1967) BERNER(1967) DOYLE(1968) KOROLEFF(1968) Brewer (1971), in DAVISON(1980) DANIELSSONet al. (1975) EMERSON(1976) LAHANN(1977) DAV1SONand HEANEY(1978) DAV1SONand HEANEY(1980) BALZER(1982) SKEI(1988) This study
are two populations of data, with two slopes. The first population of data arise from measurements from the day 7, 8 and 9. The slope of these data is -0.21 and, hence, it is not unreasonable to say that the Fe2* concentration is nearly independent of [HS-] at these low sulphide concentrations. Examining later experimental values (from day 11, Fig. 3) in the same way, linear regression analysis using all data gave the equation: p[Fe2+]tot = - 0 . 8 6 p [ H S ] + 9.18
(14)
with r = 0.90 and a significance level of 1% (n = 21). In this case, precipitation of FeS0.86 is suggested. This is close to the range of composition of amorphous FeS0.87q).92 reported by SWEENEY and KAPLAN (1973). A closer examination of the data from days 11-22 (Fig. 3) suggests that there are two data points near the origin of the figure (from the two opaque boxes at day 22, see Table 2) which could be looked upon as outliers and hence be excluded from the linear regression analysis. Without these two points regression analysis yielded a different interpretation giving an opposite picture with respect to deficiency in sulphur vs iron in the assumed precipitate i.e. the slope is -1.14. It is, however, difficult from a scientific point of view, to explain why these data should be excluded. One could instead look upon some other data points from the later part of the experiment that could be subject to exclusion, i.e. the sulphide data from the transparent box, days 15, 16 and 22. These data differ conspicuously from the data obtained from the two opaque boxes (Table 2). A probable explanation to this could be that photoautotophic sulphur bacteria ( C h r o m a t i u m and/or Chlorobium) oxidize some parts of the produced sulphide. Linear regression without these three points gave a slope close to that with all data points included (-0.83). Hence, under natural conditions one can expect to find variations in the composition of amorphous FeS due to environmental factors. For example, BERNER (1964) and RICKARD (1969) reported higher sulphur content of FeS1.1 which they
Solubility product of amorphous iron sulphide attributed to be adsorption of hydrogen sulphide. HALLBERG (1970) reported a FeS ratio close to 1 and suggested (HALLBERG, 1972) that the ratio of 1:1 may be explained by the co-existence of greigite. All these ratios were measured from precipitate compositions whereas our calculated values are based solely on solution chemistry. The divergence from the ideal composition of FeS (1:1) could be explained by errors in the analytical techniques, e.g. overestimation of iron(II) concentrations due to particulate iron species <0.45 ~ m and/or loss of sulphide due to an oxidation. A n o t h e r explanation could be that the measured deficiency of sulphur, based on solution chemistry, is charge balanced by some other anion. Such charge balances have been observed for the hydrolysis of another c o m m o n amorphous precipitate during early diagenesis, i.e. amorphous iron(III) hydroxide (e.g. Fox, 1988; CORNELL et al., 1989). These anions provide charge balance but do not seem to affect the solubility of the precipitate (Music et al., 1982; Fox, 1988). In an applied sense the knowledge of the equilibrium constant for the least stable iron sulphide is important to know because of its broad influence on the organic and inorganic chemistry in all aquatic environments with permanent or temporary anoxic conditions. It regulates to a great extent the amount of upward migration of both nutrients as well as trace and heavy metals and, thus, determines the environmental biogeochemical conditions of the examined aquatic environment. Acknowledgements--We thank Professor R. O. Hallberg
for valuable criticism and discussions during the preparation of the manuscript. Editorial handling: John Catts.
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BERNERR. A., BALDWINT. and HOLDRENG. R., JR (1979) Authigenic iron sulfides as paleosalinity indicators. J. sedim. Petrol. 3, 335-338
385
BOESENC. and POSTMAD. (1988) Pyrite formation in anoxic environments of the Baltic. Am. J. Sci. 288,575-603. BOULEGUEJ. and MICHAROG. (1978) Constantes de formation des ions polysulfures S2-, S2-, S2- en phase aqueuse. J. Fr. Hydrol., 9 fasc. 1. 25, 27-34. CLINE J. D. (1969) Spectrophotometric determination of hydrogen sulphide in natural waters. Limnol. Oceanogr. 14,454-458. CORNELE R. M., GIOVANOLIR. and SCHNEIDERW. (1989) Review of the hydrolysis of Iron(Ill) and the crystallization of amorphous Iron(III) hydroxide hydrate. J. Chem. Tech. Biotechnol. 46, 115-134. DANIELSSON L.-G., DYRSSEN D., JOHANSSON T. and NYQUIST G. (1975) The By Fjord: Chemical investigations. SNV PM 609, 1-85. DAVISONW. (1979) Soluble inorganic ferrous complexes in natural waters. Geochim. cosmochim. Acta 43, 16931696. DAVlSONW. (1980) A critical comparison of the measured solubilities of ferrous sulphide in natural waters. Geochim. cosmochim. Acta 44, 803-808. DAVISONW. and HEANEYS. I. (1978) Ferrous iron-sulphide interactions in anoxic hypolimnic waters. Limnol. Oceanogr. 23, 1194-1200. DAVISONW. and HEANEYS. l. (1980) Determination of the solubility of ferrous sulphide in a seasonally anoxic marine basin. Limnol. Oceanogr. 25,153-156. DOUABULA. A. and RILEYJ. P. (1979) The solubility of gases in distilled water and seawater--V. Hydrogen sulphide. Deep-Sea Res. 26A, 259-268. DOYLE R. W. (1968) Identification and solubility of iron sulphide in anaerobic lake sediments. Am. J. Sci. 266, 980-994. DYRSSEN D. and SIELI~NL. G. (1967) Alkalinity and total carbonate in sea water. A plea for p-T-independent data. Tellus 19, 113-121. EICHEESD()REERD. and ROSOPULOA. (1968) Methoden der Eisenbestimmung in Trink- und Betriebswasser. Vom Wasser 34, 82-96. ELLIS A. J. and GIGGENBACHW. (1971) Hydrogen sulphide ionization and sulphur hydrolysis in high-temperature solution. Geochim. cosmochim. Acta 35,247-260. EMERSONS. (1976) Early diagenesis in anaerobic lake sediments: Chemical equilibria in the interstitial waters. Geochim. cosmochim. Acta 40, 925-934. FEED W. (1977) 15ber die Bildung von Eisenbisulfid (FeS~) in L6sungen und die Entstehung der naturlichen Pyritlager. Z. Angew. Chem. 24, 97-103. Fox L. E. (1988) Solubility of colloidal ferric hydroxide. Nature 333,442-444. GOLDHABER M. B. and KAPLAN 1. R. (1975) Apparent dissociation constants of hydrogen sulphide in chloride solutions. Mar. Chem. 3, 83-104. HAEEBERGR. O. (1970) An apparatures for the continuous cultivation of sulfate reducing bacteria and its application to geomicrobial purposes. Antonie van Leeuwanhoek 36, 241-254. HALEBERGR. O. (1972) Iron and zinc sulfides formed in a continuous culture of sulfate-reducing bacteria. Neues Jb. Mineral Monatsh. II, 481-500. HALLBERG R. O., B/kGANDER L. E., ENGVALLA.-G. and SCHIPPELF. A. (1972) Method for studying geochemistry of sediment-water interface. Ambio 1, 71-72. HEANEYS, I. and DAVISONW. (1977) The determination of ferrous iron in natural waters with 2,2'bipyridyl. Limnol. Oceanogr. 22, 753-760. Iy~m N., KAKOLOWlCZW., S1LLI~.NL. G. and WARNQU1STB. (1977) Haltafall, a general program for calculating the composition of equilibrium mixtures. Talanta 14, 12611286. KJENSMO J. (1967) The development and some main features of "iron-meromictic" soft water lakes. Archs Hydrobiol. Suppl. 32, 137-312.
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KOROLEFFF. (1968) A note on the iron content of the Baltic waters. ICES. C.M. 1968/C34. LAHANN R. W. (1977) Molybdenum and iron behavior in oxic and anoxic lake water. Chem. Geol. 20, 315-323. LINDSTROM C. H. M. (1980) Transformation of iron constituents during early diagenesis--in situ studies of a Baltic Sea sediment-water interface. Contribution in Microbial Geochemistry. Department of Geology, University of Stockholm. LUTHER G. W., III. (1991) Pyrite synthesis via polysulfide compounds. Geochim. cosmochim. Acta 55, 2839-2849. MACHEL H. G. (1988) Some aspects of diagenetic sulfatehydrocarbon redox reactions. In Diagenesis of Sedimentary Sequences (ed. J. D. MARSHALL). Spec. Publs geol. Soc. Lond. 36, 15-28. MORSE J. W. and CASEY W. H. (1988) Ostwald processes and mineral paragenesis in sediments. Am. J. Sci. 288, 537-560. MORSE J. W., M1LLEROF. J., CORNWELLJ. C. and RICKARD D. (1987) The chemistry of the hydrogen sulphide and iron sulphide systems in natural waters. Earth Sci. Rev. 24, 1-42. MORTIMER C. H. (1942) The exchange of dissolved substances between mud and water in lakes. J. Ecol. 30, 147201. Music S., VERTESA., SIMMONSG. W., NAGY-CZAKOI. and LEDIHEISER H. J. (1982) M6ssbauer spectroscopic study of the formation of Fe(III) oxyhydroxides and oxides by hydrolysis of aqueous Fe(lII) salt solutions. J. Colloid Interface Sci. 85,256-266. OSTWALD W. Z. (1897) Studien fiber die Bildung und Umwandlung fester K6rper. 1. Abhandlung: lJbers/ittigung und 1Jberkfiltung. Z. Physik Chemie 22, 289-330. PANKOW J. F. and MORGAN J. J. (1980) Dissolution of tetragonal ferrous sulphide (mackinawite) in anoxic aqueous systems, 1. Implications for cycling of iron, sulfur and trace metals. Environ. Sci. Technol. 14, 183186. POHL H. A. (1954) The formation and dissolution of metal sulfides. J. Am. Chem. Soc. 76, 2182-2184. PYTKOWICZ R. M. (1969) Use of apparent equilibrium constants in chemical oceanography, geochemistry and biochemistry. Geochem. J. 3, 182-184.
PYZIK A. J. and SOMMER J. E. (1981) Sedimentary iron monosulfides: Kinetics and mechanisms of formation. Geochim. cosmochim. Acta 45,687-698. RICHARDS F. A. (1965) Chemical observations in some anoxic sulphide bearing basin sand fjords. In Proc. Second Int. Water Poll. Res. Conf., 1964, pp. 215-243. Pergamon, Oxford. RICHARDS F. A., CLINE J. D., BROENKOW W. W. and
ATK1NSONL. P. (1965) Some consequences of the decomposition of organic matter in Lake Nitinat, an anoxic fjord. Limnol. Oceanogr. 10, Suppl., R185-201. RICKARD D. T. (1969) The chemistry of iron sulphide formation at low temperatures. Stockholm Contr. Geol. 20, 67-95. SCHOONEN M. A. A. and BARNES H. L. (1988) An approximation of the second dissociation constant for H2S. Geochim. cosmochim. Acta 52, 649-654. SCHOONEN M. A. A. and BARNES H. L. (1992) Reaction forming pyrite and marcasite from solution, parts l and 2. Geochim. cosmochim. Acta 55, 1495-1514. SILLl~N L. G. and MARTELL A. E. (1964) Stability Constants of Metal Ion Complexes (2rid ed.). Spec. Publ. Chem. Soc. Lond. 17. SKEI J. M. (1988) Formation of framboidal iron sulfides in the water of a permanently anoxic fjord--Framvaren, South Norway. Mar. Chem. 23,345-352. STUMMW. and BRAUNERP. A. (1975) Chemical speciation. In Chemical Oceanography (2rid ed.) (eds J. P. RILEYand G. SKIRROW), VOL 3, pp. 173-239. Academic Press, New York. STUMM W. and MORGAN J. J. (1981) Aquatic Chemistry. Wiley-lnterscience, New York. SWEENEY R. E. and KAPLAN l. R. (1973) Pyrite framboid formation: Laboratory synthesis and marine sediments. Geology 68, 618-634. TRUDINGER P. A . , CHAMBERSL. A. and SMITHJ. W. (1985) Low-temperature sulphate reduction: Biological versus abiotical. Can. J. Earth Sci. 22, 1910-1918. VOGEL A. (1961) A Textbook of Quantitative Analysis. Longmans, London. VOLKOV I. I. and OSTROUMOVE. A. (1957) Iron sulphide concentrations in Black Sea deposits. Akad. Nauk. SSSR 116, 645-648.
Pergamon
0883-2927(94)E001 I - S
In situ determination of the apparent solubility product of amorphous
iron sulphide
LARS EmK BAGANDER a n d ROLF CARMAN Department of Geology and Geochemistry, Section of Biogeochemistry, Stockholm University, S-106 91 Stockholm, Sweden (Received 12 November 1992; accepted in revised form 14 February 1994)
Abstract--The apparent (pK~) and pKs~Bs solubility products of amorphous iron sulphide (FeS) were NBS determined experimentally, in situ, in the Baltic Sea. The apparent and pKsp equilibrium constants were 2.6 and 3.15, respectively, at a mean experimental temperature of 15.8°C. The temperature dependence of the apparent equilibrium constant (within the temperature range 12.8-18.8°C) is described by the equation: pK~ = -0.068t + 3.70, where t denotes the temperature in centigrade. The composition of the amorphous precipitate was estimated to be Fe~86. The charge of the non-stoichiometric iron sulphide compound is most likely charge balanced by some major sea salt anion, e.g. chlorine and/or fluorine. Such charge balance has been observed for colloidal amorphous iron 0II) hydroxides, another common and kineticly fast-forming precipitate in natural waters.
INTRODUCTION BECAUSE of bacterially mediated sulphate reduction, thermodynamically metastable sulphide is formed in reduced environments. At low temperatures (<100°C) this is the only known mechanism of sulphate reduction (TRUDINGER et al., 1985; MACHEL, 1988). Precipitation of iron sulphides will normally follow the qualitative Ostwald step rule for sequential reactions (OsTVv'ALD,1897; MORSE and CASEY, 1988), i,e. the thermodynamically least stable iron sulphide will precipitate first. However, the formation of sedimentary iron sulphide minerals involves a complex set of reactions in which kinetics play an important role (BERNER, 1970; PYZIK and SUMMER, 1981). Different reaction pathways of iron sulphide formation have been suggested but a general conclusion is that amorphous iron sulphide (FeS) regulates the first precipitation of iron sulphide (BERNER, 1970; MORSE and CASEY, 1988) because the precipitation proceeds rapidly (PYzm and SUMMER, 1981). This phase is unstable and will crystallize to mackinawite and/or greigite (BERNER, 1967; RICKARD, 1969; HALLBERG,1972; SWEENEYand KAPLAN,1973). Pyrite is the most stable form of iron sulphide and usually the end product in natural waters. However, the mechanisms involved in the formation of pyrite are still debated. Reactions involving thiosulphate (VOLKOV and OSTROUMOV, 1957), polysulphides (HALLBERG, 1972; PANKOW and MORGAN, 1980; LUTHER, 1991), excess H2S or elemental sulphur reacting with precursor iron sulphides (RICKARD,
1969; FELD, 1977; BERNER et al., 1979) or slight oxidation of spherical greigite (SWEENEYand KAPLAN, 1973) have been proposed for pyrite formation in natural waters. Kinetics of pyrite formation in hypoxic-anaerobic environments may vary considerably, depending on sulphate availability, presence of organic matter and reactive iron, as well as influence by oxygen or oxygen-carrying species (BERNER, 1984). As an example, BOESEN and POSTMA (1988) found high contents of iron monosulphide in the Bornholm basin with an estimated turnover time to pyrite of several hundred years. However, pyrite may form rapidly in sulphide-rich environments (BOESEN and POSTMA, 1988; SKEI, 1988). Bacterial mediated reduction of iron(III) to iron(II) is one of the most important geochemical reactions in anaerobic aquatic environments. This is because of its broad influence on the overall chemistry of these environments. For example, excess sulphide in reduced marine environments results in scavenging of iron(II), largely inhibiting formation of iron carbonates and phosphates. Hence, the apparent solubility equilibrium constant of the amorphous iron sulphide within different aquatic environments is an important tool which can be used to define geochemical controls on sediment sink-source of, for instance, the fertilizer phosphorus as well as for many t r a c e and heavy metals because of their close connection with the iron and sulphur cycle. The aim of this work was to determine the solubility product of amorphous FeS under natural conditions. Experiments were carried out in situ during
379
380
L . E . Bhgander and R. Carman
the s u m m e r on a n e a r - s h o r e Baltic s o f t - b o t t o m sediment.
250 - - O2(I-I M ) 200
--
150
--
100
-
20
(*c)-~+
16 METHODS
Field methods
50-
The in situ experiments were conducted at 10 m water depth in closed sediment-water systems consisting of Plexiglass boxes (HALLaERGet al., 1972) with a salinity of 4.83%0 (calculated from concentration of major sea salt ion within the boxes; Table 1). The boxes, with the bottom open, were gently pressed down into the sediment and were closed at the start of the experiment. One transparent and two opaque boxes were used, the former to investigate the influence of light on oxygen production and H 2S consumption. The temperature varied with the ambient water temperature and was measured with thermometers inserted through rubber stoppers into the boxes. Oxygen concentrations in the box water were measured with a YSI model 57 electrode and the redox potential with a platinum electrode. Water samples were taken with 50 ml syringes through self-sealing rubber membranes at time intervals indicated in Fig. 1A. All work under water was carried out using SCUBA equipment.
A I
0
I
12~ [l=e~Tt t~M) V
t x 8 6-4
B
2-0 1500
I
I
I
I
I
i
I
10
15
20
25
H2S(tot)(t~M)
1000
500
I
Laboratory methods
12
I
5
Day
All water samples were stored, in closed syringes, in darkness at 5°C. Analysis was performed within 1 h of sampling. Sulphide and iron analyses were performed in 100 ml incubation flasks flushed with nitrogen gas to prevent oxidation. Water samples were filtered through 0.45/~m Millipore filters directly into the flasks through self-sealing rubber membranes and the reagents were added with syringes. All glassware was cleaned with 6 M HC1 and rinsed with double-distilled and deionized water. Sulphide was analysed spectrophotometrically at a wavelength of 670 nm using the methylene blue complex method (CLINE, 1969). Dissolved Fe(II) and total dissolved Fe were determinated spectrophotometricallyat 533 nm by using the method described by EICHELSDORFERand ROSOPULO(1968) with modifications according to LINDSTRGM(1980). NO significant difference was found between concentration of Fe(II) and Fe(total), and between duplicates of both the Fe and H2S analyses. The pH of water samples was measured on the National Institute for Standards and Technology pH scale (BATES, 1964). Chloride was measured with an Orion 94-17-06 solid-state ion specific electrode. Sulphate was analysed after precipitation as BaSO4 and stabilized in the solution by gum arabic and measured spectrophotometrically at 450 nm (modified from V O G E L , 1961). Sodium, magnesium, calcium and potassium were analysed with a
Table 1. Major sea salt ion concentration in the experimental solution
Ion CINa ÷ Mg 2+ SO]Ca 2+ K÷ HCO 3
Concentration (M) 1.1 6.9 1.1 5.7 1.6 1.6 1.6
x x x x x x x
10 -1 10 -2 10-2 10-3 10-3 10-3 10 -3
F]G. 1. Experimental variations of: (A) oxygen concentration and temperature for the three systems; (B) dissolved iron(II) concentration; (C) dissolved sulphide concentration. In (B) and (C) A = opaque system, 1, + = opaque system 2 and O = transparent system. The equilibrium constants have been calculated on data between days 7 and 22.
Varian Techtron model AA-6 atomic absorption spectrophotometer according to the manual supplied by the manufacturer. Alkalinity and total carbonate concentration were determined by potentiometric titration with HCI (DvRSSEN and SILL~N, 1967).
THEORY
T h e r e are three types of equilibrium constants used to describe activities of reactants and products, t h e r m o d y n a m i c , pKs0 (e.g. SILL~N and MARTELL, 1964), a p p a r e n t , pK" (e.g. GOLDHABER and KnPLAN, 1975) and p K ~ Bs ( D n v I s o n , 1980). For t h e r m o dynamic constants the activity scale is defined so that the activity coefficient (y), that is, the q u o t i e n t activity/concentration, is equal to 1 at infinite dilution. W h e n using an a p p a r e n t constant the q u o t i e n t is equal to 1 at the ionic strength (1) of the e x p e r i m e n t a l m e d i u m . C o n s t a n t s derived by Davison's m e t h o d use the e x p e r i m e n t a l data g a t h e r e d at some ionic strength, along with individual ion activity coefficients to correct to zero ionic strength. In the usual analytical d e t e r m i n a t i o n of a specific e l e m e n t all dissolved species are included, not only the free ion. The single ion activity coefficient
Solubility product of amorphous iron sulphide depends on the ionic strength and successive extent of ion association. The extent of ion-pair and complex formation is generally not well known over the ranges of temperature, pressure and concentration that occur in nature. This is the main difficulty when using thermodynamic equilibrium constants in concentrated multi-electrolyte solutions (PYTKOWICZ, 1969). Previous authors have determined the thermodynamic solubility product (Ks0) for precipitated amorphous FeS and FeS(solid), according to the reaction FeS ~-~ Fe z+ + S2- , either by laboratory experiments (e.g. BERNER, 1967) or by theoretical calculations (SILLI~Nand MARTELL,1964). To calculate the/(so for the reaction above, the value of the second dissociation constant for H2S is needed. This constant is very small and has not yet been determined accurately (MORSEet al., 1987; SCHOONENand BARNES, 1988, 1992). The value of/(so for FeS is, thus, uncertain. At the pH of most natural sea waters (7-8) the concentration of S2- is negligible compared to the concentration of H2S and H S - and can thus be disregarded. By studying the pH dependence of metal sulphide dissolution, POnE (1954) deduced that the rate-determining step involves the H S - ion rather than S2 . This agrees with ELLIS and GIGGENBACH (1971) who assert that a reaction such as H + + FeS ~ Fe 2+ + H S - is a better explanation of the equilibrium. The apparent solubility product involves the sum of the concentration of the free and complex species (GOLDHABERand KAPEAN, 1975). However, this precludes comparison between waters of different ionic strength. Therefore, DAVISON (1980) introduced a new operational equilibrium constant to overcome some of these comparison problems, which accounts for differences in ionic strength and ionic constituency.
CALCULATION OF THE EQUILIBRIUM C O N S T A N T S OF A M O R P H O U S I R O N S U L P H I D E
Influences of complexation in calculation of the equilibrium constants of amorphous iron sulphide The subscript zero in the thermodynamic equilibrium constant pKs0 denotes that the equilibrium of the solid ML (M = metal; L = ligand) only considers the uncomplexed species M and L (SILE~N and MARTELL, 1964). Thus, Ks0 for amorphous FeS is /(so = YFe2+[Fe2+]ys2-[S2-],
(1)
where y denotes the single ion activity coefficient and [] the concentration. The total dissolved concentration of sulphide can be expressed as [H2S]tot = [H2S] + [HS-] + [Sz-] + polysulphides + thio complexes.
(2)
381
Calculations of polysulphide concentrations in the experimental solution were made based on data from BOt~LECUE and MICHARO(1978). At pH 7 the polysulphide concentration can reach a maximum of 1% of the total H2S(aq) concentration and was thus disregarded. The complexation between sulphur and heavy metals to form thio complexes is also negligible at the experimental conditions (S~UMM and BRAUNER, 1975; DAVISON, 1980). Besides free Fe 2+ ions, the complex's F e O H ÷, FeCI- and FeSO4° may affect the total dissolved concentration of Fe 2÷ (DAvISON, 1979). Thus, [Fe2+]tot--- [Fe 2+] + [FeOH +]
+ [FeCl-] + [VeSO°].
(3)
We have used selected stability constant data for iron(II) complexation to calculate the iron(II) speciation. The maximum percentage complexation under experimental conditions was around 9% (calculated by using a modified version of the equilibrium program H A L T A F A L L (INGR! et al., 1977)) of the total dissolved iron content where FeSO °, FeC1- and F e O H + constitutes 4, 4 and 1%, respectively. However, since there is a general uncertainty regarding these complexation constants (DAvISON, 1979) we assume that the influence from complexation is negligible compared with the uncertainty in the colorimetric measurements of Fe(II) (HEANEV and DAVISON, 1977). In addition, a test of the analytical method with and without C1- and SO]- ions at experimental concentrations showed that only free Fe z+ ions reacted with the iron reagent. Hence, we have assumed that Fe 2+ = Fetot.
Calculation of pK s of amorphous iron sulphide The second dissociation constant, K2, is difficult to determine. ELLIS and GIGGENaACH(1971) suggested therefore that the solubility product of iron sulphide should be defined according to the following reaction: FeS + H + = H S - + Fe 2+.
(4)
The apparent solubility product, K~, is then K ~ - [Fe2+]t°t[HS-]t°t (H+) ,
(5)
and the first apparent dissociation constant of hydrogen sulphide is (GOLDHABERand KAPEAN, 1975): ,
K x-
(H+)*[HS -] [HzS(aq)]
(6)
in which (H+) * is an empirical quantity defined by pH = - l o g (H+) *, where pH is measured on the National Institute for Standards and Technology pH
382
L . E . Bfigander and R. Carman Table 2. Measured and calculated parameters for the experimental systems. The values are given for the opaque 1, 2 and transparent systems for each day, respectively
[Fe2+]tot
[H2S]tot
Day
Temp. (°C)
pH
Q~M)
(~M)
pKs0
pK~
7
12.8
8
15.0
9
14.3
11
18.1
12
16.8
13
15.8
14
16.7
15
18.0
16
18.3
22
18.8
7.01 6.99 7.01 7.01 6.99 7.01 7.00 6.98 7.01 7.00 6.98 7.00 7.00 6.98 7.00 7.00 6.97 7.00 6.99 6.97 7.00 7.00 6.98 7.00 7.00 6.98 7.00 7.04 7.01 7.01
5.55 5.01 4.66 4.30 3.58 3.94 3.94 3.58 3.58 1.79 1.79 1.79 1.79 1.97 2.15 1.43 1.43 1.43 1.07 1.25 1.25 1.07 1.07 1.07 0.716 0.895 1.43 0.716 0.716 0.716
77.4 38.7 50.6 140.0 140.0 89.4 167.0 176.0 149.0 363.0 432.0 372.0 524.0 429.0 444.0 545.0 448.0 506.0 575.0 539.0 569.0 685.0 748.0 563.0 795.0 727.0 563.0 1330.0 1310.0 903.0
17.53 17.91 17.89 17.30 17.41 17.55 17.30 17.35 17.38 17.15 17.11 17.14 17.04 17.12 17.04 17.17 17.26 17.20 17.25 17.24 17.17 17.10 17.10 17.20 17.20 17.17 17.05 16.90 16.95 17.11
2.76 3.14 3.02 2.60 2,71 2,83 2.58 2.63 2.66 2.55 2.51 2.54 2.40 2.48 2.39 2.49 2.59 2.52 2.60 2.59 2.52 2.50 2.49 2.58 2.61 2.58 2.46 2.32 2.39 2.53
scale (BATES, 1964). The total dissolved concentration of sulphide at the experimental p H can be expressed as [H2Sltot = [H2S ] + [ H S - ] .
p K ~ Bs : K~ YFe:+YHS •
(8)
This equilibrium constant is corrected for differences in ionic strength. H o w e v e r , it is not entirely independent of ionic strength since the effect of the liquid junction potential of p H measurements has not been accounted for.
[HeS]tot
(H+) ,
- - + 1
The expression below is from DAVlSON (1980):
(7)
Rearranging (6) and (7) gives
[HS-I =
Calculation o f p K NBs of amorphous iron sulphide
(11)
Ki
RESULTS AND DISCUSSION The negative logarithm of (8) combined with (5) gives pK~ = - l o - [Fez+ ]t°--t[H2S]t°t
(9)
g (H+)* + (H+) * K~ K[ was calculated for the experimental temperatures using the equation pK~ = 2.527 - 0.169C11/3 + 1359.96/T
(10)
from GOLDHABER and KAPLAN (1975), where C1 stands for chlorinity and T for temperature in Kelvin.
A t the beginning of the experiment the sediment was oxidized to a depth of 2-3 m m (Eh > 200 mV; MORXIMER, 1942). The variations of the experimental temperature are shown in Fig. la. During the first days there was a rapid p H decrease, around 1 p H unit, in all three boxes. As a result of oxidation of organic matter CO2(g ) produced in the box cannot leave the closed system. Subsequent to that, p H stabilizes around 7 (Table 2) in all boxes, indicating a sulphate-sulphide system control of p H (BENYAAKOV, 1973). Figure 1 A - C shows the variations in temperature and concentrations of oxygen, Fe(II) and H2S. Dissolved oxygen in the water was con-
Solubility product of amorphous iron sulphide sumed within 5 d of the beginning of each experiment in all three boxes (Fig. 1A). Redox potential in the three experimental systems was measured occasionally during the experiment and used only as an environmental parameter (BAGANDER and NIEMISTO, 1978). In nalural waters with high dissolved oxygen content and near-neutral or higher pH, the concentration of soluble iron(Ill) is very low. Most of the soluble and suspended iron constituents occur as solid iron(III) oxyhydroxides. The increase in total dissolved Fe(II) during the first 4 days was, hence, due to reduction of the Fe(III) constituents in the water phase. At oxygen concentrations below the detection limit (<3/~M) dissolution of solid iron(Ill) phases released dissolved iron(II) from the sediment to the overlaying water. Between days 6 and 22, Millipore filters (0.45 ktm) had a residue of black precipitated iron sulphide, which was X-ray amorphous. Hydrogen sulphide could be detected by odour from day 6 and detected analytically from day 7. The decrease in H2S in the transparent box from day 15 (Fig. 1C) was due to oxidation by photosynthetic sulphur bacteria. The experiment continued after day 22 but Fe(II) was then below the detection limit (0.36/~M). Different solubility products of amorphous FeS were indirectly calculated from solution chemistry between day 7 [Fe(II) decrease] and 22 [limit of Fe(II) detection] (Fig. 1B). At the experimental pH, around 7 (Table 2), the HCO~ concentration was almost equal to the total carbonate concentration (deduced from STUMMand MORGAN, 1981). The concentration of NH2, H2PO4, H P O ] - and H S - (maximum 0.2 x 10 -3 M measured in earlier experiment) was too small to contribute to the overall ionic strength. The ionic strength, 0.128, of uncomplexed media, was calculated from the values in Table 1. This value is too high since the formation of ion-pairs has not been considered. To determine the maximum influence of ion-pairing we have used a modified computer version of HALTAFALL (INGRI et al., 1977) which includes a large thermodynamic database. The program uses the Davies" equation to consider the ionic strength and calculate the activity coefficients (log y). The results from these calculations are expressed in Table 3, which give the ionic strength 0.122. The activity coefficients of Fe e+ and HS- were determined by Davies' equation (e.g. Sxurar~ and MORGAN, 1981). The activity coefficient of H2S(aq) differs only slightly from unity (1.03) for sea water (DOUABULand RILEY, 1979). We have thus calculated pK~0 of amorphous FeS with YH.~S,,,,,~equal to 1.
383
Table 3. Maximum ion-pairing concentration Ion compound
% of total concentration
CI Na + Mg2+ MgSO° Ca2+ CaSO~ K+ SO] HCO~
100 99 87 8 90 9 99 76 89
Concentration (M) 1.10 x 6.90 X 9.60 x 4.60 x 1.44 x 1.40 x 1.60 x 4.30 x 1.42 x
10-I 10 -2
10 3 10 4 10 3 I0 4 10-3 10 3 10-.3
3.2-
3.0-
2.8 pK"s 2.6
:4-
2.4
2.2
'
12.0
'
I
'
14.0 16.0 Temperature (°C)
I
180
'
I
20.0
FIG. 2. Temperature dependence of the apparent equilibrium constant of amorphous iron sulphide (r = 0.75. n = 30).
pK; = -0.068t + 3.70.
(12)
with r = 0.75 (n = 30) at 1% level of significance (t denotes the temperature in celsius). Table 4 shows measured apparent solubility products from different environments together with the equilibrium constant defined by DAVISON(1980). The apparent and pK~ Bs equilibrium constant of the present study is given for the average experimental temperature (15.8°C) together with equilibrium constants from other investigations, p K ~ Bs values above 4 are probably related to greigite or mackinawite equilibria, according to data from BERNER (1967). The other values agree rather well but differences in analytical methods can partly explain some discrepancies as well as different stoichiometric compositions depending on the ionic media.
Apparent solubility constant, pK~, and Davison's equilibrium constant, p K ~ Bs
Solubility diagram
Linear regression analysis on the apparent equilibrium constants (Fig. 2) gave the equation:
The concentration HS- in the water was calculated from analytical values of [H2S]tot using Eqns (8) and
AG 9: 4-C,
L. E. B~gander and R. Carman
384
Table 4. pK~ and pKNpBs solubility products from different investigations Site
pK~
pK NBs
Lake Nitinat Cariaco Trench Skejennungen Laboratory Laboratory Baltic Sea Black Sea Byfjorden Lake Greifensee Clear Lake Esthwaite Abereiddy Kiel Bight Framvaren Baltic Sea
3.79 3.34 1.95
4.43 4.04 2.06 2.97 3.68 4.34 3.53 3.74 2.99 3.23 2.59 2.55 2.47 3.67 3.15
3.15 3.7 2.85 3.1 2.8 3.12 2.48 1.94 1.8 3.0 2.6
pH
7.1 0.6 7.8 0.7 6.8 0.002 3.7 ---)0 4.0 0.1 -0.24 7.7 0.45 -0.62 7.05 0.007 6.9 0.002 7.13 0.002 7.17 0.7 7.6 0.42 7.0 0.40 7.0 0.12
(10). Values of p[HS-] were then plotted against p[FeZ+]tot (Fig. 3). Linear regression analysis gave the equation: p[Fe2+]to, = - 0 . 6 7 p [ H S - ] + 8.44.
(13)
The values determined prior to the [Fez+ ]tot decrease in the box water are excluded. The correlation coefficient is 0.95 and significant at the 1% level (n = 30). The abscissa (Fig. 3) also could be looked upon as a time-axis since [HS-] increased with time during the experiment. Ideally, at constant temperature, amorphous FeS would precipitate when the equilibrium solubility product is reached as [HS-] increases along a p[Fe 2+ ]tot line. The concentrations of [FeZ+]tot and [HS-] would then follow a line with the slope - 1, if the mole fraction of amorphous FeS is 1. Thus, Eqn (13) indicates precipitation of a solid amorphous iron sulphide phase deficient in sulphur i.e. FeS0.67. However, a "true" equilibrium was probably not reached immediately in this in situ experiment due to a number of "uncontrolled" factors. From the data presented in Fig. 3 it appears that there
5.2
5.6 p F e 2+ 5.8
:
~, ~.
6.0-
6.2
I 5.2
4.8
'
I 4.4
I
I
4.0
3.6
pHS -
'
I 3.2
3. Solubility diagram of amorphous iron sulphide (r = 0.95, n = 30). Days 11-22 = pHS- values below 4.
FIG.
[I]
t (°C) 10 17 4 25 25 5 9 5 5 10 10 10 14 9 16
Reference RICHARDSet al. (1965) RICHARDS(1965) KJENSMO(1967) BERNER(1967) DOYLE(1968) KOROLEFF(1968) Brewer (1971), in DAVISON(1980) DANIELSSONet al. (1975) EMERSON(1976) LAHANN(1977) DAV1SONand HEANEY(1978) DAV1SONand HEANEY(1980) BALZER(1982) SKEI(1988) This study
are two populations of data, with two slopes. The first population of data arise from measurements from the day 7, 8 and 9. The slope of these data is -0.21 and, hence, it is not unreasonable to say that the Fe2* concentration is nearly independent of [HS-] at these low sulphide concentrations. Examining later experimental values (from day 11, Fig. 3) in the same way, linear regression analysis using all data gave the equation: p[Fe2+]tot = - 0 . 8 6 p [ H S ] + 9.18
(14)
with r = 0.90 and a significance level of 1% (n = 21). In this case, precipitation of FeS0.86 is suggested. This is close to the range of composition of amorphous FeS0.87q).92 reported by SWEENEY and KAPLAN (1973). A closer examination of the data from days 11-22 (Fig. 3) suggests that there are two data points near the origin of the figure (from the two opaque boxes at day 22, see Table 2) which could be looked upon as outliers and hence be excluded from the linear regression analysis. Without these two points regression analysis yielded a different interpretation giving an opposite picture with respect to deficiency in sulphur vs iron in the assumed precipitate i.e. the slope is -1.14. It is, however, difficult from a scientific point of view, to explain why these data should be excluded. One could instead look upon some other data points from the later part of the experiment that could be subject to exclusion, i.e. the sulphide data from the transparent box, days 15, 16 and 22. These data differ conspicuously from the data obtained from the two opaque boxes (Table 2). A probable explanation to this could be that photoautotophic sulphur bacteria ( C h r o m a t i u m and/or Chlorobium) oxidize some parts of the produced sulphide. Linear regression without these three points gave a slope close to that with all data points included (-0.83). Hence, under natural conditions one can expect to find variations in the composition of amorphous FeS due to environmental factors. For example, BERNER (1964) and RICKARD (1969) reported higher sulphur content of FeS1.1 which they
Solubility product of amorphous iron sulphide attributed to be adsorption of hydrogen sulphide. HALLBERG (1970) reported a FeS ratio close to 1 and suggested (HALLBERG, 1972) that the ratio of 1:1 may be explained by the co-existence of greigite. All these ratios were measured from precipitate compositions whereas our calculated values are based solely on solution chemistry. The divergence from the ideal composition of FeS (1:1) could be explained by errors in the analytical techniques, e.g. overestimation of iron(II) concentrations due to particulate iron species <0.45 ~ m and/or loss of sulphide due to an oxidation. A n o t h e r explanation could be that the measured deficiency of sulphur, based on solution chemistry, is charge balanced by some other anion. Such charge balances have been observed for the hydrolysis of another c o m m o n amorphous precipitate during early diagenesis, i.e. amorphous iron(III) hydroxide (e.g. Fox, 1988; CORNELL et al., 1989). These anions provide charge balance but do not seem to affect the solubility of the precipitate (Music et al., 1982; Fox, 1988). In an applied sense the knowledge of the equilibrium constant for the least stable iron sulphide is important to know because of its broad influence on the organic and inorganic chemistry in all aquatic environments with permanent or temporary anoxic conditions. It regulates to a great extent the amount of upward migration of both nutrients as well as trace and heavy metals and, thus, determines the environmental biogeochemical conditions of the examined aquatic environment. Acknowledgements--We thank Professor R. O. Hallberg
for valuable criticism and discussions during the preparation of the manuscript. Editorial handling: John Catts.
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