Envlron. Sci. Technol, 1992, 26, 2408-2413
Association, Anaheim, CA, 1989; Paper 89-139.1. (62) Rasmussen, R. A. SCAQS hydrocarbon collection and analyses (Part I). Report to the California Air Resources Board under Contract A6-179-32, Biospherics Research Corp., Hillsboro, OR, 1990. (63) Finlayson-Pitts, B. J.; Pitts, J. N. Atmospheric
chemistry-fundamentals and experimental techniques, 1st ed.;John Wiley and Sons: New York, 1986;pp 313-316.
(64) Gorse, R. A. Environ. Sci. Technol. 1984, 18, 500-507. (65) Wesely, M. L. Atmos. Environ. 1989, 23, 1293-1304.
Received for review April 21,1992. Revised manuscript received July 28, 1992. Accepted August 3, 1992. This work was supported by the Electric Power Research Institute under Agreement RP3189-3.
Kinetics and Mechanism of the Reaction of H2S with Lepidocrocite Stefan Pelffer,*,+ Maria dos Santos Afonso,' Bernhard Wehrli,+ and Rene Gachtert
Lake Research Laboratory of EAWAGIETH, CH-6047 Kastanienbaum, Depto de Qdmica Inorgdnica, Anantica y Qdmica Flsica, Facultad de Ciencias Exactas y Naturales, Ciudad Universitarla Pabelldn 11, 1428 Buenos Aires, Argentina The initial reaction between hydrogen sulfide and the surface of lepidocrocite was studied in the pH range between 4 and 8.6 by monitoring the change of the emf of a pH2S sensor. The rate of H2S oxidation is pseudo first order with respect to H2S and shows a strong pH dependence with a maximum at pH 7. Two rate laws were derived: R 5FeS- and >FeHS. The overall rate law for the dissolution was derived as
R = k{>FeS-l + kl>FeHS)
(2)
{>FeS-)and {>FeHSjare surface concentrations (mol rn3.
0013-936X/92/0926-2408$03.00/0
0 1992 American Chemlcal Society
pH 5.12
pH 6.03
h
EE -7 0
20
40
80
60
100
120
0
0
a
ls
10
t in min
tinmin
pH 6.94
pH 7.42
4
0
6
0
10
a
4
0
e
10
re
tinmin
t in min
Figure 1. Negative logarithm of number of moles of HzS, n(HzS), consumed due to reactlon wth lepklocrocte at different pH values.
It is the objective of this study to investigate the reaction of hydrogen sulfide itself with ferric oxide surfaces in the pH range of natural waters (pH 4-8.5) by monitoring the emf of a H2S-sensitive electrode cell. Lepidocrocite (yFeOOH) was used as a model compound since it is formed during reoxidation of ferrous iron under natural conditions (20) and is regarded to be reactive in sediments (21). Materials and M e t h o d s
Kinetic batch experiments with synthetic lepidocrocite and dissolved sulfide were performed in the pH range 4 < pH < 8.7 at a constant ionic strength (I= 0.1 M KC1) and at room temperature (26 f 1"C). Dissolved H2Swas measured in situ using a pH,S-electrode cell [Ag0,AgI/0.2 M HI/glass/solution/Ag2S,A$ with Ago and AgI in 0.2 M HI as the internal reference system, Ingold, Model Ag275-85-6329 (2291. The emf was continuously recorded using an A/D data aquisition system. Synthetic lepidocrocite was prepared by following the method of Gupta (23). Surface area was determined as 57.26 m2 g-' (BET method). The samples produced electron and X-ray diffraction patterns corresponding to lepidocrocite. All reactions were conducted in a 200-mL glass vessel, into which various electrodes could be placed by gas-tight screw plugs. The solutions were stirred with a Tefloncoated stirring bar at constant rate. The reaction solutions were prepared by deaerating 200 mL of 0.1 M KC1 with N2 and purified by bubbling through alkaline pyrogallol solution (40 g of pyrogallol and 60 g of KOH into 200 mL of deionized water) and adding appropriate amounts of reagent grade HCl and Na2S in order to establish a prechosen pH. The Na2S stock solution was iodometrically checked before each run. All chemicals were reagent grade. Recording of the emf of both, pH and pH2S electrode cell, was started just before an aliquot of deaerated 7-FeOOH stock solution wm added with a syringe. When the reaction with sulfide proceeded, protons were consumed and the subsequent increase of pH
had to be counterbalanced by using a combined Metrohm Impulsomat 614 and Dosimat 665 system, which allows pH stabilization on a chosen set point. By use of this pH-stat system, interferences of y-FeOOH with pH buffers were avoided. The volume of HC1 added to the solution during the reaction was recorded likewise. The change of ionic strength due to the addition is neglectable. The reaction vessel was washed with 2 M HC1 before each run in order to clean adsorbed y-FeOOH from the glass surface. In some experimental runs samples were taken for iron determination. A 1-mLsample was filtered through a membrane filter (0.2 pm) into 3 mL of dilute HNOB(Suprapur), using a plastic syringe, and analyzed by atomic absorption spectroscopy. We assume the detected dissolved iron to consist mainly of Fe(I1) species. Quantification of Total Dissolved Sulfide. The slope of the pH,S electrode cell was determined by an external calibration procedure (24). In contrast, the emf 8 measured in experimental solutions were related to the initial known HS activity of each experiment. This was done in order to prevent differences in response between standard calibration solutions and the reaction solutions. Thus, each pH2S value at a certain time step t was calculated as PH2St = -1% ([H2Sltot/t=oao)- (emf,=, - emft)/&
(3)
where a, is the protolysis coefficient of H2Sand S ,, is the slope of the pH2S electrode cell determined by external calibration (mV). Results
Reaction Pattern and Reaction Order. In Figure 1, the logarithm of the number of moles of HzS, log ( n (H2S,3), is plotted vs time for four different experimental pH values, but constant surface area concentration of ferric oxide (10 m2 L-l) and the same initial amount of dissolved sulfide M). The patterns are representative for all experiments performed in this study. At the lower and Environ. Sci. Technol., Vol. 26, No. 12, 1992 2409
PH 4
t
I
a
200
400
600
m~im~zao14oo
t in sec Flgure 4. Formation of ferrous iron at pH 5.7.
in the alkaline region. Between pH 5 and 8.6 an empirical rate law can be derived as follows: 5FeS-, >Fe"S, and >Fe"OH2+ to be in a steady state. We assume an innersphere coordination of sulfide with Fe(II1) followed by ~ ( p , ,to ) a(d,,) electron transfer (29),which is expected to be faster than an outer-sphere process. The So- radical is assumed to react in fast steps with the ferric oxide surface
C
Table I. Proposal for the Reaction Sequence of HS-with a Ferric Oxide Surface According to ref 1
.-0
E
-6
C
8 C 00
reversible adsorption of HS-
-7
a,
w
ki
>FeOH
Lt 3
+ HS- e >FeS- + H2O k-1
-9
v)
PH
reversible electron transfer
reversible release of the oxidized product
2? -81 -IO ; < 2
H+ ks
new surface site
4
-., 6
8
IO 10
PH
detachment of Fe2+
>Fen0H2+
.--
+ Fe2+
(IV)
to higher oxidized sulfur compounds. In Figure 5a we plotted the concentration of sulfide surface species as a function of pH under the experimental conditions given above. The surface complex formation constants for both species >FeS- and >FeSH (>FeS-)/(>FeOH)[HS-] = KF~s-= 105.3 (7) (>FeHS)/(>FeOH)[HS-I[H30+] = KFeHS = 1010.82 (8) and the acidity constant for the deprotonation of >FeSH (>FeS-)[H,O+]/(>FeHS)= K , = 10-5.5
(9)
were taken from ref 1 and the acidity constants of the mol L-l, K,2 = lepidocrocite surface (Kal = mol L-I) from ref 23. A t least phenomenologically, the pH maximum can be explained according to the pH dependency of the surface speciation of >FeSH. Application of the reaction scheme presented by ref 1 seems to be reasonable, if one takes into account that the relative contribution of each surface species to the reaction rate is weighed by the individual rate constants k and k' (cf. 1) and that the adsorption constants of HS-to the ferric oxide surface will be different for hematite and lepidocrocite. We therefore conclude that the reaction rate of H2S with the lepidocrocite surface is also proportional to the surface species >FeS- and >FeHS. According to Table I, the consumption rate of HS- in terms of the surface species >FeS is derived as ke,k2{>FeS-) d[HS-] --(10) dt k2k3 + ke,(k3 + k-&[S*-] A similar derivation can be done for the species >FeHS. Assuming steady-state concentration for S'-, proportionality between the rate and the surface concentration of >FeS- and >FeSH is obtained:
-d[HS-]/dt = k{>FeS-)+ k'(>FeSH)
(11)
In comparison to the pH dependency of the change in surface speciation, an even better fit of the experimental
Flgwe 5. (a, top) Concentration of surface species >FeS and >FeHS as a function of pH, calculated ustng the adsorption constants determined by dos Santos Afonso and Stumm ( 7 ) for hematite. (b, bottom) Logarbm of the mole fractions of the pHdependent reactants (aHs-, @Fa, and )+ , @ , to form a surface species >FeHS: upper curve, reactants are HS- and FeOH,'; lower curve, reactants are HS-, >FeOH, and H,O+; conditions for both plots, Stat = 1.6 X IO-' mol M, I = 0.1 M. m-2, [HPSIW=
data is obtained if one plots the product of the reactants vs pH. Combining eq 11 with eqs 7 and 8 yields
(124 Alternatively, HS- was allowed to react with the protonated ferric oxide surface species FeOH2+,which gives
(12b) Combining eqs 12a and 12b with eq 4, one obtains (letting
k FeOFe+ (reaction XI in Table 11). A combination of the various processes (formation of polysulfides and/or elemental sulfur, precipitation of FeS, and adsorption of Fez+)leads to the low AH+/AH2Sratios observed at pH >6. At lower pH values FeS does not form and adsorption of Fe2+decreases, providing an increased AH+/AHzS ratio.
Table 11. Ratio AHt/AHzStot for Different Products of the Reaction of H2S with yFeOOHn
(V)
AHt/ AHZStO, 3.5
(VI)
3.8
(VII) (VIII)
5.0 8.0
reaction
-
6FeOOH + 4HS- + 2H20 S42- + 6FeZt + 140H8FeOOH + 5HS- + 3Hz0 S2-+ $Fez+ + 190H2FeOOH + HS- + hzO So + 2Fe2++ 50H8FeOOH + 2HS- + 3H20 S2032-+ 8FeZt + 160H8FeOOH + HS- + 3Hz0 SO4" + 8Fe2++ 150H-
--
-
Formation of FeS 6FeOOH + 10HS- 6FeS + S42- + 80H- + 4HzO 6FeOOH + 6>FeOH + 4 HS- 6>FeO-Fet + S2-+ 80H- + 4H20
-
(IX)
15.0
(X)
0.8
(XI)
2.0
'While reactions V-IX release products into solution, FeS is formed in reactions X and XI.
Conclusion
X
x
x
x
X
0 ~
'
4
"
~
'
5
~
!
" 6
'
"
'
7
~
~
~ 8
'
"
~ Q
"
'
PH Figure 6. Ratio of the consumed protons per mole of total sulfide consumed as a function of pH.
for the reaction of H2S with lepidocrocite. In summary, both the pH dependency of the surface speciation of the two >FeS complexes and the pH dependency of the product of surface-bound and dissolved reactant strongly support the applicability of the model proposed in ref 1 to our data. Formation of Products: Stoichiometry of the Reaction. Since we did not measure the oxidation products of H2S, we estimated them by forming the ratio of consumed protons per mole of total sulfide consumed. In Table 11, the reaction of H2S with a ferric oxide is formulated for different oxidation products of HS-. Due to the wide span of ratios ranging from 1to 15 it should be possible to clearly differentiate between these products. Figure 6 shows that the measured ratios range between 0.5 and 3.5 and reveals a distinct pH dependence. Sulfate does not appear to be a major product in these initial rate experiments. This is in contrast to the findings of dos S a n d Afonso and Stumm (I),who measured mainly sulfate, thiosulfate, and traces of sulfite as oxidation products of H2S. Since our study reflects more the experimental conditions as met in Pyzik and Sommer's study (19),we may also assume elemental sulfur or polysulfides (S42-and SS2-)to be the main products, which may be further oxidized to sulfate on a longer time scale. At pH values of >6.5, a black color appeared during the experiments, indicating the formation of FeS. The lag phase in the H2S consumption, which can be observed within the first minute at higher pH values (cf. pH 6.94 and at pH 7.42 in Figure l), may be explained by this phenomenon: The reaction rate is increased by the number of Fez+ions to react with surplus HS- to FeS (e.g., 6 Fe2+in the case of S42- as a product, reaction X in Table 2412
Envlron. Sci. Technol., Vol. 26, No. 12, 1992
~
The reaction of hydrogen sulfide with y-FeOOH is a fast surface-controlled process. There is strong evidence that the reaction mechanism proposed for hematite (I) can also be applied to our data. This model perfectly explains the observed pH dependency of the reaction rate, if only the surface species >FeSH is considered to be relevant for the electron-transferprocess. Surface spectroscopictechniques will be needed to further elucidate the fate of oxidized sulfur compounds and ferrous iron at the ferric oxide "surface at pH values higher than 6.5. The observed maximum of the reaction rate at pH 7 corresponds to a pH to which anoxic, sulfide-containing sediment pore waters are usually buffered (30). In addition, the formation of FeS is favored under these conditions. Polysulfides may be expected to be at least an intermediate product of the reaction (19). Therefore the requirements for the formation of pyrite from FeS and polysulfides under natural conditions would be fulfilled (31-33). Upon further oxidation of the intermediate sulfur compounds to sulfate with ferric oxides, the anaerobic oxidation of sulfide may also play an essential role in the cycling of sulfur in sediments (34). The conclusion is therefore allowed that abiotic oxidation of H2S by (reactive) ferric oxides will play a much more important role in sediment diagenesis than believed so far. Acknowledgments
We thank A. Tessier and three anonymous reviewers for their helpful and constructive comments and criticisms. Registry No. H2S, 7783-06-4; lepidocrocite, 12022-37-6.
Literature Cited Dos Santos Afonso, M.; Stumm, W. Langmuir, in press. Einsele, W. Arch. Hydrobiol. 1936, 29, 664. Mortimer, C. H. J . Ecol. 1941,29, 280. Sholkovitz, E. R.; Copland, D. Geochim. Cosmochim. Acta 1982, 46, 393. Johnson, A. C. Geochim. Cosmochim. Acta 1986,50,2433. Tessier, A.; Carignan, R.; Dubreuil, B.; Rapin, F. Geochim. Cosmochim. Acta 1989,53, 1511. Belzile, N.; Tessier, A. Geochim. Cosmochim. Acta 1990, 54, 103. Giblin, A. E.; Likens, G. E.; White, D.; Howarth, R. W. Limnol. Oceanogr. 1990,35,852. Nealson, K. H. In Microbial Geochemistry; Krumbein, W. E., Ed.; Blackwell: Oxford, England, 1983; Chapter 4. Schindler, D. W. In Chemical Processes in Lakes; Stumm, W., Ed.; John Wiley & Sons: New York, 1985; Chapter 11.
Environ. Sci. Technol. 1992, 26, 2413-2420
Lovley, D. R.; Phillips, E. J. P.; Lonergan, D. J. Environ. Sci. Technol. 1991,25, 1062. De Vitre, R. R.; Buffle, J.; Perret, D.; Baudat, R. Geochim. Cosmochim. Acta 1988,52,1601. Wieland, E.;Wehrli, B.; Stumm, W. Geochim. Cosmochim. Acta 1988,52,1969. Stumm, W.; Wieland, E. In Aquatic Chemical KineticsReaction Rates of Processes in Natural Waters; Stumm, W., Ed.; John Wdey & Sons: New York, 1990;Chapter 13. Zinder, B.; Furrer, G.; Stumm, W. Geochim. Cosmochim. Acta 1986,50,1861. Dos Santos Afonso, M.; Morando, P. J.; Blesa, M. A.; Banwart, S.; Stumm, W. J . Colloid Interface Sci. 1990,138, 74. Sulzberger, B.; Suter, D.; Siffert,C.; Banwart, S.; Stumm, W. Mar. Chem. 1989,28,127. Rickard, T. Am. J . Sci. 1974,274,941. Pyzik, A. J.; Sommer, S. E. Geochim. Cosmochim. Acta 1981,45, 687. Schwertmann,U.; Taylor, R. M. In Minerals in the Soil Environment, 2nd ed.;Dixon, J. B., Ed.; Soil Science Society of America: Madison, WI, 1989;Chapter 8. Canfield, D. E. Geochim. Cosmochim. Acta 1989,53,619. Frevert, T.; Galster, H. Schweiz. 2.Hydrol. 1978,40,199. Gupta, S.K.Dissertation, Universitiit Bern, 1976. Peters, K.; Huber, G.; Netsch, S.; Frevert, T. GWT, GasWasserfach: WasserlAbwasser 1984,125,386.
(25) Hering, J. G.; Stumm, W. In Mineral Water-Interface Geochemistry;Reviews in Mineralogy 23;Hochella, M. F.,
White, A. F., Eds.; Mineralogical Society of America: Washington, DC, 1990; Chapter 11. (26) Baumgartner,E.; Blesa, M. A.; Maroto, A. J. G. J . Chem. Soc., Dalton Trans. 1982,9,1649. (27) Cornell, R.M.; Schindler,P. W. Clays Clay Miner. 1987, 33,347. (28) LaKind, J. S.;Stone,A. T. Geochim. Cosmochim. Acta 1989, 53,961. (29) Luther, G. W., III In Aquatic Chemical Kinetics-Reaction Rates of Processes in Natural Waters; Stumm, W., Ed.; John Wiley & Sons: New York, 1990;Chapter 6. (30) Ben-Yaakov, S.Limnol. Oceanogr. 1973,18,86. (31) Berner, R. A. Am. J . Sci. 1970,268, 1. (32) Schoonen, M. A. A.; Barnes, H. L. Geochim. Cosmochim. Acta 1991,55,1495. (33) Luther, G. W., I11 Geochim. Cosmochim. Acta 1991,55, 2839. (34) Urban, N. In Environmental Chemistry of Lakes and Reservoirs; Baker, L. A., Ed.; ACS Advances in Chemistry
Series; American Chemical Society: Washington, DC, in press. Received for review February 12, 1992. Revised manuscript received July 2,1992. Accepted July 13,1992.
A Polar High Molecular Mass Constituent of Bleached Kraft Mill Effluent Toxic to Marine Organisms
Is
Richard M. Higashi,' Gary N. Cherr, Jonathan M. Shenker,+ Jeffrey M. Macdonald,* and Donald G. Crosbyg
Bodega Marine Laboratory, University of California-Davis,
Box 247, Bodega Bay, California 94923
A high molecular mass constituent (HMM) of whole bleached kraft mill effluent (BKME),which represents the majority of toxicity to early life stages of marine animals and a plant, has been isolated and partially characterized. BKME was subjected to fractionation coupled with toxicity testing to determine the toxicity of each isolated fraction. The toxic mode of action was also tracked throughout the fractionation using echinoderm sperm motility as an indicator. While most fractions inhibited sperm motility, BKME and HMM did not. Yet, HMM exhibited most of the toxicity of BKME to echinoderm sperm, mollusc embryos, larval sole, and kelp gametophytes. HMM was soluble only in water and appeared to be free of the resin and fatty acids or chlorinated aromatic compounds that are implicated in freshwater acute toxicity of BKME to salmonid fish. Structural analyses indicate that this polar, high molecular mass constituent was devoid of aromatic structure and had other characteristics indicative of lignin breakdown products.
Introduction Bleached kraft mill effluent (BKME) is the combined aqueous waste of a major chemical pulp-making process. A mill typically generates waste water from many of its sections, including pulping, chemical recovery, evaporation and condensation, and multistage bleaching operations. Present address: Department of Biological Sciences, Florida Institute of Technology, Melbourne, FL 32901. Present address: Department of Pharmaceutical Chemistry, University of California, San Francisco, CA 94143. *Department of Environmental Toxicology, University of California, Davis, CA 95616. 0013-936X/92/0926-2413$03.0010
Thus, waste waters from acid, alkaline, chlorine oxidant, and other chemically diverse processes are sewered together to form the whole-mill discharge. The mixing and ensuing reactions of these streams lead to a final effluent that is highly complex. It consists of simple inorganic salts as well as over 250 identified organic and inorganic compounds of low molecular weight (or mass) ( I ) , with probably many more yet to be identified. In addition, an unusual property of BKME is that many of the organic constituents are of high molecular mass (>1 m a ) (2). This material is thought to consist largely of the polar breakdown product(s) of lignin, with lesser amounts of lignin at various degradation stages as well as polysaccharides (2).
The acute toxicity of BKME to a variety of aquatic organisms has been well-documented (3). Studies of sublethal effects of whole effluent range from biochemical and mutagenic to behavioral aspects (3). Partly to obtain more detail for pollution abatement, toxicity studies have been performed on selected mill process streams (cf. ref 2). However, the relevance of these studies to the final effluent toxicity is not clear ( 4 4 3 , probably due to the complex changes that take place upon mixing to form BKME. Therefore, it is essential to first identify and characterize the constituents responsible for toxicity in BKME itaelf, in order to assess environmental hazards of BKME discharges and to develop appropriate pollution control technology. The complexity of BKME has made this identification a nontrivial task. By use of chemical fractionation coupled with toxicity testing and structure identification, the acute freshwater toxicity of nonbiologically treated BKME to salmonid fish has been associated with resin and fatty acids
0 1992 American Chemical Society
Envlron. Sci. Technol., Vol. 26,No. 12, 1992 2413
