Ind. Eng. Chem. Res. 1988,27, 187-191
187
Removal of Selenate from Water by Chemical Reduction Andrew P. Murphy US.Department of the Interior, Bureau
of Reclamation, Engineering and Research Center,
Denver, Colorado 80225
A previously unknown chemical reaction in which selenate is reduced to elemental selenium may prove t o be the basis of a chemical process to remove selenate from water systems and t o control selenate in agricultural wastewater. T h e reaction occurs between selenate and ferrous hydroxide under alkaline conditions, producing magnetic iron oxides (magnetite and maghemite) and elemental selenium. T h e reaction rate reaches a maximum a t about p H 9, dropping sharply below p H 8 or above p H 10. I n the presence of excess ferrous hydroxide, magnetite is the predominant product. As the ratio of selenate to ferrous hydroxide increases, more maghemite seems t o form. The reaction may proceed in a stepwise fashion, with selenate first reduced to selenite and then to selenium. The Gibbs free energy of the selenate reduction reaction is calculated t o be -83.1 kcal/mol, assuming magnetite is the final iron oxide produced. The elemental selenium remains trapped in the iron oxide. This chemistry may find application in the recovery and refining of selenium and provide insight into the selenium cycle in the environment. Waters containing selenium pollution generally consist of the selenate and selenite anions in small concentrations relative to the major cations and anions. Although treatment steps exist to selectivelyremove selenite, no such processes exist to selectively remove selenate. This is partly due to the fact that selenate chemically resembles sulfate. Thus, to remove trace amounts of selenate, sulfate must be removed also. Without having the selectivity,this process may not compare well with the cost and benefits of total desalination. Currently used methods have achieved only limited success in removing selenate from water systems. These methods can be divided into four classes: (1)conventional desalting (Maneval et al., 1983; Sorg and Logsdon, 1978), (2) biological processes (Kauffman et al., 1985; Gersberg et ai., 1986; Doran and Alexander, 1977), (3) adsorption (Hingston et al., 1967, 1986; Ames et al., 1979; Schwertmann and Taylor, 1977), and (4) chemical reduction schemes (Baldwin et al., 1983; Marchant, 1976; Doumas, 1968). A disadvantage of conventional desalting techniques is the lack of specificity for selenate removal (Maneval et al., 1983). Thus, the cost associated with total desalination of the water must be included. In addition, a brine stream rich in selenate must be discarded. Biological processes, although capable of reducing selenate, require days rather than minutes for the reduction, achieve only partial reduction, and may generate more highly toxic compounds such as selenomethionine (Hartmanis and Stadtman, 1982). On the basis of LD6,* (*a dose which is lethal to 50% of the animals tested) for striped bass, selenomethionine is 10000 times more toxic than selenate (Jackson, 1986). Adsorption processes normally involve the formation of insoluble complexes such as the ferric selenite complex (Chau and Riley, 1965; De Carlo et al., 1981),but they do not form insoluble selenate complexes. Therefore, selenate must first be reduced to selenite. In addition, adsorption is not specific for selenate. Chemical reduction techniques have appeared in the patent literature using iron and zinc. Metallic iron (Baldwin et al., 1983) is claimed to reduce selenate under acid conditions. Metallic zinc (Marchant, 1976) is claimed to reduce selenate too, but zinc is toxic to fish and is an environmental concern. Unfortunately, both reduction processes in the patent literature add little knowledge to the understanding of selenium chemistry.
Experimental Section Materials. The sodium selenate, sodium selenite, and
carbonyliron used in the study were purchased from Sigma Chemical Corp., St. Louis, MO. All other chemicals were ACS grade. Ferrous hydroxide was prepared from either ferrous sulfate or ferrous chloride and sodium hydroxide. Instrumentation and Instrumental Procedures. The ion chromatograph (IC) used in the study was a Dionex Model 2120i with conductivity detection. The IC procedure (Oppenheimer et al., 1985) did not use a pretreatment step with barium hydroxide. The sulfate-toselenate ratio was not great, and one peak each for sulfate and selenate was always present. Samples for selenate analysis were diluted with deionized water and analyzed immediately on the ion chromatograph. Atomic absorption spectroscopy (AAS)work was performed by using a Varian Model 975 equipped with background correction. Selenium analysis by AAS consisted of dissolving the residual iron oxide solids and selenium with 1:l HN03, gentle heating, and then diluting to volume. Tests in which the pH was held constant with 0.1 N sulfuric acid were performed on either a Fisher titrimeter or Brinkmann autotitrator. A colorimetric method (Lee and Stumm, 1960) was used for the determination of ferrous in the presence of ferric iron with bathophenanthroline. Precipitates were analyzed either by nickel-filtered Cu K a X-ray diffraction (XRD) or silver targeted X-ray fluorescence. Iron powders were submitted for scanning electron microscope (SEM) examination directly from their reagent bottles. Screening Tests. Previous work by the investigator involved testing numerous chemical systems to identify whether selenate could be removed from water, either by chemical reduction or complexation. These tests were performed in loosely capped disposable plastic containers, maintained a t 25 "C with no stirring. Each container had 200 mL of deionized water, 0.020 g of sodium selenate, 0.020 g of sodium sulfate, and 2.0 g of reductant. When iron wire was used as a reductant, large lengths consisting of approximately half a meter were twisted to a small bundle to fit neatly into the containers. No pH adjustment was made. A 250-mL, three-neck, round-bottom flask was used to examine the effect of O2 on the reaction. Gas was introduced into one neck and vented from the other and a sampling tube inserted under the solution on the middle neck. The flow rate for all gases was 20 cm3/min. Ferrous Hydroxide Tests. Tests similar to the screening tests were then made with ferrous hydroxide. Ferrous hydroxide was freshly precipitated in solution for each test using the stoichiometric amount of sodium hy-
This article not subject to U.S. Copyright. Published 1987 by the American Chemical Society
188 Ind. Eng. Chem. Res., Vol. 27, No. 1,1988
l
o
O
T
= 0
(a 1
=
-
= Na2Se0,
A = Na2S04
a
b 0
18
66
T I ME ( H r s . )
Figure 1. Effect of iron on selenate removal: (a) no iron, (b) fine iron wire, (c) electrolytic reduced iron powder, and (d) reduced pentacarbonyliron.
droxide needed to form the compound from either ferrous sulfate or ferrous chloride. The ferrous hydroxide was formed in a 250-mL, three-neck, round-bottom flask. Argon was used to exclude air while the reaction progressed. The argon was introduced in an identical fashion as discussed in the previous section. The pH was then quickly adjusted to the final value required for the test. Ratios of selenate to hydroxide were determined by using one of the above autotitrators which maintained the pH a t 9 with sulfuric acid. Samples were immediately withdrawn and diluted for selenate analysis, while hydroxide consumption was measured by titration. Ratios of selenate to total selenium were obtained by running a series of identical tests ending at various times. Selenate was detehined by injecting a subsample directly into the ion chromatograph through a 0.2-pm syringe filter. Immediately after withdrawing the subsample, the remaining reactor contents were acidified with hydrochloric acid, dissolving the iron oxides and leaving a red precipitate. The aqueous phase was decanted off and discarded. The red selenium was disssolved with gentle heating and 1:l nitric acid. Total selenium on this solution was determined by AAS. Ratios of selenate to ferrous hydroxide were determined by filling 200-mL glass bottles with a known quantity of selenate ion and various amounts of ferrous hydroxide. The bottles were filled to the top to exclude oxygen and
Figure 2. SEM photos of two iron powders used. Particle sizes are larger for electrolytic than for (a) carbonyl or (b) iron.
were tightly capped. The bottles were kept a t 25 "C, protected from light with aluminum foil for a week. Next, the ferrous hydroxide was estimated based on ferrous ion by the use of bathophenanthroline. Selenate ion was determined by ion chromatography. The data for reaction rate versus pH were derived with an autotitrator, maintained at the pH required, with the beaker covered as much as possible to exclude air. The selenate values were determined by taking subsamples from the reactor identical with the method for the selenate to total selenium test.
Results and Discussion Screening Tests Leading to the Discovery of the Ferrous Hydroxide-Selenate Reaction. Selenate removal from systems containing finely divided iron proved to be the most successful. Carbonyliron (Feoderived from pentacarbonyliron) gave the best results. Figure 1shows the rate of selenate ion loss from water in the following order: carbonyliron > electrolytic iron > fine iron wire. The surface area of the different iron types may explain the difference in the reaction rates. Scanning electron microscope (SEM) imagery shows that carbonyliron has a larger surface area to mass ratio than electrolytic iron (Figure 2), which in turn has a larger surface area to mass ratio than the fine iron wire (254-pm diameter). Dissolved sulfate did not change during the test. For the system containing carbonyliron, the pH increased from 5 to 10.5 during the course of the reaction. The pH increase and lack of sulfate removal imply that the selenate removal is due to chemical reduction, rather than adsorption.
Ind. Eng. Chem. Res., Vol. 27, No. 1, 1988 189
loon
.o
Figure 4. Relationship showing 2 mol of NaOH produced per mole of NaPSeO, reduced.
T I M E (Hrr.)
Figure 3. Effect of oxygen on the reaction: (a) reaction under argon, (b) reaction under pure oxygen, and (c) reaction under air.
Selenate removal was not appreciable with carbonyliron and no oxygen (Figure 3a). Therefore, FeO alone cannot be the reductant for selenate. However, selenate removal decreases when the system was exposed to an atmosphere of pure oxygen (Figure 3b). Presumably this is due to the overwhelming oxidation of iron as 2Fe0 + 3/202 + 3H20
-
2Fe(OH),
(1)
When air was metered into the reaction, limiting the availability of oxygen, selenate removal was complete (Figure 3c). This suggests that the iron must first oxidize to ferrous hydroxide, for the reduction to proceed. Proposed Reaction Stoichiometry. The following reaction stoichiometrywas determined experimentallywith excess ferrous hydroxide present:
-
NazSeOl + 9Fe(OH), Seo + 3Fe304+ 2NaOH (magnetite)
+ 8H20 (2)
As the ratio of sodium selenate to ferrous hydroxide increases, the following reaction producing the magnetic oxide, maghemite, becomes predominant, based on XRD evidence supplemented with measurements of ferrous iron:
-+
Na2Se04+ GFe(OH), SeO
+
3Fe203+ 2NaOH 5H20 (3) (maghemite) The standard free energy of a reaction is defined as
Thus, AGO is -83.1 and -23.2 kcal/mol for reactions 2 and 3, respectively (data from Handbook of Chemistry and Physics (1976-1977)). It is assumed in this calculation that AGOse = 0, although the red form of selenium is unstable near room temperature (Cooper and Westbury, 1974) and
02
04 06 08 M O L E S Se04= REDUCED ( x 10-3 1
I O
Figure 6. Relationship between selenate and selenium.
AGOse would be slightly negative. Results on Reaction Stoichiometry. A 2 5 molar ratio of selenate to hydroxide was determined experimentaly for .the reaction (Figure 4). XRD on the iron oxide precipitate with excess ferrous hydroxide present revealed the material to be magnetite or maghemite. The XRD patterns for magnetite and maghemite are similar, and it was not possible to distinguish one from the other. Silver selenide and selenosulfate spot tests (Feigl and Anger, 1972) both gave positive results for elemental selenium. The iron oxides of the precipitate were dissolved with 6 N HC1 a t room temperature, leaving behind a red powder. X-ray fluorescence on the powder showed only selenium with a trace of iron. A 1:lmolar ratio of selenate to elemental selenium was determined experimentally for the reaction (Figure 5 ) . This supports reactions 2 and 3. Figure 5 also provides additional evidence that total selenium in the aqueous phase can be "trapped" in the iron oxide solids.
190 Ind. Eng. Chem. Res., Vol. 27, No. 1, 1988
IZ
-X = 8 . 9
\
Q = 0.4
\
J
€‘E
‘I O i
c
0.0 10
20
30
40
50
60
10
90
90
IO0
I10
120
Time imin )
Figure 7. pH and the reduction of selenate. 0
I 1
I
I
1
I
1
I
J
2
3
4
1
6
7
8
m o l e s o f Fe
(OH), ( ~ 1 0 ” )
Figure 6. Relationship showing 15 mol of Fe(OH2) used per mole of Na2Se04reduced.
A 9:l molar ratio of ferrous hydroxide to selenate was determined experimentally for the reaction (Figure 6). Thus, reaction 2 is believed the principal reaction when excess (>9 ferrous hydroxide/selenate molar ratio) molar ratio is used. It may be that reaction 3 is appropriate at lower ferrous hydroxide/selenate molar ratios. Attempts To Identify Selenite As a Possible Intermediate in the Reduction Process. Ion chromatographic data obtained for the selenate-to-hydroxide test showed no selenite peaks. Selenite was not detected in the aqueous phase over the course of the reaction at a detection limit of 1 mg/L. Several additional tests were performed to determine if selenite existed in the iron oxide precipitate. The solid ferric selenite complex (Siu and Berman, 1984) was prepared. Hydrofluoric acid was shown experimentally to dissolve the complex and give good recoveries for the selenite present. If the ferric selenite complex can be written as “Fe(OH)SeO,”, then possibly the fluoride complexes the iron, releasing the selenite as Fe(OH), Fe(OH)SeO, + 12HF 2FeFG3-+ H,SeO, + 6H+ + 4H20 (4)
+
-
When hydrofluoric acid was used to dissolve the iron oxide from the reaction and the subsequent solution was analyzed by ion chromatography, no selenite was detected. If seleniteselenium does exist in the iron oxide, it would be less than 1% of the total selenium present. In addition, spot tests for selenite (Feigl and Anger, 1972) gave negative results. Possibly the reaction does proceed via selenite and is then quickly further reduced to elemental selenium. Reduction of Selenite with Ferrous Hydroxide. When an equivalent molar amount of sodium selenite instead of sodium selenate was used and the pH of the reaction was maintained a t 9, the reaction proceeded faster than the buret could deliver acid to the system. This was less than 10 min compared to almost 2 h with selenate for a 95% reduction. The XRD of the resulting oxide confirmed magnetite or maghemite. Elemental selenium remained in the beaker after dissolving the iron oxide with acid. Reaction Rate and pH. Figure 7 shows the reaction rate is a function of pH. It is unknown at this time why
a pH of about 9 is preferred. Mining of Selenium. Solutions containing a significant fraction of selenium as selenate, such as in the soda-roast processes, are generally acidified before being treated with reducing agents (Jennings and Yannopoulos, 1974). These heated acid chloride solutions require expensive equipment with glass liners to contain the reaction. The elemental selenium produced may have sulfur as an impurity from the thiourea catalyst. The chemistry of selenate reduction with ferrous hydroxide may be useful in the alkaline soda-roast solutions. The selenium could be recovered by dissolving the iron oxides in acid. It is beyond the scope of this paper to study the feasibility of such a process. Selenium Cycling in Nature. This reaction may provide the basis for aqueous selenate and selenite fixation. Waters that are slightly alkaline may lose selenium to ferrous containing rocks by similar chemistries.
Conclusions Selenate can be reduced to elemental selenium and removed from the aqueous phase by contacting the solution with ferrous hydroxide. The two possible reactions (one forming magnetite, the other maghemite) are exothermic, since the free-energy change for both support a spontaneous reaction. The reaction stoichiometries are derived from the empirical data presented. Since data show no sulfate reduction, the use of ferrous hydroxide solids may be useful for the selective removal of selenate in high sulfate waters. Selenium could be recovered from the spent selenium containing iron oxide bed by dissolving the iron oxides in acid leaving elemental selenium. The powders obtained experimentally were the red form. This is an amorphous form of selenium (Cooper and Westbury, 1974). The optimum operating conditions are between pH 8 and 10. Under similar conditions, selenite is reduced considerably faster than selenate with ferrous hydroxide. Finally, magnetite solids formed may oxidize to maghemite.
Acknowledgment The author thanks the US. Bureau of Reclamation for financial support and David K. Mueller, U S . Geological Survey, for editorial assistance.
Znd. Eng. Chem. Res. 1988,27, 191-194 Registry No. Fe(OH),, 18624-44-7.
Literature Cited Ames, L. L.; Salter, P. F.; McGarrah, J. E.; Walker, B. A. “Selenate-Selenium Sorption On a Columbia River Basalt“. Chem. Geol. 1979,43, 287. Baldwin, R. J.; Acres, W.; Stauter, J. C.; Terrell, D. L. U S . Patent 4 405 464, 1983. Chau, Y. K; Riley, J. P. Anal. Chim. Acta 1965,33, 36-49. Cooper, W. C.; Westbury, R. A. In Selenium; Zingaro, R. A., Cooper, W. C., Ed.; Van Nostrand Reinhold: New York, 1974; pp 109-111. De Carlo, E. H.; Zeitlin, H.; Fernado, Q. Anal. Chim. 1981, 53, 1104-1107. Doran, J. W.; Alexander, M. “Microbial Transformations of Selenium”. Appl. Enuiron. Microbiol. 1977, 33, 31-37. Doumas, A. C. US.Patent 3387928, 1968. Feigl, F.; Anger, V. Spot Tests in Inorganic Analysis, 6th ed.; Elsevier: New York, 1972; pp 408-409. Gersberg, R. M.; Brenner, R.; Elkins, B. V. “Removal of Selenium Using Bacteria”. California Agricultural Technology Institute: 1986; Pub. No. CATI/860201. Handbook of Chemistry and Physics, 57th ed.; CRC: Boca Raton, FL, 1976-1977; pp D-67-D-77. Hartmanis, M. G.; Stadtman, T. C. “Isolation of a Selenium-Containing Thiolase From Clostridium Kluyuen’: Identification of the Selenium Moiety as Selenomethionine”. Proc. Natl. Acad. Sci. U.S.A. 1982, 79, 4912-4916. Hingston, F. J.; Atkinson, R. J.; Possner, A. M.; Quirk, J. P. “Specific Adsorption of Anions”. Nature (London) 1967,215,1459-1461. Hingston, F. J.; Possner, A. M.; Quirk, J. P. “Adsorption of Selenite by Geothite”. Adu. Chem. Ser. 1986, 70,82-90.
191
Jackson, T., Research Leader, Fish and Wildlife Service, Denver Field Office, Columbia National Fisheries Research Laboratory, unpublished results, 1986. Jennings, P. H.; Yannopoulos, J. C. In Selenium: Zingaro, R. A., Cooper, W. C., Eds.; Van Nostrand Reinhold New York, 1974; pp 48-59. Kauffman, J. W.; Laughlin, W. C.; Baldwin, R. A. U S . Patent 4 519 912, 1985. Lee, G. F.; Stumm, W. “Determination of Ferrous Iron in the Presence of Ferric Iron With Bathophenantholine”. J . Am. Water Works Assoc. 1960,52, 1567-1574. Maneval, J. E.; Klein, G.; Sinkovic, J. “Selenium Removal From Drinking Water by Ion Exchange”. Municipal Environmental Research Laborator.7, Office of Research and Development, EPA-81-02-5401, Cincinnati, OH, 1983. Marchant, W. N. US.Patent 3 933 636, 1976. Oppenheimer, J. A.; Eaton, A. D.; Kreft, P. H. “Speciation of Selenium in Groundwater”. Water Engineering Research Laboratory, EPA-600/S2-84-190, Cincinnati, OH, 1985. Schwertmann, U.; Taylor, R. M. In Minerals in Soil Enuironments; Dixon, J. B., Weed, S. B., Eds.; Soil Science Society of America: Madison, WI, 1977; pp 167-172. Siu, K. W.; Berman, S. S. “Determination of Selenium(1V) in Seawater by Gas Chromatography after Coprecipitation with Hydrous Iron(II1) Oxide”. Anal. Chem. 1984, 56, 1806-1808. Sorg, T. J.; Logsdon, G. S. “Treatment Technology to Meet the Interim Primary Drinking Water Regulations for Inorganics: Part 2”. J. Am. Water Works Assoc. 1978, 70, 379-393. Received for review September 17, 1986 Revised manuscript received August 25, 1987 Accepted October 3, 1987
COMMUNICATIONS A Finite Element Solution for Diffusion and Reaction in Partially Wetted Catalysts with Power-Law Kinetics T h e Galerkin finite element method (FEM) is used t o obtain approximate solutions to the problem of diffusion and reaction in partially wetted catalyst pellets for power-law reaction rate forms. For the case of a first-order reaction, it is shown t h a t effectiveness factors developed from the FEM solution are in good agreement with those obtained from a n integral equation and a least-squares solution. The effects of reaction order and other associated parameters on the concentration profiles and catalyst effectiveness factor are briefly illustrated. Fixed beds of solid catalyst with gas-liquid cocurrent flow are commonly used in commercial processes such as hydroprocessing of various petroleum feedstocks, selective partial hydrogenation of chemicals, and oxidation of dilute contaminants that may be present in aqueous process streams (Germain et al., 1979; Ramachandran and Chaudhari, 1984). Under certain process conditions for a particular reacting system, the catalyst pores can be either totally or partially liquid-filled and the catalyst surface can be covered by a combination of actively flowing liquid films and stagnant liquid films or contain dry patches that are in direct contact with the gas or possibly volatile liquid (Mills and DudukoviE, 1980). The ability to evaluate the catalyst effectiveness factor for these situations is an important step that should be performed when designing new reactors or during interpretation of performance for existing reactor systems (cf., Beaudry et al. (1985)). Over the past decade, various models for evaluation of catalyst effectiveness factors for partially wetted catalysts have 0888-5885/88/2627-0191$01.50/0
been developed (DudukoviE, 1977; DudukoviE and Mills, 1978; Goto et al., 1981; Herskowitz et al., 1979;Herskowitz, 1981a,b;Lee and Smith, 1982; Levec et al., 1980; Martinez et al., 1981; Mills and DudukoviE, 1979, 1980; Mills et al., 1981; Ramachandran and Smith, 1979; Sakornwimon and Sylvester, 1982; Tan and Smith, 1980). With only one exception (Goto et al., 1981),these models assumed linear first-order reaction kinetics which limits their application since many systems having practical significance as described by nonlinear kinetics (cf., Ramachandran and Chaudhari, 1984). For the above case where nonlinear kinetics were assumed, a finite-difference scheme with a fixed grid size was used to solve the resulting split boundary-value problem for regular slab geometry. This scheme is not necessarily the best choice for this type of problem since detailed studies on related problems described by Laplace’s equation have shown that the solution accuracy can be quite low in the vicinity of the split boundary discontinuity (cf., Whiteman, 1968). Steep 0 1988 American Chemical Society
