Environ. Sci. Technol. 1992, 26, 1922-1931
Extraction of Uranium and Molybdenum from Aqueous Solutions: A Survey of Industrial Materials for Use in Chemical Barriers for Uranium Mill Tailings Remediation Stan J. Morrison* and Robert R. Spangler
Environmental Sciences Laboratory, Grand Junction Projects Office,+P.O. Box 14000, Grand Junction, Colorado 81502 Laboratory experiments were performed to simulate the interaction of contaminated pore fluids with a variety of industrial materials. The objective was to evaluate the materials for use in a chemical barrier under a repository containing uranium mill tailings. Pore water would pass through the barrier, but contaminants would remain fixed in the solid fraction. More than 99% of the dissolved uranium in a synthetic pore fluid (initial uranium concentration of 30.0 mg/L) was extracted by the addition of hydrated lime, fly ash, barium chloride, calcium phosphate, titanium oxide, peat, and lignite. More than 96% of the molybdenum (initial molybdenum concentration of 8.9 mg/L) was extracted by ferrous sulfate, ferric oxyhydroxide, titanium oxide, peat, hematite, calcium chloride, and barium chloride. Some materials were effective only for a limited range of pH values. Extraction was caused by both precipitation (as calcium uranate, calcium molybdate, ferrous molybdate, or barium molybdate) and sorption (on ferric oxyhydroxide, hematite, calcium phosphate, peat, or titanium oxide). Chemicals that precipitate contaminant-bearing minerals are able to control solution chemistry and, therefore, have an advantage over sorbents which are subject to externally determined solution variables such as pH. On the basis of the predicted flux of pore fluid from the Monticello (Utah) uranium mill tailings, some industrial materials may be suitable for a chemical barrier at that site. ~
~~
~
~~
Introduction More than 230 million tons of uranium mill tailings are present at mill sites throughout the United States ( 1 , 2 ) . Uranium tailings result from uranium ore processing to supply fuel for civilian nuclear power plants and defense purposes. The need for nuclear fuels is likely to continue because of increased awareness of the environmental impacts of burning fossil fuels (particularly the “greenhouse effect”) and the desire to reduce oil imports. Contaminants in the uranium mill tailings piles threaten to spread into the groundwater at the mill sites. Economically attractive means are needed to treat and dispose of new tailings and to remediate existing tailings. The pump-and-treat method is the most widely used technology for restoring groundwater. However, the results of many pump-and-treat projects indicate that this technology is not an effective permanent solution for reducing contaminant concentrations in groundwater (3). In addition, costs can be prohibitive; the remediation of existing contaminated groundwater associated with the 24 Uranium Mill Tailings Remedial Action (UMTRA) sites is an estimated $1billion (4). Methods are needed that will ensure the prevention of groundwater contamination. The purpose of this study was to develop a chemical barrier for a proposed uranium mill tailings repository a t Monticello, UT. We examined chemically complex systems, similar to the pore fluid chemistry expected in the mill tailings, and simplified systems to determine the fundamental reactions responsible for contaminant ex+ Operated by Chem-Nuclear Geotech, Inc., for the U.S. Department of Energy.
1922 Environ. Scl. Technol., Vol. 26, No. 10, 1992
traction. Specific objectives of the study were to (1)screen materials that might economically extract uranium and molybdenum from uranium mill tailings pore fluids, (2) determine the extraction characteristics of the selected materials for uranium and molybdenum, and (3) define operative chemical reactions that can be used in predictive models. Primarily, low-cost industrial materials (or reagent-grade proxies) were investigated. Uranium mill tailings commonly contain a wide variety 230Th,Pb, Se, As, of contaminants, including U, Mo, 226Ra, Cu, Zn, and Cr, that potentially could be released to the groundwater. The mineralogy of the tailings (5) and groundwater flux partially govern these releases. This study was limited to two major contaminants because a great number of analyses would be required to test all potential contaminants with a variety of materials. Our selection of uranium and molybdenum was based on two factors: (1)uranium and molybdenum are the most mobile elements at uranium mill tailings sites, and (2) molybdenum is often mobile under high-pH conditions in which uranium is immobile and vice versa. Combined, these two elements are probably the most difficult to stabilize. The US. Environmental Protection Agency has proposed standards for uranium and molybdenum of 30 pCi/L (0.043 mg/L, based on natural isotopic abundances) and 0.1 mg/L, respectively, for groundwater at uranium mill tailings sites (I). Uranium, in concentrations above 0.043 mg/L, has migrated at least 4000 f t downgradient from tailings piles at Monticello, UT (6), and at least 2000 f t downgradient from the tailings pile at Riverton, WY (7). Uranium plumes have spread from most other uranium mill tailings sites, and molybdenum plumes are also known to exist (8). A chemical barrier contains materials that retard the spread of contaminants by allowing the passage of pore water while retaining the contaminants. Few studies exist that address the technology of chemical barriers. Longmire et al. (9) examined the possibility of mixing peat, limestone, or hydrated lime with uranium mill tailings to prevent contamination of underlying aquifers. They found that the addition of hydrated lime decreased the concentration of dissolved uranium and other contaminanb but increased the concentration of molybdenum. However, the addition of peat decreased dissolved uranium, molybdenum, and other contaminants. Other studies focused on the extraction of contaminants from uranium mill effluent using lime, limestone (IO), or biosorbents (11). Opitz et al. (IO) determined that uranium concentrations were reduced from 40.0 to 0.10 mg/L by adding lime to reach a final pH of 7.3. In the Barnes et al. (11)study, calcium alginate beads extracted 78% of the uranium and 74% of the molybdenum from uranium-contaminated mill tailings pore fluid. Materials and Methods Potential chemical barrier materials were evaluated by batch tests. A batch test consisted of combining the test material with an aliquot of synthetic uranium mill tailings pore fluid. The 500-mL Erlenmeyer flask containing the
0013-936X/92/0926-1922$03.00/0
0 1992 Amerlcan Chemical Society
Table I. Comparison of Synthetic Pore Fluids (SPF1 and SPFZ) with Tailings Pore Fluid (TPF") (in mg/L; 226Ra in pCi/L; Ehin mV)
SPFl SPFS TPF L1
Na
K
Ca
Mg
SOa
C1
C
As
3680 3654 3529
23 0 23
340 301 340
65
6300 6920 6300
1000 1000 1000
232 232 466b
1.00 2.00 0.30
0
64
U
V
Mo
Pb
226Ra
pH
E,,
30 30 13
1.6
8.9 8.9 9.6
2.0
20
0
0
nac
na
7.8 7.8 7.5
300 300 434
0
1.6
TPF was collected October 29, 1984, from the Monticello site. *Estimated from alkalinity measurements.
e
Not analyzed.
Table 11. Materials Used in Batch Tests
acronym
material
BACL2 CACLB CAOH2 CLINO COAL FEN03 FES04 FE3SO FE203 FE304 FLY 1 FLY2 FLY3 GOETH GYP LIGN MONT PEAT PHOSl PHOS2 SAND SAWD TIOH TI02
BaC12.2Hz0 CaCl2-2H2O Ca(OH), clinoptilolite subbituminous coal Fe(N03)3.9H20 FeSO, Fez(SO&-nHzO Fe2OS(hematite) FeSO4(magnetite) fly ash F fly ash C fluidized-bed fly ash F FeO(0H) CaS04.2Hz0 lignite Ca-montmorillonite sphagnum peat moss phosphate rock hydroxyapatite quartz sand sawdust Ti(OH), TiOz (anatase)
BET," m2/g nab na na 4.10 2.57 na na na 10.88 6.05 na na na 30.7 na 2.20 38.65 0.66c 7.41 45.16 na 2.25 0.49 7.41
description and source reagent-grade granular (Fisher B34-500) reagent-grade powder (Fisher C-79) reagent-grade powder (Aldrich 23,923-2) powder (Ward's 25519) powder (Argonne Premium Coals 202) reagent-grade crystals (Fisher 1-110) reagent-grade powder (Fisher 1-147) reagent-grade powder (Baker 2046) reagent-grade powder (Fisher I 116-3) reagent-grade powder (Am. Chem. Enterprises A-310) industrial powder (Craig, CO: sample from Pozzolanic Int.) industrial powder (Pawnee, CO: sample from Western Ash Co.) industrial powder (sample from Western Ash Co.) synthetic goethite (13) reagent-grade powdered gypsum (Aldrich 25,554-8) -100 mesh (Argonne Premium Coals 801) Clay Minerals Society source clay STx-1 garden grade (Hyponex Corp.) -100 mesh; P206,32.5% (Soda Springs, ID, Monsanto Corp.) reagent-grade powder; P206,39.4% (Aldrich 12167-74-7) granular (UNIMIN Corp., Granusil silica) industrial grade (about 85% pine, 15% hardwoods) synthetic crystals (12) reagent-grade powder (Fisher T-315)
Surface area as determined by BET analysis. Not analyzed. Dried and milled to -100 mesh.
mixture was sealed with a rubber stopper and agitated on an Eberbach Model 6140 rotating shaker table. Following agitation, the mixture was filtered through a 0.45-pm filter, and the filtrate was analyzed for uranium, molybdenum, pH, and Eh. In some experiments, the pH was adjusted using NaOH or HN03. Controls (with no additives present) confirmed that neither uranium nor molybdenum extraction was due to pH adjustments alone. Water used in all experiments was deionized to greater than 10 M!km resistivity. An Orion Ross glass combination electrode was used to measure pH. A combination platinum redox electrode and silver/silver chloride reference electrode (Orion 96-78-00) was used to measure E h . Measurements of pH and E h were usually made within 30 min after opening the sealed test vessel; however, no special precautions were taken to exclude the atmosphere during filtering or measurement. The effects of atmospheric exposure on the pH and E h measurements for the peat experiments were examined in an experiment in which 5 g of peat was agitated with 200 mL of synthetic pore fluid (SPF1) maintained under an argon atmosphere for 9 days. Four measurements each of pH and Eh were made without exposure to the atmosphere. The ranges of E h (426-534 mV) and pH (3.8-4.1) are similar to the measured E h and pH (489 mV and 4.4) of the exposed sample, indicating that the relatively high E h values of the peat experiments are not artifacts of atmospheric exposure. Synthetic pore fluids (SPF1and SPF2) were prepared from reagent-grade chemicals. SPFl represents one type of uranium mill tailings pore fluid from the Monticello site (Table I); SPFS contains concentrations of the major
cations and anions similar to those of SPF1. Periodic analyses of the synthetic pore fluids indicated that uranium and molybdenum concentrations did not change over time. The uranium concentration (30 mg/L) used in SPFl and SPF2 was greater than in the tailings pore fluid obtained from field lysimeters (13 mg/L) at Monticello to observe a broader range of extraction capacity. The batch tests comprised 24 potential extraction materials. Table I1 presents descriptions and sources of the materials (materials are referred to by an upper case threeto five-letter abbreviation). The provided mineral names were confirmed by X-ray diffraction. Hydrous titanium oxide [Ti(OH),] was synthesized from Tic& and NaOH using the method described by Inoue and Tsuji (12). The OH/Ti ratio was 5.1, which Inoue and Tsuji concluded was the optimal ratio for uranium sorption. Goethite was synthesized by the method described by Tripathi (13). Uranium concentrations were determined by standard addition laser-induced fluorometry on a Scintrex UA-3 uranium analyzer. Uranium results were verified on selected samples with inductively coupled plasma/mass spectrometry (ICP/MS). Because organics can interfere with laser-induced fluorometry, uranium from one of the peat experiments was confirmed by ICP/MS analysis. The UA-3 method detects uranium concentrations as low as 0.0005 mg/L. Our reproducibility of a standard was about 10%. The UA-3 method can detect only uranyl ions. The E h in all experiments was high enough that uranyl should be the dominant form of dissolved uranium. Molybdenum concentrations were measured by ternary complex or mercaptoacetic acid photometric methods. Multiple molybdenum analyses of SPFl and SPF2 on Envlron. Sci. Technol., Vol. 26, No. 10, 1992
1923
different days were reproducible to f3%. Because iron can interfere with molybdenum analyses, filtered FeS04 solutions spiked with known aliquots of molybdenum were analyzed. The results indicate that iron did not interfere with the molybdenum analyses. The detection limit for molybdenum was 0.02 mg/L. The computer program MINEQL (14) was used for calculations of mineral stability and adsorption. The uranium speciation model is from Tripathi (13). A complete list of the thermodynamic data used in the calculations is provided as supplementary material.
Rationale for Choice of Test Materials Materials selected for study were known to be effective for concentrating uranium into ore deposits, precipitating uranium and molybdenum from mill solutions, extracting uranium from seawater, or treating wastewater. Availability and cost also were considered in the selection process because substantial quantities would be required for a chemical barrier under a tailings repository. Reagent-grade chemicals were used in some experiments as proxies for industrial chemicals. To determine the extraction efficiency of humic materials, we chose to use sawdust, peat, lignite, and subbituminous coal. Humic materials are derived from the woody tissues of plants. Leventhal(15) described humic material as complex organic molecules with an approximate chemical formula of (C15H1608NS)nand molecular weight of 500-20 OOO. Field studies of uranium ore deposition report the ability of humic substances to extract uranium from solution (16,17). Moore (18)determined that peat, lignite, and subbituminous coal extracted 98.0, 98.4, and 99.9% each of the uranium from a 200 mg/L solution. Longmire et al. (9) showed that mobility of uranium substantially decreased by adding peat to the tailings. Hematite, goethite, magnetite, and ferric oxyhydroxide were selected to assess the extraction capabilities of a variety of oxidized iron compounds. Field observations of uranium ore deposits associated with iron oxides (19) and laboratory studies with goethite [FeO(OH)] (13) suggest that oxidized iron compounds are effective in extracting uranium from solution. Some magnetic iron oxide compounds are able to adsorb metals and have been used for waste stream treatment (20). Ferric oxides and ferric oxyhydroxides also are known to be effective for extracting molybdenum and uranium, as evidenced from laboratory adsorption experiments (21,22) and by the adsorption of molybdenum by hematite (Fe203)in soils (23)and by ferric oxyhydroxides in stream sediments (24). Ferrous sulfate was selected because of its potential to form ferrous molybdate, which is thought to limit molybdenum concentrations in some molybdenum-rich streams (25). The effectiveness of hydrated lime for extracting uranium from solution is reported in studies with hydratedlime-treated uranium mill effluents (10,26) and hydrated-lime treatment of uranium mill tailings (9). Fly ash is a waste product from coal combustion and has a significant lime (CaO) component. Because fly ash has some reactive properties similar to those of lime, it was thought that it might extract uranium from solution. Fly ash is classified in two categories, C and F; class C fly ash contains more lime than class F (271. Most marine phosphorites contain 50-300 mg/kg uranium (28) and many yield uranium as a byproduct. Some fossil bones and teeth containing phosphate minerals have accumulated uranium (28). These observations suggest that phosphate is able to extract uranium from groundwater. Titanium oxides and hydroxides have the ability to extract uranium. Numerous studies report Ti(OHI4is 1924
Envlron. Scl. Technol., Vol. 26, No. 10, 1992
Table 111. Results of Experiments" Using Lime with Other Additives
concn, mgjL U Mo 0.18 0.23 0.04
0.03 0.33
10.70'
8.00 8.07 8.00 9.40'
Eh,
pH
mV
12.10 12.45 12.49 12.50 12.32
199 194 139 132
289
other additives:
g
5 CACL2,3 GYP, 90 SAND 2 BACL2, 3 GYP, 90 SAND 5SAWD, 90SAND 5PEAT, 90 SAND 4 FE203
All experiments contained 5 g of Ca(OHI2and 200 mL of SPFl in addition to the other listed additives. Shake time was 14 days. Initial concentrations of U and Mo were 30 and 8.9 mg/L, respectively. *Refer to Table I1 for acronyms. 'Values that exceed the starting value of 8.9 are probably the result of analytical error.
capable of extracting uranium from seawater that has only 0.003 mg/L uranium (29, 30). Barium chloride and calcium chloride were selected because of the possibility of forming low-solubility barium or calcium molybdates. Merritt (26) described the precipitation of calcium molybdate from uranium mill liquors using calcium chloride. Under reduced conditions, uranium will form sparingly soluble minerals such as uraninite. Agents that require extremely reduced conditions to extract uranium and molybdenum (for example, sulfides) were not considered because we believe it would be difficult to maintain a suitable reduced state under the shallow earth conditions associated with a tailings repository.
Results Batch experiments involving a single additive (with no pH adjustment) in synthetic pore fluid were run for 14 days (Figure 1). These experiments used a relatively large concentration of the test material to maximize extraction results. Dissolved uranium concentrations were lowered from 30.0 mg/L to less than 0.5 mg/L by lignite, peat, hydrated lime [Ca(OH)2],fly ash, barium chloride (BaC12*2H20),titanium oxide (Ti02), and hydroxyapatite [Calo(OH)z(P04)6].Dissolved molybdenum concentrations were lowered from 8.9 mg/L to less than 0.4 mg/L by ferrous sulfate (FeSOJ, calcium chloride (CaC1,.2H20), ferric sulfate [Fe2(S04)B-nH20], ferric nitrate [Fe(N03)3-9H20], hematite (Fez03),peat, and barium chloride. In these experiments, only peat and barium chloride extracted significant amounts of both uranium and molybdenum (Figure 1). Synthetic pore fluid treated with hydrated lime, Pawnee class C fly ash, and fluidized-bed fly ash (CAOH2, FLY2, and FLY3 on Figure 1)yielded relatively high pH values (12.65, 12.18, and 10.85, respectively). Ferrous sulfate, ferric sulfate, ferric nitrate, subbituminous coal, peat, and calcium chloride produced acidic solutions (pH 2.18,2.82, 2.85, 3.62, 4.40, and 4.58, respectively), whereas most of the other additives produced neutral to weak alkaline (mostly near pH 8) solutions. Ehmeasurements indicated that all solutions were relatively oxidized, so the precipitation of reduced uranium or molybdenum solid phases is unlikely (Figure 1). In experiments where synthetic pore fluid was treated with mixtures of hydrated lime and other additives (Table 111),the uranium concentration was significantly depleted but the molybdenum concentration was only slightly affected by the mixtures. The pH of each of these solutions exceeded 12. The amounts of additive were varied in some experiments (Figure 2a and b). The uranium concentration decreased by 99.8%, from 30.0 to 0.07 mg/L, when only a small amount (0.75 g/L) of hydrated lime was used
800
.
a
................................................. ...................................................
14.00 I
I
12 00
1000
0 00
pH
600 4 00
2 00 0 00
I
l
I
l
l
l
15.00
I ............
30,00-.
.
.
.
.
.
.
.
.
.
........................................................... ..........................
...........................
(mg/L)
15.00-
.
.
.
.
...............................
.........................
5.00 -
........................... 0.48 0.33 0.19 0.08 0.040.03 0.020.0' i
1
l
l
l
I
Flguro 1. Resutts of batch experlments using single additlves, arranged in order of decreasing U in the final solution. The number in parentheses by each additive Is the ratio of additive to SPFl (g/L). For abbreviations, refer to Table 11. All experiments were run for 14 days. Starting concentratlons of U and Mo were 30.0 and 8.9 mg/L, respectlvely. I n the experiments with Fe,SO and FeNO,, the pH was Increased (to 2.82 and 2.85, respectively) wlth NaOH to precipitate ferric oxyhydroxides. All other experiments contained only the indicated additive and SPF1.
(Figure 2a). Fluidized-bed fly ash (FLY3), barium chloride, calcium phosphate (PHOS2), and peat also extracted uranium but required larger amounts (about 7,30,20, and 5 g/L, respectively) to lower the uranium concentration to less than 1.0 mg/L. Large amounts (to 50 g/L) of sawdust (SAWD), calcium chloride (CACLB), titanium oxide (TI02), and hematite (FE203) did not significantly lower uranium concentrations (only TI02 is shown on Figure 2a). Maximum uranium extraction in the hydrated lime series occurred at pH 10.88, which is consistent with the range of pH values (10.27-11.22) that successfully extracted uranium when fluidized-bed fly ash was added. Ferrous sulfate (FES04) required only 5 g/L to lower the molybdenum concentration by 99.4%, from 8.9 mg/L to
less than 0.02 mg/L (Figure 2b). Peat (PEAT), barium chloride (BACLS), and hematite (FE203) required more material, about 25,30, and 45 g/L, respectively, to lower molybdenum concentrations to less than 0.5 mg/L. Large amounts (to 50 g/L) of calcium phosphate (PHOSS), sawdust, hydrated lime (CAOHB), titanium oxide, and calcium chloride did not significantly lower molybdenum concentrations (only TI02 is shown on Figure 2b). The extraction efficiency of peat, hematite, titanium oxide, and phosphate for uranium and molybdenum varied with pH. All of these materials had uranium extraction minima between pH 5.5 and 7.0 (Figure 3a). Peat was effective for uranium extraction even when no pH adjustment was made; however, hematite and titanium oxide Envlron. Sci. Technol., Voi. 26, No. 10, 1992
1925
100 00
I
I
I
I
I
I
I
I
1
\ \ a
2
3
4
5
6
7
8
9
6
7
8
9
DH
0001
Ratio of Additive to SPF (g/L) 10000
I
~
i
I
I
I
I
I
I
* Value was leos lhsn delscllan !
. 3
t 0
PEAT FEZ03 TI02 PHOS2
3
2
5
4
PH
0
5
10
15
20
25
30
35
40
45
50
Ratio of Additive to SPF (giL)
Flgure 2. Effect of the amount of addklve on contaminant extractlon. The ratio of additive to synthetic pore fluid is plotted against the (a) U and (b) Mo left In solution. Initial concentrations were 30.0 and 8.9 mg/L for U and Mo, respectively. Shake times were 7 days for the FeSO, and 3 days for all others. PHOS2, and TiO,, used SPF2; all others used SPF 1.
were relatively ineffective without pH adjustments (arrows on Figure 3a point to data from experiments with no pH adjustments). Phosphate extracted uranium effectively over a wide range of pH. Peat and titanium oxide extracted molybdenum most efficiently at the low pH values; hematite and phosphate extracted molybdenum most efficiently at about pH 6 and 4, respectively (Figure 3b). Hematite lowered molybdenum concentrations to less than 1 mg/L in a broad pH range of 2.15-7.91. The adsorption of uranium and molybdenum on ferric oxyhydroxide also varies with pH. A series of experiments was run in which 1 g of ferric nitrate [Fe(N0J3.9H20], adjusted to various pH values, was agitated with SPFl for 3 days. Uranium adsorption behaved similarly to that of cations, adsorbing more strongly at high pH (Figure 4a), while molybdenum adsorption behaved similarly to that of anions, adsorbing more strongly at low pH (Figure 4b).
Discussion Use of chemical barriers for waste containment necessitates reliance on groundwater models for long-term prediction of contaminant migration. To confidently predict contaminant migration, the chemical reactions of the barrier components responsible for contaminant extraction need to be reasonably well understood and extrapolated to the subsurface environment. Both sorption and precipitation may be responsible for the extraction of uranium or molybdenum. The concentrations of contaminants on solid sorbents decrease and the solubility of contaminant-bearing solids lQ26 Environ. Scl. Technoi., Vol. 26, No. 10, 1992
Flgure 3. Effect of pH on contamlnant extraction: (a) U (initial concentratlon 30 mg/L) and (b) Mo (initial concentration 8.9 mg/L) left in solution. Experiments used 1.0 g of peat, 4.0 g of TIO, or 4.0 g of Fe203In 100 mL of SPFP; or 5 g of PHOS2 in 100 mL of SPF1. Shake tlme was 3 days. Arrows polnt to experiments that had no pH adjustment. 100
I
I
90 --
0
-
80 --
-
Yi
2
3
k
4
i
5
6 3- i
6
8
7
e
1
0
1
1
1
2
'
3
PH
Figure 4. Adsorption from SPFl on ferric oxyhydroxide [ l o g/L of Fe(NO3),.9H,O]: (a) U (inltlai concentration 30 mg/L); (b) Mo (initial concentration 8.9 mg/L). Open circles are experimental data; solid curves are model calculations.
increase as the concentrations of complexing agents increase. Ligands in the synthetic pore fluid were limited to carbonate, sulfate, chloride, and hydroxyl. Fluoride and phosphate complex with uranium (13)but were not con-
pn
-101
8
I
I
10
I
1
I
I
12
PH Figure 5. Activity diagrams. Both diagrams are drawn at T = 25 OC, P = 1 bar; no corrections were made to activity coefficients for the effect of ionic strength. (a) €,-pH diagram showlng the stability fields for Iron molybdates In equilibrium with amorphous Fe(OH), or Fe(OH),. [Mol = m (0.1 mg/L). E , and pH values from experiments that used FeSO,, FB~(SO,)~, and Fe(NO& are plotted, and the resulting Mo concentrations (mg/L) are shown. Dashed llne shows the extension of the FeMoO, stability fleld if Fe#hO,), does not form. The arrow connects two experiments that are identical except that the pH of the point at the arrow tip was Increased with NaOH. (b) Stablllty field of CaUO, for Ca concentrations UO': of lo-', lo-*, and lo-, mol/L. Point A depicts equilibrium between Ca(OH), and CaU04. log K = -24.9 was used for the reaction Ca2' -I-2H,O = CaUO, 4H'.
+
sidered because they typically do not occur in significant concentrations in sandstone-type uranium mill tailings pore fluids. Carbonate is the most important uranium complexing agent in common groundwaters. The predominant aqueous complex of molybdenum in groundBelow pH 4,HMo04- dominates. water is Contaminant Extraction by Precipitation. In some of the batch experiments, contaminants were extracted from solution by precipitation of a solid material. The conditions under which some of the precipitation reactions occur are shown on activity diagrams in Figure 5. Trends shown on these diagrams are probably correct; however, equilibrium constants for the molybdenum- and uraniumbearing solid phases used in the diagrams and for aqueous uranium species above pH 9 have not been accurately determined. Additionally, corrections were not made to activity coefficients for the effect of ionic strength, a factor that is particularly important in the high-solubility areas of the diagrams. Synthetic pore fluid turned orange with the addition of ferrous sulfate, indicating that some ferrous iron oxidized to ferric solids. Only a small amount of ferric solids formed with the addition of ferrous sulfate in contrast with the mwive brownish-red ferric solids formed when NaOH was added to solutions containing ferric sulfate, ferric chloride, or ferric nitrate. We did not identify the solids that formed from ferrous sulfate, ferric sulfate, ferric chloride, or ferric nitrate; however, they were probably similar to the ferric oxyhydroxide that was identified in precipitates formed
+
from the fast hydrolysis of ferric salts (31). It is difficult to distinguish precipitation of molybdates from coprecipitation of molybdenum with ferric solids or sorption of molybdenum on ferric solids. Considerations of molybdate mineral stability fields (Figure 5a) and estimates of the quantities of precipitated ferric solids suggest that ferrous sulfate caused precipitation of ferrous molybdate (FeMo0,) whereas ferric nitrate and ferric sulfate caused precipitation of ferric solids, which then sorbed molybdenum. Figure 5a shows Eh and pH stability fields in equilibrium with ferric hydroxide [Fe(OH)3]. Ferric hydroxide is used as an analog for the ferric solids that formed in the experiments. Precipitation of iron molybdates occurs at relatively low pH values. At high pH, ferric hydroxide is stable, which causes dissolved iron concentrations to be low and limits the formation of iron molybdate. The addition of ferric sulfate to synthetic pore fluid (without pH adjustment) did not cause significant extraction of molybdenum, even though ferric molybdate should be a stable phase (Figure 5a), which suggests that ferric molybdate did not form in our experiments. However, after the addition of NaOH, which increased the pH from 2.36 to 2.82 and caused precipitation of ferric oxyhydroxide, less than 0.1 mg/L molybdenum remained in solution (arrow on Figure 5a). Sorption of molybdenum on ferric oxyhydroxides in the experiments involving ferric sulfate or ferric nitrate is suggested since significant molybdenum extraction occurred only if ferric oxyhydroxide precipitated. Environ. Sci. Technol., Voi. 26, No. 10, 1992
1027
Table IV. Effect of Calcium Chloride on Molybdenum Extraction CaC12.2H20" added, g/L 0.0 (SPF2) 1.00
9.98 49.39 97.56 445.20 633.40
shake time, h 0
72 72 72
72 72
336
CaMoOl (powellite) SIb
dissolved Mo, mg/L
-0.09 0.19 0.96 OS' OS' OS' OS'
8.90 8.85 8.75 8.60 8.20 7.10 0.67
"After accounting for the water of hydration. bSaturation index, log (ion activity product/solubility product); negative and positive indicates undersaturation and oversaturation, respectively. Oversaturated with powellite; however, the ionic strength is above the limit to allow accurate calculation of activity coefficients (or SI).
The solution chemistry of four experiments that used ferrous sulfate concentrations ranging from 5 to 50 g/L plot within the stability field for ferrous molybdate if ferric molybdate does not form (Figure 5a). Molybdenum concentrations were low (less than 0.02 mg/L) in all four experiments. A fifth experiment, which used only 0.6 g/L of ferrous sulfate, plots outside the stability fields for the solid molybdates and had a higher molybdenum concentration (2.85 mg/L). The molybdenum concentrations in these five experiments are consistent with the formation of ferrous molybdate. In a separate experiment, reagent-grade hydrated lime added to uranyl nitrate solution (1000 mg/L uranium containing no other components) caused a yellow uraniferous solid to precipitate and a decrease in uranium concentration to 0.03 mg/L. Because the precipitate is X-ray amorphous, identification was not possible. I t is probably a calcium uranate (for example, CaUOJ as described by Tanford et al. (32). Using 0.03 mg/L as the solubility, a preliminary log K value of -24.9 was derived and the stability field for CaU04 was calculated (Figure 5b). Figure 5b shows that the stability field for CaUO, is limited to high-pH, Ca-bearing solutions. The extraction of uranium from synthetic pore fluid using hydrated lime is likely caused by precipitation of calcium uranate. Molybdenum is mobile in the high-pH environment associated with hydrated lime (Table 111). Fly ash significantly lowered uranium concentrations only in those experiments with elevated pH suggesting that, as in the situation with hydrated lime, the mechanism is the formation of calcium uranate. Although the pH of saturated lime solutions is usually between about 12.0 and 12.6 (depending in part on the calcium concentration), the most effective pH for uranium extraction was about 11in combination with either hydrated lime or fluidized-bed fly ash. This suggests that a uranyl aqueous complex (not included in the calculation for Figure 5b) is stable at pH more than 11. Treatment with CaCl2-2Hz0is thought to cause molybdenum extraction from uranium mill liquids by precipitation of powellite (CaMoO,) (26). The addition of 150 g of CaC12.2H20to 100 mL of SPF2 caused the molybdenum concentration to decrease from 8.9 to 0.67 mg/L after 336 h. Five additional experiments with CaCl2-2Hz0 additions ranging from 1to 445 g/L failed to significantly reduce molybdenum concentrations after 72 h, despite oversaturation with powellite (Table IV). Apparently, precipitation of powellite requires highly oversaturated conditions or contact times longer than 72 h. Synthetic pore fluid SPF2 is close to saturation (Table IV), suggesting that molybdenum concentrations in the Monticello 1928
Envlron. Sci. Technol., Vol. 26, No. 10, 1992
tailings pore fluid may be controlled by reactions with powellite. Barium chloride effectively extracted both uranium and molybdenum if 50 g/L or more was used (Figures 1 and 2). Because barium is not known to form minerals that are good sorbents but is known to form low-solubility solids with molybdenum and possibly with uranium, extractions were likely caused by precipitation. Barium molybdate (BaMo04)probably formed in some of the barium chloride experiments. Given the low solubility of barite (BaS04), dissolved barium was limited by precipitation of barite. Barium concentrations apparently had to rise above the concentration imposed by barite precipitation (16 g/L barium is required to combine with all the dissolved sulfate in the synthetic pore fluid) before barium molybdate or barium uranium solids could form. Contaminant Extraction by Sorption. In some of the experiments, uranium and molybdenum exhibited a relationship between extraction and pH that is similar to adsorption. Most adsorption processes have a rather narrow pH range in which adsorption increases from almost 0% to almost 100% (33). Humic compounds are responsible for the accumulation of uranium in many sedimentary-type ore deposits. Apparently, uranium bonds to oxygen-bearing functional groups such as carboxyl (COOH), carbonyl (COO), and hydroxyl (OH) (16). Sawdust, peat, lignite, and subbituminous coal were sources of humic material in this study. In our experiments that used peat or sawdust, filtered solutions were dark yellow, indicating the presence of dissolved humic acid. Because E,, values were relatively high, the extraction of uranium and molybdenum by peat and lignite was probably the result of sorption on solid humic materials rather than the result of precipitation of reduced phases. Maximum uranium sorption on peat occurred at about pH 6 (Figure 3a), a value similar to that observed for lignite and humic acid by Rozhkova et al. (34). Little is known about sorption of molybdenum on humic compounds. In our experiments with peat, molybdenum was extracted most efficiently at low pH (Figure 3b), which suggests that negatively charged molybdate anions were adsorbed to positively charged surfaces. Uranium and molybdenum extraction by titanium oxide and uranium extraction by phosphate also probably occurred by sorption because these phases have low solubilities and are effective sorbents. The relationship of uranium extraction to the amount of titanium oxide is nearly linear, which suggests a linear sorption isotherm. Extractions of uranium and molybdenum by titanium oxide are strongly pH dependent (Figure 3a and b), which also suggests sorption. The relatively low uranium extraction efficiency of titanium hydroxide was probably due to the small surface area and the small amount used (Figure 1). The results from the experiments with montmorillonite and clinoptilolite with synthetic pore fluid (Figure 1)suggest low sorption potential for both uranium and molybdenum. Clays are known to be considerably less effective for uranium adsorption than peat and iron oxyhydroxides (35). Site complexation models are useful in predicting adsorption on oxide surfaces. We applied a site complexation model developed by Davis et al. (36)to quantify the adsorption of uranium and molybdenum to ferric oxyhydroxide. In this model, two layers exist between the diffuse layer and the oxide surface. Protonation and deprotonation reactions occur in the inner layer, and all other specific adsorptions occur in the outer layer. Surfacmite complexation reactions are modeled using equilibrium
Table V. Adsorption Model Parameters reaction
log K
SOH = SO- + Ht SOH + H+ = SOH2+ SOH + Na' = SO-Na + H+ SOH + H+ + NO3- = SOH2-N03 SOH + UOzz++ 3H20 = SOH-UO,(OH), + 2H+ SOH + 3UO?+ + 8H20 = SOH-(U02)s(OH)8z-+ 8H+ 2SOH + M004~+ 2H+ = (SOHz)z-M004
-11.10" 4.80° -9.30"
7.w -3.00b -31.30b
90
--
20.00c
from Hsi and Langmuir (22) for ferric oxyhydroxide formula assumed for ferric oxyhydroxide is Fe(OH)3 capacitance of inner layer, 130 pF/cm2 capacitance of outer layer, 20 pF/cm2 surface area of ferric oxyhydroxide, 700 m2/g site concentration, 20 sites/nm2
ln
P
From Hsi and Langmuir (22). *Adsorption species are from a goethite model by Tripathi (13); log K's were modified. Determined by curve fitting. (I
mass action expressions similar to those used for aqueous speciation. However, the surfacesite complexation model also accounts for energy changes resulting from variable surface charge densities. The parameters used in our adsorption model are provided in Table V. Uranium adsorption was modeled using the same two adsorbed uranyl-hydroxyl complexes used by Tripathi (13) to model adsorption on goethite. We modified the intrinsic equilibrium constants to account for the stronger adsorption on ferric oxyhydroxide than on goethite. Equilibrium constants for uranyl hydroxyl polymers, uranyl carbonates, and all other dissolved complexes are provided in the supplementary material. The uranium model was calibrated using data from Hsi and Langmuir (22). The model matches the data reasonably well for low-carbonate systems but fails to adequately match the high-carbonate (0.01mol/L) data (Figure 6a). Hsi and Langmuir (22)were able to fit the high-carbonate data better by incorporating adsorbed uranyl carbonate complexes, but their model failed to adequately fit the low-carbonate data without readjustment of the equilibrium constants. Molybdenum adsorption was modeled using a single, adsorbed bidentate molybdenum complex (Table V). All dissolved molybdenum was modeled as MOO:- or HMo04. Equilibrium constants for these and all other dissolved complexes are provided in the supplementary material. The model was calibrated against a series of molybdenum adsorption experiments (Figure 6b). When the bidentate complex was used, the model fit the data slightly better than with only monodentate complexes but the slope was not as steep as the observed data slope. The uranium and molybdenum adsorption models were applied to the synthetic pore fluid experiments (Figure 4a and b). No surface complexes other than those given in Table V were included. Equilibrium constants for all dissolved species are provided in the supplementary material. The model predicted slightly more adsorption of uranium near pH 4 and slightly less in the presence of aqueous uranyl-carbonate complexes near pH 7.5 (Figure 4a). It also predicted stronger molybdenum adsorption above pH 7 (Figure 4b). Considering that no additional cation or anion competition effects were included, the model results were close enough to the observed data to provide confidence that, with some refinements, this type of modeling might be applicable to complex solutions. Implications for a Chemical Barrier at a Uranium Mill Tailing Repository. Federal regulations mandate that uranium mill tailings disposal sites be constructed so
4
5
6
7
6
9
10
PH
Flgure 6. Callbration of adsorption models for U and Mo on ferrlc oxyhydroxide. Model parameters are given in Table V. SolM curves are model calculated; points are experimental data. (a) U adsorptlon data from Hsl and Langmulr (ZZ), [ferric oxyhydroxide] = 1 g/L, [U] = 2.4 mg/L, total dlssolved carbon 0, and lo-*mol/L; (b) Mo adsorption, [ferric oxyhydroxide] = 1.15 g/L [as Fe(OH),], [Mol = 10.2 mg/L, [NaN03] = 0.1 m .
that groundwater is protected for at least 200 years ( I ) . The task of protecting groundwater is most difficult during the initial time period (perhaps the first 20 years) because the tailings contain greater quantities of water than in succeeding years. This water (we will refer to it as transient water) comes from existing pore water, dust control, and precipitation during repository site construction. The amount of transient water the tailings will contain at the Monticello repository is estimated to be 5 X lo6 L. Rigorous cover design and a semiarid climate will cause a significant reduction in pore water flow from the tailings over time. Therefore, a chemical barrier that effectively controls contaminant migration during the transient water stage, when flow rates are highest, would be beneficial. Precipitants have an advantage compared with sorbenta because they exert more control on the composition of the aqueous solution. For example, the presence of hydrated lime causes the pH to exceed 12 regardless of the initial fluid composition. In contrast, sorbents are only effective in restricted ranges of pH and are more sensitive to chemical perturbations. Precipitants have a disadvantage because their dissolution can cause elevated concentrations of ions (for example, the increase in pH from hydrated lime and in barium from barium chloride) that may be undesirable. Environ. Sci. Technol., Vol. 26, No. 10, 1992
1029
Table VI. Estimated Amounts of Additives Raquired to Treat the Transient Water at the Monticello Site"
material CAOH2 PEAT PEAT (pH 2-76)' FES04 BACL2 FE203 FE203 (DH5.86)' TI02 (pH 6.22)c'
amt resultant amt resultant required, U concn, required, Mo concn, tons mg/L tons mg/L 188 2503 na na 12515 na 510010d llOOIOd
0.07
0.32 0.02
0.75
0.09
nab 10012 52f103~ 1226 12515 12515 510010d 510010d
0.05
0.22 0.05 0.05 0.05 0.01 0.10
See Table I1 for description of materials. *Not applicable because the material was ineffective as an extractor. 'pH was adjusted to the given value. dCalculations are from a single experiment and, thus, these estimates are maximums.
The amounts of various materials needed to control contaminant migration from the 5 X 108 L of transient pore fluid were estimated from the data shown in Figures 2 and 3. The estimates, presented in Table VI, assume that the pore fluid at the site is similar to the synthetic pore fluid and no reactions occur that did not occur in the laboratory experiments. These estimates suggest that hydrated lime and ferrous sulfate would be a suitable combination for a chemical barrier, but there are other considerations. A layer of ferrous sulfate could be constructed over a layer of hydrated lime; however, in an open system, the ferrous sulfate dissolves and mobilized ferrous iron reprecipitates in the lime layer as ferrous hydroxide (37). Although the two-layer system effectively controls uranium and molybdenum releases, the chemical mechanisms are related to oxidation-reduction states rather than to the nonredox mechanisms suggested above. Some of the sorbents that extracted uranium and molybdenum effectively over a pH range were relatively ineffective outside that range (Figure 3). A chemical barrier that relies on sorbents will be functional only if the pH remains in the effective range. Ferric oxyhydroxide (formed from ferric nitrate, ferric chloride, or ferric sulfate), hematite, and titanium oxide are particularly attractive sorbents for a chemical barrier if a suitable pHbuffered system can be designed. Other considerationsin designing a chemical barrier that need investigation include the effect of dissolved organic acids generated from humic materials, the effect of high pH and high calcium concentrations produced by hydrated lime, the degeneration of hydrated lime by reaction with silicates to form calc-silicate minerals and with carbonate to form calcite, the effect of chemical changes on permeability, the decrease in surface area of ferric oxyhydroxides by recrystallization, and the release of potential undesirable ions such as barium. Many of the materials used in this study are available in industrial quantities. For example, iron oxide is produced as a waste material in the production of titanium oxide in the pigment industry; iron oxides, coal, lignite, phosphate, titanium oxide, and lime occur naturally and are mined; and ferrous sulfate is manufactured for fertilizer.
Conclusions Although the results may differ quantitatively, the same processes that act to sorb or precipitate uranium and molybdenum in simple aqueous solutions are apparently operative in more complex solutions resembling uranium mill tailings pore fluids. Peat, hematite, ferric oxyhydroxide, phosphate, and titanium oxide lower concen1930 Envlron. Scl. Technol., Vol. 26, No. 10, 1992
trations of both uranium and molybdenum but are sensitive to pH fluctuations. Hydrated lime, some types of fly ash, and phosphate extract uranium. Soluble ferrous materials extract molybdenum. Barium chloride extracts both uranium and molybdenum, and calcium chloride extracts molybdenum but requires large barium or calcium concentrations. No single additive extracts both uranium and molybdenum over a wide range of pH values. Thus, a chemical barrier may require two or more materials. Contaminant extraction is either by sorption or by precipitation. Sorbents can extract contaminants without affecting much else in the chemical system but are sensitive to solution variables such as pH. Precipitants need to produce elevated concentrations of one or more ions to be effective, which may produce unanticipated, undesirable effects. A chemical barrier containing hydrated lime will lower uranium concentrations in pore fluid that passes through the barrier to levels that will meet regulatory guidelines at compliance points. It is likely such a barrier will also extract metal contaminants (particularly Pb, Cu, Zn, and Cr) by causing the formation of low-solubility hydroxide minerals. For the barrier to remain effective, the hydrated lime must not be depleted, and hydrologic conditions must be such that the contaminated-tailings pore fluid makes contact with the barrier.
Acknowledgments We gratefully acknowledge Dave Emilia, Brian Wilson, and Hal Langner for their support throughout the project and Vijay Tripathi for many discussions of the chemical systems. We thank Sarah Morris for providing technical assistance.
Supplementary Material Available Thermodynamic database used for calculations in tables and figures (3 pages) w i l l appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper or microfiche (105 X 148 mm, 24x reduction, negatives) may be obtained from Microforms Office, American Chemical Society, 1155 16th St., N.W., Washington, DC 20036. Full bibliographic citation Qournal, title of article, authors' names, inclusive pagination, volume number, and issue number) and prepayment, check or money order for $10.00 for photocopy ($12.00 foreign) or $10.00 for microfiche ($11.00 foreign), are required. Canadian residents should add 7% GST. Registry No. FeS04, 7720-78-7; GEOTH, 1310-14-1; TiO,, 13463-67-7;Fe203,1317-60-8; CaC12, 10043-52-4;BaCl,, 10361-37-2; PHOS2, 12167-74-7; U, 7440-61-1; Mo, 7439-98-7.
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Received for review December 18, 1991. Revised manuscript received May 22,1992. Accepted June 15,1992. The work was funded by the U.S. Department of Energy Office of Environmental Restoration and Waste Management (DOE Contract DE-AC04-86ID12584), Monticello Remedial Action Project, and Grand Junction Projects Office's Exploratory Research and Development Program.
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