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EPA Region 5 Records Ctr.
225262
Geochemical Remediation of Extremely AlkalineGround-Water Discharges in the Lake Calumet Region
by
G.S. RoadcapW.R. Kelly
M.L. Machesky
Prepared for theUnited State Environmental Protection Agency
December 2000
Illinois State Water SurveyGround-Water Section
A Division of the Illinois Department of Natural Resources
Table of Contents
Abstract 1
Problem Definition/Background 2
Methods 4Sites 4Laboratory Experiments 5Field Experiments 6Quality Assurance and Control Summary 7
Results 8Chemical Characteristics of Site A 8Chemical Characteristics of Site B 14Experimental Results 15
pH 15Metals 20Toxicity 24
Natural Remediation at Site C 26
Conclusions 34
References 34
:•*»
List of Tables
Table 1. Water Quality Characteristics 10Table 2. Changes in concentrations of selected metals from start to end of
experiments as a function of treatment type 24Table 3. Water Quality Characteristics of Site C 28
List of Figures
I Figure 1. Conceptual ground-water flow system 3Figure 2. Location of field sites 4
|| Figure 3. Photographs of remediation experiments: a) air sparging, b) dolomiteis addition 6
Figure 4. Photographs of Site A: a) typical conditions, b) visible spring shown!*. by the arrow 9ii Figure 5. Photograph of Site B 14
Figure 6. Reduction in pH with CO2 treatment 16' Figure 7. Reduction in pH with air treatment 16'' Figure 8. Reduction in pH with acid treatment 17„_ Figure 9. Reduction in pH with dolomite treatment 17• Figure 10. Change in carbonate species concentration versus pH during CO2 gas
sparging of samples from Site A 18,, Figure 11. Initial and final aqueous silicon concentrations for different treatments 21j: Figure 12. Initial and final aqueous aluminum concentrations for different treatments . . . . 21
Figure 13. Initial and final aqueous calcium concentrations for different treatments 22Figure 14. Initial and final aqueous magnesium concentrations for different treatments ... 22
I Figure 15. MircoTox™ toxicity responses for Site A 25Figure 16. MircoTox™ toxicity responses for Site B 25
Ty Figure 17. Photographs of Site C: a) on sampling date, b) under high water conditions ... 27ji Figure 18. Sampling location and pH distribution at Site C 27
Figure 19. Calcium distribution along the flow path from the spring 32*;• Figure 20. Magnesium and silicon distribution along the flow path from the surface
runoff 32Figure 21. Distribution of metals in the sediment 33
11
Abstract
Experimental results show that in situ remediation of ponds and wetlands contaminatedby extremely alkaline ground-water discharges is possible using aeration techniques. Alkaline
-r ground-water conditions in the Lake Calumet region of southeast Chicago and northwest Indiana'. are caused by the dissolution of steel slag wastes used to fill in wetlands. These alkaline
conditions persist in ponds and wetlands dominated by ground-water inflow. Water and»t sediment samples for remediation experiments were collected from three of these sites. Site A
was a roadside ditch and was characterized by a pH value of 12.16, no dissolved oxygen, andhigh organic carbon content. Site B was a shallow pond and was characterized by a pH value
>"• of 11.16, saturated dissolved oxygen and low organic carbon content. Site C was a calcium-richi, spring with a pH value of 12.19 that dropped to 7.77 downstream due to mixing with the
atmosphere and other surface waters. Large amounts metal-rich calcite precipitate were presentf| at all three sites and the toxicities, as measured by bioassays, were 100 percent.U
Remediation strategies tested on subsamples from Sites A and B included four methodsItfcg. aimed at lowering the pH: sparging with carbon dioxide, sparging with air, adding acid, and^ adding locally-quarried crushed dolomite. The experimental results for the two sites were similar
except that the Site A water required significantly more time and amount of treatment. The* ability for the Site A water to neutralize the added remediation agents, or buffering capacity, was
roughly 10 times that of Site B. The carbon dioxide sparging effectively reduced the pH valueto 6.3 because of its reaction with water to form carbonic acid. When sediment from the siteswas added there was less pH reduction because of dissolution of calcite in the sediment. Airsparging also reduced the pH but at a rate that was two to three orders of magnitude slowerbecause of the lower partial pressure of carbon dioxide in air. Adding the sediment did notsubstantially affect the air sparging results because the pH did not drop below 8.3, the point atwhich calcite could start to dissolve. The addition of hydrochloric acid also rapidly decreasedthe pH of the water and when sediment was included there was also buffering by calcitedissolution. The addition of crushed dolomite rock reduced the pH very slowly through thedissolution of quartz grains within the crushed rock. The MicroTox™ toxicity was reducedsignificantly by some of the remediation methods. The air sparging method was the mosteffective method at reducing toxicity at Site B because it reduced the pH without dissolving thesediment. The toxicity results from Site A also showed a reduction; however, interpreting theresults was difficult because of color interferences from the reddish organic material in the water.
The alkaline ground-water discharge at Site C was naturally attenuated by enhancedinteraction with CO2 in the atmosphere during turbulent flow and by mixing with neutral pHsurface waters. Along the flow path of the discharge the pH dropped from 12.19 to 7.77 and thetoxicity to Ceriodaphnia dubia dropped from 100% to 0%. As the pH decreased, carbonate ionswere introduced which then reacted with the calcium in the discharge to form calcite. Thesecalcite sediments contained a large amount of coprecipiated metals.
Problem Definition/Background
Contamination of ground water and surface water runoff and disruption of natural flowin the wetlands of the Calumet region have adversely affected wetland quality as vital habitats
r- for fish, birds, and other wildlife in the southern end of Lake Michigan. While the point-sourceproblems of major rivers in the region have been or are being addressed, the non-point sourcesof contamination continue to affect many small wetlands and tributary streams. These small
*•* ecosystems are an important part of the limited natural habitat in this region of the Great Lakes.Contaminating source materials are widespread and of variable composition. Because removalwould be prohibitively expensive and disruptive, cleanup and restoration alternatives that meetsite-specific conditions need to be found.
) ;
A major source of contamination is the leachate produced from the vast accumulations|f of fill that were used to raise the low-lying wetland areas high enough to be developed as*i industrial or residential property. Due. to the haphazard, truckload-by-truckload way in which
the fill was dumped, the lithologic and hydraulic character of the fill is extremely variable and_, cannot be quantified into a single description for even short horizontal and vertical distances.
- The result is a man-made aquifer that has wildly varying physical, hydrogeological, and chemicalproperties. The main types of fill in the region are slag wastes from nearby steel mills, dredged
I; material from the Calumet River system, demolition debris, municipal wastes, and other" industrial wastes (Colton, 1985; Roadcap and Kelly, 1994; Kay etal., 1996). The native geology— consist of lacustrine silts and beach sands overly ing relatively impermeable glacial clay-rich till.i The sand was often quarried out before an area was filled in. At the original surface is a thin
layer of wetland soil.»«I i Figure 1 shows a conceptual model of a ground-water flow system that is discharging to
a wetland. The discharge can occur along a seepage face (Roadcap et al., 1999) or in a set of;. discrete springs that can have a somewhat regular spacing along the banks of a wetland (Duwal,j> 1994). The springs are the result of preferential ground-water flow paths that develop in
macropore-forming heterogeneities within the coarsest material of the fill. An example ofTT macropore formation is a dump load of refractory bricks that is surrounded by dump loads ofti finer-grained slag. The springs can be visually identified by the presence of channels hi the
wetland bottom leading away from the bank, flowing water during period of low surface waterlevels, patches of open water during a winter freeze, or the presence of white calcite precipitate.
The leachate discharging from the slag fill is adversely affecting the water quality ofmany of the wetlands in the Calumet region. The slag contains lime and other high-temperaturecalcium silicate minerals such as akermanite that react with water and drive up the pH of theground water to above 12 pH units by consuming hydrogen ions. The slags can also contain up
'* Akermanite (Ca^gSiA) + 6H+ => 3H2O + 2Ca2+ + Mg2+ + 2SiO2(aq) (1)
tto 50% iron and other metals, such as manganese, vanadium, and chromium, in the metallic
Infiltration ofwater and air
Ground-water flow path
Water table
Calciteprecipitation
Native soil
Native sand
Figure 1. Conceptual ground-water flow system.
form. These metals are dissolved and reprecipitated as oxides, hydroxides, carbonates or otherminerals or become coprecipitated with these minerals. The resulting contaminated groundwater can be extremely alkaline (pH >12) and low in dissolved oxygen while having highconcentrations of ammonium, dissolved solids, and heavy metals (Roadcap and Kelly, 1994).Because many of the wetlands and tributary streams in this southern Lake Michigan area arehydraulically dominated by ground-water inflow, waters in the wetlands can also be extremelyalkaline with very little dissolved oxygen or carbon dioxide even though they are exposed to theatmosphere. Further, as carbon dioxide is introduced either locally or downstream, metal-richminerals precipitate creating contaminated sediments. In sediment analyses from ten wetlandsin this area, concentrations of metals were elevated above background levels (ISWS/IEPA,unpublished data). Seven sites have highly elevated concentrations of chromium, lead, and/orzinc and four sites have highly elevated concentrations of cadmium and/or copper.
Wetlands of the region provide an important habitat for many plant and wildlife specieswhich are on state endangered and threatened lists, including the Black-crowned Night Heronand the Great Egret. However, the high quality wetland habitats and the water quality of LakeMichigan and the Calumet River System are adversely impacted by interaction withcontaminated wetlands, especially the highly alkaline wetlands. These alkaline conditions canpersist in ponds and wetlands dominated by ground-water inflow and can be extremely toxic toaquatic life. Fish kills are common when variations in ground-water inflows and wetlandoutflows produce toxic conditions. Toxicity testing at the ten sediment sites indicated that four
of the sites were significantly toxic and four sites were marginally toxic (Mike Henebry, 10/7/94written comm. with ISWS). The wetland vegetation is greatly impacted by the high dissolvedsolids of the water which gives more salt-tolerant invasive species a competitive advantage overthe native vegetation. Further, natural flow patterns have been disrupted by railroad and streetcauseways, reducing or eliminating the supply of high quality waters to the wetlands.
The purpose of this study was to examine whether in situ remediation of ponds andwetlands contaminated by extremely alkaline ground-water discharges is possible. Toaccomplish this, surface-water and sediment samples were taken from two sites for remediationexperiments in the laboratory. The physical/chemical remediation experiments were used toassess how specific treatments affected the water chemistry and toxicity. A pilot field test wasattempted at one of two sampling sites and in addition, a third site also was examined whereremediation was observed to be occurring naturally.
Methods
Field Sites
The field sites (figure 2) were selected for an examination of their remediation potential.The waters at Sites A and B consistently had high pH values. The water coming from a springat Site C had a high pH, however, the pH quickly dropped with distance from the spring. SitesA and C are in a roadside ditch along Doty Avenue West, the west frontage road for the Bishop
Figure 2. Location of field sites.
Ford Expressway (194). The ditch is directly connected to Lake Calumet through a culvertunderneath the expressway and is vegetated primarily with the invasive wetland speciesPhragmites australis (common reed). Site B is a shallow pond near the shore of Lake Calumetthat was dug into a slag pile presumably to control the local drainage. The pond is very sparselyvegetated by an unidentified leafy plant.
Laboratory Experiments
Laboratory experiments were performed on samples collected from Sites A and B.Twenty liters (L) of water and 10 kilograms (kg) of bottom sediment were collected from eachtest site. Water samples were taken using a peristaltic pump and stored on ice in 10 2-Lpolyethylene bottles. Sediment samples were collected and stored on ice hi 2-L polyethylenebottles. Water samples were taken from each site in the field for chemical analysis of a relativelyfull suite of anions, cations, and nutrients. These parameters are listed hi Table 1. Prior tocollection of water samples, the following chemical parameters were measured in the field: pH,dissolved oxygen (DO), temperature, oxidation reduction potential (Eh), and specificconductance. Water samples were also collected for toxicity tests using the Microtox™ method.Sediment samples were digested and analyzed for metals. Samples of plant tissue were alsodigested and analyzed for metals as an exploratory measure to see if plants were affecting the fateof the metals in the water. In the laboratory, the 10 2-L bottles of water from each site werecombined into a collapsible, air-tight container and resampled for full chemical analysis todetermine if any chemical changes had occurred during transport or during the two to four weeksof storage. Potential changes include mineral dissolution or precipitation due to changingtemperature conditions, biological activity, or either exchange with or isolation from theatmospheric gases. All chemical analyses were performed at the ISWS analytical laboratory.
Water and sediment samples from each site were subjected to several sets of remediationexperiments in the laboratory. Remediation strategies tested on subsamples included fourmethods aimed at lowering the pH:
• sparging with carbon dioxide (CO2)• sparging with air• addition of hydrochloric acid (HC1)• addition of crushed dolomite aggregate
For each treatment, two separate 1-L flasks were prepared; one containing 900 milliliters (mL)of ex situ surface water only and one containing 900 mL of ex situ surface water and 100 grams(g) of ex situ sediment. For each site, two additional flasks were used as controls. The controlflasks were not treated.
The pH was monitored continuously during each experiment. After the pH value reacheda steady state, the experiment was stopped and samples were taken for complete inorganicanalysis of water chemistry. Samples were also collected for MicroTox™ toxicity screening tests
.
y
to determine if the remediation schemes had a positive impact on the quality of the aquatichabitat. For the air and CO2 sparging experiments, the water was sparged with an aspirator ata constant rate until the pH had stabilized. The CO2 was sparged from a gas cylinder and the airwas sparged from the compressed air line in the ISWS labs using a moisture trap (figure 3a). Forthe acid addition experiments, 0.1 normal HC1 was titrated into the flasks which werecontinuously stirred. The titration continued until pH reached 4.0, the typical endpoint foralkalinity titrations (American Public Health Association et al., 1992), or until enough sedimenthad dissolved to neutralize additional acid. For the dolomite addition experiments, 35 g of fine-grained crushed rock that had been passed through a 0.5 millimeter (mm) sieve was added. Amagnetic stir bar was added so the flasks could be continuously stirred on a mechanical stir plate(figure 3b). The dolomite aggregate was collected from a spoil pile of cuttings from the DeepTunnel constructed along the Little Calumet River. No sulfide or other dark-colored mineralswere visible in the samples..
Figure 3. Photographs of remediation experiments: a) air sparging, b) dolomite addition
Field Experiments
A pilot test was designed to test the performance of the most successful remediationagents in a field setting. In preparation for a pilot test, the field conditions at Sites A and B andseveral other high pH springs were periodically monitored. During the course of this monitoring,the natural remediation occurring at Site C was discovered and six samples were collected forwater quality and sediment analysis.
Site A was selected as the test site because it was the only site in the region where thehigh pH springs and seeps had not dried up due to the prolonged dry period in the summer andfall of 1999. Site B would have been preferable to Site A because of the lesser amount ofrequired treatment, as discussed later in the Results section. Site A was further complicated byan attempt to clean up the ditch initiated by the Sherwin-Williams Corporation, located west ofthe site. This effort included removing the vegetation, reshaping the ditch, and adding anengineered berm down the center of the ditch that was intended to isolate the ground waterdischarge from surface runoff. The ground water would then be collected and treated offsite.
Due to flow through the berm and a lack of any surface water runoff, the chemistry in the ditchappeared to be unaffected by the presence of the berm, the reshaping of the ditch, or the removalof the vegetation. The pilot test was conducted before the system to pump water from the ditchwas operational.
Based on the laboratory experimental results discussed later, air sparging was selectedas the remediation scheme to test. Two aerators were placed hi a small section of the ditch onthe surface water side of the berm at its downstream end. Two small berms were added acrossthe ditch to separate a 4-foot by 15-foot section which would be chemically and hydraulicallyisolated from the rest of the ditch. The water depth in the treated section averaged approximatelythree inches. The aerators had a small submersible pump and a vertical discharge pipe that stuck8 inches above the water and sprayed the water downward. They were designed to aerate coolersand buckets full offish. Unfortunately, some unforseen episodic ground-water inflows into theditch inundated the treatment area and ruined the pilot test. Because the timing and amount ofthese inflows could not be predicted, it was decided not to restart the pilot test and place moreeffort on analyzing Site C.
Quality Assurance and Control Summary
A quality assurance project plan or QAPP (Locke, 1998, unpublished) was developed forthis project and submitted to the USEPA - Region 5 Office. The overall quality assuranceobjective for this study was to develop and implement field sampling and experimentallaboratory procedures that would help accomplish the study goals of testing remediationalternatives for contaminated ground-water discharges. The QAPP specifically identified:measurements to be made in the field and laboratory, all methods of analysis to be used, fieldprocedures for sampling, and methods to assess the quality of data from water sampling.Methods described in the QAPP are summarized in this report when appropriate, and additionaldata are presented below.
A total of 26 sample sets (8 fields sample, 18 experimental results) were collectedbetween October 29, 1998, and April 20, 1999, for water matrices. Methods described in theQuality Assurance and Field Methods Manual (IEPA, 1987) were used as a basis for thesampling procedures. Methods used for sample analysis hi the laboratory are cited from eitherthe USEPA or the American Public Health Association and are noted with the method detectionlimit (MDL) for each analyte in Table 1. Ten noncritical sample sets of sediment, slag, andplants were added to the project to better characterize the chemical processes that are occurring.The MDLs for the anlytes in these ten sets are listed in Tables land 2 if the measureconcentration was below the MDL. All samples were analyzed within the sample-holding tunesspecified in the QAPP.
Data precision was assessed by the collection and analysis of one duplicate field sampleand four duplicate laboratory samples. The relative percent difference (RPD) was calculated foreach sample and its duplicate. For concentrations (in the original sample) at or above ten times
the MDL, an RPD of up to 20 percent between a duplicate and original sample analysis wasacceptable. For concentrations lower than ten times the MDL, an RPD exceeding 100 percentwas allowed. Similar guidelines have been used as a general measure of analytical performanceby recent ISWS projects (e.g., Roadcap et al. 1999, Locke et al. 1997). This general measureallows for higher RPDs when concentrations are closer to the MDL. All of the water analyseswere acceptable except six; two aluminum, one copper, and three zinc analyses. Five of thesesix exceptions were at levels below their minimum concentration of concern (1 mg/L). Theremaining exception was a zinc analysis with a RPD of 146 percent. These anomalies may bedue to the presence of metal-rich colloids that are unevenly distributed in the sampled watercolumn.
Data accuracy was assessed by the collection and analysis of four sets of lab matrixspikes to test laboratory procedures. Recoveries between 50 and 150 percent were consideredacceptable. All of the matrix spike recoveries were acceptable except seven. Four exceptionswere for calcium, iron, silicon and sulfate analyses where the matrix concentration was 10 to1000 times less than the sample concentration. The remaining three exceptions occurred at lowconcentrations and included two zinc analyses from water samples, and one manganese analysisfrom a sediment sample. Field blanks to test sampling procedures were not collected due to anoversight with the unexpected need to sample the mixing waters at Site C.
Data completeness is a measure of the amount of valid data obtained from a measurementsystem compared to the amount that was expected to be obtained under normal conditions. Datacompleteness of 90 percent was defined as acceptable to meet project goals. A percentcompleteness (%C) of 96.9 was estimated for the field and laboratory chemistry data. Fluorideand nitrate ion chromatography analyses were the most problematic and accounted for nearly halfof the invalid results. The high sulfate concentrations (e.g. 1900 mg/L) at Site A made peakseparation difficult and fluoride and nitrate were not determinable.
Results
Chemical Characteristics of Site A
Site A (figure 4) is a roadside wetland/ditch characterized by a pH value of 12.2, nodissolved oxygen, and high non-volatile organic carbon (NVOC) content (Table 1). Organiccontaminants such as toluene and cresol compounds were also present. The site normally hasa dark reddish brown color (figure 4a) due to a large amount of humic compounds. When theditch is clearer, the reddish black ground-water seeps are easily visible (figure 4b). The toxiciry,as measured by the Microtox procedure, was 100 percent mortality.
The total dissolved solids (TDS) value of 5,118 mg/L is very high for a surface water andreflects a significant amount of ground-water discharge. Nearby monitoring wells have TDSvalues that range between 1,781 mg/L and 7,060 mg/L (Roadcap and Kelly, 1994) while theCalumet River has a TDS value of less than 400 mg/L (Roadcap et al., 1999). The dominant
r
Figure 4. Photographs of Site A: a) typical conditions, b) visible spring shown by the arrow.
cation was sodium at 1,788 mg/L (77 millimoles per liter or mmolal) and the dominant anionswere sulfate (SO4
2~) at 1,883 mg/L (19.7 mmolal), hydroxide ion (OH"), and carbonate ion(CO3
2~). The alkalinity, defined as the ability of a solution to neutralize acid, was 1,953 mg/L(reported as CaCO3; 19.5 mmolal). At a pH of 12.2, the hydroxide ions accounted for much ofthis alkalinity. A geochemical model of the alkalinity titration suggested that the OH"concentration was 342 mg/L (20.1 mmolal) and the carbonate concentration was 550 mg/L (9.17mmolal). The Geochemist's Workbench™ software (Bethke, 1998) was used for all of thegeochemical modeling.
Because the total sulfur analysis of 740 mg/L (report as elemental sulfur; 23.1 mmolal)cannot all be accounted for by the 1,883 mg/L (19.7 mmolal) of sulfate, there may have been asmuch as 112 mg/L (3.5 mmolal) of sulfur present in other forms such as organic sulfur,thiosulfate, or sulfide. For both sulfate and sulfide to be present in the water, the amount ofoxygen would need to be extremely low. Because oxygen can enter the water at the surfaceditch, the sulfur system is likely controlled by the sulfide oxidation reaction.
HS" + 2O2 * SO42" + H+
(2)
This reaction produces protons (H+) which will react with the hydroxide ions to lower the pH.Oxidation of the organic carbon (represented as CH2O) will compete with the sulfide for theincoming oxygen from the surface.
CH2O + O2 C02 + H2O (3)
This reaction produces CO2 which will react with water to form additional protons and lower thePH.
The introduction of oxygen and carbon dioxide into the water contributes to the spatiallyvariable chemistry along the ditch with the field-measured pH value dropping by more than 1
9
Table 1. Water Qaulity Characteristics
SampleMDL
Site AFieldLabControlC02CO2 w/ sedimentAir*Air w/ sedimentAcidAcid w/ sedimentSedimentPlants
SiteBField*LabControlControl w/ sed.*CO2C02 w/ sedimentDolomiteDolomite w/ sed.AirAir w/ sedimentAcid w/ sediment*SlagSedimentPlant LeavesPlant StemsPlant Roots
PH
12.1612.29
11.936.19
6.29
9.439.543.867.32
-
-
11.1611.1610.3810.866.22
7.10
9.099.24
8.138.237.36
---
--
Al0.042
0.8090.6800.581
0.5630.8531.2431.122
0.8610.14720045407
0.3330.4230.3360.2070.3380.177
bdbd
0.4020.0940.08199972662282280
1113
As
0.13
0.71
0.96
0.78
0.800.99
0.800.800.750.6057.4
<16.4
bdbd
bdbdbd
bdbdbdbdbdbd
<16.4<16.4<16.4<16.4<16.4
B0.04
0.75
0.90
0.810.87
0.840.630.58
0.750.8159.514.0
bdbd
0.130.55
bd
0.050.33
1.19
0.05bd
bd1467.13
54.235.5
42.5
Ba0.004
0.0410.061
0.0600.133
0.4830.0470.0650.0440.281162530.1
0.0320.0280.0240.0450.0280.044
0.0060.0200.0230.0500.122268145
25.316.8
26.8
Be0.004
bdbd
bdbdbdbdbd
0.0060.006
--
bdbdbdbdbd
0.007
bdbdbdbd
0.006----.
Ca0.10
10.911.610.755.52013.574.859.97159»»4994
32.6
27.84.1510.928.0
39.37.64
6.1015.5
16.6164»»»»
221811057234520
Cd
0.016
bdbd
bdbdbdbdbdbdbd
4.552.26
bdbdbdbdbdbdbdbdbdbdbd
23.2<2.022.22<2.022.30
Co
0.011
bd
bd0.020
bdbdbdbdbdbd
12.33.58
bdbdbdbdbdbdbdbdbdbdbd
32.9<1.3910.54.758.79
Cr0.009
0.0240.0310.044
0.0280.0440.0280.0120.0120.02263.32.11
0.0160.0210.0200.0200.017
0.0190.0230.0120.018
0.0150.0231047
12.07.09
8.90
113
Cu
0.005
0.0340.0320.0260.0480.0210.1040.2520.0240.00896.814.1
0.0330.0280.0120.0090.0090.0320.015
0.0080.0210.0090.012404
8.26
14.99.37
20.8
Fe
0.009
0.512
0.7260.7851.508
2.8980.3931.1981.3340.38021512591
bd
0.016bdbd
0.0990.306
bd
bd0.0470.0460.033»»
2205451819
10788
Hg0.06
bdbdbdbdbdbdbdbdbd--
bdbdbdbdbdbdbdbdbdbdbd----.
K4.69
23.617.619.619.526.127.516.348.158.4672812524
26.625.422.425.026.220.922.4
22.833.7
20.220.51651
581
297481303611498
Notes: All values in mg/l except sediment, slag, and plants which are in mg/kg* Duplicate sample taken »» Detection range exceeded
- Not measured int Interference from other analytes
bd, < Below detection
MDL Method detection limit
Table 1. Cont.
SampleMDL
Site A
Field
LabControl
CO2CO2 w/ sedimentAir*Air w/ sediment
AcidAcid w/ sedimentSedimentPlants
SiteB
Field*LabControlControl w/ sed.*
CO2CO2 w/ sediment
DolomiteDolomite w/ sed.AirAir w/ sedimentAcid w/ sediment*SlagSedimentPlant LeavesPlant StemsPlant Roots
Li0.012
0.0420.037
0.041
0.044
0.051
0.029bd
0.102
0.12229.1<1.52
0.034
0.023
0.0230.0230.023
0.0360.032
0.025
0.0360.0380.03620.23.995.07
<1.52<1.52
Mg0.06
0.250.34
0.26
5.29
35.2
0.510.542.097.12
271791262
0.110.09
bd0.14
0.08
0.135.46
6.23
0.11
1.150.86»»
1771
1839
18194949
Mn0.003
0.0050.012
0.052
0.181
0.502
0.0100.0200.0700.071
898169
bdbdbdbdbd
0.003
bdbdbd
0.0050.004»»
87.5
189123
2136
Mo0.023
0.047
0.053
0.052
0.052
0.0450.0490.049
0.0350.0353.33<2.91
bd0.033
bdbd
0.029
bdbdbdbdbdbd
139<2.91
<2.91<2.91
<2.91
Na0.10
1788
17131753
1710
1650
1595182317321991
1319311991
12.5
9.5310.610.38.96
8.9315.4
12.5
8.708.72
8.68<12.6<12.6
839125197
Ni0.024
bd0.028
bdbd
0.039
0.0230.0330.042
bd36.0<3.03
bdbdbdbdbdbdbdbdbdbdbd
11.4
3.55
4.08<3.03
2.68
P0.33
0.480.54
0.380.79
1.40
bd0.490.48
1.116322706
bdbdbdbdbdbdbdbdbdbdbd
2802
<41.7
13861037
1356
Pb0.075
bdbdbdbdbdbdbd
0.102bd
430<9.48
bdbdbdbdbdbdbdbdbdbdbd
44.3
<9.48<9.48
<9.48
<9.48
S0.33
740772804779784
652648773
74845804970
5.304.695.396.164.67
5.28
9.6110.6
5.205.064.58
890379
4890
1306
1933
Sb0.37
bd
bd
bdbdbdbdbdbdbd
<46.8<46.8
bdbdbdbdbdbdbdbdbdbdbd
<46.8<46.8
<46.8
<46.8
60.2
Se0.27
bdbdbdbdbdbdbdbdbd
<34.1<34.1
bdbdbdbdbdbdbdbdbdbdbd
<34.1
<34.1
<34.1<34.1
<34.1
Si0.040
77.4
86.4
83.0
86.388.7
59.554.473.9
50.8227686615
1.681.64
2.537.591.611.853.31
9.281.75
3.282.66
40441
5271
5964
1001
7085
Sn0.07
bdbdbdbdbdbdbdbdbd
<8.85<8.85
bdbdbdbdbdbd
0.07bdbdbdbd
<8.85<8.85
9.34
<8.85<8.85
Sr0.006
0.0560.062
0.062
0.1250.419
0.0390.0270.0610.33422112.6
0.131
0.1250.1020.144
0.1240.1530.1240.086
0.1170.1420.239154
282
89.344.9
43.3
Notes: All values in mg/l except sediment, slag, and plants which are in mg/kg
* Duplicate sample taken »» Detection range exceeded
- Not measured int interference from other analytes
bd, < Below detection
MDL Method detection limit
c
Table l.Cont.
SampleMDL
Site AFieldLabControlCO2CO2 w/ sedimentAir*Air w/ sedimentAcidAcid w/ sedimentSedimentPlants
SiteBField*LabControlControl w/ sed.*CO2CO2 w/ sedimentDolomiteDolomite w/ sed.AirAir w/ sedimentAcid w/ sediment*SlagSedimentPlant LeavesPlant StemsPlant Roots
Ti0.005
0.0080.0220.0130.0150.0330.0050.0220.009
bd987206
bdbdbdbdbdbdbdbdbdbdbd
39275673164711023
TI0.40
bdbdbdbdbdbdbdbdbd--
bdbdbdbdbdbdbdbdbdbdbd-----
V0.008
0.1320.1190.1210.1140.1050.0990.1540.0950.06385.511.4
0.010bdbdbd
0.0100.011
bdbd
0.012bdbd
72332.616.224.9111
Zn0.008
2.160.4570.6650.2821.195.962.88
0.2010.52542885.7
0.5791.62
0.2690.2450.9200.3840.4420.4771.181.01
0.64980.525.520867.143.6
NH4-N0.016
2.893.042.942.773.15
0.1310.0812.765.27
--
0.2100.2250.1710.3570.2250.2800.1090.2100.3880.3260.295
-----
F0.100
intintintintintintintintint--
1.020.990.871.520.900.950.870.710.880.740.30
----.
Cl0.31
1.011.000.980.800.800.800.800.100.11
--
3.323.283.497.183.803.544.769.7717.63.51208
----.
NO3-N0.09
intintintintintintintintint--
0.660.630.720.660.630.640.750.700.680.640.52
-----
SO40.510
188319051927190318341955203018101862
--
13.013.013.115.613.213.024.129.413.514.713.0
----.
NVOC Alk IDS TSS TS0.2 2.0 2.0 4.0 6.0
82.4 1953 5118 1721 683982.1 1650..
10556.5 - - - -70.0 . . . .27.6 . . . .107..
4.00 124 169 51 2204.50 116 147 5 152
-.-......
-...
. . . .
Notes: All values in mg/l except sediment, slag, and plants which are in mg/kg* Duplicate sample taken »» Detection range exceeded
- Not measured int interference from other analytes
bd, < Below detection
MDL Method detection limit
Table 1. Cont.
Sample VSS TKN total PMDL 6.0 0.070 0.04
Site AField 190 10.6 2.01Lab - 12.0 0.98ControlCO2C O 2 w / sediment . . .Air*A i r w / sediment . . .AcidAcid w / sediment . . .Sediment . . .Plants
SiteBField* 18 0.96 bdLab bd 0.74 bdControl . . .Control w / sed.* . . .CO2C O 2 w / sediment . . .Dolomite . . .Dolomite w / sed. . . .A i r . . .A i r w / sediment . . .Acid w / sediment* . . .SlagSediment . . .Plant Leaves . . .Plant Stems . . .Plant Roots . . .
DO
0.01.5-
1.30.66.4
5.3
-
--
-
8.8-
5.8
5.8
8.0
7.3
5.8
5.8
7.9
7.8
-
-
--
-
-
Sp Cond Eh(uS/cm) (volts)
6220 -273-200
-
-130-100214164
----
395 25125099.7148
-212218167313310
------
Temp(deg C)
13.639.5-
19.621.022.823.1
----
14.548.022.122.210.312.622.122.022.522.2
------
Notes: All values in mg/l except sediment, slag, and plants which are in mg/kg* Duplicate sample taken »» Detection range exceeded bd, < Below
- Not measured int interference from other analytes MDL Meth
unit away from the springs and at the surface of the water when the water is over 2 feet deep.The fact that such a high pH is maintained in the ditch suggests that the inflow of ground waterto the ditch is continuous and is not significantly diluted by the inflow of surface water andprecipitation.
The relative absence of chloride at 1.01 mg/L is very unusual for a roadside ditch wherelarge amounts of road salt are used for deicing in the winter. Arsenic and zinc were above theIEPA general-use water quality standards of 0.36 mg/L and 1.0 mg/L, respectively.
Because of the potential for the chemistry of the water samples for the experiments tovary with collection, transport and handling, a water sample was taken from the 25 L containerimmediately prior to the experiments. This initial lab sample (Table 1) showed very little changein the maj or constituents while some of the minor constituents at levels near their detection limitsvaried by more than 25 percent.
The bottom of the ditch was covered with metal-rich calcite precipitate overlyinglacustrine clays. In addition to calcium, the sediment contained large amounts of iron, aluminum,magnesium, sodium and silicon. Because the concentrations of dissolved metals in the waterwere very low, the source of metals in the precipitated sediment was likely to be particulatematter. The concentrations of some metals in the precipitate, such as aluminum, iron,magnesium, manganese, and titanium, were significantly higher than the dissolved analyseswould suggest. Cadmium, cobalt, nickel, and lead were not detected in the water but appearedin the sediment. The plants appear to be taking up greater amounts of cadmium, cobalt, copper,manganese, titanium, and zinc than were found in the precipitate. The plant material in theseanalyses were digested with acids in a microwave digester. These plant analyses are not reportedon the commonly used "percent ash basis" which would have shown concentrations roughly 20times greater.
Chemical Characteristics of Site B
Site B (figure 5) is a shallow pond with a pHvalue of 11.2, saturated dissolved oxygen levels and alow non-volatile organic carbon content (Table 1).The toxicity, as measured by the Microtox procedure,was 100 percent mortality. With a TDS of 169 mg/L,this site was 30 times more dilute than Site A.Calcium, potassium and sodium were the principalcations and hydroxide ions, which accounted for mostof the alkalinity, was the principal anion. Dissolvedmetal concentrations were low. Iron was below thedetection limit.
Figure 5. Photograph of Site B.
14
The bottom of the pond was covered by a 2-inch layer of calcite precipitate overlyingslag. The precipitate contained some metals, including iron, manganese, chromium, copper,titanium, vanadium and zinc. Most of these metals were probably derived from ground waterdischarging from the slag that surrounds the pond. The slag contained large amounts of calcium,magnesium, iron, and manganese and significant concentrations of chromium, copper, titanium,and vanadium. The plant roots also appear to be taking up each of these metals, although thetranslocation of these metals to the stems and the leaves appear to be limited.
Experimental Results
The experimental results for the two sites were similar except that the time and amountof treatment were drastically greater at Site A than at Site B. This was because at Site A theinitial pH was 1.0 unit higher and the total dissolved solids content was 30 times higher, whichgreatly increased the alkalinity and the capacity of the solution to buffer the remediation agents.The decrease in pH with treatment is shown on figures (6-9) and the final water chemistry islisted in Table 1.
pH. The carbon dioxide sparging (figure 6) effectively reduced the pH value to 6.2 atboth sites. The pH of these experiments are controlled by the carbonate buffer system (H2CO3/HCO3VCO3
2"), the same system which controls the pH of most natural waters (Drever, 1997).At a pH above approximately 10.8, the CO2 introduced in the experiment reacts with water toform CO3
2", 2 protons which lower the pH, and calcite (CaCO3)
CO2 + H2O => 2H+ + CO32' (pH>~10.8) (4)
Ca2+ + C032' => CaC03 (5)
Figure 10 shows the results of a geochemical model of the experiment at Site A where the partialpressure of CO2 gas in the air that is in contact with the system is increased to 1-atmosphere(atm) as the reaction proceeds from left to right across the figure. The initial partial pressure ofCO2 gas was estimated to be 3.07E-09 atm. When the model reached 3.0E-04 atm, the partialpressure of CO2 gas in the atmosphere (Drever, 1997), the pH was 9.5 at Site A and 8.4 at SiteB. The amount of calcite precipitated is controlled by the amount of dissolved calcium in thewater. The model predicts that almost all of the calcium will be precipitated as calcite, forming26 mg (0.26 mmoles) at Site A and 80 mg (0.80 mmoles) at Site B.
Below a pH of 10.8, significant amounts of bicarbonate ions (HCO3~) are producedbecause a sufficient concentration of H+ is now present to start converting the C03
2" to HCO3".
CO2 + H2O => H+ + HCO3- (-6.3 <pH<-10.8) (6)H+ + CO3
2' => HC03- (7)H+ + CaCO3 => Ca2+ + HCO3' (8)
The generated H+ and HCO3" ions act against each other in the calcite dissolution reaction,
15
.,l f
(Ir
10 20 30 40 50 60 70 80 90 100
Figure 6. Reduction in pH with CO2 treatment.
1000 2000 3000 4000 5000 6000 7000 8000 9000 10000Time (min)
Figure 7. Reduction in pH with air treatment.
16
5.
13
12 -
11 -
10 -
9 -
7 -
6 -
5 -
4 -
3
.A w/ sediment
10 15 20 25 301 N HCI added (ml)
Figure 8. Reduction in pH with acid treatment.
35
5000 10000 15000 20000 25000 30000 35000 40000 45000 50000Time (min)
Figure 9. Reduction in pH with dolomite treatment.
17
-
12 11 10pH
Figure 10. Change in carbonate species concentration versus pHduring CO2 gas sparging of samples from Site A.
however, because the magnitude of the increase in concentration is much greater for H+ than forHCO3", the reaction proceeds and start to dissolve the calcite. This phenomenon is referred toas the common ion effect. The calcite dissolution greatly increases as the pH drops below 8.3,the pH at which calcite is in equilibrium with water that has a CO2 gas partial pressure equal tothat of the atmosphere. Below a pH of 6.2, CO2 (as H2CO3) becomes the dominant aqueouscarbonate species so the introduced gas no longer reacts with water and the pH of theexperimental water stops dropping.
It took 10 times longer sparging with CO2 to reach a final pH for Site A as Site B. In theexperiment where sediment from Site B was added, the pH reduction was 0.9 units less becauseof buffering from dissolution of the calcite in the sediment. A much smaller amount of bufferingwas observed at Site A. The shorter treatment time for the Site A sample with sediment may bedue to the lower starting pH and a lower concentration of hydroxide ions to buffer the solution.After the sparging was stopped on the two Site A samples, the excess CO2 slowly degassed over9 days and the pH increased to 9.5.
Air sparging also effectively reduced the pH but at a rate that was approximately 100times slower than the C02 sparging (figure 7). This is because the air sparging relies on the same
18
reaction, but the atmospheric partial pressure of CO2 is only 3.0E-04. Because the amount ofCO2 is limited and cannot build up in excess, the drop in pH will be controlled by the alkalinity,i.e., the ability of a solution to neutralize an acid. At Site B the pH dropped to 8.13, which is inthe typical range for an aqueous system dominated by calcite. Because Site A had a much higheralkalinity, the pH dropped to only 9.43, essentially the same value as the degassed samplessparged with CO2. Adding the sediment did not significantly affect the results at either sitebecause the drop hi pH was not great enough to dissolve calcite in the sediment. During the airsparging 20% of the sample water was evaporated from the flask which made the remainingsolution more concentrated, assuming that only H20 was evaporated. To correct for this, all ofthe values of the parameters listed in Table 1 for the air sparging experiments, except pH, weremultiplied by 0.8.
Unsparged control samples that were allowed to interact with the atmosphere by naturalair diffusion reached the same pH values for each site as the sparged samples, but at a muchslower rate. Because we do not see a decrease in pH in the field, ground-water inflow andbuffering by the slag around the shoreline overwhelms the natural diffusion of CO2 into thewater.
The addition of hydrochloric acid rapidly decreased the pH of the water to below 4.0(figure 8) with Site A requiring approximately 10 times the amount of acid as Site B. Whensediment was included there was significant buffering by calcite dissolution with the pHdropping to around 7.34 at both sites. The calcite dissolution continued after the acid additionwas stopped, raising the pH to 8.6 at Site A and 8.3 at Site B. If the acid was added at a slowerrate, the calcite dissolution would keep pace with the acid addition and the pH would not fallbelow these higher values. An acid neutralization system was installed in a high pH ditch alongan elevated highway in Maryland (Broyer, 1994). While the system has able to reduce the pH,it did not have any success in lowering the mortality of bioassay organisms (Bruce Broyer,Maryland State Highways Administration, personal communication).
The addition of dolomite reduced the pH at Site B to 9.09 but at a very slow rate (figure9). At Site A the pH reduction was only slightly greater than a control sample so the experimentswere terminated after 2 months and not sampled. The increase in silicon at Site B suggests thepotential for the dissolution of silica (SiO2) grains in the crushed material, which is enhancedunder alkaline conditions (Stumm and Morgan, 1981).
SiO2 + 2H2O => H+ + H3SiO4- (9)
The dissolution of dolomite [CaMg(CO3)2] could also potentially reduce the pH by reacting toform calcite (CaCO3) and brucite [Mg(OH)2].
CaMg(CO3)2 + Ca2+ + OH' => 2CaCO3 + Mg(OH)2 (10)
19
A geochemical reaction model shows that only 0.2 g of either quartz or amorphous silicain the 35 g of material added could lower the pH of the Site B water to below 9.6 if smallamounts of silicate minerals are allowed to precipitate out. The combination of aqueous silica(SiO2) with cations and water could produce protons and a silicate mineral or complex.However, the particular silicate phase is difficult to predict because many of the mineralssuggested by the model would be unlikely to form in an aqueous system (Hem, 1985). Withoutsilicate precipitation, the aggregate would have to contain 5 grams of quartz per 30 grams ofdolomite to drop the pH to 9.6. To remove the effect of dolomite buffering, adding 5 g of quartzwith no dolomite caused the modeled pH to drop to 7.4. At Site A, reaction with the dolomiteaggregate reduced the pH to only 10.9, which was consistent with a geochemical modelassuming 0.2 g of quartz in the aggregate. As with Site B, adding higher amounts of quartz inthe geochemical model significantly reduced the pH further. This suggests that the extremelyhigh pH could be naturally remediated in areas where the ground water flows through an under-lying sand before it discharges. While the volume of quartz may be too small and the reactionrate too slow for remediating surface waters, this material could have potential applications asreactive barriers in ground-water systems or as a potential neutralizer for slag in soil.
Metals. If the pH of a solution open to the atmosphere decreases from 12 to near neutral,several potential reactions might be expected (Stumm and Morgan, 1991). For example, bothsilicon and aluminum are relatively soluble at high pH but insoluble at neutral pH. Thus wemight expect silicate and aluminosilicate minerals, such as amorphous SiO2, to precipitate outof solution and Al and Si concentrations to decrease. As the pH drops, especially below 8,calcite becomes less stable and may dissolve. Calcium concentrations would increase and totaldissolved inorganic carbon would decrease. The dissolution of calcite may also release metalsthat may have co-precipitated with the calcite. The solubility of metal oxides is more complex.As the pH drops from 12 to between 9 and 10, metals such as zinc, copper, and iron becomemore insoluble, but as the pH continues to drop, metal oxides are more soluble andconcentrations of metals in solution may increase. The results from most of the remediationexperiments are ambiguous with respect to these solubility relationships, with the presence ofsediment being an important influence on solution chemistry.
Silicon concentrations were extremely high at Site A, around 80 mg/L. Concentrationsdecreased significantly in the acid and air treatments, but not the CO2 treatment (table 1, figure11). In figures 11-13, the three blank bars at Site A and the one at Site B are for the samples thatwere not analyzed. The presence of sediment only made a difference for the acid treatment, withSi concentrations much lower in the presence of sediment. At Site B, Si concentrations werevery low, around 2 mg/L. This suggests Site B waters were not in contact with SiO2. However,the presence of sediment in the control produced increased concentrations of Si after 12 days (7.5mg/L). Silicon concentrations did not change appreciably for any of the treatments except thedolomite with sediment, with a final Si concentration of 9.3 mg/L.
Aluminum concentrations were slightly high compared to most other natural waters, butnot extraordinarily so. At Site A, concentrations were around 0.6 mg/L, and slightly lower at
20
ri
J
-
;
100
80 -
60 -
O>
CO
Experiment startControlControl w/ sedimentCO2
CO2 w/ sedimentDolomiteDolomite w/ sedimentAirAir w/ sedimentAcidAcid w/ sediment
40 -
20 -
Site A Site B
Fieure 11. Initial and final aaueous silicon concentrations for different treatments.
Site A Site B
Figure 12. Initial and final aqueous aluminum concentrations for different treatments.
21
250
Site A Site B
Figure 13. Initial and final aqueous calcium concentrations for different treatments.
Experiment startControlControl w/ sedimentCC-2CO2 w/ sedimentDolomiteDolomite w/ sediment
•HPJ AirAir w/ sedimentAcidAcid w/ sediment
Site A Site B
Figure 14. Initial and final aqueous magnesium concentrations for different treatments.
22
Site B (0.4 mg/L). For Site B, there was little change in Al concentrations except for the air withsediment, acid with sediment treatments and both dolomite treatments, where concentrationsdropped to below 0.1 mg/L (table 1, figure 12). For Site A, the only significant drop was for theacid with sediment treatment (0.15 mg/L). For the two air treatments, Al actually increased togreater than 1 mg/L.
Calcium concentrations at Site A were low, about 11 mg/L. Large increases (15-20times) were observed for the CO2 with sediment and acid with sediment treatments, probably dueto dissolution of calcite in the sediment (table, figure 13). However, concentrations in the airwith sediment treatment decreased. The pH in the air treatment experiments decreased to only9.4, a pH where calcite is stable. In the CO2 treatment without sediment, Ca concentrations alsoincreased significantly (5 times). The source of the Ca is probably dissolution of suspendedcolloidal material. Calcium concentrations were greater at Site B (30 mg/L). The only treatmentwith a significant increase in Ca was the acid with sediment. The CO2 with sediment treatmentincreased slightly. Most of the other treatments had a decrease in Ca, including the controls withand without sediment. The pH at the end of these experiments was greater than 8.
Magnesium concentrations were very low at both sites (< 1 mg/L). At Site A, substantialincreases were observed for the CO2 with and without sediment treatments (table 1, figure 14).There was a smaller increase for both acid treatments. This suggests Mg containing phases, suchas brucite or dolomite, were in the sediment and suspended fraction. At Site B, the largestincreases in Mg were for the dolomite treatments, suggesting dolomite was dissolving.
The behavior of Si, Al, Ca, and Mg was as expected for most of the treatments, but therewere instances where the behavior was opposite to what was expected. The relatively high finalpH values for the air and dolomite treatments apparently prevented significant dissolution ofcarbonate minerals.
One of the major concerns about remediating the water is the possibility of releasing toxicmetals from the sediments into solution. This was observed for some treatments (table 2). Therelease of metals from sediments is generally not a concern at Site B, with a few exceptions (Band Fe under one treatment). Site A has more potential for release of toxic metals, mainlybecause of the high levels of metals in the sediment. Many of the metals, such as Zn, Pb, Ni, andCd, could be in the form of sulfide precipitates, which when oxidized could potentially result intheir release. In the experiments Ba, Cu, Fe, and Mn appearred to be the most susceptible torelease. One of the treatment variables was the length of the experiment. The air treatments tookconsiderably longer than CO2 or acid treatments, and the ah" treatments generally had lowerconcentrations of metals. This is probably due at least partly to the higher final pH. But it is alsopossible that if the other treatments were allowed to sit for longer times, the metals wouldprecipitate out of solution or be adsorbed by sediments. Thus, the release of metals duringremediation might be a temporary phenomenon.
23
Table 2. Changes in concentrations of selected metals from startto end of experiments as a function of treatment type.
Treatment B Ba Cu Fe Li Mn Sr ZnControlC02
CO2 w/ sedimentSite A Air
Air w/ sedimentAcidAcid w/ sedimentControl + -Control w/ sediment + -CO2 - +CO2 w/ sediment ++
SiteB Dolomite + - -Dolomite w/sediment ++ -AirAir w/ sediment^cid_w/_sedirnent +_ - _+_
Notes: + increased by at least 2-fold,++ increased by at least 10-fold- decreased by at least half
Despite the high pH at Site A, the NVOC concentrations were very high, and becauselevels increased in the treatments containing sediment, the sediment appearred to be organic rich.The identity of the organic compounds are unknown, but they may play a role in the metalschemistry. They may chelate with metal ions or they may be substrates for microbial activity,making conditions more reducing (DO near zero in the field) which would affect metal mobility.
Several other observations can be made from the data. At Site A, potassium increased2-fold in the two acid treatments. Sodium was very high (1700-1800 mg/L) and decreased> 1 OOmg/L in the air treatment and increased > 100 mg/L in the acid with sediment. Phosphorousincreased over 2-fold in the CO2 with sediment and the acid with sediment treatments.Ammonia-N decreased significantly in the air treatments and was either oxidized or volatilized.Sulfate was very high (1900 mg/L) and did not change significantly for any of the treatments.NVOC increased in the CO2 with sediment and the acid with sediment treatments, and decreasedin the acid alone experiments.
At Site B, chloride was low (3 mg/L) and increased in the control with sediment, thedolomite with sediment treatment, and especially the air alone treatment. Sulfate was low (13mg/L) and doubled in the two dolomite treatments due to the dissolution of sulfate minerals.
Toxicity. The toxicity as measured by the MicroTox™ procedure was reducedsignificantly by some of the remediation methods (figures 15-16). This procedure uses thecolorimetric response of the luminescent bacteria Vibrio flscheri to measure the decrease in
24
c0
D5 Minutes
D10 Minutes
• 20 Minutes
Field
Sample
C02 CO2w/
Sediment
Airw/
Sediment
Acid Acid w/
Sediment
Figure 15. MircoTox™ toxicity responses for Site A.
Q5 minutes
Q10 minutes
• 20 minutes
C02w/
Sediment
Airw/
Sediment
Acid w/
Sediment
Figure 16. MircoTox™ toxicity responses for Site B.
25
metabolic activity (increase in mortality) caused by toxic substances in a water sample. AMicroTox™ response close to zero after the bacteria have been exposed to the water sample for20 minutes is desirable. A large negative value indicates overstimulation of growth. Eachsample was split into triplicates and the three results were averaged. The response to the fieldsamples from both sites was 100%, meaning that there was no metabolic activity. The high pHis the most likely cause of the high mortality although the levels of copper and zinc were abovethe lowest observable effect concentrations of 0.02 and 0.1 mg/L, respectively, for longer lengthtests (Azure, 2000).
The toxicity at Site B was substantially reduced by the CO2, air, and acid. Toxicitysamples were not collected for the dolomite test due to its long duration. In the experiments thatincluded sediment, the air sparging method appears to be the most effective method because itreduced the pH without dissolving the sediment. The increases in calcium and bicarbonate inthe CO2 experiment and calcium and chloride in the acid experiments seem unlikely to havecaused all the overstimulation (negative response). Another potential factor for overstimulationmay be the dissolution of organic material in the sediment.
The toxicity results from Site A also showed a reduction, however, the results werecomplicated by the reddish color of the organic material in the water. A color correction wasperformed on the data from this site, however, in some cases the color correction was six timesthe uncorrected change in response. Therefore, the Site A data are more qualitative thanquantitative. The overstimulation for the CO2 and acid experiments that included sediment maybe due to an increase in phosphorous and organic material from the dissolved sediment.
Natural Remediation at Site C
The high pH spring at Site C is a great example of how extremely alkaline waters can beremediated naturally. The spring is at the end of the former 119th Street right-of-way whichforms a swale that drains as much as 20 acres or more. The infiltration of some of this drainageinto the slag fill in the area above the spring is the likely source for much of the water in thespring. Ground water discharge from the at-grade landfill and the automobile junk yard adjacentmay also contribute water to the spring. Flow from the spring is ephemeral. On the day thesamples were taken, the spring discharged into the ditch along Doty Avenue West at an elevationclose to that of the surrounding land surface (figure 17). After flowing three feet, the springdischarge mixed with surface water runoff from same drainage area (figure 18). The mixingratio was approximately 40% ground water to 60% surface water. After the mixing zone, thedischarge flow became turbulent as it went down through a narrow cut and dropped roughly eightinches in elevation. The flow then entered the main portion of the ditch and combined withwater flowing from a culvert in roughly a 50/50 ratio.
The spring discharge had a pH value of 12.19 (figurel 8, table 3). In the first mixing zonethe pH dropped 0.4 pH units due to dilution with the surface runoff which had a near-neutral pHand was likely saturated in dissolved CO2. The pH continued to drop through the turbulent zone
26
-
Figure 17. Photographs of Site C:a) on sampling date,b) under high water conditions.
Surface runoff Spring
Fence
Sampling pointspH contourCalcitePrecipitation
N
Figure 18. Sampling location and pH distribution at Site C.
27
Table 3. Water Quality Charactistics of Site C
00
Distance pHSample (ft) (units) Al
C1 - DissolvedC1 - TotalC1 - Sediment
C2 - DissolvedC2 - Total
C3 - DissolvedC3 - TotalC3 - Sediment
C4 - DissolvedC4 - TotalC4 - Sediment
C5 - DissolvedC5 - Total
C6 - DissolvedC6 - Total
111
44
10
10
10
20
20
20
1
1
7
7
12.19 0.058
0.549
6094
11.78 <0.042
0.839
9.00 <0.042
0.337
6648
7.77 <0.042
0.174
13423
7.43 <0.042
0.208
8.60 0.082
0.773
As
<0.13
<0.25
<21.8
<0.13
<0.25
<0.13
<0.25
<21.8
<0.13
<0.25
<21.8
<0.13
<0.25
<0.13
<0.25
B
0.32
0.341
92.26
1.53
1.74
1.85
2.07
109
1.89
2.03
106
1.82
1.96
1.87
2.11
Ba
0.302
0.283
517
0.135
0.237
0.125
0.336
636
0.158
0.267
1602
0.294
0.305
0.085
0.072
Be
<0.004
<0.003
-
<0.004
O.003
<0.004
<0.003
-
O.004
<0.003
-
<0.004
<0.003
<0.004
<0.003
Ca
731
744
290372
311
413
133
317
284634
159
245
97243
194
202
173
198
Cd
<0.016
O.011
29.8
O.016
<0.011
<0.016
0.011
21.8
O.016
<0.011
50.3
<0.016
<0.011
<0.016
<0.01 1
Co
<0.011
<0.007
44.1
<0.011
<0.007
<0.011
<0.007
13.4
<0.011
<0.007
24.5
<0.011
<0.007
<0.011
<0.007
Cr
0.018
0.021
44.8
0.012
0.024
<0.009
0.016
64.7
<0.009
0.011
328
<0.009
0.007
<0.009
0.011
Cu
0.035
0.038
192
0.017
0.054
0.007
0.032
213
0.006
0.018
633
0.008
0.014
0.016
0.020
Fe
<0.009
0.834
46311
0.009
6.76
0.018
3.20
34156
0.012
1.84
96439
0.018
5.66
O.009
0.519
K
23.4
22.8
1237
25.8
26.1
20.1
21.7
1062
19.6
20.8
3224
16.4
16.9
24.4
26.9
Notes: All values in mg/L unless noted< Below detection- Not measured
Table 3. Continued.
10
Sample
C1 - Dissolved
C1 - Total
C1 - Sediment
C2 - Dissolved
C2 - Total
C3 - Dissolved
C3 - Total
C3 - Sediment
C4 - Dissolved
C4 - Total
C4 - Sediment
C5 - Dissolved
C5 - Total
C6 - Dissolved
C6 - Total
Li0.046
0.055
14.3
0.055
0.071
0.087
0.113
9.95
0.099
0.115
17.5
0.115
0.128
0.057
0.073
Mg
<0.06
4.92
19910
0.13
27.0
37.1
41.3
11915
39.1
39.9
16113
41.7
42.4
34.1
37.0
Mn
O.003
0.030
708
0.003
0.242
0.063
0.673
1363
0.127
0.555
1304
0.448
0.624
0.021
0.101
Mo
0.037
0.026
<3.86
0.026
O.025
O.023
<0.025
5.29
<0.023
<0.025
15.6
O.023
<0.025
<0.023
<0.025
Na
58.3
60.6
813
61.4
65.3
118
136
985
146
154
1685
172
181
58.3
63.3
Ni
O.024
0.038
31.2
<0.024
0.027
<0.024
0.033
38.9
0.026
0.038
76.4
<0.024
0.029
0.026
0.024
' P
<0.33
O.33
956
O.33
O.33
O.33
0.33
903
O.33
0.41
3677
O.33
O.33
O.33
O.33
Pb
O.075
O.063
1135
O.075
0.249
O.075
O.063
1362
<0.075
O.063
4136
O.075
O.063
0.075
O.063
S
11.6
12.2
4284
55.4
62.1
40.8
48.3
6875
40.2
44.3
8704
33.8
36.5
78.9
87.8
Sb
O.37
O.17
<62.2
O.37
0.17
0.37
O.17
O2.2
O.37
O.17
75.88
O.37
O.17
0.37
O.17
Se
O.27
0.24
<45.4
O.27
0.24
O.27
O.24
<45.4
O.27
O.24
<45.4
O.27
O.24
0.27
O.24
Si
0.62
0.995
13807
1.93
6.77
5.29
6.45
12346
5.69
5.79
24074
5.70
6.37
6.84
7.49
Sn
O.07
O.11
145
O.07
O.11
O.07
O.11
182
O.07
O.11
570
O.07
O.11
0.07
0.11
Sr
0.826
0.839
445
0.721
0.817
0.826
1.34
566
1.02
1.25
241
1.37
1.41
0.618
0.679
Notes: All values in mg/L unless noted
< Below detection
- Not measured
Table 3. Continued.
U)o
C1
C1
C1
C2
C2
C3
C3
C3
C4
C4
C4
C5
C5
C6
C6
Sample
- Dissolved
- Total
- Sediment
- Dissolved
- Total
- Dissolved
- Total
- Sediment
- Dissolved
- Total
- Sediment
- Dissolved
- Total
- Dissolved
- Total
Ti
<0.005
0.010
99.3
<0.005
0.031
O.005
0.007
540
<0.005
0.008
937
<0.005
0.005
<0.005
0.006
V
0.015
0.014
8.41
0.015
<0.003
0.013
0.006
40.6
0.013
0.011
76.6
0.011
0.014
0.024
0.016
Zn NH4-N F
0.227 0.544 4.51
0.631
4298
0.236 0.520 2.26
1.71
0.250 3.44 0.45
1.97
5267
0.236 0.753 0.29
0.982
6109
0.249 3.76 0.14
0.739
0.252 0.295 1.10
0.495
Eh Temp
Cl NO3-N SO4 0-PO4-P TDS NVOC Alk (v) (deg C)
281 2.06 20.7 0.054 1949 21.7 1617 -242 8.7
298 1.63 161 0.497 1201 17.5 389 -140 10.3
337 0.81 133 0.135 1147 10.9 268 -240 12.1
347 0.74 121 0.026 1217 10.5 315 -80 13.1
352 0.48 101.0 0.030 1320 7.7 471 0 12.8
284 2.53 244.3 0.035 1025 11.9 78 187 13.4
Notes: All values in mg/L unless noted
< Below detection
- Not measured
and into the main portion of the ditch due to the influx of atmospheric CO2 (see reaction 4) andmixing with neutral water coming from the culvert. By the time the flow had traveled about 20feet, the pH had been reduced to 7.7.
The water in the spring was actively reacting as it discharged, as evidenced by the largebuild up of calcite precipitate along the bottom of the ditch (figures 17 and 18). The spring hada very high calcium concentration (731 mg/L) that dropped along the flow path along with thedrop in pH (figure 19). The calcium reacted with the carbonate ions, formed from theintroduction of CO2 into the system, to form calcite (see reaction 5). By the tune the pH haddropped to 9.0, the concentration of calcium dissolved in the water had dropped by 82% to 133mg/L. The precipitation appears to be occurring in the water column, as evidenced by thepresence of a significant amount of calcium in suspension. As these fine suspended particles ofcalcite grow, they have a strong potential to coprecipitate metals. The opposite occurs in a moreneutral pH spring discharging from a typical limestone aquifer where degassing of excess CO2
in the ground water causes calcite to precipitate (Hem, 1985).
The concentration of other metals also decreased along the flow path, including copper,aluminum, chromium, lead, and molybdenum. The concentrations of total dissolved solids,fluoride, nitrate, and non-volatile organic carbon also dropped. The concentration of sodiumincreased along the flow path, possibly due to runoff of accumulated salts from the road saltapplied to the road or to cation ion exchange with calcium at sorption sites.
The surface runoff (C5) also contributed to the system by adding iron, magnesium,silicon, sulfate, manganese, boron, and lithium. The iron concentration was high (5.66 mg/L),but almost all of it was in the suspended material and did not dissolve downgradient in thesystem. These suspended iron particles will likely adsorb other metals in the water to theirsurfaces. When the surface runoff mixed with the high pH water from the spring, the dissolvedmagnesium and silicon concentration dropped dramatically and there was a correspondingincrease in the suspended concentration, indicating that they flocculated out as suspendedmaterial (figure 20). A geochemical model of this water shows that brucite [Mg(OH)2] has thepotential to precipitate out. This suspended material then redissolved downgradient as the pHdecreased. Ammonium was also lost in the mixing zone due to its conversion to ammonia gas,which can volatilize. The reappearance of NH4-N at sampling point C3 suggests that it may alsohave become attached to suspended material at high pH.
The sediment samples collected at sampling points Cl, C3 and C4 show a significantaccumulation of 18 different metals along the flow path (figure 21). This distribution probablyreflects the fact that the longer calcite and iron particles are suspended in the water, the moreopportunity other metals will have to adsorb or coprecipitate with them. Iron appears to havemade it farther down the flow path than the suspended calcite. Cadmium, cobalt, antimony, andtin were not found at detectable levels in the water but were found in the sediment.
31
•L
800
700
600
500
do 400
E
J5 300
200
100
0
Ca suspended
13
- 12
- 11
10
- 9
- 8
- 7
- 6
5
5 10 15 20
Distance from spring (ft)
Figure 19. Calcium distribution along the flow path from the spring.
15 200 5 10
Distance (ft)
Figure 20. Magnesium and silicon distribution along the flow path from the surface runoff.
32
r17 I
1000000
c100000
10000
Ic.o'•»-«(0
§oo1000
100 --
105 10 15
Distance from spring (ft)
20
Figure 21. Distribution of metals in the sediment.
33
The toxicity of the water was tested using the water flea Ceriodaphnia dubia. These fleasare important primary consumers in aquatic ecosystems and are an important food source forsecondary consumers. They have been shown to be at least as sensitive to many contaminantsas fish and other invertebrates. In the high pH samples Cl and C2, all five of the organismsintroduced in each of three replicates died. In sample C3 the number of organisms that died wasreduced to three. In samples C4, C5 and C6 none of the organisms died. The pH appears tocontrol the toxicity response. However, it would require extensive study to determine iforganisms were reacting solely to the pH, to the concentration of an individual constituent, ora sweet of constituents that are elevated at high pH.
Conclusions
From the data it appears as though the ah1 sparging technique may be the most feasibletechnique for remediating the alkaline ground-water discharges. Air sparging can effectivelylower the pH without the risk of dissolving the calcite and other minerals in the sediment andreleasing metals into the water. The precipitation of calcite at Site C, where sufficient airsparging was occurring naturally, removed a significant amount of metals from the water columnby coprecipitation. Air sparging is also effective in reducing the toxicity in both the laboratoryexperiments and at Site C. Finally, air sparging is considerably more practical because air is freeand does not involve the complex engineered systems to control the input of treatment agents,as is required for CO2 sparging or HC1 addition. An ideal air sparging system would use a simpleaerator that is easy to set up and maintain, if the hydrology of the spring to be treated could notbe modified to promote natural remediation. If the calcite precipitation could be confined to asmall area, it could be periodically removed to prevent further interaction with the environment.
References
American Public Health Association, American Water Works Association, and WaterEnvironment Federation. 1992. Standard Methods for the Examination of Water andWaste-water. 18th edition, Washington D.C.
Azure Environmental. 2000. http://www.azurenv.com/
Bethke, C.M. 1998. The Geochemist's Workbench Release 3.0. University of Illinois, Urbana,IL.
Broyer, B.W. 1994. Alkaline Leachate and Calcareous Tufa Originating from Slag in aHighway Embankment near Baltimore, Maryland. Transportation Research Record,Volume 1434, pp. 3- 7.
34
Colton, C.E. 1985. Industrial Wastes in the Calumet Area, 1869-1970. Research Report 001,Illinois Hazardous Waste Research and Information Center, Champaign, IL.
Drever, J.I. 1997. The Geochemistry of Natural Water. Prentice Hall, Upper Saddle River, NJ.
Duwal, K.G. 1994. Event-Based and Seasonal Precipitation Effects on Ground Water-WetlandsInteractions near Lake Calumet, Southeast Chicago, Illinois. Masters Thesis, Departmentof Geological Sciences, University of Illinois, Chicago.
Hem, J.D. 1985. Study and Interpretation of the Chemical Characteristics of Natural Water.U.S. Geological Survey Water-Supply Paper 2254.
Illinois Environmental Protection Agency. 1987. Quality Assurance and Field Methods Manual.Division of Water Pollution Control, Planning Section, IEPA, Springfield, IL.
Kay, R.T., T.K. Greeman, R.F. Duweiius, R.B. King, I.E. Nazimek, and D.M. Petrovski. 1996.Characterization of Fill Deposits in the Calumet Region of Northwestern Indiana andNortheastern Illinois: U.S. Geological Survey Water-Resources Investigations Report96-4126, Urbana, IL.
Locke, R.A., R.C. Berg, H.A. Wehrmann, M.V. Miller, and D.A. Keefer. 1997. Vulnerabilityof Illinois Nature Preserves to Potential Ground-Water Contamination, Volume I:Methodology and Initial Assessment. Illinois State Water Survey Contract Report 612.
Roadcap, G.S., and W.R. Kelly. 1994. Shallow Ground-Water Quality and Hydrogeology of theLake Calumet Area, Chicago, Illinois. Interim Report. Illinois State Water Survey.
Roadcap G.S., M.B. Wentzel, S.D. Lin, E.E. Herricks, R.K. Raman, R.L. Locke, and D.LHullinger. 1999. An Assessment of the Hydrology and Water Quality of Indian RidgeMarsh and the Potential Effects of Wetland Rehabilitation on the Diversity of WetlandPlant Communities. Illinois State Water Survey Contract Report 654.
Stumm, W., and J.J. Morgan, 1991. Aquatic Chemistry. John Wiley & Sons, New York.
35
