Environ. Sci. Technol. 3994, 28,2203-2210
Electrorestotation of Metal Contaminated Soils R. Edwln Hicks" and Sebastlan Tondorft
Department of Mechanical Engineering, Massachusetts Institute of Technology, Cambridge, Massachusetts 02 139 The removal of metals from contaminated soils using electric fields has been successfully demonstrated in the laboratory, yet field trials have yielded anomalous results. Poor performance may be attributed to interaction of the metals with naturally occurring electrolytes, humic substances, and co-disposed wastes. Immobilization of contaminants in a narrow band in the soil, analogous to isoelectric focusing, was reproduced experimentally and simulated with a mathematical model. It was shown that the focusing effect can be eliminated by controlling the pH at the cathode using a water rinse. Immobilization resulting from precipitation with carbonates and codispo&d wastes may additionally require chelating agents and control of the redox potential to effect removal. Pourbaix diagrams provide a means for rapidly identifying pH and redox conditions suitable for mobilizing metal wastes. Optimum operating conditions can then be determined using a mathematical model that incorporates the appropriate metal speciation chemistry.
rapid and is not affected by variations in the (potential. In fact, laboratory experiments with solutions of single metals usually indicate that high removal efficiencies can be achieved by electromigration (5-8). However, anomalous results have been obtained with metals even under laboratory conditions (1, 9), and the limited field trials that have been made at metal-containing sites have also produced inconsistent results (10). The purpose of this paper is to explore the effect of various site and operating conditions on the efficacy of metall removal by electromigration. In particular, the effect of substances that are either naturally present in the groundwater or that might be co-disposed along with the metal wastes are considered. The efficiency of zinc removal by electromigration under various operating conditions is predicted using a theoretical model that incorporates equilibrium chemistry, and the results compared with experimental findings. The techniques discussed can be used to predict performance under different field conditions and so help in selectingsites most amenable to this treatment technology.
Introduction
Background
In electric field restoration, pairs of electrodes are placed in the contaminated soil and a direct current (dc) potential is applied across them ( I ) . The contaminants are then transported under the action of the electric field to the electrode wells from where they can be brought to the surface. Unlike soil flushing, electrorestoration is effective in soils of low or variable permeability and does not disperse the contaminants outside the treatment zone. The electric field drives the contaminants toward the electrodes by two mechanisms. One is electroosmosis, an electrokinetic phenomenon in which the saturating liquid and dissolved substances flow toward an electrode. The electroosmotic flow rate is proportional to the product of the applied electric field strength and the ( potential at the soil-liquid interface. The value of the ( potential depends on the soil properties as well as the ionic strength and pH of the saturating liquid. A typical ( potential is about -10 to -100 mV, and in a 100 V/m electric field, water will flow toward the cathode at a velocity of around 10 cm/day. Laboratory tests with soluble organic chemicals have generally shown high removal efficiencies (2-4). The second mechanism is electromigration in which charged ions move in the electric field at a rate that is proportional to the product of the electric field strength, the charge on the ion, and the mobility of the ion ( I ) . Transport of charged ions by electromigration is generally more rapid than by electroosmosis. In cases where the ionic strength is high due to the presence of inorganic contaminants, the (potential is small and transport occurs mainly by electromigration. Removal of heavy metals might be expected to be straightforward because transport by electromigration is
For a metal to be transported by electromigration, it must be in solution and carry a charge. Furthermore, the polarity of the charge should not change as the metal migrates toward an electrode. Changes in solubility and polarity do however occur as a result of reaction with codisposed substances and substances naturally present in groundwaters and as a result of pH changes brought about by electrolysis reactions at the electrodes. Variation of pH in Saturating Liquid. The electrolysis reactions that normally occur at the electrodes can produce large changes in the pH of the pore liquid. These reactions generate hydrogen ions a t the anode
~
* To whom correspondence should be addressed.
t Present address: Dr. Johannes Heidenhain GmbH, Abteilung Laengenmessgeraete 11, Postfach 1260,83292 Traunreut, Germany.
0013-938X/94/0928-2203$04.50/0
0 1994 American Chemical Society
and hydroxyl ions at the cathode 2H,O
+ 2e
-
20H- + H,
= -0.83 V
(2)
where Eo is the standard electrode potential. The rate at which the hydrogen and hydroxyl ions are produced is fixed by the current, and this is in turn dependent on the applied voltage and the conductivity of the medium. The hydrogen and hydroxyl ions are transported into the soil, mainly by electromigration, so that the soil near the anode becomes more acidic and that the soil near the cathode becomes more alkaline with time. For a given value of the electric field, the magnitude of the pH change is limited by the amount of other ions that are present. This is because the difference in the concentration of hydrogen and hydroxyl ions at any position in the soil must be balanced by other ions to satisfy the electroneutrality condition:
czici= 0 Environ. Scl. Technol., Vol.
(3) 28, No. 12, 1994 2203
Here, zi is the charge number and Cj is the concentration of species i. The concentration of the hydrogen and hydroxyl ions is further constrained by their dissociation constant, K,:
K , = [H'IEOH-I
(4)
In practical terms, this means that where the migrating hydrogen and hydroxyl ions meet, theyreact to form water. The position of this meeting point, which is marked by a sharp change from acidic to alkaline conditions, is determined by the relative velocity of the hydrogen and hydroxyl ions. In the absence of advection (that is, for negligible electroosmosis and stagnant groundwater), the velocity of each ion can be written as the sum of a diffusion u d and migration U m term:
Here, Dj is the diffusion coefficient, Uj is the mobility of species i, F is the Faraday constant, 4 is the voltage, and 7 is the tortuosity of the soil. Because the electric field is determined by the local conductivity, the position of the jump in pH depends on the concentration and mobility of all the ions present. Effect of pH Variation on Metal Transport. The pH regulates the adsorption and desorption, precipitation and dissolution, and speciation reactions of the heavy metals. At low pH, metals tend to desorb from the soil (11)and will be present in solution as positively charged ions. At higher pH's, metal solubility is reduced by the formation of hydroxides: Mn++ nOH-
-
M(OH),
4
+ H,O
(7)
The heavy metals consequently exhibit a characteristic minimum solubility at some intermediate pH as shown in Figure 1 for zinc, trivalent chromium, and lead. Such metals will be positively charged under the acidic conditions on the anode side of the pH jump and on migrating toward the cathode will encounter the region of increasing pH and precipitate. Metals in the high pH region near the cathode might be negatively charged, migrate toward the anode, and also precipitate in the region of the pH jump. The accumulation of the metals at the pH jump, which has been likened to isoelectric focusing (1, 12), effectively prevents them from being removed at the electrodes. Effect of Other Species. Mobility can also be affected by the presence of substances which combine with the metal to form precipitates or soluble complexes. Naturally occurring substances such as carbonate salts (13) and humic substances (14)can affect the solubility of the heavy metals. Substances that promote precipitation (sulfides in tailing wastes) or that enhance solubility (complexing agents in metal cleaning wastes) might have been disposed of along with the metals or can be introduced into the groundwater as part of the treatment process (15). 2204
Envlron. Scl. Technol., Vol. 28, No. 12, 1994
1
F b. Chromium 111 In presence of 35 mgiL chloride
i
/ 0%
-3
Hydroxide only
T
I
100 mgiL sulfate
In presence of I
/I
-
(6)
At still higher pH's, the solubility of metals might increase due to the formation of soluble complexes which can be positively charged, neutral, or negatively charged like the bizincate ion: Zn2++ 30H- == HZnO;
-3
Sulfides are widely used to precipitate metals in wastewater treatment and mining operations and if present can immobilize the metals even at low pH. The solubility of zinc, for example, is reduced by some 9 orders of magnitude by as little as 0.0001 mol/L (3.2 mg/L) sulfide (Figure la). Under oxidizing conditions, the sulfide is converted to sulfate and then is not usually a problem. In the case of lead, however, 0.001 mol/L (about 100 mg/L) sulfate reduces its solubility by more than 6 orders of magnitude at pH's below 4 (Figure IC). Carbonates are present in groundwaters under oxidizing and reducing conditions. In a system not open to the atmosphere and having a total inorganic carbon concentration of 0.002 mol/L (24 mg/L), the solubility of lead is reduced by about 2 orders of magnitude at pH's below 7 (Figure IC).Some reported laboratory studies with lead are of interest in this connection. Removal efficiencies of 75-95% were found for a synthetic waste consisting of kaolinite and dilute lead nitrate (6). However, with an actual waste contaminated with 13%lead and containing shells (calcium carbonate), most of the metal accumulated as a precipitate at a point approximately midway between the electrodes (9). The solubilities of calcium zincate and calcium chromite are relatively low (16),so the increase in solubility shown in Figure la,b at the higher pH's may not occur in some soils. On the other hand, solubilities of copper and lead might be increased by naturally occurring organics (17,
1.o In c 0
30.5 .c
9 Q 0
::
U
0.0
2 -0.5 Cathode I
2
4
6
>m I
I
I
a
10
12
PH Figure 2. Effect of pH and redox potential on chromium removal by electromigration. Under oxidizing conditions (upper bold line), the chromium is present as negative complexes and is readily transported to the anode. Under reducing conditions, the chromium will precipitate on entering regions having pH > 4.5.
It?), and as little as 0.001 mol/L (35 mg/L) chloride can increase the solubility of trivalent chromium by 2-3 orders of magnitude (Figure lb). The solubility of the heavy metals can also be enhanced by introducing complexing or chelating agents into the soil (19). Effect of Redox Potential. The oxidation state of a metal significantly affects its solubility (12). Iron is considerably more soluble in the ferrous (reduced) form than in the ferric (oxidized) state, whereas in the case of chromium it is the more oxidized hexavalent form that is more soluble. Transport into the soil of oxygen generated at the anode, or other oxidizingagents that might be added to the electrode rinse, could affect the mobility of the metals by altering their oxidation state or the oxidation state of other substances. The mobility of metal wastes can be estimated using computerized equilibrium models (20,211,which provide a means of rapidly exploring the effect of pH, redox potential, and other substances on the speciation of a specific metal. The dominant species present at equilibrium can be shown graphically over the full range of pH and redox potential in the form of Pourbaix ( E H - ~ Hor predominance) diagrams (22,23). For example, Hem and Durum have used Pourbaix diagrams to show how sulfide, sulfate, and carbonate affect the solubility of lead (24). Computer programs suitable for running on a PC are available (25)for preparing Pourbaix diagrams similar to that shown in Figure 2 for 0.01 mol/L (about 500 mg/L) chromium in the presence of 0.003 mol/L (abut 100 mg/L) chloride. Locating the electrodes on the Pourbaix diagram provides a means of identifying favorable operating conditions. The two bold lines in Figure 2 trace paths followed by the contaminants between unrinsed electrodes in an oxidizing environment (top) and a reducing environment (bottom). It is seen that for the oxidizing case, hexavalent chromium should be removed at the anode, but in the reducing case, trivalent chromic oxide will precipitate on entering the high pH environment near the cathode. Successfulremoval of trivalent chromium clearly requires that either the cathode be “moved” into the low pH region or the contaminant be oxidized to its hexavalent (and highly toxic) state. Alkaline conditions might be
selected in the hexavalent case because, unlike cations, anions are not strongly adsorbed at high pH (19). These equilibrium studies provide a rough guide only and must be further interpreted in terms of the detailed composition and chemistry at the site. The actual system may not be at equilibrium (13), and the effect of background substances such as natural organic materials might not be accurately known or might not be included in the model. Control of pH and Oxidation State. Various options exist for controlling the pH ( 1 , 9, 12). In the usual case in which a low pH is desirable, the electroosmotic velocity might in some cases be sufficient to drive the acid generated at the anode all the way to the cathode. More effective control can be achieved by rinsing the cathode with water or dilute acid to remove the hydroxyl ions as they are generated and so prevent the high pH wave from entering the soil ( 1 , 26). This effectively moves the cathode into the low pH region of the Pourbaix diagram. Bubbling carbon dioxide through the cathode well has been shown to effectively maintain acid conditions throughout the soil (12). Similarly,the anode can be rinsed to promote alkaline conditions. More complicated approaches such as isolating the electrodes with membranes (27) or salt bridges (28) have also been suggested for pH control. In theory, the redox potential in the soil can also be modified. Many groundwaters have little or no contact with the atmosphere so oxygen may be depleted or absent (29). The pore liquid could be rendered more oxidizing if the anode solution, which is saturated with oxygen, is moved into the soil by electroosmosis. Further, oxidizing agents could be added to the electrode rinses and moved into the soil by electromigration. The oxidation rate of chromium (30)and many other heavy metals is slow. In some cases, adjustment of the pH, oxygen concentration, etc. might enable the oxidation reaction to proceed at a rate that is fast compared to the expected duration of the restoration process. Electrode Reactions. The electrode reaction that occurs in a multicomponent system is determined by the electrode potential E, which can be calculated using the Nernst equation (31):
RT E = Eo+-ln nF
ni ni
[oxidized speciesIui [reduced specieslui
(8)
where Eo is the standard electrode potential, n is the number of electrons transferred, and ui is the stoichiometric coefficients in the reaction oxidized species
+ ne
-
reduced species
(9)
In a complex mixture, the electrode reaction with the most positive potential will occur exclusively at the cathode and that with the most negative potential will occur at the anode. Should concentrations be such that two or more electrode potentials are the same, then these electrode reactions will occur simultaneously. In practice, the calculated potential must be corrected for the overpotential for the specific reaction and electrode material (31). The smallest concentration of the metal which will result in it plating on the electrode can be found as a function of pH by equating the electrode potentials from the Nernst equation for hydrogen evolution and for metal plating. As might be expected, metals with more positive electrode Environ. Scl. Technol., Vol. 28, No. 12, 1994
2205
Table 1. Experimental Parametern.
A Ammeter COnneCtlOn to
reference electrode
I
Voltmeter
I
-\
experiment
cell length (em)
electrical control
initial pH
Zn10.4 Zn21.5 Zn2.6 Zn27.V Zn18.8c Zn20.aC m7.10 Zn28.10 Znl.10 Zn6.11 Zn9.11 Zn15.11
20 20 20 20 20 50 20 20 20 20 20 20
100VIm 100V/rn 100V/m 100V/m 100V/m 100V/m 2mAmo 2mAmb 100Vlm 15 Vlm 40 Vlm 40 V/m
3.5 6.0 4.5 6.0
~~
ii
-Electrode housing
Flow connectold
Anode well
-/
PlexlolasC1ay n i l 1der
Carbon cathode
Figure 3. Schematic of the
......
-
HCrO,
experlmental test cell.
+ 7H'
-I 3e
E" = +1.35V (10)
Conditions under which such reactions could proceed can be delineated following the technique used above for plating. Experiments
Experiments were conducted to investigate the effect of the changes in soil pH that occur during electrorestoration on the efficiency of metal removal. Zinc was chosen as a contaminant as it shows the characteristic solubility minimum (Figure la), yet its solubility is not overly sensitive to carbonate levels, nor is its speciation affected by the redox potential. The test cell is shown in Figure 3. The soil sample saturated with the contaminated liquid is contained in the 20 cm long, 3.2 cm diameter, Plexiglas tube that is connected on each end to a 1.5 cm long electrode well containing a carbon electrode. The soil is prevented from entering the wells by filter paper supported on a stainless steel screen, which also serves as a reference electrode for measuring the potential across the soil. The electrode wells can be connected to reservoirs for supplying purge solutions and rinses and for collecting effluents resulting from electroosmotic flow. Pressure-induced flow is prevented by equalizing the static heads in the connecting tubing. Ports in the electrode wells vent the oxygen and hydrogen produced by the electrode reactions. Sample Preparation. The experiments were conducted with Georgia kaolin (Albion Sperse 100) with a particle size smaller than 60 fim. The contaminated pore water was prepared by diluting a 1000 mg/L zinc (in nitric acid) standard solution and adjusting the pH with sodium hydroxide. All the test solutions had a zinc concentration of 7.6 mmol/L (500 mg/L) and a nitrate concentration of 2200
Envlmn. Sd. Technd.. Vd. 28. No. 12, 1994
6.0 6.2 6.2 6.0 6.0 6.0
6.0
'
I
potentials are more likely to plate. Plating may be desirable as it separates the metal from the water and also reduces the amount of hydroxyl ions produced. Oxidation of co-disposed ions such as cyanides to less toxic cyanates or evencarbon dioxideandammoniaisfeasible. However, less desirable electrode reactions might occur. For example, trivalent chromium could be oxidized to the toxic hexavalent form at the anode: Cr3+-I 4H,O
~
none BPB' PR' none BPB BPB BPB+PR BPB+PR BPB+PR BPB+PR BPB+PR BPB+PR
n All experiments were run in Albion Sperse 100 clay Containing 40% by mass of a 500 mg/L zinc solution. The conductivity of the solutions at the start was in the range 0.58-0.71 Slm. BPB = bromophenol blue; PR = phenol red. In these experiments, the cathode was rinsed with t a water. ~
Reference electrode I..
~
6.0
indicator
0.2 mol/L (12 400 mg/L). The solutions used in various tests differed in their initial pH as summarized in Table 1. In most experiments,the pH was adjusted to 6 resulting in a sodium concentration of 4250 mg/L. The zinc was in solution at the start of the tests. The dry clay and zinc solution were combined in a ratio of32 by mass and thoroughly mixed. In someexperiments (Table l), pH indicators were added to visualize pH changes in the clay. The indicatorsused were bromphenol blue, which has a yellow color below pH 3 and a blue color above pH 4.8, and phenol red, which is yellow below pH 6.8 and red above pH 8.4. The mixed indicator produces a yellow color below pH 3, a green color between pH 3 and pH 6.8, and a purple color at higher pHs. The saturated clay was allowed to equilibrate for 1 day before the prepared sample was placed in the Plexiglas cylinder using avibratingtable. Thesystem stood for another day before starting the experiment. Data Acquisition. Except for two runs that were operated under constant current conditions, the experiments were run with a constant voltage across the clay as measured at the reference electrodes. Most experiments were run for about 9 days. During the experiment, the volume of fluid collected in the cathode reservoir was monitored as well as the current through the cell and the voltages across the reference and working electrodes. Samples were withdrawn periodically from the electrode wells for determination of pH and dissolved zinc concentration. Zinc was analyzed with a liquid chromatograph or a spectrophotometer. A t the end of each run, the anode and cathode solutions were analyzed for dissolved and total zinc. In addition, the clay was divided into 10sections, and each section was analyzed for its pH, average zinc concentration,and liquid fraction. The pH was estimated by placing a pH paper in contact with the clay. To determine the average dissolved zinc concentration, the pore solution was separated from a portion of a clay section by centrifugation. Another portion of the clay sample was acidifiedwith nitric acid before centrifugation for determining the total zinc concentration. The water content was determined by comparing the weight of the remaining clay section before and after drying in an oven at 105 "C.
-
Position of pH 8.4 front
c
m
I
U U
m N
0.4 Position of OH 3 front
I 0.0 0.0 Flpu~e4. stages of
0.1
0.2
0.3 0.4 Time, days
0.5
0.6
55hrs
Position of me acld and alkaline fronts durlng ttm inniai a run (experiment Zn15.11).
Results
Based on the volume of solution in the reservoirs, a small electroosmotic flow of less than one-fourth of a pore volume occurred during the first 24 h, but thereafter electroosmosis was negligible. This was expected in view of the high ionic strength of the saturating solutions. The redistribution of the zinc in the clay and transport to the electrodewells as reported below is therefore a consequence of electromigration and diffusion only. pH in Electrode Wells. The pH was found to drop to its final value of around 1at the anode and to about 12-13 at the cathode within the first 15 min of a run and then to remain essentially constant. The pH at the cathode varied slightly between 12 and 13 during the first 2 days, possibly due to the consumption of hydroxyl ions by zinc precipitation. pH of Pore Solution. The progression of the position of the pH 3 and pH 8.4 fronts with time as determined by the pH indicatorsisshowninFigure4foranappliedvoltage of 100 Vim. The relative motion and final meeting of the two fronts can be clearly seen in the photographs of Figure 5. The velocity of both fronts remained fairly constant, with that of the acid front measured at 56 cm/day, and that for the alkaline front a t 28 cm/day. Similar velocities were obtained with a 50 cm long cell a t the same overall voltage gradient. As discussed in connection with eq 5, the velocities of the progressing fronts are dependent on the electric field and the chemical composition of the system. The time taken for the two fronts to meet is plotted in Figure 6 as a function of the applied voltages. These data are for identical initial chemical composition of the saturating solutions. For voltages higher than 40 V/m, the relationship between the relative velocity and applied voltage is approximately linear. Removal Efficiency. The amount of zinc removed from the clay wasdetermined from the totalconcentration of the metal remaining in the clay after the experiment. Figure 7 shows the measured distribution of zinc and pH in the clay a t the end of a 9-day run. The steepness of the pH jumpandita influenceon thezinc distribution is clearly seen. Near the anode, the zinc is a t a relatively low concentration and is in solution. Near the cathode, the zinc is precipitated and present in amounts on the order of the initial concentration. In confirmation of the isoelectric focusingeffect discussed in the previoussection,
6 0 hrs
...
.
~
. -
. ,. .-
I
I
I
I
0,375
0.5
0,625
0.75
..
I > 0,875
Normalised distance from anode FWI~ 5. Regression Of the acld and alkaline fronts as revealed by the pH indicators (experiment Zn15.11). 5
~
~
~
~
~~
~T~
.
~~
,
~
.,
.
,
i
4!
3
II b~
I
2 ;
~
I I!
0
20
40
60
80
100
Mean electric field strength, Vlm F ~ u6.I ~ Effect of applied voltage on htime taken fa hacld and alkaline fronts to meet fw Mentical geomeiries and innial cheml-l compositions (experiments Zn27.7. Znl.10. Zn6.11. Zn9.11. and
Zn15.11). some 60% or more of the original zinc is found as a precipitate a t the position of the pH jump, where its concentration peaks a t some 10 times the initial value. This measured concentration a t the peak is probably attenuated by the finite size of the clay section used for analysis. Efficiency Enhancement. As a consequence of the focusing of the zinc at the pH jump, measured removal efficiencies were low a t from 2 to 10%. To remove the zinc a t an electrode well, the isoelectric focusing effect must be eliminated. This would occur if theelectroosmotic flowrate weresufficienttoprevent the hydroxylions from entering the soil region. For lower electroosmotic flow rates, hydroxyl ions can he prevented from entering the Envkon. Sol. Techmi.. Vol. 28. No. 12. i994
2207
14 I
I
I
i
I
I
a. pH
10 l2
I,
Model A theoretical model that was previously developed to simulate the transport process (2) has been extended to include chemical equilibrium reactions of the heavy metals (26) and used to simulate the experimental system described here. In this model, the overall transport rate is written
6 -
2
0 ' 0.0 10
I
I
I
I
0.2
0.4
0.6
0.8
I
I
I
1.o
I
where R is the molar rate of production due to chemical reactions and sorption. The advective (electroosmosis) velocity, U,, was set to zero for this application, and the diffusion and migration velocities, u d and U,, were evaluated using eq 5. The electric field term is obtained from the equation for the current density i
b. Zn concentration 0
9
8
.-0 c E
-
6
8
4
0
I\
* Dissolved zinc
c-
1%
c 0
N
0
0
9 0
2 initial concentration, C,
2
1
0.3 1
I
-13
p H \ /
r.
p--
00 0.0
0.2
'il
0
0.4
0.6
0.8
1.o
Normalized distance from anode Flgure 8. Distribution of pH and total zinc concentration on rinsing the cathode with water (time = 9 days: experiment 27.7).
soil by removing or neutralizing them as they are formed or by promoting conditions for an alternative electrode reaction that does not produce hydroxyl ions. Removing the hydroxyl (and equivalent amount of sodium) ions by simply rinsing the cathode with tap water was found to be effective. On keeping the pH of the cathode well near neutral by rinsing with tap water from a large reservoir, the zinc removal efficiencies increased to around 98%. The final pH and zinc distributions in the clay for the case of the cathode rinse are shown in Figure 8. The pH is reduced to below 2 throughout most of the soil and only approaches the maximum value of 4 near the cathode. Although a peak in the zinc distribution 2208
is still discernible at the position where the pH increases, the peak concentration now reaches only one-fifth of the initial zinc concentration.
Envlron. Sci. Technol., Vol. 28, No. 12, 1994
i = i(t) = - -7 -F dz
2
1
Cz?u,ci- -7 F
ac,
~ ~ ~ D (12) ~ az
and appropriate boundary conditions are applied at the electrodes. The reaction term included all significant species for the zinc-water system (ZnOH+,Zn(OH)a(aq),HZnO2-, and Zn022-). It was assumed that sodium and nitrate were completely ionized, sorption was negligible, and the temperature was constant at 25 "C.Adsorption isotherms and resistive heating effects can be readily included if required (26, 32). The system was assumed ideal (unit activity coefficient), so solubilities and hence mobilities may, in practice, be higher than predicted. There are no adjustable parameters in the model as used for this application. The model and solution technique have been described in more detail elsewhere (26). Here, we show solutions in Figures 9 and 10 for simulations of the experiments of Figures 7 and 8. To model the rinsed electrode, the pH of the cathode well was fixed at 7.0 and did not precisely duplicate the experimental conditions. As seen, the model accurately predicts the final distribution of zinc for operation with and without an electrode rinse. The pH is also predicted satisfactorily, although at high and low values it is overpredicted due to the assumption of unit activity coefficient for the hydrogen and hydroxyl ions. This is particularly noticeable in the case of using an electrode rinse, which introduced additional electrolytes into the system. Solutions that were obtained for other concentrations of sodium and nitrate show that the upper and lower values of the pH are set by the concentration of these background ions. To satisfy the electroneutrality condition (eq 3), the model shows that the excess hydroxyl ions in the high pH region are balanced by sodium ions and the excess hydrogen ions in the low pH region are balanced by nitrate ions. Of some interest is that a t the region where the pH jumps, the ion concentration is very low, and a sharp minimum in the conductivity occurs at this region. The electric field strength is therefore very small at all locations except a t the position of the pH jump. Any ions entering this region (mainly by diffusion) will either precipitate due to the
10
Experiment
- Model
r,
i i
d
1
0
Q0
8
c
0 .-c E c
c
b. Zn concentration
Experiment
6c
- Model
0 0
4
.-N
0.2
0.0
0.4
0.6
1.o
0.8
Normalized distance from anode Flgure 9. Comparison of the distribution of pH and zinc concentration predicted by the model with the experimental data (time = 9 days; adapted from ref 26).
0’4 0
8i
0.3
c
Ec 0 0
c .-
I
Measured Z i concentration
-’ 6
- - Calculated Zn concentration
.‘a
E .-
r. r
0.2
0.1
1 Ic
0 Measured pH
1
.
Calculated pH
0 0-4
0 0
0.0
.-
I,
0
0
c
O
0
0.4 0.6 0.8 Normalized distance from anode 0.2
-2
C 1
1.o
Flgure 10. Calculated distributlon of pH and zinc concentration for the cathode rinse case (time = 9 days; adapted from ref 26).
high pH or rapidly migrate away due to the high electric field strength (eq 5 ) . In regions removed from the focusing point, the electric field strength is small, migration velocities are low, and concentrations do not change significantly. Consequently, not only does the isoelectric focusing effect trap the metals in a well-defined region, it also effectively “switchesoff“the overall transport process.
Conclusions The presence of naturally occurring matter and codisposed substances can significantly affect the removal
of heavy metals by electrorestoration. The process itself can further affect the mobility by imposing changes in pH and redox potential on the system. These effects might not be observed in laboratory tests with single metal contaminants and relatively pure soils, but probably account for the anomalous results obtained in the field. Reproducing site conditions in the laboratory is difficult; for example, exposing soil samples to the atmosphere can alter the carbon dioxide saturation and redox potential. Equilibrium models can be used to determine the speciation of metals in complex geochemicalmixtures, and if presented as Pourbaix diagrams, are particularly useful for rapidly identifying conditions that might enhance mobility. Although these models serve as a useful guide, allowance must be made for kinetic constraints and for substances such as organic matter that might not be adequately handled. Metal speciation chemistry has been incorporated into the theoretical model for electrorestoration. In the case of the relatively simple zinc system studied here, the model accurately simulates the accumulation of the metal in the soil that occurs by isoelectric focusing. The model also clarifies the role of electroneutrality and the background electrolytes in establishing the pH profiles. Both the model and experiment show that problems related to isoelectric focusing can be prevented simply by rinsing away the hydroxyl ions generated at the cathode and results in better than 95 96 zinc removal. To simulate specific site conditions, the model must be extended to include the chemistry of soil electrolytes, organic matter, and co-disposed substances and should also include thermal effects. The accuracy of predicting field performance is limited by our ability to describe the chemistry of the complex waste-geochemical mixtures. However, enhancing techniques that are readily incorporated into the electric field process such as acidifying the soil or introducing chelates should in principle result in a sufficiently robust mobilization mechanism that nonequilibrium and organic interactions become less important. It is concluded that the excellent removal efficiencies that have been obtained in the laboratory can be achieved in the field provided mobilization problems are adequately addressed. In this respect, the model described here is expected to provide a useful tool for laboratory and pilot-scale treatability studies.
Acknowledgments Supported in part by the Office of Technology Development within DOE’S Office of Environmental Management, under the In Situ Remediation Integrated Program; the US. Environmental Protection Agency Northeast Hazardous Substance Research Center at New Jersey Institute of Technology; and Southern California Edison Company; with additional support provided by Stadtwerke Dusseldorf AG. We thank R. F. Probstein, M. Z. Sengun, R. A. Jacobs, and J. Dzenitis for their contributions and help in preparing this paper. The support of S.T. by the Deutsche Forschungsgemeinschaft is also acknowledged.
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Received for review April I , 1994. Revised manuscript received July 21, 1994. Accepted August 8, 1994." Abstract published in Advance ACS Abstracts, September 15, 1994.
