Environ. Sci. Technol. 7994, 28, 443-457
Interaction of Municipal Solid Waste Ash with Water Carl S. Kirby'it and J. Donald Rlmstidt Department of Geological Sciences, Virginia Polytechnic Institute and State University, 4044 Derring Hall, Blacksburg, Virginia 24061-0420
Municipal solid waste (MSW) incineration ash is composed mainly of glasses and common minerals. We examine the dominant chemical reactions occurring between water and MSW ash using batch reactors. The ash-water solutions are dominated by ions released by soluble salts. X-ray diffraction documents the dissolution of soluble salts and the precipitation of at least one secondary alteration mineral. Three types of reactions are identified. (1)After rapid exhaustion of soluble salts, sodium and potassium exhibit nearly steady-state behavior due to the slow release of ions from less soluble minerals and glasses. Approximately 10 kg of anhydrous salts could be precipitated from a solution contacting each metric ton of ash. (2) Calcium and sulfate concentrations are controlled by either gypsum or anhydrite equilibrium after a few hours. Iron, aluminum, and manganese concentrations rapidly equilibrate with respect to hydroxide or oxide solid phases. (3) Silicon clearly shows temperature-dependent kinetic behavior, but its rate of release into solution is slowed by back-reaction of a secondary silicate phase.
Introduction Municipal solid waste (MSW) incineration ash is composed mainly of glasses and minerals, materials that geochemists have long studied in natural materials. We apply a geochemical approach to this synthetic material which is considerably enriched in trace metals compared to average soils and rocks. Much has been written about the leaching behavior of ash. We have chosen to consider ash-water interactions as an alteration phenomenon, thus our emphasis is on the solution in contact with minerals and glasses in ash rather than the solution that exists an ash disposal site. Of course, the question of greatest environmental concern to be answered is "How much of a potentially harmful species exits the disposal site?". Once we better understand the chemistry which controls solution compositions, we will better be able to answer this question and predict the composition of leachates. There have been numerous studies of MSW ash (see refs 1-3 for reviews). These studies were preceded by an even larger number of studies on coal ashes (see 4 for a review),which have many similarities to MSW ashes. Many leaching tests were designed by regulatory agencies to characterize mainly trace element mobility. The E P tox (US.Environmental Protection Agency, EPA), TCLP (EPA), WET (California), and MWEP (EPA) procedures are various aggressive techniques used to estimate the amounts of toxic trace elements "available" for leaching. Because they are designed to simulate co-disposal of ash with uncombusted waste, the E P tox, TCLP, and WET tests probably overestimate the leachability of an ash disposed of in a monofill. A wide variety of batch and column experiments have modeled leaching behavior in Present address: Geology Department, Bucknell University, Lewisburg, PA 17837; e-mail address:
[email protected]. 0013-936X/94/0928-0443$04.50/0
0 1994 American Chemical Society
nonstandardized leaching studies (5-18). Silicon, which may be important in the precipitation of secondary mineral phases, has notably been excluded from analyses in most of these studies. Little field data are available on solutions from ashfills; three exceptions include refs 19-21. The purpose of this research was to determine the types of chemical reactions that occur when water comes into contact with MSW ash. We concentrate mainly, but not exclusively, on major elements because master variables pH, oxidation-reduction state, and concentrations of most solution components are determined by geochemically abundant elements. In turn, trace element behavior depends on pH, redox state, ionic strength, and presence of sorptive solid substrates-factors which are strongly dependent on major element geochemistry. These reactions include (1) complete dissolution of very soluble phases, (2) partial dissolution of sparingly soluble phases, (3) achievement of partial equilibrium between some of the solid phases and the solution, (4) precipitation of new phases, and (5) surface adsorption and desorption. We designed our experiments to examine reactions 1-4; we did not investigate surface reactions. We have studied a combined bottom and fly ash from a single facility, but this geochemical approach is also applicable to the leaching/weathering behaviors of other waste forms, volcanic ashes, rocks, and soils.
Methods The ash used in this study came from the University City Resource Recovery Facility (UCRRF) in Charlotte, NC, which handles approximately 90 000 kg/day of mainly residential refuse. Twin furnaces operate a t approximately 1000 "C. Fly ash is collected by an electrostatic precipitator and mixed with bottom ash on site. The combined ash is then wetted with water for dust control. All ash used in this study was bottom and fly ash combined a t the incinerator. The ash examined in this study is not a representative average sample of UCRRF ash because spatial and temporal variations due to extreme heterogeneity of the refuse stream are not addressed by our grab samples. Kirby and Rimstidt (22) describe the composition, mineralogy, and surface properties of ash from this study in greater detail. Chemical composition data are presented in Kirby and Rimstidt (22). Concentrations of major minerals and glass are presented in Table 1. Three different types of experiments were performed to determine the nature of ash-water reactions. First, the soluble salt content of the ash was determined by measuring the conductivity of successive aliquots of water that were reacted with an ash sample for a short time. The second and third experiment types involved reacting deionized water-ash mixtures in batch reactors at temperatures between 20 and 60 "C. Elevated temperatures were employed to speed up reaction kinetics and thereby simulate longer reaction times. This elevation of temperatures is also used in characterizing the long-term rates of reaction of nuclear waste forms (23). These three aqueous experiments are described further below. Environ. Sci. Technol., Vol. 28, No. 3, 1994 443
Table 1. Lists of Minerals Identified in Unreacted and Reacted Ash* mineral or amorphous material Fez03(hematite) CaC03 (calcite) NaCl (halite) Si02 (quartz) AB204 spinel (magnetite) Ti02 (rutile) CaS0~2Hz0(gypsum) Cas04 (anhydrite) =3(A1203).2(Si02) (mullite) FeO (wustite) KCl (sylvite) K~Ca(SO&-H20(syngenite) glasses
wt % in unreacted ash
present in reacted ash (exp 9 and 10)
3.7 f 1.7 3.5 f 1.9 0.5 f 0.4 2.3 f 1.0 E3.5 f 2 1.1f 1.3 1.8 f 1.9 na na na na na =I2
Yes Yes no Yes Yes Yes Yes no Yes Yes no no Yes
a Approximately 72 % of unreacted ash is amorphous material, Le., glasses of several compositions. A similar percentage of glass is in the reacted ash. Minerals below detection limit using XRD standard additions were not analyzed (na).
To determine soluble salts, 6 g of ash was rinsed and shaken with deionized water, the suspension was centrifuged, the supernatant was poured off, and its conductivity was measured. More deionized water was added to the remaining solid and occluded solution, and the process was repeated up to a cumulative solution volume of 220 mL; the experiment was stopped at this cumulative solution volume because further repetitions were beginning to dissolve less soluble phases such as calcite and gypsum to an appreciable degree. Conductivity data were converted to NaCl equivalent based on a polynomial fit of NaCl from Lobo (24). Unstirred batch experiments were conducted at a 4:l L/S ratio using 6 g of ash and 24 mL of deionized water in PTFE (Teflon) 30-mL screw-top containers a t 42 “C. This temperature was chosen as part of a planned series of 42, 75, and 100 OC experiments; not all of these experiments were continued because it was discovered that the highest temperature experiments produced different secondary solid phases than the lower temperature experiments. Constant temperature was maintained by a circulating water bath. Only one solution sample was collected from each vessel at the end of each experiment. Experiment numbers refers to numbers assigned in Table A. I in the supplementary material; experiment duration is also listed. Stirred (except as noted) semi-batch experiments were conducted a t 20, 40, and 60 OC in either water bath thermostated 1-L polyethylene bottles stirred by a magnetic stir bar or in a reciprocating shaker bath in 1-L glass Erlenmeyer flasks. Each experiment contained 17.5 g of ash and 485 mL of deionized water resulting in a 28:l L/S ratio. We refer to these experiments as semi-batch reactor experiments (see ref 25), because at each sampling time, the stirring was stopped for 5 min to allow settling, a 1525-mL aliquot was removed, and then 15-25 mL of deionized water was injected to replace the lost solution volume. A dilution correction was made before concentrations were reported. Two shaker bath experiments (nos. 5 and 8) give results from a 500-mL flask that did not allow stirring as evidenced by the settling of solids to the bottom of the flask, leaving a clear supernatant. Sample solutions from each experiment were cooled to room temperature and filtered (0.2 pm), and then conductivity and pH were recorded. Samples were acidified to approximately pH 2 using concentrated nitric acid; the 444
Environ. Sci. Technol., Vol. 28, No. 3, 1994
reported data are corrected for this dilution. Solutions were then analyzed for Si, Ca, K, S, Mg, Zn, Mn, Fe, and A1 using inductively coupled argon plasma spectrometry (ICAP) and for Pb, Cd, Cr, As, Se, Ni using graphite furnace atomic absorption spectrometry (GFAAS), We confirmed with ion chromatograph that sulfur is in the form of sulfate. Solids were sampled at the end of each experiment. Alkalinity titrations were performed on selected 28:l and 4:l L/S ratio experiments. A t the end of experiments 3,4,5,and 8, Le., after the last solution sample was collected for cation and anion analysis, the remaining solution was titrated with 0.01 N HC1. Titrations were also performed on two 120-h and two 288-h batch experiments a t 4:l L/S ratio; these four batch experiments were not part of the series of experiments listed in the supplementary material. The geochemical computer code WATEQ4F (26) was used to estimate saturation states for solutions with respect to solid phases. WATEQ4F uses a combination of the Davies equation and the WATEQ Debye-Huckel equation for the calculation of ion activity coefficients. The system was modeled as not open to atmospheric COz. Speciation was adjusted for the different temperature experiments, but pH values were recorded a t room temperature (22 f 1 “C) after quenching from experimental temperatures. Precipitation of solids was not allowed. WATEQ4F calculations were carried out on solution compositions averaged for times greater than 24 h for like experiments, e.g., for experiments 3 and 4, data gathered from the 48509-h samples were averaged together. Bicarbonate (HCOs-) values are those listed in the results section. E h values were based on redox potential measurements made with a platinum electrode on experiments 6-8 and 17 and 18. Two 40 OC, 28:l L/S ratio experiments (nos. 9 and 10) were sampled after 720 hand combined; this material was selected as a “typical” reacted ash for X-ray analysis. Solids were analyzed a Scintag powder X-ray diffractometer automated by the Scintag Diffraction Management System. Slurries of ash ground in acetone for at least 5 min were mounted on glass slides. Copper K, radiation was used with an Ni filter. The accelerating voltage was 45 kV, and the current was 35 mA. Data were recorded digitally, and peak position and intensity were determined either on screen or by using the peak-finder feature in the software. All mineral identifications were based on a manual search of the JCPDS data file guided by operator knowledge of the bulk chemistry. The heavy liquid s-tetrabromomethane (BrZHC-CHBr2, sp gr 2.96) was employed to isolate the densest phases. A simple handheld electromagnet was used to separate magnetic materials.
Results Figure 1 is an X-ray diffractogram of bulk unreacted and “typical” reacted ash. The peaks are labeled based on definite identification of mineral phases from fractions described in Kirby and Rimstidt (22) rather than from this pattern alone. Overall, the two patterns in Figure 1 are very similar, indicating little change in mineralogy during the experiment. The soluble salts halite and sylvite both dissolve during the experiment and, therefore, are not present in the reacted pattern. Open arrowheads indicate the lack of a peak in one pattern that is present in the other pattern. A t approximately 15,18, and 42 “C, three peaks appear in the reacted pattern that are not present in the unreacted pattern. These peaks document
g gypsum h hematite ha halite m magnetite mu mullite q quartz r rutile s sylvite
Reacted, Exp. #9 & 10
30
10
70
50
20, O
Flgure 1. X-ray diffractograms of unreacted and reacted ash. Open arrowheads indicate a peak is lacking in that pattern which is present in the other pattern. mt, h mt magnetite h hematite wu wiistite q quartz
10
25
40
55
conductivity experiment theoretical dilution
E
70
29," Flgure 2. X-ray diffractogram of a magnetic separate of reacted ash.
the formation of a t least one unidentified alteration mineral during the experiment. Figure 2 is a diffractogram of the reacted heavy nonmagnetic fraction, separated using s-tetrabromomethane. This fraction was examined in an effort to concentrate and identify any Pb-rich mineral phases. Though we have documented a t least two Pb-rich minerals in the unreacted ash with scanning electron microscope (SEM) (29, only quartz, magnetite, hematite, and wustite were definitely identified in this fraction; all four of these minerals are present in the original ash (22). Any Pb-rich minerals present in the unreacted ash are not in high enough abundance to be identified using powder XRD. The majority (by weight) of the minerals in this ash occur commonly in nature. Minerals identified in the unreacted and reacted ash are listed in Table 1. See Kirby and Rimstidt (22) for details on physical and chemical separations and method of quantification. Cation, anion, pH, and conductivity data from aqueous experiments are shown in the supplementary material.
40000
30
80
130
180
230
CUMULATIVE SOLUTION VOLUME, mL Figure3. Concentration of NaCl equivalent (determinedby conductivity measurement) versus cumulative solution volume.
Figure 3 shows the results from the conductivity experiment. Conductivity is reported as the concentration of NaCl equivalent versus the cumulative volume of solution. Alkalinity of solutions averaged 110 mg of HC03-/L for 28:l L/S ratio experiments and 230 mg of HC03-/L for 28:l L/S ratio experiments. Redox potentials averaged 0.39 V i 0.04 (2a).
Discussion The ash-water solutions are dominated by ions released by soluble salts. Figure 3 shows conductivity reported as the concentration of salts (NaC1 equivalent) versus the cumulative volume of solution. Halite (NaCl), sylvite (KCl), and syngenite [K2Ca(S04)yH201 are the predominant soluble salts in the ash in this study. Syngenite was not reported in Kirby and Rimstidt (22), but we have subsequently identified it in an XRD pattern of reprecipitated salts. Gypsum (CaSOq2H20),anhydrite (CaSOd), and calcite (CaC03) are relatively soluble sources of Ca2+ Environ. Sci. Technol., Vol. 28, No. 3, 1994
445
Table 2. Means and Standard Deviations (SD) of Concentrations (in mg/L) for Selected Elements from All Experiments before and after Application of Chauvinet’s Criteriona
all data N Pb Cr Cd Se Zn As Ni Fe Mn A1
mean
using Chauvinet’s criterion SD
N
mean
SD
A% mean
71 0.039 71 0.0063 71 0.0063 18 0.0075 70 0.036 18 0.0035 18 0.0055 76 0.07 92 0.008 93 9.0
0.054 66 0.025 0.017 36 0.0035 67 0.0057 0.0027 10 0.0086 65 0.0039 0.0028 38 0.0026 17 5 0.0071 0.0020 0.025 72 0.035 0.023 3 0.0036 16 31 0.0024 0.0016 0.0015 16 0.0051 0.0011 7 73 0.2 0.04 0.03 43 87 0.004 0.008 0.003 0 6.0 85 7.6 4.6 16 a N is the number of observations; A % mean is the percentage change in the mean upon application of Chauvinet’s criterion.
ions. These minerals therefore account for the majority of the observed salinity. The contribution due to soluble salts vastly overwhelms any contribution from other minerals and glasses because the rates of dissolution of soluble salts are several orders of magnitude faster than the rates of dissolution of silicate minerals and glasses. The curve labeled “theoretical dilution” in Figure 3 was calculated assuming that all of the conductivity was due to immediate dissolution of NaCl with no subsequent addition of ions to solution. Comparison of the observed data and the theoretical curve suggests that ions continue to be released from the ash to solution even after the most soluble salts dissolved. Dissolution of gypsum, anhydrite, and calcite are probably responsible for most of the conductivity after the more soluble salts dissolve. A common problem for “real-world’’ chemical experiments, i.e., experiments which involve complex systems of reactants and products, is that many variables are unconstrained. For experiments with ash, the heterogeneity of this material leads to a wide variation in solution composition. We “homogenized” our samples by using a riffle splitter (22)in order to increase the likelihood that all aliquots of ash were as similar as possible. Nevertheless, significant variations in solutions compositions occur. Explanations of this variation include (1) analytical uncertainty, (2) actual variation in the composition and mineralogy of the samples, and (3) analytical mistakes. Analytical uncertainty is unavoidable, but it is probably random and on the order of a few percent a t most for ICAP and GFAAS. Explanations 2 and 3 are systematic errors and difficult to eliminate. We have no reason to suspect that mistakes contributed to our uncertainties. We do have evidence for the heterogeneity of some samples even after our attempts at homogenization. For example, experiment 8 displayed anomalous behavior, with lead concentrations nearly an order of magnitude higher than all other experiments. Follow-up experiments did not reproduce such high lead concentrations. This lead anomaly was the most extreme example of sample heterogeneity. Such variations due to sample heterogeneity are real, and we do not wish to dismiss their importance. However, as a first approximation, the concentrations of Pb, Cr, Cd, Se, Zn, As, Ni, Fe, Mn, and A1 do not vary greatly with time, and except for Se, As, and Ni for which we have only 40 “C data, vary little with temperature. Therefore we present means and standard deviations for these elements in Table 2, which lists these statistics before and after the 446
Environ. Scl. Technol., Vol. 28, No. 3, 1994
TIME Figure 4. Schematic diagram concentration versus time behaviors observed in ash-water alteration experiments. For explanation of numbers, see General Dissolution Behaviors section.
application of Chauvinet’s criterion (28). We have not reported means and standard deviations for elements which show clear kinetic behavior (Si, Mg) nor for elements which show significant variation with L/S ratio (Ca, S, Na, K). It should also be noted that some of these elements can show significant dependence on pH (18). One further complication that we do not address is the issue of below detection limit (bdl) analysis. A small number of Fe and Zn analyses returned bdl values. Helsel (29) presents a rigorous approach to bdl analyses. We have used Chauvinet’s criterion (28)to identify data which are significantly different from the mean concentration. Chauvinet’s criterion is a means to rationalize the rejection of data, and while some may disagree with the rejection of any data, we have used this method in order to give what we consider to be a reasonable estimate of mean values and standard errors. The following steps are followed in Chauvinet’s criterion: We make N measurements of the quantity x and calculate
where xsusis the suspected outlier, tausis the number of standard deviations by which xsusdeviates from 2 , and u, is the standard deviation. We then calculate N times the probability that xBuslies outside t,,,a, n(worse than x,,,) = N.P(outside tsu,ux)
(2)
which gives n,the number of measurements expected to fall at least as far from the mean. If n is less than l/2, then xaus fails Chauvinet’s Criterion and is rejected. Any rejected outliers are in bold print in Table A.1 (supplementary material). General Dissolution Behaviors. Figure 4 is a schematic diagram of concentration versus time illustrating the types of concentration versus time behaviors observed in our experiments. Precipitation effects are ignored in this discussion for clarity. Closed symbols indicate kinetic behavior; open symbols indicate equilibrium or steadystate behavior. Four scenarios are represented in this diagram. (1)Dissolution of a phase until it is exhausted; the system does not reach equilibrium with respect to this phase. (2) Dissolution of a phase until equilibrium is attained; the concentration is independent of the remaining amount of the solid phase. (3) Dissolution of a phase that never reaches equilibrium or steady state over the course of the experiment; the concentration continues to increase over the entire experiment. (4) Rapid dissolution of a phase to exhaustion (behavior 1)and simultaneous slower dissolution of a less soluble phase toward equilibrium or
steady state (behavior 2); the concentration rises quickly at first and then much more slowly after the more soluble phase is exhausted. One cannot tell the difference between the first and second behaviors without some independent information such as a calculation of the saturation state of the solution with respect to the solid phase. Similarly, it may be difficult to tell which of the three behaviors controls solution composition if the slope is not significantly different from zero. Interpretations of concentration versus time plots may be confounded by variations in temperature and L/S ratio. In attempting to deduce reaction controls with this approach, caution must be employed to avoid drastically changing saturation states of minerals due to temperature differences and also due to dilution in semi-batch experiments. For example, the rate of calcite precipitation is known to be a function of temperature and saturation state (30),and calcite solubility is also a function of temperature (31 ) . In our experiments, silicon concentrations and dissolution rates show strong dependence on temperature. For most other elements, such temperature dependence is absent or hidden in experimental error. For example, iron concentration appears to be controlled by the equilibrium solubility of a solid phase which has little dependence on temperature. Calcium concentration also showed little temperature dependence but did show dependence on LIS ratio. Specific Behaviors: Silica. Silicon clearly exhibits kinetic behavior (behavior no. 3 from Figure 4) for the duration of our experiments, Le., the concentration of silicon increases with time throughout each experiment. Silicon is released to solution through the dissolution of uncombusted glasses, glasses formed in the incinerator, quartz, and other silicates in relatively low concentrations. The release of ions from glass can be represented by Si02(glass) + 2H,O = H,SiO,
Mn+(g,ass) + rz-H20 = M"+(aq) + nOH-
(3)
(4) where M is a metal ion in the glass matrix. An an example of silicon behavior, Figure 5a shows the concentration of silicon (not Si02 versus time for 40 O C experiments (nos. 6-12). In quartz dissolution experiments (32)and in glassy siliceous tuff dissolution experiments (33),silicon concentrations were considerably higher than ours after accounting for different LIS ratios. For example, in Henne's (33)study, the lowest amount of silicon released was 1.4 X mol of Si/g of sample in 25 "C experiments a t 48 h. In comparison, the maximum amount of silicon released in our experiments was 7 X lo4 mol of Si/g of sample. Our solutions were always well below saturation with respect to quartz and amorphous silica. There is apparently a precipitation reaction occurring in our experiments which occurs slightly more slowly than the dissolution reaction, thus allowing the silicon concentration to continue to rise, but much more slowly than in either Rimstidt and Barnes (32) or in Henne (33). We cannot compare rates directly because we did not perform a surface area determination of this ash. Precipitation of a silicate phase is likely in our 20-60 "C experiments because of the abundance of other cations (Ca, Na, Al, Mg) which can combine with silicon to form zeolite or clay minerals. We also note the appearance of two X-ray peaks between 10 and 20 OC 2 8 in Figure 1,which suggests that precipitation of at least one secondary phase has occurred in the 40 OC experiments. In analogous experiments, Crovisier et al.
'C-I
500
0
lO0OC 1
'
60°C
20°C
40°C
I
.
1500
1000
TIME, hr I
I
1
-19.5-
-
-20.5-
E
-21.5-
-22.5 4 2.8
3.0
3.2
3.4
3.6
1000/T Figure 5. (a, top) Concentration of silicon versus time for 40 O C experiments (no. 9-12). (b, bottom) Natural logarithm of the net rate constant k versus lOOO/T, showing Arrhenius relation for silicon dissolution.
(34)observed the formation of phyllosilicate (clay) minerals which limited silicon concentrations in their 60°dissolution of synthetic basaltic glass in seawater. The following equation for the formation of an illite clay
3.5H4Si0, + 1.2H'
+ 2.3Al(OH); + 0.25Mg2++
is an example of a "competing" reaction which could take up silica from solution. WATEQ4F calculations suggest that both 4:l and 28:l solutions are supersaturated with respect to illite and kaolinite [Al~Siz05(OH)41.Thus, the rates of silicon release from our experiments are apparent rates that reflect the combined rates of silicon release from primary glasses and minerals (eq 1) and the uptake of silicon into alteration minerals (such as those shown in eq 2) formed during the experiment. The net or apparent rate of release of silica release, rnet, is given by
-- r+ - r-
(6) where r+ is the rate of dissolution of silica and r- is the rate of precipitation of the secondary silicate phase. If we assume that surface area does not change significantly during an experiment and that both dissolution and precipitation reactions are zero order, then we obtain rnet
'net C- 'net (7) where k n e t is the rate constant for the combined reactions. This approximation allows us to use the Arrhenius equation
k = A exp(E,/RT)
(8)
and its linearized form In k = In A + (E,/R)(1/ !lJ (9) where k is a rate constant, A is the pre-exponential factor, Envlron. Sci. Technol., Vol. 28, No. 3. 1994 447
I
8
z l
I
60°C
I
100
0
500
1000
TIME, hr
1500
Figure 6. Concentration of sodium versus time for 20, 40, and 60 "C experiments (nos 1, 2, 6-12, and 14-18).
E, is the activation energy for reaction, R is the gas constant, and T is the temperature in Kelvin. In order to determine the apparent rates for silicon release, we fitted a line to the 24-252-h data for each temperature and took the slopes of these lines. We chose this data range because the slopes in concentration versus time space are nearly linear after 24 h, suggesting that these rates should be in effect over relatively long periods of time. In addition, since the slopes are nearly constant, the rates are also, and thus the assumption in eq 7 appears reasonable. These rates are plotted in Figure 5b versus reciprocal temperature. The three rates for 20,40, and 60 "C are fitted to a straight line in this Arrhenius plot. Assuming no change in reaction mechanism, lower landfill condition temperatures (12 "C) should result in an apparent rate of silica release of 7.7 X LO-14 mol of Si g1 S-1.
The activation energy for silica dissolution in our experiments is 43 kJ mol-'. This value represents a composite activation energy because the reaction rates are apparent rates combining the effects of dissolution of glass and precipitation of an secondary phase. Henne (33) reports E , = 15 k J mol-l for the dissolution of glassy tuff; Rimstidt and Barnes (32)report E , = 61-65 k J mol-l for the dissolution of amorphous silica. Soluble Salts. Ions predominantly from soluble salts show behavior no. 4 from Figure 4, showing negligible effect from changing temperature. The soluble salts dissolve completely by the reactions NaCl(s) = Na+ + C1-
(10)
KCl(s) = K+ + Cl-
(11)
+
K2Ca(S04),(s).H20= 2K+ + Ca2+ 2S042-+ H 2 0 (12) quickly releasing Na, K, Ca, S, and C1 to solution. As an example, the concentration of sodium is shown in Figure 6. Within minutes, the soluble salts dissolve completely, and the concentrations increase much more slowly as different solid phases such as carbonates, sulfates, silicate minerals, and silicate glasses are tapped. The behaviors of sodium and potassium correspond to the square points in Figure 4; soluble salts are quickly dissolved, and less soluble phases then slowly release sodium and potassium to solution. Chloride must exhibit similar behavior because most chloride will be found in soluble salts rather than glasses or silicate minerals, and there are no major phases to serve as sinks for chloride. Belevi et al. (6) find this behavior for chloride in batch experiments with MSW bottom ash. 448
Environ. Sci. Technol., Vol. 28, No. 3, 1994
The leaching of soluble salts is a potentially significant environmental and engineering problem in the utilization or disposal of MSW ash. Ions from these dissolved salts may contaminate groundwater near the disposal environment. These ions will also promote corrosion where ash is utilized as a building material. We estimated the weight percent (f2a) of halite as 0.5 f 0.4 and of gypsum as 1.8 f 1.9 in ash used in this study (22). Syngenite and sylvite appear as reprecipitated salts but are below the detection limit for our XRD standard addition analysis. It may be possible to recycle reprecipitated salts washed from MSW ash either a t an incinerator or a disposal facility. In Table 3, we use the dominant ions for a nominal average solution composition from a 28:l L/S ratio experiment to estimate the mass of salt which could precipitate from such a solution. The salt components that we show are not necessarily the actual phases that would precipitate from this solution; the minerals which actually precipitate will be determined by temperature-dependent phase equilibria. Rather, the salt components are analogous to a normative recalculation such as that used in igneous petrology (see ref 35). The approach used in Table 3 could be applied as a first approximation to determine the amount of salt which could be collected as a byproduct of MSW incineration, for example, as road de-icing salt, if trace metals are a t acceptably low levels in the salts. Salt content of MSW ash will vary considerably with the type of pollution control device at a particular incineration facility: facilities with stack scrubbers will have considerably higher salt contents. This simple analysis could be performed on ash from different incinerators to estimate the potential for salt recovery. Iron, Aluminum, and Manganese. Whereas silicon concentrations varied systematically with both time and temperature, these variables have little effect on iron concentration, and therefore iron concentrations appear to be controlled by an iron phase which reaches equilibrium with solution rapidly and whose solubility product has little dependence on temperature [though some concentrations from the 20" samples (experiments 1 and 2) fell under the detection limit]. Iron concentrations rapidly reach a nearly constant value, thus showing behavior no. 2 from Figure 4. For these reasons, we average iron concentrations from all experiments together. When five outliers are omitted based on Chauvinet's criterion (281, the mean iron concentration for all experiments is 0.06 f 0.08 mg/L. Mullen (19)reports similar iron concentrations in a field study of bottom ash over a 3-month period with an average value of 0.025 f 0.027 mg/L. WATEQ4F calculations on our solutions suggest that all aqueous iron is present as trivalent hydrolysis products and that all solutions are supersaturated with respect to ferrihydrite, crystalline Fe(OH)3 (SI = 1.4). Figure 7 shows the solubility of Fe3+and its hydrolysis products in equilibrium with amorphous Fe(OH)S, which is slightly more soluble than ferrihydrite. Thermodynamic data for these calculations are from Baes and Mesmer (36)and Snoeyink and Jenkins (37). Comans et al. (18) find similar supersaturation with respect to solid phases. WATEQ4F calculations should account for any inorganic complexation. Possible explanations for the apparent supersaturation include organic complexation, control by a more soluble phase, or kinetic inhibition. Organic complexation cannot be discounted because we do not have data on organics in the experimental solutions. Control by a more soluble phase would require the presence of a solid phase with a
Table 3. Section a: Nominal Average Solution Composition Leached from Ash at 25 “C for 1Day average concentration ion, ppm ( 2 8 1 LIS) Caz+ Na+ sop c1-
K+ 184
377
306
1257
Hco3-
360
110
Section b Estimation of “Normative” Salt Components Recoverable from This Solution’ salt component precipitated NaCl Cas04 KzCaS04 NaHC03 3.9 x 10-5 2.3
mol precipitated/g of ash kg of anhydrous salt component/t of ash no. of days to produce 1t of salt component (I
11
2.8 x 10-5 3.8 7
9.4 x 10-6 2.0 12
NazS04
7.2 X lo4 0.6 41
6.9 X lo4 1.0
25
Assuming 200 t/day incinerator and 20% of original mass present after incineration.
4
6
Ph
10
12
Flgure 7. Common logarithm of metal activity versus pH for SoiUtiOnS in equilibrium with respect to amorphous iron and aluminum hydroxides. Note apparent supersaturation by approximately 1 order of magnitude for each metal.
larger solubility product than ferrihydrite. Kinetic inhibition seems unlikely because iron(II1) hydroxides precipitate rapidly. Aluminum and manganese exhibit similar behavior to iron in solution and are also likely controlled by equilibrium solubility. Figure 7 shows aluminum solubility in equilibrium with amorphous aluminum hydroxide [Al(OH)31 with our data superimposed. Thermodynamic data are from Baes and Mesmer (36) and Snoeyink and Jenkins (37). Our data show aluminum concentrations approximately 1 order of magnitude higher than concentrations in equilibrium with amorphous Al(OH13. As with iron, possible explanations for the apparent supersaturation include organic complexation, control by a more soluble phase, or kinetic inhibition. From Figure 7 it can be inferred that higher and lower pH values than were observed in our experiments could result in significantly higher iron or aluminum concentrations in solution. Calcium and Sulfate. Some leaching behaviors cannot be explained solely by kinetics, exhaustion of a phase, or equilibrium. Rather, some behaviors are combinations of these three controls on solution composition. Anhydrite is present in ash immediately upon combustion, but hydrates quickly to form gypsum upon exposure to water. Ca2+and S042-are also released to solution predominantly by soluble salts (eq 13) and by release from gypsum and calcite. Gypsum and calcite react according to CaS0,.2H20(s) = Ca2++ SO:-
+
+ 2H,O
(13)
CaC03(s) H+ = Ca2++ HCO; (14) Calcium and sulfur concentrations follow behavior no. 4 in Figure 4. These concentrations rise very rapidly initially, reflecting the dissolution of soluble salts. Sub-
*?
F!
v!
0
log
n
i
ac,2+
Flgure 8. Common logarithm of sulfate activity versus calcium activity for all experiments. Univariant lines for gypsum and anhydrite are shown. Time increases toward the upper right corner.
sequently, the concentrations show little or no increase with time in either the 28:l or 4:l L/S ratio experiments. WATEQ4F calculations show for both 28:l or 4:l L/S ratio experiments after 2 days of time, solutions were supersaturated withrespect to calcite and aragonite, both CaCO3 minerals, with saturation indices (SI) of 1.3-1.7, respectively. The SI is loglo (IAPIK),where IAP is the observed ion activity product, in this case ucaz+aco3z-,and K is the equilibrium product. Undersaturated solutions have negative values for SI, supersaturated solutions have positive values, and exactly saturated solutions have an SI of 0. The SI is estimated as -0.22 for gypsum and -1.0 for anhydrite for the 28:l L/S ratio experiments and as +0.03 for gypsum and -0.21 for anhydrite for the unrinsed 4:l L/S ratio experiment. These estimated saturation states suggest that the more dilute solutions are slightly undersaturated with respect to both minerals, whereas the more concentrated solutions are in equilibrium with gypsum. Figure 8 shows univariant lines for gypsum and anhydrite in equilibrium with solution. Our data are plotted, and in general, time increases and L/S ratio decreases to the upper right corner. Calcium and sulfur enter the solution in equimolar amounts until reaching equilibrium with gypsum or anhydrite. Though the solutions are supersaturated with respect to calcite and aragonite, these phases apparently do not control calcium concentrations. Both 28:l and 4:l L/S ratio experiments are significantly undersaturated (SI = -7) with respect to portlandite [CaEnviron. Scl. Technol., Vol. 28, No. 3, 1994
449
da
1 -
2 zoJ I
Table 4. Postulated Controls on Solution Concentration for Selected Elements
.
z
8 c g z
40' unstirred
looi __
EPA drinkrng water MCL
-
-.
..
0
0
500
1000
behavior no. in Figure 4
element
20" unstirred
Si
3
K, C1 Ca, S
4 4 or 2
Fe,Al,Mn
2
Mg
2? or 3?
1500
(0H)zI which may control calcium in some instances (18). Temperature has a negligible effect on the calcium and sulfur concentrations. At L/S ratios likely to be encountered in the disposal environment (20) (approximately 0.01-0.11, calciumand sulfur should be controlled byeither gypsum or anhydrite. In bottom ash experiments, Comans et al. (18) find calcium and sulfur controlled by gypsum solubility, with calcite not controlled by calcite, though gypsum solubility should not explain calcium and sulfur concentrations a t pH values greater than approximately 10. Magnesium displays both kinetic and equilibrium behavior; the systematics for this behavior are not clear. All of the 4:l L/S ratio experiments (experiments 13 and 14) show apparent equilibrium control (behavior no. 2 in Figure 4) at magnesium concentrations below 5 mg/L. Two 40 OC 28:l LIS ratio experiments (nos. 11 and 12) also show this equilibrium behavior. In contrast, five other 40 "C 28:l L/S ratio experiments (nos. 6-10) show kinetic behavior (behavior no. 3 in Figure 4). All of the 20 and 60 OC experiments show kinetic behavior. Comans et al. (18)suggest brucite (Mg(0H)d as a controlling phase at higher pH values, but they do not suggest such a phase for lower pH conditions. Our solutions are undersaturated with respect to brucite according to WATEQ4F calculations. Trace Elements. We did not identify soluble lead salts in the solid phase in ash from this study. Henry et al. (38) found anglesite (PbS04) in MSW fly ash. Other likely sources of lead include lead chloride, cerussite (PbC03), and lead in solid solution with calcite or gypsum. SEM observation of a heavy separate (by s-tetrabromomethane) of unreacted ash shows lead-rich silicate glasses and leadrich mineral phases; however, these silicates will release lead to solution only very slowly. Figure 9 shows concentration versus time data for lead. For most experiments, lead concentrations decreased slowly with time from as much as 57 pg/L to between 4 and 30 yg/L; most experiments yielded lead concentrations near the 15 pg/L EPA maximum contaminant level (39). The 40 "C unstirred semi-batch experiment (no. 8) conducted on unrinsed ash showed unique behavior, with lead concentrations nearly an order of magnitude higher than all other experiments, Follow-up experiments did not reproduce such high lead concentrations. We attribute this high lead value to sample heterogeneity. Approximately 40% of the lead in ash from this study is acid-soluble and therefore probably sequestered in carbonates (22). Because our solutions are supersaturated with respect to calcite, very little of this fraction of the 450
Environ. Sci.
Technoi., Vol.
28, No. 3, 1994
temperature-dependent combined kinetics of glass dissolution and secondary-phase precipitation availability of soluble salts availability of soluble salts or anhydrite solubility at low LIS ratio pH-dependent hydroxide/oxyhydroxide solubility hydroxide? solubility or dissolution kinetics
_-
TIME, hr Flgure 9. Concentration of lead versus time for selected experiments. US. EPA maximum contaminant level (MCL) is 15 pg/L.
likely controlling mechanism
lead is likely released in our experiments. Neither temperature nor pH could be clearly related to lead concentrations. It is unclear as to whether solid-phase solubility or surface adsorption controls lead concentrations. Comans et al. (18) found that lead concentrations were difficult to explain using equilibrium with a solid phase. Eighmy et al. (40) suggest that the addition of Pod3-might control lead concentrations at very low levels by precipitation of chloropyromorphite [Pb5(P04)3Cll. WATEQ4F calculations suggest that our solutions are supersaturated with respect to chloropyromorphite
+
5Pb2+ 3POd3-+ C1- = Pb,(PO,),Cl(s)
(15)
(SI = 4.5 and 3.1 for 28:l and 4:l L/S experiments, respectively) with no Pod3-added to the ash. Solutions are undersaturated with respect to cerussite and anglesite (SI = -1.3 and -4, respectively). The fact that lead concentrations usually decrease suggest that kinetics play some role in lead mobility. Lead concentrations in all experiments, stirred and unstirred, deceased over time; therefore, lead either adsorbs onto surface sites, is incorporated into a solid solution phase such as a carbonate, or forms its own phase such as chloropyromorphite. Data for other trace elements are presented in Table A.1 (supplementary material). Trace elements are not the major focus of this study, so we have not focused on discerning mechanisms of trace element control. Trace elements show considerable variation in their solution behaviors, thus we have not attempted to categorize the behaviors of trace elements according to the outline shown in Figure 4. We did note however that cadmium behaves much like lead in solution, i.e., its concentration decreases with time. Also like lead, the cadmium concentration of experiment 8 was much higher than all other experiments. Conclusion We observed four basic types of solution behavior with time: equilibrium, kinetic, availability (lack of concentration change due to the exhaustion of a phase), and a combination of availability and kinetic behaviors (see Figure 4). Silicon clearly exhibits kinetic control, but the kinetics of silicon release to solution are slowed by the precipitation of a secondary phase. Sodium and potassium, associated with soluble salts, are controlled by their availability in the solid phase. Apparently the solutions do not reach saturation with respect to any phases which rapidly remove these cations. Manganese, iron, and aluminum are controlled by equilibrium solubility of solid phases. Calcium and sulfur are controlled by kinetics and availability in 28:l L/S ratio experiments. After 1-2 days in 4:l L/S ratio experiments, calcium and sulfur are
controlled by anhydrite solubility. Table 4 summarizes our conclusions concerning the controls on concentrations of the elements which are abundant in ash. Though the leaching of MSW ash in the disposal environment is clearly janalogous to a column leaching experiment, batch reactors can be successfully employed to understand some of the geochemical processes involved in the alteration of ash. This observation should add confidence to those wishing to use geochemical equilibrium computer codes as starting points for modeling the longterm stabilization of MSW ash, especially for major elements. Trace elements are more problematic; accurate forward modeling of trace elements requires further understanding of the basic controls on the concentrations in solution. Acknowledgments
C.S.K. was supported by a National Science Foundation Graduate Fellowship during part of this research. This research was funded by the U.S.EnvironmentalProtection Agency (Grant R818234-01-0) but has not been reviewed by the EPA. Thanks to Nancy Phillips and Marilyn Grender for solution analyses. Two anonymous reviewers provided insightful comments that helped to improve this manuscript. Supplementary Material Available Table A.1 giving the experiment numbers, time, and solution parameters discussed in the Methodssection (3pages) will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper or microfiche 9105 X 148mm, 24X reduction, negatives) may be obtained from Microforms Office, American Chemical Society, 1155 16th St. NW, Washington, DC 20036. Full bibliographic citation (journal, title of article, names of authors, inclusive pagination, volume number, and issue number) and prepayment, check or money order for $10.00 for photocopy ($12.00 foreign) or $10.00 for microfiche ($11.00 foreign), are required. Canadian residents should add 7 % GST.
Literature Cited 1
Lisk, D. J. Sci. Total Environ. 1988,74,39-66. Theis, T. L.; Gardner, K. H. Crit. Rev. Control 1990,21-42. Ujihara, A. M.; Gough, M. Managing Ash from Municipal Waste Incinerators; Center for Risk Management, Resources for the Future: Washington, DC, 1989. Roy, W. R.; Thiery, R. G.; Schuller, R. M.; Suloway, J. J. Environ. Geol. Notes (Ill.State Geol. Surv.) 1981,No. 96, 69. Cernuschi, S.;Giugliano, M.; de Paoli, I. Waste Manage. Res. 1990,8, 419-427. Belevi, H.; S t b p f l i , D. M.; Baccini, P. Waste Manage. Res. 1992,10,153-167. Wadge, A,; Hutton, M. Environ. Pollut. 1987, 48, 85-99. Northeim, C. M.; Gaskill, A.; Hasnen, G. A. Proceedings of the International Conference on Municipal Waste Combustion;Hollywood, FL, 1989;U.S. EPA and Environment Canada: 1989; pp 3B.39-52. Cundari, K. L.; Lauria, J. M. Waste Age 1986,November, 82-88. Francis, C. W.; White, G. H. J . Water Pollut. Control Fed. 1987,59,11, 979-986. Van der Sloot, H. A,; De Groot, G. J.; Wijkstra, J.; Leenders, P. Proceedings of the International Conference on Municipal Waste Combustion;Hollywood, FL, 1989;U.S. EPA and Environment Canada: 1989; pp 2B.1-13. Kullberg, S.; Fallman, A,-M. Proceedings of the Znternational Conferenceon Municipal Waste Combustion;Hollywood, FL, 1989;U.S. EPAand Environment Canada: 1989; pp 3B.21-37. Sawhney, B. L.; Frink, C. R. Water,Air, Soil Pollut. 1991, 57-58, 289-296. Austin, D. E.;Newland, L. W. Chemosphere 1985,14 (I), 41-51.
(15) Hwa,T. J. Environ. Monit. Assess. 1991,19(1-3), 157-164. (16) Dipietro, J. V. Master's Thesis, University of New Hampshire, 1989. (17) Gardner, K. H. Characteristicsof Leachates from Municipal Incinerator Ash Materials. Master's Thesis, Clarkson University, 1991. (18) Comans, R. N. J.; Van der Sloot, H. A.; Bonouvrie, P. A. Municipal Waste Combustion Conference;Air and Waste Management Association: Williamsburg, VA, 1993; pp 667-679. (19) Mullen, J. F. In Proceedings of the Second International Conference on Municipal Solid Waste Combustor Ash Utilization; Resource Recovery R e p o r t , Chesner Engineering: Washington, DC, 1989; pp 197-210. (20) Hjelmar, 0.Proceedings of the International Conference on Municipal Waste Combustion; Hollywood, FL, 1989; U S . EPA and Environment Canada: 1989; pp 3B.1-6. (21) Roffman, H. 2nd International Conference on Municipal Waste Combustion;Tampa, FL, 1991; Air & Waste Management Association: Pittsburgh, 1991; pp 775-794. (22) Kirby, C. S.;Rimstidt, J. D. Environ. Sci. Technol. 1993, 27, 652-660. (23) ASTM. Standard Practice for prediction of the long-term behavior of wastepackage materials including waste forms used in the geologic disposal of high-level nuclear waste; American Society for Testing and Materials: Philadelphia, 1991. (24) Lobo, V. M. Electrolyte Solutions: Literature Data On Thermodynamic and Transport Properties;University of Coimbra: Coimbra, Portugal, 1984; Vol. 1. (25) Rimstidt, J. D.;Newcomb, W. D. Geochim. Cosmochim. Acta 1993,57,1919-1934. (26) Ball, J. W.; Nordstorm, D. K. User'sManual for WATEQ4F, with Revised Thermodynamic Data Base and Test Cases for Calculating Speciation of Major, Trace, and Redox Elements in Natural Waters, U.S. Geological Survey: Reston, VA, 1991. (27) Kirby, C. S.;Rimstidt, J. D. Municipal Waste Combustion Conference; Air and Waste Management Association: Williamsburg, VA, 1993; pp 479-511. (28) Taylor, J. R. An Introduction to Error Analysis: The Study of Uncertainties in Physical Measurements; University Science Books: Mill Valley, CA, 1982; p 270. (29) Helsel, D.R. Environ. Sci. Technol. 1990,24 (12)) 17661774. (30) Burton, E. A.; Walter, L. M. Geology 1987, 15,111-114. (31) Plummer, L. N.; Busenberg, E. Geochim. Cosmochim.Acta 1982,46, 1011-1040. (32) Rimstidt, J. D.; Barnes, H. L. Geochim. Cosmochim.Acta 1980,44 (ll),1683-1699. (33) Henne, M. S. Masters Thesis, University of Nevada, 1982. (34) Crovisier, J. L.; Honnorez, J.; Eberhart, J. P. Geochim. Cosmochim. Acta 1987,51, 2977-2990. (35) Best, M. G. Igneous and Metamorphic Petrology; W. H. Freeman and Co.: New York, 1982; p 630. (36) Baes, C. F., Jr.; Mesmer, R. E. The Hydrolysis of Cations; Wiley-Interscience: New York, 1976; p 489. (37) Snoeyink, V. L.;Jenkins, D. Water Chemistry;John Wiley & Sons: New York, 1980, p 463. (38) Henry, W. M.; Barbour, R. L.; Jakobsen, R. J.; Schumacher, P. M. Inorganic compound identification of fly ash emissions from municipal incinerators. U.S. EPA Project Number 600/S3-83-095, 1983. (39) U.S. Code of Federal Regulations 141.81, July 1 ed., 1992, 704. (40) Eighmy, T. T.;Bobowski, S. F.; Ballestero, T. P.; Collins, M. R. In Proceedings of the Second International Conference on Municipal Solid Waste Combustor Ash Utilization; Resource Recovery Report, Chesner Engineering: Washington, DC, 1989; pp 275-314.
Received for review May 27, 1993. Revised manuscript received October 29, 1993.Accepted November 3, 1993.'
'Abstract published in Advance ACS Abstracts, December 15, 1993. Envlron. Sci. Technol., Vol. 28, No. 3, 1994
451
