Environ. Sci. Technol. 1904, 18, 739-742
Page, A. L.; Elseewi, A. A.; Straughan, I. R. Residue Rev. 1979, 71, 83-120. Townsend, W. N.; Hodgson, D. R. In "Ecology and Reclamation of Devasted Land";Hurtlick, R. J.; Davis, G., Eds.; Gordon: New York, 1973; Vol. 1, Paper 1-4. Ball, J. W.; Nordstrom, D. K.; Jene, E. A. Water--Resour. Invest. (U.S. Geol. Surv.) 1980, No. WRI 78-116.
(19) Larimore, R.W.;Tranquilli, J. A., Eds. Bull.-Ill. Nut. Hist. Surv. 1982, 32 (Article 4). Received for review December 13, 1982. Revised manuscript received June 30, 1983. Accepted April 2, 1984. This project was partially funded by the Illinois Department of Energy and Natural Resources (Project 90.025).
Illinois Basin Coal Fly Ashes. 2. Equilibria Relationships and Qualitative Modeling of Ash-Water Reactions William R. Roy' and Robert A. Griffln
Geochemlstry Section, Illlnols State Geological Survey, Champaign, Illlnols 61820 Alkaline and acidic Illinois Basin coal fly ash samples were each mixed with deionized water and equilibrated for about 140 days to simulate ash ponding environments. Common to both equilibrated solutions, anhydrite solubility dominated Ca2+activities, and AP+ activities were in equilibrium with both matrix mullite and insoluble aluminum hydroxide phases. Aqueous silica activities were controlled by both mullite and matrix silicates. The pH of the extract of the acidic fly ash was 4.1 after 24 h but increased to a pH value of 6.4 as the H2S04,assumed to be adsorbed to the particle surfaces, was exhausted by the dissolution of matrix iron oxides and aluminosilicates. The activities of aqueous A13+and iron, initially at high levels during the early stages of equilibration, decreased to below analytical detection limits as the result of the formation of insoluble Fe and A1 hydroxide phases. The pH of the extract of the alkaline fly ash remained above a pH value of 10 during the entire equilibration interval as a result of the hydrolysis of matrix oxides. As with the acidic system, A13+activities were controlled by amorphous aluminum hydroxide phases that began to form after about 7 days of equilibration. The proposed mechanisms and their interrelations are discussed in addition to the solubility diagrams used to deduce these relationships.
Introduction The application of chemical equilibria models can lead to useful insights into the chemistry of aqueous systems such as coal fly ash effluent. However, the resulb of such modeling must be interpreted cautiously. Discrepancies in reported values for solubility constants for some mineral phases resulting from a lack of uniform experimental procedures can make the determination of equilibrium controls difficult. Moreover, pure mineral compounds in distilled water are often used for determining thermodynamic parameters that might not be representiative of these compounds when they are associated with complex multicomponent material such as coal fly ash. However, solubility models do provide a first approximation for understanding these complex solutions. As a step toward understanding fly ash aqueous systems, the chemical analyses of laboratory extracts were treated by the thermodynamic model WATEQP (1-3). This computer program calculates the ionic strength of each solution from the input chemical data. The calculated ionic strength is used to determine single ion activity coefficie,its, via a Daviesextended Debye-Hiickel equation, which is used to convert concentration to activity. These calculated activities were plotted on solubility diagrams. These solubility diagrams were used as a basis for defining equilibria as functions of time and for proposing a number of chemical mecha0013-936X/84/0918-0739$01.50/0
nisms that take place when coal fly ash comes in contact with water.
Materials and Methods As discussed in previous work (4),the authors generated laboratory extracts using a long-term equilibration (LTE) procedure with five distinctly different types of coal fly ash. Detailed descriptions of these samples in terms of origin, mineralogy, particle size, and chemical composition are given elsewhere (4, 5). The procedure used to assess the solubility of the fly ash samples involved mixing 3400 g of fly ash with 17 L of deionized water in a 19-L reaction vessel made of Pyrex glass. These mixtures were stirred for 30 min, 3 times a week, in order to (as a first approximation) simulate ash ponding environmenb by a procedure designed by Griffin et al. (6). However, this extraction procedure was more specifically oriented toward generating a solution at chemical equilibrium with the solid waste. This procedure was done to produce a solution that may approximate the aqueous chemistry of pond effluent in settings where metastable chemicaI equilibrium conditions develop. Two of the LTE solutions discussed in Roy et al. ( 4 ) were chosen for presentation since it was believed that these two samples best represented the diversity of the extracts from Illinois Basin fly ashes. One sample (12) is an acidic fly ash, an acid C-Modic silt loam by the nomenclature of Roy and Griffin (7); the other sample (18) is an alkaline Modic silt. During the equilibration interval, a sample of each solution was collected and analyzed. Results from the chemical analysis of aqueous samples taken after about 1, 7, 37, 63, 106, and 140 days were submitted to the WATEQ2 program. After about 140 days, the experiments were terminated. Results and Discussion Solubility Relationships. After 24 h both solutions were supersaturated with respect to calcium sulfate, but both were in equilibrium with anhydrite within 7 days of extraction (Figure 1). Figure 2 illustrates the calcium carbonate equilibria for the same two fly ash extracts. Neither solution fell along the calcite boundary which underscores the first graph, which suggested that the calcium activities were controlled by the solubility of anhydrite. The alkaline extract (18)was supersaturated with respect to calcite during the entire extraction interval. The acidic extract (12) remained undersaturated with respect to calcite because of the absence of carbonate at the low pH. Calcium concentrations in highly alkaline solutions in contact with atmospheric C02 (such as 18) should be
0 1984 American Chemical Society
Envlron. Sci. Technol., Vol. 18, No. 10, 1984 739
2'01\
Superratursted
2,1i
Arnorphou%SrO, IpK = 3.071
\"'"'
Supersaturated
Quartz IpK = 4.081
'
'\3Bd
"1
Gibbrite pK = 3 2 I. A 7 7i I O H I , I
A24hr
\\
I
-
b -
16-
4, 16-
Gypsum CaSO, *
2Hz0
2.6
20-
12~1.3 A12H,SiO,
22-
2:6
210 24-
0 i8Al"
Flgure 1. Calclum sulfate equillbrla of two long-term fly ash extracts.
Undersaturated
AItlH,SiO,
28-
28-
4 1
30
Supersaturated
o
, i
i
~
i
i
b
i
b
B
"
10
"
\
11
' 12
\\\
,
13
14
PH
037 d
6F Calcite CaCO, (pK =
I
Flgure 3. Sllica and aluminum hydroxlde solubillty equilibria of two long-term fly ash extracts.
8.446)
7-i 108dO
gd 0 1 4 2 d
36dO
0 F l y ash I2
0 F l y ash I8
Undersaturated
e 7 d 0 2 4 hr
"
I
2.6
i
2.4
I 2.2
pCa2+
Flgure 2. Calcium carbonate equllibrla of two long-term fly ash extracts.
controlled by calcite solubility. Talbot et al. (8) found that aqueous calcium concentrations in a fly ash extract were controlled by carbonate and hydroxide solid phases. However, studies have demonstrated that calcite is more soluble when Mg2+(9, 10) and SO$-(11) are present in solution, as was the case in these extracts. High concentrations of SO$- cause gypsum to precipitate on the calcite surface, resulting in a masking effect and higher Ca2+ concentrations than occur when S042-is absent. These data suggested that the calcium activity in the extracts was controlled by anhydrite in the solid wastes. Similar results with an acidic high-iron fly ash were reported by Griffin et al. (6). In Figure 3, various aluminum hydroxide and silica equilibria are shown as a function of pH. The upper part of the diagram is the silica equilibria. The initially acidic solution (12) was near equilibrium with amorphous SiOz (glass) during the early stages of equilibration but came to equilibrium with cryptocrystalline SiOz after about 36 days. The alkaline solution equilibrated with quartz after initially being undersaturated with respect to all the silica species. Most fly ashes derived from eastern bituminous coals are composed largely of silica; therefore, it is probable 740
Envlron. Sci. Technol., Vol. 18, No. 10, 1984
that the solubility of the aqueous H4Si04was controlled by matrix silica species and not by the Pyrex glass reaction vessels used to equilibrate the solutions. The acidic fly ash extract (12) was also in equilibrium with either gibbsite Al(OH), or amorphous Al(OH),; the alkaline extract, 18, appeared to be in equilibrium with boehmite AlO(OH),a metastable phase of gibbsite (Figure 3) during the latter part of the extraction period. The results of solubility modeling could be interpreted as indicating that aluminum hydroxide species controlled by solubility of aqueous A13+. Although no aluminum hydroxide compounds were detected by XRD in the solid samples, such compounds could have been present in amorphous forms or could have formed as precipitates during the extraction interval. Boehmite has been found to form as an amorphous precipitate on silicate surfaces during the dissolution of feldspar (12). The source of the A13+that forms the aluminum hydroxide precipitates in the extracts could have been the aluminosilicates such as mullite detected in I8 but not in I2 by XRD. However, modeling also indicates that the alkaline extract (18) was in equilibrium with mullite throughout the equilibration interval (Figure 4) whereas the acidic sample, 12, was in equilibrium with mullite and then appeared to be going toward supersaturation after 24 h. This trend could not be described beyond 36 days since the concentration of A13+was below analytical detection limits during the latter part of the equilibration interval of (12) due to gradual increase in pH as noted by Roy et al. (4). These results could be interpreted to mean that both mullite and the aluminum hydroxide precipitates were controlling the solubility of A13+. Chemical Mechanisms. As previously discussed, the pH of the initially acidic samples became neutral after about 3 weeks of extraction. Changes in the pH of fly ash systems over time have been documented, but the chemical mechanisms responsible for these changes have not been well-defined. However, with the assistance of the equilibrium relationships predicted by WATEQZ and the chemical data included in Roy et al. (4), a number of reaction pathways have been developed that may explain the
n
Table 11. Proposed Chemical Reactions and Resulting Equilibria That Occurred during the Equilibration of the Alkaline Coal Fly Ash IS"
01 1
initial contact to 24 hours
2MO + H 2 0
M'+
-+
20H-
t
(M = Ca,Mg. . .)
(1)
3-
Supersaturated
JzO t H z O + 2J' + 2 0 H (J = Na,K.. ,)
4-
g r'
A16Siz0,, + 1 9 H 2 0 5-
6AI(OHj4- + 2H4Si040 t 6H'
+
6AI(OHl4$.
(mullits)
82
(2)
+
+
z
(3) 6A10(0H)J. t 6 H z 0 + 60H-
14)
(boehmite) 6-
(0 P
between 24 hours and 7 days
7-
Undersaturated (reactions 1 , 2 , 3 and 41
E-
CaS04
012
9-
-L
Ca*'
+
,902 -
(5)
(anhydrite)
0 1s 10-
after 7 to 37 days
,
I
I
I
5
4
6
I
E
7
I
9
1
0
1
I
1
1
1
2
1
t
PH
AI6Si20,,
Flgure 4. Mullite equilibria of two long-term fly ash extracts.
Table I. Proposed Chemical Reactions and Resulting Equilibria That Occurred during the Equilibration of the Acidic Coal Fly Ash 12" initial contact to 7 days
12H,S04
-
Fe,O, 12S0,'
ZFe3+ + 6 H 2 0 + ZFe(OH),l + 6H'
-
f + 6H+ ZFe3* + 3H,O ! '(3) - + 24"
Al,Si201, (muilite)
+
46A13* +
SiOl + H,O
18H10
-
+ 5H10
*
14)
6AI(OH),l (gibbrite)
H4SiOIo
t
18"
15)
16)
between 7 and 21 days 2Si01 + 4 H 2 0 Al6SizO1, + 13H,O Fe,O,
+ 3Ht0
161
2H,SiOIo
11
2H4SiO4' + GAIIOH),I
ZFe(OHI,J
CaS04 Ca'* (anhydrite)
t
SO?-
18)
(7)
I
6AIO(OH)J
+
t
19H20
6 H z 0 + 60H-
+
(5)
6AI(OHj4- + 2HaSiOao t 6H'
11
-L
-
6AlIOH);
13)
\\ (4)
2Si02 t 4H,O
(6)
(matrix oxides continue to dissolve via reactions 1, and 2, generating OH-)
"A single arrow (+) designates that the reaction is proceeding in the direction indicated while the double arrow (+) denotes that the reaction has reached equilibrium.
(1)
(2)
4 K 6A13+ + ZH4Si0a0
lEH*
Ca'+ t SO4'
CaSo4 3
(aluminum and iron concentrations below analytical detection limits)
(9)
"A single arrow (+) designates that the reaction is proceeding in the direction indicated while the double arrow (e) denotes that the reaction has reached equilibrium.
chemical behavior of the extracts. Table I presents the possible sequence of reactions that were operative in the initially acidic fly ash extract (12). Sulfuric acid has been observed as a condensate on particle surfaces of some fly ashes by Swaine (13)and was assumed to have been the source of acidity during the early stages of extraction in this study. The hydrogen ions released by the complete dissociation of the HzS04 (Table I, reaction 3) reacted with both the iron oxides (Fe203)and the mullite (A&SiZO13)in the solid phase. The iron oxides (only Fez03 is shown to simplify the illustration) were slightly soluble, releasing Fe3+which hydrolyzed, and eventually formed stable iron hydroxide precipitates (Table I, reaction 1). WATEQS predicted that the solution was in equilibrium
with Fe(OH), after about 7 days but not with the source iron oxides. Roy et al. (4) indicated that the Fe solution concentrations progressively decreased with time. The hydrolysis of the solution iron generated additional hydrogen ions, which may have established a cycle that constantly dissolved the iron oxides as long as there was a source of H+. As all the H+ via the H2S04was consumed, resulting in a nonacidic solution, the iron dissolution reactions terminated, resulting in solution iron concentrations below analytical detection limits. While the relatively insoluble iron oxides were dissolving, some of the H+ (via the HzS04)dissolved mullite, releasing A13+ and H4Si04,establishing equilibrium with the solid phase (Table I, reaction 4). Modeling predicted that the solutions equilibrated with both mullite and quartz after 7-21 days. However, the solution A13+ also hydrolyzed, forming aluminum hydroxide precipitates, as suggested by the computer modeling. The hydrolysis of the AP+ also resulted in hydrogen ions which, analogous to the iron cycle, could have augmented the supply of hydrogen ions from the H2S04. As the hydrogen ions from the H2S04 became exhausted, both cyclic reactions ceased, with the net result that little Fe3+or A13+persisted in solution as the equilibrium controls became established. The pH of the system may have continued to increase as a result of the same chemical reactions that produced the alkaline fly ash extract. A number of chemical reactions thay may have taken place in the alkaline fly ash extract I8 have also been developed (Table 11). The chemical mechanism responsible for the high pH of many fly ash extrads and leachates has been attributed to the hydrolysis of matrix oxides (8, 14-16) and was assumed to be the cause of the high pH of the LTE solution of 18. After 140 days, the liquid phase was undersaturated with respect to CaO, MgO, Na20, and Envlron. Sci. Technol., Vol. 18, No. 10, 1984
741
K20, indicating that the dissolution reactions had not gone to completion. In contrast to results of Talbot et al. (8) and Elseewi et al. (15),the LTE solution of I8 had not established Ca(OH), equilibrium when the experiments were terminated. The hydrolysis of mullite (Table 11) generated predominantly ANOH),, which eventually precipitated as AlO(OH), resulting in the gradual decrease in solution A1 concentrations. After about 7 days, the solution concentration of A1 was less than 1.0 mg/L and was in equilibrium with both mullite and boehmite, a metastable aluminum hydroxide phase (Table 11,reactions 3 and 4). As in the acidic I2 solution, solution H4Si04activities equilibrated with both the matrix silica and mullite after approximately 1 week of extraction. The reactions listed in Tables I and I1 do not account for all the changes in solution chemistry of the two fly ash extracts and suggest only a qualitative sense of the reaction kinetics involved. However, these models may provide a basis for anticipating the type and sequence of reactions that take place when coal fly ash comes in contact with water.
Summary and Conclusions As coal fly ash samples from Illinois Basin coals come in contact with distilled water, a number of chemical reactions may take place at different times, influencing the type and distribution of chemical constituents leached from the solid waste. In acidic fly ash extracts, H2S04, adsorbed on particle surfaces, dissolved mullite and iron oxides, releasing A1 and Fe ions that formed insoluble hydroxide precipitates as the H2S04was exhausted. In alkaline fly ash extracts, the pH of the liquid phase was dominated by the hydrolysis of oxide phases. Amorphous aluminum hydroxide phases also appeared to control solution AI concentrations. Aqueous Si concentrations in both acidic and alkaline extracts equilibrated with matrix Si02phases, and Ca2+concentrations were influenced by anhydrite solubility. In general, when laboratory extracts of Illinois Basin coal fly ashes reached steady-state conditions, they were dominated by Ca2+and S042- ions and contained intermediate levels of Mg, Na, K, and Si ions,
742
Envlron. Scl. Technol., Vol. 18, No. 10, 1984
and low concentrations of A1 and Fe ions as a consequence of the mechanisms proposed. Registry No. CaS04, 7778-18-9; CaC03, 471-34-1; SO4'-, 14808-79-8; SiOz, 7631-86-9; H2Si04, 10193-36-9; A1(OH)3, 21645-51-2;AlO(OH), 1318-23-6;Ca, 7440-70-2;Al,7429-90-5;Fe, 7439-89-6;Mg, 7439-95-4;Na, 7440-23-5;K, 7440-09-7;anhydrite, 14798-04-0; calcite, 13397-26-7; gypsum, 13397-24-5; gibbsite, 14762-49-3.
Literature Cited Truesdell, A. H.; Jones, B. F. J. Res. U.S. Geol. Surv. 1974, 2, 233-248. Plummer, N. L.; Jones, B. F.; Truesdell, A. H. WaterResour. Invest. (US.Geol. Surv.) 1976, No. 76-3. Ball, J. W.; Nordstrom, D. K.; Jene, E. A. Water-Resour. Invest. (US.Geol. Surv.) 1980, No. WRI 78-116. Roy, W . R.; Griffin, R. A.; Dickerson, D. R.; Schuller, R. M. Environ. Sci. Technol., preceding paper in this issue. Suloway, J. J.; Roy, W. R.; Skelly, T. M.; Dickerson, D. R.; Schuller, R. M.; Griffin, R. A. Environ. Geol. Notes (Ill. State Geol. Surv.) 1983, No. 105. Griffin, R. A,; Schuller, R. M.; Suloway, J. J.; Shimp, N. F.; Childers, W. F.; Shiley, R. H. Environ. Geol. Notes (Ill. State Geol. Surv.) 1980, No. 89. Roy, W . R.; Griffin, R. A. J. Environ. Qual. 1982, 11, 563-568. Talbot, R. W.; Anderson, M. A,; Adren, A. W. Environ. Sci. Technol. 1978,12,1056-1062. Hassett, J. J.; Jurinak, J. J. Soil Sci. SOC.Am. h o c . 1971, 35, 403-406. Berner, R. A. Geochim. Cosmochim. Acta 1975,39,489-504. Akin, G. W.; Lagerwerff, J. V. Geochim. Cosmochim.Acta 1965,39, 353-360. Tchoubar, C.; Oberlin, A. J. Microsc. (Paris) 1963, 2, 415-432. Swaine, D. J. Proc. Annu. Conf. Trace Subst. Environ. Health, 11th 1977, 107-116. Shannon, D. G.; Fine, L. 0. Environ. Sci. Technol. 1974, 8, 1026-1028. Elseewi, A. A.; Page, A. L.; Grimm, S. R. J.Environ. Qual. 1980,9,424-427. Theis, T . L.; Wirth, J. L. Environ. Sci. Technol. 1977,11, 1096-1 100.
Received for review December 13, 1982. Revised manuscript received June 30, 1983. Accepted April 2, 1984.
