Dissolution Kinetics of Synthetic Zeolite NaP1 and Its Implication to

(11) Cama, J.; Metz, V.; Ganor, J. The effect of pH and temperature on kaolinite dissolution rate under acidic conditions. Geochim. Cosmochim. Acta 20...
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Environ. Sci. Technol. 2005, 39, 4871-4877

Dissolution Kinetics of Synthetic Zeolite NaP1 and Its Implication to Zeolite Treatment of Contaminated Waters J O R D I C A M A , †,‡,§ C A R L E S A Y O R A , † XAVIER QUEROL,† AND J I W C H A R G A N O R * ,‡ Department of Environmental Geology, Institute of Earth Sciences “Jaume Almera”, CSIC, Lluı´s Sole´ i Sabarı´s s/n, Barcelona 08028, Catalonia, and Department of Geological and Environmental Sciences, Ben-Gurion University of the Negev, P.O. Box 653, Beer-Sheva 84105, Israel

The effect of pH on the dissolution kinetics of NaP1 zeolite, which was produced from the alkaline treatment of coal fly ash and may be used for decontamination of acid mine waters, is studied. The sample contains considerable amounts of accessory phases that partly dissolve during the experiment. Therefore, the dissolution rate was estimated during a stage in which the Al/Si ratio was equal to that of NaP1 (0.6). The release rate of these elements is controlled by the dissolution of the zeolite itself during this stage. The dissolution rate of NaP1 slows down with increasing pH in the acidic range, becomes constant at an intermediate pH, and increases with increasing pH in the basic range. The observed changes in rates were described using a rate law based on a surface speciation model. Using this rate law, we calculated the half-life of NaP1 to be about 2 years at near neutral pH and less than 10 days at pH below 3. For the utilization of NaP1 in the treatment of wastewaters or acid mine waters, these short half-lives bear two implications: (1) The treated waters must be kept at near neutral pH, and NaP1 should be added periodically to the treated waters in order to compensate for zeolite loss. (2) In water treatment applications that require a relatively short reaction time, the zeolite removed from the effluents should be kept dry in order to avoid its decomposition and the consequent release of the adsorbed metal to the environment.

Introduction The disposal of the large quantities of ash generated from the combustion of coal used in electric power plants throughout the world is a serious environmental problem. Much research was, therefore, devoted to the transformation of this so-called fly ash into useful materials in general, and into zeolite in particular (e.g., ref 1, and references therein). Because of their high cation exchange capacity, both natural * Corresponding author phone: ++972 8 6472651; fax: ++972 8 6472997; e-mail: [email protected]. † Institute of Earth Sciences. ‡ Ben-Gurion University of the Negev. § E-mail: [email protected]. 10.1021/es0500512 CCC: $30.25 Published on Web 05/26/2005

 2005 American Chemical Society

and synthetic zeolites are suitable materials for the treatment of different wastes, such as acid mine waters and NH4-rich waste on farmland (ref 2, and references therein). Although numerous patents and technical articles have proposed different methods of synthesizing different zeolites from fly ash (2), and many papers examined the ionic exchange properties of the synthesized zeolite, not much is known on the stability of zeolite. This absence of knowledge is critical because the durability of zeolite is a key parameter that must be taken into consideration while using it for treating contaminated waters. Considering the wide variability of zeolites, scarce data on their solubility and dissolution rates are available in the literature. The dissolution kinetics of heulandite (a natural zeolite) was studied by Ragnarsdottir (3), who proposed a surface speciation model explaining the pH-rate dependence of its dissolution. Ragnarsdottir et al. (4) examined the morphology and chemistry of the dissolved minerals’ surfaces and showed depletion in A1 in all of the samples, particularly those reacted at pH < 5. In addition, they found that Na and K were depleted below the depth of Al depletion, whereas Ca was enriched, indicating that the nonframe ions behave independently of the framework ions during dissolution. Although the removal of the framework ions was controlled by the dissolution reaction, the removal of the nonframe ions was controlled by ion exchange between Na/K and Ca. Using calorimetric data, Bowers et al. (5) and Bowers and Burns (6) constructed activity-activity diagrams defining the stability fields of clinoptilolite and heulandite with respect to other minerals at low temperature. The present study concentrates on the dissolution kinetics of NaP1 zeolite, which is produced from the alkaline treatment of fly ash and may be used for the decontamination of acid mine waters (e.g., ref 7). The cation exchange capacity of NaP1 was measured in previous studies (7, 8) at different pH conditions. Preliminary data on the dissolution kinetics of NaP1 at pH 3 (9) showed that the presence of amorphous phases, traces of NaOH, and proton consumption due to dissolution govern the solution pH and influence the release of Al and Si from the sample. The present study focuses on the effect of pH on the dissolution rate of NaP1.

Experimental Section The morphology of the zeolite samples was examined using a scanning electron microscope (SEM). Their mineralogical compositions were characterized by X-ray diffraction (XRD). The specific surface areas of the samples were measured by the Brunauer-Emmett-Teller (BET) method (10) using 5-point N2 adsorption isotherms, with a previous degassing of 24 h at approximately 140 °C. Three NaP1 samples were used in the present study: a commercial NaP1 sample supplied by Industrias Quimicas del Ebro (IQE) (henceforth, commercial NaP1); a raw NaP1 sample that was synthesized from fly ash (henceforth, raw NaP1); and a sub sample of raw NaP1 that was further pretreated in order to shorten the time required to achieve stoichiometric dissolution (henceforth, pretreated NaP1). Raw NaP1 was synthesized from fly ash by microwaveassisted hydrothermal alkaline activation (8). In addition to NaP1 (Na6(Si10Al6O32)‚2H2O), this sample contains a significant amount of amorphous phases as well as quartz (SiO2), mullite (Al6Si2O13), NaOH (which was used as a reagent in the production of the zeolite from the fly ash), and traces of calcite (CaCO3), tobermorite (Ca5Si6O16(OH)2) and magnetite (Fe3O4). The grain size distribution of raw NaP1 was VOL. 39, NO. 13, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 1: Experimental Conditions and Steady-State Average Concentration of the Experimentsa

experiment

flow rate

input pH

pH spc

(mL min-1) pret-NaP1-25-4 raw-NaP1-25-2 pret-NaP1-25-3 pret-NaP1-25-5-B pret-NaP1-25-2 raw-NaP1-25-1 pret-NaP1-25-6C pret-NaP1-25-5 pret-NaP1-25-6B pret-NaP1-25-6 raw-NaP1-25-5 raw-NaP1-25-3 raw-NaP1-25-4 raw-NaP1-25-6b comm-NaP1-25-3 comm-NaP1-25-1 comm-NaP1-25-2

0.12 0.12 0.11 0.10 0.07 0.06 0.13 0.10 0.13 0.13 0.05 0.05 0.05 0.03 0.05 0.11 0.05

Al spc

Si spc

Al/Si

initial sample mass

(µM) 2.04 3.04 3.01 3.03 3.11 3.01 3.56 4.05 4.59 5.04 7.53 9.10 11.82 2.53 2.29 3.00 3.00

2.07 3.02 3.08 3.16 3.21 3.47 3.70 4.24 4.80 5.61 7.49 9.03 11.74 3.13 2.96 3.19 3.47

55 27 29 85 94 205 31 7.7 0.9 0.2 1.3 6.1 72 244 1310 90 243

initial mass NaP1

mass NaP1 spc

(mol g-1 s-1)

(g) 92 46 54 128 149 325 46 14 5.6 6.2 22 11 114 422 796 65 251

0.60 0.59 0.53 0.66 0.63 0.63 0.67 0.54 0.15 0.03 0.06 0.54 0.63 0.58 1.64 1.38 0.97

0.201 0.095 0.105 0.187 0.215 0.472 0.191 0.187 0.191 0.191 0.500 0.514 0.500 0.503 0.500 0.088 0.504

0.078 0.049 0.041 0.073 0.084 0.245 0.074 0.073 0.074 0.074 0.260 0.267 0.260 0.261 0.340 0.060 0.342

dissolution rate

0.045 0.011 0.017 0.017 0.040 0.087 0.005 0.068 0.071 0.072 0.210 0.216 0.129 0.031 0.248 0.022 0.223

-4.2E-10 -8.8E-10 -5.9E-10 -1.2E-09 -4.1E-10 -3.6E-10 -2.2E-10 -3.6E-11 -1.7E-11 -1.8E-11 -8.9E-12 -4.2E-12 -7.1E-11 -6.7E-10 -2.0E-10 -2.5E-10 -7.1E-11

a 52 wt % of NaP1 in sample raw-NaP1; BET surface area ) 18.1 m2g-1; 39 wt % of NaP1 in sample pret-NaP1; BET surface area ) 63.4 m2g-1; 69 wt % of NaP1 in sample comm-NaP1; BET surface area ) 10 m2g-1. b Cation exchange: input solution with [Na] ) 1000 µM. c sp ) stoichiometric period or steady state in nonstoichiometric experiments.

determined by sieving. Approximately 8, 20, and 32 wt % of the grains were coarser than 149, 53, and 25 µm, respectively. The remaining 40% were smaller than 25 µm. The BETdetermined initial specific surface area of raw NaP1 was 18 ( 2 m2 g-1. Pretreated NaP1 was produced by pretreating raw NaP1 in a batch reactor using 0.001 N HNO3 at 70 °C for three weeks, during which most of the remaining NaOH and ultra fine particles were dissolved. The solution in the reactor was replaced periodically with fresh solution. The pH was measured daily and decreased from 10 to 4 as the NaOH dissolved. The BET-determined initial surface area of pretreated NaP1 was 63 ( 6 m2 g-1 and that of commercial NaP1 was 10 ( 1 m2 g-1. Experiments were carried out using Lexan nonstirred flowthrough reactors (ca. 35 mL in volume) fully immersed in a thermostatic water-bath held at a constant temperature of 25 ( 0.1 °C. Details of the experimental procedures can be found in refs 11 and 12. The experimental conditions of each of the experiments are compiled in Table 1. Input solutions were prepared by mixing reagent grade 1M HNO3 and double deionized water. The input solution of one experiment (raw-NaP1-25-6) also contains 1000 µM Na. Total Al and Si were analyzed either colorimetrically with a UV-visible spectrophotometer, using the catechol violet method (13) and molybdate blue method (14), respectively, or using inductively coupled plasma atomic emission spectroscopy (ICP-AES). Na was analyzed using ICP-AES. The uncertainties in the measurements were (5%. The pH values of the input and output solutions were measured at 25 °C on an unstirred aliquot of solution using a Ross combination electrode with a reported accuracy of (0.02 pH units ((4.5% in H+ activities).

Results and Discussion The variations with time of output pH and Al and Si concentrations in a representative flow-through experiment with raw NaP1 are shown in Figure 1a. Although the input pH was 3 throughout the experiment, the initial output pH was 7.6 and it decreased with time toward the input value. This change in pH reflects mainly the dissolution of the dry NaOH and exchange between the channel ions and protons, and to a lesser extent proton consumption by dissolution reactions. The complex variations in Al and Si concentrations 4872

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FIGURE 1. Variation of pH and concentrations of Si and Al (a) and of Al/Si ratio and accumulated loss of mass (b) in a representative dissolution experiment with raw NaP1. The shaded area represents the stoichiometric period (see text). Note that the pH is initially high due to the presence of NaOH in the sample and it decreases with time toward the input value (pH 3) as the NaOH is consumed. with time reflect the dissolution of the zeolite itself and that of other phases in the sample (9), and may be described

using a simple mass balance approach by

dCj,out mzeoSzeo ) Rzeoνj,zeo + dt V

∑R

otherνj,other

motherSother V

q (C - Cj,inp) (1) V j,out where R is the dissolution rate (mole m-2 s-1), Cj,inp and Cj,out are concentrations of species j (Al or Si) in the input and the output solutions, respectively (mole L-1), νj is the stoichiometric coefficient of j in the dissolution reaction, t is time (sec), m is the total mass of the dissolving phase (g), S is the specific reactive surface area of the mineral (m2 g-1), V is the volume of the cell (L), q is the fluid volume flux through the system (L sec-1), and the subscripts zeo and other refer to the zeolite itself and to other dissolving phases, respectively. Figures 1a and b show that both the Al and Si concentrations approached a maximum at 1000 h when the Al/Si ratio approached 0.6, which is the stoichiometric ratio of NaP1 zeolite. Thereafter, both the Si and Al output concentrations decreased, but the Al/Si ratio remained stoichiometric for another 400 h (shaded area in Figure 1a and b). We suggest that the stoichiometric release of Si and Al during this period indicates that the dissolution of NaP1 is dominant during this stage and that the contribution of other phases to the release of Al and Si is small. This suggestion will be verified later on. If indeed the release of Al and Si is controlled by the dissolution of NaP1, then the second term in the right-hand side of eq 1 is negligible, and eq 1 becomes

dCj,out mzeoSzeo q ) Rzeoνj,zeo - (Cj,out - Cj,inp) dt V V

(2)

Therefore, the dissolution rate of the zeolite during the stoichiometric period may be calculated if its mass and its specific surface area are known. Using the Rietveld refinement of X-ray powder diffraction (15, 16), Garrido (17) quantified the phase content and found that raw NaP1 contains 52 ( 3 wt % zeolite. This value is supported by the following observations. Accumulated mass loss during the dissolution experiment with raw NaP1 was estimated based on the total amount of Al and Si that was released to solution assuming that only zeolite was dissolved (Figure 1b). Although the accumulated mass loss at the end of the experiment was only about 50% of the initial mass of the raw sample, SEM observations (Figure 2c) and XRD measurements of the powder that was retrieved from the dissolution experiment indicated that the concentration of NaP1 was below the detection level (ca. 5%). This last observation strongly supports the results of the Rietveld phase analysis that the zeolite concentration in raw NaP1 is close to 52 wt %. The amount of zeolite in pretreated NaP1 and commercial NaP1 was found to be 39 wt % and 69 wt %, respectively, using the Rietveld method. The mass balance calculation indicates that the amount of Al and Si that was released to solution during dissolution is similar to the amount of Al and Si that is expected to be released to solution due to the dissolution of the zeolite itself. As an example, we will examine experiment NaP1-25-1. The initial mass of raw NaP1 in this experiment was 0.4715 g. On the basis of the Rietveld phase analysis, the zeolite concentration in raw NaP1 is close to 52 wt %. The initial mass of zeolite in this experiment was at least 0.233 g, taking an error of 5% in the estimation of the percentage of the zeolite. The final mass of zeolite in this experiment was below the detection limit of both XRD and SEM, that is, significantly less than 0.012 g. Therefore, at least 0.221 g of zeolite was dissolved during this experiment, leading to the release of at least 1013 micromoles of Al and 1689 micromoles of Si. The integrated amounts of released Al and Si during the experiment were 1057 micromoles of Al

FIGURE 2. SEM images of raw NaP1 before and after a dissolution experiment. The spherical shape of the zeolite particles (A) is due to their formation on the surface of spherical fly ash particles, which emerge in places in which the zeolite is cracked (B). After dissolution (C), the zeolite disappeared and the sample is composed mainly of mullite fibers and fly ash. and 1819 micromoles of Si. Therefore, at least 93% of the Si and 96% of the Al are from the dissolution of the zeolite. The result of these calculations verifies the above suggestion that the contribution of other phases to the release of Al and Si is small. Because the reactive surface areas of minerals (eq 2) are generally unknown, the common practice in experimental kinetics is to normalize the dissolution rate to the surface area of the pure mineral, which is measured by the BET method. In the present study, the specific surface area of raw NaP1 increased from the initial value of 18 m2 g-1 to values between 120 and 170 m2 g-1 after the dissolution experiments. Because most or all of the the zeolite had been dissolved during the course of the dissolution experiments, the high specific surface area at the end of the experiments represents the surface area of the other phases in the sample. Because about 50 wt % of raw NaP1 is composed of these phases, one expects that the initial surface area of raw NaP1 would be at VOL. 39, NO. 13, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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least 60 m2 g-1. The smaller initial surface area may be explained by an increase in the surface area of the other phases during dissolution or by exposure of these phases due to the dissolution of the zeolite. The second explanation is supported by the SEM observation showing that the zeolite in raw NaP1 coats the fly ash (Figure 2). As can be seen in Figure 2b, the coating is not complete, and therefore the initial surface area is a result of contributions from zeolite and the other phases. As a result, the BET surface area of raw NaP1 cannot be used as a proxy for the reactive surface area of the zeolite itself. Although the total reactive surface area of the zeolite (mzeoSzeo in eq 2) is unknown, it is clear that it would decrease as the mass of the zeolite decreases. Therefore, we normalized the dissolution rates by the mass of the zeolite, instead of normalizing it by the product of the mass and the specific surface area as appears in eq 2. Normalizing the rate to the mass of the sample is always better than not normalizing the rate at all. Using this normalization, we can describe the change in Al and Si concentrations with time by

dCj,out mzeo q ) ratezeoνj,zeo - (Cj,out - Cj,inp) dt V V

(3)

where ratezeo is the mass-normalized zeolite dissolution rate (mole g-1 s-1). The input solutions in the experiments contained only pure dilute HNO3 or NaOH and did not contain Al or Si. Therefore, rearranging eq 3 it becomes

ratezeo )

dCj,out V q + C dt mzeoνj,zeo mzeoνj,zeo j,out

(4)

It is important to note that the Al and Si concentrations are not constant during the stoichiometric period (e.g., Figure 1), and therefore, in contrast to the common practice in interpreting flow-through experiments, one cannot presuppose that the first term in the right-hand side of eq 4 equals 0. The time derivative of the concentration (dC/dt) in eq 4 was approximated from the slope of the concentration versus time between the previous time (i - 1) and the following time (i + 1), that is dCj,out

∆Cj,out ) ≈ dt ∆t 3(Ci-1ti-1 + Ci ti + Ci+1ti+1) - (Ci-1 + Ci + Ci+1)(ti-1 + ti + ti+1) 3(ti-12 + ti2 + ti+12) - (ti-1 + ti + ti+1)2

(5) In all of the experiments with raw and pretreated NaP1, the change in concentration with time during the stoichiometric period was relatively slow, and therefore the contribution of the first term in the right-hand side of eq 4 was smaller than the uncertainty. In other words, during the stoichiometric period the system was in a quasi steady-state. The slow decrease in concentration during this period was the result of the decrease in the mass of zeolite (Figure 1b). In three of the experiments with raw and pretreated NaP1 that were conducted at pH range of 4.8-7.6 (raw-NaP125-5, pretreated-NaP1-25-6 and pretreated-NaP1-25-6B), the dissolution was nonstoichiometric throughout the experiments because of the precipitation of secondary Al phases. Because of the slow dissolution rate in these experiments, the loss in mass was slower and therefore the rate of dissolution was calculated at steady state based on the concentrations of Si. It is important to note that one cannot rule out the possibility that some of the steady-state release of Si in these nonstoichiometric experiments is due to dissolution of other phases. Therefore, the calculated dissolution rate at near neutral pH is an upper bound of the dissolution rate of the zeolite itself. 4874

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FIGURE 3. The effect of pH on the dissolution rate of NaP1. The solid line is the best-fit curve of the nonlinear regression of the proposed rate law (eq 11) to all of the measured rate data obtained using raw NaP1 and pretreated NaP1. The dotted lines are the predictions of the rates of the three reaction paths. Throughout all of the experiments with commercial NaP1, the dissolution was nonstoichiometric, and the amount of aluminum released was higher than expected. The dissolution rates at these experiments were calculated at steady-state only based on Si release, and as before they are upper bounds for the dissolution rates of the zeolite itself. The initial mass of the zeolite in each experiment was calculated from the product of the measured starting mass and the percentage of the zeolite in the sample, which was calculated using the Rietveld method. Following each stage (i.e., the time between the replacements of two sequential output bottles), the remaining mass of zeolite was updated based on the release of Al and Si, assuming stoichiometric dissolution. The average dissolution rates during the stoichiometric period (or steady-state) of all of the experiments were calculated using eq 4 and are compiled in Table 1. There is a reasonable agreement between the dissolution rates of NaP1 observed in experiments with raw NaP1 and those with commercial NaP1 (Figure 3). This agreement supports our arguments that the observed dissolution rate is controlled dominantly by the dissolution of the zeolite itself. Likewise, the rate in the experiment that was conducted with an input solution that contains 1000 µM Na was similar to the dissolution rates in the other experiments at the same pH. Figure 3 shows that the dissolution rates of NaP1 slow with increasing pH in the acidic range, become constant at an intermediate pH, and increase with increasing pH in the basic range. This pattern of change with pH in the dissolution rate of oxides and silicates was interpreted classically by three reaction mechanisms: proton-promoted, hydroxyl-promoted, and water-promoted, which dominate the reaction rate under acidic, basic, or near neutral conditions, respectively (18). In this paper, we examine the pH effect on the dissolution rate of NaP1 zeolite under low ionic strength conditions by performing the experiments using different concentrations of HNO3 solutions. The drawback of such an experimental setting is that modifying the acid concentration changes the H+ and the ionic strength. It may be possible to vary the pH and keep the ionic strength relatively constant by adding a high concentration of salt. However, this is not suitable for experiments under low ionic strength. We assumed that under low ionic strength conditions the acid effect is solely due to the proton effect. The agreement between the dissolution rate in the experiment, which was conducted with an input solution containing 1000 µM NaCl, and the dissolution rates in the other experiments at the same pH supports the above assumption. As was done in previous studies, we modeled the dependence of the dissolution rate on pH using a surface speciation model (e.g., refs 19-21). The model is based on the following assumptions: (1) the NaP1 dissolution rate is controlled by three independent parallel reaction paths that represent a

pH-independent water attack of the surface and a pHdependent breakdown of bonds near a protonated surface site (represented by >Sp-OH2+ in eq 6 below) and a deprotonated surface site (represented by >Sd-O-) in eq 7 below); (2) The rate of bond breaking at the protonated and the deprotonated sites is linearly proportional to the surface concentration of the respective site, whereas the rate of the water attack on the neutral sites is constant and therefore pH-independent; (3) The adsorption and desorption of the protons on each of the two surface sites may be described by a simple independent Langmuir adsorption isotherm (eqs 8 and 9); and (4) The total amount of reactive surface sites per mass unit of zeolite is approximately constant during the experiment, and it does not change as the zeolite is dissolved. The last approximation, which is required in order to express the rates in mass-normalized units, cannot be justified theoretically. However, as will be shown later, the reactivity of the zeolite remained constant within error throughout the experiment, indicating that the approximation is practically valid. Writing the protonation reaction as Kprot

>Sp-OH + H+ 798 > Sp-OH2+

(6)

and the deprotonation reaction as Kdeprot

>Sd-O- + H+ 798 > Sd-OH

(7)

The surface coverage of the two sites may be described by

X>Sp-OH2+ ) Fprot X>Sd-O- )

KprotaH+ 1 + KprotaH+

(8)

X>Sd-OH KdeprotaH+ 1 ) F ) KdeprotaH+ KdeprotaH+ deprot 1 + KdeprotaH+ 1 Fdeprot (9) 1 + KdeprotaH+

where Fprot and Fdeprot are the maximum surface coverage of the protonated and the deprotonated sites, respectively, Kprot and Kdeprot are the equilibrium constants of the respective reactions and aH+ is the activity of protons in solution. If steady-state of the reaction intermediates is maintained, the rate (ratei) of the reaction path (i) is proportional to the surface coverage of the respective site Xi (22)

ratei ) ki‚Xi nr

(10)

where ki (s-1) is the rate coefficient of this path and nr (mol g-1) is the total amount of reactive surface sites in 1 g of zeolite. The overall dissolution rate is the sum of the three reaction paths

KprotaH+ rate ) kprotFprot + nr 1 + KprotaH+ kdeprotFdeprot

1 + kwater (11) 1 + KdeprotaH+

Using all of the dissolution rate data of raw and pretreated NaP1, the coefficients k1 ) kprotFprotnr, k2 ) kdeprotFdeprotnr, Kprot, and Kdeprot were calculated from a nonlinear regression of eq 11 using least squares (Figure 3). The solid line in Figure 3 is the best-fit curve that is described by

1300aH+ + 1 + 1300aH+ 1 3.3 × 10-7 + 7 × 10-12 (12) 15 1 + 2.9 × 10 ‚aH+

ratezeo ) 1.03 × 10-10

The dotted lines are plots of the three terms in eq 12, that is, of the pH dependencies of the three reaction paths. The protonated reaction path dominates the overall rate below pH 5, whereas the deprotonated path dominates the overall rate above pH 11. Although the fitting of the experimental data to the protonated reaction path is constrained by a dense array of data points, the fitting to the deprotonated reaction path is constrained poorly. Therefore, different combinations of the two coefficients in the second term of eq 12 yield similar curves that describe the experimental data adequately. As we noted above, the dissolution rate data at near neutral pH represent upper bounds for the dissolution rate of the zeolite. Therefore, the value of kwater (7 × 10-12 mol g-1 s-1), the rate coefficient that controls the overall rate at near neutral pH is also an upper bound for the rate coefficient. It is important to note that the ability of the rate law to predict the dissolution rate below pH 5 and above pH 10 is not influenced by the uncertainty in the determination of kwater. Our proposed mechanism is similar to that proposed by Ragnarsdottir (3) for the dissolution of heulandite. In their mechanism, the two active sites that control the rate are positively charged aluminol (>Al-OH2+) and negatively charge silanol (>Si-O-). The results of the present study neither support nor contradict this suggestion. Following the study of Ragnarsdottir (3), we based our model on two independent surface sites. Alternatively, we could model the experimental data adequately using a single amphoteric site. Because the study of the surface complexation of NaP1 is beyond the scope of the present paper, and the major conclusions regarding the stability of NaP1 would be identical regardless of the exact surface speciation, this alternative model will not be discussed further. Its predictive power should be examined before using eq 12 to predict the stability of NaP1 during its possible usage for treating acid effluents. Equation 12 was based on the fitting of stoichiometric and quasi steady-state data, which were usually derived during a late stage of the dissolution experiments, after a significant portion of the zeolite in the sample was dissolved. It is very common that mineral dissolution rates decrease with time as a result of the extinction of highly reactive fine particles (e.g., ref 23) and of the decrease in the ratio of reactive to unreactive sites (e.g., ref 24). An increase in mineral dissolution rate with time is less common but yet was observed in coated minerals (e.g., ref 25, 26, 27). In addition, eq 12 was based on the assumption that the total amount of reactive surface sites per mass unit of zeolite is approximately constant. Therefore, one can argue that the rate law of eq 12 may underestimate or overestimate the dissolution rates of bulk NaP1. To demonstrate that the rate law of eq 12 can be used to predict the entire dissolution of NaP1, we substitute eq 12 for ratezeo in eq 3 and simulated the expected change in Al and Si concentrations in the experiments with raw NaP1. The input data for each simulation include the initial mass of raw NaP1 (Table 1), the percentage of zeolite in raw NaP1 (52 ( 3 wt %), the concentration of the input solution (0 M), the stoichiometric coefficient of the element (6 for Al and 10 for Si), and matrixes describing the changes with time of the pH and the flow rate. The calculations were conducted using time steps of 1080 s (0.3 h). After each time step, the remaining mass of zeolite was updated based on the calculated release of Al or Si. Figure 4 compares the simulated changes in Al and Si concentrations with time (lines) to the actual measurements VOL. 39, NO. 13, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 5. Effect of pH on the half-life of NaP1 (solid line) and heulandite (dots).

FIGURE 4. Comparison of predictions (solid curves) to observations (data points) of the variation of Si (a) and Al (b) concentrations in an experiment with raw NaP1. Shaded areas are the uncertainties of the simulation. (symbols). The shaded areas represent the uncertainties in the results of the simulation that were estimated using the the Gaussian error propagation method (28). The two major sources for uncertainties are the uncertainty in the rate law of eq 12, which is estimated to be (20%, and the uncertainty in the initial mass of the zeolite ((3 wt %). Initially, the relative uncertainties in the results of the simulation are relatively small (about 20%). Because of the uncertainty in the initial mass of the zeolite, the relative uncertainty in the mass of zeolite that is left in the system increases with time as most of the zeolite dissolved, and as a result the relative uncertainties in the results of the simulation increase to above 60%. Figure 4 shows that there is a general agreement between the predicted (line) and the observed changes in Si and Al concentrations with time. Notwithstanding this agreement, the simulation significantly underestimates the concentration of Si during the first 230 h of the experiment and the Al concentration between 500 and 1000 h. These underestimations are expected because the simulation predicts the change in Al and Si concentrations due to the dissolution of the zeolite itself, which constitutes only about 50 wt % of the sample. During the first 230 h of the experiment, Si concentrations were much higher than predicted by the simulation, probably due to the fast dissolution of the amorphous phase (e.g., fly ash material). During this period, the observed change in Al concentration was similar to the change that is predicted by the simulation of the zeolite dissolution. Because the pH during the first 230 h ranges between 7.6 and 4.3, it is suggested that the dissolution of the Al-rich phases was inhibited initially due to the low solubility of Al-phases at near neutral pH. During the next several hundred hours, Al concentrations were much higher than predicted by the simulation, apparently due to the fast dissolution of the Al-rich amorphous phase. This suggestion is supported by the high Al/Si ratio during this period (Figure 1b). It is important to note that the simulation does not contain any adjustable parameter, that is, the simulated changes in concentrations are not fitted to the observed ones. The parameters in the equation were derived using only the stoichiometric (or steady states) stages of the experiments. Therefore, the good general agreement between the simulation and all of the observations indicates that the changes in Si and Al concentrations with time are controlled mostly by the dissolution of zeolite, that the dependence of the zeolite dissolution rate on pH may be predicted using the rate law 4876

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of eq 12, and that the reactivity of the zeolite remained relatively constant throughout the experiment. Because we could not determine a reliable proxy for the reactive surface area of the zeolite, the rate law of eq 12 was determined from dissolution rates that were normalized by the mass of the zeolite, and therefore the rates that are obtained using eq 12 are first order with respect to the zeolite mass. This characteristic of the rates enables the direct calculation of the half-life of the zeolite from the equation

t1/2 )

ln(2) ratezeoM

(13)

where t1/2 is the half-life of the zeolite and M is its molecular weight (1308 g mol-1). Figure 5 shows the pH dependence of the half-life of NaP1, which was calculated by substituting the rate law of eq 12 into eq 13. At near neutrral pH (pH 6.5-9.5), where the rates are the slowest, the half-life of NaP1 is above 800 days (above 2 years). The half-life decreases sharply below pH 6 and it is less than 10 days below pH 3. These half-lives have several implications that one should take into account when planning to use NaP1 as an adsorbent for metal from acid waters. (1) The acidity should be kept at near neutral conditions; otherwise, the zeolite would be decomposed and the adsorbed metal would be released back into the environment. (2) Even if the treated waters are kept at near neutral pH, the half-life of the zeolite is short in terms of remediation systems, and therefore NaP1 should be replaced periodically in order to compensate for zeolite loss. (3) If the zeolite is removed after decontaminating the water, as was proposed by ref 7, the removed zeolite should be kept dry in order to avoid its decomposition and the consequent release of the adsorbed metal to the environment. It would be of interest to compare its stability to that of other zeolites, taking into account the fast decomposition rate of NaP1. Figure 5 compares the pH dependence of the half-life of NaP1 to that of heulandite, which we calculated using the experimental data of Ragnarsdottir (3). The heulandite seems to be more stable than NaP1. For example, at pH 2 the half-life of NaP1 is about one-third that of heulandite. The observed difference may be related to differences in the size of the particles that were used in the two studies. Although the size range of the heulandite was 75-125 µm (3), most of the raw NaP1 particles were smaller than 53 µm. Based on the SEM observations (Figure 2), it seems that the big particles are fly ash particles that are covered by the zeolite. Under near neutral conditions, the half-life of heulandite is 1 order of magnitude longer than that of NaP1. Although it is possible that NaP1 is less stable than heulandite to water attack, the difference in particle size can be the main factor. Note that the calculated dissolution rate at near neutral pH is an upper bound of the dissolution rate of the zeolite itself and therefore the

difference between the stability of NaP1 and heulandite under near neutral conditions may be less than apparent in Figure 5.

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Acknowledgments

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This research was supported by the Spanish Government Project 02-04055-C02. We thank Esther Shani, Esteban Sanz, and Javier Pe´rez for technical assistance and Rafa Bartrolı´, Merce` Cabanes, Josep Elvira, and Xavier Alcover for analytical support. Montse Marsal helped us to obtain the SEM images. We also thank three anonymous reviewers for their thorough review that improved the quality of the manuscript.

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Received for review January 10, 2005. Revised manuscript received April 9, 2005. Accepted April 15, 2005. ES0500512

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