Spontaneous Precipitation of Struvite from ... - ACS Publications

Spontaneous Precipitation of Struvite from Synthetic Wastewater Solutions A. N. Kofina and P. G. Koutsoukos* Department of Chemical Engineering, University of Patras and Institute of Chemical Engineering and High-Temperature Chemical Processes, P.O. Box 1414, Patras, GR-26500, Greece Received June 17, 2004;

CRYSTAL GROWTH & DESIGN 2005 VOL. 5, NO. 2 489-496

Revised Manuscript Received November 10, 2004

ABSTRACT: Phosphorus recovery from wastewater by precipitation in the form of crystalline struvite is an attractive option contributing toward sustainable development. Struvite or magnesium ammonium phosphate hexahydrate (MgNH4PO4‚6H2O) is a crystalline salt, which may be used either as a fertilizer or as raw material for the production of phosphorus. The supersaturation of wastewater with respect to the target mineral is the key parameter determining the extent and the rate of recovery of struvite. In the present work, we investigated the stability domain of struvite in synthetic wastewater (SWW) and report on the kinetics of precipitation of the respective salt forming spontaneously. The investigation was carried out in aqueous solutions with compositions typically encountered in municipal wastewaters. The supersaturation with respect to struvite was varied through the appropriate variation of the concentrations of the Mg2+, NH4+, and PO43- ions, for which the stoichiometric molar ratio of 1:1:1 was observed. All experiments were done at 25 °C and at a constant solution pH 8.50 in a stirred batch reactor closed to the atmosphere. The component salt concentrations were selected so that the respective supersaturated solutions were not stable. Precipitation of struvite from the supersaturated solutions prepared in SWW was initiated spontaneously past the lapse of well-defined induction times. The measurements of the induction times showed that the stability range of the supersaturated solutions in SWW was very narrow. The induction times, preceding the formation of struvite, were inversely proportional to the solution supersaturation and followed the dependence predicted by the classical nucleation theory. It was thus possible to calculate a surface energy of 15 mJ m-2 for the struvite nuclei forming. The rates of the struvite precipitating past the end of the induction period were measured at a constant driving force by the addition of stoichiometric titrant solutions throughout the precipitation process, using solution pH as a master variable monitored by a glass electrode sensor. The rates of precipitation measured from the titrants addition showed a parabolic dependence on the solution supersaturation. The high order (>1) of dependence of the rates on the solution supersaturation suggested a surface diffusion-controlled mechanism. Introduction According to the Urban Waste Water Treatment Directive,1 the EU legislation concerning the disposal of liquid wastewater in the natural environment has become stricter. Thus, not only dumping of untreated wastewaters and sewage sludge in the aquatic natural systems is forbidden, but also the maximum allowable limits of phosphorus and nitrogen have been further reduced, in an effort to minimize the risk of eutrophication of natural waters.2,3 Phosphorus is an important product of the chemical industry with a broad spectrum of applications in everyday life and widely encountered in consumer products and commodities. The removal of phosphorus from wastewater is therefore imperative, and important technological developments have been done in this direction. As consumption increases however, recovery of phosphorus is an important option, which is expected to contribute to sustainable development through saving essential raw materials.2,4 Recovery of phosphorus from wastewater in the form of crystalline struvite (MgNH4PO4‚6H2O) is a promising option, as it may be used as a fertilizer. This aspect has attracted research and technological interest over the recent years because of the relatively low solubility of struvite and the low content of heavy metals of this salt. The release of both ammonia and phosphorus which * To whom correspondence should be addressed. E-mail: pgk@ iceht.forth.gr.

may become available to plants5 is a source of important nutrients. Struvite precipitation is possible in cases in which the activities of the Mg2+, NH4+, and PO43- ions exceed the respective solubility product. The extent of struvite precipitation and the characteristics of the precipitating solid depend on the solution pH, supersaturation, temperature, and presence of foreign ions. A number of studies have been conducted concerning the precipitation of struvite in aqueous media.6-9 For the most part, the majority of the published works related to struvite precipitation were conducted either in simple electrolyte solutions or the researchers did not provide adequate information to correlate the thermodynamics of the solutions with kinetics results. In the present work, we have investigated the process of spontaneous precipitation of struvite in synthetic wastewater (SWW) solution at 25 °C and pH 8.50 at constant pH and solution supersaturation. The induction times preceding the onset of struvite formation in SWW were measured as a function of the solution supersaturation with respect to struvite, and the stability diagrams for the formation of struvite at these conditions were constructed. Moreover, the rates of the precipitation were measured as a function of the solution supersaturation, thus obtaining for the first time mechanistic information for the precipitation of struvite in SWW.

10.1021/cg049803e CCC: $30.25 © 2005 American Chemical Society Published on Web 01/12/2005

490 Crystal Growth & Design, Vol. 5, No. 2, 2005

Kofina and Koutsoukos

Experimental Procedures All experiments were carried out at 25 °C in a 250 mL double-walled Pyrex vessel thermostated by water circulation through a constant-temperature bath. Stock solutions of magnesium sulfate heptahydrate and dihydrogen ammonium phosphate were prepared from corresponding crystalline solids MgSO4‚7H2O and NH4H2PO4 (Merck, reagent-grade) using triply distilled water. Reagent grade chemicals and triply distilled water were used to prepare the synthetic wastewater solution. The chemical composition of SWW2-3,10-12 is shown in Tables 1 and 2, respectively. It should be noted that in this model solution, the glucose added corresponds to a COD value of 100 ppm. This medium was used in the present work, for the investigation of struvite precipitation. All stock solutions of the salts used and the SWW solution were filtered through membrane filters (0.22 µm Millipore) before the experiment. The magnesium stock solution was standardized with EDTA titrations. The sodium hydroxide solutions (Merck, Titrisol) were standardized against potassium hydrogen phthalate solutions prepared fresh from the respective crystalline solid, dried overnight at 105 °C. The ammonium phosphate stock solution was standardized with potentiometric titration with standard sodium hydroxide solution. The SWW was used as a solvent for the preparation of the supersaturated solutions. In all solutions, the stoichiometry Mg/NH4/PO4 ) 1:1:1 was observed. The supersaturated solutions were prepared in the reactor, by rapidly mixing SWW and NH4H2PO4. Next, the solution pH was adjusted to 8.50, by the addition of the appropriate amount of a standard solution of sodium hydroxide, followed by the addition of the appropriate volume of stock magnesium sulfate solution. Finally, the solution pH was readjusted to 8.50 as needed. The homogeneity of the solution was ensured by stirring with a Teflon-coated stirring bar and a magnetic stirrer.8 A constant flow of water vapor saturated nitrogen was maintained over the solution throughout the precipitation process to avoid atmospheric carbon dioxide intrusion. A combination glass/ saturated calomel electrode standardized before and after each experiment with standard buffer13 solutions was used for the solution pH measurements. The experiments were done at least in triplicate for the assessment of repeatability. The precipitation reaction may be described by eq 1:

Mg2+ + NH4+ + H2PO4- S MgNH4PO4‚6H2O + 2H+ (1) According to eq 1, during the precipitation of struvite, protons are released into the solution. Solution pH drop exceeding 0.005 pH units triggered the addition of titrant solutions from the computer-driven motorized, electrical burets. Standard sodium hydroxide solution was added in the experiments in which the solution pH was maintained (pH-stat experiments). For the experiments in which constant supersaturation was maintained, the titrant solutions consisted of the stock solutions of magnesium and phosphate, the SWW solution as solvent, and NaOH solution as needed to adjust the working solution pH. The concentrations were calculated from the mass balance equations for all ions as follows. Assuming that the aqueous solution contains solutions of concentrations: x1 M MgSO4‚7H2O, x2 M NH4H2PO4, x3 M NaOH, and x4 M SWW, the titrant solution should contain y1 M MgSO4‚7H2O, y2 M NH4H2PO4, y3 M NaOH, and y4 M SWW to maintain the activities of all ions constant. Assuming dm moles of MgNH4PO4 precipitated, in the total volume of the supersaturated solution, V, and dV the respective titrant solution volume added by each buret, the requirement for constant Mg activity imposes the mass balance eq 2, from which the concentration of magnesium in the titrant, y1, is obtained:

x1V + y1dV - dm dm x1 ) w 2x1 ) y1 w y1 ) 2x1 + c V + 2dV dV (2)

Figure 1. Experimental set up for the investigation of the kinetics of struvite formation in SWW at constant solution pH and/or supersaturation. where dm/dV ) c is a measure of the amount of solid precipitated per unit volume. This parameter was determined empirically, by preliminary experiments. Similarly, the requirement of titrant solution addition for constant phosphate activity is

y2 ) 2x2 + c

(3)

Because of the fact that two protons are released to the solution during the precipitation of each mole of solid, sodium hydroxide titrant is needed with a concentration y3:

y3 ) 2x3 + 2c

(4)

Similarly, the requirement for the SWW is

y4 ) 2x4 - 2c

(5)

The titrant solutions were therefore prepared as follows:

titrant 1: (2x1 + c)M MgSO4‚7H2O + (2x4 - 2c)M SWW titrant 2: (2x2 + c)M NH4H2PO4 + (2x3 + 2c)M NaOH + (2x4 - 2c)M SWW In our experiments, the value of c ) 20 was obtained from preliminary experiments in which it was found adequate to keep the solution supersaturation constant. The time lapsed between the preparation of the supersaturated solutions (including pH adjustment) and the first addition of the titrant solution was taken as the induction time, τ. The experimental setup used for the kinetics measurements is schematically shown in Figure 1. Both before the initiation of the salt precipitation (induction time) and past the start of the salt precipitation, an adequate number of samples were withdrawn and filtered through 0.22 µm membrane filters. The filtrates were analyzed for magnesium by atomic absorption spectrometry (Perkin-Elmer AAnalyst 300) and for phosphate spectrophotometrically (Hitachi U-2001UV/Vis) by the vanadomolybdate complex formation method.14 At the end of each experiment, the solids were collected on the 0.22 µm filter, dried at room temperature, and characterized by powder X-ray diffraction (Philips 1830/ 40) and scanning electron microscopy (SEM, JEOL JSM 5200 and LEO VP-35 FEM). Thermogravimetric analysis of the precipitated solids was carried out (TGA Q50, Thermal Analysis) under nitrogen and air gas flow. Finally, specific surface area measurements were made by N2 adsorption (BET method, Micromeritics, Gemini III 2375).

Results and Discussion One of the main tasks of the present work was the determination of the formation curve of the struvite

Struvite Precipitation from Synthetic Wastewater

Crystal Growth & Design, Vol. 5, No. 2, 2005 491

Table 1. Composition of Synthetic Wastewater, for the Experiments in which the Spontaneous Precipitation of Struvite Was Investigated at Constant Solution pH component

concentration (M)

glucose NaHCO3 NaCl NaNO3

5.167 × 10-4 17.86 × 10-3 10 × 10-3 0.5882 × 10-3

system in a complex medium such as the SWW, i.e., the experimental determination of the domain of stability of the solutions supersaturated with respect to struvite. The driving force for the formation of struvite in aqueous supersaturated solutions is the difference between the chemical potentials, ∆µ, of the salt in the supersaturated solution µs from the corresponding value at equilibrium, µ∞:

∆µ ) µ∞ - µs ) [µo∞ + kT ln(aMg2+‚aNH4+‚aPO43-)1/3 ∞ ] [µos + kT ln(aMg2+‚aNH4+‚aPO43-)1/3 s ] (6) Assuming that the chemical potentials of the standard states in the supersaturated solution and at equilibrium are equal, the difference in chemical potentials is

(aMg2+‚aNH4+‚aPO43-)1/3 s kT ∆µ ) kT ln )ln Ω (7) 1/3 3 (aMg2+‚aNH4+‚aPO43-)∞ where k is the Boltzmann’s constant and T is the absolute temperature. The logarithmic term is the supersaturation ratio Ω given by

Ω)

aMg2+‚aNH4+‚aPO43K0s

(8)

where K0s is the thermodynamic solubility product of struvite. Ω is a measure of the deviation of the system from equilibrium and a measure of the driving force for the precipitation. For Ω ) 1, the solution is saturated (equilibrium), for Ω > 1 the solution is supersaturated and precipitation may occur, while for Ω < 1 the solution is undersaturated and dissolution may take place. The relative supersaturation σ was defined as

σ ) Ω1/3 - 1

(9)

The activities of the ionic species in solution and the supersaturation ratio Ω were calculated by the MINEQL+ chemical equilibrium modeling software15 taking into account all chemical equilibria together with mass balance and electroneutrality conditions. The equilibrium and the stability constants used for the calculations are summarized in Table 3. The calculations were done by successive approximations for the ionic strength, I, while activity coefficients were calculated from the extended form of the Debye-Hu¨ckel equation proposed by Davies.16,17 The results of this program indicate what thermodynamically can happen. The value of log K0s ) 13.26 for the logarithm for the thermodynamic solubility product of struvite was taken from literature values corrected for ionic strength.

Table 2. Composition of Synthetic Wastewater, for the Experiments in which the Spontaneous Precipitation of Struvite Was Investigated at Constant Solution Supersaturation component

concentration (M)

glucose NaHCO3 NaCl NaNO3 Na2SO4

5.167 × 10-4 17.86 × 10-3 10 × 10-3 0.5882 × 10-3 1.200 × 10-2

Table 3. Equilibria Involved in the Computation of the Solution Speciation and the Respective Thermodynamic Constants at 25 °C equilibrium

log K

ref

H+ + PO43- S HPO42H+ + HPO42- S H2PO4H+ + H2PO4- S H3PO4 Na+ + PO43- S NaPO42Na+ + HPO42- S NaHPO4Na+ + H2PO4- S NaH2PO4 Na+ + NaPO42- S Na2PO4H+ + Na2PO4- S Na2HPO4 Mg2+ + PO43- S MgPO4Mg2+ + HPO42- S MgHPO4Mg2+ + H2PO4- S MgH2PO4 Na+ + NO3- S NaNO3 H+ + CO32- S HCO3H+ + HCO3- S H2CO3 H2O + CO2 S H2CO3 Na+ + CO32- S NaCO3Na+ + HCO3- S NaHCO3 Mg2+ + CO32- S MgCO3 Mg2+ + HCO3- S MgHCO3+ NH3 + H+ S NH4+ NH3 + Mg2+ S MgNH3+2 2NH3 + Mg2+S Mg(NH3)2+2 3NH3 + Mg2+ S Mg(NH3)3+2 H+ + SO42- S HSO4Na+ + SO42- S NaSO4NH4+ + SO42- S NH4SO4Mg2+ + SO42- S MgSO4 H+ + OH- S H2O Na+ + OH- S NaOH Mg2+ + OH- S MgOH+ Na+ + Cl- S NaCl Mg2+ + Cl- S MgCl+ Mg2+ + PO4-3 + NH4 + 6‚H2O S MgNH4PO4‚6H2O

12.375 7.198 2.148 1.43 1.07 0.3 1.16 10.73 4.8 2.80 0.45 -0.55 10.329 6.352 -1.466 1.27 -0.25 2.92 1.01 9.244 0.24 0.2 -0.3 (I ) 2) 1.99 0.73 1.03 2.26 13.997 0.1 2.6 -0.5 0.6 13.26

18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 18 19

Typical time profiles for the solution de-supersaturation monitored through the reduction of the reacting magnesium and phosphate components are shown in Figure 2. From these profiles for each experiment, it was possible to calculate the phosphorus recovery achieved in comparison to the initially present phosphorus and in relation to the equilibrium concentration (saturation) with respect to struvite. At constant solution supersaturation, the activities of all ions in solution were kept constant. Typical profiles of the reacting components as a function of time in this case are shown in Figure 3: The experimental conditions and the kinetic results obtained from the experiments at constant pH and constant supersaturation are summarized in Tables 4 and 5, respectively. The induction time preceding the onset of crystallization was found to be inversely proportional to the solution supersaturation as can be seen in Figure 4, in which the driving force for precipitation as a function of the induction time is shown.

492

Crystal Growth & Design, Vol. 5, No. 2, 2005


Figure 2. Total magnesium (a) and total phosphate (b) change as a function of time for the spontaneous precipitation of struvite in SWW at constant pH; pH 8.50, 25 °C.

Figure 3. Total magnesium (a) and total phosphate (b) change as a function of time for the spontaneous precipitation of struvite in SWW at constant pH; pH 8.50, 25 °C. Table 4. Experimental Conditions: Supersaturation, Relative Supersaturation, Induction Time, τ, and Subsequent Growth Rates, R, for the Spontaneous Precipitation of Struvite in SWW at Constant pH 8.5 and 25 °C CMg ) CNH4 ) CPO4 (× 10-3 M)

Ω

σ

τ (s)

R (× 10-6 mol min-1)

3.00 2.90 2.80 2.75 2.70 2.60 2.50 2.40 2.30 2.20

4.2953 3.8636 3.6141 3.5975 3.3342 3.2359 2.7352 2.5409 2.1183 2.0797

0.6255 0.5691 0.5346 0.5322 0.4939 0.4791 0.3985 0.3645 0.2843 0.2764

252 348 543 630 708 1050 1254 1332 1482 2466

4.5 3.8 3.4 2.1 1.9 1.6 1.2 0.9 0.7 0.5

B (log Ω)2

(10)

where Asp and B constants according to the classical nucleation theory are

B)

βυm2γs3 (2.303kT)3

CMg) CNH4 ) CPO4 (× 10-3 M)

Ω

σ

τ (s)

R (× 10-6 mol min-1)

4.20 4.10 4.00 3.90 3.80 3.60 3.50 3.40 3.30

6.2520 6.0390 5.6880 5.0460 4.9320 4.4157 3.9530 3.7325 3.6475

0.8422 0.8210 0.7851 0.7152 0.7022 0.6460 0.5812 0.5512 0.5393

235 500 750 1015 1150 1750 1840 2100 4150

30.3 20.7 11.8 5.8 5.6 4.2 2.7 2.3 1.7

a

The dependence of the measured induction time on the solution supersaturation in a logarithmic form is given by20

log τ ) log Asp +

Table 5. Experimental Conditions, Supersaturation, Relative Supersaturation, Induction Time, and Subsequent Growth Rates for the Spontaneous Precipitation of Struvite at Constant Supersaturationa

(11)

Aqueous medium: SWW pH 8.5, 25 °C.

In eq 11, β is a shape factor (for cubes is 32), υm is the molecular volume of struvite () molar volume/(Avogadro’s number × density × number of ions in formula unit) ) 7.99 × 10-23 cm3), and γs is the surface energy of the solid that is forming. As may be seen, the plot according to eq 10 consists of two linear parts.20 The change of the slope of the straight lines is considered to correspond to the transition from a homogeneous to heterogeneous nucleation process. As may be seen from Figure 5 the threshold between homogeneous and heterogeneous precipitation processes is at Ω ≈ 3.18 for constant pH experiments and at Ω ≈ 4.45 for experiments at constant supersaturation. From



Figure 4. Stability diagram for struvite in SWW. pH 8.50, 25 °C; (9) constant supersaturation; (b) constant solution pH.

Figure 5. Plot of the logarithm of induction time preceding the spontaneous precipitation of struvite in SWW as a function of (log Ω)-2. pH 8.50, 25 °C; (b) constant supersaturation; ([) constant solution pH.

the experimental results (Tables 4 and 5) and the slope of the steeper part of the above lines (Figure 5), a value of surface energy equal to 15 mJ m-2 was calculated for the struvite. This difference may be ascribed to the effect of the SO42- present in excess in the case of the solutions used in the constant supersaturation method. In this case, the presence of the additional SO42- ions in the supersaturated solutions originate from the Na2SO4 background electrolyte added to maintain the ionic strength of the solution during the addition of the titrant solutions. As may be seen from the kinetics measurements summarized in Table 5, the initial rates depend strongly on the solution supersaturation. The dependence of the rate of spontaneous precipitation of struvite, Rp, on the solution supersaturation may be expressed by powerlaw equations such as

Rp ) kpσn

(12)

In eq 12 kp is the apparent rate constant, σ is the relative supersaturation, and n is the apparent order of the precipitation. It should be noted that the rates of precipitation are initial rates and for each supersaturation they were calculated from the slopes of the time

Figure 6. Plot of the rates of struvite precipitation as a function of the solution supersaturation, pH 8.5, 25 °C; (b) constant solution pH; (9) constant solution supersaturation.

Figure 7. Powder X-ray diffractιpatterns from (a) struvite precipitated spontaneously from SWW at constant pH/constant supersaturation; (b) reference pattern file no. 15-762 for synthetic struvite (reference JCPDS).

profiles of the titrant solutions addition. Plots of the rates of precipitation of struvite as a function of the relative supersaturation showed a high order dependence (>1) over the range of supersaturations investigated, as may be seen in Figure 6. This dependence is indicative of a surface-controlled mechanism.21,22 It should be noted that the precipitation process is a complex process involving primary and secondary nucleation and agglomeration. Secondary nucleation however is not very likely in the system investigated as there is evidence that measurements of the number of particles formed following the initial stages of precipitation remained constant. These measurements were done in preliminary experiments using laser light scattering. Moreover, from the fact that the rates measured at constant supersaturation were found proportional to the total surface area precipitated it may also be concluded that agglomeration does not play a critical role in the overall precipitation rates. In all cases, the mineral phase precipitated was struvite23 as may be seen from the powder X-ray diffraction spectrum shown in Figure 7. Thermogravimetric analysis of the precipitated solids, done under nitrogen and air gas flow, did not show any significant differences. The weight loss as function of temperature is shown in Figure 8.

494


Figure 8. Thermogravimetric analysis curves of struvite precipitated spontaneously from SWW at constant supersaturation: heating rate 5 °C/min, nitrogen atmosphere.

As may be seen, the TGA curve shows a first weight loss of 46.39% between 80 and 120 °C, which corresponds the release of the six crystalline water molecules of MgNH4PO4‚6H2O (calculated 44.03%). A second weight loss of 6.7% (or 53.09% total), taking place at temperatures between 120 and 400 °C, which corresponds the release of the ammonia (calculated 7.33 or 51.36% total loss) was observed. The morphology of the precipitated crystals is shown in the scanning electron micrographs in Figure 9. As may be seen, the precipitated struvite crystals are prismatic crystals elongated along the c axis. The EDX


spectrum showed that the only elements present in the solid are magnesium and phosphorus. In Figure 10 the morphology of struvite crystals obtained by spontaneous precipitation in SWW at constant supersaturation in different stages is shown. During the first minutes past the initiation of the precipitation a relatively small number of crystals was formed of 10 µm size (Figure 10a). In the next 25 min the number of crystals increased and so did the respective size, which reached 30 µm (Figure 10b). In the next stages and up to 90 min, the crystals formed became longer than 30 µm (Figure 10c,d) and increased in thickness. The BET surface area of the precipitated struvite was found equal to 150 m2/g, regardless of the method of precipitation (constant pH/constant supersaturation). The high value of the SSA measured was not surprising according to the observed morphology (see for example porous structure details in Figure 9). The experiments at constant pH-variable supersaturation have shown that when this parameter is controlled a phosphorus recovery corresponding to 60% of the initially present phosphorus is feasible and 75% with respect to the solubility of struvite. At conditions of constant supersaturation, it is possible to precipitate continuously struvite from a solution simulating SWW. Conclusions The kinetics of the spontaneous precipitation of struvite was investigated in aqueous solutions simulating, from the point of view of chemical compositions, SWW. The conditions of the present investigation were

Figure 9. (a-c) Morphology of struvite crystals precipitated in SWW at constant pH/supersaturation; pH 8.50, 25 °C; (d) EDX microanalysis corresponding to crystals shown in panels a and b.



Figure 10. Struvite crystals formed by spontaneous precipitation in SWW at conditions of constant solution supersaturation; pH 8.50, 25 °C; (a) 25 min, (b) 45 min, (c) 60 min, (d) 90 min.

limited to equimolar concentrations of Mg2+, NH4+, and PO43- ions in the supersaturated solutions. At 25 °C and for pH 8.50 white, orthorhombic crystalline struvite was precipitated spontaneously in all experiments. The stability diagram for the struvite system was measured both from experiments in which spontaneous precipitation was done at constant pH and at sustained solution supersaturation. The measured induction times as a function of the solution supersaturation were found to reduce rapidly, suggesting narrow limits of stability of the respective supersaturated solutions. Differences found in the stability determined between the solutions investigated at constant pH and at constant supersaturation were ascribed to the presence of the sulfate ions in solutions, which have a retarding effect in the initiation of the precipitation process. This is reflected in the different surface energies calculated according to the classical nucleation theory. From the dependence of induction time on the supersaturation of the solution, it was found that the threshold supersaturation for the homogeneous/heterogeneous nucleation was Ω ) 3.18 for the pH-stat experiments and Ω ) 4.45 for the constant supersaturation experiments. The rates of precipitation measured, however, increased with supersaturation and showed a parabolic dependence on the solution relative supersaturation, suggesting a surface diffusion mechanism. The kinetics measurements done both at constant pH and at constant supersaturation did not show any differentiation with respect to the dependence of the measured rates as a function of the solution supersaturation.

Acknowledgment. The authors acknowledge support from the General Secretariat for research and Technology through the PENED 2001 programme (Project No. 01E∆413). References (1) UWWTD, Council of the European Communities, Directive Concerning the Collection, Treatment and Discharge of Urban Wastewater from Certain Industrial Sectors (91/271/ EEC) Official J L, 135/40, 1991. (2) Doyle, J. D.; Parsons, S. A. Water Res. 2002, 36, 3925-3940. (3) Jaffer, Y.; Clark, T. A.; Pearce, P.; Parsons, S. A. Water Res. 2000, 36, 1834-1842. (4) Munch, E. V.; Barr, K. Water Res. 2000, 35, 151-159. (5) Li, X. Z.; Zhao, Q. L. Ecol. Eng. 2003, 20, 171-181. (6) Bouropoulos, N. Ch.; Koutsoukos, P. G. J. Cryst. Growth 2000, 213, 381-388. (7) Babic-Ivancic, V.; Kontrec, J.; Kralj, D.; Brecevic, L. Croat. Chem. Acta 2002, 75, 86-106. (8) Golubev, S. V.; Pokrovsky, O. S.; Savenco, V. S. J. Cryst. Growth 2001, 223, 550-556. (9) Stratful, I.; Scrimshaw, M. D.; Lester, J. N. Water Res. 2001, 35, 4191-4199. (10) Battistoni, P.; Pavan, M.; Prisciandaro, M.; Cecchi, F. Water Res. 2000, 34, 3033-3041. (11) Battistoni, P.; De Angelis, A.; Pavan, M.; Prisciandaro, M.; Cecchi, F. Water Res. 2001, 35, 2167-2178. (12) Nelson, N. O.; Mikkelsen, R. D.; Hesterberg, D. L. Bioresour. Technol. 2003, 89, 229-236. (13) Bates R. G. pH Determination; Wiley: New York, 1973. (14) APHA, AWWA and WPCF. Standard Methods for the Examination of Water and Wastewater, 15th ed.; American Public Health Association: Washington, DC, 1980. (15) Mineql+ A Chemical Equilibrium Modeling System, Version 4.0, Environmental Research Software, Hallowell, ME, 1998.

496


(16) Nancollas, G. H. Interactions in Electrolyte Solutions; Elsevier: Amsterdam, 1966. (17) Davies, C. W. Ion Association; Butterworth: Washington, 1962. (18) Martell, A.; Smith, R.; Motekaitis, R. NIST Critically Selected Stability Constants of Metal Complexes, Database Version 5.0 NIST Standard Reference Database 46, Texas A&M University, College Station, TX, 1998. (19) Ohlinger, K. N.; Young, T. M.; Schroeder, E. D. Water Res. 1998, 32, 3607-3614. (20) Nielsen, A. E.; Shonel, O. J. Cryst. Growth 1971, 11, 233242.

Kofina and Koutsoukos (21) Mullin, J. W. Crystallization, 3rd ed.; Butterworth-Heinemann: Oxford, 1993. (22) Sangwal, K. Kinetics and Mechanisms of Crystal Growth, In Elementary Crystal Growth; Sangwal, K., Ed., Saan Publishers: Lublin, 1994; pp 83-172. (23) Crystallographica Search-Match, version 2.0.3.1; Ammonium magnesium phosphate hydrate Struvite, syn (PDF Number 15-762), A computer database; Cryosystems: Oxford, 1996-1999.

CG049803E

Spontaneous Precipitation of Struvite from ... - ACS Publications

Recommend Documents